https://www.sciencemadness.org/smwiki/api.php?action=feedcontributions&user=Ave369&feedformat=atomSciencemadness Wiki - User contributions [en]2024-03-28T09:57:25ZUser contributionsMediaWiki 1.25.1https://www.sciencemadness.org/smwiki/index.php?title=Potassium_chlorate&diff=15597Potassium chlorate2024-02-03T18:31:03Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Potassium chlorate<br />
| Reference = <br />
| IUPACName = Potassium chlorate<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Berthollet's salt<br>Potcrate<br>Potassium chlorate(V)<br />
<!-- Images --><br />
| ImageFile = KClO3 by No Tears Only Dreams Now.jpg<br />
| ImageSize = 300<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Crystals of potassium chlorate made from bleach.<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White crystalline solid<br />
| BoilingPt = <br />
| BoilingPtC = 400<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 2.32 g/cm<sup>3</sup><br />
| Formula = KClO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 122.55 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 356<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 3.13 g/100 ml (0 °C)<br>4.46 g/100 ml (10 °C)<br>8.15 g/100 ml (25 °C)<br>13.21 g/100 ml (40 °C)<br>53.51 g/100 ml (100 °C)<br />
| SolubleOther = Reacts with [[sulfuric acid]]<br>Slightly soluble in liq. [[ammonia]], [[glycerol]]<br>Almost insoluble in [[acetone]], [[alcohol]]s<br>Insoluble in alkanes, halocarbons<br />
| Solubility1 = 1 g/100 ml (20 °C)<br />
| Solvent1 = glycerol<br />
| VaporPressure = ~ 0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Monoclinic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = -289.9 kJ/mol<br />
| DeltaHc = <br />
| DeltaHf = −391.2 kJ/mol<br />
| Entropy = 142.97 J·mol<sup>-1</sup>·K<sup>-1</sup><br />
| HeatCapacity = 100.25 J·mol<sup>-1</sup>·K<sup>-1</sup><br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/yWwzE9J/potassium-chlorate-sa.pdf.html Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = 1,870 mg/kg (oral, rat)<br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium chloride]]<br>[[Potassium hypochlorite]]<br>[[Potassium chlorite]]<br>[[Potassium perchlorate]]<br />
}}<br />
}}<br />
'''Potassium chlorate''', also known as '''Berthollet's salt''', is the inorganic chemical compound with the formula '''[[potassium|K]][[Chlorine|Cl]][[Oxygen|O]]<sub>3</sub>''', and is the potassium salt of [[chloric acid]]. It is a strong oxidizing agent.<br />
<br />
==Properties==<br />
===Chemical===<br />
Potassium chlorate is a powerful oxidizer.<br />
<br />
If [[sulfuric acid]] is added to potassium chlorate, [[chloric acid]] is formed:<br />
<br />
:2 KClO<sub>3</sub> + H<sub>2</sub>SO<sub>4</sub> → 2 HClO<sub>3</sub> + K<sub>2</sub>SO<sub>4</sub><br />
<br />
The chloric acid decomposes immediately to perchloric acid and [[chlorine dioxide]]. The latter will spontaneously ignite any combustible material (sugar, paper, dust). As a fun project, various sweets, often gummy bears (jelly babies) are dropped into molten potassium chlorate to produce intense pink flames and, if in a narrow container, a screaming sound.<br />
<br />
Potassium chlorate will decompose if heated in the presence of a catalyst, usually [[manganese dioxide]], releasing [[oxygen]] and leaving [[potassium chloride]] behind. This effect can be taken advantage of in order to quickly add oxygen to a system.<br />
<br />
:2 KClO<sub>3(s)</sub> → 3 O<sub>2(g)</sub> + 2 KCl<sub>(s)</sub><br />
<br />
If the heating occurs without a catalyst, [[potassium perchlorate]] will be formed (although in practice, this is difficult to do):<br />
<br />
:4 KClO<sub>3</sub> → 3 KClO<sub>4</sub> + KCl<br />
<br />
The temperature should be controlled, such that the compound just melts. Too strong heating leads to decomposition with loss of oxygen.<br />
<br />
===Physical===<br />
Potassium chlorate is a transparent to white salt that precipitates as well-formed, lustrous crystals which have a silky texture and are moderately soluble in water and poorly soluble in glycerol. Similar to [[potassium nitrate]], it is not hygroscopic, making it useful as an oxidizer for pyrotechnics. In modern times, however, the use of potassium chlorate in pyrotechnics has declined strongly, because many compositions with potassium chlorate are unstable and sensitive, leading to unacceptable risk of early ignition.<br />
<br />
Unlike hypochlorites, potassium chlorate has no odor, so a 'bleachy' smell is an indication of impure samples.<br />
<br />
==Availability==<br />
Potassium chlorate was available in the past as a fruit growth fertilizer as well as weed killer, but in recent years it has become restricted, due to its powerful oxidizing properties, as it was used in many terror bombings.<br />
<br />
It can also be found in safety match heads, where it is mixed with [[sulfur]] and glue, though one would need a large amount of safety matches. It's much cheaper to make it from the electrolysis of KCl (see below).<br />
<br />
==Preparation==<br />
Potassium chlorate can be prepared by boiling bleach ([[sodium hypochlorite]] or [[calcium hypochlorite]] solution), for about 10-20 minutes, which causes the hypochlorite to disproportionate into chlorate and chloride. Since sodium chloride is less soluble than the chlorate, it will crystallize, while the chlorate will remain in solution. Crystallization begins at about the same time that the bleach reaches one third of the original volume. This chlorate solution is then added to an equivalent amount of saturated [[potassium chloride]] solution, to precipitate the potassium chlorate. Carefully cooling the solution to about 0 degrees Celsius will yield more product. The flat, shiny crystals should then be filtered out and washed multiple times with ice cold water.<ref>http://www.youtube.com/watch?v=JtxQT7aVDeg</ref><br />
<br />
If you already have access to another chlorate, for example, a [[sodium chlorate]] weed killer, preparation of potassium chlorate is very easy. Make a saturated solution of the weed killer at high temperature, cool down the solution to precipitate any sodium chloride adulterant, and perform a metathesis as written above.<br />
<br />
Potassium chlorate can also be produced more efficiently via electrolysis of a saturated solution of [[potassium chloride]] with inert electrodes. This procedure is known as [[alkali chlorate cell]]. A detailed procedure of the process can be found on both the Sciencemadness board as well as in other places online, like YT. Woelen's page also details this process very well.<ref>http://woelen.homescience.net/science/chem/exps/miniature_chlorate_cell/index.html</ref><br />
<br />
Recrystallization of KClO<sub>3</sub> is easy, as it is very soluble in hot water but sparingly soluble in freezing water<br />
<br />
==Projects==<br />
*Screaming jelly baby (gummy bear)<br />
*Burning hearts<br />
*Cockroach cremation<br />
*[[Flash powder]] <br />
*[[Chloric acid]] and [[perchloric acid]] synthesis<br />
<br />
==Handling==<br />
===Safety===<br />
When mixed with combustible materials, even those normally slightly flammable (such as dust and lint), it will burn vigorously in combination and the fires are extremely hard to put out, as the chlorate provides the oxygen for the fire. [[Sulfur]] and red [[phosphorus]], should be avoided in pyrotechnic compositions containing potassium chlorate, as well as any acidic salts, as these mixtures are shock and friction sensitive and prone to spontaneous deflagration (in the safety head matches, such mixture is stabilized with glue). Molten potassium chlorate will ignite any combustible material and can burn even through standard lab safety clothing.<br />
<br />
===Storage===<br />
Potassium chlorate should be stored in closed containers and away from any organic sources, as well as strong acidic vapors. Since it is not hygroscopic, it is not necessary to keep it air tight.<br />
<br />
===Disposal===<br />
Potassium chlorate can be neutralized with a reducing agent, such as [[sodium metabisulfite]], [[sodium bisulfite]], [[sodium sulfite]] or a mixture of sulfuric acid and [[Ammonium iron(II) sulfate|ferrous ammonium sulfate]]. The resulting products should be neutralized with a base and safely poured down the drain.<ref>http://www.oocities.org/capecanaveral/campus/5361/chlorate/destroy.html</ref><br />
<br />
==References==<br />
<references /><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=10727 Chlorate from pool chem?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=24082 Potassium Chlorate from Electrolysis - yay :)]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Potassium compounds]]<br />
[[Category:Chlorates]]<br />
[[Category:Oxidizing agents]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sodium_chlorate&diff=15596Sodium chlorate2024-02-03T18:18:21Z<p>Ave369: /* Availability */</p>
<hr />
<div>{{distinguish|Sodium chlorite}}<br />
{{Chembox<br />
| Name = Sodium chlorate<br />
| Reference =<br />
| IUPACName = Sodium chlorate<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Asex<br>Polybor-chlorate<br>Sodium chlorate(V)<br />
<!-- Images --><br />
| ImageFile = Sodium chlorate crystals.jpg<br />
| ImageSize = 300<br />
| ImageAlt =<br />
| ImageName =<br />
| ImageCaption = Sodium chlorate crystals<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White crystalline solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 2.49 g/cm<sup>3</sup> (15 °C)<br> 2.54 g/cm<sup>3</sup> (20.2 °C)<br />
| Formula = NaClO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 106.44 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 248<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 79 g/100 ml (0 °C)<br>89 g/100 ml (10 °C)<br>105.7 g/100 ml (25 °C)<br>125 g/100 ml (40 °C)<br>220.4 g/100 ml (100 °C)<br />
| SolubleOther = Soluble in [[glycerol]], [[hydrazine]], [[methanol]]<br> Slightly soluble in liq. [[ammonia]], [[ethanol]]<br>Sparingly soluble in [[acetone]]<br>Insoluble in hydrocarbons<br />
| Solubility1 = 14.7 g/100 g<br />
| Solvent1 = ethanol<br />
| Solubility2 = 16 g/100 g (25 °C)<br />
| Solvent2 = ethylene glycol<br />
| Solubility3 = 20 g/100 g (15.5 °C)<br />
| Solvent3 = glycerol<br />
| Solubility4 = 66 g/100 g (25 °C)<br />
| Solvent4 = hydrazine<br />
| Solubility5 = 51.35 g/100 g (25 °C)<br />
| Solvent5 = methanol<br />
| VaporPressure = ~ 0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Cubic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = -275 kJ/mol<br />
| DeltaHc = <br />
| DeltaHf = -365.4 kJ/mol<br />
| Entropy = 129.7 J·mol<sup>-1</sup>·K<sup>-1</sup><br />
| HeatCapacity = 104.6 J·mol<sup>-1</sup>·K<sup>-1</sup><br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/Z8VLpWn/sodium-chlorate-sa.pdf.html Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = 6,500 mg/kg (rat, oral)<br>700 mg/kg (dog, oral)<br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Harmful<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Sodium hypochlorite]]<br>[[Sodium chlorite]]<br>[[Sodium perchlorate]]<br />
}}<br />
}}<br />
'''Sodium chlorate''' is an inorganic compound, comprised of equal numbers of sodium cations and [[chlorate]] anions, giving it the fomula '''NaClO<sub>3</sub>'''. It is a very powerful oxidizer.<br />
<br />
==Properties==<br />
===Chemical===<br />
It is a strong oxidizing agent, easily supplying oxygen to combustibles. It decomposes above 300 °C yielding oxygen and [[sodium chloride]].<br />
<br />
:2 NaClO<sub>3</sub> → 2 NaCl + 3 O<sub>2</sub><br />
<br />
Sodium chlorate will react with potassium chloride to precipitate [[potassium chlorate]]:<br />
<br />
: KCl + NaClO<sub>3</sub> → NaCl + KClO<sub>3</sub><br />
<br />
===Physical===<br />
Sodium chlorate is a colorless or white crystalline solid with a cubic crystal structure. It is soluble in water, methanol, glycerol, hydrazine and slightly soluble in ethanol and ammonia. Because sodium chlorate is [[Hygroscopy|hygroscopic]], potassium chlorate is often preferred for use as an oxidizer.<br />
<br />
Sodium chlorate crystals are chiral, and both enantiomorphic (R/L) forms crystallize from solution in roughly equal proportions. Seeding with a selected enantiomorphic form leads to enrichment in that form. The presence of D-glucose or other chiral sugars skews the proportion of the two forms during crystallization. This is only a solid state effect due to the crystal structure, and in solution no enantiomers or optical activity can be observed. Big crystals can be sorted by hand using polarizing filters into their respective enantiomorphs.<ref>''Gmelins Handbuch der anorganischen Chemie, Natrium'', Verlag Chemie GmbH, Berlin, 8th edition '''1928''', p. 401</ref><br />
<br />
==Availability==<br />
It can be bought as "weed killer" at a hardware store, or it can be bought online. Some brands may be adulterated with sodium chloride; in this case, the weed killer cannot be used as a reagent as is, but can be used as a raw material to prepare [[potassium chlorate]]. Many countries, however, have banned sodium chlorate weed killers.<br />
<br />
Its sale is banned in the EU.<br />
<br />
==Preparation==<br />
Sodium chlorate can be produced by boiling [[bleach]], which causes it to disproportionate into sodium chlorate and sodium chloride.<br />
<br />
A more efficient way of producing sodium chlorate is via the electrolysis of a supersaturated sodium chloride solution with an appropriate anode at ~5 volts DC.<br />
<br />
Although the exact reactions are very complex, the basic overall equation is:<br />
<br />
: NaCl + 3 H<sub>2</sub>O → NaClO<sub>3</sub> + 3 H<sub>2</sub><br />
<br />
==Projects==<br />
*Preparation of [[potassium chlorate]]<br />
*Make a dry chemical oxygen generator: Heat is generated by oxidation of a small amount of iron powder mixed with the sodium chlorate, and the reaction consumes less oxygen than is produced. [[Barium peroxide]] is used to absorb the chlorine which is a minor product in the decomposition.<ref>http://en.wikipedia.org/wiki/Sodium_chlorate#cite_note-8</ref> An ignitor charge is activated by pulling on the emergency mask. Similarly, the [http://en.wikipedia.org/wiki/SolidOx_(welding) Solidox] welding system used pellets of sodium chlorate mixed with combustible fibers to generate oxygen.<br />
<br />
==Handling==<br />
===Safety===<br />
Powerful oxidizer! Fire hazard! Keep away from any flammables.<br />
<br />
Due to its oxidative nature, sodium chlorate can be very toxic if ingested. The oxidative effect on [http://en.wikipedia.org/wiki/Hemoglobin hemoglobin] leads to [http://en.wikipedia.org/wiki/Methaemoglobin methaemoglobin] formation, which is followed by [http://en.wikipedia.org/wiki/Denaturation_(biochemistry) denaturation] of the [http://en.wikipedia.org/wiki/Globin globin] protein and a [http://en.wikipedia.org/wiki/Cross-link cross-linking] of [http://en.wikipedia.org/wiki/Erythrocyte erythrocyte]membrane proteins with resultant damage to the membrane enzymes. This leads to increased permeability of the membrane, and severe [http://en.wikipedia.org/wiki/Hemolysis hemolysis]. The denaturation of hemoglobin overwhelms the capacity of the [http://en.wikipedia.org/wiki/Glucose-6-phosphate_dehydrogenase G6PD] [http://en.wikipedia.org/wiki/Metabolic_pathway metabolic pathway]. In addition, this enzyme is directly denatured by chlorate reducing its activity.<br />
<br />
Avoid contact with concentrated acids like sulfuric acid due to formation of highly reactive [[chloric acid]].<br />
<br />
===Storage===<br />
Sodium chlorate should be stored in closed bottles, away from any flammable materials and strong acids. Since it's hygroscopic, it should be kept in a dry place.<br />
<br />
===Disposal===<br />
Sodium chlorate can be neutralized with acidified sodium or potassium metabisulfite.<br />
<br />
==Gallery==<br />
<gallery widths="200" position="center" columns="4" orientation="none"><br />
Sodium chlorate polarized.jpg|Crystals of sodium chlorate between two polarizing filters<br />
</gallery><br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=15903 Sodium chlorate production questions]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Sodium compounds]]<br />
[[Category:Chlorates]]<br />
[[Category:Oxidizing agents]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Hydrogen_peroxide&diff=15083Hydrogen peroxide2023-07-23T12:58:23Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Hydrogen peroxide<br />
| Reference =<br />
| IUPACName = Hydrogen peroxide<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Dioxidane<br>Oxidanyl<br>Perhydroxic acid<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Very light blue liquid (conc.)<br>Colorless liquid (dil.)<br />
| BoilingPt = <br />
| BoilingPtC = 150.2<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 1.450 g/cm<sup>3</sup> (20 °C, pure)<br>1.11 g/cm<sup>3</sup> (20 °C, 30% aq. solution)<br />
| Formula = H<sub>2</sub>O<sub>2</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 34.0147 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = −0.43<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Slightly sharp<br />
| pKa = 11.75<br />
| pKb = <br />
| Solubility = Miscibility<br />
| SolubleOther = Reacts with ketones<br>Soluble in [[alcohol]], [[ether]]<br>Insoluble in [[benzene]], [[chloroform]], [[petroleum ether]], [[toluene]]<br />
| Solvent = <br />
| VaporPressure = 5 mmHg (30 °C)<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = −187.80 kJ/mol<br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/iq5Srtl/hydrogen-peroxide-30-sa.pdf.html Sigma-Aldrich] (30%)<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Water]]<br />
}}<br />
}}<br />
'''Hydrogen peroxide''' ('''H<sub>2</sub>O<sub>2</sub>''') is a mostly clear, blue-ish liquid with similar melting and boiling points to water. It is a powerful and versatile [[Oxidizing agent|oxidizer,]] but can act as a reducing agent in some circumstances. It also acts as a very weak acid (pK<sub>a</sub> = 11.6), forming hydrated peroxide salts (such as [[sodium peroxide]] octahydrate) with alkalis in aqueous solution.<br />
<br />
==Properties==<br />
===Chemical===<br />
Hydrogen peroxide disproportionates into [[water]] and [[oxygen]] gas. This happens rapidly at high temperatures or when a catalyst, such as [[manganese dioxide]] or [[potassium iodide]], is added and this is often used to produce oxygen gas in a home chemistry setting:<br />
<br />
: H<sub>2</sub>O<sub>2</sub> → H<sub>2</sub>O + ½ O<sub>2</sub><br />
<br />
Hydrogen peroxide can be used as an oxidizer, and may enhance the oxidizing capabilities on mixing. For example, mixtures of [[sulfuric acid]] and hydrogen peroxide not only will react faster than the acid alone, but will also react with organic compounds, sometimes explosively. Depending on the ratio of peroxide and sulfuric acid, there are several types: [[piranha solution]] (min H<sub>2</sub>SO<sub>4</sub>:H<sub>2</sub>O<sub>2</sub> ratio 3:1), [[peroxymonosulfuric acid]] (H<sub>2</sub>SO<sub>4</sub>:H<sub>2</sub>O<sub>2</sub> molar ratio of 1:1), [[peroxydisulfuric acid]] (H<sub>2</sub>SO<sub>4</sub>:H<sub>2</sub>O<sub>2</sub> molar ratio of 1:2).<br />
<br />
Hydrogen peroxide is dangerous as it can cause explosions when in contact with combustible materials in high concentration. An acid mixture of hydrogen peroxide and [[hydrochloric acid]] behaves like an oxidizing acid, similarly to nitric acid, and reacts with non-reactive metals such as [[copper]]:<br />
<br />
: Cu + 2 HCl + H<sub>2</sub>O<sub>2</sub> → CuCl<sub>2</sub> + 2 H<sub>2</sub>O<br />
<br />
This acid mixture alone undergoes a number of slow reactions with various products, which include oxygen and chlorine ''in statu nascendi'' and [[hypochlorous acid]], which makes it a powerful oxidizer that can be used as a much safer alternative to concentrated nitric acid for dissolving metals. However, you should never store this "green acid" as it invariably decomposes over time.<br />
<br />
Hydrogen peroxide will reduce [[potassium permanganate]] is reduced to Mn<sub>2+</sub>:<br />
<br />
: 2 KMnO<sub>4</sub> + 3 H<sub>2</sub>O<sub>2</sub> → 2 MnO<sub>2</sub> + 2 KOH + 2 H<sub>2</sub>O + 3 O<sub>2</sub><br />
<br />
Reaction with hypochlorites, like that with [[sodium hypochlorite]] will also produce oxygen, more specifically singlet oxygen.:<br />
<br />
: NaOCl + H<sub>2</sub>O<sub>2</sub> → NaCl + H<sub>2</sub>O + O<sub>2</sub><br />
<br />
This route is a safe way or neutralizing bleach, without producing any unpleasant smells.<br />
<br />
Reaction with [[borax]] and [[sodium hydroxide]] leads to [[sodium perborate]]:<br />
<br />
: Na<sub>2</sub>B<sub>4</sub>O<sub>7</sub> + 2 NaOH + 4 H<sub>2</sub>O<sub>2</sub> → 2 Na<sub>2</sub>B<sub>2</sub>O<sub>4</sub>(OH)<sub>4</sub> + H<sub>2</sub>O<br />
<br />
Reaction with [[ozone]] forms trioxidane.<br />
<br />
It converts [[carboxylic acid]]s into peroxy acids<br />
<br />
Hydrogen peroxide reacts with [[acetone]] to form the highly dangerous [[acetone peroxide]]. It also forms stable [[adduct]]s with [[urea]] ([[urea peroxide]]), sodium carbonate ([[sodium percarbonate]]) and other compounds.<br />
<br />
Hydrogen peroxide will oxidize thioethers to sulfoxides:<br />
<br />
: R-S-R' + H<sub>2</sub>O<sub>2</sub> → R-S(=O)-R' + H<sub>2</sub>O<br />
<br />
Hydrogen peroxide solution containing ferric ion, is called [[Fenton's reagent]], and can be used to safely neutralize many organic compounds. <br />
<br />
===Physical===<br />
Hydrogen peroxide is tinted slightly blue in high concentrations. It has boiling and melting points similar to water, but can be concentrated by [[fractional crystallization]]. Concentrations of peroxide 30% and above are considered concentrated. <br />
<br />
==Availability==<br />
Hydrogen peroxide is available readily as a disinfectant in pharmacies and grocery stores, but may only be obtained easily in low concentrations often as 3% or 6% solutions.<br />
<br />
Higher concentration peroxide is sold for animal disinfectants, pool/spa treatments, hair bleaching. Some may contain other chemicals, including stabilizers, so read the label first.<br />
<br />
In the EU, the sale of hydrogen peroxide solutions with a concentration higher than 12% is restricted for private individuals.<br />
<br />
==Preparation==<br />
Hydrogen peroxide can be prepared by reacting concentrated [[sulfuric acid]] and [[barium peroxide]].<br />
<br />
: BaO<sub>2</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>O<sub>2</sub> + BaSO<sub>4</sub><br />
<br />
The insoluble [[barium sulfate]] is filtered from the mixture.<br />
<br />
Due to the availability of low concentrations of peroxide worldwide, concentrated hydrogen peroxide solutions are often prepared by evaporating the water from the peroxide, making sure not to boil the solution (as this will break down the peroxide).<br />
<br />
One of the earliest large scale production of hydrogen peroxide involves the electrolysis of a solution of [[sulfuric acid]] of 50% concentration, or [[ammonium bisulfate]] solution. The resulting hydrogen peroxide solution has a concentration of 30%.<br />
<br />
While sulfuric acid is more readily available, the ammonium bisulfate process is cheaper and has a higher cell efficiency. The electrolysis is carried out in stoneware tanks with platinum electrodes; conversion of bisulfate to the persulfate takes place at the anode. After hydrolysis of the persulfate (with steam) in an evaporator, the resulting dilute aqueous solution of H<sub>2</sub>O<sub>2</sub> is separated from the bisulfate and further distilled in a stoneware distillation column. The resulting solution is approximately 30 w/o H<sub>2</sub>O<sub>2</sub>. Both the cathode liquor (after purification) and the bisulfate from the evaporator (and separator) are recycled back to the cells.<br />
<br />
Industrially, hydrogen peroxide is produced via the anthraquinone process (2-ethylanthraquinone gives the best performance, as it has better selectivity), where the anthraquinone can be hydrogenated in the presence of a [[palladium]] catalyst to 9,10-anthrahydroquinone, which, upon oxidation with oxygen, reverts back to anthraquinone releasing hydrogen peroxide. Cody's Lab made a [https://www.youtube.com/watch?v=Mt1itiHT6wU video] on making hydrogen peroxide via this route.<br />
<br />
==Projects==<br />
*Make [[sulfuric acid]]<br />
*Make peroxides<br />
*Cleaning [[silver]] metal<br />
*Make [[piranha solution]]<br />
*Make [[Fenton's reagent]]<br />
*Make copper etching solution<br />
*Make singlet oxygen<br />
<br />
==Handling==<br />
===Safety===<br />
As it is an oxidizer, high concentrations of hydrogen peroxide can ignite or detonate combustible or explosive materials. Lower concentrations are much safer, but regardless of concentration, poisonous. Concentrated hydrogen peroxide leaves weak itching or stinging burns and a white discoloration on skin.<br />
<br />
===Storage===<br />
Hydrogen peroxide solutions are best stored in cold dark places, such as a fridge. High concentration peroxides are metastable and will slowly build-up pressure, so it's recommended to open the bottles from time to time, to release the pressure, and store the solutions with ample headroom in the bottles. [[Phosphoric acid]] and [[Ethylenediaminetetraacetic acid|EDTA]] are added as stabilizers.<br />
<br />
NEVER store hydrogen peroxide near volatile organic compounds, such as [[acetone]], as there is a risk of forming [[acetone peroxide]].<br />
<br />
Concentrated solutions of hydrogen peroxide are known to slowly react with glass. This does not harm the glass significantly, but it does harm the peroxide: the reaction yields sodium hydroxide, which can provoke its decomposition. It is preferable to use polyethylene to store such solutions.<br />
<br />
===Disposal===<br />
Hydrogen peroxide can be decomposed by adding a catalyst, such as [[manganese dioxide]] or iron oxides (ordinary rust will do). This method however should not be used to neutralize concentrated peroxide (>30%) as the decomposition will generate lots of heat and can lead to explosion, and the peroxide should first be diluted. The explosion that crippled the Kursk submarine for example, occured when the peroxide that leaked from a torpedo entered in contact with some rust. It's recommended to not be poured down the drain, as it will quickly decompose in the sewage and may pose an explosion hazard. Adding a sulfide, such as [[lead(II) sulfide]], will result in [[lead(II) sulfate]] and water.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=15881 Hydrogen Peroxide purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16726 Hydrogen peroxide concentration by freezing I need to make 12% h2o2 from 3%]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=126644 Hydrogen peroxide maximum safe concentration]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=157743 Storage of 35% H2O2]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Hydrogen compounds]]<br />
[[Category:Oxygen compounds]]<br />
[[Category:Acids]]<br />
[[Category:Weak acids]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Peroxides]]<br />
[[Category:Inorganic peroxides]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Materials available as food grade]]<br />
[[Category:Essential reagents]]<br />
[[Category:Irritants]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Anthocyanin&diff=15067Anthocyanin2023-07-21T20:03:54Z<p>Ave369: /* Preparation */</p>
<hr />
<div>'''Anthocyanin''' is a naturally occurring pH indicator that behaves similarly to litmus. It is actually a whole family of related organic compounds; the most useful particular anthocyanins are '''cyanidin''' (found in red cabbage, red wine) and '''chrysanthemin''' (found in hibiscus). Cyanidin is better soluble in [[water]], chrysanthemin in [[ethanol]], but their indicator properties are identical.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
Anthocyanins are solid dye-like substances, not very stable. They are usually stored as solutions (chrysanthemin, dissolved in ethanol, is the best for long-term storage).<br />
<br />
=== Chemical ===<br />
Cyanidin and chrysanthemin are useful pH indicators, with properties similar to [[litmus]]. In neutral pH, they are violet-blue. In low pH, they are red (very pronounced red in strongly acidic solutions, pink in weakly acidic ones). In high pH, they are green (bluish-green in weakly basic solutions, yellowish-green in strongly basic ones). In addition, extremely strong bases cause anthocyanins to slowly degrade, turn yellow and lose their indicator properties; concentrated alkalis first turn anthocyanins sickly green as normal, then in seconds or minutes they turn yellow. <br />
<br />
[[File:Anthocyanin_scale.png]]<br />
<br />
Aqueous solutions of anthocyanins aren't terribly stable and degrade in a matter of days. Ethanol solutions of chrysanthemin are stable.<br />
<br />
== Availability ==<br />
Anthocyanins can be extracted from common food products. Cyanidin is usually leached with water from red cabbage. The more desirable anthocyanin, chrysanthemin, is obtained from hibiscus herbal tea, which is marketed under names such as "Sorrel", "Red Zinger", "Karkade", "Italian Tea", "Flor de Jamaica", etc. in various countries.<br />
<br />
== Preparation ==<br />
Cyanidin is usually prepared in the form of cabbage water, by infusing water on red cabbage leaves. Such an indicator is useful but inferior to chrysanthemin, which is leached by ethanol from hibiscus herbal tea. Do not use isopropanol: while it does dissolve chrysanthemin, the solubility is significantly lower and you won't have any deep or bright colors.<br />
<br />
Once an anthocyanin solution is prepared, one has to adjust its pH. Cabbage water usually does not need such an adjustment, it is already purple. Hibiscus, on the other hand, also contains acids which make its extract reddish-pink. To adjust the pH of the alcohol solution, one should drop a basic solution such as [[sodium hydroxide]] or [[ammonia]] into it until it becomes purple-blue. Important: before you adjust the pH, filter the infusion off the hibiscus plant matter. Acids that the plant matter contains neutralize the base and ruin your work!<br />
<br />
You can keep the dissolved form of anthocyanin, but test paper strips can also be made from it. In this, chrysanthemin is also superior: the alcoholic solution dries on paper much faster and gives a more uniform color. The best paper for the strips is filter paper, it gives the brightest colors with concentrated solutions. It is recommended to drench paper strips in the solution while it is still acidic, then neutralize it: the neutral form of anthocyanin is sparsely soluble in ethanol and forms a colloid, so the neutralizing agent acts as a mordant and results in brighter color of the strips.<br />
<br />
You can also make acidic and basic test paper in addition to the universal one. This is recommended if you use normal paper for strips, which gives more pastel colors. Acidic test paper is pink or red rather than violet, it shows no color change with acids, but lets you identify bases easier. Basic test paper is green rather than violet, it shows no color change in bases (other than the yellow discoloration in very strong alkali) but lets you identify acids easier. Acidic test paper is easier to make: just skip the step of pH adjustment with chrysanthemin. To make basic paper, drop more base into the solution until it's green. However, do not overdo it: if your indicator solution or paper turns yellow, it is ruined and no longer good! That's why ammonia is preferable to nonvolatile bases for adjusting pH: when the test strips dry, a nonvolatile alkali can get too concentrated on the paper and ruin it, making it yellow and non-functional.<br />
<br />
One package of herbal tea or one head of cabbage is enough for hundreds, maybe a thousand, of test paper strips. This makes this "DIY litmus" essentially free.<br />
<br />
== Projects ==<br />
* Make DIY litmus substitute paper strips! Bajillions and metric craptons of paper strips! Free pH testing strips for everyone!<br />
<br />
== Handling ==<br />
=== Safety ===<br />
Dear Cthulhu forbid, what safety? Anthocyanins are edible!<br />
<br />
=== Storage ===<br />
Aqueous solutions of anthocyanins are very perishable. However, ethanol solutions and test paper have a long shelf life and do not require any special precautions to store.<br />
<br />
=== Disposal ===<br />
<br />
No special disposal is required. Any form of anthocyanin can be disposed of just like common garbage, into the trash bin or down the drain.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2464 Isolation of flavonoids (anthocyanidins ) from red cabbage juice]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Organic compounds]]<br />
[[Category:Aromatic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]<br />
[[Category:Edible chemicals]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Anthocyanin&diff=15066Anthocyanin2023-07-21T20:02:17Z<p>Ave369: /* Preparation */</p>
<hr />
<div>'''Anthocyanin''' is a naturally occurring pH indicator that behaves similarly to litmus. It is actually a whole family of related organic compounds; the most useful particular anthocyanins are '''cyanidin''' (found in red cabbage, red wine) and '''chrysanthemin''' (found in hibiscus). Cyanidin is better soluble in [[water]], chrysanthemin in [[ethanol]], but their indicator properties are identical.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
Anthocyanins are solid dye-like substances, not very stable. They are usually stored as solutions (chrysanthemin, dissolved in ethanol, is the best for long-term storage).<br />
<br />
=== Chemical ===<br />
Cyanidin and chrysanthemin are useful pH indicators, with properties similar to [[litmus]]. In neutral pH, they are violet-blue. In low pH, they are red (very pronounced red in strongly acidic solutions, pink in weakly acidic ones). In high pH, they are green (bluish-green in weakly basic solutions, yellowish-green in strongly basic ones). In addition, extremely strong bases cause anthocyanins to slowly degrade, turn yellow and lose their indicator properties; concentrated alkalis first turn anthocyanins sickly green as normal, then in seconds or minutes they turn yellow. <br />
<br />
[[File:Anthocyanin_scale.png]]<br />
<br />
Aqueous solutions of anthocyanins aren't terribly stable and degrade in a matter of days. Ethanol solutions of chrysanthemin are stable.<br />
<br />
== Availability ==<br />
Anthocyanins can be extracted from common food products. Cyanidin is usually leached with water from red cabbage. The more desirable anthocyanin, chrysanthemin, is obtained from hibiscus herbal tea, which is marketed under names such as "Sorrel", "Red Zinger", "Karkade", "Italian Tea", "Flor de Jamaica", etc. in various countries.<br />
<br />
== Preparation ==<br />
Cyanidin is usually prepared in the form of cabbage water, by infusing water on red cabbage leaves. Such an indicator is useful but inferior to chrysanthemin, which is leached by ethanol from hibiscus herbal tea.<br />
<br />
Once an anthocyanin solution is prepared, one has to adjust its pH. Cabbage water usually does not need such an adjustment, it is already purple. Hibiscus, on the other hand, also contains acids which make its extract reddish-pink. To adjust the pH of the alcohol solution, one should drop a basic solution such as [[sodium hydroxide]] or [[ammonia]] into it until it becomes purple-blue. Important: before you adjust the pH, filter the infusion off the hibiscus plant matter. Acids that the plant matter contains neutralize the base and ruin your work!<br />
<br />
You can keep the dissolved form of anthocyanin, but test paper strips can also be made from it. In this, chrysanthemin is also superior: the alcoholic solution dries on paper much faster and gives a more uniform color. The best paper for the strips is filter paper, it gives the brightest colors with concentrated solutions. It is recommended to drench paper strips in the solution while it is still acidic, then neutralize it: the neutral form of anthocyanin is sparsely soluble in ethanol and forms a colloid, so the neutralizing agent acts as a mordant and results in brighter color of the strips.<br />
<br />
You can also make acidic and basic test paper in addition to the universal one. This is recommended if you use normal paper for strips, which gives more pastel colors. Acidic test paper is pink or red rather than violet, it shows no color change with acids, but lets you identify bases easier. Basic test paper is green rather than violet, it shows no color change in bases (other than the yellow discoloration in very strong alkali) but lets you identify acids easier. Acidic test paper is easier to make: just skip the step of pH adjustment with chrysanthemin. To make basic paper, drop more base into the solution until it's green. However, do not overdo it: if your indicator solution or paper turns yellow, it is ruined and no longer good! That's why ammonia is preferable to nonvolatile bases for adjusting pH: when the test strips dry, a nonvolatile alkali can get too concentrated on the paper and ruin it, making it yellow and non-functional.<br />
<br />
One package of herbal tea or one head of cabbage is enough for hundreds, maybe a thousand, of test paper strips. This makes this "DIY litmus" essentially free.<br />
<br />
== Projects ==<br />
* Make DIY litmus substitute paper strips! Bajillions and metric craptons of paper strips! Free pH testing strips for everyone!<br />
<br />
== Handling ==<br />
=== Safety ===<br />
Dear Cthulhu forbid, what safety? Anthocyanins are edible!<br />
<br />
=== Storage ===<br />
Aqueous solutions of anthocyanins are very perishable. However, ethanol solutions and test paper have a long shelf life and do not require any special precautions to store.<br />
<br />
=== Disposal ===<br />
<br />
No special disposal is required. Any form of anthocyanin can be disposed of just like common garbage, into the trash bin or down the drain.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2464 Isolation of flavonoids (anthocyanidins ) from red cabbage juice]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Organic compounds]]<br />
[[Category:Aromatic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]<br />
[[Category:Edible chemicals]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Propanol&diff=15010Propanol2023-06-09T11:13:10Z<p>Ave369: /* Safety */</p>
<hr />
<div>{{Stub}}<br />
{{Chembox<br />
| Name = Propanol<br />
| Reference = <br />
| IUPACName = Propan-1-ol<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = 1-Hydroxypropane<br>Ethylcarbinol<br>n-PrOH<br>n-Propanol<br>n-Propyl alcohol<br>Propionic alcohol<br>Propionyl alcohol<br>Propionylol<br>Propyl alcohol<br>Propylic alcohol<br>Propylol<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless liquid<br />
| BoilingPt = <br />
| BoilingPtC = 97.2<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 0.803 g/mL<br />
| Formula = C<sub>3</sub>H<sub>8</sub>O<br />
| HenryConstant = <br />
| LogP = 0.329<br />
| MolarMass = 60.10 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = −126<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Alcoholic<br />
| pKa = 16<br />
| pKb = -2<br />
| Solubility = Miscible<br />
| SolubleOther = Miscible with [[alcohol]]s, [[ester]]s, [[ether]]s, glycols, [[acetone]]<br />
| Solvent = <br />
| VaporPressure = 21.0 mmHg (25 °C)<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = −2,021.56–−2,021.06 kJ/mol<br />
| DeltaHf = −302.79–−302.29 kJ/mol<br />
| Entropy = 192.8 J·K<sup>−1</sup>·mol<sup>−1</sup><br />
| HeatCapacity = 143.96 J·K<sup>−1</sup>·mol<sup>−1</sup><br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = 371 °C (700 °F; 644 K)<br />
| ExploLimits = 2.2% - 13.7%<br />
| ExternalMSDS = [https://www.docdroid.net/b1shZwz/propanol-sa.pdf.html Sigma-Aldrich]<br />
| FlashPt = 22 °C (72 °F; 295 K)<br />
| LD50 = 2,800 mg/kg (rabbit, oral)<br>6,800 mg/kg (mouse, oral)<br>1,870 mg/kg (rat, oral)<br />
| LC50 = <br />
| MainHazards = Flammable<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Ethanol]]<br>[[Butanol]]<br />
}}<br />
}}<br />
'''Propanol''', also known as '''1-propanol''', '''propan-1-ol''', '''1-propyl alcohol''', '''n-propyl alcohol''', or '''n-propanol''', is a primary alcohol with the formula '''CH<sub>3</sub>CH<sub>2</sub>CH<sub>2</sub>OH'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
Oxidation of 1-propanol with [[chromic acid]] yields [[propionic acid]].<br />
<br />
===Physical===<br />
Propanol is a colorless liquid with an alcohol like smell. It is miscible in water and many organic solvents.<br />
<br />
==Availability==<br />
Delete this section if not applicable<br />
<br />
==Preparation==<br />
1-Propanol is manufactured by catalytic hydrogenation of propionaldehyde.<br />
<br />
==Projects==<br />
*Make n-propyl iodide<br />
<br />
==Handling==<br />
===Safety===<br />
Propanol is thought to be similar to ethanol in its effects on human body, but 2-4 times more potent. However, the toxicity of propanol is comparable to that of [[methanol]] (LD50 1870 mg/kg, as compared to 1500 mg/kg of methanol and 7060 mg/kg of ethanol), so it should never be consumed.<br />
<br />
===Storage===<br />
In closed bottles away from any heat source.<br />
<br />
===Disposal===<br />
Propanol can be safely burned.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13477 Cheap Propanol?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=63718 n-propanol to propionaldehyde]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Organic compounds]]<br />
[[Category:Alcohols]]<br />
[[Category:Primary alcohols]]<br />
[[Category:Solvents]]<br />
[[Category:Polar solvents]]<br />
[[Category:Volatile chemicals]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sulfuric_acid&diff=14960Sulfuric acid2023-03-27T21:08:38Z<p>Ave369: /* Storage */</p>
<hr />
<div>{{Chembox<br />
| Name = Sulfuric acid<br />
| Reference =<br />
| IUPACName = Sulfuric acid<br />
| PIN = Sulfuric acid<br />
| SystematicName = Sulfuric acid<br />
| OtherNames = Battery acid<br>Dihydrogen sulfate<br>Oil of vitriol<br>Spirit of vitriol<br>Sulphuric acid<br />
<!-- Images --><br />
| ImageFile = Smw1.png<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Structure of sulfuric acid<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = OS(=O)(=O)O<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless oily liquid<br />
| BoilingPt = <br />
| BoilingPtC = 337<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (above 300 °C slowly decomposes)<br />
| Density = 1.84 g/cm<sup>3</sup><br />
| Formula = H<sub>2</sub>SO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 98.079 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 10<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless (air above it may feel dry due to its strong hygroscopicity)<br />
| pKa = −3;1.99<br />
| pKb = <br />
| Solubility = Miscible<br />
| SolubleOther = Reacts with [[amine]]s<br>Miscible with [[alcohol]]s<br>Immiscible with hydrocarbons<br />
| Solvent = <br />
| VaporPressure = 0.001 mmHg (20 °C)<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = −814 kJ·mol<sup>−1</sup><br />
| Entropy = 157 J·mol<sup>−1</sup>·K<sup>−1</sup><br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.fishersci.com/msdsproxy%3FproductName%3DA300700LB%26productDescription%3DSULFURIC%2BAC%2BACS%2B700LB%26catNo%3DA300-700LB%26vendorId%3DVN00033897%26storeId%3D10652 FisherSci]<br />
| FlashPt = Non-flammable<br />
| LD50 = 2.140 mg/kg (rat, oral)<br />
| LC50 = 50 mg/m<sup>3</sup> (guinea pig, 8 hr)<br>510 mg/m<sup>3</sup> (rat, 2 hr)<br>320 mg/m<sup>3</sup> (mouse, 2 hr)<br>18 mg/m<sup>3</sup> (guinea pig)<br />
| MainHazards = Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Sulfurous acid]]<br>[[Sulfur trioxide]]<br />
}}<br />
}}<br />
'''Sulfuric acid''' (alternative spelling '''sulphuric acid'''), represented by the molecular formula '''H<sub>2</sub>SO<sub>4</sub>''', is one of the most important [[acid]]s in chemistry and the most important chemical to industries in the world. It is the strongest easily available acid, with a [[Measures of acidity|pK<sub>a</sub>]] of -3.<br />
<br />
==Properties==<br />
===Chemical properties===<br />
Sulfuric acid is a diprotic acid, and thus it is able to give away two protons (H<sup>+</sup>). It first dissociates to form [[hydronium]] and hydrogen sulfate/bisulfate ions, with a pK<sub>a</sub> of -3, indicative of a strong acid:<br />
<br />
: H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → H<sub>3</sub>O + HSO<sub>4</sub><sup>−</sup><br />
<br />
The second dissociation forms sulfate and another hydronium ion from a hydrogen sulfate ion. It has a pKa of 1.99, indicative of a mid-strength acid, and occurs like this: <br />
<br />
: HSO<sub>4</sub><sup>−</sup> + H<sub>2</sub>O ⇌ H<sub>3</sub>O<sup>+</sup> + SO<sub>4</sub><sup>2-</sup><br />
<br />
Concentrated sulfuric acid also has a strong oxidizing effect, converting nonmetals such as [[carbon]] and [[sulfur]] to [[carbon dioxide]] and [[sulfur dioxide]], respectively, reducing sulfuric acid into sulfur dioxide and water in the process.<br />
<br />
: 2 H<sub>2</sub>SO<sub>4</sub> + C → CO<sub>2</sub> + SO<sub>2</sub> + H<sub>2</sub>O + H<sub>2</sub>SO<sub>4</sub><br />
<br />
: 2 H<sub>2</sub>SO<sub>4</sub> + S → 2 SO<sub>2</sub> + H<sub>2</sub>O + H<sub>2</sub>SO<sub>4</sub><br />
<br />
This property is useful for producing large amounts of sulfur dioxide for use as a reducing agent if water is continually removed. Heat accelerates this process.<br />
<br />
Sulfuric acid is sufficiently strong enough to protonate [[nitric acid]], forming the nitronium ion, which can be used in a nitration mixture to make [[alkyl nitrate]]s.<br />
<br />
In organic chemistry, sulfuric acid is the most practical acid in most cases where a source of H<sub>3</sub>O<sup>+</sup> ions are needed as it introduces the least amount of water. Organic compounds are often easily attacked by the nucleophiles left behind by the dissociation of acids such as HCl which leaves Cl<sup>-</sup> ions behind which can easily attack many organic compounds. However, the [[sulfate]] ions left behind by the dissociation of sulfuric acid are far less reactive than the ions left behind by most acids, it allows to protonate the reaction mixture without causing undesired side reactions in most cases.<br />
<br />
When concentrated, it is strongly [[hygroscopy|hygroscopic]] and has strong dehydrating properties. It can break down most organic molecules containing OH<sup>-</sup> groups to use them to form water, leaving only the carbon behind. This property is exploited in the famous [http://youtu.be/w6lfq7BOCik "black snake" demonstration], where sulfuric acid dehydrates [[sucrose]] (table sugar), forming water with the hydrogen and oxygen atoms and leaving amorphous carbon behind.<br />
<br />
===Physical properties===<br />
[[File:H2so4boil.jpg|thumb|left|342px|Boiling point of H2SO4 VS concentration]]<br />
<br />
Sulfuric acid is an oily liquid at room temperature. It is colorless but often has a very light yellow color when slightly contaminated with iron or carbon from organic matter like dust. Even very small amounts of dissolved organic matter can change the color of concentrated sulfuric acid to pale yellow or pink, red, brown, and even black. It is commonly sold diluted at around 35% w/w with water as car battery acid and concentrated between 95% and 98% w/w as drain cleaner.<br />
<br />
Sulfuric acid's boiling point raises with the concentration as described in this figure to the right. An [[azeotrope]] forms at 98% w/w.<br />
<br />
At room temperature, sulfuric acid does not fume and has no smell. However, due to its hygroscopicity, closed bottles of conc. sulfuric acid may "smell" harsh, a consequence of inhaling the very dry air from the bottle. Solutions of sulfuric acid may have a weak acidic odor, especially at temperatures higher than room temperature, as a consequence of the solvent vapors carrying tiny amounts of H<sub>2</sub>SO<sub>4</sub> droplets in the air. Hot sulfuric acid is known to fume profusely and smells like a mix of burnt matches and pure pain (this is because of its partial decomposition when hot; the smells correspond to sulfur dioxide and trioxide respectively).<br />
<br />
==Sources and concentration==<br />
===OTC availability===<br />
Sulfuric acid is a commonly used chemical for lead-acid batteries and drain cleaning. Battery acid can often be found at an auto store or a department store and is approximately 33-35% sulfuric acid by weight. This is sufficient for most amateur chemists. If more concentrated sulfuric acid is desired, one can look in hardware stores for drain cleaner, which can be over 90% sulfuric acid by weight. For safety purposes, this concentration of sulfuric acid may have a dye in it. Other forms of sulfuric acid may be contaminated with various chemicals and will appear yellow, black, red.<br />
<br />
For some amateurs, it can be hard to find concentrated sulfuric acid, with acid drain cleaners being banned (as a result of [[wikipedia:Acid_throwing|acid throwing]] or illicit drug manufacture) or very contaminated in some countries.<br />
<br />
As of 2021, concentrated sulfuric acid over 15% is not available in the EU for private individuals, and all conc. sulfuric acid drain cleaners are restricted for professional use only. So far, it's unclear how this affects lead-acid batteries, which require acid in conc. higher than 15%. In certain other countries, 30-36% battery acid is OTC but drain cleaner acid is forbidden; if you happen to live in one of these countries, concentrating sulfuric acid is a must.<br />
<br />
===Concentration===<br />
The most well-tested method of concentrating sulfuric acid is described in a sub-article: [[Boiling the Bat]].<br />
<br />
* If you have technical grade sulfuric acid of concentrations from 80% to 94%, it can be converted to the pure compound by Zintl-Karyakin distillation. This process yields sulfuric acid of the highest quality and of concentration above the azeotrope. However, it is demanding in terms of glassware and very risky if performed at home. To perform this distillation, you need [[chromium trioxide]] or a dichromate salt (any will do, ''except ammonium'': [[ammonium dichromate]] will decompose on heating, and you'll have green murky acid contaminated with chromium (III) oxide and chromium sulfate) that will work as an azeotrope breaker. Add the H<sub>2</sub>SO<sub>4</sub>-Cr(VI) mixture to a round-bottom flask, pour the acid in and connect it to an air-cooled condenser. Put thermal insulation ([[asbestos]], rockwool) on the flask and start heating it. Discard the first few grams of the distillate, until its density reaches 1.84; collect every drop after that. This gives pure sulfuric acid with a concentration above 98%. Beware of any spillage of hexavalent chromium, it's a carcinogen! If such a spillage occurs, neutralize it with any reducing solution such as [[sodium thiosulfate]], [[ascorbic acid]] or [[glucose]].<br />
* Simple distillation of conc. drain cleaner sulfuric acid can work on some products, as hot sulfuric acid is oxidizing enough on its own that it will break down many organic contaminants.<ref>https://www.youtube.com/watch?v=4DUGRWjdNLI</ref> Similar to above, discard the first distillate fractions, and only keep the one with a density value of 1.84. This process however, may not work on all drain cleaners, so verify first.<br />
<br />
It is possible to further concentrate sulfuric acid by adding [[sulfur trioxide]], which reacts with the remaining water to form pure sulfuric acid. Sulfur trioxide can continue to be added to the solution to form [[oleum]], which fumes in air to form sulfuric acid droplets. When an equimolar concentration of sulfuric acid and sulfur trioxide is added, it forms [[pyrosulfuric acid]], which is a solid at room temperature. Sulfur trioxide can easily be obtained through the pyrolysis of certain salts, like anhydrous [[copper(II) sulfate]], [[iron(II) sulfate]], [[sodium pyrosulfate]] or [[potassium persulfate]].<br />
<br />
==Preparation==<br />
Sulfuric acid is industrially produced from sulfur, oxygen and water via the conventional contact process (DCDA), lead chamber process<ref>https://www.youtube.com/watch?v=7SDHeTcOXtI</ref> or the wet sulfuric acid process (WSA). The general way these processes work is by burning sulfur to obtain sulfur dioxide, which is oxidized to sulfur trioxide with the help of a catalyst, which in turn is dissolved in concentrated sulfuric acid, to form [[oleum]], which can be further concentrated into and eventually pyrosulfuric acid. The latter two products can be diluted using dil. sulfuric acid into conc. sulfuric acid. Diluted sulfuric acid is preferred instead of pure water, as the dilution is highly exothermic, while the reaction between sulfur trioxide with water is exothermic enough that the resulting sulfuric acid turns into a dense mist. The overall process can be written as:<br />
<br />
: S + O<sub>2</sub> → SO<sub>2</sub><br />
: SO<sub>2</sub> + ½ O<sub>2</sub> → SO<sub>3</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>S<sub>2</sub>O<sub>7</sub><br />
: H<sub>2</sub>S<sub>2</sub>O<sub>7</sub> + H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → 3 H<sub>2</sub>SO<sub>4</sub><br />
<br />
Each of the three main processes have their own advantages and disadvantages, but in general they work better at large scale, and for the average hobby chemist, while possible to reproduce them at smaller scale, it requires quite a lot of work to make the installation work properly. As such, working with volatile corrosive substances that melt your face off is quite an interesting project, if one were to try.<br />
<br />
There are many other routes to obtain sulfuric acid, most will produce diluted or mildly concentrated solutions, which can be concentrated to obtain more concentrated acid:<br />
<br />
*Absorbtion of sulfur dioxide in hydrogen peroxide: hydrogen peroxide will oxidize sulfur dioxide to sulfur trioxide, which reacts immediately with water to form sulfuric acid. Since this reaction is exothermic, an ice bath should be used. If an excess of SO<sub>2</sub> is used, warming the resulting solution to room temperature will cause some of the dissolved gas to boil off as the solution warms.<ref>https://www.youtube.com/watch?v=okvvD3-DF9U</ref><br />
<br />
: H<sub>2</sub>O<sub>2</sub> + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub><br />
<br />
While very easy to do, this reaction consumes hydrogen peroxide, and since H<sub>2</sub>O<sub>2</sub> is usually available OTC only as solutions from 3% up to 30%, the resulting sulfuric acid will be diluted, requiring further concentration.<ref>https://www.youtube.com/watch?v=mQMj5ier1lY</ref><br />
<br />
*Oxidation of SO<sub>2</sub> with conc. nitric acid: Similar to the reaction above with H<sub>2</sub>O<sub>2</sub>, conc. nitric acid can be used to oxidize sulfur dioxide directly to sulfuric acid, producing [[nitrogen dioxide]] as side product:<ref>https://www.youtube.com/watch?v=okvvD3-DF9U</ref><br />
<br />
: 2 HNO<sub>3</sub> + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub> + 2 NO<sub>2</sub><br />
<br />
The advantage of this reaction over the one with hydrogen peroxide, is that the nitrogen dioxide can be used to determine when the reaction is complete: when there is not more brown gas being produced, all the nitric acid has been consumed in the reaction. Main disadvantage of this route is that conc. nitric acid is a bit harder to acquire than sulfuric acid, and if one needs conc. sulfuric acid to obtain nitric acid, this route is not suitable. A modification of this reaction can be used, where the resulting nitrogen dioxide gets separated from the reaction, reacted with water to regenerate nitric acid, and then re-added in the reaction flask, to further oxidize the sulfur dioxide. Any nitric oxide produced from the side reaction between sulfur dioxide and nitrogen dioxide, can be reoxidized into nitrogen dioxide by injecting air in the mixture. <br />
<br />
*Ozone oxidation of sulfur dioxide: Ozone will oxidize sulfur dioxide into sulfur trioxide. This in turn reacts with water to form sulfuric acid. Ozone can be easily made by exposing oxygen to strong UV light, like that one produced by commercial ozone generators or low/high pressure mercury-vapor lamps. If atmospheric air is used, nitrogen dioxide may be produced as side product. This route is attractive since it uses cheap reagents, and while mercury UV lamps are somewhat difficult to properly operate, it's extremely easy to build a contraption where a continuous mixture of sulfur dioxide-oxygen is irradiated by strong UV light in a quartz tube, which produces sulfur trioxide directly. <br />
<br />
: 3 O<sub>2</sub> + hv → 2 O<sub>3</sub><br />
: SO<sub>2</sub> + O<sub>3</sub> → SO<sub>3</sub> + O<sub>2</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
<br />
*Electrolysis of aq. [[copper(II) sulfate]]: In a beaker, a concentrated solution of copper(II) sulfate is added. For cathode, a copper wire is added in the solution, at the bottom, and connected to the negative terminal of a power source, while for anode, a graphite electrode is added in the upper part of the solution, and connected to the positive terminal of the power source. During the process, the copper ions gets deposited on the copper electrode, while oxygen and hydrogen are produced at the carbon electrode. Overall, the reaction is as follows: <br />
<br />
: CuSO<sub>4</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub> + Cu + ½ O<sub>2</sub><br />
<br />
The resulting dil. solution of sulfuric acid is purified by filtering it, then concentrated by boiling it. This yields crude conc. H<sub>2</sub>SO<sub>4</sub>, which is distilled off to obtain the pure acid. The process is much easier than other electrochemical routes, as it's clean and relative quickly. Instead of graphite, other electrodes, like lead dioxide, titanium, platinum, or platinum on titanium can also be used.<ref>https://www.youtube.com/watch?v=5dUSF9Gl0xE</ref><ref>https://www.youtube.com/watch?v=ZRYtAquxffE</ref><br />
<br />
*Electrolysis of sulfate salt: This route involves electrolysis of a solution of a soluble sulfate salt, like [[magnesium sulfate]] or even [[ammonium sulfate]], using a diaphragm, which can either be either a classical ion-exchange diaphragm or a flower pot. <ref>https://www.youtube.com/watch?v=6BThiJpbBJQ</ref> The process yields dirty and diluted H<sub>2</sub>SO<sub>4</sub>, which requires purification and concentration.<ref>https://www.youtube.com/watch?v=b2wTha6Z-fA</ref><br />
<br />
*Pyrolysis of pyrosulfates: thermal decomposition of solid pyrosulfates yields sulfate and sulfur trioxide. The resulting sulfur trioxide is absorbed in crushed ice to form sulfuric acid. Further addition of sulfur trioxide yields conc. acid, and if SO<sub>3</sub> keeps getting added, it will convert into oleum, and eventually pyrosulfuric acid. The latter two products can be further diluted to concentrated sulfuric acid, by adding diluted sulfuric acid. For this process, [[sodium pyrosulfate]] is the best material, as it decomposes at a relative low temperature (460 °C) compared to other pyrosulfates, and the compound itself can be made by dehydrating [[sodium bisulfate]], which is readily and cheaply available:<br />
<br />
: 2 NaHSO<sub>4</sub> → Na<sub>2</sub>S<sub>2</sub>O<sub>7</sub> + H<sub>2</sub>O<br />
: Na<sub>2</sub>S<sub>2</sub>O<sub>7</sub> → Na<sub>2</sub>SO<sub>4</sub> + SO<sub>3</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
<br />
In theory, transition metal sulfates can also be used for this process, but since they decompose at higher temperatures, the resulting sulfur trioxide will partially decompose to sulfur dioxide and oxygen, which may lower the overall yield.<br />
<br />
*Copper chloride process: in an aqueous solution of [[copper(II) chloride]], sulfur dioxide is bubbled through. This reacts with the CuCl<sub>2</sub> from the aq. solution to form dil. sulfuric acid, HCl and CuCl:<br />
<br />
: 2 CuCl<sub>2</sub> + 2 H<sub>2</sub>O + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub> + 2 CuCl + 2 HCl<br />
<br />
CuCl precipitates out of the solution. By injecting air in the suspension, the CuCl gets reoxidized to CuCl<sub>2</sub>, which can be reused. Sulfur dioxide is reinjected in the solution, which restarts the reaction, then the process gets repeated, until no more SO<sub>2</sub> can absorb in the reaction solution. The yield of this process is not great, unless one uses kg amounts of reagents. Likewise, the oxidation of Cu(I) to Cu(II) using air is very slow, taking many hours, which limits the efficiency of the overall process.<br />
<br />
*Electrobromine process: involves the reaction of elemental sulfur with elemental [[bromine]], using a graphite anode and copper metal cathode. In a beaker, where elemental sulfur is added at the bottom, the two electrodes are introduces, with the graphite electrode resting on the sulfur bed, while the copper anode is only partially submerged in the electrolyte solution. A solution of 5 M [[hydrobromic acid]] is used as electrolyte. When the process is activated, the HBr gets oxidized to bromide ions, which in term convert to elemental bromine, that sink to the bottom, reacting with the sulfur bed to yield disulfur dibromide, which hydrolyzes in water to yield sulfuric acid and HBr, the latter rising back to the anode, where it gets converted back to bromine, and the process repeats. It's important to keep the Cu electrode as high as possible, to prevent the bromide ions from reacting with the elemental bromine, as this yields tribromide ionds, which do not react with the sulfur, and instead just get reduced back into bromide ions, wasting electricity. Eventually, after 1-2 days, the process is almost complete. The solution is filtered off, and the resulting HBr is distilled to be recycled, while the sulfuric acid is concentrated and purified by distillation. The yield of this process is not great, and as it uses bromine, which is highly corrosive and toxic. Likewise, the graphite electrodes get used up very quickly in the reaction. The sulfur bed may break apart during the process, and stirring may be required to break it apart and allow it to settle back. Stop the process and remove the electrodes, before stirring the suspension, and once the sulfur settles back, reintroduce the electrodes, and restart the process. Alternatively, one can a solid piece of sulfur instead of powder, as this shouldn't rise, though this may affect the speed of the reaction, as bulk sulfur reacts slower than powdered sulfur. A porous separating membrane, like a glass fiber cloth may be used to pin the sulfur bed down, while allowing the bromine to diffuse through it to reach the sulfur, though this hasn't been tested so far.<ref>https://www.youtube.com/watch?v=6ms6xbPhdVs</ref><br />
<br />
==Projects==<br />
* Preparation of metal sulfates<br />
* Preparation of nitro compounds through [[nitration]]<br />
* The dehydration of [[sucrose]] to produce elemental [[carbon]]<br />
* [[Esterification]]s that require a dehydrating agent, such as that of [[ethyl acetate]], [[methyl salicylate]], etc.<br />
* Making simple [[rayon]] fibers with [[Schweizer's reagent]] and [[cellulose]]<br />
* Producing other concentrated acids by the reaction of sulfuric acid with an anhydrous salt, such as in the production of fuming [[nitric acid]] and glacial [[acetic acid]]<br />
<br />
==Handling==<br />
===Safety===<br />
[[File:Corrosive.png|thumb|right|Corrosive]] While low concentration sulfuric acid is relatively safe to work with (under 40% w/w)), concentrated sulfuric acid (over 90% w/w) is extremely corrosive and dangerous. It does not only causes chemical burns, it also causes burns by dehydration of organic materials (like skin), destroying the molecules to form water with the -OH groups in them. Safety measures should be taken and all skin should be covered when working with concentrated sulfuric acid.<br />
<br />
When heating sulfuric acid, it is important to DO NOT OVERFILL THE FLASK. Concentrated sulfuric acid's volume increases by nearly 16% between 0 and 330°C, an overfilled flask will spill its content. Also, sulfuric acid, even diluted, tends to bump when it boils, accumulating heat to release a violent burst of steam from time to time. The use of boiling chips reduces this phenomenon, but there is no way to stop it completely. It is advised to take measures to prevent spills, an anti-splash adapter with ground glass joint being a very convenient option.<br />
<br />
Hot concentrated sulfuric acid may decompose to form sulfur dioxide and sulfur trioxide, which are toxic and corrosive, respectively. It fumes profusely when hot, the fumes consist of sulfuric acid droplets and a SOx mix. These fumes are very dangerous and a known lung carcinogen.<br />
<br />
When carrying glass bottles of sulfuric acid and you worry there's a risk you might break it, a good tip would be to carry it in a (plastic) bucket, partially filled with sand.<br />
<br />
===Storage===<br />
Sulfuric acid should be stored in closed bottles. While glass bottles, being inert, are good for storing concentrated sulfuric acid, concentrated (80-98%) sulfuric acid is often stored in PE (more specifically UDPE or UHDPE) bottles, as PE is not brittle, so in the event you drop the bottle on a hard surface, it will not shatter and splash conc. sulfuric all over the place. Unfortunately, PE bottles are sensitive to light and will degrade over the years if exposed to sunlight, so they must be stored in a dark place away from UV light, like a cupboard. Commercial PE bottles used for conc. sulfuric acids tend to be colored, which helps to limit degradation from strong light and oxygen. However, if you plan to store the acid for more that several years, it's recommended to use glass bottles.<br />
<br />
Long-term storage of concentrated sulfuric acid may lead to it absorbing water from air and becoming less concentrated. When this happens, the acid needs to be "re-freshened" by distilling unnecessary water off it. If the acid acquired a black or brown color during storage, it needs to be decarbonized: add several drops of concentrated H2O2 to it before distilling off water, the dark color will disappear during heating.<br />
<br />
===Disposal===<br />
Sulfuric acid can be neutralized with any base or carbonate, preferably [[calcium hydroxide]] or carbonate.<br />
<br />
Concentrated sulfuric acid, like any concentrated acid, should be first strongly dilute it in a large volume of water before neutralizing it with a base. Another method would be to add it in an acid-resistant container with a lid and slowly add solid calcium hydroxide/carbonate or sodium bicarbonate chunks and close the lid to limit splashing. Wait until it stopped fizzing then keep adding until it no longer reacts. Be careful, as the thicker the solution becomes, the stronger the foaming gets.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6911 Sulfuric Acid Production: Revisited]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2824 H2SO4 by the Lead Chamber Process - success]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=64535 I will now be building and testing my new Batparatus!]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=3722 cleaning sulfuric acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13313 Sulfuric Acid at Home]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19117 Concentrating dilute sulphuric acid(battery acid) without distillation]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=91332 Sulfuric acid from gypsum using diaphragm cell]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14857 Sulfuric acid purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14570 sulfuric acid turned black]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=61920 Distilling Sulfuric Acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=65331 Sulfuric acid in NZ]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14291 Should I get rid of my H2SO4?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13726 sulfuric acid accident]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=62863 Sulfuric acid storage]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13964 HDPE as a storage for Sulfuric Acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13148 Safely Storing H2SO4 (35%)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6217 Storage for Sulfuric Acid (H2SO4)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=25679 Sulfuric Acid and LDPE issue]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Oxoacids]]<br />
[[Category:Sulfur oxoacids]]<br />
[[Category:Sulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Materials unstable in basic solution]]<br />
[[Category:Things that can kill you very quickly]]<br />
[[Category:Hygroscopic compounds]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Essential reagents]]<br />
[[Category:DEA List II chemicals]]<br />
[[Category:Catalysts]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sulfuric_acid&diff=14959Sulfuric acid2023-03-27T21:00:02Z<p>Ave369: /* OTC availability */</p>
<hr />
<div>{{Chembox<br />
| Name = Sulfuric acid<br />
| Reference =<br />
| IUPACName = Sulfuric acid<br />
| PIN = Sulfuric acid<br />
| SystematicName = Sulfuric acid<br />
| OtherNames = Battery acid<br>Dihydrogen sulfate<br>Oil of vitriol<br>Spirit of vitriol<br>Sulphuric acid<br />
<!-- Images --><br />
| ImageFile = Smw1.png<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Structure of sulfuric acid<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = OS(=O)(=O)O<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless oily liquid<br />
| BoilingPt = <br />
| BoilingPtC = 337<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (above 300 °C slowly decomposes)<br />
| Density = 1.84 g/cm<sup>3</sup><br />
| Formula = H<sub>2</sub>SO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 98.079 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 10<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless (air above it may feel dry due to its strong hygroscopicity)<br />
| pKa = −3;1.99<br />
| pKb = <br />
| Solubility = Miscible<br />
| SolubleOther = Reacts with [[amine]]s<br>Miscible with [[alcohol]]s<br>Immiscible with hydrocarbons<br />
| Solvent = <br />
| VaporPressure = 0.001 mmHg (20 °C)<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = −814 kJ·mol<sup>−1</sup><br />
| Entropy = 157 J·mol<sup>−1</sup>·K<sup>−1</sup><br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.fishersci.com/msdsproxy%3FproductName%3DA300700LB%26productDescription%3DSULFURIC%2BAC%2BACS%2B700LB%26catNo%3DA300-700LB%26vendorId%3DVN00033897%26storeId%3D10652 FisherSci]<br />
| FlashPt = Non-flammable<br />
| LD50 = 2.140 mg/kg (rat, oral)<br />
| LC50 = 50 mg/m<sup>3</sup> (guinea pig, 8 hr)<br>510 mg/m<sup>3</sup> (rat, 2 hr)<br>320 mg/m<sup>3</sup> (mouse, 2 hr)<br>18 mg/m<sup>3</sup> (guinea pig)<br />
| MainHazards = Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Sulfurous acid]]<br>[[Sulfur trioxide]]<br />
}}<br />
}}<br />
'''Sulfuric acid''' (alternative spelling '''sulphuric acid'''), represented by the molecular formula '''H<sub>2</sub>SO<sub>4</sub>''', is one of the most important [[acid]]s in chemistry and the most important chemical to industries in the world. It is the strongest easily available acid, with a [[Measures of acidity|pK<sub>a</sub>]] of -3.<br />
<br />
==Properties==<br />
===Chemical properties===<br />
Sulfuric acid is a diprotic acid, and thus it is able to give away two protons (H<sup>+</sup>). It first dissociates to form [[hydronium]] and hydrogen sulfate/bisulfate ions, with a pK<sub>a</sub> of -3, indicative of a strong acid:<br />
<br />
: H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → H<sub>3</sub>O + HSO<sub>4</sub><sup>−</sup><br />
<br />
The second dissociation forms sulfate and another hydronium ion from a hydrogen sulfate ion. It has a pKa of 1.99, indicative of a mid-strength acid, and occurs like this: <br />
<br />
: HSO<sub>4</sub><sup>−</sup> + H<sub>2</sub>O ⇌ H<sub>3</sub>O<sup>+</sup> + SO<sub>4</sub><sup>2-</sup><br />
<br />
Concentrated sulfuric acid also has a strong oxidizing effect, converting nonmetals such as [[carbon]] and [[sulfur]] to [[carbon dioxide]] and [[sulfur dioxide]], respectively, reducing sulfuric acid into sulfur dioxide and water in the process.<br />
<br />
: 2 H<sub>2</sub>SO<sub>4</sub> + C → CO<sub>2</sub> + SO<sub>2</sub> + H<sub>2</sub>O + H<sub>2</sub>SO<sub>4</sub><br />
<br />
: 2 H<sub>2</sub>SO<sub>4</sub> + S → 2 SO<sub>2</sub> + H<sub>2</sub>O + H<sub>2</sub>SO<sub>4</sub><br />
<br />
This property is useful for producing large amounts of sulfur dioxide for use as a reducing agent if water is continually removed. Heat accelerates this process.<br />
<br />
Sulfuric acid is sufficiently strong enough to protonate [[nitric acid]], forming the nitronium ion, which can be used in a nitration mixture to make [[alkyl nitrate]]s.<br />
<br />
In organic chemistry, sulfuric acid is the most practical acid in most cases where a source of H<sub>3</sub>O<sup>+</sup> ions are needed as it introduces the least amount of water. Organic compounds are often easily attacked by the nucleophiles left behind by the dissociation of acids such as HCl which leaves Cl<sup>-</sup> ions behind which can easily attack many organic compounds. However, the [[sulfate]] ions left behind by the dissociation of sulfuric acid are far less reactive than the ions left behind by most acids, it allows to protonate the reaction mixture without causing undesired side reactions in most cases.<br />
<br />
When concentrated, it is strongly [[hygroscopy|hygroscopic]] and has strong dehydrating properties. It can break down most organic molecules containing OH<sup>-</sup> groups to use them to form water, leaving only the carbon behind. This property is exploited in the famous [http://youtu.be/w6lfq7BOCik "black snake" demonstration], where sulfuric acid dehydrates [[sucrose]] (table sugar), forming water with the hydrogen and oxygen atoms and leaving amorphous carbon behind.<br />
<br />
===Physical properties===<br />
[[File:H2so4boil.jpg|thumb|left|342px|Boiling point of H2SO4 VS concentration]]<br />
<br />
Sulfuric acid is an oily liquid at room temperature. It is colorless but often has a very light yellow color when slightly contaminated with iron or carbon from organic matter like dust. Even very small amounts of dissolved organic matter can change the color of concentrated sulfuric acid to pale yellow or pink, red, brown, and even black. It is commonly sold diluted at around 35% w/w with water as car battery acid and concentrated between 95% and 98% w/w as drain cleaner.<br />
<br />
Sulfuric acid's boiling point raises with the concentration as described in this figure to the right. An [[azeotrope]] forms at 98% w/w.<br />
<br />
At room temperature, sulfuric acid does not fume and has no smell. However, due to its hygroscopicity, closed bottles of conc. sulfuric acid may "smell" harsh, a consequence of inhaling the very dry air from the bottle. Solutions of sulfuric acid may have a weak acidic odor, especially at temperatures higher than room temperature, as a consequence of the solvent vapors carrying tiny amounts of H<sub>2</sub>SO<sub>4</sub> droplets in the air. Hot sulfuric acid is known to fume profusely and smells like a mix of burnt matches and pure pain (this is because of its partial decomposition when hot; the smells correspond to sulfur dioxide and trioxide respectively).<br />
<br />
==Sources and concentration==<br />
===OTC availability===<br />
Sulfuric acid is a commonly used chemical for lead-acid batteries and drain cleaning. Battery acid can often be found at an auto store or a department store and is approximately 33-35% sulfuric acid by weight. This is sufficient for most amateur chemists. If more concentrated sulfuric acid is desired, one can look in hardware stores for drain cleaner, which can be over 90% sulfuric acid by weight. For safety purposes, this concentration of sulfuric acid may have a dye in it. Other forms of sulfuric acid may be contaminated with various chemicals and will appear yellow, black, red.<br />
<br />
For some amateurs, it can be hard to find concentrated sulfuric acid, with acid drain cleaners being banned (as a result of [[wikipedia:Acid_throwing|acid throwing]] or illicit drug manufacture) or very contaminated in some countries.<br />
<br />
As of 2021, concentrated sulfuric acid over 15% is not available in the EU for private individuals, and all conc. sulfuric acid drain cleaners are restricted for professional use only. So far, it's unclear how this affects lead-acid batteries, which require acid in conc. higher than 15%. In certain other countries, 30-36% battery acid is OTC but drain cleaner acid is forbidden; if you happen to live in one of these countries, concentrating sulfuric acid is a must.<br />
<br />
===Concentration===<br />
The most well-tested method of concentrating sulfuric acid is described in a sub-article: [[Boiling the Bat]].<br />
<br />
* If you have technical grade sulfuric acid of concentrations from 80% to 94%, it can be converted to the pure compound by Zintl-Karyakin distillation. This process yields sulfuric acid of the highest quality and of concentration above the azeotrope. However, it is demanding in terms of glassware and very risky if performed at home. To perform this distillation, you need [[chromium trioxide]] or a dichromate salt (any will do, ''except ammonium'': [[ammonium dichromate]] will decompose on heating, and you'll have green murky acid contaminated with chromium (III) oxide and chromium sulfate) that will work as an azeotrope breaker. Add the H<sub>2</sub>SO<sub>4</sub>-Cr(VI) mixture to a round-bottom flask, pour the acid in and connect it to an air-cooled condenser. Put thermal insulation ([[asbestos]], rockwool) on the flask and start heating it. Discard the first few grams of the distillate, until its density reaches 1.84; collect every drop after that. This gives pure sulfuric acid with a concentration above 98%. Beware of any spillage of hexavalent chromium, it's a carcinogen! If such a spillage occurs, neutralize it with any reducing solution such as [[sodium thiosulfate]], [[ascorbic acid]] or [[glucose]].<br />
* Simple distillation of conc. drain cleaner sulfuric acid can work on some products, as hot sulfuric acid is oxidizing enough on its own that it will break down many organic contaminants.<ref>https://www.youtube.com/watch?v=4DUGRWjdNLI</ref> Similar to above, discard the first distillate fractions, and only keep the one with a density value of 1.84. This process however, may not work on all drain cleaners, so verify first.<br />
<br />
It is possible to further concentrate sulfuric acid by adding [[sulfur trioxide]], which reacts with the remaining water to form pure sulfuric acid. Sulfur trioxide can continue to be added to the solution to form [[oleum]], which fumes in air to form sulfuric acid droplets. When an equimolar concentration of sulfuric acid and sulfur trioxide is added, it forms [[pyrosulfuric acid]], which is a solid at room temperature. Sulfur trioxide can easily be obtained through the pyrolysis of certain salts, like anhydrous [[copper(II) sulfate]], [[iron(II) sulfate]], [[sodium pyrosulfate]] or [[potassium persulfate]].<br />
<br />
==Preparation==<br />
Sulfuric acid is industrially produced from sulfur, oxygen and water via the conventional contact process (DCDA), lead chamber process<ref>https://www.youtube.com/watch?v=7SDHeTcOXtI</ref> or the wet sulfuric acid process (WSA). The general way these processes work is by burning sulfur to obtain sulfur dioxide, which is oxidized to sulfur trioxide with the help of a catalyst, which in turn is dissolved in concentrated sulfuric acid, to form [[oleum]], which can be further concentrated into and eventually pyrosulfuric acid. The latter two products can be diluted using dil. sulfuric acid into conc. sulfuric acid. Diluted sulfuric acid is preferred instead of pure water, as the dilution is highly exothermic, while the reaction between sulfur trioxide with water is exothermic enough that the resulting sulfuric acid turns into a dense mist. The overall process can be written as:<br />
<br />
: S + O<sub>2</sub> → SO<sub>2</sub><br />
: SO<sub>2</sub> + ½ O<sub>2</sub> → SO<sub>3</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>S<sub>2</sub>O<sub>7</sub><br />
: H<sub>2</sub>S<sub>2</sub>O<sub>7</sub> + H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → 3 H<sub>2</sub>SO<sub>4</sub><br />
<br />
Each of the three main processes have their own advantages and disadvantages, but in general they work better at large scale, and for the average hobby chemist, while possible to reproduce them at smaller scale, it requires quite a lot of work to make the installation work properly. As such, working with volatile corrosive substances that melt your face off is quite an interesting project, if one were to try.<br />
<br />
There are many other routes to obtain sulfuric acid, most will produce diluted or mildly concentrated solutions, which can be concentrated to obtain more concentrated acid:<br />
<br />
*Absorbtion of sulfur dioxide in hydrogen peroxide: hydrogen peroxide will oxidize sulfur dioxide to sulfur trioxide, which reacts immediately with water to form sulfuric acid. Since this reaction is exothermic, an ice bath should be used. If an excess of SO<sub>2</sub> is used, warming the resulting solution to room temperature will cause some of the dissolved gas to boil off as the solution warms.<ref>https://www.youtube.com/watch?v=okvvD3-DF9U</ref><br />
<br />
: H<sub>2</sub>O<sub>2</sub> + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub><br />
<br />
While very easy to do, this reaction consumes hydrogen peroxide, and since H<sub>2</sub>O<sub>2</sub> is usually available OTC only as solutions from 3% up to 30%, the resulting sulfuric acid will be diluted, requiring further concentration.<ref>https://www.youtube.com/watch?v=mQMj5ier1lY</ref><br />
<br />
*Oxidation of SO<sub>2</sub> with conc. nitric acid: Similar to the reaction above with H<sub>2</sub>O<sub>2</sub>, conc. nitric acid can be used to oxidize sulfur dioxide directly to sulfuric acid, producing [[nitrogen dioxide]] as side product:<ref>https://www.youtube.com/watch?v=okvvD3-DF9U</ref><br />
<br />
: 2 HNO<sub>3</sub> + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub> + 2 NO<sub>2</sub><br />
<br />
The advantage of this reaction over the one with hydrogen peroxide, is that the nitrogen dioxide can be used to determine when the reaction is complete: when there is not more brown gas being produced, all the nitric acid has been consumed in the reaction. Main disadvantage of this route is that conc. nitric acid is a bit harder to acquire than sulfuric acid, and if one needs conc. sulfuric acid to obtain nitric acid, this route is not suitable. A modification of this reaction can be used, where the resulting nitrogen dioxide gets separated from the reaction, reacted with water to regenerate nitric acid, and then re-added in the reaction flask, to further oxidize the sulfur dioxide. Any nitric oxide produced from the side reaction between sulfur dioxide and nitrogen dioxide, can be reoxidized into nitrogen dioxide by injecting air in the mixture. <br />
<br />
*Ozone oxidation of sulfur dioxide: Ozone will oxidize sulfur dioxide into sulfur trioxide. This in turn reacts with water to form sulfuric acid. Ozone can be easily made by exposing oxygen to strong UV light, like that one produced by commercial ozone generators or low/high pressure mercury-vapor lamps. If atmospheric air is used, nitrogen dioxide may be produced as side product. This route is attractive since it uses cheap reagents, and while mercury UV lamps are somewhat difficult to properly operate, it's extremely easy to build a contraption where a continuous mixture of sulfur dioxide-oxygen is irradiated by strong UV light in a quartz tube, which produces sulfur trioxide directly. <br />
<br />
: 3 O<sub>2</sub> + hv → 2 O<sub>3</sub><br />
: SO<sub>2</sub> + O<sub>3</sub> → SO<sub>3</sub> + O<sub>2</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
<br />
*Electrolysis of aq. [[copper(II) sulfate]]: In a beaker, a concentrated solution of copper(II) sulfate is added. For cathode, a copper wire is added in the solution, at the bottom, and connected to the negative terminal of a power source, while for anode, a graphite electrode is added in the upper part of the solution, and connected to the positive terminal of the power source. During the process, the copper ions gets deposited on the copper electrode, while oxygen and hydrogen are produced at the carbon electrode. Overall, the reaction is as follows: <br />
<br />
: CuSO<sub>4</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub> + Cu + ½ O<sub>2</sub><br />
<br />
The resulting dil. solution of sulfuric acid is purified by filtering it, then concentrated by boiling it. This yields crude conc. H<sub>2</sub>SO<sub>4</sub>, which is distilled off to obtain the pure acid. The process is much easier than other electrochemical routes, as it's clean and relative quickly. Instead of graphite, other electrodes, like lead dioxide, titanium, platinum, or platinum on titanium can also be used.<ref>https://www.youtube.com/watch?v=5dUSF9Gl0xE</ref><ref>https://www.youtube.com/watch?v=ZRYtAquxffE</ref><br />
<br />
*Electrolysis of sulfate salt: This route involves electrolysis of a solution of a soluble sulfate salt, like [[magnesium sulfate]] or even [[ammonium sulfate]], using a diaphragm, which can either be either a classical ion-exchange diaphragm or a flower pot. <ref>https://www.youtube.com/watch?v=6BThiJpbBJQ</ref> The process yields dirty and diluted H<sub>2</sub>SO<sub>4</sub>, which requires purification and concentration.<ref>https://www.youtube.com/watch?v=b2wTha6Z-fA</ref><br />
<br />
*Pyrolysis of pyrosulfates: thermal decomposition of solid pyrosulfates yields sulfate and sulfur trioxide. The resulting sulfur trioxide is absorbed in crushed ice to form sulfuric acid. Further addition of sulfur trioxide yields conc. acid, and if SO<sub>3</sub> keeps getting added, it will convert into oleum, and eventually pyrosulfuric acid. The latter two products can be further diluted to concentrated sulfuric acid, by adding diluted sulfuric acid. For this process, [[sodium pyrosulfate]] is the best material, as it decomposes at a relative low temperature (460 °C) compared to other pyrosulfates, and the compound itself can be made by dehydrating [[sodium bisulfate]], which is readily and cheaply available:<br />
<br />
: 2 NaHSO<sub>4</sub> → Na<sub>2</sub>S<sub>2</sub>O<sub>7</sub> + H<sub>2</sub>O<br />
: Na<sub>2</sub>S<sub>2</sub>O<sub>7</sub> → Na<sub>2</sub>SO<sub>4</sub> + SO<sub>3</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
<br />
In theory, transition metal sulfates can also be used for this process, but since they decompose at higher temperatures, the resulting sulfur trioxide will partially decompose to sulfur dioxide and oxygen, which may lower the overall yield.<br />
<br />
*Copper chloride process: in an aqueous solution of [[copper(II) chloride]], sulfur dioxide is bubbled through. This reacts with the CuCl<sub>2</sub> from the aq. solution to form dil. sulfuric acid, HCl and CuCl:<br />
<br />
: 2 CuCl<sub>2</sub> + 2 H<sub>2</sub>O + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub> + 2 CuCl + 2 HCl<br />
<br />
CuCl precipitates out of the solution. By injecting air in the suspension, the CuCl gets reoxidized to CuCl<sub>2</sub>, which can be reused. Sulfur dioxide is reinjected in the solution, which restarts the reaction, then the process gets repeated, until no more SO<sub>2</sub> can absorb in the reaction solution. The yield of this process is not great, unless one uses kg amounts of reagents. Likewise, the oxidation of Cu(I) to Cu(II) using air is very slow, taking many hours, which limits the efficiency of the overall process.<br />
<br />
*Electrobromine process: involves the reaction of elemental sulfur with elemental [[bromine]], using a graphite anode and copper metal cathode. In a beaker, where elemental sulfur is added at the bottom, the two electrodes are introduces, with the graphite electrode resting on the sulfur bed, while the copper anode is only partially submerged in the electrolyte solution. A solution of 5 M [[hydrobromic acid]] is used as electrolyte. When the process is activated, the HBr gets oxidized to bromide ions, which in term convert to elemental bromine, that sink to the bottom, reacting with the sulfur bed to yield disulfur dibromide, which hydrolyzes in water to yield sulfuric acid and HBr, the latter rising back to the anode, where it gets converted back to bromine, and the process repeats. It's important to keep the Cu electrode as high as possible, to prevent the bromide ions from reacting with the elemental bromine, as this yields tribromide ionds, which do not react with the sulfur, and instead just get reduced back into bromide ions, wasting electricity. Eventually, after 1-2 days, the process is almost complete. The solution is filtered off, and the resulting HBr is distilled to be recycled, while the sulfuric acid is concentrated and purified by distillation. The yield of this process is not great, and as it uses bromine, which is highly corrosive and toxic. Likewise, the graphite electrodes get used up very quickly in the reaction. The sulfur bed may break apart during the process, and stirring may be required to break it apart and allow it to settle back. Stop the process and remove the electrodes, before stirring the suspension, and once the sulfur settles back, reintroduce the electrodes, and restart the process. Alternatively, one can a solid piece of sulfur instead of powder, as this shouldn't rise, though this may affect the speed of the reaction, as bulk sulfur reacts slower than powdered sulfur. A porous separating membrane, like a glass fiber cloth may be used to pin the sulfur bed down, while allowing the bromine to diffuse through it to reach the sulfur, though this hasn't been tested so far.<ref>https://www.youtube.com/watch?v=6ms6xbPhdVs</ref><br />
<br />
==Projects==<br />
* Preparation of metal sulfates<br />
* Preparation of nitro compounds through [[nitration]]<br />
* The dehydration of [[sucrose]] to produce elemental [[carbon]]<br />
* [[Esterification]]s that require a dehydrating agent, such as that of [[ethyl acetate]], [[methyl salicylate]], etc.<br />
* Making simple [[rayon]] fibers with [[Schweizer's reagent]] and [[cellulose]]<br />
* Producing other concentrated acids by the reaction of sulfuric acid with an anhydrous salt, such as in the production of fuming [[nitric acid]] and glacial [[acetic acid]]<br />
<br />
==Handling==<br />
===Safety===<br />
[[File:Corrosive.png|thumb|right|Corrosive]] While low concentration sulfuric acid is relatively safe to work with (under 40% w/w)), concentrated sulfuric acid (over 90% w/w) is extremely corrosive and dangerous. It does not only causes chemical burns, it also causes burns by dehydration of organic materials (like skin), destroying the molecules to form water with the -OH groups in them. Safety measures should be taken and all skin should be covered when working with concentrated sulfuric acid.<br />
<br />
When heating sulfuric acid, it is important to DO NOT OVERFILL THE FLASK. Concentrated sulfuric acid's volume increases by nearly 16% between 0 and 330°C, an overfilled flask will spill its content. Also, sulfuric acid, even diluted, tends to bump when it boils, accumulating heat to release a violent burst of steam from time to time. The use of boiling chips reduces this phenomenon, but there is no way to stop it completely. It is advised to take measures to prevent spills, an anti-splash adapter with ground glass joint being a very convenient option.<br />
<br />
Hot concentrated sulfuric acid may decompose to form sulfur dioxide and sulfur trioxide, which are toxic and corrosive, respectively. It fumes profusely when hot, the fumes consist of sulfuric acid droplets and a SOx mix. These fumes are very dangerous and a known lung carcinogen.<br />
<br />
When carrying glass bottles of sulfuric acid and you worry there's a risk you might break it, a good tip would be to carry it in a (plastic) bucket, partially filled with sand.<br />
<br />
===Storage===<br />
Sulfuric acid should be stored in closed bottles. While glass bottles, being inert, are good for storing concentrated sulfuric acid, concentrated (80-98%) sulfuric acid is often stored in PE (more specifically UDPE or UHDPE) bottles, as PE is not brittle, so in the event you drop the bottle on a hard surface, it will not shatter and splash conc. sulfuric all over the place. Unfortunately, PE bottles are sensitive to light and will degrade over the years if exposed to sunlight, so they must be stored in a dark place away from UV light, like a cupboard. Commercial PE bottles used for conc. sulfuric acids tend to be colored, which helps to limit degradation from strong light and oxygen. However, if you plan to store the acid for more that several years, it's recommended to use glass bottles.<br />
<br />
Long-term storage of concentrated sulfuric acid may lead to it absorbing water from air and becoming less concentrated. When this happens, the acid needs to be "re-freshened" by distilling unnecessary water off it.<br />
<br />
===Disposal===<br />
Sulfuric acid can be neutralized with any base or carbonate, preferably [[calcium hydroxide]] or carbonate.<br />
<br />
Concentrated sulfuric acid, like any concentrated acid, should be first strongly dilute it in a large volume of water before neutralizing it with a base. Another method would be to add it in an acid-resistant container with a lid and slowly add solid calcium hydroxide/carbonate or sodium bicarbonate chunks and close the lid to limit splashing. Wait until it stopped fizzing then keep adding until it no longer reacts. Be careful, as the thicker the solution becomes, the stronger the foaming gets.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6911 Sulfuric Acid Production: Revisited]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2824 H2SO4 by the Lead Chamber Process - success]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=64535 I will now be building and testing my new Batparatus!]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=3722 cleaning sulfuric acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13313 Sulfuric Acid at Home]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19117 Concentrating dilute sulphuric acid(battery acid) without distillation]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=91332 Sulfuric acid from gypsum using diaphragm cell]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14857 Sulfuric acid purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14570 sulfuric acid turned black]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=61920 Distilling Sulfuric Acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=65331 Sulfuric acid in NZ]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14291 Should I get rid of my H2SO4?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13726 sulfuric acid accident]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=62863 Sulfuric acid storage]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13964 HDPE as a storage for Sulfuric Acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13148 Safely Storing H2SO4 (35%)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6217 Storage for Sulfuric Acid (H2SO4)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=25679 Sulfuric Acid and LDPE issue]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Oxoacids]]<br />
[[Category:Sulfur oxoacids]]<br />
[[Category:Sulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Materials unstable in basic solution]]<br />
[[Category:Things that can kill you very quickly]]<br />
[[Category:Hygroscopic compounds]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Essential reagents]]<br />
[[Category:DEA List II chemicals]]<br />
[[Category:Catalysts]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sulfuric_acid&diff=14958Sulfuric acid2023-03-27T20:32:06Z<p>Ave369: /* Storage */</p>
<hr />
<div>{{Chembox<br />
| Name = Sulfuric acid<br />
| Reference =<br />
| IUPACName = Sulfuric acid<br />
| PIN = Sulfuric acid<br />
| SystematicName = Sulfuric acid<br />
| OtherNames = Battery acid<br>Dihydrogen sulfate<br>Oil of vitriol<br>Spirit of vitriol<br>Sulphuric acid<br />
<!-- Images --><br />
| ImageFile = Smw1.png<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Structure of sulfuric acid<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = OS(=O)(=O)O<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless oily liquid<br />
| BoilingPt = <br />
| BoilingPtC = 337<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (above 300 °C slowly decomposes)<br />
| Density = 1.84 g/cm<sup>3</sup><br />
| Formula = H<sub>2</sub>SO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 98.079 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 10<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless (air above it may feel dry due to its strong hygroscopicity)<br />
| pKa = −3;1.99<br />
| pKb = <br />
| Solubility = Miscible<br />
| SolubleOther = Reacts with [[amine]]s<br>Miscible with [[alcohol]]s<br>Immiscible with hydrocarbons<br />
| Solvent = <br />
| VaporPressure = 0.001 mmHg (20 °C)<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = −814 kJ·mol<sup>−1</sup><br />
| Entropy = 157 J·mol<sup>−1</sup>·K<sup>−1</sup><br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.fishersci.com/msdsproxy%3FproductName%3DA300700LB%26productDescription%3DSULFURIC%2BAC%2BACS%2B700LB%26catNo%3DA300-700LB%26vendorId%3DVN00033897%26storeId%3D10652 FisherSci]<br />
| FlashPt = Non-flammable<br />
| LD50 = 2.140 mg/kg (rat, oral)<br />
| LC50 = 50 mg/m<sup>3</sup> (guinea pig, 8 hr)<br>510 mg/m<sup>3</sup> (rat, 2 hr)<br>320 mg/m<sup>3</sup> (mouse, 2 hr)<br>18 mg/m<sup>3</sup> (guinea pig)<br />
| MainHazards = Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Sulfurous acid]]<br>[[Sulfur trioxide]]<br />
}}<br />
}}<br />
'''Sulfuric acid''' (alternative spelling '''sulphuric acid'''), represented by the molecular formula '''H<sub>2</sub>SO<sub>4</sub>''', is one of the most important [[acid]]s in chemistry and the most important chemical to industries in the world. It is the strongest easily available acid, with a [[Measures of acidity|pK<sub>a</sub>]] of -3.<br />
<br />
==Properties==<br />
===Chemical properties===<br />
Sulfuric acid is a diprotic acid, and thus it is able to give away two protons (H<sup>+</sup>). It first dissociates to form [[hydronium]] and hydrogen sulfate/bisulfate ions, with a pK<sub>a</sub> of -3, indicative of a strong acid:<br />
<br />
: H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → H<sub>3</sub>O + HSO<sub>4</sub><sup>−</sup><br />
<br />
The second dissociation forms sulfate and another hydronium ion from a hydrogen sulfate ion. It has a pKa of 1.99, indicative of a mid-strength acid, and occurs like this: <br />
<br />
: HSO<sub>4</sub><sup>−</sup> + H<sub>2</sub>O ⇌ H<sub>3</sub>O<sup>+</sup> + SO<sub>4</sub><sup>2-</sup><br />
<br />
Concentrated sulfuric acid also has a strong oxidizing effect, converting nonmetals such as [[carbon]] and [[sulfur]] to [[carbon dioxide]] and [[sulfur dioxide]], respectively, reducing sulfuric acid into sulfur dioxide and water in the process.<br />
<br />
: 2 H<sub>2</sub>SO<sub>4</sub> + C → CO<sub>2</sub> + SO<sub>2</sub> + H<sub>2</sub>O + H<sub>2</sub>SO<sub>4</sub><br />
<br />
: 2 H<sub>2</sub>SO<sub>4</sub> + S → 2 SO<sub>2</sub> + H<sub>2</sub>O + H<sub>2</sub>SO<sub>4</sub><br />
<br />
This property is useful for producing large amounts of sulfur dioxide for use as a reducing agent if water is continually removed. Heat accelerates this process.<br />
<br />
Sulfuric acid is sufficiently strong enough to protonate [[nitric acid]], forming the nitronium ion, which can be used in a nitration mixture to make [[alkyl nitrate]]s.<br />
<br />
In organic chemistry, sulfuric acid is the most practical acid in most cases where a source of H<sub>3</sub>O<sup>+</sup> ions are needed as it introduces the least amount of water. Organic compounds are often easily attacked by the nucleophiles left behind by the dissociation of acids such as HCl which leaves Cl<sup>-</sup> ions behind which can easily attack many organic compounds. However, the [[sulfate]] ions left behind by the dissociation of sulfuric acid are far less reactive than the ions left behind by most acids, it allows to protonate the reaction mixture without causing undesired side reactions in most cases.<br />
<br />
When concentrated, it is strongly [[hygroscopy|hygroscopic]] and has strong dehydrating properties. It can break down most organic molecules containing OH<sup>-</sup> groups to use them to form water, leaving only the carbon behind. This property is exploited in the famous [http://youtu.be/w6lfq7BOCik "black snake" demonstration], where sulfuric acid dehydrates [[sucrose]] (table sugar), forming water with the hydrogen and oxygen atoms and leaving amorphous carbon behind.<br />
<br />
===Physical properties===<br />
[[File:H2so4boil.jpg|thumb|left|342px|Boiling point of H2SO4 VS concentration]]<br />
<br />
Sulfuric acid is an oily liquid at room temperature. It is colorless but often has a very light yellow color when slightly contaminated with iron or carbon from organic matter like dust. Even very small amounts of dissolved organic matter can change the color of concentrated sulfuric acid to pale yellow or pink, red, brown, and even black. It is commonly sold diluted at around 35% w/w with water as car battery acid and concentrated between 95% and 98% w/w as drain cleaner.<br />
<br />
Sulfuric acid's boiling point raises with the concentration as described in this figure to the right. An [[azeotrope]] forms at 98% w/w.<br />
<br />
At room temperature, sulfuric acid does not fume and has no smell. However, due to its hygroscopicity, closed bottles of conc. sulfuric acid may "smell" harsh, a consequence of inhaling the very dry air from the bottle. Solutions of sulfuric acid may have a weak acidic odor, especially at temperatures higher than room temperature, as a consequence of the solvent vapors carrying tiny amounts of H<sub>2</sub>SO<sub>4</sub> droplets in the air. Hot sulfuric acid is known to fume profusely and smells like a mix of burnt matches and pure pain (this is because of its partial decomposition when hot; the smells correspond to sulfur dioxide and trioxide respectively).<br />
<br />
==Sources and concentration==<br />
===OTC availability===<br />
Sulfuric acid is a commonly used chemical for lead-acid batteries and drain cleaning. Battery acid can often be found at an auto store or a department store and is approximately 33-35% sulfuric acid by weight. This is sufficient for most amateur chemists. If more concentrated sulfuric acid is desired, one can look in hardware stores for drain cleaner, which can be over 90% sulfuric acid by weight. For safety purposes, this concentration of sulfuric acid may have a dye in it. Other forms of sulfuric acid may be contaminated with various chemicals and will appear yellow, black, red.<br />
<br />
For some amateurs, it can be hard to find concentrated sulfuric acid, with acid drain cleaners being banned (as a result of [[wikipedia:Acid_throwing|acid throwing]] or illicit drug manufacture) or very contaminated in some countries.<br />
<br />
As of 2021, concentrated sulfuric acid over 15% is not available in the EU for private individuals, and all conc. sulfuric acid drain cleaners are restricted for professional use only. So far, it's unclear how this affects lead-acid batteries, which require acid in conc. higher than 15%.<br />
<br />
===Concentration===<br />
The most well-tested method of concentrating sulfuric acid is described in a sub-article: [[Boiling the Bat]].<br />
<br />
* If you have technical grade sulfuric acid of concentrations from 80% to 94%, it can be converted to the pure compound by Zintl-Karyakin distillation. This process yields sulfuric acid of the highest quality and of concentration above the azeotrope. However, it is demanding in terms of glassware and very risky if performed at home. To perform this distillation, you need [[chromium trioxide]] or a dichromate salt (any will do, ''except ammonium'': [[ammonium dichromate]] will decompose on heating, and you'll have green murky acid contaminated with chromium (III) oxide and chromium sulfate) that will work as an azeotrope breaker. Add the H<sub>2</sub>SO<sub>4</sub>-Cr(VI) mixture to a round-bottom flask, pour the acid in and connect it to an air-cooled condenser. Put thermal insulation ([[asbestos]], rockwool) on the flask and start heating it. Discard the first few grams of the distillate, until its density reaches 1.84; collect every drop after that. This gives pure sulfuric acid with a concentration above 98%. Beware of any spillage of hexavalent chromium, it's a carcinogen! If such a spillage occurs, neutralize it with any reducing solution such as [[sodium thiosulfate]], [[ascorbic acid]] or [[glucose]].<br />
* Simple distillation of conc. drain cleaner sulfuric acid can work on some products, as hot sulfuric acid is oxidizing enough on its own that it will break down many organic contaminants.<ref>https://www.youtube.com/watch?v=4DUGRWjdNLI</ref> Similar to above, discard the first distillate fractions, and only keep the one with a density value of 1.84. This process however, may not work on all drain cleaners, so verify first.<br />
<br />
It is possible to further concentrate sulfuric acid by adding [[sulfur trioxide]], which reacts with the remaining water to form pure sulfuric acid. Sulfur trioxide can continue to be added to the solution to form [[oleum]], which fumes in air to form sulfuric acid droplets. When an equimolar concentration of sulfuric acid and sulfur trioxide is added, it forms [[pyrosulfuric acid]], which is a solid at room temperature. Sulfur trioxide can easily be obtained through the pyrolysis of certain salts, like anhydrous [[copper(II) sulfate]], [[iron(II) sulfate]], [[sodium pyrosulfate]] or [[potassium persulfate]].<br />
<br />
==Preparation==<br />
Sulfuric acid is industrially produced from sulfur, oxygen and water via the conventional contact process (DCDA), lead chamber process<ref>https://www.youtube.com/watch?v=7SDHeTcOXtI</ref> or the wet sulfuric acid process (WSA). The general way these processes work is by burning sulfur to obtain sulfur dioxide, which is oxidized to sulfur trioxide with the help of a catalyst, which in turn is dissolved in concentrated sulfuric acid, to form [[oleum]], which can be further concentrated into and eventually pyrosulfuric acid. The latter two products can be diluted using dil. sulfuric acid into conc. sulfuric acid. Diluted sulfuric acid is preferred instead of pure water, as the dilution is highly exothermic, while the reaction between sulfur trioxide with water is exothermic enough that the resulting sulfuric acid turns into a dense mist. The overall process can be written as:<br />
<br />
: S + O<sub>2</sub> → SO<sub>2</sub><br />
: SO<sub>2</sub> + ½ O<sub>2</sub> → SO<sub>3</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>S<sub>2</sub>O<sub>7</sub><br />
: H<sub>2</sub>S<sub>2</sub>O<sub>7</sub> + H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → 3 H<sub>2</sub>SO<sub>4</sub><br />
<br />
Each of the three main processes have their own advantages and disadvantages, but in general they work better at large scale, and for the average hobby chemist, while possible to reproduce them at smaller scale, it requires quite a lot of work to make the installation work properly. As such, working with volatile corrosive substances that melt your face off is quite an interesting project, if one were to try.<br />
<br />
There are many other routes to obtain sulfuric acid, most will produce diluted or mildly concentrated solutions, which can be concentrated to obtain more concentrated acid:<br />
<br />
*Absorbtion of sulfur dioxide in hydrogen peroxide: hydrogen peroxide will oxidize sulfur dioxide to sulfur trioxide, which reacts immediately with water to form sulfuric acid. Since this reaction is exothermic, an ice bath should be used. If an excess of SO<sub>2</sub> is used, warming the resulting solution to room temperature will cause some of the dissolved gas to boil off as the solution warms.<ref>https://www.youtube.com/watch?v=okvvD3-DF9U</ref><br />
<br />
: H<sub>2</sub>O<sub>2</sub> + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub><br />
<br />
While very easy to do, this reaction consumes hydrogen peroxide, and since H<sub>2</sub>O<sub>2</sub> is usually available OTC only as solutions from 3% up to 30%, the resulting sulfuric acid will be diluted, requiring further concentration.<ref>https://www.youtube.com/watch?v=mQMj5ier1lY</ref><br />
<br />
*Oxidation of SO<sub>2</sub> with conc. nitric acid: Similar to the reaction above with H<sub>2</sub>O<sub>2</sub>, conc. nitric acid can be used to oxidize sulfur dioxide directly to sulfuric acid, producing [[nitrogen dioxide]] as side product:<ref>https://www.youtube.com/watch?v=okvvD3-DF9U</ref><br />
<br />
: 2 HNO<sub>3</sub> + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub> + 2 NO<sub>2</sub><br />
<br />
The advantage of this reaction over the one with hydrogen peroxide, is that the nitrogen dioxide can be used to determine when the reaction is complete: when there is not more brown gas being produced, all the nitric acid has been consumed in the reaction. Main disadvantage of this route is that conc. nitric acid is a bit harder to acquire than sulfuric acid, and if one needs conc. sulfuric acid to obtain nitric acid, this route is not suitable. A modification of this reaction can be used, where the resulting nitrogen dioxide gets separated from the reaction, reacted with water to regenerate nitric acid, and then re-added in the reaction flask, to further oxidize the sulfur dioxide. Any nitric oxide produced from the side reaction between sulfur dioxide and nitrogen dioxide, can be reoxidized into nitrogen dioxide by injecting air in the mixture. <br />
<br />
*Ozone oxidation of sulfur dioxide: Ozone will oxidize sulfur dioxide into sulfur trioxide. This in turn reacts with water to form sulfuric acid. Ozone can be easily made by exposing oxygen to strong UV light, like that one produced by commercial ozone generators or low/high pressure mercury-vapor lamps. If atmospheric air is used, nitrogen dioxide may be produced as side product. This route is attractive since it uses cheap reagents, and while mercury UV lamps are somewhat difficult to properly operate, it's extremely easy to build a contraption where a continuous mixture of sulfur dioxide-oxygen is irradiated by strong UV light in a quartz tube, which produces sulfur trioxide directly. <br />
<br />
: 3 O<sub>2</sub> + hv → 2 O<sub>3</sub><br />
: SO<sub>2</sub> + O<sub>3</sub> → SO<sub>3</sub> + O<sub>2</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
<br />
*Electrolysis of aq. [[copper(II) sulfate]]: In a beaker, a concentrated solution of copper(II) sulfate is added. For cathode, a copper wire is added in the solution, at the bottom, and connected to the negative terminal of a power source, while for anode, a graphite electrode is added in the upper part of the solution, and connected to the positive terminal of the power source. During the process, the copper ions gets deposited on the copper electrode, while oxygen and hydrogen are produced at the carbon electrode. Overall, the reaction is as follows: <br />
<br />
: CuSO<sub>4</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub> + Cu + ½ O<sub>2</sub><br />
<br />
The resulting dil. solution of sulfuric acid is purified by filtering it, then concentrated by boiling it. This yields crude conc. H<sub>2</sub>SO<sub>4</sub>, which is distilled off to obtain the pure acid. The process is much easier than other electrochemical routes, as it's clean and relative quickly. Instead of graphite, other electrodes, like lead dioxide, titanium, platinum, or platinum on titanium can also be used.<ref>https://www.youtube.com/watch?v=5dUSF9Gl0xE</ref><ref>https://www.youtube.com/watch?v=ZRYtAquxffE</ref><br />
<br />
*Electrolysis of sulfate salt: This route involves electrolysis of a solution of a soluble sulfate salt, like [[magnesium sulfate]] or even [[ammonium sulfate]], using a diaphragm, which can either be either a classical ion-exchange diaphragm or a flower pot. <ref>https://www.youtube.com/watch?v=6BThiJpbBJQ</ref> The process yields dirty and diluted H<sub>2</sub>SO<sub>4</sub>, which requires purification and concentration.<ref>https://www.youtube.com/watch?v=b2wTha6Z-fA</ref><br />
<br />
*Pyrolysis of pyrosulfates: thermal decomposition of solid pyrosulfates yields sulfate and sulfur trioxide. The resulting sulfur trioxide is absorbed in crushed ice to form sulfuric acid. Further addition of sulfur trioxide yields conc. acid, and if SO<sub>3</sub> keeps getting added, it will convert into oleum, and eventually pyrosulfuric acid. The latter two products can be further diluted to concentrated sulfuric acid, by adding diluted sulfuric acid. For this process, [[sodium pyrosulfate]] is the best material, as it decomposes at a relative low temperature (460 °C) compared to other pyrosulfates, and the compound itself can be made by dehydrating [[sodium bisulfate]], which is readily and cheaply available:<br />
<br />
: 2 NaHSO<sub>4</sub> → Na<sub>2</sub>S<sub>2</sub>O<sub>7</sub> + H<sub>2</sub>O<br />
: Na<sub>2</sub>S<sub>2</sub>O<sub>7</sub> → Na<sub>2</sub>SO<sub>4</sub> + SO<sub>3</sub><br />
: SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub><br />
<br />
In theory, transition metal sulfates can also be used for this process, but since they decompose at higher temperatures, the resulting sulfur trioxide will partially decompose to sulfur dioxide and oxygen, which may lower the overall yield.<br />
<br />
*Copper chloride process: in an aqueous solution of [[copper(II) chloride]], sulfur dioxide is bubbled through. This reacts with the CuCl<sub>2</sub> from the aq. solution to form dil. sulfuric acid, HCl and CuCl:<br />
<br />
: 2 CuCl<sub>2</sub> + 2 H<sub>2</sub>O + SO<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub> + 2 CuCl + 2 HCl<br />
<br />
CuCl precipitates out of the solution. By injecting air in the suspension, the CuCl gets reoxidized to CuCl<sub>2</sub>, which can be reused. Sulfur dioxide is reinjected in the solution, which restarts the reaction, then the process gets repeated, until no more SO<sub>2</sub> can absorb in the reaction solution. The yield of this process is not great, unless one uses kg amounts of reagents. Likewise, the oxidation of Cu(I) to Cu(II) using air is very slow, taking many hours, which limits the efficiency of the overall process.<br />
<br />
*Electrobromine process: involves the reaction of elemental sulfur with elemental [[bromine]], using a graphite anode and copper metal cathode. In a beaker, where elemental sulfur is added at the bottom, the two electrodes are introduces, with the graphite electrode resting on the sulfur bed, while the copper anode is only partially submerged in the electrolyte solution. A solution of 5 M [[hydrobromic acid]] is used as electrolyte. When the process is activated, the HBr gets oxidized to bromide ions, which in term convert to elemental bromine, that sink to the bottom, reacting with the sulfur bed to yield disulfur dibromide, which hydrolyzes in water to yield sulfuric acid and HBr, the latter rising back to the anode, where it gets converted back to bromine, and the process repeats. It's important to keep the Cu electrode as high as possible, to prevent the bromide ions from reacting with the elemental bromine, as this yields tribromide ionds, which do not react with the sulfur, and instead just get reduced back into bromide ions, wasting electricity. Eventually, after 1-2 days, the process is almost complete. The solution is filtered off, and the resulting HBr is distilled to be recycled, while the sulfuric acid is concentrated and purified by distillation. The yield of this process is not great, and as it uses bromine, which is highly corrosive and toxic. Likewise, the graphite electrodes get used up very quickly in the reaction. The sulfur bed may break apart during the process, and stirring may be required to break it apart and allow it to settle back. Stop the process and remove the electrodes, before stirring the suspension, and once the sulfur settles back, reintroduce the electrodes, and restart the process. Alternatively, one can a solid piece of sulfur instead of powder, as this shouldn't rise, though this may affect the speed of the reaction, as bulk sulfur reacts slower than powdered sulfur. A porous separating membrane, like a glass fiber cloth may be used to pin the sulfur bed down, while allowing the bromine to diffuse through it to reach the sulfur, though this hasn't been tested so far.<ref>https://www.youtube.com/watch?v=6ms6xbPhdVs</ref><br />
<br />
==Projects==<br />
* Preparation of metal sulfates<br />
* Preparation of nitro compounds through [[nitration]]<br />
* The dehydration of [[sucrose]] to produce elemental [[carbon]]<br />
* [[Esterification]]s that require a dehydrating agent, such as that of [[ethyl acetate]], [[methyl salicylate]], etc.<br />
* Making simple [[rayon]] fibers with [[Schweizer's reagent]] and [[cellulose]]<br />
* Producing other concentrated acids by the reaction of sulfuric acid with an anhydrous salt, such as in the production of fuming [[nitric acid]] and glacial [[acetic acid]]<br />
<br />
==Handling==<br />
===Safety===<br />
[[File:Corrosive.png|thumb|right|Corrosive]] While low concentration sulfuric acid is relatively safe to work with (under 40% w/w)), concentrated sulfuric acid (over 90% w/w) is extremely corrosive and dangerous. It does not only causes chemical burns, it also causes burns by dehydration of organic materials (like skin), destroying the molecules to form water with the -OH groups in them. Safety measures should be taken and all skin should be covered when working with concentrated sulfuric acid.<br />
<br />
When heating sulfuric acid, it is important to DO NOT OVERFILL THE FLASK. Concentrated sulfuric acid's volume increases by nearly 16% between 0 and 330°C, an overfilled flask will spill its content. Also, sulfuric acid, even diluted, tends to bump when it boils, accumulating heat to release a violent burst of steam from time to time. The use of boiling chips reduces this phenomenon, but there is no way to stop it completely. It is advised to take measures to prevent spills, an anti-splash adapter with ground glass joint being a very convenient option.<br />
<br />
Hot concentrated sulfuric acid may decompose to form sulfur dioxide and sulfur trioxide, which are toxic and corrosive, respectively. It fumes profusely when hot, the fumes consist of sulfuric acid droplets and a SOx mix. These fumes are very dangerous and a known lung carcinogen.<br />
<br />
When carrying glass bottles of sulfuric acid and you worry there's a risk you might break it, a good tip would be to carry it in a (plastic) bucket, partially filled with sand.<br />
<br />
===Storage===<br />
Sulfuric acid should be stored in closed bottles. While glass bottles, being inert, are good for storing concentrated sulfuric acid, concentrated (80-98%) sulfuric acid is often stored in PE (more specifically UDPE or UHDPE) bottles, as PE is not brittle, so in the event you drop the bottle on a hard surface, it will not shatter and splash conc. sulfuric all over the place. Unfortunately, PE bottles are sensitive to light and will degrade over the years if exposed to sunlight, so they must be stored in a dark place away from UV light, like a cupboard. Commercial PE bottles used for conc. sulfuric acids tend to be colored, which helps to limit degradation from strong light and oxygen. However, if you plan to store the acid for more that several years, it's recommended to use glass bottles.<br />
<br />
Long-term storage of concentrated sulfuric acid may lead to it absorbing water from air and becoming less concentrated. When this happens, the acid needs to be "re-freshened" by distilling unnecessary water off it.<br />
<br />
===Disposal===<br />
Sulfuric acid can be neutralized with any base or carbonate, preferably [[calcium hydroxide]] or carbonate.<br />
<br />
Concentrated sulfuric acid, like any concentrated acid, should be first strongly dilute it in a large volume of water before neutralizing it with a base. Another method would be to add it in an acid-resistant container with a lid and slowly add solid calcium hydroxide/carbonate or sodium bicarbonate chunks and close the lid to limit splashing. Wait until it stopped fizzing then keep adding until it no longer reacts. Be careful, as the thicker the solution becomes, the stronger the foaming gets.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6911 Sulfuric Acid Production: Revisited]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2824 H2SO4 by the Lead Chamber Process - success]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=64535 I will now be building and testing my new Batparatus!]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=3722 cleaning sulfuric acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13313 Sulfuric Acid at Home]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19117 Concentrating dilute sulphuric acid(battery acid) without distillation]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=91332 Sulfuric acid from gypsum using diaphragm cell]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14857 Sulfuric acid purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14570 sulfuric acid turned black]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=61920 Distilling Sulfuric Acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=65331 Sulfuric acid in NZ]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=14291 Should I get rid of my H2SO4?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13726 sulfuric acid accident]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=62863 Sulfuric acid storage]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13964 HDPE as a storage for Sulfuric Acid]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13148 Safely Storing H2SO4 (35%)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6217 Storage for Sulfuric Acid (H2SO4)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=25679 Sulfuric Acid and LDPE issue]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Oxoacids]]<br />
[[Category:Sulfur oxoacids]]<br />
[[Category:Sulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Materials unstable in basic solution]]<br />
[[Category:Things that can kill you very quickly]]<br />
[[Category:Hygroscopic compounds]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Essential reagents]]<br />
[[Category:DEA List II chemicals]]<br />
[[Category:Catalysts]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Boiling_the_Bat&diff=14957Boiling the Bat2023-03-27T20:25:33Z<p>Ave369: </p>
<hr />
<div>{{main|Sulfuric acid}}<br />
<br />
Fresh car battery acid (the Bat) is the main source of sulfuric acid by many. However, it is diluted sulfuric acid, and many experiments require concentrated. So you have to boil it down to concentrated. This process is best done in two steps.<br />
<br />
Step first involves pouring the Bat in a flask and boiling it, preferably with boiling chips. For the most part, water will come out and sulfuric acid remain. However, this process can only go until the acid reaches the concentrations of 70-80%, when the noxious SOx gases and acid mist will start to emerge. If you have access to a fume hood, you can continue the process. Otherwise, you need a special distilling setup to distill remaining water off.<br />
<br />
The setup must have two condensers: one air-cooled (such as the nose of the retort) and one water-cooled (such as a Liebig). The Liebig serves to condense the SOx gases and water vapor into dilute sulfuric acid, thwarting any possible damage done by them. The boiling flask should be insulated with asbestos cord, to prevent refluxing of sulfuric acid on the retort walls.<br />
<br />
You pour your half-boiled (70-80%) sulfuric acid in the distillation retort, put in some boiling chips and assemble the apparatus. Once you start heating, you will notice white mist in the nose of the retort and droplets of very dilute sulfuric acid coming out as distillate. Ignore it.<br />
<br />
The cue you should watch for is when oily waves of sulfuric acid appear on the inner walls of the retort nose instead of the dewey drops of water. From this on, you should check the strength of the distillate with any method. When it starts to char matches, it is a good sign: it means that at least 95% sulfuric acid is in the pot. If you can measure the density of the distillate, it is the best method: stop distilling when the distillate will start coming with a density of at least 1.8 g/cm3.<br />
<br />
After that, turn off the apparatus, remove the asbestos and wait for it to cool. The liquid in the pot is sulfuric acid with a concentration of 96-98%, useful for most reactions.<br />
<br />
The second step can also be used to "re-freshen" old sulfuric acid that turned brown and weak from improper storage. Add several drops of concentrated hydrogen peroxide to get rid of the dark color before you start heating.<br />
<br />
Use only inert boiling chips such as silica gel. Chips like pieces of red brick tend to partially dissolve in the acid and contaminate it with ions of iron, making the acid ranging in color from castor oil to soy sauce. Avoid getting any organics into the pot: it will make the acid black. To clear your acid from organic blackness, add hydrogen peroxide in it and re-dry it as described above.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=64535 I will now be building and testing my new Batparatus!]<br />
<br />
[[Category:How-to]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Urea_nitrate&diff=14956Urea nitrate2023-03-27T18:27:47Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Urea nitrate<br />
| Reference =<br />
| IUPACName = Urea nitrate<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Carbamide nitrate<br>Urea hydronitrate<br>Urea mononitrate<br>Uronium nitrate<br />
<!-- Images --><br />
| ImageFile = Urea_nitrate_crystals.jpg<br />
| ImageSize = 270<br />
| ImageAlt = <br />
| ImageName =<br />
| ImageCaption = Large urea nitrate crystals <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 1.69 g/cm<sup>3</sup><br />
| Formula = CH<sub>5</sub>N<sub>3</sub>O<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 123.068 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 163<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 15 g/100 ml<br />
| SolubleOther = Soluble in [[ethanol]]<br>Insoluble in hydrocarbons<br />
| Solvent = <br />
| Solubility1 = 3.7 g/100 g (30 °C)<ref>Tokuoka, M.; Morooka, H.; Bull. Agr. Chem. Soc. Japan; vol. 10; (1934); p. 127 - 129</ref><br />
| Solvent1 = acetone<br />
| Solubility2 = 50.5 g/100 g (30 °C)<ref>Tokuoka, M.; Morooka, H.; Bull. Agr. Chem. Soc. Japan; vol. 10; (1934); p. 127 - 129</ref><br />
| Solvent2 = ethanol<br />
| Solubility3 = 11.65 g/100 g (30 °C)<ref>Tokuoka, M.; Morooka, H.; Bull. Agr. Chem. Soc. Japan; vol. 10; (1934); p. 127 - 129</ref><br />
| Solvent3 = methanol<br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Monoclinic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = 525.89 kJ/kmol<br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = 3,400 m/s<br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.pfaltzandbauer.com/MSDS/U01150%20%20SDS%20%20031017.pdf Pfaltz&Bauer]<br />
| FlashPt = <br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Explosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Ammonium nitrate]]<br>[[Guanidinium nitrate]]<br>[[Methylammonium nitrate]]<br />
}}<br />
}}<br />
'''Urea nitrate''' is the nitrate salt of the organic base [[urea]]. It is an explosive material used in various applications, but it has gained a bad reputation due to its use in many terrorists plots, such as the World Trade Center bombing in 1993.<br />
<br />
Urea nitrate should not be confused with ''[[nitrourea]]'', which is a nitroamine and not a nitrate salt.<br />
<br />
Although the name "urea nitrate" is technically incorrect since urea is the free base, the proper name for this compound, '''uronium nitrate''' is very rarely used.<br />
<br />
==Properties==<br />
===Chemical===<br />
Urea nitrate is unstable in basic solution. It also hydrolyzes in solution, which is quite acidic (urea being a very weak base); hot water causes the hydrolysis to proceed more fully, and nitric acid can be distilled back from the solution if heated enough.<br />
<br />
===Physical===<br />
Urea nitrate is a white solid, sparingly soluble in water (less so than urea itself). Its solubility in cold water is significantly less than in warm water.<br />
<br />
===Explosive===<br />
Urea nitrate is an explosive with an average detonation velocity of 4,000 m/s (depending on the purity, it can be between 3,400 m/s and 4,700 m/s). Its destructive properties are similar to that of [[ammonium nitrate]]-based explosives.<br />
<br />
==Availability==<br />
Due to being an explosive material, the sale of urea nitrate is restricted.<br />
<br />
==Preparation==<br />
Urea nitrate can be prepared by reacting urea with nitric acid. The reaction is exothermic, so it's best to do it at low temperatures.<br />
<br />
Urea nitrate can be prepared via a [[double replacement]] by combining urea with a nitrate salt and concentrated hydrochloric acid in water. Gentle heat, as from a water bath, should be added to bring the urea and nitrate into solution before the HCl is added. The turbid solution clears as the HCl is added and a copious precipitate of urea nitrate is obtained when the solution is cooled below 0 Celsius. The remaining solution contains a small amount of product and the chloride salt of the cation whose nitrate was used.<ref>[http://www.angelfire.com/empire/megapyro6/Pyro/UN.html Making Urea Nitrate without Nitric Acid]</ref><br />
<br />
==Projects==<br />
*Make blasting charges<br />
*Prepare Urea nitrate from assorted nitrates, recovering the chloride in each case.<br />
<br />
==Handling==<br />
===Safety===<br />
Urea nitrate appears to be a flame retardant. Filter paper soaked in it and dried will char in a butane lighter flame but not catch fire.<br />
<br />
===Storage===<br />
Storage should be limited to very small amounts, as larger amounts may invite visits by ''Les Gendarmes''.<br />
<br />
===Disposal===<br />
Casual tests are planned to evaluate using urea nitrate as a house plant food.<br />
<br />
==Gallery==<br />
<gallery widths="200" position="center" columns="4" orientation="none"><br />
Urea_nitrate.jpg|Small crystals of urea nitrate<br />
</gallery><br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13470 Synthesizing UN with calcium nitrate]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=10362 Urea Nitrate Procedure Mechanism]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Organic compounds]]<br />
[[Category:Nitrated organic compounds]]<br />
[[Category:Nitrates]]<br />
[[Category:Energetic materials]]<br />
[[Category:High explosives]]<br />
[[Category:Secondary explosives]]<br />
[[Category:Things that can kill you very quickly]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Carbonic_acid&diff=14954Carbonic acid2023-03-26T18:41:21Z<p>Ave369: /* Physical */</p>
<hr />
<div>{{Chembox<br />
| Name = Carbonic acid<br />
| Reference =<br />
| IUPACName = Carbonic acid<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Acid of air<br>Aerial acid<br>Carbon dioxide solution<br>Dihydrogen carbonate<br>Hydroxymethanoic acid<br />
<!-- Images --><br />
| ImageFile = Carbonic_acid.jpg<br />
| ImageSize = 270<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Solution of carbon dioxide in water<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid/colorless gas (pure compound)<br>Unstable liquid (aq. solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 1.668 g/cm<sup>3</sup><br />
| Formula = H<sub>2</sub>CO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 62.03 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| Odor = Odorless<br />
| pKa = 3.6 (p''K''<sub>a1</sub> for H<sub>2</sub>CO<sub>3</sub> only)<br>6.3 (p''K''<sub>a1</sub> including CO<sub>2</sub>(aq))<br>10.32 (p''K''<sub>a2</sub>)<br />
| pKb = <br />
| Solubility = Only stable in solution<br />
| SolubleOther = Soluble in [[glycerol]]<br>Insoluble in [[ethanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://beta-static.fishersci.com/content/dam/fishersci/en_US/documents/programs/education/regulatory-documents/sds/chemicals/chemicals-c/S25234.pdf FisherScientific] (aq. sol.)<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Acetone]]<br>[[Urea]]<br>[[Dimethyl carbonate]]<br />
}}<br />
}}<br />
'''Carbonic acid''' is a very unstable, weak acid formed in small concentrations with dissolution of its anhydride, [[carbon dioxide]], in water. Any solution containing it also contains molecular carbon dioxide dissolved in water.<br />
<br />
==Properties==<br />
=== Physical ===<br />
<br />
Anhydrous carbonic acid cannot be isolated under normal conditions. Only in extreme conditions it becomes stable and isolable in the form of a white solid (below 220 K) or a colorless gas (above 220 K); when transferred to normal conditions, it remains stable in absence of water vapor but decomposes rapidly on contact with it<ref>https://doi.org/10.1002%2F%28SICI%291521-3773%2820000303%2939%3A5%3C891%3A%3AAID-ANIE891%3E3.0.CO%3B2-E</ref>. Solutions of carbonic acid typically resemble soda water or mineral water (and usually are soda water or mineral water).<br />
<br />
=== Chemical ===<br />
<br />
Carbonic acid is a weak diprotic acid. Its first dissociation gives H+ and the hydrocarbonate anion, HCO<sub>3</sub><sup>-</sup>. The second dissociation gives another H+ and the carbonate anion, CO<sub>3</sub><sup>2-</sup>. <br />
<br />
Solutions of carbonic acid and carbon dioxide react with alkaline bases, forming carbonate salts.<br />
<br />
== Sources and production ==<br />
<br />
Solutions of carbon dioxide containing carbonic acid are sold as carbonated water.<br />
<br />
Such a solution can also be made by reacting any carbonate salt with any acid stronger than carbonic acid. It can also be made from [[water]] and [[carbon dioxide]] using a soda siphon or carbonator.<br />
<br />
==Projects==<br />
*Make [[calcium bicarbonate]]<br />
*Aq. solution of CO<sub>2</sub><br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Solutions of carbonic acid are absolutely safe. They can be consumed by humans with no ill effects, providing no other chemicals are present in the solution. They are well known for the peculiar taste and refreshing effect provided by bubbles of carbon dioxide.<br />
<br />
=== Storage ===<br />
Solutions of carbonic acid can be stored in closed plastic bottles. You know, the same way you store soda. Because they are soda.<br />
<br />
===Disposal===<br />
Carbonic acid does not require any special disposal and can be poured down the drain.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Weak acids]]<br />
[[Category:Carbonates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials unstable in basic solution]]<br />
[[Category:Edible chemicals]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Carbonic_acid&diff=14953Carbonic acid2023-03-26T18:40:20Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Carbonic acid<br />
| Reference =<br />
| IUPACName = Carbonic acid<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Acid of air<br>Aerial acid<br>Carbon dioxide solution<br>Dihydrogen carbonate<br>Hydroxymethanoic acid<br />
<!-- Images --><br />
| ImageFile = Carbonic_acid.jpg<br />
| ImageSize = 270<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Solution of carbon dioxide in water<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid/colorless gas (pure compound)<br>Unstable liquid (aq. solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 1.668 g/cm<sup>3</sup><br />
| Formula = H<sub>2</sub>CO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 62.03 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| Odor = Odorless<br />
| pKa = 3.6 (p''K''<sub>a1</sub> for H<sub>2</sub>CO<sub>3</sub> only)<br>6.3 (p''K''<sub>a1</sub> including CO<sub>2</sub>(aq))<br>10.32 (p''K''<sub>a2</sub>)<br />
| pKb = <br />
| Solubility = Only stable in solution<br />
| SolubleOther = Soluble in [[glycerol]]<br>Insoluble in [[ethanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://beta-static.fishersci.com/content/dam/fishersci/en_US/documents/programs/education/regulatory-documents/sds/chemicals/chemicals-c/S25234.pdf FisherScientific] (aq. sol.)<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Acetone]]<br>[[Urea]]<br>[[Dimethyl carbonate]]<br />
}}<br />
}}<br />
'''Carbonic acid''' is a very unstable, weak acid formed in small concentrations with dissolution of its anhydride, [[carbon dioxide]], in water. Any solution containing it also contains molecular carbon dioxide dissolved in water.<br />
<br />
==Properties==<br />
=== Physical ===<br />
<br />
Anhydrous carbonic acid cannot be isolated under normal conditions. Only in extreme conditions it becomes stable and isolable in the form of a colorless gas; when transferred to normal conditions, it remains stable in absence of water vapor but decomposes rapidly on contact with it<ref>https://doi.org/10.1002%2F%28SICI%291521-3773%2820000303%2939%3A5%3C891%3A%3AAID-ANIE891%3E3.0.CO%3B2-E</ref>. Solutions of carbonic acid typically resemble soda water or mineral water (and usually are soda water or mineral water).<br />
<br />
=== Chemical ===<br />
<br />
Carbonic acid is a weak diprotic acid. Its first dissociation gives H+ and the hydrocarbonate anion, HCO<sub>3</sub><sup>-</sup>. The second dissociation gives another H+ and the carbonate anion, CO<sub>3</sub><sup>2-</sup>. <br />
<br />
Solutions of carbonic acid and carbon dioxide react with alkaline bases, forming carbonate salts.<br />
<br />
== Sources and production ==<br />
<br />
Solutions of carbon dioxide containing carbonic acid are sold as carbonated water.<br />
<br />
Such a solution can also be made by reacting any carbonate salt with any acid stronger than carbonic acid. It can also be made from [[water]] and [[carbon dioxide]] using a soda siphon or carbonator.<br />
<br />
==Projects==<br />
*Make [[calcium bicarbonate]]<br />
*Aq. solution of CO<sub>2</sub><br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Solutions of carbonic acid are absolutely safe. They can be consumed by humans with no ill effects, providing no other chemicals are present in the solution. They are well known for the peculiar taste and refreshing effect provided by bubbles of carbon dioxide.<br />
<br />
=== Storage ===<br />
Solutions of carbonic acid can be stored in closed plastic bottles. You know, the same way you store soda. Because they are soda.<br />
<br />
===Disposal===<br />
Carbonic acid does not require any special disposal and can be poured down the drain.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Weak acids]]<br />
[[Category:Carbonates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials unstable in basic solution]]<br />
[[Category:Edible chemicals]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Carbonic_acid&diff=14952Carbonic acid2023-03-26T18:39:21Z<p>Ave369: /* Physical */</p>
<hr />
<div>{{Chembox<br />
| Name = Carbonic acid<br />
| Reference =<br />
| IUPACName = Carbonic acid<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Acid of air<br>Aerial acid<br>Carbon dioxide solution<br>Dihydrogen carbonate<br>Hydroxymethanoic acid<br />
<!-- Images --><br />
| ImageFile = Carbonic_acid.jpg<br />
| ImageSize = 270<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Solution of carbon dioxide in water<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid (pure compound)<br>Unstable liquid (aq. solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 1.668 g/cm<sup>3</sup><br />
| Formula = H<sub>2</sub>CO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 62.03 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| Odor = Odorless<br />
| pKa = 3.6 (p''K''<sub>a1</sub> for H<sub>2</sub>CO<sub>3</sub> only)<br>6.3 (p''K''<sub>a1</sub> including CO<sub>2</sub>(aq))<br>10.32 (p''K''<sub>a2</sub>)<br />
| pKb = <br />
| Solubility = Only stable in solution<br />
| SolubleOther = Soluble in [[glycerol]]<br>Insoluble in [[ethanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://beta-static.fishersci.com/content/dam/fishersci/en_US/documents/programs/education/regulatory-documents/sds/chemicals/chemicals-c/S25234.pdf FisherScientific] (aq. sol.)<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Acetone]]<br>[[Urea]]<br>[[Dimethyl carbonate]]<br />
}}<br />
}}<br />
'''Carbonic acid''' is a very unstable, weak acid formed in small concentrations with dissolution of its anhydride, [[carbon dioxide]], in water. Any solution containing it also contains molecular carbon dioxide dissolved in water.<br />
<br />
==Properties==<br />
=== Physical ===<br />
<br />
Anhydrous carbonic acid cannot be isolated under normal conditions. Only in extreme conditions it becomes stable and isolable in the form of a colorless gas; when transferred to normal conditions, it remains stable in absence of water vapor but decomposes rapidly on contact with it<ref>https://doi.org/10.1002%2F%28SICI%291521-3773%2820000303%2939%3A5%3C891%3A%3AAID-ANIE891%3E3.0.CO%3B2-E</ref>. Solutions of carbonic acid typically resemble soda water or mineral water (and usually are soda water or mineral water).<br />
<br />
=== Chemical ===<br />
<br />
Carbonic acid is a weak diprotic acid. Its first dissociation gives H+ and the hydrocarbonate anion, HCO<sub>3</sub><sup>-</sup>. The second dissociation gives another H+ and the carbonate anion, CO<sub>3</sub><sup>2-</sup>. <br />
<br />
Solutions of carbonic acid and carbon dioxide react with alkaline bases, forming carbonate salts.<br />
<br />
== Sources and production ==<br />
<br />
Solutions of carbon dioxide containing carbonic acid are sold as carbonated water.<br />
<br />
Such a solution can also be made by reacting any carbonate salt with any acid stronger than carbonic acid. It can also be made from [[water]] and [[carbon dioxide]] using a soda siphon or carbonator.<br />
<br />
==Projects==<br />
*Make [[calcium bicarbonate]]<br />
*Aq. solution of CO<sub>2</sub><br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Solutions of carbonic acid are absolutely safe. They can be consumed by humans with no ill effects, providing no other chemicals are present in the solution. They are well known for the peculiar taste and refreshing effect provided by bubbles of carbon dioxide.<br />
<br />
=== Storage ===<br />
Solutions of carbonic acid can be stored in closed plastic bottles. You know, the same way you store soda. Because they are soda.<br />
<br />
===Disposal===<br />
Carbonic acid does not require any special disposal and can be poured down the drain.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Weak acids]]<br />
[[Category:Carbonates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials unstable in basic solution]]<br />
[[Category:Edible chemicals]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Regulatory_quirks&diff=14951Regulatory quirks2023-03-24T09:04:08Z<p>Ave369: /* Ethanol */</p>
<hr />
<div>{{Stub}}<br />
In an effort to limit the availability of various hazardous chemicals or to limit their impact, various laws have been passed over the years with the purpose to restrict or regulate various chemicals or lab equipment with various degrees of success. However, very often, due to a variety of reasons ranging from badly phrased text, laws written without consulting experts in the domain, bad science or simply a desire to have the laws passed for political gain, many such laws are either incomplete, cannot be properly enforced, or have loopholes that often allow for the restriction to be circumvented with little inconvenience. This page will present examples of various '''regulatory quirks''' found for many chemicals and lab/scientific equipment in the laws of many countries, as well as the unforeseen or negative effects of said legislation.<br />
<br />
==Chemicals==<br />
===[[Acetyl chloride]]===<br />
Although [[acetic anhydride]] is classified as List II Drug precursor in most of the world, and its sale is strictly regulated or restricted, it can be very easily made by distilling a mixture of anhydrous [[sodium acetate]] (easily prepared) and acetyl chloride. Unlike acetic anhydride, acetyl chloride is not generally classified as drug precursor, so there are fewer restrictions to it (may still be classified as hazardous chemical due to its ability to release harmful HCl vapors in presence of moisture). However, acetyl chloride itself can be used instead of acetic anhydride in the manufacturing of many illicit substances, with various success. Only in some countries, like Australia, Russian Federation, acetyl chloride is also classified as drug precursor, and it's not easy to acquire.<br />
<br />
As of 2021, in most EU countries it's nearly impossible to acquire acetyl chloride, even in small amounts, as most sellers now require filling in an EUD. It's unclear if this applies to other acetyl halides.<br />
<br />
===[[Ammonium sulfamate]]===<br />
Ammonium sulfamate is no longer an accepted herbicide in the EU since 2008<ref>https://www.allotment-garden.org/garden-diary/1989/ammonium-sulphamate-weed-killer-banned/</ref> due to the Irish Rapporteur not receiving testing on dogs for said chemical and thus the compound did not receive the license to be allowed as herbicide.<ref>https://web.archive.org/web/20091113071628/http://www.pesticides.gov.uk/garden.asp?id=1997</ref> However, ammonium sulfamate is still legally allowed to be used as compost accelerator. Since both products are found in the same section of most stores (gardening) or online stores just a click away, nothing will stop anyone who knew that the compost accelerator can also be used as herbicide to, well, use it as herbicide.<br />
<br />
===[[Barium]] salts===<br />
Barium and his compounds, like most heavy metal compounds are often classified as hazardous, however, unlike other heavy metals, the restriction on its compounds varies significantly depending on the individual compound and differs from country to country. In UK, all barium salts, except for [[barium sulfate]], [[barium carbonate]] and barium silicofluoride/hexafluorosilicate, are classified as poisons and restricted from sale to private individuals without a permit.<ref>https://www.gov.uk/government/publications/licensing-for-home-users-of-explosives-precursors/licensing-for-home-users-of-poisons-and-explosive-precursors</ref> However, since barium carbonate is not excepted from the restriction, it's possible to make all the other barium compounds from the carbonate, which defeats the entire purpose of the restriction! Likewise, since other insoluble and non-toxic barium compounds, like barium titanate, are not excepted, this means that said insoluble and non-toxic barium compound is actually considered to be toxic. Natural minerals of barium, like Gurimite (barium vanadate) would also be classified as poison by this regulation. Barium metal is also not included, but since it's a metal and not a salt, so its status remains unclear (unless another piece of legislature classifies it as hazardous/poisonous material), and while much more expensive than most barium compounds, it can also be used to make any barium compound.<br />
<br />
===[[Benzaldehyde]]===<br />
Although benzaldehyde is classified as List I precursor in many countries, not all countries classify this substance as drug precursor, even though toluene, which is usually a precursor for making benzaldehyde, is classified as List II precursor in those countries' legislation. Likewise, bitter almond oil, which is essentially crude benzaldehyde, is not covered in all countries as drug precursor. In US for example, bitter almond oil is classified as drug precursor, while in Canada is not. In most of EU countries, benzaldehyde is not classified as drug precursor, though it's sale may be monitored by drug agencies as unclassified drug precursor.<br />
<br />
===[[Benzyl alcohol]]===<br />
Although [[benzaldehyde]], [[benzyl chloride]], [[toluene]] are classified as drug precursors of different types in most countries that enforce the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, [[benzyl alcohol]], which may be used as precursor to synthesize all the three mentioned compounds, is not classified as drug precursor in any country, and is readily available without any restrictions.<br />
<br />
In some EU countries, chemical sellers may ask an EUD over a certain amount of product, but this appears to be mostly local, as other EU countries do not require any paperwork for (almost) any amount.<br />
<br />
===[[Binary explosive]]s===<br />
In US, under ATF regulations, binary explosives are not classified as explosive materials, ''when kept as separate components''. However, the moment they are mixed, the resulting mixture is considered explosive material.<ref>https://www.atf.gov/explosives/binary-explosives</ref> This means that one can potentially have a "just add water" bomb-making material without breaking the law.<br />
<br />
===[[Chlorate]]s/[[Perchlorate]]s===<br />
In EU, the sale of sodium and potassium chlorate/perchlorates is forbidden to private persons and only companies are allowed to handle products containing more than 40% chlorate/perchlorate and an explosive permit for working with these 4 compounds must be obtained. However, there are no mentioned limits for how much <40% chlorate/perchlorate containing material one may purchase unless specified in said country's legislation, meaning one may buy a large amount of dilute product and concentrate the desired chlorate/perchlorate. Likewise, keeping perchlorate mixed with an inert solid and then concentrating it by extracting it with a selective solvent, like acetone in this case, in which perchlorate is very soluble, circumvents the need to keep the perchlorate at a low concentration in the first place.<br />
<br />
The restriction also does not seem to cover the other chlorates and perchlorates, at least at directive level, meaning that all the other metal chlorates/perchlorates can be acquired without a problem, unless the member country specifically restricts said compound. This also seems to include perchloric acid (which can be used to make any perchlorate), unless local laws classify it as explosive precursor, and, unless classified as explosive material, [[ammonium perchlorate]] (explosive on its own) is exempted from the EU-wide restriction. Other energetic compounds like [[guanidinium perchlorate]] also seem to fall in this issue. In the Federal Republic of Germany, guanidinium perchlorate and guanidinium picrate were classified as explosive materials, and thus their use was restricted, and presumably, the same regulation was carried after the German reunification.<ref>Sprengstoffgesetz, annex I, part 1, no. 33, 34 and part 2,no. 2.5, mixtures 4,5</ref><br />
<br />
Chlorates are a decomposition product of hypochlorite salts aka bleach. Thus, any old bottle of bleach will contain a significant (though not large) amount of chlorates, specifically sodium chlorate. The directive does not mention any exception for accidental/side production of chlorates/perchlorates in any form, and while the amount of chlorate produced is usually below 40% of the solution's mass, there is no mention if the precipitated chlorate (like adding KCl to the decomposed bleach) which is very easy to separate from the reaction mass is considered part of the bottle or solution.<br />
<br />
===[[Ethanol]]===<br />
In almost all countries, food-grade distilled ethanol is subjected to taxes (excise duty on alcohol), which drives up the cost of liquors and rectified spirits. To avoid these taxes, concentrated ethanol is denatured (turned poisonous, bad-tasting, foul-smelling or nauseating) by adding various additives, making it unfit for human consumption. This practice is also done for lab-grade ethanol, while non-denatured lab-grade ethanol has a similar price to the food-grade rectified spirit (even though lab-grade non-denatured alcohol is clearly labeled as unfit for consumption). Some countries, like Bulgaria, Hungary, Romania allow the production of small amounts of home distilled liquor up to a certain amount or sometimes said practices are ignored by the authorities if the deed is not worth pursuing legally due to low resources, willful ignorance, nepotism or more often local tradition regarding cultural alcohol production.<br />
<br />
In some countries, ethanol with a concentration over 70% is a regulated substance that cannot be purchased over the counter, and even denatured alcohol is not available. Homebrewing, if done clandestinely or without profit, is usually ignored. So the main source of relatively pure ethanol in these countries is either the black market, homebrewing or buying from a homebrewer.<br />
<br />
===[[Ether]]s===<br />
All drug precursor lists (DEA List e.g.) list [[diethyl ether]] as List II precursor (used in the manufacturing of controlled substances). However, no other ether is present on the list, even though for this intended purpose (organic extractions), other ethers are just as good. [[Diisopropyl ether]], while extremely hazardous on its own since it rapidly builds up explosive peroxides over the course of a few months, can be easily made from near anhydrous isopropanol, which, unlike ethanol, is easily and cheaply available in high concentration at most hardware stores. Since it's very easy to synthesize, there's no need to store it and can be burned after use, thus eliminating the peroxide threat completely.<br />
<br />
[[Tetrahydrofuran]] can also be used instead of diethyl ether for many reactions (including Grignard reactions) and in most Western countries there aren't restrictions on its use or possession, though in a few countries, like Russia, its possession and use are restricted. However, there is no mention of THF derivatives being included on the drug precursor list in Russia or other countries that regulate THF, meaning that unless specified in another list/annex, 2-methyltetrahydrofuran could be used instead, and is also suitable for Grignard reactions (though this derivative compound is a bit more expensive compared to regular THF and it's not miscible with water).<br />
<br />
===[[Hydrogen iodide]]/[[Hydroiodic acid]]===<br />
Hydrogen iodide/Hydroiodic acid are classified as List I precursor in US by the DEA, meaning their sale is regulated. However, HI can be easily made by adding conc. phosphoric acid to an iodide salt, both readily available reagents and not classified as precursors. Also, since HI is unstable in the presence of air and or light, the acid itself can be stored separately in the form of salt and acid, without breaking any laws, essentially as a quick DIY HI kit.<br />
<br />
===[[Hydrogen peroxide]]===<br />
In the EU, all hydrogen peroxide in concentrations higher than 12% is forbidden to be sold to private individuals. However, [[sodium percarbonate]], which contains 32.5% H<sub>2</sub>O<sub>2</sub> by weight is not restricted at all. Granted, extracting the peroxide from the percarbonate is not easy, but it does appear that the restriction only covers liquids containing H<sub>2</sub>O<sub>2</sub>. This also applies to other hydrogen peroxide adducts like [[urea peroxide]] (anhydrous form contains 34% H<sub>2</sub>O<sub>2</sub> by weight, while the tetrahydrate 22% H<sub>2</sub>O<sub>2</sub> by weight).<br />
<br />
Likewise, the directive does not place a limit on the amount of <12% hydrogen peroxide one can acquire, meaning one can concentrate the 10-12% peroxide to a more concentrated solution (though naturally the resulting volume is smaller), meaning one can convert 3x1 L bottles of 10% H<sub>2</sub>O<sub>2</sub> into 1 L H<sub>2</sub>O<sub>2</sub> 30% for example.<br />
<br />
The directive does not restrict any precursor chemicals that could be used to prepare hydrogen peroxide (both diluted and concentrated), such as [[sodium perborate]], metal peroxides ([[barium peroxide]], [[zinc peroxide]], etc.), sodium/potassium peroxodisulfate, etc.<br />
<br />
===[[Iodine]]===<br />
Iodine is classified as DEA List I precursor in US, meaning its sale to private individuals is restricted. However, elemental iodine can be made easily from any iodide salt (which are not classified as drug precursors) by simply reacting the iodide with a mixture of [[sulfuric acid]] and [[hydrogen peroxide]], or just with [[Potassium peroxymonosulfate|Oxone]] (while sulfuric acid is List II precursor and conc. hydrogen peroxide may be classified as explosive precursor, Oxone is not classified as any precursor in US or any other country and can be easily bought from most swimming pool stores).<br />
<br />
Likewise, given that elemental iodine is not easy to store, keeping the iodide salt and the oxidizer/acid separately is also a good way of storing iodine for long term and will also ignore any legal problems associated with the possession of the element.<br />
<br />
In most countries outside the US, iodine is not classified as drug precursor, though restrictions on its transport or storage may exist.<br />
<br />
===[[Methanol]]===<br />
In many countries, like Italy, Austria, Bulgaria or Romania, methanol is classified as poison and the sale of the compound is regulated or restricted to private individuals. However, methanol-based fuels, which consist mainly of methanol with a few additives added, do not appear to be subjected to the same restrictions. RC fuels containing methanol also don't appear to be restricted. Occasionally, some technical alcohol products used for cleaning are also almost pure methanol (with some dye added) and such products are sometimes sold in auto part shops or sometimes even hardware stores, depending on the country.<br />
<br />
===[[Nitrate]]s===<br />
Some nitrates, like [[potassium|potassium nitrate]] and [[sodium nitrate]] are classified as monitored explosive precursors, since they can be used for the production of concentrated nitric acid, [[black powder]] or crude home-made bombs. [[Calcium nitrate]], calcium ammonium nitrate and [[magnesium nitrate]] are also included in the same category (in the EU as per Regulation No 98/2013), while nitrates like [[ammonium nitrate]] and [[urea nitrate]] are classified as explosives. Since all these nitrates are or have been used in the past as fertilizers, they are more tightly monitored than other nitrates. However, many other nitrates may be used to either prepare nitric acid or the previously mentioned nitrates, albeit at a higher cost. [[Copper(II) nitrate]], [[aluminium nitrate]] and [[zinc nitrate]]s may be used to produce nitric acid and alkali nitrates via thermal decomposition and double replacement respectively.<br />
<br />
Iron(III) nitrate has recently being considered as replacement for the potentially explosive ammonium nitrate, either pure iron nitrate or obtained by adding iron(II) sulfate to ammonium nitrate, and since it has a higher nitrogen and oxygen content per mole than other fertilizers, it appears an attractive solution.<ref>https://share-ng.sandia.gov/news/resources/news_releases/ied_fertilizer/</ref> However, it's still possible to obtain alkali nitrates and ammonium nitrate from iron nitrate, by double replacement with an alkali hydroxide (NaOH, KOH and Ca(OH)<sub>2</sub>) and double salt replacement with [[ammonium oxalate]], both easily available or easy to make without restrictions. Granted, the resulting nitrates are impure and wet, and thus require further purification.<br />
<br />
While organic and other explosive nitrates, like [[urea nitrate]] and [[hydrazine nitrate]] are clearly classified as explosive materials, the status for other nitrates is less clear. [[Guanidinium nitrate]], officially classified as rocket fuel and gas generator fuel, it's not always included in the explosive category, or at the very least is indirectly via the definition of what an explosive material is in the respective country/location.<br />
<br />
===[[Nitric acid]]===<br />
In the EU, the sale of nitric acid in concentrations above 3% without a license/permit is forbidden, as per Regulation (EU) No 98/2013. The reason as to why this concentration was chosen is not given, and such diluted acid is unsuitable even for basic chemistry reactions, as one would need very large amounts of 3% nitric acid to obtain any useful amount of reaction product, and extracting the nitrate/nitrating product from such a large volume of water is very intensive and time consuming, even for small amounts, never mind the possibility of hydrolysis of product during the extraction process, which may contaminate the final product, requiring further purification, and thus a lower yield.<br />
<br />
Given how low the 3% limit set by said regulation is, there is no mention of any exceptions or other legal effects for any accidental or side production of a more concentrated acid solution. Hydrolysis of [[copper(II) nitrate]], which can be easily done by simply strongly heating a concentrated solution of said salt will produce lots of nitric acid, and for concentrated solutions, the 3% limit can be easily exceeded. This is also possible for other transition metal nitrates, and even for magnesium nitrate, which decomposes at relative low temperatures to release NO<sub>2</sub> gas which can be used to produce nitric acid. Capturing nitrogen dioxide fumes in water (aka scrubbing a very toxic gas in a liquid), usually from various oxidation reactions or even decomposition or other nitrates/nitrites will produce a diluted nitric acid solution, which can very easily be concentrated above 3% by simply injecting more NO<sub>2</sub> in water.<br />
<br />
===[[Oxalic acid]]===<br />
In UK, oxalic acid is classified as poison and any product containing >10% oxalic acid requires a poison license.<ref>https://www.gov.uk/government/publications/licensing-for-home-users-of-explosives-precursors/licensing-for-home-users-of-poisons-and-explosive-precursors</ref> However, the restriction doesn't apply to oxalate salts or natural sources of oxalates (like rhubarb), even though many soluble oxalates like [[ammonium oxalate]] and [[sodium oxalate]] are just as poisonous as oxalic acid. Also, since oxalate salts aren't banned, it's very easy to make oxalic acid from said salts, just by adding cold concentrated phosphoric or relative concentrated sulfuric acid to a supersaturated solution of oxalate, filter the precipitated phosphate/sulfate and recrystallize the oxalic acid from the filtrate.<br />
<br />
===[[Pentaerythritol tetranitrate]]===<br />
Pentaerythritol tetranitrate (PETN) is a common explosive material used in both military and civilian applications due to its great performances. Although the commerce of PETN explosive materials is strongly regulated, PETN is also used as vasodilator drug to treat certain heart conditions, such as for management of angina. The drug Lentonitrat is described as being "pure PETN", and it is sold in many countries, under different brand names.<ref>https://www.ndrugs.com/?s=lentonitrat</ref>.<br />
<br />
===[[Phosphide]]s===<br />
In most countries, the sale and use of phosphides are well regulated, since contact with water and/or acids will release the highly poisonous and potentially pyrophoric [[phosphine]] (and diphosphane) gas. However, the restriction doesn't appear to apply to all phoshpides the same. In some countries, like UK, aluminium and magnesium phosphides are clearly classified as poisons and a license is required to handle them.<ref>https://www.gov.uk/government/publications/licensing-for-home-users-of-explosives-precursors/licensing-for-home-users-of-poisons-and-explosive-precursors</ref> However, other phoshpides aren't included in the same list, nor does there is any mention of a blanket restriction on all phosphide compounds. Other countries have similar restrictions, though the exact ban on phosphides may be indirect, as in "compounds that may generate poisonous gases in contact with water, acids or bases", which can also refer to other non-phosphide compounds.<br />
<br />
===[[Phosphorus]]===<br />
Except for the US, EU (above 100 g) and a few Asian countries, the sale of red phosphorus is usually not restricted (though regulations on its transport or storage may exist given its flammability hazard). White phosphorus on the other hand is almost always strongly regulated, restricted or outright banned in most countries due to its pyrophoricity/fire hazard and toxicity. White phosphorus however can be easily made from the red variety by simply heating it in a tube, in an oxygen-less atmosphere, meaning that potentially one can store the white form as the red form without any fire risk normally associated with the former.<br />
<br />
Also, if one were to produce elemental phosphorus from phosphates by reducing them with [[carbon]] or metal powders (Al, Mg, etc.) in the presence of silicon dioxide/sand in a kiln, the crude product obtained is white phosphorus, meaning that it's much easier to obtain the more dangerous form of phosphorus than the safer one. This also has the bonus of skipping any obstacles in acquiring the white form, not to mention neither of the mentioned precursors are restricted (though in case of aluminium and magnesium powders some regulations may exist).<br />
<br />
===[[Phosphorus triiodide]]===<br />
Although it's made from two DEA List I chemicals (phosphorus and iodine) and upon hydrolysis (which can also occur by simply leaving the compound in open air), PI<sub>3</sub> releases another List I chemical (hydroiodic acid), phosphorus triiodide is curiously not listed in the DEA List of chemicals. However, PI<sub>3</sub>'s status is covered by the same legislation that covers phosphorus halides and individuals normally cannot purchase it.<br />
<br />
===[[Pinacolyl alcohol]]===<br />
Despite displaying low toxicity, pinacolyl alcohol is included in the List of Schedule 2 substances of the Chemical Weapons Convention, since it's a binary precursor for the nerve agent Soman. Even though the way this class of nerve agents are produced is by simply mixing methylphosphonyl difluoride with an alcohol, pinacolyl alcohol is the only acyclic alcohol included in this category and no other alcohol involved in the production of nerve agents is included (like [[isopropanol]], which is used to make sarin). This is possible because pinacolyl alcohol, unlike isopropanol, has very little uses in chemistry and industry, so basically it "drew the short straw".<br />
<br />
=== [[Potassium permanganate]] ===<br />
In many Western countries, this compound is not always a restricted substance (in US it's DEA List II chemical), but in some former Soviet block countries and the Russian Federation it is included in the lists of drug precursors in because it is sometimes used to make methcathinone. In Russia, it can be sold freely as a low grade chemical, with the concentration of potassium permanganate below 50% (the rest being either inert or [[potassium ferrate]]). Higher grade potassium permanganate is regulated. Obtaining pure potassium permanganate from the impure grade can be easily done via recrystallization from water, though this requires one to purchase relative significant amounts of impure potassium permanganate, which may draw attention.<br />
<br />
The restriction/drug precursor classification applies only to potassium permanganate and sometimes to the rarer and more expensive [[sodium permanganate]], and doesn't cover other permanganates, like [[barium permanganate]], [[calcium permanganate]], [[lithium permanganate]], etc., which can be easily used to prepare potassium permanganate. Unless of course, the legal text clearly specifies other permanganates (e.g. salts of the permanganate acid).<br />
<br />
===[[Prussian blue]]===<br />
Although Prussian blue is not classified as illegal compound, hazardous, poison or even precursor, assuming it's even included in any classification, its use in many common products has declined over the years, [[Chemophobia|since it's "cyanide"]]. As such, overzealous authorities may consider it true cyanide salt/poison and treat it as such, even though the compound is inert to most reagents and even stomach digestion.<br />
<br />
In many places where cyanides are regulated or restricted, there is a mention in the text of the law that ferrocyanides are exempted from the restrictions. However, this doesn't stop overzealous or incompetent authorities from mistaking ferrocyanides with cyanide salts, especially if said persons enforcing the law are poorly trained or corrupt.<br />
<br />
=== [[Sulfuric acid]] ===<br />
In many countries, concentrated sulfuric acid is a controlled substance or otherwise unavailable, and sulfuric acid based drain openers are unavailable as well. Dilute (~30%) sulfuric acid, on the other hand, is often available over the counter as (car) battery acid. Therefore, concentrating this acid in home conditions is required. This resulted in the practice of [[boiling the Bat]].<br />
<br />
In some countries where concentrated sulfuric acid is either monitored or restricted, there doesn't appear to be any regulation over the drain cleaner type, even though there are conc. H<sub>2</sub>SO<sub>4</sub> drain cleaners that can be >95% and have no additives or dye added.<ref>http://richenza.cz/de/professionelle-reinigungsprodukte/371-cleamen-420-sanitaerabfall-1-liter.html</ref> Such products can be used as lab-grade conc. sulfuric acid, while the concentrated form, but with dye added, can be cleaned very easily by just heating the acid to almost near the boiling point of the acid, which causes all organic material to break down.<ref>https://www.youtube.com/watch?v=4DUGRWjdNLI</ref> Any regulations seem to cover mostly lab-grade sulfuric acid, with drain-cleaner types being less monitored.<br />
<br />
Concentrated sulfuric acid is often restricted since it can be used to prepare concentrated nitric acid from a nitrate salt, or anhydrous/concentrated HCl gas or conc. hydrochloric acid (drug precursor). However, [[sodium bisulfate]], cheaply available as pH lowering swimming pool chemical can replace conc. H<sub>2</sub>SO<sub>4</sub> in these cases, though in case of conc. HNO<sub>3</sub>, the resulting acid is not as concentrated as it could be. [[Phosphorus pentoxide]], which is not classified as precursor nor usually restricted (though a bit harder to find and somewhat expensive), can be used as a strong(er) desiccant instead of conc. H<sub>2</sub>SO<sub>4</sub> and can even be used to prepare a more concentrated acid. Likewise, one can obtain sulfur trioxide from thermal decomposition of sodium bisulfate (cheaper) or by dehydrating conc. sulfuric acid with P<sub>2</sub>O<sub>5</sub> (albeit wasteful and expensive).<br />
<br />
==Chemistry equipment==<br />
===[[Erlenmeyer flask]] (CORRECTED)===<br />
In a ridiculous decision to limit drug manufacturing, Texas has restricted the sale of Erlenmeyer flasks and a permit is required to purchase them.<ref>https://www.dps.texas.gov/RSD/Precursor/Laws/index.htm</ref> The restriction does not appear to cover flasks with a similar function and appearance, like [[Florence flask]]s, [[fleaker]]s or [[Büchner flask|side-arm flask]]s.<br />
<br />
As of late 2019, there is no more need for a permit and it appears that restrictions on glassware have been loosened.<ref>https://legiscan.com/TX/text/SB616/2019</ref><br />
<br />
===[[Heating mantle]]===<br />
In the United States, the sale of all heating mantles with a volume of 22 liters (5.8 US gallons) are monitored, as these devices are placed in the DEA Special Surveillance List, in the equipment section. There are no mentions if heating mantles with a volume smaller or larger are also included nor is any explanation given as to why this exact volume was selected for special surveillance by the DEA.<ref>https://web.archive.org/web/20110420054619/http://www.deadiversion.usdoj.gov/chem_prog/advisories/surveillance.htm</ref><br />
<br />
==Other==<br />
===1986 California Proposition 65===<br />
Proposition 65 (formally titled The Safe Drinking Water and Toxic Enforcement Act of 1986) is a California law passed by direct voter initiative in 1986 by a 63%–37% vote. Its goals are to protect drinking water sources from toxic substances that cause cancer and birth defects and to reduce or eliminate exposures to those chemicals generally, such as consumer products, by requiring warnings in advance of those exposures. The warning label uses the following phrase or a slight variation of it:<br />
<br />
: WARNING: This product contains chemicals known to the State of California to cause cancer and birth defects or other reproductive harm.<br />
<br />
The legislation includes a list of various chemicals that are considered to be suspicious of causing cancer, and thus any product containing them contains a label with a warning.<br />
<br />
The law has been criticized for causing "over-warning" or "meaningless warnings", and this risk has even been recognized by a California court.<ref>http://ag.ca.gov/prop65/pdfs/G035101.pdf</ref><ref>http://www.law.com/?id=1144672792347&slreturn=20200203132953</ref> There is no penalty for posting an unnecessary warning sign and to the extent that warnings are vague or overused, they may not communicate much information to the end user. As such, a recurring joke that everything in California causes cancer was born, which makes fun of the nanny-state-like regulation.<br />
<br />
While the list of suspected chemicals does include substances that are actually harmful and confirmed to cause cancer (acrylamide, aflatoxins, [[asbestos]], nitrosamines, etc.), many many chemicals included on the list aren't even suspected of causing cancer (aloe vera extract), usually either due to containing minute traces of a compound that is suspected or shown to cause cancer in lab rats at high or very high doses (like the aloin from aloe vera) or just suspected from non-reproducible tests. Some poisonous chemicals that don't leave long-lasting injuries if treated immediately ([[hydrogen cyanide]], [[carbon monoxide]], etc.) are also included, even though the evidence to suggest they may cause cancer is not clear ([[nitrous oxide]], [[sulfur dioxide]], e.g.) or has not been reproduced independently. Many chemicals included on the list have not been confirmed to cause cancer in labs or have been used in treating some forms of cancer, either in research or medically ([[dichloroacetic acid]], Fluorouracil, [[lithium carbonate]]/citrate). Some of the mentions on the list include odd descriptions, like in the case of [[ethylene glycol]] "when ingested", without a clear indication why only ingestion is considered to cause cancer risk and not inhalation or skin permeation, or [[formaldehyde]] as "gas", with no mention of the more common available form (solution, formalin), for [[vanadium pentoxide]] "only orthorhombic" with no other vanadium compounds enumerated. Hormones are also included on the list with a special mention ("synthetic"), but this does not apply to all (see testosterone). Wood dust is also included in the list. Many chemicals on the list were eventually removed, even though allyl chloride, is classified as Group C (a possible human carcinogen) by the EPA, while Chloramphenicol is known to affect the bone marrow and increase the risk of leukemia. Some chemicals are also being ping-ponged, like Bisphenol A (BPA), which was removed from the list on April 19, 2013, and was relisted on May 11, 2015. <ref>https://oehha.ca.gov/proposition-65/proposition-65-list</ref><ref>https://en.wikipedia.org/wiki/California_Proposition_65_list_of_chemicals</ref><br />
<br />
===Distillation===<br />
In most countries, as well as US, it is not required to posses a permit for owning a distillation still or apparatus. It is also permitted to distill almost any liquid, like water, essential oils, hydrocarbon solvents, etc. While it's legal to distill alcohols like methanol, propanol, iropropanol, butanol, etc., distillation of ethanol requires a distillation permit. However, while this is true for distilling food-grade ethanol, distilling denatured ethanol is a complicated matter, as in some countries there is no distinction between the distillation of food-grade ethanol and the denatured variety, while in others there is. In the US, for distilling ethanol fuel, a Federal Fuel Alcohol Permit is required. This permit however, does not cover the production of food-grade ethanol.<ref>https://www.ttb.gov/industrial/alcohol-fuel</ref> For distilling consumable alcohol, a Federal Distilled Spirits Permit is necessary.<ref>https://www.ttb.gov/spirits/spirits-permits.shtml</ref> Neither permits can be used to distill the other type of ethanol, so if you need to distill non-food grade ethanol and you have the Federal Distilled Spirits Permit, you will have to apply for the other permit to legally distill the technical alcohol, and if you want to distill both types, you will need both type of permits. It's unclear what happens if the resulting food-grade distilled ethanol ends up unfit for consumption due to an accident during the distillation or intentional denaturation after, as when that happens, the ethanol can only be used as as solvent or as fuel. However, since the distillation permit only covers the distillation activity for which the permit was released after filing the necessary paperwork and passing whatever inspection or certification is required, failed distillations should be classified as business loss and not penalized.<br />
<br />
In many countries around the world, there are legal rules regarding the use and ownership of alcoholic beverage distillation equipment, which do not cover other form of distillation equipment, mainly since such items are not certified for alcohol distillation. As such, the following alternative distillation equipment are classified as:<br />
<br />
*Chemistry distillation glassware parts, either acquired as separate pieces (boiling flask, condenser, etc.) or as full distillation kit are legal to acquire, posses and use without restrictions, though you may not legally distill food-grade ethanol.<br />
<br />
*[[Rotary evaporator|Rotavaps]] are legal to own and use, but you may not distill ethanol with it.<br />
<br />
*Although commercial water distillers can be used for distilling alcohol, it is still not legal to do so, since it's the distillation activity that's regulated by the law, and not a specific distillation apparatus. <strike>Not that anyone will ask...</strike><br />
<br />
*A supersimple distillation still can be made by simply taking a large pot, placing the desired liquid inside, then placing a jar or flask in the middle of the pot, covering the pot with its lid turned upside down. The resulting liquid condenses on the inside of the lid (works better if you add some ice on its outside curve), which then pours in the receiving flask.<ref>https://www.youtube.com/watch?v=ZZyzpDJtK5s</ref> Since this isn't a distillation apparatus in its official sense, and it's basically just a simple hack of a typical kitchen pot (which works by refluxing its volatile liquids), it's unclear how this would be legally classified, though since you can collect the resulting distilled liquid, it may be still classified as distillation. Also, there is no mention of any legal restrictions if said "distillation" still is left outside in the strong sunlight, which due to heat from the sun will cause the alcohol to condense inside and collect in the flask. Any restrictions on this route would involve the legal text to mention "any heating source" or "distillation time" (if there is one) in some way, to eliminate any ambiguities.<br />
<br />
===Hazardous chemicals===<br />
Many countries will classify various chemicals that are proven to cause harm as "hazardous" and may even be restricted if they're deemed too harmful. However, the exact term "hazardous" is itself poorly defined, and varies from country to country. While some definitions include phrases like "proven to cause immediate harm" or "recognized to be dangerous to health and environment", many definitions will rely on various lists of reagents deemed hazardous, like the EU's REACH list or chemicals. This however, leads to the problem of following the law too literal and only considering the substances that are included on the list as being hazardous, and often tends to ignore substances that are very similar, yet similar or even more harmful than the listed substance. Occasionally it might lead to cases of ignoring obscure harmful substances from being identified since they didn't raise certain red flags at a glance.<br />
<br />
Expired reagents, even those that are non-toxic substances, are often lumped together as hazardous chemicals, mainly since lab wastes are all treated as hazardous, even when their only "crime" is that they passed their expiration date. Depending on the country, this may also apply to reagents that are also available as OTC products and don't need special disposal and standard solutions of said reagents. This will sometimes lead to tragicomical situations where chemical disposal teams are called in with full protection gear, police arriving at the site and area evacuated just to remove some expired sodium chloride, copper sulfate, glutamic acid, etc. or some bottle of a diluted solution with "scary symbols" which turns out to be just some old standard solution. Sometimes even the bomb squad are called in. The end result of this is lots of taxpayer money being wasted and pointless scare. Sometimes, not even an expert is called in to verify if there are actual hazardous or highly poisonous substances at the location, or sometimes just an "expert" is called who doesn't seem to put much effort into verifying the content of the location. More often than not, such behavior arises from the legislation regarding chemicals, which is not always very clear and the people who enforce it receive very little training if any.<br />
<br />
===Reagent expiration===<br />
Most reagents have noted an expiration date on their label, which is often encountered on many organic reagents, and even on inorganic ones. This is because many chemical compounds will undergo chemical changes over time, either on their own, or after prolonged air/sunlight exposure, or may absorb moisture or carbon dioxide from air. The role of the expiration date is to provide a reasonable time period in which the reagent can be used safely without causing anomalies during the chemical reaction or even spoiling the reaction altogether. However, many reagents that normally do not suffer any changes in most storage conditions (like most common salts: NaCl, K<sub>2</sub>SO<sub>4</sub>, EDTA, etc.) also come with an expiration date, even though they should remain practically unchanged from when they were added in the bottle in the supplier's warehouse. While it's possible for some reagents to absorb water from air and thus become damp over time, this is more likely to happen if the reagent is not kept in proper storage conditions (after all, one can keep said reagent bottle in a desiccator far longer than its expiration date and it won't absorb any moisture and thus remain unchanged, while keeping it a few weeks in a very damp place is sufficient to turn the reagent into a mushy paste or a solid block). Once a reagent has passed its expiration date, it can no longer be used in professional labs, even though it didn't suffer any changes or wasn't contaminated, and even if it wasn't open since it was purchased. This is very much true for analytical reagents, which can no longer be used in chemical analysis once they've reached the expiration date and have to be disposed of, even if they weren't open or have shown any signs of degradation/contamination. This leads to a hoarding problem, since the expired reagents, even though they're still technically good, can no longer be used (at least for that purpose), so if you used, for example, small amounts of reagent from a 1 kg bottle, and after the expiration date you still have >900 g of unused reagent, you end up discarding much more than you used, which translates in lost money and resources.<br />
<br />
For reagents used in synthesis, this can be avoided, by simply recrystallizing the salt from a concentrated solution, or for liquids a simple distillation is sufficient to generate a "fresh" reagent, since the resulting product is no longer the old one, and it's usually more pure. However, while this method is widely used by "poorer" labs, in some circles it's looked down upon, since there are concerns about contamination during the preparation of new product. Likewise, most reagent labels show the value of traces present in the product, which, in the event of a recrystallization or distillation, will change. While this isn't a problem for reagents used in synthesis, this practice is not allowed for reagents used in analysis, since, while one would have to reanalyze the traces present in the new product, an official certification is required for the newly produced analytical reagent to be valid.<br />
<br />
===Waste disposal===<br />
Most regulations forbid the release of lab-grade reagents down the drain or in the municipal sewage system. The laws are written in such way, that even dumping certain harmless lab-grade reagents, like acetic, citric, tartaric acids and their salts, sodium hydroxide, various alcohols and esters, that not only encountered as OTC products but are also disposed of by being poured down the drain, may be considered illegal. However, so far nobody has been prosecuted for dumping chemical compounds whose OTC equivalent can be poured down the drain (only a few cases of certain academics being prosecuted for dumping expired lab-grade solvents down the drain have been known<ref>https://archive.seattletimes.com/archive/?date=20070308&slug=uwprof08m0</ref><ref>https://blogs.sciencemag.org/pipeline/archives/2007/03/08/how_not_to_do_it_more_diethyl_ether_now_with_extra_hardware</ref>), so there is no clear answer (so far) what is legal to pour down the drain and what is not, especially since the legal interpretations of waste disposal differ from country to country.<br />
<br />
While pouring volatile chemicals down the drain is not legally allowed, there are few regulations regarding allowing evaporation of said reagents in open air. While air quality monitoring is present in many cities and administrations, it's only when the concentration of a certain chemical species increases above a certain limit that there will be legal problems. As such, one may allow a small amount of volatile liquid on a tray on the roof during the day to allow maximum evaporation. The vapors will rise in the air and eventually break down under the influence of UV light and oxygen to less harmful or harmless products. Granted this method is not a great idea for disposing of foul-smelling chemicals, especially for those that can be sensed at very low concentrations.<br />
<br />
Most countries forbid throwing expired medication down the drain or even in trash, though unless said medication becomes toxic after expiring, there are ''other'' ways of disposing of said medication without directly breaking the law. However, this is not necessary, as pharmacies around the world will collect expired medication for free, which then get sent to proper hazardous waste disposal facilities.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
<br />
[[Category:Legal and societal issues]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sodium_persulfate&diff=14950Sodium persulfate2023-03-24T08:55:04Z<p>Ave369: /* Chemical */</p>
<hr />
<div>{{Chembox<br />
| Name = Sodium persulfate<br />
| Reference =<br />
| IUPACName = Sodium peroxydisulfate<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Sodium peroxodisulfate<br>Sodium peroxodisulphate<br>Sodium peroxydisulphate<br />
<!-- Images --><br />
| ImageFile = Sodium peroxydisulfate sample in watchglass.jpg<br />
| ImageSize = 300<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Sodium persulfate sample<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 2.59 g/cm<sup>3</sup><br>(Loose bulk density: 1.12 g/cm<sup>3</sup>)<br />
| Formula = Na<sub>2</sub>S<sub>2</sub>O<sub>8</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 238.10 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 180<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = (decomposes)<br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 55.6 g/100 ml (20 °C)<br />
| SolubleOther = <br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [https://www.docdroid.net/SfeHLp8/sodium-persulfate-sa.pdf Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = 920 mg/kg (rat, female, oral)<br>930 mg/kg (rat, male, oral)<br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Ammonium persulfate]]<br>[[Potassium persulfate]]<br />
}}<br />
}}<br />
'''Sodium persulfate''' is the chemical compound with the formula '''Na<sub>2</sub>S<sub>2</sub>O<sub>8</sub>''', the sodium salt of [[peroxydisulfuric acid]]. It is almost non-[[hygroscopy|hygroscopic]] and has good shelf-life.<br />
<br />
Sodium persulfate is preferred by many chemists over other persulfate salts, as it has good solubility in water and long shelf-life, while also having better performance than other persulfate salts.<br />
<br />
==Properties==<br />
===Chemical===<br />
Sodium persulfate is an oxidizer. It is a common oxidizing agent in Elbs persulfate oxidation or Boyland–Sims oxidation.<br />
<br />
Heating it in solution to 80-90 Celsius for 2-3 minutes in presence of MnO<sub>2</sub> leads to a decomposition reaction catalyzed by manganese dioxide:<br />
<br />
:2 Na<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + 2H<sub>2</sub>O → 4 NaHSO<sub>4</sub> + O<sub>2</sub><br />
<br />
This reaction can be used as a qualitative test for the peroxodisulfate anion: just heat it with MnO<sub>2</sub> and check the pH with a test strip. Low pH means positive test.<br />
<br />
===Physical===<br />
Sodium persulfate is a white solid, soluble in water.<br />
<br />
==Availability==<br />
Sodium persulfate can be bought from various electronic shops, as etching powder for zinc and printed circuit boards, as well as for pickling of copper and some other metals.<br />
<br />
==Preparation==<br />
The salt is prepared by the electrolytic oxidation of [[sodium bisulfate]]:<br />
<br />
:2 NaHSO<sub>4</sub> → Na<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + H<sub>2</sub><br />
<br />
Oxidation is conducted at a [[platinum]] anode.<br />
<br />
==Projects==<br />
*PCB etchant<br />
*Make permanganates<br />
*Make [[benzaldehyde]] from [[benzyl alcohol]]<br />
*Elbs persulfate oxidation<br />
*Boyland–Sims oxidation<br />
<br />
==Handling==<br />
===Safety===<br />
Sodium persulfate is an oxidizer and forms combustible mixtures with organic materials such as paper.<br />
<br />
===Storage===<br />
In closed plastic or glass bottles.<br />
<br />
===Disposal===<br />
Can be neutralized by heating it, which causes it to decompose.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=10137 Sodium Persulphate from Ammonium Persulphate]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=17732 Preparation of Sodium persulfate]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=8113 Preparing Na2S2O8]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=30694 Would Persulfate adequately oxidize bromide?]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Sodium compounds]]<br />
[[Category:Persulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Materials that react with water]]<br />
[[Category:PCB etchants]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Potassium_persulfate&diff=14949Potassium persulfate2023-03-24T08:53:49Z<p>Ave369: /* Chemical */</p>
<hr />
<div>{{Chembox<br />
| Name = Potassium persulfate<br />
| Reference =<br />
| IUPACName =<br />
| PIN =<br />
| SystematicName = Potassium peroxydisulfate<br />
| OtherNames = Dipotassium peroxodisulfate<br>Potassium perdisulfate<br>Potassium peroxydisulfate<br />
<!-- Images --><br />
| ImageFile = Potassium peroxydisulfate bottle and sample.jpg<br />
| ImageSize = 300<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Potassium persulfate sample and original bottle<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = [O-]S(=O)(=O)OOS(=O)(=O)[O-].[K+].[K+]<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 2.477<br />
| Formula = K<sub>2</sub>S<sub>2</sub>O<sub>8</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 270.322 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 186-250<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = (decomposes)<br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 1.75 g/100 mL (0 °C)<br>4.49 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in conc. [[hydrogen peroxide]]<ref>Skogareva; Ippolitov; Russian Journal of Inorganic Chemistry; vol. 43; nb. 11; (1998); p. 1668 - 1670</ref>, [[hydrogen fluoride]]<ref>Fredenhagen, H.; Z. Anorg. Chem.; vol. 242; (1939); p. 23 - 32</ref><br>Insoluble in [[benzene]], [[chloroform]], [[ethanol]], [[methanol]]<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Triclinic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = 1,697.30 kJ/mol (at 25 °C)<br />
| DeltaHc = <br />
| DeltaHf = -1,916 kJ/mol<ref>Balej, Jan; Zeitschrift fur Physikalische Chemie; vol. 224; nb. 6; (2010); p. 883 - 892</ref><br />
| Entropy = <br />
| HeatCapacity = -84.0±3 J·mol<sup>-1</sup>·K<sup>-1</sup> (calculated, at 25 °C)<br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Non-explosive<br />
| FrictionSens = Non-explosive<br />
| DetonationV = Non-explosive<br />
| REFactor = Non-explosive<br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [http://www.sciencelab.com/msds.php?msdsId=9927234 ScienceLab]<br />
| FlashPt = Non-flammable<br />
| LD50 = 802 mg/kg (oral, rat)<br />
| LC50 = <br />
| MainHazards = Oxidant<br>Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = [[Sodium persulfate]]<br>[[Ammonium persulfate]]<br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium peroxymonosulfate]]<br>[[Sodium persulfate]]<br />
}}<br />
}}<br />
'''Potassium persulfate''', also known as '''potassium peroxydisulfate''' is a chemical compound used as an oxidizing agent, with the formula '''K<sub>2</sub>S<sub>2</sub>O<sub>8</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
When potassium persulfate is heated in a 50% aqueous solution of [[sulfuric acid]], [[hydrogen peroxide]] results, which, due to the high temperature, distills from the solution. The H<sub>2</sub>O<sub>2</sub> obtained this way has a concentration of 40-60%. This method was previously used in the manufacturing of hydrogen peroxide on industrial scale before being replaced by the quinone process.<br />
<br />
Potassium persulfate will react with [[silver nitrate]] to form [[silver(I,III) oxide]] (silver peroxide):<ref>Marshall, H.; Journal of the Chemical Society; vol. 59; (1891); p. 775</ref><br />
:K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + AgNO<sub>3</sub> → Ag<sub>4</sub>O<sub>4</sub> + K<sub>2</sub>SO<sub>4</sub> + SO<sub>x</sub> + NO<sub>x</sub> + O<sub>x</sub><br />
<br />
Potassium peroxydisulfate oxidizes [[acetone]] in the presence of diluted [[sulfuric acid]] and [[silver]] metal to [[acetic acid]], releasing carbon dioxide:<ref>Broensted, J. N.; Z. phys. Ch.; vol. 102; (1922); p. 169 - 207</ref><br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + (CH<sub>3</sub>)<sub>2</sub>CO → CH<sub>3</sub>COOH + KHSO<sub>4</sub> + CO<sub>2</sub><br />
<br />
Reaction with [[nitric acid]] gives off [[oxygen]] and [[ozone]] fumes, and nitrogen as byproduct.<ref>Baeyer, A.; Villiger, V.; Ber.; vol. 34; (1901); p. 853 - 862</ref><br />
<br />
Heating it in solution or slush to 80-90 Celsius for 2-3 minutes in presence of MnO<sub>2</sub> leads to a decomposition reaction catalyzed by manganese dioxide:<br />
<br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + 2H<sub>2</sub>O → 4 KHSO<sub>4</sub> + O<sub>2</sub><br />
<br />
This reaction can be used as a qualitative test for the peroxodisulfate anion: just heat it with MnO<sub>2</sub> and check the pH with a test strip. Low pH means positive test.<br />
<br />
===Physical===<br />
Potassium persulfate is a white crystalline solid, poorly soluble in water. It decomposes if heated to temperatures over 125°C. It has a density of 2.477 g/cm<sup>3</sup>.<br />
<br />
==Availability==<br />
Potassium persulfate can be bought from chemical suppliers and eBay.<br />
<br />
==Preparation==<br />
The most common way to synthesize potassium persulfate is via the [[electrolysis]] of a cold solution [[potassium bisulfate]] in [[sulfuric acid]], at a high current density<ref>F. Feher, "Potassium Peroxydisulfate", Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 392.</ref>:<br />
: 2 KHSO<sub>4</sub> → K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + H<sub>2</sub><br />
<br />
[[Tantalum]] electrodes can be used in this reaction.<ref>Gmelin, vol. O: MVol.2; 46, page 276 - 278</ref><br />
<br />
Bubbling elemental [[fluorine]] through an aqueous solution of KHSO<sub>4</sub> or K<sub>2</sub>SO<sub>4</sub> will also yield potassium peroxydisulfate. The reaction also works in the absence of water.<ref>Fichter, F.; Humpert, K.; Helvetica Chimica Acta; vol. 9; (1926); p. 521 - 525</ref><ref>Mueller, E.; Zeitschrift fuer Elektrochemie; vol. 10; (1904); p. 753 - 776</ref><br />
<br />
It can also be prepared by adding KHSO<sub>4</sub> to a solution of the more soluble salt [[Ammonium persulfate|ammonium peroxydisulfate]]. Potassium persulfate will precipitate from this reaction.<br />
<br />
==Projects==<br />
*Make silver peroxide<br />
*Make hydrogen peroxide<br />
*PCB etchant<br />
<br />
==Handling==<br />
===Safety===<br />
Potassium peroxydisulfate is a strong oxidizer and should be handled with care.<br />
<br />
Potassium persulfate will slowly decompose in wet environment.<br />
<br />
===Storage===<br />
Potassium persulfate should be stored in closed bottles, away from any acidic vapors or organic materials.<br />
<br />
If kept in a dark and dry environment, it will remain stable for years.<ref>Elbs, K.; Neher, P.; Chemiker-Zeitung; vol. 45; (1921); p. 1113 - 1114</ref><br />
<br />
===Disposal===<br />
Heating the salt will cause it to decompose.<br />
<br />
==Gallery==<br />
<gallery widths="250" position="center" columns="4" orientation="none"><br />
Potassium persulfate sample.jpg|A sample of potassium persulfate<br />
</gallery><br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=11022 Persulfate cell - Please translate passage (German)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13955 Potassium Persulfate for SO3 synth?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16912 Potassium Persulfate Oxidizer]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=41198 Potassium Persulfate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Potassium compounds]]<br />
[[Category:Persulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Materials that react with water]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Potassium_persulfate&diff=14948Potassium persulfate2023-03-24T08:52:34Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Potassium persulfate<br />
| Reference =<br />
| IUPACName =<br />
| PIN =<br />
| SystematicName = Potassium peroxydisulfate<br />
| OtherNames = Dipotassium peroxodisulfate<br>Potassium perdisulfate<br>Potassium peroxydisulfate<br />
<!-- Images --><br />
| ImageFile = Potassium peroxydisulfate bottle and sample.jpg<br />
| ImageSize = 300<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Potassium persulfate sample and original bottle<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = [O-]S(=O)(=O)OOS(=O)(=O)[O-].[K+].[K+]<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 2.477<br />
| Formula = K<sub>2</sub>S<sub>2</sub>O<sub>8</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 270.322 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 186-250<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = (decomposes)<br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 1.75 g/100 mL (0 °C)<br>4.49 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in conc. [[hydrogen peroxide]]<ref>Skogareva; Ippolitov; Russian Journal of Inorganic Chemistry; vol. 43; nb. 11; (1998); p. 1668 - 1670</ref>, [[hydrogen fluoride]]<ref>Fredenhagen, H.; Z. Anorg. Chem.; vol. 242; (1939); p. 23 - 32</ref><br>Insoluble in [[benzene]], [[chloroform]], [[ethanol]], [[methanol]]<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Triclinic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = 1,697.30 kJ/mol (at 25 °C)<br />
| DeltaHc = <br />
| DeltaHf = -1,916 kJ/mol<ref>Balej, Jan; Zeitschrift fur Physikalische Chemie; vol. 224; nb. 6; (2010); p. 883 - 892</ref><br />
| Entropy = <br />
| HeatCapacity = -84.0±3 J·mol<sup>-1</sup>·K<sup>-1</sup> (calculated, at 25 °C)<br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Non-explosive<br />
| FrictionSens = Non-explosive<br />
| DetonationV = Non-explosive<br />
| REFactor = Non-explosive<br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [http://www.sciencelab.com/msds.php?msdsId=9927234 ScienceLab]<br />
| FlashPt = Non-flammable<br />
| LD50 = 802 mg/kg (oral, rat)<br />
| LC50 = <br />
| MainHazards = Oxidant<br>Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = [[Sodium persulfate]]<br>[[Ammonium persulfate]]<br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium peroxymonosulfate]]<br>[[Sodium persulfate]]<br />
}}<br />
}}<br />
'''Potassium persulfate''', also known as '''potassium peroxydisulfate''' is a chemical compound used as an oxidizing agent, with the formula '''K<sub>2</sub>S<sub>2</sub>O<sub>8</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
When potassium persulfate is heated in a 50% aqueous solution of [[sulfuric acid]], [[hydrogen peroxide]] results, which, due to the high temperature, distills from the solution. The H<sub>2</sub>O<sub>2</sub> obtained this way has a concentration of 40-60%. This method was previously used in the manufacturing of hydrogen peroxide on industrial scale before being replaced by the quinone process.<br />
<br />
Potassium persulfate will react with [[silver nitrate]] to form [[silver(I,III) oxide]] (silver peroxide):<ref>Marshall, H.; Journal of the Chemical Society; vol. 59; (1891); p. 775</ref><br />
:K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + AgNO<sub>3</sub> → Ag<sub>4</sub>O<sub>4</sub> + K<sub>2</sub>SO<sub>4</sub> + SO<sub>x</sub> + NO<sub>x</sub> + O<sub>x</sub><br />
<br />
Potassium peroxydisulfate oxidizes [[acetone]] in the presence of diluted [[sulfuric acid]] and [[silver]] metal to [[acetic acid]], releasing carbon dioxide:<ref>Broensted, J. N.; Z. phys. Ch.; vol. 102; (1922); p. 169 - 207</ref><br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + (CH<sub>3</sub>)<sub>2</sub>CO → CH<sub>3</sub>COOH + KHSO<sub>4</sub> + CO<sub>2</sub><br />
<br />
Reaction with [[nitric acid]] gives off [[oxygen]] and [[ozone]] fumes, and nitrogen as byproduct.<ref>Baeyer, A.; Villiger, V.; Ber.; vol. 34; (1901); p. 853 - 862</ref><br />
<br />
Heating it in solution or slush to 80-90 Celsius for 2-3 minutes in presence of MnO<sub>2</sub> leads to a decomposition reaction catalyzed by manganese dioxide:<br />
<br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + 2H<sub>2</sub>O → 4 KHSO<sub>4</sub> + O<sub>2</sub><br />
<br />
This reaction can be used as a qualitative test for the peroxodisulfate anion: just heat it with MnO<sub>2</sub> and check the pH with a test strip.<br />
<br />
===Physical===<br />
Potassium persulfate is a white crystalline solid, poorly soluble in water. It decomposes if heated to temperatures over 125°C. It has a density of 2.477 g/cm<sup>3</sup>.<br />
<br />
==Availability==<br />
Potassium persulfate can be bought from chemical suppliers and eBay.<br />
<br />
==Preparation==<br />
The most common way to synthesize potassium persulfate is via the [[electrolysis]] of a cold solution [[potassium bisulfate]] in [[sulfuric acid]], at a high current density<ref>F. Feher, "Potassium Peroxydisulfate", Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 392.</ref>:<br />
: 2 KHSO<sub>4</sub> → K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + H<sub>2</sub><br />
<br />
[[Tantalum]] electrodes can be used in this reaction.<ref>Gmelin, vol. O: MVol.2; 46, page 276 - 278</ref><br />
<br />
Bubbling elemental [[fluorine]] through an aqueous solution of KHSO<sub>4</sub> or K<sub>2</sub>SO<sub>4</sub> will also yield potassium peroxydisulfate. The reaction also works in the absence of water.<ref>Fichter, F.; Humpert, K.; Helvetica Chimica Acta; vol. 9; (1926); p. 521 - 525</ref><ref>Mueller, E.; Zeitschrift fuer Elektrochemie; vol. 10; (1904); p. 753 - 776</ref><br />
<br />
It can also be prepared by adding KHSO<sub>4</sub> to a solution of the more soluble salt [[Ammonium persulfate|ammonium peroxydisulfate]]. Potassium persulfate will precipitate from this reaction.<br />
<br />
==Projects==<br />
*Make silver peroxide<br />
*Make hydrogen peroxide<br />
*PCB etchant<br />
<br />
==Handling==<br />
===Safety===<br />
Potassium peroxydisulfate is a strong oxidizer and should be handled with care.<br />
<br />
Potassium persulfate will slowly decompose in wet environment.<br />
<br />
===Storage===<br />
Potassium persulfate should be stored in closed bottles, away from any acidic vapors or organic materials.<br />
<br />
If kept in a dark and dry environment, it will remain stable for years.<ref>Elbs, K.; Neher, P.; Chemiker-Zeitung; vol. 45; (1921); p. 1113 - 1114</ref><br />
<br />
===Disposal===<br />
Heating the salt will cause it to decompose.<br />
<br />
==Gallery==<br />
<gallery widths="250" position="center" columns="4" orientation="none"><br />
Potassium persulfate sample.jpg|A sample of potassium persulfate<br />
</gallery><br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=11022 Persulfate cell - Please translate passage (German)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13955 Potassium Persulfate for SO3 synth?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16912 Potassium Persulfate Oxidizer]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=41198 Potassium Persulfate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Potassium compounds]]<br />
[[Category:Persulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Materials that react with water]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Potassium_persulfate&diff=14947Potassium persulfate2023-03-24T08:51:20Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Potassium persulfate<br />
| Reference =<br />
| IUPACName =<br />
| PIN =<br />
| SystematicName = Potassium peroxydisulfate<br />
| OtherNames = Dipotassium peroxodisulfate<br>Potassium perdisulfate<br>Potassium peroxydisulfate<br />
<!-- Images --><br />
| ImageFile = Potassium peroxydisulfate bottle and sample.jpg<br />
| ImageSize = 300<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Potassium persulfate sample and original bottle<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = [O-]S(=O)(=O)OOS(=O)(=O)[O-].[K+].[K+]<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 2.477<br />
| Formula = K<sub>2</sub>S<sub>2</sub>O<sub>8</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 270.322 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 186-250<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = (decomposes)<br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 1.75 g/100 mL (0 °C)<br>4.49 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in conc. [[hydrogen peroxide]]<ref>Skogareva; Ippolitov; Russian Journal of Inorganic Chemistry; vol. 43; nb. 11; (1998); p. 1668 - 1670</ref>, [[hydrogen fluoride]]<ref>Fredenhagen, H.; Z. Anorg. Chem.; vol. 242; (1939); p. 23 - 32</ref><br>Insoluble in [[benzene]], [[chloroform]], [[ethanol]], [[methanol]]<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Triclinic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = 1,697.30 kJ/mol (at 25 °C)<br />
| DeltaHc = <br />
| DeltaHf = -1,916 kJ/mol<ref>Balej, Jan; Zeitschrift fur Physikalische Chemie; vol. 224; nb. 6; (2010); p. 883 - 892</ref><br />
| Entropy = <br />
| HeatCapacity = -84.0±3 J·mol<sup>-1</sup>·K<sup>-1</sup> (calculated, at 25 °C)<br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Non-explosive<br />
| FrictionSens = Non-explosive<br />
| DetonationV = Non-explosive<br />
| REFactor = Non-explosive<br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [http://www.sciencelab.com/msds.php?msdsId=9927234 ScienceLab]<br />
| FlashPt = Non-flammable<br />
| LD50 = 802 mg/kg (oral, rat)<br />
| LC50 = <br />
| MainHazards = Oxidant<br>Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = [[Sodium persulfate]]<br>[[Ammonium persulfate]]<br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium peroxymonosulfate]]<br>[[Sodium persulfate]]<br />
}}<br />
}}<br />
'''Potassium persulfate''', also known as '''potassium peroxydisulfate''' is a chemical compound used as an oxidizing agent, with the formula '''K<sub>2</sub>S<sub>2</sub>O<sub>8</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
When potassium persulfate is heated in a 50% aqueous solution of [[sulfuric acid]], [[hydrogen peroxide]] results, which, due to the high temperature, distills from the solution. The H<sub>2</sub>O<sub>2</sub> obtained this way has a concentration of 40-60%. This method was previously used in the manufacturing of hydrogen peroxide on industrial scale before being replaced by the quinone process.<br />
<br />
Potassium persulfate will react with [[silver nitrate]] to form [[silver(I,III) oxide]] (silver peroxide):<ref>Marshall, H.; Journal of the Chemical Society; vol. 59; (1891); p. 775</ref><br />
:K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + AgNO<sub>3</sub> → Ag<sub>4</sub>O<sub>4</sub> + K<sub>2</sub>SO<sub>4</sub> + SO<sub>x</sub> + NO<sub>x</sub> + O<sub>x</sub><br />
<br />
Potassium peroxydisulfate oxidizes [[acetone]] in the presence of diluted [[sulfuric acid]] and [[silver]] metal to [[acetic acid]], releasing carbon dioxide:<ref>Broensted, J. N.; Z. phys. Ch.; vol. 102; (1922); p. 169 - 207</ref><br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + (CH<sub>3</sub>)<sub>2</sub>CO → CH<sub>3</sub>COOH + KHSO<sub>4</sub> + CO<sub>2</sub><br />
<br />
Reaction with [[nitric acid]] gives off [[oxygen]] and [[ozone]] fumes, and nitrogen as byproduct.<ref>Baeyer, A.; Villiger, V.; Ber.; vol. 34; (1901); p. 853 - 862</ref><br />
<br />
Heating it in solution or slush to 80-90 Celsius for 2-3 minutes in presence of MnO<sub>2</sub> leads to a decomposition reaction catalyzed by manganese dioxide:<br />
<br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + 2H<sub>2</sub>O → 2 K<sub>2</sub>SO<sub>4</sub> + 2H<sub>2</sub>SO<sub>4</sub> + O<sub>2</sub><br />
<br />
This reaction can be used as a qualitative test for the peroxodisulfate anion: just heat it with MnO<sub>2</sub> and check the pH with a test strip.<br />
<br />
===Physical===<br />
Potassium persulfate is a white crystalline solid, poorly soluble in water. It decomposes if heated to temperatures over 125°C. It has a density of 2.477 g/cm<sup>3</sup>.<br />
<br />
==Availability==<br />
Potassium persulfate can be bought from chemical suppliers and eBay.<br />
<br />
==Preparation==<br />
The most common way to synthesize potassium persulfate is via the [[electrolysis]] of a cold solution [[potassium bisulfate]] in [[sulfuric acid]], at a high current density<ref>F. Feher, "Potassium Peroxydisulfate", Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 392.</ref>:<br />
: 2 KHSO<sub>4</sub> → K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + H<sub>2</sub><br />
<br />
[[Tantalum]] electrodes can be used in this reaction.<ref>Gmelin, vol. O: MVol.2; 46, page 276 - 278</ref><br />
<br />
Bubbling elemental [[fluorine]] through an aqueous solution of KHSO<sub>4</sub> or K<sub>2</sub>SO<sub>4</sub> will also yield potassium peroxydisulfate. The reaction also works in the absence of water.<ref>Fichter, F.; Humpert, K.; Helvetica Chimica Acta; vol. 9; (1926); p. 521 - 525</ref><ref>Mueller, E.; Zeitschrift fuer Elektrochemie; vol. 10; (1904); p. 753 - 776</ref><br />
<br />
It can also be prepared by adding KHSO<sub>4</sub> to a solution of the more soluble salt [[Ammonium persulfate|ammonium peroxydisulfate]]. Potassium persulfate will precipitate from this reaction.<br />
<br />
==Projects==<br />
*Make silver peroxide<br />
*Make hydrogen peroxide<br />
*PCB etchant<br />
<br />
==Handling==<br />
===Safety===<br />
Potassium peroxydisulfate is a strong oxidizer and should be handled with care.<br />
<br />
Potassium persulfate will slowly decompose in wet environment.<br />
<br />
===Storage===<br />
Potassium persulfate should be stored in closed bottles, away from any acidic vapors or organic materials.<br />
<br />
If kept in a dark and dry environment, it will remain stable for years.<ref>Elbs, K.; Neher, P.; Chemiker-Zeitung; vol. 45; (1921); p. 1113 - 1114</ref><br />
<br />
===Disposal===<br />
Heating the salt will cause it to decompose.<br />
<br />
==Gallery==<br />
<gallery widths="250" position="center" columns="4" orientation="none"><br />
Potassium persulfate sample.jpg|A sample of potassium persulfate<br />
</gallery><br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=11022 Persulfate cell - Please translate passage (German)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13955 Potassium Persulfate for SO3 synth?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16912 Potassium Persulfate Oxidizer]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=41198 Potassium Persulfate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Potassium compounds]]<br />
[[Category:Persulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Materials that react with water]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Potassium_persulfate&diff=14946Potassium persulfate2023-03-24T08:49:41Z<p>Ave369: /* Chemical */</p>
<hr />
<div>{{Chembox<br />
| Name = Potassium persulfate<br />
| Reference =<br />
| IUPACName =<br />
| PIN =<br />
| SystematicName = Potassium peroxydisulfate<br />
| OtherNames = Dipotassium peroxodisulfate<br>Potassium perdisulfate<br>Potassium peroxydisulfate<br />
<!-- Images --><br />
| ImageFile = Potassium peroxydisulfate bottle and sample.jpg<br />
| ImageSize = 300<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Potassium persulfate sample and original bottle<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = [O-]S(=O)(=O)OOS(=O)(=O)[O-].[K+].[K+]<br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = 2.477<br />
| Formula = K<sub>2</sub>S<sub>2</sub>O<sub>8</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 270.322 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 186-250<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = (decomposes)<br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 1.75 g/100 mL (0 °C)<br>4.49 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in conc. [[hydrogen peroxide]]<ref>Skogareva; Ippolitov; Russian Journal of Inorganic Chemistry; vol. 43; nb. 11; (1998); p. 1668 - 1670</ref>, [[hydrogen fluoride]]<ref>Fredenhagen, H.; Z. Anorg. Chem.; vol. 242; (1939); p. 23 - 32</ref><br>Insoluble in [[benzene]], [[chloroform]], [[ethanol]], [[methanol]]<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Triclinic<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = 1,697.30 kJ/mol (at 25 °C)<br />
| DeltaHc = <br />
| DeltaHf = -1,916 kJ/mol<ref>Balej, Jan; Zeitschrift fur Physikalische Chemie; vol. 224; nb. 6; (2010); p. 883 - 892</ref><br />
| Entropy = <br />
| HeatCapacity = -84.0±3 J·mol<sup>-1</sup>·K<sup>-1</sup> (calculated, at 25 °C)<br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Non-explosive<br />
| FrictionSens = Non-explosive<br />
| DetonationV = Non-explosive<br />
| REFactor = Non-explosive<br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [http://www.sciencelab.com/msds.php?msdsId=9927234 ScienceLab]<br />
| FlashPt = Non-flammable<br />
| LD50 = 802 mg/kg (oral, rat)<br />
| LC50 = <br />
| MainHazards = Oxidant<br>Irritant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = [[Sodium persulfate]]<br>[[Ammonium persulfate]]<br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium peroxymonosulfate]]<br>[[Sodium persulfate]]<br />
}}<br />
}}<br />
'''Potassium persulfate''', also known as '''potassium peroxydisulfate''' is a chemical compound used as an oxidizing agent, with the formula '''K<sub>2</sub>S<sub>2</sub>O<sub>8</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
When potassium persulfate is heated in a 50% aqueous solution of [[sulfuric acid]], [[hydrogen peroxide]] results, which, due to the high temperature, distills from the solution. The H<sub>2</sub>O<sub>2</sub> obtained this way has a concentration of 40-60%. This method was previously used in the manufacturing of hydrogen peroxide on industrial scale before being replaced by the quinone process.<br />
<br />
Potassium persulfate will react with [[silver nitrate]] to form [[silver(I,III) oxide]] (silver peroxide):<ref>Marshall, H.; Journal of the Chemical Society; vol. 59; (1891); p. 775</ref><br />
:K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + AgNO<sub>3</sub> → Ag<sub>4</sub>O<sub>4</sub> + K<sub>2</sub>SO<sub>4</sub> + SO<sub>x</sub> + NO<sub>x</sub> + O<sub>x</sub><br />
<br />
Potassium peroxydisulfate oxidizes [[acetone]] in the presence of diluted [[sulfuric acid]] and [[silver]] metal to [[acetic acid]], releasing carbon dioxide:<ref>Broensted, J. N.; Z. phys. Ch.; vol. 102; (1922); p. 169 - 207</ref><br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + (CH<sub>3</sub>)<sub>2</sub>CO → CH<sub>3</sub>COOH + KHSO<sub>4</sub> + CO<sub>2</sub><br />
<br />
Reaction with [[nitric acid]] gives off [[oxygen]] and [[ozone]] fumes, and nitrogen as byproduct.<ref>Baeyer, A.; Villiger, V.; Ber.; vol. 34; (1901); p. 853 - 862</ref><br />
<br />
Heating it in solution or slush in presence of MnO<sub>2</sub> leads to a decomposition reaction catalyzed by manganese dioxide:<br />
<br />
:2 K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + 2H<sub>2</sub>O → 2 K<sub>2</sub>SO<sub>4</sub> + 2H<sub>2</sub>SO<sub>4</sub> + O<sub>2</sub><br />
<br />
===Physical===<br />
Potassium persulfate is a white crystalline solid, poorly soluble in water. It decomposes if heated to temperatures over 125°C. It has a density of 2.477 g/cm<sup>3</sup>.<br />
<br />
==Availability==<br />
Potassium persulfate can be bought from chemical suppliers and eBay.<br />
<br />
==Preparation==<br />
The most common way to synthesize potassium persulfate is via the [[electrolysis]] of a cold solution [[potassium bisulfate]] in [[sulfuric acid]], at a high current density<ref>F. Feher, "Potassium Peroxydisulfate", Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 392.</ref>:<br />
: 2 KHSO<sub>4</sub> → K<sub>2</sub>S<sub>2</sub>O<sub>8</sub> + H<sub>2</sub><br />
<br />
[[Tantalum]] electrodes can be used in this reaction.<ref>Gmelin, vol. O: MVol.2; 46, page 276 - 278</ref><br />
<br />
Bubbling elemental [[fluorine]] through an aqueous solution of KHSO<sub>4</sub> or K<sub>2</sub>SO<sub>4</sub> will also yield potassium peroxydisulfate. The reaction also works in the absence of water.<ref>Fichter, F.; Humpert, K.; Helvetica Chimica Acta; vol. 9; (1926); p. 521 - 525</ref><ref>Mueller, E.; Zeitschrift fuer Elektrochemie; vol. 10; (1904); p. 753 - 776</ref><br />
<br />
It can also be prepared by adding KHSO<sub>4</sub> to a solution of the more soluble salt [[Ammonium persulfate|ammonium peroxydisulfate]]. Potassium persulfate will precipitate from this reaction.<br />
<br />
==Projects==<br />
*Make silver peroxide<br />
*Make hydrogen peroxide<br />
*PCB etchant<br />
<br />
==Handling==<br />
===Safety===<br />
Potassium peroxydisulfate is a strong oxidizer and should be handled with care.<br />
<br />
Potassium persulfate will slowly decompose in wet environment.<br />
<br />
===Storage===<br />
Potassium persulfate should be stored in closed bottles, away from any acidic vapors or organic materials.<br />
<br />
If kept in a dark and dry environment, it will remain stable for years.<ref>Elbs, K.; Neher, P.; Chemiker-Zeitung; vol. 45; (1921); p. 1113 - 1114</ref><br />
<br />
===Disposal===<br />
Heating the salt will cause it to decompose.<br />
<br />
==Gallery==<br />
<gallery widths="250" position="center" columns="4" orientation="none"><br />
Potassium persulfate sample.jpg|A sample of potassium persulfate<br />
</gallery><br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=11022 Persulfate cell - Please translate passage (German)]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13955 Potassium Persulfate for SO3 synth?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16912 Potassium Persulfate Oxidizer]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=41198 Potassium Persulfate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Potassium compounds]]<br />
[[Category:Persulfates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Materials that react with water]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=HF&diff=14944HF2023-03-23T07:09:05Z<p>Ave369: Redirected page to Hydrofluoric acid</p>
<hr />
<div>#REDIRECT[[Hydrofluoric acid]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Lead&diff=14943Lead2023-03-23T07:08:44Z<p>Ave369: /* Projects */</p>
<hr />
<div>{{Infobox element<br />
<!-- top --><br />
|image name=Lead_Ingot.JPG<br />
|image alt=<br />
|image size=<br />
|image name comment=A freshly cast lead ingot<br />
|image name 2=<br />
|image alt 2=<br />
|image size 2=<br />
|image name 2 comment=<br />
<!-- General properties --><br />
|name=Lead<br />
|symbol=Pb<br />
|pronounce=<br />
|pronounce ref=<br />
|pronounce comment=<br />
|pronounce 2=<br />
|alt name=Plumbum (Latin)<br />
|alt names=<br />
|allotropes=<br />
|appearance=Metallic gray<br />
<!-- Periodic table --><br />
|above=[[Tin|Sn]]<br />
|below=Fl<br />
|left=[[Thallium]]<br />
|right=[[Bismuth]]<br />
|number=82<br />
|atomic mass=207.2(1)<br />
|atomic mass 2=<br />
|atomic mass ref=<br />
|atomic mass comment=<br />
|series=<br />
|series ref=<br />
|series comment=<br />
|series color=<br />
|group=14<br />
|group ref=<br />
|group comment=(carbon group)<br />
|period=6<br />
|period ref=<br />
|period comment=<br />
|block=p<br />
|block ref=<br />
|block comment=<br />
|electron configuration= [Xe] 4f<sup>14</sup> 5d<sup>10</sup> 6s<sup>2</sup> 6p<sup>2</sup><br />
|electron configuration ref=<br />
|electron configuration comment=<br />
|electrons per shell=2, 8, 18, 32, 18, 4<br />
|electrons per shell ref=<br />
|electrons per shell comment=<br />
<!-- Physical properties --><br />
|physical properties comment=<br />
|color=Metallic gray<br />
|phase=Solid<br />
|phase ref=<br />
|phase comment=<br />
|melting point K=600.61<br />
|melting point C=327.46<br />
|melting point F=621.43<br />
|melting point ref=<br />
|melting point comment=<br />
|boiling point K=2022<br />
|boiling point C=1749<br />
|boiling point F=3180<br />
|boiling point ref=<br />
|boiling point comment=<br />
|sublimation point K=<br />
|sublimation point C=<br />
|sublimation point F=<br />
|sublimation point ref=<br />
|sublimation point comment=<br />
|density gplstp=<br />
|density gplstp ref=<br />
|density gplstp comment=<br />
|density gpcm3nrt=11.34<br />
|density gpcm3nrt ref=<br />
|density gpcm3nrt comment=<br />
|density gpcm3nrt 2=<br />
|density gpcm3nrt 2 ref=<br />
|density gpcm3nrt 2 comment=<br />
|density gpcm3nrt 3=<br />
|density gpcm3nrt 3 ref=<br />
|density gpcm3nrt 3 comment=<br />
|density gpcm3mp=10.66<br />
|density gpcm3mp ref=<br />
|density gpcm3mp comment=<br />
|density gpcm3bp=<br />
|density gpcm3bp ref=<br />
|density gpcm3bp comment=<br />
|molar volume=<br />
|molar volume unit =<br />
|molar volume ref=<br />
|molar volume comment=<br />
|triple point K=<br />
|triple point kPa=<br />
|triple point ref=<br />
|triple point comment=<br />
|triple point K 2=<br />
|triple point kPa 2=<br />
|triple point 2 ref=<br />
|triple point 2 comment=<br />
|critical point K=<br />
|critical point MPa=<br />
|critical point ref=<br />
|critical point comment=<br />
|heat fusion=4.77<br />
|heat fusion ref=<br />
|heat fusion comment=<br />
|heat fusion 2=<br />
|heat fusion 2 ref=<br />
|heat fusion 2 comment=<br />
|heat vaporization=179.5<br />
|heat vaporization ref=<br />
|heat vaporization comment=<br />
|heat capacity=26.65<br />
|heat capacity ref=<br />
|heat capacity comment=<br />
|heat capacity 2=<br />
|heat capacity 2 ref=<br />
|heat capacity 2 comment=<br />
|vapor pressure 1=978<br />
|vapor pressure 10=1088<br />
|vapor pressure 100=1229<br />
|vapor pressure 1 k=1412<br />
|vapor pressure 10 k=1660<br />
|vapor pressure 100 k=2027<br />
|vapor pressure ref=<br />
|vapor pressure comment=<br />
|vapor pressure 1 2=<br />
|vapor pressure 10 2=<br />
|vapor pressure 100 2=<br />
|vapor pressure 1 k 2=<br />
|vapor pressure 10 k 2=<br />
|vapor pressure 100 k 2=<br />
|vapor pressure 2 ref=<br />
|vapor pressure 2 comment=<br />
<!-- Atomic properties --><br />
|atomic properties comment=<br />
|oxidation states='''4''', 3, '''2''', 1<br />
|oxidation states ref=<br />
|oxidation states comment=(2 and 4 are most common)<br />
|electronegativity=1.87 (+2)<br />
|electronegativity ref=<br />
|electronegativity comment=<br />
|ionization energy 1=715.6<br />
|ionization energy 1 ref=<br />
|ionization energy 1 comment=<br />
|ionization energy 2=1450.5<br />
|ionization energy 2 ref=<br />
|ionization energy 2 comment=<br />
|ionization energy 3=3081.5<br />
|ionization energy 3 ref=<br />
|ionization energy 3 comment=<br />
|number of ionization energies=<br />
|ionization energy ref=<br />
|ionization energy comment=<br />
|atomic radius=175<br />
|atomic radius ref=<br />
|atomic radius comment=<br />
|atomic radius calculated=<br />
|atomic radius calculated ref=<br />
|atomic radius calculated comment=<br />
|covalent radius=146±5<br />
|covalent radius ref=<br />
|covalent radius comment=<br />
|Van der Waals radius=202<br />
|Van der Waals radius ref=<br />
|Van der Waals radius comment=<br />
<!-- Miscellanea --><br />
|crystal structure=<br />
|crystal structure prefix=<br />
|crystal structure ref=<br />
|crystal structure comment=Face-centered cubic (fcc) <br />
|crystal structure 2=<br />
|crystal structure 2 prefix=<br />
|crystal structure 2 ref=<br />
|crystal structure 2 comment=<br />
|speed of sound=<br />
|speed of sound ref=<br />
|speed of sound comment=<br />
|speed of sound rod at 20=<br />
|speed of sound rod at 20 ref=<br />
|speed of sound rod at 20 comment=<br />
|speed of sound rod at r.t.=1190<br />
|speed of sound rod at r.t. ref=<br />
|speed of sound rod at r.t. comment=(annealed)<br />
|thermal expansion=<br />
|thermal expansion ref=<br />
|thermal expansion comment=<br />
|thermal expansion at 25=28.9<br />
|thermal expansion at 25 ref=<br />
|thermal expansion at 25 comment=<br />
|thermal conductivity=35.3<br />
|thermal conductivity ref=<br />
|thermal conductivity comment=<br />
|thermal conductivity 2=<br />
|thermal conductivity 2 ref=<br />
|thermal conductivity 2 comment=<br />
|thermal diffusivity=<br />
|thermal diffusivity ref=<br />
|thermal diffusivity comment=<br />
|electrical resistivity=<br />
|electrical resistivity unit prefix=<br />
|electrical resistivity ref=<br />
|electrical resistivity comment=<br />
|electrical resistivity at 0=<br />
|electrical resistivity at 0 ref=<br />
|electrical resistivity at 0 comment=<br />
|electrical resistivity at 20=208×10<sup>-9</sup><br />
|electrical resistivity at 20 ref=<br />
|electrical resistivity at 20 comment=<br />
|band gap=<br />
|band gap ref=<br />
|band gap comment=<br />
|Curie point K=<br />
|Curie point ref=<br />
|Curie point comment=<br />
|magnetic ordering=Diamagnetic<br />
|magnetic ordering ref=<br />
|magnetic ordering comment=<br />
|tensile strength=<br />
|tensile strength ref=<br />
|tensile strength comment=<br />
|Young's modulus=16<br />
|Young's modulus ref=<br />
|Young's modulus comment=<br />
|Shear modulus=5.6<br />
|Shear modulus ref=<br />
|Shear modulus comment=<br />
|Bulk modulus=46<br />
|Bulk modulus ref=<br />
|Bulk modulus comment=<br />
|Poisson ratio=0.44<br />
|Poisson ratio ref=<br />
|Poisson ratio comment=<br />
|Mohs hardness=1.5<br />
|Mohs hardness ref=<br />
|Mohs hardness comment=<br />
|Mohs hardness 2=<br />
|Mohs hardness 2 ref=<br />
|Mohs hardness 2 comment=<br />
|Vickers hardness=<br />
|Vickers hardness ref=<br />
|Vickers hardness comment=<br />
|Brinell hardness=38–50<br />
|Brinell hardness ref=<br />
|Brinell hardness comment=<br />
|CAS number=7439-92-1<br />
|CAS number ref=<br />
|CAS number comment=<br />
<!-- History --><br />
|naming=<br />
|predicted by=<br />
|prediction date ref=<br />
|prediction date=<br />
|discovered by=<br />
|discovery date ref=<br />
|discovery date=in the Middle East (7000 BCE)<br />
|first isolation by=<br />
|first isolation date ref=<br />
|first isolation date=<br />
|discovery and first isolation by=<br />
|named by=<br />
|named date ref=<br />
|named date=<br />
|history comment label=<br />
|history comment=<br />
<!-- Isotopes --><br />
|isotopes=<br />
|isotopes comment=<br />
|engvar=<br />
}}<br />
[[File:Weathered lead pieces.jpg|thumb|220x220px|Weathered lead pieces with various lead oxides on the outer surface.]]<br />
[[File:Remelted weathered lead pieces.jpg|thumb|220x220px|The same lead pieces pictured earlier, re-melted to show fresh surfaces.]]<br />
'''Lead''' is a chemical element with symbol '''Pb''' and atomic number 82. It is a very heavy and dense metal with a variety of uses, and is well-known for its toxicity.<br />
<br />
==Properties==<br />
===Chemical===<br />
Lead is resistant to certain acids such as [[sulfuric acid]] but will react with hot [[nitric acid]] to form [[lead(II) nitrate]], one of very few water-soluble lead compounds. Hot [[hydrochloric acid]] can also be used to convert lead into the poorly soluble [[lead(II) chloride]]. It will react very quickly with [[peracetic acid]] to form soluble [[lead(II) acetate]] and insoluble basic lead acetates. Lead will react with [[chlorine]] gas at elevated temperatures to produce the oily yellow liquid [[lead(IV) chloride]].<br />
<br />
Freshly cut lead will oxidize in air. Lead compounds span a wide range of colors, and the pigments [[lead carbonate|white lead]], [[Lead(II,IV) oxide|red lead]], and [[Lead(II) chromate|chrome yellow]] are all derived from it. Solutions can be tested for lead by adding a few drops of [[potassium iodide]] solution, which forms a bright yellow precipitate of [[lead(II) iodide]]. [[Sodium sulfide]] can also be used, precipitating black [[lead sulfide]].<br />
<br />
===Physical===<br />
Lead is a soft, malleable, and dense post-transition metal. Metallic lead has a bluish-silver color after being freshly cut, but it soon tarnishes to a dull grayish color when exposed to air. Lead has one of the lowest thermal and electrical conductivity of all metals. It is usually quickly identified by its high density and rather low melting point at 327 degrees Celsius. <br />
<br />
==Availability==<br />
Lead is available for sale as bars or ingots, in various purities. Oftentimes it is alloyed with [[antimony]] for hardness. Dissolving this alloyed lead in [[nitric acid]] will remove the antimony, as it precipitates as white [[antimony(III) oxide]] while the lead goes into solution. Certain wheels weights are made of lead or lead-antimony alloy (those of purer lead are very soft).<br />
<br />
Lead-acid batteries, such as those used in cars, contain lead and lead oxide. Extracting the lead from these batteries is not easy, and it's quite a dangerous method. You will need to remove any leftover electrolyte, then rip apart the outer plastic coating, which is very tough, then rip apart the lead plates from their cells. During this process, the plates may short, which can be dangerous is flammable gasses are generated from the battery. After removing the lead plates, heat them with a blowtorch in a metal can to melt away the lead from them, then recast the lead to remove any impurities.<ref>https://www.youtube.com/watch?v=SgGhNfJfSK0</ref><ref>https://www.youtube.com/watch?v=-FR4sEmY-54</ref><br />
<br />
Many items made in the earlier 20th century are a good source of lead, either pure or as alloy: old water pipes are a good bulk source; some car battery cable contacts were made of lead; very old hard drives tend to have counterweights made of lead; solders contain lead-tin alloy; scuba diving weight belts. The standard firearm bullets and shotgun pellets are also made of lead. Finally, lead fishing weights are widely available in outdoor or department stores. Round lead seals, used for sealing various electronic terminals, are also a good source of lead. You can occasionally find them on the ground, most often in the dirt around railway and phone control cabinets or electrical transformer stations.<br />
<br />
==Preparation==<br />
Lead can be prepared by reducing one of its oxides with [[lead sulfide]] or from soluble lead ions via electrowinning.<br />
<br />
==Projects==<br />
*[[Lead(II) acetate]] synthesis<br />
*Lead dioxide synthesis<br />
*[[Lead(II,IV) oxide|Lead tetroxide]] synthesis<br />
*Make [[sodium nitrite]]<br />
*Lead electrodes<br />
*Lead battery<br />
*Lead casting<br />
*Retort for making [[HF]]<br />
<br />
==Handling==<br />
===Safety===<br />
While lead is resistant to chemical attacks, it will rapidly oxidize into compounds that are extremely toxic to living beings. Lead poisoning is one of the most studied form of heavy metal poisoning in medicine and the nasty effects are well understood. It is not absolutely necessary to wear gloves while handling pieces of the metal, as long as hands are washed thoroughly after handling it. Lead in the metallic form is not absorbed through the skin at all. Soluble lead compounds, however, require more protection, and organolead compounds are the most dangerous as they are far more bioavailable and easily absorbed than any other source of lead. Because of its low melting point, lead is sometimes a popular use in home casting, although this has become less popular due to its toxicity. Contrary to popular belief, lead does not fume much when it is melted. At the temperatures that lead melts at, its vapor pressure is highly insignificant. Because of this, it is not necessary to wear a respirator while melting lead, though it is very necessary when working with lead dust, as this can be inhaled and absorbed through the lungs. It is not recommended to heat lead much higher than its melting point, as this may cause it to fume.<br />
<br />
Most lead compounds are poorly soluble in water, but [[lead(II) acetate]] and [[lead(II) nitrate]] are quite soluble and therefore are very toxic.<br />
<br />
===Storage===<br />
Since it does not form volatile compounds under standard conditions, it's not necessary to be stored in special containers. To prevent it from oxidizing, lead may be stored in a closed bottle under inert atmosphere, [[carbon dioxide]] is best. It's recommended to avoid storing it underwater or in any other liquids, as it will slowly oxidize, since there is some oxygen dissolved in liquid, and some lead oxide may flake off and contaminate the liquid.<br />
<br />
Lead alloys, like Pb-Sb resist oxidation and can remain relative lustrous for months up to one year, as long as they're kept in a place with low humidity and no corrosive vapors are present in the storage place.<br />
<br />
===Disposal===<br />
Lead scraps can be taken to metal recycling facilities. Scraps of metal can also be collected and re-cast into fresh pieces. Lead compounds should be converted to insoluble forms, before being taken to a hazardous waste facility.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=24104 recovery of Pb metal]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=29674 Getting lead metal from fishing weights]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=10534 Extracting lead from car batteries]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=63604 Lead metal from a lead-acid battery]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=20175 RECOVERY OF LEAD FROM CAR BATTERY SCRAPS.]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2342 lead from batteries]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2277 Separation of tin and lead]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=8245 Pyrophoric lead and tin]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=12635 Detecting lead in brass?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=8593 Removing lead stuck to Stainless Steel]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=62452 Success at Turning Lead into Gold]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=74200 Lead recovery from glass, theory and practic.]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13620 Lead Safety Question]<br />
<br />
[[Category:Elements]]<br />
[[Category:Metals]]<br />
[[Category:Post-transition metals]]<br />
[[Category:Poor metals]]<br />
[[Category:P-block]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Heavy metal toxicants]]<br />
[[Category:Carcinogenic]]<br />
[[Category:Neurotoxins]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Ammonium_nitrate&diff=14942Ammonium nitrate2023-03-22T21:26:35Z<p>Ave369: /* Projects */</p>
<hr />
<div>{{Chembox<br />
| Name = Ammonium nitrate<br />
| Reference = <br />
| IUPACName = Ammonium nitrate<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Ammonium nitricum<br>Ammonium saltpeter<br>Azanium nitrate<br>Nitrate of ammonia<br />
<!-- Images --><br />
| ImageFile = Ammonium nitrate.JPG<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Prills of store-bought ammonium nitrate which could be found in fertilizer or cold packs.<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = 210<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 1.725 g/cm<sup>3</sup> (20 °C)<br />
| Formula = NH<sub>4</sub>NO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 80.043 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 169.6<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 118 g/100 ml (0 °C)<br>150 g/100 ml (20 °C)<br>297 g/100 ml (40 °C)<br>410 g/100 ml (60 °C)<br>576 g/100 ml (80 °C)<br>1024 g/100 ml (100 °C)<br />
| SolubleOther = Soluble in [[acetone]], anh. [[ammonia]]<br>Insoluble in [[diethyl ether]], [[toluene]]<br />
| Solubility1 = 391 g/100 ml (25 °C)<br />
| Solvent1 = ammonia<br />
| Solubility2 = 3.8 g/100 ml (20 °C)<br />
| Solvent2 = ethanol<br />
| Solubility3 = 17.1 g/100 ml (20 °C)<br />
| Solvent3 = methanol<br />
| Solubility4 = 23.35 g/100 ml (25 °C)<br />
| Solvent4 = pyridine<br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Trigonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Very low<br />
| FrictionSens = Very low<br />
| DetonationV = 5,270 m/s<br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/lHcr5oK/ammonium-nitrate-sa.pdf Sigma-Aldrich]<br />
| FlashPt = <br />
| LD50 = 2,085–5,300 mg/kg (rats, oral)<br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Explosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Ammonium nitrite]]<br>[[Ammonium perchlorate]]<br />
}}<br />
}}<br />
'''Ammonium nitrate''' is an inorganic compound with the chemical formula '''NH<sub>4</sub>NO<sub>3</sub>'''. This white crystalline salt finds use as an oxidizer and as a reagent in the production of [[nitric acid]] and [[ammonia]].<br />
<br />
==Properties==<br />
===Chemical===<br />
Ammonium nitrate, like some other [[ammonium]] compounds, is unstable both at high temperatures and at high pH. At high temperatures it decomposes into [[nitrous oxide]] and water.<br />
<br />
: NH<sub>4</sub>NO<sub>3</sub> → N<sub>2</sub>O + 2 H<sub>2</sub>O<br />
<br />
In a very useful double replacement reaction, ammonium nitrate can be reacted with [[sodium hydroxide]] or [[sodium carbonate]] in somewhat differing procedures to yield [[sodium nitrate]], another useful oxidizer, and [[ammonia]] gas (with in the latter case [[carbon dioxide]]), which can be channeled into cold water to produce a solution.<br />
<br />
: NH<sub>4</sub>NO<sub>3</sub> + NaOH → NaNO<sub>3</sub> + NH<sub>3</sub> + H<sub>2</sub>O<br />
<br />
When sodium hydroxide is used, the chemicals should be added in dry form to a round-bottom or [[Erlenmeyer flask]] with a stopper and gas tubing ready to channel the ammonia. Upon the addition of a small amount of water or ice a vigorous reaction begins. If sodium carbonate, which is cheaper and safer, is used instead, the two reactants must be combined in solution and the solution heated until the reaction is complete, when ammonia can no longer be smelled.<br />
<br />
[[Nitric acid]] is produced when ammonium nitrate is reacted with either [[Sulfuric acid|sulfuric]] or [[hydrochloric acid]]s, along with an ammonium salt.<br />
<br />
: H<sub>2</sub>SO<sub>4</sub> + NH<sub>4</sub>NO<sub>3</sub> → HNO<sub>3</sub> + NH<sub>4</sub>HSO<sub>4</sub><br />
<br />
In general, however, alkali nitrates are preferred, as the separation is better, thus the ammonium nitrate is usually converted into sodium or potassium nitrate first.<br />
<br />
A mixture of molten ammonium nitrate, [[urea]] and silica gel (added to desensitize the AN), will yield [[guanidinium nitrate]].<ref>https://www.youtube.com/watch?v=QK_zsRBlkt0</ref><br />
<br />
Ammonium nitrate can decompose very violently if mixed with a fuel such as diesel or kerosene and hit by a shock wave. This mixture is a high explosive called [[ANFO]] (ammonium nitrate-fuel oil). In addition to being very dangerous, detonation of ANFO is very likely to be heard very far away and will almost certainly lead to trouble with the law, so it is not advised.<br />
<br />
=== Physical ===<br />
[[File:Ammonium nitrate NH4NO3 recrystallized from water by MrHomeScientist.jpg|250px|thumb|left|Some very nice looking ammonium nitrate crystal needles.]]<br />
Ammonium nitrate appears typically as a hard off-white white crystalline solid or as prills, such as those found in instant cold packs. The dissolution of ammonium nitrate in water is very endothermic, making it useful for cooling baths. Such a cooling bath is convenient in that, because of the salt's high solubility, it can be re-cooled many times during a synthesis by adding more ammonium nitrate. While it is already very soluble in cold water, ammonium nitrate is extremely soluble in boiling water, at 1024 g/100 ml. These factors make re-crystallization difficult, as instead of crystals you are more likely to end up with a large brick of wet ammonium nitrate. Boiling a solution completely to dryness is very dangerous, as rapid decomposition will begin if the solid heats too much, producing [[nitrous oxide]], a potent oxidizer. While only mildly hygroscopic, it is very difficult to dry ammonium nitrate after crystallization, and it is nearly impossible to do without the use of an oven or desiccator (however, leaving it overnight on a 50 C heating radiator does the trick and produces a material dry enough to be ground in a mortar and stored).<br />
<br />
=== Explosive ===<br />
Pure ammonium nitrate is not explosive under standard conditions, and neither strong heating nor impact will cause it to explode under normal circumstances, but if an explosive charge is used, it will detonate. However, when mixed with oxidizable compounds, like [[aluminium]] powder or [[nitromethane]], its sensitivity greatly increases, and a strong impact is enough to cause it to detonate, such as the impact from a gun. The detonation velocity of pure ammonium nitrate is 2500 m/s.<br />
<br />
==Availability==<br />
Due to the use of ammonium nitrate in explosives, the availability of this chemical varies greatly with the country it is being purchased in. It is easily obtainable in often nearly pure form in cheap instant cold packs at pharmacies in the United States. Larger amounts can sometimes be found as a fertilizer, either by itself or with [[calcium nitrate]]. In some countries, ammonium nitrate prill fertilizers are mixed with [[calcium sulfate]] which makes the material unsuitable for making explosives. In many other countries, however, it is not only not sold, but even restricted or altogether illegal due to fear of its use in bombings. In the EU, ammonium fertilizers tend to be restricted above a certain concentration.<br />
<br />
==Preparation==<br />
Ammonium nitrate can be produced by the combination of [[nitric acid]] and [[ammonia]] solution, though these two reagents are more often prepared from ammonium nitrate than being used to make it.<br />
<br />
: HNO<sub>3</sub> + NH<sub>3</sub> → NH<sub>4</sub>NO<sub>3</sub><br />
<br />
It can also be prepared by mixing two supersaturated solutions of [[calcium nitrate]] and [[ammonium sulfate]]. The resulting slush is filtrated and the liquid is cooled until ammonium nitrate begins to crystallize.<br />
<br />
==Projects==<br />
*Make [[nitric acid]]<br />
*Make [[ammonia]] and aq. ammonia<br />
*Make [[nitrous oxide]] (laughing gas)<br />
*Make [[Tannerite]]<br />
*Make [[guanidinium nitrate]]<br />
*Cooling bath<br />
<br />
==Safety==<br />
===Handling===<br />
Ammonium nitrate is a sensitive oxidizer, though not as sensitive as chlorates or perchlorates, and mixtures of it with organic compounds pose a great danger in substantial amounts. If mixed with bases, toxic and irritating ammonia gas is given off, and in high heat ammonium nitrate decomposes to form [[nitrous oxide]], which can be a dangerous airborne oxidizer in enclosed spaces. Ammonium nitrate must never come in contact with with [[nitromethane]], as this will form the very sensitive [[ANNM]], which may set off the ammonium nitrate if detonated. <br />
<br />
While not a matter of physical safety, the purchase, possession, purification or synthesis of ammonium nitrate is often frowned upon by law enforcement, and may require a license in some jurisdictions.<br />
<br />
===Storage===<br />
Ammonium nitrate should be stored in closed containers or bags as it is mild hygroscopic and away from any reducing agents or bases.<br />
<br />
===Disposal===<br />
Ammonium nitrate can be safely released in the environment, in small quantities, except water bodies.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=23482 Ammonium nitrate purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13181 A Unique Synthesis of Ammonium Nitrate]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19578 ammonium nitrate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Ammonium compounds]]<br />
[[Category:Nitrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Essential reagents]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Ammonium_nitrate&diff=14941Ammonium nitrate2023-03-22T21:26:15Z<p>Ave369: /* Physical */</p>
<hr />
<div>{{Chembox<br />
| Name = Ammonium nitrate<br />
| Reference = <br />
| IUPACName = Ammonium nitrate<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Ammonium nitricum<br>Ammonium saltpeter<br>Azanium nitrate<br>Nitrate of ammonia<br />
<!-- Images --><br />
| ImageFile = Ammonium nitrate.JPG<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Prills of store-bought ammonium nitrate which could be found in fertilizer or cold packs.<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = 210<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 1.725 g/cm<sup>3</sup> (20 °C)<br />
| Formula = NH<sub>4</sub>NO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 80.043 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 169.6<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 118 g/100 ml (0 °C)<br>150 g/100 ml (20 °C)<br>297 g/100 ml (40 °C)<br>410 g/100 ml (60 °C)<br>576 g/100 ml (80 °C)<br>1024 g/100 ml (100 °C)<br />
| SolubleOther = Soluble in [[acetone]], anh. [[ammonia]]<br>Insoluble in [[diethyl ether]], [[toluene]]<br />
| Solubility1 = 391 g/100 ml (25 °C)<br />
| Solvent1 = ammonia<br />
| Solubility2 = 3.8 g/100 ml (20 °C)<br />
| Solvent2 = ethanol<br />
| Solubility3 = 17.1 g/100 ml (20 °C)<br />
| Solvent3 = methanol<br />
| Solubility4 = 23.35 g/100 ml (25 °C)<br />
| Solvent4 = pyridine<br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Trigonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Very low<br />
| FrictionSens = Very low<br />
| DetonationV = 5,270 m/s<br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/lHcr5oK/ammonium-nitrate-sa.pdf Sigma-Aldrich]<br />
| FlashPt = <br />
| LD50 = 2,085–5,300 mg/kg (rats, oral)<br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Explosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Ammonium nitrite]]<br>[[Ammonium perchlorate]]<br />
}}<br />
}}<br />
'''Ammonium nitrate''' is an inorganic compound with the chemical formula '''NH<sub>4</sub>NO<sub>3</sub>'''. This white crystalline salt finds use as an oxidizer and as a reagent in the production of [[nitric acid]] and [[ammonia]].<br />
<br />
==Properties==<br />
===Chemical===<br />
Ammonium nitrate, like some other [[ammonium]] compounds, is unstable both at high temperatures and at high pH. At high temperatures it decomposes into [[nitrous oxide]] and water.<br />
<br />
: NH<sub>4</sub>NO<sub>3</sub> → N<sub>2</sub>O + 2 H<sub>2</sub>O<br />
<br />
In a very useful double replacement reaction, ammonium nitrate can be reacted with [[sodium hydroxide]] or [[sodium carbonate]] in somewhat differing procedures to yield [[sodium nitrate]], another useful oxidizer, and [[ammonia]] gas (with in the latter case [[carbon dioxide]]), which can be channeled into cold water to produce a solution.<br />
<br />
: NH<sub>4</sub>NO<sub>3</sub> + NaOH → NaNO<sub>3</sub> + NH<sub>3</sub> + H<sub>2</sub>O<br />
<br />
When sodium hydroxide is used, the chemicals should be added in dry form to a round-bottom or [[Erlenmeyer flask]] with a stopper and gas tubing ready to channel the ammonia. Upon the addition of a small amount of water or ice a vigorous reaction begins. If sodium carbonate, which is cheaper and safer, is used instead, the two reactants must be combined in solution and the solution heated until the reaction is complete, when ammonia can no longer be smelled.<br />
<br />
[[Nitric acid]] is produced when ammonium nitrate is reacted with either [[Sulfuric acid|sulfuric]] or [[hydrochloric acid]]s, along with an ammonium salt.<br />
<br />
: H<sub>2</sub>SO<sub>4</sub> + NH<sub>4</sub>NO<sub>3</sub> → HNO<sub>3</sub> + NH<sub>4</sub>HSO<sub>4</sub><br />
<br />
In general, however, alkali nitrates are preferred, as the separation is better, thus the ammonium nitrate is usually converted into sodium or potassium nitrate first.<br />
<br />
A mixture of molten ammonium nitrate, [[urea]] and silica gel (added to desensitize the AN), will yield [[guanidinium nitrate]].<ref>https://www.youtube.com/watch?v=QK_zsRBlkt0</ref><br />
<br />
Ammonium nitrate can decompose very violently if mixed with a fuel such as diesel or kerosene and hit by a shock wave. This mixture is a high explosive called [[ANFO]] (ammonium nitrate-fuel oil). In addition to being very dangerous, detonation of ANFO is very likely to be heard very far away and will almost certainly lead to trouble with the law, so it is not advised.<br />
<br />
=== Physical ===<br />
[[File:Ammonium nitrate NH4NO3 recrystallized from water by MrHomeScientist.jpg|250px|thumb|left|Some very nice looking ammonium nitrate crystal needles.]]<br />
Ammonium nitrate appears typically as a hard off-white white crystalline solid or as prills, such as those found in instant cold packs. The dissolution of ammonium nitrate in water is very endothermic, making it useful for cooling baths. Such a cooling bath is convenient in that, because of the salt's high solubility, it can be re-cooled many times during a synthesis by adding more ammonium nitrate. While it is already very soluble in cold water, ammonium nitrate is extremely soluble in boiling water, at 1024 g/100 ml. These factors make re-crystallization difficult, as instead of crystals you are more likely to end up with a large brick of wet ammonium nitrate. Boiling a solution completely to dryness is very dangerous, as rapid decomposition will begin if the solid heats too much, producing [[nitrous oxide]], a potent oxidizer. While only mildly hygroscopic, it is very difficult to dry ammonium nitrate after crystallization, and it is nearly impossible to do without the use of an oven or desiccator (however, leaving it overnight on a 50 C heating radiator does the trick and produces a material dry enough to be ground in a mortar and stored).<br />
<br />
=== Explosive ===<br />
Pure ammonium nitrate is not explosive under standard conditions, and neither strong heating nor impact will cause it to explode under normal circumstances, but if an explosive charge is used, it will detonate. However, when mixed with oxidizable compounds, like [[aluminium]] powder or [[nitromethane]], its sensitivity greatly increases, and a strong impact is enough to cause it to detonate, such as the impact from a gun. The detonation velocity of pure ammonium nitrate is 2500 m/s.<br />
<br />
==Availability==<br />
Due to the use of ammonium nitrate in explosives, the availability of this chemical varies greatly with the country it is being purchased in. It is easily obtainable in often nearly pure form in cheap instant cold packs at pharmacies in the United States. Larger amounts can sometimes be found as a fertilizer, either by itself or with [[calcium nitrate]]. In some countries, ammonium nitrate prill fertilizers are mixed with [[calcium sulfate]] which makes the material unsuitable for making explosives. In many other countries, however, it is not only not sold, but even restricted or altogether illegal due to fear of its use in bombings. In the EU, ammonium fertilizers tend to be restricted above a certain concentration.<br />
<br />
==Preparation==<br />
Ammonium nitrate can be produced by the combination of [[nitric acid]] and [[ammonia]] solution, though these two reagents are more often prepared from ammonium nitrate than being used to make it.<br />
<br />
: HNO<sub>3</sub> + NH<sub>3</sub> → NH<sub>4</sub>NO<sub>3</sub><br />
<br />
It can also be prepared by mixing two supersaturated solutions of [[calcium nitrate]] and [[ammonium sulfate]]. The resulting slush is filtrated and the liquid is cooled until ammonium nitrate begins to crystallize.<br />
<br />
==Projects==<br />
*Make [[nitric acid]]<br />
*Make [[ammonia]] and aq. ammonia<br />
*Make [[nitrous oxide]] (laughing gas)<br />
*Make [[Tannerite]]<br />
*Make [[guanidinium nitrate]]<br />
<br />
==Safety==<br />
===Handling===<br />
Ammonium nitrate is a sensitive oxidizer, though not as sensitive as chlorates or perchlorates, and mixtures of it with organic compounds pose a great danger in substantial amounts. If mixed with bases, toxic and irritating ammonia gas is given off, and in high heat ammonium nitrate decomposes to form [[nitrous oxide]], which can be a dangerous airborne oxidizer in enclosed spaces. Ammonium nitrate must never come in contact with with [[nitromethane]], as this will form the very sensitive [[ANNM]], which may set off the ammonium nitrate if detonated. <br />
<br />
While not a matter of physical safety, the purchase, possession, purification or synthesis of ammonium nitrate is often frowned upon by law enforcement, and may require a license in some jurisdictions.<br />
<br />
===Storage===<br />
Ammonium nitrate should be stored in closed containers or bags as it is mild hygroscopic and away from any reducing agents or bases.<br />
<br />
===Disposal===<br />
Ammonium nitrate can be safely released in the environment, in small quantities, except water bodies.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=23482 Ammonium nitrate purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13181 A Unique Synthesis of Ammonium Nitrate]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19578 ammonium nitrate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Ammonium compounds]]<br />
[[Category:Nitrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Essential reagents]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Ammonium_nitrate&diff=14932Ammonium nitrate2023-03-20T06:33:31Z<p>Ave369: /* Physical */</p>
<hr />
<div>{{Chembox<br />
| Name = Ammonium nitrate<br />
| Reference = <br />
| IUPACName = Ammonium nitrate<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Ammonium nitricum<br>Ammonium saltpeter<br>Azanium nitrate<br>Nitrate of ammonia<br />
<!-- Images --><br />
| ImageFile = Ammonium nitrate.JPG<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageCaption = Prills of store-bought ammonium nitrate which could be found in fertilizer or cold packs.<br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = White solid<br />
| BoilingPt = <br />
| BoilingPtC = 210<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 1.725 g/cm<sup>3</sup> (20 °C)<br />
| Formula = NH<sub>4</sub>NO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 80.043 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = 169.6<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = 118 g/100 ml (0 °C)<br>150 g/100 ml (20 °C)<br>297 g/100 ml (40 °C)<br>410 g/100 ml (60 °C)<br>576 g/100 ml (80 °C)<br>1024 g/100 ml (100 °C)<br />
| SolubleOther = Soluble in [[acetone]], anh. [[ammonia]]<br>Insoluble in [[diethyl ether]], [[toluene]]<br />
| Solubility1 = 391 g/100 ml (25 °C)<br />
| Solvent1 = ammonia<br />
| Solubility2 = 3.8 g/100 ml (20 °C)<br />
| Solvent2 = ethanol<br />
| Solubility3 = 17.1 g/100 ml (20 °C)<br />
| Solvent3 = methanol<br />
| Solubility4 = 23.35 g/100 ml (25 °C)<br />
| Solvent4 = pyridine<br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = Trigonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = Very low<br />
| FrictionSens = Very low<br />
| DetonationV = 5,270 m/s<br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/lHcr5oK/ammonium-nitrate-sa.pdf Sigma-Aldrich]<br />
| FlashPt = <br />
| LD50 = 2,085–5,300 mg/kg (rats, oral)<br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Explosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Ammonium nitrite]]<br>[[Ammonium perchlorate]]<br />
}}<br />
}}<br />
'''Ammonium nitrate''' is an inorganic compound with the chemical formula '''NH<sub>4</sub>NO<sub>3</sub>'''. This white crystalline salt finds use as an oxidizer and as a reagent in the production of [[nitric acid]] and [[ammonia]].<br />
<br />
==Properties==<br />
===Chemical===<br />
Ammonium nitrate, like some other [[ammonium]] compounds, is unstable both at high temperatures and at high pH. At high temperatures it decomposes into [[nitrous oxide]] and water, while in highly basic aqueous conditions ammonia is evolved, leaving the nitrate ion in solution.<br />
<br />
[[Nitric acid]] is produced when ammonium nitrate is reacted with either [[Sulfuric acid|sulfuric]] or [[hydrochloric acid]]s, along with an ammonium salt.<br />
<br />
In a very useful double replacement reaction, ammonium nitrate can be reacted with [[sodium hydroxide]] or [[sodium carbonate]] in somewhat differing procedures to yield [[sodium nitrate]], another useful oxidizer, and [[ammonia]] gas (with in the latter case [[carbon dioxide]]), which can be channeled into cold water to produce a solution. When sodium hydroxide is used, the chemicals should be added in dry form to a round-bottom or [[Erlenmeyer flask]] with a stopper and gas tubing ready to channel the ammonia. Upon the addition of a small amount of water or ice a vigorous reaction begins. If sodium carbonate, which is cheaper and safer, is used instead, the two reactants must be combined in solution and the solution heated until the reaction is complete, when ammonia can no longer be smelled.<br />
<br />
Ammonium nitrate can decompose very violently if mixed with a fuel such as diesel or kerosene and hit by a shock wave. This mixture is a high explosive called [[ANFO]] (ammonium nitrate-fuel oil). In addition to being very dangerous, detonation of ANFO is very likely to be heard very far away and will almost certainly lead to trouble with the law, so it is not advised.<br />
<br />
=== Physical ===<br />
[[File:Ammonium nitrate NH4NO3 recrystallized from water by MrHomeScientist.jpg|250px|thumb|left|Some very nice looking ammonium nitrate crystal needles.]]<br />
Ammonium nitrate appears typically as a hard off-white white crystalline solid or as prills, such as those found in instant cold packs. The dissolution of ammonium nitrate in water is very endothermic, making it useful for cooling baths. While it is already very soluble in cold water, ammonium nitrate is extremely soluble in boiling water, at 1024 g/100 ml. These factors make re-crystallization difficult, as instead of crystals you are more likely to end up with a large brick of wet ammonium nitrate. Boiling a solution completely to dryness is very dangerous, as rapid decomposition will begin if the solid heats too much, producing [[nitrous oxide]], a potent oxidizer. While only mildly hygroscopic, it is very difficult to dry ammonium nitrate after crystallization, and it is nearly impossible to do without the use of an oven or desiccator (however, leaving it overnight on a 50 C heating radiator does the trick and produces a material dry enough to be ground in a mortar and stored).<br />
<br />
=== Explosive ===<br />
Pure ammonium nitrate is not explosive under standard conditions, and neither strong heating nor impact will cause it to explode under normal circumstances, but if an explosive charge is used, it will detonate. However, when mixed with oxidizable compounds, like [[aluminium]] powder or [[nitromethane]], its sensitivity greatly increases, and a strong impact is enough to cause it to detonate, such as the impact from a gun. The detonation velocity of pure ammonium nitrate is 2500 m/s.<br />
<br />
==Availability==<br />
Due to the use of ammonium nitrate in explosives, the availability of this chemical varies greatly with the country it is being purchased in. It is easily obtainable in often nearly pure form in cheap instant cold packs at pharmacies in the United States. Larger amounts can sometimes be found as a fertilizer, either by itself or with [[calcium nitrate]]. In some countries, ammonium nitrate prill fertilizers are mixed with [[calcium sulfate]] which makes the material unsuitable for making explosives. In many other countries, however, it is not only not sold, but even restricted or altogether illegal due to fear of its use in bombings.<br />
<br />
==Preparation==<br />
Ammonium nitrate can be produced by the combination of [[nitric acid]] and [[ammonia]] solution, though these two reagents are more often prepared from ammonium nitrate than being used to make it.<br />
<br />
It can also be prepared by mixing two supersaturated solutions of [[calcium nitrate]] and [[ammonium sulfate]]. The resulting slush is filtrated and the liquid is cooled until ammonium nitrate begins to crystallize.<br />
<br />
==Projects==<br />
*Make [[nitric acid]]<br />
*Make [[ammonia]] and aq. ammonia<br />
*Make [[nitrous oxide]] (laughing gas)<br />
*Make [[Tannerite]]<br />
*Make [[guanidinium nitrate]]<br />
<br />
==Safety==<br />
===Handling===<br />
Ammonium nitrate is a sensitive oxidizer, though not as sensitive as chlorates or perchlorates, and mixtures of it with organic compounds pose a great danger in substantial amounts. If mixed with bases, toxic and irritating ammonia gas is given off, and in high heat ammonium nitrate decomposes to form [[nitrous oxide]], which can be a dangerous airborne oxidizer in enclosed spaces. Ammonium nitrate must never come in contact with with [[nitromethane]], as this will form the very sensitive [[ANNM]], which may set off the ammonium nitrate if detonated. <br />
<br />
While not a matter of physical safety, the purchase, possession, purification or synthesis of ammonium nitrate is often frowned upon by law enforcement, and may require a license in some jurisdictions.<br />
<br />
===Storage===<br />
Ammonium nitrate should be stored in closed containers or bags as it is mild hygroscopic and away from any reducing agents or bases.<br />
<br />
===Disposal===<br />
Ammonium nitrate can be safely released in the environment, in small quantities, except water bodies.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=23482 Ammonium nitrate purification]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13181 A Unique Synthesis of Ammonium Nitrate]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19578 ammonium nitrate]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Ammonium compounds]]<br />
[[Category:Nitrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Readily available chemicals]]<br />
[[Category:Essential reagents]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=PH_indicator&diff=14931PH indicator2023-03-19T20:31:58Z<p>Ave369: /* Phenolphthalein */</p>
<hr />
<div>{{DISPLAYTITLE:pH indicator}}<br />
{{stub}}<br />
'''pH indicators''' are compounds that change color depending on the pH of the solution they are in. This property of theirs is widely used in labs to measure pH.<br />
<br />
== Common indicators ==<br />
<br />
=== [[Litmus]] ===<br />
Litmus is a complex mixture of plant dyes, discovered back in the days of alchemists. Litmus turns red in low pH, purple in neutral pH, and blue in high pH. Modern litmus paper strips are no longer based on natural dyes, they are impregnated with litmoid, or 7-hydroxyphenoxazone, which is produced artificially. Litmoid is available as paper strips, solution or dry solid.<br />
<br />
=== [[Universal indicator]] ===<br />
The most common modern indicator strips are impregnated with a mixture of dyes which allows very precise measuring of pH. Generally, universal indicator strips are red in low pH, yellow or green in neutral pH and blue in high pH, but there is a hue corresponding to every integer meaning of pH. Universal indicator is available as paper strips. A somewhat crude version of universal indicator can be prepared by mixing thymolphthalein and methyl red.<br />
<br />
=== [[Anthocyanin]] ===<br />
An easily obtained replacement for litmus, anthocyanin can be extracted from various food products, but the best one comes from hibiscus herbal tea (marketed under names such as "Flor de Jamaica", "Karkade", "Italian Tea" or "Red Zinger"), extracted with [[ethanol]]. It is red in very low pH, pink in low pH, purple in neutral pH, green in high pH and greenish-yellow in very high pH. It can also be made from red cabbage, but such an extract is inferior, because cabbage contains a lot of water, which goes into the extract and makes it less stable. Paper strips can be easily made from an alcohol extract of anthocyanin, or it can be used as a solution.<br />
<br />
=== [[Phenolphthalein]] ===<br />
An indicator for bases, the colorless phenolphthalein turns rose, pink or crimson in high pH. Extremely high pH causes it to become colorless again. Several decades ago, it was in common use as an over-the-counter laxative, which guaranteed easy access to this indicator for amateur chemists. Currently, it is no longer sold in drugstores. Phenolphthalein is usually available as a solution or crystalline solid, rarely paper strips.<br />
<br />
Very similar to phenolphthalein are its closest relatives, thymolphthalein, naphtholphthalein and cresolphthalein. These colorless indicators turn blue, cyan and red respectively in high pH. They do not lose color in extremely high pH.<br />
<br />
=== [[Curcumin]] ===<br />
An over-the-counter replacement for phenolphthalein, the yellow curcumin turns red in high pH. It can be extracted from inexpensive spices such as turmeric or "curry mix". Do not use curcumin in solutions that contain the borate ion, because it gives a false positive with curcumin. Curcumin can be used as a liquid extract, or paper strips can be made from it.<br />
<br />
== Uncommon indicators ==<br />
<br />
=== [[Methyl orange]] ===<br />
An indicator for acids, methyl orange changes its color to red in very low pH. It is available as paper strips, solution or solid.<br />
<br />
=== [[Crystal violet]] ===<br />
Also an indicator for acids, crystal violet allows distinguishing strong, mid-strength and weak acids. Strong acids make this intensely violet dye turn yellow, mid-strength or very dilute strong acids turn it green or blue, and weak acids do not change its color. It usually comes as a solid.<br />
<br />
=== [[Bromocresol purple]] ===<br />
This indicator is unusual in that it is somewhat acidic itself, and orange-yellow in its pure form. However, when the pH turns to neutral or higher, it turns purple. It usually comes as a solid. Bromocresol purple is very useful in acid-base titrations, because its pronouncedly purple form comes to existence at pH precisely 7, which allows for very precise measurements. First it turns crimson, and when you add more base, it becomes purple. The crimson color, corresponding to pH 7, signals the end of the titration.<br />
<br />
==See also==<br />
*[[pH test strip]]<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=1372 Homemade Ph Indicator from purple cabbage]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=74719 Extraction of pH indicator from black pellargonium]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16782 Mulberry pH indicator]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=8748 homemade natural indicator]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19902 PH Strips shelf life]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=10189 Check for pH in non aqueous media]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6716 pH indicator for portland cement.]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=11558 pH indicators and dyes]<br />
<br />
[[Category:PH indicators]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=PH_indicator&diff=14930PH indicator2023-03-19T20:24:19Z<p>Ave369: /* Common indicators */</p>
<hr />
<div>{{DISPLAYTITLE:pH indicator}}<br />
{{stub}}<br />
'''pH indicators''' are compounds that change color depending on the pH of the solution they are in. This property of theirs is widely used in labs to measure pH.<br />
<br />
== Common indicators ==<br />
<br />
=== [[Litmus]] ===<br />
Litmus is a complex mixture of plant dyes, discovered back in the days of alchemists. Litmus turns red in low pH, purple in neutral pH, and blue in high pH. Modern litmus paper strips are no longer based on natural dyes, they are impregnated with litmoid, or 7-hydroxyphenoxazone, which is produced artificially. Litmoid is available as paper strips, solution or dry solid.<br />
<br />
=== [[Universal indicator]] ===<br />
The most common modern indicator strips are impregnated with a mixture of dyes which allows very precise measuring of pH. Generally, universal indicator strips are red in low pH, yellow or green in neutral pH and blue in high pH, but there is a hue corresponding to every integer meaning of pH. Universal indicator is available as paper strips. A somewhat crude version of universal indicator can be prepared by mixing thymolphthalein and methyl red.<br />
<br />
=== [[Anthocyanin]] ===<br />
An easily obtained replacement for litmus, anthocyanin can be extracted from various food products, but the best one comes from hibiscus herbal tea (marketed under names such as "Flor de Jamaica", "Karkade", "Italian Tea" or "Red Zinger"), extracted with [[ethanol]]. It is red in very low pH, pink in low pH, purple in neutral pH, green in high pH and greenish-yellow in very high pH. It can also be made from red cabbage, but such an extract is inferior, because cabbage contains a lot of water, which goes into the extract and makes it less stable. Paper strips can be easily made from an alcohol extract of anthocyanin, or it can be used as a solution.<br />
<br />
=== [[Phenolphthalein]] ===<br />
An indicator for bases, the colorless phenolphthalein turns rose, pink or crimson in high pH. Extremely high pH causes it to become colorless again. Several decades ago, it was in common use as an over-the-counter laxative, which guaranteed easy access to this indicator for amateur chemists. Currently, it is no longer sold in drugstores. Phenolphthalein is usually available as a solution or crystalline solid, rarely paper strips.<br />
<br />
Very similar to phenolphthalein are its closest relatives, thymolphthalein and cresolphthalein. These colorless indicators turn blue and red respectively in high pH. They do not lose color in extremely high pH.<br />
<br />
=== [[Curcumin]] ===<br />
An over-the-counter replacement for phenolphthalein, the yellow curcumin turns red in high pH. It can be extracted from inexpensive spices such as turmeric or "curry mix". Do not use curcumin in solutions that contain the borate ion, because it gives a false positive with curcumin. Curcumin can be used as a liquid extract, or paper strips can be made from it.<br />
<br />
== Uncommon indicators ==<br />
<br />
=== [[Methyl orange]] ===<br />
An indicator for acids, methyl orange changes its color to red in very low pH. It is available as paper strips, solution or solid.<br />
<br />
=== [[Crystal violet]] ===<br />
Also an indicator for acids, crystal violet allows distinguishing strong, mid-strength and weak acids. Strong acids make this intensely violet dye turn yellow, mid-strength or very dilute strong acids turn it green or blue, and weak acids do not change its color. It usually comes as a solid.<br />
<br />
=== [[Bromocresol purple]] ===<br />
This indicator is unusual in that it is somewhat acidic itself, and orange-yellow in its pure form. However, when the pH turns to neutral or higher, it turns purple. It usually comes as a solid. Bromocresol purple is very useful in acid-base titrations, because its pronouncedly purple form comes to existence at pH precisely 7, which allows for very precise measurements. First it turns crimson, and when you add more base, it becomes purple. The crimson color, corresponding to pH 7, signals the end of the titration.<br />
<br />
==See also==<br />
*[[pH test strip]]<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=1372 Homemade Ph Indicator from purple cabbage]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=74719 Extraction of pH indicator from black pellargonium]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16782 Mulberry pH indicator]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=8748 homemade natural indicator]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19902 PH Strips shelf life]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=10189 Check for pH in non aqueous media]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=6716 pH indicator for portland cement.]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=11558 pH indicators and dyes]<br />
<br />
[[Category:PH indicators]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=PH_test_strip&diff=14929PH test strip2023-03-19T13:26:55Z<p>Ave369: /* DIY pH strip */</p>
<hr />
<div>{{DISPLAYTITLE:pH test strip}}<br />
{{Stub}}<br />
[[File:PH paper and litmus paper by NurdRage.jpg|thumb|350px|Blue litmus paper (left) and a pH test strip (right)]]<br />
A '''pH test strip''', also called '''pH strip''', '''pH test kit''', '''strip paper''', '''universal indicator test kit''' among other dozen names, is a type of testing material which, when immersed in a solution, displays smooth color changes over a pH value range from 0 to 14, to indicate the acidity or alkalinity of solutions.<br />
<br />
==General==<br />
pH strips come in various forms, the most common form involves a plastic case with a colored palette indicating the color of the indicator at different pH values. Some have three rolls of pH paper, each for different pH range (acidic, neutral, basic), while others have a series of dots on a paper strip, which display the pH level via various indicators.<br />
<br />
==Availability==<br />
pH test strips can be bought from most hardware stores, as swimming pool water pH test kits. They can also be bought online.<br />
<br />
==DIY pH strip==<br />
A simple ph test kit can be made by carefully adding bits of prepared [[universal indicator]] on a strip of paper and using a printed image of the indicator at various pH levels as reference. If you don't have access to that, you can substitute a more common indicator such as [[anthocyanin]].<br />
<br />
==See also==<br />
*[[Litmus]]<br />
*[[pH indicator]]<br />
*[[Universal indicator]]<br />
*[[Anthocyanin]] for easy and cheap DIY test strips<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=19902 PH Strips shelf life]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=66203 Red Cabbage pH juice - can it be stored?]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=1372 Homemade Ph Indicator from purple cabbage]<br />
<br />
[[Category:Lab equipment]]<br />
[[Category:Paperware]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Litmus&diff=14928Litmus2023-03-19T13:19:40Z<p>Ave369: /* Properties */</p>
<hr />
<div>{{Stub}}<br />
'''Litmus''' is a water soluble mixture of different dyes extracted from lichens. Litmus absorbed onto filter papers is one of the oldest and most known form of [[pH indicator]].<br />
<br />
==Composition==<br />
Litmus contains between 10 to 15 different dyes, with 7-hydroxyphenoxazone being the most important.<br />
<br />
==Properties==<br />
Litmus is a black powder, soluble in [[water]] and some organic solvents, which changes color in acidic environment to red and blue in basic solution. At pH 7 the solution is purple. Litmus, both as solution and as paper strips, is as versatile as [[universal indicator]] strips, but it offers less precise measurements of pH. Still, if you don't need to know the exact pH, litmus strips can be used instead of universal indicator strips.<br />
<br />
Litmus paper strips come in three varieties: red, neutral (violet) and blue. Red strips are made from an acidified litmus solution, they are red by default; the inverse is true with blue litmus strips. Red and blue strips are used when you already know if you are testing for an acid or a base, and the color change is more pronounced with them (this is useful with old paper strips, which turn very pale with time). Red and blue litmus paper strips can usually be converted to neutral litmus paper strips by soaking them in water and drying.<br />
<br />
==Availability==<br />
Litmus powder and litmus paper can be purchased from chemical suppliers and online. It can be found on eBay and Amazon.<br />
<br />
==Isolation==<br />
Litmus can be isolated from certain lichens, though you'll need an appreciable amount of them to extract useful quantities of litmus.<br />
<br />
==Projects==<br />
*[[pH indicator]]<br />
*Dye<br />
<br />
==Handling==<br />
===Safety===<br />
Litmus has low toxicity and no special protection or safety measures are required when handling it. However, it can stain, so wear disposable gear.<br />
<br />
===Storage===<br />
In closed bottles.<br />
<br />
===Disposal===<br />
No special disposal is required. Discard it as you wish.<br />
<br />
==See also ==<br />
* [[Anthocyanin]], a natural dye mix similar in properties to litmus, and much easier to obtain from OTC products, which can be used instead of it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=7357 Indicators]<br />
<br />
[[Category:Materials with no specific chemical formula]]<br />
[[Category:Organic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Litmus&diff=14927Litmus2023-03-19T13:18:42Z<p>Ave369: /* Properties */</p>
<hr />
<div>{{Stub}}<br />
'''Litmus''' is a water soluble mixture of different dyes extracted from lichens. Litmus absorbed onto filter papers is one of the oldest and most known form of [[pH indicator]].<br />
<br />
==Composition==<br />
Litmus contains between 10 to 15 different dyes, with 7-hydroxyphenoxazone being the most important.<br />
<br />
==Properties==<br />
Litmus is a black powder, soluble in [[water]] and some organic solvents, which changes color in acidic environment to red and blue in basic solution. At pH 7 the solution is purple. Litmus, both as solution and as paper strips, is as versatile as [[universal indicator]] strips, but it offers less precise measurements of pH. Still, if you don't need to know the exact pH, litmus strips can be used instead of universal indicator strips.<br />
<br />
Litmus paper strips come in three varieties: red, neutral (violet) and blue. Red strips are made from an acidified litmus solution, they are red by default; the inverse is true with blue litmus strips. Red and blue strips are used when you already know if you are testing for an acid or a base, and the color change is more pronounced with them. Red and blue litmus paper strips can usually be converted to neutral litmus paper strips by soaking them in water and drying.<br />
<br />
==Availability==<br />
Litmus powder and litmus paper can be purchased from chemical suppliers and online. It can be found on eBay and Amazon.<br />
<br />
==Isolation==<br />
Litmus can be isolated from certain lichens, though you'll need an appreciable amount of them to extract useful quantities of litmus.<br />
<br />
==Projects==<br />
*[[pH indicator]]<br />
*Dye<br />
<br />
==Handling==<br />
===Safety===<br />
Litmus has low toxicity and no special protection or safety measures are required when handling it. However, it can stain, so wear disposable gear.<br />
<br />
===Storage===<br />
In closed bottles.<br />
<br />
===Disposal===<br />
No special disposal is required. Discard it as you wish.<br />
<br />
==See also ==<br />
* [[Anthocyanin]], a natural dye mix similar in properties to litmus, and much easier to obtain from OTC products, which can be used instead of it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=7357 Indicators]<br />
<br />
[[Category:Materials with no specific chemical formula]]<br />
[[Category:Organic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Litmus&diff=14926Litmus2023-03-19T13:16:40Z<p>Ave369: /* Handling */</p>
<hr />
<div>{{Stub}}<br />
'''Litmus''' is a water soluble mixture of different dyes extracted from lichens. Litmus absorbed onto filter papers is one of the oldest and most known form of [[pH indicator]].<br />
<br />
==Composition==<br />
Litmus contains between 10 to 15 different dyes, with 7-hydroxyphenoxazone being the most important.<br />
<br />
==Properties==<br />
Litmus is a black powder, soluble in [[water]] and some organic solvents, which changes color in acidic environment to red and blue in basic solution. At pH 7 the solution is purple. Litmus, both as solution and as paper strips, is as versatile as [[universal indicator]] strips, but it offers less precise measurements of pH. Still, if you don't need to know the exact pH, litmus strips can be used instead of universal indicator strips.<br />
<br />
Litmus paper strips come in three varieties: red, neutral and blue. Red strips are made from an acidified litmus solution, they are red by default; the inverse is true with blue litmus strips. Red and blue strips are used when you already know if you are testing for an acid or a base, and the color change is more pronounced with them. Red and blue litmus paper strips can usually be converted to neutral litmus paper strips by soaking them in water and drying.<br />
<br />
==Availability==<br />
Litmus powder and litmus paper can be purchased from chemical suppliers and online. It can be found on eBay and Amazon.<br />
<br />
==Isolation==<br />
Litmus can be isolated from certain lichens, though you'll need an appreciable amount of them to extract useful quantities of litmus.<br />
<br />
==Projects==<br />
*[[pH indicator]]<br />
*Dye<br />
<br />
==Handling==<br />
===Safety===<br />
Litmus has low toxicity and no special protection or safety measures are required when handling it. However, it can stain, so wear disposable gear.<br />
<br />
===Storage===<br />
In closed bottles.<br />
<br />
===Disposal===<br />
No special disposal is required. Discard it as you wish.<br />
<br />
==See also ==<br />
* [[Anthocyanin]], a natural dye mix similar in properties to litmus, and much easier to obtain from OTC products, which can be used instead of it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=7357 Indicators]<br />
<br />
[[Category:Materials with no specific chemical formula]]<br />
[[Category:Organic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Litmus&diff=14925Litmus2023-03-19T13:16:28Z<p>Ave369: </p>
<hr />
<div>{{Stub}}<br />
'''Litmus''' is a water soluble mixture of different dyes extracted from lichens. Litmus absorbed onto filter papers is one of the oldest and most known form of [[pH indicator]].<br />
<br />
==Composition==<br />
Litmus contains between 10 to 15 different dyes, with 7-hydroxyphenoxazone being the most important.<br />
<br />
==Properties==<br />
Litmus is a black powder, soluble in [[water]] and some organic solvents, which changes color in acidic environment to red and blue in basic solution. At pH 7 the solution is purple. Litmus, both as solution and as paper strips, is as versatile as [[universal indicator]] strips, but it offers less precise measurements of pH. Still, if you don't need to know the exact pH, litmus strips can be used instead of universal indicator strips.<br />
<br />
Litmus paper strips come in three varieties: red, neutral and blue. Red strips are made from an acidified litmus solution, they are red by default; the inverse is true with blue litmus strips. Red and blue strips are used when you already know if you are testing for an acid or a base, and the color change is more pronounced with them. Red and blue litmus paper strips can usually be converted to neutral litmus paper strips by soaking them in water and drying.<br />
<br />
==Availability==<br />
Litmus powder and litmus paper can be purchased from chemical suppliers and online. It can be found on eBay and Amazon.<br />
<br />
==Isolation==<br />
Litmus can be isolated from certain lichens, though you'll need an appreciable amount of them to extract useful quantities of litmus.<br />
<br />
==Projects==<br />
*[[pH indicator]]<br />
*Dye<br />
<br />
==Handling==<br />
===Safety===<br />
Litmus has low toxicity and no special protection or safety measures are required when handling it. However, it can stain, so wear disposable gear.<br />
<br />
===Storage===<br />
In closed bottles.<br />
<br />
===Disposal===<br />
No special disposal is required. Discard it as you wish.<br />
<br />
===See also ===<br />
* [[Anthocyanin]], a natural dye mix similar in properties to litmus, and much easier to obtain from OTC products, which can be used instead of it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=7357 Indicators]<br />
<br />
[[Category:Materials with no specific chemical formula]]<br />
[[Category:Organic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Litmus&diff=14924Litmus2023-03-19T13:11:36Z<p>Ave369: /* Properties */</p>
<hr />
<div>{{Stub}}<br />
'''Litmus''' is a water soluble mixture of different dyes extracted from lichens. Litmus absorbed onto filter papers is one of the oldest and most known form of [[pH indicator]].<br />
<br />
==Composition==<br />
Litmus contains between 10 to 15 different dyes, with 7-hydroxyphenoxazone being the most important.<br />
<br />
==Properties==<br />
Litmus is a black powder, soluble in [[water]] and some organic solvents, which changes color in acidic environment to red and blue in basic solution. At pH 7 the solution is purple. Litmus, both as solution and as paper strips, is as versatile as [[universal indicator]] strips, but it offers less precise measurements of pH. Still, if you don't need to know the exact pH, litmus strips can be used instead of universal indicator strips.<br />
<br />
Litmus paper strips come in three varieties: red, neutral and blue. Red strips are made from an acidified litmus solution, they are red by default; the inverse is true with blue litmus strips. Red and blue strips are used when you already know if you are testing for an acid or a base, and the color change is more pronounced with them. Red and blue litmus paper strips can usually be converted to neutral litmus paper strips by soaking them in water and drying.<br />
<br />
==Availability==<br />
Litmus powder and litmus paper can be purchased from chemical suppliers and online. It can be found on eBay and Amazon.<br />
<br />
==Isolation==<br />
Litmus can be isolated from certain lichens, though you'll need an appreciable amount of them to extract useful quantities of litmus.<br />
<br />
==Projects==<br />
*[[pH indicator]]<br />
*Dye<br />
<br />
==Handling==<br />
===Safety===<br />
Litmus has low toxicity and no special protection or safety measures are required when handling it. However, it can stain, so wear disposable gear.<br />
<br />
===Storage===<br />
In closed bottles.<br />
<br />
===Disposal===<br />
No special disposal is required. Discard it as you wish.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=7357 Indicators]<br />
<br />
[[Category:Materials with no specific chemical formula]]<br />
[[Category:Organic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Litmus&diff=14923Litmus2023-03-19T13:08:04Z<p>Ave369: /* Properties */</p>
<hr />
<div>{{Stub}}<br />
'''Litmus''' is a water soluble mixture of different dyes extracted from lichens. Litmus absorbed onto filter papers is one of the oldest and most known form of [[pH indicator]].<br />
<br />
==Composition==<br />
Litmus contains between 10 to 15 different dyes, with 7-hydroxyphenoxazone being the most important.<br />
<br />
==Properties==<br />
Litmus is a black powder, soluble in [[water]] and some organic solvents, which changes color in acidic environment to red and blue in basic solution. At pH 7 the solution is purple. Litmus, both as solution and as paper strips, is as versatile as [[universal indicator]] strips, but it offers less precise measurements of pH. Still, if you don't need to know the exact pH, litmus strips can be used instead of universal indicator strips.<br />
<br />
==Availability==<br />
Litmus powder and litmus paper can be purchased from chemical suppliers and online. It can be found on eBay and Amazon.<br />
<br />
==Isolation==<br />
Litmus can be isolated from certain lichens, though you'll need an appreciable amount of them to extract useful quantities of litmus.<br />
<br />
==Projects==<br />
*[[pH indicator]]<br />
*Dye<br />
<br />
==Handling==<br />
===Safety===<br />
Litmus has low toxicity and no special protection or safety measures are required when handling it. However, it can stain, so wear disposable gear.<br />
<br />
===Storage===<br />
In closed bottles.<br />
<br />
===Disposal===<br />
No special disposal is required. Discard it as you wish.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=7357 Indicators]<br />
<br />
[[Category:Materials with no specific chemical formula]]<br />
[[Category:Organic compounds]]<br />
[[Category:Biologically-derived compounds]]<br />
[[Category:PH indicators]]<br />
[[Category:Dyes]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Prospectors&diff=14922Prospectors2023-03-19T00:29:33Z<p>Ave369: </p>
<hr />
<div>'''Prospectors''' are people who practice amateur chemistry for the purpose of extracting [[gold]] from old electronics. In some countries they are common and considered a nuisance to both amateur chemists (since their activities are often illegal and draw law enforcement attention to all home chemists) and old hardware aficionados (since they barbarically destroy rare and valuable vintage devices, not knowing their true value). There may be gold in them thar IBM AT's, but the machines themselves are worth much more in working condition.<br />
<br />
Gold chemistry is not prohibited at all on Sciencemadness, but it may be illegal according the laws of certain countries that regulate gold ownership and circulation. <br />
<br />
== See also ==<br />
<br />
* [[Kewls]]<br />
* [[Cooks]]<br />
<br />
[[Category:Whimsy]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Gold&diff=14921Gold2023-03-19T00:27:40Z<p>Ave369: /* Availability */</p>
<hr />
<div>{{Infobox element<br />
<!-- top --><br />
|image name=<br />
|image alt=<br />
|image size=<br />
|image name comment=<br />
|image name 2=<br />
|image alt 2=<br />
|image size 2=<br />
|image name 2 comment=<br />
<!-- General properties --><br />
|name= Gold<br />
|symbol= Au<br />
|pronounce=<br />
|pronounce ref=<br />
|pronounce comment=<br />
|pronounce 2=<br />
|alt name=Aurum<br />
|alt names=<br />
|allotropes=<br />
|appearance= Metallic dark yellow<br />
<!-- Periodic table --><br />
|above= [[Silver|Ag]]<br />
|below= Rg<br />
|left= [[Platinum]]<br />
|right= [[Mercury]]<br />
|number= 79<br />
|atomic mass= 196.966569(5)<br />
|atomic mass 2=<br />
|atomic mass ref=<br />
|atomic mass comment=<br />
|series= Transition metal<br />
|series ref=<br />
|series comment=<br />
|series color=<br />
|group= <br />
|group ref=<br />
|group comment=11<br />
|period= 6<br />
|period ref=<br />
|period comment=<br />
|block=d<br />
|block ref=<br />
|block comment=<br />
|electron configuration= [Xe] 4f<sup>14</sup> 5d<sup>10</sup> 6s<sup>1</sup><br />
|electron configuration ref=<br />
|electron configuration comment=<br />
|electrons per shell= 2, 8, 18, 32, 18, 1<br />
|electrons per shell ref=<br />
|electrons per shell comment=<br />
<!-- Physical properties --><br />
|physical properties comment=<br />
|color= Metallic yellow<br />
|phase= Solid<br />
|phase ref=<br />
|phase comment=<br />
|melting point K=1337.33<br />
|melting point C=1064.18<br />
|melting point F=1947.52<br />
|melting point ref=<br />
|melting point comment=<br />
|boiling point K=3243<br />
|boiling point C=2970<br />
|boiling point F=5378<br />
|boiling point ref=<br />
|boiling point comment=<br />
|sublimation point K=<br />
|sublimation point C=<br />
|sublimation point F=<br />
|sublimation point ref=<br />
|sublimation point comment=<br />
|density gplstp=<br />
|density gplstp ref=<br />
|density gplstp comment=<br />
|density gpcm3nrt=19.30<br />
|density gpcm3nrt ref=<br />
|density gpcm3nrt comment=<br />
|density gpcm3nrt 2=<br />
|density gpcm3nrt 2 ref=<br />
|density gpcm3nrt 2 comment=<br />
|density gpcm3nrt 3=<br />
|density gpcm3nrt 3 ref=<br />
|density gpcm3nrt 3 comment=<br />
|density gpcm3mp=17.31<br />
|density gpcm3mp ref=<br />
|density gpcm3mp comment=<br />
|density gpcm3bp=<br />
|density gpcm3bp ref=<br />
|density gpcm3bp comment=<br />
|molar volume=<br />
|molar volume unit =<br />
|molar volume ref=<br />
|molar volume comment=<br />
|triple point K=<br />
|triple point kPa=<br />
|triple point ref=<br />
|triple point comment=<br />
|triple point K 2=<br />
|triple point kPa 2=<br />
|triple point 2 ref=<br />
|triple point 2 comment=<br />
|critical point K=<br />
|critical point MPa=<br />
|critical point ref=<br />
|critical point comment=<br />
|heat fusion=12.55<br />
|heat fusion ref=<br />
|heat fusion comment=<br />
|heat fusion 2=<br />
|heat fusion 2 ref=<br />
|heat fusion 2 comment=<br />
|heat vaporization=342<br />
|heat vaporization ref=<br />
|heat vaporization comment=<br />
|heat capacity=25.418<br />
|heat capacity ref=<br />
|heat capacity comment=<br />
|heat capacity 2=<br />
|heat capacity 2 ref=<br />
|heat capacity 2 comment=<br />
|vapor pressure 1= 1646<br />
|vapor pressure 10= 1814<br />
|vapor pressure 100= 2021<br />
|vapor pressure 1 k= 2281<br />
|vapor pressure 10 k= 2620<br />
|vapor pressure 100 k= 3078<br />
|vapor pressure ref=<br />
|vapor pressure comment=<br />
|vapor pressure 1 2=<br />
|vapor pressure 10 2=<br />
|vapor pressure 100 2=<br />
|vapor pressure 1 k 2=<br />
|vapor pressure 10 k 2=<br />
|vapor pressure 100 k 2=<br />
|vapor pressure 2 ref=<br />
|vapor pressure 2 comment=<br />
<!-- Atomic properties --><br />
|atomic properties comment=<br />
|oxidation states= 5, '''3''', 2, '''1''', −1, −2, −3<br />
|oxidation states ref=<br />
|oxidation states comment=(an amphoteric oxide)<br />
|electronegativity=2.54<br />
|electronegativity ref=<br />
|electronegativity comment=<br />
|ionization energy 1=890.1<br />
|ionization energy 1 ref=<br />
|ionization energy 1 comment=<br />
|ionization energy 2=1980<br />
|ionization energy 2 ref=<br />
|ionization energy 2 comment=<br />
|ionization energy 3=<br />
|ionization energy 3 ref=<br />
|ionization energy 3 comment=<br />
|number of ionization energies=<br />
|ionization energy ref=<br />
|ionization energy comment=<br />
|atomic radius=144<br />
|atomic radius ref=<br />
|atomic radius comment=<br />
|atomic radius calculated=<br />
|atomic radius calculated ref=<br />
|atomic radius calculated comment=<br />
|covalent radius=136±6<br />
|covalent radius ref=<br />
|covalent radius comment=<br />
|Van der Waals radius=166<br />
|Van der Waals radius ref=<br />
|Van der Waals radius comment=<br />
<!-- Miscellanea --><br />
|crystal structure=<br />
|crystal structure prefix=<br />
|crystal structure ref=<br />
|crystal structure comment= Face-centered cubic (fcc)<br />
|crystal structure 2=<br />
|crystal structure 2 prefix=<br />
|crystal structure 2 ref=<br />
|crystal structure 2 comment=<br />
|speed of sound=<br />
|speed of sound ref=<br />
|speed of sound comment=<br />
|speed of sound rod at 20=<br />
|speed of sound rod at 20 ref=<br />
|speed of sound rod at 20 comment=<br />
|speed of sound rod at r.t.=2030<br />
|speed of sound rod at r.t. ref=<br />
|speed of sound rod at r.t. comment=<br />
|thermal expansion=<br />
|thermal expansion ref=<br />
|thermal expansion comment=<br />
|thermal expansion at 25=14.2<br />
|thermal expansion at 25 ref=<br />
|thermal expansion at 25 comment=<br />
|thermal conductivity=318<br />
|thermal conductivity ref=<br />
|thermal conductivity comment=<br />
|thermal conductivity 2=<br />
|thermal conductivity 2 ref=<br />
|thermal conductivity 2 comment=<br />
|thermal diffusivity=<br />
|thermal diffusivity ref=<br />
|thermal diffusivity comment=<br />
|electrical resistivity=<br />
|electrical resistivity unit prefix=<br />
|electrical resistivity ref=<br />
|electrical resistivity comment=<br />
|electrical resistivity at 0=<br />
|electrical resistivity at 0 ref=<br />
|electrical resistivity at 0 comment=<br />
|electrical resistivity at 20=22.14<br />
|electrical resistivity at 20 ref=<br />
|electrical resistivity at 20 comment=<br />
|band gap=<br />
|band gap ref=<br />
|band gap comment=<br />
|Curie point K=<br />
|Curie point ref=<br />
|Curie point comment=<br />
|magnetic ordering= Diamagnetic<br />
|magnetic ordering ref=<br />
|magnetic ordering comment=<br />
|tensile strength=120<br />
|tensile strength ref=<br />
|tensile strength comment=<br />
|Young's modulus=79<br />
|Young's modulus ref=<br />
|Young's modulus comment=<br />
|Shear modulus=27<br />
|Shear modulus ref=<br />
|Shear modulus comment=<br />
|Bulk modulus=180<br />
|Bulk modulus ref=<br />
|Bulk modulus comment=<br />
|Poisson ratio=0.4<br />
|Poisson ratio ref=<br />
|Poisson ratio comment=<br />
|Mohs hardness=2.5<br />
|Mohs hardness ref=<br />
|Mohs hardness comment=<br />
|Mohs hardness 2=<br />
|Mohs hardness 2 ref=<br />
|Mohs hardness 2 comment=<br />
|Vickers hardness=188–216<br />
|Vickers hardness ref=<br />
|Vickers hardness comment=<br />
|Brinell hardness=188–245<br />
|Brinell hardness ref=<br />
|Brinell hardness comment=<br />
|CAS number=7440-57-5<br />
|CAS number ref=<br />
|CAS number comment=<br />
<!-- History --><br />
|naming=from Latin ''aurum'' (gold)<br />
|predicted by=<br />
|prediction date ref=<br />
|prediction date=<br />
|discovered by=<br />
|discovery date ref=<br />
|discovery date=~6000 BCE (Middle East)<br />
|first isolation by=<br />
|first isolation date ref=<br />
|first isolation date=<br />
|discovery and first isolation by=<br />
|named by=<br />
|named date ref=<br />
|named date=<br />
|history comment label=<br />
|history comment=<br />
<!-- Isotopes --><br />
|isotopes=<br />
|isotopes comment=<br />
|engvar=<br />
}}<br />
'''Gold''' is a chemical element with the symbol '''Au''' and atomic number 79. It is a transitional metal, part of Group 11, the same group as [[silver]] and [[copper]]. It's well known for its corrosion resistance and its high economic value. Gold is mainly used in jewels, electronics, catalyst and as exchange.<br />
<br />
The symbol Au comes from the latin name of gold, '''aurum''', and derivatives of this term are used in many countries as designation for gold, most often in Romance-speaking countries.<br />
<br />
==Properties==<br />
===Chemical===<br />
Gold is very resistant to acid and alkali attacks and does not react with oxygen or halogens at standard conditions. However a mixture of [[hydrochloric acid]] and [[nitric acid]] known as [[aqua regia]] will dissolve gold.<br />
<br />
:Au + HNO<sub>3</sub> + 4 HCl → HAuCl<sub>4</sub> + NO + 2 H<sub>2</sub>O<br />
<br />
Gold can also be dissolved by cyanides, such as [[sodium cyanide]], a process used in gold extraction, when the gold concentration is low. [[Mercury]] dissolves gold forming an amalgam.<br />
<br />
Gold resists the attack of molten [[sodium hydroxide]], however, at temperatures above 700 °C, there is visible corrosion of the metal, and traces of gold flakes and gold oxide can be observed in the alkali melt. Small amounts of metallic [[sodium]] have also been observed, which rapidly form an alloy with the gold, which is stable enough that it doesn't readily react with water or acids.<ref>http://pubs.acs.org/doi/pdf/10.1021/ja01601a004</ref><br />
<br />
Gold is unaffected by concentrated (40%) [[hydrofluoric acid]] at standard conditions.<ref>https://www.youtube.com/watch?v=Ri8heWPz5zY</ref><br />
<br />
===Physical===<br />
[[File:Gold.jpg|thumb|250px|Small sample of gold foil]]<br />
Gold is a bright yellow dense, soft, malleable and ductile metal. Very pure gold (24 carat) is soft enough to be dent by biting it, a practice occasionally seen in gold diggers and Olympic athletes, who traditionally bit their gold medals. Gold is the most malleable of all metals, one gram can be beaten into a sheet of 1 square meter. It has high thermal and electric conductivity, properties that gives it many uses in electronics. Its density of 19.3 g/cm<sup>3</sup> is slightly higher than that of [[tungsten]] and [[uranium]].<br />
<br />
==Availability==<br />
Gold can be found in nature as nuggets, either pure or mixed with [[silver]] or platinum group metals. During the Gold Rush, very large nuggets were dug up from the rivers. Nowadays, nuggets tend to be rarer, instead grain sized gold is more often found, as previous extraction methods focused on large nuggets. Extracting gold from gold-rich soil/sand is very intensive and may not be 100% legal depending on where you live.<br />
<br />
Gold can be extracted from jewelry, but doing so often destroys jewelry that would cost more than the gold it is made of. Gold bullions and coins are also a source of gold, albeit an expensive one.<br />
<br />
Gold leaf, used in food decorations are also a source of gold, albeit the quantity is small and it's usually a gold alloy.<br />
<br />
However the most sought source of gold are scrap electronics. Extracting gold from old electronics such as finger and socket contacts, pins, CPUs, RAM chips, board plating, adjustable switches, etc. is one of the most known aspects of amateur chemistry (see [[Prospectors]]). Usually the older the electronic device is, the more gold it has. Extracting the gold is done by various methods: dissolving the copper circuit with a PCB etchant, such as [[Iron(III) chloride|ferric chloride]] and collecting the gold foil by filtering the solution, which is later purified by dissolving it in [[aqua regia]] and melted; dissolving the boards in cyanide solution, reducing the gold cyanide compound and melting the powder; dissolving the gold with [[mercury]] and extracting the gold; electrochemical separation. The amount of gold obtained is low, but it's a cheap source.<br />
<br />
Gold itself is usually found uncombined in nature, but when found as a chemical compound, it is most often combined with tellurium, in the form of calaverite and krennerite (two different polymorphs of AuTe<sub>2</sub>), petzite (Ag<sub>3</sub>AuTe<sub>2</sub>) and sylvanite (AgAuTe<sub>4</sub>).<br />
<br />
==Preparation==<br />
Gold can be reduced from its salts by reducing it with a reducing compound. Since gold sits close to the bottom of the reactivity scale, any common metal will reduce it to its elemental form. In case of [[chloroauric acid]], [[sodium metabisulfite|sodium]] or [[potassium metabisulfite]] are commonly used as reducing agents, as they're cheap and readily available.<br />
<br />
==Projects==<br />
*Gold plating<br />
*Make gold colloids<br />
*Gold electrode in water electrolysis<br />
*Catalyst<br />
*Make jewelry<br />
*Element collection<br />
<br />
==Handling==<br />
===Safety===<br />
Pure gold is non-toxic and it's even used in medical implants. On the other hand, most gold compounds (especially the salts) are toxic and they should be handled with proper protection.<br />
<br />
===Storage===<br />
No special storage is required for bulk and powdered gold. Though given the value of gold, it's best to keep it in a hidden place or a safe.<br />
<br />
===Disposal===<br />
Due to gold's price and rarity, it's best to try and recycle as much gold as possible.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=41465 chemistry of GOLD, known reactions and techniques]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=63606 Gold Recovery Computer Scrap Complete Process]<br />
<br />
[[Category:Elements]]<br />
[[Category:Metals]]<br />
[[Category:Transition metals]]<br />
[[Category:Precious metals]]<br />
[[Category:D-block]]<br />
[[Category:Minerals]]<br />
[[Category:Solids]]<br />
[[Category:Noble metals]]<br />
[[Category:Inert chemicals]]<br />
[[Category:Coinage metals]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Prospectors&diff=14920Prospectors2023-03-19T00:24:43Z<p>Ave369: Created page with "'''Prospectors''' are people who practice amateur chemistry for the purpose of extracting gold from old electronics. In some countries they are common and considered a nuisanc..."</p>
<hr />
<div>'''Prospectors''' are people who practice amateur chemistry for the purpose of extracting gold from old electronics. In some countries they are common and considered a nuisance to both amateur chemists (since their activities are often illegal and draw law enforcement attention to all home chemists) and old hardware aficionados (since they barbarically destroy rare and valuable vintage devices, not knowing their true value). There may be gold in them thar IBM AT's, but the machines themselves are worth much more in working condition.<br />
<br />
Gold chemistry is not prohibited at all on Sciencemadness, but it may be illegal according the laws of certain countries that regulate gold ownership and circulation. <br />
<br />
== See also ==<br />
<br />
* [[Kewls]]<br />
* [[Cooks]]<br />
<br />
[[Category:Whimsy]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Chloric_acid&diff=14426Chloric acid2022-04-26T14:17:59Z<p>Ave369: /* Chemical */</p>
<hr />
<div>{{Chembox<br />
| Name = Chloric acid<br />
| Reference = <br />
| IUPACName = Chloric acid<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Chloric(V) acid<br>Hydrogen trioxochlorate(V)<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless unstable liquid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 1 g/mL, solution (approximate)<br />
| Formula = HClO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 84.45914 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| pKa = -1<br />
| pKb = <br />
| Solubility = >40 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in [[methanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = Pyramidal<br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/Kr8EB6J/chloric-acid-35-sa.pdf.html Sigma-Aldrich] (35% aq. sol.)<br />
| FlashPt = <br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Hydrochloric acid]]<br>[[Hypochlorous acid]]<br>[[Chlorous acid]]<br>[[Perchloric acid]]<br />
}}<br />
}}<br />
'''Chloric acid''' is an unstable, strong, oxidizing acid that can exist in aqueous solutions up to 40%. It is one of the oxoacids of chlorine. Its formula is '''HClO<sub>3</sub>'''.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
Chloric acid solutions are colorless liquids. Their density is similar to that of water. Concentrated chloric acid (over 30%) has a pungent smell, because of its disproportionation and release of chlorine and chlorine dioxide. <br />
<br />
Chloric acid is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, chloric acid solutions decompose to give a variety of products, for example:<br />
:8 HClO<sub>3</sub> → 4 HClO<sub>4</sub> + 2 H<sub>2</sub>O + 2 Cl<sub>2</sub> + 3 O<sub>2</sub><br />
:3 HClO<sub>3</sub> → HClO<sub>4</sub> + H<sub>2</sub>O + 2 ClO<sub>2</sub><br />
<br />
=== Chemical ===<br />
<br />
Chloric acid is a strong oxidizer. Concentrated solutions can light organic materials on fire. It is also a strong acid. It reacts with bases to form chlorate salts.<br />
<br />
The reaction between chloric acid and metals can take one of the two possible courses. With active metals (sodium to aluminum), little or no reduction of the acid occurs, and the metal reacts with evolution of hydrogen. With less active metals (iron to copper), chloric acid acts like an oxidizing acid and dissolves the metal without evolving any gas, producing chlorates, chlorides and possibly oxides<ref>*[https://scholarworks.uni.edu/cgi/viewcontent.cgi?article=7036&context=pias Reaction of Chloric Acid with Metals]</ref>.<br />
<br />
When one tries to prepare an overconcentrated solution of this acid (over 40% under vacuum, over 30% by normal pressure boiling), it disproportionates to give a variety of products, which always include [[perchloric acid]], and gases such as [[chlorine]] and [[chlorine dioxide]] are released. Because of this, chloric acid can serve as a precursor chemical to perchloric acid.<br />
<br />
== Preparation ==<br />
<br />
The easiest way to prepare this acid is reacting [[barium chlorate]] with [[sulfuric acid]]. Concentrations should be stoichiometrically calculated to avoid accidentally preparing an overconcentrated solution that will decompose instantly.<br />
<br />
== Projects ==<br />
<br />
* Make [[perchloric acid]]<br />
* Make various exotic chlorates for your fireworks, i.e. strontium chlorate for red, calcium chlorate for orange, etc.)<br />
<br />
== Handling ==<br />
<br />
=== Safety ===<br />
<br />
Chloric acid is corrosive, similarly to nitric acid of the same concentrations. It is also known to react violently with oxidizable organic materials. The products of its decomposition are toxic gases. Its ability to set flammable materials on fire is comparable to fuming nitric acid.<br />
<br />
=== Storage ===<br />
<br />
Solutions up to 30% can be stored in bottles of amber glass, with ample headroom for gases evolving when the acid decomposes. The acid should be kept cool to avoid decomposition. 31-40% solutions are not advised to store.<br />
<br />
=== Disposal ===<br />
<br />
One should not dispose of chloric acid directly into the environment. A reducing agent such as [[sodium metabisulfite|sodium]] or [[potassium metabisulfite]] can be used to neutralize it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13746 chloric acid solution becomes yellow ...produces a peculiar noise]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2904 Chloric Acid]<br />
<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Chlorates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials stable only in solution]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Chloric_acid&diff=14425Chloric acid2022-04-26T14:17:37Z<p>Ave369: /* References */</p>
<hr />
<div>{{Chembox<br />
| Name = Chloric acid<br />
| Reference = <br />
| IUPACName = Chloric acid<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Chloric(V) acid<br>Hydrogen trioxochlorate(V)<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless unstable liquid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 1 g/mL, solution (approximate)<br />
| Formula = HClO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 84.45914 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| pKa = -1<br />
| pKb = <br />
| Solubility = >40 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in [[methanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = Pyramidal<br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/Kr8EB6J/chloric-acid-35-sa.pdf.html Sigma-Aldrich] (35% aq. sol.)<br />
| FlashPt = <br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Hydrochloric acid]]<br>[[Hypochlorous acid]]<br>[[Chlorous acid]]<br>[[Perchloric acid]]<br />
}}<br />
}}<br />
'''Chloric acid''' is an unstable, strong, oxidizing acid that can exist in aqueous solutions up to 40%. It is one of the oxoacids of chlorine. Its formula is '''HClO<sub>3</sub>'''.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
Chloric acid solutions are colorless liquids. Their density is similar to that of water. Concentrated chloric acid (over 30%) has a pungent smell, because of its disproportionation and release of chlorine and chlorine dioxide. <br />
<br />
Chloric acid is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, chloric acid solutions decompose to give a variety of products, for example:<br />
:8 HClO<sub>3</sub> → 4 HClO<sub>4</sub> + 2 H<sub>2</sub>O + 2 Cl<sub>2</sub> + 3 O<sub>2</sub><br />
:3 HClO<sub>3</sub> → HClO<sub>4</sub> + H<sub>2</sub>O + 2 ClO<sub>2</sub><br />
<br />
=== Chemical ===<br />
<br />
Chloric acid is a strong oxidizer. Concentrated solutions can light organic materials on fire. It is also a strong acid. It reacts with bases to form chlorate salts.<br />
<br />
The reaction between chloric acid and metals can take one of the two possible courses. With active metals (sodium to aluminum), little or no reduction of the acid occurs, and the metal reacts with evolution of hydrogen. With less active metals (iron to copper), chloric acid acts like an oxidizing acid and dissolves the metal without evolving any gas, producing chlorates, chlorides and possibly oxides.<br />
<br />
When one tries to prepare an overconcentrated solution of this acid (over 40% under vacuum, over 30% by normal pressure boiling), it disproportionates to give a variety of products, which always include [[perchloric acid]], and gases such as [[chlorine]] and [[chlorine dioxide]] are released. Because of this, chloric acid can serve as a precursor chemical to perchloric acid.<br />
<br />
== Preparation ==<br />
<br />
The easiest way to prepare this acid is reacting [[barium chlorate]] with [[sulfuric acid]]. Concentrations should be stoichiometrically calculated to avoid accidentally preparing an overconcentrated solution that will decompose instantly.<br />
<br />
== Projects ==<br />
<br />
* Make [[perchloric acid]]<br />
* Make various exotic chlorates for your fireworks, i.e. strontium chlorate for red, calcium chlorate for orange, etc.)<br />
<br />
== Handling ==<br />
<br />
=== Safety ===<br />
<br />
Chloric acid is corrosive, similarly to nitric acid of the same concentrations. It is also known to react violently with oxidizable organic materials. The products of its decomposition are toxic gases. Its ability to set flammable materials on fire is comparable to fuming nitric acid.<br />
<br />
=== Storage ===<br />
<br />
Solutions up to 30% can be stored in bottles of amber glass, with ample headroom for gases evolving when the acid decomposes. The acid should be kept cool to avoid decomposition. 31-40% solutions are not advised to store.<br />
<br />
=== Disposal ===<br />
<br />
One should not dispose of chloric acid directly into the environment. A reducing agent such as [[sodium metabisulfite|sodium]] or [[potassium metabisulfite]] can be used to neutralize it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13746 chloric acid solution becomes yellow ...produces a peculiar noise]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2904 Chloric Acid]<br />
<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Chlorates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials stable only in solution]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Chloric_acid&diff=14424Chloric acid2022-04-26T14:17:15Z<p>Ave369: /* Storage */</p>
<hr />
<div>{{Chembox<br />
| Name = Chloric acid<br />
| Reference = <br />
| IUPACName = Chloric acid<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Chloric(V) acid<br>Hydrogen trioxochlorate(V)<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless unstable liquid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 1 g/mL, solution (approximate)<br />
| Formula = HClO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 84.45914 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| pKa = -1<br />
| pKb = <br />
| Solubility = >40 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in [[methanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = Pyramidal<br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/Kr8EB6J/chloric-acid-35-sa.pdf.html Sigma-Aldrich] (35% aq. sol.)<br />
| FlashPt = <br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Hydrochloric acid]]<br>[[Hypochlorous acid]]<br>[[Chlorous acid]]<br>[[Perchloric acid]]<br />
}}<br />
}}<br />
'''Chloric acid''' is an unstable, strong, oxidizing acid that can exist in aqueous solutions up to 40%. It is one of the oxoacids of chlorine. Its formula is '''HClO<sub>3</sub>'''.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
Chloric acid solutions are colorless liquids. Their density is similar to that of water. Concentrated chloric acid (over 30%) has a pungent smell, because of its disproportionation and release of chlorine and chlorine dioxide. <br />
<br />
Chloric acid is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, chloric acid solutions decompose to give a variety of products, for example:<br />
:8 HClO<sub>3</sub> → 4 HClO<sub>4</sub> + 2 H<sub>2</sub>O + 2 Cl<sub>2</sub> + 3 O<sub>2</sub><br />
:3 HClO<sub>3</sub> → HClO<sub>4</sub> + H<sub>2</sub>O + 2 ClO<sub>2</sub><br />
<br />
=== Chemical ===<br />
<br />
Chloric acid is a strong oxidizer. Concentrated solutions can light organic materials on fire. It is also a strong acid. It reacts with bases to form chlorate salts.<br />
<br />
The reaction between chloric acid and metals can take one of the two possible courses. With active metals (sodium to aluminum), little or no reduction of the acid occurs, and the metal reacts with evolution of hydrogen. With less active metals (iron to copper), chloric acid acts like an oxidizing acid and dissolves the metal without evolving any gas, producing chlorates, chlorides and possibly oxides.<br />
<br />
When one tries to prepare an overconcentrated solution of this acid (over 40% under vacuum, over 30% by normal pressure boiling), it disproportionates to give a variety of products, which always include [[perchloric acid]], and gases such as [[chlorine]] and [[chlorine dioxide]] are released. Because of this, chloric acid can serve as a precursor chemical to perchloric acid.<br />
<br />
== Preparation ==<br />
<br />
The easiest way to prepare this acid is reacting [[barium chlorate]] with [[sulfuric acid]]. Concentrations should be stoichiometrically calculated to avoid accidentally preparing an overconcentrated solution that will decompose instantly.<br />
<br />
== Projects ==<br />
<br />
* Make [[perchloric acid]]<br />
* Make various exotic chlorates for your fireworks, i.e. strontium chlorate for red, calcium chlorate for orange, etc.)<br />
<br />
== Handling ==<br />
<br />
=== Safety ===<br />
<br />
Chloric acid is corrosive, similarly to nitric acid of the same concentrations. It is also known to react violently with oxidizable organic materials. The products of its decomposition are toxic gases. Its ability to set flammable materials on fire is comparable to fuming nitric acid.<br />
<br />
=== Storage ===<br />
<br />
Solutions up to 30% can be stored in bottles of amber glass, with ample headroom for gases evolving when the acid decomposes. The acid should be kept cool to avoid decomposition. 31-40% solutions are not advised to store.<br />
<br />
=== Disposal ===<br />
<br />
One should not dispose of chloric acid directly into the environment. A reducing agent such as [[sodium metabisulfite|sodium]] or [[potassium metabisulfite]] can be used to neutralize it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13746 chloric acid solution becomes yellow ...produces a peculiar noise]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2904 Chloric Acid]<br />
*[https://scholarworks.uni.edu/cgi/viewcontent.cgi?article=7036&context=pias Reaction of Chloric Acid with Metals]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Chlorates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials stable only in solution]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Chloric_acid&diff=14423Chloric acid2022-04-26T14:16:37Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Chloric acid<br />
| Reference = <br />
| IUPACName = Chloric acid<br />
| PIN = <br />
| SystematicName = <br />
| OtherNames = Chloric(V) acid<br>Hydrogen trioxochlorate(V)<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless unstable liquid<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 1 g/mL, solution (approximate)<br />
| Formula = HClO<sub>3</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 84.45914 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| pKa = -1<br />
| pKb = <br />
| Solubility = >40 g/100 ml (20 °C)<br />
| SolubleOther = Soluble in [[methanol]]<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = Pyramidal<br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = <br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/Kr8EB6J/chloric-acid-35-sa.pdf.html Sigma-Aldrich] (35% aq. sol.)<br />
| FlashPt = <br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br>Corrosive<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Hydrochloric acid]]<br>[[Hypochlorous acid]]<br>[[Chlorous acid]]<br>[[Perchloric acid]]<br />
}}<br />
}}<br />
'''Chloric acid''' is an unstable, strong, oxidizing acid that can exist in aqueous solutions up to 40%. It is one of the oxoacids of chlorine. Its formula is '''HClO<sub>3</sub>'''.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
Chloric acid solutions are colorless liquids. Their density is similar to that of water. Concentrated chloric acid (over 30%) has a pungent smell, because of its disproportionation and release of chlorine and chlorine dioxide. <br />
<br />
Chloric acid is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, chloric acid solutions decompose to give a variety of products, for example:<br />
:8 HClO<sub>3</sub> → 4 HClO<sub>4</sub> + 2 H<sub>2</sub>O + 2 Cl<sub>2</sub> + 3 O<sub>2</sub><br />
:3 HClO<sub>3</sub> → HClO<sub>4</sub> + H<sub>2</sub>O + 2 ClO<sub>2</sub><br />
<br />
=== Chemical ===<br />
<br />
Chloric acid is a strong oxidizer. Concentrated solutions can light organic materials on fire. It is also a strong acid. It reacts with bases to form chlorate salts.<br />
<br />
The reaction between chloric acid and metals can take one of the two possible courses. With active metals (sodium to aluminum), little or no reduction of the acid occurs, and the metal reacts with evolution of hydrogen. With less active metals (iron to copper), chloric acid acts like an oxidizing acid and dissolves the metal without evolving any gas, producing chlorates, chlorides and possibly oxides.<br />
<br />
When one tries to prepare an overconcentrated solution of this acid (over 40% under vacuum, over 30% by normal pressure boiling), it disproportionates to give a variety of products, which always include [[perchloric acid]], and gases such as [[chlorine]] and [[chlorine dioxide]] are released. Because of this, chloric acid can serve as a precursor chemical to perchloric acid.<br />
<br />
== Preparation ==<br />
<br />
The easiest way to prepare this acid is reacting [[barium chlorate]] with [[sulfuric acid]]. Concentrations should be stoichiometrically calculated to avoid accidentally preparing an overconcentrated solution that will decompose instantly.<br />
<br />
== Projects ==<br />
<br />
* Make [[perchloric acid]]<br />
* Make various exotic chlorates for your fireworks, i.e. strontium chlorate for red, calcium chlorate for orange, etc.)<br />
<br />
== Handling ==<br />
<br />
=== Safety ===<br />
<br />
Chloric acid is corrosive, similarly to nitric acid of the same concentrations. It is also known to react violently with oxidizable organic materials. The products of its decomposition are toxic gases. Its ability to set flammable materials on fire is comparable to fuming nitric acid.<br />
<br />
=== Storage ===<br />
<br />
Solutions up to 30% can be stored in bottles of amber glass, with ample headroom for gases evolving when the acid decomposes. 31-40% solutions are not advised to store.<br />
<br />
=== Disposal ===<br />
<br />
One should not dispose of chloric acid directly into the environment. A reducing agent such as [[sodium metabisulfite|sodium]] or [[potassium metabisulfite]] can be used to neutralize it.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13746 chloric acid solution becomes yellow ...produces a peculiar noise]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=2904 Chloric Acid]<br />
*[https://scholarworks.uni.edu/cgi/viewcontent.cgi?article=7036&context=pias Reaction of Chloric Acid with Metals]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Acids]]<br />
[[Category:Mineral acids]]<br />
[[Category:Strong acids]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Chlorates]]<br />
[[Category:Unstable materials]]<br />
[[Category:Materials stable only in solution]]<br />
[[Category:Corrosive chemicals]]<br />
[[Category:Liquids]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Carbon_dioxide&diff=14419Carbon dioxide2022-04-10T17:39:54Z<p>Ave369: /* Physical */</p>
<hr />
<div>{{Chembox<br />
| Name = Carbon dioxide<br />
| Reference =<br />
| IUPACName = Carbon dioxide<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Carbon oxide<br>Carbon(IV) oxide<br>Carbonic acid gas<br>Carbonic anhydride<br>Carbonic oxide<br>Dry ice (solid phase)<br />
<!-- Images --><br />
| ImageFile = Carbon dioxide.png<br />
| ImageSize = 250<br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Colorless gas<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = <br />
| Density = 1,562 kg/m<sup>3</sup> (solid at 1 atm and −78.5 °C)<br>1,101 kg/m<sup>3</sup> (liquid at saturation −37 °C)<br>1.977 kg/m<sup>3</sup> (gas at 1 atm and 0 °C)<br />
| Formula = CO<sub>2</sub><br />
| HenryConstant = <br />
| LogP = 0.83<br />
| MolarMass = 44.01 g/mol<br />
| MeltingPt = <br />
| MeltingPtC = −56.6<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = (Triple point at 5.1 atm)<br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = ''760 mm Hg''<br>171 ml/100 ml at 0 °C<br>88 ml/100 ml at 20 °C<br>80 ml/100 ml at 25 °C <br>36 ml/100 ml at 60 °C<br />
| SolubleOther = Soluble in organic solvents<br />
| Solubility1 = 8.2 ml/g (20 °C)<br />
| Solvent1 = acetone<br />
| Solubility2 = 2.71 ml/g (20 °C)<br />
| Solvent2 = benzene<br />
| Solubility3 = 6.3 ml/g (20 °C)<br />
| Solvent3 = diethyl ether<br />
| Solubility4 = 3.6 ml/g (20 °C)<br />
| Solvent4 = ethanol<br />
| Solubility5 = 2.8 ml/g (20 °C)<br />
| Solvent5 = heptane<br />
| Solubility6 = 4.1 ml/g (20 °C)<br />
| Solvent6 = methanol<br />
| Solubility7 = 7.4 ml/g (20 °C)<br />
| Solvent7 = methyl acetate<br />
| Solubility8 = 3.0 ml/g (20 °C)<br />
| Solvent8 = toluene<br />
| Solubility9 = 2.31 ml/g (20 °C)<br />
| Solvent9 = xylene<br />
| VaporPressure = 5.73 MPa (20 °C)<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = [https://www.docdroid.net/xhBOse4/carbon-dioxide-sa.pdf.html Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Asphyxiant<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Carbon monoxide]]<br />
}}<br />
}}<br />
'''Carbon dioxide''', one of the more common gasses in earth's atmosphere, is a molecule containing one [[carbon]] atom and two [[oxygen]] atoms ('''CO<sub>2</sub>'''). It is related to [[carbon monoxide]].<br />
<br />
==Properties==<br />
===Chemical===<br />
Carbon dioxide is soluble in water, making [[carbonic acid]]. This is not very stable, existing as an equilibrium between carbonic acid and carbon dioxide/water, and cannot be concentrated.<br />
<br />
:CO<sub>2</sub> + H<sub>2</sub>O ⇌ H<sub>2</sub>CO<sub>3</sub><br />
<br />
However, this limited temporary solubility can still be taken advantage of in converting hydroxides to carbonates, such as in the treatment of [[Calcium hydroxide|limewater]] with carbon dioxide to precipitate calcium carbonate. Carbon dioxide can also be used as an acid to oxidize some species, such as the oxidation of [[potassium manganate]] to [[Potassium permanganate|permanganate]].<br />
<br />
Carbon dioxide is released in many chemical reactions, usually because of the acidification of a [[carbonate]].<br />
<br />
[[Magnesium]], [[titanium]], [[potassium]], and other reactive metals will reduce carbon dioxide to [[carbon]]. This can be achieved by burning said metals in an atmosphere of carbon dioxide.<br />
<br />
===Physical===<br />
Carbon dioxide is a colorless gas that is odorless unless at very high concentration (at which it has a sharp, sour smell of freshly opened soda pop). It cannot be isolated in liquid form at room temperature and atmospheric pressure, but sublimes. In solid form, it is called "dry ice", and used for chilling experiments. Carbon dioxide is a relatively cheap inert gas, and although for most chemicals it may cause reaction, it is a good storage gas for most reactive metals, except the metals from the s-block and lanthanides.<br />
<br />
==Availability==<br />
Carbon dioxide is available by generation from common materials (like [[sodium bicarbonate]]), from the air, in canisters, or in the form of dry ice, which may be sold at grocery stores and wedding planner entities. Compressed CO<sub>2</sub> tanks can also be found in many aquarium stores.<br />
<br />
Dry ice may also be produced by placing the end of a fire extinguisher(which often contains carbon dioxide) into a bag and releasing the pressurized gas. Fresh dry ice is formed within the confines of the bag.<br />
<br />
==Preparation==<br />
Acidification of a [[carbonate]] or [[bicarbonate]] releases carbon dioxide, which is a cheap way to generate it.<br />
<br />
: CaCO<sub>3</sub> + H<sub>2</sub>SO<sub>4</sub> → CaSO<sub>4</sub> + CO<sub>2</sub> + H<sub>2</sub>O<br />
<br />
Burning organic materials will also generate carbon dioxide, but will also generate soot, water vapor, formaldehyde, aromatic compounds, so it's best to bubble the resulting smoke into a liquid or an adsorption column that will remove the contaminants.<br />
<br />
==Projects==<br />
*Cooling baths<br />
*Carbonates<br />
*Make [[calcium bicarbonate]] solution<br />
*Gas nozzles<br />
*Inert atmosphere (not useful for bases, block A metals, most organometallic compounds)<br />
<br />
==Handling==<br />
===Safety===<br />
Carbon dioxide is not directly poisonous, like [[carbon monoxide]], but it will lower the pH of the blood to dangerous levels if it is inhaled in excess. Carbon dioxide<br />
<br />
The extreme cold of solid carbon dioxide can immediately cause reverse thermal burns if it comes into contact with skin.<br />
<br />
===Storage===<br />
Compressed carbon dioxide should be stored in cold places, away from heat. Dry ice should be stored in an insulated container, polystyrene box, dewar, as well away from heat sources. Both forms should be stored in an area with good ventilation, as CO<sub>2</sub> may build-up and pose an asphyxiation hazard. The gas is heavier than air and may pour down and accumulate in low areas.<br />
<br />
===Disposal===<br />
Carbon dioxide can be released in open air. Avoid releasing it in an enclosed space.<br />
<br />
==References==<br />
<references /><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=15185 Co2 in beverages?]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Carbon compounds]]<br />
[[Category:Oxides]]<br />
[[Category:Carbon oxides]]<br />
[[Category:Inorganic acid anhydrides]]<br />
[[Category:Inert atmospheres]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Barium_peroxide&diff=14389Barium peroxide2022-02-13T17:04:40Z<p>Ave369: </p>
<hr />
<div>{{Chembox<br />
| Name = Barium peroxide<br />
| Reference =<br />
| IUPACName = Barium peroxide<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Barium binoxide<br>Barium dioxide<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Grey-white crystalline (anhydrous)<br>Colorless solid (octahydrate) <br />
| BoilingPt = <br />
| BoilingPtC = 800<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 5.68 g/cm<sup>3</sup> (anhydrous)<br>2.292 g/cm<sup>3</sup> (octahydrate)<br />
| Formula = BaO<sub>2</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 169.33 g/mol (anhydrous)<br>313.45 (octahydrate)<br />
| MeltingPt = <br />
| MeltingPtC = 450<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = ''anhydrous''<br> 0.091 g/100 ml (20 °C)<hr>''octahydrate''<br>0.168 g/100 ml<br />
| SolubleOther = Reacts with acids<br>Insoluble in hydrocarbons<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = 6<br />
| CrystalStruct = Tetragonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [https://www.docdroid.net/JJosUxN/barium-peroxide-sa.pdf Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Barium oxide]]<br />
}}<br />
}}<br />
'''Barium peroxide''' is the inorganic compound with the formula '''BaO<sub>2</sub>'''. It is the barium salt of [[hydrogen peroxide]].<br />
<br />
==Properties==<br />
===Chemical===<br />
Barium peroxide can be used to produce highly concentrated [[hydrogen peroxide]] via its reaction with conc. [[sulfuric acid]].<br />
<br />
: BaO<sub>2</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>O<sub>2</sub> + BaSO<sub>4</sub><br />
<br />
The insoluble [[barium sulfate]] is filtered from the mixture.<br />
<br />
===Physical===<br />
Barium peroxide is a white solid, insoluble in water.<br />
<br />
==Availability==<br />
Barium peroxide is sold by lab suppliers. So far there aren't any sellers on eBay and Amazon.<br />
<br />
==Preparation==<br />
Barium peroxide can be made by the reversible reaction of O<sub>2</sub> with [[barium oxide]]. The peroxide forms around 500 °C and oxygen is released above 820 °C.<br />
<br />
: 2 BaO + O<sub>2</sub> ⇌ 2 BaO<sub>2</sub><br />
<br />
A different, aqueous preparation is performed through a metathesis reaction with barium nitrate and hydrogen peroxide:<br />
<br />
: Ba(NO<sub>3</sub>)<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> ⇌ BaO<sub>2</sub> + 2HNO<sub>3</sub><br />
<br />
Cool the mixed solution to near freezing, and the crystallohydrate BaO<sub>2</sub>*8H<sub>2</sub>O will precipitate. Calcine the crystallohydrate carefully at the temperature of 100-120 degrees Celsius to convert it to the anhydrous salt.<br />
<br />
==Projects==<br />
*Flash powders and fireworks<br />
*Make concentrated [[hydrogen peroxide]]<br />
<br />
==Handling==<br />
===Safety===<br />
Barium peroxide is corrosive and an oxidizer. Proper protection must be worn when handling this compound.<br />
<br />
===Storage===<br />
Barium peroxide should be kept in closed plastic or glass bottles.<br />
<br />
===Disposal===<br />
Can be neutralized by adding diluted sulfuric acid then iron oxide, to decompose the hydrogen peroxide.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=888 Barium Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=42263 Barium Peroxide -> Hydrogen Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16714 KMnO4 and BaO2 reaction]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13965 Synthesis of BaO2 from Ba(NO3)2]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Barium compounds]]<br />
[[Category:Oxides]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Peroxides]]<br />
[[Category:Inorganic peroxides]]<br />
[[Category:Insoluble compounds]]<br />
[[Category:Irritants]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Barium_peroxide&diff=14388Barium peroxide2022-02-13T17:02:28Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Barium peroxide<br />
| Reference =<br />
| IUPACName = Barium peroxide<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Barium binoxide<br>Barium dioxide<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Grey-white crystalline (anhydrous)<br>Colorless solid (octahydrate) <br />
| BoilingPt = <br />
| BoilingPtC = 800<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 5.68 g/cm<sup>3</sup> (anhydrous)<br>2.292 g/cm<sup>3</sup> (octahydrate)<br />
| Formula = BaO<sub>2</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 169.33 g/mol (anhydrous)<br>313.45 (octahydrate)<br />
| MeltingPt = <br />
| MeltingPtC = 450<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = ''anhydrous''<br> 0.091 g/100 ml (20 °C)<hr>''octahydrate''<br>0.168 g/100 ml<br />
| SolubleOther = Reacts with acids<br>Insoluble in hydrocarbons<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = 6<br />
| CrystalStruct = Tetragonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [https://www.docdroid.net/JJosUxN/barium-peroxide-sa.pdf Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Barium oxide]]<br />
}}<br />
}}<br />
'''Barium peroxide''' is the inorganic compound with the formula '''BaO<sub>2</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
Barium peroxide can be used to produce highly concentrated [[hydrogen peroxide]] via its reaction with conc. [[sulfuric acid]].<br />
<br />
: BaO<sub>2</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>O<sub>2</sub> + BaSO<sub>4</sub><br />
<br />
The insoluble [[barium sulfate]] is filtered from the mixture.<br />
<br />
===Physical===<br />
Barium peroxide is a white solid, insoluble in water.<br />
<br />
==Availability==<br />
Barium peroxide is sold by lab suppliers. So far there aren't any sellers on eBay and Amazon.<br />
<br />
==Preparation==<br />
Barium peroxide can be made by the reversible reaction of O<sub>2</sub> with [[barium oxide]]. The peroxide forms around 500 °C and oxygen is released above 820 °C.<br />
<br />
: 2 BaO + O<sub>2</sub> ⇌ 2 BaO<sub>2</sub><br />
<br />
A different, aqueous preparation is performed through a metathesis reaction with barium nitrate and hydrogen peroxide:<br />
<br />
: Ba(NO<sub>3</sub>)<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> ⇌ BaO<sub>2</sub> + 2HNO<sub>3</sub><br />
<br />
Cool the mixed solution to near freezing, and the crystallohydrate BaO<sub>2</sub>*8H<sub>2</sub>O will precipitate. Calcine the crystallohydrate carefully at the temperature of 100-120 degrees Celsius to convert it to the anhydrous salt.<br />
<br />
==Projects==<br />
*Flash powders and fireworks<br />
*Make concentrated [[hydrogen peroxide]]<br />
<br />
==Handling==<br />
===Safety===<br />
Barium peroxide is corrosive and an oxidizer. Proper protection must be worn when handling this compound.<br />
<br />
===Storage===<br />
Barium peroxide should be kept in closed plastic or glass bottles.<br />
<br />
===Disposal===<br />
Can be neutralized by adding diluted sulfuric acid then iron oxide, to decompose the hydrogen peroxide.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=888 Barium Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=42263 Barium Peroxide -> Hydrogen Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16714 KMnO4 and BaO2 reaction]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13965 Synthesis of BaO2 from Ba(NO3)2]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Barium compounds]]<br />
[[Category:Oxides]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Peroxides]]<br />
[[Category:Inorganic peroxides]]<br />
[[Category:Insoluble compounds]]<br />
[[Category:Irritants]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Barium_peroxide&diff=14387Barium peroxide2022-02-13T17:01:43Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Barium peroxide<br />
| Reference =<br />
| IUPACName = Barium peroxide<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Barium binoxide<br>Barium dioxide<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Grey-white crystalline (anhydrous)<br>Colorless solid (octahydrate) <br />
| BoilingPt = <br />
| BoilingPtC = 800<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 5.68 g/cm<sup>3</sup> (anhydrous)<br>2.292 g/cm<sup>3</sup> (octahydrate)<br />
| Formula = BaO<sub>2</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 169.33 g/mol (anhydrous)<br>313.45 (octahydrate)<br />
| MeltingPt = <br />
| MeltingPtC = 450<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = ''anhydrous''<br> 0.091 g/100 ml (20 °C)<hr>''octahydrate''<br>0.168 g/100 ml<br />
| SolubleOther = Reacts with acids<br>Insoluble in hydrocarbons<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = 6<br />
| CrystalStruct = Tetragonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [https://www.docdroid.net/JJosUxN/barium-peroxide-sa.pdf Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Barium oxide]]<br />
}}<br />
}}<br />
'''Barium peroxide''' is the inorganic compound with the formula '''BaO<sub>2</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
Barium peroxide can be used to produce highly concentrated [[hydrogen peroxide]] via its reaction with conc. [[sulfuric acid]].<br />
<br />
: BaO<sub>2</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>O<sub>2</sub> + BaSO<sub>4</sub><br />
<br />
The insoluble [[barium sulfate]] is filtered from the mixture.<br />
<br />
===Physical===<br />
Barium peroxide is a white solid, insoluble in water.<br />
<br />
==Availability==<br />
Barium peroxide is sold by lab suppliers. So far there aren't any sellers on eBay and Amazon.<br />
<br />
==Preparation==<br />
Barium peroxide can be made by the reversible reaction of O<sub>2</sub> with [[barium oxide]]. The peroxide forms around 500 °C and oxygen is released above 820 °C.<br />
<br />
: 2 BaO + O<sub>2</sub> ⇌ 2 BaO<sub>2</sub><br />
<br />
A different, aqueous preparation is performed through a metathesis reaction with barium nitrate and hydrogen peroxide:<br />
<br />
: Ba(NO<sub>3</sub>)<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> ⇌ BaO<sub>2</sub> + 2HNO<sub>3</sub><br />
<br />
Cool the mixed solution to near freezing, and the crystallohydrate of BaO<sub>2</sub> will precipitate. Calcine the crystallohydrate carefully at the temperature of 100-120 degrees Celsius to convert it to the anhydrous salt.<br />
<br />
==Projects==<br />
*Flash powders and fireworks<br />
*Make concentrated [[hydrogen peroxide]]<br />
<br />
==Handling==<br />
===Safety===<br />
Barium peroxide is corrosive and an oxidizer. Proper protection must be worn when handling this compound.<br />
<br />
===Storage===<br />
Barium peroxide should be kept in closed plastic or glass bottles.<br />
<br />
===Disposal===<br />
Can be neutralized by adding diluted sulfuric acid then iron oxide, to decompose the hydrogen peroxide.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=888 Barium Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=42263 Barium Peroxide -> Hydrogen Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16714 KMnO4 and BaO2 reaction]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13965 Synthesis of BaO2 from Ba(NO3)2]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Barium compounds]]<br />
[[Category:Oxides]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Peroxides]]<br />
[[Category:Inorganic peroxides]]<br />
[[Category:Insoluble compounds]]<br />
[[Category:Irritants]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Barium_peroxide&diff=14386Barium peroxide2022-02-13T17:01:23Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Barium peroxide<br />
| Reference =<br />
| IUPACName = Barium peroxide<br />
| PIN =<br />
| SystematicName =<br />
| OtherNames = Barium binoxide<br>Barium dioxide<br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Grey-white crystalline (anhydrous)<br>Colorless solid (octahydrate) <br />
| BoilingPt = <br />
| BoilingPtC = 800<br />
| BoilingPt_ref = <br />
| BoilingPt_notes = (decomposes)<br />
| Density = 5.68 g/cm<sup>3</sup> (anhydrous)<br>2.292 g/cm<sup>3</sup> (octahydrate)<br />
| Formula = BaO<sub>2</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = 169.33 g/mol (anhydrous)<br>313.45 (octahydrate)<br />
| MeltingPt = <br />
| MeltingPtC = 450<br />
| MeltingPt_ref = <br />
| MeltingPt_notes = <br />
| Odor = Odorless<br />
| pKa = <br />
| pKb = <br />
| Solubility = ''anhydrous''<br> 0.091 g/100 ml (20 °C)<hr>''octahydrate''<br>0.168 g/100 ml<br />
| SolubleOther = Reacts with acids<br>Insoluble in hydrocarbons<br />
| Solvent = <br />
| VaporPressure = ~0 mmHg<br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = 6<br />
| CrystalStruct = Tetragonal<br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = Non-explosive<br />
| ExternalMSDS = [https://www.docdroid.net/JJosUxN/barium-peroxide-sa.pdf Sigma-Aldrich]<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Barium oxide]]<br />
}}<br />
}}<br />
'''Barium peroxide''' is the inorganic compound with the formula '''BaO<sub>2</sub>'''.<br />
<br />
==Properties==<br />
===Chemical===<br />
Barium peroxide can be used to produce highly concentrated [[hydrogen peroxide]] via its reaction with conc. [[sulfuric acid]].<br />
<br />
: BaO<sub>2</sub> + H<sub>2</sub>SO<sub>4</sub> → H<sub>2</sub>O<sub>2</sub> + BaSO<sub>4</sub><br />
<br />
The insoluble [[barium sulfate]] is filtered from the mixture.<br />
<br />
===Physical===<br />
Barium peroxide is a white solid, insoluble in water.<br />
<br />
==Availability==<br />
Barium peroxide is sold by lab suppliers. So far there aren't any sellers on eBay and Amazon.<br />
<br />
==Preparation==<br />
Barium peroxide can be made by the reversible reaction of O<sub>2</sub> with [[barium oxide]]. The peroxide forms around 500 °C and oxygen is released above 820 °C.<br />
<br />
: 2 BaO + O<sub>2</sub> ⇌ 2 BaO<sub>2</sub><br />
<br />
A different, aqueous preparation is performed through a metathesis reaction with barium nitrate and hydrogen peroxide:<br />
<br />
: BaNO(<sub>3</sub>)<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> ⇌ BaO<sub>2</sub> + 2HNO<sub>3</sub><br />
<br />
Cool the mixed solution to near freezing, and the crystallohydrate of BaO<sub>2</sub> will precipitate. Calcine the crystallohydrate carefully at the temperature of 100-120 degrees Celsius to convert it to the anhydrous salt.<br />
<br />
==Projects==<br />
*Flash powders and fireworks<br />
*Make concentrated [[hydrogen peroxide]]<br />
<br />
==Handling==<br />
===Safety===<br />
Barium peroxide is corrosive and an oxidizer. Proper protection must be worn when handling this compound.<br />
<br />
===Storage===<br />
Barium peroxide should be kept in closed plastic or glass bottles.<br />
<br />
===Disposal===<br />
Can be neutralized by adding diluted sulfuric acid then iron oxide, to decompose the hydrogen peroxide.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=888 Barium Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=42263 Barium Peroxide -> Hydrogen Peroxide]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=16714 KMnO4 and BaO2 reaction]<br />
*[http://www.sciencemadness.org/talk/viewthread.php?tid=13965 Synthesis of BaO2 from Ba(NO3)2]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Barium compounds]]<br />
[[Category:Oxides]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Peroxides]]<br />
[[Category:Inorganic peroxides]]<br />
[[Category:Insoluble compounds]]<br />
[[Category:Irritants]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sodium_ferrate&diff=14001Sodium ferrate2021-06-02T04:57:45Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Sodium ferrate<br />
| Reference = <br />
| IUPACName = Sodium ferrate<br />
| PIN = Sodium ferrate<br />
| SystematicName = Sodium ferrate (VI)<br />
| OtherNames = <br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Red-purple (solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = <br />
| Formula = Na<sub>2</sub>FeO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = <br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| pKa = <br />
| pKb = <br />
| Solubility = Very soluble<br />
| SolubleOther = Reacts with various organic solvents<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = None<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium ferrate]]<br />
}}<br />
}}<br />
'''Sodium ferrate''' is a compound with a formula of '''Na<sub>2</sub>FeO<sub>4</sub>'''. It is a very elusive sodium salt of ferric acid. Ferric acid is extremely unstable and does not exist under normal conditions in any way, shape, form or concentration. Its salts also tend to be unstable, sodium ferrate in particular.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
<br />
Those chosen few who have seen sodium ferrate as a dry solid, describe it similarly to [[potassium ferrate]]: a dark crystalline solid that dissolves in water to form red-purple solutions. Sodium ferrate is extremely soluble, and this is the reason why it is so hard to isolate it as a pure solid: it cannot be displaced from the solution by excess NaOH. Other methods of crystallization, such as boiling the solution down, are too harsh for the unstable ferrate ion, and tend to decompose it completely.<br />
<br />
=== Chemical ===<br />
<br />
Sodium ferrate is a very strong oxidizer, stronger and more reactive than [[potassium ferrate]]. Generally, however, their properties are similar.<br />
<br />
== Availability ==<br />
<br />
This is an exceptionally rare chemical, it is usually not stocked by any major suppliers, including Sigma-Aldrich.<br />
<br />
== Preparation ==<br />
<br />
It is relatively easy to prepare an aqueous solution of sodium ferrate: the same two methods that are used for synthesizing potassium ferrate, namely the electrolytic method and the hypochlorite method, will work here. However, the resulting red-purple solution is more or less a dead end: there's no way to turn it into the pure solid in an amateur setting.<br />
<br />
Another, non-aqueous way to prepare this compound involves a precursor chemical, sodium ferrate (IV), or sodium hypoferrate Na<sub>4</sub>FeO<sub>4</sub>. This crystalline solid is stable when anhydrous, but extremely unstable towards disproportionation in solution. It disproportionates into sodium ferrate, sodium hydroxide and iron (III) hydroxide instantly on contact with water.<br />
<br />
Sodium hypoferrate is synthesized in crucibles with the following reaction:<br />
<br />
:8 Na<sub>2</sub>O + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub><br />
<br />
This requires blowing hot oxygen or air through the crucible. If you cannot do that, you can use sodium peroxide for that:<br />
<br />
: 8 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub> + 3O<sub>2</sub><br />
<br />
Excess oxygen will cause a side reaction that directly leads to formation of solid sodium ferrate:<br />
<br />
: 4 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>2</sub>FeO<sub>4</sub><br />
<br />
The resulting solid is a mixture of sodium hypoferrate and sodium ferrate. If you conduct this synthesis with both sodium peroxide and blowing hot oxygen, it is possible for the second reaction to dominate, resulting in a working method that can be used to synthesize solid sodium ferrate of reasonable purity.<br />
<br />
==Projects==<br />
*Oxidize organic compounds<br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Sodium ferrate is non-toxic. The products of its decomposition aren't toxic either. However, dry sodium ferrate should not come in contact with flammable organic compounds. <br />
<br />
=== Storage ===<br />
<br />
Dry sodium ferrate should be stored in a dark place, without access to air (it reacts with carbon dioxide in the air). Ideally, it should be kept under vacuum or inert gas.<br />
<br />
Sodium ferrate solutions are very perishable and cannot be stored for any longer period of time.<br />
<br />
=== Disposal ===<br />
<br />
Sodium ferrate solutions can just be poured into the ground or drain. Contact with any organics causes the ferrate to be quickly reduced and decomposed.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[https://www.sciencemadness.org/whisper/viewthread.php?tid=62978 Overlooked Ferrates]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Sodium compounds]]<br />
[[Category:Ferrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Unstable materials]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sodium_ferrate&diff=14000Sodium ferrate2021-06-02T04:57:30Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Sodium ferrate<br />
| Reference = <br />
| IUPACName = Sodium ferrate<br />
| PIN = Sodium ferrate<br />
| SystematicName = Sodium ferrate (VI)<br />
| OtherNames = <br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Red-purple (solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = <br />
| Formula = Na<sub>2</sub>FeO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = <br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| pKa = <br />
| pKb = <br />
| Solubility = Very soluble<br />
| SolubleOther = Reacts with various organic solvents<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = None<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium ferrate]]<br />
}}<br />
}}<br />
'''Sodium ferrate''' is a compound with a formula of '''Na<sub>2</sub>FeO<sub>4</sub>'''. It is a very elusive sodium salt of ferric acid. Ferric acid is extremely unstable and does not exist under normal conditions in any way, shape, form or concentration. Its salts also tend to be unstable, sodium ferrate in particular.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
<br />
Those chosen few who have seen sodium ferrate as a dry solid, describe it similarly to [[potassium ferrate]]: a dark crystalline solid that dissolves in water to form red-purple solutions. Sodium ferrate is extremely soluble, and this is the reason why it is so hard to isolate it as a pure solid: it cannot be displaced from the solution by excess NaOH. Other methods of crystallization, such as boiling the solution down, are too harsh for the unstable ferrate ion, and tend to decompose it completely.<br />
<br />
=== Chemical ===<br />
<br />
Sodium ferrate is a very strong oxidizer, stronger and more reactive than [[potassium ferrate]]. Generally, however, their properties are similar.<br />
<br />
== Availability ==<br />
<br />
This is an exceptionally rare chemical, it is usually not stocked by any major suppliers, including Sigma-Aldrich.<br />
<br />
== Preparation ==<br />
<br />
It is relatively easy to prepare an aqueous solution of sodium ferrate: the same two methods that are used for synthesizing potassium ferrate, namely the electrolytic method and the hypochlorite method, will work here. However, the resulting red-purple solution is more or less a dead end: there's no way to turn it into the pure solid in an amateur setting.<br />
<br />
Another, non-aqueous way to prepare this compound involves a precursor chemical, sodium ferrate (IV), or sodium hypoferrate Na<sub>4</sub>FeO<sub>4</sub>. This crystalline solid is stable when anhydrous, but extremely unstable towards disproportionation in solution. It disproportionates into sodium ferrate, sodium hydroxide and iron (III) hydroxide instantly on contact with water.<br />
<br />
Sodium hypoferrate is synthesized in crucibles with the following reaction:<br />
<br />
:8 Na<sub>2</sub>O + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub><br />
<br />
This requires blowing hot oxygen or air through the crucible. If you cannot do that, you can use sodium peroxide for that:<br />
<br />
: 8 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub> + 3O<sub>2</sub><br />
<br />
Excess oxygen will cause a side reaction that directly leads to formation of solid sodium ferrate:<br />
<br />
: 4 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>2</sub>FeO<sub>4</sub><br />
<br />
The resulting solid is a mixture of sodium hypoferrate and sodium ferrate. If you conduct this reaction with both sodium peroxide and blowing hot oxygen, it is possible for the second reaction to dominate, resulting in a working method that can be used to synthesize solid sodium ferrate of reasonable purity.<br />
<br />
==Projects==<br />
*Oxidize organic compounds<br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Sodium ferrate is non-toxic. The products of its decomposition aren't toxic either. However, dry sodium ferrate should not come in contact with flammable organic compounds. <br />
<br />
=== Storage ===<br />
<br />
Dry sodium ferrate should be stored in a dark place, without access to air (it reacts with carbon dioxide in the air). Ideally, it should be kept under vacuum or inert gas.<br />
<br />
Sodium ferrate solutions are very perishable and cannot be stored for any longer period of time.<br />
<br />
=== Disposal ===<br />
<br />
Sodium ferrate solutions can just be poured into the ground or drain. Contact with any organics causes the ferrate to be quickly reduced and decomposed.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[https://www.sciencemadness.org/whisper/viewthread.php?tid=62978 Overlooked Ferrates]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Sodium compounds]]<br />
[[Category:Ferrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Unstable materials]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sodium_ferrate&diff=13999Sodium ferrate2021-06-02T04:56:22Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Sodium ferrate<br />
| Reference = <br />
| IUPACName = Sodium ferrate<br />
| PIN = Sodium ferrate<br />
| SystematicName = Sodium ferrate (VI)<br />
| OtherNames = <br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Red-purple (solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = <br />
| Formula = Na<sub>2</sub>FeO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = <br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| pKa = <br />
| pKb = <br />
| Solubility = Very soluble<br />
| SolubleOther = Reacts with various organic solvents<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = None<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium ferrate]]<br />
}}<br />
}}<br />
'''Sodium ferrate''' is a compound with a formula of '''Na<sub>2</sub>FeO<sub>4</sub>'''. It is a very elusive sodium salt of ferric acid. Ferric acid is extremely unstable and does not exist under normal conditions in any way, shape, form or concentration. Its salts also tend to be unstable, sodium ferrate in particular.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
<br />
Those chosen few who have seen sodium ferrate as a dry solid, describe it similarly to [[potassium ferrate]]: a dark crystalline solid that dissolves in water to form red-purple solutions. Sodium ferrate is extremely soluble, and this is the reason why it is so hard to isolate it as a pure solid: it cannot be displaced from the solution by excess NaOH. Other methods of crystallization, such as boiling the solution down, are too harsh for the unstable ferrate ion, and tend to decompose it completely.<br />
<br />
=== Chemical ===<br />
<br />
Sodium ferrate is a very strong oxidizer, stronger and more reactive than [[potassium ferrate]]. Generally, however, their properties are similar.<br />
<br />
== Availability ==<br />
<br />
This is an exceptionally rare chemical, it is usually not stocked by any major suppliers, including Sigma-Aldrich.<br />
<br />
== Preparation ==<br />
<br />
It is relatively easy to prepare an aqueous solution of sodium ferrate: the same two methods that are used for synthesizing potassium ferrate, namely the electrolytic method and the hypochlorite method, will work here. However, the resulting red-purple solution is more or less a dead end: there's no way to turn it into the pure solid in an amateur setting.<br />
<br />
Another, non-aqueous way to prepare this compound involves a precursor chemical, sodium ferrate (IV), or sodium hypoferrate Na<sub>4</sub>FeO<sub>4</sub>. This crystalline solid is stable when anhydrous, but extremely unstable towards disproportionation in solution. It disproportionates into sodium ferrate, sodium hydroxide and iron (III) hydroxide instantly on contact with water.<br />
<br />
Sodium hypoferrate is synthesized in crucibles with the following reaction:<br />
<br />
:8 Na<sub>2</sub>O + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub><br />
<br />
This requires blowing hot oxygen or air through the crucible. If you cannot do that, you can use sodium peroxide for that:<br />
<br />
: 8 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub> + 3O<sub>2</sub><br />
<br />
Excess oxygen will cause a side reaction that directly leads to formation of solid sodium ferrate:<br />
<br />
: 4 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>2</sub>FeO<sub>4</sub><br />
<br />
The resulting solid is a mixture of sodium hypoferrate and sodium ferrate. If you conduct this reaction with both sodium peroxide and blowing hot oxygen, it is possible to synthesize solid sodium ferrate of reasonable purity.<br />
<br />
==Projects==<br />
*Oxidize organic compounds<br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Sodium ferrate is non-toxic. The products of its decomposition aren't toxic either. However, dry sodium ferrate should not come in contact with flammable organic compounds. <br />
<br />
=== Storage ===<br />
<br />
Dry sodium ferrate should be stored in a dark place, without access to air (it reacts with carbon dioxide in the air). Ideally, it should be kept under vacuum or inert gas.<br />
<br />
Sodium ferrate solutions are very perishable and cannot be stored for any longer period of time.<br />
<br />
=== Disposal ===<br />
<br />
Sodium ferrate solutions can just be poured into the ground or drain. Contact with any organics causes the ferrate to be quickly reduced and decomposed.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[https://www.sciencemadness.org/whisper/viewthread.php?tid=62978 Overlooked Ferrates]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Sodium compounds]]<br />
[[Category:Ferrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Unstable materials]]</div>Ave369https://www.sciencemadness.org/smwiki/index.php?title=Sodium_ferrate&diff=13998Sodium ferrate2021-06-02T04:53:11Z<p>Ave369: /* Preparation */</p>
<hr />
<div>{{Chembox<br />
| Name = Sodium ferrate<br />
| Reference = <br />
| IUPACName = Sodium ferrate<br />
| PIN = Sodium ferrate<br />
| SystematicName = Sodium ferrate (VI)<br />
| OtherNames = <br />
<!-- Images --><br />
| ImageFile = <br />
| ImageSize = <br />
| ImageAlt = <br />
| ImageName = <br />
| ImageFile1 = <br />
| ImageSize1 = <br />
| ImageAlt1 = <br />
| ImageName1 = <br />
| ImageFile2 = <br />
| ImageSize2 = <br />
| ImageAlt2 = <br />
| ImageName2 = <br />
| ImageFile3 = <br />
| ImageSize3 = <br />
| ImageAlt3 = <br />
| ImageName3 = <br />
| ImageFileL1 = <br />
| ImageSizeL1 = <br />
| ImageAltL1 = <br />
| ImageNameL1 = <br />
| ImageFileR1 = <br />
| ImageSizeR1 = <br />
| ImageAltR1 = <br />
| ImageNameR1 = <br />
| ImageFileL2 = <br />
| ImageSizeL2 = <br />
| ImageAltL2 = <br />
| ImageNameL2 = <br />
| ImageFileR2 = <br />
| ImageSizeR2 = <br />
| ImageAltR2 = <br />
| ImageNameR2 = <br />
<!-- Sections --><br />
| Section1 = {{Chembox Identifiers<br />
| 3DMet = <br />
| Abbreviations = <br />
| SMILES = <br />
}}<br />
| Section2 = {{Chembox Properties<br />
| AtmosphericOHRateConstant = <br />
| Appearance = Red-purple (solution)<br />
| BoilingPt = <br />
| BoilingPtC = <br />
| BoilingPt_ref = <br />
| BoilingPt_notes = Decomposes<br />
| Density = <br />
| Formula = Na<sub>2</sub>FeO<sub>4</sub><br />
| HenryConstant = <br />
| LogP = <br />
| MolarMass = <br />
| MeltingPt = <br />
| MeltingPtC = <br />
| MeltingPt_ref = <br />
| MeltingPt_notes = Decomposes<br />
| pKa = <br />
| pKb = <br />
| Solubility = Very soluble<br />
| SolubleOther = Reacts with various organic solvents<br />
| Solvent = <br />
| VaporPressure = <br />
}}<br />
| Section3 = {{Chembox Structure<br />
| Coordination = <br />
| CrystalStruct = <br />
| MolShape = <br />
}}<br />
| Section4 = {{Chembox Thermochemistry<br />
| DeltaGf = <br />
| DeltaHc = <br />
| DeltaHf = <br />
| Entropy = <br />
| HeatCapacity = <br />
}}<br />
| Section5 = {{Chembox Explosive<br />
| ShockSens = <br />
| FrictionSens = <br />
| DetonationV = <br />
| REFactor = <br />
}}<br />
| Section6 = {{Chembox Hazards<br />
| AutoignitionPt = Non-flammable<br />
| ExploLimits = <br />
| ExternalMSDS = None<br />
| FlashPt = Non-flammable<br />
| LD50 = <br />
| LC50 = <br />
| MainHazards = Oxidizer<br />
| NFPA-F = <br />
| NFPA-H = <br />
| NFPA-R = <br />
| NFPA-S = <br />
}}<br />
| Section7 = {{Chembox Related<br />
| OtherAnions = <br />
| OtherCations = <br />
| OtherFunction = <br />
| OtherFunction_label = <br />
| OtherCompounds = [[Potassium ferrate]]<br />
}}<br />
}}<br />
'''Sodium ferrate''' is a compound with a formula of '''Na<sub>2</sub>FeO<sub>4</sub>'''. It is a very elusive sodium salt of ferric acid. Ferric acid is extremely unstable and does not exist under normal conditions in any way, shape, form or concentration. Its salts also tend to be unstable, sodium ferrate in particular.<br />
<br />
== Properties ==<br />
=== Physical ===<br />
<br />
Those chosen few who have seen sodium ferrate as a dry solid, describe it similarly to [[potassium ferrate]]: a dark crystalline solid that dissolves in water to form red-purple solutions. Sodium ferrate is extremely soluble, and this is the reason why it is so hard to isolate it as a pure solid: it cannot be displaced from the solution by excess NaOH. Other methods of crystallization, such as boiling the solution down, are too harsh for the unstable ferrate ion, and tend to decompose it completely.<br />
<br />
=== Chemical ===<br />
<br />
Sodium ferrate is a very strong oxidizer, stronger and more reactive than [[potassium ferrate]]. Generally, however, their properties are similar.<br />
<br />
== Availability ==<br />
<br />
This is an exceptionally rare chemical, it is usually not stocked by any major suppliers, including Sigma-Aldrich.<br />
<br />
== Preparation ==<br />
<br />
It is relatively easy to prepare an aqueous solution of sodium ferrate: the same two methods that are used for synthesizing potassium ferrate, namely the electrolytic method and the hypochlorite method, will work here. However, the resulting red-purple solution is more or less a dead end: there's no way to turn it into the pure solid in an amateur setting.<br />
<br />
Another, non-aqueous way to prepare this compound involves a precursor chemical, sodium ferrate (IV), or sodium hypoferrate Na<sub>4</sub>FeO<sub>4</sub>. This crystalline solid is stable when anhydrous, but extremely unstable towards disproportionation in solution. It disproportionates into sodium ferrate, sodium hydroxide and iron (III) hydroxide instantly on contact with water.<br />
<br />
Sodium hypoferrate is synthesized in crucibles with the following reaction:<br />
<br />
:8 Na<sub>2</sub>O + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub><br />
<br />
This requires blowing hot oxygen or air through the crucible. If you cannot do that, you can use sodium peroxide for that:<br />
<br />
: 8 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> → 4 Na<sub>4</sub>FeO<sub>4</sub> + 3O<sub>2</sub><br />
<br />
Excess oxygen will cause a side reaction that directly leads to formation of solid sodium ferrate:<br />
<br />
: 4 Na<sub>2</sub>O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> + O<sub>2</sub> → 4 Na<sub>2</sub>FeO<sub>4</sub><br />
<br />
The resulting solid is a mixture of sodium hypoferrate and sodium ferrate.<br />
<br />
==Projects==<br />
*Oxidize organic compounds<br />
<br />
== Handling ==<br />
=== Safety ===<br />
<br />
Sodium ferrate is non-toxic. The products of its decomposition aren't toxic either. However, dry sodium ferrate should not come in contact with flammable organic compounds. <br />
<br />
=== Storage ===<br />
<br />
Dry sodium ferrate should be stored in a dark place, without access to air (it reacts with carbon dioxide in the air). Ideally, it should be kept under vacuum or inert gas.<br />
<br />
Sodium ferrate solutions are very perishable and cannot be stored for any longer period of time.<br />
<br />
=== Disposal ===<br />
<br />
Sodium ferrate solutions can just be poured into the ground or drain. Contact with any organics causes the ferrate to be quickly reduced and decomposed.<br />
<br />
==References==<br />
<references/><br />
===Relevant Sciencemadness threads===<br />
*[https://www.sciencemadness.org/whisper/viewthread.php?tid=62978 Overlooked Ferrates]<br />
<br />
[[Category:Chemical compounds]]<br />
[[Category:Inorganic compounds]]<br />
[[Category:Sodium compounds]]<br />
[[Category:Ferrates]]<br />
[[Category:Oxidizing agents]]<br />
[[Category:Unstable materials]]</div>Ave369