Difference between revisions of "Sulfuric acid"

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−	{{distinguish|Sulfurous acid}}	+	[[File:Smw1.png|thumb]]
−	{{chembox	+	'''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.
−	| Watchedfields = changed	+
−	| verifiedrevid = 477003658	+
−	| ImageFile2 = Sulfuric-acid-2D-dimensions.svg	+
−	| ImageAlt2 = S=O bond length = 142.2 pm, <br>S-O bond length = 157.4 pm, <br>O-H bond length = 97 pm	+
−	| ImageSize2 = 150	+
−	| ImageFileL1 = Sulfuric-acid-Givan-et-al-1999-3D-vdW.png	+
−	| ImageCaptionL1 = Space-filling model	+
−	| ImageFileR1 = Sulfuric-acid-Givan-et-al-1999-3D-balls.png	+
−	| ImageCaptionR1 = Ball-and-stick model	+
−	| ImageFile3 = Sulphuric acid 96 percent extra pure.jpg	+
−	| ImageSize3 = 140px	+
−	| IUPACName = Sulfuric acid	+
−	| OtherNames = Oil of vitriol	+
−	| Section1 = {{Chembox Identifiers	+
−	| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}	+
−	| ChemSpiderID = 1086	+
−	| UNII_Ref = {{fdacite|correct|FDA}}	+
−	| UNII = O40UQP6WCF	+
−	| KEGG_Ref = {{keggcite|correct|kegg}}	+
−	| KEGG = D05963	+
−	| InChI = 1/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)	+
−	| InChIKey = QAOWNCQODCNURD-UHFFFAOYAC	+
−	| ChEBI_Ref = {{ebicite|correct|EBI}}	+
−	| ChEBI = 26836	+
−	| SMILES = OS(=O)(=O)O	+
−	| ChEMBL_Ref = {{ebicite|correct|EBI}}	+
−	| ChEMBL = 572964	+
−	| StdInChI_Ref = {{stdinchicite|correct|chemspider}}	+
−	| StdInChI = 1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)	+
−	| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}	+
−	| StdInChIKey = QAOWNCQODCNURD-UHFFFAOYSA-N	+
−	| CASNo = 7664-93-9	+
−	| CASNo_Ref = {{cascite|correct|CAS}}	+
−	| RTECS = WS5600000	+
−	| EINECS = 231-639-5	+
−	| UNNumber = 1830	+
−	}}	+
−	| Section2 = {{Chembox Properties	+
−	| Formula = {{chem|H|2|S|O|4}}	+
−	| MolarMass = 98.079 g/mol	+
−	| Appearance = Clear, colorless, odorless liquid	+
−	| Density = 1.84 g/cm<sup>3</sup>, liquid	+
−	| Solubility = miscible	+
−	| MeltingPtC = 10	+
−	| BoilingPtC = 337	+
−	| BoilingPt_notes = When sulfuric acid is above {{convert|300|C|F}}, it will decompose slowly	+
−	| Viscosity = 26.7 [[Poise|cP]] (20&nbsp;°C)	+
−	| pKa = −3, 1.99	+
−	| VaporPressure = 0.001 mmHg (20°C)<ref name=PGCH/>	+
−	}}	+
−	| Section4 = {{Chembox Thermochemistry	+
−	| DeltaHf = −814&nbsp;kJ·mol<sup>−1</sup><ref name=b1>{{cite book| author = Zumdahl, Steven S.|title =Chemical Principles 6th Ed.| publisher = Houghton Mifflin Company| year = 2009| isbn = 0-618-94690-X|page=A23}}</ref>	+
−	| Entropy = 157&nbsp;J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1/>	+
−	}}	+
−	| Section7 = {{Chembox Hazards	+
−	| ExternalMSDS = [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc03/icsc0362.htm ICSC 0362]	+
−	| EUIndex = 016-020-00-8	+
−	| FlashPt = Non-flammable	+
−	| EUClass = {{Hazchem C}}<ref>{{cite web|url=http://www.msdsauthoring.com/sulfuric_acid_nugentec_ghs_msds.pdf|title=NuGenTec Material Safety Datasheet-Sulfuric acid}}</ref><ref>{{cite web|url=http://www.ilo.org/dyn/icsc/showcard.display?p_lang=en&p_card_id=0362|title=Sulfuric acid IPCS|quote=The substance is harmful to aquatic organisms.(ENVIRONMENTAL DATA)}}</ref>	+
−	| NFPA-H = 3	+
−	| NFPA-F = 0	+
−	| NFPA-R = 2	+
−	| NFPA-S = W	+
−	| RPhrases = {{R35}}	+
−	| SPhrases = {{S1/2}} {{S26}} {{S30}} {{S45}}	+
−	| GHSPictograms = {{GHS corrosion}}	+
−	| GHSSignalWord = '''Danger'''	+
−	| HPhrases = {{H-phrases|314}}	+
−	| PPhrases = {{P-phrases|260|264|280|301+330+331|303+361+353|363|304+340|305+351+338|310|321|310|405|501}}	+
−	| TLV-TWA = 1 mg/m<sup>3</sup>	+
−	| TLV-STEL = 2 mg/m<sup>3</sup>	+
−	| TLV = 15 mg/m<sup>3</sup> (IDLH)	+
−	| PEL = TWA 1 mg/m<sup>3</sup><ref name=PGCH>{{PGCH|0577}}</ref>	+
−	| IDLH = 15 mg/m<sup>3</sup><ref name=PGCH/>	+
−	| REL = TWA 1 mg/m<sup>3</sup><ref name=PGCH/>	+
−	| LD50 = 2140 mg/kg (rat, oral)<ref name=IDLH>{{IDLH|7664939|Sulfuric acid}}</ref>	+
−	| 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)<ref name=IDLH/>	+
−	| LCLo = 87 mg/m<sup>3</sup> (guinea pig, 2.75 hr)<ref name=IDLH/>	+
−	}}	+
−	| Section8 = {{Chembox Related	+
−	| Function = [[strong acid]]s	+
−	| OtherFunctn = [[Selenic acid]]<br />[[Hydrochloric acid]]<br />[[Nitric acid]]<br />[[Chromic acid]]	+
−	| OtherCpds = [[Sulfurous acid]]<br />[[Peroxymonosulfuric acid]]<br />[[Sulfur trioxide]]<br />[[Oleum]]	+
−	}}	+
−	}}	+
−	'''Sulfuric acid''' ([[Sulfur#Spelling and etymology|alternative spelling]] '''sulphuric acid''') is a highly [[corrosive]] [[strong acid|strong]] [[mineral acid]] with the [[molecular formula]] [[Hydrogen|H<sub>2</sub>]][[sulfate|SO<sub>4</sub>]]. It is a pungent-ethereal, colorless to slightly yellow viscous liquid which is soluble in [[water]] at [[Miscibility|all concentrations]].<ref name="ds">{{cite web|url=http://www.arkema-inc.com/msds/01641.pdf|work=arkema-inc.com|title=Sulfuric acid safety data sheet|quote=Clear to turbid oily odorless liquid, colorless to slightly yellow.}}</ref> Sometimes, it is dyed dark brown during production to alert people to its hazards.<ref>{{cite web|url=http://chemicalland21.com/industrialchem/inorganic/SULFURIC%20ACID.htm|work=chemicalland21.com|title=Sulfuric acid|quote=Colorless (pure) to dark brown, oily, dense liquid with acrid odor.}}</ref> The historical name of this acid is '''oil of vitriol'''.<ref>{{cite book|title=sulfuric acid|year=2010|publisher=Encyclopædia Britannica|url=http://www.britannica.com/EBchecked/topic/572815/sulfuric-acid}}</ref>	+	==Properties==
		+	===Chemical properties===
		+	Sulfuric acid is a diprotic acid, it is able to give away two protons (H<sup>+</sup>). It first dissociates to form [[hydronium]] and [[hydrogen sulfate]], with a pKa of -3, indicative of a strong acid:
−	Sulfuric acid is a [[diprotic acid]] and shows different properties depending upon its concentration. Its corrosiveness on other materials, like [[metals]], [[tissue (biology)|living tissues]] or even [[stone]]s, can be mainly ascribed to its [[strong acid|strong acidic nature]] and, if concentrated, [[Dehydration reaction|strong dehydrating]] and [[Oxidizing agent|oxidizing]] properties. Sulfuric acid at a high [[concentration]] can cause very serious damage upon contact, since not only does it cause [[chemical burn]]s via [[hydrolysis]], but also [[burn#By depth|secondary thermal burns]] through [[Dehydration reaction|dehydration]].<ref name="OA"/><ref name=TB>{{cite web|url=http://www.basf.ca/group/corporate/ca/en_GB/function/conversions:/publishdownload/content/sustainability/employees/occupational-medicine/responsible-care-files/BASF_medGuidelines_E015_Sulfuric_acid_C.pdf|title=BASF Chemical Emergency Medical Guidelines - Sulfuric acid (H2SO4)|publisher=BASF Chemical Company|date=2012|accessdate=December 18, 2014}}</ref> It can lead to [[blindness|permanent blindness]] if splashed onto [[eyes]] and irreversible damage if swallowed.<ref name="OA"/> Accordingly, safety precautions should be strictly observed when handling it. Moreover, it is [[hygroscopic]], readily absorbing [[water vapour]] from the [[air]].<ref name="ds"/>	+	:H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → H<sub>3</sub>O + HSO<sub>4</sub><sup>−</sup>
−	Sulfuric acid has a wide range of applications including [[drain cleaner|domestic acidic drain cleaner]],<ref name="dc"/> [[electrolyte]] in [[lead–acid battery|lead-acid batteries]] and various [[cleaning agent]]s. It is also a central substance in the [[chemical industry]]. Principal uses include [[mineral processing]], [[fertilizer]] manufacturing, [[Oil refinery|oil refining]], [[wastewater processing]], and [[chemical synthesis]]. It is widely produced with different methods, such as [[contact process]], [[wet sulfuric acid process]] and some other methods.	+	The second dissociation forms sulfate and another hydronium ion from a hydrogen sulfate ion. It has a pKa of 1.99, indicative of a weak acid, and occurs like this:
−	==History==	+	: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>
−	[[File:Dalton's-sulphuric-acid.jpg|left|thumb|[[John Dalton]]'s 1808 sulfuric acid molecule shows a central [[sulfur]] atom bonded to three oxygen atoms, or [[sulfur trioxide]], the [[anhydride]] of sulfuric acid.]]	+
−	The study of [[vitriol]], a category of glassy minerals from which the acid can be derived, began in [[Classical antiquity|ancient times]]. [[Sumerians]] had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician [[Dioscorides]] (first century AD) and the Roman naturalist [[Pliny the Elder]] (23–79 AD). [[Galen]] also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of [[Zosimos of Panopolis]], in the treatise ''Phisica et Mystica'', and the [[Leyden papyrus X]].<ref name="Karpenko">Karpenko, Vladimir and Norris, John A. (2001). [http://www.chemicke-listy.cz/docs/full/2002_12_05.pdf Vitriol in the history of Chemistry], [[Charles University in Prague|Charles University]]</ref>	+
−	[[Alchemy and chemistry in medieval Islam|Persian alchemists]] [[Jābir ibn Hayyān]] (c. 721 – c. 815 AD), [[Muhammad ibn Zakariya al-Razi|Razi]] (865 – 925 AD), and [[Jamal Din al-Watwat]] (d. 1318, wrote the book ''Mabāhij al-fikar wa-manāhij al-'ibar''), included vitriol in their mineral classification lists. [[Ibn Sina]] focused on its medical uses and different varieties of vitriol.<ref name="Karpenko"/>	+	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. This property is useful for producing large amounts of sulfur dioxide for use as a reducing agent if water is continually removed.
−	Sulfuric acid was called "oil of vitriol" by medieval European alchemists because it was prepared by roasting "green vitriol" ([[iron (II) sulfate]]) in an iron [[retort]].  There are references to it in the works of [[Vincent of Beauvais]] and in the ''Compositum de Compositis'' ascribed to Saint [[Albertus Magnus]]. A passage from [[Pseudo-Geber]]´s ''Summa Perfectionis'' was long considered to be the first recipe for sulfuric acid, but this was a misinterpretation.<ref name="Karpenko"/>	+	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 nitrates|alkyl nitrates]].
−	In the seventeenth century, the German-Dutch chemist [[Johann Glauber]] prepared sulfuric acid by burning [[sulfur]] together with [[Potassium nitrate|saltpeter]] ([[potassium nitrate]], {{chem|KNO|3}}), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to {{chem|SO|3}}, which combines with water to produce sulfuric acid. In 1736, [[Joshua Ward]], a London pharmacist, used this method to begin the first large-scale production of sulfuric acid.
−	In 1746 in Birmingham, [[John Roebuck]] adapted this method to produce sulfuric acid in [[lead]]-lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the [[lead chamber process]] or "chamber process", remained the standard for sulfuric acid production for almost two centuries.<ref name=b1/>	+	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.
−	Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist [[Joseph Louis Gay-Lussac]] and British chemist John Glover improved concentration to 78%. However, the manufacture of some [[dye]]s and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by [[dry distillation|dry distilling]] minerals in a technique similar to the original [[alchemy|alchemical]] processes. [[Pyrite]] (iron disulfide, {{chem|FeS|2}}) was heated in air to yield iron(II) sulfate, {{chem|FeSO|4}}, which was oxidized by further heating in air to form [[iron(III) sulfate]], Fe<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>, which, when heated to 480&nbsp;°C, decomposed to [[iron(III) oxide]] and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid.<ref name=b1/>	+	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.
−	In 1831, British [[vinegar]] merchant Peregrine Phillips patented the [[contact process]], which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method.<ref name=b1>{{cite book|author=Philip J. Chenier|title=Survey of industrial chemistry|url=http://books.google.com/books?id=KlziQA-yx3gC&pg=PA28|accessdate=23 December 2011|date=1 April 2002|publisher=Springer|isbn=978-0-306-47246-6|pages=28–}}</ref>	+	===Physical properties===
		+	[[File:H2so4boil.jpg|thumb|342px|Boiling point of H2SO4 VS concentration]]
−	==Physical properties==	+	Sulfuric acid is a oily liquid at room temperature. It is colorless, but often has a very light yellow color when slightly contaminated with iron ions. 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.
−	===Grades of sulfuric acid===	+	Sulfuric acid's boiling point raises with the concentration as described in this figure to the right. An [[azeotrope]] forms at 98% w/w.
−	Although nearly 99% sulfuric acid can be made, the subsequent loss of {{chem|link=sulfur trioxide|SO|3}} at the boiling point brings the concentration to 98.3% acid. The 98% grade is more stable in storage, and is the usual form of what is described as "concentrated sulfuric acid." Other concentrations are used for different purposes. Some common concentrations are:<ref name="Columbia">{{Cite book|chapter = sulfuric acid|url = http://www.encyclopedia.com/topic/sulfuric_acid.aspx|title = The Columbia Encyclopedia|edition = 6th|year = 2009|accessdate = 2010-03-16}}</ref><ref name="EB11">{{Cite book|chapter = Sulphuric acid|title = [[Encyclopædia Britannica Eleventh Edition|Encyclopædia Britannica]]|edition = 11th|year = 1910–1911|volume = 26|pages = 65–69}}</ref>	+
−	{| class="wikitable"	+
−	|-	+
−	! Mass fraction<br />H<sub>2</sub>SO<sub>4</sub>	+
−	! Density<br />(kg/L)	+
−	! Concentration<br />(mol/L)	+
−	! Common name	+
−	|-	+
−	| 10% || 1.07 || align=center|~1 || dilute sulfuric acid	+
−	|-	+
−	| 29–32% || 1.25–1.28 || align=center|4.2–5 || battery acid<br />(used in [[Lead–acid battery|lead–acid batteries]])	+
−	|-	+
−	| 62–70% || 1.52–1.60 || align=center|9.6–11.5 || chamber acid<br />fertilizer acid	+
−	|-	+
−	| 78–80% || 1.70–1.73 || align=center|13.5–14 || tower acid<br />Glover acid	+
−	|-	+
−	| 98% || 1.83 || align=center|~18 || concentrated sulfuric acid	+
−	|}	+
−	"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the [[lead chamber process]], chamber acid being the acid produced in lead chamber itself (<70% to avoid contamination with [[nitrosylsulfuric acid]]) and tower acid being the acid recovered from the bottom of the Glover tower.<ref name="Columbia"/><ref name="EB11"/> They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid (the modern equivalent of chamber acid, used in many [[titration]]s) is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80&nbsp;°C (176&nbsp;°F) or higher.<ref name="EB11"/>	+	==Sources, production, and concentration==
		+	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 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 and will appear ye
−	Sulfuric acid reacts with its anhydride, {{chem|SO|3}}, to form {{chem|H|2|S|2|O|7}}, called ''[[pyrosulfuric acid]]'', ''fuming sulfuric acid'', ''Disulfuric acid'' or ''oleum'' or, less commonly, ''Nordhausen acid''. Concentrations of oleum are either expressed in terms of % {{chem|SO|3}} (called % oleum) or as % {{chem|H|2|SO|4}} (the amount made if {{chem|H|2|O}} were added); common concentrations are 40% oleum (109% {{chem|H|2|SO|4}}) and 65% oleum (114.6% {{chem|H|2|SO|4}}). Pure {{chem|H|2|S|2|O|7}} is a solid with melting point of 36&nbsp;°C.	+	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. Some brand of battery acids however are quite pure and easy to find almost anywhere. To concentrate it from 35 to 98%, one can boil it down to evaporate the water. It is advised to use glassware with anti-splash adapters, as sulfuric acid tends to bump a lot while it boils. The sulfuric acid will be quite concentrated when the temperature reaches around 300°C. The maximum reachable is 98.3% w/w, a point at which the boiling point will be 327°C. Extreme care must be taken when boiling sulfuric acid. The hazards and safety tips section must be read. Near the end, the acid will decompose into various [[sulfur oxides]].
−	Pure sulfuric acid has a vapor pressure of <0.001 torr at 25&nbsp;°C and 1 torr at 145.8&nbsp;°C,<ref name="OEHHA">{{Cite book|chapter = Sulfuric acid|url = http://oehha.ca.gov/air/chronic_rels/pdf/sulfuric.pdf|title = Determination of Noncancer Chronic Reference Exposure Levels Batch 2B December 2001|year = 2001|accessdate = 2012-10-01}}</ref> and 98% sulfuric acid has a <1 mmHg vapor pressure at 40&nbsp;°C.<ref name="Rhodia">{{Cite web|url = http://www.rhodia.com/our_company/businesses/documents/Sulfuric_Acid_98.pdf|title = Sulfuric Acid 98%|year = 2009|accessdate = 2014-07-02|work=rhodia.com}}</ref>	+	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 are added, it forms [[pyrosulfuric acid]], which is a solid at room temperature.
−	Pure sulfuric acid is a viscous clear liquid, like oil, and this explains the old name of the acid ('oil of vitriol').	+	==Projects==
		+	* Producing metal sulfates
		+	* Producing nitro compounds through [[nitration]]
		+	* The dehydration of [[sucrose]] to produce elemental [[carbon]]
		+	* E[[Esterification|sterifications]] that require a dehydrating agent, such as that of [[methyl salicylate]]
		+	* Making simple [[rayon]] fibers with [[schweizer's reagent|Schweizer's reagent]] and [[cellulose]]
		+	* Producing other concentrated acids by reaction of sulfuric acid with an anhydrous salt, such as in the production of fuming [[nitric acid]] and glacial [[acetic acid]]
−	Commercial sulfuric acid is sold in several different purity grades. Technical grade {{chem|H|2|SO|4}} is impure and often colored, but is suitable for making fertilizer. Pure grades, such as [[United States Pharmacopeia]] (USP) grade, are used for making [[pharmaceutical]]s and [[dye]]stuffs. [[Analytical chemistry|Analytical]] grades are also available.	+	==Handling==
		+	===Safety===
		+	[[File:Corrosive.png|thumb|left|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.
−	There are nine hydrates known, but three of them were confirmed to be tetrahydrate (H<sub>2</sub>SO<sub>4</sub>·4H<sub>2</sub>O), hemihexahydrate (H<sub>2</sub>SO<sub>4</sub>·{{frac|6|1|2}}H<sub>2</sub>O) and octahydrate (H<sub>2</sub>SO<sub>4</sub>·8H<sub>2</sub>O).	+	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 phenomena, 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.
−	===Polarity and conductivity===	+	Hot concentrated sulfuric acid may decompose to form sulfur dioxide and sulfur trioxide, which are toxic and corrosive, respectively.
−	{| class="wikitable sortable" align=right	+
−	|+colspan=2|Equilibrium of anhydrous sulfuric acid<ref name = greenwood/>	+
−	!Species	+
−	!mMol/kg	+
−	|-	+
−	|{{chem|HSO|4|-}}	+
−	|15.0	+
−	|-	+
−	|{{chem|H|3|SO|4|+}}	+
−	|11.3	+
−	|-	+
−	|{{chem|H|3|O|+}}	+
−	|8.0	+
−	|-	+
−	|{{chem|H|S|2|O|7|-}}	+
−	|4.4	+
−	|-	+
−	|{{chem|H|2|S|2|O|7}}	+
−	|3.6	+
−	|-	+
−	|{{chem|H|2|O}}	+
−	|0.1	+
−	|}	+
−	[[Anhydrous]] {{chem|H|2|SO|4}} is a very [[Chemical polarity|polar]] liquid, having a [[dielectric constant]] of around 100. It has a high [[electrical conductivity]], caused by dissociation through [[protonation|protonating]] itself, a process known as [[autoprotolysis]].<ref name = greenwood>{{Greenwood&Earnshaw}}</ref>	+
−	: 2 {{chem|H|2|SO|4}} {{eqm}} {{chem|H|3|SO|4|+}} + {{chem|HSO|4|-}}	+
−	The [[equilibrium constant]] for the autoprotolysis is<ref name = greenwood/>	+
−	:K<sub>ap</sub>(25&nbsp;°C)= [{{chem|H|3|SO|4|+}}] [{{chem|HSO|4|-}}] = {{val|2.7|e=-4}}	+
−	The comparable [[Self-ionization of water|equilibrium constant for water]], K<sub>w</sub> is 10<sup>−14</sup>, a factor of 10<sup>10</sup> (10 billion) smaller.	+	===Storage===
−		+	Sulfuric acid should be stored in closed bottles.
−	In spite of the viscosity of the acid, the effective [[Molar conductivity|conductivities]] of the {{chem|H|3|SO|4|+}} and {{chem|HSO|4|-}} ions are high due to an intra-molecular proton-switch mechanism (analogous to the [[Grotthuss mechanism]] in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.	+
−		+
−	==Chemical properties==	+
−		+
−	===Reaction with water and dehydrating property===	+
−	[[File:Sulphuric acid on a piece of towel.JPG|thumb|right|Drops of concentrated sulfuric acid rapidly dehydrate a piece of cotton towel.]]	+
−		+
−	Because the [[hydration reaction]] of sulfuric acid is highly [[exothermic reaction|exothermic]], dilution should always be performed by adding the acid to the [[Properties of water|water]] rather than the water to the acid.<ref>[http://www.cleapss.org.uk/attachments/article/0/SSS22.pdf?Secondary/Science/Student%20Safety%20Sheets/ Consortium of Local Education Authorities for the Provision of Science Equipment -STUDENT SAFETY SHEETS 22 Sulfuric(VI) acid]</ref> Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the ''acid'' is the limiting reagent. This reaction is best thought of as the formation of [[hydronium]] ions:	+
−		+
−	: {{chem|H|2|SO|4}} + {{chem|H|2|O}} → {{chem|H|3|O|+}} + HSO<sub>4</sub><sup>−</sup> &nbsp;&nbsp; K<sub>1</sub> = 2.4{{e|6}} &nbsp; (strong acid)	+
−		+
−	: {{chem|HSO|4|-}} + {{chem|H|2|O}} → {{chem|H|3|O|+}} + {{chem|SO|4|2-}} &nbsp;&nbsp;&nbsp; {{chem|K|2}} = 1.0{{e|-2}}&nbsp;<ref>{{cite web|url=http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/acidity.htm |title=Ionization Constants of Inorganic Acids |publisher=.chemistry.msu.edu |accessdate=2011-05-30}}</ref>	+
−		+
−	{{chem|HSO|4|-}} is the ''[[bisulfate]]'' anion and {{chem|SO|4|2-}} is the ''[[sulfate]]'' anion. K<sub>1</sub> and K<sub>2</sub> are the [[acid dissociation constant]]s.	+
−		+
−	Because the hydration of sulfuric acid is [[thermodynamic]]ally favorable and the affinity of it for [[water (molecule)|water]] is sufficiently strong, sulfuric acid is an excellent dehydrating agent. Concentrated sulfuric acid has a very powerful [[Dehydration reaction|dehydrating]] property, removing water ([[Water|H<sub>2</sub>O]]) from other [[Chemical compound|compounds]] including [[sugar]] and other [[carbohydrate]]s and producing [[carbon]], heat, steam, and a more dilute acid containing increased amounts of [[hydronium]] and [[bisulfate]] ions.	+
−		+
−	In [[laboratory]], this is often demonstrated by mixing [[table sugar]] (sucrose) into sulfuric acid. The sugar changes from white to dark brown and then to black as carbon is formed. A rigid column of black, porous carbon will emerge as well. The carbon will smell strongly of [[caramel (aroma)|caramel]] due to the heat generated.<ref>[http://www.youtube.com/watch?v=UcpodCsTxtc Sulphuric acid on sugar cubes chemistry experiment 8. Old Version]. YouTube. Retrieved on 2011-07-18.</ref>	+
−		+
−	:C<sub>12</sub>H<sub>22</sub>O<sub>11</sub> (white sucrose) + sulfuric acid → 12 C<sub>(black graphitic foam)</sub> + 11 H<sub>2</sub>O (steam) + sulfuric acid/water mixture	+
−		+
−	Similarly, mixing [[starch]] into concentrated sulfuric acid will give elemental [[carbon]] and water as absorbed by the sulfuric acid (which becomes slightly diluted). The effect of this can be seen when concentrated sulfuric acid is spilled on paper which is composed of [[cellulose]]; the cellulose reacts to give a [[combustion|burnt]] appearance, the [[carbon]] appears much as soot would in a fire.	+
−	Although less dramatic, the action of the acid on [[cotton]], even in diluted form, will destroy the fabric.	+
−		+
−	:({{chem|C|6|H|10|O|5}})''n'' + sulfuric acid → 6''n'' C + 5''n'' {{chem|H|2|O}}	+
−		+
−	The reaction with [[copper(II) sulfate]] can also demonstrate the dehydration property of sulfuric acid. The blue crystal is changed into white powder as water is removed.	+
−	:CuSO<sub>4</sub>·5H<sub>2</sub>O (blue crystal) + sulfuric acid → CuSO<sub>4</sub> (white powder) + 5 H<sub>2</sub>O	+
−		+
−	===Acid-base properties===	+
−	As an acid, sulfuric acid reacts with most [[base (chemistry)|bases]] to give the corresponding sulfate. For example, the blue [[copper]] salt [[copper(II) sulfate]], commonly used for [[electroplating]] and as a [[fungicide]], is prepared by the reaction of [[copper(II) oxide]] with sulfuric acid:	+
−		+
−	: CuO (s) + {{chem|H|2|SO|4}} (aq) → {{chem|CuSO|4}} (aq) + {{chem|H|2|O}} (l)	+
−		+
−	Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with [[sodium acetate]], for example, displaces [[acetic acid]], {{chem|CH|3|COOH}}, and forms [[sodium bisulfate]]:	+
−		+
−	:{{chem|H|2|SO|4}} + {{chem|CH|3|COONa}} → {{chem|NaHSO|4}} + {{chem|CH|3|COOH}}	+
−		+
−	Similarly, reacting sulfuric acid with [[potassium nitrate]] can be used to produce [[nitric acid]] and a precipitate of [[potassium bisulfate]]. When combined with [[nitric acid]], sulfuric acid acts both as an acid and a dehydrating agent, forming the [[nitronium ion]] {{chem|NO|2|+}}, which is important in [[nitration]] reactions involving [[electrophilic aromatic substitution]]. This type of reaction, where protonation occurs on an [[oxygen]] atom, is important in many [[organic chemistry]] reactions, such as [[Fischer esterification]] and dehydration of alcohols.	+
−		+
−	[[File:Structure of protonated sulfuric acid.png|right|thumb|200px|Solid state structure of the [D<sub>3</sub>SO<sub>4</sub>]<sup>+</sup> ion present in [D<sub>3</sub>SO<sub>4</sub>]<sup>+</sup>[SbF<sub>6</sub>]<sup>−</sup>, synthesized by using [[deuterium|D]]F in place of HF. (see text)]]	+
−		+
−	When allowed to react with [[superacid]]s, sulfuric acid can act as a base and be protonated, forming the [H<sub>3</sub>SO<sub>4</sub>]<sup>+</sup> ion. Salt of [H<sub>3</sub>SO<sub>4</sub>]<sup>+</sup> have been prepared using the following reaction in liquid [[hydrogen fluoride|HF]]:	+
−		+
−	: ((CH<sub>3</sub>)<sub>3</sub>SiO)<sub>2</sub>SO<sub>2</sub> + 3 HF + SbF<sub>5</sub> → [H<sub>3</sub>SO<sub>4</sub>]<sup>+</sup>[SbF<sub>6</sub>]<sup>−</sup> + 2 (CH<sub>3</sub>)<sub>3</sub>SiF	+
−		+
−	The above reaction is thermodynamically favored due to the high [[bond enthalpy]] of the Si–F bond in the side product. Protonation using simply [[fluoroantimonic acid|HF/SbF<sub>5</sub>]], however, have met with failure, as pure sulfuric acid undergoes [[molecular autoionization|self-ionization]] to give [H<sub>3</sub>O]<sup>+</sup> ions, which prevents the conversion of H<sub>2</sub>SO<sub>4</sub> to [H<sub>3</sub>SO<sub>4</sub>]<sup>+</sup> by the HF/SbF<sub>5</sub> system:<ref name="InorgChem">{{cite book	+
−	| title = Inorganic Chemistry, 3rd Edition	+
−	| chapter = Chapter 16: The group 16 elements	+
−	| author1 = Housecroft, Catherine E.	+
−	| author2 = Sharpe, Alan G.	+
−	| publisher = Pearson	+
−	| year = 2008	+
−	| isbn = 978-0-13-175553-6	+
−	| page = 523	+
−	}}</ref>	+
−		+
−	: 2 H<sub>2</sub>SO<sub>4</sub> {{eqm}} [H<sub>3</sub>O]<sup>+</sup> + [HS<sub>2</sub>O<sub>7</sub>]<sup>−</sup>	+
−		+
−	===Reactions with metals and strong oxidizing property===	+
−	Dilute sulfuric acid reacts with metals via a single displacement reaction as with other typical [[acid]]s, producing [[hydrogen]] gas and [[salt]]s (the metal sulfate). It attacks reactive metals (metals at positions above [[copper]] in the [[reactivity series]]) such as [[iron]], [[aluminium]], [[zinc]], [[manganese]], [[magnesium]] and [[nickel]].	+
−	: Fe (s) + {{chem|H|2|SO|4}} (aq) → {{chem|H|2}} (g) + {{chem|FeSO|4}} (aq)	+
−		+
−	However, concentrated sulfuric acid is a [[oxidizing agent|strong oxidizing agent]]<ref name="OA">{{cite web|url=http://www.dynamicscience.com.au/tester/solutions/chemistry/sulfuricacid1.html|title=Sulfuric acid – uses|work=dynamicscience.com.au}}</ref> and does not react with metals in the same way as other typical [[acid]]s. [[Sulfur dioxide]], [[water]] and SO<sub>4</sub><sup>2−</sup> ions are evolved instead of the [[hydrogen]] and [[salt]]s.	+
−	: 2 H<sub>2</sub>SO<sub>4</sub> + 2 e<sup>−</sup> → SO<sub>2</sub> + 2 H<sub>2</sub>O + SO<sub>4</sub><sup>2−</sup>	+
−		+
−	It can oxidize non-active metals such as [[tin]] and [[copper]], depending upon the temperature.	+
−	:Cu + 2 H<sub>2</sub>SO<sub>4</sub> → SO<sub>2</sub> + 2 H<sub>2</sub>O + SO<sub>4</sub><sup>2−</sup> + Cu<sup>2+</sup>	+
−		+
−	[[Lead]] and [[tungsten]], however, are resistant to sulfuric acid.	+
−		+
−	===Reactions with non-metals===	+
−	Hot concentrated sulfuric acid oxidizes non-metals such as [[carbon]]<ref>{{cite book|author1=Kinney, Corliss Robert |author2=Grey, V. E. |title=Reactions of a Bituminous Coal with Sulfuric Acid|year=1959|publisher=Pennsylvania State University|url=https://web.anl.gov/PCS/acsfuel/preprint%20archive/Files/03_2_BOSTON_04-59_0169.pdf}}</ref> (as bituminous coal) and [[sulfur]].	+
−	:C + 2 H<sub>2</sub>SO<sub>4</sub> → CO<sub>2</sub> + 2 SO<sub>2</sub> + 2 H<sub>2</sub>O	+
−	:S + 2 H<sub>2</sub>SO<sub>4</sub> → 3 SO<sub>2</sub> + 2 H<sub>2</sub>O	+
−		+
−	===Reaction with sodium chloride===	+
−	It reacts with [[sodium chloride]], and gives [[hydrogen chloride]] [[gas]] and [[sodium bisulfate]]:	+
−		+
−	:NaCl + H<sub>2</sub>SO<sub>4</sub> → NaHSO<sub>4</sub> + HCl	+
−		+
−	===Electrophilic aromatic substitution===	+
−	Benzene undergoes [[electrophilic aromatic substitution]] with sulfuric acid to give the corresponding [[sulfonic acid]]s:<ref>{{cite web|url =http://web.archive.org/web/20080706063639/http://www.chem.ucalgary.ca/courses/351/Carey/Ch12/ch12-4.html| title = Reactions of Arenes. Electrophilic Aromatic Substitution|author = Carey, F. A. |work = On-Line Learning Center for Organic Chemistry|publisher = [[University of Calgary]]|accessdate = 27 January 2008}}</ref>	+
−		+
−	:[[File:BenzeneSulfonation.png|250px]]	+
−		+
−	==Occurrence==	+
−	[[File:Rio tinto river CarolStoker NASA Ames Research Center.jpg|thumb|left|250px|[[Rio Tinto (river)|Rio Tinto]] with its highly acidic water]]	+
−		+
−	Pure sulfuric acid is not encountered naturally on Earth in anhydrous form, due to its great [[Hygroscopy|affinity for water]]. Dilute sulfuric acid is a constituent of [[acid rain]], which is formed by atmospheric [[Redox|oxidation]] of [[sulfur dioxide]] in the presence of [[water (molecule)|water]] – i.e., oxidation of [[sulfurous acid]]. Sulfur dioxide is the main byproduct produced when sulfur-containing fuels such as coal or oil are burned.	+
−		+
−	Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called [[acid mine drainage]] (AMD) or acid rock drainage (ARD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic streams. The oxidation of [[pyrite]] (iron sulfide) by molecular oxygen produces iron(II), or {{chem|Fe|2+}}:	+
−		+
−	:2 {{chem|FeS|2}} (s) + 7 {{chem|O|2}} + 2 {{chem|H|2|O}} → 2 {{chem|Fe|2+}} (aq) + 4 {{chem|SO|4|2-}} (aq) + 4 {{chem|H|+}}	+
−		+
−	The {{chem|Fe|2+}} can be further oxidized to {{chem|Fe|3+}}:	+
−		+
−	:4 {{chem|Fe|2+}} + {{chem|O|2}} + 4 {{chem|H|+}} → 4 {{chem|Fe|3+}} + 2 {{chem|H|2|O}}	+
−		+
−	The {{chem|Fe|3+}} produced can be precipitated as the [[hydroxide]] or [[hydrous iron oxides|hydrous oxide]]:	+
−		+
−	:{{chem|Fe|3+}} (aq) + 3 {{chem|H|2|O}} → {{chem|Fe(OH)|3}} (s) + 3 {{chem|H|+}}	+
−		+
−	The iron(III) ion ("ferric iron") can also oxidize pyrite:	+
−		+
−	:{{chem|FeS|2}} (s) + 14 {{chem|Fe|3+}} + 8 {{chem|H|2|O}} → 15 {{chem|Fe|2+}} (aq) + 2 {{chem|SO|4|2-}} (aq) + 16 {{chem|H|+}}	+
−		+
−	When iron(III) oxidation of pyrite occurs, the process can become rapid. [[pH]] values below zero have been measured in ARD produced by this process.	+
−		+
−	ARD can also produce sulfuric acid at a slower rate, so that the [[acid neutralizing capacity]] (ANC) of the aquifer can neutralize the produced acid. In such cases, the [[total dissolved solids]] (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.	+
−		+
−	Sulfuric acid is used as a defence by certain marine species, for example, the phaeophyte alga ''Desmarestia munda'' (order [[Desmarestiales]]) concentrates sulfuric acid in cell vacuoles.<ref name='Pelletreau'>{{cite journal|first = K.|last = Pelletreau|author2=Muller-Parker, G. |journal = Marine Biology|year = 2002|volume = 141|issue=1|pages=1–9|doi=10.1007/s00227-002-0809-6|title = Sulfuric acid in the phaeophyte alga Desmarestia munda deters feeding by the sea urchin Strongylocentrotus droebachiensis}}</ref>	+
−		+
−	===Extraterrestrial sulfuric acid===	+
−		+
−	====Venus====	+
−	Sulfuric acid is produced in the upper atmosphere of [[Venus]] by the [[Sun]]'s [[photochemistry|photochemical]] action on [[carbon dioxide]], [[sulfur dioxide]], and [[water]] vapor. [[Ultraviolet]] [[photon]]s of wavelengths less than 169&nbsp;[[nanometre|nm]] can [[photodissociation|photodissociate]] carbon dioxide into [[carbon monoxide]] and atomic [[oxygen]]. Atomic oxygen is highly reactive. When it reacts with sulfur dioxide, a trace component of the Venusian atmosphere, the result is [[sulfur trioxide]], which can combine with water vapor, another trace component of Venus's atmosphere, to yield sulfuric acid. In the upper, cooler portions of Venus's atmosphere, sulfuric acid exists as a liquid, and thick sulfuric acid clouds completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45–70&nbsp;km above the planet's surface, with thinner hazes extending as low as 30&nbsp;km and as high as 90&nbsp;km above the surface. The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain.	+
−		+
−	The atmosphere exhibits a sulfuric acid cycle. As sulfuric acid rain droplets fall down through the hotter layers of the atmosphere's temperature gradient, they are heated up and release water vapor, becoming more and more concentrated. When they reach temperatures above 300&nbsp;°C, sulfuric acid begins to decompose into sulfur trioxide and water, both in the gas phase. Sulfur trioxide is highly reactive and dissociates into sulfur dioxide and atomic oxygen, which oxidizes traces of carbon monoxide to form carbon dioxide. Sulfur dioxide and water vapor rise on convection currents from the mid-level atmospheric layers to higher altitudes, where they will be transformed again into sulfuric acid, and the cycle repeats.	+
−		+
−	====Europa====	+
−	Infrared spectra from [[NASA]]'s [[Galileo (spacecraft)|Galileo mission]] show distinct absorptions on [[Jupiter]]'s moon [[Europa (moon)|Europa]] that have been attributed to one or more sulfuric acid hydrates. Sulfuric acid in solution with water causes significant [[freezing-point depression]] of water's [[melting point]], down to {{convert|210|K|°C}}, and this would make more likely the existence of liquid solutions beneath Europa's icy crust.The interpretation of the spectra is somewhat controversial. Some planetary scientists prefer to assign the spectral features to the sulfate ion, perhaps as part of one or more minerals on Europa's surface.<ref>{{cite journal |first=T. M. |last=Orlando |first2=T. B. |last2=McCord |first3=G. A. |last3=Grieves |title=The chemical nature of Europa surface material and the relation to a subsurface ocean |journal=[[Icarus (journal)|Icarus]] |volume=177 |year=2005 |issue=2 |pages=528–533 |doi=10.1016/j.icarus.2005.05.009 |bibcode=2005Icar..177..528O}}</ref>	+
−		+
−	==Manufacture==	+
−	{{Main|Contact process|Wet sulfuric acid process}}	+
−		+
−	Sulfuric acid is produced from [[sulfur]], oxygen and water via the conventional [[contact process]] (DCDA) or the [[wet sulfuric acid process]] (WSA).	+
−		+
−	===Contact process===	+
−	{{main|Contact process}}	+
−	In the first step, sulfur is burned to produce sulfur dioxide.	+
−	: S (s) + {{chem|O|2}} (g) → {{chem|SO|2}} (g)	+
−		+
−	This is then oxidized to sulfur trioxide using oxygen in the presence of a [[vanadium(V) oxide]] [[catalyst]]. This reaction is reversible and the formation of the sulfur trioxide is exothermic.	+
−	: 2 {{chem|SO|2}} (g) + {{chem|O|2}} (g) {{eqm}} 2 {{chem|SO|3}} (g) (in presence of {{chem|V|2|O|5}})	+
−		+
−	The sulfur trioxide is absorbed into 97–98% {{chem|H|2|SO|4}} to form [[oleum]] ({{chem|H|2|S|2|O|7}}), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.	+
−		+
−	: {{chem|H|2|SO|4}} (l) + {{chem|SO|3}} (g)→ {{chem|H|2|S|2|O|7}} (l)	+
−		+
−	: {{chem|H|2|S|2|O|7}} (l) + {{chem|H|2|O}} (l) → 2 {{chem|H|2|SO|4}} (l)	+
−		+
−	Note that directly dissolving {{chem|SO|3}} in water is not practical due to the highly [[Exothermic reaction|exothermic]] nature of the [[Chemical reaction|reaction]] between sulfur trioxide and water. The reaction forms a corrosive aerosol that is very difficult to separate, instead of a liquid.	+
−	: {{chem|SO|3}} (g) + {{chem|H|2|O}} (l) → {{chem|H|2|SO|4}} (l)	+
−		+
−	===Wet sulfuric acid process===	+
−	{{main|Wet sulfuric acid process}}	+
−	In the first step, sulfur is burned to produce sulfur dioxide:	+
−	: S(s) + {{chem|O|2}}(g) → {{chem|SO|2}}(g)	+
−		+
−	or, alternatively, [[hydrogen sulfide]] ({{chem|H|2|S}}) gas is incinerated to {{chem|SO|2}} gas:	+
−	: 2 {{chem|H|2|S}} + 3 {{chem|O|2}} → 2 {{chem|H|2|O}} + 2 {{chem|SO|2}} (−518&nbsp;kJ/mol)	+
−	This is then oxidized to sulfur trioxide using oxygen with [[vanadium(V) oxide]] as [[catalyst]].	+
−	: 2 {{chem|SO|2}} + {{chem|O|2}} → 2 {{chem|SO|3}} (−99&nbsp;kJ/mol) (reaction is reversible)	+
−		+
−	The sulfur trioxide is hydrated into sulfuric acid {{chem|H|2|SO|4}}:	+
−	: {{chem|SO|3}} + {{chem|H|2|O}} → {{chem|H|2|SO|4}}(g) (−101&nbsp;kJ/mol)	+
−		+
−	The last step is the condensation of the sulfuric acid to liquid 97–98% {{chem|H|2|SO|4}}:	+
−	: {{chem|H|2|SO|4}}(g) → {{chem|H|2|SO|4}}(l) (−69&nbsp;kJ/mol)	+
−		+
−	===Other methods===	+
−	Another method is the less well-known metabisulfite method, in which [[metabisulfite]] is placed at the bottom of a beaker, and 12.6 molar concentration [[hydrochloric acid]] is added. The resulting gas is bubbled through [[nitric acid]], which will release brown/red vapors. The completion of the reaction is indicated by the ceasing of the fumes. This method does not produce an inseparable mist, which is quite convenient.	+
−		+
−	Sulfuric acid can be produced in the laboratory by burning sulfur in air and dissolving the gas produced in a [[hydrogen peroxide]] solution.	+
−		+
−	: SO<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub>	+
−		+
−	Prior to 1900, most sulfuric acid was manufactured by the [[lead chamber process]].<ref>{{cite journal |first=Edward M. |last=Jones |title=Chamber Process Manufacture of Sulfuric Acid |journal=Industrial and Engineering Chemistry |year=1950 |volume=42 |issue=11 |pages=2208–2210 |doi=10.1021/ie50491a016 }}</ref> As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.	+
−		+
−	In early to mid nineteenth century "vitriol" plants existed, among other places, in [[Prestonpans]] in Scotland, [[Shropshire]] and the [[Lagan Valley]] in County Antrim Ireland where it was used as a bleach for linen. Early bleaching of linen was done using milk but this was a slow process and the use of vitriol sped up the bleaching process.	+
−		+
−	==Uses==	+
−	[[File:2000sulphuric acid.PNG|thumb|right|300px|Sulfuric acid production in 2000]]	+
−	Sulfuric acid is a very important commodity chemical, and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength.<ref>{{cite book |last=Chenier |first=Philip J. |title=Survey of Industrial Chemistry |pages=45–57 |publisher=John Wiley & Sons |location=New York |year=1987 |isbn=0-471-01077-4 }}</ref> World production in 2004 was about 180 million [[tonne]]s, with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%.<ref>{{cite book|author=Davenport, William George and King, Matthew J.|title=Sulfuric acid manufacture: analysis, control and optimization|url=http://books.google.com/books?id=tRAb2CniRG4C|accessdate=23 December 2011|year=2006|publisher=Elsevier|isbn=978-0-08-044428-4|pages=8, 13}}</ref> Most of this amount (~60%) is consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and [[antifreeze]], as well as in various processes such as oil well acidicizing, aluminium reduction, paper sizing, water treatment. About 6% of uses are related to [[pigment]]s and include paints, [[Enamel paint|enamels]], printing inks, coated fabrics and paper, and the rest is dispersed into a multitude of applications such as production of explosives, [[cellophane]], acetate and viscose textiles, lubricants, non-ferrous metals and batteries.<ref>{{Greenwood&Earnshaw2nd|page=653}}</ref>	+
−		+
−	===Industrial production of chemicals===	+
−	The major use for sulfuric acid is in the "wet method" for the production of [[phosphoric acid]], used for manufacture of [[phosphate]] [[fertilizer]]s. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as [[fluorapatite]], though the exact composition may vary. This is treated with 93% sulfuric acid to produce [[calcium sulfate]], [[hydrogen fluoride]] (HF) and [[phosphoric acid]]. The HF is removed as [[hydrofluoric acid]]. The overall process can be represented as:	+
−		+
−	: {{chem|Ca|5|F(PO|4|)|3}} + 5 {{chem|H|2|SO|4}} + 10 {{chem|H|2|O}} → 5 {{chem|CaSO|4|·2 H|2|O}} + HF + 3 {{chem|H|3|PO|4}}	+
−		+
−	[[Ammonium sulfate]], an important nitrogen fertilizer, is most commonly produced as a byproduct from [[Coke (fuel)|coking plants]] supplying the iron and steel making plants. Reacting the [[ammonia]] produced in the thermal decomposition of [[coal]] with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.	+
−		+
−	Another important use for sulfuric acid is for the manufacture of [[aluminium sulfate]], also known as paper maker's alum. This can react with small amounts of soap on [[paper pulp]] fibers to give gelatinous aluminium [[carboxylate]]s, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making [[aluminium hydroxide]], which is used at [[water treatment]] plants to [[filter (water)|filter]] out impurities, as well as to improve the taste of the [[water]]. [[Aluminium sulfate]] is made by reacting [[bauxite]] with sulfuric acid:	+
−		+
−	: 2 {{chem|Al|O|(OH)}} + 3 {{chem|H|2|SO|4}} → {{chem|Al|2|(SO|4|)|3}} + 4 {{chem|H|2|O}}	+
−		+
−	Sulfuric acid is also important in the manufacture of [[dye]]stuffs solutions.	+
−		+
−	===Sulfur-iodine cycle===	+
−	The [[sulfur-iodine cycle]] is a series of thermo-chemical processes used to obtain [[hydrogen]]. It consists of three chemical reactions whose net reactant is [[water]] and whose net products are hydrogen and [[oxygen]].	+
−	:{|	+
−	|-	+
−	| 2 {{chem|H|2|SO|4}} → 2 {{chem|SO|2}} + 2 {{chem|H|2|O}} + {{chem|O|2}} || &nbsp;&nbsp;&nbsp; || (830&nbsp;°C)	+
−	|-	+
−	| {{chem|I|2}} + {{chem|SO|2}} + 2 {{chem|H|2|O}} → 2 HI + {{chem|H|2|SO|4}} || &nbsp;&nbsp;&nbsp; || (120&nbsp;°C)	+
−	|-	+
−	| 2 HI → {{chem|I|2}} + {{chem|H|2}} || &nbsp;&nbsp;&nbsp; || (320&nbsp;°C)	+
−	|}	+
−		+
−	The sulfur and [[iodine]] compounds are recovered and reused, hence the consideration of the process as a cycle. This process is [[endothermic]] and must occur at high temperatures, so energy in the form of heat has to be supplied.	+
−		+
−	The sulfur-iodine cycle has been proposed as a way to supply hydrogen for a [[hydrogen economy|hydrogen-based economy]]. It does not require [[hydrocarbons]] like current methods of [[steam reforming]]. But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it.	+
−		+
−	The sulfur-iodine cycle is currently being researched as a feasible method of obtaining hydrogen, but the concentrated, corrosive acid at high temperatures poses currently insurmountable safety hazards if the process were built on a large scale.	+
−		+
−	===Industrial cleaning agent===	+
−	{{Main|Cleaning agent}}	+
−	Sulfuric acid is used in large quantities by the [[iron]] and [[steel]]making industry to remove oxidation, [[rust]] and [[Fouling|scaling]] from rolled sheet and billets prior to sale to the [[automobile]] and [[major appliances]] industry.{{citation needed|date=September 2011}} Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid{{clarify|reason="What is it, exactly? Is it still the same acid, dirty, reacted, or what?"|date=February 2015}} with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous [[sulfur dioxide]] ({{chem|SO|2}}) and [[sulfur trioxide]] ({{chem|SO|3}}) which are then used to manufacture "new" sulfuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where sulfuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases.	+
−		+
−	===Catalyst===	+
−	Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of [[cyclohexanone oxime]] to [[caprolactam]], used for making [[nylon]]. It is used for making [[hydrochloric acid]] from [[salt]] via the [[Mannheim process]]. Much {{chem|H|2|SO|4}} is used in [[petroleum]] refining, for example as a catalyst for the reaction of [[isobutane]] with [[isobutylene]] to give [[isooctane]], a compound that raises the [[octane rating]] of [[gasoline]] (petrol).	+
−		+
−	===Electrolyte===	+
−	[[File:Acidic drain cleaner containing sulfuric acid (sulphuric acid).jpg|right|thumb|Acidic [[drain cleaner]]s usually contain sulfuric acid at a high concentration which turns a piece of [[pH paper]] red and chars it instantly, demonstrating both the strong acidic nature and dehydrating property.]]	+
−		+
−	Sulfuric acid acts as the electrolyte in [[lead-acid batteries|lead-acid (car) batteries]] (lead-acid accumulator):	+
−		+
−	At [[anode]]:	+
−	:{{chem|Pb}} + {{chem|SO|4}}<sup>2−</sup> {{unicode|⇌}} {{chem|PbSO|4}} + 2 e<sup>−</sup>	+
−		+
−	At [[cathode]]:	+
−	:{{chem|PbO|2}} + 4 H<sup>+</sup> + {{chem|SO|4}}<sup>2−</sup> + 2 e<sup>−</sup> {{unicode|⇌}} {{chem|PbSO|4}} + 2 H<sub>2</sub>O	+
−		+
−	[[File:Acidic drain opener.JPG|right|thumb|An acidic [[drain cleaner]] can be used to dissolve grease, hair and even tissue paper inside water pipes.]]	+
−		+
−	Overall:	+
−	:{{chem|Pb}} + {{chem|PbO|2}} + 4 H<sup>+</sup> + 2 {{chem|SO|4}}<sup>2−</sup> {{unicode|⇌}} 2 {{chem|PbSO|4}} + 2 H<sub>2</sub>O	+
−		+
−	===Domestic uses===	+
−	Sulfuric acid at high concentrations is frequently the major ingredient in [[drain cleaner|acidic drain cleaners]]<ref name="dc">{{cite web|url=http://www.staplesdisposables.com/uploads/products/B470FF98A27F414881DB3FE1A1116C93.pdf|title=Sulphuric acid drain cleaner|work=herchem.com }}</ref> which are used to remove [[lipids|grease]], [[hair]], [[tissue paper]], etc. Similar to their [[drain opener|alkaline versions]], such drain openers can dissolve fats and proteins via [[hydrolysis]]. Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.	+
−		+
−	==Safety==	+
−		+
−	===Laboratory hazards===	+
−	[[File:Sulfuric acid burning tissue paper.jpg|thumb|left|Drops of 98% sulfuric acid char a piece of tissue paper instantly. Carbon is left after the dehydration reaction staining the paper black.]]	+
−		+
−	Sulfuric acid is capable of causing very severe burns, especially when it is at high [[concentrations]]. In common with other corrosive [[acids]] and [[alkali]], it readily decomposes [[proteins]] and [[lipids]] through [[amide hydrolysis|amide]] and [[ester hydrolysis]] upon contact with [[Tissue (biology)|living tissues]], such as [[skin]] and [[flesh]]. In addition, it exhibits a strong [[Dehydration reaction|dehydrating property]] on [[carbohydrates]], liberating extra [[heat]] and causing [[burn#By depth|secondary thermal burns]].<ref name="OA"/><ref name=TB/> Accordingly, it rapidly attacks the [[cornea]] and can induce [[blindness|permanent blindness]] if splashed onto [[eye]]s. If ingested, it damages [[internal organs]] irreversibly and may even be fatal.<ref name="ds"/>  [[Protective equipment]] should hence always be used when handling it. Moreover, its [[oxidizing|strong oxidizing property]] makes it highly corrosive to many [[metal]]s and may extend its destruction on other materials.<ref name="OA"/> Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable [[strong acids]], such as [[hydrochloric acid]] and [[nitric acid]].	+
−	<div style="float: right; margin-left: 1.0 em">[[File:Hazard C.svg|70px]] [[File:Dangclass8.png|70px]]	+
−	</div>	+
−		+
−	Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5&nbsp;M are labeled "CORROSIVE", while solutions greater than 0.5&nbsp;M but less than 1.5&nbsp;M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1&nbsp;M, 10%) will char paper if left in contact for a sufficient time.	+
−		+
−	The standard first aid treatment for acid spills on the skin is, as for other [[corrosive|corrosive agents]], irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.	+
−		+
−	Preparation of the diluted acid can also be dangerous due to the heat released in the dilution process. The concentrated acid is always added to water and not the other way around, to take advantage of the relatively high [[heat capacity]] of water. Addition of water to concentrated sulfuric acid leads to the dispersal of a sulfuric acid [[aerosol]] or worse, an [[explosion]]. Preparation of solutions greater than 6&nbsp;M (35%) in concentration is most dangerous, as the heat produced may be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (such as an ice bath) are essential.	+
−		+
−	On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.	+
−		+
−	===Industrial hazards===	+
−	Although sulfuric acid is non-flammable, contact with metals in the event of a spillage can lead to the liberation of [[hydrogen]] gas. The dispersal of acid aerosols and gaseous sulfur dioxide is an additional hazard of fires involving sulfuric acid.	+
−		+
−	The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent [[pulmonary edema]] if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the [[respiratory tract]] are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent.<ref>[http://www.ncbi.nlm.nih.gov/pubmed/3479642 Lung cancer mortality in workers exposed to sulfuric acid mist and other acid mists.]</ref> In the United States, the [[permissible exposure limit]] (PEL) for sulfuric acid is fixed at 1&nbsp;mg/m<sup>3</sup>: limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to [[vitamin B12 deficiency]] with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show [[demyelination]], loss of [[axon]]s and [[gliosis]].	+
−		+
−	==Legal restrictions==	+
−	International commerce of sulfuric acid is controlled under the [[United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances|United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988]], which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances.<ref name=incb>[http://web.archive.org/web/20080227224025/http://www.incb.org/pdf/e/list/red.pdf Annex to Form D ("Red List")], 11th Edition, January 2007 (p. 4). [[International Narcotics Control Board]]. [[Vienna, Austria]].</ref>	+
−		+
−	==See also==	+
−	* [[Aqua regia]]	+
−	* [[Piranha solution]]	+
−	* [[Sulfuric acid poisoning]]	+
−	* [[Sulfur oxoacid]]	+
	==References==		==References==
−	{{Reflist|30em}}	+	===Relevant Sciencemadness threads===
−	==External links==	+	[[Category:Strong acids]]
−	{{Commons category|Sulfuric acid}}	+	[[Category:Acids]]
−	* {{ICSC|0362|03}}	+	[[Category:Inorganic compounds]]
−	* [http://www.periodicvideos.com/videos/mv_sulfuric_acid.htm Sulfuric acid] at ''[[The Periodic Table of Videos]]'' (University of Nottingham)	+
−	* [http://www.cdc.gov/niosh/npg/npgd0577.html NIOSH Pocket Guide to Chemical Hazards]	+
−	* [http://www.cdc.gov/niosh/topics/sulfuric-acid/ CDC – Sulfuric Acid – NIOSH Workplace Safety and Health Topic]	+
−	* [http://ptcl.chem.ox.ac.uk/MSDS/SU/sulfuric_acid_concentrated.html External Material Safety Data Sheet]	+
−	*Calculators: [http://www.aim.env.uea.ac.uk/aim/surftens/surftens.php surface tensions], and [http://www.aim.env.uea.ac.uk/aim/density/density_electrolyte.php densities, molarities and molalities] of aqueous sulphuric acid	+
−	* [http://www2.iq.usp.br/docente/gutz/Curtipot_.html Sulfuric acid analysis – titration freeware]	+
−	* Process flowsheet of sulfuric acid manufacturing by [http://www.inclusive-science-engineering.com/manufacture-of-h2so4-by-chamber-process/manufacture-of-h2so4-by-chamber-process-2/ lead chamber process]	+
−		+
−	{{Use dmy dates|date=August 2010}}	+
−	{{Hydrogen compounds}}	+
−	{{sulfur compounds}}	+
−	{{Sulfates}}	+
−		+
−	{{Authority control}}	+
−		+
−	{{DEFAULTSORT:Sulfuric Acid}}	+
−	[[Category:Alchemical substances]]	+
−	[[Category:Equilibrium chemistry]]	+
−	[[Category:Hydrogen compounds]]	+
−	[[Category:Inorganic solvents]]	+
−	[[Category:Mineral acids]]	+
−	[[Category:Oxidizing acids]]	+
	[[Category:Oxidizing agents]]		[[Category:Oxidizing agents]]
−	[[Category:Photographic chemicals]]
	[[Category:Sulfates]]		[[Category:Sulfates]]
−	[[Category:Sulfur oxoacids]]	+	[[Category:Chemical compounds]]
−	[[Category:Sulfur]]	+	[[Category:Things that can kill you very quickly]]
−	[[Category:Dehydrating agents]]	+	[[Category:Mineral acids]]
		+	[[Category:Materials unstable in basic solution]]
		+	[[Category:Corrosive chemicals]]

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Revision as of 08:30, 3 August 2015

Sulfuric acid (alternative spelling sulphuric acid), represented by the molecular formula H₂SO₄, is one of the most important acids in chemistry and the most important chemical to industries in the world. It is the strongest easily available acid, with a pK_a of -3.

Properties

Chemical properties

Sulfuric acid is a diprotic acid, it is able to give away two protons (H⁺). It first dissociates to form hydronium and hydrogen sulfate, with a pKa of -3, indicative of a strong acid:

H₂SO₄ + H₂O → H₃O + HSO₄⁻

The second dissociation forms sulfate and another hydronium ion from a hydrogen sulfate ion. It has a pKa of 1.99, indicative of a weak acid, and occurs like this:

HSO₄⁻ + H₂O ⇌ H₃O⁺ + SO₄^2-

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. This property is useful for producing large amounts of sulfur dioxide for use as a reducing agent if water is continually removed.

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 nitrates.

In organic chemistry, sulfuric acid is the most practical acid in most cases where a source of H₃O⁺ 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^- 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.

When concentrated, it is strongly hygroscopic and has strong dehydrating properties. It can break down most organic molecules containing OH^- groups to use them to form water, leaving only the carbon behind. This property is exploited in the famous "black snake" demonstration, where sulfuric acid dehydrates sucrose (table sugar), forming water with the hydrogen and oxygen atoms and leaving amorphous carbon behind.

Physical properties

Sulfuric acid is a oily liquid at room temperature. It is colorless, but often has a very light yellow color when slightly contaminated with iron ions. 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.

Sulfuric acid's boiling point raises with the concentration as described in this figure to the right. An azeotrope forms at 98% w/w.

Sources, production, and concentration

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 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 and will appear ye

For some amateurs, it can be hard to find concentrated sulfuric acid, with acid drain cleaners being banned(as a result of acid throwing or illicit drug manufacture) or very contaminated in some countries. Some brand of battery acids however are quite pure and easy to find almost anywhere. To concentrate it from 35 to 98%, one can boil it down to evaporate the water. It is advised to use glassware with anti-splash adapters, as sulfuric acid tends to bump a lot while it boils. The sulfuric acid will be quite concentrated when the temperature reaches around 300°C. The maximum reachable is 98.3% w/w, a point at which the boiling point will be 327°C. Extreme care must be taken when boiling sulfuric acid. The hazards and safety tips section must be read. Near the end, the acid will decompose into various sulfur oxides.

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 are added, it forms pyrosulfuric acid, which is a solid at room temperature.

Projects

• Producing metal sulfates
• Producing nitro compounds through nitration
• The dehydration of sucrose to produce elemental carbon
• Esterifications that require a dehydrating agent, such as that of methyl salicylate
• Making simple rayon fibers with Schweizer's reagent and cellulose
• Producing other concentrated acids by reaction of sulfuric acid with an anhydrous salt, such as in the production of fuming nitric acid and glacial acetic acid

Handling

Safety

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.

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 phenomena, 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.

Hot concentrated sulfuric acid may decompose to form sulfur dioxide and sulfur trioxide, which are toxic and corrosive, respectively.

Storage

Sulfuric acid should be stored in closed bottles.

References

Relevant Sciencemadness threads