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elementcollector1
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Sounds interesting, but remember that CaCl2 is soluble: A much better sodium derivative would be the carbonate/bicarbonate. The sulfate does not work
for this process because, surprisingly enough, calcium sulfate has roughly the same solubility as the hydroxide.
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AJKOER
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I came up with an alternate idea to prepare NaOH from Fe, HCl, NH4OH and Na2CO3.
Step 1. Prepare Fe(OH)2 from, as an example, the reaction of NH4OH and FeCl2. The latter from:
Fe + 2 HCl --> FeCl2 + H2 (g)
Step 2. React Na2CO3 with Fe(OH)2:
Na2CO3 + Fe(OH)2 --> FeCO3 (s) + 2 NaOH
where on this last reaction, one needs an increased solubility of Ferrous hydroxide in Na2CO3 to move this reaction to the right. However, more likely
the solubility is reduced with increasing OH- concentration and starting with an excess of Sodium carbonate may not be sufficient.
Perhaps better cheaper/available amphoteric alternatives from Al(OH)3, Cu(OH)2, Zn(OH)2 or Pb(OH)2. For example (see http://gcsechemistryhelp.tumblr.com/post/33791502799/in-my-c... ):
3 Na2CO3 + 2 Al(OH)3 —> 6 NaOH + Al2(CO3)3
Now, the particular issue with this synthesis is the unstable nature of Aluminum carbonate decomposing at 120 C into Al(OH)3 and CO2 in the presence
of water. As both of these decomposition products could react with the newly formed NaOH, this could present a challenge (requires a low temperature
synthesis). In general even without the decomposition of the Aluminum carbonate, any unreacted Al(OH)3 (need some excess Na2CO3) could be attacked by
the newly created NaOH forming NaAl(OH)4 reducing NaOH yield.
[Edited on 19-1-2013 by AJKOER]
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S.C. Wack
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Use the hypochlorite to make ferrate...
An old process (Loewig, DE 1650, 1877) for NaOH production not involving lime uses ferric oxide from iron ore. It is heated with sodium carbonate to
form sodium ferrate. Apparently this can be washed with cold water and not be decomposed, removing excess carbonate. The addition of boiling water
gives strong NaOH solution and Fe2O3 again, which can be reused. This solution is much stronger than that which can be produced by lime, without
evaporation.
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chemicalmixer
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I think I have a better method to make lye, starting with baking soda:
First, heat the soda to convert it to washing soda:
2NaHCO3 --> Na2CO3 + H2O + CO2
Now, make a saturated solution of washing soda in DH2O, and make an electrochemical cell with it, by applying DC electricity to the solution using two
nickel, or nickel-plated electrodes.
Nickel(III) oxide-hydroxide is not amphoteric, and thus not soluble in caustic solutions. This is why nickel iron batteries work (which use
concentrated KOH as an electrolyte). Ni(III) oxide-hydroxide also conducts current much better than other similar metal oxides. An iron cathode would
also work instead of Ni, but Ni must be used as the anode. Also, nickel is one of the easier metals that can be successfully plated by the home
experimenter using common chemicals, thus an iron substrate could be heavily plated with nickel prior to use as an anode here.
A permeable partition could be used to keep the CO/CO2 formed at the anode from reacting with the sodium ions, but this is probably unnecessary. A pH
meter could help determine when all of the carbonate has been converted to hydroxide.
[Edited on 19-1-2013 by chemicalmixer]
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Manifest
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Quote: Originally posted by chemicalmixer  | I think I have a better method to make lye, starting with baking soda:
First, heat the soda to convert it to washing soda:
2NaHCO3 --> Na2CO3 + H2O + CO2
Now, make a saturated solution of washing soda in DH2O, and make an electrochemical cell with it, by applying DC electricity to the solution using two
nickel, or nickel-plated electrodes.
Nickel(III) oxide-hydroxide is not amphoteric, and thus not soluble in caustic solutions. This is why nickel iron batteries work (which use
concentrated KOH as an electrolyte). Ni(III) oxide-hydroxide also conducts current much better than other similar metal oxides. An iron cathode would
also work instead of Ni, but Ni must be used as the anode. Also, nickel is one of the easier metals that can be successfully plated by the home
experimenter using common chemicals, thus an iron substrate could be heavily plated with nickel prior to use as an anode here.
A permeable partition could be used to keep the CO/CO2 formed at the anode from reacting with the sodium ions, but this is probably unnecessary. A pH
meter could help determine when all of the carbonate has been converted to hydroxide.
[Edited on 19-1-2013 by chemicalmixer] |
If he doesn't have Sodium Hydroxide, I highly doubt that he has DH2O
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12AX7
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| Quote: Originally posted by S.C. Wack | Use the hypochlorite to make ferrate...
An old process (Loewig, DE 1650, 1877) for NaOH production not involving lime uses ferric oxide from iron ore. It is heated with sodium carbonate to
form sodium ferrate. Apparently this can be washed with cold water and not be decomposed, removing excess carbonate. The addition of boiling water
gives strong NaOH solution and Fe2O3 again, which can be reused. This solution is much stronger than that which can be produced by lime, without
evaporation. |
Hmm. Ferrate(III), aka ferrite. Just to be clear! 
Chrome might also be a candidate, though Na2Cr2O3 itself may be too soluble (and also lead to chromate!).
Tim
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S.C. Wack
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Your interjection is not clear. I'm obviously talking about +6 ferrate as it's always been known. My post was not a commentary on the patent and the
ease of (or lack of) oxidation in that way.
...whether they are making Na2Fe2O4 or what, I wasn't recommending going that way...just mentioning it since hypochlorite came up...
[Edited on 20-1-2013 by S.C. Wack]
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AJKOER
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| Quote: Originally posted by S.C. Wack | Use the hypochlorite to make ferrate...
An old process (Loewig, DE 1650, 1877) for NaOH production not involving lime uses ferric oxide from iron ore. It is heated with sodium carbonate to
form sodium ferrate. Apparently this can be washed with cold water and not be decomposed, removing excess carbonate. The addition of boiling water
gives strong NaOH solution and Fe2O3 again, which can be reused. This solution is much stronger than that which can be produced by lime, without
evaporation. |
Here is a reference that confirms the above Fe2O3 and Na2CO3 fusion reaction producing a ferrite ion described as:
Na2C03 + Fe203 → 2NaFe02 + C02
Link: http://www.sumobrain.com/patents/wipo/Apparatus-method-produ...
The reaction apparently occurs at 'bright red heat' in a rotating furnace (see http://www.lenntech.com/chemistry/caustic-soda.htm ).
The hydrolysis reaction is probably given by:
NaFeO2 + 2 H2O = Fe(OH)3 + NaOH
The water is hot steam at 900 C and the Fe(OH)3 also breaks down releasing Fe2O3.
This fusion synthesis may not be practical for the home.
---------------------------------------------------------------------------------
I have, however, have been looking at wet room temperature ferrate production via FeCl3 and Na2CO3.H2O2 (Sodium percarbonate) plus a weak hydroxide.
The Na2CO3. H2O2 is substitute candidate for NaClO + a strong base which are ingredients in the classic hypochlorite ferrate synthesis.
[Edited on 20-1-2013 by AJKOER]
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Random
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| Quote: Originally posted by AJKOKER |
I have, however, have been looking at wet room temperature ferrate production via FeCl3 and Na2CO3.H2O2 (Sodium percarbonate) plus a weak hydroxide.
The Na2CO3. H2O2 is substitute candidate for NaClO + a strong base which are ingredients in the classic hypochlorite ferrate synthesis.
[Edited on 20-1-2013 by AJKOER] |
You know, ferrates are so unstable species that they only exist in very basic solutions, easily decompose to ferric oxides and stuff like that. So if
you want to attempt synthesis of some oxidant, you would do better with wet synthesis of permanganate. Unfortunatelly, as it's been said already on
this forum, yields are very small.
Ferrates are best synthesized using dry fusion method as I heard, maybe use molten NaOH and oxidant such as nitrate, chlorate? But yeah, then again
you would be better using MnO2.
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12AX7
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That's what I thought -- ferrate(VI) is so difficult to create, and so unstable, that it can't possibly be formed at high temperature from such simple
reagents as soda ash and oxygen.
Tim
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AJKOER
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OK, per this reference ( http://www.lenntech.com/chemistry/caustic-soda.htm ), here is an old process for making NaOH. To quote:
"To make very strong caustic, zinc oxide is often used to remove the sulphide from the tank liquors : -
Na2S+ ZnO + H20 = 2NaOH + ZnS
The precipitated zinc sulphide is settled out, before evaporating the caustic liquor. By calcining the zinc sulphide, the zinc is reconverted to
oxide."
Note, water is consumed in the above reaction permitting the formation of concentrated Sodium hydroxide.
---------------------------------------------
So to create the NaOH, one apparently needs Sodium sulfide and Zinc oxide. But, how does one make Na2S at home? Here is one possible route starting
with Sulfur:
2 Al + 3 S --> Al2S3
Then add the Aluminum sulfide to an aqueous solution of Na2CO3. Expected reactions:
Al2S3 + 6 H2O --> 2 Al(OH)3 + 3 H2S
Na2CO3 + H2S --> NaHS + NaHCO3 (see http://books.google.com/books?id=rpzrIZW-OcEC&pg=PA387&a... )
Followed by boiling or aeration to oxidize the Sodium hydrogen sulfide (see www.allreactions.com/index.php/group-1a/natrium/sodium-hydro... ):
2 NaHS (solution) --> Na2S + H2S↑ (boiling)
[Caution: Hydrogen sulfide is quite a toxic gas. Upon deadening ones' sense of smell, it is also an insidious poison with a delayed mortality effect.
Use ventilation and appropriate safety measures]
2n NaHS (solid) + (n - 1)O2 = 2 H2O + (2n - 4) NaOH + 2 Na2(Sn) [100—250° С]
Attention should be paid to avoid the presence of any unreacted H2S or Sulfur which could react with any created NaOH (forming NaHS, Na2S and/or
Na2S2O3). The presence of polysulfides themselves may not present a significant issue (forming an insoluble Zinc polysulfide).
Lastly, a reference on Zinc oxide creation (link: http://www.imm.ac.cn/journal/ccl/1506/150630-733-03-0359-p4.... ). To quote:
"The simple approach (precipitation—heat treatment) to fabrication of ZnO prickly spheres by
dehydration of the precursor obtained via chemical reaction between Zn(CH3COO)2·2H2O
and NH3·H2O in the presence of surfactant (SDS) was reported. Compared with other
methods the reaction condition was considerably moderate and the temperature was lower.
Moreover, prickly spheres could be obtained in high yield."
[Edited on 22-1-2013 by AJKOER]
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elementcollector1
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Distilled water? I can get that from the local Safeway for 99 cents. I hope you aren't thinking of deuterium oxide...
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chemicalmixer
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Distilled water would ideal, but not necessary. How about actually contemplating what I've said, instead of nitpicking with your BS,
non-relavent criticism? Electrolysis of a Na2CO3 or NaHCO3 solution with a Ni anode seems a hell of a lot more pratical for the average home
experimenter than building a lime kiln or making sodium ferrate.
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AJKOER
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| Quote: Originally posted by AJKOER | OK, per this reference ( http://www.lenntech.com/chemistry/caustic-soda.htm ), here is an old process for making NaOH. To quote:
"To make very strong caustic, zinc oxide is often used to remove the sulphide from the tank liquors : -
Na2S+ ZnO + H20 = 2NaOH + ZnS
The precipitated zinc sulphide is settled out, before evaporating the caustic liquor. By calcining the zinc sulphide, the zinc is reconverted to
oxide."
Note, water is consumed in the above reaction permitting the formation of concentrated Sodium hydroxide.
|
Actually, per this reaction (see http://www.allreactions.com/index.php/group-1a/natrium/sodiu... ):
2 NaOH (conc. 60%) + H2O + ZnO = Na2[Zn(OH)4] (90°С )
one should be mindful of the reaction temperature (keep below 90 C) and/or the NaOH concentration (under 60%) in the presence of any excess ZnO per
the synthesis:
Na2S + ZnO + H20 --> 2NaOH + ZnS (s)
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LanthanumK
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I tried the electrolysis method and after quite a lot of current only came up with a little NaOH. It is not the most efficient method but it does
work.
hibernating...
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elementcollector1
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Depends on how long, how much current in amps, and what kind of cell you used. I assume the electrolysis of Na2CO3/NaHCO3 works because the carbon
dioxide produced leaves the solution, thus shifting the reaction in favor of NaOH. This should work even in a 1-cell, concentrated apparatus, so I'll
give this a shot when I get the time (and when I can find my darn nickel scraps!)
Titration of your NaOH, or boiling it down, weighing it, and comparing it to the starting Na salt (assuming the NaOH is now pure) would be a good
method of determining efficiency.
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Manifest
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LOL. Sorry, it was late at night when I posted that....
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Manifest
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| Quote: Originally posted by chemicalmixer |
Distilled water would ideal, but not necessary. How about actually contemplating what I've said, instead of nitpicking with your BS,
non-relavent criticism? Electrolysis of a Na2CO3 or NaHCO3 solution with a Ni anode seems a hell of a lot more pratical for the average home
experimenter than building a lime kiln or making sodium ferrate. |
I'm sorry. Again, it was late at night....
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science_guy1
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I read that sodium carbonate decomposes to sodium oxide at high temperatures.1 The sodium oxide could then be reacted with water to
produce lye.
I propane torched a crucible containing 30g of sodium carbonate for ten minutes but I failed to produce sodium oxide. Perhaps a furnace would work.
Any thoughts on this method? I have not seen it discussed.
Also I saw a youtube video of someone electrolyzing molten NaCl and producing sodium metal. You could then throw the sodium metal in water to make
lye.
Footnotes:
1: "In practice, ordinary glass is formed by melting sand (quartz crystals) with sodium carbonate and calcium carbonate. The carbonates decompose to
the oxides plus carbon dioxide." - General Chemistry 9th ed, D. Ebbing
[Edited on 4-4-2013 by science_guy1]
[Edited on 4-4-2013 by science_guy1]
[Edited on 4-4-2013 by science_guy1]
[Edited on 5-4-2013 by science_guy1]
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elementcollector1
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Oh, for science's sake...
You're going to need a temperature much higher than what a propane torch can provide.
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AJKOER
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Here is a theoretically interesting preparation for NaOH with practical issues as it involves working with Chloramine (toxic and potentially explosive
vapor, see MSDS at http://www.guidechem.com/dictionary/10599-90-3.html ).
NH3 (g) + NaOCl ---> NaOH + NH2Cl (g) (see http://www.buzzle.com/articles/mixing-bleach-and-ammonia.htm... )
More precisely, slowly pass ammonia gas (or, in the form of drops) into Sodium hypochlorite solution. Chloramine fumes should be liberated due to the
exothermic reaction, leaving largely aqueous NaOH.
To address NH2Cl vapor, use a fumehood, or lead it into dilute H2O2. Reactions:
NH2Cl + H2O <--> HOCl + NH3
HOCl + H2O2 --> HCl + H2O + O2
However, using Chlorine bleach (NaOCl, NaCl, Na2CO3, NaOH,..see Chlorox product list at http://www.clorox.com/products/clorox-regular-bleach/ ), means that the NaOH product would be contaminated with salt and Sodium carbonate. Adding
some NaOH to the final product solution and cooling should separate out the impurities as NaOH is about 3 and 8 times more soluble, respectively, than
NaCl and Na2CO3.
Yet another issue with this reaction is that the Chloramine must be completely removed by warming from the solution. If not, some Hydrazine (described
by Wikipedia as being 'highly toxic', see http://en.wikipedia.org/wiki/N2H4 ) can be formed:
NH2Cl + NH3 --> N2H4 + HCl
When the addition of NH3 to NaOCl is no longer reactive, adding more NaOCl to restart the process could also remove any formed Hydrazine (although
this may prove to be unwise as the reaction between any formed N2H4 and NaOCl may prove to be too energetic even with a small addition of bleach, see
the MSDS for pure Hydrazine at https://docs.google.com/viewer?a=v&q=cache:kHchCn9SNMkJ:... ).
Note, this preparation for Sodium hydroxide must be repeated several fold just to create a dilute impure solution of NaOH as the starting
concentration of NaOCl (with NaCl and Na2CO3) is generally low, and given the exothermic nature of this reaction with toxic fume production (a
significant safety issue), this dilution is actually desirable. Hence, this preparation is not a recommended practical path for several reasons, but I
thought its associated issues should be properly addressed.
[EDIT] Chloramine has recently been incorporated in certain districts as a desirable water purification agent. I will not take sides on this issue
here, but I believe, I have witnessed some down playing (white washing?) of the associated toxic and mutagenic properties of NH2Cl (as support, note
the sparse MSDS by Guide Chemical cited above, which is one of the few remaining sample reports mentioning any negative issues). As such, readers
should be conscious of this controversy when reading up Chloramine, and what I believe are valid toxicity issues, which are perhaps, IMHO, more
recently not as properly put forward.
[Edited on 10-4-2013 by AJKOER]
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ElectroWin
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if your'e going to use NaHCO3 or Na2CO3, then you might as well skip all the sulphur and use lime to steal the carbonate:
Na2CO3 (aq) + Ca(OH)2 --> 2 NaOH (aq) + CaCO3 (s)
the CaCO3 is only slightly soluble. not sure what to do if you need much higher purity.
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annaandherdad
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ktw_100, you're getting some great advice from these guys on how to make NaOH, and I'm all for doing something for its educational value. But NaOH is
a cheap chemical, easily available from the hardware store as drain cleaner or from chemical supply places if you need purity. The amount of work
involved in making it for yourself is hardly worth it, unless it's just for fun.
Any other SF Bay chemists?
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Random
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How concentrated NaOH solution can we get with baking soda reaction with lime?
I mean NaHCO3 will react with Ca(OH)2 to form CaCO3 and Na2CO3.
Na2CO3 will proceed to react with Ca(OH)2 to form CaCO3 and NaOH.
Once NaOH is formed, could we basically dissolve more Na2CO3 and Ca(OH)2 in already formed solution to obtain even higher concentrations?
What is the limit where NaOH could basically push solubility of other two reactants downwards?
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blogfast25
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| Quote: Originally posted by Random | How concentrated NaOH solution can we get with baking soda reaction with lime?
I mean NaHCO3 will react with Ca(OH)2 to form CaCO3 and Na2CO3.
Na2CO3 will proceed to react with Ca(OH)2 to form CaCO3 and NaOH.
Once NaOH is formed, could we basically dissolve more Na2CO3 and Ca(OH)2 in already formed solution to obtain even higher concentrations?
What is the limit where NaOH could basically push solubility of other two reactants downwards? |
Sodium bicarbonate isn't very soluble in water, so it's not a great choice.
With the slaked lime/washing soda method you can probably go to 30 % (or slightly higher) of NaOH. But isolating the NaOH as solid lye... that's the
REAL challenge.
As others have surely remarked here: the slaked lime/soda (or potash) reaction is a nice little demonstration of a displacement reaction but lye is so
cheap and so OTC (see soap making sites e.g.) that it's not worth trying to prepare at home in any significant quantities.
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