| Pages:
1
2 |
jimmyboy
Hazard to Others
Posts: 235
Registered: 1-3-2004
Location: Texas
Member Is Offline
Mood: No Mood
|
|
Improvised Sodium Hydroxide
I was interested in the old process of making lye since there seems to be a shortage or Red Devil Lye where i am - no clue why - i guess just another
freedom slowly being taken away - has anyone tried the process of pouring a hot sodium carbonate into dry quicklime? any worthwhile yields? I read
about the process in Ullman's so im guessing there is some fact to it. The mixture of calcium and sodium salts would be a mess though. I was
thinking maybe after i could extract the newly made lye with alcohol leaving the calcium behind - anyone have any experiences with this? on another
point - I hear that red devil has a ton of impurities - maybe alcohol could be a way to purify it - maybe not - mercury salts and plenty of other
toxins are alcohol soluble - just a few ideas i have been thinking about 
|
|
|
Magpie
lab constructor
Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline
Mood: Chemistry: the subtle science.
|
|
Do a search on Red Devil lye. Polverone has recently discovered why it will no longer be available.
The single most important condition for a successful synthesis is good mixing - Nicodem
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Er, well the old method of potash lye is to line a barrel with straw, lime and ashes, in that order. Percolate water through the pile and K2CO3 in
the ash dissolves, reacts as Ca(OH)2 + K2CO3 = CaCO3 + 2KOH, which proceeds because Ca(OH)2 is uniquely semisoluble while CaCO3 is rather insoluble.
(Most hydroxides are less soluble than the carbonates.)
Of course the reaction works in aqueous solution and with any other alkali ion (sodium, etc.). Boil the remaining solution to yield solid lye.
Tim
|
|
|
jimmyboy
Hazard to Others
Posts: 235
Registered: 1-3-2004
Location: Texas
Member Is Offline
Mood: No Mood
|
|
red devil gone - liability issues surrounding meth production?? thats just stupid as all getout - what next - no glassware because we might mix
something? well i guess its time to figure out a good electrocell and make it at home or go this way - whichever is cheapest 
|
|
|
prole
Hazard to Self
Posts: 94
Registered: 4-8-2005
Member Is Offline
Mood: No Mood
|
|
If they haven't yet banned NaCl, then perhaps a little electrolysis is in order:
2NaCl + 2H2O---------------->H2 + Cl2 + 2NaOH
-prole
|
|
|
Marvin
National Hazard
Posts: 995
Registered: 13-10-2002
Member Is Offline
Mood: No Mood
|
|
Making NaOH is a lot easier than making KOH by the lime process.
Lots of boiling in steel containers and several successive stages with fresh portions of lime are advised. Dilute solutions help also. Then
presumably you need to crystalise out the remaining sodium or potassium carbonate from the hydroxide.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
You can also make it my fire but I'm not sure how effective it is, and it probably involves carbon dioxide anyway.
I'm thinking 2NaCl + H2O <==> Na2O + 2HCl. It definetly happens when you have molten salt and/or vapor, you can smell the acidity in the
exhaust fumes. But I don't know if the product is Na2O, NaOH or Na2CO3 (produced by the equilibrium 2NaCl + CO2 + H2O <==> Na2CO3 + 2HCl,
driven by CO2 exhaust gasses).
Another reason it may proceed is due to silicate formation, sequestering the soda in a glass melt.
Tim
|
|
|
jimmyboy
Hazard to Others
Posts: 235
Registered: 1-3-2004
Location: Texas
Member Is Offline
Mood: No Mood
|
|
molten salt ? - a little much for lye huh? looking for cheap and simple - i also saw a patent for the Lowig process - combing ferric oxide and sodium
carbonate at high temp forming ferrate which will break down to hydroxide - way too much hassle
|
|
|
BromicAcid
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
Making potassium hydroxide should be a somewhat simple affair. Simply electrolyze a concentrated solution of KCl with the anode right at the surface
so only a little chlorine gets absorbed then after electrolysis stops giving chlorine heat to dryness, this removes the water and decomposes any
hypochlorite and some chlorate. Then extract with methanol or ethanol, KOH being somewhat soluble in these, finally distill off the alcohol and
you've got somewhat pure KOH, this would work for sodium too but NaOH is less soluble in alcohol. At least it sounds good on paper.
|
|
|
jimmyboy
Hazard to Others
Posts: 235
Registered: 1-3-2004
Location: Texas
Member Is Offline
Mood: No Mood
|
|
hmm isnt chlorine heavier than air and it would sink into the water? not sure - i was reading about them using asbestos diaphragms for the
electrolytic cells - was gonna try that route (well after the boiling quicklime/soda combination) - i dont think anyone sells asbestos anymore maybe a
water permeable ceramic instead anything to keep the chlorine away from the cathode side of the cell
[Edited on 4-11-2005 by jimmyboy]
[Edited on 4-11-2005 by jimmyboy]
|
|
|
Marvin
National Hazard
Posts: 995
Registered: 13-10-2002
Member Is Offline
Mood: No Mood
|
|
What Tim is describing sounds more like the Solvay process for making sodium carbonate and calcium chloride out of calcium carbonate and sodium
chloride. A solution of sodium chloride in ammonia and water has carbon dioxide blown in causing precipitation of sodium bicarbonate.
I don't see why you need methanol to seperate the results of the electrolysis and I can see disadvantages to adding a flammable solvent to
something that could contain chlorates... Just fractionally crystalise from water.
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
Long ago I have produced NaOH from Ca(OH)2 and Na2CO3. It was an experiment in my chemistry set!
Na2CO3 was dissolved in water as a dilute solution, Ca(OH)2 was added (a bit more than theoretically necessary) and agitated (shaken, stirred etc.)
for some time.
Then it was filtered from the CaCO3.
If the NaOH solution is found to attack the filter paper too much, decantation with washing of the precipitate with fresh water is the method of
choice.
The NaOH solution is tested for presence of residual carbonate by taking a small sample and acidifying it. If CO2 is produced, then it either
wasn't agitated long enough or not enough Ca(OH)2 was used.
The process is very easy and clean as long as the Na2CO3 solution is dilute. With concentrated solution, it becomes a thick sludge with the Ca(OH)2,
making filtration or decantation impossible.
The more dilute the solution is the easier the process will be, but at the price of a more dilute NaOH solution.
Ca(OH)2 will be the only impurity in the NaOH solution, but it contains less than a gram per liter of Ca(OH)2. The NaOH solution is pure enough for
all purposes.
Boiling down the NaOH solution (if it has to be done at all) must be done in an iron or nickel dish, glass is attacked considerably, introducing
impurities into the NaOH.
This process is also very useful for preparation of KOH from K2CO3.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
| Quote: | Originally posted by Marvin
What Tim is describing sounds more like the Solvay process |
Nah, I mean actually displacing chlorine with oxygen (or another anion, especially one which forms a solid rather than liquid or gaseous phase
compound). Steam is probably bigger in the process (thus producing hydroxide rather than oxide) than oxygen because there's no Cl2 odor, only
HCl, and there's very little free O2 in a well-tuned propane burner such as I use for most of my heating. But O is more electronegative, so it
should technically work!
Tim
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
More on the carbonate method, would porcelain evaporating dishes be OK to concentrate KOH filtrate? I boiled 30% KOH solution once in a
"stainless" pot and it began to oxidize the metal! My KOH was contaminated with FeOH3 but the metal was attacked only superficially. MAybe
this was a local hot spot.
Fellow molecular manipulator
|
|
|
Magius
Harmless
Posts: 20
Registered: 9-6-2004
Location: Green Bay
Member Is Offline
Mood: Constrianed
|
|
Practacality of Na2CO3 + Ca(OH)2
Tim, could you give us some more specific's on the Ca(OH)2 + Na2CO3 method? How much dillution did you have to use for the solution of the
Calcium Hydroxide? IIRC, isnt it only sparingly soluble even at high temperature?
Oh, and is the CaCO3 solid easier to filter then the infamously difficult CaSO4?
Wait for it...
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
I haven't tried it, BTW.
CaSO4 is easy to precipitate if you do it slowly and possibly when warm. CaCO3 is less soluble so probably finer and harder to filter but that
depends what kind of apparatus you have to do it.
In the old days they just heaped lime and ashes into a barrel and let it percolate, can't take too long but Ca(OH)2 certainly isn't
massively soluble.
Tim
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
The Ca(OH)2 doesn't dissolve, of course!
Just a tiny bit dissolves, this bit reacts with the Na2CO3, precipitating CaCO3 and allowing another tiny bit of Ca(OH)2 to dissolve, and so on.
I don't know what kind of dilution I used, I didn't have a scale back then.
CaCO3 is definately easier to filter than CaSO4, but you have to boil the solution in order for the CaCO3 to agglomerate into larger particles.
Otherwise filtration will be just as difficult as with CaSO4.
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
My KOH experience
Well to save money on incredible UPS charges I went ahead with the KOH production from gardening hydrated lime and technical Potassium Carbonate. I
was operating at little less than 1 Mole per liter so it is reasonably dilute. I mixed vigorously at 49 degrees C for 3 hours and the
"milk" eventually threw down the CaCO3 as a concrete like mass. I filtered the KOH with no attack of the filter paper to get a perfectly
pure filtrate. The filtrate is very slippery to the fingers and taste VERY ACRID I am now using a stainless bowl that I bought at the Salvation Army for $0.49! I am using low heat so not to oxidize the metal. I did not take
any specific gravity measurements but if I reduce the fluid volumeto one third I should have a KOH solution at a concentration of 30ish percent. If
KOH concentration increases I know potassium carbonate solubility decreases but does anyone know if this evaporationg would deposit any Ca(OH)2 that
might be in solution? It looks about as clear as distilled water now, but you never know until you do it I guess.
[Edited on 11/12/2005 by chloric1]
Fellow molecular manipulator
|
|
|
neutrino
International Hazard
Posts: 1583
Registered: 20-8-2004
Location: USA
Member Is Offline
Mood: oscillating
|
|
If it really is as basic as your claim, any calcium hydroxide should have already been forced out of solution due to the common ion effect.
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
Well that is my thinking. It is DEFINATELY basic. my 800 to 900 ml of solution should yield anywhere from 104 to 108 grams of KOH. NOt bad for $0.30
worth of chemicals!
Fellow molecular manipulator
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
Check if the reaction was complete by taking a small sample of your KOH solution and acidifying with HCl. If it doesn't fizz, the reaction was
complete.
Garden Ca(OH)2 is often impure with CaCO3 from reaction with aerial CO2. Use an excess.
And keep your KOH solution in a tightly closed bottle, or it will turn into K2CO3 solution again because of CO2 uptake.
[Edited on 12-11-2005 by garage chemist]
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
I have a beaker cover for my 1000ml beaker. I got the solution covered now. But, while it was evapoarting for about four hours it took
impurities(chromate?) from the stainless as I had feared. THe KOH is piss yellow. Don't want to extract with expensive alcohol. I did a flash
evaporation in glazed porcelain dish witha MAPP gas torch and I did not notice any corrosive effects. I need to repeat the experiment and evaporate
in porcelain dishes. I just wish I could find one for more than 125 ml.
P.S. The lime is stated to be only 95% pure. I used a 5 to 10 % excess
[Edited on 11/12/2005 by chloric1]
[Edited on 11/12/2005 by chloric1]
Fellow molecular manipulator
|
|
|
jimmyboy
Hazard to Others
Posts: 235
Registered: 1-3-2004
Location: Texas
Member Is Offline
Mood: No Mood
|
|
yet another good way - but requires alot of heat is just decompose sodium carbonate -- it will form molten sodium oxide then just put in water to make
the hydroxide .. probably alot of energy though - need a good furnace to reach a high temp
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
As I recall, sodium carbonate decomposes *very* slowly, if at all.
For instance, smelting sodium or potassium by using charcoal and a melt involves the molten carbonate, not the oxide.
Tim
|
|
|
franklyn
International Hazard
Posts: 3026
Registered: 30-5-2006
Location: Da Big Apple
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by jimmyboy
yet another good way - but requires alot of heat is just decompose sodium carbonate -- |
Old tricks are the best tricks
Quicklime CaO is produced by calcinating lime CaCO3 to drive out the CO2
In the same way Soda Ash , Sodium Carbonate Na2CO3 , Washing Soda or
just plain old laundry detergent , Duh , roasted on a stove will give you
Caustic Soda Na2O , hydrating this gives you lye NaOH.
See section 5 reactivity ->
http://www.kencro.ca/PDF/SODASH.pdf
See section 10 stability ->
http://www.commercialaquaticsupplies.com/Templates/msds%20fi...
Because cooking is expensive this is no longer an industrial process.
See quoted below ->
http://en.wikipedia.org/wiki/Sodium_hydroxide
"An older method for sodium hydroxide production was the LeBlanc process,
which produced sodium carbonate, followed by roasting to create carbon dioxide
and sodium oxide. This method is no longer used, but it helped to establish
sodium hydroxide as an important commodity chemical."
Pure Sodium Hydroxide is the by product of the electrolysis of brine to produce
Chlorine. Enviornmental consequences of chlorinated hydrocarbons reduced the
demand for elemental Chlorine and consequently this has impacted the cheap
availability of lye. See ->
http://digicoll.library.wisc.edu/cgi-bin/EcoNatRes/EcoNatRes...
.
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
