BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
Sodium hydroxide from baking soda
Hello everyone, first post!
I wonder if anyone has made NaOH from baking soda. Heat (a few hundred C?) makes NaHCO3 give off the CO2. But some of the NaOH may react with NaHCO3,
making some sodium carbonate. I've tested this by heating some baking soda with a propane torch. Sure enough, ph paper was very dark blue, and the
product gave that "slimey" feel on my fingers.
Alternatively, this post is also a question about purifying NaOH which has "turned" on air exposure, contaminating it with carbonate.
Supposedly, aqueous NaOH/Na2CO3 mixed with calcium hydroxide will change the sodium carbonate back to hydroxide and precipitate chalk. (I was thinking
of starting with chalk to make the CaOH).
It seems to me that recrystalizing NaOH from any of this will also deposit the small quantities of carbonate left in solution, since it would be
saturated with carbonate.
Anyone have any thoughts on this?
[Edited on 9/18/2006 by BeerChloride]
|
|
|
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
Thats probably just Na2CO3. The best way is get that and react with Ca(OH)2 to ppt chalk and get NaOH. Make sure CO2 doesn't get in so cover up the
container while the reaction proceeds.
|
|
|
jimmyboy
Hazard to Others
Posts: 235
Registered: 1-3-2004
Location: Texas
Member Is Offline
Mood: No Mood
|
|
there is already a thread on this under sodium hydroxide..
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
I tried to search for such a post and couldn't find it. ??
|
|
|
chromium
Hazard to Others
Posts: 284
Registered: 27-6-2005
Member Is Offline
Mood: reactive
|
|
Sodium carbonate does not decompose by heating. (at least if temperature is less than 1000C)
Preparation of NaOH:
https://www.sciencemadness.org/talk/viewthread.php?tid=4730#...
[Edited on 18-9-2006 by chromium]
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
Thanks for that link chromium, and you're right - you need >1000 C to directly decompose Na2CO3. But bicarbonate decomposes at only 270 C. In the
thread you gave I think garagechemist had it right.
I know sodium hydroxide might not interest many here, but what I like is the USP grades and cheapness of baking soda and lime. Incidentally I just got
a pound of pickling lime (Ca(OH)2) for $2.50. USP!! ??
I'm going to do a few little tests....
|
|
|
Odyssèus
Hazard to Self
Posts: 50
Registered: 7-3-2006
Member Is Offline
Mood: No Mood
|
|
Bicarbonate decomposes to carbonate.
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
How about, at least as a 'polishing' step after Na2CO3 + Ca(OH)2
Na2CO3 + Ba(OH)2 => 2 NaOH + BaCO3
Barium carbonate is 1/2 to 1/4 as soluble as CaCO3, while Ba(OH)2 is much more soluble than Ca(OH)2
Then
BaCO3 + H2O (as a stream of superheated steam) => Ba(OH)2 + CO2
The flowing stream of gas carries away the CO2 as it is produced, driving the reaction to the right.
Note that you can use NaHCO3 and Ca(OH)2 (or Ba(OH)2) in water and boil the mix to drive off CO2, converting the somewhat soluble alkaline earth
bicarbonate to the much less soluble carbonate. This isn't a brief 5 minute boil, should be much longer time; the steam boiling off helps carry away
the CO2 and concentrates the NBaOH solution. Don't run in glass or silicate ceramic, use stainless steel.
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
Ok, I set up some tests to measure the CO2 given off by acidifying a sample of the cooked sodium bicarbonate. First I tested the procedure with
uncooked bicarbonate, and I can measure down to a few milligrams of CO2. My hypothesis was that if NaOH is an intermediate in the cooking of NaHCO3
into Na2CO3, then incomplete conversion might leave some amount of NaOH.
I found the cooked product to be 98% Na2CO3 - the experimental error was at least a few percent, so I am convinced that cooking gives sodiuim
carbonate (like Odysseus said..). So for the record:
Heat (>270 C):
2NaHCO3 --------> Na2CO3 +
H2O + CO2
So I'm thinking of allowing a good settling to occur and decanting, rather than messing with filtration.
What I'm realizing, though, is that even if I get a reasonably pure NaOH solution, but the reality of going from THAT to some type of solid is
actually not so pleasant (fighting carbonate re-formation, boiling down, scraping crust, NaOH damage to glass, etc...). For some reason I don't like
boiling away liquids. 
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
Part 2: the calcium hydroxide conversion to NaOH
I spent several hours doing solubility product ion system calculations, and had to take an aspirin, but the results are interesting. First, I
calculated the ion concentrations resulting from saturated NaOH/Na2CO3. (PURIFYING NaOH FROM CARBONATE CONTAMINATION). At room temp,
Ksp(NaOH) = 770.1 = [Na][OH]
Ksp(Na2CO3) = 90.7 = [Na]^2 [CO3]
[Na]=[OH] + 2[CO3] (charge balance)
This system gives the solute contents to be
NaOH: 98.9%
Na2CO3: 1.11%
Then, I calculated the ion concentrations for all four species of ion Na+, OH-, Ca+2, CO3-2, assuming saturation for calcium hydroxide, calcium
carbonate, and sodium hydroxide. IN OTHER WORDS, you add calcium hydroxide to aqueous sodium carbonate - how much carbonate and calcium exist in
solution?
I wonder if I did this correctly. Here's the equations:
NaOH: 770.1=[Na][OH]
CaCO3: 3.4E-9=[Ca][CO3]
Ca(OH)2: 8E-6=[Ca][OH]^2
[Na]+2[Ca] = 2[CO3]+[OH]
The results of these were:
[OH]=27.44 M
[Na]=28.06 M
[Ca]=1.06E-8 M
[CO3]=0.321 M
CONCLUSIONS: Purifying NaOH/Na2CO3 by forming solution at room temp gives a 98.9% NaOH, 1.1% Na2CO3 solution. But making it from carbonate using
Ca(OH)2 gives a solution of only 97.0% NaOH, and still 3.0% carbonate. Not all the carbonate is dragged into precipitated chalk, because of the
significant solubility of sodium carbonate. Also, calcium levels are way too low to even worry about as an impurity. (Note: in reality, suspended
particles of calcium compounds will be a primary factor there).
I hope someone could verify these things / check my calculations? (I know, I probably wouldn't want to either..)
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
I believe that an excess of Ca(OH)2 would be used, to help force the reaction in the direction of NaOH.
I do know that process engineers did a lot of work on this method, balancing purity, convertion efficiency, and work-up costs (more dilute solutions
give better conversion and purity but cost more to evaporate). I suspect you're not alone in your headache.
You are correct, suspended particles will be important, which is why filtration is needed at some point.
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
A search some time ago showed that some of the most detailed information available to me on this is in two IEC articles: 33, 204 (1941) and 48, 408
(1956).
Attachment: naoh_via_lime.zip (2.6MB)
This file has been downloaded 664 times
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
Cool - thaks for those papers, SCWhack. I'll take a look at them.
Mr. not_important: the calculations I did were based on excess Ca(OH)2, and it doesn't really matter how much, as long as it's saturated which it will
be for any solid Ca(OH)2 present. The same conditions exist for CaCO3, and I'm ASSUMING saturated condition for NaOH. You say dilute solutions give
purer NaOH? That seems plausible. As far as filtering, I'm betting that a cool solution allowed to really settle will yield a clearer solution than
filtering (at least for any equipment I have), or at least be a worthy trade-off.
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
All I can tell you is that every chem book that gives any detail on NaOH from Na2CO3 and Ca(OH) states that a definate excess of Ca(OH)2 is used to
force the reaction. They also talk about the trade-off between running more dilute for higher purity and more concentrated for better economics.
When fancy filter grear wasn't available, old time lab textbooks often combined settling with filtering, decanting through a filter whose purpose it
to capture the small amount of solids brought along. As the filter wasn't handling much solids, it could be fairly fine.
(Mr. ?)
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
No offense intended (Mr.). I just feel strange calling you "not_important"! Anyway, you're right, you need excess Ca(OH)2, but actually the papers
provided by SCWhack confirm that it doesn't really matter to what degree in excess. Those papers also emphasize not filtering, but of course that's
talking 1950's. You're also right about the trade-off between purity and concentration.
I did a calculation implied in one of the papers for the percent purity. It's actually approximately linear (theoretically), and agrees with my
previous calculation above. At room temp, the line is defined by the two points [NaOH]=0, %NaOH=100% and [NaOH]=27 M, %NaOH=97%. The impurity is
sodium carbonate. For temperature above 75 C, the point at [NaOH]=27 M has %NaOH=95.6%. I think this is right, but I need to go back over those papers
because there were some conditions which dropped the purity down to 90% or so. So, you could say the purity is directly proportional to how dilute
your final solution is.
I'm still trying to think of easy, simple ways to evaporate or boil off that solution!
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
When I made KOH, couldn't get that but had K2CO3, I used stainless steel pots for evaporation. Find one with a lid that fits fairly well, and add a
connector for attaching to a vacuum source. For those with just a single screw holding the handle, removing that and enlarging the hole worked well.
Used a grommet and a pair (inside/outside) of washers to get a seal. If the lid doesn't fit very well, you can make a gasket from either soft solid
silicone or closed-cell foam, or try faking a gasket with a bead of silicone rubber - let it cure without the lid in place.
Put the hydroxide solution in the pot, put the lid in place, and pull a vacuum on it. Doesn't need a strong one, I used a recirculating aspirator rig
with an underpowered fountain pump; the reduced pressure is to pull the steam and any leaked-in air (CO2) out of the pot. Heat the pot, if you
condense the steam taken off you can get a pretty good idea of how much concentration has been done.
An improvement is to mount a disk on the inside if the lid so that there is no line-of-sight with the take-off port, this is to help prevent splatters
from the boiling solution from hitting that hardware. Several nested stainless steel meshes might do instead, the tea strainer and kitchen drain
filter types are about the right range of sizes.
This works to get a strongly concentrated solution. For getting to the solid you need something else to finish the job. I used a copper pan that I
silver plated on the inside, you do pick up a little CO2 but if you start with the concentrated solution it goes pretty quickly, adding more of the
concentrated solution as water boiled off. Finish off by taking it to the hydroxide's fusion point and above, pour into metals molds to form sticks or
lumps. Watch out for splatters.
|
|
|
BeerChloride
Harmless
Posts: 47
Registered: 17-9-2006
Location: Alabama, USA
Member Is Offline
Mood: Dunno
|
|
I like your suggestions, thank you. I was thinking about using some sort of vacuum system.
Do you happen to know a reference for reaction constants for sodium hydroxide with various metals?
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
