Sciencemadness Discussion Board - Making Sodium Hydroxide (lye) ??

Making Sodium Hydroxide (lye) ??

ktw_100 - 21-4-2006 at 10:34

Anybody know how to make sodium hydroxide? I am a biodioesel maker, and would like to make it vs buying it! I have heard of it being done with DC voltage and electrodes in salt water, but that's about it.... Or making from wood ash.

Thanks

garage chemist - 21-4-2006 at 11:10

Best way would be from slaked lime (calcium hydroxide) and sodium carbonate solution.

Ca(OH)2 + Na2CO3 -----> 2 NaOH + CaCO3

An excess of Ca(OH)2 is agitated for several hours with a solution of sodium carbonate. It is filtered or decanted from the resulting chalk. There you have your lye solution.
Boiling it down is messy and dangerous, and hot NaOH solution attacks glass containers.

rot - 21-4-2006 at 13:31

Yes, you can make it from salt water:
First, you make a saturated salt solution. (35.9g/100mL).
Then you electrolyse this solution with graphite electrodes.
At the anode, chlorine gas will evolve, At the cathode Sodium Hydroxide and Hydrogen gas will form. you can just keep running electricity through it until no more chlorine gas evolves, then boil the solution down. Care must be taken to prevent the chlorine gas from mixing with the dissolve lye, as this will react to make sodium hypochlorite (bleach).
This can be done by a salt bridge.

mrjeffy321 - 21-4-2006 at 13:49

In Industry, there are suppose to make NaOH electrolytically by passing an electric current through a Sodium Chloride Solution.
Through the use of a special membrane, NaOH solution is allowed to exit the cell and be collected while the other products remain behind.
In a diagram I have in one of my chemistry books, it shows two graphite anodes in an NaCl solution surrounded by a wire screen cathode with an asbestos diaphram in between. NaOH is produced at the cathode and drips out of the cell while the Cl2 gas bubbles out without reacting with the water. This is called a "diaphragm cell".

Another process, called a "Mercury Cell", Hg serves as the cathode, which reduced Na+ ions to Na metal which then dissolves in the Hg. The Hg-Na amalgam is then exposed to pure water (not NaCl solution) and the Na will react to produce NaOH and H2 gas. This procedure is suppose to produce much purer NaOH.

"rot", How could you set up a salt bridge in the electrolytic cell in order to produce NaOH? I never quite understood the idea of a salt bridge any further than it is used to keep two solutions seperate during an electrochemical reaction.

bereal511 - 21-4-2006 at 20:37

Well, the point of the salt bridge is just to keep the chlorine gas evolving in the electrolysis seperate from the sodium hydroxide being produced, which would form sodium hypochlorite.

I've been interested in thermochemically decomposing sodium carbonate using a large fresnel lens to produce sodium oxide, which would of course react with water to form sodium hydroxide. At the maximum focal length, a fresnel lens should be able to reach temperatures of 1500 degrees celcius on a sunny day. I'm not sure how feasible this approach would be, especially because white sodium carbonate would barely absorb the rays of the sun. The best way I can think of using this approach would be to heat a black substrate (graphite? metal oxide?) with the sodium carbonate over it. I'll probably be testing this soon enough when I can get enough money to buy the lens from this webstore: http://www.goldmine-elec-products.com/prodinfo.asp?number=G1...

$50 doesn't seem bad does it? For 40" x 31"?

12AX7 - 21-4-2006 at 20:44

Don't count on it- sodium compounds are volatile, you'll mostly boil it off. Electricity, by resistance, arc, or else flame heat, is a lot easier and a lot less weather-dependent than solar heat.

Tim

rot - 21-4-2006 at 23:07

Quote:

Originally posted by mrjeffy321
"rot", How could you set up a salt bridge in the electrolytic cell in order to produce NaOH? I never quite understood the idea of a salt bridge any further than it is used to keep two solutions seperate during an electrochemical reaction.

I've never understood it completely either, but as far as I know it's just a tube filled with electrolyte (in this case NaCl solution) connecting the two beakers containing NaCl Solution, to prevent the chlorine gas from mixing with the lye.

The_Davster - 21-4-2006 at 23:18

Yeah, thats all they are. I've used them a couple times for electrolysis, huge cell resistance, makes your reaction take much longer compared to an undifided cell.

mrjeffy321 - 22-4-2006 at 00:28

Quote:

At least it is not just me. Seriously, my only "experience" with them is in diagrams inside a text book, in all my practical experience with electrochemical cells, I have never seen one, much less used one. If the "salt bridge" is made of salt, why wouldnt it dissolve?...we assume it is made of the soluble salt which is being used as the electrolyte. Anyway, but that isnt important.

Decomposing Sodium Carbonate into Sodium Oxide using a giant frensel lens would be pretty energy intensive considering how hot you can get the focal point...especially with such a huger lens. You might just vaporize anything in the path of the light.
But if you could indeed get it to work successfully, that would be an excellent method to use since the Na2O would form NaOH in solution, but you might end up decomposing the Sodium Oxide with that intense heat while making it.

chochu3 - 22-4-2006 at 05:46

salt bridge will not dissolve because the solutions in each cell is saturated.

rot - 22-4-2006 at 07:45

Then the bridge will dissolve, because as electrolysis continues the Sodium Chloride concentration drops, so the solution is not saturated anymore, so the bridge will dissolve.

bereal511 - 22-4-2006 at 16:55

Hmmm, I'd like to see if I could at least calculate the direct energy output of the area of the lens. One could always change the range of the focal length to reduce the energy output of the focal point, then calculate the amount of energy absorbed by the heating medium underneath, and find the temperature range of the medium to prevent sodium oxide decomposition/boiling, no? As long as I can stay in the 800 - 1000 degree range, I hope I'll be fine.

I'll probably do some testing next month.

Chris The Great - 22-4-2006 at 21:43

A salt bridge is a tube filled with a salt solution, not a piece of solid salt. Each end is plugged with a permeable membrane that prevents the two solution from mixing but allows the ions to transfer charge.

Poor Mans Salt Bridge

FloridaAlchemist - 24-4-2006 at 06:20

For a temporary salt bridge, use a filter paper (Chromotography) strip soaked in 1M Potassium Nitrate.
Put each reagent and electrode into separate beakers and insert the bridge between them

When the bridge starts to dry out, just add more 1M Potassium Nitrate from a dropper to wet it.

Making Sodium Hydroxide (lye) ??

ktw_100 - 24-4-2006 at 11:29

Thanks for the input....
I would assume that the chlorine gas is toxic, even at low levels? Just how much is generated when making the NaOH? Also, the electrodes I assume could be made from 1.5V battery cell posts - I think they are the right material.

What would be the best DC voltage to run on the cathode and plate, and how long a process is it? Does it need to be really 'flat' DC, or very well filtered?

Anything else I should be aware of?

Cloner - 25-4-2006 at 00:29

you end up electrolysing all the chloride into chlorine, so it's quite a heap. Electrolysis is slow, howver, so if done outside it should be OK. Alternatively, you can lead the chlorine through 'something plentiful and OTC' that reacts with it in a train of 'washing bottles' and end up with a lot less smell.

In electrolysis, I'd say the ideal voltage is the voltage at which you see a nice amount of bubbles. Just be sure that your current is not so big that you fry your power source.

rot - 30-4-2006 at 11:46

The danger of chlorine is not that high, chlorine is detectable by the humon nose in quantities far under the lethal limit, plus the fact that chlorine evolution is very slow. Do it outside and you'll be fine. Don't do it inside though, I had once ran a cell next to my computer (used as power supply

) within about 30 minutes I got a headache even though I couldn't smell any chlorine. I'm sure it was because of the chlorine though. I turned off the cell and opened a window and headache was gone

Making Sodium Hydroxide - again

ktw_100 - 1-5-2006 at 08:11

I THINK I have some Sodium Hydroxide (?)

I saturated some distilled water with rock salt. (35.9g/100ml). Then put equal amounts of this into two separate glass containers. Containers joined with a homemade salt bridge (cloth stuffed inside some good plastic tubing, saturated with salt solution). Applied 12VDC to carbon electrodes placed in the cells.

After a couple of days, I have white crystalline stuff in both containers, as electrolysis taking place. At the cathode is Sodium Hydroxide. What is the crystalline stuff at the anode - Sodium Chloride?

I did measure the pH of the Sodium Hydroxide - seems to be around 9 or so at the moment.

(PS: The salt bridge didn't dry out, or need any resoaking, due to capillary action keeping the cloth inside the tube wet.)

Also, when it comes to de-watering the Sodium Hydroxide, as mentioned, it can be done by distilling the water off. However, Sodium Hydroxide is hygroscopic, is it not? Therefore, I will never get all the water out of it.

I am thinking of making a saturated solution of NaOH/H2O, and using this as part of the process of producing biodiesel, versus using crystalline form of dry NaOH that I have been using.

From what I understand, if NaOH has been exposed to air, the best thing when making biodiesel with it is to increase the amount necessary for the process by 25% (maybe off topic here, but I don't know if I can get these questions answered on the biodiesel newsgroups - perhaps there is someone here making biodiesel?)

Appreciate any info already supplied! This is fun stuff!

12AX7 - 1-5-2006 at 09:16

pH 9 isn't much at all, you need pH up around 12-13 before you have any percentage of NaOH, IIRC.

The crust is most definetly a salt crust. As you electrolyze water, you remove it, and the salt loses solubility. Also, even if it is still soluble in the solution, brine likes to crust up around its container anyway.

Tim

mrjeffy321 - 1-5-2006 at 10:11

I agree with 12AX7, with a pH of only 9, you really have not made very much Sodium Hydroxide.
Even small concentrations of NaOH will yeild very high pH values (on the order to 12, 13, ...14). If you evaporate the solution now, you will get a little bit of NaOH but most of the NaCl still remains.

Just for illustrative purposes, say you had a 500 mL solution with a pH of 9.0, this would mean that your [OH-] concentration (the same as your NaOH concentration) would be roughly 1 E-5 Molar, or about 5 E-6 moles of NaOH.
NaOH has a molar mass of 40.00 g/mol, so you would get (in theory) .0002 grams of NaOH by evaporating off the water.

neutrino - 1-5-2006 at 13:22

Remember that small amounts of NaOH in contact with the air will react rapidly with atmospheric CO2 to form (bi)carbonates, thus neutralizing the hydroxide.

There was a thread on biodiesel around here somewhere. Search around a little.

Making Sodium Hydroxide (lye)??

ktw_100 - 1-5-2006 at 13:38

Thanks for the info ..... I will let the process finish, and try to scrape off the brine from around the edges, and see what I have. I guessed at ph 9 using a wide range indicator. That was yesterday morning, so 24 hours have passed. My liquid solution is about half of what I put in there to begin with. I thought it was evaporation!

so if I get the NaOH out of it as soon as it is done, and then 'dry' it as best I can, I assume I will get a bit of NaOH anyway. If I had put 35.9 grams of salt in the solution, would I get 35.9 g of NaOH (in a perfect world)?

And still wondering about the non-sodium hydroxide side of the process. Is the precipitate also NaOH, or what?

Thanks for the help ...

woelen - 2-5-2006 at 10:39

If you are using a salt bridge, then at the positive side, chloride is converted to chlorine, and chloride ions from the salt bridge are moving out of it, towards the positive electrode.

At the same time, sodium ions move to the other side. The net result is that the salt bridge becomes depleted of ions. Of course, the mechanic capillary effect and diffusion keep it filled with ions, so the liquid remains conductive.

Net result will be that at the anode part, the liquid remains salty and only chlorine gas escapes. At the cathode side, certainly some hydroxide will be formed, but it will be very hard to obtain it in a pure state. You always have NaCl over there as well, at high concentration.

elementcollector1 - 16-1-2013 at 10:29

What about a drying tube, plugged or covered at both ends with filter paper, and filled with a conductive solution of your choice?

AJKOER - 18-1-2013 at 10:08

Here is questionable path to NaOH (or Ca(OH)2) that some may find interesting and others may find insane (appropriate for Sciencemadness perhaps) on many levels. However, if you are short on ingredients, here is not very cost effective method that requires Hypochlorous acid. Note, one path to HOCl would be to add a weak acid (acetic from vinegar, ascorbic from Vitamin C, critic...) to Chlorine Bleach (essentially NaOCl and NaCl and a touch of NaOH) and distill (stop when you have obtained 1/2 of the starting solution which contains most of the HOCl and Cl2O gas). There are other paths as well (search on Sciencemadness and the internet).

Now, to quote from "A comprehensive treatise on inorganic and theoretical chemistry", by Joseph William Mellor, page 256:

"According to F. von Tiesenholt, there is a reversible reaction: NaCl + HOCl = NaOH + Cl2, and with 2 or 3 grams. of anhydrous calcium chloride dissolved in the smallest possible quantity of hypochlorous acid, there is an energetic development of chlorine, and a formation of calcium hydroxide; some chlorate is formed at the same time. According to J. L. Gay Lussac, hypochlorous acid at about 100° attacks metal chlorides with the evolution of chlorine and oxygen, and the formation of chlorates; according to A. W. Williamson and J. Kolb, the metal chlorides are not attacked in the cold, and on warming the chlorates are formed. Hypochlorous acid precipitates the higher oxides from the chlorides of manganese, tin, lead, iron, cobalt, and nickel; and copper oxychloride from cupric chloride. Silver chloride decomposes the acid catalytically."

Link: http://books.google.com/books?ei=hWP3UP78GoLi0gHd14GYCw&...

{EDIT} Interestingly, the following equation is given in a recent reference, confirming Mellor:

Cl2 + -OH <-----> Cl- + HOCl

See "Handbook of Detergents: Production", by Uri Zoller abd Paul Sosis, page 443. Link: http://books.google.com/books?id=dXn3aB1DKk4C&pg=PA443&a...

I would, nevertheless, not be surprised, if one does better making Ca(OH)2 from CaCl2 (as moving the reaction in the direction of expelling chlorine may be easier), and if this is the case, react the Ca(OH)2 with NaCl as detailed above to form NaOH. In either case, one must avoid Iron and heavy metals presence as these are suggested catalysts for the production of chlorates.

With my suspicions noted, one questionable embodiment of the process would be to spray a small amount of HOCl evenly over dry NaCl (maintaining an excess of NaCl so assume a high estimate of the Hypochlorous acid strength) and lightly heat in the open air (best in a fumehood or outside) to vent any toxic and irritating Chlorine fumes.

NaCl + HOCl --> NaOH + Cl2 (g)

However, I questioned whether the reaction can be really stopped at this point. For one, local concentrations may still result in some undesirable side reaction including:

NaOH + HOCl --> NaOCl + H2O

3 NaOCl --> 2 NaCl + NaClO3

NaOCl + HOCl --> NaClO2 + HCl

NaClO2 + HOCl --> NaClO3 + HCl

As such, this probably not a practical or particular easy to execute route, however, the basic starting chemicals are easy to obtain (salt and/or CaCl2, bleach and vinegar). Note, my side reactions listed are not unique to my suggested synthesis, and actually, may be of even greater concern in an electrolysis approach where control over concentration, pH and temperature is perhaps more difficult to reduce disproportionation into chlorate.

[Edited on 19-1-2013 by AJKOER]

elementcollector1 - 18-1-2013 at 10:42

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.

AJKOER - 18-1-2013 at 21:30

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]

S.C. Wack - 19-1-2013 at 11:27

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.

chemicalmixer - 19-1-2013 at 12:51

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]

Manifest - 19-1-2013 at 15:03

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

12AX7 - 19-1-2013 at 15:18

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

S.C. Wack - 19-1-2013 at 15:53

Quote: Originally posted by 12AX7

Hmm. Ferrate(III), aka ferrite. Just to be clear!

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]

AJKOER - 19-1-2013 at 17:28

Quote: Originally posted by S.C. Wack

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]

Random - 20-1-2013 at 04:39

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.

12AX7 - 20-1-2013 at 08:06

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

AJKOER - 21-1-2013 at 18:12

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]

elementcollector1 - 22-1-2013 at 09:45

Quote: Originally posted by Manifest

If he doesn't have Sodium Hydroxide, I highly doubt that he has DH2O

Distilled water? I can get that from the local Safeway for 99 cents. I hope you aren't thinking of deuterium oxide...

chemicalmixer - 22-1-2013 at 15:06

Quote: Originally posted by Manifest

If he doesn't have Sodium Hydroxide, I highly doubt that he has DH2O

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.

AJKOER - 23-1-2013 at 05:21

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)

LanthanumK - 23-1-2013 at 10:01

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.

elementcollector1 - 23-1-2013 at 10:04

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.

Manifest - 25-1-2013 at 03:25

Quote: Originally posted by elementcollector1

Quote: Originally posted by Manifest

If he doesn't have Sodium Hydroxide, I highly doubt that he has DH2O

Distilled water? I can get that from the local Safeway for 99 cents. I hope you aren't thinking of deuterium oxide...

LOL. Sorry, it was late at night when I posted that....

Manifest - 25-1-2013 at 03:35

Quote: Originally posted by chemicalmixer

Quote: Originally posted by Manifest

If he doesn't have Sodium Hydroxide, I highly doubt that he has DH2O

I'm sorry. Again, it was late at night....

science_guy1 - 4-4-2013 at 13:08

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]

elementcollector1 - 4-4-2013 at 14:43

Oh, for science's sake...
You're going to need a temperature much higher than what a propane torch can provide.

AJKOER - 9-4-2013 at 18:51

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]

ElectroWin - 17-6-2013 at 09:11

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.

annaandherdad - 18-6-2013 at 08:38

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.

Random - 18-6-2013 at 15:07

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?

blogfast25 - 19-6-2013 at 03:47

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.

ScienceSquirrel - 19-6-2013 at 04:07

Your limiting reactant is calcium hydroxide which is not very soluble at all.
Getting solid sodium hydroxide is the real challenge as Blogfast says.
It never crystallises, it just turns from being a very strong solution to being a melt of sodium hydroxide that contains water.
This stuff is truly evil, as it can spatter molten droplets that will cause horrendous burns.
You can buy it for about EUR 10 for a kilogram, 99.9% pure for soap making, etc.

ElectroWin - 20-6-2013 at 18:29

but if you boil NaHCO3 solution, it loses half of its CO2 giving Na2CO3 solution. then add the lime, let settle, decant, and continue boiling down the liquors until concentrated

[Edited on 2013-6-21 by ElectroWin]

blogfast25 - 21-6-2013 at 04:01

Bear in mind that NaOH attracts CO2: if you're not careful you end up with what you started from!

annaandherdad - 21-6-2013 at 10:46

I'm wondering about what science_guy1 said, about decomposing Na2CO3 into Na2O+CO2 at high temperature. Wikipedia says Na2CO3 melts at 851C and boils at 1633C. It doesn't say anything about decomposition.

Nevertheless, I've always assumed that part of what is in wood ashes is Na2O or K2O, and that that gives lye when treated with water. It certainly wouldn't make lye if it were the carbonate.

Again according to wikipedia, CaCO3 produces a vapor pressure of CO2 of one atmosphere at about 900C. I wonder what the number is for Na2CO3.

blogfast25 - 21-6-2013 at 11:27

Quote: Originally posted by annaandherdad

I'm wondering about what science_guy1 said, about decomposing Na2CO3 into Na2O+CO2 at high temperature. Wikipedia says Na2CO3 melts at 851C and boils at 1633C. It doesn't say anything about decomposition.

Nevertheless, I've always assumed that part of what is in wood ashes is Na2O or K2O, and that that gives lye when treated with water. It certainly wouldn't make lye if it were the carbonate.

Again according to wikipedia, CaCO3 produces a vapor pressure of CO2 of one atmosphere at about 900C. I wonder what the number is for Na2CO3.

It's not clear what Na2CO3 decomposes to at extremely high temperatures. But that doesn't appear to be a very practical way of preparing Na2O, unlike CaO that is obtained on a rather grand scale from 'burning' limestone.

What you find in wood ashes is K2CO3, aka 'pot ash' (or 'potash'), from which the name potassium is derived. It's only a couple of percent (or so) but it was once a prized source of potassium before potassium salts began to be mined. Wood ash would be collected in pots, the potassium carbonate then leached out of the ash with hot water.

[Edited on 21-6-2013 by blogfast25]

Dr.Bob - 21-6-2013 at 13:37

"What you find in wood ashes is K2CO3, aka 'pot ash' (or 'potash'), from which the name potassium is derived. It's only a couple of percent (or so) but it was once a prized source of potassium before potassium salts began to be mined. Wood ash would be collected in pots, the potassium carbonate then leached out of the ash with hot water."

That is true, but wood being a complex mixture and fire being a vague process, you actually get a complex mixture of K2CO3, Na2CO3, KOH, NaOH, NaCl, KCl, and much more in real life. That is a plenty basic enough mixture to make some simple soap with grease if heated long enough, but it would not work as well for making biodiesel or doing much real chemistry.

And as others have pointed out, getting a soln. of NaOH into a solid, dry form is not easy. So while electrolysis and other methods are fine for experimental purposes, and even small scale reagent synthesis, they are tough to use for preparative amounts of chemicals, especially if needed in commercial amounts, like for making biodiesel, which requires as dry a source of base as possible, preferably NaOMe, but NaOH will do if you don't mind using more of it, and lower yields.

annaandherdad - 21-6-2013 at 14:15

Thanks, Bob, in the meantime I looked up the thermal decomposition of Na2CO3. It seems that you only get some decomposition into Na2O + CO2 after the carbonate is melted, and then only a small percentage of Na2O appears in the melt. That's because the Na2O vaporizes, or comes off as Na + O2, as it is produced. The bottom line seems to be that thermal decomposition of sodium carbonate is not a good way to make Na2O, no matter how hot you are willing to go.

Reference:
The Thermal Decomposition.of Sodium Carbonate by the Effusion Method

Ketil Motzfeldt
J. Phys. Chem., 1955, 59 (2), pp 139–147

blogfast25 - 22-6-2013 at 05:21

Quote: Originally posted by Dr.Bob

That is true, but wood being a complex mixture and fire being a vague process, you actually get a complex mixture of K2CO3, Na2CO3, KOH, NaOH, NaCl, KCl, and much more in real life. That is a plenty basic enough mixture to make some simple soap with grease if heated long enough, but it would not work as well for making biodiesel or doing much real chemistry.

KOH and NaOH? I doubt that VERY much. They would not survive the CO2 rich environment in which they are generated. Na2CO3 is probably a minority constituent. NaCl and KCl as contaminants only.

Fire is not really a 'vague' process: what so 'vague' about complete incineration?

Biodiesel (transesterification of triglycerides) requires KOH or NaOH: carbonates aren't alkaline enough.

ElectroWin - 22-6-2013 at 07:03

Quote: Originally posted by blogfast25

KOH and NaOH? I doubt that VERY much. They would not survive the CO2 rich environment in which they are generated. Na2CO3 is probably a minority constituent. NaCl and KCl as contaminants only.

yes, Na2CO3 and K2CO3 appear in wood ashes as minor constituents; when fresh, though, much wood ash has plenty of CaO, which is effective at stealing the carbonate when you leach it

blogfast25 - 22-6-2013 at 07:15

Quote: Originally posted by ElectroWin

yes, Na2CO3 and K2CO3 appear in wood ashes as minor constituents; when fresh, though, much wood ash has plenty of CaO, which is effective at stealing the carbonate when you leach it

I was recently involved in a minor research project with the aim of reducing the alkalinity of paper/paper pulp/cardboard fly ash for safe and cheap disposal.

I found more than 30 % of CaCO3 in the ash and yes, some CaO (by then Ca(OH)2, of course) in the water leachate of said ash. But the amount was very small and appeared to be due to too high incineration temperature. I doubt very much if wood ash obtained from 'normal' wood fires contains much quick lime at all: CaCO3 only starts losing CO2 above 800 C and it's a slow process at that temperature. And in paper/cardboard, cheap CaCO3 is likely used as a cheap filler, so there's more of it.

[Edited on 22-6-2013 by blogfast25]

ElectroWin - 8-7-2013 at 14:43

yes, laser-printer paper especially has CaCO3 added, as a filler. but i got good results from leaching paper i had burned in a wood furnace. the leachate had a strong soapy feel to it and made decent lye.

but i also sealed the ashes from moisture and atmosphere after burning, to help keep them fresh. When fresh, they are white. after moisture gets to them, they turn a dull gray.

as for burning temperatures and time, the paper i burned got to a red-orange to orange hot. 800 C is only about a cherry red.

[Edited on 2013-7-08 by ElectroWin]

bfesser - 8-7-2013 at 16:41

Interesting. I was under the impression that <a href="http://en.wikipedia.org/wiki/Bentonite" target="_blank">bentonite</a> <img src="../scipics/_wiki.png" />, rather than CaCO3, was used as a filler in paper. I've never observed fizzing, as would be expected with CaCO3, from wetting of copy paper with acidic solutions:

CaCO3(s) + 2 H3O+(aq) → Ca2+(aq) + 3 H2O(l) + CO2(g)

After burning, I would think the bentonite would leave various K, Na, Ca, Al oxides, hydroxides, and carbonates. I'm struggling to find a reference for the use of bentonite in paper, however. My searches keep turning up sellers of bentonite for use in papermaking, but no explanation.

[edit]
Evidently, either or both are used along with TiO2. So now you've got Na, Al, K, Ca, Ti and who knows what else?<a href="http://www.paperonweb.com/A1056.htm" target="_blank">

Quote:

Inorganic portion consisting of mainly filling and loading material such as calcium carbonate, clay, titanium oxide etc may be 0 - 30% of paper. <img src="../scipics/_ext.png" />

</a>It's not a pretty website, but the claim is within reason.

[Edited on 7/12/13 by bfesser]

ElectroWin - 11-7-2013 at 14:37

ok! bubbles, or the lack, thereof, in acid solution, would seem to be definitive.

it frustrates me to be wrong about what the filler material is, but i stand by my observations, that (1) fresh ashes are white; (2) there is an exothermic reaction upon adding the ashes to water; and (3) the leachate gets a soapy feel to it and makes good lye.

[Edited on 2013-7-11 by ElectroWin]

Random - 11-7-2013 at 15:48

I think it depends on how hot the fire was. Needless to say I have dissolved aluminium in wood ash/Ca(OH)2 mixture in water. There was big excess of both compounds for their water solubility. Almost like a paste.

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: removed unnecessary quote(s)]

[Edited on 7/12/13 by bfesser]

bfesser - 11-7-2013 at 20:26

ElectroWin, I never said you were wrong. Calcium carbonate is a filler component, but it may not be the only component. In fact, this composition would seem to support your observations. I have no doubt that burning copy paper produces plentiful CaO, Na2O/NaOH, KOH, etc.

blogfast25 - 12-7-2013 at 00:17

Quote: Originally posted by bfesser

I have no doubt that burning copy paper produces plentiful CaO, Na2O/NaOH, KOH, etc.

A recent analysis of fly ash from paper/cardboard incineration for a client showed it to contain only small amounts of CaO. Ordinary burning of copying paper doesn't generate enough heat to convert CaCO₃ to CaO, I think. Never mind Na₂O...

[Edited on 12-7-2013 by blogfast25]

bfesser - 12-7-2013 at 09:27

Thanks, blogfast25. Any chance you could share the analysis—if it isn't confidential or proprietary? Even a heavily <a href="http://en.wikipedia.org/wiki/Sanitization_(classified_information)" target="_blank">redacted</a> <img src="../scipics/_wiki.png" /> copy could prove insightful

ElectroWin - 12-7-2013 at 09:44

also it isnt the fly ash but the bottom ash that i collected and processed

experimenter_ - 28-4-2016 at 02:14

I have tried to electrolyse concentrated NaHCO3 solution on carbon electrodes. It gave the typical H2/O2 mix (it "bangs" with flame), not CO2 (or very few). Also the pH in the solution did not rise at all after few hours of electrolysis. The electrodes corroded heavily.

- Probably Na2CO3 or K2CO3 is needed for this to work.
- The bubbling CO2 might react with the NaOH and convert it back to carbonate. So maybe cell separation is required.

XeonTheMGPony - 28-4-2016 at 06:49

Quote: Originally posted by experimenter_

I have tried to electrolyse concentrated NaHCO3 solution on carbon electrodes. It gave the typical H2/O2 mix (it "bangs" with flame), not CO2 (or very few). Also the pH in the solution did not rise at all after few hours of electrolysis. The electrodes corroded heavily.

- Probably Na2CO3 or K2CO3 is needed for this to work.
- The bubbling CO2 might react with the NaOH and convert it back to carbonate. So maybe cell separation is required.

If your goal is to make sodium hydroxide go the mercury route IMO, I did this, only I reacted the Cl2 and the H2 to make hydrochloric acid while I was at it.

I been thinking of modifying the cell some how to recover pure sodium as as well, will post one of these days on it if I suceed

experimenter_ - 29-4-2016 at 05:32

Quote: Originally posted by XeonTheMGPony

If your goal is to make sodium hydroxide go the mercury route IMO, I did this, only I reacted the Cl2 and the H2 to make hydrochloric acid while I was at it.

Cool. How did you combine the gases and what were the results? I have been trying something similar:
http://www.sciencemadness.org/talk/viewthread.php?tid=65883

My interest in the NaHCO3 electrolysis experiment was to find a way to capture atmospheric CO2. I would not be interested in the lye produced except that it would capture CO2 from the air.

I'm not interested in the mercury way also since I do the experiments only for amusement. I have tried gallium but it seems that it cannot form a liquid alloy with sodium. Maybe lithium electrolysis with the Ga cathode has more chances of working.

XeonTheMGPony - 29-4-2016 at 06:34

The way I did it was very dangerous back then! I used simple water flash back arresters, and a simple premixing tube that lead into a big jar and spark plug, then that ran to ice water.

the spark was started then the cell, it seemed to work but again back then it was mucking about and no way to test the hcl concentration.

The cell its self how ever worked exceedingly well, good strong hydroxide! Just ran it till I couldn't make hydrogen any more.

Your UV issues is you're using UV-B, You need UV-C Air cleaner UV bulbs make this, or if you buy a "Plasma cell Ozone generator" for hot tubs it is just a UV-C based generator, care fully dismantle and repurpos it.

UV-C is to be considered a serious health risk, shield body and eyes from radiation and any thing you wish to live for that matter, and it will produce ozone in the immediate area.

https://www.youtube.com/watch?v=MtygiCwnEzw < Hydrogen flame in Chlorine

I am sure now days I could do better, but so you can surely do even better! I used a spark plug but with some thinking you can surely find a better method. but despite how haphazard my rig was it did work! it ran like a little blow torch inside the jar and exhausted though the ice water.

Side note: I totally forgot about that system that I made! that was when I was 15! It has been 18 years now and till I found this thread I had never remembered it, lol all this time I was wondering where I could get high purity HCL LOL life is weird like that!
[Edited on 29-4-2016 by XeonTheMGPony]

[Edited on 29-4-2016 by XeonTheMGPony]

[Edited on 29-4-2016 by XeonTheMGPony]

experimenter_ - 30-4-2016 at 05:16

Some new experiments showed that with A.C. current (50Hz) a gas that is not flammable at all is produced. However, the rate of production was very small; even at 2A I wasn't able to capture enough of the gas to test it. I think that a lower frequency of the A.C. current might result in a higher evolution rate.

Romix - 30-6-2016 at 22:42

Electrolyse NaCl, separating cathode from anode, so that electrons can flow through.
You'll get, NaOH formed on the cathode.

Be about the price of buying bulk. + saves time.

But if you have free electricity, then worth it.

[Edited on 1-7-2016 by Romix]

experimenter_ - 1-7-2016 at 04:18

According to this source:

Na2CO3 = Na2O + CO2 (t>1000°C)

I don't know if it worths the trouble to find a furnace for 1000°C but it seems a possible route.

Afterwards, just dissolve the Na2O in water and get NaOH.

[Edited on 1-7-2016 by experimenter_]

experimenter_ - 1-7-2016 at 14:30

Tried to heat baking soda in a canthal wire "furnace" for few minutes.
The soda melted. The pH after dissolving it in water was 14. It had the typical "oily" feeling of NaOH upon touching.

Added some acid and it fizzed. Obviously not all the soda was converted to NaOH; more time in the heat is required.

So, this reaction does work, at least qualitatively.

Geocachmaster - 1-7-2016 at 15:05

I recently made a dilute solution of sodium hydroxide- by placing a steel can containing sodium carbonate on some coals in a fire and leaving it there for ~15 minutes, and then combining the contents of the can with water. The resulting solution was a fairly strong basic liquid which reacted with aluminum (slowly) to produce hydrogen (I did the classic "pop" test). There was obviously a lot of sodium carbonate contamination. I did this on a whim and I am planning to try again with actual measurements and a gas torch instead of a wood fire. I'll post any of my results.

Metacelsus - 1-7-2016 at 16:09

Concentrated sodium carbonate by itself will react with aluminum. There was likely little to no hydroxide present.

CharlieA - 1-7-2016 at 17:52

Based on the Kb of carbonate ion ( l.8E-4), the pH would be 10. I don't think that you would need HO(-1) to get that pH. Good job confirming the evolution of H2, and not just assuming (you now what that means!

) that the gas evolved was H2.

ElizabethGreene - 13-7-2016 at 10:50

A tangent:
I've read that both paper and Stiff Salty Gelatin can work for separating a NaOH electrolytic cell, as long as you change the Hydroxide-side water regularly to keep it from becoming too caustic.

These are readings, not practical experience. Would you like me to conduct an experiment?