Sciencemadness Discussion Board

mystery reaction...?

497 - 6-10-2007 at 21:36

I have this old childrens "chemistry" book that mentions producing hydrogen by dissolving sodium carbonate (washing soda) in water and adding aluminum foil. i thought this was interesting so i tried it, and it does work although a rather slow reaction. i found it proceeds at a much more reasonable rate when the reaction container was placed in boiling water. the reaction consumes the Al foil and produces a purplish clumpy precipitate (im guessing its clumpy because of the polymer coating on the Al foil.) after being filtered the remaining clear liquid only contains a fairly small amount of Na2CO3. the mystery of it all, is what exact reaction is taking place? i have searched the internet and found no reference to it. does anyone know what it could be? once i have time, i'll have to do some more detailed testing. one reason im so interested is it seems to be a cheap and easy method of producing hydrogen with fai

Sauron - 6-10-2007 at 21:45

Aluminum and not just foil, dissolves in bases. Sodium carbonate being a salt of a very weak acid is a strong base.

Try it with NaOH and you will get same reaction.

Dunno about your purple ppt.

unionised - 7-10-2007 at 01:28

"(im guessing its clumpy because of the polymer coating on the Al foil.)"
What polymer coating?

My guess is that the ppt is the mixture of other metals etc that were present in the Al alloy the foil was made from. Try filtering some off, washing it with water, wetting it with vinegar, leaving it exposed to air, and seeing if you get the blue/green colour of copper compounds.

Antwain - 7-10-2007 at 03:05

Originally posted by Sauron
Aluminum and not just foil, dissolves in bases. Sodium carbonate being a salt of a very weak acid is a strong base.

Try it with NaOH and you will get same reaction.

Dunno about your purple ppt.

Try it with NaOH and Al foil and you will get a steam explosion included for free :D

nitroglycol - 7-10-2007 at 03:59

I think the reaction would be as follows:

Na2CO3 + 2Al + 3H2O --> 2NaAlO2 + 3H2 +CO2

Thus, if you want pure hydrogen, this is not the best way to go. Sure, you could pass it through a solution of NaOH, but then if you have a solution of NaOH anyway, you could react that with the aluminum directly and not have to worry about CO2.

[Edited on 7/10/2007 by nitroglycol]

chemkid - 7-10-2007 at 07:17

Using sodium hydroxide wouldn't make much of a difference because

2Na2CO3 + H2O --> 2NaOH + CO2

and then add aluminum to the resulting NaOH solution...

2NaOH + 2Al + 6H2O --> 2NaAl(OH)4 + 3H2

Apparently the reaction is extremely violent, i would try tiny quantities first and work up.


[Edited on 7-10-2007 by chemkid]

chemkid - 7-10-2007 at 08:19

I tried this out. I started with a dilute sodium hydroxide solution and added aluminum. I pushed it to the bottom and after a few minutes it rose again indicating that gas bubbles were forming on it, however they were too small to see. So i added straight up sodium hydroxide (i used sodium carbonate before) and a large piece of aluminum which caused visible gas bubbles. I was using rather dilute solutions however the reaction was not terribly violent, as some websites warned. I would agree with 497 the reaction is reasonably slow. I am pretty sure the gas was hydrogen because it was very poorly soluble.


497 - 7-10-2007 at 09:52

hmmm i tend to agree with nitroglycol, Na2CO3 + 2Al + 3H2O --> 2NaAlO2 + 3H2 +CO2 seems the most likely. as to sodium hydroxide, i have used it before and it does surely work, but this reaction is definitely different, much less heat evolved. and 2Na2CO3 + H2O --> 2NaOH + CO2 definitely does not work, at least at room temp, i never get any CO2 from dissolving the carbonate. And the whole reason i like this reaction is that it avoids NaOH which in my experience is a messy reaction with lots of heat and NaOH mist. Also NaOH is not very easy to find now because of all the stupid regulations on meth precoursers. impurities in the Al do seem to be the most likely cause for the unusual precipitate. ill have to find some good pure Al and try that. thank you for all the good info.

chemkid - 7-10-2007 at 10:05

If you are using washing soda ( a fine powder of sodium carbonate) upon mixing with water you should get a transluecent to opaque solution that slowly clears. Upon close inspection you will see tiny rising bubbles of CO2. I will try to aquire some lime water to test.

The amended formula: Na2CO3 + H2O --> 2NaOH + CO2


[Edited on 7-10-2007 by chemkid]

497 - 7-10-2007 at 10:22

i kindof doubt it, since by my logic adding about 10 (106 per mol) grams of carbonate would give me about a 20th of a mol of CO2. i know for sure that that cloudiness does not contain a full liter of CO2. also NaOH absorbs CO2 effectively, and forms Na2CO3 which doesnt decompose until high temps. if it was NaOH i think it would react much more violently with the Al also... maybe im confused..

chemkid - 7-10-2007 at 10:33

I will test to see if a gas is being emiited by the reaction. I could very well be mistaken, as i am readily prone to being.:P


497 - 7-10-2007 at 10:45

yeah well... i think we have that in common. about the polymer coating, i suspect there is one because i read it somewhere (vary possibly wrong) but also noticed that one side of the Al foil is different, less shiny. maybe a very thin nonstick coating. probably varies from brand to brand. im not sure how pure it is either, so im going to have to get some pure Al powder. that will hopefully give a faster reaction too.

chemkid - 7-10-2007 at 10:58

I have read that sometimes Teflon is applied but that would really stop the reaction. As soon as i finish this damn essay i will do some more testing and give us a solution.


chemkid - 7-10-2007 at 12:06

I apologize, i was entirely wrong. I don't know where i read that information but it was clearly false. After testing with large amounts of reagents and a balloon on the flask i have established that most definitely did not evoke a gas. Thank you for correcting me.


post #100

497 - 7-10-2007 at 13:01

it would make sense that a teflon coating would stop the reaction, but i think its only coated on one side (if at all) so the other side still is able to react.. not sure about that though.

16MillionEyes - 7-10-2007 at 14:42

It seems to me that the reaction shouldn't produce any CO2 at all. If I'm correct the reaction goes as follows:
CO3^2-(aq) + H2O(l) --> HCO3-(aq) + OH-(aq)
2Al(s) + 6OH-(aq) --> 2AlO3^3- (aq) + 3H2(g)
The reason why this reaction is slow is indeed the coating (Al2O3) that prevents it from reacting directly but once all of it has dissolved the reaction should pick up and considering the reaction would be getting warm by now it would react even faster. This, however, is slowed down by the first reaction and is probably why it doesn't go as fast as it would otherwise.
Anyhow, I know someone else mentioned a similar reaction but forming the AlO2- instead the AlO3^-3 but I don't know for sure so any insights on the actual product would be helpful. Also, I'm fairly positive CO2 doesn't form as this would require that the carbonate anion protonate twice and then decompose to water and CO2 and being that a regular aqueous solution of bicarbonate ions doesn't give any noticeable gas (even though itself is slightly basic) this reaction shouldn't either as proton source would be the same for both.
Try doing a stoichiometrical reaction and precipitate out with some acetic acid (again stoichiometrically) this would help us determine what really forms.

497 - 7-10-2007 at 15:28

what you said makes sense, it was actually the first thing i suspected but the fact that the amount of ppt produced was quite alot more the the amount of Al used made me unsure. the ppt is almost gel like, gray purple, and apparently not very soluable if at all. i have yet to test it further. as i said before, it is very possible that the unusual properties are due to impurities. but you are correct, i (or someone else) need to do an accurately measured stoichiometrical reaction. the information is very much appreciated!

[Edited on 7-10-2007 by 497]

UnintentionalChaos - 7-10-2007 at 18:03

We did something very similar in Inorganic Chem and it forced me to look up all about amphoteric oxides...My knowledge has grown immensely since then...I thought I knew a lot then. I didn't and I bet I don't know that much now either, but definetly more than before.

In aqueous solution, NaOH dissociates almost completely into sodium and hydroxide ions.

NaOH (aq) <-> Na+ (aq) + OH- (aq)

The equilibrium lies very far to the right and in terms of aqueous reactions, it can be assumed in almost all cases that it lies all the way to the right.

Aluminum is an extremely reactive metal, but is normally passivated by a quickly forming and very tough oxide layer. Aluminum oxide, however, is what is called an amphoteric oxide. It can act both as a base and an acid in reactions. When in a strong solution of NaOH, the aluminum oxide protective layer will behave as an acid and react with the base forming sodium aluminate.

Al2O3 + 2NaOH -> 2NaAlO2 +H2O

In aqueous solution, it is probably better to assume that is exists as NaAl(OH)4.

Now that there is no protective oxide coating, the aluminum reacts quite energetically with the water.

H2O <-> OH- + H+

2Al (s) + 6OH- (aq) + 6H+ (aq) -> 2Al(OH)3 (s) + 3H2 (g)

Aluminum is being oxidized to aluminum hydroxide and is reducing hydrogen ions to hydogen gas.

Since the solution has lots more NaOH present, all the Al(OH)3 dissolves to form more sodium aluminate and exposes more metal surface.

Of course, the formation of the aluminate is an equilibrium reaction and only works to a degree. When the amount of NaOH starts to run low, the reaction pushes back towards the left and gelationous aluminum hydroxide drops out of solution, trapping large quatities of water in it which makes it look like a lot more than it is.

Aluminum foil is not completely pure aluminum and generally contains a percent or two of iron and maybe some other metals. The other metals thaht do not react with NaOH will exist as tiny spots of black in the hydroxide coloring it grayish. I can't quite explain the purple, but I've had reddish tan from iron oxide contamination. If you want to dissolve the precipitate, add lye or a strong acid. Some decent muriatic acid or lye should dissolve it quite nicely.

With sodium carbonate, you get

Na2CO3 (aq) <-> 2Na+ (aq) + CO3 (-2) (aq)

The equilibrium lies to the right, but not as far to the right as NaOH.


H2O <-> OH- (aq) + H+ (aq)

This occurs all the time in water, but the ions are balanced so the solution is neutral.

CO3 (2-) (aq) + H+ (aq) <-> HCO3- (aq)

HCO3- (aq) + H+ (aq) <-> H2CO3 (aq)

The first equilibrium is fairly far to the right, but the second is much further to the left. Since the number of hydrogen ions here and hydroxides must be equal, lets say that a little more than one out of two on average gets trapped as HCO3- or H2CO3. Since the hydrogen is tied up, the hydroxide ion is free to attack the aluminum. If the hydrogen isn't trapped, it is effectively neutralized by its partner OH. In NaOH, each sodium ion comes with a premade hydroxide ion. In sodium carbonate, 2 sodium come with around one hydroxide ion. In baking soda solution, a lot of sodium produces very few free hydroxides since it comes with the HCO3- ion premade. They understandably are strongest to weakest base, since bases produce free OH- ions in solution.

^This is something of an oversimplification since in sodium carbonate and bicarbonate, even the carbonate or bicarbonate ion isn't always released and may remain stuck to the sodium ion.
For somewhat complex reasons that 12AX7 explained in another thread, aluminum carbonate can't exist, so the carbonate ions from your solution stay trapped in solution, balanced by sodium ions. I do not think any appreciable amount of CO2 should be released, but I may be wrong.

I've mixed strong lye (NaOH) solution and aluminum foil in a soda bottle to fill a balloon and it got hot enough to soften and deform the plastic within a few minutes.

Sorry if you can't understand most of that lingo...I can't help it, but it will be a good oppurtunity to read up on the topic, which is always a good thing to do.

[Edited on 10-7-07 by UnintentionalChaos]

497 - 7-10-2007 at 22:10

ok i mostly understand what youre saying, but what is the final product when using carbonate? i guess what im not sure about is where the carbon is going. thanks, its nice to hear from someone who knows the subject well, im learning... ;)

not_important - 8-10-2007 at 02:37

Carbon will partially go off as CO2, boiling a solution of Na2CO3 will result in its slowing going more alkaline as CO2 is lost to the atmosphere; the remainder will be left as carbonate. So taking the equations already given

Na2CO3 (aq) <=> 2Na+ (aq) + CO3 (-2) (aq)
H2O <=> OH- (aq) + H+ (aq)
NaOH (aq) <=> Na+ (aq) + OH- (aq) mostly on the right side

CO3 (2-) (aq) + H+ (aq) <=> HCO3- (aq)
HCO3- (aq) + H+ (aq) <=> H2CO3 (aq)
H2CO3 <=> H2O + CO2 (aq)
CO2 (aq) <=> CO2 (g)

In a boiling solution the steam coming off help remove CO2 from the solution. It still doesn't go very much in that direction, the Na2CO3 solution will start out with a pH of 11.5 to 11.8, and slowly drift up towards 13 which means about a 0,01% to 0,1% NaOH solution.

Antwain - 8-10-2007 at 06:46

Originally posted by 497
ok i mostly understand what youre saying, but what is the final product when using carbonate? i guess what im not sure about is where the carbon is going. thanks, its nice to hear from someone who knows the subject well, im learning... ;)

bicarbonate and CO2 ;)

497 - 8-10-2007 at 10:18

so if i understand correctly, when the reaction is finished it ends up with carbonate and/or bicarbonate in solution with NaAl02 as ppt. hmm so the Na is going to be mostly used up in the NaAlO2 the water is going to have mostly straight carbonate ions? and if the water was distilled they would all boil off as CO2? ill have to test the H2 for CO2... anyone know a fairly sensitive test for it, i dont have litmus paper?

UnintentionalChaos - 8-10-2007 at 11:11

The precipitate is actually just aluminum hydroxide.

Al(OH)3 (s) + OH- (aq) <-> Al(OH)4- (aq)

The equilibrium is quite far to the left unless you have a huge excess of hydroxide ions, as in NaOH solution to push it to the right.

The water should still have its sodium carbonate floating around, although as not_important said, it will lose a very small amount of CO2 to make NaOH while boiling.

If you dry the entire reaction mixture out, grind the (now small amount of) solid up and leach with water several times, you should recover just about all of your original NaCO3. The NaCO3 (or NaOH) could be considered catalytic in promoting the reaction of aluminum metal with water to produce aluminum hydroxide and hydrogen gas via a sodium aluminate intermediate. Since sodium carbonate is a weak base compared to NaOH, not much sodium aluminate will be present at any given time, but that will still allow for small amounts of CO2 to be emitted.

You could make some limewater to test for CO2. Look for hydrated lime or slaked lime at agricultural stores and mix some with water, then filter to make a clear solution (only a tiny bit dissolves) but will form a white precipitate if CO2 containing gas mixtures are bubbled through it. Something like barium hydroxide solution is probably easier to prepare, but harder to obtain and a whole lot more toxic.

497 - 8-10-2007 at 13:11

hmm i like that scenario. id much rather deal with Al hydroxide than sodium aluminate. anyone know a use for Al hydroxide? i know its an intermediate for Al production, so it could be refined back to Al, but that would be impossible without a factory and a few million bucks...

Antwain - 8-10-2007 at 15:26

If you wash it well you could use it to make other aluminum salts by reacting it with HCl, H2SO4 etc. The only other thing I have ever used it for is column chromatography, but I expect that is beyond your needs and equipment at the moment. It probably also has to be specially prepared for that use.

16MillionEyes - 8-10-2007 at 15:46

The purple precipitate is very strange though. What's your source for Na2CO3? I've done the Al and OH- many times before and I tell you no purple precipitate forms and being that the OH- is the "active" reactant your reaction shouldn't either. I think I'll try doing the same exact set up and see if I get any purple precipitate but I highly doubt it. I just can't see where that stuff would come from.

497 - 8-10-2007 at 17:22

i agree it is a little surprising. i think it must come from the Al foil but i may be wrong. as far as my exact setup goes, it was pretty crude. well.. actually very very crude... i realllly need to get some good glassware... it consisted of a quart mason jar with maybe 2 inches of water in the bottom, in it was dissolved maybe 3 to 5 grams Arm and Hammer brand washing soda (about a year or two old.) then i added a 2 inch by 18 inch piece of Reynolds Wrap brand Al foil (made into a spiral that would fit in the jar.) the jar is set in a pot of water on the stove, simmering. it starts bubbling, increasing steadily. only after several minutes does a dark grey-blue gel like ppt start forming in noticeable amounts, eventually the foil disappears and the reaction stops. i filtered it through a paper towel, and squeezed out as much water from the ppt as i could. the remaining liquid was clear, when evaporated in an oven it crystallized out some white ppt, which im assuming is NaCO3. i cant tell if it crystallized out the same amount as i added, it looks roughly the same but i never measured it. the purplish Al hydroxide becomes lighter colored as it dries. as soon as i have time i'll try it again, this time measured and hopefully with some more pure Al. anyone know where i can get some other than the internet? have fun trying to reproduce the purple ppt.

16MillionEyes - 8-10-2007 at 18:14

Ahhhh, well, now that you mention it that way then it seems very likely that your purple precipitate is some sort of impurity from the Arm and Hammer washing soda. From what I found just now it contains all sorts of things raging from enzymes to Bentonite clay so I really find it very likely to get unexpected results like that. In other words, you're basically shooting in the dark here.
You should try getting some Na2CO3 from somewhere else with reliable at least discernible purity. One easy OTC way that I'd recommend is to get regular baking soda and oven heat at 250C or so for a while, this will get you relatively pure Na2CO3. Once you've done this then try again, I'm sure no purple precipitate will form this time.

497 - 8-10-2007 at 18:27

hehh that makes so much more sense. the reason i never considered it is i seemed to remember seeing that it was pure. i apparently remembered wrong, i should have checked. thanks that clears things up alot.

497 - 8-10-2007 at 18:34

im going to go try it right now, ill get back you you in a while.

UnintentionalChaos - 8-10-2007 at 19:04

I've used arm and hammer washing soda before. It is fairly pure, probably with a bit of crud in it since the stuff I have dissolves to a pale tan solution. I use it for washing certain stains off glassware (used extensively while doing curcumin extractions). The hotter the oven, the faster the baking soda (which is quite pure and very wonderfully does not form hydrates like sodium carbonate) decomposes. Spread a thin layer, breaking up lumps and ramp up the oven. I give it an hour at 500F which is massive overkill most likely, but does the job well.

497 - 8-10-2007 at 19:14

hmm yeah my oven only goes to 450F but it should work. its cooking right now. so if its not the washing soda where could the damn purple color come from???

UnintentionalChaos - 8-10-2007 at 19:23

Give it a shot with a different brand of aluminum foil. I usually get grayish aluminum hydroxide from very finely divided metal dispersed in it, but purple seems odd.

chemkid - 9-10-2007 at 11:19

Have you been able to replicate the purple precipitant?


[Edited on 9-10-2007 by chemkid]

16MillionEyes - 9-10-2007 at 17:22

With my set up certainly not, I did obtain the grayish precipitate from the NaAlO2 though. My sources were pharmaceutical grade NaHCO3 decomposed to Na2CO3 and regular aluminum foil. The reactants were left over night and then heated for about 5 minutes the day after. About 8 hours later I came back to find the precipitate on the bottom however a quick pH test revealed that the solution is still mostly basic (~10-11) and thus suggests that the reaction is still not done. I think that these results, nonetheless, prove the purple precipitate is not from the reaction itself but from some sort of impurity in his arm and hammer reactant.

497 - 9-10-2007 at 17:57

well i just finished trying it again, this time with food grade NaHCO3 decomposed to Na2CO3, different brand Al foil. other than that the setup was the same as before. got the normal grey ppt. must have been the washing soda that made things go differently. there was also much less ppt than before. it reacted quickly for a few minutes then slowed down until it almost came to a stop. there as excess Al foil. according to Unintentional Chaos's explanation the Na2CO3 is not used up in the reaction, so what would have stopped it? after it slowed down i added more water but it had little effect. but youre saying its making NaAlO2 which would explain the stop in the reaction.

UnintentionalChaos - 9-10-2007 at 18:52

Some degree of it does get tied up in solution as sodium aluminate and some simply gets trapped in the precipitate's matrix. Did you use identical quantities of washing soda and foil as before? Adding more water would actually slow it down since the formation of sodium aluminate is an equilibrium reaction pushed foward by a concentrated solution.

Use the original aluminum foil with the new sodium carbonate and see if you get purple. Simple elimination of possible factors.

arm and hammer + brand 1 = purple
NaHCO3 derived + brand 2 = gray
NaHCO3 derived + brand 1 =?
arm and hammer + brand 2 =?

One of those ? should be purple and the other should be gray. Test one of them.

Did you consider the fact that the new foil may be coated differently than the first foil and this may be stopping the reaction?

[Edited on 10-9-07 by UnintentionalChaos]

497 - 9-10-2007 at 20:41

well as far the purple color goes i am pretty sure its from the arm and hammer. i might do some more on that later, but we'll see. after finishing the the reaction with baking soda derived Na2CO3, filtering out the grey sludge and leftover Al. something interesting is that the clear liquid began to precipitate a very fine white ppt. so i put it in a ice water and precipitated as much as i could but it was too fine to be easily filtered. then i boiled off most of the water and filtered out the now clumpy ppt. so i decided to try the same steps on straight NaCO3 to hopefully help determine if the reaction changes things. so i dissolved roughly the same amount of NaCO3 as i used in the earlier reaction (i dont have the time or energy to measure it right now) and am cooling it, but its not precipitating... maybe because of impurities gained in the reaction, im not sure. i'll go boil the unreacted solution and how it precipitates that way. thanks for all the help and suggestions btw.

16MillionEyes - 11-10-2007 at 19:21

Well, my take in it is since the reaction of the Al is directly dependent on the presence of OH- anions once all the CO3(2-) has converted itself (due to change in equlibirum) to HCO3- the reaction will stop almost completely (or too slow to be noticeable) due to the reduced amount of OH- present in solution. Also if you consider the equation it balances out rather nicely:
Na +2CO3(2-) + 2H2O + Al ---> 2HCO3- + NaAlO2(s) + H2(g)
A quick test of the presence of CO3(2-) in solution confirms its presence (CO2 gas) and it would follow that it's from the HCO3-. In other words we're getting a NaHCO3 solution with NaAlO2 as precipitate. I would also discard Al(OH)3 as the source of this precipitate due to its amphoteric nature it would dissolve in a strong base and it does not.
Anyway, I hope this helps.

497 - 11-10-2007 at 19:54

well is there a good test to dertermine if it is NaAlO2? and/or a good way to test if the dissolved Na compound is bicarb or carbonate? also after the reaction completed and was filtered the liquid had a pH ~11.5 if thats any help.

497 - 11-10-2007 at 20:10

i boiled to the straight Na2CO3, it didnt begin to ppt untill there was very little water left (the post reaction solution began to ppt soon after it began boiling.) it formed a crust on the surface of the water, which did not occur with the other solution. ill try the same thing with NaHCO3 as soon as i have time.

[Edited on 12-10-2007 by 497]

16MillionEyes - 12-10-2007 at 18:16

You know, I have to retract what I said about NaAlO2. What I did report was in fact true in the sense that the precipitate didn't dissolve right away (as I would expect from Al(OH)3) in the presence of a strong base but after a day or so the precipitate had completely dissolved. As UnintentionalChaos said, the reaction depends on an equilibrium and this base might have thrown it off to the right but it still doesn't answer whether what forms in the precipitate is Al(OH)3 or NaAlO2.
Your reported pH is very similar to mine but I'd like to know just how much water is in your solution. In my case water is very limited so it's very probable that highly concentrated NaHCO3 is giving the relatively high pH.
I really can't think of a good way of determining whether there are CO3 or HCO3 ions in solution other than taking in consideration the basic properties of the CO3 over the HCO3. I still believe that a stoichiometric result is the best choice and would be very elucidating at this point so perhaps you would like try that. I would try it myself but I have no balance.:(

497 - 12-10-2007 at 19:07

yeah i lack a decent scale also... hmmm actually i might be able to get ahold of my grandma's gold scale :D. it should be accurate enough. on another note, i doubt the post reaction solution is HCO3 mainly because the pH of a solution of HCO3 similar to the quantity of CO3 i started out with in the earlier reaction was only about 8.5. the pH of the post reaction liquid was about 11.5. i got a pH of about 12 with a similar aount of straight CO3 dissolved. but i find it interesting that neither the known CO3 or HCO3 ppt out at room temp while the mystery post reaction solution does... maybe my quantities weren't as close as i had hoped. i think i'll finally have time to do a somewhat more accurate expirement this weekend.

Ozone - 12-10-2007 at 21:46

IIRC, some AlĀ° materials are Mn alloys (beer cans, for example). I have observed these becoming pink (purple at high c) under oxidative conditions...

Maybe the Al is not a pure as you think?

just pondering,


[Edited on 12-10-2007 by Ozone]

497 - 13-10-2007 at 00:34

certainly possible. but after switching to Na2CO3 made from NaHCO3 i dont get the purple ppt so i suspect the Arm and Hammer is the culprit. i would still like to use so good pure Al though, for good measure.

so i did a little crude testing of the soluble reaction product. well it is not very soluble at all. i added ~30 ml water to a teaspoon of it, just clouds up the water, eventually settles, very little dissolves (dissolves much better at boiling temp). i also tested carbonate and bicarbonate. a teaspoon of carbonate completely dissolves in 30 ml at room temp, a teaspoon of bicarb nearly completely dissolves in 30 ml. so what could it be? maybe impurities? just a plain old white powder, you couldn't tell it from bicarb or carbonate by looking.

[Edited on 12-10-2007 by 497]

Antwain - 13-10-2007 at 09:57

You should also be aware that sodium aluminate is rather easily hydrolysed. Hence diluting a solution of it will result in the precipitation of hydrated aluminum hydroxide.

497 - 13-10-2007 at 10:04

well i have never heard that. interesting...

497 - 13-10-2007 at 13:32

could it be that my mystery white powder is Al(OH)3? it is insoluble at room temp as far as i can tell. 30 ml of water with a teaspoon of it reads a pH of about 10.5. pH of a teaspoon of carbonate in 30 ml is 11.5 and a teaspoon of bicarb is 9.5.

[Edited on 13-10-2007 by 497]

497 - 13-10-2007 at 15:01

i have managed to acquire the balance, the good news is it can weigh up to 12 kg.. bad news is it has 1 g resolution :(. but it'll do.

so what should my molar ratios be? i was thinking .5m Al, but how much Na2CO3, since its not just simple Al to Al(OH)3? well i suppose i will just have to try something.

chemkid - 14-10-2007 at 16:03

flame test will distinguish between sodium aluminate and aluminum hydroxide.


497 - 14-10-2007 at 19:40

how exactly would one do that?

[Edited on 14-10-2007 by 497]

UnintentionalChaos - 14-10-2007 at 20:29

Problem there is that real tiny traces of sodium will discolor the flame if it is aluminum hydroxide....If any aluminate is present, washing it until there would be no sodium carbonate traces left will convert all of it to aluminum hydroxide and makes a flame test useless.

[Edited on 10-14-07 by UnintentionalChaos]

Antwain - 14-10-2007 at 23:18

You could titrate it against an acid. I can't find anything which I could use to tell you the pH of aluminium hydroxide, but from memory the hydroxide is not stable (will form aluminates or aluminium salts) outside the range ~4.5-9.5, while the aluminate is "strongly alkaline"

In one sense it is a moot point, since if you take pure hydrated aluminium hydroxide [Al(OH)3(H2O)3]3- and start adding OH- ions they will start to replace the water ligands sequentially, and conversely if you add H3O+ ions the hydroxide ligands will be replaced by water ligands. For this reason I have always found aluminium compounds to be aesthetically displeasing. Obtaining a 'pure' hydrated aluminium compound is damn near impossible. The alums are the only aluminium species I know of that crystalise even remotely decently.

And what the fuck is with the spell checker thinking that "aluminium" is "aluminum". I mean, sure, the americans can't pronounce it properly, but even they know how to spell it last time I checked. No offense to any americans here :P

497 - 15-10-2007 at 00:08

hah well... i find that whole naming thing funny....

"By 1812, Davy had settled on aluminum, which, as other sources note, matches its Latin root. He wrote in the journal Chemical Philosophy: "As yet Aluminum has not been obtained in a perfectly free state." But the same year, an anonymous contributor to the Quarterly Review, a British political-literary journal, objected to aluminum and proposed the name aluminium, "for so we shall take the liberty of writing the word, in preference to aluminum, which has a less classical sound.""

while you're at it why dont you change platinum to platinium and molybdenum to molybdenium??? :D

so... no definitive way to identify Al hydroxide? i cant say i could disagree with you about the aesthetically displeasing aspect of it. these damn alumiNUM compounds really are beginning to annoy me..

not_important - 15-10-2007 at 07:41

It's likely Al(OH)3. The sodium aluminates are only stable under very alkaline conditions, CO2 will split them up. The freshly precipitate hydroxide is highly absorbent, both compounds that interact with the -OH groups such as dyes and metal ions such as Na will be absorbed onto the precipitate; and it has a high surface area. This means that simple tests, such as checking for sodium, are just showing that something was absorbed in a mechanical and/or hydrogen bonding fashion. That also means that the white colour of Al(OH)3 may be modified by anything coloured in the reaction mix

Because Al compounds hydrolyse so readily you can assume that any white precipitate is going to be either a basic salt or Al(OH)3, if in moderately alkaline conditions then Al(OH)3 is the choice.

497 - 15-10-2007 at 11:11

thanks! that clears things up. i just did a larger scale measured test. more on that later, but your discription of Al(OH)3 fits. grey clumpy ppt, absorbent. it doesnt dry well at all... 300F in the oven for an hour and it was still a little damp...

chemkid - 21-10-2007 at 14:33

Aluminium and HCl gray insoluble precipitant

I reacted aluminium foil with muriatic acid until no more aluminium would react. The reaction formed a gray foam of bubbles. I left the beaker out overnight with a watch glass over it to allow any remaining aluminium flakes to react. In the morning i found there was still a gray precipitant. It measured a very low pH (cabbage indicator solution and pH paper). The precipitant did not dissolve with water or HCl. The precipitant seemed to float on the solution but i am not sure of that. Filtration yielded a gray filtrate.

To the best of my knowledge aluminium chloride is yellow and soluble in water. I don't think it was aluminium because that would have reacted upon the addition of more acid. Perhaps an impurity in the metal. Any insight on what the precipitant may be and if this is normal would be appreciated. Furthermore , why does the reaction of aluminium and HCl turn gray?

Sorry polverone if the intent wasn't for me just to post over here, but this thread isn't terribly helpful considering it discusses the reaction of aluminium with NaOH and sodium carbonate. Furthermore, even the information presented about aluminium hydroxide does nothing to explain why that would form over aluminium chloride. Are you proposing that this mysterious precipitant may also be aluminium hydroxide?


12AX7 - 21-10-2007 at 17:36

Commercial metals are not analytically pure or anything. The gray stuff is most likely silicon and other impurities, present in the 1-2% range.


Antwain - 21-10-2007 at 22:01

As 12AX7 says, impurities. Or as I call it, "crap". You always get that from Aluminium foil of any brand. I think if you filter it properly you will find that the particles are just really small, and that the distillate is actually clear. Best way to do it is actually to let it settle, maybe for days, in a beaker and then pour off as much solution as you can without disturbing the 'mud' at the bottom.

If its aluminium chloride you are after then good luck, I mean it! I was never able to make this to my exacting standards. As it concentrates, it hydrolises- which both gasses you with HCl and leads to a product which is not AlCl3. If you come up with a way to work that out I would be interested to know.

not_important - 21-10-2007 at 22:09

The most common foil alloys - the 1000, 3000 and 8000 series - contain between 0.5% and 1.5% iron, 0.1% and 0.7% silicon and 0.02% to 1.5% manganese. Up to 0.2% copper may be added when additional strength is required.

1000 series alloys are nearly pure aluminium, around 99% and in some cases 99,5 or better; the digits following the 1 indicate the fractional purity above 99%, 1030 is +99,3% Al. The 3000 series has manganese as its main alloying element, while the 8000 series uses lithium and is fairly rare. Almost all aluminium (alumium for you Davy fans) alloys have a few tenths of a percent each of silicon and iron.

Contrary to a popular belief, cooking foil is generally not coated with any polymer layer. Even the fluorocarbon plastics break down some at the temperatures the foil is exposed to during use, most other plastics go sooner; in any case not something you want coating your food. The differing appearance of the two sides of foil is due to the manufacturing process, the differences in forces applied and rolling/sliding contact with the equipment.

12AX7 - 22-10-2007 at 05:26

I believe the "nonstick" foil is anodized, FWIW.


not_important - 22-10-2007 at 05:57

Reynolds makes a non-stick foil that has a silicone release agent on the matte side (USP 6696511 - had to look that up out of curiosity) that was put on the market in 2002 or 2003. The coating can cause problems in recycling and using methane from landfills, so some areas are considering banning such products. Higher temperatures do cause off-gassing, using such foil as heat containment around glassware can cause your lab to acquire a thin coating of silicones.

Any additional processing adds to the cost, cheap generic food foil is likely to be just plain Al and thus the best for most uses in the home lab. Foil intended for construction or industrial applications has a better chance of not being a 1000 series alloy, but also is more likely to have information on which alloy it is made of.

woelen - 22-10-2007 at 06:53

To the best of my knowledge aluminium chloride is yellow and soluble in water. I don't think it was aluminium because that would have reacted upon the addition of more acid. Perhaps an impurity in the metal. Any insight on what the precipitant may be and if this is normal would be appreciated. Furthermore , why does the reaction of aluminium and HCl turn gray?

AlCl3 is a white powder, the hydrated salt AlCl3.6H2O also is white, somewhat transparent, crystalline. A solution of AlCl3 in hydrochloric acid is clear and colorless, not yellow.
As others have stated, the grey stuff is crap, impurity in the metal foil. I have reagent grade Al-needles and dissolving these in acid give perfectly clear and colorless solutions. I also have pyrotechnics grade Al-powder, and that gives dark grey crap as well (carbon??).

Making anhydrous AlCl3 is not only difficult, as Antwain writes, it is near impossible for a moderately equipped homelab. Anhydrous AlCl3 is nasty stuff. It heavily fumes in air, and is very corrosive. Not something you can make at home.
Hydrated AlCl3 on the other hand should not be that difficult. If you dissolve aluminium in a large excess amount of 10% HCl (use colorless acid, not the yellow stuff sold as muriatic acid) and let the grey stuff settle (which may take a long time, at least several hours), then the clear liquid can be allowed to evaporate in a petri dish. Don't heat it, just put it in a nice dry and warm place (e.g. above a heat radiator, with a paper tissue loosely covering the dish in order to prevent dust from entering it). After a few days you should have crystals of solid AlCl3.6H2O.

chemkid - 22-10-2007 at 11:16

That sounds excellent woelen. I think i will try that with a desiccator. (i have found that a simple coffee can and calcium chloride is extremely effective). Currently i am making aluminium hydroxide though. I have been able to filter the impurities so i will continue on that route. Perhaps i will acquire some high purity aluminium as well.

Thank you.

Antwain - 23-10-2007 at 05:23

A much cheaper way of doing that would be to buy alum from a garden shop and hit it with just about any base, say sodium carbonate or hydroxide

chemkid - 23-10-2007 at 11:31

Are your refering to aluminum sodium sulfate or one of the many other "Alums"?

woelen - 23-10-2007 at 12:15

The most common alum is KAl(SO4)2.12H2O, plain potassium alum.

Aluminium ion gives a hydroxide on addition of sodium carbonate. Aluminium carbonate cannot be precipitated from water. It is hydrolysed at once.

When you use sodium hydroxide, then you must adjust the amount carefully. If you add too much hydroxide, then the aluminium hydroxide redissolves again, giving aluminates. Aluminiumhydroxide is amphoteric and can act as base, and as acid. On addition of excess sodium hydroxide it acts as acid. When you try to evaporate a solution of an aluminate, then, however, it hydrolyses again, giving aluminium hydroxide.

kclo4 - 23-10-2007 at 17:45

Hmm, sorry if i missed a conclusion on this already, but it is only releasing H2 and NOT CO2. Because it produces Sodium Bicarbonate :)


chemkid - 23-10-2007 at 18:11

what is "it"? (if your talking about the sodium carbonate and water discussion a while back, i was totally confused and sodium carbonate does not react with water)


kclo4 - 23-10-2007 at 20:44

Yeah, earlier.
The reaction between sodium carbonate, water, and Al does not make CO2
it makes NaHCO3 instead.

497 - 24-10-2007 at 15:27

im pretty sure it ended up with the same amount of Na2CO3 as it started with. so no it doesnt make CO2 (or very much anyway). but it'd be easy to test if you wanted to know for sure.

[Edited on 24-10-2007 by 497]

woelen - 25-10-2007 at 12:58

Al, Na2CO3 and water will produce H2 and Al(OH)3, the net reaction being

2Al + 6H2O --> 2Al(OH)3 + 3H2

The carbonate only is used catalytically, it is required to make some OH(-) by means of hydrolysis, which in turn causes the Al to form H2 and hydroxide.

This reaction, however, will be painfully slow. Maybe a lot of heating of Al in a concentrated solution of Na2CO3 will give a somewhat higher reaction rate.

497 - 25-10-2007 at 20:33

oh yes its quite slow at room temp. but at 95C it goes quite nicely, ~6g Al reacted (not quite completely) in an hour or two. concentration of Na2CO3 was about 65g in 1 liter. i think the concentration could be higher and would probably go even faster.

[Edited on 25-10-2007 by 497]

ShadowWarrior4444 - 2-7-2008 at 12:57

I have recently been testing various gas generation mixtures, and came across something that I am not entirely sure of.

The mixture was simply Draino Kitchen Crystals and aluminum foil, with a bit of water. Very crude, however the reaction vessel and subsequent gas containment and piping were rigorously air-tight.

Draino adds NaNO3 to their product to produce NH3 in normal operation, however I have found that the gas produced from the reaction with foil contained very little NH3. Only on standing do the reaction byproducts emit a detectable amount of ammonia.

The main observation of interest is this, though:
16 hours after the completion of the reaction a white precipitate with quite a bit of volume formed from the previously clear liquid, (discounting the grey iron/silicon impurities.) I suspect this white precipitate to be aluminum hydroxide, so then what does it form from? The hydrolysis of sodium aluminate? The clear solution remained heavily basic throughout, and no evaporation took place.

UnintentionalChaos - 2-7-2008 at 16:30

It forms from the absorbtion of carbon dioxide by the solution. Carbonic acid being a stronger acid than aluminum hydroxide displaces it from sodium aluminate forming sodium carbonate and the aluminum hydroxide precipitates.

ShadowWarrior4444 - 2-7-2008 at 16:51

Originally posted by UnintentionalChaos
It forms from the absorbtion of carbon dioxide by the solution. Carbonic acid being a stronger acid than aluminum hydroxide displaces it from sodium aluminate forming sodium carbonate and the aluminum hydroxide precipitates.

Hmmm, there shouldnt have been much CO2 in the chamber--it was under positive hydrogen pressure the entire time. I suppose I'll run another test after flushing the vessel with difluoroethane ('heavier' than CO2.)

This may also be a useful way to make alumina powder for crucible and abrasive manufacture at home. Disolve the foil, let the impurities settle, decant and expose to CO2. I wonder if the NaOH can be recovered as well. Probably with the use of CaO and a bit of heat. Can Ca(OH)2 be used to dissolve aluminum in a similar fashion?

12AX7 - 2-7-2008 at 17:21

Incorrect, aluminic acid is slightly stronger than carbonic. I have crystals of sodium aluminate (originally the hydrate, but now dehydrated) which are quite stable in air in regards to hydrolysis.

Probably, alumina hydrate was supersaturated and only precipitated after some time.

If you had an excess of aluminum, continued corrosion certainly occurs, where the basic sodium aluminate solution (or acidic aluminum chloride solution, or...) is essentially catalyst, yielding more white goop.

Ca(OH)2 is certainly basic enough, but don't expect to make calcium aluminate (the cement).


ShadowWarrior4444 - 2-7-2008 at 19:58

I am fairly certain that the aluminum was not in excess, if anything the sodium hydroxide was.

It may be likely that alumina hydrate was supersaturated, however it did not precipitate on cooling, only after standing 16 hours at room temperature; this leads me away from saturation theories and more toward hydrolysis theories. Could sodium aluminate hydrolyze over time in the presence of a strong base?

The precipitation pattern may be of interest as well, it seemed to have a geological appearance. A bit like a porous rock, or a tangle of vines. Upon jarring the container slightly, the formation broke up into fine white powder which settled.

[Edited on 7-2-2008 by ShadowWarrior4444]

12AX7 - 3-7-2008 at 06:33

Hey, alumina hydrate is a refratory compound, it could take a while to form even a small crystal lattice. Or something.

Crystallizing a considerable time after cooling is *exactly* what supersaturation can do.