I am just at the beginning of being able to interpret reactions using electron transfers and ionic equations. But I now have a question.
OBJECT: Elemental Aluminium and 5M NaOH to produce 1 mole of Aluminium hydroxide. I expected a gelatinous precipitate upon addition of Sulphuric
acid.
RESULTS: Gelatinous Aluminium Hydroxide precipitaed, but upon pouring off excess liquid the gelatinous precipitate dissolved into white crystalline
precipitate which is now very mushy and drying outside.
EXTRA OBSERVATION: I used 3 moles of NaOH to 1 mole of Al, I should have used a 1:1 ratio but I originally did this stupid equation: 3NaOH + Al =
Al(OH)3 + 3Na I forgot that the water changed this whole equation to produce NaAl(OH)4- and so is in a negatively
charged state until the sulphuric acid makes it a completely neutral Ion. So I have excess Sodium Hydroxide I know.
Question: Excess Na+OH- re-dissolves Al(OH)3 into Al(OH)4- . As does an excess of
H2SO4. However the addition of NH4OH precipitates Al(OH)3 but an excess does not re-dissolve it. Has this
something to do with the fact that Ammonium Hydroxide deprotonates in water but not the sodium hydroxide? Just want to understand the re-dissolving
from the gelatinous state, to the white crystaline state to absolutely dissolved back into solution.
(Assuming that I have understood things correctly so far)
[Edited on 1-5-2014 by CHRIS25]DraconicAcid - 1-5-2014 at 09:11
Ammonia (it's not actually ammonium hydroxide) is a weak base, and a solution of it will have a much lower concentration of hydroxide than a solution
of sodium hydroxide.
When aluminum hydroxide dissolves in a basic solution, the reaction is Al(OH)3(s) + OH-(aq) <=>
Al(OH)4-(aq). In a solution of sodium hydroxide, there's enough hydroxide around to drive the equilibrium to the right,
dissolving the aluminum hydroxide. In a solution of ammonia, there isn't, and the reaction stays on the left.CHRIS25 - 1-5-2014 at 09:44
"""and a solution of it will have a much lower concentration of hydroxide"""" thankyou Draconic, this explains it.
Thought I would add that I will be reacting the left over 2 moles of NaOH and Al(OH)4 in solution, precipitating this and melting to get
the aluminium oxide.
[Edited on 1-5-2014 by CHRIS25]blogfast25 - 1-5-2014 at 11:02
Aluminium hydroxide from Al is best prepared the other way around: dissolve Al in required amount of dilute (about 30 % I'd say) sulphuric acid, plus
20 % excess, acc:
This dissociation is almost 100 % complete and as a result [OH<sup>-</sup>] almost equals the nominal concentration of the base.
But when ammonia dissolves in water an equilibrium is established:
NH3(aq) + H2O(l) < === > NH4+(aq) + OH-(aq)
Acc. acid/base theory:
K<sub>B</sub> = ([NH4+]] x [OH-]) / [NH3] with K<sub>B</sub> the base equilibrium constant for NH3.
A little reworking yields the following approximate formula for dilute weak bases:
[OH-] = SQRT (K<sub>B</sub> x C<sub>NH3</sub> with
C<sub>NH3</sub> the nominal concentration of the NH3 solution.
So let’s crunch some numbers for C<sub>NH3</sub> = 0.1 M. We know that K<sub>B</sub> = 10<sup>-4.75</sup>
(literature), so [OH-] = √( K<sub>B</sub> x C<sub>NH3</sub> = √ ( 10<sup>-4.75</sup> x 0.1 ) = √ (10<sup>-5.75</sup> = 1.33 x 10<sup>-3</sup> = 0.00133 M
For a comparative solution of NaOH 0.1 M, we’d have [OH-] ≈ 0.1 M or about 75 times higher!
[Edited on 1-5-2014 by blogfast25]CHRIS25 - 1-5-2014 at 11:18
Oh the advantages of common sense.... this is brilliant. Thankyou.
And for this extra information, just seen it. 5M NaOH took 2 hours to dissolve 27g Al pipe found on the beach without any heat at all except the
exothermic generated, using sulphuric acid is so much slower I see and I have to heat it, even now it looks like it will be quite some hours, a much
slower reaction, and cleaner.
[Edited on 2-5-2014 by CHRIS25]CHRIS25 - 3-5-2014 at 07:27
You can see here the mess up I made: 1.5moles of sulphuric acid and 1mole aluminium. Only needed 85mLs But used 105mLs of 17.6M acid mixed in with
300mLs water to make a 5M solution, I later added another 50mLs water. (I know that concentrated sulphuric acid is a strong oxidizer and would not
dissolve the metal). I left the reaction on gentle heat and returned to check - as you see here I have aluminium sulphate precipitated, and although
I have not weighed it I have well over 3/4 of undissolved aluminium. I really did not expect this at all since the reaction was on heat yesterday for
three hours and hardly anything dissolved, I thought going away for two hours would not matter since there was still 400mLs of solution in the beaker.
It's not been a successful few days for me, though I do have lovely Tin chloride crystals now.
Ok, should have checked solubility of the sulphate, that explains it, 300 not enough especially on cooling, so add more water.
[Edited on 3-5-2014 by CHRIS25]blogfast25 - 3-5-2014 at 10:39
Chris25:
What has happened there is perfectly normal. Aluminium sulphate is highly soluble but also forms a very high hydrate, something silly like 18 H2O!
When that starts crystallising it takes with it lots of water and the whole thing tends to solidify. Been there, done that.
So you need to use simply more water and more H2SO4. A stoichiometric amount of acid will always cause the reaction to stall at some point because the
acid concentration become very low after a while.
I've dissolved 'tons' of Al scrap in H2SO4 solutions, w/o problems, to prepare alum (the potassium aluminium sulphate double salt.
Of course you can just add water to redissolve the aluminium sulphate you now have and filter off the unreacted metal.CHRIS25 - 3-5-2014 at 11:23
Yes so happy that I am not the only one, yes I read about that hydrated state After this happened. I prepared also the Potassium Aluminium sulphate
for a crystal garden from the same source of Al pipe. Sp anyway, now I know. Thanks Blogfast. Out of curiosity, since you have prepared so much
alum, what do you do with it if I may ask?blogfast25 - 3-5-2014 at 12:13
Out of curiosity, since you have prepared so much alum, what do you do with it if I may ask?
I used to sell it. Now I prepare it from aluminium sulphate and potassium sulphate: dissolve both together in the right ratio in boiling water, allow
to cool, it then crystallises beautifully. Pronto, Alum's your uncle!
Alum is one of the oldest Al compounds known to man and used since MA as a dye mordant. Some miles away from where I live there used to be the 'alum
works' (alum found in shallow underground). But no more forest because they used all the wood to fire the boilers!
[Edited on 3-5-2014 by blogfast25]CHRIS25 - 3-5-2014 at 13:08
I think I will try that and see how different they come out from the jar that I already have. Hey have you seen this: http://www.youtube.com/watch?v=l_USYub3djYblogfast25 - 3-5-2014 at 13:13
Nice crystal! (Understatement of the week)
To compare homemade alum with commercial grade, recrystallize the homemade. Take two parts of alum and dissolve it in one part of deionised water.
Heat and stir until all is fully dissolved. Then allow to cool overnight. Better purity that way...
with
C<sub>NH3</sub> the nominal concentration of the NH3 solution.