[9-92-6-19]

Earlier yesterday, my lab group and I were tasked with finding the mass of Al in a sample of alum; my group and I drew up a procedure and concluded that the best course of action would be to dissolve the alum in a minimum amount of water, and then to precipitate the sulfate ions out of solution via adding an excess of Ba(NO₃)₂.

The precipitation worked well, however the vacuum filtration apparatus failed; precipitate flowed through the filter. We tried several types of filter paper, and even tried gravity filtration. Despite these efforts, my lab group (and all other lab groups that used a BaX_y compound) failed.

Yet there was one group that used Ca(NO₃)₂. Interestingly, this was the only group that the filtration worked properly. The group ended with approximately 5% error from the stoichiometrically ideal amount of Al from this precipitates, so I would presume most of the sulfate was precipitated out of solution (of course there were more steps following precipitating the sulfate, though).

None of this makes any sense to me. BaSO₄ is much more insoluble than CaSO₄. Why would CaSO₄ work while all other groups that used Ba fail?

[Arkcon]

Do you have some other observations on your precipitate, that you can compare with others, that might show a difference between BaSO₄ and CaSO₄?  I'm talking about observations that we get all the time from people writing up observations here on the Chemical Forums -- color, consistency, ... things like that?  This may be why the class was split into groups.

[9-92-6-19]

Group 1 (CaSO₄): "Upon mixing Ca(NO₃)₂ a whitish precipitate formed with a slight brown tinge. Mass of precipitate: 3.2 grams"

Group 2 (BaSO₄): "Mixing Ba(NO₃)₂ forms a dense white precipitate.

Group 3 (BaSO₄): "Once the Ba(NO₃)₂ a bright white precipitate came out of solution."

Group 4 (BaSO₄): "Once the Ba(NO₃)₂ a bright white precipitate with a cream-like opaqueness was created."

Groups 2, 3, and 4 could not mass the precipitate due to the aforementioned filtration problems.

[Borek]

It would help to compare pH of solutions added to precipitate sulfates. How they were prepared? Were they acidified?

[9-92-6-19]

Unfortunately, I don't have the exact procedure that every group followed, but they all should have simply dissolved KAl(SO₄)₂·12H₂O in water.

I do, however, have the pH of all the solutions before the precipitate was formed.

Group 1: "The pH was found to be 4.5."

Group 2: "After dissolved the alum into water the pH was found to be 4.6."

Group 3: "The dissolved alum resulted in a pH drop of 2.7 from neutral water."

Group 4: "Based on the following equations:

Al³⁺ + 6H₂O  Al(H₂O)₆³⁺

SO₄^2- + H₂O  HSO₄^- + OH^-

A equilibrium is created and results in the following hydrolyzed reaction:

Al(H₂O)₆³⁺ + 3OH^-  Al(OH)₃(H₂O)₃ + 3H₃O⁺

Thus the solution is predictad to be acidic.

Before addition of alum: 7.0
After addition of alum: 4.3 "