The task at hand is to complete the chemical equation and and determine the whether the products form a precipitate from aqueous solution.
Reactants: NaBr + Al2(SO4)2
I am really lost with Al2(SO4)2. As the compound must have a neutral charge, the net charge of the anions must equal the net charge of the cations. As (SO4)2 has a net charge of 4-, with each SO4 having a charge of 2-, Al2 must have a net charge of 4-. This really doesn't make much sense to me, as the aluminum cation has a charge of 3+, and 2 aluminum cations would logically have a charge of 6+, which is unequal to the 4- net charge of (SO4)2. The only diatomic cation my class has been introduced to is the mercurous cation, Hg2 2+, but it seems to me that the only way the compound Al2(SO4)2 makes sense is if Al2 is a diatomic cation with a charge of 4+. Is this so?
As for the products of the reaction, assuming Al2 is a diatomic cation with a 4+ charge, I think they would be Na2(SO4), which I know to be water-soluble, and Al2Br4. I know ionic bromide compounds are generally water-soluble, with the exception of silver, Hg2 2+, and lead salts, but is a bromide salt of a diatomic aluminum cation water-soluble?
Thank you for your time.