Question

The molecular formula of aluminium sulphate is:
(a) $A{l_2}{\left( {S{O_4}} \right)_3}$
(b) $AlS{O_4}$
(c) $A{L_3}{\left( {S{O_4}} \right)_2}$
(d) $A{L_2}{\left( {S{O_4}} \right)_4}$

Answer 1

Hint: The charges of a molecule should add up to zero, to have a net charge of zero. Since it is a neutral compound. We have to balance the molecular formula so that we have a neutral charge.

Complete Step by step answer: Starting from the basic, we know that symbol for aluminium is $Al$and for sulphate, we won't find it in the periodic table since it is a polyatomic ion. So, sulphate is $SO_4^{2 - }$we can get this from some common polyatomic ions table.

Aluminium is a metal and sulphate ions are made up of non-metals, therefore we have an ionic compound. Now, we need to think about the charges on aluminium and when we look into the periodic table we can get that It has a charge of 3+ also we know the charge of sulphate is 2-. Since the net charge has to zero, for that we can use something called the criss-cross method to make that happen. We move the charge of sulphate toward aluminium and vice-versa.
Now, we get $A{l_2}{\left( {S{O_4}} \right)_3}$ we used a parenthesis for sulphate ions since it is a polyatomic ion.
We can see now that the charges add to zero. $2 \times 3 + 3 \times \left( { - 2} \right) = 0$
i.e., we have 2 aluminium with charge 3+ and 3 sulphate ions with charge 2-.

Hence this is the correct formula for aluminium sulphate, option (a) is correct .

Note: We can see that the options (c) and (d) have given the symbol of aluminium as AL, we cannot use the capital letter for the second alphabet for an element in the periodic table. Therefore, clearly that two options are wrong right away.