Acids and Bases: Strong Acids

The strong acid assumption is that the acid (HA) in solution dissociates completely into ions (i.e., H⁺ and A^–).

This assumption is useful because we can work out the end [H⁺] and pH simply from the concentration of acid that we start with and we don’t have to worry about any equilibrium constants. An ICE table can be useful to see what’s going on (but probably unnecessary):

So, if I have a bottle of HCl solution with 0.4 mol dm^-3 HCl written on it, it is actually a solution containg 0.4 mol dm^-3 H⁺ and 0.4 mol dm^-3 Cl^–. There will be very few molecules of undissociated HCl left. The pH of this solution is -log(0.4) = 0.40

We use the same approach for strong diprotic acids such as H₂SO₄. Here, an ICE table is a good reminder of the fact that complete dissociation of H₂SO₄ produces 2 H⁺ ions.

So, if I have a bottle of H₂SO₄ solution with 0.4 mol dm^-3 H₂SO₄ written on it, applying the strong acid approximation suggests that I actually have a solution containg 0.8 mol dm^-3 H⁺ and 0.4 mol dm^-3 SO₄^2-. The pH of this solution would be -log(0.8) = 0.10

(The strong acid assumption actually gets a little shaky here. In reality, the first dissociation of H₂SO₄ is complete but the second dissociation of H₂SO₄ is much more partial. A solution of H₂SO₄ therefore actually contains mostly H⁺ and HSO₃^–, so there is a lower [H⁺] and higher pH than expected)