A2-Level Acids and Bases

Up until now, we have just linked pH to H⁺ ion concentration in a solution. This is fine for calculating the pH of an acidic solution, as it is only the concentration of H⁺_(aq) ions we need to worry about.

If we only have an alkaline solution, however, there is no H⁺ ion concentration to use for a pH calculation – this concentration needs to be found based upon the concentration of hydroxide (OH^-_(aq)) ions present. For example, if you have a solution of 1.0 mol dm^-3 sodium hydroxide, NaOH, there is no H⁺ ion concentration to use for calculating pH. Instead, we need to find the H⁺_(aq) concentration based on the hydroxide ion concentration.

In water, a very small percentage of molecules dissociate to release H⁺ and OH^- ions. It’s only a small percentage that dissociate but, because the number of water molecules present in a solution is so high, there is still a concentration of H⁺ ions. There is an equilibrium between the water, the hydroxide ions and protons present.

Here, the water is acting as a weak acid and, even though the position of equilibrium massively favours the backward reaction, a K_a expression can still be applied.

As already mentioned, the actual number of water molecules dissociating is so small the overall concentration of water molecules remains effectively the same – this means we can consider it to be ‘constant’.

The K_a of a weak acid is also constant (at a given temperature), this means K_a and [H₂O] can be treated as two constants that are never going to change at a particular temperature. These are multiplied together to give a new ‘combined’ constant, K_w. This constant is called the ionic product of water and has a value of 1x10^-14 at 298K.

This gives us a new expression:

When dealing with alkaline solutions, the concentration of the alkali can be used to find the concentration of H⁺ ions present by using the above K_w expression. The concentration of H+ ions can then be used to find the pH (pH = -log₁₀[H⁺_(aq)]).