The strength of a bronsted acid, HA, in aqueous solution is given in term of the acidity constant, Ka (which is also known as the acid ionization constant), where:
and a(X) is the activity of species X. The activity of a pure liquid is 1, and in dilute solution the approximation a(H2O) = 1 is made. Also, it is often convenient in dilute solutions, where the total ion concentration is less than 10-3 moldm-3, to substitute the concentration of a species, [X], for its activity.
Hence the acidity constant, Ka, can be defined, as can the pKa.
The acidity constant is usually reported as pKa for convenience, as Ka can span many orders of magnitude.
In water, some degree of autoionization (self-ionisation of water molecules) occurs, and this can be described in terms of the autoprotolysis constant, Kw. Under standard conditions, pKw = 14.00.
It is often the amount of H3O+ that is present which is the indication of acid strength. This is reported in terms of its concentration, as pH, which is defined as -log10[H3O+].
Substances with pKa<0 are classified as strong acids, as the proton transfer equilibrium lies in favour of donation to water. Substances with pKa>0 are known as weak acids, and now the proton transfer equilibrium lies in favour of protonated HA. The conjugate bases of strong acids are weak bases, and the conjugate bases of weak acids are strong bases.
The weaker the acid, the stronger the conjugate base.
Polyprotic Acids: These acids can donate more than one proton. Separate Ka‘s can be defined for the succesive loss of each of the protons.
It is harder to remove the second proton, than to remove the first. (ie: second deprotonation require more energy.) This may be represented by Ka1 < Ka2: which reflects the more difficult removal of the second proton from what is then a negatively charged species.
There are two main reasons that the second deprotonation is more difficult than the first:
- as implied above, there is an attractive electrostatic interaction which disfavours the removal of the positively charged proton from the negatively charged anion.
- the removal of another proton produces a doubly negatively charged species, which will induce more ordering in a surrounding dipolar solvent than a singly charged species will. This is disfavoured on entropic grounds.
Any acid stronger than H3O+ in water donates a proton to H2O. Any base stronger than OH– in water accepts a proton from H2O. Therefore, no acid stronger than H3O+ and no base stronger than OH– can survive in H2O. This is the effect called solvent leveling.
Water is said to have a leveling effect that controls the pH range achievable in aqueous solution. A pKa<0 implies an acid which has been leveled, whereas a pKa>14 implies a base which has been leveled. The width of the range of available pH’s is given by the pKw value.
Different solvents have different leveling ranges. Hence, pH’s may be achieved in other solvents which are not possible in aqueous solution.