Some substances like water are unique in their ability of acting both as an acid and a base. We have seen this in case of water in section 7.10.2.
In presence of an acid, HA it accepts a proton and acts as the base while in the presence of a base, B– it acts as an acid by donating a proton. In pure water, one H2O molecule donates proton and acts as an acid and another water molecules accepts a proton and acts as a base at the same time.
The following equilibrium exists:
H2O(I) + H2O(l) ⇔ H3O+(aq) + OH-(aq)
Acid base conjugate acid conjugate base
The dissociation constant is represented by,
K = [H3O+] [OH-] / [H2O]
The concentration of water is omitted from the denominator as water is a pure liquid and its concentration remains constant.
[H2O] is incorporated within the equilibrium constant to give a new constant, Kw, which is called the ionic product of water.
Kw = [H+] [OH-]
The concentration of H+ has been found out experimentally as 1.0 × 10–7 M at 298 K. And, as dissociation of water produces equal number of H+ and OH– ions, the concentration of hydroxyl ions,
[OH–] = [H+] = 1.0 × 10–7 M. Thus, the value of Kw at 298K,
Kw = [H3O+][OH-] = (1 × 10-7)2 = 1 × 10-14 M2
The value of Kw is temperature dependent as it is an equilibrium constant.
The density of pure water is 1000 g / L and its molar mass is 18.0 g /mol. From this the molarity of pure water can be given as,
[H2O] = (1000 g /L)(1 mol/18.0 g) = 55.55 M.
Therefore, the ratio of dissociated water to that of undissociated water can be given as:
10-7/ (55.55) = 1.8 × 10-9or ~ 2 in 10-9 (thus, equilibrium lies mainly towards undissociated water)
We can distinguish acidic, neutral and basic aqueous solutions by the relative values of the H3O+ and OH– concentrations:
Acidic: [H3O+] > [OH-]
Neutral: [H3O+] = [OH-]
Basic: [H3O+] < [OH-]