The Dissociation of Water

$H$₂O can act as both a proton donor and acceptor for itself. A proton can be transferred from one water molecule to another, resulting in the formation of one hyroxide ion ($OH-$) and one hydronium ion ($H$₃O⁺).

$2H$₂O(l) H₃O⁺(aq) + OH^-(aq)

This is called the autoionization or dissociation of water. This equilibrium can also be expressed as:

$H$₂O(l) H⁺(aq) + OH^-(aq)

In the above equilibrium, water acts as both an acid and a base. The ability of a species to act as either an acid or a base is known as amphoterism.

The concentrations of $H$₃O⁺ and $OH-produced\; by\; the\; dissociation\; of\; water\; are\; equal.\; The\; corresponding\; equilibrium\; expression\; for\; this\; would\; be:$

In pure water at $25oC$, [H₂O] = 55.5 M. Why?

This value is relatively constant in relation to the very low concentration of H⁺ and OH^- (1 x 10^-7 M).

Therefore,

Rearranging gives:

Where $K$_W designates the product (55.5 M)$K$_C and is called the ion-product constant for water. At $25oC,$ K$$_W is equal to $10-14$.

The ion-product constant always remains constant at equilibrium (as the name implies). Consequently, if the concentration of either $H+$ or $OH-$ rises, then the other must fall to compensate. In acidic solutions, $[H+]\; >\; [OH-]$, and in basic solutions .

The pH Scale

For example, a neutral solution at $25oC$ contains equal concentrations of $H+$ ions and $OH-$ ions, where $[H+]\; =\; 10-7$M. Thus, the pH of the solution is obtained by:

From this it is obvious that at $25oC$:

Notice that the pH of a solution measures the concentration of dissociated protons, and not the total concentration of acid in a solution.

The negative log scale is useful for measuring other minute quantities, for example to measure $[OH-]$:

Knowing this, we obtain the following useful expression: