Introduction

Water molecules can transfer protons between themselves in a process called self-ionization. This fundamental concept in chemistry helps us understand acid-base equilibria and forms the foundation for calculating pH and pOH in aqueous solutions.

The Self-ionization Process

In pure water, molecules continuously undergo a reversible reaction where a proton transfers from one water molecule to another. This process produces hydronium ($\mathrm{H_3O^+}$) and hydroxide ($\mathrm{OH^-}$) ions:

$$2\mathrm{H_2O}(l) \rightleftharpoons \mathrm{H_3O^+}(aq) + \mathrm{OH^-}(aq)$$

At 25°C, this equilibrium is characterized by the water ionization constant ($K_w$):

$$K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14}$$

Temperature Effects on Water Self-ionization

The self-ionization of water is an endothermic process ($\Delta H > 0$). Temperature changes affect the equilibrium in the following ways:

Despite these changes in $K_w$, pure water remains neutral at all temperatures because $[\mathrm{H_3O^+}] = [\mathrm{OH^-}]$.

The Ka-Kb Relationship

The water ionization constant connects the acid dissociation constant ($K_a$) and base dissociation constant ($K_b$) through the relationship:

$$K_w = K_a \times K_b = 1.0 \times 10^{-14} \text{ (at 25°C)}$$

For a weak acid HA and its conjugate base A⁻, we can write:

Acid ionization:

$$\mathrm{HA}(aq) + \mathrm{H_2O}(l) \rightleftharpoons \mathrm{A^-}(aq) + \mathrm{H_3O^+}(aq)$$ $$K_a = \frac{[\mathrm{A^-}][\mathrm{H_3O^+}]}{[\mathrm{HA}]}$$

Base ionization:

$$\mathrm{A^-}(aq) + \mathrm{H_2O}(l) \rightleftharpoons \mathrm{HA}(aq) + \mathrm{OH^-}(aq)$$ $$K_b = \frac{[\mathrm{HA}][\mathrm{OH^-}]}{[\mathrm{A^-}]}$$

This leads to important relationships:

$$K_a = \frac{10^{-14}}{K_b} \text{ and } K_b = \frac{10^{-14}}{K_a}$$

Strength Relationships

The strength of an acid is inversely related to its conjugate base:

Acid Strength	$K_a$ Value	$pK_a$ Value	Conjugate Base Strength
Strong	High	Low	Weak
Weak	Low	High	Strong

Sample Problem

Question: In a hydrofluoric acid solution at 25°C, $[\mathrm{H_3O^+}] = 0.0700$ mol/L. Calculate $[\mathrm{OH^-}]$.

Solution:

$$[\mathrm{OH^-}] = \frac{K_w}{[\mathrm{H_3O^+}]} = \frac{1.0 \times 10^{-14}}{0.0700} = 1.43 \times 10^{-13} \text{ mol/L}$$