In the study of acid-base chemistry, understanding the constants associated with weak acids and bases is crucial. Weak acids, characterized as monoprotic, release one hydrogen ion (H⁺) when dissolved in water. This process can be represented by the ionization reaction where the weak acid (HA) donates an H⁺ to water (H₂O), forming hydronium ions (H₃O⁺) and the conjugate base (A^-). The reaction can be summarized as:

HA (aq) + H₂O (l) ⇌ H₃O⁺ (aq) + A^- (aq)

To quantify the strength of a weak acid, we use the acid dissociation constant (K_a), which is derived from the equilibrium expression. The equilibrium expression for this reaction is formulated as:

K_a = \frac{[H₃O^+-]}{[HA]}

In this expression, the concentrations of the products (H₃O⁺ and A^-) are multiplied together, while the concentration of the weak acid (HA) is in the denominator. It is important to note that when constructing equilibrium expressions, pure solids and liquids are not included, which is why water (H₂O) is omitted from the expression.

Understanding these concepts allows for a deeper comprehension of how weak acids behave in solution and how their strengths can be compared through their respective K_a values.