Acid dissociation constant

Acids and bases
Acceptor number Acid Acid–base reaction Acid–base homeostasis Acid strength Acidity function Amphoterism Base Buffer solutions Dissociation constant Donor number Equilibrium chemistry Extraction Hammett acidity function pH Proton affinity Self-ionization of water Titration Lewis acid catalysis Frustrated Lewis pair Chiral Lewis acid
Acid types
Brønsted–Lowry Lewis Mineral Organic Oxide Strong Superacids Weak Solid
Base types
Brønsted–Lowry Lewis Organic Oxide Strong Superbases Non-nucleophilic Weak
v t e

In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted ${\displaystyle K_{a}}$) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction

${\displaystyle {\ce {HA <=> A^- + H^+}}}$

known as dissociation in the context of acid–base reactions. The chemical species HA is an acid that dissociates into A⁻, called the conjugate base of the acid, and a hydrogen ion, H⁺.^[a] The system is said to be in equilibrium when the concentrations of its components do not change over time, because both forward and backward reactions are occurring at the same rate.^[1]

The dissociation constant is defined by^[b]

${\displaystyle K_{\text{a}}=\mathrm {\frac {[A^{-}][H^{+}]}{[HA]}} ,}$ or by its logarithmic form

${\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log _{10}K_{\text{a}}=\log _{10}{\frac {{\ce {[HA]}}}{[{\ce {A^-}}][{\ce {H+}}]}}}$

where quantities in square brackets represent the molar concentrations of the species at equilibrium.^[c][2] For example, a hypothetical weak acid having K_a = 10⁻⁵, the value of log K_a is the exponent (−5), giving pK_a = 5. For acetic acid, K_a = 1.8 x 10⁻⁵, so pK_a is about 5. A higher K_a corresponds to a stronger acid (an acid that is more dissociated at equilibrium). The form pK_a is often used because it provides a convenient logarithmic scale, where a lower pK_a corresponds to a stronger acid.

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