Acid dissociation constant
equilibrium constant as a measure of acid strength in solution
In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted
K
a
{\displaystyle K_{a}}
) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction
HA
↽
−
−
⇀
A
−
+
H
+
{\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+. 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.
The dissociation constant is defined by
K
a
=
[
A
−
]
[
H
+
]
[
H
A
]
,
{\displaystyle K_{\text{a}}=\mathrm {\frac {[A^{-}][H^{+}]}{[HA]}} ,}
or by its logarithmic form
p
K
a
=
−
log
10
K
a
=
−
log
10
[
A
−
]
[
H
+
]
[
HA
]
=
log
10
[
HA
]
[
A
−
]
[
H
+
]
{\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log _{10}K_{\text{a}}=-\log _{10}{\frac {[{\ce {A^-}}][{\ce {H+}}]}{{\ce {[HA]}}}}=\log _{10}{\frac {{\ce {[HA]}}}{[{\ce {A^-}}][{\ce {H+}}]}}}
where quantities in square brackets represent the molar concentrations of the species at equilibrium. For example, a hypothetical weak acid having Ka = 10−5, the value of log Ka is the exponent (−5), giving pKa = 5. For acetic acid, Ka = 1.8 x 10−5, so pKa is 4.7. A lower Ka corresponds to a weaker acid (an acid that is less dissociated at equilibrium). The form pKa is often used because it provides a convenient logarithmic scale, and a lower pKa corresponds to a stronger acid.
Theoretical background
The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the pKa value is directly proportional to the standard Gibbs free energy change for the reaction.
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