# How does equilibrium affect the strength of an acid?

Jan 2, 2015

Effect of equilibrium
One of the ways of determining the strength of an acid is its pKa value.

${p}_{{K}_{a}}$ of an acid is defined as $- \log \left({K}_{a}\right)$ where ${K}_{a}$ is the equilibrium constant of the dissociation reaction.

Any acid will dissociate in the following manner

$H A r i g h t \le f t h a r p \infty n s {A}^{-} + {H}^{+}$

And ${K}_{a}$ for the reaction is defined as

${K}_{a} = \frac{\left[{A}^{-}\right] \cdot \left[{H}^{+}\right]}{\left[H A\right]}$

$p {K}_{a} = - {\log}_{10} {K}_{a}$

The larger the value of pKa, the smaller the extent of dissociation at any given pH—that is, the weaker the acid. A weak acid has a pKa value in the approximate range −2 to 12 in water. Acids with a pKa value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable. In other words the denominator of the above equation is so small that the value of ${K}_{a}$ becomes high.

Weak acids have considerably high concentrations of molecules that are not dissociated. Hence the value of ${K}_{a}$ becomes small.

pKa values for strong acids can, however, be estimated by theoretical means or by extrapolating from measurements in non-aqueous solvents in which the dissociation constant is smaller, such as acetonitrile and dimethylsulfoxide.

Source:Acid Dissociation