Acid Dissociation Constant Definition: Ka

What Is an Acid Dissociation Constant, or Ka in Chemistry?

The acid dissociation constant or Ka is an equilibrium constant that can be used to gauge the strength of an acid in a solution.
The acid dissociation constant or Ka is an equilibrium constant that can be used to gauge the strength of an acid in a solution. Maartje van Caspel / Getty Images

The acid dissociation constant is the equilibrium constant of the dissociation reaction of an acid and is denoted by Ka. This equilibrium constant is a quantitative measure of the strength of an acid in a solution. Ka is commonly expressed in units of mol/L. There are tables of acid dissociation constants, for easy reference. For an aqueous solution, the general form of the equilibrium reaction is:

HA + H2O ⇆ A- + H3O+

where HA is an acid which dissociates in the conjugate base of the acid A- and a hydrogen ion that combines with water to form the hydronium ion H3O+. When the concentrations of HA, A-, and H3O+ no longer change over time, the reaction is at equilibrium and the dissociation constant may be calculated:

Ka = [A-][H3O+] / [HA][H2O]

where the square brackets indicate concentration. Unless an acid is extremely concentrated, the equation is simplified by holding the concentration of water as a constant:

HA ⇆ A- + H+

Ka = [A-][H+]/[HA]

The acid dissociation constant is also known as the acidity constant or acid-ionization constant.

Relating Ka and pKa

A related value is pKa, which is the logarithmic acid dissociation constant:

pKa = -log10Ka

Using Ka and pKa To Predict Equilibrium and Strength of Acids

Ka may be used to measure the position of equilibrium:

  • If Ka is large, the formation of the products of the dissociation is favored.
  • If Ka is small, the undissolved acid is favored.

Ka may be used to predict the strength of an acid:

  • If Ka is large (pKa is small) this means the acid is mostly dissociated, so the acid is strong. Acids with a pKa less than around -2 are strong acids.
  • If Ka is small (pKa is large), little dissociation has occurred, so the acid is weak. Acids with a pKa in the range of -2 to 12 in water are weak acids.

Ka is a better measure of the strength of an acid than pH because adding water to an acid solution doesn't change its acid equilibrium constant, but does alter the H+ ion concentration and pH.

Ka Example

The acid dissociation constant, Ka of the acid HB is:

HB(aq) ↔ H+(aq) + B-(aq)

Ka = [H+][B-] / [HB]

For the dissociation of ethanoic acid:

CH3COOH(aq) + H2O(l) = CH3COO-(aq) + H3O+(aq)

Ka = [CH3COO-(aq)][H3O+(aq)] / [CH3COOH(aq)]

Acid Dissociation Constant From pH

The acid dissociation constant may be found it the pH is known. For example:

Calculate the acid dissociation constant Ka for a 0.2 M aqueous solution of propionic acid (CH3CH2CO2H) that is found to have a pH value of 4.88.

To solve the problem, first write the chemical equation for the reaction. You should be able to recognize propionic acid is a weak acid (because it's not one of the strong acids and it contains hydrogen). It's dissociation in water is:

CH3CH2CO2H + H2 ⇆ H3O+ + CH3CH2CO2-

Set up a table to keep track of the initial conditions, change in conditions, and equilibrium concentration of the species. This is sometimes called an ICE table:

Initial Concentration 0.2 M 0 M 0 M
Change in Concentration -x M +x M +x M
Equilibrium Concentration (0.2 - x) M x M x M

x = [H3O+

Now use the pH formula:

pH = -log[H3O+]

-pH = log[H3O+] = 4.88

[H3O+ = 10-4.88 = 1.32 x 10-5

Plug in this value for x for solve for Ka:

Ka = [H3O+][CH3CH2CO2-] / [CH3CH2CO2H]

Ka = x2 / (0.2 - x)

Ka = (1.32 x 10-5)2 / (0.2 - 1.32 x 10-5)

Ka = 8.69 x 10-10