Themes > Science > Chemistry > Inorganic Chemistry > Acids and Bases > Acids and Bases Index > pH of a Buffer


Computing the pH of a buffer solution is a fairly straightforward application of the equilibrium constant expression. A buffer is a solution of a weak acid and its conjugate base, so

  • HA(aq) < = > H+(aq) + A-(aq)
  • Ka = [H+][A-]/[HA]
If you want to know the pH of a buffer solution, just solve for the [H+] concentration above
[H+] = Ka*([HA]/[A-])

There are a couple of simplifications that you can almost always make in the above equation for [H+]. First, since both the acid and its conjugate base are weak, the concentrations of HA and A- are very close to their initial values, [HA]0 and [A-]0: there's no need to do the usual equilibrium calculation. Second, since the volume of the solution is the same for both the acid and the conjugate base, the volumes in the concentration cancel out, leaving just the moles of acid and moles of base added.

[H+] = Ka*(nHA/nA-)

Example: a buffer solution is made from 0.100 moles of acetic acid and 0.200 moles of sodium acetate in 1 liter of water. If Ka for acetic acid is 1.8*10-5, what is the pH of the buffer?

Solution: Use the equation above to compute the pH. Sodium acetate dissolves completely in water, so the amount of Ac- in solution is 0.200 moles

  • [H+] = Ka*(nHA/nA-)
  • [H+] = 1.8*10-5*(0.100/0.200)
  • [H+] = 9.0*10-6
  • pH = -log10([H+]) = -log10(9.0*10-6) = 5.0


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