Acids and Bases

Buffer Solutions

Buffer Solutions

A buffer solution is a type of solution able to resist a change in pH due to the addition of a little amount of acid or alkali or due to dilution.

How buffers work

A buffer solution is a mix of acid and alkali: the acid neutralises the alkali and the alkali neutralises the acid. However, a strong acid and a strong alkali cannot be used as they would react with one another thereby not acting as an effective buffer. Neither can a very weak acid and a very weak alkali as they will not react enough.

Instead, acids and bases should be chosen that are strong enough to react with H3O+ and OHions but not with one another. A weak acid and its conjugate base are an ideal mixture, for example NH4Cl and NH3.

Resisting pH change: adding acid and alkali

Taking the example above, this mixture will react with an acid and an alkali as follows:

  • NH4+(aq) + OH(aq) ? NH3(aq) + H2O(l)
  • NH3(aq) + H3O+(aq) ? NH4+(aq) + H2O(l)

Resisting pH change: dilution

With dilution, the weak acid and weak base are both able to dissociate more in order to compensate:

  • NH4+(aq) + H2O(l) ? NH3(aq) + H3O+(aq)
  • NH3(aq) + H2O(l) ? NH4+(aq) + OH(aq)

Other buffer solutions

A buffer a mixture does not have to be a weak acid and its conjugate base. Instead, you can use any mixture of a weak acid and a weak base.

A substance which is able to act as a weak acid and a weak base can also behave like a buffer. For example sodium hydrogencarbonate:

  • HCO3(aq) + H3O+(aq) ? CO2(g) + 2H2O(l)
  • HCO3(aq) + OH(aq) ? CO32-(aq) + H2O(l)

Amino acids can also act like buffers, 2-aminopropanoic acid for example:

  • CH3CH(NH2)COOH(aq) + OH+(aq) ? CH3CH(NH2)COO(aq) + H2O(l)
  • CH3CH(NH2)COOH(aq) + H3O+(aq) ? CH3CH(NH3+)COOH(aq) + H2O(l)

Calculating the pH of buffer solutions

If the buffer solution is composed of a weak acid and its conjugate base then it is possible to calculate its pH using the following method.

Take the mixture: CH3COOH / CH3COONa

Therefore: Ka = ([CH3COO][H3O+]) / [CH3COOH]

So: [H3O+] = Ka[CH3COOH] / [CH3COO]

In general it is possible to calculate the concentration of H3O+ ions from:

[H3O+] = Ka[acid] / [base]

To create a buffer solution of a specific pH the concentrations of the acid and base relate to a particular ratio. This can be worked out by expressing the pH in terms of the concentrations of the reactants:

[H3O+] = Ka[acid] / [base]

So: log10[H3O+] = log10 Ka + log10([acid] / [base])

So: – log10[H3O+] = -log10 Ka + log10([base] / [acid])

So: pH = pKa + log10([base] / [acid])

Limiting changes in pH with a buffer solution

It is possible to quantitatively show how a buffer solution is able to limit changes in pH.

Take the following mixture: 0.6M HClO (Ka = 3.7 x 10-8 M) and 0.2M NaClO with a pH of 7.0

When 0.01 moles of HCl are added to 100cm3 of the buffer:

ClO(aq)

+

H3O+(aq)

?

HClO(aq)

+

H2O(l)

Initially

0.02

0.01

0.06

Finally

0.01

0.07

Therefore: the [base] / [acid] ratio = 0.01 / 0.07

pH = pKa + log10([base] / [acid]) = 6.58 (a unit change of 0.4)

If 0.01 moles of HCl were added to 100cm3 of pure water then the resulting solution would have a pH of 1.0, a change of 6 units.

When 0.01 moles of NaOH are added to 100cm3 of the buffer:

HClO(aq)

+

OH(aq)

?

ClO(aq)

+

H2O(l)

Initially

0.06

0.01

0.02

Finally

0.05

0.03

Therefore: the [base] / [acid] ratio = 0.03 / 0.05

pH = pKa + log10([base] / [acid]) = 7.21 (a unit change of 0.2)

If 0.01 moles of NaOH were added to 100cm3 of pure water then the resulting solution would have a pH of 13.0.

Therefore, when a small amount of acid or alkali is added to a buffer solution the pH does not change by much.

However, if too much of an acid or an alkali is added then the buffering capacity of a buffer solution can be exceeded leading to a massive pH change. In the example above this would be over 0.02 moles of HCl or over 0.06 moles of NaOH.

A buffer solution is most effective at resisting a pH change if the concentrations of acid and alkali are equal. If there is more acid than base or more base than acid in a buffer solution then it cannot resist change on the addition of an acid or base respectively.

Natural buffers

A number of biological systems reply on a relatively constant pH:

  • Blood: the pH must be kept around 7.4 for which hydrocarbonate ions are used to maintain this level.
  • Tears: the pH must also be kept around 7.4 and amino acids are used to maintain this level.