3.4.3 Acids and Bases - Buffer action
Students should:
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Buffers
These are compounds or mixtures that resist change in pH on addition of small amounts of acid or base.
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To give this statement perspective, consider the effect of adding 0.1 cm3 of 0.1 mol dm-3 of a strong monobasic acid to 25 cm3 of a neutral water solution at pH 7. Moles of acid in 0.1 cm3 of 0.1 mol dm-3 acid = 0.01 moles New volume of solution = 25.1 cm3 Moles of hydrogen ions = molarity x volume (litres) = 1 x 10-5 moles Therefore molarity of hydrogen ions = moles/volume = 3.98 x 10-4 Therefore pH = -log[H+] = 3.40 Hence, the addition of 0.1 cm3 of acid has resulted in a fall in pH of 3.60 units. |
There are two types of buffer system:
- 1 Acidic buffers
- 2 Basic buffers
Acidic buffers
These operate below pH 7 and consist of a weak acid mixed with a salt of the weak acid. A typical mixture would be ethanoic acid and sodium ethanoate.
The weak acd only partially dissocates in solution:
CH3COOH 
CH3COO- + H+
This means that the solution contains mainly undissociated ethanoic acid molecules. If hydroxide ions, base, is added to this solution they remove the hydrogen ions from the right hand side, pulling the equilibrium to the right. But as there is an excess of the ethanoic acid, there is always enough to provide more hydrogen ions and the pH is kept approximately the same.
The salt of the weak acid dissociates 100% to the right hand side:
CH3COONa
CH3COO- + Na+
Addition of hydrogen ions to this solution get absorbed by the large amount of ethanoate ions and make more ethanoic acid molecules without affecting the overall amount of hydrogen ions present. The pH remains relatively unchanged.
In summary, the mixture provides a sink for hydrogen ions (the excess ethanoate ions) and a source of ethanoic acid to provide hydrogen ions in the event of additon of a base.
Basic buffers
These operate above pH 7 and consist of a weak base mixed with a salt of the weak base. A typical mixture would be ammonia and ammonium chloride. Their mode of action is very similar to that of the acidic buffer.
The ammonia solution dissociates very little:
NH3 + H2O
NH4+ + OH-
The ammonium salt dissociates 100%
NH4Cl
NH4+ + Cl-
The first equilibrium provides a source of ammonia to make more hydroxide ions if an acid is added. The second equilibrium provides a 'sink' of ammonium ions to absorb any hydroxide ions added.
The buffer law
Derivation of the buffer law is recommended as this will ensure that all of the signs are right (if carried out correctly of course).
For an acidic buffer system
Start with the weak acid equilibrium:
CH3COOH
CH3COO- + H+
And the equilibrium law expression:
If we take logs throughout we get
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log Ka = log [H+] + log
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[CH3COO-]
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[CH3COOH]
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Now change the signs throughout (multiply through by -1) and this gives:
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pKa = pH - log
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[CH3COO-]
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[CH3COOH]
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This is the buffer law expression (or at least one of them - the others being variations on this)
For calculations we simply need to know the concentration of the weak acid and the salt and if we know the value of pKa we can calculate pH (and vice versa)
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Example:: Calculate the pH of a buffer containing 8.2g of sodium ethanoate in 100 cm3 of 0,25M ethanoic acid (pKa = 4.75) 8.2g sodium ethanoate = 0.1 moles (Mr = 82) 0.1 moles in 100 cm3 = 1M solution Therefore: [CH3COOH] = 0.25 M, and [CH3COO-] = 1 M Substituting into the buffer law expression
This gives:
Therefore
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