Buffer solutions are the ones that resist a change in its pH on adding a small amount of an acid or a base.
There are two kinds of commonly used buffer-solutions
(i) A weak acid and a soluble ionic salt of the weak acid e.g. acetic acid and sodium acetate; CH3 COOH + CH3COONa and,
(ii) A weak base and a soluble ionic salt of the weak base e.g. ammonium hydroxide and ammonium chloride ; NH4OH + NH4Cl.
Buffer action in general is defined as the ability of the buffer solution to resist the changes in pH value when a small amount of an acid or a base is added to it.
Mechanism of Buffering Action
To understand the mechanism of buffer action we can take the example of an acidic buffer that is made up of a weak acid like acetic acid and its sodium salt Sodium acetate. In this acidic buffer, the solution will contain equimolar amounts of acetic acid and sodium acetate. Usually, a large number of sodium ions (Na+), acetate ions (CH3COO–) and undissociated acetic acid molecules are present.
The salt exists completely as ions.
Here the buffer will consist of both acid (CH3COOH) and its conjugate base (CH3COO–). If we add a small quantity of acid the hydrogen ions will be removed by the conjugate base (CH3COO–). It is represented as follows:
H+ (aq) + CH3COO– (aq) ↔ CH3COOH (aq)
Here the ethanoic acid will only be slightly dissociated in the form CH3COOH. What it means is that it will not contribute any H+ ion. Therefore, the pH of the resulting solution will remain more or less constant. The added H+ ions are also removed due to which there is no appreciable decrease in pH.
The reaction comes to a completion as CH3COOH is a weak acid whose ions have a strong tendency to form non-ionized CH3COOH molecules. On the other hand, if we add a strong base the OH– ion gets neutralized by the reaction with the acid in the buffer,
CH3COOH (aq) + OH– (aq) → CH3COO– (aq) + H2O (l)
We can also consider that the OH– ion can react with the H+ ion in order to form water. The OH– ions are that are added are removed wherein the acid equilibrium shifts to the right to replace the H+ ions that are exhausted. This results in a minor change in the pH value.
Alternatively, if we add a drop of NaOH, the OH– ions react with the free acid to give undissociated water molecules. The extra OH– ions of the base are neutralized. As a result, the pH of the solution remains the same. This condition or the resistance offered by the pH when a base is added is known as reserve acidity. This is mainly due to CH3COOH.
If we add a strong base, the acid present in the buffer neutralizes the hydroxide ions (OH–).
So we can say when acid or base is added its effect is practically balanced and the pH of the solution is always constant.