Contents
What is a Buffer solution?
A buffer solution is one which maintains its pH fairly constant even upon the addition of small amounts of acid or base.
It is often necessary to maintain a certain pH of a solution in laboratory and industrial processes.
Buffer solutions help to maintain pH.A buffer solution resists a change in its pH. We can add a small amount of an acid or base to a buffer solution and the pH will change very little.
Two common types of buffer solutions are :
1. Acid Buffers
It is a weak acid together with a salt of the same acid with a strong base.
E.g., CH3COOH + CH3COONa.
2.Basic Buffers
It is a weak base together with salt with a strong acid.
E.g.,NH4OH + NH4Cl.
Let us understand buffer action by taking an example of a common buffer system consisting of a solution of acetic acid and sodium acetate (CH3COOH/CH3COONa).
CH3COOH ⇋ H+ + CH3COO–
CH3COONa → Na+ + CH3COO–
The salt completely ionises, provides the common ions CH3COO– in excess. The common ion effect suppresses the ionisation of acetic acid. This reduces the concentration of H+ ions which means that pH of the solution is raises. Thus, a 0.1 M acetic acid solution has a pH of 2.87 but a solution of 0.1 M acetic acid and 0.1 M sodium acetate has a pH of 4.74. Thus 4.74 is the pH of the buffer. On addition of 0.01 mole NaOH the pH changes from 4.74 to 4.83, while on the addition of 0.01 mole HCl the pH changes from 4.74 to 4.66. Obviously the buffer solution maintains fairly constant pH .
How does a buffer operate?
A buffer solution containing equimolar amounts (0.10 M) of acetic acid and sodium acetate has pH 4.74.
Let’s understand how the addition of a small amount of HCl or NaOH to the buffer solution affects its pH.
The equilibrium is as follows
CH3COOH ⇋ CH3COO– + H+ ……….(1)
The buffer solution has a large excess of CH3COO– ions produced by complete ionisation of
sodium acetate,
CH3COONa→ CH3COO– + Na+ ………..(2)
Upon addition of HCl, the increase of H+ ions associate with excess of acetate ions to form CH3COOH. Thus the added H+ ions are
neutralised and the pH of the buffer solution remains virtually unchanged. However owing to the increased concentration of CH3COOH, the equilibrium (1) shifts slightly to the right to increase H+ions. This explains the marginal increase of pH of the buffer solution on addition of HCl
Upon addition of NaOH additional OH– ions combine with H+ ions of the buffer to form water molecules. As a result the equilibrium (1) shifts to the right to produce more and more H+ ions till practically all the excess OH– ions are neutralised and the original buffer pH restored. However, a new equilibrium system is set up in which [CH3COOH] is lower than it was in the original buffer. Consequently [H+] is also slightly less and pH slightly higher than the buffer pH values.
Operation of a Basic buffer as NH4OH/NH4Cl can also be explained on the same lines as of an acid buffer upon addition of HCl the H+ ions combine with OH– ions of the buffer to form water molecules.The equilibrium,
NH4OH → NH4+ + OH −
is shifted to the right till all the additional H+ ions are neutralised and the original buffer pH restored.When NaOH is added to the buffer solution, OH– ions associate with excess of NH4+ ions to form unassociated NH4OH. Thus the pH of the buffer is maintained approximately constant.
Reference: Essentials of Physical Chemistry by Arun Bahl, B.S. Bahl , G.D. Tuli
Written by Uma Maheswari P