Buffer Solutions

BUFFER SOLUTIONS

It is often necessary to maintain a certain pH of a solution in laboratory and industrial processess. This is achieved with the help of buffer solutions, buffer systems or simply buffers.

A buffer solution is one which maintains its pH fairly constant even upon the addition of small amounts of acid or base.

In other words, a buffer solution resists (or buffers) a change in its pH. That is, 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 :

       i.        a weak acid together with a salt of the same acid with a strong base. These are called Acid buffers. (e.g.,) CH₃COOH + CH₃COONa.

     ii.            a weak base and its salt with a strong acid. These are called Basic buffers. (e.g.,) NH₄OH + NH₄Cl.

Buffer action : Let us illustrate buffer action by taking example of a common buffer system consisting of a solution of acetic acid and sodium acetate (CH₃COOH/CH₃COONa).

CH₃COOH < --- --- > H⁺ + CH₃COO^-

CH₃COONa < -- --- > Na⁺       +       CH₃COO^-

since the salt is completely ionised, it 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 raised.

It is stated that a buffer solution containing equimolar amounts (0.10 M) of acetic acid and sodium acetate has pH 4.74. Now we proceed to discuss how the addition of a small amount of HCl or NaOH to the buffer solution affects its pH.

The pH of the buffer is governed by the equilibrium

CH₃ COOH < -- --- > CH₃ COO^-  + H⁺

The buffer solution has a large excess of CH₃COO^- ions produced by complete ionisation of sodium acetate,

CH₃ COONa -- -- -- > CH₃ COO^-+   Na⁺

1. Addition of HCl. Upon the addition of HCl, the decrease of H⁺ ions is counteracted by association with the excess of acetate ions to form unionised CH₃COOH. Thus the added H⁺ ions are neutralised and the pH of the buffer solution remains 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.

2. Addition of NaOH. When NaOH is added to the buffer solution, the additional OH^- ions combine with CH₃COOH to give CH₃COO^- and H₂O. Thus pH of the buffer solution is maintained almost constant. The buffer NH₄OH/NH₄Cl can also be explained on the same lines as of an acid buffer upon addition of HCl the H⁺ ions combine with NH₄OH to form NH₄⁺ and H₂O. pH is retained. Similarly when NaOH is added, the OH^- ions combine with NH₄⁺ ions present in the buffer solution to give NH₄OH and hence pH is maintained.

Henderson equation : The pH of an acid buffer can be calculated from the dissociation constant, K_a, of the weak acid and the concentrations of the acid and the salt used.

The dissociation expression of the weak acid, HA, may be represented As

HA < -- -- -- > H⁺ + A^-

K_a = [H⁺][A^-] / [HA]

[H⁺] = [HA] K_a  / [A^-]       ……………… 1

The weak acid is only slightly dissociated and its dissociation is further depressed by the addition of the salt (Na⁺ A^- ) which provides A^- ions (Common ion effect). As a result the equilibrium concentration of the unionised acid is nearly equal to the initial concentration of the acid. The equilibrium concentration [A^- ] is presumed to be equal to the initial concentration of the salt added since it is completely dissociated. Thus we can write the equation (1) as

[H⁺] = K_a ([acid] / [salt] )   ….. ..2

where [acid] is the initial concentration of the added acid and [salt] that of the salt used.

Taking negative logs of both sides of the equation (2), we have

-log[H⁺] = -logK_a - log ([acid]/[salt])

-log[H⁺] = pH and -logK_a = pK_a

pH = pK_a + log ([salt]/[acid])

This relationship is called the Henderson-Hasselbalch equation or simply Henderson equation.

In a similar way, the Henderson-Hasselbalch equation for a basic buffer can be derived. This can be stated as :

^{pOH = pK}b ^{+ log} ([salt]/[acid])

Significance of the Henderson-Hasselbalch equation. With its help

1. The pH of a buffer solution can be calculated from the initial concentrations of the weak acid and the salt provided K_a is given.

However, the Henderson-Hasselbalch equation for a basic buffer will give pOH and its pH can be calculated as (14 - pOH).

2. The dissociation constant of a weak acid (or weak base) can be determined by measuring the pH of a buffer solution containing equimolar concentrations of the acid (or base) and the salt.

Since, [salt] = [acid],   log ([salt]/[acid]) = log 1 =0

pK_a = pH

The measured pH, therefore, gives the value of pK_a of the weak acid.

Likewise we can find the pK_b of a weak base by determining the pOH of equimolar basic buffer.

3. A buffer solution of desired pH can be prepared by adjusting the concentrations of the salt and the acid added for the buffer.

It is noteworthy that buffer solutions are most effective when the concentrations of the weak acid (or weak base) and the salt are about equal. This means that pH is close to the value of pK_a of the acid (or pK_b of the base).