Buffer solution

Certain solutions, such as that of ammonium acetate, have a tendency to resist any change in its hydronium ion concentration or pH, whenever a small amount of a strong acid or a strong base is added to it. This property of a solution is known as buffer action.

Buffer Solution

A solution which resists any change of pH when a small amount of a strong acid or a strong base is added to it, is called a buffer solution or simply as a buffer.

Alternatively, a buffer solution may be defined as a solution whose pH value does not change appreciably upon the addition of small amounts of a strong acid, base and/or water from outside.

Thus, buffers have reserve acidity and reserve alkalinity.

Buffer solutions usually consist of a mixture of a weak acid and its salt with a strong base e.g., CH₃COOH and CH₃COONa, or that of a weak base and its salt with a strong acid e.g., NH₄OH and NH₄Cl. The solution of any salt of a weak acid and a weak base e.g., ammonium acetate, also shows buffering property.

Types of Buffers

There are two types of buffers, acid buffer and basic buffer.

Acid buffer

A buffer solution containing a large amounts of a weak acid, and its

salt with a strong base, is termed as an acid buffer. Such buffer solutions have pH on the acidic side i.e., pH is less than 7 at 298 K. The pH of an acid buffer is given by the equation.

where K_a is the acid dissociation constant of the weak acid.

Basic buffer

A buffer solution containing relatively large amounts of a weak base and its salt with a strong acid, is termed as a basic buffer. Such buffers have pH on the alkaline side i.e., pH is higher than 7 at 298 K.

The pH of a basic buffer is given by the equation:

where K_b is the base dissociation constant of the weak base.

These equations are called Henderson equation.

Buffer-capacity and Buffer-range

The effectiveness of any buffer is described in terms of its buffer capacity. It is defined as, 'the number of equivalents of a strong acid (or a strong base) required to change the pH of one litre of a buffer solution by one unit, keeping the total amount of the acid and the salt in the buffer constant'.

The buffer capacity of a buffer is maximum when acid to salt or base to salt ratio is equal to 1 i.e., it contains equal number of moles of acid (or base) and the salt. All buffer solutions remain effective over a small pH range: this pH-range is characteristic of the buffer and is termed as the buffer-range.

For the two types of buffers, it is given by

Buffer range in pH units

Acid buffer: pK_a -1 to pK_a +1

Basic buffer: (pK_w - pK_b) -1 to (pK_w - pK_b) +1

Consider acetic acid - sodium acetate buffer, an acid buffer. The acid dissociation constant (K_a) of acetic acid is 1.84 x 10^-5. Therefore, pK_a for acetic acid is 4.74. Then, the buffer range of an acetic acid - sodium acetate buffer is,

pH = (pK_a) - 1 to (pK_a) + 1

= 4.74 1 to 4.74 + 1

= 3.74 to 5.74

Thus, the acetic acid - sodium acetate buffer will act as an effective buffer over the pH range 3.74 to 5.74.

The pH of a buffer solution depends only on the ratio of the concentrations of the salt and the acid, or salt and the base. It does not depend on the individual concentration. Since, the ratio remains the same even when the solution is diluted. However, at higher dilutions, buffers do not remain effective as, they are not able to resist a change in the pH value due to the addition of a strong acid or a strong base.

Mechanism of Buffer Action

The buffering action of buffer solutions can be explained in terms of the Bronsted-Lowry concept of acids and bases as follows:

Mechanisms

Action of an acid buffer

An acid buffer contains relatively large amounts of a weak acid (HA) and its salt with a strong base (say NaA). The buffer solution thus contains large concentration of HA and A^- (due to the dissociation of the salt), apart from H₃O⁺ and OH^-.

An addition of small amount of a strong acid causes the reaction,

to proceed in such a direction that an equivalent amount of A^- combines with H₃O⁺ to give the same amount of undissociated weak acid, HA. Thus, the added acid is picked up by the anions (from the salt) present in large concentrations in the buffer. As long as the added strong acid is in smaller amounts, the changes in the concentrations of salt and that of the weak acid, (HA) are small. Therefore, the acid to salt ratio does not change appreciably by the addition of strong acid to the buffer solution. As a result no noticeable change is seen in the pH value of the buffer. Addition of a strong base to an acid buffer on the other hand causes the reaction

to proceed in the forward direction, resulting in the formation of an equivalent amount of the salt at the cost of the buffer acid. As long as the added base is in small amounts, the ratio of weak acid to salt remains virtually unchanged. As a result, no observable change in the pH value is seen.

Action of basic buffer

A basic buffer contains a weak base (BOH), and its salt with strong acid (BX). The buffer solution thus contains large amounts of the weak base BOH, and the cation B⁺ (coming from the dissociation of the salt BX), in addition to H₃O⁺ and OH^-.

The addition of an acid or a base to the basic buffer causes the following reactions:

proceeds in the forward direction. It is clear that the addition of an acid or a base to any buffer solution does cause a change in the concentrations of the buffer acid (or base) and the salt. But, because of the relatively much larger concentrations of these in the buffer solution, for all practical purposes, the ratio, [Salt] / [Acid] or [Salt] / [Base] remains constant. Hence, the pH does not change.

Problem

12. Calculate the pH of a buffer solution containing 0.2 mole of NH₄Cl and 0.1 mole of NH₄OH per litre. K_b for NH₄OH = 1.85 x 10^-5.

Solution

According to Henderson's equation:

Applications of Buffers

Buffers find extensive applications in a variety of fields.

In biochemical systems

pH plays a very significant role in biochemical reactions. For example, the blood in our bodies is buffered at a pH value of 7.36-7.42 due to bicarbonate - carbonic acid buffer. A mere change of 0.2 pH units can cause death. Certain enzymes get activated only at certain definite pH values.

Agriculture

The pH of the soil is very important for having proper crop yield. The soils get buffered due to the presence of salts such as carbonates, bicarbonates, phosphates and organic acids. The choice of fertilizers depends upon pH of the soil.

Industry

Practically all industries use buffers in one process or the other. Major industries, which employ buffers are paper, dyes, ink, paints and drugs industries.

Analytical chemistry

Buffers find extensive use in analytical chemistry, viz., both in qualitative and quantitative analysis. For example, qualitative analysis of Group III and Group IV is done in solutions buffered by NH₄Cl + NH₄OH. Buffers are used in the removal of interfering radicals such as phosphate, oxalate, borate and fluoride etc. The control of pH is very important in the field of food preservation