Buffer

What is Buffer in Chemistry?

A solution whose pH is not altered to any great extent by the addition of small quantities of either an acid or base is called buffer solution.

Buffer is also defined as the solution of reserve acidity or alkalinity which resists change of pH upon the addition of a small amount of acid or alkali.

Many chemical reactions are carried out at a constant pH. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood, and bicarbonate also acts as a buffer in the ocean.

Table of Contents

• Recommended videos
• Characteristics of buffer solution
• Types of buffer solutions
• Buffer Action
• Hendersion’s Equation (pH of buffer)
• Buffer capacity
• Applications of Buffer in chemistry
• FAQs

Recommended videos

Characteristics of buffer solution

(i) It has a definite pH.

(ii) Its pH does not change on standing for long periods of time.

(iii) Its pH does not change on dilution.

(iv) Its pH is slightly changed by the addition of small quantity of an acid or base.

Types of buffer solutions

(a) Acidic Buffer:

It is formed by the mixture of weak acid and its salt with a strong base.

Examples: (i) CH₃COOH + CH₃COONa, (ii) HCN + NaCN, (iii) Boric acid + Borax etc.

(b) Basic Buffer:

It is formed by the mixture of a weak base and its salt with strong acid.

Examples: (i) NH₄OH + NH₄Cl, (ii) NH₄OH + NH₄NO3, (iii) Glycine + Glycine hydrochloride

(c) Simple Buffer:

It is formed by a mixture of acid salt and normal salt of a polybasic acid,

example Na₂HPO₄ + Na₃PO₄

Or a salt of weak acid and a weak base. Example: CH₃COONH₄

Buffer Actions

(a) Acidic Buffer:

It is the mixture of CH₃COOH and CH₃COONa in aqueous solution.

CH₃COOH ⇋ CH₃COO^– + H⁺ (incomplete dissociation)

CH₃COONa → CH₃COO^– + Na⁺ (complete dissociation)

H₂O ⇋ H⁺ + OH^– (incomplete dissociation)

Action of acid: when a drop of stong acid (HCl) is added in the above buffer solution H+ ions combine with CH3COO- ions to form feebly ionised CH3COOH. Whose ionisation is further suppressed due to common ion effect. So pH of the solution unaltered.

Action of base: when a drop of strong base (NaOH) is added to the above buffer solution it react with free acid to form undissociated water molecules. So pH of the solution unaltered.

CH₃COOH + OH^– ⇋ CH₃COO^– + H₂O

(b) Basic Buffer:

It is the mixture of NH₄OH and NH₄Cl in aqueous solution.

NH₄OH ⇋ NH₄⁺ + OH^– (incomplete dissociation)

NH₄Cl → NH₄⁺ + Cl^– (complete dissociation)

H₂O ⇋ H⁺ + OH^– (incomplete dissociation)

Action of acid: when a drop of HCl is added, the added H+ ions combine with NH4OH to form undissociated water molecules. So the pH of buffer is unaffected.

NH₄OH + OH^–   ⇋  NH₄⁺ + H₂O

Action of base: when a drop of NaOH is added, the added OH^– ions combine with NH₄⁺ ions to form feebly ionised NH₄OH. It is further suppressed due to common ion effect. So the pH of buffer is unaffected.

Hendersion’s Equation (pH of buffer)

(a) Acidic Buffer:

It is a mixture of CH₃COOH and CH₃COONa

CH₃COOH ⇋ CH₃COO^– + H⁺

CH₃COONa → CH₃COO^– +  Na⁺

By the law of chemical equilibrium,  K_a = {[CH₃COO^–] [H⁺]} / [CH₃COOH]

∴ [H⁺] = {K_a [CH₃COOH]} / [CH₃COO^–]

Taking negative log both sides, we obtain that

– log[H⁺] = – log K_a – log {[CH₃COOH]/[CH₃COO^–]}

pH = pK_a + log {[CH₃COO^–]/[CH₃COOH]}

pH = pK_a + log {[salt] / [acid]}

This equation is known as Hendersion’s Equation

Where, K_a = dissociation constant

(b) Basic Buffer:

It is a mixture of NH₄OH and NH₄Cl

NH₄OH ⇋ NH₄⁺ + OH^–

NH₄Cl → NH₄⁺ + Cl^–

By the law of chemical equilibrium,   K_b = {[NH₄⁺] [OH^–]} / [NH₄OH]

∴ [OH^–] = {K_b [NH₄OH]} / [NH₄⁺]

Taking negative log both sides, we obtain that

– log [OH^–] = – log K_b – log {[NH₄OH] / [NH₄⁺]}

pOH = pK_b + log { [NH₄⁺] / [NH₄OH]}

pOH = pK_b + log {[salt] / [base]}

This equation is known as Hendersion’s Equation

Where, K_b = dissociation constant

pH + pOH = 14

Buffer capacity

Buffer capacity is defined as the number of moles of acid or base added in one litre of solution as to change the pH by unity.

Buffer capacity (Φ) = No. of moles of acid or base added to 1 litre solution/change in pH

Φ = ∂b /∂(pH)

Where ∂b – No. of moles of acid or base added to 1 litre

∂(pH) – change in pH

Applications of Buffer in chemistry

(i) Buffers are used in industrial processes such as manufacture of paper, dyes, inks, paints, drugs, etc.

(ii) Buffers are also employed in agriculture, dairy products and preservation of various types of foods and fruits.

(iii) It is used to determine the pH with the help of indicators.

(iv) Blood is the natural buffer, it maintenance of pH is essential to sustain life because enzyme catalysis is pH sensitive process. The normal pH of blood plasma is 7.4.

(v) For the removal of phosphate ion in the qualitative inorganic analysis after the second group using CH₃COOH + CH₃COONa buffer.