Buffer Solution

For several purposes, we need solutions which should have constant pH. Many reactions, particularly the biochemical reactions, are to be carried out at a constant pH. But it is observed that solutions and even pure water (pH = 7) cannot retain the constant pH for long. If the solution comes in contact with air, it will absorb CO₂ and becomes more acidic. If the solution is stored in a glass bottle, alkaline impurities dissolve from glass and the solution becomes alkaline.

A solution whose pH is not altered to any great extent by the addition of small quantities of either an acid (H⁺ ions) or a base (OH^- ions) is called the buffer solution. It can also be defined as a solution of reserve acidity or alkalinity which resists change of pH upon the addition of small amount of acid or alkali.

Refer to the following video for buffer solutions

So, a buffer solution can be defined as a solution which resists a change in its pH when such a change is caused by the addition of a small amount of acid or base.

This does not mean that the pH of the buffer solution does not change (we make this assumption while doing numerical problems). It only means that the change in pH would be less than the pH that would have changed for a solution that is not a buffer.

There are three types of buffer solutions:

weak acid–salt buffer
weak base–salt buffer and
salt buffer

General characteristics of a buffer solution

It has a definite pH, i.e., it has reserve acidity or alkalinity.
Its pH does not change on standing for long.
Its pH does not change on dilution.
Its pH is slightly changed by the addition of small quantity of an acid or a base.

Buffer solutions can be obtained:

by mixing a weak acid with its salt with a strong base, eg; Acetic Acid + Soidum Acetate, Boric acid + Borax, Phthalic acid + Potassium acid phthalate.
by mixing a weak base with its salt with a strong acid, e.g; (a) PNH₄OH + NH₄Cl , (b) Glycine + Glycine hydrochloride
by a solution of ampholyte. The ampholytes or amphoteric electrolytes are the substances which show properties of both an acid and a base. Proteins and amino acids are the examples of such electrolytes.
by a mixture of an acid salt and a normal salt of a polybasic acid, e.g., Na₂HPO₄ + Na₃PO₄, or a salt of weak acid and a weak base, such as CH₃COONH₄.

The first and second type are also called acidic and basic buffers respectively.

Explanation of buffer action

Acidic buffer:

Consider the case of the solution of acetic acid containing sodium acetate. Acetic acid is feebly ionised while sodium acetate is almost completely ionised.

The mixture thus contains CH₃COOH molecules, CH₃COO^- ions, Na⁺ ions, H⁺ ions and OH^- ions. Thus, we have the following equilibria in solution:

Feebly ionised	Completely ionised	Very feebly ionised
CH₃COOH $\rightleftharpoons$ H⁺ + CH₃COO^-	CH₃COONa $\rightleftharpoons$ Na⁺ + CH₃COC^-	H₂O $\rightleftharpoons$ H⁺ + OH^-

When a drop of strong acid, say HCl, is added, the H⁺ ions furnished by HCl combine with CH₃COO^- ions to form feebly ionised CH₃COOH whose ionisation is further suppressed due to common ion effect.

Thus, there will be a very slight effect in the overall H⁺ ion concentration or pH value.

When a drop of NaOH is added, it will react with free acid to form undissociated water molecules.

CH₃COOH + OH^- $\rightleftharpoons$ CH₃COO^- + H₂O

Thus, OH^- ions furnished by a base are removed and pH of the solution is practically unaltered.

Buffer capacity

The property of buffer solution to resist alteration in its pH value is known as buffer capacity. It has been found that if the ratio [Salt]/[Acid] or [Salt]/[Base] is unity, the pH of a particular buffer does not change at all. Buffer capacity is defined quantitatively as number of moles of acid or base added in one litre of solution as to change the pH by unity,

i.e..

$Buffer\ Capacity = \frac{Change\ in\ concentration\ or\ number\ of\ moles\ of\ acid\ or\ base\ added\ to\ litre\ of\ buffer\ }{Chnage\ in\ pH}$

$\phi =\frac{\Delta B}{\Delta pH}$

where δb → number of moles of acid or base added to 1 litre solution and δ(pH) → change in pH.

Buffer capacity is maximum:

When [Salt] = [Acid], i.e., pH = pK_a for acid buffer
When [Salt] = [Base], i.e., pOH = pK_b for base buffer under above conditions, the buffer is called efficient.

Utility of buffer solutions in analytical chemistry

Buffers are used:

To determine the pH with the help of indicators.
For the removal of phosphate ion in the qualitative inorganic analysis after second group using CH₃COOH + CH₃COONa buffer.
For the precipitation of lead chromate quantitatively in gravimetric analysis, the buffer, CH₃COOH + CH₃COONa, is used.
For precipitation of hydroxides of third group of qualitative analysis, a buffer, NH₄Cl + NH₄OH, is used.
A buffer solution of NH₄Cl, NH₄OH, and (NH₄)₂CO₃is used for precipitation of carbonates of fifth group in qualitative inorganic analysis.
The pH of intracellular fluid, blood is naturally maintained. This maintenance of pH is essential to sustain life because, enzyme catalysis is pH sensitive process. The normal pH of blood plasma is 7.4.
Following two buffers in the blood help to maintain pH (7.4).

(a) Buffer of carbonic acid (H₂CO₃ and NaHCO₃)

(b) Buffer of phosphoric acid (H₂P0₄, HPO^2-)

Buffers are used in industrial processes such as manufacture of paper, dyes, inks, paints, drugs, etc.
Buffers are also employed in agriculture, dairy products and preservation of various types of foods and fruits.
Buffer solutions are necessary to keep the correct pH for enzymes in many organisms to work.
A buffer of carbonic acids (H₂CO₃) and bicarbonate (HCO₃⁻) is present in blood plasma to maintain a pH between 7.35 and 7.45.
Buffer solutions are used to increase shelf life of drugs

Henderson's Equation or henderson-Hasselbalch equation
(pH of buffer solution):

Refer to the following video for Henderson's Equation

(i) Acidic buffer:

It consists of a mixture of weak acid and its salt (strong electrolyte). The ionisation of the weak acid, HA, can be shown by the equation

HA $\rightleftharpoons$ H⁺ + A^-

Applying law of mass action,

K_a = H⁺A^-/[HA]

It can be assumed that concentration of A^- ions from complete ionisation of the salt BA is too large to be compared with concentration of A^- ions from the acid HA.

BA $\rightleftharpoons$ B⁺ + A^-

Thus, [HA] = Initial concentration of the acid as it is feebly ionised in presence of common ion

and [A^-] = Initial concentration of the salt as it is completely ionised.

$[H^+] = K_a \times \frac{[Acid]}{[Salt]}$

Taking logarithm and reversing sign,

$-log\ [H^+] =log\ {[Salt]}-log\ [Acid]-log\ K_a$

$\Rightarrow\ pH =log\ K_a+log\ \frac{[Salt]}{[Acid]}$

This is known as Henderson's equation.

When

$\frac{[Salt]}{[Acid]}=10$ , then

pH = 1 + pK_a

and when [Salt]/[Acid] = 0.1 then

pH = pK_a -1

So weak acid may be used for preparing buffer solutions having pH values lying within the ranges pK_a + 1 and pK_a -1. The acetic acid gas a pK_a of about 4.8; it may, therefore, be used for making buffer solutions with pH values lying roughly within the ranges 3.8 to 5.8.

(ii) Basic Buffer:

It consists of a weak base and its salt with strong acid. Ionization of a weak base, BOH, can be represented by the equation.

BOH $\rightleftharpoons$ B⁺ + OH^-

Applying law of mass action,

K_b = [B⁺][OH^-]/[BOH]

or [OH^-] = K_b [BOH]/[B⁺]

As the salt is completely ionized, it can be assumed that whole of B⁺ ion concentration comes from the salt and contribution of weak base to B⁺ ions can be ignored.

BA $\rightleftharpoons$ B⁺ + A^- (Completely ionised)

So [OH^-]= K_b[Base]/[Salt]

or pOH = log[Salt]/[Base] log K_b

or pOH = pKb + log[Salt]/[Base]

Knowing pOH, pH can be calculated by the application of the formula.

pH + pOH = 14

(iii) Salt Buffer

The fundamental principle behind a buffer action is the fact that on adding an acid the system consumes the H⁺ ion added to produce a weak acid and on adding a base, it consume the OH^– ion added to produce a weak base. This ensures that H⁺ or OH^– ion added is consumed and the weak acid or the weak base produced gives less H⁺ or OH^– ion as they are weak.

Based on this principle, a solution of a salt of a weak acid and weak base is also called as a buffer. Let us take the example of CH₃COONH₄. It dissociates as,
CH₃COONH₄ → CH₃COO^– + NH₄⁺.

When H⁺ ion is added, CH₃COO^– ion consumes it to give CH₃COOH. When OH^– ion is added, CH₃COO^- ion consumes it to give NH₄OH. Hence it acts as a buffer.

Test Your Knowledge

Q.1 Which of the following pairs represent an acidic buffer?

a. NH₄OH +NH₄Cl b. CH₃COOH + NH₄OH c. CH₃COOH +CH₃COONa

Q.2 . Buffer solutions are not used for..

a. qualitative analysis of mixture

b. qualitative analysis of mixtures

c. digestion of food

d. purification of water

Q.3 . Which of the following equations represents the correct form of Henderson's Equation?

a. $pH =log\ K_a+log\ \frac{[Salt]}{[Acid]}$

b. $pH =log\ K_a+log\ \frac{[Acid]}{[Salt]}$

c. $pH =log\ K_a-log\ \frac{[Salt]}{[Acid]}$

d. $pH =log\ K_b-log\ \frac{[Salt]}{[Acid]}$

Q.4 . Buffer solution can not be formed by

a. mixing a weak acid with its salt with a strong base.

b. by mixing a weak base with its salt with a strong acid.

c. by a solution of protic solvents.

d. by a solution of ampholyte.

Q.5. Which of the following buffer is present in our blood plasma?

a. Acetic acid + Sodium Acetate

b. Carbonic acid + bicarbonate

c. Boric acid + Borax,

d. Phthalic acid + Potassium acid phthalate

Answer key:

Q.1	Q.2	Q.3	Q.4	Q.5
c	d	a	d	b

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