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-  H2\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,


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:

  1. When [Salt] = [Acid], i.e., pH = pKa for acid buffer

  2. When [Salt] = [Base], i.e., pOH = pKb 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)2COis 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, HPO2-)

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

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

​Refer to the folloiwng 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,

Ka =   H+A/[HA] 

It can be assumed that concentration of A- ions from complete ionisation of the saltBA 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.


\frac{[Salt]}{[Acid]}=10 , then

pH = 1 + pKa

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

pH = pKa -1

So weak acid may be used for preparing buffer solutions having pH values lying within the ranges pKa + 1 and pKa -1. The acetic acid gas a pKa 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,

Kb =  [B+][OH-]/[BOH]                      

or      [OH-] = Kb [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-]= Kb[Base]/[Salt]                         

or    pOH = log[Salt]/[Base]  log Kb

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,

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 folloiwng 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


You can also refer to 

 IIT JEE  Syllabus 

Reference Books for IIT JEE

Chemical Equilibrium

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