Buffer solution - Ionic Equilibrium | Physical Chemistry PDF Download

Buffer:

A solution that resists change in pH value upon addition of small amount of strong Acid or Base (less than 1%) is called Buffer solution.

Characteristics of Buffer Solution:

• Constant pH range

• Resist any change

• Addition of small amount of acid, base dilution and temperature

Types of buffer solutions

(A)       Simple Buffer solution 

(B)       Mixed Buffer solution 

(A)       Simple buffer solution: A solution of weak Acid and weak base in water e.g. CH₃COONH₄, NH₄CN.

Buffer action of simple buffer solution

Consider a simple buffer solution of CH₃COONH₄, since it is a salt which dissociate completely.

            CH₃COONH₄→ CH₃COO^– + NH₄⁺

If a strong acid such as HCl is added then

            HCl→H⁺ + Cl^–

The H⁺ ions from the added acid (HCl) combine with CH₃COO^– ions of form CH₃COOH_, which is a weak acid so will not further ionized.

Thus, there is a rise in H⁺ ion concentration and pH remains constant

CH₃COO^– + H⁺ CH₃COOH (Weak acid)

(B)  Mixed Buffer solutions:

Acidic buffer solution:

An acidic buffer solution consists of solution of a weak acid and its alt with strong base. The best example is mixture of solution of acetic acid and its salt with strong base. (CH₃COONa).

                        CH₃COOH  CH₃COO^– + H⁺

                                C₁                        -              -

At equation     C₁ – C₁α             C₁α          C₁α

                        CH₃COOH → CH₃COO^– + Na⁺

                                C₂                     -             -

                                -                 C₂            C₂

New equilibrium

                        CH₃COOH   CH₃COO^– + H⁺

                        C₁ - C₁α₁               C₁α¹          C₁α¹

                        C₁a¹ + C₂≈C₂

                        1 - α¹ ≈ 1

                        Kα = C₂α¹

                      
                   pH = -log [H⁺]

This equation is called Handerson equation.

Where C₂ = Concentration of common ion

            C₁ = Concentration of weak acid

Example: Find pH of solution contains 0.1 M (100 ml) CH₃COOH and mixed with 0.1 M, 50 ml NaOH. K_a = 10^–5.

Solution: Here, an acid and base is mixed. So, firstly acid – Base Reaction take place.

MIlimoles of CH₃COOH = 10 mmol

Milimoles of NaOH = 5 mmol

CH₃COOH + NaOH→ CH₃COO^– Na⁺

      10                5       

   5/150              -                     5/150

Now, solution is having weak acid CH₃COOH and it’s salt having common ion. So, this solution is buffer solution Þ Acidic Buffer

pH = pKa + log

pH = pKa + log

pH = 5

Basic Buffer Solution

A basic buffer solution consists of a mixture of a weak base and its salt with strong acid. The best known example is a mixture of NH₄OH and NH₄Cl.

NH₄OH   NH₄⁺ + OH^–                         (weakly ionised)

NH₄Cl    → NH₄⁺ + Cl^–                      (highly ionized)

When a fewe drops of Base (NaOH) are added, the OH^– ions from NaOH combine with NH₄⁺ to form feebly ionized NH₄OH thus there is no rise in the concentration of OH^– ions and hence the pH value Remains constant.

NH₄⁺ + OH^–→ NH₄OH

pH of basic buffer solution 

                                    NH₄OH  NH₄⁺ + OH^–

                                        C₁                     -           -

                                    C₁ – C₁aα      C₁α      C₁α

                                    NH₄Cl   →  NH₄⁺ + Cl^–

                                        C₂            -          -

                                        -               C₂           C₂

At new equilibrium

                                    NH₄OH   NH₄⁺ + OH^–

                                    C₁ – Cα¹          C₁α¹    C₁α¹

          K_b =  

1 - α¹≈ 1

C₁α¹ + C₂ ≈ C₂

            K_b = C₂α¹

          [OH^–] = C₁α¹ =

Example: 0.1M, 100 ml NH₄Cl reacts  with 0.1M, 20 ml Ca(OH)_2. Calculate the pH of solution if k_b is 10^–5.

Solution :     2 NH₄Cl + Ca(OH)₂→ 2 NH₄OH + CaCl₂

Milimoles        0.1 × 100                  0.1 × 20             

                        10                              2

                        6                                -                  4

Concentration 6                                                   4

                        120                                            120

Now, solution is having weak base (NH₄OH) and its salt NH₄Cl which is having ion. Hence this solution acts as buffer

            pOH = pK_b +

            pOH = 5.17

            pH = 8.83

Example: 0.1 M, 100 ml NH₄OH is titrated with 0.1 M, H₂SO₄ solution. Calculate the pH if kb = 10^–5

1. Nothing is added

2. 25 ml, H₂SO₄ is added

3. 0.1 M, 20 ml NaOH is mixed in resulting solution for made in part b.

4. Find the pH if 0.1M, 10 ml H₂SO₄ is added in part b. 

Solution

a)         NH₄OH  NH₄⁺ OH^–

            C₁ – C₁α       C₁α C₁α

            Kb =

After neglection

                        k_b = C₁α²

                        α = 10^–2

                        pH = 11

b)         When an acid H₂SO₄ is added with NH₄OH (a Base) then acid-Base Reaction takes place.

                        2 NH₄OH + H₂SO₄→ (NH₄)₂SO₄

milimole               10               2.5                –

                               5               –                2.5

            pOH = 5 + log

            pH = 14 – pOH

            pH = 9

c)         In part b we are having a weak base and its salt when NaOH is added it will React with (NH₄)₂SO₄.

                        (NH₄)₂SO₄ + 2 NaOH→ 2 NH₄OH + Na₂SO₄

                             2.5                  2

Milimoles            1.5                   -                   2

Concentration             1.5/145                -                   2/145

Now, the solution is having weak base and it’s salt, hence Buffer

For (NH₄)₂SO₄→ 2 NH₄⁺ + SO₄^–

Hence, pOH = 5 +     

pH = 4.632

In part b, when H₂SO₄ is added it will react with NH₄OH.

2NH₄OH + H₂SO₄→ (NH₄)₂ SO₄ + H₂O

      5                1

      3                –                  1

Total milimoles of (NH₄)₂ SO₄ = 7

Total milimoles of NH₄OH = 3

pOH = 5 + 

pH = 8.63

Buffer Capacity or Buffer Index:

It is the measure of power to resist change in pH of Buffer Solution.

It is defined as the number of moles of acid or Base require by one litre of a Buffer solution for changing its pH by one unit.

Buffer Capacity =

Buffer Range:

A solution can act as buffer solution only if ratio of concentration of salt to base or acid is from 1 to 10.

pH = pKa + 1

pH = pKa – 1

So, pH range is pKa+ 1