Buffer Solutions | Ionic Equilibrium Class 11

Buffer Solutions

Resisting Change in pH | Ionic Equilibrium

1. What is a Buffer Solution?

                    Definition: A solution which resists any change in its pH value even when small amounts of strong acid or base are added to it, or when it is diluted.
                

This resistance is known as Buffer Action.

2. Types of Buffer Solutions

A. Acidic Buffer

A mixture of a Weak Acid and its Salt with a Strong Base.

Example: $CH_3COOH$ (Weak Acid) + $CH_3COONa$ (Salt).
pH Range: Generally $pH < 7$.
Mechanism:
- Added $H^+$ is consumed by the conjugate base ($CH_3COO^-$) to form weak acid.
- Added $OH^-$ is consumed by the weak acid ($CH_3COOH$) to form water.

B. Basic Buffer

A mixture of a Weak Base and its Salt with a Strong Acid.

Example: $NH_4OH$ (Weak Base) + $NH_4Cl$ (Salt).
pH Range: Generally $pH > 7$.

3. Henderson-Hasselbalch Equation

Used to calculate the pH of buffer solutions.

For Acidic Buffer:

$$ pH = pK_a + \log \frac{[\text{Salt}]}{[\text{Acid}]} $$

For Basic Buffer:

$$ pOH = pK_b + \log \frac{[\text{Salt}]}{[\text{Base}]} $$ $$ pH = 14 - pOH $$

Note: $[\text{Salt}]$ refers to the concentration of the conjugate ion.

4. Buffer Capacity ($\beta$)

The ability of a buffer to resist pH change. It is defined as the number of moles of strong acid or base required to change the pH of 1 liter of buffer solution by 1 unit.

$$ \beta = \frac{dn}{d(pH)} $$

Maximum Buffer Capacity: Occurs when $[\text{Salt}] = [\text{Acid}]$ (or $[\text{Base}]$). In this case, $pH = pK_a$ (or $pOH = pK_b$).

Buffer Range: $pH = pK_a \pm 1$. Effective when ratio of Salt:Acid is between 1:10 and 10:1.

5. Natural Buffers

Blood Buffer: Human blood maintains a pH of approx 7.4 using the Carbonic Acid - Bicarbonate buffer system ($H_2CO_3 / HCO_3^-$).

Practice Quiz

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