5.6.5 Buffers

What are Buffer Solutions?

A buffer solution resists changes in pH when small amounts of acid or base are added or when diluted slightly. Buffers are essential in many chemical and biological systems, as they help maintain a stable pH environment.

Types of Buffer Solutions

Buffers come in two main types: acidic buffers and basic buffers.

Acidic Buffers:

These contain a weak acid and its salt (such as ethanoic acid and sodium ethanoate).
They maintain a pH below 7 and resist pH changes upon the addition of small amounts of acid or base.

Basic Buffers:

These consist of a weak base and its salt (such as ammonia and ammonium chloride), maintaining a pH above 7.

Mechanism of Buffering Action

Buffering action relies on the equilibrium between a weak acid (or base) and its conjugate base (or conjugate acid).

When Acid is Added

Additional $\text{H}^+$ ions are introduced.
The conjugate base (e.g., $\text{A}^-$ ) reacts with $\text{H}^+$ ions, neutralizing them and preventing a significant pH drop.
The pH remains nearly constant.

When Base is Added

Additional $\text{OH}^-$ ions are introduced.
The $\text{OH}^-$ ions react with $\text{H}^+$ ions in the solution, lowering the $\text{H}^+$ concentration.
The equilibrium shifts, dissociating more of the weak acid to restore $\text{H}^+$ , thus keeping the pH stable.

Applications of Buffer Solutions

Biological Systems: Blood contains buffers to maintain a pH of around 7.4.
Industrial Processes: Buffers stabilize pH in products like shampoos and lotions.
Chemical Analysis: Buffers are used in titrations and electrochemical measurements to maintain consistent pH.

Calculating the pH of Acidic Buffers

For a buffer solution consisting of a weak acid ( $\text{HA}$ ) and its conjugate base ( $\text{A}^-$ ), the pH can be calculated using the Henderson-Hasselbalch equation:

\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)

Where:

$[\text{A}^-]$ is the concentration of the conjugate base (salt).
$[\text{HA}]$ is the concentration of the weak acid.
$\text{p}K_a = -\log K_a$ (the dissociation constant of the weak acid).

Example Calculations

pH of a Buffer Made by Mixing a Weak Acid and Its Salt

infoNote

Example: A buffer is prepared by dissolving 0.012 mol of sodium ethanoate in 100, cm $^3$ of 0.052 mol dm $^{-3}$ ethanoic acid.

$K_a$ for ethanoic acid is $1.74 \times 10^{-5}$ , mol dm $^{-3}$

Step 1: Determine $[\text{H}^+]$

[\text{H}^+] = \frac{K_a \times [\text{HA}]}{[\text{A}^-]}

= \frac{(1.74 \times 10^{-5}) \times (5.20 \times 10^{-3})}{0.012}

= :success[7.54 \times 10^{-6} \, \text{mol dm}^{-3}]

Step 2: Calculate pH

\text{pH} = -\log(7.54 \times 10^{-6}) = :success[5.12]

pH of a Buffer Prepared by Adding a Salt Solution to a Weak Acid Solution

infoNote

Example: A buffer is made by mixing $50.0 \, \text{cm}^3$ of a $0.0417 \, \text{mol dm}^{-3}$ salt solution with $150 \, \text{cm}^3$ of $0.0204 \, \text{mol dm}^{-3}$ of the weak acid.

$\text{p}K_a$ for the acid is $5.12$

Step 1: Convert $\text{p}K_a$ to $K_a$

K_a = 10^{-5.12}

= :success[7.58 \times 10^{-6}]

Step 2: Calculate $[\text{H}^+]$

[\text{H}^+] = \frac{K_a \times [\text{HA}]}{[\text{A}^-]}

= \frac{(7.58 \times 10^{-6}) \times (3.06 \times 10^{-3})}{2.085 \times 10^{-3}}

= :success[1.11 \times 10^{-5}]

Step 3: Find pH

\text{pH} = -\log(1.11 \times 10^{-5}) = :success[4.95]

Buffers Simplified Revision Notes for A-Level AQA Chemistry