Practical Example: Calculating Buffer pH

This section provides a practical example of how to calculate the pH of a buffer solution, which is a common type of question in A Level Chemistry buffer calculations.

Example: Ethanoic acid buffer

Given:

• 125 cm³ of 0.8 mol dm⁻³ ethanoic acid (CH3COOH)
• 75 cm³ of 1 mol dm⁻³ sodium ethanoate (CH3COONa)
• Ka of ethanoic acid = 1.74 × 10⁻⁵ mol dm⁻³

Step 1: Calculate new concentrations after mixing

[CH3COOH] = (0.8 × 125) / 200 = 0.5 mol dm⁻³ [CH3COO⁻] = (1 × 75) / 200 = 0.375 mol dm⁻³

Step 2: Use the equation [H+] = Ka × [HA] / [A-]

[H+] = (1.74 × 10⁻⁵ × 0.5) / 0.375 = 2.32 × 10⁻⁵ mol dm⁻³

Step 3: Calculate pH

pH = -log(2.32 × 10⁻⁵) = 4.63

Highlight: This example demonstrates how to prepare a buffer solution with a specific pH and how to calculate pH of buffer solution given molarity and volume.

Example: If you needed to prepare a pH 4 buffer solution, you would adjust the ratio of acid to conjugate base to achieve this pH.

Understanding these calculations is crucial for mastering buffers in A Level Chemistry OCR and other exam boards.

Understanding Buffer Solutions

Buffer solutions play a crucial role in maintaining stable pH levels in various chemical and biological systems. This section explores the fundamental concepts of buffer solutions and their mechanisms.

Definition: A buffer solution is a system that minimizes pH changes when small amounts of an acid or alkali are added.

The key components of a buffer solution are:

1. A weak acid (HA)
2. Its conjugate base (A-)

Highlight: Buffer solutions require a large reservoir of both HA and A- to effectively resist pH changes.

The mechanism of buffer action can be understood in two scenarios:

1. When acid is added:

   • H+ ions are introduced to the solution
   • The large reservoir of A- reacts with excess H+
   • The equilibrium shifts to the left
   • [H+] and pH are restored

2. When alkali is added:

   • OH- ions are introduced
   • H+ + OH- → H2O reaction occurs
   • [H+] decreases
   • HA dissociates
   • The equilibrium shifts to the right
   • [H+] and pH are restored

Example: In a buffer solution containing acetic acid (CH3COOH) and acetate ion (CH3COO-), adding a small amount of HCl would cause the acetate ions to neutralize the added H+ ions, maintaining the pH.