The pH scale is a logarithmic scale used to measure the acidity or alkalinity of a solution. It ranges from 0 to 14, where:

• pH < 7 — the solution is acidic.
• pH = 7 — the solution is neutral.
• pH > 7 — the solution is basic (alkaline).

The pH scale is based on the concentration of hydrogen ions in solution. The lower the pH, the higher the hydrogen ion concentration.

The pH of a solution is defined as the negative logarithm (base 10) of the hydrogen ion concentration:

$$ \text{pH} = -\log [\text{H}^+] $$

Similarly, the pOH of a solution is the negative logarithm of the hydroxide ion concentration:

$$ \text{pOH} = -\log [\text{OH}^-] $$

The relationship between pH and pOH at 25°C is given by:

$$ \text{pH} + \text{pOH} = 14 $$

In pure water or any aqueous solution at 25°C:

$$ [\text{H}^+] \times [\text{OH}^-] = 1 \times 10^{-14} $$

Hence, knowing either hydrogen ion or hydroxide ion concentration allows calculation of the other.

Example 1: An orange juice is found to have a pH of 3.4. Find its pOH.

Using the relationship: $$ \text{pH} + \text{pOH} = 14 $$ $$ \text{pOH} = 14 - 3.4 = 10.6 $$ The pOH of the orange juice is 10.6.

Example 2: Find the pH and pOH of the following solutions:

1. 0.001 M HCl solution

$$ [\text{H}^+] = 0.001 = 1 \times 10^{-3} $$ $$ \text{pH} = -\log [\text{H}^+] = 3 $$ $$ \text{pOH} = 14 - 3 = 11 $$

2. 0.001 M H₂SO₄ solution

Each molecule of H₂SO₄ produces 2 H⁺ ions: $$ [\text{H}^+] = 2 \times 0.001 = 0.002 = 2 \times 10^{-3} $$ $$ \text{pH} = -\log(2 \times 10^{-3}) = 2.70 $$ $$ \text{pOH} = 14 - 2.70 = 11.30 $$

3. 0.01 M NaOH solution

$$ [\text{OH}^-] = 0.01 = 1 \times 10^{-2} $$ $$ \text{pOH} = -\log [\text{OH}^-] = 2 $$ $$ \text{pH} = 14 - 2 = 12 $$

Example 3: A sample of acid was found to have a pH of 6.3. Find:

1. pOH
2. [H⁺]
3. [OH⁻]

(a) Using the relationship: $$ \text{pH} + \text{pOH} = 14 $$

(b) Hydrogen ion concentration: $$ [\text{H}^+] = 10^{-\text{pH}} $$

(c) Hydroxide ion concentration: $$ [\text{OH}^-] = 10^{-\text{pOH}} $$

Example 4: Calculate the pH of a solution where \( [\text{H}^+] = 2.5 \times 10^{-2} \, \text{mol/dm}^3 \).

Example 5: The hydroxide ion concentration of a solution is \( 4.0 \times 10^{-3} \, \text{mol/dm}^3 \). Calculate the pH.

The pH of a solution is a vital chemical property that affects biological processes, industrial operations, and environmental balance. Below are five major areas where pH plays an important role:

• 1. In Agriculture: The pH of soil affects the solubility and availability of essential nutrients for plant growth. Example: Most crops grow best in slightly acidic soils (pH 6–7).
• 2. In the Human Body: The pH of blood must be maintained around 7.4 for normal body function. Deviation can cause health problems such as acidosis or alkalosis.
• 3. In Industry: pH control is important in manufacturing products such as soap, paper, textiles, and pharmaceuticals to ensure quality and safety.
• 4. In Food Preservation: The pH of food affects its shelf life and taste. Example: Vinegar (acidic, pH ≈ 3) helps prevent bacterial growth in pickled foods.
• 5. In Water Treatment: Maintaining the correct pH in drinking water and swimming pools prevents corrosion and promotes effective disinfection.