pH (Grade 12 NSC Matric Physical Sciences): Revision Notes

pH

What is pH?

pH is a measurement that tells us how acidic or basic (alkaline) a solution is. The term pH stands for "power of hydrogen" and was first introduced by Danish biochemist Søren Peter Lauritz Sørensen in 1909.

The pH scale ranges from 0 to 14 and works as follows:

• pH < 7: The solution is acidic
• pH = 7: The solution is neutral (pure water)
• pH > 7: The solution is basic or alkaline

Many everyday items have specific pH values. For example, popular soft drinks are quite acidic, with pH values between 2.5 and 3.7. This explains why these drinks can be harmful to tooth enamel over time.

Understanding pH helps us make sense of the world around us. From the citric acid in lemon juice (pH 3.14) to the sodium hydroxide in caustic soda (pH 13), pH measurements help us classify and understand different substances.

pH calculations

The pH of a solution can be calculated using logarithmic equations. The concentration of hydrogen ions [H⁺] and hydronium ions [H₃O⁺] can be used interchangeably in these calculations.

Key formula:

• $\text{pH} = -\log[\text{H}^+]$
• $\text{pH} = -\log[\text{H}_3\text{O}^+]$

The brackets [ ] represent concentration in mol.dm⁻³.

This table shows the relationship between pH values and ion concentrations. Notice how as pH increases (becomes more basic), the hydrogen ion concentration [H⁺] decreases while the hydroxide ion concentration [OH⁻] increases.

pH indicator strips like these are practical tools for quickly measuring pH in laboratories and everyday situations.

Water ionisation and Kw

Water can act as both an acid and a base through a process called auto-protolysis or auto-ionisation. This means water molecules can transfer protons between themselves.

The equation for water ionisation is: $\text{H}_2\text{O}(\ell) + \text{H}_2\text{O}(\ell) \rightleftharpoons \text{H}_3\text{O}^+(\text{aq}) + \text{OH}^-(\text{aq})$

In this reaction:

• One water molecule acts as an acid (proton donor)
• Another water molecule acts as a base (proton acceptor)
• Conjugate pairs are formed: H₂O/OH⁻ and H₂O/H₃O⁺

The equilibrium constant for this process is called Kw:

• $K_w = [\text{H}_3\text{O}^+][\text{OH}^-]$
• At 25°C: $K_w = 1 \times 10^{-14}$

This relationship is crucial for understanding pH calculations and the behaviour of aqueous solutions.

Salt hydrolysis

When salts dissolve in water, they can affect the pH of the solution. The final pH depends on the strength of the original acid and base that formed the salt.

The general rules for salt hydrolysis are:

Strong acid + strong base

• Result: Neutral solution (pH ≈ 7)
• Example: HCl + NaOH → NaCl + H₂O
• The salt does not significantly affect water's pH

Weak acid + strong base

• Result: Basic solution (pH ≈ 9)
• Example: CH₃COOH + NaOH → CH₃COONa + H₂O
• The salt makes the solution slightly basic

Strong acid + weak base

• Result: Acidic solution (pH ≈ 5)
• Example: HCl + NH₃ → NH₄Cl + H₂O
• The salt makes the solution slightly acidic

Indicators

Indicators are chemical compounds that change colour depending on the pH of the solution they're in. They are essential tools for determining when chemical reactions reach completion, especially in titrations.

Different indicators work best at different pH ranges and show distinct colour changes:

Common indicators and their uses:

Bromothymol blue:

• pH range: 6.0 - 7.6
• Colours: Yellow (acidic) → Green (neutral) → Blue (basic)
• Best for: Strong acid + strong base titrations

Phenolphthalein:

• pH range: 8.3 - 10.0
• Colours: Colourless (acidic) → Faint pink (neutral) → Pink (basic)
• Best for: Weak acid + strong base titrations

Bromocresol green:

• pH range: 3.8 - 5.4
• Colours: Yellow (acidic) → Green (neutral) → Blue (basic)
• Best for: Strong acid + weak base titrations