Acids, Bases and pH

Introduction

As-salamu alaykum, students. I am Dr. Fatima Malik. In this crucial A Level topic, we move beyond the simple definitions of acids and bases you learned in O Level/IGCSE. We will explore the Brønsted-Lowry theory, which provides a more sophisticated understanding of proton transfer reactions. This is the foundation for quantifying acidity using the pH scale and understanding the vital difference between strong and weak acids.

The concepts of pH, buffers, and acid-base equilibria are not just abstract theories; they are fundamental to life and industry. From the delicate pH balance in our blood that keeps us alive, to the industrial processes at ENGRO Fertilisers in Daharki or the quality control of pharmaceuticals at GSK in Karachi, a deep understanding of this chapter is essential. Mastering these principles will equip you with the analytical skills required for success in Paper 2, Paper 3 titrations, and the more complex problems in Paper 4.

Core Theory

The Brønsted-Lowry theory defines acid-base interactions in terms of proton (H⁺ ion) transfer.

* Brønsted-Lowry Acid: A proton (H⁺) donor.

* Brønsted-Lowry Base: A proton (H⁺) acceptor.

Consider the reaction of ethanoic acid with water:

CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq)

Acid 1 Base 2 Base 1 Acid 2

Here, CH₃COOH donates a proton, so it's the acid. H₂O accepts the proton, so it's the base. In the reverse reaction, H₃O⁺ donates a proton (acid) and CH₃COO⁻ accepts it (base). Pairs of species like CH₃COOH/CH₃COO⁻ and H₃O⁺/H₂O are called conjugate acid-base pairs.

Strong vs. Weak Acids and Bases

The strength of an acid depends on its degree of dissociation in water.

* Strong acids (e.g., HCl, HNO₃, H₂SO₄) fully dissociate. For 0.1 mol dm⁻³ HCl, [H⁺] is also 0.1 mol dm⁻³.

HCl(aq) → H⁺(aq) + Cl⁻(aq)

* Weak acids (e.g., CH₃COOH, HCN) only partially dissociate, establishing an equilibrium.

CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)

The pH Scale and K_w

The pH scale is a convenient logarithmic measure of [H⁺]:

pH = -log₁₀[H⁺] and conversely, [H⁺] = 10⁻pH

Water itself undergoes auto-ionisation: H₂O(l) ⇌ H⁺(aq) + OH⁻(aq).

The equilibrium constant for this is the ionic product of water, K_w.

K_w = [H⁺][OH⁻] = 1.0 x 10⁻¹⁴ mol²dm⁻⁶ (at 298 K)

In neutral water, [H⁺] = [OH⁻] = 1.0 x 10⁻⁷ mol dm⁻³, so pH = 7.

We can also define pK_w = -log₁₀(K_w) = 14. Thus, pH + pOH = 14.

Calculations for Weak Acids: K_a and pK_a

For a generic weak acid, HA ⇌ H⁺ + A⁻, the acid dissociation constant, K_a, is:

K_a = ([H⁺][A⁻]) / [HA]

A larger K_a value means a greater extent of dissociation and a stronger acid. We often use pK_a = -log₁₀(K_a). A smaller pK_a indicates a stronger acid.

To calculate the pH of a weak acid, we make two key assumptions:

1. The dissociation of the acid is so small that [HA] at equilibrium is approximately the same as the initial concentration: [HA]eqm ≈ [HA]initial.
2. The contribution of H⁺ ions from the auto-ionisation of water is negligible.

This means [H⁺] ≈ [A⁻], so the K_a expression simplifies to:

K_a ≈ [H⁺]² / [HA]initial

From this, we can find [H⁺] and then calculate the pH.

Buffer Solutions

A buffer solution is a solution that resists changes in pH upon the addition of small amounts of acid or alkali. It consists of a mixture of a weak acid and its conjugate base (e.g., CH₃COOH and CH₃COONa) or a weak base and its conjugate acid (e.g., NH₃ and NH₄Cl).

Consider an acidic buffer (CH₃COOH / CH₃COO⁻):

CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)

The reservoir of weak acid (CH₃COOH) and conjugate base (CH₃COO⁻) is high.

* Add acid (H⁺): The conjugate base reacts with the added H⁺. Equilibrium shifts left.

CH₃COO⁻(aq) + H⁺(aq) → CH₃COOH(aq)

* Add alkali (OH⁻): The weak acid reacts with the added OH⁻. Equilibrium shifts right to replace the H⁺ consumed.

CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l)

The pH of a buffer can be calculated using the Henderson-Hasselbalch equation:

pH = pK_a + log₁₀([A⁻] / [HA])

Titration Curves & Indicators

Plotting pH against the volume of titrant added gives a titration curve.

* Strong Acid - Strong Base: Equivalence point at pH 7. Steep vertical section.

* Weak Acid - Strong Base: Equivalence point at pH > 7 (due to hydrolysis of the salt formed, e.g., CH₃COO⁻). A buffer region exists before the equivalence point.

* Strong Acid - Weak Base: Equivalence point at pH < 7 (due to hydrolysis of the salt formed, e.g., NH₄⁺).

An indicator is a weak acid (HIn) where HIn and its conjugate base (In⁻) have different colours. It changes colour over a specific pH range. A suitable indicator must have its colour change range (pKin ± 1) fall within the steep vertical section of the titration curve. For a weak acid-strong base titration (equivalence point pH ≈ 8-9), phenolphthalein (range 8.3-10.0) is suitable. For a strong acid-weak base titration (equivalence point pH ≈ 5-6), methyl red (range 4.2-6.3) is suitable.