🧪 Relative Strengths of Acids and Bases

The strength of an acid or base is determined by the extent to which it ionises in aqueous solution. A strong acid donates protons completely (or nearly completely) to water, producing high concentrations of hydronium ions (\(H_3O^+\)). A weak acid ionises only partially, establishing an equilibrium between the molecular form and ions. The same principle applies to bases: strong bases dissociate fully, while weak bases reach an equilibrium.

Here, HA is the acid, \(A^-\) its conjugate base, and \(H_3O^+\) the conjugate acid of water. The equilibrium constant for this reaction is the acid‑ionisation constant (\(K_a\)):

Similarly, for a base \(B\) reacting with water:

A larger \(K_a\) (or \(K_b\)) indicates a stronger acid (or base). Often, we use the p\(K_a\) = –log \(K_a\) scale: the smaller the p\(K_a\), the stronger the acid.

📌 Factors Affecting Acid Strength

Several intrinsic properties determine how readily an acid donates a proton:

• Bond strength (H–A): Weaker bonds break more easily, favouring ionisation and increasing acidity. This dominates when comparing acids within the same group of the periodic table (e.g., HF, HCl, HBr, HI).
• Polarity of the H–A bond: A more polar bond makes the hydrogen more partially positive, easing proton loss. This is the primary factor when comparing elements in the same period (e.g., CH₄, NH₃, H₂O, HF).
• Atomic size of A: Larger atoms form longer, weaker bonds, enhancing acidity (e.g., HI > HBr > HCl > HF).
• Stability of the conjugate base (A⁻): If the conjugate base is stabilised by resonance, electronegativity, or solvation, the acid is stronger.

🔁 Predicting the Direction of Acid‑Base Reactions

A fundamental rule governs any acid‑base equilibrium: the reaction always favours the side with the weaker acid and weaker base. In other words, the proton (\(H^+\)) will always bind to the stronger base. Consequently, the equilibrium lies toward the formation of the weaker acid‑base pair.

Examples:

• Hydrochloric acid (HCl, strong) reacts completely with water: HCl + H₂O → H₃O⁺ + Cl⁻ (equilibrium far right).
• Acetic acid (CH₃COOH, weak) establishes an equilibrium: CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻ (left‑favoured).
• Ammonia (NH₃, weak base) with water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (equilibrium left, because OH⁻ is a strong base and NH₄⁺ is a stronger acid than water).

🧮 Understanding pH, pOH, pKa, pKb, and Related Constants

The acidity of a solution is measured by pH = –log[H₃O⁺]. Similarly, pOH = –log[OH⁻]. At 25 °C, water self‑ionises: Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴, and pH + pOH = 14.

For any conjugate pair (HA / A⁻), the relationship is:

📊 Table of Common Strong and Weak Acids (with Ka values)

⚖️ Relationship Between Ka and Kb for Conjugate Pairs

For a weak acid HA and its conjugate base A⁻, the product of their equilibrium constants equals the ion product of water:

This allows calculation of one constant if the other is known. For example, the acetate ion (CH₃COO⁻) has \(K_b = K_w / K_a(\text{acetic acid}) = 10^{-14} / (1.8\times10^{-5}) = 5.6\times10^{-10}\).

📚 Summary of Key Equations & Relations