How Acid-Base Reactions Transfer Protons Between Molecules

Blood pH 7.35 — A Margin of 0.4 Units Between Life and Death

Human arterial blood is maintained at pH 7.35–7.45. A drop to pH 7.1 causes cardiac arrhythmia. A rise to pH 7.8 triggers convulsions. The body devotes elaborate biochemical machinery — three distinct buffer systems — to holding blood pH within this narrow range. Acid-base chemistry, at its core, governs not just laboratory reactions but the survival of every living cell. Understanding it begins with the simplest possible view: acid-base reactions are proton transfer events.

Three complementary theories define acids and bases at increasing levels of generality. The Arrhenius definition limits acids to substances that release H⁺ in water and bases to substances that release OH⁻. The Brønsted-Lowry definition, more broadly applicable, defines acids as proton donors and bases as proton acceptors. The Lewis definition, broadest of all, defines acids as electron-pair acceptors and bases as electron-pair donors.

Brønsted-Lowry Theory and Conjugate Pairs

In the Brønsted-Lowry framework, every acid-base reaction involves a proton transfer from an acid (HA) to a base (B):

HA + B ⇌ A⁻ + BH⁺

The acid HA donates a proton and becomes its conjugate base A⁻. The base B accepts a proton and becomes its conjugate acid BH⁺. The stronger the acid, the weaker its conjugate base, and vice versa.

• Hydrochloric acid (HCl) is a strong acid; its conjugate base Cl⁻ is an extremely weak base with negligible tendency to recapture a proton.
• Acetic acid (CH₃COOH) is a weak acid; its conjugate base acetate (CH₃COO⁻) is a moderately strong base.
• Water acts as both acid and base (it is amphoteric): it accepts a proton from HCl and donates a proton to ammonia (NH₃).
• The autoionisation of water — H₂O + H₂O ⇌ H₃O⁺ + OH⁻ — has an equilibrium constant K_w = 1.0 × 10⁻¹⁴ at 25°C, giving pH 7.00 for pure neutral water.

The pH Scale

pH is defined as pH = −log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in moles per litre. The scale typically runs from 0 to 14, though values outside this range are possible for concentrated strong acids and bases.

Strong and Weak Acids

Strong acids dissociate completely in water; weak acids dissociate partially. This distinction is quantified by the acid dissociation constant Ka.

For a weak acid HA ⇌ H⁺ + A⁻:

Ka = [H⁺][A⁻] / [HA]

• Acetic acid has Ka = 1.8 × 10⁻⁵; only about 1.3% of acetic acid molecules dissociate in a 0.1 M solution.
• Hydrocyanic acid (HCN) has Ka = 6.2 × 10⁻¹⁰ — a very weak acid; solutions of HCN are nearly pH 5–6.
• The six common strong acids — HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄ — are essentially fully dissociated at all practical concentrations.
• pKa = −log₁₀(Ka); lower pKa means stronger acid. Formic acid (pKa 3.74) is a stronger acid than acetic acid (pKa 4.75).

Neutralisation and Titration

When a strong acid meets a strong base in stoichiometric amounts, the reaction produces water and a salt: HCl + NaOH → NaCl + H₂O, ΔH ≈ −57 kJ/mol. The driving force is the formation of water from H⁺ and OH⁻ ions.

Acid-base titration exploits neutralisation analytically. A solution of known concentration (the titrant) is added to a solution of unknown concentration (the analyte) until the equivalence point — where moles of acid equal moles of base. An indicator changes colour near the equivalence point: phenolphthalein shifts from colourless to pink around pH 8.2–10; methyl orange shifts from red to yellow around pH 3.1–4.4.

Buffer Solutions

Buffers resist pH change upon addition of acid or base. They consist of a weak acid and its conjugate base (or weak base and conjugate acid) in comparable concentrations.

The Henderson-Hasselbalch equation describes buffer pH: pH = pKa + log([A⁻]/[HA]). A buffer is most effective when [A⁻] ≈ [HA], giving pH ≈ pKa. Blood's bicarbonate buffer maintains pH 7.4 through the equilibrium CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻, regulated by the lungs (controlling CO₂) and kidneys (controlling HCO₃⁻). This coupled physiological-chemical system is among the most elegant buffer mechanisms in nature.