Acids, Bases & Buffers

Introduction to Acid-Base Chemistry in Biology

Acid-base chemistry is central to biological function. The activity of enzymes, the folding of proteins, the transport of gases, and the function of virtually every biological molecule depend on maintaining precise hydrogen ion concentrations. The pH of arterial blood (7.35–7.45) is maintained within this narrow range by sophisticated buffer systems and physiological compensation mechanisms.

Brønsted-Lowry Theory

According to the Brønsted-Lowry definition, an acid is a proton (H⁺) donor and a base is a proton acceptor. This is more general than the Arrhenius definition and applies in all solvents.

Every acid has a conjugate base formed when it donates a proton:

• Acetic acid (CH₃COOH) donates H⁺ → conjugate base acetate (CH₃COO⁻)
• Carbonic acid (H₂CO₃) donates H⁺ → bicarbonate (HCO₃⁻)
• Ammonium (NH₄⁺) donates H⁺ → ammonia (NH₃)

Strong acids (e.g., HCl, H₂SO₄) dissociate completely in water; [H⁺] ≈ [acid]₀.

Weak acids (e.g., acetic acid, carbonic acid) establish an equilibrium with their conjugate base.

Acid Dissociation Constant (Ka) and pKa

For a weak acid HA ⇌ H⁺ + A⁻, the equilibrium constant is:

Ka = [H⁺][A⁻] / [HA]

The pKa = −log₁₀(Ka). A lower pKa means a stronger acid (higher tendency to donate protons).

Biologically important pKa values:

• Carbonic acid (H₂CO₃): pKa₁ = 6.1 (effective, including dissolved CO₂)
• Phosphoric acid (H₃PO₄): pKa₂ = 7.2 (the relevant form at physiological pH)
• Histidine side chain (imidazole): pKa ≈ 6.0
• Aspartate/glutamate side chains: pKa ≈ 3.7–4.1
• Lysine ε-amino: pKa ≈ 10.5
• Cysteine thiol: pKa ≈ 8.3

The Henderson-Hasselbalch Equation

Rearranging the Ka expression gives the Henderson-Hasselbalch equation:

pH = pKa + log₁₀([A⁻]/[HA])

or equivalently: pH = pKa + log₁₀([Base]/[Acid])

This equation is enormously useful because it allows calculation of:

1. The pH of a buffer given the ratio of conjugate acid to base
2. The ionisation state of a molecule at a given pH
3. The ratio needed to achieve a target pH

Key insight: When [A⁻] = [HA] (equal concentrations of acid and conjugate base), log(1) = 0, so pH = pKa at the half-equivalence point.

A buffer is most effective (has maximum buffer capacity) within ±1 pH unit of its pKa.

The Bicarbonate Buffer System

The most important buffer in blood is the bicarbonate system:

CO₂(dissolved) + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻

Using an effective pKa of 6.1 and applying Henderson-Hasselbalch:

pH = 6.1 + log([HCO₃⁻] / [CO₂])

Normal values in arterial blood:

• [HCO₃⁻] ≈ 24 mEq/L
• [CO₂] (dissolved) ≈ 1.2 mEq/L (= 0.03 × PaCO₂ of 40 mmHg)
• Ratio = 24/1.2 = 20
• pH = 6.1 + log(20) = 6.1 + 1.3 = 7.4 ✓

Why is this system effective despite pKa = 6.1 being 1.3 units below blood pH? Because both components are physiologically regulated: CO₂ by the lungs (ventilation) and HCO₃⁻ by the kidneys (reabsorption and generation). This makes the bicarbonate system an open buffer with virtually unlimited capacity.

Other Physiological Buffer Systems

Phosphate buffer: H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻, pKa = 7.2. Most important in the intracellular fluid and urine. Concentration too low in plasma to be a major blood buffer.

Protein buffers: histidine residues (pKa ≈ 6.0) and the terminal amino groups of haemoglobin are the major intracellular and blood protein buffers. Haemoglobin accounts for about one-third of blood buffering capacity. Deoxyhaemoglobin is a weaker acid than oxyhaemoglobin (Haldane effect), facilitating CO₂ transport.

Titration Curves

A titration curve plots pH vs. equivalents of base (or acid) added. The curve for a weak acid has a characteristic sigmoid shape:

• At the start (all HA), pH is low
• At the half-equivalence point: pH = pKa (inflection point of the curve)
• At the equivalence point: all converted to A⁻; pH jumps sharply
• The buffer region (relatively flat portion) spans ±1 pH unit around the pKa

For polyprotic acids (e.g., phosphoric acid with three pKa values: 2.1, 7.2, 12.4), the titration curve shows three distinct buffer regions.

Clinical Acid-Base Disorders

The four primary acid-base disorders and their immediate causes:

Metabolic acidosis (↓ pH, ↓ HCO₃⁻): excess acid production (lactic acidosis, ketoacidosis, renal failure) or loss of bicarbonate (diarrhoea). Respiratory compensation: hyperventilation ↓ PaCO₂ (Kussmaul breathing in DKA).

Metabolic alkalosis (↑ pH, ↑ HCO₃⁻): loss of acid (vomiting, gastric suction) or gain of HCO₃⁻ (excess antacid). Respiratory compensation: hypoventilation ↑ PaCO₂.

Respiratory acidosis (↓ pH, ↑ PaCO₂): hypoventilation from COPD, opioid overdose, neuromuscular disease. Renal compensation: ↑ HCO₃⁻ reabsorption and H⁺ excretion.

Respiratory alkalosis (↑ pH, ↓ PaCO₂): hyperventilation from anxiety, hypoxia, salicylate poisoning. Renal compensation: ↓ HCO₃⁻ reabsorption.

Anion Gap

The anion gap (AG) = [Na⁺] − ([Cl⁻] + [HCO₃⁻]). Normal: 8–12 mEq/L (representing unmeasured anions: albumin, phosphate, sulphate).

Elevated AG metabolic acidosis: accumulation of unmeasured acids (lactate, ketones, uraemic acids, salicylate, methanol metabolites). Mnemonic: MUDPILES (Methanol, Uraemia, DKA, Propylene glycol, Isoniazid, Lactic acidosis, Ethylene glycol, Salicylates).

Normal AG (hyperchloraemic) metabolic acidosis: loss of HCO₃⁻ (diarrhoea) or impaired H⁺ excretion (renal tubular acidosis).

The Kidney's Role in Acid-Base Balance

The kidneys regulate pH more slowly than the lungs (hours to days vs. minutes) but with greater precision. Mechanisms:

1. HCO₃⁻ reabsorption: proximal tubule reabsorbs ~85% of filtered HCO₃⁻ via carbonic anhydrase (CA) — CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻; H⁺ secreted into lumen combines with HCO₃⁻ to reform CO₂
2. New HCO₃⁻ generation: in the collecting duct, H⁺ is excreted with phosphate (titratable acidity) or ammonium (NH₄⁺), generating new HCO₃⁻
3. Ammonium excretion: glutamine → NH₄⁺ + HCO₃⁻ (new); NH₄⁺ excreted in urine; HCO₃⁻ returned to blood — major mechanism in chronic acidosis