Acid-Base Chemistry & pH

pH is one of the most tightly regulated variables in the human body. Blood pH must stay within 7.35–7.45 — a range so narrow that deviations of even 0.1 pH units cause measurable physiological effects, and deviations of 0.3–0.4 units can be fatal. What is pH? pH = −log₁₀[H⁺] — the negative logarithm of the hydrogen ion (proton) concentration. This means: - Lower pH = more H⁺ = more acidic (e.g. stomach acid pH 1–2) - Higher pH = less H⁺ = more alkaline (e.g. pancreatic juice pH 8) - pH 7 is neutral; blood at pH 7.4 is very slightly alkaline - Because the scale is logarithmic, a change of 1 pH unit represents a 10-fold change in [H⁺]. This is why even small pH changes represent large changes in actual acid concentration. Brønsted-Lowry definitions: - Acid — a proton (H⁺) donor. Example: HCl (hydrochloric acid), H₂CO₃ (carbonic acid), lactic acid - Base — a proton acceptor. Example: HCO₃⁻ (bicarbonate), NH₃ (ammonia), OH⁻ (hydroxide) - Conjugate base — what remains after an acid donates its proton. Example: HCl → H⁺ + Cl⁻ (Cl⁻ is the conjugate base) Strong vs weak acids: - Strong acids fully dissociate in water (e.g. HCl, H₂SO₄) — release all their H⁺. Very low pH. - Weak acids partially dissociate (e.g. acetic acid, carbonic acid, lactic acid). They establish an equilibrium between dissociated and undissociated forms. Biologically important — most of the body's acids are weak acids, which allows fine-tuned buffering.

A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Buffers are essential for life — without them, even minor variations in metabolic acid production would cause lethal pH swings. How a buffer works: A buffer consists of a weak acid and its conjugate base (or a weak base and its conjugate acid) in equilibrium: HA ⇌ H⁺ + A⁻ When acid (H⁺) is added: H⁺ + A⁻ → HA — the base "mops up" the extra protons When base (OH⁻) is added: OH⁻ + HA → A⁻ + H₂O — the acid donates protons to neutralise the base In both cases, the pH changes much less than it would in an unbuffered solution. The Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]) This equation describes the relationship between pH, the dissociation constant (pKa), and the ratio of conjugate base to acid. A buffer works best when [A⁻] ≈ [HA] — i.e. when pH ≈ pKa. The body's three main buffer systems: 1. Bicarbonate buffer (most important in blood): CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ pKa = 6.1, but works effectively at blood pH 7.4 because the system is open — the lungs continuously remove CO₂. This makes bicarbonate the dominant extracellular buffer despite its pKa being far from 7.4. 2. Phosphate buffer (important in urine and intracellularly): H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻ (pKa = 6.8 — closer to physiological pH). Used in kidneys to carry excess acid in urine. 3. Protein buffer (haemoglobin and plasma proteins): Amino acid side chains (histidine residues particularly) can accept or donate protons. Haemoglobin is a particularly important buffer within red blood cells.

The bicarbonate buffer is the body's most important extracellular buffering system — it accounts for over 50% of buffering capacity in blood. Its power lies not just in the chemistry but in the fact that both components are physiologically controlled: The system: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ This reaction is catalysed by carbonic anhydrase — one of the fastest enzymes known. It is found abundantly in red blood cells, kidney tubule cells, and the stomach. CO₂ as an acid: CO₂ dissolves in blood → forms carbonic acid (H₂CO₃) → immediately dissociates to H⁺ and HCO₃⁻. This is why rising blood CO₂ acidifies the blood and falling CO₂ makes it more alkaline. Dual physiological control: - Lungs regulate CO₂ (and therefore carbonic acid) — faster response (seconds to minutes) - Breathe faster → blow off more CO₂ → less carbonic acid → pH rises - Breathe slower → retain CO₂ → more carbonic acid → pH falls - Kidneys regulate HCO₃⁻ — slower but more powerful and precise (hours to days) - Retain more HCO₃⁻ → raises pH - Excrete more HCO₃⁻ → lowers pH - Can also generate new HCO₃⁻ and excrete H⁺ in urine This dual control system means the body can compensate for acid-base disturbances with great precision: - Respiratory acidosis (high CO₂) → kidneys retain HCO₃⁻ to compensate - Metabolic acidosis (low HCO₃⁻) → lungs hyperventilate to blow off CO₂ - Respiratory alkalosis (low CO₂) → kidneys excrete HCO₃⁻ - Metabolic alkalosis (high HCO₃⁻) → lungs hypoventilate to retain CO₂

Acid-base disorders are classified by two parameters: direction (acidosis vs alkalosis) and mechanism (respiratory vs metabolic). Respiratory acidosis (pH↓, CO₂↑, HCO₃⁻ compensatory↑): Caused by hypoventilation — CO₂ builds up in blood. - Acute: opiate overdose (respiratory depression), COPD exacerbation, severe asthma, Guillain-Barré - Chronic: COPD (body adapts over years; HCO₃⁻ rises as renal compensation) Compensation: Kidneys retain HCO₃⁻ (takes 3–5 days) Respiratory alkalosis (pH↑, CO₂↓, HCO₃⁻ compensatory↓): Caused by hyperventilation — CO₂ is blown off. - Anxiety/panic attacks (most common cause in young people) - Altitude (low O₂ drives faster breathing) - Pregnancy (progesterone stimulates breathing) - Early salicylate (aspirin) poisoning — directly stimulates respiratory centre Compensation: Kidneys excrete HCO₃⁻ Metabolic acidosis (pH↓, HCO₃⁻↓, CO₂ compensatory↓): Low bicarbonate — either acid is being produced or bicarbonate is lost. - High anion gap acidosis — acid accumulation: DKA (ketones), lactic acidosis (sepsis, shock), uraemia (kidney failure), salicylate toxicity, methanol/ethylene glycol poisoning - Anion gap = Na⁺ − (Cl⁻ + HCO₃⁻). Normal = 8–12 mmol/L. Raised when unmeasured anions (ketones, lactate) accumulate. - Normal anion gap acidosis — bicarbonate loss: severe diarrhoea (losing HCO₃⁻), renal tubular acidosis, IV normal saline excess Compensation: Lungs hyperventilate (Kussmaul breathing) — immediate Metabolic alkalosis (pH↑, HCO₃⁻↑, CO₂ compensatory↑): High bicarbonate — either acid is lost or base is gained. - Prolonged vomiting (losing HCl from gastric acid) - Diuretics (loop/thiazide — cause Cl⁻ and H⁺ loss) - Excess antacids or bicarbonate Compensation: Lungs hypoventilate Reading an ABG (Arterial Blood Gas): 1. Is pH acidotic (<7.35) or alkalotic (>7.45)? 2. Is CO₂ in the right direction to explain it? → respiratory cause 3. Is HCO₃⁻ in the right direction to explain it? → metabolic cause 4. Is the "other" component compensatory (moving opposite to normalise pH)? 5. Is compensation appropriate or is there a mixed disorder?

pH has a profound effect on enzyme function — and therefore on virtually every biochemical process in the body. This is why maintaining blood pH within 7.35–7.45 is so critical. Why pH affects enzymes: Enzyme activity depends on the precise 3D structure of the protein — particularly the active site. Amino acid side chains in the active site must be in the correct ionisation state to bind substrate and catalyse the reaction. Changing pH alters the ionisation of these side chains (histidine, aspartate, glutamate, lysine, cysteine are particularly sensitive) → the active site changes shape → enzyme activity is lost. pH optima: Every enzyme has a characteristic pH optimum — the pH at which it is maximally active: - Pepsin (stomach protease) — pH optimum 1.5–2.0. Functions in the highly acidic stomach - Salivary amylase — pH optimum 6.7–7.0. Rapidly inactivated by stomach acid - Pancreatic enzymes (trypsin, lipase, amylase) — pH optimum 7.5–8.5. The pancreas secretes bicarbonate to neutralise stomach acid in the duodenum, creating the right environment - Cytoplasmic enzymes (most metabolic enzymes) — pH optimum ~7.2–7.4 (intracellular pH) - Lysosomal enzymes — pH optimum ~4.5–5.0 (lysosomes are acidic compartments) Clinical consequences of pH on enzyme function: - In severe acidosis (pH <7.1): enzyme function across the body is compromised — cardiac muscle weakens, neurological function deteriorates, coagulation fails - In severe alkalosis (pH >7.6): enzyme hyperactivity, overexcitable nerves and muscles (tetany, seizures), cardiac arrhythmias - Aspirin overdose causes a complex mixed acid-base picture — respiratory alkalosis (aspirin directly stimulates the respiratory centre) followed by metabolic acidosis (aspirin and its metabolites are acids). The resulting disruption of cellular enzymes causes hyperthermia, coma, and death if untreated.