Thermodynamics & Chemical Equilibrium

Chemistry is fundamentally the science of energy and matter. In the health sciences, a working knowledge of chemical thermodynamics is indispensable: it explains why metabolic reactions proceed spontaneously in one direction, how the body maintains physiological pH, and why drugs bind their targets with measurable affinity. This lesson covers Gibbs free energy, chemical equilibrium, Le Chatelier's principle, and the chemistry of pH and biological buffers.

Thermodynamics describes the energy changes that accompany chemical and physical processes. The first law -- conservation of energy -- tells us that the total energy of the universe is constant; energy can be converted from one form to another but cannot be created or destroyed. The enthalpy change of a reaction (deltaH) reflects heat exchanged with the surroundings at constant pressure. Exothermic reactions (deltaH < 0) release heat; endothermic reactions (deltaH > 0) absorb heat. However, enthalpy alone does not determine whether a reaction will occur spontaneously.

The second law introduces entropy (S), a measure of the dispersal of energy and matter. Spontaneous processes increase the total entropy of the universe (deltaS_universe > 0). Many biological processes involve a decrease in local order (e.g., protein unfolding), but the overall entropy of the universe still increases because heat is released to the surroundings.

Gibbs Free Energy -- the master variable of biochemical thermodynamics -- combines enthalpy and entropy into a single criterion for spontaneity at constant temperature and pressure: deltaG = deltaH - TdeltaS. A negative deltaG (exergonic reaction) indicates spontaneous reaction under the stated conditions; a positive deltaG (endergonic) is non-spontaneous. When deltaG = 0, the system is at equilibrium. The standard Gibbs free energy change (deltaGdegrees) refers to conditions where all reactants and products are at 1 mol/L activity; the biochemical standard (deltaGdegrees') uses pH 7 as the reference. deltaGdegrees relates to the equilibrium constant Keq by: deltaGdegrees = -RT ln Keq, where R = 8.314 J mol-1 K-1 and T is absolute temperature. A large Keq (strongly product-favoured) corresponds to a large negative deltaGdegrees. In cells, the actual deltaG depends on both deltaGdegrees and the prevailing concentrations of reactants and products; the cell actively maintains metabolites far from equilibrium to drive biosynthetic and energy-transducing reactions in the desired direction.

ATP hydrolysis (ATP + H2O -> ADP + Pi) has a deltaGdegrees' of approximately -30.5 kJ/mol under standard biochemical conditions, but the actual deltaG in a living cell is closer to -50 kJ/mol because ATP/ADP ratios are maintained far from equilibrium. This large negative deltaG drives otherwise unfavourable endergonic reactions by coupling -- for example, the phosphorylation of glucose by hexokinase (deltaGdegrees' ~= +16.7 kJ/mol for glucose phosphorylation alone) is rendered spontaneous when coupled to ATP hydrolysis (net deltaGdegrees' ~= -16.7 kJ/mol).

Chemical Equilibrium describes the state in which the forward and reverse rates of a reaction are equal and no net change in concentration occurs. For a general reaction aA + bB ? cC + dD, the equilibrium constant Keq = [C]^c[D]^d / [A]^a[B]^b. A large Keq (>>1) means the reaction lies far to the right (product-favoured); a small Keq (<<1) means it lies to the left (reactant-favoured).

Le Chatelier's Principle states that if an equilibrium system is subjected to a stress -- a change in concentration, pressure, or temperature -- the system will shift to partially counteract that stress. For example: adding substrate to an enzyme-catalysed reaction shifts the equilibrium toward product formation; increasing CO2 in the blood shifts the bicarbonate equilibrium (CO2 + H2O ? H+ + HCO3-) to the right, lowering pH. This principle underpins respiratory physiology: hyperventilation lowers pCO2, shifting the equilibrium left, raising blood pH (respiratory alkalosis).

pH and the Henderson-Hasselbalch Equation. The pH scale is a logarithmic measure of hydrogen ion activity: pH = -log[H+]. Normal arterial blood pH is 7.35-7.45. An acid is a proton donor (Br?nsted-Lowry definition); a base is a proton acceptor. The acid dissociation constant Ka = [H+][A-]/[HA]; pKa = -log Ka. A weak acid is only partially dissociated; the ratio of its conjugate base to acid at any pH is described by the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]).

Biological Buffers resist changes in pH by absorbing or releasing protons. The most important physiological buffer is the bicarbonate system: H2CO3 ? H+ + HCO3- (pKa ~= 6.1). Although the pKa is not at 7.4, the bicarbonate buffer is supremely effective because both components are independently regulated: the lungs control pCO2 (and hence H2CO3) and the kidneys control HCO3- reabsorption. The phosphate buffer system (pKa ~= 6.8) is the dominant intracellular buffer. Proteins, via histidine residues (pKa ~= 6.0), also buffer blood. Understanding acid-base chemistry is essential for interpreting blood gases, managing diabetic ketoacidosis (excess ketoacid production), and understanding how drugs such as aspirin (a weak acid) distribute across biological membranes based on pH gradients and the Henderson-Hasselbalch equation.