Thermodynamics — Energy and Entropy

Thermodynamics is the study of energy — how it flows, transforms, and determines whether chemical reactions will occur spontaneously. In medicine and biology, thermodynamic principles explain why ATP powers cellular work, why proteins fold, and why drugs bind to their targets. The three laws of thermodynamics: First Law — Energy is conserved. Energy cannot be created or destroyed, only converted from one form to another. The total energy of the universe is constant. In biological terms: the energy in glucose is converted to ATP, heat, and mechanical work — nothing is lost or gained. Second Law — Entropy always increases. In any spontaneous process, the total disorder (entropy) of the universe increases. This is why hot objects cool, why iron rusts, and why mixed gases don't spontaneously separate. Third Law — At absolute zero (0 K, −273.15°C), the entropy of a perfect crystal is zero. This sets the reference point for all entropy calculations.

Enthalpy (H) measures the heat content of a system at constant pressure — the conditions of most biological and chemical reactions. Change in enthalpy (ΔH): ΔH = H(products) − H(reactants) Exothermic reactions: ΔH is negative — heat is released to the surroundings. Example: burning glucose (cellular respiration), neutralisation reactions, combustion. The products have less energy than the reactants. Endothermic reactions: ΔH is positive — heat is absorbed from the surroundings. Example: photosynthesis, dissolving ammonium nitrate in water, cooking an egg. Standard enthalpy of formation (ΔHf°): the heat change when one mole of a compound is formed from its elements in their standard states. Tables of ΔHf° values allow calculation of ΔH for any reaction using Hess's Law. Hess's Law: the total enthalpy change of a reaction is the same regardless of the pathway taken — only the initial and final states matter. This allows calculation of ΔH for reactions that can't be measured directly. Calorimetry: enthalpy changes are measured experimentally using a calorimeter. The heat absorbed or released is calculated from: q = mcΔT where m = mass, c = specific heat capacity, ΔT = temperature change.

Entropy (S) measures the degree of disorder or randomness in a system. More precisely, entropy measures the number of possible microstates — the more ways a system can be arranged, the higher its entropy. Factors that increase entropy: - Dissolving a solid into solution (more dispersal of particles) - Heating a substance (more kinetic energy, more possible states) - Gas formation from solids or liquids (gases have far more freedom of movement) - Increasing the number of moles of gas in a reaction - Mixing two substances together Standard entropy change (ΔS°): ΔS° = ΣS°(products) − ΣS°(reactants) Positive ΔS° = increased disorder (more probable) Negative ΔS° = increased order (less probable) Biological significance: Protein folding involves a decrease in entropy — the unfolded protein has many possible configurations; the folded form has only one. This seems to violate the second law, but the folding releases heat that increases entropy in the surroundings more than the decrease in the protein itself. The universe's entropy still increases. The hydrophobic effect (non-polar groups clustering together in water) is largely entropy-driven — it releases water molecules from ordered hydration shells, massively increasing their entropy.

Gibbs free energy (G) combines enthalpy and entropy into a single value that predicts whether a reaction will occur spontaneously at constant temperature and pressure — the conditions of all biological systems. The Gibbs equation: ΔG = ΔH − TΔS Where T is the absolute temperature in Kelvin. Interpreting ΔG: ΔG < 0 (negative): reaction is spontaneous (exergonic) — it releases free energy ΔG > 0 (positive): reaction is non-spontaneous (endergonic) — it requires energy input ΔG = 0: system is at equilibrium — no net change The four possible combinations: 1. ΔH negative, ΔS positive → ΔG always negative → always spontaneous 2. ΔH positive, ΔS negative → ΔG always positive → never spontaneous 3. ΔH negative, ΔS negative → spontaneous only at low temperatures (when −TΔS is small) 4. ΔH positive, ΔS positive → spontaneous only at high temperatures (when TΔS > ΔH) ATP and coupled reactions: Many biological reactions have positive ΔG and cannot occur spontaneously. Cells couple these reactions to ATP hydrolysis (ΔG = −30.5 kJ/mol) so the overall ΔG becomes negative. This is how muscle contraction, active transport, and biosynthesis are powered. ΔG° and equilibrium: ΔG° = −RT ln K Where K is the equilibrium constant. A large K (products favoured) corresponds to a large negative ΔG°.

Thermodynamics has direct clinical relevance: Metabolic rate and calorimetry: The body's energy expenditure can be measured by direct calorimetry (measuring heat output) or indirect calorimetry (measuring O₂ consumption and CO₂ production). Basal metabolic rate (BMR) is the energy required to maintain basic physiological functions at rest. Clinical nutrition calculations rely on thermodynamic principles. Fever: Fever increases body temperature (T). Since ΔG = ΔH − TΔS, an increase in T shifts the spontaneity of reactions. Some enzyme-catalysed reactions become more or less favourable. High fevers (>40°C) can denature proteins — the thermal energy exceeds the energy stabilising the folded structure. Drug binding thermodynamics: The binding of a drug to its receptor involves changes in enthalpy (non-covalent bonds formed) and entropy (changes in freedom of movement). The best drugs often have favourable contributions from both. Drug designers now routinely measure ΔH and ΔS of binding — "enthalpy-entropy compensation" is a major challenge in drug design. Sterilisation: Autoclaving (steam under pressure, 121°C) kills microorganisms because the high temperature makes protein denaturation thermodynamically favourable — the free energy change for unfolding becomes negative at high temperatures. Cold chain: Vaccines and many drugs must be stored at low temperatures because chemical degradation reactions (hydrolysis, oxidation) are thermodynamically spontaneous but kinetically slow at low T. Raising T accelerates these reactions exponentially.