Principles of Biochemical Thermodynamics

Gibbs Free Energy

For a process at constant temperature T and pressure, the Gibbs free energy change is defined as ΔG = ΔH - TΔS, where ΔH is the enthalpy change and ΔS is the entropy change. This relation links heat flow, disorder, and whether a reaction can proceed without external work at fixed T and P.

The sign of ΔG classifies the direction of spontaneity: a negative ΔG means the process is spontaneous as written - exergonic, releasing free energy. A positive ΔG means the forward reaction is non-spontaneous unless energy is supplied - endergonic. At equilibrium, ΔG = 0 for that reaction under the prevailing conditions.

Standard vs Physiological Free Energy

ΔG° (standard Gibbs energy change) refers to defined standard states: typically all solutes at 1 M (except H⁺, often taken as 1 M in physical chemistry tables), pure liquids, gases at 1 bar, and commonly 25 °C (298 K). In metabolism, we almost always use the biochemical standard, denoted ΔG°', where pH = 7.0 (so [H⁺] = 10^-7 M in the biochemical standard state) and water is assigned activity corresponding to 55.5 M. Tabulated values for pathways are therefore ΔG°' values unless stated otherwise.

Actual driving force depends on concentrations through the reaction quotient Q. The key equation is:

Here R is the gas constant and T is absolute temperature. Q matters in cells because metabolites are rarely at 1 M: a reaction with slightly positive ΔG°' can still have negative ΔG if product concentrations are kept low and substrate concentrations high, and the opposite can reverse flux even when ΔG°' is negative on paper.

Coupled Reactions

Cells make endergonic steps feasible by coupling them to strongly exergonic processes. The most common currency is ATP hydrolysis (or sometimes GTP, ion gradients, or group-transfer from high-potential donors). Thermodynamically, the overall process must have ΔG < 0 for the combined net reaction.

A classic example is hexokinase-catalyzed glucose phosphorylation: direct phosphorylation of glucose by Pi is endergonic, but coupling to ATP → ADP + P_i makes the net conversion of glucose to glucose-6-phosphate favorable under cellular conditions when enzyme binding and catalysis prevent futile hydrolysis.

Equilibrium and Free Energy

At equilibrium, ΔG = 0, which links standard free energy to the equilibrium constant:

When K_eq > 1 (products favored at standard concentration ratios), ΔG°' is negative. When Q = K_eq, the reaction quotient matches equilibrium and ΔG = 0 - there is no net driving force until concentrations change.

What You Will Learn in This Course

This five-chapter course connects thermodynamics, ATP chemistry, amino acids, and protein conformation:

1. Biochemical Thermodynamics (this page) - ΔG, ΔG°', Q, and equilibrium
2. ATP & Phosphoryl Group Transfer Potential - energy currency and transfer ladders
3. Amino Acid Ionization & pI - charge, pK_a, and isoelectric point
4. Secondary Structure & Torsion Angles - φ, ψ, and backbone geometry
5. Ramachandran Plot & Folding - allowed conformations and stability

Reinforce concepts with the Bioenergetics Game or consolidate topics in the Study Guide.