Chemical thermodynamics is the study of energy transformations in chemical systems. It provides the theoretical foundation for predicting whether reactions occur spontaneously, the extent of reactions at equilibrium, and the energy changes associated with chemical processes.
Fundamental Concepts
In thermodynamics, a system is the part of the universe under study, which can be open, closed, or isolated, while the surroundings are everything else. State functions such as P, V, T, U, H, S, and G depend only on the current state of the system, not the path taken to reach it. The First Law of Thermodynamics states that energy cannot be created or destroyed, only converted between forms: ΔU = q + w, where ΔU is the change in internal energy, q is heat, and w is work.
Enthalpy (H)
Enthalpy is defined as H = U + PV, and at constant pressure, ΔH = qp, meaning the heat absorbed or released equals the change in enthalpy. Exothermic reactions release heat (ΔH < 0), while endothermic reactions absorb heat (ΔH > 0). The standard enthalpy of formation (ΔHf°) is the enthalpy change when one mole of a compound is formed from its elements in their standard states, and Hess’s Law states that ΔH for a reaction is the sum of ΔHf° of products minus reactants. Bond dissociation energies can estimate reaction enthalpies: ΔH ≈ Σ(bond energies broken) - Σ(bond energies formed).
Entropy (S)
Entropy is a measure of disorder or randomness in a system, and the Second Law states that the total entropy of the universe always increases for a spontaneous process. Gases have higher entropy than liquids, which have higher entropy than solids, and dissolving a solute increases entropy. Standard molar entropy (S°) values are positive and increase with temperature and molecular complexity. ΔS° for a reaction equals ΣS°(products) - ΣS°(reactants).
Gibbs Free Energy (G)
Gibbs free energy combines enthalpy and entropy: G = H - TS, and at constant temperature and pressure, ΔG = ΔH - TΔS. When ΔG < 0, the process is spontaneous (exergonic); ΔG = 0 indicates equilibrium; and ΔG > 0 means the process is non-spontaneous (endergonic). The standard Gibbs free energy change (ΔG°) is related to the equilibrium constant by ΔG° = -RT ln K, so a large negative ΔG° corresponds to K >> 1. Note that ΔG is not the reaction rate — thermodynamics predicts spontaneity, not kinetics.
Temperature Dependence of Spontaneity
When ΔH is negative and ΔS is positive, the reaction is spontaneous at all temperatures. When ΔH is positive and ΔS is negative, the reaction is non-spontaneous at all temperatures. When ΔH is negative and ΔS is negative, the reaction is spontaneous at low temperatures where |TΔS| < |ΔH|. When ΔH is positive and ΔS is positive, the reaction is spontaneous at high temperatures where TΔS > ΔH. The temperature at which a reaction becomes spontaneous is T = ΔH/ΔS, where ΔG = 0.
Chemical Equilibrium
At equilibrium, the forward and reverse reaction rates are equal and ΔG = 0. The equilibrium constant K is temperature-dependent according to the van’t Hoff equation: d(ln K)/dT = ΔH°/RT2. Le Chatelier’s Principle states that a system at equilibrium responds to disturbances in concentration, pressure, or temperature by shifting to counteract the change.
Applications
Chemical thermodynamics is used for predicting reaction feasibility in industrial chemical synthesis; calculating the maximum work obtainable from fuel cells and batteries; understanding biochemical energy transduction, such as how ATP hydrolysis drives endergonic reactions; and designing processes with optimal temperature and pressure conditions for maximum yield.
