The reaction of chloromethane with hydroxide ion at 30 °C has a ΔG° value of −21.7 kcal/mol. What is the equilibrium constant for the reaction?

Gibbs Free Energy (ΔG°) is a thermodynamic potential that measures the maximum reversible work obtainable from a thermodynamic system at constant temperature and pressure. A negative ΔG° value indicates that a reaction is spontaneous under standard conditions, meaning it can proceed without external energy input. In this case, the ΔG° of -21.7 kcal/mol suggests that the reaction of chloromethane with hydroxide ion is favorable.

The equilibrium constant (K) is a dimensionless number that expresses the ratio of the concentrations of products to reactants at equilibrium for a given reaction at a specific temperature. It is derived from the Gibbs Free Energy change, where a negative ΔG° corresponds to a K value greater than 1, indicating that products are favored at equilibrium. The relationship between ΔG° and K is given by the equation ΔG° = -RT ln(K), where R is the gas constant and T is the temperature in Kelvin.

Temperature plays a crucial role in chemical reactions, influencing both the rate and the position of equilibrium. In this context, the reaction occurs at 30°C, which must be converted to Kelvin for calculations. The temperature affects the kinetic energy of molecules, thereby impacting the likelihood of successful collisions between reactants. Understanding how temperature interacts with Gibbs Free Energy and the equilibrium constant is essential for predicting reaction behavior.