For a reaction with a negative ΔG^∘ value, which of the following statements is false? (a) The e

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For a reaction with a negative ΔG^∘ value, which of the...

For a reaction with a negative $\Delta G^{\circ}$ value, which of the following statements is false? (a) The equilibrium constant $K$ is greater than one, (b) the reaction is spontaneous when all the reactants and products are in their standard states, and (c) the reaction is always exothermic.

Key Concepts

Exothermic vs. Exergonic

Exothermic reactions release heat and are characterized by a negative enthalpy change (?H), whereas exergonic reactions have a negative Gibbs free energy change (?G). A reaction can be exergonic (spontaneous) without being exothermic if the entropy change (?S) is sufficiently positive to drive the reaction, meaning that negative ?G does not always imply the reaction is exothermic.

Equilibrium Constant

The equilibrium constant (K) is a measure of the relative concentrations of reactants and products at equilibrium. The relationship between Gibbs free energy and the equilibrium constant is given by the equation ?G° = -RT ln K, which shows that a negative ?G° corresponds to a K value greater than one, indicating that the products are favored at equilibrium.

Gibbs Free Energy

Gibbs free energy is a thermodynamic potential that determines the maximum amount of reversible work obtainable from a chemical reaction at constant temperature and pressure. A negative ?G° indicates that the reaction can occur spontaneously under standard conditions, meaning the free energy of the products is lower than that of the reactants.

Spontaneity of a Reaction

Spontaneity in chemical reactions refers to whether a reaction can occur without the input of additional energy. A reaction with a negative ?G° is considered spontaneous in its standard state, making it thermodynamically favorable. It is important to note that spontaneity is about the energy change and not necessarily the reaction rate.

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