ΔG°

Definition

ΔG° is the standard Gibbs free energy change, which represents the maximum amount of useful work that can be extracted from a spontaneous chemical reaction under standard conditions. It is a fundamental concept in chemical equilibria and the determination of equilibrium constants.

5 Must Know Facts For Your Next Test

1. The sign of ΔG° determines the spontaneity of a reaction: negative values indicate a spontaneous process, while positive values indicate a non-spontaneous process.
2. The magnitude of ΔG° indicates the driving force of the reaction, with more negative values corresponding to greater spontaneity and energy release.
3. ΔG° is related to the equilibrium constant (K) through the equation: ΔG° = -RT ln K, where R is the universal gas constant and T is the absolute temperature.
4. Reactions with a large, negative ΔG° value tend to have a large equilibrium constant, meaning the products are favored at equilibrium.
5. ΔG° can be used to predict the direction of a reaction and the extent to which it will proceed towards equilibrium.

Review Questions

• Explain how the sign and magnitude of ΔG° relate to the spontaneity and driving force of a chemical reaction.
  • The sign of ΔG° determines the spontaneity of a reaction. A negative value of ΔG° indicates that the reaction is spontaneous and will occur naturally, while a positive value indicates a non-spontaneous reaction that requires an input of energy. The magnitude of ΔG° reflects the driving force of the reaction - the more negative the value, the greater the spontaneity and energy release. Reactions with large, negative ΔG° values have a strong tendency to proceed towards the products, while those with positive ΔG° values are less favorable and require an input of energy to occur.
• Describe the relationship between ΔG° and the equilibrium constant (K) of a chemical reaction.
  • The standard Gibbs free energy change (ΔG°) and the equilibrium constant (K) of a reaction are related through the equation ΔG° = -RT ln K, where R is the universal gas constant and T is the absolute temperature. This relationship allows us to determine the equilibrium constant of a reaction from its ΔG° value, or vice versa. Reactions with a large, negative ΔG° value will have a correspondingly large equilibrium constant, meaning the products are strongly favored at equilibrium. Conversely, reactions with a positive ΔG° value will have a small equilibrium constant, indicating the reactants are more stable under standard conditions.
• Explain how ΔG° can be used to predict the direction and extent of a chemical reaction towards equilibrium.
  • The sign and magnitude of ΔG° provide valuable information about the spontaneity and driving force of a chemical reaction. If ΔG° is negative, the reaction will occur spontaneously and proceed towards the products. The more negative the value of ΔG°, the greater the spontaneity and the further the reaction will proceed towards equilibrium. Conversely, if ΔG° is positive, the reaction is non-spontaneous and will not occur naturally. In this case, an input of energy is required to drive the reaction forward. Additionally, the relationship between ΔG° and the equilibrium constant (K) allows us to predict the position of equilibrium - reactions with large, negative ΔG° values will have correspondingly large equilibrium constants, indicating the products are strongly favored at equilibrium.

Related terms

Gibbs Free Energy:

A thermodynamic potential that combines the internal energy, pressure-volume work, and entropy of a system to determine the spontaneity and maximum work potential of a chemical process.

Standard Conditions: The reference state where temperature is 25°C (298 K) and pressure is 1 atm (101.325 kPa), used to tabulate thermodynamic data like ΔG°.

Equilibrium Constant (K): The ratio of the concentrations of the products to the reactants at equilibrium, which is related to the Gibbs free energy change of the reaction.