ΔG

Definition

ΔG, or Gibbs free energy change, is a thermodynamic quantity that represents the maximum amount of non-expansion work that can be extracted from a closed system under constant temperature and pressure conditions. It is a crucial concept in understanding chemical equilibria and the spontaneity of chemical reactions.

5 Must Know Facts For Your Next Test

1. The sign of ΔG determines the spontaneity of a reaction: a negative ΔG indicates a spontaneous reaction, while a positive ΔG indicates a non-spontaneous reaction.
2. ΔG is calculated as ΔG = ΔH - TΔS, where ΔH is the change in enthalpy and ΔS is the change in entropy.
3. At equilibrium, ΔG = 0, meaning that the system has reached a state where no further change will occur without the input of external energy.
4. ΔG is directly related to the equilibrium constant (K) of a reaction through the equation ΔG = -RT ln K, where R is the universal gas constant and T is the absolute temperature.
5. The magnitude of ΔG reflects the driving force of a reaction, with larger negative values indicating a more favorable and spontaneous process.

Review Questions

• Explain how the sign of ΔG determines the spontaneity of a chemical reaction.
  • The sign of ΔG, the Gibbs free energy change, determines the spontaneity of a chemical reaction. A negative ΔG indicates that the reaction is spontaneous and will occur naturally, as the system is moving towards a more stable state with a lower free energy. Conversely, a positive ΔG means the reaction is non-spontaneous and will not occur naturally without the input of external energy to drive the process forward.
• Describe the relationship between ΔG, ΔH, and ΔS, and how this relationship can be used to predict the spontaneity of a reaction.
  • The relationship between ΔG, ΔH, and ΔS is given by the equation ΔG = ΔH - TΔS, where T is the absolute temperature. The spontaneity of a reaction is determined by the relative magnitudes and signs of ΔH and ΔS. A negative ΔH (exothermic reaction) and a positive ΔS (increase in disorder) will both contribute to a negative ΔG, making the reaction spontaneous. Conversely, a positive ΔH (endothermic reaction) and a negative ΔS (decrease in disorder) will both contribute to a positive ΔG, making the reaction non-spontaneous.
• Explain how ΔG is related to the equilibrium constant (K) of a chemical reaction, and discuss the significance of this relationship.
  • The Gibbs free energy change, ΔG, is directly related to the equilibrium constant, K, of a chemical reaction through the equation ΔG = -RT ln K, where R is the universal gas constant and T is the absolute temperature. This relationship is significant because it allows us to predict the spontaneity and extent of a reaction at equilibrium. A negative ΔG indicates a spontaneous reaction with a large equilibrium constant, while a positive ΔG indicates a non-spontaneous reaction with a small equilibrium constant. The magnitude of ΔG also reflects the driving force of the reaction, with larger negative values corresponding to more favorable and spontaneous processes.

Related terms

Enthalpy (ΔH): The total energy released or absorbed during a chemical reaction, including the energy required to break and form chemical bonds.

Entropy (ΔS): A measure of the disorder or randomness of a system, which increases as a system moves towards a more disordered state.

Spontaneity: The tendency of a process to occur naturally without the input of external energy, as determined by the sign of the Gibbs free energy change.