Energy Changes in Chemical Reactions - SS1 Chemistry Past Questions and Answers - page 5

Spontaneity in chemical reactions refers to the tendency of a reaction to occur without any external influence. It is determined by the energy changes that take place during the reaction. Gibbs free energy (ΔG) is a thermodynamic quantity that measures the energy available to do useful work in a system and is directly related to the spontaneity of a reaction.

 

The relationship between spontaneity and Gibbs free energy is given by the following equation:

 

ΔG = ΔH - TΔS

 

Where:

- ΔG is the change in Gibbs free energy

- ΔH is the change in enthalpy

- T is the temperature in Kelvin

- ΔS is the change in entropy

 

For a reaction to be spontaneous at a given temperature, the change in Gibbs free energy must be negative (ΔG < 0). This indicates that the reaction is energetically favourable, and the reaction will proceed spontaneously in the forward direction. Conversely, if ΔG is positive (ΔG > 0), the reaction is nonspontaneous, and an input of energy is required for the reaction to occur.

 

The significance of Gibbs free energy in determining spontaneity lies in its ability to account for both enthalpy and entropy changes in a system. ΔH represents the change in heat energy, while TΔS represents the change in the energy associated with the disorder or randomness of the system (entropy). The combination of these factors determines the overall change in Gibbs free energy.

Equilibrium in a chemical reaction refers to a state where the forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products over time. The position of equilibrium is determined by the Gibbs free energy (ΔG) of the system.

 

At equilibrium, the Gibbs free energy change (ΔG) is zero (ΔG = 0). This means that the system is at its lowest possible energy state and is in a state of dynamic balance. The relationship between ΔG and equilibrium is given by the following equation:

 

ΔG = ΔG° + RT ln(Q)

 

Where:

- ΔG is the change in Gibbs free energy

- ΔG° is the standard Gibbs free energy change

- R is the gas constant

- T is the temperature in Kelvin

- Q is the reaction quotient

 

If ΔG is negative (ΔG < 0), the reaction proceeds in the forward direction to reach equilibrium, favouring the formation of products. If ΔG is positive (ΔG > 0), the reaction proceeds in the reverse direction to reach equilibrium, favouring the formation of reactants. If ΔG is zero (ΔG = 0), the system is at equilibrium, and there is no net change in the concentrations of reactants and products.

 

The position of equilibrium is determined by the relative magnitudes of the forward and reverse Gibbs free energy changes. If the forward Gibbs free energy change is lower in magnitude than the reverse Gibbs free energy change (ΔGf < Gr), the equilibrium position will favour the reactants. Conversely, if the forward Gibbs free energy change is higher in magnitude than the reverse Gibbs free energy change (ΔGf > Gr), the equilibrium position will favour the products.