Chemical Thermodynamics - SS3 Chemistry Past Questions and Answers - page 2

Enthalpy, Entropy, and Gibbs Free Energy in Chemical Thermodynamics:

1.    Enthalpy (H):

Enthalpy is a thermodynamic property representing a system's total heat content at constant pressure. It accounts for the internal energy (U) of a system and the energy required to overcome the pressure-volume work during a process. Mathematically, enthalpy is expressed as:

H = U + PV

where H is the enthalpy, U is the internal energy, P is the pressure, and V is the volume.

2.    Entropy (S):

Entropy is a measure of the degree of randomness or disorder in a system. It represents the distribution of energy among different energy states. In an isolated system, entropy tends to increase over time due to the tendency of energy to be distributed more randomly. Entropy is denoted by the symbol S.

3.    Gibbs Free Energy (G):

Gibbs free energy is a thermodynamic potential that combines the effects of both enthalpy and entropy to predict the spontaneity and feasibility of a chemical reaction under given conditions. The Gibbs free energy change (ΔG) for a reaction is calculated as:

ΔG = ΔH - TΔS

where ΔG is the change in Gibbs free energy, ΔH is the change in enthalpy, ΔS is the change in entropy, and T is the temperature in Kelvin.

Significance in Determining Spontaneity and Feasibility of Chemical Reactions:

The concepts of enthalpy, entropy, and Gibbs free energy are essential for understanding the spontaneity and feasibility of chemical reactions. For a reaction to occur spontaneously (without the need for external intervention), the change in Gibbs free energy (ΔG) must be negative (ΔG < 0).

1.    ΔG < 0 (Negative ΔG): If ΔG is negative, the reaction is spontaneous, and the system proceeds in the forward direction under the given conditions. The reaction is thermodynamically favourable, as it releases energy and increases the system's disorder (entropy). The reaction is exergonic, meaning it can do work on the surroundings.

2.    ΔG > 0 (Positive ΔG): If ΔG is positive, the reaction is nonspontaneous, and the system does not proceed in the forward direction under the given conditions. The reaction is thermodynamically unfavourable, as it requires an input of energy to occur. The reaction is endergonic, meaning it requires energy from the surroundings to proceed.

3.    ΔG = 0: If ΔG is zero, the reaction is at equilibrium, and the system is in a state of minimum free energy. The reaction proceeds neither in the forward nor the reverse direction.

Examples:

1.    Combustion of Methane (CH4):

●     CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

●     ΔG < 0: The negative ΔG indicates that the combustion of methane is a spontaneous and exergonic reaction, releasing energy.

2.    Melting of Ice:

●     H2O(s) → H2O(l)

●     ΔG < 0: The negative ΔG shows that the melting of ice is a spontaneous process at temperatures above the melting point.

3.    Photosynthesis:

●     6CO2(g) + 6H2O(l) → C6H12O6(aq) + 6O2(g)

●     ΔG > 0: The positive ΔG indicates that photosynthesis is a non-spontaneous process that requires an input of energy.

In conclusion, enthalpy, entropy, and Gibbs free energy are critical concepts in chemical thermodynamics. They provide valuable insights into the spontaneity and feasibility of chemical reactions, allowing us to predict and understand the direction of reactions and their energy changes under different conditions. These concepts have significant applications in various fields, including chemistry, biology, environmental science, and engineering.