Enthalpy, Entropy, and Gibbs Free Energy - SS3 Chemistry Lesson Note

Enthalpy, entropy, and Gibbs free energy are fundamental concepts in thermodynamics and play a crucial role in understanding chemical reactions and physical processes. Let's delve into each of these concepts and explore their significance:

1.    Enthalpy (H):

Enthalpy is a thermodynamic property representing a system's total heat content at constant pressure. It is denoted by the symbol "H." In chemistry, most reactions occur at constant pressure, making enthalpy a particularly useful quantity for studying chemical reactions. The enthalpy change (ΔH) of a reaction is defined as the difference between the enthalpy of the products and the enthalpy of the reactants.

●     If ΔH is positive, the reaction is endothermic, meaning it absorbs heat from the surroundings.

●     If ΔH is negative, the reaction is exothermic, meaning it releases heat to the surroundings.

Enthalpy is an extensive property, which means it depends on the quantity of the substance involved.

2.    Entropy (S):

Entropy is a measure of the degree of disorder or randomness in a system. It is denoted by the symbol "S" and is also an extensive property. Entropy can be thought of as the spreading out of energy and matter within a system.

●     If the entropy of a system increases (ΔS > 0), the system becomes more disordered.

●     If the entropy of a system decreases (ΔS < 0), the system becomes more ordered.

Entropy is a key factor in determining the spontaneity of a process. The Second Law of Thermodynamics states that in any spontaneous process, the total entropy of the universe always increases.

3.    Gibbs Free Energy (G):

Gibbs free energy is a thermodynamic potential that combines both enthalpy and entropy to predict the spontaneity and maximum work obtainable from a system. It is denoted by the symbol "G."

●     ΔG = ΔH - TΔS

where:

ΔG = Change in Gibbs free energy

ΔH = Change in enthalpy

ΔS = Change in entropy

T = Temperature (in Kelvin)

Gibbs free energy helps us determine the direction of a reaction or a process:

●     If ΔG < 0, the reaction is spontaneous in the forward direction (exergonic).

●     If ΔG > 0, the reaction is nonspontaneous in the forward direction (endergonic).

●     If ΔG = 0, the system is at equilibrium.

For a reaction to be spontaneous, it must have both a negative ΔH (exothermic) and a positive ΔS (increase in disorder).

Overall, these three thermodynamic properties provide critical insights into the behaviour of chemical systems. They help us understand the energy changes associated with reactions, the tendency of processes to occur spontaneously, and the conditions under which reactions reach equilibrium. This knowledge is invaluable in various fields of chemistry and is essential for rationalising and predicting the behaviour of matter under different conditions.