Spontaneity and Equilibrium - SS3 Chemistry Lesson Note
Spontaneity and equilibrium are fundamental concepts in chemistry, particularly in the study of thermodynamics and chemical reactions. Let's explore these concepts in more detail:
1. Spontaneity:
Spontaneity refers to the tendency of a process to occur without external intervention or assistance. In the context of chemical reactions, it indicates whether a reaction will proceed on its own under specific conditions.
a. Spontaneous Processes:
A spontaneous process occurs naturally, driven by the system's internal energy and without the need for a continuous input of energy. The spontaneity of a process is determined by the change in Gibbs free energy (ΔG). As mentioned in the previous note, the Gibbs free energy is a combination of enthalpy (ΔH) and entropy (ΔS) changes:
ΔG = ΔH - TΔS
where:
ΔG = Change in Gibbs free energy
ΔH = Change in enthalpy
ΔS = Change in entropy
T = Temperature (in Kelvin)
● If ΔG < 0, the process is spontaneous in the forward direction (exergonic), and the reaction proceeds spontaneously.
● If ΔG > 0, the process is nonspontaneous in the forward direction (endergonic), and the reaction will not proceed spontaneously.
● If ΔG = 0, the system is at equilibrium, and the reaction is at a dynamic balance.
Spontaneous processes tend to move towards lower energy states and higher entropy (greater disorder). It's important to note that a spontaneous process may occur slowly or rapidly, depending on the activation energy barrier that must be overcome.
b. Activation Energy:
Even though a process may be spontaneous, some reactions may have a high activation energy barrier that needs to be surpassed before the reaction can proceed. Activation energy is the minimum energy required for reactant molecules to collide effectively and lead to the formation of products. Catalysts are substances that lower the activation energy and increase the rate of a reaction without being consumed themselves.
Equilibrium:
Chemical equilibrium is a state reached in a reversible reaction when the rate of the forward reaction is equal to the rate of the reverse reaction. At equilibrium, the concentrations of reactants and products remain constant over time, but the system is still dynamic with both forward and reverse reactions occurring simultaneously.
a. Dynamic Equilibrium:
Dynamic equilibrium is achieved when the rates of the forward and reverse reactions become equal. At this point, the net change in concentrations of reactants and products is zero. The reaction does not stop; instead, the forward and reverse reactions proceed at the same rate, leading to a constant concentration of all species involved.
b. Equilibrium Constant (K):
For a general reversible reaction:
aA + bB ⇌ cC + dD
The equilibrium constant, K, is defined as the ratio of the concentrations of products to reactants, each raised to their respective stoichiometric coefficients:
K = [C]c [D]d / [A]a [B]b
K is a constant at a specific temperature and represents the position of the equilibrium.
● If K > 1, the equilibrium lies towards the product side, and products are favoured.
● If K < 1, the equilibrium lies towards the reactant side, and reactants are favoured.
● If K = 1, the concentrations of reactants and products are approximately equal, and the reaction is close to completion.
Understanding spontaneity and equilibrium is crucial for predicting the behaviour of chemical systems, designing chemical processes, and interpreting the results of chemical reactions. These concepts are foundational in many branches of chemistry, including chemical kinetics, thermodynamics, and chemical equilibrium.