Recap of Equilibrium Concepts - SS3 Chemistry Lesson Note
Chemical equilibrium is a crucial concept in chemistry that describes the state of a reversible chemical reaction where the rate of the forward reaction is equal to the rate of the reverse reaction. In such a state, the concentrations of reactants and products remain constant over time, but the system is dynamic, with both forward and reverse reactions occurring simultaneously. Here's a recap of the key equilibrium concepts in chemistry:
Dynamic Equilibrium:
Dynamic equilibrium is achieved when the rates of the forward and reverse reactions become equal. At this point, the system no longer appears to change on a macroscopic scale, even though the reactions continue at the molecular level. The concentrations of reactants and products reach a constant value, but the reaction has not stopped; it is merely proceeding in both directions at the same rate.
Equilibrium Constant (K):
The equilibrium constant, denoted as K, is a quantitative measure of the position of the equilibrium for a particular reaction at a given temperature. 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
The magnitude of K provides valuable information about the extent to which a reaction has reached equilibrium:
● If K > 1, the equilibrium lies towards the product side, indicating products are favoured.
● If K < 1, the equilibrium lies towards the reactant side, indicating reactants are favoured.
● If K = 1, the concentrations of reactants and products are approximately equal, and the reaction is close to completion.
Reaction Quotient (Q):
The reaction quotient, denoted as Q, is a measure of the relative concentrations of reactants and products at any given point during a reaction, regardless of whether the reaction has reached equilibrium. It is calculated using the same formula as the equilibrium constant, but with concentrations at a specific moment in time:
Q = [C]c [D]d / [A]a [B]b
By comparing the reaction quotient (Q) to the equilibrium constant (K), we can determine the direction in which the reaction will proceed to reach equilibrium:
● If Q = K, the reaction is at equilibrium.
● If Q < K, the reaction proceeds forward (towards the products) to reach equilibrium.
● If Q > K, the reaction proceeds in the reverse direction (towards the reactants) to reach equilibrium.
Le Chatelier's Principle:
Le Chatelier's Principle is a guiding principle that helps predict how a system at equilibrium responds to changes in temperature, pressure, or concentration of reactants or products. It states that if an external stress is applied to a system at equilibrium, the system will adjust in such a way as to partially offset the effect of the stress and restore equilibrium.
For example:
● If the concentration of a reactant or product is increased, the equilibrium will shift in the direction that consumes or reduces that component, respectively.
● If the temperature is increased in an exothermic reaction (heat is a product), the equilibrium will shift in the direction that consumes heat (the forward reaction).
● If the pressure is increased in a gaseous reaction involving a change in moles of gas, the equilibrium will shift to the side with fewer moles of gas.
Understanding these equilibrium concepts is crucial for analysing and predicting the behaviour of chemical systems. Equilibrium principles are widely applied in various areas of chemistry, including chemical reactions, industrial processes, and environmental studies, enabling scientists to optimise reaction conditions and design efficient chemical processes.