Chemical Thermodynamics - SS2 Chemistry Past Questions and Answers - page 3

Spontaneity refers to the tendency of a chemical reaction to occur without any external influence. A spontaneous reaction proceeds on its own, releasing energy and progressing toward equilibrium. The spontaneity of a reaction depends on two main factors: enthalpy (ΔH) and entropy (ΔS).

If a reaction has a negative ΔH (exothermic, releases heat) and a positive ΔS (increase in disorder), it will be spontaneous at all temperatures. For example, the combustion of gasoline (exothermic) and the dissolution of sugar in water are spontaneous processes.

Conversely, a reaction with a positive ΔH (endothermic, absorbs heat) and a negative ΔS (decrease in disorder) is non-spontaneous at all temperatures. An example is the decomposition of water into hydrogen and oxygen gas.

However, if a reaction has a positive ΔH and a positive ΔS, it can still be spontaneous at high temperatures. The increase in temperature allows the entropy term to dominate, driving the reaction forward. An example is the Haber-Bosch process, where nitrogen and hydrogen react to form ammonia.

Chemical equilibrium refers to the state in a reversible reaction where the forward and reverse reactions occur at equal rates. At equilibrium, the concentrations of reactants and products remain constant, although the individual molecules continue to undergo reactions.

The principles associated with equilibrium include the law of mass action, which states that the ratio of the concentrations of reactants and products at equilibrium is constant and can be expressed as the equilibrium constant (K). Le Chatelier's principle is also crucial, stating that a system at equilibrium will respond to changes in concentration, pressure, or temperature to reestablish equilibrium.

Changes in concentration: If the concentration of a reactant or product is increased, the equilibrium will shift in the direction that consumes or reduces that substance. Conversely, if the concentration is decreased, the equilibrium will shift in the direction that produces or increases that substance. For example, increasing the concentration of CO2 in the reaction CO2 + H2O ⇌ H2CO3 will shift the equilibrium to the right, favouring the formation of carbonic acid.

Changes in pressure: For reactions involving gases, changes in pressure can affect the equilibrium position. An increase in pressure shifts the equilibrium towards the side with fewer moles of gas, while a decrease in pressure favours the side with more moles of gas. The reaction N2(g) + 3H2(g) ⇌ 2NH3(g) illustrates this principle.

Changes in temperature: Altering the temperature affects the equilibrium position differently depending on whether the reaction is exothermic or endothermic. Increasing the temperature favours the endothermic reaction, while decreasing the temperature favours the exothermic reaction. An example is the equilibrium between nitrogen dioxide (NO2) and dinitrogen tetroxide (N2O4).