Heterogeneous equilibrium

Definition

Heterogeneous equilibrium refers to a state of balance in a chemical reaction where the reactants and products exist in different phases, such as solids, liquids, and gases. This concept is crucial for understanding how different physical states can affect the rates of reactions and the conditions under which they occur. In this type of equilibrium, the concentrations of the reactants and products do not change over time, despite being present in varying states.

5 Must Know Facts For Your Next Test

1. In heterogeneous equilibrium, the concentrations of solids and liquids do not appear in the equilibrium expression because their activities are defined as 1.
2. The presence of different phases allows for unique reactions; for example, a solid can react with a gas or liquid to reach equilibrium.
3. The position of heterogeneous equilibrium can be affected by changes in temperature, pressure, and concentration of reactants or products.
4. Heterogeneous equilibria often involve reactions such as dissolution, where a solid solute dissolves in a liquid solvent.
5. Common examples include reactions involving gases reacting with solids or liquids, like the formation of gaseous products from solid reactants.

Review Questions

• How does heterogeneous equilibrium differ from homogeneous equilibrium?
  • Heterogeneous equilibrium involves reactants and products in different phases, while homogeneous equilibrium consists of all substances in the same phase. For example, a reaction where a solid reacts with a gas represents heterogeneous equilibrium due to the different states involved. In contrast, a reaction between gases alone would be classified as homogeneous. This distinction impacts how concentrations are expressed in equilibrium constants and how changes in conditions affect each type of equilibrium.
• What role does Le Chatelier's Principle play in understanding heterogeneous equilibria?
  • Le Chatelier's Principle is crucial for predicting how a heterogeneous equilibrium will respond to changes in conditions. When a change occurs, such as an increase in pressure or concentration of a gaseous reactant, the system will shift to minimize that change by favoring either the formation of reactants or products. This principle helps explain how systems reach new equilibria under various external influences, making it essential for understanding chemical dynamics.
• Evaluate how temperature changes can affect the position of heterogeneous equilibrium and provide an example.
  • Temperature changes can significantly shift the position of heterogeneous equilibrium depending on whether the reaction is endothermic or exothermic. For instance, consider the decomposition of calcium carbonate (CaCO₃) into calcium oxide (CaO) and carbon dioxide (CO₂). If heat is added to this system (endothermic reaction), it shifts to favor the formation of products at higher temperatures. Conversely, removing heat from an exothermic reaction would shift equilibrium towards reactants. Understanding these shifts is vital for optimizing chemical processes and reactions.

Related terms

Le Chatelier's Principle: A principle stating that if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system will shift in a direction that counteracts the change.

Dynamic Equilibrium: A state in which the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of reactants and products.

Equilibrium Constant (K): A numerical value that expresses the ratio of the concentrations of products to reactants at equilibrium, each raised to the power of their coefficients in the balanced chemical equation.