Equilibrium pressures

Definition

Equilibrium pressures refer to the partial pressures of gases in a chemical reaction at equilibrium, which remain constant over time as the rates of the forward and reverse reactions are equal. Understanding equilibrium pressures is essential for calculating equilibrium concentrations, as they are directly related to the concentration of reactants and products through the ideal gas law and the equilibrium constant expression. This concept is crucial in predicting how changes in conditions like temperature or volume affect the system's behavior.

5 Must Know Facts For Your Next Test

1. At equilibrium, the total pressure of a gaseous mixture can be determined by adding the partial pressures of each individual gas present in the reaction.
2. The equilibrium constant can be expressed in terms of partial pressures, where $$K_p$$ = $$rac{P_{products}}{P_{reactants}}$$ at a given temperature.
3. Changes in temperature will affect equilibrium pressures; increasing temperature typically favors endothermic reactions, shifting the equilibrium position.
4. If the volume of a gas reaction container is decreased, the system will adjust by shifting towards the side with fewer moles of gas to relieve pressure.
5. Equilibrium pressures can also be affected by adding or removing reactants or products, causing a shift in the reaction to restore balance.

Review Questions

• How do changes in temperature influence equilibrium pressures in a reaction?
  • Changes in temperature can significantly influence equilibrium pressures by favoring either exothermic or endothermic reactions. For instance, increasing the temperature generally shifts the equilibrium position toward endothermic reactions, leading to an increase in product formation and thus altering their partial pressures. Conversely, decreasing temperature favors exothermic reactions, potentially increasing reactant pressures. Therefore, understanding how temperature affects these pressures is crucial for predicting system behavior.
• Discuss how Le Chatelier's Principle can be used to predict changes in equilibrium pressures when a system is disturbed.
  • Le Chatelier's Principle states that if an external change is applied to a system at equilibrium, the system will adjust to counteract that change and restore balance. For example, if pressure is increased by decreasing volume, the equilibrium will shift towards the side with fewer moles of gas to reduce pressure. This principle helps predict how adding or removing reactants or products will affect equilibrium pressures, allowing for better control over chemical reactions.
• Evaluate how calculating equilibrium pressures can aid in industrial processes and chemical manufacturing.
  • Calculating equilibrium pressures is vital in industrial processes as it allows chemists and engineers to optimize reaction conditions for maximum yield. By understanding how factors like temperature and pressure affect the position of equilibrium, they can manipulate these conditions to favor product formation. Additionally, accurately predicting these pressures helps manage resources effectively, reduces waste, and ensures safety in chemical manufacturing. This evaluation highlights the importance of applying theoretical knowledge to practical applications.

Related terms

Le Chatelier's Principle:

A principle stating that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change.

Ideal Gas Law: An equation of state for an ideal gas, expressed as PV = nRT, where P is pressure, V is volume, n is moles of gas, R is the gas constant, and T is temperature.

Equilibrium constant (K): A numerical value that expresses the ratio of concentrations or partial pressures of products to reactants at equilibrium at a specific temperature.