key term - Partial pressures

Definition

Partial pressures refer to the individual pressures exerted by each gas in a mixture of gases. In the context of equilibrium, understanding how these partial pressures interact helps in predicting how changes in temperature or pressure can shift the balance of a chemical reaction, impacting the concentrations of reactants and products.

5 Must Know Facts For Your Next Test

1. The partial pressure of a gas can be calculated using the formula: $$P_{i} = X_{i} imes P_{total}$$, where $$P_{i}$$ is the partial pressure, $$X_{i}$$ is the mole fraction of the gas, and $$P_{total}$$ is the total pressure of the gas mixture.
2. In equilibrium reactions involving gases, changes in total pressure can affect the position of equilibrium by shifting it towards the side with fewer moles of gas.
3. When temperature increases in an exothermic reaction, it can lead to a decrease in the partial pressures of products as the reaction shifts to favor reactants.
4. The use of partial pressures is especially important in gaseous equilibria, where gas concentration can be related directly to pressure through the ideal gas law.
5. Calculating and manipulating partial pressures is critical when using Le Chatelier's Principle to predict how equilibrium will shift in response to changes in external conditions.

Review Questions

• How do partial pressures relate to Dalton's Law and what implications does this have for gas mixtures in equilibrium?
  • Partial pressures are directly linked to Dalton's Law, which states that the total pressure of a gas mixture is equal to the sum of its individual partial pressures. This means that understanding each component's contribution helps determine how gas mixtures behave at equilibrium. For example, if you know the partial pressures of each gas in a reaction mixture, you can predict how changes in one component will affect the overall system.
• In what way do changes in total pressure impact the equilibrium position when dealing with gaseous reactions?
  • Changes in total pressure can significantly shift the position of equilibrium for gaseous reactions. According to Le Chatelier's Principle, increasing total pressure will drive the reaction towards the side with fewer moles of gas, thereby reducing the overall system pressure. Conversely, decreasing total pressure will favor the side with more moles of gas. Understanding this relationship allows chemists to manipulate conditions to favor desired products.
• Evaluate how temperature changes affect both partial pressures and equilibrium constants in exothermic reactions.
  • In exothermic reactions, an increase in temperature typically results in a decrease in product formation as the system shifts towards reactants to absorb excess heat. This shift leads to lower partial pressures for products and influences the equilibrium constant (Kp), which decreases with increased temperature. By evaluating these effects together, one can predict how temperature variations will alter both partial pressures and overall reaction dynamics, affecting industrial applications and laboratory experiments.

Related terms

Dalton's Law:

A principle stating that the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of each individual gas.

Equilibrium Constant (Kp): A ratio that expresses the relationship between the partial pressures of products and reactants at equilibrium for a given reaction.

Le Chatelier's Principle:

A principle stating that if a dynamic equilibrium is disturbed by changing the conditions, the system responds by counteracting the change to re-establish equilibrium.