STUDENT SLIDES - Chapter 9 - FA24 DW

Reaction Rates and Equilibrium

Biological Importance

  • Enzymes facilitate complex biological reactions.

  • Biological reactions often run in both forward and reverse directions.

  • Aim of this chapter: understand reactions at equilibrium.

Chapter Outline

  • 9.1 Reaction Rates and Energy Diagrams

  • 9.2 Reactions at Equilibrium

  • 9.3 Using Equilibrium Expressions

  • 9.4 Le Châtelier’s Principle

9.1 Reaction Rates and Energy Diagrams

  • Factors affecting reaction rates:

    • Concentration of reactants

    • Surface area of solid reactants

    • Temperature

    • Presence of a catalyst

    • Nature of reactants

Collision Theory

  • Successful reaction requires:

    • Sufficient energy during collision.

    • Correct molecular orientation.

Reaction Energy Diagram

  • Illustrates energy changes during a reaction.

  • Catalysts lower activation energy to increase reaction rate.

Reaction Rate Definition

  • Measures the speed at which reactants convert to products.

  • Units: molarity/time (e.g., M/s).

9.2 Reactions at Equilibrium

  • Equilibrium Condition:

    • Forward and reverse reactions occur at equal rates.

  • Equilibrium Expressions:

    • Relation of concentrations of reactants and products.

    • K = [C][D] / [A][B].

Dynamic Nature of Equilibrium

  • Concentrations remain constant but reactions continue.

9.3 Using Equilibrium Expressions

  • Calculating K:

    • Insert equilibrium concentrations into K expression.

    • K value indicates reaction favorability.

9.4 Le Châtelier’s Principle

  • Stress Response:

    • System at equilibrium shifts to relieve stress (concentration, pressure, temperature).

  • Direction of Shift:

    • Increase in concentration of a reactant shifts right.

    • Increase in concentration of a product shifts left.

Effects of Temperature

  • Heating endothermic reactions shifts right.

  • Cooling endothermic reactions shifts left.

  • Heating exothermic reactions shifts left.

  • Cooling exothermic reactions shifts right.

Catalysts and Equilibrium

  • Catalysts increase rate of reaction without altering equilibrium position.

Main Points for Solving Reaction Rates and Equilibrium Problems

Key Concepts to Understand:

  • Reaction Rates: Speed at which reactants convert to products; affected by factors such as concentration, temperature, surface area, and catalysts.

  • Collision Theory: Successful reactions require sufficient energy during collision and correct molecular orientation.

  • Equilibrium: Occurs when forward and reverse reactions happen at equal rates; dynamic in nature with constant concentrations of reactants and products.

  • Equilibrium Expressions: K = [C][D] / [A][B]; used to relate concentrations of reactants and products.

  • Le Châtelier’s Principle: System shifts in response to stresses (changes in concentration, pressure, temperature) to restore equilibrium.

Essential Points Before Solving Problems:

  1. Identify the Reaction: Know the reactants and products involved.

  2. Understand K: Be able to calculate K from equilibrium concentrations and understand its significance for the favorability of the reaction.

  3. Recognize Factors Affecting Rates: Consider how changes in concentration, temperature, and catalysts may impact the reaction rate.

  4. Apply Le Châtelier’s Principle: Predict the shift in equilibrium based on changes to the system.

While I cannot directly create or display diagrams, I can guide you on how to create your own diagrams of catalysts and label them effectively. Here’s how you can illustrate a catalyst in a reaction energy diagram:

  1. Draw the Reaction Energy Diagram:

    • X-Axis: Label the x-axis as 'Reaction Progress'.

    • Y-Axis: Label the y-axis as 'Energy'.

  2. Plot the Energy Levels:

    • Reactants: Start with a horizontal line indicating the energy level of the reactants. Label it "Reactants".

    • Products: Draw another horizontal line at a different level to represent the products. Label it "Products".

    • Activation Energy Without Catalyst: Draw a peak (the transition state) between the reactants and products, showing the activation energy required for the reaction to proceed without a catalyst. Label this peak "Activation Energy (No Catalyst)".

    • Activation Energy With Catalyst: Draw a lower peak to represent the activation energy with a catalyst. Label this "Activation Energy (With Catalyst)".

  3. Add a Horizontal Arrow: Connect the reactants to the products, showing the overall energy change of the reaction. Label the arrow as "Overall Energy Change (ΔE)".

  4. Catalyst Label: On the lower peak, you might write "Catalyst (Lowering Activation Energy)" to indicate its role in reducing the activation energy.

Additional Elements

  • Color Coding: Consider using different colors for reactants, products, and activation energy levels to make your diagram clearer.

  • Legend: You can add a small legend to explain any colors or symbols used in your diagram.

This format will effectively illustrate how catalysts work in terms of energy changes during a chemical reaction.

Key Formulas for Reaction Rates and Equilibrium

  1. Reaction Rate (R):

    • R = \frac{\Delta [A]}{\Delta t}(Rate of change of concentration over time)

  2. Equilibrium Constant (K):

    • K = \frac{[C][D]}{[A][B]}(For a general reaction A + B ⇌ C + D)

  3. Le Châtelier’s Principle Expression:

    • If stress is applied, equilibrium shifts to counteract the stress.

    • Changes:

      • Increase in reactant concentration: shifts to the right (forward reaction).

      • Increase in product concentration: shifts to the left (reverse reaction).

  4. Calculating K:

    • Insert equilibrium concentrations into K expression to find the value.

  5. Temperature Effect on K:

    • K changes with temperature; specific values depend on the reaction type (endothermic or exothermic).

  6. Catalyst Effect:

    • Catalysts do not affect K; they only speed up the rate to reach equilibrium without affecting its position.

Main Phrases for Reaction Rates and Equilibrium

  1. If reactant concentration increases:

    • Shifts equilibrium to the right (forward reaction).

  2. If product concentration increases:

    • Shifts equilibrium to the left (reverse reaction).

  3. Increase in pressure (for gaseous reactions):

    • Shifts equilibrium towards the side with fewer gas molecules.

  4. Decrease in pressure:

    • Shifts equilibrium towards the side with more gas molecules.

  5. Heating an endothermic reaction:

    • Shifts equilibrium to the right (producing more products).

  6. Cooling an endothermic reaction:

    • Shifts equilibrium to the left (producing more reactants).

  7. Heating an exothermic reaction:

    • Shifts equilibrium to the left (producing more reactants).

  8. Cooling an exothermic reaction:

    • Shifts equilibrium to the right (producing more products).

  9. Addition of a catalyst:

    • Increases the rate of reaction but does not alter the position of equilibrium.