3.1.9 - Rates

3.1.9 Rates Physical Chemistry

AS Review: Key Concepts

Rate of Reaction

  • Definition: Measure of how the concentration of a reagent or product changes over time.

  • Units for rate: mol dm⁻³ s⁻¹

Effect of Concentration on Rate

  • Increasing reactant concentration results in more molecules in a volume, leading to more successful collisions and an increased reaction rate.

Activation Energy

  • Definition: Minimum energy required to initiate a reaction by breaking chemical bonds.

Catalysts

  • Function: Increase the rate of reaction by lowering the activation energy through alternative reaction pathways without undergoing any permanent change themselves.

Enthalpy Profile for Exothermic Reaction

  • Illustrated with and without a catalyst to show the difference in activation energies.

Boltzmann Distribution

  • At higher temperatures, the curve shifts, indicating that more particles have energy greater than or equal to the activation energy.

Recall from Year 1

Summary of Key Concepts

  • Rate of Reaction: Change in concentration of reactant or product per unit time.

  • Collision Theory: Particles must collide with sufficient energy and correct orientation for a successful reaction.

  • Initial Rate of Reaction: Rate at the start (t=0).

3.1.9.1 Rates of Reaction

  • Define key terms such as order of reaction and rate constant.

  • Analyze how changing concentration affects reaction rates based on various orders (0, 1, 2).

  • Focus on calculating rate constants, their units, and how they vary with temperature.

  • Apply the Arrhenius equation for calculations.

Reaction Order

Definitions

  1. Zero Order:

    • Relationship: Rate = [A]⁰ (No effect of [A] on rate).

  2. First Order:

    • Relationship: Rate = [A]¹ (Rate doubles when concentration doubles).

  3. Second Order:

    • Relationship: Rate = [A]² (Rate increases quadratically with concentration).

Order of Reaction in Experiments

  • Using the example: 2A + B + 3C → products.

  • Determining the rate equation based on experimental data and orders derived from concentration changes.

  • Overall order of reaction equals the sum of individual orders.

Analyzing Rate from Experimental Data

Sample Data Table

  • Experiments comparing initial rates and concentrations to derive the order.

    • Example: As [A] doubles, rate doubles.

  • Conclude that the reaction is first order with respect to [A].

Example Tasks for Order Determination

  • Use experimental data to formulate and justify rate equations.

  • Understand how to relate experimental changes (e.g., doubling concentrations) to changes in reaction rate.

Rate Equations and Rate Constants

  • Rate = k[A]²[B] (Second order in A, first order in B).

  • Determine units for rate constant based on the overall order of the reaction.

Arrhenius Equation

  • Understand that k = Ae⁻(Ea/RT) links temperature, rate constant, and activation energy:

    • A = Pre-exponential factor

    • Ea = Activation energy

  • ln k = ln A - (Ea/RT) for linearization of the equation.

Key Graph Analysis

Concentration vs. Rate Graphs

  1. Zero Order: No effect of concentration on the rate (flat line).

  2. First Order: Linear relationship (linearly increasing).

  3. Second Order: Exponential curve showing rate increase.

Concentration-Time Graphs

Initial Rate Measurements

  • Measure the gradient at t=0 to find initial rate.

Rate Determining Step (RDS)

Concepts

  • The slowest step in a mechanism determines the rate of the overall reaction.

  • The balanced equation does not provide information regarding RDS.

Applications of the Rate Equation

  • Practical examples on finding rate constants through experiments and data analysis.

  • Task review to solidify understanding of rate constants, activation energy, and Arrhenius applications.

Summary of Learning Outcomes

  • Recognize and derive orders from concentration and time graphs.

  • Understand the components and application of the Arrhenius equation in reaction kinetics.