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CA_Lesson_1_A+Model+for+Reaction+Rates Half

A Model for Reaction Rates

  • Focus Question: How do you determine how fast a reaction is going?

  • Unit 3 Module 15

New Vocabulary

  • Reaction Rate: Change in concentration of a reactant or product per unit of time, expressed as mol/(L.s).

  • Collision Theory: Atoms, ions, and molecules must collide to react.

  • Activated Complex: A temporary, unstable arrangement of atoms that may form products or revert to reactants.

  • Activation Energy: The minimum energy required for reacting particles to form an activated complex, enabling the reaction to proceed.

Review Vocabulary

  • Energy: The ability to do work or produce heat; exists in two forms: potential energy and kinetic energy.

Expressing Reaction Rates

  • The Greek letter delta (Δ) indicates a change in quantity.

  • Equation for Average Rate: Average rate = [C] at t2 - [C] at t1 / (t2 - t1)

  • Reaction rate measures change in concentration per unit of time, generally expressed as mol/(L·s).

Determining Reaction Rates

  • Reaction rates are measured experimentally as concentrations change over time.

  • Graphically, the progression from reactants to products can be observed.

Example Calculation

  • Reaction: CO (G) + NO2 (G) → CO2 (G) + NO (G)

  • To monitor reaction rates, calculate production rates with positive values indicating increased concentration.

Example Problem: Calculate Average Reaction Rates

  • Problem Statement: In a reaction between butyl chloride (C4H9Cl) and water, [C4H9Cl] decreases from 0.220 M to 0.100 M in 4.00 s. Calculate average rate.

  • Known Values:

    • t1 = 0.00 s, t2 = 4.00 s

    • [C4H9Cl] at t1 = 0.220 M

    • [C4H9Cl] at t2 = 0.100 M

  • Rate Calculation: Average reaction rate = (0.100 M - 0.220 M) / (4.00 s - 0.00 s) = -0.0300 mol/(L·s)

    • Positive value calculated suggests a reasonable decrease in concentration.

Collision Theory

  • States that reactions require collisions between particles with the right orientation and minimum energy to occur successfuly.

  • Colliding particles must have sufficient kinetic energy and correct orientation to react.

Factors Influencing Collision Success

  • Activation Energy: Minimum energy needed to start a chemical reaction; some reactions require very little (like rusting).

  • Fast reactions are characterized by low activation energy; slow reactions require high activation energy.

Energy Profiles

  • Illustrated by energy changes during reactions with exothermic reactions having lower product energy than reactants, while endothermic reactions have higher product energy.

Catalysts

  • Catalyst: Substance that accelerates reactions by lowering activation energy but is not consumed.

  • Inhibitor: Slows down reactions or prevents them from occurring.

  • Heterogeneous Catalyst: Different physical state than the reaction it catalyzes.

  • Homogeneous Catalyst: Same physical state as the reaction it catalyzes.

Nature of Reactants

  • Concentration: Higher concentration increases collision frequency, accelerating reactions.

  • Surface Area: Increased surface area leads to more effective collisions.

  • Temperature: Raising temperature increases average kinetic energy, facilitating more collisions with enough energy to surpass the activation energy.

Reaction Rate Laws

  • Rate Law: Explains relationship between reaction rates and reactant concentrations at a given temperature.

  • First-Order Reactions: Rate is directly proportional to one reactant’s concentration.

  • Overall Reaction Order: Sum of individual orders from rate law.

Instantaneous Reaction Rates and Mechanisms

  • Instantaneous Rate: Rate at a specific time derived from the rate law and concentrations.

  • Complex Reaction: Consists of multiple elementary steps, with one rate-determining step being the slowest, limiting overall reaction speed.