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.