Reaction Rates and Laws

Kinetics Definition:

1. Studies reaction rates – How fast or slow reactions occur.
2. Explains influencing factors – Temperature, concentration, catalysts, etc.
3. Compares reaction speeds – Fast (methane explosion) vs. slow (hydrogen peroxide breakdown).
4. Analyzes reaction mechanisms – Steps and conditions affecting speed.

Steps to Measure the Rate of a Reaction:

1. Observe Concentrations – Track changes in reactant and product concentrations over time. The reactants decrease while the products increase.
2. Use Rate Formulas – Calculate using:
   • Rate = Δ[Reactant] / t
   • Rate = Δ[Product] / t
   • Maintain the correct ratio of reactants to products based on the balanced equation.
3. Ensure Correct Units – Express rate in mol/L·s (Ms⁻¹) and adjust for different time units (seconds, hours, etc.).

Reaction Rate and Equilibrium:

1. Equilibrium Concept – A reaction reaches equilibrium when the forward and reverse reaction rates are equal, keeping reactant and product concentrations constant.
2. Graphical Representation – The rate is shown as the slope of a concentration vs. time graph.
3. Types of Reaction Rates:
   • Average Rate – Change in concentration over a time interval (slope between two points).
   • Instantaneous Rate – The rate at a specific moment, found using calculus (derivative).

Factors Affecting Reaction Rate:

1. Concentration – Higher reactant concentration increases collision frequency, speeding up the reaction.
2. Temperature – Higher temperature increases kinetic energy, leading to more successful collisions.
3. Surface Area – More exposed surface (e.g., powdered solids) allows more collisions, increasing reaction speed.
4. Catalyst Presence – Lowers activation energy, providing an alternative reaction pathway without being consumed.
5. Pressure (for gases) – Higher pressure increases molecule density, leading to more frequent collisions.

1.What is a Rate Law?

• Describes how reaction rate depends on reactant concentrations.
• General form: R = k[A]^ⁿ[B]^ᵐ, where:
  • R = reaction rate
  • k = rate constant
  • [A], [B] = reactant concentrations
  • n, m = reaction orders

2. Reaction Order

• Reaction order (n, m) tells how concentration affects the rate:
  • If n = 1, doubling [A] doubles the rate (first-order).
  • If n = 2, doubling [A] quadruples the rate (second-order).
  • Overall reaction order = sum of individual reactant orders.

3. Determining Rate Law Experimentally

• Run experiments with different concentrations.
• Compare rates to find reaction order for each reactant.
• Example: If doubling [A] makes rate 4×, A is second order (n = 2).

4. Understanding the Rate Constant (k)

• k is a proportionality constant that changes with temperature.
• Units of k depend on reaction order:
  • Zeroth order: M/s
  • First order: s⁻¹
  • Second order: M⁻¹s⁻¹

The reaction order, denoted as n, describes how concentration affects the reaction rate. While reaction orders are often integers (e.g., n = 0, 1, 2), some reactions, like Michaelis–Menten kinetics, can have fractional orders.

Exponential Decay and Chemical Reactions

• The rate of a reaction can be compared to population decay, where the rate of change is proportional to the quantity present.
• The reaction rate for A → B follows:
• Rate = k[A]ⁿ
  • where k is the rate constant and n is the reaction order.

Integrated Rate Laws

• Integrated rate laws express concentration as a function of time. The order of the reaction determines the form of the equation:

Reaction Mechanisms and Rate Laws: A Summary

What is a Mechanism?

• A mechanism describes the step-by-step process of a reaction.
• Most reactions occur in multiple steps, known as elementary steps.
• The sum of all elementary steps gives the overall balanced chemical equation.

Elementary Steps and Intermediates

• Each elementary step involves reactants forming intermediates or products.
• Intermediates are species that appear in one step and are consumed in another.
• Catalysts are present at the start and regenerated at the end.

Rate-Determining Step (RDS)

• The slowest step in a reaction mechanism.
• Limits the overall reaction speed.
• The rate law is derived from the reactants of the RDS using their stoichiometric coefficients.

• Unit 5 Lesson 1