Unit 5: Chemical Kinetics
Reaction Rates
What is a Reaction Rate?
The reaction rate measures how quickly reactants are converted into products over time.
Usually expressed as a change in concentration of a substance (Δ) per unit time (Δt).
Instantaneous Reaction Rate:
The rate of a reaction at a specific moment in time.
Calculated using the slope of the tangent to the concentration vs. time graph at that point.
Reaction Rate Expression:
Key Points:
Reactants have negative rates because their concentrations decrease.
Products have positive rates because their concentrations increase.
The rate of change for all species (reactants and products) is proportional.
Introduction to Rate Law
Definition:
A rate law expresses how the reaction rate depends on the concentration of reactants.
General form: Rate=k[A]^m[B]^n
k: Rate constant (depends on temperature).
[A],[B]: Concentrations of reactants.
m, n: Reaction orders (must be determined experimentally).
Reaction Order:
The sum of exponents m+nm + nm+n in the rate law determines the overall reaction order.
Determining Rate Law:
Only experimental data can determine the exact rate law and reaction orders.
Use initial rates or concentration-time data to calculate m, n, k
Concentration Changes Over Time
Graphical Analysis:
The effect of concentration on reaction rate can be visualized in a concentration vs. time graph:
Steeper slope: Faster reaction (higher initial concentration of reactants).
Shallower slope: Slower reaction as reactants are consumed.
Collision Theory
Definition:
Chemical reactions occur when reactant particles collide effectively.
Requirements for an Effective Collision:
Orientation: Particles must collide in the correct orientation to break and form bonds.
Collision Energy: Particles must have enough kinetic energy to overcome the activation energy (Ea).
Factors Affecting Reaction Rate:
Surface Area: Increased surface area allows more collisions (e.g., powdered solids react faster than large chunks).
Temperature: Higher temperature increases particle kinetic energy and collision frequency.
Concentration: More particles in a given volume lead to more frequent collisions.
Catalyst: Lowers the activation energy (Ea) without being consumed in the reaction.
Reaction Energy Profile
Reaction Coordinate Diagram:
Shows the energy changes during a reaction.
Key Features:
Peaks: Transition states (highest energy points).
Troughs: Intermediates (short-lived species between steps).
Activation Energy (Ea): Energy barrier that must be overcome for the reaction to proceed.
Energy Change (ΔH):
ΔH>0: Endothermic (energy absorbed).
ΔH<0: Exothermic (energy released).
Reaction Mechanisms
Definition:
A sequence of elementary steps that describe how a reaction occurs.
Molecularity:
The number of reactant particles involved in an elementary step:
Unimolecular: Involves one molecule.
Bimolecular: Involves two molecules.
Termolecular: Involves three molecules (rare due to low probability).
Rate-Determining Step:
The slowest step in the mechanism dictates the overall reaction rate.
The rate law is determined by the reactants in the slow step.
Intermediates and Catalysts:
Intermediate: Formed and consumed during the reaction (does not appear in the rate law).
Catalyst: Introduced at the beginning, reappears at the end (lowers Ea).
Steady-State Approximation
Definition:
Assumes the concentration of intermediates remains constant during the reaction.
Used to simplify rate law expressions for multistep reactions.
Multistep Reaction Energy Profile
Diagram Features:
Each peak represents a transition state.
Each trough corresponds to an intermediate.
The highest peak is the rate-determining step.
Energy Analysis:
Endothermic Reaction: Products are higher in energy than reactants.
Exothermic Reaction: Products are lower in energy than reactants.
Example:
For A→GA :
If the energy of G>AG, the reaction is endothermic.
If the energy of G<AG, the reaction is exothermic.