Daley_Student_Abbrev_Unit_III_2025_Chem152
Chemical Kinetics
Overview: Study of rates and mechanisms of chemical reactions.
Example Reaction:
Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
Understanding a Reaction
To study a reaction, we need to observe changes occurring during the reaction.
Key Questions:
What compounds are consumed?
How much of each compound is consumed?
Which compounds are formed?
How much of each compound is formed?
How fast is the reaction?
How do variables like concentration, pressure, and temperature affect the reaction?
Reaction Rates
Definition: The speed of a chemical reaction, indicating how fast products are formed or reactants are consumed.
Collision Model: Reactions occur due to collisions between particles (atoms, ions, molecules).
Factors Affecting Reaction Rate:
Concentration of Reactants:
Increased concentration generally increases reaction rate due to more frequent collisions.
Temperature:
Higher temperatures lead to faster particle movement and frequency of collisions.
Approx. 10°C increase can double the reaction rate.
Nature of Reactants:
Smaller molecules react faster than larger ones.
Gases react faster than liquids, which react faster than solids.
Powdered solids react faster due to increased surface area.
Examining a Reaction: Defining Reaction Rate
Example Reaction:
2 NO2 (g) + heat ⟶ 2 NO(g) + O2 (g)
Concentration Over Time:
[conc] 0 Time [NO2]
Reaction Rate Expressions
Rate must be expressed as “+” valued.
Rate = ± Δ[conc] / Δt
Rate formation of NO: Rate = Δ[NO] / Δt
Rate formation of O2: Rate = Δ[O2] / Δt
Rate consumption of NO2: Rate = - Δ[NO2] / Δt
Example Reaction Problems
Problem 46:
For the reaction: 2 NO2 (g) → 2 NO(g) + O2 (g):
(a) If the rate of consumption of NO2 is 0.5 M/min, what is the rate of O2 formation?
(b) If the rate of NO formation is 0.75 M/min, what is the rate of NO2 consumption?
(c) General formula for rates of consumption and formation for this reaction.
Problem 47:
For the reaction: N2 (g) + 3 H2 (g) → 2 NH3 (g):
(a) Identify compounds present after 1 N2 reacts.
(b) Calculate rate of N2 and H2 consumption and NH3 formation over the first 2 minutes.
(c) Calculate the general rate of reaction.
Relative Reaction Rates
General Rule: All reaction rates decrease over time, and comparative rates must be over the same time interval.
Relative Rates: Some aspects of kinetics utilize reaction stoichiometry.
∆ = change in concentration (final - initial)
Rate Law
Definition: A mathematical relationship showing how each reactant influences reaction rate.
Rate Law General Form:
Rate = k [A]^x [B]^y
Exponents (x, y) must be determined experimentally.
Involves order with respect to each reactant and overall order (sum of exponents).
Rate Constant (k)
Depends on Temperature: Report with temperature (e.g. "k at 20 °C").
Units of k depend on overall reaction order:
Order 0: M/s
Order 1: 1/s
Order 2: 1/(M s)
Order 3: 1/(M² s)
Determine Units of k
(a) For Rate = k[A][B]: k would have units of M⁻¹ s⁻¹.
(b) For Rate = k[A]²[B]: k would have units of M⁻³ s⁻¹.
Reaction Order Effects
Effects on the Rate:
Order 0: Concentration change has no effect.
Order 1: Doubling concentration doubles the rate; halving cuts the rate in half.
Order 2: Doubling triples the rate; halving cuts the rate in half.
Skills to Master
Write a reaction mechanism from a Rate Law and vice versa.
Calculate relative rates and overall rates.
Determine a rate law from experimental data with different methods.
Draw and interpret a reaction coordinate diagram.
Assess the impact of temperature and/or catalysts on reaction rate.
Methods to Determine Rate Law
Method of Initial Rates (MOIR):
Systematically vary reactant concentration to observe their effect on the initial reaction rate.
Integrated Rate Law Method (IRL):
Collect concentration vs time data to generate specific plots; the plot yielding a straight line indicates order with respect to that reactant.
Initial Rate vs. Instantaneous Rate
Initial Rate: The slope of the tangent to the curve at time = 0.
Instantaneous Rate: The slope of the tangent at a specific time (e.g. at t = 350s).
The average rate of formation during a time interval is the slope derived from concentration measurements.
Reaction Rate Law Example
For the reaction: 2 NO(g) + 2 H2 (g) → N2 (g) + 2 H2O(g)
Rate Law Formulation: rate = k·[NO]ⁿ·[H2]ᵐ
Trials can be set up to solve for orders n and m, and the rate constant k.