Lecture 5 – Rxn Mechanisms Reaction Mechanisms and Elementary Reactions Reaction Mechanism: a series of steps that lead from reactants to products Elementary Reaction: single step in reaction mechanism For elementary reactions, rate law exponents = stoichiometric coefficients. Elementary reactions can be Unimolecular, Bimolecular and Trimolecular. (Trimolecular = very rare)

Types • Unimolecular – A  P • Bimolecular – A + A  P – A + B  P • Trimolecular – A + B + C  P – A + A + B  P – A + A + A  P

Practice Write rate laws for each elementary step and identify as unimolecular, bimolecular, and trimolecular. 2 2 NO  N2O2 Rate = k[NO] bimolecular O3  O2 + O Rate = k[O3] unimolecular NO + O3  NO2 + O2 Rate = k[NO][O3] bimolecular 2 2 Cl + N2  Cl2 + N2 Rate = k[Cl] [N2] trimolecular

Reaction Mechanisms and Elementary Reactions Intermediates: Compounds produced and used in a mechanism. They do not appear in the overall chemical equation. Rate Determining Step: slowest step that determines the rate of the overall reaction Catalysts: provide lower Ea for a reaction – Not part of reactants or products

Example The following mechanism has been proposed for the reaction of NO with H2 to form N2O and H2O: (1) NO(g) + NO(g)  N2O2(g) (2) N2O2(g) + H2(g)  N2O(g) + H2O(g) a) Write the rate law for each elementary step. b) Show the overall reaction for this mechanism. c) Identify the intermediate(s) of this mechanism. d) What is the observed rate law for the reaction if step #1 is the slowest step?

Practice Thallium(I) is oxidized by cerium(IV) as follows: Tl+ + 2 Ce4+  Tl3+ + 2 Ce3+ The elementary steps, in the presence of Mn(II), are as follows: Step 1: Ce4+ + Mn2+  Ce3+ + Mn3+ Step 2: Ce4+ + Mn3+  Ce3+ + Mn4+ Step 3: Tl+ + Mn4+  Tl3+ + Mn2+ (a) Identify any intermediates, catalysts, and the rate-determining step if the rate law is rate = k[Ce4+][Mn2+]. (b) Explain why the reaction is slow without the catalyst.

Practice (Not required for notes page) The diagram shown here represents a two-step mechanism. (a) Write the equation for each step and the overall reaction. (b) Identify the intermediate and catalyst. The color codes are A = green and B = red.

Collision Theory Reactants must collide to react. Must collide with enough energy. Must collide with correct orientation. Activation Energy (Ea) = minimum energy needed for a reaction occur.

Arrhenius Equation

Convert Arrhenius equation to its linear form.

Arrhenius Equation with 2 Ts

Example A reaction has a rate constant of 0.0117 s-1 at 400 K and 0.689 s-1 at 450 K. What is the activation energy in kJ for the reaction? What is the value of the rate constant at 425K?

Practice 2 The activation energy for the reaction: N2O (g)  N2(g) + O(g) is 2.4 × 10 kJ/mol at 600 K. Calculate the percentage of the increase in rate from 600 K to 606 K. Comment on your results.

Molecular Energy Diagrams

Molecular Energy Diagrams 1. What is the highest kinetic energy value for the molecule shown? 2. Is the molecule accelerating or decelerating at the given points (1-3)?

Potential Energy Diagram

Practice 1. The reaction between A and B is determined to be a fairly fast reaction and only slightly exothermic. Which of the following potential energy surfaces best fits this description?

2. A particular reaction was found to have forward and reverse activation energies of 60 and 140 kJ mol-1, respectively. The energy change for the reaction is…