Ch 13 Reaction Mechanisms

I. Reaction MechanismsA. Mechanism = the series of discrete elementary steps leading from reactants to

products1) The slowest of these elementary steps determines how fast the overall

reaction can be2) The study of kinetics helps us determine what the mechanism is3) Example: NO2 + CO NO + CO2

a) Rate = k[NO2]2

b) The balanced equation doesn’t tell us how it happens, but the rate data tells us two NO2 molecules seem to be involved in the slowest step

c) We think the mechanism for this reaction is:1. NO2 + NO2 NO3 + NO

2. NO3 + CO NO2 + CO2

B. Definitions:1. Intermediate = molecule produced in one step of a mechanism and used up in

another step; it is not isolated as a final productFor the following mechanism, NO3 is an intermediate

1. NO2 + NO2 NO3 + NO

2. NO3 + CO NO2 + CO2

2. Elementary Step = reaction whose rate law is written directly from its molecularitya) Molecularity = how many molecules collide in the reaction

i. Unimolecular = only one molecule is involved in the stepii. Bimolecular = two molecules collide in the stepiii. Termolecular = three molecules collide in the step (this is very rare

because of the low probability)

b) Since the elementary steps are as simple as possible, we don’t need to analyze experimental data to find their rate laws. We can simply write the rate law directly from the elementary step equation.

1. NO2 + NO2 NO3 + NO rate = k[NO2]2

2. NO3 + CO NO2 + CO2 rate = k[NO3][CO]

3. Rate Determining Step (rds) = the slowest step in a mechanisma) One step has to be the slowest.b) The overall reaction can’t go faster than that one step, so it

determines the rate of the reactionc) The overall reaction rate will be the same as that of the r.d.sd) Example: NO2 + CO NO + CO2

i. Step 1 is the slowest stepii. Overall rate = k[NO2]2

e) Like the narrow point of a funnel, the rds determines the rate

C. Mechanisms can’t be proved, only disproved1) Since we can’t actually see individual molecules react, we can’t prove the

exact steps they go through2) When chemists suggest a mechanism, it must meet two criteria

a) The individual steps must add up to the observed reaction equation 1. NO2 + NO2 NO3 + NO

2. NO3 + CO NO2 + CO2

NO2 + CO NO + CO2

b) The proposed mechanism must match the observed rate lawObserved rate law is rate = k[NO2]2

Step 1 would give this rate law if it is r.d.s.

c) If either of these criteria aren’t met, the proposed mechanism is wrong

3) Example: 2 NO2 + F2 2NO2F

Step 1 NO2 + F2 NO2F + F (slow)

Step 2 F + NO2 NO2F (fast)

Observed rate law: rate = k[NO2][F2]

II. A Model for Chemical Kinetics

A. The Collision Model = Molecules must collide to react1) Reactions are faster at higher concentrations

(rate laws show this; more collisions)2) Reactions are faster at higher temperatures

(molecules are moving faster)

B. Activation Energy1) Chemical reactions happen much slower than predicted based on how

many collisions occur2) Only some of the collisions must lead to reaction

3) Activation Energy = the minimum energy required of a collision to cause reactiona) 2 BrNO 2 NO + Br2

b) The Br—NO bond energy is 243 kJ/molc) To break that bond, the collision must provide at least that much

energy

C. Potential Energy Diagrams = plot of energy over the course of a reaction

1. Reactant is on left side and the product is on the right side2. Activation Energy (Ea) is the amount of energy from the reactant to the

highest point (Transition State)3. Transition State (Activated Complex) = high energy species with a

structure intermediate between the reactant and the product. It is never isolated.

More Potential Energy Diagrams

1. NO2 + NO2 NO3 + NO

2. NO3 + CO NO2 + CO2

NO2 + CO NO + CO2

Rate = k[NO2]2

D. Factors Affecting the Reaction Rate1. Increased Temperature allows moremolecules to collide with the minimumamount of energy to react

2. Collision Frequency (z) depends on what kind and how many molecules are reacting

3. Molecular Orientation = how the colliding molecules are oriented during the collision also determines whether or not they will react

The Orientation Factor (p) is the fraction of collisions with the correct orientation for reaction. It is always between 01.

E. The Arrhenius Equation relates all of the factors to the rate constant k

1. Taking the natural log of each side gives us another form of the equation that gives a linear plot. lnk vs. 1/T gives straight line with slope = -Ea/R and intercept = ln(A)

k = rate constantA = frequency factor (combines z and

p)Ea = activation energy

T = temperature in KelvinsR = gas constant = 8.3145 J/K.mol

ln(A)T1

REln(k) a

2. Example: 2 N2O5 4 NO2 + O2 Ea?

3. For only 2 temperatures, the Arrhenius Equation can be rewritten:

4. Example: CH4 + 2 S2 CS2 + 2 H2S Ea?

T(oC) T(K) 1/T(K) k (s-1) ln(k)

20 293 3.41x10-3 2.0x10-5 -10.82

30 303 3.30x10-3 7.3x10-5 -9.53

40 313 3.19x10-3 2.7x10-4 -8.22

50 323 3.10x10-3 9.1x10-4 -7.00

60 333 3.00x10-3 2.9x10-3 -5.84

211

2 11lnTTR

Ekk a

k (L/mol.s) T (oC) T (K)

1.1 = k1 550 823 = T1

6.4 = k2 625 898 = T2

III. CatalysisA. Catalyst = compound that speeds up a chemical reaction without being

consumed itself1) It is not always possible to speed up a reaction by increasing the

concentrations or increasing the temperatures (ex: living things)2) Catalysts work by lowering the activation energy of a reaction

3) Heterogenous Catalyst: present in a different phase than the reacting molecules.a) This is often a solid not dissolved in the reaction solutionb) Reactants Adsorb on the surface to react (a sponge absorbs)c) Example: H2C=CH2 + Pt + H2 CH3CH3

4. Homogeneous Catalyst = Present in the same phase as the reacting molecules.A. Gaseous catalyst in a gas phase reactionB. Ozone Depletion is caused by Cl catalyst from chlorofluorocarbons in the

atmosphereO + O3 2 O2 (uncatalyzed = slow)CCl2F2 CClF2 + Cl (source of catalyst)

Cl + O3 ClO + O2 (elem. Step #1)O + ClO Cl + O2 (elem. Step #2)O + O3 2 O2 (catalyzed = fast)

B. Enzymes = biological catalysts1. Most reactions needed by an organism happen too slowly to be of use2. Enzymes are proteins that have been developed to catalyze almost

every biological reaction3. Examples: Carboxypeptidase A cleavage of other proteins

Sucrase cleavage of complex sugars

HO

OH

H

OH

CH2

CH CO2- NH2

+ C

NH2NH

COCHR

NH

C O

Zn2+

OC

O

Carboxypeptidase A

Sucrase