1. Ch 13 Reaction Mechanisms
Mechanism = the series of discrete elementary steps leading from reactants to products
The slowest of these elementary steps determines how fast the overall reaction can be
The study of kinetics helps us determine what the mechanism is
Example: NO CO NO CO2
Rate = k[NO2]2
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
We think the mechanism for this reaction is:
NO NO NO NO
NO CO NO CO2

2. Definitions:
Intermediate = molecule produced in one step of a mechanism and used up in another step; it is not isolated as a final product
For the following mechanism, NO3 is an intermediate
NO NO NO NO
NO CO NO CO2
Elementary Step = reaction whose rate law is written directly from its molecularity
Molecularity = how many molecules collide in the reaction
Unimolecular = only one molecule is involved in the step
Bimolecular = two molecules collide in the step
Termolecular = three molecules collide in the step (this is very rare because of the low probability)

3. 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.
NO NO NO NO rate = k[NO2]2
NO CO NO CO2 rate = k[NO3][CO]
Rate Determining Step (rds) = the slowest step in a mechanism
One step has to be the slowest.
The overall reaction can’t go faster than that one step, so it determines the rate of the reaction
The overall reaction rate will be the same as that of the r.d.s
Example: NO CO NO CO2
Step 1 is the slowest step
Overall rate = k[NO2]2
Like the narrow point of a funnel, the rds determines the rate

4. Mechanisms can’t be proved, only disproved
Since we can’t actually see individual molecules react, we can’t prove the exact steps they go through
When chemists suggest a mechanism, it must meet two criteria
The individual steps must add up to the observed reaction equation
NO NO NO NO
NO CO NO CO2
NO CO NO CO2
The proposed mechanism must match the observed rate law
Observed rate law is rate = k[NO2]2
Step 1 would give this rate law if it is r.d.s.
If either of these criteria aren’t met, the proposed mechanism is wrong
Example: 2 NO F NO2F
Step 1 NO F NO2F F (slow)
Step 2 F NO NO2F (fast)
Observed rate law: rate = k[NO2][F2]

5. A Model for Chemical Kinetics
The Collision Model = Molecules must collide to react
Reactions are faster at higher concentrations
(rate laws show this; more collisions)
2) Reactions are faster at higher temperatures
(molecules are moving faster)
Activation Energy
Chemical reactions happen much slower than predicted based on how many collisions occur
Only some of the collisions must lead to reaction
Activation Energy = the minimum energy required of a collision to cause reaction
2 BrNO NO Br2
The Br—NO bond energy is 243 kJ/mol
To break that bond, the collision must provide at least that much energy

6. Potential Energy Diagrams = plot of energy over the course of a reaction
Reactant is on left side and the product is on the right side
Activation Energy (Ea) is the amount of energy from the reactant to the highest point (Transition State)
Transition State (Activated Complex) = high energy species with a structure intermediate between the reactant and the product. It is never isolated.

7. More Potential Energy Diagrams
Rate = k[NO2]2
1. NO NO NO NO
2. NO CO NO CO2
NO CO NO CO2

8. D. Factors Affecting the Reaction Rate
Increased Temperature allows more
molecules to collide with the minimum
amount 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 01.

9. 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 constant
A = frequency factor (combines z and p)
Ea = activation energy
T = temperature in Kelvins
R = gas constant = J/K.mol

10. Example: N2O NO O2 Ea?
For only 2 temperatures, the Arrhenius Equation can be rewritten:
Example: CH S CS 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
k (L/mol.s)
T (oC)
T (K)
1.1 = k1
550
823 = T1
6.4 = k2
625
898 = T2

11. Catalysis This is often a solid not dissolved in the reaction solution
Catalyst = compound that speeds up a chemical reaction without being consumed itself
It is not always possible to speed up a reaction by increasing the concentrations or increasing the temperatures (ex: living things)
Catalysts work by lowering the activation energy of a reaction
Heterogenous Catalyst: present in a different phase than the reacting molecules.
This is often a solid not dissolved in the reaction solution
Reactants Adsorb on the surface to react (a sponge absorbs)
Example: H2C=CH Pt + H CH3CH3

12. Enzymes = biological catalysts
Homogeneous Catalyst = Present in the same phase as the reacting molecules.
Gaseous catalyst in a gas phase reaction
Ozone Depletion is caused by Cl catalyst from chlorofluorocarbons in the atmosphere
O O O2 (uncatalyzed = slow)
CCl2F CClF Cl (source of catalyst)
Cl O ClO O2 (elem. Step #1)
O ClO Cl O2 (elem. Step #2)
O O O2 (catalyzed = fast)
Enzymes = biological catalysts
Most reactions needed by an organism happen too slowly to be of use
Enzymes are proteins that have been developed to catalyze almost every biological reaction
Examples: Carboxypeptidase A cleavage of other proteins
Sucrase cleavage of complex sugars

13. Carboxypeptidase A

14. Sucrase