Chapter 17Reaction Kinetics

17-1 The Reaction Process

How did How did you you

meet? meet?

Can you Can you remember remember

the first time the first time you ever you ever made a made a friend?friend?

What had to What had to happen happen

before the before the friendship friendship

could could begin?begin?

Eye Eye ContactContact

Mutual Mutual FriendFriend

Accidentally Accidentally Bumped into Bumped into each othereach other

Collision Theory• In order for a reaction to occur particles

must collide in:

1. A specific orientation and

2. with enough energy

Activation Energy

• The amount of energy required for a reaction to occur

Activation Energy• Activation energy - the amount of

energy the particles must have when they collide to force a reaction to occur.

Activation Energy

Products

Reactants

Reaction Pathways

The products have less

energy than the reactants. The rxn released

energy (heat) = exothermic

∆H will be negative since energy has left

the system

Reaction Pathways

The products have more energy than

the reactants. The rxn

absorbed energy (heat)

= endothermic

∆H will be positive since energy has been added to the

system

Practice

• Draw and label the energy diagram for a reaction in which ΔE = 30 kJ/mol, Ea = 40 kJ/mol. Place reactants at energy level zero. Indicate determined values of ΔEforward, ΔEreverse & Ea’

Reaction Mechanisms

• Step-by-step sequence of rxns in order to obtain a final product

Proposed Mechanism for Ozone Depletion via Free Chlorine Atoms Created by Decomposition of CFCs

Step 1) Cl + O3 → ClO + O2

Step 2) 2 ClO → ClOOClStep 3) ClOOCl → ClOO + Cl

Step 4) ClOO → Cl + O2

Proposed Mechanism for Ozone Depletion via Free Chlorine Atoms Created by Decomposition of CFCs

Step 1) Cl + O3 → ClO + O2

Step 2) 2 ClO → ClOOClStep 3) ClOOCl → ClOO + Cl

Step 4) ClOO → Cl + O2

Mechanisms

Intermediates

overall rxn

Mechanisms

SlowSlow

FastFast

FastFast

FastFast

Rate Determining StepRate Determining Step

overall rxn

Catalysts vs. Intermediatesoverall rxn

Catalysts appear 1st as a reactant and then as a product during a mechanism.

Intermediates appear 1st as a product and then as a reactant during a mechanism.

Chapter 17Reaction Kinetics

17-2 Reaction Rate

How can we increase the rate of a reaction?

1. Increase Surface Area

2. Increase Temperature

3. Increase Concentration

4. Increase in Pressure

5. Add a Catalyst

Surface Area

• Increase the surface area allows for a greater chance for effective collision

Temperature

• An increase in temperature will cause particles to move at a higher velocity resulting in more effective collisions

Concentration

• An increase in concentration will also cause an increase in the chance that effective collisions will occur

Pressure

• Increasing the pressure of a gas system will cause more frequent collisions

Catalysts• Adding a catalyst lowers the amount of

activation energy required

Catalysts

Reactants

Catalyst

Slow

Rate Determining Step

Rate Laws

Rate = k[HBr][ORate = k[HBr][O22]]

• An equation that relates the rxn rate and the concentration of reactants

Rate Laws

• If no mechanism is given, then…

2H2H22 + 2NO + 2NO N N22 + 2H + 2H22OO

Rate = k[HRate = k[H22]]22[NO][NO]22

Rate Orders• 0, 1st and 2nd order rates• Order is dependent upon what will yield a

straight line0 order

2nd order

ln [

reac

tant

s] [re

acta

nts]

1st order 1/[r

eact

ants

]

Rate Orders

• 1st order: reaction rate is directly proportional to the concentration of that reactant

• 2nd order: reaction rate is directly proportional to the square of that reactant

• 0 order: rate is not dependant on the concentration of that reactant, as long as it is present.

For Individual Components:For Individual Components:

Rate Orders

• Overall reaction orders is equal to the sum of the reactant orders.

• Always determined experimentally!

For Overall Order:For Overall Order:

Calculating for k

A + 2B A + 2B C C

Rate = k[A][B]Rate = k[A][B]22

Experiment Initial [A] Initial [B] Rate of Formation of

C

1 0.20 M 0.20 M 2.0 x 10-4 M/min

2 0.20 M 0.40 M 8.0 x 10-4 M/min

3 0.40 M 0.40 M 1.6 x 10-3 M/min

What is the value of k, the rate What is the value of k, the rate constant?constant?

Calculating for k

Rate = k[A][B]Rate = k[A][B]22

Experiment Initial [A] Initial [B] Rate of Formation

of C

1 0.20 M 0.20 M 2.0 x 10-4 M/min

2 0.20 M 0.40 M 8.0 x 10-4 M/min

3 0.40 M 0.40 M 1.6 x 10-3 M/min

2.0 x 10-4 = k[0.20][0.20]2

2.0 x 10-4 = k(0.008)k = 2.50 x 10-2 min-1 M-2

Practice1. In a study of the following reaction:

2Mn2O7(aq) → 4Mn(s) + 7O2(g)

When the manganese heptoxide concentration was changed from 7.5 x 10-5 M to 1.5 x 10-4 M, the rate increased from 1.2 x 10-4 to 4.8 x 10-4. Write the rate law for the reaction.

2. For the reaction:

A + B → C

When the initial concentration of A was doubled from 0.100 M to 0.200 M, the rate changed from 4.0 x 10-5 to 16.0 x 10-5. Write the rate law & determine the rate constant for this reaction.

Rate = k[Mn2O7]2

Rate = k[A]2

Constant = 4.0 x 10-3 M/s

More Practice3. The following reaction is first order:

CH3NC(g) → CH3CN(g)

The rate of this reaction is 1.3 x 10-4 M/s when the reactant concentration is 0.040 M. Predict the rate when [CH3NC] = 0.025.

4. The following reaction is first order:

(CH2)3(g) → CH2CHCH3 (g)

What change in reaction rate would you expect if the pressure of the reactant is doubled?

New Rate = 8.1 x 10-5 M/s

An increase by a factor of 2

Even More Practice5. The rate law for a single step reaction that forms one product, C is R

= k[A][B]2. Write the balanced reaction of A & B to form C.

6. The rate law of a reaction is found to be R = k[X]3. By what factor does the rate increase if the concentration of X is tripled?

7. The rate of reaction, involving 2 reactants, X & Z, is found to double when the concentration of X is doubled, and to quadruple when the concentration of Z is doubled. Write the rate law for this reaction.

A + 2B → C

R = k[X][Z]2

The rate will increase by a factor of 27