Chapter 14 Chemical Kinetics

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```					Chapter 14 Chemical Kinetics

Kinetics – study of the speed or rate at which a
reaction occurs
(see handout for “collision theory” for key parts
of collision theory and factors that affect speed
of reaction – this is section 14.4)

Section 14.1
Reaction Rate
Consider the following reaction
2N2O5(g)  4NO2(g) + O2(g)

As the reaction proceeds, the amount of N2O5,
NO2, and O2 is recorded and plotted on the
graph.
Number of moles

time

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Reaction rate can be defined in two ways:
a. ∆(moles products)
∆time

b. - ∆(moles reactants)
∆time

Why do we use the negative sign for the
reaction rate in terms of reactants?

When expressing reaction rate, you must also
take into account the stoichiometry of the reaction.
We want reaction rate to be equivalent no matter
which reactant or product we study. Therefore, we
will write reaction rates with the coefficients in
mind.
Ex: Write the term for the reaction rate of the
following chemical equations in terms of each
reactant and product.
1. 2N2O5(g)  4NO2(g) + O2(g)

2
2. 2HI(g)  H2(g) + I2(g)

Reaction rates are often calculated in terms of
concentration rather than absolute number of
moles. This works well when the volume of a
container is fixed.

For example, for the reaction
2N2O5(g)  4NO2(g) + O2(g)

The graph in terms of concentration will be
similar as the graph in terms of number of
moles. We will consider just the change in
dinitrogen pentoxide.

How does the rate of the reaction change
during the course of the reaction?
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Two ways to calculate reaction rates:
1. average rate over a given period of time.

2. instantaneous rate

Practice Problem:
The isomerization of methyl isonitrile, CH3NC, to
acetonitrile, CH3CN, was studied in the gas phase
at 215°C and the following data were obtained:

Time (s) [CH3NC] (M)         Average rate of
reaction (M/s)
0         0.0165
2000      0.0110
5000      0.00591
8000      0.00314
12000     0.00074

Graph the data above. Draw a tangent to the line at
t = 4000. Determine the instantaneous rate at this
time.
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Section 14.2
rate law – equation that shows how the rate of a
reaction is dependent on the concentrations of
reactants

How to calculate rate law?
 you need to study the initial reaction rate
(instantaneous reaction rate at t = 0)
 vary the concentrations of reactants and
products
 determine how the concentration affects the
rate
Example:
Consider the reaction
NH4+ + NO2-  N2 + 2H2O

Trial     Initial        Initial      Initial rate
[NH4+]         [NO2-]       (M/s)
1          0.01           0.2        5.4 x 10-7
2          0.02           0.2       10.8 x 10-7
3          0.04           0.2       21.5 x 10-7
4           0.2          0.02       10.8 x 10-7
5           0.2          0.04       21.6 x 10-7
6           0.2          0.06       32.4 x 10-7
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According to the data, how is the rate of reaction
affected by:
1. concentration of NH4+

2. concentration of NO2-

Expressing this as a rate law:

Determination of k (rate constant):

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Practice example:
Consider the following reaction:
2ClO2 + 2OH-  ClO3- + ClO2- + H2O

Trial        [ClO2]        [OH-]         Initial rate
(M)           (M)           (M/s)
1            0.06          0.03          0.0248
2            0.02          0.03          0.00276
3            0.02          0.09          0.00828

1. Determine the rate law.

2. Determine the value of the rate constant.

3. How are the exponents in the rate law
related to the coefficients in the balanced
chemical equation?

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Important terms:
1. reaction order

2. overall reaction order

Note: A reaction does not have to be just first or
second order in terms of a specific reactant. It can
be 3rd, 4th, etc.
Example Problem:
Consider the following reaction: A + 2B  C

rate law:   rate = k[A]1/2[B]0

What is the affect on the rate of the reaction:
a. when concentration of A is doubled?

b. when concentration of B is doubled?

Notice: you cannot determine the rate law
of a reaction based solely on the balanced
chemical equation.
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Section 14.3
You should be able to tell if a reaction is first,
second, or zero order in respect to a reactant based
on the graph of the reactant over a period of time.
(see handout – important kinetic relationships)

Section 14.5
balanced chemical equations: give the mole ratios
of reactants required to complete a chemical
reaction and the mole ratios of products produced

reaction mechanism: describes the mechanism
(steps) by which a reaction occurs
two types of mechanisms:
1. elementary steps

different molecularities:
a. unimolecular

b. bimolecular

c. termolecular

2. multistep mechanisms
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When determining a reaction mechanism, two
things must be true:
1. steps must add up to the overall reaction
2. the rate law proposed by the reaction
mechanism must match experimental
observations

Ex #1:
Consider the following reaction mechanism:
A+BC              (slow)
C + 2B  D         (fast)

Overall reaction: A + 3B  D
rate law = [A][B]     (bimolecular process)

1.     Does this reaction fit the criteria to be a
reaction mechanism?

2.     What is the rate determining step of this
reaction? Why?

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3.      If the rate law of the second step was
found to be rate = k[B]2, what would this
tell you about the second step?

Ex #2: It is harder to determine the rate law when
the first step is not the rate determining step.
Consider the following reaction mechanism

NO(g) + Br2(g)  NOBr2                (fast)

NOBr2(g) + NO(g)  2NOBr              (slow)

rate law of overall reaction = k[NO]2[Br2]

1. Write the balanced chemical equation for the
overall reaction.

2. Write the rate law for each of the above
elementary steps.

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3. The rate law for an overall reaction must not
include the intermediates (must only include
reactants). Rewrite this rate law based on the
reactants.

4. What does the rate law and the reaction
mechanism tell you about the molecularity of
the processes?

Note: While you can determine the rate law of
an elementary step, you cannot tell the rate law of
a reaction by just looking at its balanced
chemical equation. Why?

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Practice Problem #1
The following mechanism is proposed for the
reaction of NO with H2 to form N2O and H2O:

NO + NO  N2O2
N2O2 + H2  N2O + H2O

a. show that the elementary steps of the proposed
mechanism add to provide a balanced equation for
the reaction.

b. write a rate law for each elementary step in the
mechanism.

c. Identify the intermediates.

d. the observed rate law is rate = k[NO]2[H2]. If
the proposed mechanism is correct, what can we
conclude about the relative speeds of the first and
second steps?

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