Chemistry 12
Notes on Chemical Kinetics
Chemistry 12 – Unit 3 – chapter 6 – Chemical Kinetics
Introduction to chapter 6
Kinetics:
how fast reactions go, and mechanisms, the paths molecules take in going from reactants to products.
Reaction Rates
A rate of a reaction is a rate at which a product is formed or a reactant is consumed during the reaction.
The rate of a reaction is usually expressed in terms of a change in concentration of one of the reactants or
products per unit time.
The main feature of a rate is that it must be able to be measured physically.
Methods of Measuring Reaction Rates
for gas product …
for ion product …
for coloured product …
for energy changes …
average reaction rate = change in concentration / elapsed time
Average Rate = Δ (change in) concentration
Δ (change in) time
r=
C2 – C1
t2 – t1
r=
C
t
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Chemistry 12
Notes on Chemical Kinetics
Practice Problem 1:
CaCO3 (s) + 2HCl (aq)
CO2 (g) + H2O (l) + CaCl2 (aq)
What are the rates of reaction with respect to the various reactants and products? The rate of reaction with
respect to HCl (rate of consumption) is 2.0 x 10-4 mol/L.s.
Answers
Rate of reaction of HCl:
Rate of reaction of CO2:
Rate of reaction of H2O:
6.1 Practice (p.350): 1,
6.1 Questions (p.361): 1, 2, 3, 4
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Chemistry 12
Notes on Chemical Kinetics
6.1 Determining the Average reaction Rate using Graphical Data

The average rate of reaction between two time points is equal to the slope of the secant line
(Secant line is a line that intersects two points on a curve)

The slope of a line is calculated by: ∆y (concentration) =
∆x (time)
Rate A = ∆ [A]
∆t
Instantaneous reaction rate: rate of a chemical reaction at a single point in time
 Slope of the tangent to the curve at a particular instant in time ( tangent is a straight line that
touches a curve at a single point and does not cross through the curve
HOMEWORK: Practice p.360 #1-2
Questions p. 361 #1-4
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Chemistry 12
Notes on Chemical Kinetics
6.2 – Factors Affecting Reaction Rate
The Six Factors Affecting Reaction Rate

Chemical nature of reactants

Catalyst

Concentration of reactants (solutions)

Surface area

Temperature

Pressure (gases)
Chemical Nature of Reactants
Concentration of Reactant
Temperature
Surface Area
Catalyst
Read 6.3 and answer questions p. 365 1, 2, 4
Problem Set #2 Due in two days
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Chemistry 12
Notes on Chemical Kinetics
6.3 – Collision Theory and Rate of Reaction
THE COLLISION THEORY AND ACTIVATION ENERGY
Effective vs ineffective collisions the molecules must collide so that the right atoms come into contact. No
bond can form if the molecules collide with the wrong orientation or with insufficient energy.
The orientation of collision
Reactions involving collisions between two species…
Activation Energy
Even if the species are orientated properly, you still won't get a reaction unless the particles collide with a
certain minimum energy called the activation energy of the reaction.
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Chemistry 12
Notes on Chemical Kinetics
Temperature:
The Boltzmann distribution is a thermodynamic equation that tells us what fraction of the molecules have a
certain amount of energy. As you know, at higher temperatures the average kinetic energy of the molecules
increases. Therefore, at higher temperatures more molecules have an energy greater than the activation energy-as shown in the figure below.
Theoretical Effect of Chemical Nature of Reactants

faster reactions … lower activation energies

slower reactions … higher … activation energies
Theoretical Effect of Catalysts

catalyst speeds up a chemical reaction without being consumed

catalyst provides alternative lower energy mechanism from same reactants to same products
… new intermediate steps have lower Ea

heterogeneous catalyst … different state than reactants e.g. Solid platinum catalyzes the
reaction between hydrogen gas and oxygen gas.

homogeneous catalyst … same state than reactants e.g. Cobalt (II) ions in solution catalyze the
reaction between aqueous tartrate ions and a solution of hydrogen peroxide.
Four criteria must be satisfied in order for something to be classified as catalyst.

Catalysts increase the rate of reaction.

Catalysts are not consumed by the reaction.

A small quantity of catalyst should be able to affect the rate of reaction for a large amount of
reactant.
Read 6.3 and answer questions p. 372 1, 2, 5
Problem Set #2
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Chemistry 12
Notes on Chemical Kinetics
6.3 – Rate Laws and Order of Reaction
Reaction Rate – The measure of how quickly a chemical reaction will proceed
Rate Constant (k) – A proportional constant for a chemical reaction which depends on temperature
(must be held constant)
Rate Law Equation – A mathematical equation which equates the rate of a reaction to the initial
concentrations of each reactant raised to some power.
Example:
2A +
B
+
C → 2D
Rate = k [A]x [B]y …
Note: If the exponent of the concentration of a reactant in the rate law is 0, then that reactant has no
effect on the rate of the reaction and is not included in the rate law equation.
In the example above the concentration of [C] has no effect on the overall rate of reactions and is not
included in the rate law equation.
Determining the Rate Law of a Reaction from Experimental Data
1) Observe experimental results in which the concentration of all but one reactant remained
constant.
2) Find the relationship between the change in concentration and the change of the reactions rate.
This relationship is a direct result of the exponential function of the reactant which is changing
in concentration.
3) Repeat steps one and two for any other reactants in the rate determining step.
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Chemistry 12
Notes on Chemical Kinetics
For example, let's use the method of initial rates to determine the rate law for the following
reaction:
whose rate law has the form:
rate = k[C3H6O]p[Br2]q
Solve for the unknowns p and q first…
I could’ve use x and y but why be boring
Using the following initial rates data, it is possible to calculate the order of the reaction for both
bromine and acetone:
To calculate the order of the reaction for bromine, notice that experiments 1 and 2 hold the
concentration of acetone constant while doubling the concentration of bromine. The initial rate of the
reaction is unaffected by the increase in bromine concentration, so the reaction is zero order in
bromine. Therefore q = 0.
Solve:….
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Chemistry 12
Notes on Chemical Kinetics
Initial rates data for the bromination of acetone
By similar reasoning, we can conclude that because the rate of reaction doubled when the
concentration of acetone was doubled (experiments 1 and 3) the reaction must be first order in
acetone. Therefore a = 1.
Solve:
…remember p=1, q=0
To calculate the value of k: (using data from trial 3)
p
q
rate = k[C3H6O] [Br2]
6.5 Practice (p.380, 381): 1, 2, 3, 4, 5
6.5 Questions (p. 382): 1, 2. 3
Problem Set #3 Due in two days
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Chemistry 12
Notes on Chemical Kinetics
6.6 ORDERS OF REACTION AND MECHANISMS
Reaction mechanisms
In any chemical change, some bonds are broken and new ones are made. Quite often, these changes are too
complicated to happen in one simple stage. Instead, the reaction may involve a series of small changes one
after the other.
The rate-determining step
The overall rate of a reaction (the one which you would measure if you did some experiments) is controlled by
the rate of the slowest step. The faster step(s) is in a sense waiting around for the slow step to happen.
The slow step of a reaction is known as the rate-determining step.
Rules for proposing reaction mechanisms … properties of a reaction mechanism
Reaction mechanisms are only “best guesses” at the behaviour of molecules … 3 rules for proposing a
mechanism:



Each step is elementary. There are no more than 3 reactant molecules.
The slowest or rate-determining step has a molecularity consistent with the rate equation
Elementary steps add up to the overall equation.
Example:
4 HBr(g) + O2(g) → 2H2O(g) + 2Br2(g)
• the rate law equation has been proven experimentally as:
r= k [HBr][O2]
Theorized reaction:
HBr(g) + O2 → HOOBr(g)
HOOBr(g) + HBr(g) → 2HOBr(g)
HOBr(g) + HBr(g) →H2O(g) + Br2(g)
HOBr(g) + HBr(g) →H2O(g) + Br2(g)}
(slow)
(fast)
(fast)
(fast)
____________________________
 Each step is elementary. There are no more than 3 reactant molecules.
The most was 2 molecules.
 The slowest or rate-determining step has a molecularity consistent with the rate equation
HBr(g) + O2 → HOOBr(g)
(slow) reactants matches with
r= k [ HBr ] [ O2 ]
 Elementary steps add up to the overall equation.
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Chemistry 12
Notes on Chemical Kinetics
4 HBr(g) +O2(g) → 2H20(g) +2Br2
The potential energy diagram appears as:
PROBLEMS and SOLUTIONS
1. Identify the intermediates and the catalysts (if any) in the following 3-step mechanism.
Solution for Problem 1
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Chemistry 12
Notes on Chemical Kinetics
2. Consider the reaction below.
2NO(g) + 2H2 (g)
N2 (g) + 2H2O(g)
The experimentally determined rate law is:
Rate = k[NO]2[H2]
A chemist proposes the mechanism below for the reaction.
Step 1:
2NO(g) + H2 (g)
N2O(g) + H2O(g)
(slow)
Step 2:
N2O(g) + H2 (g)
N2 (g) + H2O(g)
(fast)
Determine whether the proposed mechanism is reasonable.
Solution for Problem 2
Rules for mechanism:
1)
.
2)
.
3)
Adding the 2 steps
.
Step 1:
2NO(g) + H2 (g)
N2O(g) + H2O(g)
Step 2:
N2O(g) + H2 (g)
N2 (g) + H2O(g)
The reaction mechanism seems reasonable because neither step has more than 3 reactant molecules;
the rate law for the overall reaction is the same as the rate law for the rate-determining step; and the
elementary steps add up to the overall equation.
6.4 Practice (p.389, 390): 1abcd, 2ab, 3
6.4 Questions (p.387): 1, 3, 5 - 8
Problem Set #3
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Chemistry 12
Notes on Chemical Kinetics
SCH 4U1-____
Name:________________________
Chemistry 12: PROBLEM SET #2 - Measuring Reaction Rates
1.
A chemist wishes to determine the rate of reaction of zinc with hydrochloric acid. The equation
for the reaction is:
Zn(s) + 2HCl(aq)
H2(g) + ZnCl2(aq)
A piece of zinc is dropped into 1.00 L of 0.100 M HCl and the following data were obtained:
Time
of Zinc
0s
4.0 s
8.0 s
12.0 s
16.0 s
20.0 s
Mass of Zinc
Moles of Zinc
Molar Concentration
0.016 g
0.014 g
0.012 g
0.010 g
0.008 g
0.006 g
a) Using proper technique, graph the change of concentration during the 20 seconds.
b)
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Chemistry 12
Notes on Chemical Kinetics
b) Calculate the overall average Rate of Reaction in moles of Zinc per litre consumed per
second.
Answer___________________
c) Calculate the instantaneous Rate of Reaction in moles of Zinc per litre of solution consumed
per second at 15 seconds. Show your work on the graph.
Answer___________________
2.
When magnesium is reacted with dilute hydrochloric acid (HCl), a reaction occurs in which
hydrogen gas and magnesium chloride is formed.
a) Write a balanced formula equation for this reaction.
_____________________________________________________________________
b) If the rate of consumption of magnesium is 5.0 x 10-9 mol/L.s, find the rate of
consumption of HCl in moles/L.s.
Answer___________________
c) If the rate of consumption of magnesium is 5.0 x 10-9 mol/L.s, find the rate of production
of H2 in mol/L.s.
Answer___________________
3.
Given the reaction:
CO2(g)
+
colourless
NO (g)
CO(g)
colourless
+
colourless
NO2(g)
brown
Suggest a method which could be used to monitor the rate of this reaction.
Why wouldn’t total pressure be a good way to monitor the rate of this reaction?
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Chemistry 12
4.
Notes on Chemical Kinetics
Equal volumes of Fe2+(aq) and C2O42-(aq) are separately reacted with 0.10 M MnO4-(aq) and the
following data were obtained:
Reactant
Concentration
Temperature
Time for complete reaction
Fe2+
0.20 M
25°C
1.6 s
C2O42-
0.40 M
35°C
17.0 s
Compare the rates of reaction
Explain why these results are obtained in terms of factors that affect the number of collisions
and effective collisions.
5.
The longer the time of reaction, the ____________________________ the rate of reaction.
6.
On the following set of axes, draw the shape of the curve you would expect if you plotted the
[HCl] vs. Time, starting immediately after the two reactants are mixed. The equation for the
reaction is:
Mg(s)
+
2HCl(aq)

H2(g) +
MgCl2(aq)
[HCl]
Explain how you got that particular shape. Be
detailed.
Time
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Chemistry 12
7.
Notes on Chemical Kinetics
Give some examples of situations where we might want to increase the rate of a particular
reaction.
_____________________________________________________________________
_____________________________________________________________________
8.
Give some examples of situations where we might want to decrease the rate of a particular
reaction.
_____________________________________________________________________
_____________________________________________________________________
9.
Give two reasons why water is effective at putting out fires. Use concepts learned in this unit so
far.
_____________________________________________________________________
_____________________________________________________________________
10. Consider the rate of the following reaction:
Fe(s) + 2HCl(aq)
H2(g) + FeCl2(aq)
a) Is rate dependent on temperature? ____________________. Explain your answer.
_____________________________________________________________________
b) Is rate dependent on surface area? ___________________. Explain your answer.
_____________________________________________________________________
11. Consider the rate of the following reaction:
2NaOCl(aq)
2NaCl(aq) + O2(g)
a) Is rate dependent on temperature? ____________________. Explain your answer.
_____________________________________________________________________
b) Is rate dependent on surface area? ____________________. Explain your answer.
_____________________________________________________________________
c) Is rate dependent on [NaOCl]? _______________________. Explain your answer.
_____________________________________________________________________
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Chemistry 12
Notes on Chemical Kinetics
CHEMICAL KINETICS PROBLEM SET #3
1. The reaction A + B —> C was studied using the initial rate method with the following results:
[A] (M)
0.030
0.060
0.060
[B] (M)
0.030
0.030
0.090
Initial Rate (M/s)
0.75 x 10-4
3.0 x 10-4
3.0 x 10-4
a) What is the order with respect to A? __________________________________
b) What is the order with respect to B? __________________________________
c) What is the value of the rate constant (including units)? __________________________________
d) What is the rate equation? __________________________________
2. The reaction A + B —> C was studied using the initial rate method with the following results: (note: negative orders
and orders with fractions/decimals are possible)
[A] (M)
0.030
0.060
0.060
[B] (M)
0.030
0.060
0.030
Initial Rate (M/s)
0.30 x 10-4
2.4 x 10-4
9.6 x 10-4
a) What is the order with respect to B? __________________________________
b) What is the order with respect to A? __________________________________
c) What is the numerical value of the rate constant (including units)? __________________________________
3. The experimental rate law for the reaction 2 A + B —> 3 C + D is:
Rate = k [A] [B]3
a) What is the overall order of this reaction? __________________________________
b) If the concentration of A is tripled, what happens to the reaction rate? __________________________________
c) If the concentration of B is doubled, what happens to the reaction rate? __________________________________
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Chemistry 12
Notes on Chemical Kinetics
4. The initial rate of the reaction 2A + 2B ---> C + D is determined for different initial conditions, with the results listed
in the following table:
Experiment #
[A] , M
[B] , M
Initial rate, M/s
1
0.185
0.133
3.35 x 10-4
2
0.185
0.266
1.35 X 10-3
3
0.370
0.133
6.75 X 10-4
Find the full rate law for this reaction. Give all proofs.
5. The following data are for the next questions.
A(gas) + 2 B(gas) + 3 C(gas) ------> Z(gas) + 2 Y(gas) Temp = 50.0 oC
Experiment
1
2
3
4
5
[A]
0.10
0.10
0.20
0.20
0.050
[B]
0.020
0.03
0.02
0.02
0.01
[C]
0.040
0.04
0.04
0.16
0.08
Rate of reaction
10 M/hr
15 M/hr
80 M/hr
160 M/hr
??
Solve the unknown rate of experiment 5. Show your work.
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Chemistry 12
Chemistry 12
Notes on Chemical Kinetics
Unit 3 – Chapter 6 - Chemical Kinetics
Chemistry 12 Problem Set #4 - Reaction Mechanisms
1. It is known that compounds called chlorofluorocarbons (C.F.C.s) (eg. CFCl3) will break up in the
presence of ultraviolet radiation, such as found in the upper atmosphere, forming single chlorine atoms:
CFCl3  CFCl2 + Cl
The Cl atoms then react with ozone (O3) as outlined in the following mechanism.
Step 1: Cl + O3  ClO + O2
Step 2: ClO + O  Cl + O2 (single "O" atoms occur naturally in the atmosphere.)
a) Write the equation for the overall reaction. (Using steps 1 and 2)
b) What is the catalyst in this reaction?
c) Identify an intermediate in this reaction
d) Explain how a small amount of chlorofluorocarbons can destroy a large amount of ozone.
e) What breaks the bond in the CFCl3 and releases the free Cl atom?
2. Consider the following mechanism:
Step 1: H2O2 + I-  H2O + IO(slow)
Step 2: H2O2 + IO-  H2O + O2 + I- (fast)
a) Give the equation for the overall reaction.
b) What acts as a catalyst in this mechanism?
c) What acts as an intermediate in this mechanism?
3. What is meant by the rate determining step in a reaction mechanism?
4. What is meant by a reaction mechanism?
5. Given the reaction:
4HBr + O2  2H2O + 2Br2
a) Would you expect this reaction to take place in a single step?
Why or why not?
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Chemistry 12
Notes on Chemical Kinetics
b) This reaction is thought to take place by means of the following mechanism:
Step 1: HBr + O2  HOOBr
(slow)
Step 2: HBr + HOOBr  2HOBr
(fast)
Step 3: 2HBr + 2HOBr  2H2O + 2Br2
(fast)
c) Identify the two intermediates
d) A catalyst is discovered which increases the rate of Step 3. How will this affect the rate
of the overall reaction?
Explain your answer.
e) A catalyst is discovered which increases the rate of Step 1. How will this affect the rate
of the overall reaction?
Explain your answer.
f) Which step has the greatest activation energy?
g) How many "bumps" will the potential energy diagram for the reaction mechanism have?
h) Which step is called the rate determining step in this mechanism?
i) On the set of axes below, draw the shape of the curve you might expect for the reaction
in this question. The overall reaction is exothermic! Make sure you get the "bumps" the
correct relative sizes.
Potential
Energy
(kJ)
Progress of Reaction
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Chemistry 12
Notes on Chemical Kinetics
6. The equation for an overall reaction is:
I- + OCl-  IO- + Cla) The following is a proposed mechanism for this reaction. One of the species has been left out.
Determine what that species is and write it in the box. Make sure the charge is correct if it has one!
Step 1: OCl- + H2O  HOCl + OH( fast )
Step 2: I- + HOCl  IOH + Cl( slow )
Step 3: IOH + OH-  ______ + H2O ( fast )
b) Which species in the mechanism above acts as a catalyst?
c) Which three species in the mechanism above are intermediates?
d) Step ______________ is the rate determining step.
e) On the set of axes below, draw the shape of the curve you might expect for the reaction in this question.
The overall reaction is endothermic! Make sure you get the "bumps" the correct relative sizes.
Potential
Energy
(kJ)
Progress of Reaction
7. A certain chemical can provide a reaction with an alternate mechanism having a greater activation energy.
What will happen to the rate of the reaction when this chemical is added?
Explain your answer.
8. The following overall reaction is fast at room temperature:
H+ + I- + H2O2  H2O + HOI
A student proposes the following two-step mechanism for the above reaction:
Step 1 : H+ + H+ + H2O2  H4O2 2+
Step 2 : H4O2 2+ + I-  H2O + HOI + H+
Would you agree or disagree with this proposed mechanism? (i.e. Is this mechanism reasonable?) Explain
your answer.
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Chemistry 12
Notes on Chemical Kinetics
9. Consider the following reaction:
CO + NO2  CO2 + NO
a) The first step in each of two proposed reaction mechanisms for the above reaction is listed below. If
each proposed reaction mechanism consists of only two steps, determine the second step for each
mechanism.
Proposed Mechanism One:
Step 1: 2NO2  NO3 + NO
(slow)
Step 2: __________________________________________ (fast)
Proposed Mechanism Two:
Step 1: 2NO2  N2O4
(fast)
Step 2: __________________________________________ (slow)
b) Experimental data show that the rate of the reaction is not affected by a change in the [CO]. Which of
these two mechanisms would be consistent with these data?
Explain your answer.
10.) Questions for Diagram #1
1. Which letter represents the activation
energy for the forward reaction? A, B,
C, or D
2. What is the value of the activation energy
in kJ for the forward reaction? 120 kJ, 240
kJ, 160 kJ, or none of the above
3. What letter represents the H for the
forward reaction, and what is it's value?
B: 240 kJ, B: 120 kJ, A: 160 kJ, or D: 120 kJ
4. The forward reaction is exothermic. True
or False
Answers for Diagram #1 … 1:A; 2:160 kJ;
3:D:120 kJ; 4:False
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Chemistry 12
Notes on Chemical Kinetics
Course Code: SCH 4U1- _____Name: _____________________________________
Date: ______________
CHEMICAL KINETICS PROBLEM SET #5
Deriving Rate Laws from Mechanisms
An accident or road construction on a highway slows all the traffic because limitations imposed by
the accident and constructions apply to all cars on that road. The narrow passing stretch limits the
speed of the traffic. If several steps are involved in an overall chemical reaction, the slowest step
limits the rate of the reaction. Thus, a slow step is called a rate-determining step. Note the presence
of the fluorine atom, F, an intermediate in equation ii below. F is an intermediate because it is both
created and destroyed in the mechanism and does not appear in the overall equation. The following
example illustrates the method of deriving rate laws from the proposed mechanism.
Problem 1:
If the reaction 2 NO2 + F2  2 NO2F follows the mechanism,
i. NO2 (g) + F2 (g)  NO2F(g) + F(g)
ii. NO2 (g) + F(g)  NO2F(g)
(slow)
(fast)
What is the rate law?
Solution
Since step i is the rate-determining step, the rate law is
rate = k [NO2] [F2]
Since both NO2 and F2 are reactants, this is the rate law for the reaction.
Determining the Reasonableness of a Possible Reaction Mechanism
Problem 2
Consider the reaction below.
CO(g) + NO2 (g) → CO2 (g) + NO(g)
It has been experimentally determined that this reaction takes place
according to the rate law, rate = k[NO2]2. Therefore, a possible
mechanism which this reaction takes place is:
Step 1:
Step 2:
2 NO2 (g) → NO3 (g) + NO(g)
NO3 (g) + CO(g) → NO2 (g) + CO2 (g)
(slow)
(fast)
Determine whether the proposed mechanism is reasonable.
Solution
\
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Chemistry 12
PROBLEM # 3:
Notes on Chemical Kinetics
What is the rate law? Consider the following mechanism as O3 decomposes.
Given the overall reaction shown
2 O3 → 3 O2
and the following mechanism,
O3 → O2 + O2(slow)
O2- + O3 → 2O2
(fast)
What would the observed rate law for the reaction be?
Solution
PROBLEM # 4:
Consider the above reaction again.
CO(g) + NO2 (g) → CO2 (g) + NO(g)
It has been experimentally determined that this reaction takes place
according to the rate law, rate = k[NO2]2. Therefore, a possible
mechanism by which this reaction takes place is:
Step 1:
Step 2:
2 NO3 (g) → 2 NO2 (g) + O2 (g)
NO3 (g) + CO(g) → NO2 (g) + CO2 (g)
(slow)
(fast)
Determine whether the proposed mechanism is reasonable.
Solution and Discussion:
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