Chemistry 12
Practice Problems
Chpt. 17, 18, 19
1.
Consider the following reaction,
2 A + B + 2D
>
E+F
with the following mechanism,
(1)
2A+B
<>
C
(2)
C+2D
>
E+F
(a) Derive a general rate law for this reaction using the steady state approximation.
(b) Write a simpler rate law that is valid when [D] is sufficiently small so that,
k2[D]2 << k –1 .
(c) Write a simpler rate law that is valid when [D] is sufficiently small so that,
k2[D]2 >> k –1 .
2.
The rate constant for the decomposition of gaseous azomethane,
C2 H6 N2 (g)
>
N2 (g) + C2 H6 (g)
Is 40.8 min-1 at 425OC. Find the number of moles azomethane and the number of moles of
nitrogen 0.0500 minutes after 2.00 g of azomethane is introduced. Assume that the
decomposition reaction is first order and that the reverse reaction can be ignored (Note that the
first order integrated rate equation can be expressed in terms of the number of moles of the
reacting species instead of the concentration).
3.
Given the following reaction,
A+2B
>
C+2D
And the following data,
Experiment
[A] (mol/L)
[B] (mol/L)
1
2
3
4
0.10
0.30
0.30
0.40
0.10
0.30
0.10
0.20
Initial Rate
(mol/L min)
2.0 x 10-4
6.0 x 10-4
2.0 x 10-4
4.0 x 10-4
Where the rate is defined,
Rate = -[A]/ t .
(a) What is the rate law?
(b) What is the value of the rate constant?
4.
The rate constant of a certain reaction is tripled when the temperature is increased
from 10.0OC to 30.0OC. Calculate the activation energy for this reaction.
5.
Consider the following reaction.
2 N2O5 (g)
>
2 N2O4 (g) + O2 (g)
Given that,
Ea = 103 kJ/mol
and,
k = 0.0900 min-1 at 328 K
calculate the value of k at 338 K.
6.
Consider the following reaction.
2 NO2 (g) + F2 (g) > 2 NO2F (g)
The mechanism is,
(1)
NO2 + F2
>
NO2F + F
(slow)
(2)
NO2 + F
>
NO2F
(fast)
What is the rate law for this reaction?
7.
(a)
Rate = k[NO2]2[F2]
(b)
Rate = k[NO2][F2]
(c)
Rate = k[NO2][F]
(d)
Rate = [ NO2F]/ t
Consider the reaction for the decomposition of ozone,
2 O3 (g)
>
3 O2 (g)
The mechanism for the reaction is,
(1)
O3
<>
(2)
O3 + O
O2 + O
>
2 O2
(a)
Use the steady state approximation to derive a general rate law for this reaction.
(b)
Derive a simpler rate law that is valid when k-1[O2] >> k 2[O3]
(c)
How is the rate of decomposition of ozone affected by an increase in the concentration
of oxygen?
8.
An evacuated chamber is filled with NOCl gas. The NOCl decomposes according to the
following equation.
2 NOCl (g)
>
2 NO (g) + Cl2 (g)
If the initial pressure of the pure NOCl is 0.7500 atm and the final pressure in the chamber
after the gases have reached equilibrium is 0.7860 atm calculate Kp for the reaction.
For problems 9-11 consider the following reaction.
CO2 (s)
<>
CO2 (g)
Assume that the volume of the container is constant and that the volume of the solid is
negligible compared to the volume of the gas.
9.
10.
11.
If the system is at equilibrium and some additional CO2 (s) is added,
(a)
the concentration of CO2 (g) will increase.
(b)
the concentration of CO2 (g) will decrease.
(c)
the concentration of CO2 (g) will remain the same.
If the system is at equilibrium and some additional CO2 (g) is added,
(a)
the mass of CO2(s) will increase.
(b)
the mass of CO2(s) will decrease.
(c)
the mass of CO2(s) will remain the same
Given that the reaction is endothermic, an increase in temperature will cause
(a)
the concentration of CO2(g) to increase.
(b)
the concentration of CO2(g) to decrease
(c)
no change in the concentration of CO2(g)
12.
Consider the following reaction.
CO (g) + Cl2 (g) >
COCl2 (g)
The mechanism is believed to be,
(1)
Cl2
<> 2Cl
(fast equilibrium)
(2)
Cl + CO <> COCl
(fast equilibrium)
(3)
COCl + Cl2 > COCl2 + Cl
(slow)
(4)
2 Cl
(fast)
>
Cl2
Assuming that the mechanism is correct, derive the rate law for this reaction.
13.
Antimony pentachloride decomposes according to the following equation.
SbCl5 (g) → SbCl3 (g) + Cl2 (g)
Suppose that 10.0 g of antimony pentachloride is placed in a 5.00 L container and allowed to
establish equilibrium. Calculate the concentration of all species at equilibrium.
K = 2.51 x 10-2
14.
Consider the following reactions.
Na2O (s) → 2 Na (l) + ½ O2 (g)
Na2O2 (s) → 2 Na (l) + O2 (g)
K = 2 x 10-25
K = 5 x 10-29
Calculate the equilibrium constant for the following reaction.
Na2O (s) + ½ O2 (g) →
Na2O2 (s)
15.
The following reaction,
NO (g) + O3 (g)
→
NO2 (g) + O2 (g)
was studied by performing two experiments. In the first experiment the rate of
disappearance of NO was followed in the presence of a large excess of O3 (g).
([O3] = 1.0 x 1014 molecules/cm3) The results are as follows.
Time (ms)
[NO] (molecules/cm3)
0
100.
500.
700.
1000.
6.0 x 108
5.0 x 108
2.4 x 108
1.7 x 108
0.99 x 108
In the second experiment the rate of disappearance of O3 was followed in the presence of a
large excess of NO ([NO] = 2.0 x 1014). The results are as follows.
Time (ms)
[O3] (molecules/cm3)
0
50.
100.
200.
300.
10.0 x 109
8.4 x 109
7.0 x 109
4.9 x 109
3.4 x 109
(a)
What is the order with respect to each reactant?
(b)
What is the effective rate constant for each experiment?
(c)
What is the overall rate law?
(d)
What is the rate constant associated with the overall reaction?
16.
Consider the following reaction.
→
A (g) + B (g)
C (g) + D (g)
In one experiment the concentration of A (g) was monitored in the presence of a large
excess of B (g) ( [B]o = 0.750 M ) .
Time (s)
0.00
1.00
5.00
7.00
10.00
[A] (mol/L)
6.00 x 10-6
5.00 x 10-6
2.40 x 10-6
1.70 x 10-6
0.99 x 10-6
In a second experiment the concentration of B (g) was monitored in the presence of a large
excess of A (g) ( [A]o = 1.50 M ) .
Time (s)
0.00
0.50
1.00
2.00
3.00
17.
[B] (mol/L)
1.00 x 10-4
0.84 x 10-4
0.70 x 10-4
0.49 x 10-4
0.34 x 10-4
(a)
Determine the differential rate law for this reaction.
(b)
Determine the value of the rate constant for this reaction.
Consider the following reaction.
2 CO2 (g)
<>
2 CO (g) + O2 (g)
K = 2.0 x 10-6
If 2.00 mol of CO2 is placed in a 5.0 L container calculate the concentrations of all species
when the system has come to equilibrium.
18.
Consider the following reaction.
N2 (g) + O2 (g)
<>
2 NO (g)
Kp = 0.050
Calculate the partial pressure of all species at equilibrium given the following initial
pressures. (You will need to use the method of successive approximations)
PoN2 = 0.80 atm
PoO2 = 0.20 atm
19.
Consider the following reaction,
H2 (g) + Br2 (g) < > 2 HBr (g) ,
Kp= 2.4 x 108 at 425OC
Suppose that H2 (g) and Br2 (g) are pumped into a container at 425 OC. Given that the
initial partial pressure of H2 (g) is 1.00 atm and that the initial partial pressure of Br2 (g) is
2.00 atm calculate the partial pressure of all species once equilibrium is attained. Assume
that the temperature is maintained at 425 OC.
20.
Consider the following reaction
2 NOCl (g)
<>
2 NO (g) + Cl2 (g) , K = 1.6 x 10-5
Suppose that 0.050 mol of NOCl is pumped into a 50.0 L container. Calculate the
concentration of all species (to two significant figures) once the system has come to
equilibrium. (You will need to use the method of successive approximations)
21.
Consider the decomposition of NH3 (g).
2 NH3 (g)
→
N2 (g) + 3 H2 (g)
The reaction is catalyzed on a surface of either tungsten or osmium. The activation
energies for the catalyzed and uncatalyzed reactions are given below.
W
Os
No catalyst
22.
Ea = 163 kJ/mol
Ea = 197 kJ/mol
Ea = 335 kJ/mol
(a)
Which surface is a better catalyst for the decomposition of NH3?
(b)
Assuming a temperature of 298 K, how many times faster is the reaction on the
W surface compared to when the reaction is uncatalyzed?
(c)
The reaction is first order in NH3 at low concentrations of NH3 but ultimately
becomes 0 order in NH3 as the concentration of NH3 is increased. Provide an
explanation for this observation.
Consider the following reaction.
NO2 (g) + CO (g) → CO2 (g) + NO (g)
The mechanism is as follows.
(1)
2 NO2
(2)
NO3 + CO
<>
>
NO3 + NO
NO2 + CO2
(a)
Derive the general rate law for this reaction using the steady state mechanism.
(b)
Derive a simpler rate law that is valid during the initial phase of the reaction
when it can be assumed that [NO] is negligible. Which is step is the rate
limiting step during the initial phase of the reaction.