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Homework #1
Chapter 15 Chemical Kinetics
8. Arrhenius Equation
π = π΄ππΈπ
π πβ Therefore, k depends only on temperature. The rate of the reaction depends on all of these items (a-d).
14. a) πππ
πΏβπ b) πππ
πΏβπ c) 1
π
d) πΏ
πππβπ e) πΏ2
πππ2βπ
15. Rate has units of πππ
πΏβπ therefore, the units of the rate constant must compensate for any
missing/additional units.
πΏ1
2β
πππ1
2β βπ
18. a) General Rate Law π ππ‘π = π[ππ]π₯[πΆπ2]π¦
Experiment [NO]o (M) [Cl2]o (M) Initial Rate ( πππ
πΏβπππ)
1 0.10 0.10 0.18
2 0.10 0.20 0.36
3 0.20 0.20 1.45
From experiments 1 and 2 ([NO]o is kept constant) as the initial concentration of Cl2 is
doubled, the initial rate of the reaction doubles (0.36
0.18=2.0); therefore, the reaction is first
order with respect to Cl2. From experiments 2 and 3 ([Cl2]o is kept constant) as the initial concentration of NO is
doubled, the initial rate of the reaction quadruples(1.45
0.36=4.0); therefore, the reaction is
second order with respect to NO. π ππ‘π = π[ππ]2[πΆπ2]
b)
π ππ‘π = π[ππ]2[πΆπ2]
0.18 πππ
πΏβπππ= π(0.10 π)2(0.10 π)
π = 180 πΏ2
πππ2βπππ
19. a) General Rate Law
π ππ‘π = π[πΌβ]π₯[ππΆπβ]π¦ Experiment [I-]o (M) [OCl-]o (M) Initial Rate (πππ
πΏβπ )
1 0.12 0.18 7.91Γ10-2
2 0.060 0.18 3.95Γ10-2
3 0.030 0.090 9.88Γ10-2
4 0.24 0.090 7.91Γ10-2
From experiments 1 and 2 ([OCl-]o is kept constant) as the initial concentration of I- is
doubled, the initial rate of the reaction doubles (7.91Γ10β2
3.95Γ10β2=2.0); therefore, the reaction is
first order with respect to I-. π ππ‘π = π[πΌβ][ππΆπβ]π¦
2
You must find the order of OCl- the long way. Divide the 2 equations (in which the [OCl-] changes) by each other.
7.91 Γ 10β2 ππππΏβπ
9.88 Γ 10β3 ππππΏβπ
=π(0.12)(0.18 π)π¦
π(0.030)(0.090 π)π¦
2.0 = 2.0π¦ π¦ = 1
π ππ‘π = π[πΌβ][ππΆπβ] b)
π ππ‘π = π[πΌβ][ππΆπβ]
7.91 Γ 10β2 πππ
πΏβπ = π(0.12 π)(0.18 π)
π = 3.7πΏ
πππβπ
c)
π ππ‘π = π[πΌβ][ππΆπβ] = (3.7 πΏ
πππβπ )(0.15 π)(0.15 π) = 0.082 πππ
πΏβπ
20. General Rate Law
π ππ‘π = π[π2π5]π₯ Experiment [N2O5]o (M) Initial Rate (πππ
πΏβπ )
1 0.0750 8.90Γ10-4
2 0.190 2.26Γ10-3
3 0.275 3.26Γ10-3
4 0.410 4.85Γ10-3
You must solve this problem the long way.
8.90 Γ 10β4 πππ
πΏβπ = π(0.0750 π)π₯
2.26 Γ 10β3 πππ
πΏβπ = π(0.190 π)π₯
Divide the 2 equations by each other.
8.90 Γ 10β4 ππππΏβπ
2.26 Γ 10β3 ππππΏβπ
=π(0.0750 π)π₯
π(0.190 π)π₯
0.394 = 0.395π₯ π₯ = 1
π ππ‘π = π[π2π4]
8.90 Γ 10β4 πππ
πΏβπ = π(0.0750 π)
π = 0.0119 1
π
21. a) General Rate Law
π ππ‘π = π[πππΆπ]π₯
Experiment [NOCl]o (ππππππ’πππ
ππ3 ) Initial Rate (ππππππ’πππ
ππ3βπ )
1 3.0Γ1016 5.98Γ104
2 2.0Γ1016 2.66Γ104
3 1.0Γ1016 6.64Γ103
4 4.0Γ1016 1.06Γ105
From experiments 2 and 3, as the initial concentration of NOCl is doubled, the initial rate
of the reaction quadruples (2.66Γ104
6.64Γ103=4.0); therefore, the reaction is second order with
respect to NOCl.
π ππ‘π = π[πππΆπ]2
b) 5.98 Γ 104 ππππππ’πππ
ππ3βπ = π(3.0 Γ 1016 ππππππ’πππ
ππ3 )2
3
π = 6.6 Γ 10β29 ππ3
ππππππ’πππ βπ
c) π = 6.6 Γ 10β29 ππ3
ππππππ’πππ βπ (6.022Γ1023 ππππππ’πππ
1 πππ)( 1 πΏ
103 ππ3) = 4.0 Γ 10β8 πΏ
πππβπ
22. a) General Rate Law
π ππ‘π = π[π»π]π₯[πΆπ]π¦
Experiment [Hb]o (ππππ
πΏ) [CO]o (ππππ
πΏ) Initial Rate (ππππ
πΏβπ )
1 2.21 1.00 0.619
2 4.42 1.00 1.24
3 4.42 3.00 3.71
From experiments 1 and 2 ([CO]o is kept constant) as the initial concentration of Hb is
doubled, the initial rate of the reaction doubles( 1.24
0.619=2.00); therefore, the reaction is
first order with respect to Hb. You cannot use the shortcut to calculate the order with respect to [CO]
Find 2 equations where the concentration of Hb is constant
1.24 ππππ
πΏβπ = π (4.42
ππππ
πΏ) (1.00
ππππ
πΏ)
π¦
3.71 ππππ
πΏβπ = π (4.42
ππππ
πΏ) (3.00
ππππ
πΏ)
π¦
Divide the 2 equations by each other.
1.24 πππππΏβπ
= π(4.42 πππππΏ
)(1.00 πππππΏ
)π¦
3.71 πππππΏβπ = π(4.42 ππππ
πΏ )(3.00 πππππΏ )
π¦
0.33 =(1.00 ππππ
πΏ)
π¦
(3.00 πππππΏ )
π¦ = (0.33)π¦
π¦ = 1
The reaction is 1st order with respect to CO b) π ππ‘π = π[π»π][πΆπ]
c) π ππ‘π = π[π»π][πΆπ]
0.619 ππππ
πΏβπ = π (2.21
ππππ
πΏ) (1.00
ππππ
πΏ)
π = 0.280 πΏ
ππππβπ
d) π ππ‘π = 0.280 πΏ
ππππβπ [π»π][πΆπ] = (0.280 πΏ
ππππβπ ) (3.36 ππππ
πΏ)(2.40 ππππ
πΏ) = 2.26 ππππ
πΏβπ
23. a) General
π ππ‘π = π[πΆππ2]π₯[ππ»β]π¦
Experiment [ClO2]o (M) [OH-]o (M) Initial Rate (πππ
πΏβπ )
1 0.0500 0.100 5.75Γ10-2
2 0.100 0.100 2.30Γ10-1
3 0.100 0.0500 1.15Γ10-1
From experiments 1 and 2 ([OH-]o is kept constant) as the initial concentration of ClO2 is
doubled, the initial rate of the reaction quadruples(2.30Γ10β1
5.75Γ10β2=4.00); therefore, the
reaction is second order with respect to ClO2. From experiments 2 and 3 ([ClO2]o is kept constant) as the initial concentration of OH- is
doubled, the initial rate of the reaction doubles(2.30Γ10β1
1.15Γ10β1=2.00); therefore, the reaction is
first order with respect to OH-. π ππ‘π = π[πΆππ2]2[ππ»β]
4
Find k
5.75 Γ 10β2 πππ
πΏβπ = π(0.0500 π)2(0.100 π)
π = 230.πΏ2
πππ2βπ
π ππ‘π = 230.πΏ2
πππ2βπ [πΆππ2]2[ππ»β]
b) π ππ‘π = 230. πΏ2
πππ2βπ (0.175 π)2(0.0844 π) = 0.594 πππ
πΏβπ
30. a) Because the plot of 1
[π΄] verse time was linear the rate law is a second order rate law.
π ππ‘π = π[π΄]2
βπ[π΄]
ππ‘= π[π΄]2
1
[π΄]2π[π΄] = βπππ‘
β«1
[π΄]2π[π΄]
[π΄]
[π΄]Β°
= β« βπππ‘π‘
0
β1
[π΄]+
1
[π΄]Β°= βππ‘ + 0
1
[π΄]= ππ‘ +
1
[π΄]Β° (integrated rate law)
The rate constant is the slope of the line π = 3.60 Γ 10β2 πΏ
πππβπ
π ππ‘π = 3.60 Γ 10β2 πΏ
πππβπ [π΄]2 (rate law)
b) The half-life is the amount of time it takes for Β½ of the original concentration to react.
1
[π΄]= ππ‘ +
1
[π΄]Β°
1
0.5(2.80 Γ 10β3 π)= (3.60 Γ 10β2
πΏ
πππβπ ) π‘ +
1
2.80 Γ 10β3 π
π‘ = 9920 π
c) 1
[π΄]= ππ‘ +
1
[π΄]Β°
1
7.00 Γ 10β4 π= (3.60 Γ 10β2
πΏ
πππβπ ) π‘ +
1
2.80 Γ 10β3 π
π‘ = 29,800 π
31. a) Because the plot of ln[A] verses t is linear, this is a 1st order rate law. π ππ‘π = π[π΄]
βπ[π΄]
ππ‘= π[π΄]
1
[π΄]π[π΄] = βπππ‘
β«1
[π΄]π[π΄]
[π΄]
[π΄]Β°
= β β« πππ‘π‘
0
ππ[π΄] β ππ[π΄]Β° = βππ‘ + 0 ππ[π΄] = βππ‘ + ππ[π΄]Β° (integrated rate law)
The rate constant is the negative slope of the line π = 2.97 Γ 10β2 1
πππ
π ππ‘π = 2.97 Γ 10β2 1
πππ[π΄] (rate law)
b) ππ[π΄] = βππ‘ + ππ[π΄]Β°
5
ππ (1
2(2.00 Γ 10β2 π)) = β (2.97 Γ 10β2
1
πππ) π‘ + ππ(2.00 Γ 10β2 π)
π‘ = 23.3 πππ c)
ππ[π΄] = βππ‘ + ππ[π΄]Β°
ππ(2.50 Γ 10β3 π) = β (2.97 Γ 10β2 1
πππ) π‘ + ππ(2.00 Γ 10β2 π)
π‘ = 70.0 πππ
32. a) Because the plot of [C2H5OH] verses t is linear, this is a 0th order rate law
π ππ‘π = π
βπ[πΆ2π»5ππ»]
ππ‘= π
π[πΆ2π»5ππ»] = βπππ‘
β« π[πΆ2π»5ππ»]πΆ2π»5ππ»
[πΆ2π»5ππ»]Β°
= β β« πππ‘π‘
0
[πΆ2π»5ππ»] β [πΆ2π»5ππ»]Β° = βππ‘ + 0 [πΆ2π»5ππ»] = βππ‘ + [πΆ2π»5ππ»]Β° (integrated rate law)
The rate constant equals the negative slope of the line; π = 4.00 Γ 10β5 πππ
πΏβπ
π ππ‘π = 4.00 Γ 10β5 πππ
πΏβπ (rate law)
b) [πΆ2π»5ππ»] = βππ‘ + [πΆ2π»5ππ»]Β°
1
2(1.25 Γ 10β2 π) = β(4.00 Γ 10β5 πππ
πΏβπ )π‘ + (1.25 Γ 10β2 π)
π‘ = 156 π
c) Since this is a 0th order rate law (which means that the rate is independent of concentration), the amount of time that it takes to decompose the total amount of C2H5OH is just double the half-life.
π‘ = 2(156 π ) = 312 π Or
0 = β(4.00 Γ 10β5 πππ
πΏβπ )π‘ + (1.25 Γ 10β2 π)
π‘ = 312 π 33. Plot [H2O2] vs. t, ln[H2O2] vs. t, and [H2O2]-1 vs. t
Since the plot of ln[H2O2] vs. t results in a linear relationship, the reaction is 1st order.
0
0.5
1
1.5
0 1000 2000 3000 4000
[H2
O2
]
time (s)
-3.5
-2.5
-1.5
-0.5
0.5
0 2000 4000
ln[H
2O
2]
time (s)
0
5
10
15
20
25
0 2000 4000
[H2
O2
]^-1
time (s)
6
If you do not have a graphing calculator look at the spacing between the evenly spaces times to determine the order.
Time (s) [H2O2] Difference
between points ln[H2O2]
Difference between points
[H2O2]-1 Difference
between points
0 1.00 0.59-1.00=-0.41 0.00 -0.53-0.00=-0.53 1.00 1.7-1.00=0.7
600 0.59 -0.22 -0.53 -0.46 1.7 1.0
1200 0.37 -0.15 -0.99 -0.5 2.7 1.8
1800 0.22 -0.09 -1.5 -0.5 4.5 3.3
2400 0.13 -0.05 -2.0 -0.5 7.7 4.3
3000 0.082 -2.5 12
Since the all the differences between ln[H2O2] are all approximately -0.5 this will be a straight line and the reaction is 1st order. Differential Rate Law
π ππ‘π = π[π»2π2] Integrated Rate Law
ππ[π»2π2] = βππ‘ + ππ[π»2π2]Β° The slope of the line will equal βk.
π ππππ =πππ π
ππ’π=
ππ(0.050 π)βππ(1.00 π)
3,600 π β0 π = β8.3 Γ 10β4 1
π = βπ
π = 8.3 Γ 10β4 1
π
Calculate the concentration of H2O2 after 4000. s.
ππ[π»2π2] = β (8.3 Γ 10β4 1
π ) (4,000 π ) + ππ(1.00 π)
[π»2π2] = 0.036 π 35. Plot
[A] vs. t, ln[A] vs. t, and [A]-1 vs. t
Since the plot of [A]-1 vs. t is linear, this is a 2nd order rate law
If you do not have a graphing calculator look at the spacing between the evenly spaces times to determine the order. You must have at least 3 points.
Time (s) [NO2] Difference
between points ln[NO2]
Difference between points
[NO2]-1 Difference
between points
0 0.500 0.25-0.50=-0.25 -0.693 -1.39-0.693=-0.69 2.00 4.00-2.00=2.00
9.00Γ103 0.250 -0.076 -1.39 -0.36 4.00 1.75
1.80Γ104 0.174 -1.75 5.75
This problem it is a little hard to determine. (I will not give you something this hard on an exam.) The difference between the difference between the [NO2] points is 0.17 over a concentration
range of 0.326 making the difference between differences 52% (0.17
0.326100% = 52%) of the
0
0.1
0.2
0.3
0.4
0.5
0.6
0 5000 10000 15000
[NO
2]
time (s)
0
2
4
6
8
0 5000 10000 15000
[NO
2]^
-1
time (s)
-2
-1.5
-1
-0.5
0 5000 10000 15000
ln[N
O2
]
time (s)
7
range. The difference between the difference between the ln[NO2] is 0.33 over a ln[NO2] range
of -1.06 making the difference between differences 31% (0.33
1.06100% = 33%) of the range. The
difference between the difference between [NO2]-1 is 0.25 over a range of 3.75 making the
difference between differences 6% (0.25
3.75100% = 6%)of the range. The graph [NO2]-1 has the
smallest difference between differences, when you take the size of the range into account, the order is 2nd order. Differential Rate Law
π ππ‘π = π[ππ2]2
Integrated Rate Law 1
[ππ2]= ππ‘ +
1
[ππ2]Β°
The slope of the line is the rate constant
π ππππ =πππ π
ππ’π=
1
0.174 πβ
1
0.500 π
1.80Γ104 π β0 π = 2.08 Γ 10β4 πΏ
πππβπ = π
1
[ππ2]= (2.08 Γ 10β4
πΏ
πππβπ ) π‘ +
1
[ππ2]Β°
1
[ππ2]= (2.08 Γ 10β4
πΏ
πππβπ ) (2.70 Γ 104 π ) +
1
0.500 π= 7.62
πΏ
πππ
[ππ2] = 0.131 π 36. a) Plot [O] vs. t, ln[O] vs. t, and [O]-1 vs. t
Since ln[O] vs. t is linear, the rate expression is 1st order with respect to O.
If you do not have a graphing calculator look at the spacing between the evenly spaces times to determine the order.
Time (s) [O] Difference
between points ln[O]
Difference between points
[O]-1 Difference
between points
0 5.0Γ109 1.9Γ109-5.0Γ109
=-3.1Γ109 22 21-22=-1.0 2.0Γ10-10 5.2Γ10-10-2.0Γ10-10
= 3.2Γ10-10
0.010 1.9Γ109 -1.2Γ109 21 -1.0 5.2Γ10-10 9.8Γ10-10
0.020 6.8Γ108 -4.3Γ108 20. -1.0 1.5Γ10-9 2.5Γ10-9
0.030 2.5Γ108 19 4.0Γ10-9
Since the all the differences between ln[H2O2] are all approximately -1.0 this will be a straight line and the reaction is 1st order. Differential Rate Law
b) Overall Rate Law π ππ‘π = π[π][ππ2]
Because the concentration of [NO2] >>[O] this can be turned into a pseudo rate law π ππ‘π = πβ²[π]
0.E+00
2.E+09
4.E+09
6.E+09
0 0.01 0.02 0.03 0.04
[O]
time
19
20
21
22
23
0 0.02 0.04
ln[O
]
time
0.E+00
1.E-09
2.E-09
3.E-09
4.E-09
5.E-09
0 0.02 0.04
[O]^
-1
time
8
πβ² = π[ππ2] Determine k' using the 1st order integrated rate law
ππ[π] = βπβ²π‘ + ππ[π]Β° Therefore, the slope of the line is βk'
π ππππ =πππ π
ππ’π=
ππ(2.5Γ108 ππ‘πππ
ππ3 )βππ(5.0Γ109 ππ‘πππ
ππ3 )
3.0Γ10β2 π β0 π = β1.0 Γ 102 1
π = βπβ²
πβ² = 1.0 Γ 102 1
π
Find k πβ² = π[ππ2]
(1.0 Γ 1021
π ) = π[ππ2]
π = 1.0 Γ 10β11 ππ3
ππππππ’πππ βπ
38. a) The order of the reaction with respect to A is 2nd order because the plot of [A]-1 vs. t is
linear. The integrated rate law is.
1
[π΄]= ππ‘ +
1
[π΄]Β°
Therefore, the [A]0-1 is the y intercept. [A]0
-1 = 10 πΏ
πππ and [A]0 = 0.1 M
b) 1
[π΄]= ππ‘ +
1
[π΄]Β°
π = π ππππ =πππ π
ππ’π=
70 πΏπππ
β20 πΏπππ
6 π β0 π = 10 πΏ
πππβπ
1
[π΄]= (10
πΏ
πππβπ ) (9 π ) + 10
πΏ
πππ= 100
πΏ
πππ
[π΄] = 0.01 π c) Calculate the time for the 1st half-life ([A] 0.1 0.05)
1
0.05 π= (10
πΏ
πππβπ ) π‘ +
1
0.1 π
π‘ = 1 π Calculate the time for the 2nd half-life
1
0.025 π= (10
πΏ
πππβπ ) π‘ +
1
0.05 π
π‘ = 2 π Calculate the time for the 3rd half-life
1
0.0125 π= (10
πΏ
πππβπ ) π‘ +
1
0.025 π
π‘ = 4 π 39. a) Plot [NO] vs. t, ln[NO] vs. t, and, [NO]-1 vs. t
Since the plot of ln[NO] vs. t is linear, the rate law is 1st order with respect to NO.
0.E+00
2.E+08
4.E+08
6.E+08
8.E+08
0 500 1000 1500
[NO
]
time
18
18.5
19
19.5
20
20.5
0 500 1000 1500
ln[N
O]
time
0.E+00
5.E-09
1.E-08
2.E-08
0 500 1000 1500
[NO
]^-1
time
9
If you do not have a graphing calculator look at the spacing between the evenly spaces times to determine the order.
Time (s) [NO] Difference
between points ln[NO]
Difference between points
[NO]-1 Difference
between points
0 6.0Γ108 2.4Γ108-6.0Γ108
=-3.6Γ108 20 19-20=-1.0 1.7Γ10-9 4.2Γ10-9-1.7Γ10-9
= 2.5Γ10-9
500 2.4Γ108 -1.4Γ108 19 -1.0 4.2Γ10-9 5.8Γ10-9
1000 9.9Γ107 18 1.0Γ10-8
Since the all the differences between ln[NO] are all approximately -1.0 this will be a straight line and the reaction is 1st order.
Plot [O3] vs. t, ln[O3] vs. t, and [O3]-1 vs. t
Since plot of ln[O3] vs. t is linear, the rate law is 1st order with respect to O3
If you do not have a graphing calculator look at the spacing between the evenly spaces times to determine the order.
Time (s) [O3] Difference
between points ln[O3]
Difference between points
[O3]-1 Difference
between points
0 1.0Γ1010 7.0Γ109-1.0Γ1010
=-3.0Γ109 23.0 22.7-23.0=-0.3 1.0Γ10-10 1.4Γ10-10-1.0Γ10-10
= 4.0Γ10-11
100 7.0Γ109 -2.1Γ109 22.7 -0.4 1.4Γ10-10 6.0Γ10-11
200 4.9Γ109 -1.5Γ109 22.3 -0.4 2.0Γ10-10 9.0Γ10-11
300 3.4Γ109 21.9 2.9Γ10-10
Since the all the differences between ln[O3] are all approximately -0.4 this will be a straight line and the reaction is 1st order. b) Rate Law
π ππ‘π = π[ππ][π3] c) When the [NO] data was taken the concentration of [O3]>>[NO]. Therefore, a pseudo
rate law can be written. π ππ‘π = πβ²[ππ] πβ² = π[π3]
Since the reaction was first order with respect to [NO] the slope of the line equals -k'
πβ² = βπ ππππ =πππ π
ππ’π=
ππ(9.9Γ107 ππππππ’πππ
ππ3 )βππ(6.0Γ108 ππππππ’πππ
ππ3 )
1000. ππ β0 ππ = 0.0018 1
ππ = 1.8 1
π
When the [O3] data was taken the concentration of [NO]>>[O3]. Therefore, a pseudo rate law can be written.
π ππ‘π = πβ²β²[π3] πβ²β² = π[ππ]
Since the reaction was first order with respect to [O3] the slope of the line equals -k''
0.E+00
5.E+09
1.E+10
2.E+10
0 200 400
[O3
]
Time
21.5
22
22.5
23
23.5
0 200 400
ln[O
3]
Time
0.E+00
1.E-10
2.E-10
3.E-10
4.E-10
0 100 200 300 400
[O3
]^-1
Time
10
πβ²β² = βπ ππππ =πππ π
ππ’π=
ππ(3.4 Γ 109 ππππππ’πππ ππ3 ) β ππ(1.0 Γ 1010 ππππππ’πππ
ππ3 )
300. ππ β 0 ππ = 0.0036
1
ππ = 3.6
1
π d) You can solve this using either of the two data sets. In both data sets the
reactions are pseudo 1st order system. Overall rate = pseudo 1st order rate Using data set 1 [O3] held constant
π[ππ][π3] = πβ²[ππ]
π =πβ²
[π3]=
1.8 1π
1.0 Γ 1014 ππππππ’πππ ππ3
= 1.8 Γ 10β14 ππ3
ππππππ’πππ βπ
= 1.8 Γ 10β17 ππ3
ππππππ’πππ βππ
Using data set 2 [NO] held constant π[ππ][π3] = πβ²β²[π3]
π =πβ²β²
[ππ]=
3.6 1π
2.0 Γ 1014 ππππππ’πππ ππ3
= 1.8 Γ 10β14 ππ3
ππππππ’πππ βπ
= 1.8 Γ 10β17 ππ3
ππππππ’πππ βππ
43. ππ[ππ»3] = βππ‘ + ππ[ππ»3]π ππ(0.250 π) = βπ(120. π ) + ππ(1.00 π)
π = 0.0116 1
π
ππ[ππ»3] = β (0.0116 1
π ) π‘ + ππ[ππ»3]π
ππ(0.350 π) = β (0.0116 1
π ) π‘ + ππ(2.00 π)
π‘ = 151 π 47. a) Calculate k
ππ[π΄] = βππ‘ + ππ[π΄]π ππ(0.250[π΄]π) = βπ(320. π ) + ππ[π΄]π
ππ (0.250[π΄]π
[π΄]π) = βπ(320. π )
π = 0.00433 1
π
Calculate time for 1st half life
ππ ( 1
2[π΄]π) = β (0.00433
1
π ) π‘ + ππ[π΄]π
ππ ( 12[π΄]π
[π΄]π) = β (0.00433
1
π ) π‘
π‘ = 159 π Calculate time for the 2nd half life
ππ ( 1
4[π΄]π) = β (0.00433
1
π ) π‘ + ππ (
1
2[π΄]π)
ππ ( 14[π΄]π
12
[π΄]π
) = β (0.00433 1
π ) π‘
π‘ = 159 π b)
ππ( 0.100[π΄]π) = β(0.00433 1
π )π‘ + ππ[π΄]π
ππ ( 0.100[π΄]π
[π΄]π) = β (0.00433
1
π ) π‘
11
π‘ = 532 π 48. Since the time of the half-life doubles every successive half-life, the reaction is a 2nd order
reaction
a) 1
[π΄]= ππ‘ +
1
[π΄]π
1
0.05 π= π(10.0 πππ) +
1
0.10 π
π = 1.0 πΏ
πππβπππ
1
[π΄]= (1.0
πΏ
πππβπππ) (80.0 πππ) +
1
0.10 π= 90.
πΏ
πππ
[π΄] = 0.011 π
b)
1
[π΄]= (1.0 πΏ
πππβπππ)(30.0 πππ) +
1
0.10 π= 40. πΏ
πππ
[π΄] = 0.025 π 52. The problem wants you to find the time when [π΄] = 4.00[π΅]. Known
ππ΄ = 4.50 Γ 10β4 1
π
ππ΅ = 3.70 Γ 10β3 1
π
Both 1st order reactions Initially
[π΄]π = [π΅]π ππ[π΄] = βππ‘ + ππ[π΄]π ππ[π΅] = βππ‘ + ππ[π΅]π
Plug in[π΅]π for[π΄]π and[π΄] = 4.00[π΅] for[π΄]and solve for [π΅] ππ(4.00[π΅]) = βππ΄π‘ + ππ[π΅]π
4.00[π΅] = πβππ΄π‘+ππ[π΅]π
[π΅] =πβππ΄π‘+ππ[π΅]π
4.00
Plug into equation for [π΅] and solve for t ππ[π΅] = βππ΅π‘ + ππ[π΅]π
ππ (πβππ΄π‘+ππ[π΅]π
4.00) = βππ΅π‘ + ππ[π΅]π
ππ(πβππ΄π‘+ππ[π΅]π) β ππ(4.00) = βππ΅π‘ + ππ[π΅]π
βππ΄π‘ + ππ[π΅]π β ππ(4.00) = βππ΅π‘ + ππ[π΅]π βππ΄π‘ β ππ(4.00) = βππ΅π‘ βππ(4.00) = (βππ΅ + ππ΄)π‘
π‘ =βππ(4.00)
βππ΅ + ππ΄=
βππ(4.00)
β3.70 Γ 10β3 1π
+ 4.50 Γ 10β4 1π
= 427 π
54. The concentrations of B and C are so much larger than A, therefore, during the course of the
reaction they do not effectively change making this a pseudo 1st order reaction. The rate of the pseudo 1st order reaction is equal to the rate of the overall reaction times [π΅]2. a) πππ‘π = π[π΄][π΅]2 = πβ²[π΄] πβ² = π[π΅]2
ππ[π΄] = βπβ²π‘ + ππ[π΄]π ππ(3.8 Γ 10β3 π) = βπβ²(8.0 π ) + ππ(1.0 Γ 10β2 π)
12
πβ² = 0.12 1
π
πβ² = π[π΅]2
π =πβ²
[π΅]2=
0.12 1π
(3.0 π)2= 0.013
πΏ2
πππ2βπ
b)
ππ[π΄] = β(0.12 1
π )π‘ + ππ[π΄]π
ππ(0.0050 π) = β (0.12 1
π ) π‘ + ππ(1.0 Γ 10β2 π)
π‘ = 5.8 π c)
ππ[π΄] = β(0.12 1
π )π‘ + ππ[π΄]π
ππ[π΄] = β (0.12 1
π ) (13.0 π ) + ππ(1.0 Γ 10β2 π)
[π΄] = 0.0021 π d) The concentration of C is the [πΆ]π β 2β[π΄]π β[π΄] = 1.0 Γ 10β2 π β 0.0021 π = 0.008 π
[πΆ] = [πΆ]π β 2β[π΄] = 2.0 π β 2(0.008 π) = 2.0 π This should not be a surprising answer because in order to have a pseudo 1st order reaction with respect to A, the concentration of C must be essentially constant.
56. a) Elementary Step: An individual reaction in a proposed reaction mechanism. The rate law
can be written from the coefficients in the balanced equation b) Molecularity: The number of reactant molecules (or free atoms) taking part in an
elementary reaction. This is also the number of species that must collide in order to produce the reaction represented by an elementary reaction.
c) Reaction Mechanism: The pathway that is proposed for an overall reaction and accounts for the experimental rate law.
d) Intermediate: A species that is produced and consumed during a reaction but does not appear in the overall chemical equation.
e) Rate Determining Step: The elementary reaction that governs the rate of the overall reaction. This is the slowest elementary reaction that occurs in a mechanism.
59. For elementary reactions the rate is just equal to the rate constant times the concentration of
the reactants. a) π ππ‘π = π[πΆπ»3ππΆ] b) π ππ‘π = π[π3][ππ] c) π ππ‘π = π[π3] d) π ππ‘π = π[π3][π] e) π ππ‘π = π[ πΆ6
14 ] 60. The rate law found in problem 33 was
π ππ‘π = π[π»2π2] In order for this to be the rate, the first step would have to be the rate determining step. Overall Reaction 2H2O2 2H2O + O2
13
61. π ππ‘π = π[πΆ4π»9π΅π] Overall Reaction C4H9Br + 2H2O Br- + C4H9OH + H3O+ Intermediates C4H9
+ and C4H9OH2+
63. The rate of a reaction is dependent on the slowest step. a) π ππ‘π = π[ππ][π2] The mechanism is not consistent with rate law.
b) π ππ‘π = π2[ππ3][ππ] NO3 is an intermediate, therefore, it needs to be eliminated from the rate equation. Use equilibrium to eliminate
π1[ππ][π2] = πβ1[ππ3]
[ππ3] =π1[ππ][π2]
πβ1
π ππ‘π = π2[ππ3][ππ] = π2[ππ]π1[ππ][π2]
πβ1=
π2π1[ππ]2[π2]
πβ1= π[ππ]2[π2]
The mechanism is consistent with rate law. c) π ππ‘π = π[ππ]2 The mechanism is not consistent with rate law.
d) In order for a mechanism to be plausible the elementary reactions must add up to the overall reaction. These elementary reactions do not add up to the overall reaction. In fact the middle reaction is not even balanced.
65. π ππ‘π = π3[π΅πβ][π»2π΅ππ3
+] H2BrO3
+ is an intermediate therefore, needs to be eliminated from the rate law. Use equilibrium to eliminate the intermediate.
π2[π»π΅ππ3][π»+] = πβ2[π»2π΅ππ3+]
[π»2π΅ππ3+] =
π2[π»π΅ππ3][π»+]
πβ2
π ππ‘π = π3[π΅πβ][π»2π΅ππ3+] =
π2π3[π΅πβ][π»π΅ππ3][π»+]
πβ2
HBrO3 is an intermediate therefore, needs to be eliminated from the rate law. Use equilibrium to eliminate the intermediate.
π1[π΅ππ3β][π»+] = πβ1[π»π΅ππ3]
[π»π΅ππ3] =π1[π΅ππ3
β][π»+]
πβ1
π ππ‘π =π2π3[π΅πβ][π»π΅ππ3][π»+]
πβ2=
π1π2π3[π΅πβ][π΅ππ3β][π»+]2
πβ1πβ2= π[π΅πβ][π΅ππ3
β][π»+]2
68. a) π ππ‘π = π3[πΆππΆπ][πΆπ2]
COCl is an intermediate, therefore, it needs to be eliminated from the rate law. Use equilibrium to eliminate the intermediate.
π2[πΆπ][πΆπ] = πβ2[πΆππΆπ]
[πΆππΆπ] =π2[πΆπ][πΆπ]
πβ2
π ππ‘π = π3[πΆππΆπ][πΆπ2] =π3π2[πΆπ][πΆπ][πΆπ2]
πβ2
14
Cl is an intermediate, therefore, it needs to be eliminated from the rate law. Use equilibrium to eliminate the intermediate.
π1[πΆπ2] = πβ1[πΆπ]2
[πΆπ] =π1
12β
[πΆπ2]1
2β
πβ1
12β
π ππ‘π =π3π2[πΆπ][πΆπ][πΆπ2]
πβ2=
π3π2[πΆπ][πΆπ2]π1
12β
[πΆπ2]1
2β
πβ2πβ1
12β
=π3π2π1
12β
[πΆπ][πΆπ2]3
2β
πβ2πβ1
12β
π ππ‘π = π[πΆπ][πΆπ2]3
2β b) COCl and Cl are intermediates
69. a) MoCl5- is an intermediate
b) π[ππ2
β]
ππ‘= π2[ππ3
β][πππΆπ5β]
Use the steady state approximation to eliminate the intermediate, rate of formation = rate of consumption
π1[πππΆπ62β] = πβ1[πππΆπ5
β][πΆπβ] + π2[ππ3β][πππΆπ5
β]
[πππΆπ5β] =
π1[πππΆπ62β]
πβ1[πΆπβ] + π2[ππ3β]
π[ππ2β]
ππ‘= π2[ππ3
β][πππΆπ5β] =
π1π2[ππ3β][πππΆπ6
2β]
πβ1[πΆπβ] + π2[ππ3β]
70. Need to calculate the rate of decomposition of O3
βπ[π3]
ππ‘= π1[π][π3] + π2[π][π3] β πβ1[π2][π][π]
Note: Rate of decomposition=ways O3 used up β ways O3 forms O is an intermediate therefore need to remove it from the rate law.
Use the same steady state approximation to solve for [O] π1[π][π3] = πβ1[π2][π][π] + π2[π][π3]
[π] =π1[π][π3]
πβ1[π2][π] + π2[π3]
βπ[π3]
ππ‘= π1[π][π3] + π2[π][π3] β πβ1[π2][π][π]
βπ[π3]
ππ‘= π1[π][π3] +
π1π2[π][π3]2
πβ1[π2][π] + π2[π3]β
π1πβ1[π2][π]2[π3]
πβ1[π2][π] + π2[π3]
β
π[π3]
ππ‘=
π1[π][π3](πβ1[π2][π] + π2[π3]) + π1π2[π][π3]2 β π1πβ1[π2][π]2[π3]
πβ1[π2][π] + π2[π3]
βπ[π3]
ππ‘=
π1πβ1[π2][π]2[π3] + π1π2[π][π3]2 + π1π2[π][π3]2 β π1πβ1[π2][π]2[π3]
πβ1[π2][π] + π2[π3]
β
π[π3]
ππ‘=
2π1π2[π][π3]2
πβ1[π2][π] + π2[π3]
72. a) As the activation energy decreases the rate of the reaction goes up.
b) As the temperature increases the molecules/atoms move faster and more collisions occur, therefore, the faster the rate of the reaction.
c) As the frequency of collisions increases the rate of the reaction goes up.
15
d) As the orientation of collisions is more specific the rate of the reaction goes down.
79.
ππ (π1
π2) =
πΈπ
π (
1
π2β
1
π1)
ππ (π1
3.52 Γ 10β7 πΏπππβπ
) =186 ππ½
πππ
0.0083145 ππ½πππβπΎ
(1
555 πΎβ
1
645 πΎ)
π1 = 9.75 Γ 10β5 πΏ
πππβπ
80.
ππ (π1
π2) =
πΈπ
π (
1
π2β
1
π1)
ππ (7.2 Γ 10β4 1
π
1.7 Γ 10β2 1π
) =πΈπ
8.3145 π½πππβπΎ
(1
720. πΎβ
1
660. πΎ)
πΈπ = 210,000 π½
πππ
ππ (π1
π2) =
πΈπ
π (
1
π2β
1
π1)
ππ (π1
1.7 Γ 10β2 πΏπππβπ
) =210,000 π½
πππ
8.3145 π½πππβπΎ
(1
720. πΎβ
1
598 πΎ)
π1 = 1.3 Γ 10β5 1
π
After each half-life the number of moles of iodoethane is halved. Therefore after 3 half-lives
there is 81 the number of moles of iodoethane that there was initially. For a gas, PV=nRT,
therefore, since the volume, gas constant, and temperature are constant π1
π1=
π2
π2
894 π‘πππ
π₯=
π218
π₯
π2 = 112 π‘πππ
81.
ππ (π1
π2) =
πΈπ
π (
1
π2β
1
π1)
ππ (7π₯
π₯) =
5.40 Γ 104 π½πππ
8.3145 π½πππβπΎ
(1
295 πΎβ
1
π1)
π1 = 324 πΎ = 51β
83. a)
ππ(π) = βπΈπ
π (
1
π) + ππ(π΄)
Therefore, when the ln(k) is plotted vs. 1
πthe slope of the line equalsβ
πΈπ
π
βπΈπ
π = β
πΈπ
8.3145 π½πππβπΎ
= β1.10 Γ 104 πΎ
πΈπ = 9.15 Γ 104 π½
πππ= 91.5
ππ½
πππ
b) When the ln(k) is plotted vs. T1 the y intercept of the line equals Aln
ππ(π΄) = 33.5
π΄ = 3.54 Γ 1014 1
π
c)
ππ(π) = βπΈπ
π (
1
π) + ππ(π΄)
16
ππ(π) = β91,50 π½
πππ
8.3145 π½πππβπΎ
(1
298 πΎ) + ππ (3.54 Γ 1014
1
π ) = β3.43
π = 0.0324 1
π
84. ππ(π) = βπΈπ
π (
1
π) + ππ(π΄)
Therefore if you plot the ln(k) vs. 1
π the slope of the line will be β
πΈπ
π
π ππππ =πππ π
ππ’π=
ππ(3.5 Γ 10β5 1π ) β ππ(4.9 Γ 10β3 1
π )
1298 πΎ
β1
338 πΎ
= β12,400 πΎ
β12,400 πΎ =πΈπ
π =
πΈπ
0.0083145 ππ½πππβπΎ
πΈπ = 103 ππ½
πππ
86. a) b)
c)
-12
-10
-8
-6
-4
-2
0
0.0028 0.003 0.0032 0.0034
ln(k
)
T^-1
17
87.
The second graph is a two-step problem. 1st: reactants intermediates 2nd: intermediates products. Since the second reaction has the greater activation energy it will be the slower reaction and will determine the rate of the reaction. Therefore, we label the activation energy of the 2nd step.
88.
The activation energy for the reverse reaction is the -ΞE + Ea.
β(β216 ππ½
πππ) + 125
ππ½
πππ= 341
ππ½
πππ
90. A catalyst increases the reaction rate because it provides another pathway with a lower
activation energy for the reaction to proceed by. Homogeneous catalysts are present in the same phase as the reactants and heterogeneous catalyst are present in a different phase from the reactants. The uncatalyzed and catalyzed reactions most likely will have different rate laws because they have different pathways that they occur by.
94. a) The catalyst is the species that is initially added to the reaction but not used in the
overall reaction. Therefore NO is the catalyst. b) An intermediate is a species that is formed in one of the steps of the reaction but not
one of the reactants or products. Therefore NO2 is an intermediate.
c)
ππ(π) = βπΈπ
π (
1
π) + ππ(π΄)
ππ(ππππ‘) β ππ(ππ’ππππ‘) = βπΈπ(πππ‘)
π (
1
π) + ππ(π΄) + β
πΈπ(π’ππππ‘)
π (
1
π) β ππ(π΄)
ππ (ππππ‘
ππ’ππππ‘) =
1
π π(βπΈπ(πππ‘) + βπΈπ(π’ππππ‘)) = 0.848
ππ (ππππ‘
ππ’ππππ‘) =
1
(8.3145 π½πππβπΎ)(298 πΎ)
(β11,900 π½ + 14,000 π½) = 0.848
ππππ‘
ππ’ππππ‘= 2.3