CHAPTER 6 CHEMICAL KINETICS
Reflect on Your Learning
(Page 358)
1. (a) Some possibilities include increasing temperature, using a higher concentration of reactants, increasing the
surface area of the reactants if the system is heterogeneous, or adding a catalyst.
(b) Examples are, respectively: cooking food, using concentrated acid, starting a campfire with twigs, and using
industrial catalysts.
2. Answers will vary but the mechanism of the reaction involves collisions of octane and oxygen molecules in which
single carbon–carbon and single carbon–hydrogen bonds are broken in succession down the hydrocarbon chain with
reformation of bonds to form carbon dioxide and water molecules.
3. Some reactant molecules have complex structures and strong chemical bonds which make reaction more difficult.
Other reactions involve simple ions or molecules with weak or unstable chemical bonds in which most collisions
result in reaction. Some reactions produce large amounts of heat that accelerate further reaction.
Try This Activity: Slowing the Browning Process
(Page 359)
(a)
(b)
(c)
(d)
(e)
Apples react with oxygen from the air.
The lemon juice is an antioxidant that slows the reaction.
Temperature is the variable.
Reducing the concentration of a reactant is investigated.
Lemon juice, refrigeration, and removal of oxygen should all reduce rate.
6.1 RATE OF REACTION
PRACTICE
(Page 361)
Understanding Concepts
∆c
1. r = ∆t
(0.200 mol/L)
= 40 s
r = 0.0050 mol/(L•s)
∆c
2. r = ∆t
(0.60 mol/L)
= 5 min
r = 0.12 mol/(L•min)
PRACTICE
(Page 364)
Understanding Concepts
1. Rate of reaction with respect to
+2 consumption is 2.0 101 mol Fe+2 /(L•min)
Fe(aq)
(aq)
+ consumption is 3.2 101 mol H+ /(L•min)
H(aq)
(aq)
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Chemical Kinetics
193
+2 production is 4.0 102 mol Mn+2 /(L•min)
Mn(aq)
(aq)
+3 production is 2.0 101 mol Fe+3 /(L•min)
Fe(aq)
(aq)
H2O(l) production is 1.6 101 mol H2O(l)/(L•min)
[O2(g)]
5
2. (a) = 2.0 102 mol O2(g)/(L•s)
t
4
= 2.5 102 mol O2(g)/(L•s)
[H2O(g)]
6
(b) = 2.0 10–2 mol H2O(g)/(L•s)
t
4
= 3.0 10–2 mol H2O(g)/(L•s)
3. (a) Rate of production of NO2:
[NO2 ]
+ = 4.0 10–3 mol NO2/(L•s)
t
(b) Rate of consumption of NO:
[NO]
– = 4.0 10–3 mol NO/L•s
t
Rate of consumption of O2:
[O2]
– = 2.0 10–3 mol O2/(L•s)
t
4. 2 NaHCO3 + H2SO4 → 2 CO2(g) + 2 H2O(l) + Na2SO4
(a) Rate of consumption of NaHCO3:
∆m
= 3.25 g NaHCO3/20 s
∆t
∆m
= 0.16(3) g NaHCO3/s
∆t
1 mol
(b) nNaHCO = 3.25g 3
84.0 g
nNaHCO = 0.0387 mol
3
1 mol H2SO4
nH SO = 0.0387 mol NaHCO3 2 mol NaHCO3
2
4
nH SO = 0.0193 mol H2SO4
2
4
98.1 g
mH SO = 0.0193 mol 2
4
1 mol
mH SO = 1.89 g H2SO4
2
4
Rate of consumption of H2SO4:
1.89 g H2SO4
∆m
= 20 s
∆t
∆m
= 0.095 g H2SO4/s
∆t
nH SO
0.0193 mol H2SO4
2
4
(c) = ∆t
20 s
nH SO
2
4
= 9.7 104 mol H2SO4/s
∆t
nCO
0.0387 mol CO2
(d) 2 = ∆t
20 s
nCO
2 = 1.9 103 mol CO2/s
∆t
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Chapter 6
Copyright © 2003 Nelson
PRACTICE
(Page 365)
Understanding Concepts
5. Pressure or volume (of gases), colour, conductivity, or pH (of solutions) could all be used to measure a reaction rate.
6. (a) colour (Permanganate is purple.)
(b) pressure or volume of hydrogen gas
7. Concentrations of reactants are highest at the beginning of a reaction, and rate depends on concentrations.
8.
Change of Concentration
Concentration of Product (mol/L)
Over Time
[HCl]
[Cl2]
[CH4 ]
Reaction Progress
9. (a)
0.8
Concentration (mol/L)
0.7
product
0.6
0.5
0.4
0.3
0.2
0.1
0
50
100
150
Time (s)
∆cproduct
0.70 mol/L – 0 mol/L
(b) = ∆t
60 s – 0 s
∆cproduct
= 0.011(7) mol/(L•s)
∆t
∆cproduct
0.70 mol/L – 0.40 mol/L
(c) = ∆t
60 s – 20 s
∆cproduct
= 0.0075 mol/(L•s)
∆t
Copyright © 2003 Nelson
Chemical Kinetics
195
∆cproduct
(d) r = ∆t
(0.40 – 0.13) mol/L
= (20 – 0) s
∆cproduct
= 0.014 mol/(L•s)
∆t
SECTION 6.1 QUESTIONS
(Page 366)
Understanding Concepts
1. (a) CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
Rate of consumption of methane:
–[CH4 ]
8.0 mol
= t
2.00 L 3.2 s
–[CH4 ]
= 1.2(5) mol/(L•s)
t
(b) Rate of consumption of oxygen:
–[O2 ]
= 2 1.2(5) mol/(L•s)
t
1
–[O2 ]
= 2.5 mol/(L•s)
t
(c) Rate of production of carbon dioxide:
+[CO2 ]
= 1 1.2(5) mol/(L•s)
t
1
+[CO2 ]
= 1.2(5) mol/(L•s)
t
(d) Rate of production of water vapour:
+[H2O]
= 2 1.2(5) mol/(L•s)
t
1
+[H2O]
= 2.5 mol/(L•s)
t
2. (a)
Concentration of Ethanal During
Thermal Decomposition
0.40
Concentration (mol/L)
0.35
0.30
0.25
0.20
0.15
[C2H4O(g)]
0.10
0.05
0
500
1000
1500
2000
Time (s)
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Chapter 6
Copyright © 2003 Nelson
(0.360 – 0.180) mol/L
(b) rate(t = 0 to 420 s) = (420 – 0) s
= 4.3 10–4 mol/(L•s)
(0.180 – 0.090) mol/L
(c) rate(t = 420 to 1250 s) = (1250 – 420) s
= 1.1 10–4 mol/(L•s)
(d) Rate falls as the concentration of the reactant(s) falls.
(e) Extrapolation yields a value of about 2000 s.
(f) From slopes of the appropriate tangents, approximate rates are:
rate(c = 0.20 mol/L) = 2.8 10–4 mol/(L•s)
rate(c = 0.10 mol/L) = 8.3 10–5 mol/(L•s)
Making Connections
3. (a) (Answers will vary.) Some possibilities include increasing temperature, using a higher concentration of reactants,
increasing the surface area of the reactants if the system is heterogeneous, or adding a catalyst.
(b) Examples are, respectively: baking a cake, using pure oxygen gas in an oxyacetylene torch, dust explosions, and
using a catalyst disc in contact lens solution.
(c) (Answers will vary, depending on individual student responses.)
6.2
FACTORS AFFECTING REACTION RATE
PRACTICE
(Page 371)
Understanding Concepts
1. The five factors are: the chemical nature of reactants (gold vs. sodium in water); temperature (burning wood); surface
area (steel wool burning in pure oxygen); concentration (magnesium into dilute or concentrated acid); catalysis
(manganese dioxide in hydrogen peroxide).
2. Only one factor — surface area — applies to heterogeneous systems (e.g., grain dust explodes whereas grain smoulders).
3. The rate would double or triple for each increase of 10ºC; for a 20ºC rise, the rate would increase by a factor of
roughly 4 to 10 times.
4. Increased surface area increases reaction rate. The lump of coal has a very small surface area, compared to that of a
similar mass of coal dust.
5. Possible leaks of these gases would increase concentrations of reactants, thus increasing the possible rate of combustion of fuels.
Making Connections
6. Enzymes are key catalysts in metabolic reactions. “Poisoning” of these catalysts can seriously upset body chemistry.
7. BHT is an antioxidant or inhibitor (negative catalyst) that decreases the rate of food decay reactions.
SECTION 6.2 QUESTIONS
(Page 371)
Understanding Concepts
1. (a) chemical nature of reactant
(b) temperature
(c) catalysis
(d) surface area
(e) concentration
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