13
13.1
13.2
13.3
13.4
1
Rates of Chemical
Reactions
Rates of Chemical Reactions
Expressions of Reaction Rates in Terms of
Rates of Changes in Concentrations of
Reactants or Products
Methods of Measuring Reaction Rates
Factors Affecting Reaction Rates
Chemical Kinetics
A study of
(1)
reaction rates
(2)
the factors affecting reaction rates
(3)
reaction mechanisms
(the detailed steps involved in reactions)
2
Explosive reactions
2H2(g) + O2(g)  2H2O(l)
3
Vigorous reactions
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
Potassium reacts with
water vigorously
4
Very rapid reactions
Formation of insoluble salts
+
−
Ag (aq) + Cl (aq) AgCl(s)
5
Very rapid reactions
Formation of insoluble bases
Fe3+(aq) + 3OH−(aq) Fe(OH)3(s)
6
Very rapid reactions
Acid-alkali neutralization reactions
H+(aq)
7
−
+ OH (aq) H2O(l)
Q.1
+
−
Ag (aq) + Cl (aq) AgCl(s)
Fe3+(aq) + 3OH−(aq) Fe(OH)3(s)
H+(aq)
−
+ OH (aq) H2O(l)
All involve oppositely charged ions
8
Rapid or moderate reactions
Displacement reactions of metals : Zn(s) + 2Ag+(aq)  Zn2+(aq) + 2Ag(s)
9
Rapid or moderate reactions
Displacement reactions of metals : Zn(s) + 2Ag+(aq)  Zn2+(aq) + 2Ag(s)
Displacement reactions of halogens : Cl2(aq) + 2Br(aq)  2Cl(aq) + Br2(aq)
10
Slow reactions
Fermentation of glucose
C6H12O6(aq)  2C2H5OH(aq) + 2CO2(g)
11
Slow reactions
2MnO4(aq) + 5C2O42(aq) + 16H+(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
12
Very slow reactions
Rusting of iron
4Fe(s) + 3O2(g) + 2nH2O(l)  2Fe2O3 · nH2O(s)
13
Extremely slow reactions
+
2+
CaCO3(s) + 2H (aq)  Ca (aq) + CO2(g) + H2O(l)
Before corrosion
14
After corrosion
Two Ways to Express Reaction Rates
1. Average rate
2. Instantaneous rate
(rate at a given instant)
15
Averagerate of reaction
Total changein amount of a product or a reactant

Total time taken for the change to occur
Amount is usually expressed in
Concentration
Mass
Volume
Pressure
16
mol dm−3
g
cm3 or dm3
atm
Q.2 0.36 g of magnesium reacted with 50.0 cm3
of 1.0 M hydrochloric acid to give 360 cm3 of
hydrogen under room conditions.
The reaction was completely in 90 seconds.
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
0.36 g
3
1
(a) Average rate 
 4.0  10 g s
90 s
17
Q.2 0.36 g of magnesium reacted with 50.0 cm3
of 1.0 M hydrochloric acid to give 360 cm3 of
hydrogen under room conditions.
The reaction was completely in 90 seconds.
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
360 cm3
(b) Averagerate 
 4.0 cm3 s1
90 s
18
2.(c) Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
0.36 g
No. of moles of Mg 
 0.015 mol
1
24.3 g mol
No. of moles of HCl  1.0 mol dm3  0.0500dm3  0.0500mol
Mg is the limiting reactant
No. of moles of HCl reacted  2  0.015 mol  0.030 mol
Decrease in concentration of HCl(aq) in 90 s

0.030mol
3

0.60
mol
dm
0.0500dm3
0.60 mol dm-3
Averagerate 
 6.7  10-3 mol dm-3 s1
90 s
19
2.(d) Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Rate of reaction
Rate of reaction
= 2
w.r.t. HCl(aq)
w.r.t. MgCl2(aq)
Increase in concentration of MgCl2(aq) in 90 s
1
  0.60 mol dm 3  0.30 mol dm -3
2
0.30 mol dm-3
Averagerate 
 3.3 10-3 mol dm-3 s1
90 s
20
2. Instantaneous rate
The rate at a particular instant of the
reaction is called the instantaneous rate.
For the chemical reaction
aA + bB  cC + dD
Instantaneous rate
 d[A] 1
 d[B] 1
d[C] 1
d[D] 1

( )
( )
( )
( )
dt
a
dt b
dt c
dt d
[X] = molarity of X
21
2. Instantaneous rate
The rate at a particular instant of the
reaction is called the instantaneous rate.
For the chemical reaction
aA + bB  cC + dD
Instantaneous rate
 d[A] 1
 d[B] 1
d[C] 1
d[D] 1

( )
( )
( )
( )
dt
a
dt b
dt c
dt d
Units : mol dm3 s1, mol dm3 min1, mol dm3 h1…etc.
22
Graphical Representation of Reaction
Rates – Rate curves
A rate curve is a graph plotting the amount of
a reactant or product against time.
23
Consider the reaction
A
(reactant)
24

B
+ C
(product)
At any time t, the instantaneous rate of the
reaction equals the slope of the tangent to the
curve at that point.
The greater the slope, the higher the rate of the
reaction.
25
-ve slope of curve of reactant A
 [A]  with time
26
+ve slope of curve of product B
 [B]  with time
27
The rate at t0 is usually the fastest and is called
the initial rate.
The curve is the steepest with the greatest
slope at time t0.
28
The rate of the reaction gradually  as the
reaction proceeds.
Flat curve
 reaction completed
29
Concentration of product Z
(mol dm−3)
Q.3
30
X + Y  2Z
C
B
A
Time of reaction (min)
Concentration of product Z
(mol dm−3)
1 5.4 mol dm 3
Averagerate  
 0.39 mol dm 3 min1
2
7 min
31
X + Y  2Z
C
B
A
Time of reaction (min)
Concentration of product Z
(mol dm−3)
X + Y  2Z
32
C
B
Instantane ous rate at A
1 (6.0- 0.0) mol dm3
 
2
(1.6- 0.0) min
 1.9 mol dm3 min1
A
1.6
Time of reaction (min)
Concentration of product Z
(mol dm−3)
X + Y  2Z
33
C
5.1
B
2.7
Instantane ous rate at B
1 (5.1- 2.7) mol dm 3
 
2
(3.0- 1.0) min
 0.6 mol dm3 min1
A
Time of reaction (min)
Concentration of product Z
(mol dm−3)
X + Y  2Z
34
C
B
Instantane ous rate at C  0
A
Time of reaction (min)
Methods of Measuring Reaction Rates
A. Physical measurements
1. Continuous measurements
2 Initial rate measurements
(Clock reactions)
B. Chemical measurements (Titration)
35
1. Continuous measurements
Experiment is done in ONE take.
The reaction rates are determined by
measuring continuously a convenient property
which is directly proportional to the
concentration of any one reactant or product
of the reaction mixture.
Properties to be measured : –
Gas volume / Gas pressure / Mass /
Color intensity / Electrical conductivity
36
1.1 Measurement of large volume changes
Examples:
(1) CaCO3(s) + 2HCl(aq)
 CaCl2(aq) + H2O(l) + CO2(g)
(2) Zn(s) + H2SO4(aq)
 ZnSO4(aq) + H2(g)
(3) 2H2O2(aq)  2H2O(l) + O2(g)
37
1.1 Measurement of large volume changes
Temperature is
kept constant
38
A typical laboratory set-up for measuring the
volume of gas formed in a reaction
Volume of gas formed (cm3)
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)
39
dV
slope 
 rate
dt
Time of reaction (min)
Q.4
(2) Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)
H2(g) is sparingly soluble in water while
CO2 is quite soluble in water.
Volume
of CO2
Rate 
Rate 
40
Sigmoid curve
1.2 Measurement of small volume changes
- Dilatometry
Capillary tube
Liquid phase reaction
mixture
CH3COOH(l) + CH3CH2OH(l)  CH3COOCH2CH3(l) + H2O(l)
41
1.3 Measurement of mass changes
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
42
The cotton wool plug is to allow the escape of CO2(g) but
to prevent loss of acid spray due to spurting.
stopwatch
cotton wool plug
limestone pieces
of known mass
measured volume of
standard
hydrochloric acid
electronic
balance
43
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)
 CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
Which reaction is more suitable to be followed
by mass measurement ?
Hydrogen is a very light gas.
The change in mass of the reaction mixture
may be very small.
The electronic balance used in the school
laboratory may not be sensitive enough to
detect the small change.
44
Loss of mass (m)
mfinal = total mass loss
dm
slope 
 rate
dt
time
mfinal - mt
45
mfinal = mfinal – m0 (∵ m0 = 0)
d[H  ]
slope 
=  rate 2
dt
time
1.4 Colorimetry
∵ colour intensity  [coloured species]
d(colourintensity)
rate  
dt
46
H2O2(aq) + 2H+(aq) + 2I(aq)  I2(aq) + 2H2O(l)
colour intensity  as reaction proceeds
CH3COCH3(aq) + I2(aq)
 CH3COCH2I(aq) + H+(aq) + I(aq)
Br2(aq) + HCOOH(aq)
 2H+(aq) + 2Br(aq) + CO2(g)
2MnO4(aq) + 16H+(aq) + 5C2O42(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
colour intensity  as reaction proceeds
47
48
cuvettes
A colorimeter
49
Yellow
light
Yellow
filter
Blue solution
Complementary colours
50
Red  Cyan
Pairs of opposite colours are complementary
colours
51
Red  Cyan
Green  Magenta
Pairs of opposite colours are complementary
colours
52
Red  Cyan
Green  Magenta
Blue  Yellow
CMYK
Pairs of opposite colours are complementary
colours
53
When mixed in the proper proportion,
complementary colours produce a neutral color
(grey, white, or black).
54
I0
I
I0 = intensity before absorption
I = intensity after absorption
55
I0
I
% transmitta nce     100%
 I0 
56
I
I 
Absorbance  log10  0 
I
If I = I0 ,
If I = 0 ,
%T = 100%
%T = 0%
A = log101 = 0
A  log10  
zero absorption
complete absorption
A = bC
Beer’s law
57
A
Deviation at higher
concentrations
A calibration curve is first
constructed for AC conversion
C
58
Q.5
[I2]
d[I 2 ]
slope 
 rate
dt
time
A
dA
slope 
 rate
dt
59
time
1.5 Measurement of electrical
conductivity
Na+OH(aq) + CH3COOH(aq)  CH3COONa+(aq) + H2O(l)
∵ conducting mobility : OH > CH3COO
∴ conductivity  as the rx proceeds
60
1.5 Measurement of electrical
conductivity
2MnO4(aq) + 16H+(aq) + 5C2O42(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
∵ total number of ions 
∴ electrical conductivity  as the rx proceeds
61
1.6 Measurement of pressure changes
d(PT )
rate  
dt
PT = total pressure of the reaction
mixture
62
Q.6
(i)
(ii)
2NO(g) + 2H2(g)  N2(g) + 2H2O(g)
3H2(g) + N2(g)  2NH3(g)
At fixed V and T, PT  n
In both reactions,
n  as the reactions proceed
 PT  as the reactions proceed
63
suction
flask
dilute hydrochloric acid
pressure sensor
magnesium ribbon
to data-logger
interface and computer
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
64
A(g) + B(g)  products
65
A chemical clock is a complex mixture of
reacting chemical compounds in which the
concentration of one or more components
exhibits periodic changes.
In cases where one of the reagents has a
visible color, crossing a concentration
threshold can lead to an abrupt color change in
a reproducible time lapse.
66
2. Initial Rate Measurements-Clock Reactions
1. A set of experiments is done in which all
reaction conditions but one are kept constant.
2–
S2O3 (aq)
67
+
+ 2H (aq)  SO2(aq) + H2O(l) + S(s)
Experiment
[S2O32(aq)] / M
[H+(aq)] / M
1
2
0.10
0.08
1
1
3
0.04
1
4
0.02
1
2. Initial Rate Measurements-Clock Reactions
2–
S2O3 (aq)
+
+ 2H (aq)  SO2(aq) + H2O(l) + S(s)
yellow
precipitate
2. The time taken for the reaction to arrive at a
particular point at the early stage of the
reaction is measured.
68
The beaker containing the reaction mixture is
placed over a cross marked on a white tile.
69
As more sulphur forms, the reaction mixture
becomes more cloudy.
70
The cross becomes more and more difficult to
see and finally disappears.
71
2–
+
S2O3 (aq) + 2H (aq)  SO2(aq) + H2O(l) + S(s)
Average rate in the early stage
=
yellow
precipitate
Amount of S required to blot out the mark
Time taken to blot out the mark
Since the amount of S required to blot out
the mark is a constant,
1
Average 
rate
time taken to ‘blot out’ the mark
72
1
Average 
rate
time taken to ‘blot out’ the mark
The average rate of reaction is inversely
proportional to the time taken to ‘blot out’ the
mark.
The faster is the reaction, the shorter is the
time taken for the mark to disappear.
73
dS
dt
S
slope  average rate 
t
slope  initialrate 
amount of S
If S and t are small(early stage)
dS ΔS

dt Δt
time
74
dS ΔS

dt Δt
Since S is a constant
dS ΔS 1


dt Δt t
75
Initial rate  k[S2O32(aq)]x[H+(aq)]y
Since HCl is in large excess,
[H+(aq)]y  constant at the early stage
Initial rate  k[S2O32(aq)]x[H+(aq)]y  k’[S2O32(aq)]x
ΔS 1
 Initialrate 

Δt t
1
2
''
 k [S2O3 (aq)] x
t
76
77
Time taken (t)
to mask the
mark / s
Expt.
[S2O32(aq)]
[H+(aq)]
1
0.10
1
10
2
0.08
1
13
3
0.04
1
25
4
0.02
1
50
(M)
(M)
1
t
/ s1
Q.7
1
t
1 ''
2
x
 k [S2O3 (aq)]
t
Linear  x = 1
[S2O32(aq)]
78
Other Examples of Clock Reactions : 5I(aq) + IO3(aq) + 6H+(aq)  3I2(aq) + 3H2O(l)
Small and fixed amounts of S2O32(aq) and starch are
added to the reaction mixtures in all runs.
I2(aq) + 2S2O32(aq)  2I(aq) + S4O62(aq)
(fixed) (fixed)
I2(aq)
+
(excess)
starch  deep blue complex
(fixed)
Time taken for the reaction mixture to turn deep blue
is measured.
79
Other Examples of Clock Reactions : 5I(aq) + IO3(aq) + 6H+(aq)  3I2(aq) + 3H2O(l)
I2(aq) + 2S2O32(aq)  2I(aq) + S4O62(aq)
(fixed) (fixed)
I2(aq)
+
(excess)
starch  deep blue complex
(fixed)
By changing the concentration of any one of the
reactants, deep blue colour will appear in different
time lapses  a chemical clock !
Halloween clock
80
Other Examples of Clock Reactions : 5Br(aq) + BrO3(aq) + 6H+(aq)  3Br2(aq) + 3H2O(l)
OH
OH
Br
+
(fixed)
Br2
+
(excess)
81
Br
3Br2
(fixed)
Br
methyl red  colourless
(fixed)
Advantages of physical measurements
1. Suitable for fast reactions.
2. Small sample size
3. More accurate than chemical method
(titration)
4. No interruption  continuous measurements
5. Can be automated.
82
Disadvantages of physical measurements
1. More sophisticated
2. More expensive
3. More specific – only suit a limited number of
reactions.
83
B. Chemical Measurements (Titration Methods)
1. Start a reaction with all reaction conditions
but one fixed.
2. Withdraw and quench fixed amounts of the
reaction mixture at different times.
84
Quenching methods:
Temperature 
• Cooling the reaction mixture rapidly in ice.
• Diluting the reaction mixture with a
sufficient amount of cold water or an
appropriate solvent. Concentration 
• Removing one of the reactants or the
catalyst (if any) by adding another
reagent.
85
B. Chemical Measurements (Titration Methods)
1. Start a reaction with all reaction conditions
but one fixed.
2. Withdraw and quench fixed amounts of the
reaction mixture at different times.
3. Titrate the quenched samples to determine
the concentration of one of the reactants or
products.
86
CH3COCH3 + I2
H+ as catalyst
CH3COCH2I + HI
Q.8
The reaction is quenched by adding to it NaHCO3(aq)
that removes the catalyst.
HCO3(aq) + H+(aq)  H2O(l) + CO2(g)
87
CH3COCH3 + I2
H+ as catalyst
CH3COCH2I + HI
Q.9
Titrated with standard solution of Na2S2O3(aq) using
starch as indicator (added when the end point is near)
2
2

2S2O3 (aq) + I2(aq)  S4O6 (aq) + 2I (aq)
Colour change at the end point : deep blue to colourless
88
CH3COCH3 + I2
H+ as catalyst
CH3COCH2I + HI
Q.10
The excess S2O32(aq) would react with H+ to give a
cloudy mixture with a pungent smell.
S2O32(aq) + 2H+(aq)  S(s) + SO2(g) + H2O(l)
89
Advantages of titrimetric method
1. Only simple apparatus are required.
2. Can be applied to a great variety of slow
reactions.
90
Disadvantages of physical measurements
1. Not suitable for fast reactions.
It takes time to withdraw samples and
perform titration.
2. Reactions are disturbed – NOT continuous
3. Time consuming – NOT automated
91
Factors Affecting
Reaction Rates
92
Collision Theory
No reaction
Sufficient K.E.
Incorrect orientation
93
Collision Theory
No reaction
Correct orientation
Insufficient K.E.
94
Collision Theory
Sufficient K.E.
Correct orientation
Effective collision
95
Collision Theory
Activation energy
Bond breaking and bond forming occur at the
same time
Ea < B.E.(s) of the bond(s) to be broken
96
Collision Theory
Activation energy
Higher Ea
 more K.E. required for effective collision
 slower reaction
97
Collision Theory
Activation energy
Lower Ea
 less K.E. required for effective collision
 faster reaction
98
Collision Theory
Activation energy
Rate of reaction depends on Ea which in turn
depends on the nature of reactants.
E.g. K is more reactive than Mg
99
Factors Affecting Reaction Rates
concentration
100
particle size
pressure
catalyst
temperature
light
Effect of concentration
•
101
e.g. Reaction between Mg and HCl
Effect of concentration
(a) 2.0 M HCl
(b) 1.0 M HCl
(c) 0.5 M HCl
Reaction rate:
(a) > (b) > (c)
102
Effect of concentration
Time for reaction to
complete: t1 < t2 < t3
Higher [HCl(aq)]
 Faster reaction
103
[X] 
 Reactant particles are more crowded
 Collision frequency 
 Number of effective collisions 
 Reaction rate 
104
For the reaction
aA + bB  cC + dD
Rate  k[A]x[B]y
where x and y are the orders of reaction
with respect to A and B
k is the rate constant
units  mol dm3 s1/(mol dm3)x+y
105
For the reaction
aA + bB  cC + dD
Rate  k[A]x[B]y
x and y can be  integers or fractional
x  y is the overall order of reaction.
x, y can ONLY be determined experimentally.
106
Effect of pressure
Only applicable to reactions involving gaseous
reactants.
107
Pressure 
 Reactant particles are more crowded
 Collision frequency 
 No. of effective collisions 
 Rate of reaction 
108
Effect of temperature
Applicable to ALL reactions
109
T
 K.E. of particles 
 Collision frequency  (minor effect) and
No. of particles with K.E. > Ea  (major effect)
 No. of effective collisions 
 Rate of reaction 
110
Rate
Rate of reaction 
exponentially with temperature
Rate  e
 Ea
RT
In general, a 10oC  in T
doubles the rate.
T / C
111
Effect of particle size
For a fixed volume of solid,
Smaller particle size  greater surface area
112
CaCO3(aq) + 2H+(excess)  CaCl2(aq) + H2O(l) + CO2(g)
Rate involving
powdered solid
reactant is higher
Reason: higher
chance of contact
between reactant
particles
113
Q.11
0.5 g powder
0.5 g granule
114
Effect of Catalyst
A catalyst is a substance that alters the rate
of a chemical reaction by providing an
alternative reaction pathway with a different
activation energy.
A positive catalyst speeds up a reaction by
providing an alternative reaction pathway
with a lower Ea.
A negative catalyst slows down a reaction
by providing an alternative reaction pathway
with a higher Ea.
115
Effect of Catalyst
Catalysts remain chemically unchanged at the
end of reactions.
116
H2O2(aq)
MnO2 as catalyst
2H2O(l) + O2(g)
Physical measurement
117
2H2O(l) + O2(g)
Volume of gas formed (cm3)
H2O2(aq)
MnO2 as catalyst
118
Time of reaction (min)
Titrimetric method (Q.12)
H2O2(aq)
MnO2 as catalyst
2H2O(l) + O2(g)
Pipette samples at different times
Remove MnO2(s) by filtration
Titrate with MnO4(aq)/H+(aq)
5H2O2(aq) + 2MnO4(aq) + 6H+(aq)
 2Mn2+(aq) + 8H2O(l) + 5O2(g)
119
Q.13
[H2O2]
Without MnO2
With MnO2
time
120
Effect of light
Light with specific frequency (E  h) can
provide sufficient energy to break a particular
chemical bond in a reactant leading to a
photochemical reaction.
Br – Br
h
Br + Br
C6H14 + Br  C6H13 + HBr
   C6H13Br…
121
Autocatalysis
Catalysis in which the product acts as the
catalyst of the reaction
2MnO4(aq) + 16H+(aq) + 5C2O42(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
CH3COCH3(aq) + I2(aq)
 CH3COCH2I(aq) + H+(aq) + I(aq)
122
Q.14
[MnO4]
Rate 
Sigmoid curve
Rate 
123
time
The END
124
13.1 Rates of Chemical Reactions (SB p.5)
Back
In a chemical reaction, a total of 0.18 g of carbon
dioxide gas is given out in 1 minute at room
temperature. What is its average rate in mol s–1 for that
time interval?
Answer
Number of moles of CO2 =
0.18 g
(12.0  16.0  2) g mol - 1
= 0.0041 mol
0.0041mol
Average rate =
60 s
= 6.83 × 10–5 mol s–1
125
13.1 Rates of Chemical Reactions (SB p.5)
In the uncatalyzed decomposition of hydrogen
peroxide solution into water and oxygen at room
conditions, the volume of oxygen given out in 20
hours is 5 cm3. What is its average rate in mol s–1 for
that time interval?
2H2O2(l)  2H2O(l) + O2(g)
(Molar volume of gas at room temperature and
pressure= 24.0 dm3 mol–1)
Answer
126
13.1 Rates of Chemical Reactions (SB p.5)
Back
Number of moles of O2
5 cm3
=
24 000 cm3 mol 1
= 2.08 × 10–4 mol
2.08  10 -4 mol
Average rate =
(20  60  60) s
= 2.89 × 10–9 mol s–1
127
13.1 Rates of Chemical Reactions (SB p.6)
The change in concentration
of reactant X in a chemical
reaction is illustrated in the
graph on the right.
128
13.1 Rates of Chemical Reactions (SB p.6)
With the use of the graph, calculate
(a)
the initial rate of the reaction;
(b)
the average rate for the time interval from
the 1st to the 2nd minute;
(c)
the instantaneous rate at the 3rd minute.
(Give your answers in mol dm–3 min–1.)
Answer
129
13.1 Rates of Chemical Reactions (SB p.6)
(a) Initial rate
=
Slope of the tangent
to the curve at t0
3
(0.100

0.160)
mol
dm
=
(1.2  0) min
= -0.05 mol dm-3 min-1
130
13.1 Rates of Chemical Reactions (SB p.6)
(b) Average rate
(0.080  0.110) mol dm 3
=
(2  1) min
= -0.03 mol dm-3 min-1
131
13.1 Rates of Chemical Reactions (SB p.6)
Back
(c) Instantaneous rate at the
3rd minute
=
Slope of the tangent to
the curve at the 3rd
minute
3
= (0.046  0.077) mol dm
(3.5  2) min
= -0.021 mol dm-3 min-1
132
13.1 Rates of Chemical Reactions (SB p.8)
(a) In the hydrolysis of an ester at a constant temperature
of 398 K, the concentration of the ester decreases from
1 mol dm–3 to 0.75 mol dm–3 in 4 minutes. What is its
average rate in mol dm–3 s–1 for that time interval?
Answer
(a) Average rate at 398 K
= –(1 – 0.75) mol dm-3  (4  60) s
= –0.001 04 mol dm-3 s-1
133
13.1 Rates of Chemical Reactions (SB p.8)
(b) The graph on the right shows the
change in concentration of
a reactant in a chemical reaction.
134
13.1 Rates of Chemical Reactions (SB p.8)
With the use of the graph above, calculate
(i)
the initial rate of the reaction;
(ii)
the average rate for the time interval from the 20th to
the 30th second;
(iii) the instantaneous rate at the 10th second.
Answer
135
13.1 Rates of Chemical Reactions (SB p.8)
Back
(i)
Initial rate = (0.02 - 0.01) mol dm
(0  10) s
-3
= -1  10-3 mol dm-3 s-1
-3
(
0.009
0.006)
mol
dm
(ii) Average rate =
(20  30) s
= -3  10-4 mol dm-3 s-1
(0.018 - 0.013) mol dm -3
(iii) Instantaneous rate =
(0  10) s
= -5  10-4 mol dm-3 s-1
136
13.2 Expressions of Reactions Rates in Terms of Rates of Changes in
Concentrations of Reactants or Products (SB p.10)
Back
Haemoglobin (Hb) binds with carbon monoxide
according to the following equation:
4Hb + 3CO  Hb4(CO)3
Express the rate of the reaction in terms of the rate of
change in concentration of any one of the reactants or
the product.
The rate of the reaction is expressed as:
Rate 
137
Answer
d [Hb 4 (CO)3 ]
1 d [Hb]
1 d [CO]
 
 
dt
4
dt
3
dt
13.2 Expressions of Reactions Rates in Terms of Rates of Changes in
Concentrations of Reactants or Products (SB p.10)
Back
Express the rate of the following reaction in terms of the
rate of change in concentration of any one of the reactants
or the product.
2H2(g) + O2(g)  2H2O(l)
d [O 2 (g)]
1 d [H2 O(l)]
1 d [H2 (g)]




Rate =
2
dt
2 dt
dt
138
Answer
13.3 Methods of Measuring Reaction Rates (SB p.11)
Alkaline hydrolysis of ethyl ethanoate (an ester) using
sodium hydroxide solution is represented by the
following equation:
CH3CO2CH2CH3(l) + NaOH(aq)
 CH3CO2Na(aq) + CH3CH2OH(aq)
The rate of the reaction can be followed by titrating
small volumes of the reaction mixture with standard
dilute hydrochloric acid at successive five-minute
intervals.
139
13.3 Methods of Measuring Reaction Rates (SB p.11)
(a) Suggest a method to quench the reaction mixture
so that the concentration of sodium hydroxide
solution can be determined accurately. Explain
briefly why this method can be used.
Answer
(a) The reaction mixture can be quenched by pipetting a
sample of the reaction mixture into a conical flask
containing ice water. The cooling and dilution of the
reaction mixture decrease the reaction rate sufficiently for
chemical analysis.
140
13.3 Methods of Measuring Reaction Rates (SB p.11)
(b) Explain why the change in concentration of
sodium hydroxide solution but not that of ethyl
ethanoate is measured in order to determine the
rate of the above reaction.
Answer
(b) Sodium hydroxide is a strong alkali that reacts with strong
mineral acids almost instantaneously. Therefore, the
titration of sodium hydroxide solution and dilute
hydrochloric acid provides accurate experimental results.
141
13.3 Methods of Measuring Reaction Rates (SB p.11)
Answer
(c) Explain which option, A or B, is a reasonable set of
experimental results for the above titration.
Option A
Time after mixing (min) Volume of HCl added
at the end point (cm3)
5
10
10
8
Option B
142
Time after mixing (min)
Volume of HCl added
at the end point (cm3)
5
8
10
10
13.3 Methods of Measuring Reaction Rates (SB p.11)
(c) Sodium hydroxide is a reactant of the hydrolysis. As the
reaction proceeds, the concentration of sodium hydroxide
in the reaction mixture decreases with time, and hence
the amount of dilute hydrochloric acid used in the titration.
Thus, option A is a reasonable set of experimental results.
143
13.3 Methods of Measuring Reaction Rates (SB p.11)
(d) Name a suitable indicator for the titration.
Answer
(d) Methyl orange / Phenophthalein
Back
144
13.3 Methods of Measuring Reaction Rates (SB p.13)
A student recorded the following experimental results
for the reaction of zinc and dilute hydrochloric acid.
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Time
(min)
0.0
Volume
of H2(g)
produced
(cm3)
0
145
1.0 2.0 3.0
4.0
5.0
6.0 7.0 8.0
9.0
15
38
40
41
42
26
33
42
42
13.3 Methods of Measuring Reaction Rates (SB p.13)
(a) Plot a graph of volume of hydrogen gas produced
against time.
Answer
(a)
146
13.3 Methods of Measuring Reaction Rates (SB p.13)
(b) Describe the change in the rate of the reaction
using your graph in (a).
Answer
(b) As shown in the graph in (a), the volume of hydrogen
gas given out at the beginning of the reaction (e.g. in
the time interval between the 1st and the 2nd minute) is
greater than that near the end of the reaction (e.g. in
the time interval between the 6th and the 7th minute).
Therefore, the rate of the reaction decreases with time.
147
13.3 Methods of Measuring Reaction Rates (SB p.13)
(c) Explain how you can measure the initial rate of the
reaction graphically.
Answer
(c) The initial rate can be found by determining the slope of
the tangent to the curve at time zero.
148
13.3 Methods of Measuring Reaction Rates (SB p.13)
Back
(d) Determine graphically the rate of the reaction at
the 5th minute. State the unit.
Answer
(d) From the graph in (a),
rate of reaction
= slope of the tangent to the curve at the 5 minute
(46  34) cm3
=
(8  2) min
= 2 cm3 min-1
149
13.3 Methods of Measuring Reaction Rates (SB p.15)
Back
Suggest an experimental method for determining the rate
of each of the following reactions:
(a) S2O82–(aq) + 2I–(aq)  2SO42–(aq) + I2( aq)
(b) CH3COOCH3(aq) + I2(aq)
 CH3COOCH2I(aq) + HI(aq)
(c) 2MnO4–(aq) + 5C2O42–(aq) + 16H+(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l) + H+(aq)
(a) Colorimetric measurement / titration
(b) Colorimetric measurement
150(c) Colorimetric mesurement / titration
Answer
13.4 Factors Affecting Reaction Rates (SB p.17)
Explain why sawdust burns explosively in pure
oxygen but slowly in air.
A higher concentration of oxygen increases
the rate of combustion.
Back
151
Answer
13.4 Factors Affecting Reaction Rates (SB p.21)
(a) List THREE factors that affect the rate of a chemical
reaction.
Answer
(a) Concentration of reactants / pressure /
temperature / surface area / catalyst /
light (any 3)
152
13.4 Factors Affecting Reaction Rates (SB p.21)
(b) The figure below shows the laboratory set-up for
measuring the change in mass of the reaction mixture
with time in the course of the reaction:
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
153
13.4 Factors Affecting Reaction Rates (SB p.21)
A certain mass of calcium carbonate was added to 50 cm3 of
2.0 M hydrochloric acid at 20°C. Carbon dioxide was
allowed to escape and the mass of the reaction mixture was
measured at regular time intervals. The results were
expressed as the loss of mass with respect to time. The
experiment was carried out with one change of condition at
a time:
(i) using 1.0 M hydrochloric acid in place of 2.0 M
hydrochloric acid.
(ii) carrying out the reaction at 30°C.
(iii) using powdered calcium carbonate of the same mass.
154