1. Rates of Chemical Reactions13
Rates of Chemical Reactions
13.1 Rates of Chemical Reactions
13.2 Expressions of Reaction Rates in Terms of Rates of Changes in Concentrations of Reactants or Products
13.3 Methods of Measuring Reaction Rates
13.4 Factors Affecting Reaction Rates

2. Chemical Kinetics A study of (1) reaction rates
(2) the factors affecting reaction rates
(3) reaction mechanisms
(the detailed steps involved in reactions)

3. Explosive reactions
2H2(g) + O2(g)  2H2O(l)

4. Potassium reacts with water vigorously
Vigorous reactions
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
Potassium reacts with water vigorously

5. Very rapid reactions Ag+(aq) + Cl−(aq) AgCl(s)
Formation of insoluble salts
Ag+(aq) + Cl−(aq) AgCl(s)

6. Very rapid reactions Formation of insoluble bases
Fe3+(aq) + 3OH−(aq) Fe(OH)3(s)

7. Very rapid reactions Acid-alkali neutralization reactions
H+(aq) + OH−(aq) H2O(l)

8. All involve oppositely charged ions
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

9. Rapid or moderate reactions
Displacement reactions of metals : -
Zn(s) + 2Ag+(aq)  Zn2+(aq) + 2Ag(s)

10. 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)

11. Slow reactions Fermentation of glucose
C6H12O6(aq)  2C2H5OH(aq) + 2CO2(g)

12. Slow reactions 2MnO4(aq) + 5C2O42(aq) + 16H+(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l)

13. Very slow reactions Rusting of iron
4Fe(s) + 3O2(g) + 2nH2O(l)  2Fe2O3 · nH2O(s)

14. Extremely slow reactions
CaCO3(s) + 2H+(aq)  Ca2+(aq) + CO2(g) + H2O(l)
Before corrosion
After corrosion

15. Two Ways to Express Reaction Rates
1. Average rate
2. Instantaneous rate (rate at a given instant)

16. Amount is usually expressed in
Concentration
mol dm−3
Mass g
Volume
cm3 or dm3
Pressure
atm

17. Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Q 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)
(a)

18. Q. 2. 36 g of magnesium reacted with 50. 0 cm3. of 1
Q 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)
(b)

19. 2.(c) Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) Mg is the limiting reactant
Decrease in concentration of HCl(aq) in 90 s

20. Rate of reaction w.r.t. HCl(aq)
2.(d)
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) =
Rate of reaction w.r.t. HCl(aq)
Rate of reaction w.r.t. MgCl2(aq)
2 
Increase in concentration of MgCl2(aq) in 90 s

21. For the chemical reaction aA + bB  cC + dD
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
[X] = molarity of X

22. For the chemical reaction aA + bB  cC + dD
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
Units : mol dm3 s1, mol dm3 min1, mol dm3 h1…etc.

23. Graphical Representation of Reaction Rates – Rate curves
A rate curve is a graph plotting the amount of a reactant or product against time.

24. Consider the reaction
A  B C (reactant) (product)

25. 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.

26. -ve slope of curve of reactant A
 [A]  with time

27. +ve slope of curve of product B
 [B]  with time

28. 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.

29. The rate of the reaction gradually  as the reaction proceeds.
Flat curve  reaction completed

30. Concentration of product Z (mol dm−3)
Q.3
X + Y  2Z
Time of reaction (min)
Concentration of product Z (mol dm−3) A B C

31. Concentration of product Z (mol dm−3)
Time of reaction (min)
Concentration of product Z (mol dm−3) A B C
X + Y  2Z

32. Concentration of product Z (mol dm−3)
X + Y  2Z
Time of reaction (min)
Concentration of product Z (mol dm−3) A B C
1.6

33. Concentration of product Z (mol dm−3)
X + Y  2Z
Time of reaction (min)
Concentration of product Z (mol dm−3) A B C
5.1
2.7

34. Concentration of product Z (mol dm−3)
X + Y  2Z
Time of reaction (min)
Concentration of product Z (mol dm−3) A B C

35. Methods of Measuring Reaction Rates
A. Physical measurements
1. Continuous measurements
2 Initial rate measurements (Clock reactions)
B. Chemical measurements (Titration)

36. 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

37. 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)

38. 1.1 Measurement of large volume changes
A typical laboratory set-up for measuring the volume of gas formed in a reaction
Temperature is kept constant

39. Volume of gas formed (cm3)
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)
Time of reaction (min)
Volume of gas formed (cm3)

40. 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 
Sigmoid curve

41. 1.2 Measurement of small volume changes - Dilatometry
Liquid phase reaction mixture
Capillary tube
CH3COOH(l) + CH3CH2OH(l)  CH3COOCH2CH3(l) + H2O(l)

42. 1.3 Measurement of mass changes
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)

43. measured volume of standard hydrochloric acid
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

44.  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.

45. =  rate 2 Loss of mass (m) mfinal = total mass loss time mfinal - mt
mfinal = mfinal – m0
(∵ m0 = 0)
=  rate 2

46. 1.4 Colorimetry
∵ colour intensity  [coloured species]

47. colour intensity  as reaction proceeds
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

49. cuvettes
A colorimeter

50. Complementary colours
Yellow light
Yellow filter
Blue solution
Complementary colours

51. Red  Cyan
Pairs of opposite colours are complementary colours

52. Red  Cyan Green  Magenta
Pairs of opposite colours are complementary colours

53. Red  Cyan Green  Magenta Blue  Yellow CMYK
Pairs of opposite colours are complementary colours

54. When mixed in the proper proportion, complementary colours produce a neutral color (grey, white, or black).

55. I0 = intensity before absorption
I = intensity after absorption

56. I0 I If I = I0 , If I = 0 , %T = 100% %T = 0% A = log101 = 0
zero absorption
complete absorption

57. A = bC
Beer’s law

58. Deviation at higher concentrations
A calibration curve is first constructed for AC conversion

59. time
[I2]
Q.5
time A

60. 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

61. 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

62. 1.6 Measurement of pressure changes
PT = total pressure of the reaction mixture

63. n  as the reactions proceed  PT  as the reactions proceed
Q.6
(i) 2NO(g) + 2H2(g)  N2(g) + 2H2O(g)
(ii) 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

64. Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
to data-logger interface and computer
pressure sensor
magnesium ribbon
suction flask
dilute hydrochloric acid
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)

65. A(g) + B(g)  products

66. 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.

67. Initial Rate Measurements-Clock Reactions
1. A set of experiments is done in which all reaction conditions but one are kept constant.
S2O32–(aq) + 2H+(aq)  SO2(aq) + H2O(l) + S(s)
Experiment
[S2O32(aq)] / M
[H+(aq)] / M 1
0.10 2
0.08 3
0.04 4
0.02

68. Initial Rate Measurements-Clock Reactions
S2O32–(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.

69. The beaker containing the reaction mixture is placed over a cross marked on a white tile.

70. As more sulphur forms, the reaction mixture becomes more cloudy.

71. The cross becomes more and more difficult to see and finally disappears.

72. Average rate in the early stage
S2O32–(aq) + 2H+(aq)  SO2(aq) + H2O(l) + S(s)
yellow precipitate
Average rate in the early stage
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,
Averagerate 1
time taken to ‘blot out’ the mark 

73. Averagerate 1
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.

74. amount of S
If S and t are small(early stage)
time

75. Since S is a constant

76. Since HCl is in large excess, [H+(aq)]y  constant at the early stage
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

77. Expt.
[S2O32(aq)] (M)
[H+(aq)] (M)
Time taken (t)
to mask the
mark / s
/ s1 1
0.10 10 2
0.08 13 3
0.04 25 4
0.02 50

78. Q.7
Linear  x = 1
[S2O32(aq)]

79. Other Examples of Clock Reactions : -
5I(aq) + IO3(aq) + 6H+(aq)  3I2(aq) H2O(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) starch  deep blue complex
(excess) (fixed)
Time taken for the reaction mixture to turn deep blue is measured.

80. Other Examples of Clock Reactions : -
5I(aq) + IO3(aq) + 6H+(aq)  3I2(aq) H2O(l)
I2(aq) + 2S2O32(aq)  2I(aq) + S4O62(aq)
(fixed) (fixed)
I2(aq) starch  deep blue complex
(excess) (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

81. Other Examples of Clock Reactions : -
5Br(aq) + BrO3(aq) + 6H+(aq)  3Br2(aq) H2O(l)
3Br2
(fixed) (fixed)
Br methyl red  colourless
(excess) (fixed)

82. Advantages of physical measurements
Suitable for fast reactions.
Small sample size
More accurate than chemical method (titration)
No interruption  continuous measurements
Can be automated.

83. Disadvantages of physical measurements
More sophisticated
More expensive
More specific – only suit a limited number of reactions.

84. B. Chemical Measurements (Titration Methods)
1. Start a reaction with all reaction conditions but one fixed.
Withdraw and quench fixed amounts of the reaction mixture at different times.

85. Cooling the reaction mixture rapidly in ice.
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.

86. B. Chemical Measurements (Titration Methods)
1. Start a reaction with all reaction conditions but one fixed.
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.

87. H+ as catalyst
CH3COCH3 + I 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)

88. 2S2O32(aq) + I2(aq)  S4O62(aq) + 2I(aq)
H+ as catalyst
CH3COCH3 + I CH3COCH2I + HI
Q.9
Titrated with standard solution of Na2S2O3(aq) using starch as indicator (added when the end point is near)
2S2O32(aq) + I2(aq)  S4O62(aq) + 2I(aq)
Colour change at the end point : deep blue to colourless

89. H+ as catalyst
CH3COCH3 + I 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)

90. Advantages of titrimetric method
Only simple apparatus are required.
Can be applied to a great variety of slow reactions.

91. Disadvantages of physical measurements
Not suitable for fast reactions.
It takes time to withdraw samples and perform titration.
Reactions are disturbed – NOT continuous
Time consuming – NOT automated

92. Factors Affecting Reaction Rates

93. Collision Theory
No reaction
Sufficient K.E.
Incorrect orientation

94. Collision Theory
No reaction
Correct orientation Insufficient K.E.

95. Effective collision Collision Theory Sufficient K.E.
Correct orientation
Effective collision

96. 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

97. Collision Theory Activation energy Higher Ea
 more K.E. required for effective collision
 slower reaction

98. Collision Theory Activation energy Lower Ea
 less K.E. required for effective collision
 faster reaction

99. 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

100. Factors Affecting Reaction Rates
concentration
particle size
pressure
catalyst
temperature
light

101. Effect of concentration
e.g. Reaction between Mg and HCl

102. Effect of concentration
2.0 M HCl
(b) 1.0 M HCl
(c) 0.5 M HCl
Reaction rate:
(a) > (b) > (c)

103. Effect of concentration
Time for reaction to complete: t1 < t2 < t3
Higher [HCl(aq)]
 Faster reaction

104. [X] 
 Reactant particles are more crowded
 Collision frequency 
 Number of effective collisions 
 Reaction rate 

105. Rate  k[A]x[B]y For the reaction aA + bB  cC + dD
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

106. Rate  k[A]x[B]y For the reaction aA + bB  cC + dD
x and y can be  integers or fractional
x  y is the overall order of reaction.
x, y can ONLY be determined experimentally.

107. Effect of pressure
Only applicable to reactions involving gaseous reactants.

108. Pressure 
Reactant particles are more crowded
Collision frequency 
No. of effective collisions 
Rate of reaction 

109. Effect of temperature
Applicable to ALL reactions

110. 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 

111. Rate of reaction  exponentially with temperature
T / C
Rate of reaction  exponentially with temperature
In general, a 10oC  in T doubles the rate.

112. Effect of particle size
For a fixed volume of solid,
Smaller particle size  greater surface area

113. Rate involving powdered solid reactant is higher
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

114. Q.11
0.5 g powder
0.5 g granule

115. 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.

116. Effect of Catalyst
Catalysts remain chemically unchanged at the end of reactions.

117. MnO2 as catalyst
H2O2(aq) H2O(l) + O2(g)
Physical measurement

118. Volume of gas formed (cm3)
MnO2 as catalyst
H2O2(aq) H2O(l) + O2(g)
Time of reaction (min)
Volume of gas formed (cm3)

119. Titrimetric method (Q.12)
MnO2 as catalyst
H2O2(aq) H2O(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)

120. Q.13
time
[H2O2]
With MnO2
Without MnO2

121. Br – Br h Br + Br C6H14 + Br  C6H13 + HBr    C6H13Br…
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…

122. 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)

123. Q.14
time
[MnO4]
Rate 
Sigmoid curve
Rate 

124. The END

125. 13.1 Rates of Chemical Reactions (SB p.5)
Back
Example 13-1A
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 =
= mol
Average rate =
= 6.83 × 10–5 mol s–1

126. 13.1 Rates of Chemical Reactions (SB p.5)
Example 13-1B
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? H2O2(l)  2H2O(l) + O2(g) (Molar volume of gas at room temperature and pressure= 24.0 dm3 mol–1)
Answer

127. Example 13-1B Back 13.1 Rates of Chemical Reactions (SB p.5)
Number of moles of O2 =
= 2.08 × 10–4 mol
Average rate =
= 2.89 × 10–9 mol s–1

128. 13.1 Rates of Chemical Reactions (SB p.6)
Example 13-1C
The change in concentration of reactant X in a chemical reaction is illustrated in the graph on the right.

129. 13.1 Rates of Chemical Reactions (SB p.6)
Example 13-1C
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

130. Example 13-1C 13.1 Rates of Chemical Reactions (SB p.6) Initial rate
= Slope of the tangent to the curve at t0 =
= mol dm-3 min-1

131. Example 13-1C 13.1 Rates of Chemical Reactions (SB p.6)
(b) Average rate =
= mol dm-3 min-1

132. Example 13-1C Back 13.1 Rates of Chemical Reactions (SB p.6)
(c) Instantaneous rate at the 3rd minute
= Slope of the tangent to the curve at the 3rd minute =
= mol dm-3 min-1

133. 13.1 Rates of Chemical Reactions (SB p.8)
Check Point 13-1
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
Average rate at 398 K
= –(1 – 0.75) mol dm-3  (4  60) s
= – mol dm-3 s-1

134. 13.1 Rates of Chemical Reactions (SB p.8)
Check Point 13-1
The graph on the right shows the change in concentration of a reactant in a chemical reaction.

135. Check Point 13-1 Answer With the use of the graph above, calculate
13.1 Rates of Chemical Reactions (SB p.8)
Check Point 13-1
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

136. Check Point 13-1 Back 13.1 Rates of Chemical Reactions (SB p.8)
(i) Initial rate =
= -1  10-3 mol dm-3 s-1
Average rate =
= -3  10-4 mol dm-3 s-1
Instantaneous rate =
= -5  10-4 mol dm-3 s-1

137. 13. 2 Expressions of Reactions Rates in Terms of Rates of Changes in
13.2 Expressions of Reactions Rates in Terms of Rates of Changes in Concentrations of Reactants or Products (SB p.10)
Example 13-2
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.
Answer
The rate of the reaction is expressed as:

138. Check Point 13-2 Answer Back
13.2 Expressions of Reactions Rates in Terms of Rates of Changes in Concentrations of Reactants or Products (SB p.10)
Check Point 13-2
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)
Answer
Rate =

139. 13.3 Methods of Measuring Reaction Rates (SB p.11)
Example 13-3A
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.

140. 13.3 Methods of Measuring Reaction Rates (SB p.11)
Example 13-3A
(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.

141. 13.3 Methods of Measuring Reaction Rates (SB p.11)
Example 13-3A
(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.

142. 13.3 Methods of Measuring Reaction Rates (SB p.11)
Answer
Example 13-3A
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 8
Option B
Time after mixing (min)
Volume of HCl added at the end point (cm3) 5 8 10

143. Example 13-3A 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.

144. Example 13-3A Answer (d) Name a suitable indicator for the titration.
13.3 Methods of Measuring Reaction Rates (SB p.11)
Example 13-3A
(d) Name a suitable indicator for the titration.
Answer
(d) Methyl orange / Phenophthalein
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145. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Volume of H2(g) produced (cm3)
13.3 Methods of Measuring Reaction Rates (SB p.13)
Example 13-3B
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
1.0
2.0
3.0
4.0
5.0
6.0
7.0
8.0
9.0
Volume of H2(g) produced (cm3) 15 26 33 38 40 41 42

146. 13.3 Methods of Measuring Reaction Rates (SB p.13)
Example 13-3B
(a) Plot a graph of volume of hydrogen gas produced against time.
Answer
(a)

147. 13.3 Methods of Measuring Reaction Rates (SB p.13)
Example 13-3B
(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.

148. 13.3 Methods of Measuring Reaction Rates (SB p.13)
Example 13-3B
(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.

149. 13.3 Methods of Measuring Reaction Rates (SB p.13)
Back
Example 13-3B
(d) Determine graphically the rate of the reaction at the 5th minute. State the unit.
Answer
From the graph in (a),
rate of reaction
= slope of the tangent to the curve at the 5 minute =
= 2 cm3 min-1

150. Check Point 13-3 Answer Back
13.3 Methods of Measuring Reaction Rates (SB p.15)
Back
Check Point 13-3
Suggest an experimental method for determining the rate of each of the following reactions:
(a) S2O82–(aq) + 2I–(aq)  2SO42–(aq) + I2( aq)
CH3COOCH3(aq) + I2(aq)
 CH3COOCH2I(aq) + HI(aq)
2MnO4–(aq) + 5C2O42–(aq) + 16H+(aq)
 2Mn2+(aq) + 10CO2(g) + 8H2O(l) + H+(aq)
Answer
Colorimetric measurement / titration
Colorimetric measurement
Colorimetric mesurement / titration

151. 13.4 Factors Affecting Reaction Rates (SB p.17)
Let's Think 1
Explain why sawdust burns explosively in pure oxygen but slowly in air.
Answer
A higher concentration of oxygen increases the rate of combustion.
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152. 13.4 Factors Affecting Reaction Rates (SB p.21)
Check Point 13-4
(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)

153. 13.4 Factors Affecting Reaction Rates (SB p.21)
Check Point 13-4
(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)

154. 13.4 Factors Affecting Reaction Rates (SB p.21)
Check Point 13-4
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.