1. Rate of Reactions

2. Fast or Slow Reactions Extremely slow reactions
Iron rusting
Limestone weathering
Extremely fast reactions
Explosion

3. Measuring Rate of Reactions
Some rate of reactions have detectable change with respect to time
Changes that are observable like
When a volume of gas is given off
When this is a change in mass during the reaction
When there are temperature changes
When there are colour changes
When a precipitate forms
When there are pH changes

4. Speed=
Speed= = =
= km/h
= km/h

5. Collision theory
Theory that explains how factors affect rate of reaction
Before reaction can occur, particles must first collide
Particles collide so that bonds are broken and new bonds can form

6. Collision theory
Collision of particles must be effective collisions to produce result
Effective collisions are collisions which
produce enough energy to overcome energy of activation
correct orientation

7. Collision theory
Activation energy is the minimum energy the colliding particles must overcome so that reaction occur
In order for particles to overcome the activation energy, several factors are involved

8. How factors affect rate of reaction according to collision theory
Size of reactant for solid reactant
Temperature of reactant mixture
Concentration of solution reactant
Presence of catalyst
Pressure on the reaction

9. Concentration Rate of reaction increases with increasing concentration
Higher concentration, more reacting particles are present
Greater probability of an effective collision
Faster rate of reaction

10. Concentration – Same no. of moles
Amt of product formed
Higher concentration
Lower concentration
Time/s

11. Concentration – Different no. of moles
Amt of product formed
Higher concentration
Lower concentration
Time/s

12. Speed of reaction doubles when the temperature rises by 10 C
Rate of reaction increases with increasing temperature
High temperature, particles have greater heat energy
Particles move faster with greater kinetic energy
Leading to more collisions between particles
Increased probability of effective collision
Reactions take place faster
Speed of reaction doubles when the temperature rises by 10 C

13. Temperature Amt of product formed Higher temperature Lower temperature
Time

14. Catalyst Presence of catalyst increases rate of reaction
(Presence of inhibitors decreases rate of reaction)
Catalysts lower activation energy of reactants
Aids the formation of unstable intermediate products
Increases probability of formation of products
Faster rate of reaction

15. POTENTIAL ENERGY DIAGRAM
A B AB HEAT P O T E I A L N R G Y
ACTIVATION
ENERGY
WITHOUT
CATALYST
LOWER ACTIVATON
ENERGY PATH
WITH CATALYST
A B
REACTANTS
PRODUCTS AB
REACTION COORDINATE

16. Catalyst
Energy
Absence of catalyst Ea Ea
Use of catalyst
Time/s

17. Catalyst Amt of product formed Use of catalyst Absence of catalyst
Time/s

18. Catalyst
Definition: A substance which increases the rate of a chemical reaction by providing an alternative pathway with a lower activation energy but remains unchanged at the end of the reaction

19. Pressure Rate of reaction increases with increasing pressure
Higher pressures, reacting particles are closer together
Increasing concentration per unit volume
Greater probability of an effective collision
Faster rate of reaction

20. Pressure
Amt of product formed
Higher pressure
Lower pressure
Time/s

21. Particle Size Rate of reaction increases when particle size decreases
Smaller particles has greater surface area than larger particles of the same mass
Greater surface area for collision by another reacting particle
Greater probability of an effective collision
Faster rate of reaction

22. HOW PARTICLE SIZE AFFECTS CHEMICAL REACTION RATE!

23. Particle size Amt of product formed Smaller particle size
Larger particle size
Time/s

24. time taken for reaction
A reaction is fast , the time taken for the reaction is short .
A reaction is slow, the time taken for the reaction is long .
The rate of reaction depends to the speed of reaction .
If a reaction is fast, its rate of reaction is high .
If a reaction is slow, its rate of reaction is low .
The rate of reaction is inversely proportional with time .
Rate ά
time taken for reaction
Rate of reaction = change in quantity of product / reactant
time taken
For gas product , rate of reaction = volume of gas
time
From a graph , the average rate of reaction = the gradient

25. HOW TO DETERMINE RATE OF REACTION FROM GRAPH
AVERAGE RATE OF REACTION
INSTANTANEOUS RATE OF REATION

26. AVERAGE RATE OF REACTION
From X min to Y min
On the X min (2nd)
( from 2nd to 1st)
For first x (3) min
( from 0 to 3rd )
For the whole exp

27. INSTANTANEOUS RATE OF REACTION ( rate of reaction at that time)
Draw tangent to the graph
Rate = Y/X Y X

28. b) Example from the graph, determine:
i) The rate of reaction at 120 s
Instantaneous rate of reaction
= Draw tangent to the graph

29. = 56 – 20 = cm3 s-1
222-18

30. Changes to the graph

31. Changes to the CURVE part of graph
Use positive catalyst
Increase temperature
Increase total surface area
Volume of gas II I
III
Use negative catalyst
Decrease temperature
Decrease TSA
Time

32. Changes to the graph

33. d) Curve I represents the result of the experiment using excess zinc powder and 50cm3 of 1.0 moldm-3 dilute hydrochloric acid
Volume of gas/ cm3
Use positive catalyst
Increase temperature of reactant II I
III
Lower concentration of hydrochloric acid
Time/s

34. (h) Time(second) 30 60 90 120 Burette reading(cm3) x 49.5 y 33.5 z
Time (seconds) 30 60 90
120
150
180
210
240
Burette reading (cm3)
49.5
33.5
23.5
16.0
10.5
5.0
2.0
Volume CO2 (cm3) 16 26 39
44.5
47.5
Time(second) 30 60 90
120
Burette reading(cm3) x
49.5 y
33.5 z
23.5
Total volume of gas(cm3)
x-x
0.00
x-y
16.00
x-z
26.00

35. Connect the point without using ruler!
Volume of CO2, cm3
Connect the point without using ruler!
Not all the point is connected
Time , s

36. It’s must be smooth graph
Volume of CO2 cm3
Cannot make this graph
Straight line
It’s must be smooth graph
Time s

37. Average Rate Of reaction
The average rate of reaction in the first 90 seconds.
= The total volume of gas released in the first 90 seconds
Time taken
Time (seconds) 30 60 90
120
150
180
210
240
Burette reading (cm3)
49.5
33.5
23.5
16.0
10.5
5.0
2.0
Volume CO2 (cm3) 16 26 39
44.5
47.5
33.5÷90=0.372 cm3s-1
unit =

38. i(ii) = The average rate of reaction in the whole experiment.
Time (seconds) 30 60 90
120
150
180
210
240
Burette reading (cm3)
49.5
33.5
23.5
16.0
10.5
5.0
2.0
Volume CO2 (cm3) 16 26 39
44.5
47.5
The average rate of reaction in the whole experiment.
= The total volume of gas released in the whole experiment
Time taken
47.5÷180= cm3s-1 =

39. Analysis of Data Cannot take directly at x
Rate of reaction at t second = gradient AB
= p/q cm3 s-1
Total volume of Hydrogen gas/cm3
Tangent is a line that touch just 1 point of graph in order to calculate gradient B
Tangent p A q
Cannot take directly at x t
Time (second)

40. Tangent
Cannot touch more than 2 points because each of point has a different gradient
Only touch 1 point of curve

41. tangent α

42. Analysis of data Each of point has a different gradient!
Total Volume of CO2(cm3) F D E
Rate of reaction at t1 = gradient AB B
Rate of reaction at t2 = gradient CD C
Rate of reaction at t3 = gradient EF
Each of point has a different gradient! A t1 t2 t3
Time (second)

43. Two methods to calculate tangent:
Total volume of Hydrogen gas/cm3
number of small boxes × value of
1 small unit box B
Tangent Y A X
Time (second)

44. First Method Total volume of Hydrogen gas/cm3 Tangent Time (second)
Gradient of graph:
Total volume of Hydrogen gas/cm3
m = ΔY ΔX B
m = Y2-y1 y2
X2-x1
Tangent A y1 x1 x2
Time (second)

45. Analysis of Data Total volume of Hydrogen gas/cm3 Tangent
Rate of reaction at t second = gradient AB
= p/q cm3 s-1
Total volume of Hydrogen gas/cm3 B
Tangent p A q t
Time (second)

46. Analysis of data Total Volume of CO2(cm3)
Rate of reaction at t1 = gradient AB B
Rate of reaction at t2 = gradient CD C
Rate of reaction at t3 = gradient EF A t1 t2 t3
Time (second)

47. The Rate Law: Reactant Concentration and Rate
The rate of a chemical reaction depends on several
factors ….
Relating Reactant Concentrations and Rate
Consider the general reaction below.
aA bB products
This reaction occurs at constant temperature

48. The reactant fomulars are represented by A and B
The stoichiometric coefficients are represented by
a and b
In this section we will look at reaction rates that are not
affected by concentrations of products
In general, the rate of a reaction increases when the
concentrations of reactants increases.
The dependence of of the rate of a reaction on the con-
centration is given by:

49. This relationship can be expressed in a general
equation called the rate law equation
For any reaction, the rate law equation expresses the
relationship between the concentrations of the
reactants and the rate of the equation.

50. The letter k represents the proportionality constant
called the rate constant
There is a different rate constant for each reaction at
any given temperature
The exponents m and n must be determined by exp-
eriment.
The do not necessarily correspond to the coefficients
of their reactants
They are usually 1 or 2, but values of 0, 3 even fractions
can occur

51. HOW CONCENTRATION EFFECTS REACTION RATES (REACTION ORDER)
REACTIONS WITH RATE EQUATIONS HAVING n = 0 ARE ZERO ORDER REACTIONS. THOSE WITH n = 1 ARE FIRST ORDER AND THOSE WITH n = 2 ARE SECOND ORDER.
IN ZERO ORDER REACTIONS, CHANGING THE CONCENTRATION OF THE REACTANT HAS NO EFFECT ON THE RATE.
IN FIRST ORDER REACTIONS, RATE CHANGES ONE FOR ONE WITH CONCENTRATION CHANGE. FOR EXAMPLE, DOUBLING CONCENTRATION DOUBLES THE RATE.

52. HOW CONCENTRATION EFFECTS REACTION RATES (REACTON ORDER)
IN SECOND ORDER REACTIONS, RATE CHANGES RELATIVE TO THE SQUARE OF THE CONCENTRATION CHANGE. FOR EXAMPLE, DOUBLING THE CONCENTRATION OF THE REACTANT RESULTS IN THE RATE INCREASING TIMES.
THIS KNOWLEDGE OF HOW RATE CHANGES WITH CONCENTRATION DEPENDING ON THE ORDER LETS US FIND REACTION ORDERS BY AN EXPERIMENTAL PROCESS CALLED “METHOD OF INITIAL RATES”
REACTION ORDERS MUST BE DETERMINED EXPERIMENTALLY. THEY CAN NEVER BE DETERMINED FROM THE CHEMICAL EQUATION.

53. SLOPE OF A TANGENT LINE TO AN AMOUNT VS. TIME GRAPH = RATE
REACTION RATES
A  B C M O L E S A
GRAPH 1 M O L E S A
GRAPH 2
RXN RATE?
FORWARD OR REVERSE?
FORWARD RXN
RATE = CONSTANT
RXN RATE?
RATE = 0
TIME
TIME
RXN RATE?
FORWARD OR REVERSE?
FORWARD RXN
RATE = VARIABLE M O L E S A M O L E S A
RXN RATE ?
FORWARD OR
REVERSE RXN?
REVERSE RXN
RATE = VARIABLE
GRAPH 3
GRAPH 4
TIME
TIME
SLOPE OF A TANGENT LINE TO AN AMOUNT VS. TIME GRAPH = RATE

54. HOW CONCENTRATION EFFECTS REACTION RATES (INITIAL RATES)
USING THE METHOD OF INITIAL RATES REQUIRES THAT A REACTION BE RUN AT SERIES OF DIFFERENT STARTING CONCENTRATIONS AND THE RATE BE DETERMINED FOR EACH.
GIVEN THE FOLLOWING DATA FOR THE REACTION
A  B + C
(TABLE 1)
EXPT [A] RATE (M/SEC)
x x 10 -1
x x
x x
AS CONCENTRATION OF A DOUBLES, RATE DOUBLES. THE REACTION IS FIRST ORDER IN REACTANT A
RATE = k[A]1 OR RATE = k[A]

55. HOW CONCENTRATION EFFECTS REACTION RATES (INITIAL RATES)
FOR THE REACTION: A B  C D
(TABLE 2)
x x 10 – x 10 -1
x x x
x x x 10 –1
USING EXPT 1 AND 2, [A] DOUBLES AND [B] IS CONSTANT. THE DOUBLING OF THE RATE IS THEREFORE CAUSED BY REACTANT A AND THE ORDER WITH RESPECT TO A IS FIRST.
USING EXPT 1 AND 3, [A] IS CONSTANT AND [B] IS DOUBLED. THE FOUR TIMES RATE INCREASE IS THEREFORE CAUSED BY REACTANT B AND THE ORDER WITH RESPECT TO B IS SECOND.
RATE = k[A]1[B]2 OR RATE = k[A][B]2

56. HOW CONCENTRATION EFFECTS REACTION RATES (RATE CONSTANTS)
FROM INITIAL RATES DATA TABLE 1,
RATE = k [A] , k CAN BE CALCULATED BY SUBSTITUTING ANY DATA SERIES INTO THE RATE EQUATION, FOR EXAMPLE, FROM EXPT 1 ON TABLE 1
[A] = 1 x 10 –3 M , RATE = 4 x 10 –1 M/SEC
4 x 10 –1 M/SEC = k (1 x 10 –3 M )
k = 1 x 102 SEC-1 OR 1 x 102 / SEC

57. HOW CONCENTRATION EFFECTS REACTION RATES (RATE CONSTANTS)
FROM INITIAL RATES DATA TABLE 2,
RATE = k [A] x [B]2, k CAN BE CALCULATED BY SUBSTITUTING ANY DATA SERIES INTO THE RATE EQUATION, FOR EXAMPLE, FROM EXPT 1 ON TABLE 2
[A] = 1 x 10 –3 M , [B] = 1 x 10 –3 M
RATE = 4 x 10 –1 M/SEC
4 x 10 –1 M/SEC = k (1 x 10 –3 M ) x (1 x 10 –3 M )2
k = 4 x 105 M-1 SEC-1 OR 4 x 105 / M x SEC

58. Let us see how the initial rates method works 2N2O3(g) 2NO2(g) + O2(g)
The general rate law equation for this reaction is:
Rate = k [N2O3]m
To determine the value of m, a chemist performs three experiments. A different initial concentration of [N2O35]0
Is used for each experiment. The 0 represents t=0

59. Initial rate (mol/(L.s))
Experiment
Initial [N2O3]0 (mol/L)
Initial rate (mol/(L.s)) 1
0.010
4.8 x 10 -6 2
0.020
9.6 x 10 -6 3
0.030
1.5 x 10-5
Value of m can be determined with at least two different methods
by inspection
rate law equation

60. 1. By inspection
When the [N2O5] is doubled expts 1 and 2 doubles
the rate also doubles
when [N2O5] is tripled, (expts 1 and 3) the rate triples
this indicates a first-order relationship as follows
Rate = k [N2O5]1

61. 2. Compare rate law equation using ratios
this method very useful when relation between conc and rate are not immediately obvious from data
Write the rate expressions for expts 1 and 2 as follows:
Rate1 = k [0.010]m
= 4.8 x 10-6 mol/(L.s)
Rate2 = k [0.020]m
= 9.6 x 10-6 mol/(L.s)

62. Create a ratio to compare the two rates
Rate1 = k(0.010 mol/L)m = 4.8 x 10-6 mol/L.s
Rate k(0.020 mol/L)m x 10-6 mol/L.s
Since k is a constant at constant temp you can cancel out
k(0.010 mol/L)m = 4.8 x 10-6 mol/L.s
k(0.020 mol/L)m x 10-6 mol/L.s
(0.5)m = 0.5
m = 1 (by inspection)

63. Determining the Rate Constant
Once you know the rate law equation for a reaction, you can calculate the rate constant using results from any of the experiments
Rate = k [N2O5]1
You can use data from any of the three experiments to calculate k
4.8 x 10-6 mol/(L.s) = k(0.010 mol/L)

64. k = 4.8 x 10-6 mol/ (L.s)
0.010 mol/L
= 4.8 x 10-4 s-1

65. TEMPERATURE & REACTION RATE
AT ANY TEMPERATURE THE MOLECULES IN A SYSTEM HAVE A DISTRIBUTION OF KINETIC ENERGIES (SIMILAR TO THE DISTRIBUTION OF THE SPEEDS OF CARS ON A HIGHWAY).
AS THE TEMPERATURE INCREASES, THE AVERAGE KINETIC ENERGY OF THE MOLECULES INCREASE (THEY MOVE FASTER AT HIGHER TEMPERATURES) AND THEREFORE COLLIDE WITH EACHOTHER MORE FREQUENTLY AND HIT HARDER WHEN THEY DO COLLIDE.
AS A RESULT, REACTION RATE INCREASES WITH TEMPERATURE.
REACTION RATE
TEMPERATURE

66. KINETIC ENERGY DISTRIBUTION CURVE
TEMPERATURE = T1 N U M B E R O F L S
AVERGE ENERGY
MOLECULES
LOW ENERGY
MOLECULES
HIGH ENERGY
MOLECULES
KINETIC ENERGY

67. KINETIC ENERGY DISTRIBUTION CURVE
TEMP 2 > TEMP 1
AT HIGHER
TEMPERATURES,
MOLECULES HAVE
HIGHER ENERGIES
ON AVERAGE N U M B E R O F L S
TEMPERATURE = T1
TEMPERATURE = T2
KINETIC ENERGY

68. TEMPERATURE & REACTION RATE
IN ORDER TO REACT, MOLECULES MUST COLLIDE WITH SUFFICIENT ENERGY. THIS MININIUM ENERGY FOR REACTION IS CALLED ACTIVATION ENERGY.
AS THE TEMPERATURE OF A SYSTEM IS INCREASED, THE NUMBER OF MOLECULES WITH THE NECESSARY ENERGY FOR REACTION (THE ACTIVATION ENERGY) INCREASES.

69. KINETIC ENERGY DISTRIBUTION CURVEM B E R O F L S
TEMPERATURE = T2
TEMPERATURE = T1
ACTIVATION ENERGY
MOLECULES WITH
SUFFICIENT ENERGY
TO REACT AT T2
MOLECULES WITH
SUFFICIENT ENERGY
TO REACT AT T1
KINETIC ENERGY
AS TEMPERATURE , REACTION RATE

70. TEMPERATURE & REACTION RATE
THE ENERGY CHARACTERISTICS OF A CHEMICAL REACTION CAN BE SHOWN ON A POTENTIAL ENERGY DIAGRAM.
THIS GRAPH SHOWS THE ENERGY STATE OF THE SYSTEM AS REACTANTS PROCEED THROUGH THE ACTIVATED COMPLEX TO FORM THE PRODUCTS. THE ACTIVATED COMPLEX IS THE INTERMEDIATE STATE (MOLECULAR FORM) WHICH REACTANTS GO THROUGH AS THEY CONVERT INTO THE PRODUCTS.
THE ACTIVATION ENERGY IS THE ENERGY REQUIRED TO FORM THE INTERMEDIATE ACTIVATED COMPLEX MOLECULE.

71. POTENTIAL ENERGY DIAGRAM
A B AB HEAT P O T E I A L N R G Y
ACTIVATED
COMPLEX
ACTIVATION
ENERGY
EXOTHERMIC
 H = (-)
ENERGY IS
RELEASED
REACTANTS
A B
 H
PRODUCTS AB
REACTION COORDINATE

72. SURFACE AREA & REACTION RATE
REACTANT THAT IS INTERIOR CANNOT BE ATTACKED UNTIL THE
EXTERIOR REACTANT IS CONSUMED. THE REACTION RATE IS SLOW!

73. SURFACE AREA & REACTION RATE
WHEN SURFACE AREA IS INCREASED MORE REACTANTS ARE EXPOSED
TO EACHOTHER SIMULTANEOUSLY AND THE REACTION IS RAPID.

74. REACTION MECHANISMS
ALL REACTIONS, NO MATTER HOW SIMPLE FOLLOW A PARTICULAR REACTION PATHWAY CALLED A MECHANISM. IT IS A SERIES OF STEPS WHICH LEAD TO THE FORMATION OF PRODUCTS.
DURING THE PROCESS OFTEN MOLECULES CALLED INTERMEDIATES ARE FORMED AND SUBSEQUENTLY CONSUMED.
THE SLOWEST STEP IN THE SERIES OF STEPS THAT MAKE UP THE MECHANISM IS CALLED THE RATE DETERMINING STEP.
THE SPEED OF THE RATE DETERMINING STEP DEPENDS ON THE COMPLEXITY OF THE STEP (NUMBER OF MOLECULES INVOLVED CALLED THE MOLECULARITY) AND THE BOND STRENGTHS OF THE REACTING COMPONENTS

75. REACTION MECHANISM FOR;
2 NO O2  2 NO2
STEP I NO  N2O2 (fast)
STEP II N2O2 + O2  2 NO2 (slow) * NO
NO2 O2 NO
NO2
N2O2 INTERMEDIATE

76. IN THE REACTION MECHANISM FOR:
REACTION MECHANISMS *
UNLIKE OVERALL REACTIONS, THE REACTON ORDERS FOR THE STEPS IN A MECHANISM ARE BASED ON THE NUMBER OF MOLECULES REACTING IN THAT STEP.
IN THE REACTION MECHANISM FOR:
2 NO O2  2 NO2
STEP I NO  N2O2 (fast)
STEP II N2O2 + O2  2 NO2 (slow)
SINCE THERE IS ONLY ONE N2O2 AND ONE O2 IN THE RATE DETERMING STEP (THE SLOW STEP), THE RATE EQUATION FOR THE REACTION IS:
RATE = k [N2O2] x [O2]
THE REACTION IS FIRST ORDER IN BOTH REACTANTS