1) The document discusses kinetic chemistry, including reaction rates, rate constants, reaction orders, and factors that affect reaction rates such as concentration, temperature, surface area, and catalysts. 2) It provides examples of determining reaction orders based on changes in reaction rate with changes in concentration. Reaction orders can be fractional, zero, or higher than one. 3) The rate constant k depends on temperature according to the Arrhenius equation, and increases with higher temperature based on the activation energy of the reaction.

1. Chapter 5
Kinetic Chemistry
PhD. Đặng Văn Hân
Office: 112B2 or 804H3 Building
Email: dvhan@hcmut.edu.vn
Faculty of Chemical Engineering
Department of Inorganic Technology
General Chemistry

2. 2
Outline
1. Chemical Thermodynamics vs. Kinetics
2. Reaction Rate (rrxn)
3. Reaction Rate Constant (k)
4. Reaction Orders
5. Main Effects on the Reaction Rate
6. Catalyst and Mechanism

3. Thermodynamics vs. Kinetics
 THERMODYNAMICS
Predicts direction and ‘driving force’ of
chemical reactions based ONLY on the
properties of reactants and products.
 KINETICS
Predicts rate of chemical reaction depend on
the pathway from reactants to products.
Domain of
thermodynamics
(Initial & Final States)
Reaction progress
Energy
Domain of
kinetics
Reactants
Products
- The rate of chemical reactions;
- How to control or influence the rate (lower – faster);
- The mechanism of reaction.

4. Reaction Rate
Reaction Rate: is expressed as the
concentration of reactant consumed or the
concentration of product formed per unit time
t = 0 s t = 30 s t = 60 s
 Determination of Reaction Rates:
Law of
Mass Action
Con. Change
in the time
Consider rxn: A → B
 UNIT of Reaction Rates: M/s or mol.l-1s-1

5. 5
 4 major factors strongly affect on the reaction rate:
Concentration of reactant
 Temperature
 Surface area
 Catalyst
Coefficients
 Coefficients in a chemical equation: are the simplest numbers used
to balance chemical equations and are placed in front of a chemical
symbols or formula. Example: 2H2 + 1O2 → 2H2O

6. The average rate
Definition: Is the concentration change of reactants or products that occurs
in the unit of time.
C
[mol
L
-1
]
aA + bB = cC
  CA  CB  CC
CA < 0 CB < 0 CC > 0
-
1
a
CA = -
1
b
CB = +
1
c
CC
𝑟A = -
𝐶𝐴

 The average rate of A, B, C:
𝑟C = +
𝐶𝐶

𝑟B = -
𝐶𝐵

 The average reaction rate:
𝑟rxn = -
1
𝑎
𝐶𝐴

= -
1
𝑏
𝐶𝐵

= +
1
𝑐
𝐶 𝐶

rrxn =
r𝐴
𝑎
=
r𝐵
𝑏
=
r𝐶
𝑐

7. The instantaneous Rate
C
[mol
L
-1
]
aA + bB = cC
  CA  CB  CC
dCA < 0 dCB < 0 dCC > 0
-
1
a
dCA = -
1
b
dCB = +
1
c
dCC
rA = -
𝑑𝐶𝐴
d
 Instantaneous rate of A, B, C:
rB = +
𝑑𝐶𝐶
d
rB = -
𝑑𝐶𝐵
d
 The instantaneous rate of reaction:
𝑟𝑟xn = -
1
𝑎
𝑑𝐶𝐴
d
= -
1
𝑏
𝑑𝐶𝐵
d
= +
1
𝑐
𝑑𝐶𝐶
d
Definition: Is the reaction rate at any given point in time.
The instantaneous rate at a given time
corresponds to the slope of a line tangent
to the concentration-versus-time curve.

8. 8
Reaction Rate Based on Mass-Action Law
Law of Mass Action (M. Guldberg and P. Waage)
aA (g) + bB (g) = cC (g) + dD (g)
At T = const., consider the simple homogeneous reaction:
Reaction rate: r = k[A]a[B]b
With:  r: reaction rate
 k: reaction rate constant
1. Reaction nature 2. Temperature 3. Catalysts
Correct solution for
simple reactions or
for each step of a
complex reaction

9. Simple vs. Complex Reaction
Simple Reaction: only undergoes 1 stage (step)
 Each step is called a simple reaction
 ∑ Steps (simple reactions): reaction mechanism
H2 (g) + I2(g) = 2HI(g) k1
In complex reaction, the overall
reaction rate is determined by the
lowest rate of simple rxn.
Complex Reaction: undergoes many stages (steps)
CH4 (g) + Cl2(g)
𝒉𝒗
CH3Cl(g) + HCl (g)
Cl2
ℎ𝑣
Cl* + Cl* k1
CH4 + Cl* → CH3
* + HCl k2
CH3
* + Cl2 → CH3Cl + Cl* k3
CH3
* + Cl* → CH3Cl k4

10. 1
Example 1: Choose the CORRECT statement. The mechanism of
complex reaction describes:
2NO2 (g) + F2 (g) → 2NO2F (g)
Can be explained through 2 simple reactions:
NO2 (g) + F2 (g)  NO2F (g) + F (g) (low)
NO2 (g) + F (g)  NO2F (g) (fast)
The rate formula of this reaction will be depicted as:
A. 𝑟 = 𝑘
𝐶𝑁𝑂2𝐹
2
𝐶𝑁𝑂2
2 .𝐶𝐹2
B. 𝑟 = 𝑘
𝐶𝑁𝑂2
2
.𝐶𝐹2
𝐶𝑁𝑂2𝐹
2 C. 𝑟 = 𝑘𝐶𝑁𝑂2
2
𝐶𝐹2
D. 𝑟 = 𝑘𝐶𝑁𝑂2
𝐶𝐹2
Complex reaction

11. 2N2O5 (g) = 4NO2 (g) + O2 (g)
N2O5 = N2O3 + O2 (1); low reaction rate
→ r1 = k1.[N2O5]
N2O5 + N2O3 = 4NO2 (2); high reaction rate
→ r2 = k2.[N2O5].[N2O3]
The 1st stage determines the overall rate → rrxn = v1 = k1.[N2O5]
Example 2:
There are two successive stages:
→ The 1st reaction order
Complex reaction

12. Complex reaction
Example 3 (5.4): The chemical reaction, 2NO(g) + Br2(g)  2NOBr(g) proceeds
as the following elementary steps:
Step 1: NO(g) + Br2(g) ⇌ NOBr2(g) (fast)
Step 2: NOBr2(g) + NO(g)  2NOBr(g) (low)
The rate law expression will be:
A. rate = k[NO][Br2] B. rate = k[NO] C. rate = k[Br2]2. D. rate = k[NO]2[Br2]
Slow stage → rL = kL[NO][NOBr2]
Solution:
unstable
Fast stage → K =
[NOBr2]
[NO][Br2]
→ [NOBr2] = K[NO][Br2]
r = k[NO]2[Br2]
D: correct

13. 13
The general reaction rate
Consider the homogeneous rxn: aA + bB = cC + dD
The general reaction rate: r = k.CA
m .CB
n
 m+n: the overall reaction order which can be a negative value, integer,
fraction or zero. Values of m & n are calculated by experiments
 m: the reaction order of A, m the reaction order of B.
Simple reactions
n = a ; m = b
Complex reactions n  a ; m  b
 1st and 2nd reaction orders are common, whereas zero or 3rd reaction
order is uncommon and reaction orders larger 3 hardly occur.

14. 14
Reaction Orders
The 3rd order respects to NO and the overall rxn = 3
The 1st order respects to NO2 and F2; the overall rxn order = 2
The 1st order for H2O2 and I-; Zero order for H+; the overall rxn = 2
3NO (g) → N2O (g) + NO2 (g) r = k[NO]3
2NO2 (g) + F (g) → 2NO2F (g) r = k[NO2] [F2]
H2O2 (l) + 3I- (l) + 2H+ (aq) → 2H2O (l) + I3
- (l)
r = k[H2O2] [I-]

15. 15
Example: Choose the CORRECT statement. Consider reaction:
2NO (g) + O2 (g) = 2NO2 (g)
The formula of the overall reaction rate is 𝐫 =
𝟏
𝟐
𝐱
𝐝[𝐍𝐎𝟐]
𝐝𝐭
= 𝐤 𝐍𝐎 𝟐 𝐎𝟐
So, we can conclude that:
1) The 1st order for O2 and the 2nd order for NO
2) Complex reaction
3) The overall reaction order is 3
4) The above reaction rate is the average reaction rate
A. 2, 3 and 4 B. 1, 2 and 3 C. 1, 3 and 4 D. Only 1 and 3
Reaction Orders

16. 16
The Effects of Reaction Order on Reaction Rate
Consider the simple rxn: A  B + C rrxn = k𝑪𝑨
𝒏
 Rxn order n = 0 → rrxn = k[A]0 = k → the rxn rate unchanges as the
concentration of reactant A change in time;
 Rxn order n = 1 → rrxn = k[A]1 = kA → the rxn rate doubles as the
concentration of reactant A increases double times;
 Rxn order n = n → rrxn = k[A]n → the rxn rate increases 2n times as the
concentration of reactant A doubles;

17. 17
Determination of Reaction Orders and Rates
1. CHEMICAL METHODS:
Using the quantative analysis: No. of analytic samples are large  titration
2. PHYSICAL METHODS:
Using some method to calculate the change of reactant com. in times:
spectrometer, pH, pressure, currency, turbidity measurements, etc.
H2SO4 + Na2S2O3 = H2O + Na2SO4 + S + SO2
The general-reaction rate: r = k.CA
m .CB
n
The overall reaction rate (m+n) must be experimentally
calculated even not only based on reaction mechanism:

18. 18
Reaction-Order Cal. based on [Reactant] Change
r1=k[A1]m[B1]n
r2=k[A2]m[B2]n
𝑟2
𝑟1
= [
𝐴2
𝐴1
]𝑚
[
𝐵2
𝐵1
]𝑛
In experiments: The change of reactant A or B
Example: A + B  C
We have: rrxn = kCA
mCB
n
𝑘 =
𝑟𝑟𝑥𝑛1
[𝐴]1
2
[𝐵]1
=
1.5𝑥10−6 𝑀. 𝑠−1
(1.0𝑥10−2𝑀)2(1.0𝑥10−2𝑀)
= 1.5 𝑀−2𝑠−1
r2/r1= 2 = [A2/A1]m[B2/B1]n = 1m 2n  n=1
r3/r1 = 4 = 2m1n  m=2
rrxn = kCA
2CB
Exp. 1 & 2:
Exp. 2 & 3:

19. 19
The Rate Constant (k)
 Unit:
[k] =
[𝑚𝑜𝑙]
[𝑙.s]
.[
[𝑚𝑜𝑙]
[𝑙]
]−(𝑛+𝑚)
= [mol/l] (1-orders) .[s]-1
 The zero reaction order: k = M.s-1 or mol.l-1.s-1
[k] = [mol/l](1-orders) .[s]-1 = M1-orders.s-1
 The 1st reaction order: k = s-1
 The 2nd reaction order: k = M-1.s-1 or l.mol-1s-1
 The nth reaction order: k = M1-n.s-1 or mol1-n.ln-1.s-1

20. 20
RT
E
e
A
k
*
.


 But, independent on concentration of reactants
The activation energy (J): is the
minimum energy necessary for
reaction occurrence.
 Nature of reaction and Temp.
 The activation energy
 Catalysts
Frequency factor
(measure a favorable collision)
k depends on:
Arrhenius Equation:
The Rate Constant (k)

21. 21
Based on Arrhenius Eq., At T1 → we had k1
T2 → What is the value of k2?
RT
E
e
A
k
*
.


ln
k2
k1
= −
E∗
R
(
1
T2
−
1
T1
)
The Effect of Temperature on Rate Constant (k)

22. 22
Main effects on the reaction rate
 Stirring, light, …
The general reaction rate: rrxn = k.CA
n .CB
m
RT
E
e
A
k
*
.


with
 Nature reaction.
 [Reactants]  rrxn 
 Temperature
 Catalysts
 Surface area (homogeneous rxn): S rrxn
 Solvent (solution reaction).
rrxn depends on

23. 23
The effects of concentration
0. THE ZERO REACTION ORDER
Consider a simple rxn: A  Products
Reaction rate: k
kC
kC
dt
dC
A
n
A
A




 0
rxn
r
 


t
0
[A]
]
[ 0
A
t
A kt
dC
kt
C
C o
A
A 


Integrate from 0 (corresponding to 0 s and
𝐶𝐴
0
) to t (corresponding to t s and CA )

24. 24
Half-life Time of Reaction (t1/2)
HALF-LIFE TIME, t1/2, is the amount of time required for
the reactant to be reduced to exactly half of its starting
concentration ([A]t=1/2 = ½ [A]t=0 )
In the case of the zero order:
and CA,t = 1/2 𝑪𝑨
𝒐
k
C
t
o
A
2
2
1 
kt
C
C o
A
A 



25. 25
1. THE 1ST REACTION ORDER
Consider a simple rxn: A  Products
Reaction rate: A
n
A
A
kC
kC
dt
dC




rxn
r
 


t
0
[A]
]
[ 0
A
A
C
dC
A
t
kt     kt
t 

 0
A
ln
A
ln
 
 
kt
t










0
A
A
ln
The half-life time of 1st rxn order ONLY depends on rate constant (k)
Half-life time:
k
k
t
693
,
0
)
2
ln(
2
1 

The effects of concentration

26. The GRAPH FORM of the 1st RXN ORDER
0
C
ln
ln A
A kt
C 


Corresponds to the graph: y = ax + b
Determination of Rate Constant base on Graph
k value is the slope this equation (*)
k
(*)

27. 27
2. The 2nd REACTION ORDER
Consider a simple rxn: 2A  Products
Reaction rate:
2
rxn
r A
n
A
A
kC
kC
dt
dC




 

t
0
[A]
]
[ 0
2
A
A
C
dC
A
t
kt kt
C A
A


0
C
1
1
Half-life time:
0
2
1
1
A
kC
t 
The effects of concentration

28. 28
Consider a simple rxn : A + B  Products
Reaction rate: B
A
B
A
C
kC
dt
dC
dt
dC





rxn
r
After integration:
B
A
A
B
B
A C
C
C
C
C
C
kt
0
0
0
0
ln
1


The effects of concentration
2. The 2nd REACTION ORDER

29. 29
Consider a simple rxn: 3A  Products
Reaction rate:
3
rxn
r A
n
A
A
kC
kC
dt
dC




 


t
0
[A]
]
[ 0
3
A
A
C
dC
A
t
kt









 2
2
0
1
1
2
1
A
A C
C
kt
The effects of concentration
3. The 3rd REACTION ORDER

30. 30
Short Summary

31. 31
The effect of Temperature
 Van’t Hoff Principle
 As the temperature increase of 10oC, the rxn rate increases 2 - 4 times
4
2
10


 
T
T
k
k

 This principle is ONLY CORRECT in small or mediate temp. ranges
k2
k1
= e
−
E∗
R
(
1
T2
−
1
T1
)
At T1 → we had k1
T2 → What is the value of k2?
 Based on Arrhenius Eq.:
𝛾: is the temperature
coefficient
𝛾n
=
kT+10n
kT
General

32. 32
Example 1: The decomposition rxn of N2O5, given
5
30
10
6
.
3
0



C
k
7
0
10
9
.
7
0



C
k C
k 0
100
 
 
  7
10
100
10
0
100
10
3
7
5
0
3
10
0
3
10
9
.
7
86
.
3
86
.
3
86
.
3
10
9
.
7
10
6
.
3
0
0
0
0
0

















C
C
C
C
C
k
k
k
k
k


and . Calculate
We have:
The effect of Temperature

33. 33
Example 2: (5.31) A chemical reaction was terminated after 3 hours at
20oC. At what temperature will the reaction be terminated after 20 min? Given
that the temperature coefficient of the reaction is 3.
A. at 30oC B. at 40oC C. at 50oC D. at 60oC
We have:
The effect of Temperature
𝛾n
=
kT2
kT1
=
k20+10n
k20
=
t2
t1
=
180
20
= 9
n = 2 ⇢ T2 = 20 + 2*10 = 40oC B: Correct

34. Reaction Mechanism
Once molecules collide they may react together or they
may not. So, the primary requirement for a reaction to occur
is that:
1. The reactants must collide and
interact with each other.
2. The molecules must have sufficient
energy to initiate the reaction.
3. The molecules must have the
proper orientation.
Reaction: 2NOBr → 2NO + Br2
Reaction
⇢No Reaction

35. 35
Catalyst
Catalyst is a substance that change the rate of chemical rxn or cause the rxn
occurrence
• Participate in chemical interactions in them intermediate steps
• After reaction, catalysts usually restores and remains the similar amount
as well as chemical properties;
• ONLY CHANGE rxn rate and times, UNCHANGES the equilibrium
constant
 The roles of catalysts:
Reduce the activation energy in reaction through
mechanism changes → increase the reaction rate
 Properties:

36. 36
 2 common types
of catalysts:
1. Homogeneous Catalysts
2. Heterogeneous Catalysis
Catalysts & Reactants have same phases
Catalysts & Reactants have different phases
 Selective properties:
 Each catalyst is ONLY suitable for a specific reaction
 With the SAME REACTANT and DIFFERENT CATALYSTS, they can form
VARIOUS PRODUCTS
Examples: 𝐶2𝐻5𝑂𝐻
𝑇, 𝐶𝑢/𝑍𝑛
𝐶𝐻3𝐶𝐻𝑂 𝐶2𝐻5𝑂𝐻
𝑇, 𝐴𝑙2𝑂3
𝐶2𝐻4
Catalyst

37. 37
Catalytic Mechanism
Example: Consider reaction: A + B = AB
 Without a catalyst:
A + B → A…B → AB 𝑬𝟏
∗
, low
 In presence of catalyst K:
A + K → A…K → AK 𝑬𝟐
∗
, fast
AK + B → AK-B → AB + K 𝑬𝟑
∗
, fast
Where: 𝐸2
∗
and 𝐸3
∗
< 𝐸1
∗
, so the reaction rate increase in presence of catalyst
A + B + K → AB + K
The overall rxn: