Chemical kinetics Dr Satyabrata si

CHEMICAL KINETICS
Dr Satyabrata Si
KIIT Deemed to be University
Email: satyabrata.si@gmail.com
IMTH-103 CHEMISTRY-I

• Rate of reaction
Factors influencing the rate of reaction:
Concentration
Temperature
Pressure
Solvent
Light
Catalyst
• Kinetics of various reactions
Zero order reactions
1st order reactions
2nd order reactions
• Theories of reaction rates
Collision
Absolute reaction rate theory
• Catalysis
Homogeneous catalysis
Heterogeneous catalysis
Factors affecting catalytic processes
Syllabus

Chemical Kinetics
 Studies the rate of a chemical reactions & rate laws
a) Shows time needed for a given amount of the product
b) Shows amount of product in a given amount of time
c) Shows how to control the reaction
d) Guide towards the mechanism of a reaction
 Factors affecting the rate of a reaction
(Temperature, pressure, concentration and catalyst)
 Studies the mechanism of a chemical reaction
a) Predict products of similar reactions
b) Better understand the reaction
c) Accurately manipulate the reaction for a desired result
d) Organize and simplify the study of chemistry
Chemical kinetics is defined as the branch of chemistry which deals with the
study of the rate of chemical reactions and their mechanism.

Rate of Reaction
A B
“The rate of reactions is defined as the change in concentration of
any of reactant or products per unit time”

Rate of Reaction
takenTime
producedBofAmount
(r)reactionof
takenTime
consumedAofAmount
(r)reactionof


Rate
Rate
dt
dx
(r)reactionof Rate

Rate of Reaction (r) = Rate of disappearance of A
= =
= Rate of appearance of B
= =
dt
A][d-
dt
Bd ][
Rate of Reaction
A B
Unit of Rate
Concentration/Time (Mole/litre)/sec mol l-1 s-1

dt
B][

dt
A][-

A + B C + D
Rate of Reaction
Rate of Reaction (r) = Rate of disappearance of A =
= Rate of disappearance of B =
= Rate of appearance of C =
= Rate of appearance of D =
dt
A][d-
dt
Cd ][
dt
B][d-
dt
Dd ][

nA + mB pC + qD
Rate of Reaction & Stoichiometry
Rate of Reaction (r) = Rate of disappearance of A =
= Rate of disappearance of B =
dt
A
n
][d1
dt
Cd
p
][1
dt
B
m
][d1
dt
Dd
q
][1

Determination of Rate
N2 + 3H2 2NH3
Zn + H2SO4 ZnSO4 + H2
H2 + I2 2HI
dt
]d[H
dt
]d[ZnSO
dt
]SOd[H
-
dt
d[Zn]
- 2442
rate
dt
]d[NH
2
1
dt
]d[H
3
1
-
dt
]d[N
- 322
rate
dt
d[HI]
2
1
dt
]d[I
-
dt
]d[H
- 22
rate

Average Rate & Instantaneous Rate
NO(g)CO)(NO)( 22  ggCO
t
[CO]-
rateousInstantane t
d
d

t
[CO]-
t
[CO]-
rateAverage
d
d




At any time the instantaneous rate is equal to the slope of a straight line
drawn tangent to the curve at that time
[CO] (mol/liter) 0.100 0.067 0.050 0.040 0.033
Time (sec) 0 10 20 30 40

Factors influencing the rate of reaction
• Nature of reactants & products
e.g., solution vs gas phase
• Concentration
Usually increase as reactant increases.
• Temperature
Usually faster at higher temperature
• Pressure
Increases as pressure increases
• Solvent
• Light
• Catalyst
Speed chemical reactions

Law Mass Action & Rate of Reaction
Law of mass action, first proposed by Guldberg & Waage
in 1867 and states that the rate at which a substance
reacts is proportional to its active mass, i.e., molar
concentration and the rate of a chemical reaction is
directly proportional to the product of the active masses or
molar concentrations of the reactants.
nA + mB Product
Rate = k [A]n [B]m
Rate  [A]n [B]m
K is the rate constant
(Velocity constant, Velocity co-efficient or Specific reaction rate)

Rate Laws/Rate Equation
It is an expression showing the relationship between the
reaction rate and the concentrations of reactants
nA + mB Product
Rate = k [A]n [B]m
Rate  [A]n [B]m
K is the rate constant
(Velocity constant, Velocity co-efficient or Specific reaction rate)

Rate Constant
A + B Product
Rate = k [A] [B]
Rate  [A] [B]
If [A] = [B] = 1, then
Rate = k  1  1 = k
Thus, rate constant of a reaction may be defind as the rate of
reaction when the concentration of each of the reactants is
unity at a given temperature
Characteristics of k
• Different value for different reactions.
• A measure of rate of reaction.
• Independent of reactant concentration.
• Varies with change in temperature.

More Examples
]H[k[NO]RateO2HN2HNO2
]k[NORateO2NO2NO
]][IHk[Rate2HIIH
]Ok[NRateO4NOO2N
]k[CO][NORateNOCONOCO
2
2
222
2
222
2222
522252
222






Order of a Reaction
The sum of the powers of concentrations in the rate law
nA + mB Product
Rate = k [A]m [B]n
Rate  [A]m [B]n
Order of the reaction = m + n
m + n = 0, a zero order reaction
m + n = 1, a first order reaction
m + n = 2, a second order reaction
m + n = 3, a third order reaction

Molecularity of a Reaction
The number of reactant molecules involved in a reaction
A ProductUnimolecular reactions
Bimolecular reactions A + B Product
Termolecular reactions A + B + C Product
The decomposition of N2O5
N2O5 NO2 + NO3
N2O5 + NO3 3NO2 + O2
2N2O5 4NO2 + O2
(Slow)
(Fast)

Order of a reaction Molecularity of a reaction
Sum of the power of the
concentration terms in the rate
low.
Number of reacting species
involved in a simple reaction.
Experimentally determined. A theoretical concept
Can have fractional value. Always a whole number
Can have zero value. Do not have zero value
Can be changed with reaction
conditions.
Can not be changed with
reaction conditions.
For a complex reaction, the
slowest step gives the order of
the reaction
For a complex reaction, each
step has its own molecularity.

Zero Order Reaction
The rate is independent of reactant concentrations.
A Product
Initial conc. a 0
Final conc. a-x x
0
][
][d-
Ak
dt
A
Rate 
kxak
dt
xa
dt
dx


 0
)(
)(d-
x = kt or
  kdtdx
t
x
k 
Unit of k = concentration per unit time

Examples of Zero Order Reaction
Photochemical reactions:
Heterogeneous reactions:
2HCl(g))(Cl)( 22   Sunlight
ggH
22
Pt
2
22
Pt
3
22
Au
2
1
N
3HN2
IH2
OON
NH
HI




First Order Reaction
Rate is determined by the change of only one concentration term
A Product
Initial conc. a 0
Final conc. a-x x
][
][d-
Ak
dt
A
Rate 
)(
)(d-
xak
dt
xa
dt
dx



kdt
xa


dx

-ln (a-x) = kt + I, the constant of integration
 

kdt
xa
dx
Unit of k = (time)-1
If t = 0 and x = 0
I = -ln a
kt
xa
a


ln
First Order Reaction
xa
a
t
k

 ln
1
xa
a
t
k

 log
303.2

Examples of First Order Reaction
Decomposition of N2O5 in CCl4 solution:
tVV
V
k
OON





10
22
1
252
log
t
2.303
2NO
Problem:
From the following data for the decomposition of N2O5 in CCl4
solution at 48 C, show that it is a first order reaction.
t (mins) 10 15 20 25 
VCO2
6.30 8.95 11.40 13.5 34.75

Decomposition of H2O2 in aqueous solution:
xa
a
OH



10
2
Pt
22
log
t
2.303
k
OOH
Problem:
The catalysed decomposition of H2O2 in aqueous solution is
followed by titrating equal volume of sample solutions with
KMnO4 solution at different time interval give the following
results. Show that the reaction is a first order reaction.
t (mins) 0 5 15 25 45
VKMnO4
37 29.8 19.6 12.3 5.0

Acid Hydrolysis of a Ester:
Vt-V
V-V
log
t
2.303
k
OHHCH
0
10
523
Acid
2523



 COOHCHOHCOOCCH
Problem:
The following data was obtained on hydrolysis of methyl
acetate at 25 C in 0.2N HCl. Show that it is a first order
reaction.
t (mins) 0 75 119 180 
Valkali 19.24 24.20 26.60 29.32 42.03

Inversion of Cane sugar:






rr
rr
log
t
303.2
k
OHCOHCOHOHC
t
0
10
61266126
Acid
2112212
Problem:
The optical rotation of sucrose in 0.9N HCl at different time
interval is given in the following table. Show that it is first order
reaction.
t (mins) 0 7.18 18 27.05 
Rotation +24.09 +21.4 +17.7 +15 -10.74

Pseudo-Order Reaction
The experimental order which is not the actual one observed.
A + B (excess) Product
][
]][[
'
AkRate
BAkRate


][, '
BkkWhere 
Acid Hydrolysis of a Ester:
Inversion of Cane sugar:
OHHCCOOHCHOHHCOOCCH 523
Acid
2523 
61266126
Acid
2112212 OHCOHCH  OOHC

Second Order Reaction
Rate is determined by the change of two concentration term
2A Product
Initial conc. a 0
Final conc. a-x x
2
][
][d-
Ak
dt
A
Rate 
dt
dx
= = k(a-x)2
dt
xa ][d- 
= kdt2
)(
dx
xa 

= kt + I
xa 
1
= kdt  2
)( xa
dx
If t = 0 and x = 0
I = 1/a = kt +
xa 
1
a
1
)(
1
xaa
x
t
k


Unit of k = (conc.)-1 (time)-1

A + B Product
Initial conc. a b 0
Final conc. a-x b-x x
kdt
xbxa
dx
xbxak
dt
dx
BAk
dt
dx
dt
Bd
dt
Ad





))((
))((
]][[
][][
 )
)(
log
)(
303.2
,
xba
xab
bat
k
gIntegratin





A + B Product
Initial conc. a a 0
Final conc. a-x a-x x
kdt
xa
dx
xakxaxak
dt
dx
BAk
dt
dx
dt
Bd
dt
Ad





2
2
)(
)())((
]][[
][][
 xaa
x
t
k
gIntegratin

 log
303.2
,

Examples of Second Order Reaction
Alkali Hydrolysis of Ester:
)(t
1
k
OHHCNa 523523
xaa
x
COONaCHOHHCOOCCH



Problem:
A gram mole of ethyl acetate was hydrolysed with a gram
mole of NaOH and was studied by titrating 25ml of the
reaction mixture at different time interval against a standard
acid. Show that the reaction is of the second order.
t (mins) 0 4 6 10 15 20
Vacid 8.04 5.3 4.58 3.5 2.74 2.22

Examples of Second Order Reaction
Thermal decompostion of Acetaldehyde:
)2(t
1
k
2CO2CH2
COCH
00
0
4
C520
3
4
C520
3
t
t
PPP
PP
CHOCH
CHOCH



 
 


Problem:
The thermal decompostion of acetaldehyde was studied at
518 C. Starting with the initial pressure of 363 mm of Hg, the
following results were obtained at different time interval. Show
that the reaction is of second order.
t (sec) 42 73 105 190
p (mm) 34 54 74 114
)(t
1
k
xaa
x



Half-Life of a Reaction (t1/2 or t0.5)
Half-life is another expression for reaction rate and is defined
as the time required for the concentration of a reactant to
decrease to half its initial value

Half-Life of a Zero-Order Reaction
As, [A] at t1/2 is one-half of the original [A],
[A]t = 0.5 [A]0.
k
A
t
k
AA
t
k
AA
k
x
t
t
x
k
t
2
][
][5.0][
][][
0
5.0
0
00
5.0
0







[A]t = 0.5 [A]0.
Half-Life of a First-Order Reaction
NOTE: For a first-order process, the half-life does not depend on [A]0 and
is inversely proportional to k .
k
t
kk
t
t
k
A
A
t
k
A
A
t
k
xa
a
t
k
693.0
3010.0
303.2
2log
303.2
2log
303.2
][5.0
][
log
303.2
][
][
log
303.2
log
303.2
2/1
2/1
2/1
0
0
2/1
0








Half-Life of a Second-Order Reaction
0
2/1
0000
2/1
00
0
][
1
][
1
][
2
][
1
][5.0
1
][
1
][
1
][][
][][
)(
)(
1
Ak
t
AAAA
kt
AAAA
AA
xaa
x
kt
xaa
x
t
k









[A]t = 0.5 [A]0.
NOTE: For a second-order process, the half-life is inversely
proportional to both k and [A]0.

Order Rate Law Integrated Rate Law Half-Life Straight line Plot
0 r = k[A]0
1 r = k[A]1
2 r = k[A]2
n
k[A]rate
Product

nA
aktx)(a
]A[kt]A[ 0


alnktx)ln(a
][ln][ln 0

 AktA
a
1
kt
x)(a
1
][
1
][
1
0



A
kt
A
2k
a
2k
[A]
5.0
0
5.0


t
t
k
0.693
5.0 t
ka
1
k[A]
1
5.0
0
5.0


t
t
tvsx)(a
s][

tvA
tvsx)ln(a
s][ln

tvA
tvs
x)(a
1
s
][
1

tv
A
Summarising the all……….
At Time = 0, Concentration of reactant = [A]0 or a
At Time = t, Concentration of reactant = [A] or (a-x)

Determination of the Order of a Reaction
 Using integrated rate expression
• A hit-and-trial method
• Calculating the constant k
 Using half-life period
 Graphical method
Linear Fitting of rate expression:
• Zero-order, [A] vs t
• 1st-order, ln[A] vs t
• 2nd-order, 1/[A] vs t













2
1
1
2
log
log
1
A
A
t
t
norder
A
t n
 
nth,
][
1
1

0th Order n=0
Rate = k[A]0
[A] = - kt + [A]o
1st Order n=1
Rate = k[A]1
ln[A] = - kt + ln[A]o
2th Order n=2
Rate = k[A]2
1/[A] = kt + 1/[A]o
[A]
Time, t
slope = k
n
k[A]rate
Product

nA
ln[A]
Time, t
slope = k
1/[A]
Time, t
slope = k

Determination of the Order of a Reaction
 Van’t Hoff’s Differential Method
 Ost wald’s Isolation Method
 
 21
21
21
21
2
2
1
1
2
2
1
1
loglog
loglog
loglogloglog
loglogloglogloglog
CC
dt
dC
dt
dC
n
CCn
dt
dC
dt
dC
Cnk
dt
dC
Cnk
dt
dC
kC
dt
dC
kC
dt
dC
kC
dt
dC
nn
n











































CBA nnnn
oductCBA

 Pr

Problems
Q1. In the reduction of nitric oxide, 50% of reaction was completed in 108 sec
when initial pressure was 336 mm Hg and in 147 sec initial pressure was 288 mm
Hg. Find the order of the reaction.
Q2. In the thermal decomposition of a gaseous substance, the time taken for the
decomposition of half of the reactant was 105 min when the the initial pressure was
750 mm and 950 min when the initial pressure was 250mm. Find the order of the
reaction.
Q3. The half-life for the thermal decompostion of phosphine at three different
pressures are given below: Find the order of the reaction.
P (mm Hg) 707 79 3.5
t0.5 84 84 84
Q4. For the reaction between gaseous chlorine and nitric oxide, it is found that
doubling the concentration of both reactants increases the rate eight times, but
doubling the chlorine concentrationalone doubles the rate. What is the order of
reaction with respect to nitric oxide and chlorine?
2NO + Cl2 = 2NOCl
O2HN2HNO2 222 

Problems
Q5. The half-life of a chemical reaction at a particular concentration is 50 min.
When the concentration is doubled, the half-life become 100 mins. Find out the
order of the reaction.
When the concentration is doubled, the half-life remains 50 mins. Find out the
When the concentration is doubled, the half-life become 25 mins. Find out the
Q8. Compound A decomposes to form B and C is a first order reaction . At 25 C
the rate constant for the reaction is 0.45 s-1 . What is the half-life of A at this
temperature.
Q9. The half-life of a susbstance in a first order reaction is15 minutes. Calculate
the rate constant.
Q10. For a certain first order reaction, half-life is 100 sec. How long will it take for
the reaction to be completed 75% ?
Q11. 50% of a first order reaction is completed in 23 min. Calculate the time
required to complete 90% of the reaction.

Temperature Effect on Reaction Rate
• A measure of the average kinetic energy of the molecules in a sample.
• At any temperature there is a wide distribution of kinetic energies.
• At higher temperatures, a larger population of molecules has higher energy.
• As the temperature increases, the fraction of molecules that can overcome the
activation energy barrier increases.
• As a result, the reaction rate increases.

The ratio of rate constants of a reaction at two different
temperatures differing by 10 degree is known as
temperature coefficient. The value of temperature coefficient
is generally 2 to 3.
Temperature Effect on Reaction Rate
Temperature Coefficient = 32
298
308
25
35
to
k
k
k
k



 Increases the rate of a reaction
 Initiate a reaction
Temperature Coefficient:

Arrhenius Equation
In 1889 Arrhenius developed a mathematical relationship
between k, T and Ea, which is known as Arrhenius equation.
where, A is an experimentally determined quantity
Ea is the activation energy
R is the gas constant
T is temperature in Kelvin
Taking the natural logarithm of both sides,
y = mx + b k is determined experimentally
at several temperatures.

Arrhenius Equation
Calculation of Ea:
Ea can be calculated from the slope of a plot of (ln k) vs (1/T).
A
RT
E
k
A
RT
E
k
a
a
log
303.2
log
lnln

 If k1 and k2 corresponds to T1 and T2,





 

21
12
1
2
303.2
log
TT
TT
R
E
k
k a
R
Ek
Slope a




T
1
ln

Theories of Reaction Rates
Collision Theory
Absolute Reaction Rate Theory
Or
(Transition State Theory)

Collision Theory
According to this theory, a chemical reaction takes place only by
collisions between the reacting molecules.
Energy barrier
 The molecules must collide with sufficient kinetic energy.
Ea = Activation Energy

 The molecules must collide with correct orientation.
Collision Theory
Thus, only molecules colloid with kinetic energy
greater than activation energy and with correct
orientation can cause reaction.

Collision Theory & Reaction Rates:
A + B Product
zpfRate 
Where, f = fraction molecules having sufficient kinetic energy
p = probable fraction of molecules with effective orientation
z = collision frequency

Limitation of the Collision Theory
• Applicable to simple gaseous reaction
• Good for simple bimolecular reaction
• Only kinetic energy is considered
• Silent for bond cleavage and bond formation
• No method to calculate Probability factor or Steric factor

Absolute Reaction Rate Theory
The transition state represents the point of
highest free energy for a reaction step.
Developed by Henry Erying in 1935 and postulated that during collision, the
reactant molecules form a transition state or activated complex which
decomposes to give the products.
In summary,
• Kinetic energy is converted into potential energy
• Interpenetration of electron clouds
• Rearrangement of valence electrons
• Formation of an activated complex or transition state
C-BAAB  CBCA

C-BAAB  CBCA
Reaction coordinate
Potentialenergy
CA B 
CB A
C-BA 
Ea
E
Activated Complex or
Transition state
Reactant
Product
Reaction Energy Diagram for an Exothermic Reaction

Reaction coordinate
Potentialenergy
CA B 
CB A
C-BA 
Ea
E
Transition state
Reactant
Product
C-BAAB  CBCA
Reaction Energy Diagram for an Endothermic Reaction

C-BAAB  CBCA
Reaction coordinate
Potentialenergy
CA B 
CB A
C-BA 
Ea
E
Transition state
Reactant
Product
Ecat
Activation Energy and Catalysis

Some Problems
Q1. The value of rate constant for the decomposition of N2O5 were
determined at several temperatures. A plot of ln k vs 1/T gave a
straight line of which the slope was found to be -1.2  104 K. Calculate
the activation energy of the reaction.
Q2. The value of rate constant for the decomposition of ethyl iodide is
1.6  10-5 s-1 at 327 ºC and 6.36  10-3 s-1 at 427 ºC. Calculate the
activation energy for the reaction. (R = 8.314 JK-1 mol-1)
Q3. The rate of a particular reaction quadruples when the temperature
changes from 293K to 313K. Calculate the activation energy for the
same reaction.
Q4. The activation energy of a non-catalyzed reaction at 37 ºC is 200
kcal mol-1 and the activation energy of the same reaction when
catalyzed by an enzyme is 6 kcal mol-1. Calculate the ratio of the rate
constant of the catalysed and the non-catalysed reaction. Assume
frequency factor to be same in both cases. (R = 1.987 cal).

Consecutive Reactions
CBA 21 kk

At t = 0, [A]0 0 0
At t = t, [A] [B] [C]
[B]k
dt
d[C]
[B]k-[A]k
dt
d[B]
[A]k
dt
d[A]-
[C][B][A]]A[
2
21
1
0




[A]
[B]
[C]

Parallel or Side Reactions
]A[k
dt
]A[d
]A)[kk(
dt
]A[d
]A[k]A[krr
dt
]A[d
]A[kr
]A[kr
'
21
2121
22
11








2
1
2
1
2
1
2
1
2
1
22
11
k
k
r
r
k
k
]A[k
]A[k
r
r
]A[kr
]A[kr





Riversible or Opposing Reactions
A B
kf
kb
At t = 0, [A]0 0
At t = t, [A] [B]
eq
eq0
fb
eq0feqb
eqbeq0f
eq
b0f
bf
bf
x
)x]A([
kk
)x]A([kxk
xk)x]A([k0
xx&0
dt
dx
,mequliberiuAt
xk)x]A([k
dt
dx
]B[k]A[k
dt
dx
]B[k]A[k
dt
dx
dt
]B[d
dt
]A[d










Chemical Kinetics & Catalysis
A catalyst is a substance which alters the rate of a chemical reaction,
itself remaining chemically unchanged at the end of the reaction.
The process is called catalysis.

Positive Catalyst Negative Catalyst
Homogeneous Catalysis Heterogeneous Catalystis
Enzyme Catalysis
Catalyst
Catalysis

Homogeneous Catalysis
In homogeneous catalysis, the catalyst is in the same phase
as the reactants and is evenly distributed throughout.
22
-
22
-
523
-
2523
4261266126422112212
2423
22
322
OO2H][IO2H
]/OH[HOHHCCOOHCH]/OH[HOHHCOOCCH
]SO[HOHCOHC]SO[HOHOHC
][ICOCH][ICHOCH
[NO]2CO[NO]O2CO
[NO]2SO[NO]OSO2







Catalysis in Gas Phase
Catalysis in Solution Phase

Heterogeneous Catalyst
In heterogeneous catalysis, the catalyst is in different phase
from the reactants. Also known as Contact Catalysis.
][MnO3O2KCl][MnOKClO
][AlClHClCOCHHC][AlClCOClCHHC
[Pt]OO2H[Pt]O2H
[Ni]CH-CH[Ni]HCHCH
[Fe]2NH[Fe]3HN
[Pt]2SO[Pt]OSO2
2223
33563366
2222
33222
322
322






Some Features:
Surface area Activation energy
Promoter action Catalytic poisons
Catalysis with Gaseous Reactants
Catalysis with Liquid Reactants
Catalysis with Solid Reactants

Characteristic of Catalytic Reactions
 Remains unchanged in mass and chemical composition
 A small quantity is generally needed
 More effective when finely divided
 It is very specific
 In general, it cannot initiate a reaction
 It does not affect the equilibrium position
 Depends on temperature
23
Cu
52
222
OAl
52
HCHOCHOHHC
OHCHCHOHHC 32

 
OH2OH
reactionNoOH
2
blackPt
22
RT
22
 


Catalytic Promoter Catalytic Poison
A substance which, though itself
not a catalyst, promotes the activity
of a catalyst
A substance which, destroy the
activity of a catalyst to accelerate
a reaction
3
Fe
22 2NH3H N
Promoter: Mo or Al2O3
• Change of lattice spacing
• Increase of peaks and cracks
3
Fe
22 2NH3H N
Poison: H2S
• Preferential adsorption
• Chemical reaction
OHCH2HCO 3
OCr&ZnO
2
32
 
3
Pt
22 2SOOSO 
Poison: As2O3

Autocatalyst
One of the product itself acts as a catalyst for the same reaction
]/OH[HOHHCCOOHCH]/OH[HOHHCOOCCH -
523
-
2523


Negative Catalyst/Inhibitor
A substance, which reduces the rate of a reaction
22223 2ClO2H4COCl3O4 CHCl Inhibitor: 2% ethanol
22424424224 10COO8HSOK2MnSOSOHOC5H2 KMnO
23 3H2As2 AsH
2222 OO2H2 OH Inhibitor: Glycerol
fuelofCombustion Inhibitor: Pb(C2H5)4

Reaction coordinate
Potentialenergy
CA B 
CB A
C-BA 
Ea
E
Transition state
Reactant
Product
Ecat
Activation Energy and Catalysis

Theories of Catalysis
1. Intermediate Compound Formation Theory
CABBAC
ACCA
ABBA C



HClAlClCHHC]AlCl[][CHHC
)teIntermedia(]AlCl[][CHAlClClCH
HClCHHCClCHHC
NOSOSONO
ate)(Intermedi2NOO2NO
2SOOSO2
33564366
4333
356
AlCl
366
322
22
3
NO
22
3


 






2. Adsorption Theory
This theory explains the mechanism of a reaction catalyzed by a solid catalyst.
The catalyst function by adsorption of the reacting molecules on its surface
Step I: Adsorption of reactant molecules.
Step II: Formation of activated complex.
Step III: Decomposition of activated complex.
Step IV: Desorption of products.

2. Adsorption Theory
33
Ni
222 CHCHHCHCH 
Step I: Adsorption of Hydrogen molecules.
Step II: Broken of H-H bond.
Step III: Formation of activated complex.
Step IV: Decompostion of the activated
complex and Desorption of
ethane molecules.

Active Centers on Catalyst surface

Adsorption Theory & Catalytic Activity
 Metals in a state of fine subdivision or colloidal
form are rich in free valence bonds and hence they
are more efficient catalysts than the metal in lumps
 Catalytic poisoning occurs because the so-called
poison blocks the free valence bonds on its
surface by preferential adsorption or by chemical
combination.
 A promoter increases the valence bonds on the
catalyst surface by changing the crystal lattice and
thereby increasing the active centers.

Acid-Base Catalysis
Homogeneous catalytic reaction catalyzed by acids or bases,
or both acids and bases are known as acid-base catalysts.
OHONNONH
NitramideofionDecomposit
OHHCCOOHCHOHHCOOCCH
EsterofHydrolysis
OHCOHCOHOHC
sugarCaneofInversion
22
Base
22
523
Acid/Base
2523
61266126
Acid
2112212

 


Enzyme Catalysis
• Enzymes are catalysts in biological systems.
• The substrate fits into the active site of the
enzyme much like a key fits into a lock.
The catalysis brought about by enzymes is known as
enzyme catalysis
E + S ES P + E

Characteristics of Enzyme Catalysis
• Most efficient catalysts known
• Marked by absolute specificity
• Maximum at optimum temperature
• Maximum at optimum pH
• Greatly affected by inhibitors
• Greatly enhanced by Activator or Coenzyme
Inversion of cane sugar by Invertase present in yeast
Conversion of glucose into ethanol by Zymase present in yeast
Hydrolysis of urea by Urease present in soya bean
61266126
Invertase
2112212 OHCOHCOH  OHC
252
Zymase
6126 2COOHH2C  OHC
23
Urease
22 CO2NHH   
NHCONH

Chemical kinetics Dr Satyabrata si

Chemical kinetics Dr Satyabrata si

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