2. Content
• Rate of reaction
• Factors which affect rate of reaction
• Rate law & Order of reaction
• Mechanism of reaction
3. CHEMICAL KINETICS
• Rate/speed of a chemical reaction ---- how chemical
process occurs.
• We study chemical kinetics, in point of industrial prospect---
-- determined rate of reaction by practically
• Besides information about the speed at which reactions
occur, kinetics also sheds light on the reaction mechanism
(exactly how the reaction occurs).Reactant Reactant Product
4. Rate of Reaction (ROR)
Rate = Speed -- Reaction speed (how fast does a chemical
reaction proceed)
• Change in concentration (in aqueous phase) or pressure (in
gaseous phase) of reactant/product with respect to time --
called Rate of Reaction
FACTORS AFFECT THE ROR:
• Temperature
• Reactant concentration
• Pressure of gaseous reactants or products
5. Rate of reaction (ROR)
For the reaction A --- B there are two ways of measuring
rate:
The speed at which the reactants disappear
The speed at which the products appear
i. Very Fast Reaction: Those chemical reaction which
completed within fraction of time. (like Ionic reactions)
NaCl + AgNO3 -- AgCl↓ + NaNO3
White ppt
Reaction proceed within microseconds. So, ROR cannot be
determined.
ii. Very Slow Reaction: Those chemical reaction which
completed in very long time duration.
Eg. Rusting of Iron
6. Rate of reaction (ROR)
For the reaction A --- B there are two ways of measuring
rate:
The speed at which the reactants disappear
The speed at which the products appear
i. Very Fast Reaction: Those chemical reaction which
completed within fraction of time. (like Ionic reactions)
NaCl + AgNO3 -- AgCl↓ + NaNO3
White ppt
Reaction proceed within microseconds. So, ROR cannot be
determined.
ii. Very Slow Reaction: Those chemical reaction which
completed in very long time duration.
Eg. Rusting of Iron
7. Rate of reaction (ROR)
iii. Moderate Reaction: Those chemical reaction which
completed in finite time.
Eg. All Molecular reactions
N2 + 3H2 ------ 2NH3
Reaction takes finite time to proceed. So, ROR can be
determined.ROR Measure
Reactant disappear Product appear
8. Rate of Reaction (ROR):
How to calculate Rate/ speed of any chemical reactions
by:
• Average rate of reaction
• Instantaneous rate of reaction
9. 2A + 3B --------- 4C
If, Initial time => t = 0; Conc. a = a0 Conc. = 0
Final time => t = t; Conc. a = at (aFinal - aInitial) Conc. = X
Average ROR =
– ½ [∆A]/∆t = – ⅓ [∆B]/∆t = + ¼ [∆C]/∆t
Average ROR = Rate of change in concentration
Rate of change in time
Average Rate of Reaction
(ROR):
10. 2X + Y --- 3Z
Average ROR =
– ½ [∆X]/∆t = – [∆Y]/∆t = + ⅓ [∆Z]/∆t
Rate of appearance (ROA): Product
ROA of Z = [∆Z]
∆ t
Rate of disappearance (ROD): Reactant
ROD of X = [∆X] ROD of Y = [∆Y]
∆ t ∆ t
Average Rate of Reaction
(ROR):
11. Example: N2 + 3H2 ------ 2NH3
Average ROR =
– [∆N2]/∆t = – ⅓ [∆H2]/∆t = + ½ [∆NH3]/∆t
Rate of appearance (ROA): Product
ROA of NH3 = [∆NH3]
∆ t
Rate of disappearance (ROD): Reactant
ROD of N2 = [∆N2] ROD of H2 = [∆H2]
∆ t ∆ t
Average Rate of Reaction
(ROR):
12. Unit of ROR = Rate of change in
concentration
Rate of change in
time
= Mol/Lit/sec => Mol L-1
s-1
In gaseous reaction
-1
Unit of ROR
13. Average ROR =
– ½ [∆N2O5]/∆t = + ¼ [∆NO2]/∆t = + [∆ O2]/∆t
Rate of appearance = [∆NO2] = 6Mol/L/s
∆ t
ROR = + ¼ [∆NO2]/∆t = + ¼ ˟ 6 = 1.5 Mol/L/sec
ROR = – ½ [∆N2O5]/∆t
1.5 = – ½ [∆N2O5]/∆t
3 Mol/L/sec = [∆ N2O5]
∆ t
Q. 2N2O5 ----------- 4NO2 + O2
If ROA of NO2 is 6Mol/L/sec. then,
find (i) ROR, and (ii) ROD of N2O5.
22. Factors of Rate of
Reaction (ROR)
Factors are affect the Rate of chemical reaction (how
fast does a reaction proceed)
FACTORS AFFECT THE ROR:
• Temperature
• Reactant concentration
• Pressure of gaseous reactants or products
• Action of catalyst
• Nature of reactant
23. i). Nature of Reactant
i. Very Fast Reaction: (like Ionic reactants)
NaCl + AgNO3 ---- AgCl↓ + NaNO3
Reaction proceed within microseconds. So, ROR cannot be
determined.
ROR ---- High
ii. Moderate Reaction: (Covalent reactants)
N2 + 3H2 ---- 2NH3
Reaction takes finite time to proceed. So, ROR can be
determined.
24. Rate of Reaction (ROR)
Homogenous reactants > Heterogenous reactants
Chemical nature α ROR
ROR α Stability of Product α 1
Stability of Reactant
i). Nature of Reactant
25. Chemical nature α ROR
ROR α Stability of Product α 1
Stability of Reactant
CH3COOH ----- CH3COO- + H+
HCOOH ----- HCOO- + H+
H-COO- > CH3-COO-
ROR more stable ROR less stable
i). Nature of Reactant
26. Rate of Reaction (ROR)
Gas > Liquid > Solid
Pressure of gaseous reactants or products Due to increase number of
collision of molecules
Chemical nature α ROR
ii). Physical state of
Reactant
27. Rate of Reaction (ROR) ----- (mostly in
Solid state)
Surface area of a solid reactant More area for
reactants to be in contact
ROR α Surface area of reactant
Eg: Sugar powder/solute (small molecules)
easily dissolve in milk/solvent/water than
sugar cubes. [Due to more surface area of
molecules].
iii). Surface area of
Reactant
28. Rate of Reaction (ROR) --(only in Photochemical/Photosensitive
reaction)
Chemical reactions, that occur on exposure to visible radiation are called
Photochemical reactions.
Some reactions are Photosensitive (light sensitive) ----
Photochemical reaction
H2 (g) + Cl2 (g) ------- 2HCl (g)
iv). Intensity of light
hʋ
29. The rate of a photochemical reactions is affected by the intensity of light
Temperature has little effect on photochemical reactions.
Quantum yield or quantum efficiency of a photochemical reactions;
hʋ = number of reactant molecules in a given time / number of photon
(quanta) of light absorbed ill the same time
ROR α Intensity of light α Number of Photons
iv). Intensity of light
30. Rate of Reaction (ROR)
Sets increase/decrease alternate path from reactant to products with lower
activation energy
2SO2 + O2 ------- 2SO3
ROR α Catalyst
(Concentration of NO)
v). Catalyst
NO
31. Rate of Reaction (ROR)
Generally, ROR increase on increasing temperature
On 10ᵒ C rise in temperature, ROR increases 2-3 times (Temperature
coefficient).
ROR α Temperature
Arrhenius equation
vi). Temperature
32. Rate of Reaction (ROR)
At higher temperature, reactant molecules have more kinetic energy,
move faster, and collide more often and with greater energy.
Collision energy: when two chemicals react, their molecules have to
collide with each other (in a particular orientation) with sufficient energy
for the reaction to take place.
Kinetic theory: Increasing temperature means the molecules move faster.
ROR α Temperature
vi). Temperature
33. Rate of Reaction (ROR)
Generally, ROR increase on increasing concentration of reactants
(reactant molecules will collide)
ROR may also decrease on increasing concentration of reactants
ROR remains unaffected
ROR may also depend upon products concentration
ROR may depend on one or all or none of reactant concentration -- Rate
Law (depend on concentration of reactant)
ROR α Concentration of Reactant
vii). Concentration
34. Rate Law
Rate law explains the dependence of :
• Rate of reaction (ROR) on concentration of reactants or
products in reaction
• This dependence is very complex as it changes with
change in concentration of reactants or products
• We only study the simplest form of dependence
All steps are determined by only Experimental method
35. 2A + B ------ 2C
Reactants Product
r = K [A]1 [B]2
Experimental method (ROR):
If, [A] ---- Double ROR ------- 2 times increase
ROR α [A]1
If, [B] ---- Double ROR ------- 4 times increase
ROR α [B]2
[⁖ ROR = r]
Rate Law of Expression
⁖ where as, K = Rate constant; [A] = Concentration of A reactant; [B] =
Concentration of B reactant; ROR = Rate of reaction; r = Rate law.
36. 2A + B ------ 2C
Reactants Product
r = K [A]2 [B]1 This method followed
in only single step
reactions.
Guldberg & Wagge method:
If, [A] ---- Double ROR ------- 4 times increase
ROR α [A]2
If, [B] ---- Double ROR ------- 2 times increase
ROR α [B]1
[⁖ ROR = r]
Law of Mass Action
(as according stoichiometric coefficient)
⁖ where as, K = Rate constant; [A] = Concentration of A reactant; [B] = Concentration
of B reactant; ROR = Rate of reaction; r = Rate law.
37. 2A + B ------ 2C
Reactants Product
r = K [A]x [B]y
Experimental method (ROR):
r α [A]x
r α [B]y
Rate Law of Expression
⁖ where as, K = Rate constant; [A] = Concentration of A reactant (Mol/Lit); [B] =
Concentration of B reactant (Mol/Lit); r = Rate law/ Rate of reaction; x = Order of
reaction with respect to A; y = Order of reaction with respect to B.
38. 2A + B ------ 2C
Reactants Product
n = x + y
Order of Reaction (denoted as “n”)
n = overall order of reaction/ Total order of reaction/ order of reaction
Order of Reaction
(Can be negative or positive or zero or fractional)
⁖ where as, n = Order of reaction; x = Order of reaction with respect to A; y =
Order of reaction with respect to B.
39. 2A + B ------ 2C
Reactants Product
Rate Law --- by experimental method (ROR):
K = Rate constant/ velocity constant/ specific reaction rate (per unit)
r α [A]x [B]y
If, concentration of all reactants or products in rate law expression is unity (1)
[A] = [B] = 1
K = Specific reaction rate (per unit)
r = K [A]x [B]y
r = K
40. Unit of Rate Constant (K)
Unit of ROR = Mol/Lit/sec => Mol L-1 s-1
ROR = K [A]x [B]y
K = ROR
[A]x [B]y
Unit of K = Mol/Lit (⁖ x + y =
n)
Sec
[Mol/Lit]x [Mol/Lit]y
= [Mol/Lit/sec] => [Mol/Lit]1-n => [Mol/Lit]1-n
41. Unit of Order of Reaction
(n) & Rate Constant (K)
Order of Reaction (n) Unit of K [Mol/Lit]1-n /sec]
Zero order of reaction [Mol/Lit] /sec
First order of reaction sec-1
Second order of reaction [Mol/Lit]-1 /sec
Half order of reaction [Mol/Lit]3/2 /sec
42. K = 7 ˟ 10-4 Lit2 Mol-2 sec-1
[Mol/Lit]1-n /sec = [Mol/Lit]-2 /sec
1 – n = -2
Third order of reaction
Q. If rate constant is 7 ˟ 10-4 Lit2 Mol-2 sec-
1 of any chemical reaction. Find order of
reaction?
43. Q. 2A + 3B ----- 4C
If order of reaction with respect to A & B is 2 & -1. Write rate law
expression and calculate order of reaction? What is the effect on rate;
when,
(i). Concentration of A is doubled alone
(ii). Concentration of B is halved alone
(iii). Concentration of A & B is doubled
(iv). Volume of container increases 3 times
r = K [A]x [B]y
r = K [A]2 [B]-1
n = x + y
n = 2 + (-1) = 1 (First order of reaction)
K = sec-1
(i). r = K [A]2 [B]-1
r’ = K [2A]2 [B]-1
r’ = K [2A]2 [B]-1 => r’ = 22 => r’ = 4r
r K [A]2 [B]-1 r
Rate increases 4 times
44. Q. 2A + 3B ----- 4C
If order of reaction with respect to A & B is 2 & -1. Write rate law
expression and calculate order of reaction? What is the effect on rate;
when,
(i). Concentration of A is doubled alone
(ii). Concentration of B is halved alone
(iii). Concentration of A & B is doubled
(iv). Volume of container increases 3 times
(ii). r = K [A]2 [B]-1
r’ = K [A]2 [B/2]-1
r’ = K [2A]2 [B/2]-1 => r’ = 1 -1 => r’ = 2
r K [A]2 [B]-1 r 2 r
Rate increases 2 times
(iii). r = K [A]2 [B]-1
r’ = K [2A]2 [2B]-1
r’ = K [2A]2 [2B]-1 => r’ = (2)2 ˟ (2)-1 => r’ = 4 = 2
r K [A]2 [B]-1 r r 2
Rate increases 2 times
45. Q. 2A + 3B ----- 4C
If order of reaction with respect to A & B is 2 & -1. Write rate law
expression and calculate order of reaction? What is the effect on rate;
when,
(i). Concentration of A is doubled alone
(ii). Concentration of B is halved alone
(iii). Concentration of A & B is doubled
(iv). Volume of container increases 3 times
(iv). r = K [A]2 [B]-1
Volume α 1/Concentration
[A] = Number of Moles
Volume of Solution (in Lit)
Volume increase 3 times --- Concentration decreases 3 times
r = K [A]2 [B]-1
r’ = K [A/3]2 [B/3]-1
r’ = K [A/3]2 [B/3]-1 => r’ = [1/3]2 [1/3]-1 => r’ = 3
r K [A]2 [B]-1 r r 9
Rate decreases 3 times
(Due to number of moles per volume
=> Molarity)
46. Note:
1. Stoichiometric coefficients having nothing to do with order
2N2O5 -------- 4NO2 + O2
By Expt.: r = K [N2O5]1
2. In Rate law, concentration terms of products may be present
O3 --------- O2
By Expt.: r = K [O3]2 [O2]-1
3. In Rate law, concentration terms of some reactant may be present
NO2 + CO -------- NO + CO2
By Expt.: r = K [NO2]2 [CO]0 => r = K
[NO2]2
47. Note:
4. In Rate law expression, concentration term of catalyst may be present.
Catalyst ---- depend on reactant/product of reaction
2SO2 + O2 -------- 4SO3
By Expt.: r = K [O2]2 [NO]1
NO
48. Simple/ Elementary single
step reaction:
At slowest (single) step, -- Rate of Reaction -- Rate determined
3A + 2B -- C
Step1: 2A + B -- P
Step2: P + A -- Q
Step3: Q + B -- C
Final reaction: 3A + 2B -- C
By Expt.: r = K [A]2 [B]1
Example:
H2 + I2 ----- 2HI
r = K [H2]1 [I2]1
Single step Reaction
Guldbery & Wagge
Method
Law of Mass Action
49. Simple/ Elementary single
step reaction:
At slowest (single) step, -- Rate of Reaction -- Rate determined
3A + 2B -- C
Step1: 2A + B -- P
Step2: P + A -- Q
Step3: Q + B -- C
Final reaction: 3A + 2B -- C
By Expt.: r = K [A]2 [B]1
Example:
H2 + I2 ----- 2HI
r = K [H2]1 [I2]1
By Expt.: It’s not single step reaction
Single step Reaction
Guldbery & Wagge
Method
Law of Mass Action
50. Arrhenius Equation
Effect of temperature on Rate of Reaction (ROR)
Rate α Temperature
Rate = K [conc]n
Generally, ROR increases on increasing temperature (Approximately their
dependency of K on T)
In most the reaction, when temperature increases 10ᵒC, ROR increases 2-3
times (Temperature coefficient).
r = K [A]x [B]y
r α K
51. Arrhenius Equation
Effect of temperature on Rate of Reaction (ROR)
Temperature coefficient
In most the reaction, when temperature increases 10ᵒC, ROR increases 2-3
times (Temperature coefficient => Ratio of two rate constant).
⁖ Standard, t = 25ᵒC and t+10=35 ᵒC
Temperature coefficient = K(t+10ᵒC)
K(tᵒC)
Temperature 10ᵒC 20ᵒC 30ᵒC 40ᵒC
Rate constant K 2K 4K 8K
ROR r 2r 4r 8r
r α K
52. Q. In a chemical reaction,
temperature coefficient is 5. if rate at
10ᵒC is x. find the rate constant at
70ᵒC?
Temper
ature
10ᵒC 20ᵒC 30ᵒC 40ᵒC 50ᵒC 60ᵒC 70ᵒC
Rate x 2x 4x 8x 16x 32x 64x
r α K
53. Arrhenius Equation:
Accurate dependency of ‘K’ on
‘T’
Arrhenius equation
[R = Universal gas constant = 8.314 joule/mole K]
Exponentially Increases
Rateconstant(K)
Temperature
(Joule/mole)
(constant)
(universal
constant)
(Kelvin)
54. Collision energy: when two chemicals react, their molecules have
to collide with each other (in a particular orientation) with
sufficient energy for the reaction to take place.
H2 + I2 ---- 2HI
For effective collision, (i) Orientation collision; and
(ii) Energy barrier
Arrhenius Equation:
Accurate dependency of ‘K’ on
‘T’
55. Arrhenius Equation:
Accurate dependency of ‘K’ on
‘T’
(i) Orientation collision:
Effective collision ----- Head to head collision
(ii) Energy barrier
Activation energy=
ER require energy to cross
the energy barrier to reach
At transition state
ER = ET - ER
Transition
state
56. K = A . e-Ea/RT
Rate α Temperature
Temperature α Energy cross molecules
(FOM)
Rate α e-Ea/RT
K α e-Ea/RT
K = e-Ea/RT
Arrhenius equation
Arrhenius Equation
57. K = A.e-Ea/RT
ln K = ln A + ln e-Ea/RT
ln K = ln A – Ea/RT logee
ln K = ln A – Ea/RT
y = mx + c
ln K = – Ea/R (1/T) + ln A
loge K = – Ea/R (1/T) + ln A (⁖loge K = 2.303 log10 K)
2.303 log10 K = 2.303 log10 A – Ea/RT
loge K = loge A – Ea/RT
Arrhenius Equation
log10 K = log10 A – Ea/2.303RT
58.
59. If, K1 ------- T1; K2 ------- T2 (Assume T2 > T1)
log10 K1 = – Ea/2.303RT1 + log10 A -------- (i)
log10 K2 = – Ea/2.303RT2 + log10 A -------- (ii)
(Assume Ea & A are temperature independent)
Sub (ii) by (i)
log10 (K2 - K1) = Ea/2.303RT2 – Ea/2.303RT1
log10 (K2/K1) = Ea/2.303R (1/T1 - 1/T2)
log10 (K2/K1) = Ea/2.303R (T2 - T1/T1T2)
Arrhenius Equation:
Ratio of two rate constant at two different
temperature
60. Arrhenius Equation:
Exception
• On increasing temperature, rate may decrease and it may not
follow Arrhenius equation.
Example: (i) Bacterial decomposition
Ti = Inversion temperature
(ii) Oxidation of NO
2NO + O2 --- 2NO
(iii) Explosion
Explosion reaction
Temperature
Rate
Temperature
Rate
Temperature
Rate
- ve
temperature
coefficient
65. Mechanism of
Reaction
On the basis of mechanism we have two types of
reactions:
(i) Simple / elementary / single type / single step
reaction
(ii) Complex / multistep reaction
Fast reaction
Slow reaction
Moderate reaction
66. Single step reaction
2A + 2B -------- 3AB
A-------A B--------B a------b
•There is no intermediate
•Activation energy:
Rate law: r = K [A] [B]
Experimentally method
Guldberg and Wagge method : Law of mass action
Single step reaction
67. 1. H2 + I2 ------ 2HI
Reactant ------ Product
A ----- PRODUCT
Molecularity = 1 (Molecules in reactant)
r = K [A]
2. A + B ------ > PRODUCT
Molecularity = 2
r = K [A] [B]
Single step reaction
68. 3. A + 2B ---- Product
Molecularity = 3
r = K [A] [B]
Maximum molecularity ===== single step reaction
• Simultaneous collision
• Proper collision
• Cross energy barrier
Single step reaction
69. Complex step reaction
Multi step reaction:
R -------- P
R - A
A- B
B - P
R - P
• There is intermediate
• Such reactions occurs in several steps, where each steps is elementary
• Molecularity of complex reaction is not define
Molecularity of each step can be defined
• Over all rate is given by slowest step of complex reaction --- rate
determining step (rds) ------ slowest step