Chemical Kinetics (M) PDF

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Chapter
H Chemical Kinetics
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M
Introduction
] Chemical Kinetics is the branch of chemistry that
deals with
+ →

a) Rate of reactions
c) Very slow reactions :The chemical reactions
b) Factors influencing the rate of reaction
which complete in very long time are called
c) Reaction mechanism very slow reactions.
] Thermodynamics tells about only the feasibility Ex : Rusting of Iron in presence of air and
of a chemical reaction where as chemical kinetics moisture
tells about the velocity of reaction.
Kinetics studies not only help us to determine + + ¡
→ ¡

the speed (or) rate of a chemical reaction but also ] It is not possible to determine the rates of very
describe the conditions by which the reaction rates fast and very slow reactions by conventional
can be altered. methods.But the rates of reactions with moderate
Ex : Thermodynamic data indicates the diamond speed can be determined.
shall convert to graphite but in reality the Reaction Rate (OR) Rate of Reaction :
conversion rate is so slow that the change is not ] The decrease in the concentration of the reactant
perceptible at all. per unit time (or) increase in the concentration of
] Based on the velocity of chemical reactions, the the product per unit time is called rate of a reaction.
reactions are classified into three types. ] The rate of a reaction measured with respect to
a) Very fast (or) instantaneous reactions : The the decrease in the concentration of the reactants .
chemical reactions which are completed within the ] The rate of the reaction measured with respect to
fraction of seconds are called as very fast reactions. the increase in the concentration of the products .
Ex : ] The rate of a reaction at any particular instant of
1) Neutralization between strong acids and strong time during the course of a reaction is the rate of
bases. change of concentration of a reactant (or) a product
at that instant of time.

+
→ +
Ex :
2) Precipitation reactions 1) A → B
+ → ↓ + −∆ ⎡⎣ ⎤⎦ ∆ ⎡⎣ ⎤⎦ − ⎡⎣ ⎤⎦ ⎡ ⎤
= =+ = =+ ⎣ ⎦
∆ ∆
3) Explosive reactions :
Explosion of T.N.T (Tri nitro toulene) 2) + →
b) Moderate reactions : The chemical reactions which
∆ ⎡⎣ ⎤⎦ ∆ ⎡ ⎤ ∆ ⎡⎣ ⎤⎦
are completed in mesurable time are called as rate = − =− ⎣ ⎦ =+
moderate reactions. ∆ ∆ ∆
Ex : (or)
1) Inversion of cane sugar
∆ ⎡⎣ ⎤⎦ ∆ ⎡⎣ ⎤⎦ ∆ ⎡ ⎤⎦

+
→
+
rate = − =− =+ ⎣

∆ ∆ ∆

3) pP + qQ → rR + sS
2) Combustion of hydrogen (or) coal [under
normal conditions].
121 Chemical Kinetics

∆ ⎡⎣ ⎤⎦ ∆ ⎡⎣ ⎤⎦ ∆ ⎡⎣ ⎤⎦ d P
rate = − =− =+ rinst = = slope
[P]
∆ ∆ ∆ dt
d [P ]
∆ ⎡⎣ ⎤⎦ dt
Concentration of products
= + [P2 ]
∆ ∆[P ]
[P1 ] ∆P P − P1
4) − ( ) + − ( ) +
+ ( ) → ∆t rav = = 2
∆t (t2 −t1 )
( ) +

t1 t2 t time
∆ ⎡ − ⎤ ∆ ⎡ − ⎤
⎢
⎣ ⎥
⎦ ⎢ ⎦⎥ Fig : Average rate of reaction
= − =− ⎣
∆ ∆
EXAMPLE-1
∆ ⎡
+ ⎤ ∆ ⎣⎡ ⎦⎤ ∆ ⎣⎡
⎦⎤
=− ⎣⎢ ⎦⎥ = = . The decomposition of N2O5 in CCl4 at 318K has
∆ ∆ ∆ been studied by monitoring the concentration
of N 2 O 5 in the solution. Initially the
5)
( ) + ( ) →
( ) concentration of N2O5 is 2.33 mol L1 and after
Rate of reaction = 184 minutes, it is reduced to 2.08 mol L1. The
reaction takes place according to the equation
∆ ⎡⎣
⎤⎦ ∆ ⎡⎣ ⎤⎦ ∆ ⎡⎣
⎤⎦
− =− = ( ) → ( ) ( ) .
∆ ∆ ∆
Calculate the average rate of this reaction in
6)
( ) →
( ) + ( ) terms of minutes? What is the rate of production
∆ ⎡⎣
⎤⎦ ∆ ⎡⎣
⎤⎦ ∆ ⎡⎣ ⎤⎦ of NO2 during the period?
Rate of reaction = − = = Sol: Average rate
∆ ∆ ∆
⎢( )−
] The rate of a reaction varies exponentially with ⎧⎪ ∆ ⎣⎡ ⎤⎦ ⎫⎪ ⎡ − ⎤
= ⎨− ⎬=− ⎥
time of the reaction. ∆ ⎢ ⎥
⎪⎩ ⎪⎭ ⎣ ⎦
] The concentration of the reactants in a reaction
varies exponentially with time. −
= × −
] No reaction takes place with uniform rate
throughout the course of the reaction. It may be understand that
] The rates of chemical reactions differ from one
another, since the number and the nature of the ∆ ⎡⎣ ⎤⎦
−
=
∆ = × −
bonds are different in the different substances
(reactants or products or both) ∆ ⎡⎣ ⎤⎦
= × × − − −
] The units for the rate of the reaction is mol.
∆
lit1 . sec1
] when the concentration of gases is expressed in = × − − −
terms of their partial pressures, then the units of
the rate will be atm s1 EXAMPLE-2
[R ]0
N2 + 3H2 ® 2NH3 the rate of disappearance of
nitrogen is 0.02 mol L1s1. What is the rate of
−∆ R appearance of ammonia?
Concentration of reactants
rav =
[R1 ] (t2 −t1 )
∆[R ]
d [R ] − ⎡⎣ ⎤⎦ ⎣⎡
⎦⎤ ⎡⎣
⎤⎦
[R2 ] Sol: = ; = 0.02
∆t
−d R
rinst = =−slope
dt
⎣⎡
⎤⎦
− −
dt
=
t

t1 t2 time
Fig : Instantaneous rate of reaction The rate of appearance of ammonia = − −
Chemical Kinetics 122

EXAMPLE-3 H H
| |
A2B is an ideal gas, which decomposes according H − C − H + 2 O = O → O = C = O + 2 O− H
|
to the equation → . At start, the H
(4) (4) (4) (4)
initial pressure is 100mm of Hg and after 5
] Reaction (a) is faster than reaction (b)
minutes, the pressure is 120mm of Hg. What is
the average rate of decomposition of A 2 B ? b) Concentration of the reactants : Except zero order
Assume T and V are constant. reactions, for all other reactions the rate depends
Sol: The decomposition reaction of gaseous A2B is on the concentration of the reactants.
] rate α (concentration of the reactants) n or
given as → +

100 0 0 initial reaction − ∝ (or) − =

100 2x 2x x final reaction
100 2x + 2x + x = 100 + x = 120mm n may be any simple value including zero.
x = 20 mm or 2x = 40 mm ] For gaseous reactants.
The decrease in pressure of reactant rate ∝ (pressure of the reactants)n
substance A2B in 5mm is 40mm. −
The rate of decomposition of − ∝ =

A2B = = − = − ] Chemical reactions occur due to the collisions
among the reacting molecules. Hence greater the
Factors affecting the Rate of reaction : number of these molecules in unit volume, greater
a) Nature of the reactants : Reactions between ionic will be the possibility of their collisions and higher
substances take place much faster than the will be the rate of reaction.
reactions occuring between covalent ] Eg : When zinc pieces are added to dilute HCl,
substances.Because in ionic reactions there is no chemical reaction takes place slowly liberating H2
breaking and making of bonds. gas. But the same reaction is rapid by taking
concentrated HCl.
( ) + ( ) → ( ) ↓ + ( )
] From the given below graph it is clear that, the
] The reactions between covalent molecules involve rate of reaction gradually decreases with time
the breaking (cleavage) and the making because of the decrease in the concentration of
(formation) of covalent bonds. reacting substances with time.

+
Eg :

( ) +

( ) ⎯⎯⎯
→
rate

( ) +
( )
] Reactions which involve lesser bond rearrange-
time
ments are rapid at room temperature than those
c) Effect of temperature on the reaction Rate : The
which involve more bond rearrangements.
rate of a reaction increases with increase in the
a) + → temperature.
] In most cases a rise of 10°C in temperature
b)
+ → +

generally doubles the specific rate of the reaction.
] Reaction (a) involves breaking of 6 bonds and ] Increasing the temperature of the substance
formation of six bonds. increases the fraction of molecues, which collide
O with energies greater than activation energy (Ea).
|| Hence increases the rate of reaction.
2 (N = O ) + O = O → 2 N − O ] The ratio of two specific rates measured at
(4) (2) (6)
temperature that differ by 10°C is called the
] Reaction (b) involves breakage of 8 bonds and
Temperature co-efficient
formation of eight bonds.

+
] Temperature co-efficient =

(Normally ° = ° ° + = ° )

For some reactions the ratio approaches the value
of 3.
] Hydrogen and oxygen combine very slowly at

( ) + ( ) →
( )
ordinary temperature but rapidly at high
temperature. According to Arrhenius, this reaction can take
] A piece of coal will burn rapidly in air when the place only when a molecule of hydrogen and a
temperature is raised sufficiently. molecule of iodine collide to form an unstable
] Arrhenius equation for temperature dependence intermediate. It exists for a very short time and
of a rate constant. then break up to form two molecules of hydrogen
iodide.
−
= H H H
I I I
K = Rate constant, + +
A = Arrhenius frequency constant H I H H
I I
Ea = Activation energy,
R= Gas constant Intermediate
T = absolute temperature Activated Complex
−
C
=
Potential Energy
Taking ln on both sides

−
A
=

+ → () H2 +I
∆H B
2HI
at temperature T1 equation (1) is
Reaction coordinate
−
= + →( ) Above Graph showing plot of potential energy Vs
reaction coordinate.
at temperature T2 equation (1) is
−
= + →( )

(since A is constant for a given reaction)
From equations (2) and (3)

− = −
t Energy of
Fraction of molecules
This area shows

activation fraction of additional
⎡ ⎤ (t +10 ) molecules which
= ⎢ − ⎥ This area react at (t+10)
⎣ ⎦ shows fraction
of molecules
reacting at t
⎡ ⎤ Kinetic energy
= ⎢ − ⎥
⎣ ⎦
d) Effect of Catalyst :
] i) Presence of positive catalyst: The function of a
] slope = − =− positive catalyst is to lower down the activation

energy.
(R=gas constant)

The grater the decrease in the activation energy Ex : Acid catalysed hydrolysis reactions of esters,
higher will be the rate of reaction the rate is proportional to the concentration of the
In the presence of a catalyst, the reaction acid catalyst.
follows a path of lower activation energy. ] Catalyst increases the rate of reaction by making
In this condition, a large number of reacting an alternate path of low activation energy for
molecules are able to cross over the energy barrier reactant molecules.
and thus the rate of reaction increases.
Ex : ⎯⎯⎯⎯
→ +

Reaction path
without catalyst EXAMPLE-4
At 27°C and 37°C, the rates of a reaction are given
as 1.6 × 102 mol L1s1 and 3.2 × 102 mol L1 s1.
Ea
Reaction Calculate the energy of activation for the given
path with
Potential energy Ea catalyst reaction.
Sol: The ratio of specific rates at two different
Reactants temperatures are given as,
The effect of a
catalyst is to lower
Product ⎡ − ⎤
the energy of activation = ⎢ ⎥
⎣ ⎦
Reaction coordinate
A catalyst changes the Substituting the values and molar
reaction path ( Positive catalyst )
gas constant R
ii) Presence of negative catalyst : A negative catalyst
×
increases the activation energy of reaction by = ×
forming a new intermediate of high energy, i.e., × ×
by changing the reaction mechanism. Due to Energy of activation = = −
increases activation energy,, some active molecules
become inactive, therefore, rate of reaction
decreases EXAMPLE-5
Let kp denote presence of catalyst and ka The temperature coefficient of a reaction is 2
denote absence of catalyst.
and the rate of reaction at 25°C is 3 mol L1
− − min1. Calculate the rate at 75°C
= ; =
Sol: According the Arrhenius, for every 10°C rise in
Deviding eq.(1) by eq. (2) we get the temperature, the rate of reaction is doubled.
Rise in temperature = 75°C 25°C = 50°C

=
( − ) = ∆ = 5 × 10°C
The increase in the rate of the reaction
= 25 = 32 times
⎡ ∆ ⎤ The rate of reaction
= ⎢ ⎥
⎦ 75°C = 3 × 32= 94 mol L1 min1
⎣
] Catalyst alters Rate Law :
i) rate of reaction ii) path of reaction ] The equation which relates the rate of the reaction
iii) activation energy iv) threshold energy and the concentration of the reactants is known
v) rate constant as rate equation or rate law.
] Catalyst does not alter Rate law is the expression in which reaction
i) ∆G of reaction rate is given in terms of molar concentration of
ii) energy of reactants and energy of products reactants with each term raised to some power,
iii) ∆H iv) ∆S v) Kc which may (or) may not be same as the
] In some reactions, the rate of the reaction is directly stoichiometric coefficient of the reacting species
proportional to the concentration of the catalyst. in a balanced chemical equation.

] nA + mB → Products
ZZZZ
X

+ YZZZ
Z

rate r ∝ [A]n [B]m
rate = K[A]n [B]m
] Rate equation is obtained experimentally. ⎣⎡
⎦⎤
Rate Constant : Rate = = ⎡⎣
⎤⎦ ⎡⎣ ⎤⎦ − ⎡⎣
⎤⎦

] The rate constant of reaction becomes equal to the
rate of the reaction when the concentration of all ⎡ ⎤ ⎡ ⎤
Rate = ⎢ ⎥−⎢ ⎥
the reactants are unity, hence the rate constant is ⎣ ⎦ ⎣ ⎦
also known as the specific reaction rate.
] Rate law equation involving side reactions.
] A + B → Products.
k1
] If the initial concentration of B is taken in large Th227
227
excess than A, then rate depends on A only. Ac
Fr223
] nA + mB → products k2
rate r ∝ [A]n [B]m
rate = K[A]n [B]0 Rate of formation of = ⎡ ⎤
⎣⎢ ⎦⎥
] 'K' is called rate constant or specific rate or rate
per unit concentrations of the reactants. = ⎡ ⎤
Rate of formation of
⎣⎢ ⎦⎥

=
⎡⎣ ⎤⎦
∴ = ( + ) ⎡⎣⎢ ⎤
⎦⎥
Characteristics of Rate constant :
Units of K = mole1− n Ln −1 sec−1
(i) Rate constant is a measure of the rate of the
reaction. Greater the value of k, faster is the
i) + → = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ reaction. Similarly, smaller the value of k indicates
slower the reaction.
ii) → + = ⎡⎣ ⎤⎦
(ii) The value of k depends on the nature of the
reactants. It is a characterstic constant for a
iii)
( ) →
( ) + ( ) = ⎡⎣
⎤⎦ and particular reaction at a fixed temeperature.
Different reactions have different values of k.

( ) + ( ) →
( ) = ⎡⎣
⎤⎦ ⎡⎣ ⎤⎦ (iii) The value of the rate constant of a reaction does
not depend on the concentrations of the reactants.
iv)
( ) →
( ) + ( ) = ⎡⎣
⎤⎦ (iv) The value of rate constant for a particular reaction
changes with temeperature.

( ) +
( ) → (v) The units of rate constant depend on the order of
v)
reaction.

( ) +

( ) Order of the reaction :
] The sum of the powers of the concentration terms
r = ⎡⎣

⎤⎦ ⎡⎣
⎤⎦ of reactants in the rate equation is called order of
the reaction.
vi)
+ → +
] Order of reaction may be zero (or) fraction (or)
negative (or) a whole number (or) integer
Rate = ⎡⎣
⎤⎦ ⎡⎣ ⎤⎦ ] Order of the reaction can be determined by
experimental method only.
vii)

+
→

+

] For elementary reactions order can be obtained

from stoichiometric equation.
Rate = ⎡⎣

⎤⎦ ⎡⎣
⎤⎦
Ex 1 : ¡ + ¢ + £ → products
] Rate law equation for reversible reaction

¡ ¢ £ Let pi be the initial pressure of A and p t the
= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
total pressure of the reaction mixture at time t
order = ¡ + ¢ + £

then =
Ex 2 : + → products ( − )
Examples :
Rate = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
a) ( ) → ( ) + ( )
Order = + = → +
b)
Zero order reactions:
(i) The reaction rate is independent of the c)
→
+
concentration of the reactants
(ii) Some examples of zero order reactions are d) Acid Hydrolysis of ester.
eg : (I) H 2 (g ) + Cl2 (g ) ⎯⎯
hv
→ 2HCl (g )

+

+
(¡ ) ⎯⎯⎯
→
(II) 2NH 3 (g ) ⎯⎯→
M0
hv
N 2 (g ) + 3H 2 (g )
(III) 2HI ⎯⎯
Au
→ H 2 + I2

+

(iii) A → product
e)
→ +

Rate = k[A]°
f) Hydrogenation of ethene is an example of first
− ¡
(iv) = = order reaction.

Where
( ) +
( ) →
( )
C0 = Initial concentration of reactant
Rate = ⎡⎣
⎤⎦
C = Concentration of reactant at t time
x = Concentration of product at t time g) Disintegration of radio active elements
a
(v) Half-life period (t1/ 2 ) = ⇒ (t1/ 2 ∝ a ) →
+
2k
Where a = initial concentration Rate = k[Ra]
(vi) Unit of rate constant : mol lit1 s1 Second order reactions :
First order reactrions:
a) 2A → products = ⎡⎣ ⎤⎦
1) A → Products = ⎡⎣ ⎤⎦
b) A+B → products = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
¡
2) Equation for rate:

= ( − ¡ ) c) Units for rate constant :

Equation for rate constant : =
d) Half life time : ∝ (or) =
−¡
3) Units for rate constant: sec-1
Examples :
4) Half life time : ∝ 1) → +
2) → +
=
alkaline hydrolysis of ester
5) Let us consider a typical first order gas phase 3)

+
→
reaction ( ) → ( ) + ( ) .
+


→ + Rate = K [O 3 ] [O 2 ]
−1
4) 2
→
5)
+
⎯⎯⎯ Order with respect to oxygen is 1
Total Order is 1
6)
→
+
For first order Growth Kinetics
Third order reactions: ] It it used in population growth and bacteria
a) 3A → products (or) multiplication
b) 2A+B → products (or)
c) A+B+C → products +¡
= log

= ⎡⎣ ⎤⎦ (or)
When a is initial population and ( + ¡ ) is
r = K3[A]2[B]1 (or) r = K3[A]1[B]1[C]1
Units for rate constant : lit2.mole-2.sec-1. population after time t
nth Order Reactions:
Half life time : ∝ ] Units for rate constant : litn-1.mole1-n.sec-1. or
(atm)1-n sec-1
Examples : When the order of reaction is n
a) + →
⎡ − − ⎤
b) + → = ⎢ ⎥
− ⎢⎣ − ⎥⎦
c) + → +
(aq) (aq) (aq) (aq)
Fractional order reactions: Half life : ∝
−
a)
+ →

Pseudo Unimolecular reactions :
Rate = ⎡⎣
⎤⎦ ⎡⎣ ⎤⎦ = ] The reactions with molecularity greater than or
equal to 2 but order is one are called Pseudo
b) + →
unimolecular or Pseudo first order reactions.
Rate = k [CO ] [Cl 2 ]
2 1/ 2
order = 2.5 Eg :
1. Hydrolysis of ethyl acetate in acid medium
c) → +
2. Inversion of cane sugar
Rate = ⎡⎣ ⎤⎦ = 3. Hydrolysis of Acetic anhydride
d) CHCl3 + Cl2 → CCl 4 + HCl Molecularity of the reactions:

] The sequence of the elementary steps of a chemical
Rate = ⎡⎣
⎤⎦ ⎡⎣ ⎤⎦ order = 1.5
reaction is known as reaction mechanism.
e) Reaction between H2 and D2
] The slowest elementary step of the reaction is
Rate =
called rate determining step or rate limiting step.
order = 1.5 ] The number of atoms (or) molecules or ions
f) Conversion of para hydrogen to ortho hydrogen participating in the slowest step is called
at high temperature
molecularity.
Rate ∝
( ) order = 1.5 ] Molecularity cannot be zero (or) fraction. It is
always a whole number (or) integer.
Negative order reactions:
] Molecularity is obtained from reaction
] Conversion of ozone into oxygen.
mechanism.

Eg :
→ +
Molecularity Order
2. Molecularity is 2. Order can become

→
+
always whole negative or fraction
+ → trimolecular number or integer. or whole number.
It cannot be zero
] The probability that more than three molecules can or fraction or
collide and react simultaneously is very small. negative.
Hence, the molecularity greater than three is not 3. This is theoretically 3. This is determined
observed. It is, therefore, evident that elucidated from experimentally.
complex reactions involving more than three the mechanism of
molecules in the stoichiometric equation must take the reaction.
place in more than one step. 4. Molecularity 4. Order usually does
does not exceed 3. not exceed 3.
+ +
→
+ ( ) +
Methods of Determination of order of Reaction
1) Trail and Error method or Integrated from of rate
This reaction which apparently seems to be of equation method
tenth order but it is actually a second order
Zero order : →
reaction. This shows that the reaction takes place
r↑
in several steps. Which step controls the rate of ¡ − ( − ¡ )
x = kt (or) = =
the overall reaction. or slow reaction that is called
as rate determining step involves just two species. First order : → ⎯⎯
t
→
Consider the decomposition of hydrogen
=
peroxide which is catalysed by iodide ion in an ( − ¡ )
alkaline medium.
a
− log

⎯⎯⎯⎯⎯⎯⎯⎯→
+ (a − x )

The rate equation for this reaction is found to

be ⎯⎯
t
→
− ⎡⎣
⎤⎦
Rate = = ⎡⎣
⎤⎦ ⎡ − ⎤
⎢⎣ ⎥⎦
This reaction is first order with respect to both
H2O2 and I. Evidences suggest that this reaction
takes place in two steps.
Difference Between Molecularity and order of
Reaction:
Molecularity Order Second order : →
1. The number of 1. The sum of the powers
¡
atoms or-ions or of the concentration = ×
( − ¡ )
molecules terms of the reactants
participating in the in the rate reactants in
rate determining the rate equation is
step of the reaction called order of the x
↑
or activated complex reaction. a (a − x )
of the reaction is
called molecularity. ⎯⎯
t
→
Second order : + → EXAMPLE-7
A first order reaction is 20% complete in 10min.
( − ¡ ) log
b (a − x )
How long it takes to complete 80%?
= a (b − x )
( − ) ( − ¡ ) Sol: Applying first order integral rate equation, rate
2) Half-time (t1/2) method ⎯⎯
t
→ constant K is given as,

⎛ ⎞ ⎛ − = ⇒
⎞ −¡ −
⎜ ⎟ ⎜ ⎟
∝ ; ⎜ ⎟=⎜ ⎟
− ⎜ ⎟ ⎝ ⎠ = −
⎝ ⎠
where n = order of reaction ( = )
3) Vant Hoff Differential method Time required for 80% completion of the first
− order reaction,
=

=
For two initial concentrations C1, C2 we have −
− −
= =
= =
⎛ − ⎞ ⎛ − ⎞
⎜ ⎟ − ⎜ ⎟ EXAMPLE-8
⎝ ⎠ ⎝ ⎠
∴ =
( − ) The initial concentration of ethyl acetate is 0.85
4) Ostwalds Isolation method mol L1. Following the acid catalysed hydrolysis,
This method is useful to determine the order the conentration of ester after 30min and 60min
w.r.t each reactant of a reaction separately by of the reaction are respectively 0.8 and 0.754 mol
L1. Calculate the rate constant and pseudo rate
taking other reactants in excess quantity.
constant.
A + B + C ⇒ products
Sol: Acid catalysed ester follows pseudo first order
Then order with respect to A is n A kinetics. The rate constant k is given as
Order with respect to B is n B
=
−¡
Order with respect to C is n C
Overall order of the reaction = nA + nB + nC = = × − −
( )
EXAMPLE-6
= = × − −
75% of a first order reaction is completed in
30min. Caluculate (a) half life, (b) rate constant
The rate constant (k) is the product of pseudo first
and (c) time required for 99.9% completion of
the reaction. order rate constant ( )
Sol: Time required for 75% completion is 2 half lifes = and concentration of water. (concentration of
30min.
water = − )
⎛ ⎞
(a) Half-life ⎜ ⎟ =15 min = ⎡⎣
⎤⎦
⎝ ⎠
Substituting the values,
(b) ( ) = = = −
− = ⎡
× ⎣ ⎤⎦
(c) Time required for 99.9% of the reaction (t)
Pseudo rate constant = ′

= = = − − −
−¡ = ×

EXAMPLE-9 a) Collisions must occur between the molecules of
reacting gases for a reaction to occur.
For a reaction A+2B → products, when B is taken
b) All collisions do not lead to the formation of
in excess, then the rate law expression and order
products. (Only fruitfull collisions leads to
is
formation of products)
Sol: For the reaction rate law expression is
c) The minimum amount of energy possessed by the
colliding molecules to the formation of products
∴ Order =1 or reaction to occur is known as threshold energy.
d) The energy possessed by the molecules at STP is
known as normal energy or internal energy.
EXAMPLE-10
e) Normal energy possessed by normal molecules is
The following data were obtained during the always less than threshold energy.
first order thermal decomposition of N2O5(g) f) The energy to be gained by the normal molecules
at constant volume : during the collision to convert into products is
known as activation energy or energy of activation.
( ) → ( ) + ( ) g) Activation energy = Threshold energy - energy of
S.NO Time(s) Total Pressure/(atm) normal colliding molecules.
1. 0 0.5 h) Activation energy increases, the rate of the reaction
2. 100 0.512 decreases.
Calculate the rate constant. i) No. of binary collisions per unit time (Z) is
Sol: Let the pressure of N2O5(g) decreases by 2x atm. −
As two moles of N2O5 decompose to give two =
moles of N2O4(g) and one mole of O2(g), the
pressure of N2O4(g) increase by 2x atm and that
= πσ
of O2(g) increases by x atm. πµ
( )→ ( ) + ( ) σAB = collision diameter ;
start t = 0.5 atm 0 atm 0 atm µ = reduced mass
At time t (0.5 2x) atm 2x atm x atm −
j) Specific rate = or = −
= + +
Where P= probability factor
=( − ¡)+ ¡ + ¡ = +¡ ] Activation energy of HI decomposition reaction
52.8K.J/mole.
¡ = − ] For 2NO2 → 2NO+O2 the activation energy is
111K.J/mole. So decomposition of NO2(g) is slower
= − ¡
than decomposition of HI(g).
] The collisions in which molecules collide with
= − (
− )= −
sufficient kinetic energy (called threshold energy)
= = and proper orientation, so as to facilitate breaking
of bonds between reacting species and formation
= − × =
of new bonds to form products are called as
effective collisions. Where as improper orientation

= = makes them simply bounce back and no products

are formed.
For example, formation of methanol from
= × = × − −
bromomethane.
Collision Theory of reaction rates or Kinetic Molecular
theory :
] Collision theory was proposed by Arrhenius and
developed by Max Trautz and William Lewis.
] The main postulates of collision theory are
] The fraction of activated collisions is smaller than ] Reactions with lower activation energy are fast and
the total number of collisions. with higher activation energy are slow.
] Actual rate of reaction is much smaller than the ] For ionic reactions, the energy of activation is
rate of the reaction calculated on the basis of the negligibly small and hence they are instantaneous.
normal collisions. ] For covalent reactions, the energy of activation is
Concept of Activation Energy: The difference high and the reactions are time consuming.
] In the presence of a catalyst alters the path, with a
between the energy barrier (i.e., threshold energy)
new path of low activation energy, the time
ET and the energy of normal molecules ER is required for a covalent reaction is also low.
called activation energy Ea. ] Increasing the concentration of reactants increases
the rate. This is because of the increase in the
= − collision frequency and increase in number of
reactant molecules crossing the energy barrier.
Collision Frequency (Z) :
] Total number of collisions which occur among the
reacting molecules per second per unit volume is
called collision frequency. Its value is given by
z= 2 π vσ 2 n 2
v = average velocity
σ = molecular diameter in cm
n = number of molecules per cc.
∴ Rate (k ) = Ze − Ea / RT
Threshold energy Transition State theory :
] According to this theory, the bimolecular reaction
between two molecules A 2 and B2 passes through
Potential energy R the formation of activated complex which then
Energy of colliding ∆H decomposes to yield the product AB, as shown
molecules
P below
Reaction coordinate + → ( ) →
Exothermic reaction ∆H is −ve ( ¡ )
Threshold energy The constant A has unit of time 1 and is

constant for a given reaction
At very high temperature rate becomes equal
Energy of to frequency factor, i.e., k = A.
Potential energy activation P Photochemical reaction :
] Reactions which take place by the absorption of
R ∆H
radiations of suitable wavelength
Energy of colliding molecules

Eg :
( ) + ( ) ⎯⎯⎯→
( )
Reaction coordinate
Endothermic Re action ∆H is + ve ] Photosynthesis of carbohydrates in plants takes
place in presence of chlorophyll and sunlight
] If activation energy of forward reaction (Eaf) is less
than that of the backward reaction (E ab), the +
⎯⎯⎯→
+

reaction is exothermic.
] The free energy change of a photochemical
] The heat of the reaction, ∆
= − reaction may not be negative.

] In the synthesis of carbohydrates and formation ⎡⎣ ⎤⎦ = ⎣⎡ ⎦⎤ . Thus,
of HCl. ∆G is +ve.
Quantum Efficiency OR Quantum yield ⎡⎣ ⎤⎦
=
] Number of molecules participating in ⎡⎣ ⎤⎦
photochemical reaction with absorption of quanta
is called Quantum Efficiency. It is expressed as = = But =

φ=

Chemiluminescence: ∴ = =
] Emission of light in a chemical reaction at ordinary
temperatures is called chemiluminescence. OTHER IMPORTANT RELATIONS
Fluorescence : Zero reactions:
] The absorption of energy and instantaneous
¡ ¡
reemitting of the energy is called fluorescence. (i) = (ii) = (iii) = ×
Phosphorescence :
] The continuous glow of some substances even First order reactions:
after the cutting of source of light is called
phosphore-scence (i) = (v) =
Eg : ZnS
Bioluminescence : (ii) = (vi) =
] The phenomenon of chemiluminescence exhibited (iii) = (vii) =
by certain living organisms is called
Bioluminescence (iv) = (viii) =
Eg : Light emission by fire flies. Second order reactions
ADDITIONAL INFORMATION (i) = × (ii) = ×
Expression for the amount left after n half lives
(iii) = × (iv) = ×
⎡⎣ ⎤⎦ To determine the order w.r.t . each reactant, follow
After one half-life, amount left =
the following rules:
After two half-lives, amount left K
K
< K

⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
= × =
d
E
After three half-lives , amount left Y ,

⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ d
= × = , ,
Z
,
In general, amount left after n half lives
Some typical linear plots for the reactions of different
⎡⎣ ⎤⎦
= orders :

(a) Plots of rate vs concentrations = ( )

No.of half-lives =
1st Order
zero order
Average life of a first order reaction : Rate Rate
] Average life of a first order reaction is the time in
which the concentration of the reactant is reduced Conc Conc
to 1/e of its original concentration i.e.,

(c) Plots of half-lives vs concentration
3rd Order
∞ − )
2nd Order
(
Rate Rate
zero order
1st Order
(Conc ) (Conc )
2 3
t1/ 2 t1/ 2
(b) Plots from integrated rate equations Conc

Conc
1st Order
Zero Order
2nd Order 3rd Order

Conc A log A
t1/ 2 t1/ 2
Time (t ) Time (t )
1/ a 1/ a 2
1 3rd Order
1 2nd Order
[A] [A] 2
Time (t ) Time (t )

CONCEPTUAL BITS
6. The chemical reaction occurring between
Rate and factors influencing rate
1. Which of the following reactions occurs at covalent molecules involve
measurable rate? 1) breaking of existing bonds
1) Reaction between H+ and OH ions in aqueous 2) formation of new bonds
solution. 3) evolution of heat energy
2) Reaction between AgNO3 and NaCl aqueous 4) 1 & 2
solutions.
3) Hydrolysis of methyl acetate. 7. In a reaction 2A +B → A2B, the reactant A will
4) Reaction between hydrogen and oxygen gases disappear at
at room temperature. 1) half the rate at which B disappears
2. Which of the following reaction is spontaneous 2) the same rate at which B disappears
at room temperature ? 3) the same rate at which A2B is formed
1) I2 + H2O → 2) 2H2 + O2 → 2H2O 4) twice the rate at which B disappears
3) N2 + O2 → 2NO 4) 2HCl → H2 + Cl2 8. For the reaction A → B ; following curves
3. Among the following slowest reaction under represent reaction
identical conditions is
1) 2) 3) 4)
1) H+ + OH → H2O
2) 2KMnO4+5H2C2O4+3H2SO4
The correct curves are
→ K2SO4 + 10CO2+2MnSO4 +8H2O 1) 1,2, only 2) 2,3 only
3) 2KMnO 4 +10FeSO 4 +8H 2 SO 4 → K 2 SO 4 +
3) 1,4 only 4) 3,4 only
2MnSO4+5Fe2(SO4) 3 +8H2O
9. If the first order reaction involves gaseous
4) AgNO3(aq)+NaCl(aq) → AgCl(s)+NaNO3(aq)
reactants & gaseous products, the units of its rate
4. The rate of reaction for N2+3H2 → 2NH3 may be are
represented as 1) atm 2) atm.sec
3) atm.sec-1 4) atm2sec-2
d [N 2 ] 1 d [H2 ] d ⎡NH3 ⎤⎦
1) r = − =− =+ ⎣
dt 3 dt dt 10. The reactions 2NO+O2 → 2NO2, 2CO + O2 →
d [N 2 ] 1 d [H2 ] 1 ⎣⎡
d NH 3 ⎦⎤
2CO2 look to be identical, yet the first is faster
2) r = − = =+ than the second. The reason is that
dt 3 dt 2 dt
1) The first reaction has lower enthalpy change
d [N 2 ] d [H2 ] 1 ⎣⎡
d NH 3 ⎦⎤ than the second.
3) r = − =3 =+
dt dt 2 dt 2) The first reaction has lower internal energy
d [N 2 ] 1 d [H2 ] d ⎡NH3 ⎦⎤ change than the second.
4) r = − = = +2 ⎣
dt 3 dt dt 3) The first reaction has lower activation energy
5. (A) : The rate of a chemical reaction decreases as than the second.
the reaction proceeds. 4) The first reaction has higher activation energy
(R) : The reactant concentration remains than the second.
constant as the reaction proceeds. 11. Burning of coal is represented as
The correct answer is (TSM-2015) C(s) +O2(g) → CO2(g). The rate of this reaction is
1) Both (A) and (R) are true and (R) properly increased by
explanation (A) 1) decrease in the concentration of oxygen
2) Both (A) and (R) are true (R) does not 2) powdering the lumps of coal
explanation (A) 3) decreasing the temperature of coal
3) (A) is true, but (R) is true 4) providing inert atmosphere
4) (A) is false, but (R) is true

12. For a hypothetical reaction ; A → L, the rate 18. For 1/2X2+Y2 → XY2, relative rates of species
given as
expression is, rate = −
− ¡ − ¢ ¡¢
1) = = =+
1) Negative sign represents that rate is negative.
2) Negative sign pertains to the decrease in the
− ¡ − ¢ ¡¢
concentrations of reactant. 2) = − = =+
3) Negative sign indicates the attractive forces
between reactants.
− − ¡ − ¢ ¡¢
4) All of the above are correct. 3) = = =+

13. For the reaction ; 2HI → H2+I2, the expression
¡ + ¢ ¡¢
−

4) = − = =+
represents

1) The rate of formation of HI 19. For N2+3H2 → 2NH3, rates of dis-appearance of
2) The rate of disappearance of HI N 2 and H 2 and rate of appearance of NH 3
3) The instantaneous rate of the reaction respectively are a, b and c, then
4) The average rate of reaction 1) a > b > c 2) a < c < b
3) a = b > c 4) a = b = c
14. For 3A → xB, is found to be 2/3rd of .
20. The reaction CH 3 COOC 2 H 5 +H 2 O ⎯⎯→
H +
Then, the value of x is (PMT-2002) products is a/an --------proces.

1) 1.5 2) 3
A) Instantaneous
3) 1/2 4) 2
B) Spontaneous
15. For a reaction A → 2B, as time proceeds C) Moderately slow
1) [A] ↓ but [B] ↑ . Then correct statement (s) is /are
1) A & C 2) B & C
2) Rate of disappearance of A ↓ , but that of rate
3) A only 4) C only
of appearance of B ↑ .
F
3) Rate of disappearance of A ↑ , but that of rate 21. From the graph the value and the value of
W
of appearance of B ↓ . rate of reaction at X respectively are called
4) Rate with respect to A and B ramain same.
16. At 298 K,l atm, among
A) 2H2+O2 → 2H2O B) H2+Cl2 → 2HCl
C) N2+O2 → 2NO
D) H 2SO4 +KOH → products, correct order of
reaction rates is
1) D > A > C > B 2) D < A < B < C
1) Average rate and instantaneous rate
3) D > B > A > C 4) D > B = C > A
2) Instantaneous rate and average rate
17. In which of the following cases, rate of dis- 3) Average rate only
appearance of any reactant at a given instant 4) Instantaneous rate only
equals to rate of appearance of any product 22. The rate of reaction which does not involve
1) H2+F2 → 2HF gases, does not depend upon
2) N2+3H2 → 2NH3 1) temperature
3) PCl5 → PCl3+Cl2 2) concentration
1 3) pressure
4) H2+ O → H2O 4) catalyst
2 2
23. The specific rate constant of a reaction depends 32. The rate constant K1 of a reaction is found to be
on the double that of rate constant K 2 of another
1) concentration of the reactant reaction. The relationship between corres-
2) time ponding activation energies of the two reactions
3) temperature E1 and E2 can be represented as
4) concentration of the product 1) E1 > E2 2) E1 < E2
3) E1 = E2 4) E1 = 2E2
24. A catalyst accelerates the reaction, because
1) it brings the reactants closer 33. The Arrhenius equation expressing the effect of
2) it lowers the activation energy temperature on the rate constant of reaction is
3) it changes the heat of reaction − E a / RT
1) K = Ea / RT 2) K = A e
4) it increases the activation energy
− E / RT
3) K = log e − Ea / RT 4) K = e a
25. The unit of rate constant depends on
1) number of reactants 34. Activation energy depends on
2) concentration terms 1) pressure
3) order of reaction 2) concentration of reactants
4) molecularity of reaction 3) concentration of products
4) nature of reactants
26. The temperature coefficient of a reaction is
1) the rate constant at a fixed temperature 35. A catalyst in a chemical reaction does not change
2) the ratio of rate constants at two 1) Average energy of reactants or products
temperatures 2) Enthalpy of the reaction
3) the ratio of rate constants at two different 3) Activation energy of the reaction
temperatures differing by 10°C 4) Both 1 and 2
4) the ratio of rate constants at two pressures. 36. In general the rate of a given reaction can be
increased by all the factors except
27. If concentration of reactants is made x times,
1) Increasing the temperature
the rate constant k becomes
2) Increasing the concentration of reactants
1) ek/x 2) k/x 3) unchanged 4) x/k
3) Increasing the activation energy
28. The temperature coefficient of most of the 4) Using a positive catalyst
reactions lies between
37. The effect of temperature on a reaction rate for
1) 1 & 3 2) 2 & 3 3) 1 & 4 4) 2 & 4
which Ea is zero is given by
K t +10 1) with increase of temperature rate increases
29. For a reaction, = x . When temperature is 2) with increase of temperature rate decreases
Kt
3) rate is independent of temperature
increased from 100C to 1000C, rate constant (K) 4) reaction never occurs
increased by a factor of 512. Then, value of x is
38. (A) : The rate law of a reaction cannot be
1) 1.5 2) 2.5 3) 3 4) 2
predicted from its balanced chemical
(Kerala CET-2002)
equation, but must be determined
30. Increase of temperature will increase the reaction experimentally only.
rate due to (Manipal-2002) (R) : The order of a reaction is always an integer
1) increase of number of effective collisions like 0, 1, 2 and 3.
2) increase of mean free path The correct answer is (TSM-2015)
3) increase of number of molecules 1) (A) and (R) are correct, (R) is not the correct
4) increase of number of collisions explanation of (A)
31. The rate constants of a reaction at 300K & 280K 2) (A) is correct but (R) is not correct
respectively are K1 & K2. Then (Kerala CET-2002) 3) (A) is not correct but (R) is correct
4) (A) and (R) are correct, (R) is the correct
1) K1 = 20K2 2) K2 = 4K1
explanation of (A)
3) K1 = 4K2 4) K1=0.5K2

39. In the graph drawn between log K and I/T, 45. Which of the following does not affect the rate
intercept equals to (TNCET-2002) of reaction?
1) Amount of the reactants taken
−
1) 2) log A 2) Physical state of the reactants

3) ∆H of reaction
3) ln A 4) (log A) / 2.303 4) Size of the vessel
40. In the Arrhenius equation, the Boltzmann factor
46. The rate expression gives the relation between
− respresents the ------- of the molecules rate of reaction and

possessing energy in excess of activation energy 1) conc. of reactants 2) conc. of products
1) number 2) fraction 3) rate constant 4) rate law
3) weight 4) percentage 47. Rate of a reaction can be expressed by Arrhenius
41. Activation energies for different reactions are equation k = A.e -E a/RT. In this equation, E a
given below represents (M-2009)
1) A → products, Ea = 14 k.Cal 1) The energy above which not all the colliding
2) B → products, Ea= 16k.Cal molecules will react.
3) C → products, Ea=12k.Cal 2) The energy below which colliding molecules
4) D → products, Ea=10k.Cal will not reacts.
3) The total energy of the reacting molecules at
If the temperature increases by 10°C for which
reactions the temperature coeffcitents are temperature T.
maximum and minimum respectively. 4) The fraction of molecules which energy greater
1) a & b 2) b & c 3) b & d 4) d & b than the activation energy of the reaction.
42. Which of the following parameters of a chemical
1
reaction are increased when a catalyst is used? 48. For N2O5 (in CCl4)®2NO2+ O2, K = 6 × 104 s1
2
1) Rate & activation energy
at at 350K and K = 1.2 × 103 s1at 360 K. Then,
2) Rate constant & enthalpy
when temperature is changed to 380 K, value of
3) Enthalpy & time duration
K (in s1)
4) Rate & rate constant
1) 1.2 × 103 2) 2.4 × 103
43. For an exothermic chemical process occurring in 3) 4.8 × 104 4) 4.8 × 103
two steps as
i)A+B → X (slow) ii)X → AB (fast) Collision theory
The progress of the reaction can be best
ZZX
Z
49. For a reversible reaction, A YZZ
Z B, which one
described by
of the following statements is wrong from the
given energy profile diagram?
1) 2)
3) 4) All are correct

1) Activation energy of forward reaction is
44. Which of the following influence the rate of greater than that of backward reaction.
reaction 2) The threshold energy is less than that of
A) Nature of reactants activation energy.
B) Concentration of reactants 3) The forward reaction is endothermic.
C) Temperature 4) Activation energy of forward reaction is equal
D) Molecularity to the sum of heat of reaction and the activation
1) A, B 2) B, C, D 3) C, D 4) A, B, C energy of backward reaction.
50. In a reaction, threshold energy is equal to 57. The population of activated molecules can be
1) activation energy increased by
1) increase in temperature
2) normal energy of the reactants
2) using a catalyst
3) activation energy +energy of reactants 3) increase of concentration of reactants
4) activation energy- energy of reactants 4) All
51. The value of activation energy for a chemical 58. Consider an endothermic reaction X → Y with
reaction is primarily depends on the activation energies Eb and Ef for
1) temperature the backward and forward reactions, respec-
tively. In general
2) nature of the reacting species 1) Eb < Ef 2) Eb > Ef
3) the collision frequency
3) Eb = Ef 4) no definite relation
4) concentration of the reacting species
59. An endothermic reaction A → B has an activation
52. Wrong statement among the following is energy as xkJ.mol-1 of A. If energy change of the
1) effective collisions are more if activation energy reaction is ykJ, the activation energy of the
is less reverse reaction is
1) x 2) x y 3) x + y 4) y x
2) zero order reaction proceeds at a constant rate
independent of concentration or time 60. Which of the following explains the increase of
3) reactions with highest rate constant values have reaction rate by a catalyst?
1) Catalyst provides the necessary energy to the
lowest activation energies colliding molecules to cross the barrier.
4) if initial concentration increases half life 2) Catalyst decreases the rate of backward
decreases in zero order reaction so that the rate of forward reaction
increases.
53. For the reaction A + B ⇔ C + D , the forward 3) Catalyst decreases the enthalpy change of the
reaction is exothermic. The activation energy of reaction.
formation of A + B is ------- that for the formation 4) Catalyst provides an alternative path of lower
of C + D activation energy.
1) equal to 2) less than
61. The plot of log k vs helps to calculate
3) greater than 4) double
1) Energy of activation
54. Collision theory satisfactorily explains 2) Rate constant of the reaction
1) First order reaction 3) Order of the reaction
2) Zero order reaction 4) Energy of activation as well as the frequency
3) Bimolecular reaction factor
4) Any order reaction 62. The activation energy of a reaction can be
determined by (CET Karnataka-2001)
55. According to collision theory of reaction rates, 1) changing the concentration of reactants
the activation energy is 2) evaluating rate constant at standard
1) the energy gained by the molecule on colliding temperature
with other molecules. 3) evaluating rate constants at two different
2) the energy that molecule should possess in temperatures
order to undergo reaction. 4) by doubling conc. of reactants
3) the energy it should possess so that it can enter Order, molecularity and half - life
into an effective collision. 63. Which of the following is correct?
4) the energy it has to acquire so that it can enter 1) Molecularity of a reaction is always same as
into an effective collision. the order of reaction.
2) In some cases molecularity of the reaction is
56. Increase in the concentration of the reactants same as the order of reaction.
leads to the change in 3) Molecularity of the reaction is always more
1) Heat of reaction 2) Activation energy than order of reaction.
3) Collision frequency 4) Threshold energy 4) Molecularity never be equal to order.

64. If the rate for the chemical reaction is expressed 72. In the following sequence of reaction
as Rate = K[A] [B] n, then
M ⎯⎯
K1
→ N ⎯⎯
K2
→ O ⎯⎯
K3
→ P : K1 < K2 < K3,
1) order of reaction is one
then the rate determining step is (M-2004)
2) order of reaction is n
1) M → N 2) N → O 3) O → P 4) M → P
3) order of reaction is 1+n
4) order of reaction is 1-n 73. ----------- of a reaction cannot be determined
65. The rate equation for the reaction 2A+B → C is experimentally.
1) order 2) rate
found to be : rate = k[A] [B]. The correct statement
3) rate constant 4) molecularity
in relation to this reaction is
1) unit of k must be sec-1 74. The rate expression for the reaction A(g)+B(g)
2) value of k is independent of the initial → C(g) is rate = kCA2 CB1/2. What changes in the
concentrations of A and B inital concentrations of A and B will cause the
3) rate of formation of C is twice the rate of rate of reaction to increase by a factor of eight ?
disappearance of A 1) CA × 2 ; CB × 2 2) CA × 2 ; CB × 4
4) t1/2 is a constant 3) CA × 1 ; CB × 4 4) CA × 4 ; CB × 1
66. Which of the following statements is correct 75. For a reaction pA+qB → products, the rate law
regarding order of reaction ? expression is r = k [A]1 [B]m then
1) first order reaction should be bimolecular 1) (p + q) = (l + m)
2) order of reaction must be positive 2) (p + q) > (l + m)
3) order depends upon stoichiometry 3) (p + q) may or may not be equal to (l + m)
4) order is determined by experimental results. 4) (p + q) ≠ (l + m)
67. If the rate of gaseous reaction is independent of 76. For H2+Cl2 ⎯⎯ X
→ 2 HCl, rate law is given by
pressure, the order of reaction is
R = K. Then, X is (AIIMS-02)
1) 0 2) 1 3) 2 4) 3
1) Pt 2) Ni 3) h υ 4) water
68. For the reaction H 2 +Br 2 → 2HBr , the rate
dc
expression is, rate = K[H 2 ] [Br 2 ] 1/2 which 77. If both & specific rate have same units, then
statement is true about this reaction (IIT) dt
1) The reaction is of second order. rate law is (PMT-2003)
2) Order of the reaction is 3/2. 1) R = K[A]2 2) R = K[A]1/2
3) The unit of K is sec-1. 3) R = K[a]-2 4) R = K
4) Molecularity of the reaction is 2.
78. For A + B → C + D, when [A] alone is doubled,
69. For the following elementary step rate gets doubled. But, when [B] alone is
(CH3)3 CBr(aq) → (CH3) C+(aq)+Br-(aq) the increased by 9 times, rate gets tripled. Then,
molecularity is order of reaction is (Karnataka-2003)
1) Zero 2) 1 3) 2 4) fractional 1) 3/4 2) 3/2 3) 4/9 4) 2
70. The units of rate constant for the reaction
obeying rate expression, r = k[A] [B] 2/3 is
79. Rate law for 2A+B → C+D from following data:
1) mole-2/3 lit2/3 time -1 2) mole2/3 lit-2/3 time-1 (Kerala CET-2004)
3) mole-5/3 lit5/3 time-1 4) mole2/3 lit2/3 time-1 S No [A](M) [B](M) Rate (M/s)
1. 0.01 0.01 2.5
71. Two gases A and B are in a container. The 2. 0.01 0.02 5
experimental rate law for the reaction between 3. 0.03 0.02 45
them has been found to be rate=k[A]2[B]. Predict 1) r = K[A]1/3 [B] 2) r = K[A]2 [B]
the effect on the rate of the reaction when the 3) r = K[A] [B]1/3 4) r = K[A]2/3 [B]1/3
partial pressure of each reactant is doubled
1) the rate is doubled 80. Which of the following relation is correct for a
2) rate becomes four times first order reaction? (k=rate constant; r = rate of
3) the rate becomes six times reaction; c = conc. of reactant ) (M-2004)
1) k = r × c 2 2) k = r × c 3) k = c/r 4) k = r/c
4) the rate becomes eight times

91. Which of the following is a first order reaction
81. of a first order reaction depends on
1) 2N2O5 → 4NO2+O2
(AFMC-03)
2) 2H2O2 → 2H2O+O2
1) time 2) concentration 3) temperature 4) All
82. Which of the following is correct for a first order 3) CH3COOC2H5+H2O ⎯⎯⎯

→ products
reaction? (K=rate constant t1/2 = half- life)
4) All the above
1) t1/2 = 0.693 × K 2) k.t1/2 = 1/0.693
3) k.t1/2 = 0.693 4) 6.93 × k × t1/2=1 92. If a reaction obeys the following equation
83. The half life for a given reaction was doubled as k = 2.303/t log a/a-x the order of the reaction will
the initial concentration of the reactant was be
doubled. The order of the reaction is 1) zero 2) one 3) two 4) three
1) Zero 2) 1st 3) 2nd 4) 3rd
93. The rate constant for a reaction is 2.05x 10-5 mole
84. The inversion of cane sugar into glucose and
fructose is lit-1.sec-1. The reaction obeys ... order
1) 1st order 2) 2nd order 1) First 2) Second 3) Zero 4) Half
3) 3rd order 4) zero order 94. A reaction involves two reactants. The rate of
85. The half-life of a first order reaction is reaction is directly proportional to the
1) independent of the initial concentration of the
concentration of one of them and inversely
reactant.
2) directly proportional to the initial proportional to the concentration of the other.
concentration of the reactant. The overall order of reaction will be
3) inversely proportional to the initial 1) One 2) Two
concentration of the reactant. 3) Zero 4) fractional
4) directly proportional to the square of the initial
concentration of the reactant. 95. In the reaction 2A+B → Products, the order w.r.t
86. The hydrolysis of ester in the presence of alkali A is found to be one and w.r.t B equal to 2.
solution is a .......order reaction. Concentration of A is doubled and that of B is
1) 1 2) 2 3) 0 4) 3 halved, the rate of reaction will be
1) Doubled 2) Halved
87. CH4+Cl2 ⎯⎯⎯ → CH3Cl+HCl; the order of this
reaction is 3) Remain unaffected 4) Four times
1) 0 2) 1 3) 2 4) 3 96. While studying the decomposition of gasesous

88. RCOOR+H2O ⎯⎯⎯→ RCOOH+ROH follows
N2O5, it is observed that a plot of logarithm of
..... reaction kinetics (Karnataka-01) its partial pressure versus time is linear. The
1) 2nd order 2) unimolecular kinetic parameter obtained from this observation
3) pseudo unimolecular 4) zero order is
89. (A) : Molecularity has no meaning for a complex 1) Specific rate 2) Reaction rate
reaction. 3) Energy of activation 4) Molecularity
(R) : The overall molecularity of a complex 97. The correct expression for the rate constant for
reaction is equal to the molecularity of the
reactions of zero order is
slow step.
1) Both A and R are ture, R is the correct 1) k-[A0]/2t 2) k=1/t{[A0] -[A]}
explanation of A 3) k=1/t {[A]-{A0]} 4) k=2.303/t log {[A0]-[A]}
2) Both A & R are true, R is not correct explanation 98. If a is the initial concentration of the reactant,
of A
the time taken for completion of the reaction, if
3) A is true, R is false
4) A is false, R is true it is of zero order, will be
90. 2A → B+C would be a zero order reaction when 1) a/k 2) a/2k 3) 2a/k 4) k/a
rate of reaction (CBSE-02) 99. The slowest step of a particular reaction is found
1) is directly proportional [A] to be 1/2 X2+Y2 → XY2
2) is directly proportional [A]2
The order of the reaction is
3) is independent of change of [A]
4) is independent of [B] & [C] 1) 2 2) 3 3) 3.5 4) 1.5

100. For the reaction A → B, the rate law expression 105. For the reaction system : 2NO (g) + O 2(g) →
is: rate =K[A]. Which of the following statements 2NO2(g) volume is suddenly reduced to half its
is incorrect? value by increasing the pressure on it. If the
1) The reaction follows first order kinetics. reaction is of first order with respect to O2 and
2) The t 1/2 of reaction depends on initial second order with respect to NO, the rate of
concentration of reactants. reaction will
3) K is constant for the reaction at a constant 1) diminish to one -eighth of its initial value
temperature. 2) increase to eight times of its initial value
4) The rate law provides a simple way of 3) increase to four times of its initial value
predicting the concentration. 4) diminish to one-fourth of its initial value
101. What are the units of the rate constant of reaction
in which the half-life is doubled by halving the
initial concentration of reactants. 1) 3 2) 2 3) 2 4) 1 5) 2
1) M-s-1 2) M-1s-1 3) sec-1 4) M-2s-1 6) 4 7) 4 8) 3 9) 3 10) 3
11) 2 12) 2 13) 3 14) 4 15) 1
102. Which of the following represents the
16) 3 17) 3 18) 2 19) 2 20) 2
expression for 3/4 th life of 1st order reaction
21) 1 22) 3 23) 3 24) 2 25) 3
1) 2) 26) 3 27) 3 28) 2 29) 4 30) 1
31) 3 32) 2 33) 2 34) 4 35) 4
36) 3 37) 3 38) 2 39) 2 40) 2
3) 4)
41) 3 42) 4 43) 2 44) 3 45) 4
103. The rate law for a reaction between the 46) 1 47) 2 48) 4 49) 2 50) 3
substances A and B given by Rate = K[A]m[B]n. 51) 2 52) 4 53) 3 54) 3 55) 4
On doubling the concentration of A and halving 56) 3 57) 4 58) 1 59) 2 60) 4
61) 4 62) 3 63) 2 64) 3 65) 2
the concentration of B, the ratio of the new rate
66) 4 67) 1 68) 2 69) 2 70) 1
to the earlier rate of the reaction will be as
71) 4 72) 1 73) 2 74) 2 75) 3
1) (m + n) 2) ( n m )
72) 3 77) 4 78) 2 79) 2 80) 4
1 81) 4 82) 3 83) 1 84) 1 85) 1
3) 2(m n) 4) (m+n)
2 86) 2 87) 1 88) 3 89) 3 90) 3
104. The formation of gas at the surface of tungsten 91) 4 92) 2 93) 3 94) 3 95) 2
due to adsorption is ----- order reaction . 96) 1 97) 2 98) 1 99) 4 100) 2
1) 0 2) 1 101) 2 102) 3 103) 3 104) 1 105) 2
3) 2 4) Insufficient data
Exercise-1 (L.S)
Reaction rates and influencing factors 2. What is the rate of the following reaction for
1. Consider the following reaction 2A → B
N2(g)+3H2(g) → 2NH3(g).
⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
The rate of the reaction in terms of N2 at T (k) is 1) − 2) − 3) − 4)

⎡⎣ ⎤⎦
mole. lit-1.sec-1.What is the value 3. For the reaction 4NH3+5O2 → 4NO + 6H2O, the

rate of reaction with respect to NH 3 is
⎡⎣
⎤⎦ 2 × 10-3 Ms-1. Then the rate of the reaction with
of − (in mole.lit -1.sec -1) at the same respect to oxygen is __ Ms-1.

temperature? 1) 2 × 10-3 2) 1.5 × 10-3
1) 0.02 2) 50 3) 0.06 4) 0.04 3) 2.5 × 10-3 4) 3 × 10-3
4. Concentration of a reactant A is changed from 11. The activation energy of a reaction is 58.3 kJ/
0.044 M to 0.032M is 25 minutes, the average rate mole. The ratio of the rate constants at 305 K and
of the reaction during this interval is 300 K is about (R = 8.3J-1k mol-1) (Antilog 0.1667
1) 0.0048 mole/lit/min = 1.468)
2) 0.00048 mole/lit/sec 1) 1.25 2) 1.75 3) 1.5 4) 2.0
3) 4.8 × 10-4 mole/lit/min 12. Decomposition of NH3 on gold surface follows
4) 0.0048 mole/lit/sec zero order kinetics. If rate constant K is
5 × 10-4 M-s-1, rate of formation of N2 will be
5. In the reaction A → 2B, the concentration of A 1) 10-3 M - s -1 2) 2.5 × 10-4 M - s-1
falls from 1.0M to 0.9982M in one minute what -4
3) 5 × 10 M-s -1 4) Zero
is the rate in moles litre1 sec1
K t+10
1) 1.8 × 103 2) 3.0 × 105 13. For X → Y, = 3. If the rate constant at 300
3) 3.6 × 103 4) 6.0 × 105
Kt
K is Q min-1, at what temperature rate constant
6. The rate of formation of SO3 in the reaction becomes 9Q min-1?
2SO2+O2 → 2SO3 is 100g min-1. Hence, 1) 47°C 2) 320°C
rate of disappearance of O2 is 3) 280K 4) ×
1) 50 g min1 2) 100 g min1 14. At 300K rate constant for A → products at
3) 20 g min 1 4) 40 g min1
t = 50 min is 0.02 s1, then rate constant at t = 75
7. 1dm3 of 2 M CH3COOH is mixed with 1 dm3 of min & 310 k will be (in s1)
3M ethanol to form ester. The decreases in the
⎛ ⎞
intial rate if each sloution is diluted with an 1) 2) ⎜ ⎟ 3) 0.04 4) 0.04 × 25
equal volume of water would be ⎝ ⎠
1) 2 times 2) 4 times 15. The rate constant of first order reaction at
3) 0.25 times 4) 0.5 times 27°C is 103 min-1. The temperature coefficient
of this reaction is 2. What is the rate constant (in
8. The rate law of the reaction, RCl + NaOH →
min1) at 17°C for this reaction ?
ROH + NaCl is given by Rate= k [RCl]. The rate
1) 103 2) 5 × 104 3) 2 × 103 4) 102
of this reaction
A) is doubled by doubling the concentration of Collision theory
NaOH. 16. The minimum energy required for molecules to
B) is halved by reducing the concentration of enter into chemical reaction is
RCl to half. 1) Kinetic energy 2) Potential energy
C) is increased by increasing the temperature of 3) Threshold energy 4) Activation energy
the reaction .
17. In the equilibrium reaction A+B ⇔ C+D,the
D) is unaffected by change in temperature.
activation energy for the forward reaction is 25
Which is correct?
kCals. mole-1 and that of the backward reaction
1) A & B 2) B & C 3) C & D 4) B & D
is 15kCal. mole-1 . Which one of the following
9. The rate of reaction becomes 2 times for every statements is correct ?
10°C rise in temperature. How many times the 1) it is an exothermic process
rate of reaction will increase when temperature 2) it is an endothermic process
is increased from 30°C to 80°C ? 3) it is a reaction for which ∆H = 0
1) 16 2) 32 3) 64 4) 28 4) it is a sublimation process
10. An endothermic reaction A → B has an Order, molecularity, Half - life
activation energy 15 k Cal/mole and the heat of 18. Sucrose decomposes in acid solution into glucose
reaction is 5 kcal / mole. The activation energy and fructose according to the first order rate law,
of the reaction B → A is with t1/2= 3.00 hours . What fraction of sample of
1) 20 kCal/mole 2) 15 kCal/mole sucrose remains after 8 hours?
3) 10 kCal/mole 4) zero 1) 1.158M 2) 0.518M 3) 0.158M 4) 3.182M

19. For an elementary reaction, 2A+B → C+D, the 27. [A] (M) [B] (M) Initial rate (Ms-1)
active mass of B is kept constant but that of A is 0.4 0.3 2 ´ 10-3
tripled. The rate of reaction will 0.8 0.3 0.8 ´ 10-2
1) decrease by 3 times 2) increase by 9 times 1.2 0.9 0.54 ´ 10-1
3) increase by 3 times 4) decrease by 6 times From the above data the rate law for the equation
20. For a chemical reaction Y2+2Z → Product, rate A+B ® products is equal to
1) K[A][B] 2) K[A]2[B]2
controlling step Y+1/2Z → Q. If the 2
3) K[A] [B] 4) K[A][B]2
concentration of Z is doubled, the rate of reaction
will 28. If the initial concentration is reduced to
1) Remain the same 2) Become four times 1/4th of the initial value of a zero order reaction,
3) Become 1.414 times 4) Become double the half-life of the reaction
1) remain constant 2) becomes 1/4th
21. In a reaction A → B, when the concentration of
3) becomes double 4) becomes fourfold
reactant is made 8 times, the rate got
doubled. The order of reaction is ¡
29. If = k[H3O+] n and rate become 100 times
1) 1/3 2) 1 3) 1/2 4) 2
22. The rate of reaction A+2B → Products is given when pH changes from 2 to 1. Hence order of
reactions is
1) 1 2) 2 3) 3 4) 0
by − . If B is present in large

30. The initial concentration of cane sugar in
excess, the order of reaction is presence of an acid was reduced from 0.20 to
1) 3 2) 2 3) 1 4) zero 0.10M in 5 hours and to 0.05M in 10 hours, value
23. For the reaction 2A+B → Products, it is found that of K? (in hr1)
doubling the concentration of both reactants 1) 0.693 2) 1.386 3) 0.1386 4) 3.465
increases the rate by a factor of 8. But doubling 31. 50% completion of a first order reaction takes
the concentration of B alone, only doubles the place in 16 minutes. Then fraction that would
rate. What is the order of the reaction w.r.t to A? react in 32 minutes from the beginning
1) 2 2) 3 1) 1/2 2) 1/4 3) 1/8 4) 3/4
3) 0 4) 1
32. The time needed for the completion of 2/3 of a
24. The increase in rate constant of a reaction is more 1st order reaction, when rate constant is
when the temperature increases from 4.771 ´ 102 min1 is
1) 290K-300K 2) 300K-310K 1) 23.03min 2) 2.303min
3) 310K-320K 4) 320K-330K 3) 6.93min 4) 69.3min
25. The initial rates for gaseous reaction 33. The rate constant of a first order reaction is 0.0693
A+3B → AB3 are given below min-1. What is the time (in minutes) required for
[A] (M) [B] (M) Rate (M sec1) reducing an initial concentration of 20mole. lit-1
0.1 0.1 0.002 to 2.5 mole . lit1 ? (M-2008)
0.2 0.1 0.002 1) 40 2) 10 3) 20 4) 30
0.3 0.2 0.008
34. The half life of a first order reaction is 100
0.4 0.3 0.018
seconds. What is the time required for 90%
Order of reaction is
completion of the reaction?
1) zero 2) three 3) one 4) two
1) 100sec. 2) 200sec.
26. Based on the following data for a reaction what 3) 333sec. 4) 500sec.
is its order (A → products)
35. For the reaction 2N2O5 ® 4NO2 +O2, rate &
Conc.A 2M 0.2M 0.02M 0.00
rate constant are 1.02 × 104 mol lit1 sec1 &
Time in min 0 10 20 ∝ 3.4 ´ 10-5 sec-1 respectively. Then the conc. of
1) 1st 2) 2nd
N2O5 at that time will be
3) 3rd 4) zero
1) 3M 2) 4M 3) 1M 4) 1.5M
36. The half life periods of four reactions labelled 46. For a first order reaction with half life of 150
by A,B,C &D are 30sec, 4.8 min, 180 sec and seconds, the time taken for the concentration of
16 min, respectively. The fastest reaction is the reactant to fall from M/10 to M/100 will be
1) A 2) B 3) C 4) D approximately
37. 3/4th of first order reaction was completed in 1) 1500s 2) 500s
32min, 15/16th part will be completed in 3) 900s 4) 600s
1) 24 min 2) 64 min 3) 16 min 4) 32 min 47. A reaction which is of first order w.r.t reactant A,
38. Initial concentration of the reactant is 1.0M. The has a rate constant is 6 min-1. If we start with
concentration becomes 0.9M, 0.8M and 0.7M in [A] = 0.5 mol.L-1, when would [A] reach the value
2 hours, 4 hours and 6 hours respectively. Then of 0.05 mol.L-1.
the order of reaction is 1) 0.384 min 2) 15min
1) 2 2) 1 3) zero 4) 3 3) 20min 4) 3.84min
39. Half-life periods for a reaction at initial
48. 99% of a 1st order reaction completed in 2.303
concentration of 0.1M and 0.01M are 5 and 50
minutes. What is the rate constant and half-life
minutes respectively. Then the order of
of the reaction
reaction is
1) 2.303 and 0.3010 2) 2 and 0.3465
1) zero 2) 1 3) 2 4) 3
3) 2 and 0.693 4) 0.3010 and 0.693
40. For a first order reaction t0.75 is 138.6 sec. Its
specific rate constant is (in s1) 49. In the case of a first order reaction, the ratio of
1) 102 2) 104 3) 105 4) 106 the time required for 99.9% completion of the
reaction to its half-life is nearly
41. 20% first order reaction is completed in 50
1) 1 2) 10 3) 20 4) 8
minutes. Time required for the completion of
60% of the reaction is ..........min 50. Out of 300g substance [decomposes as per
1) 100 2) 150 3) 262 4) 205 1st order], how much will remain after 18hr?
(t0.5 = 3hr) (CBSE-2000)
42. In a first order reaction, 20% reaction is
1) 4.6gm 2) 5.6gm
completed in 24 minutes. The percentage of
reactant remaining after 48 minutes is 3) 9.2gm 4) 6.4gm
1) 60 2) 64 3) 81 4) 80 51. 75% of a first order process is completed in
43. A first order reaction is half completed in 45 30min. The time required for 93.75% completion
minutes. How long does it need for 99.9% of the of same process (in hr.) (Karnataka-2001)
reaction to be completed ? 1) 1 2) 120
1) 20hours 2) 10hours 3) 2 4) 0.25
3) 7 1/2hours 4) 5 hours
52. For a first order reaction at 27°C, the ratio of time
44. For a first order reaction A → B, the reaction rate required for 75% completion to 25% completion
at reactant concentration of 0.01 M is found to of reaction is
be 2.0 ´ 10-5 mol L-1s-1. The half-life period of 1) 3.0 2) 2.303
the reaction is 3) 4.8 4) 0.477
1) 220s 2) 30s
53. The half-life period of a first order chemical
3) 374s 4) 347s.
reaction is 6.93 minutes. The time required for
45. 99% of a first order reaction was completed in the completion of 99% of the chemical reaction
32min. When will 99.9% of the reaction will be (log 2 = 0.301)
complete? 1) 23.03 minutes 2) 46.06 minutes
1) 50min 2) 46min 3) 460.6minutes 4) 230.3 minutes
3) 49min 4) 48min
54. In a first order reaction, the concentration of the
reactant, decreases from 0.8 M to 0.4 M in 15
minutes. Then 0.1 M becomes 0.025 M in 1) 3 2) 2 3) 1 4) 3 5) 2
6) 3 7) 3 8) 2 9) 2 10) 3
1) 7.5 minutes 2) 15 minutes
3) 30 minutes 4) 60 minutes 11) 3 12) 3 13) 1 14) 3 15) 2
16) 3 17) 2 18) 3 19) 2 20) 3
55. The half life of a first order reaction is 100 21) 1 22) 3 23) 1 24) 1 25) 4
seconds at 280K. If the temperature coefficient 26) 1 27) 3 28) 2 29) 2 30) 3
is 3.0 its rate constant at 290 K is s1 is (M-2014) 31) 4 32) 1 33) 4 34) 3 35) 1
1) 2.08× 103 2) 2.08× 102 36) 1 37) 2 38) 3 39) 3 40) 1
3) 6.93 × 103 4) 6.93× 102 41) 4 42) 2 43) 3 44) 4 45) 4
46) 2 47) 1 48) 2 49) 2 50) 1
51) 1 52) 3 53) 2 54) 3 55) 2
Hints
14. At 300k → k = 0.02 sec1
− ⎛ −
⎞ 31. 1 ⎯⎯⎯⎯→ ⎯⎯⎯⎯
→
1. Rate = = ⎜ ⎟ 310 k → k = 2 × 0.02 sec1
⎝ ⎠ = 0.04 sec1
∴ k is independent on time 32. k =
−¡
2. Rate = − = 15. ∴ At 27°C → k = 103 min1
∴ At 17°C
33. k =
3. Rate of reaction with respect 16. Conceptual −¡
any reactant is same i.e. 17. ∆H = − 34. k50% = k90%
2 × 103 M.Sec1

18. ⎯⎯⎯⎯→ ⎯⎯
→ ⎯⎯
→ =
−
4. Average rate = 19. r1 = k[A]2 [B]1

r2 = k[3A]2 [B] 35. Rate = k [N2O5]1
−
5. Average rate = 20. r1 = k[y] [z]1/2 36. Lesser the half life. Higher is

the rate of reaction.
r2 = k[y] [2z]1/2
−
6. =
21. rate ∝ 37. completion means

7. Initial rate = k[CH3COOH]
22. ∴ rate = k[A]1 [B]0 2T1/2 = 32 min
[C2H5OH]
8. Conceptual 23. Rate = K[A]x [B]y (1)
∴ completion means 4T1/2
9. The increases in the rate of 24. Conceptual
reaction = 2n 25. Conceptual ∴ t = 4T1/2
10. ∆H = − 26. Conceptual 38. For zero order reaction change
in conc. is same in equal
⎛ ⎞
⎡ − ⎤ 27. = ⎜⎜ ⎟⎟ interval of time.
11. = ⎢
⎣ ⎦
⎥ ⎝ ⎠
39. ∝
12. For zero order, rate = k ∝ −
28. −
= 5 × 104 M. Sec1 40. T0.75 = 2T0.5
13. 300 ↔ Q 29. pH decreases from 2 to 1. So 41. k20% = k60%
310 ↔ 3Q [H+] is increase by 10 times.
320 ↔ 9Q =
30. 0.2 ⎯⎯⎯⎯→ 0.1 ⎯⎯⎯⎯→ 0.05;

42. 100 80 64
52. =
− −

43. t99.9% = 10 × t50%

53. =

44. Rate = k[A] & =

45. k99% = k99.9% → 0.4

54. 0.8 ⎯⎯⎯⎯ 15 min.
= 0.2

15 min.
= 0.05 0.1

46. 15 min.
−¡ 15 min.
0.025

47. t =
−¡
55. k =

48. = & T1/2 = 280 k → 0.0693
−¡
290 k → ?
49. t99.9% = 10t50%
50. 300 ⎯⎯
→ 150 ⎯⎯
→ 75 ⎯⎯
→ 37.5 ⎯⎯
→ 18.75
→ 9.375 ⎯⎯
⎯⎯ → 4.6875 gm.
51. t75% = 2t50% & t93.75% = 4t50%
Exercise-2 (W.S-1)
Reaction rates and influencing factors 3. For the reaction N2 + 3H2 → 2NH3, the rate of
1. In the reaction, 2A+B → 2C+D, the rate of disappearance of H2 is 0.01 molelit1 min1. The
disappearance of A is 2.6 ´ 10-2M. S-1. Then the rate of appearance of NH3 would be
rate of disappearance of B and the rate of 1) 0.001 mol. lit-1 min-1
appearance of C and D are respectively 2) 0.02 mol. lit-1 min-1
1) 5.2 × 10-2; 5.2 × 10-2; 2.6 × 10-2 3) 0.007 mol. lit-1 min-1
2) 1.3 × 10-2; 2.6 × 10-2; 1.3 × 10-2 4) 0.002 mol. lit-1 min-1
3) 2.6 × 10-2; 2.6 × 10-2; 2.6 × 10-2 4. Observe the following reaction 2A + B → C.
4) 2.6 × 10-2; 5.2 × 10-2; 5.2 × 10-2 The rate of formation of C is 2.2 ´ 103 mol L1
2. In a reaction N 2 + 3H2 → 2NH3, the rate of ⎡⎣ ⎤⎦
min1. What is the value of − (in mol.
appearance of ammonia is 2.5 ´ 10-4 Msec-1, the
1
L min ? 1)
rate of disappearance of N2 will be
1) 7.5 × 10-4 2) 1.25 × 10-4 1) 2.2 × 103 2) 1.1 × 103
3) 4.4 × 10 3 4) 5.5 × 103
3) 5 × 10-4 4) 2.5 × 10-4

5. Rate of the reaction A → products is 0.5 3 × 10-3 mol.lit-1.min-1. What is the value of
M-s-1 at time t = 15sec, then rate at time t = 5sec
⎡ ⎤
will be (in M - s1) − ⎣ ⎦ in mol.lit-1. min-1? (M-2006)
1) 0.15 2) 0.164 3) > 0.5 4) (0.5)2
6. For the reaction A+3B → 2C + D, which one of 1) 3 × 10-3 2) 9 × 10-3

the following is not correct? 3) 10-3 4) 1.5 × 10-3
1) Rate of disappearance of A = Rate of formation
of D 12. For a gaseous reaction, the rate expression is
k [A] [B]. If the volume of the reaction vessel is
2) Rate of formation of C = × Rate of reduced to 1/4 th of the initial volume, the
disappearance of B reaction rate, relating to orginal rate will be ---
-times
3) Rate of formation of D = × Rate of 1) 10 2) 8 3) 1/10 4) 16
disappearance of B 13. 2SO2+O2 ⇔ 2SO3, if the volume of the reaction
4) Rate of disappearance of A = 2 × Rate of vessel is doubled, the rate of forward reaction
formation of C will be
7. For A → Products, pressures of A t = 0 & t = 10 1) 1/4 th of initial value
min respectively are 760 mm & 740 mm, rate of 2) 1/8 th of initial value
reaction during this time interval is (in mm/min) 3) 4 times of its initial value
4) 8 times of its initial value
1) 2 2) 200 3) 20 4) 14. The rate of reaction for A®products is
×
10 mol.lit-1. min-1 at time t1= 2 minutes. What will
8. For A → Products (order = 1), rate of reaction is be the rate (in mol.lit-1, min-1 )at time t2 = 12
4 ´ 102M5 when [A] = 0.5M, then, rate constant minutes ?
(k) of the reaction is (in sec1) 1) > 10 2) < 10 3) 10 4) 20
1) 0.08 2) 0.02 3) 4 ×102 4) 0.05 15. For xA → yB (gaseous phase) reaction, rate of
9. In the reaction 2NO + O2 → 2NO2, if the rate of reaction is doubled when [A] is quadrapled, then
disappearance of O2 is 16 gm min-1, then the rate rate law is given by
of appearance of NO2 is 1) Rate = k[A]2 2) Rate = k[A]1/2
1/2 1/2 4) Rate = k[B]1/2
1) 90 gm. min1 2) 46 gm. min1 3) Rate = k[A] [B]
1
3) 28gm. min 4) 32 gm. min1
16. For an elementary reaction 2A(g)+B(g) ® C(g),
10. In the formation of sulphur trioxide by the initial rate w.r.t. A is 0.04 M.s.1. when volume
of vessel is doubled, final rate w.r.t B will be (in
contact process, 2SO2(g)+ O2(g) R 2SO3(g) The
M.s1)
rate of reaction is expressed as 1) 0.04 2) 0.08 3) 0.01 4) 0.1
⎡⎣ ⎤⎦ − 17. Rate constant (k) for xA → products is 0.04M-
− M sec 1 . The rate of dis-
min-1. When [A] changes form 0.08 M to 0.02M,
appearance (in mol L s1) of SO2 will be ratio of initial to final rate is
1) 5.0 × 10-4 mol L s1 1) 1/4 2) 1/16 3) 1 4) 4
2) 2.25 × 10-4 mol L s 18. For the first order reaction A → product. When
3) 3.75 × 10-4 mol L s1 the concentration of A is 2. 5 x 10 -2 M the
activation energy is 20k. Cal/mole. If the conc.
4) 50.0 × 10-4 mol L s1
of [A] is doubled, at same temperature, the
11. Observe the following reaction : A(g) + 3B (g) activation energy becomes equal to
1) 40k.cal/mole 2) 10k.cal/mole
⎡ ⎡⎣ ⎤⎦ ⎤
−
→ 2C(g) The rate of this reaction ⎢⎢ − ⎥⎥ is
⎣ ⎦ 3) 20k.cal/mole 4) k.cal/mole

19. In a reaction vessel, the reactants have attained 27. For a reaction the following mechanism has been
a potential energy of 80k.Cal/mole and unable proposed
to give even traces of products. The threshold 2A+B → D+E
energy of the reaction may be A+B → C+D (slow) A+C → E(fast)
1) 60k.Cal 2) 70k.Cal
3) 40k.Cal 4) 100k.Cal The rate law of expression for the reaction is
20. For a reaction, K = 2 ´ 1013 e30000/RT. When log K 1) r = k[A]2[B] 2) r = k[A][B]
3) r = k[A]2 4) r = k[A][C]
(y-axis) is plotted against 1/T(x-axis), slope of line
will be ...Cal
28. For the reaction, A → products the rate of
30000 −30000 −30000 −30000 reaction is 7.5 x 10-4 mole. lit-1 sec-1. If the
1) 2) 3) 4)
4.6 4.6 2.303 46 concentration of A is 0.5 mole lit-1, the rate
21. N2(g)+3H2(g) U 2NH3(g)+22k.Cal. constant is
The activation energy for the forward reaction 1) 1.5 × 103sec1 2) 2.5 × 105sec1
is 50k.Cal. What is the activation energy for the 4
3) 3.75 × 10 sec 1 4) 8 × 104sec1
backward reaction? (2007-M)
1) 72k.Cal 2) 28k.Cal 29. A first order reaction takes 40 min for 30%
3) +28k.Cal 4) +72k.Cal decomposition. Calculate half life.
1) 17.69min 2) 7.77min
22. For xA+yB → products, change of concentration 3) 77.69min 4) 27.69min
of either A or B brings no change in reaction rate.
Then 30. A + B → products, the kinetic data of the reaction
1) x = y = 0 2) x = 0 ; y ≠ 0 is
A (mole lit-1) B(mole lit-1) Rate (m.lit-1Sec-1)
3) x ≠ 0; y = 0 4) Order of reaction = 0
1) 0.5 0.5 2x10-4
23. A+ B → products is an elementary reaction. 2) 0.5 1.0 1.99x10-4
When excess of A is taken in this reaction, then 3) 1.0 0.5 2.01x10-4
the molecularity and order are respectively
1) 2 and 2 2) 2 and 1 The order of the reaction is
3) 1 and 2 4) 1 and 1 1) one 2) zero
3) two 4) fractional
24. For the reaction A+B → products, it is found that
order of A is 1 and order of B is 1/2. When 31. The conversion of A to B follows second order
concentrations of both A&B are increased four kinetics. Doubling the conc. of A will increase
times the rate will increase by a factor the rate of formation of B by a factor
1) 6 2) 8 3) 4 4) 16 1) 4 2) 2 3) 1/4 4) 1/2
25. For 2NH 3 ⎯⎯⎯ → N +3H , rate w.r.t N is 32. nA → products is a first order reaction with
2 2 2
2 × 10 M-min -1 , then rate w.r.t N 2 after 20
3 k=10min -1 . If the reaction is started with
minutes will be (in M-min1) [A] = 0.4M, after .....min, [A] would become 0.04M
1) 2x10-3 2) >2x10-3 1) 4.606 2) 2.303
3) 10-4 4) <2x10-3 3) 0.2303 4) 6.99
26. For a given reaction of first order, it takes 20min Au
33. 2HI → H2+I2 following nth order kinetics. If rate
for the concentration to drop from 1.0M to 0.6M.
The time required for the concentration to drop 1 d[HI]
from 0.6M to 0.36M will be of the reaction is given by rate = − = K.
2 dt
1) more than 20 minutes and concentration of HI drops from 1.5M to 0.3M
2) less than 20 minutes in 7.3sec, value of K is
3) equal to 20 minutes 1) 8 × 10-2M-s-1 2) 8 ×10-2M-1- s-1
4) infinity 3) 8 × 10-2M-2-s- 4) 8 × 10-2s-1

34. Observe the following data regarding 41. For the rate of the reaction
→ N +3H 2H2O2 → 2H2O2 + O2, r = K[H2O2] it is
2NH3 ⎯⎯ 2 2
Pressure (in atm) : 5 10 20 1) Zero order reaction
Half life (min) : 3.6 1.8 0.9 2) First order reaction
The unit of K is 3) Second order reaction
1) min-1 2) atm-min-1 4) Third order reaction
3) (atm-min) -1 4) atm-2 min-1
42. 75% of a first order reaction completed in 32
35. For 2A+B+C → Products, rate law is given by rate minutes then, the time required to complete
= k[A] [B] 2 & rate constant (k) is 99.9% of a reaction in seconds is
2 x 10-6M-2-s-1. Then rate of the reaction become 1) 9600 2) 96,500
2 x10-9M-s-1 only when 3) 16 4) 1660
[A] [B] [C]
1) 0.1M 0.1M 0.2M 43. t 1/4 can be taken as the time taken for the
2) 0.2M 0.1M 0.2M concentration of a reactant to drop to 3/4 of its
3) 0.1M 0.1M any value initial value. If the rate constant for a first order
4) 0.2M 0.2M 0.1M reaction is k, the t1/4 can be written as
36. For A → B , a first order reaction, k=10 -2 1) 0.10/k 2) 1.41k 3) 0.29/k 4) 0.75/k
min-1. If initial concentration of A is 1M, rate of 44. Rate of a reaction, A+B → Product is given as a
the reaction after 2.31 hrs will be (in function of different initial concentrations of A
M-min-1) and B
1) 2.5 × 10-4 2) 5 × 10-3 [A]|mol.L-1 [B]|mol.L-1 Rate[mol.L-1.min-1]
3) 2.5 × 10-1 4) 2.5 × 10-3 0.01 0.01 0.005
0.02 0.01 0.010
37. A(g) → B(g) is first order reaction. The initial 0.01 0.02 0.005
concentration of A is 0.2 mol.lit -1 . After The half-life of in the reaction is
10 minutes the concentration of B is found to be 1) 0.693min 2) 1.386min
0.18 mol.lit-1. The rate constant (in min-1) for 3) 2.079min 4) 0.3465min
the reaction is (M-2008) 45. In the first order reaction 10% of the reaction is
1) 0.2303 2) 2.303 completed in 10hours. The percentage of the
3) 0.693 4) 0.01 reaction completed in 20 hours is
1) 20% 2) 19% 3) 38% 4) 81%
38. Rate expression for xA+yB → products is Rate=
k[A]m[B]n. Units of K w.r.t A&B respectively are 46. For the first order reaction, half life is 14s. The
s-1 and M-1. s-1,
when concentrations of A & B time required for the initial concentration to
reduce to 1/8th of its value is
are increased by 4 times, then
1) 28s 2) 42s 3) (14) 3s 4) (14) 2s
1) Rf = 16 Ri 2) Ri = 16 Rf
3) Rf = 8 Ri 4) Rf = 64 Ri 47. The time taken for the completion of 90% of a
first order reaction is `t min. What is
39. For SO2Cl2(g) → SO2(g)+Cl2(g) time required for the time (in sec) taken for the completion of 99%
99.9% completion of the process is found to be of the reaction? (M-2005)
100sec, then specific rate? 1) 2t 2) t/30 3) 120t 4) 60t
1) 6.93 × 10-3M.s-1 2) 6.93 × 10-2M.s-1 48. The half life for the reaction.
3) 6.93 × 10-3s-1 4) 6.93 × 10-2M-1s-1
N2O5 ® 2NO2+ O2 is 24hrs. at 300C. Starting
40. A reaction rate is found to depend upon the two
concentration terms. The order of the reaction is with 10g of N2O5 how many grams of N2O5 will
1) 1 2) 3 remain after a period of 96 hours?
3) 2 4) cant be predicted 1) 1.25g 2) 0.63g 3) 1.77g 4) 0.5g

49. The half life of a first order reaction A → B+C is 52. The temperature coefficient of a reaction is 2.5.
10 minutes . The concentration of A would be If its rate constant at T1K is 2.5 × 103 s1, the rate
reduced to 10% of the original concentration in constant at T2K is s1 is (T2 > T1) (M-2013)
1) 10 minutes 2) 33minutes 1) 1.0× 103 2) 6.25 × 103
3) 90 minutes 4) 70 minutes 3) 1.0× 102 4) 6.25 × 102
50. In a first order reaction, 0.16 moles of reactant 55. The rate of reaction A ® products is 10 mole/lit/
decreases to 0.02 moles in 144 minutes, its half- min at time (t1) = 2min. What will be the rate in
life is mole/lit/min at time (t2) 12 min (APM-2018)
1) 24min 2) 12min 3) 72min 4) 48min 1) more than 10 2) 10
3) less than 10 4) 20
51. For a first order reaction we have K=100s-1. The
time for completion of 50% reaction is
1) 10-2s 2) 4x10-5s
3) 6.93x10-3s 4) 7x10-5s
1) 2 2) 2 3) 3 4) 3 5) 3
52. The product of half life (t1/2) and the square of 6) 4 7) 1 8) 1 9) 2 10) 1
initial concentration of the reactant
11) 2 12) 4 13) 2 14) 2 15) 2
(a) is constant. Then the order of reaction is
16) 3 17) 3 18) 3 19) 4 20) 4
1) 2 2) 3 3) 0 4) 1
21) 4 22) 4 23) 2 24) 2 25) 1
53. For C12H22O11[sucrose](aq) ⎯⎯⎯
→ products, the 26) 3 27) 2 28) 1 29) 3 30) 2
31) 1 32) 3 33) 1 34) 2 35) 3
concentration of sucrose changes from
36) 4 37) 1 38) 4 39) 2 40) 4
0.06 M to 0.03 M in 30 minutes. Then,
41) 2 42) 1 43) 2 44) 2 45) 2
concentration of sucrose at the end of 60 minutes
46) 2 47) 3 48) 2 49) 2 50) 4
will be
51) 3 52) 2 53) 2 54) 2 55) 3
1) zero 2) 0.015 M 3) 0.09 M 4) 0.12 M
Hints
¡
− − 15. Rate ∝
1. Rate = = 7. Rate =

−
16. =
8. Rate = k[A]
−
2. = 17. Conceptual
⎛ − ⎞
9. = ⎜ ⎟ 18. Conceptual
⎝ ⎠ 19. Ea = ET ER
− ⎛ −
⎞
3. = ⎜ ⎟
⎝ ⎠
− 20. log k = log A
10. =

−
4. =
− ∴ slope =
11. =−
5. As time goes, the rate of
reaction decreases 21. ∆H = −
12. R = k[A] [B]
i.e. rate at t = 5 must be greater
13. r1 = k[SO2]2 [O2] 22. No change in reaction rate
than rate at t = 15 sec
i.e. > 0.5 M. sec1 show zero order
⎡⎣ ⎤⎦ 23. (n reactants A & B are
− − r2 = × =
6. Rate = = required. So the molecularity

is two.
14. As time goes, the rate of
If A is excess then the order
= = reaction decreases.
becomes one.

24. r1 = K[A] [B]1/2
43. k =
r2 = K[A] [B]1/2
−
r2 = K[A] [B] × 4 × 41/2
25. Conceptual 44. Rate = K[A]
26. Conceptual
27. rate = k[A] [B] =

28. k = 45. 100

29. k = 46. 1 ⎯⎯⎯→

⎯⎯⎯
→

⎯⎯⎯
→
−¡
30. In all cases rate is constant 47. K90% = K99%
∴ It is zero order reaction

=

31. r1 = k[A]2
48. 96 hrs = 4t1/2

32. k =
−¡
49. =

−
33. k =

K=
−¡
34. As ∝

50. k =
−¡
35. rate = k [A] [B]2
36. 2.31 hrs = 138.6 min = 2T1/2 =

∴ =
51. t50% =

37. 1M - - - - - 0.5 M - - - - - 0.25 M

52. t1/2 ∝ ⇒= × −
∴ Rate = k[A] = 102 × 0.25 −
= 2.5 × 103 M. min1
53. t1/2 = 30 min
38. rf = K[A] [B]2 60 min = 2t1/2

39. k = =
−¡ 54.
40. Conceptual
41. Conceptual
55. =

42. K75% = K99.9%

Exercise-3 (W.S-2)
8. H2O2 decomposes with first order kinetics in a
1. For N2+3H2 U 2NH3 rate of disappearance of H2
3 lit. container. If the pressure developed in
is 0.01 M. min-1. The amount of NH3 formed at
10min. is 380 mm, the average rate at 270C is
that instant will be
1) 0.01M.min-1 2) 0.002M.min-1
1) -0.02mol 2) zero -1
3) 0.05M.min 4) 0.06M.min-1
3) 0.1132g 4) 0.17g

9. For H2 + Cl2 ⎯⎯⎯→ 2HCl, rate of reaction at t=0
2. For A → B, [A] changed from 4.4 ´ 10-2 M to 3.2 ´

−∆ is given by rate = − [H2] = 6x 103 M-min-1,

10-2M in 25 min. Now will be
∆ then rate of the reaction at t = 30min is
1) 4.8 × 10-4M.min-1 2) 4.8 × 10-4M.s-1
-3 -1 −
3) 9.6 × 10 M.s 4) 2.4 × 10-4M.min-1 1) [Cl2] = 2 × 10-4M-min-1

3. A gaseous phase reaction A2 → B+0.5C shows an
increase in pressure from 100 mm to 120mm in
2) [HCl]=12 × 10-3M-min-1

5 min. Now, − should be
−
3) [Cl2] = 6 × 10-3 M -min-1
1)8mm-min-1 2)4mm-min-1
3) 16mm-min-1 4) 2mm-min-1 4) Zero
4. For N 2O 5 → 2NO 2 + 1/2O 2, it is found that 10. For SO 2 Cl 2(g) → SO 2(g) +Cl 2(g) , pressure of
SO2Cl2changed from 5 atm to 4 atm in 10 min.
− Then, pressure of SO2Cl2at the end of 30 minutes
[N2O5] = k1[N2O5], [NO2]=k2[N2O5];
will be (in atm)
1) 2.56 2) 3.56 3) 4.56 4) 5.56

[O2]= k3[N2O5], then
11. In the process 2N2O5(g) → 4NO2(g)+O2(g) at t = 10
1) k1 = 2k2 = 3k3 2) 2k1 = 4k2 = k3 rate of reaction w.r.t N2O5, NO2 & O2 respectively
3) 2k1 = k2 = 4k3 4) k1 = k2 = k3 are
N2O5 N O2 O2
5. In the process nA (g) → mB(g), rate of
1) 500mm/min 400mm/min 200mm/min
disappearance of A is 5 x 10-3M-min-1& rate of
appearance of B is 10-2M- min-1 at same instant. 2) 1000mm/min 1000mm/min 500mm/min
Then values of n & m respectively are 3) 1000mm/min 2000mm/min 4000mm/min
1) 2;3 2) 1;2 3) 2;1 4) 4;3 4) 400mm/min 400mm/min 400mm/min
12. K,A&Eaof a process at 25°C respectively are
6. For the process 2A → products, rate
5 ´ 10-4s-1,6 ´ 1014s1 & 108 kJ/mol. Then the value
of reaction w.r.t A at 10 th second is
of rate constant as time → ∞ will be (in s1)
2 x 10-2M-s-1 then rates of same process at 5th &
15 th seconds (order ≠ 0) respectively are 1) 1.2 × 1018 2) zero
(in M/s) 3) 6 × 1014 4) 5 × 10-4
1) 10-1 & 4 × 10-2 2) 2.7 × 10-2 & 1.6 × 10-2 13. The process 2A+B → C taking place in two steps:
3) 1.6 × 10-2 & 2.7 × 10-2 4) 2 × 10-2 & 2 × 10-2 1:2A → D
7. For N 2 +O 2 → 2NO, initially N 2 &O 2 are at 2 : D+B → C (slow), then rate of reaction gets
pressure 500mm & 700 mm at t = 0. If the pressure 1) doubled when [B] is halved
of N2is 480 mm at t = 20 min average rate of 2) doubled when [A] is doubled
reaction is (in mm - min-1) 3) doubled when [B] is doubled
1) 5 2) 1 3) 2 4) 4 4) quadrapled when [A] is doubled

14. Thermal decomposition for N2O5(g) follows 1st −
21. For 2H2O2 → 2H2O+O2. if [H2O2] = k1[H2O2]

order kinetics as per N2O5(g) → 2NO2(g)+ O2(g).

= k [H O ]
2 2 2 ; =k3[H2O2]
Initial pressure of N2O5 is 100 mm. After 10
minutes, the pressure developed is 1) k1 = k2 = k3 2) k1= k2 = 2k3
130 mm, what is rate constant? 3) 2k1 = 2k2 = k3 4) k1= k2 = 4k3
1) 0.693min-1 2) 0.025min-1
hυ
3) 2.303x10-4min-1 4) 0.03min-1 22. For CH4+Cl2 → product, rate of reaction at t= 0 is
15. For So 2 Cl 2(g) → SO 2(g) +Cl 2(g) , pressures of −
SO2Cl2at t=0 and t=20 minutes respectively are given by rate = [CH 4 ] = 2 ´ 10 -4

700 mm & 350mm. When log (P0/p) is plotted M- min-1, then rate of the reaction at t = 20 min is
against time (t), slope equals to (in M-min-1)
1) 1.505 x10-2s-1 2) 1.202 x10-3min-1
−
3) 1.505 x10-2min-1 4) 0.3465 min-1 1) rate = [CH4] = 10-4

16. Decomposition of NH3 on Pt surface follows
−
zero order kinetics. If the initial pressure is 2) rate = [Cl2] = 2 × 10-4

4 atm, the product of t1/2 and k equals to
1) 2 atm 2) 4 atm 3) 16 atm 4) 1.414 atm −
3) rate = [CH4] =4 ×10-4
17. Thermal decomposition of HI(g) on Gold surface
follows zero order kinetics. With initial pressure −
1000mm. If the pressure after 20 sec is 960 mm 4) rate = [Cl2] = zero

w.r.t HI(g), rate constant is (in mm sec-1)
1) 40 2) 20 23. For the process X(g) → products, (order ≠ o), rates
of disappearances of x at t=0, t=50 sec & t
3) 2 4) =30 sec respectively are p, q & r mol/lit/sec then
1) p < q < r 2) r < p < q
18. For 2SO2+O2 → 2SO , rate of disappearance of
3 3) q > r > p 4) p > r > q
SO2 is 4x10 M- s-1 at t=10sec. Then, the amount
-3
of SO3 formed & amount of O2 consumed at t=10 24. A → B, K1 = 0.693 sec-1 C → D, K2 = 0.693 min-1 If
sec respectively are t1 and t2 are half lives of two reactions, then
1) t1=t2 2) t1= 60t2
1) 0.1g;0.1g 2) 0.1g;0.2g
3) t2 = 60t1 4) t2=2.303t1
3) 0.016g;0.064g 4) 0.32g;0.064g
19. For 2A → B, [A] changed from 0.08M to 0.04M in 25. A → B and C → D are first order reactions. Ratio
of t99.9% values is 4:1, then ratio of rate constants
K1 to K2 is
100 seconds. Now
1) 4:1 2) 2:1 3) 1:1 4) 1:4
1) 2 × 10-4M-s-1 2) 4 × 10-4M-s-1 26. For a first order reaction temperature coefficient
3) 8 × 10-4M-s-1 4) 1.2 × 10 M-s-1 is 2. If the value of K at 310K is 2x10-2min-1, t1/2
of at 300K will be (in min)
20. For N2O4(g) → 2NO2(g), pressure is found to be
1) 69.3 2) 23.03 3) 46.06 4) 69.1
increased form 700mm to 800mm in 10min. Then
−
27. For a first order process, value is - 23.03, then
with respect to N2O4 is

value of K/A is
1) 10mm/min 2) 70mm/min 1) 102.303 2) 10-10
3) 80mm/min 4) 150mm/min 3) 10-23.03 4) 1010

28. For 2H2O2 → 2H2O2+O2, t0.5 = 0.301 hr. When 36. 900ml of pure & dry O2 is subjected to O2 seilent
[H2O2] at t= 0 is 0.5M, initial rate is electric discharge, so that after a time 10 min.
1) 2.303 M.h-1 2) 1.151M.h-1 volume of ozonized oxygen is found to be 870
3) 4.606M.h -1 4) 0.031M.h-1 ml. Now, average rate of reaction in this interval
is (in ml/min)

29. For 2NH3(g) ⎯⎯→ products follows zero order 1) 3 2) 9
3) 90 4) 60
kinetics. If t1/2 at p=4 atm is 25 sec, t1/2 at p=16
atm will be (in sec) 37. For a first order process A → B, rate constant k1
1) 6.25 2) 625 3) 100 4) (25)1/4 380.693 min-1 & another first order process
C → D, k2 = x min-1. If 99.9% of C → D requires
30. The rate constant of a reaction at 200K is 10 times
less than rate constant at 400K. Value of Ea? time same as 50% of reaction A → B, value of x?
[R= gas constant] (in min-1) .
1) 1842.4R 2) 921.2R 3) 460.6R 4) 230.3R 1) 0.0693 2) 6.93
3) 23.03 4) 13.86
31. The chemical reaction 2O3 → 3O2 proceeds as
follows: 38. 90% of first order process X → Y is completed in
a time equals to 99% of another first order
O +O&O+O ⎯⎯⎯⎯ 2O the rate
O3 ⎯⎯⎯ → 2 3 → 2 process Y → Q. If K value of Y → Q is 0.09 sec-1,
law expression should be K value of X → Y will be (in sec-1)
1) r = K[O3]2 2) r = K[O3]2[O2]-1 1) 0.27 2) 0.3 3) 0.03 4) 0.045
3) r = K[O3][O2]2 4) r = K[O3][O2]2
32. For producing the effective collisions the 39. For C12H22O11 (Sucrose)(aq) ⎯⎯⎯
→ products,
colliding molecules must have the concentration of sucrose changes from 0.06
1) a certain minimum amount of energy M to 0.03 M in 30 minutes. Then, concentration
2) energy equal to or greater than threshold of sucrose at the end of 60 munutes will be
energy 1) zero 2) 0.015 M
3) proper orientation 3) 0.09 M 4) 0.12M
4) both threshold energy and proper orientation
33. Depletion of ozone occurs as:
2O3 → 3O2 1) 3 2) 1 3) 1 4) 3 5) 2
6) 2 7) 2 8) 2 9) 3 10) 1

Z
ZZZX
Step1 : O3 YZZZ
Z O2+(O)(fast) 11) 1 12) 4 13) 3 14) 2 15) 3
16) 1 17) 3 18) 4 19) 1 20) 1
ZZZXZ
Step2 : O3 YZZZ Z 2O2(slow) 21) 2 22) 2 23) 4 24) 3 25) 4
What is the order of the reaction 26) 1 27) 2 28) 2 29) 3 30) 2
1) 1 2) 2 3) 3 4) 0.5 31) 2 32) 4 33) 1 34) 3 35) 3
36) 2 37) 2 38) 4 39) 2
34. A hypothetical reaction A2+B2 → 2AB follows
the mechanism as given below:
A2 U A+A(fast)
A+B2 → AB+B (slow)
A+B → AB(fast)
The order of the over all reaction is:
1) 2 2) 1 3) 1.5 4) 0
35. For 2A(g) → B(g) +3C(g), Pressure of A at t=0 is

500mm, then total pressure of A,B & C at t=10
min is
1) 100mm 2) 400mm 3) 540mm 4) 250mm

Hints
8. PV = nRT
⎡
⎤⎦
1. − ⎣ =
× = × ×
+ ⎡⎣
⎤⎦ ⎛ − ⎣⎡
⎦⎤ ⎞ ×
∴ = ⎜ ⎟= =
⎜ ⎟
⎝ ⎠
∆
= − = = = 0.02 M min1

= × − 9. H2 + Cl2 ⎯⎯⎯ 2HCl is a zero order reaction
→
= − ∴ Rate is independent of conc. of H2 and Cl2
⎡
⎤⎦ ⎡ ⎤⎦
∆ ⎛ − ⎞ × − − × − ∴− ⎣ =− ⎣ = 6 × 103 M. min2
2. − =⎜ ⎟=
∆ ⎝ ⎠ 10. 5 atm to 4 atm is 20% reaction completion in 10min.
= × − −
⎯⎯⎯⎯

→ ⎯⎯⎯⎯
→
⎯⎯⎯⎯
→

3. → + 11. The rate of reaction w.r.t any reactant at any time
100 0 0 will be same.
100 x x 0.5x 12. k is independent on time (5 ×104 sec1)
−¡+¡+ ¡ = 13. Slow step is the rate determining step
x = 40 ⎡⎤
= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ = ⎣ ⎦
∆ ⎡ ⎤
∴− ⎣ ⎦ = = − ⎡⎣ ⎤⎦
∆
⇒ = ⎡⎣ ⎤⎦
4. → +
= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
− ⎡⎣ ⎤⎦ ⎣⎡ ⎦⎤ ⎣⎡ ⎦⎤
= = = ∴ rate gets doubled when [B] is doubled.

14. → +
⎡⎣ ⎤⎦ = ⎡⎣ ⎤⎦ = ⎡⎣ ⎤⎦
100 0 0 Initial
∴ = = ¡
100 x 2x After time t
5. →
∴ −¡+ ¡+ ¡=
⎡ − ⎡⎣ ⎤⎦ ⎤ ⎡⎣ ⎤⎦ ¡= ⇒ x = 20
= ⎢ ⎥=
⎢⎣ ⎥⎦
− ⎡⎣ ⎤⎦
= = = −
− ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ × −
= = = =
−

= = = = −

n = 1, m = 2
6. As the time goes the rate of reaction decreased due 15. From data =
to the decrease in the conc. of reactant
At 5th sec rate > 2 × 102 Ms1 = = = −

At 15th sec rate < × − −

⎣⎡ ⎤⎦ ( − ) − −
∴ =
7. = = = =


= × ⎡ − ⎡⎣
⎤⎦ ⎤ ⎡⎣
⎤⎦ ⎡⎣ ⎤⎦
21. = ⎢ ⎥= =
⎣⎢ ⎥⎦

∴ = =
⎡⎣
⎤⎦ = ⎡⎣
⎤⎦ = ⎡⎣
⎤⎦
= = × − −
16. For zero order reaction ∴ = =
22. As time goes rate of reaction decreases
= 23. Rate of reaction w.r.t any reactant is same.

24. For a 1st order reaction

× = = =
= = =

17. For zero order, rate is equal to rate constant.
∆ − = = =
∴ = =
∆
− ∴ =
= = −
⎛ ⎡ ⎤⎦ ⎞ 25. For 1st order, ∝

⎣⎡ ⎦⎤ − ⎣⎡ ⎦⎤
− ⎜ ⎣

18. ⎟= =
⎜ ⎟
⎝ ⎠

= =
⎡ ⎤⎦ ⎡⎣ ⎤⎦
∴ ⎣ = = × − −
26. At 310k, k = 2 × 102
= 4 × 103 × 80 = 0.32 gms At 300k, k = 102
⎡⎣ ⎤⎦ ⎡ − ⎡⎣ ⎤⎦ ⎤
= ⎢ ⎥= × × − = = =
−
⎢⎣ ⎥⎦
− − −
= × 27. = −

= × − × =
− −
= = =−
∆ ⎡⎣ ⎤⎦ −
= = × − −
19.
−
=
−
∆ ⎡ ⎤ ⎛ −∆ ⎡⎣ ⎤⎦ ⎞ ×
∴ ⎣ ⎦= ⎜ ⎟=
∆ ⎜ ∆ ⎟
⎝ ⎠ 28. = =

= × − −
20. N2O4 → 2NO2 = ⎡⎣

⎤⎦ = × = −
(g) (g)
700 0 Initial 29. t1/2 is directly proportional to initial conc. for
700 x 2x Final order reaction.
700 = x + 2x = 800 ∴ = =
x = 100
Then At P = 16, = × =
∆
∴ = = −
∆

⎡ − ⎤ 35. → +
30. = ⎢ ⎥
⎣ ⎦
(g) (g) (g)
⎡ ⎤ 500 0 0 Initial
=
⎢⎣ × ⎥
⎦ 5002x x 3x After time t
∴ At t, Total presure = 500 2x + x + 3x

= = = 500 + 2x
×
i.e > 500 mm.
⎡ ⎤ ⎡ ⎤
31. U + ( ) = ⎣ ⎦ ⎣ ⎦ 36. →
⎡⎣ ⎤⎦
900 0 Initial
⎡ ⎤
⇒ ⎡⎣ ⎤⎦ = ⎣ ⎦ 900 3x 2x After time t
⎡⎣ ⎤⎦ 900 3x + 2x = 870
x = 30 ml
= ⎡⎣ ⎦⎤ ⎡⎣ ⎤⎦ + ( ) U
Vol. of O reacted = ¡ = × =
⎡ ⎤ − −
= ⎡⎣ ⎤⎦ ⎣ ⎦ = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ Average rate = ⎡ ⎤ = = −
⎡⎣ ⎤⎦ ⎣ ⎦
32. Conceptual ×
1
33. U + ( )
37. = ⇒ =

⎡ ⎤ ⎡ ⎤ ⎡ ⎤
= ⎣ ⎦⎣ ⎦ ⇒ = ⎣ ⎦
⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ = ⇒ =

= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ + ( ) U at t1 = t2 Then

⎡ ⎤ − = −
= ⎡⎣ ⎤⎦ ⎣ ⎦ = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
⇒ = = =
⎡⎣ ⎤⎦
∴ Order of reaction = 2 - 1 = 1
38. For →
34. + → +
= = =
= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ ... (1)
For → reaction
⎡⎤
U + = ⎣ ⎦ ⇒ ⎡⎣ ⎤⎦ = ⎡⎣ ⎤⎦ Then 10t1/2
⎣⎡ ⎦⎤

(→ )
= ( → )
⎡⎣ ⎤⎦ = ⎡⎣ ⎤⎦
From equation (1)
= × =
(→ )
= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
∴ = = = −
= ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
→ 0.03 M ⎯⎯⎯⎯→ 0.015M

38. 0.06M ⎯⎯⎯⎯

∴ Order = + =
60 min

Previous Questions
EAMCET
1. In a first order reaction the concentration of the 6. For a first order reaction at 27°C the ratio of time
reactant decrease from 0.6 M to 0.3 M in 15 required for 75% completion to 25% completion
minutes. The time taken for the concentration of reaction is (2009-E)
to change 0.1 M to 0.025 M in minutes is. 1) 3.0 2) 2.303 3) 4.8 4) 0.477
1) 1.2 2) 12 3) 30 4) 3 (2014-E) 7. The rate constant of a first order reaction is 0.0693
2. Which of the following plots is correct for a first min1. What is the time required (in minutes) for
order reaction (2013-E) reducing an initial concentration of 20 mol lit1
to 2.5 mol lit1 ? (2009-M)
1) 40 2) 30 3) 20 4) 10
Log(a-x)
8. For a reversible reaction ⇔ which one of

Log(a-x)
1) 2)
the following statements is wrong from the given
Time Time
energy profile diagram. (2008-E)
(a-x)
E
Log(a-x)
3) 4) B
A
Time Time
3. Which one of the following statements is correct

for the reaction ? (E-2012) Reaction coordinate
1) Activation energy of forward reaction is greater

⎯⎯→

than backward reaction.
( ) ( ) 2) The forward reaction is endothermic
1) Order is two but molecularity is one 3) The threshold energy is less than that of
2) Order is one but molecularity is two activation energy
3) Order is one but molecularity is one 4) The energy of activation of forward reaction is
4) Order is two but molecularity is two equal to the sum of heat of reaction and the
energy of activation of backward reaction.
4. Match the following : (2011-E)
LIST-I LIST - II 9. For the reaction → which one of
A) Arrhenius equation i) free energy change the following is not correct? (M-2010)
B) Slowest step in a ii) conc1 . time1 1) Rate of disappearance of A= Rate of formation
reaction mechanism
C) Rate constant of iii) conct n . time1 of D
a II order
2) Rate of formation of Rate of
D)The possibility of a iv) Rate determining
reaction step disappearance of B
−
v)
3) Rate of formation of = × Rate of
1) A-v, B-i, C-iii, D-iv 2) A-v, B-iv, C-iii, D-ii
3) A-v, B-iv, C-ii, D-i 4) A-iii, B-iv, C-ii, D-i disappearance of B
5. The plot of log K vs 1/T yields a straight line. 4) Rate of disappearance of = × Rate of
The slope of the line would be equal to(E-2010) formation of C

1) − 2) −

1) 3 2) 3 3) 4 4) 3 5) 2

3) 4) 6) 3 7) 2 8) 3 9) 4


JEE MAINS
1. The time for half life period of a certain reaction 5. For the non-stoichiometric reaction
A ® products is 1 hour. When the initial
→ , the following kinetic data were
concentration of the reactant A is 2.0 mol L1,
how much time does it take for its concentration obtained in three separate experiments, all at
to come from 0.50 to 0.25 mol L1 if it is a zero 298 K (AIEEE-2014 )
order reaction ? (AIEEE - 2010 ) / / / rate
&
1) 4 h 2) 0.5 h 3) 0.25 h 4) 1h

2. The rate of chemical reaction doubles for every mol lit −1 sec −1
10C rise of temperature. If the temperature is 0.1M 0.1M 1.2 × 10−3
raised by 50C, the rate of the reaction increases
by about (AIEEE - 2011) 0.1M 0.2M 1.2 × 10−3
1) 24 times 2) 32 times
0.2M 0.1M −3
2.4 × 10
3) 64 times 4) 10 times
3. A reactant (A) forms two products ⎯⎯⎯

→ ,
The rate law for the formation of C is

→ , Activation
Activation energy , ⎯⎯⎯ 1) = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ 2) = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦

energy . If , then K and K are

1 2
3) = ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦ 4) = ⎡⎣ ⎤⎦
related as (AIEEE -2011)

1) = 2) = 6. For a reaction → , rate of disappearance
of A is related to the rate of appearance of B

3) = 4) = by the expression (AIEEE-2008 )
4. Consider the reaction, ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
1) − =
( )
( ) → ( )
(− )
( )
⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
The rate equation for this reaction is , 2) − =

⎡⎣ ⎤⎦ ⎡⎣
⎤⎦
⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
Which of these mechanisms is/are consistent 3) − =
with this rate equation ? (AIEEE - 2010)
A) +
→
+ + − + + +
− ( ) ⎡⎣ ⎤⎦ ⎡⎣ ⎤⎦
4) − =

+ +
− →
+ + − + ( )
B)
U
+ +
− (fast eqilibrium) 1) 3 2) 2 3) 2 4) 4 5) 4
6) 2
+
− → − +
+ + ( )
1) (B) only 2) Both (A) and (B)
3) Neither (A) nor (B) 4) (A) only

1. For the reaction 3. The half-life period of a 1st order reaction is 60
5Br (aq) + BrO3 (aq) + 6H+ (aq) minutes. What percentage will be left over after
® 3Br2 (aq) + 3H2O (l) 240 minutes ? [KCET-2016]
1) 6.25% 2) 4.25% 3) 5% 4) 6%
⎡%U2− ⎤
⎣⎢ 3 ⎦⎥ 4. If the half life of a first order reaction is 60 min,
If, = 0.05 mol L1 min1,
W the approximate time in min, required to
complete 90% of the reaction is (log 2 = 0.3)
[%U 2 ] [TSM (I)-2106]
in mol L1 min1 is [TS Engg-2015 ]
W 1) 200 2) 240 3) 50 4) 100
1) 0.01 2) 0.05 3) 0.5 4) 0.01 5. A particular reaction has a rate constant of 1.3 ×
2. The time required for a first order reaction to t 1/2
complete 90% is l. What is the time required to 102 s1. Find out the ratio of
t 99.9
completed 99% of the same reaction
[AP Engg-2015] [TSM (III)-2016]
1) 2t 2) 3t 3) t 4) 4t 1) 10 2) 1/10 3) 100 4) 1
1) 4 2) 1 3) 1 4) 1 5) 2

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121 Chemical Kinetics

Fig : Instantaneous rate of reaction The rate of appearance of ammonia = −  −

Chemical Kinetics 122

123 Chemical Kinetics

(Normally ° = ° ° + = ° )

Taking ln on both sides

This area shows

Chemical Kinetics 124

125 Chemical Kinetics

] For elementary reactions order can be obtained

Chemical Kinetics 126

127 Chemical Kinetics

d) CHCl3 + Cl2 → CCl 4 + HCl Molecularity of the reactions:

Chemical Kinetics 128

The rate equation for this reaction is found to

Chemical Kinetics 130

Threshold energy The constant A has unit of time 1 and is

Chemical Kinetics 132

133 Chemical Kinetics

(b) Plots from integrated rate equations Conc

2nd Order 3rd Order

Chemical Kinetics 134

135 Chemical Kinetics

Then, the value of x is (PMT-2002) products is a/an --------proces.

137 Chemical Kinetics

3) 4) All are correct

139 Chemical Kinetics

Chemical Kinetics 140

141 Chemical Kinetics

143 Chemical Kinetics

Chemical Kinetics 146

45. k99% = k99.9%  → 0.4

 =  0.2

=  0.05 0.1

51. t75% = 2t50% & t93.75% = 4t50%

147 Chemical Kinetics

6. For the reaction A+3B → 2C + D, which one of 1) 3 × 10-3 2) 9 × 10-3

Chemical Kinetics 148

149 Chemical Kinetics

Chemical Kinetics 150

151 Chemical Kinetics

  

30. In all cases rate is constant 47. K90% = K99%

∴ It is zero order reaction

37. 1M - - - - - 0.5 M - - - - - 0.25 M

Chemical Kinetics 152

153 Chemical Kinetics

Chemical Kinetics 154

35. For 2A(g) → B(g) +3C(g), Pressure of A at t=0 is

155 Chemical Kinetics

Chemical Kinetics 156

⎛  ⎡ ⎤⎦ ⎞ 25. For 1st order,  ∝

20. N2O4 → 2NO2  =  ⎡⎣

157 Chemical Kinetics

 → 0.03 M ⎯⎯⎯⎯→ 0.015M

Chemical Kinetics 158

8. For a reversible reaction ⇔ which one of

Fig : Instantaneous rate of reaction The rate of appearance of ammonia = − −

(Normally ° = ° ° + = ° )

Threshold energy The constant A has unit of time 1 and is

45. k99% = k99.9% → 0.4

= 0.2

= 0.05 0.1

⎛ ⎡ ⎤⎦ ⎞ 25. For 1st order, ∝

20. N2O4 → 2NO2 = ⎡⎣

→ 0.03 M ⎯⎯⎯⎯→ 0.015M

energy . If , then K and K are

of A is related to the rate of appearance of B