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1 | 3405-3408 | 22)
log kk
E
T
T
T T
2
1
2
1
1
2
=2 303
−
a
R
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Example 3 10
Example 3 10
Example 3 |
1 | 3406-3409 | kk
E
T
T
T T
2
1
2
1
1
2
=2 303
−
a
R
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Solution
Example 3 10
Example 3 10
Example 3 10
Example 3 |
1 | 3407-3410 | 10
Example 3 10
Example 3 10
Example 3 10
Example 3 |
1 | 3408-3411 | 10
Example 3 10
Example 3 10
Example 3 10
Example 3 |
1 | 3409-3412 | 10
Example 3 10
Example 3 10
Example 3 9
Example 3 |
1 | 3410-3413 | 10
Example 3 10
Example 3 9
Example 3 9
Example 3 |
1 | 3411-3414 | 10
Example 3 9
Example 3 9
Example 3 9
Example 3 |
1 | 3412-3415 | 9
Example 3 9
Example 3 9
Example 3 9
Example 3 |
1 | 3413-3416 | 9
Example 3 9
Example 3 9
Example 3 9
Rationalised 2023-24
82
Chemistry
A catalyst is a substance which increases the rate of a reaction without
itself undergoing any permanent chemical change |
1 | 3414-3417 | 9
Example 3 9
Example 3 9
Rationalised 2023-24
82
Chemistry
A catalyst is a substance which increases the rate of a reaction without
itself undergoing any permanent chemical change For example, MnO2
catalyses the following reaction so as to increase its rate considerably |
1 | 3415-3418 | 9
Example 3 9
Rationalised 2023-24
82
Chemistry
A catalyst is a substance which increases the rate of a reaction without
itself undergoing any permanent chemical change For example, MnO2
catalyses the following reaction so as to increase its rate considerably 2KClO3
MnO2
2 KCl + 3O2
The word catalyst should not be used when the added substance
reduces the rate of raction |
1 | 3416-3419 | 9
Rationalised 2023-24
82
Chemistry
A catalyst is a substance which increases the rate of a reaction without
itself undergoing any permanent chemical change For example, MnO2
catalyses the following reaction so as to increase its rate considerably 2KClO3
MnO2
2 KCl + 3O2
The word catalyst should not be used when the added substance
reduces the rate of raction The substance is then called inhibitor |
1 | 3417-3420 | For example, MnO2
catalyses the following reaction so as to increase its rate considerably 2KClO3
MnO2
2 KCl + 3O2
The word catalyst should not be used when the added substance
reduces the rate of raction The substance is then called inhibitor The
action of the catalyst can be explained by intermediate complex theory |
1 | 3418-3421 | 2KClO3
MnO2
2 KCl + 3O2
The word catalyst should not be used when the added substance
reduces the rate of raction The substance is then called inhibitor The
action of the catalyst can be explained by intermediate complex theory According to this theory, a catalyst participates in a chemical reaction by
forming temporary bonds with the reactants resulting in an intermediate
complex |
1 | 3419-3422 | The substance is then called inhibitor The
action of the catalyst can be explained by intermediate complex theory According to this theory, a catalyst participates in a chemical reaction by
forming temporary bonds with the reactants resulting in an intermediate
complex This has a transitory existence and decomposes to yield products
and the catalyst |
1 | 3420-3423 | The
action of the catalyst can be explained by intermediate complex theory According to this theory, a catalyst participates in a chemical reaction by
forming temporary bonds with the reactants resulting in an intermediate
complex This has a transitory existence and decomposes to yield products
and the catalyst It is believed that the catalyst provides an
alternate pathway or reaction mechanism by
reducing the activation energy between
reactants and products and hence lowering
the potential energy barrier as shown in
Fig |
1 | 3421-3424 | According to this theory, a catalyst participates in a chemical reaction by
forming temporary bonds with the reactants resulting in an intermediate
complex This has a transitory existence and decomposes to yield products
and the catalyst It is believed that the catalyst provides an
alternate pathway or reaction mechanism by
reducing the activation energy between
reactants and products and hence lowering
the potential energy barrier as shown in
Fig 3 |
1 | 3422-3425 | This has a transitory existence and decomposes to yield products
and the catalyst It is believed that the catalyst provides an
alternate pathway or reaction mechanism by
reducing the activation energy between
reactants and products and hence lowering
the potential energy barrier as shown in
Fig 3 11 |
1 | 3423-3426 | It is believed that the catalyst provides an
alternate pathway or reaction mechanism by
reducing the activation energy between
reactants and products and hence lowering
the potential energy barrier as shown in
Fig 3 11 It is clear from Arrhenius equation (3 |
1 | 3424-3427 | 3 11 It is clear from Arrhenius equation (3 18)
that lower the value of activation energy faster
will be the rate of a reaction |
1 | 3425-3428 | 11 It is clear from Arrhenius equation (3 18)
that lower the value of activation energy faster
will be the rate of a reaction A small amount of the catalyst can catalyse
a large amount of reactants |
1 | 3426-3429 | It is clear from Arrhenius equation (3 18)
that lower the value of activation energy faster
will be the rate of a reaction A small amount of the catalyst can catalyse
a large amount of reactants A catalyst does
not alter Gibbs energy, DG of a reaction |
1 | 3427-3430 | 18)
that lower the value of activation energy faster
will be the rate of a reaction A small amount of the catalyst can catalyse
a large amount of reactants A catalyst does
not alter Gibbs energy, DG of a reaction It
catalyses the spontaneous reactions but does
not catalyse non-spontaneous reactions |
1 | 3428-3431 | A small amount of the catalyst can catalyse
a large amount of reactants A catalyst does
not alter Gibbs energy, DG of a reaction It
catalyses the spontaneous reactions but does
not catalyse non-spontaneous reactions It is
also found that a catalyst does not change the equilibrium constant of
a reaction rather, it helps in attaining the equilibrium faster, that is, it
catalyses the forward as well as the backward reactions to the same
extent so that the equilibrium state remains same but is reached earlier |
1 | 3429-3432 | A catalyst does
not alter Gibbs energy, DG of a reaction It
catalyses the spontaneous reactions but does
not catalyse non-spontaneous reactions It is
also found that a catalyst does not change the equilibrium constant of
a reaction rather, it helps in attaining the equilibrium faster, that is, it
catalyses the forward as well as the backward reactions to the same
extent so that the equilibrium state remains same but is reached earlier Though Arrhenius equation is applicable under a wide range of
circumstances, collision theory, which was developed by Max Trautz
and William Lewis in 1916 -18, provides a greater insight into the
energetic and mechanistic aspects of reactions |
1 | 3430-3433 | It
catalyses the spontaneous reactions but does
not catalyse non-spontaneous reactions It is
also found that a catalyst does not change the equilibrium constant of
a reaction rather, it helps in attaining the equilibrium faster, that is, it
catalyses the forward as well as the backward reactions to the same
extent so that the equilibrium state remains same but is reached earlier Though Arrhenius equation is applicable under a wide range of
circumstances, collision theory, which was developed by Max Trautz
and William Lewis in 1916 -18, provides a greater insight into the
energetic and mechanistic aspects of reactions It is based on kinetic
theory of gases |
1 | 3431-3434 | It is
also found that a catalyst does not change the equilibrium constant of
a reaction rather, it helps in attaining the equilibrium faster, that is, it
catalyses the forward as well as the backward reactions to the same
extent so that the equilibrium state remains same but is reached earlier Though Arrhenius equation is applicable under a wide range of
circumstances, collision theory, which was developed by Max Trautz
and William Lewis in 1916 -18, provides a greater insight into the
energetic and mechanistic aspects of reactions It is based on kinetic
theory of gases According to this theory, the reactant molecules are
3 |
1 | 3432-3435 | Though Arrhenius equation is applicable under a wide range of
circumstances, collision theory, which was developed by Max Trautz
and William Lewis in 1916 -18, provides a greater insight into the
energetic and mechanistic aspects of reactions It is based on kinetic
theory of gases According to this theory, the reactant molecules are
3 4 |
1 | 3433-3436 | It is based on kinetic
theory of gases According to this theory, the reactant molecules are
3 4 1 Effect of
Catalyst
3 |
1 | 3434-3437 | According to this theory, the reactant molecules are
3 4 1 Effect of
Catalyst
3 5 Collision
3 |
1 | 3435-3438 | 4 1 Effect of
Catalyst
3 5 Collision
3 5 Collision
3 |
1 | 3436-3439 | 1 Effect of
Catalyst
3 5 Collision
3 5 Collision
3 5 Collision
3 |
1 | 3437-3440 | 5 Collision
3 5 Collision
3 5 Collision
3 5 Collision
3 |
1 | 3438-3441 | 5 Collision
3 5 Collision
3 5 Collision
3 5 Collision
Theory of
Theory of
Theory of
Theory of
Theory of
Chemical
Chemical
Chemical
Chemical
Chemical
Reactions
Reactions
Reactions
Reactions
Reactions
Fig |
1 | 3439-3442 | 5 Collision
3 5 Collision
3 5 Collision
Theory of
Theory of
Theory of
Theory of
Theory of
Chemical
Chemical
Chemical
Chemical
Chemical
Reactions
Reactions
Reactions
Reactions
Reactions
Fig 3 |
1 | 3440-3443 | 5 Collision
3 5 Collision
Theory of
Theory of
Theory of
Theory of
Theory of
Chemical
Chemical
Chemical
Chemical
Chemical
Reactions
Reactions
Reactions
Reactions
Reactions
Fig 3 11: Effect of catalyst on activation
energy
log k2 =
a
1
1
2
1
1
log
2 |
1 | 3441-3444 | 5 Collision
Theory of
Theory of
Theory of
Theory of
Theory of
Chemical
Chemical
Chemical
Chemical
Chemical
Reactions
Reactions
Reactions
Reactions
Reactions
Fig 3 11: Effect of catalyst on activation
energy
log k2 =
a
1
1
2
1
1
log
2 303
E
k
T
T
R
=
1
5
1
1
1
1
209000 J mol L
log 1 |
1 | 3442-3445 | 3 11: Effect of catalyst on activation
energy
log k2 =
a
1
1
2
1
1
log
2 303
E
k
T
T
R
=
1
5
1
1
1
1
209000 J mol L
log 1 60 10
600 K
700K
2 |
1 | 3443-3446 | 11: Effect of catalyst on activation
energy
log k2 =
a
1
1
2
1
1
log
2 303
E
k
T
T
R
=
1
5
1
1
1
1
209000 J mol L
log 1 60 10
600 K
700K
2 303 8 |
1 | 3444-3447 | 303
E
k
T
T
R
=
1
5
1
1
1
1
209000 J mol L
log 1 60 10
600 K
700K
2 303 8 314 J mol L K
log k2 = – 4 |
1 | 3445-3448 | 60 10
600 K
700K
2 303 8 314 J mol L K
log k2 = – 4 796 + 2 |
1 | 3446-3449 | 303 8 314 J mol L K
log k2 = – 4 796 + 2 599 = – 2 |
1 | 3447-3450 | 314 J mol L K
log k2 = – 4 796 + 2 599 = – 2 197
k2 = 6 |
1 | 3448-3451 | 796 + 2 599 = – 2 197
k2 = 6 36 × 10–3
s–1
Rationalised 2023-24
83
Chemical Kinetics
assumed to be hard spheres and reaction is postulated to occur when
molecules collide with each other |
1 | 3449-3452 | 599 = – 2 197
k2 = 6 36 × 10–3
s–1
Rationalised 2023-24
83
Chemical Kinetics
assumed to be hard spheres and reaction is postulated to occur when
molecules collide with each other The number of collisions per
second per unit volume of the reaction mixture is known as
collision frequency (Z) |
1 | 3450-3453 | 197
k2 = 6 36 × 10–3
s–1
Rationalised 2023-24
83
Chemical Kinetics
assumed to be hard spheres and reaction is postulated to occur when
molecules collide with each other The number of collisions per
second per unit volume of the reaction mixture is known as
collision frequency (Z) Another factor which affects the rate of
chemical reactions is activation energy (as we have already studied) |
1 | 3451-3454 | 36 × 10–3
s–1
Rationalised 2023-24
83
Chemical Kinetics
assumed to be hard spheres and reaction is postulated to occur when
molecules collide with each other The number of collisions per
second per unit volume of the reaction mixture is known as
collision frequency (Z) Another factor which affects the rate of
chemical reactions is activation energy (as we have already studied) For a bimolecular elementary reaction
A + B ® Products
rate of reaction can be expressed as
a /
AB
Rate
Z
e E
RT
−
=
(3 |
1 | 3452-3455 | The number of collisions per
second per unit volume of the reaction mixture is known as
collision frequency (Z) Another factor which affects the rate of
chemical reactions is activation energy (as we have already studied) For a bimolecular elementary reaction
A + B ® Products
rate of reaction can be expressed as
a /
AB
Rate
Z
e E
RT
−
=
(3 23)
where ZAB represents the collision frequency of reactants, A and B
and e
-Ea /RT represents the fraction of molecules with energies equal to
or greater than Ea |
1 | 3453-3456 | Another factor which affects the rate of
chemical reactions is activation energy (as we have already studied) For a bimolecular elementary reaction
A + B ® Products
rate of reaction can be expressed as
a /
AB
Rate
Z
e E
RT
−
=
(3 23)
where ZAB represents the collision frequency of reactants, A and B
and e
-Ea /RT represents the fraction of molecules with energies equal to
or greater than Ea Comparing (3 |
1 | 3454-3457 | For a bimolecular elementary reaction
A + B ® Products
rate of reaction can be expressed as
a /
AB
Rate
Z
e E
RT
−
=
(3 23)
where ZAB represents the collision frequency of reactants, A and B
and e
-Ea /RT represents the fraction of molecules with energies equal to
or greater than Ea Comparing (3 23) with Arrhenius equation, we can
say that A is related to collision frequency |
1 | 3455-3458 | 23)
where ZAB represents the collision frequency of reactants, A and B
and e
-Ea /RT represents the fraction of molecules with energies equal to
or greater than Ea Comparing (3 23) with Arrhenius equation, we can
say that A is related to collision frequency Equation (3 |
1 | 3456-3459 | Comparing (3 23) with Arrhenius equation, we can
say that A is related to collision frequency Equation (3 23) predicts the value of rate constants fairly
accurately for the reactions that involve atomic species or simple
molecules but for complex molecules significant deviations are
observed |
1 | 3457-3460 | 23) with Arrhenius equation, we can
say that A is related to collision frequency Equation (3 23) predicts the value of rate constants fairly
accurately for the reactions that involve atomic species or simple
molecules but for complex molecules significant deviations are
observed The reason could be that all collisions do not lead to the
formation of products |
1 | 3458-3461 | Equation (3 23) predicts the value of rate constants fairly
accurately for the reactions that involve atomic species or simple
molecules but for complex molecules significant deviations are
observed The reason could be that all collisions do not lead to the
formation of products 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 |
1 | 3459-3462 | 23) predicts the value of rate constants fairly
accurately for the reactions that involve atomic species or simple
molecules but for complex molecules significant deviations are
observed The reason could be that all collisions do not lead to the
formation of products 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 For
example,
formation
of
methanol from bromoethane depends
upon the orientation of reactant
molecules
as
shown
in
Fig |
1 | 3460-3463 | The reason could be that all collisions do not lead to the
formation of products 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 For
example,
formation
of
methanol from bromoethane depends
upon the orientation of reactant
molecules
as
shown
in
Fig 3 |
1 | 3461-3464 | 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 For
example,
formation
of
methanol from bromoethane depends
upon the orientation of reactant
molecules
as
shown
in
Fig 3 12 |
1 | 3462-3465 | For
example,
formation
of
methanol from bromoethane depends
upon the orientation of reactant
molecules
as
shown
in
Fig 3 12 The proper orientation of
reactant molecules lead to bond
formation
whereas
improper
orientation
makes
them
simply
bounce back and no products are
formed |
1 | 3463-3466 | 3 12 The proper orientation of
reactant molecules lead to bond
formation
whereas
improper
orientation
makes
them
simply
bounce back and no products are
formed To account for effective collisions,
another factor P, called the probability
or steric factor is introduced |
1 | 3464-3467 | 12 The proper orientation of
reactant molecules lead to bond
formation
whereas
improper
orientation
makes
them
simply
bounce back and no products are
formed To account for effective collisions,
another factor P, called the probability
or steric factor is introduced It takes into account the fact that in a
collision, molecules must be properly oriented i |
1 | 3465-3468 | The proper orientation of
reactant molecules lead to bond
formation
whereas
improper
orientation
makes
them
simply
bounce back and no products are
formed To account for effective collisions,
another factor P, called the probability
or steric factor is introduced It takes into account the fact that in a
collision, molecules must be properly oriented i e |
1 | 3466-3469 | To account for effective collisions,
another factor P, called the probability
or steric factor is introduced It takes into account the fact that in a
collision, molecules must be properly oriented i e ,
a /
AB
Rate
Z
e E
RT
P
−
=
Thus, in collision theory activation energy and proper orientation of
the molecules together determine the criteria for an effective collision
and hence the rate of a chemical reaction |
1 | 3467-3470 | It takes into account the fact that in a
collision, molecules must be properly oriented i e ,
a /
AB
Rate
Z
e E
RT
P
−
=
Thus, in collision theory activation energy and proper orientation of
the molecules together determine the criteria for an effective collision
and hence the rate of a chemical reaction Collision theory also has certain drawbacks as it considers atoms/
molecules to be hard spheres and ignores their structural aspect |
1 | 3468-3471 | e ,
a /
AB
Rate
Z
e E
RT
P
−
=
Thus, in collision theory activation energy and proper orientation of
the molecules together determine the criteria for an effective collision
and hence the rate of a chemical reaction Collision theory also has certain drawbacks as it considers atoms/
molecules to be hard spheres and ignores their structural aspect You
will study details about this theory and more on other theories in your
higher classes |
1 | 3469-3472 | ,
a /
AB
Rate
Z
e E
RT
P
−
=
Thus, in collision theory activation energy and proper orientation of
the molecules together determine the criteria for an effective collision
and hence the rate of a chemical reaction Collision theory also has certain drawbacks as it considers atoms/
molecules to be hard spheres and ignores their structural aspect You
will study details about this theory and more on other theories in your
higher classes * Threshold energy = Activation Energy + energy possessed by reacting species |
1 | 3470-3473 | Collision theory also has certain drawbacks as it considers atoms/
molecules to be hard spheres and ignores their structural aspect You
will study details about this theory and more on other theories in your
higher classes * Threshold energy = Activation Energy + energy possessed by reacting species Fig |
1 | 3471-3474 | You
will study details about this theory and more on other theories in your
higher classes * Threshold energy = Activation Energy + energy possessed by reacting species Fig 3 |
1 | 3472-3475 | * Threshold energy = Activation Energy + energy possessed by reacting species Fig 3 12: Diagram showing molecules having proper and
improper orientation
Rationalised 2023-24
84
Chemistry
Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
3 |
1 | 3473-3476 | Fig 3 12: Diagram showing molecules having proper and
improper orientation
Rationalised 2023-24
84
Chemistry
Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
3 7
What will be the effect of temperature on rate constant |
1 | 3474-3477 | 3 12: Diagram showing molecules having proper and
improper orientation
Rationalised 2023-24
84
Chemistry
Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
3 7
What will be the effect of temperature on rate constant 3 |
1 | 3475-3478 | 12: Diagram showing molecules having proper and
improper orientation
Rationalised 2023-24
84
Chemistry
Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
3 7
What will be the effect of temperature on rate constant 3 8
The rate of the chemical reaction doubles for an increase of 10K in absolute
temperature from 298K |
1 | 3476-3479 | 7
What will be the effect of temperature on rate constant 3 8
The rate of the chemical reaction doubles for an increase of 10K in absolute
temperature from 298K Calculate Ea |
1 | 3477-3480 | 3 8
The rate of the chemical reaction doubles for an increase of 10K in absolute
temperature from 298K Calculate Ea 3 |
1 | 3478-3481 | 8
The rate of the chemical reaction doubles for an increase of 10K in absolute
temperature from 298K Calculate Ea 3 9
The activation energy for the reaction
2 HI(g) ® H2 + I2 (g)
is 209 |
1 | 3479-3482 | Calculate Ea 3 9
The activation energy for the reaction
2 HI(g) ® H2 + I2 (g)
is 209 5 kJ mol–1 at 581K |
1 | 3480-3483 | 3 9
The activation energy for the reaction
2 HI(g) ® H2 + I2 (g)
is 209 5 kJ mol–1 at 581K Calculate the fraction of molecules of reactants
having energy equal to or greater than activation energy |
1 | 3481-3484 | 9
The activation energy for the reaction
2 HI(g) ® H2 + I2 (g)
is 209 5 kJ mol–1 at 581K Calculate the fraction of molecules of reactants
having energy equal to or greater than activation energy Summary
Summary
Summary
Summary
Summary
Chemical kinetics is the study of chemical reactions with respect to reaction
rates, effect of various variables, rearrangement of atoms and formation of
intermediates |
1 | 3482-3485 | 5 kJ mol–1 at 581K Calculate the fraction of molecules of reactants
having energy equal to or greater than activation energy Summary
Summary
Summary
Summary
Summary
Chemical kinetics is the study of chemical reactions with respect to reaction
rates, effect of various variables, rearrangement of atoms and formation of
intermediates The rate of a reaction is concerned with decrease in concentration
of reactants or increase in the concentration of products per unit time |
1 | 3483-3486 | Calculate the fraction of molecules of reactants
having energy equal to or greater than activation energy Summary
Summary
Summary
Summary
Summary
Chemical kinetics is the study of chemical reactions with respect to reaction
rates, effect of various variables, rearrangement of atoms and formation of
intermediates The rate of a reaction is concerned with decrease in concentration
of reactants or increase in the concentration of products per unit time It can
be expressed as instantaneous rate at a particular instant of time and average
rate over a large interval of time |
1 | 3484-3487 | Summary
Summary
Summary
Summary
Summary
Chemical kinetics is the study of chemical reactions with respect to reaction
rates, effect of various variables, rearrangement of atoms and formation of
intermediates The rate of a reaction is concerned with decrease in concentration
of reactants or increase in the concentration of products per unit time It can
be expressed as instantaneous rate at a particular instant of time and average
rate over a large interval of time A number of factors such as temperature,
concentration of reactants, catalyst, affect the rate of a reaction |
1 | 3485-3488 | The rate of a reaction is concerned with decrease in concentration
of reactants or increase in the concentration of products per unit time It can
be expressed as instantaneous rate at a particular instant of time and average
rate over a large interval of time A number of factors such as temperature,
concentration of reactants, catalyst, affect the rate of a reaction Mathematical
representation of rate of a reaction is given by rate law |
1 | 3486-3489 | It can
be expressed as instantaneous rate at a particular instant of time and average
rate over a large interval of time A number of factors such as temperature,
concentration of reactants, catalyst, affect the rate of a reaction Mathematical
representation of rate of a reaction is given by rate law It has to be determined
experimentally and cannot be predicted |
1 | 3487-3490 | A number of factors such as temperature,
concentration of reactants, catalyst, affect the rate of a reaction Mathematical
representation of rate of a reaction is given by rate law It has to be determined
experimentally and cannot be predicted Order of a reaction with respect to
a reactant is the power of its concentration which appears in the rate law
equation |
1 | 3488-3491 | Mathematical
representation of rate of a reaction is given by rate law It has to be determined
experimentally and cannot be predicted Order of a reaction with respect to
a reactant is the power of its concentration which appears in the rate law
equation The order of a reaction is the sum of all such powers of concentration
of terms for different reactants |
1 | 3489-3492 | It has to be determined
experimentally and cannot be predicted Order of a reaction with respect to
a reactant is the power of its concentration which appears in the rate law
equation The order of a reaction is the sum of all such powers of concentration
of terms for different reactants Rate constant is the proportionality factor in
the rate law |
1 | 3490-3493 | Order of a reaction with respect to
a reactant is the power of its concentration which appears in the rate law
equation The order of a reaction is the sum of all such powers of concentration
of terms for different reactants Rate constant is the proportionality factor in
the rate law Rate constant and order of a reaction can be determined from rate
law or its integrated rate equation |
1 | 3491-3494 | The order of a reaction is the sum of all such powers of concentration
of terms for different reactants Rate constant is the proportionality factor in
the rate law Rate constant and order of a reaction can be determined from rate
law or its integrated rate equation Molecularity is defined only for an elementary
reaction |
1 | 3492-3495 | Rate constant is the proportionality factor in
the rate law Rate constant and order of a reaction can be determined from rate
law or its integrated rate equation Molecularity is defined only for an elementary
reaction Its values are limited from 1 to 3 whereas order can be 0, 1, 2, 3 or
even a fraction |
1 | 3493-3496 | Rate constant and order of a reaction can be determined from rate
law or its integrated rate equation Molecularity is defined only for an elementary
reaction Its values are limited from 1 to 3 whereas order can be 0, 1, 2, 3 or
even a fraction Molecularity and order of an elementary reaction are same |
1 | 3494-3497 | Molecularity is defined only for an elementary
reaction Its values are limited from 1 to 3 whereas order can be 0, 1, 2, 3 or
even a fraction Molecularity and order of an elementary reaction are same Temperature dependence of rate constants is described by Arrhenius equation
(k = Ae–Ea/RT) |
1 | 3495-3498 | Its values are limited from 1 to 3 whereas order can be 0, 1, 2, 3 or
even a fraction Molecularity and order of an elementary reaction are same Temperature dependence of rate constants is described by Arrhenius equation
(k = Ae–Ea/RT) Ea corresponds to the activation energy and is given by the
energy difference between activated complex and the reactant molecules, and A
(Arrhenius factor or pre-exponential factor) corresponds to the collision frequency |
1 | 3496-3499 | Molecularity and order of an elementary reaction are same Temperature dependence of rate constants is described by Arrhenius equation
(k = Ae–Ea/RT) Ea corresponds to the activation energy and is given by the
energy difference between activated complex and the reactant molecules, and A
(Arrhenius factor or pre-exponential factor) corresponds to the collision frequency The equation clearly shows that increase of temperature or lowering of Ea will
lead to an increase in the rate of reaction and presence of a catalyst lowers the
activation energy by providing an alternate path for the reaction |
1 | 3497-3500 | Temperature dependence of rate constants is described by Arrhenius equation
(k = Ae–Ea/RT) Ea corresponds to the activation energy and is given by the
energy difference between activated complex and the reactant molecules, and A
(Arrhenius factor or pre-exponential factor) corresponds to the collision frequency The equation clearly shows that increase of temperature or lowering of Ea will
lead to an increase in the rate of reaction and presence of a catalyst lowers the
activation energy by providing an alternate path for the reaction According to
collision theory, another factor P called steric factor which refers to the orientation
of molecules which collide, is important and contributes to effective collisions,
thus, modifying the Arrhenius equation to
a /
ZAB
e E
RT
k
P
|
1 | 3498-3501 | Ea corresponds to the activation energy and is given by the
energy difference between activated complex and the reactant molecules, and A
(Arrhenius factor or pre-exponential factor) corresponds to the collision frequency The equation clearly shows that increase of temperature or lowering of Ea will
lead to an increase in the rate of reaction and presence of a catalyst lowers the
activation energy by providing an alternate path for the reaction According to
collision theory, another factor P called steric factor which refers to the orientation
of molecules which collide, is important and contributes to effective collisions,
thus, modifying the Arrhenius equation to
a /
ZAB
e E
RT
k
P
Rationalised 2023-24
85
Chemical Kinetics
3 |
1 | 3499-3502 | The equation clearly shows that increase of temperature or lowering of Ea will
lead to an increase in the rate of reaction and presence of a catalyst lowers the
activation energy by providing an alternate path for the reaction According to
collision theory, another factor P called steric factor which refers to the orientation
of molecules which collide, is important and contributes to effective collisions,
thus, modifying the Arrhenius equation to
a /
ZAB
e E
RT
k
P
Rationalised 2023-24
85
Chemical Kinetics
3 1
From the rate expression for the following reactions, determine their
order of reaction and the dimensions of the rate constants |
1 | 3500-3503 | According to
collision theory, another factor P called steric factor which refers to the orientation
of molecules which collide, is important and contributes to effective collisions,
thus, modifying the Arrhenius equation to
a /
ZAB
e E
RT
k
P
Rationalised 2023-24
85
Chemical Kinetics
3 1
From the rate expression for the following reactions, determine their
order of reaction and the dimensions of the rate constants (i) 3NO(g) ® N2O (g)
Rate = k[NO]2
(ii) H2O2 (aq) + 3I– (aq) + 2H+ ® 2H2O (l) +
3I
Rate = k[H2O2][I-]
(iii) CH3CHO (g) ® CH4 (g) + CO(g)
Rate = k [CH3CHO]3/2
(iv) C2H5Cl (g) ® C2H4 (g) + HCl (g)
Rate = k [C2H5Cl]
3 |
1 | 3501-3504 | Rationalised 2023-24
85
Chemical Kinetics
3 1
From the rate expression for the following reactions, determine their
order of reaction and the dimensions of the rate constants (i) 3NO(g) ® N2O (g)
Rate = k[NO]2
(ii) H2O2 (aq) + 3I– (aq) + 2H+ ® 2H2O (l) +
3I
Rate = k[H2O2][I-]
(iii) CH3CHO (g) ® CH4 (g) + CO(g)
Rate = k [CH3CHO]3/2
(iv) C2H5Cl (g) ® C2H4 (g) + HCl (g)
Rate = k [C2H5Cl]
3 2
For the reaction:
2A + B ® A2B
the rate = k[A][B]2 with k = 2 |
1 | 3502-3505 | 1
From the rate expression for the following reactions, determine their
order of reaction and the dimensions of the rate constants (i) 3NO(g) ® N2O (g)
Rate = k[NO]2
(ii) H2O2 (aq) + 3I– (aq) + 2H+ ® 2H2O (l) +
3I
Rate = k[H2O2][I-]
(iii) CH3CHO (g) ® CH4 (g) + CO(g)
Rate = k [CH3CHO]3/2
(iv) C2H5Cl (g) ® C2H4 (g) + HCl (g)
Rate = k [C2H5Cl]
3 2
For the reaction:
2A + B ® A2B
the rate = k[A][B]2 with k = 2 0 × 10–6 mol–2 L2 s–1 |
1 | 3503-3506 | (i) 3NO(g) ® N2O (g)
Rate = k[NO]2
(ii) H2O2 (aq) + 3I– (aq) + 2H+ ® 2H2O (l) +
3I
Rate = k[H2O2][I-]
(iii) CH3CHO (g) ® CH4 (g) + CO(g)
Rate = k [CH3CHO]3/2
(iv) C2H5Cl (g) ® C2H4 (g) + HCl (g)
Rate = k [C2H5Cl]
3 2
For the reaction:
2A + B ® A2B
the rate = k[A][B]2 with k = 2 0 × 10–6 mol–2 L2 s–1 Calculate the initial
rate of the reaction when [A] = 0 |
1 | 3504-3507 | 2
For the reaction:
2A + B ® A2B
the rate = k[A][B]2 with k = 2 0 × 10–6 mol–2 L2 s–1 Calculate the initial
rate of the reaction when [A] = 0 1 mol L–1, [B] = 0 |
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