CHEMISTRY MASTER: +2 UNIT 6 PAGE- 3

COLLISION THEORY OF REACTION RATE

In order to explain the effect of temperature on reaction rates, a theory known as Collision Theory of Reaction Rates was proposed. The salient features of the theory are as follows :

(i) A reaction occurs only when the reactant molecules encounter together i.e., undergo collisions with one another.

(ii) All collisions between the reacting molecules are not effective in producing a chemical change. Only a fraction of total number of collisions are effective and lead to a reaction resulting in the formation of products.

(iii) The collisions between the reacting molecules are effective only when they acquire a definite amount of energy. The minimum amount of energy which must be possessed by the reacting molecules to make effective collisions is called threshold energy.

Effective collisions are those collisions which lead to the formation of products. The number of effective collisions are governed by the following two factors.

(i) Energy barrier : The collisions between the reacting molecules cannot be effective until they possess a definite minimum amount of energy called threshold energy. This means that the collisions between two reacting molecules will be effective and will lead to the formation of products only when they possess energy greater than or equal to the threshold energy. The distribution of energy among colliding molecules at a particular temperature T is shown in Fig.

The distribution of energy among colliding molecules at temperature T.

The shaded area abcd represents the fraction of those molecules whose energy is greater than or equal to the threshold energy and are capable of making effective collisions.

(ii) Orientation barrier : The reactant molecules must collide with favourable orientation in order to facilitate the breaking of old bonds and formation of new bonds. The correct orientation ensures direct contact between atoms of the molecules involved in collisions and makes collisions effective. For example, let us consider the following reaction.

A₂ + B₂ ® 2 AB

The orientation of molecules A₂ and B₂ leading to effective collisions are shown in Fig.

Orientation of reacting molecules A₂and B₂ leading to ineffective collisions.

ACTIVATION ENERGY

An energy barrier exists for every reaction and reactants can change into products only when they acquire sufficient energy to overcome this barrier. This barrier corresponds to the threshold energy of the reaction. Only those collisions in which reactant molecules possesses energy greater than or equal to the threshold energy , lead to the formation of products. This is why energy in one form or the other from an external source has to be supplied to the reactants in order to initiate a chemical reaction. For example, a mixture of oxygen and a fuel does not burn unless a flame is applied. The flame provides necessary energy to cross the energy barrier and the fuel starts burning.

Thus, reactants need a definite amount of energy to cross the energy barrier and to take part in the reaction leading to the formation of products. This energy is called Activation Energy and may be defined as follows :

Activation energy may be defined as the excess energy that the react molecules (having energy less than the threshold energy) must acquire in order to cross the energy barrier and to change into products.

Thus,

Activation Energy = Threshold energy - Average energy possessed by :

The activation energy is diagrammatically shown in Fig.

Illustration of Activation Energy and Energy Barrier in a Reaction

Each reaction has a definite value of activation energy(E_a). The value of activation energy decides the fraction of total number of collisions which are effective. When the activation energy of a reaction is low, a large number of molecules are able to cross the energy barrier and change into products. Consequently, the rate of the reaction is high. On the other hand, the reaction having higher value of activation energy proceeds at a slow rate.

Thus,

For fast reactions : activation energy are low.

For slow reactions : activation energies are high.

ACTIVATED COMPLEX OR TRANSITION STATE

A chemical reaction involves the breaking of the old bonds and the formation of new bonds. The breaking of bonds involves absorption of energy while during the formation of bonds energy is released. A chemical reaction occurs only when the reactants possess sufficient activation energy to cross the energy barrier.

When two molecules having necessary energy of activation approach each other, they attain an intermediate configuration before they change into products. This configuration possesses a higher energy as compared to the sum of energies of reactants. The intermediate configuration corresponding to maximum potential energy of the system is called activated complex or transition state. The activated complex then changes into products and excess energy is released.

The formation of activated complex and the course of reaction can be understood by considering the following reaction :

A + B - C ® A ….B….C ® A-B + C

Reactants Activated Complex Products

(Initial state) (Final state)

A plot of the change in potential energy of the system, as a function of the distance between atoms A and B (this distance can be termed as course of reaction or progress of reaction) is shown in Fig.

The formation of activated complex during the course of reaction A + B - C ® A-B + C

The energy of activation Ea of the reaction is given by :

Ea = E_{activated Complex}- E_reactants

and the energy change in the reaction DE is given by

DE = E_products - E_reactants

where E_{activated Complex ,} E_reactantss , E_products represent the potential energies of activated complex, reactants and products respectively.

Arrhenius Equation

The rate constants for most reactions increase with increase of temperature. In order to determine a quantitative relationship between rate constant and temperature, Arrhenius proposed the following equation :

………(1)

This equation is called Arrhenius equation, in which A is a constant known as frequency factor. The factor is related to the number of binary collisions per second per litre. Ea is the activation energy which represents the minimum energy that the reacting molecules must possess before undergoing a reaction. T is the absolute temperature and R the gas constant. The two quantities ‘A’ and ‘Ea’ are collectively called the Arrhenius parameters.

Taking logarithm, Eq.(1) may be written as :

Converting to common logarithm, we get :

When log k is plotted against 1/T , we get a straight line .

Plot of log k vs (1/T)

The intercept of this line is equal to log A and the slope is equal to

Thus, Ea = - ( 2.303 R) Slope

Alternately, the value of E_a can also be determined by measuring the rate constant of a reaction at two different temperatures. If k₁ and k₂ are the rate constants for the reaction at temperatures T₁ and T₂, then :

Subtracting Eq.(2) from Eq.(3), we get :

……(4)

Problems

5. Show that the time required for the completion of three-fourth of a first order reaction is twice the time required for completion of the half reaction.

6. The rate constant of the first order decomposition of N₂O₅ at 320 K is 3 x 10^-² min^-¹. Calculate half-life of the reaction.

7. Thermal decomposition of a compound is of first order. If 50% of the compound is decomposed in 120 minutes, how long will it take for 90% of the compound to decompose ?

8. Calculate the half-life of a first order reaction, where specific rate constant is :

a) 200 s^-¹. b) 2 min^-¹ c) 5 year^-¹

9. A first order reaction is 90% complete in 50 minutes. Calculate the half-life of the reaction.

10. Rate constant of a first order reaction is 5.78 x 10^-⁵ s^-¹. What percentage of the initial reactant will react in 10 hours.

11. A reaction that is first order with respect to reactant A has a rate constant 6 min^-¹. If we start with[A] = 0.5 mol L^-¹, when would [A] reach a value of 0.05 mol L^-¹ ?

12. If the half life of a first order with respect to A is 2 min, how long will it take [A] to reach :

a) 25% of its initial concentration.

b) 10% of its initial concentration ?

13. The half-life of a first order reaction is 100 sec. Calculate its rate constant.

14. A first order reaction has a rate constant 10^-² sec^-¹. Calculate the half life period of this reaction.

15. The rate constant for a first order reaction is 1.54 x 10^-³ s^-¹. Calculate its half time.

16. The half life period for the homogeneous gaseous reaction SO₂Cl₂ ® SO₂ + Cl₂ , which obeys first order kinetics, is 8.0 minutes. How long will it take for the concentration of SO₂Cl₂ to be reduced to 1 % of the initial value ?

17. In an enzyme solution , sucrose undergoes fermentation. If 0.1 M solution of sucrose is reduced to 0.05 M in 10 hours and to 0.025 M in 20 hours, what is the order of the reaction and what is the rate constant ?

18. Calculate the number of a-particles emitted per second by one gram of pure ThO₂ if the half-life of ²³²Th is 1.39 x 10¹⁰ years.

19. The rate for decomposition of NH₃ on platinum surface is zero order. What are the rate of production of N₂ and H₂ if k = 2.5 x 10^-⁴ M s^-¹.

20. The decomposition of dimethyl ether leads to the formation of CH₄, H₂ and CO and the reaction rate is given by : Rate = k [CH₃OCH₃]^3/2 . The rate of reaction is followed by an increase in pressure in a closed vessel and the rate can also be expressed in terms of the partial pressure of dimethyl ether ie., rate = k [ P_CH3OCH3]^3/2. If the pressure is measured in bar and time in minutes, then what are the units of rate and rate constants ?

21. A reaction is second order with respect to a reactant. How is the rate of reaction affected if the concentration of the reactant is : (i) doubled (ii) reduced to ½ ?

22. In a pseudo first order hydrolysis of ester in water the following results were obtained :

Time (s)	0	30	60	90
[Ester] M	0.55	0.31	0.17	0.085

i) Calculate the average rate of the reaction between the time interval 30 to 60 seconds.

ii) Calculate the pseudo first order rate constant for the hydrolysis of ester.

23. A reaction is first order in A and second order in B .

i) Write the differential rate equation.

ii) How is the rate affected when concentration of B is tripled ?

iii) How is the rate affected when the concentration of both A and B is doubled ?

24. In a reaction between A and B , the initial rate of reaction was measured for different initial concentrations of A and B as given below :

[A] M	0.20	0.20	0.40
[B] M	0.30	0.10	0.05
r_o M /s	5.07 x 10^-⁵	5.07 x 10^-⁵	7.6 x 10^-⁵

What is the order of the reaction with respect to A and B ?

25. Reaction between NO₂ and F₂to NO₂F takes place by the following mechanism :

Write the rate expression for the reaction.

26. The following results have been obtained during the kinetic studies of the reaction :

2 A + B ® C + D

Expt	[A] M	[B] M	Initial rate of formation of D M/min
I	0.1	0.1	6 x 10^-³
II	0.3	0.2	7.2 x 10^-²
III	0.3	0.4	2.88 x 10^-¹
IV	0.4	0.1	2.40 x 10^-²

Determine the rate law and rate constant for the reaction.

27. The reaction between A and B is first order with respect A and zero order with respect to B. Fill in the blanks in the following TABLE.

Expt	[A] M	[B] M	Initial rate : M/min
I	0.1	0.1	2 x 10^-²
II	……..	0.2	4.0 x 10^-²
III	0.4	0.4	………..
IV	………	0.2	2.0 x 10^-²

28. The rate constant for a first order reaction is 60 s^-¹. How much time will take to reduce the initial concentration of the reactant to its 1/16 th value ?

29. The rates of most reactions double when their temperature is raised from 298 k to 308 K. Calculate their activation energy.

30. The value of rate constant for the decomposition of nitrogen pentoxide,

N₂O₅(g) ® N₂O₄(g) + (1/2) O₂(g)

is 3.46 x 10^-⁵ at 25⁰C and 4.87 x 10^-³ at 65⁰C.Calculate the activation energy for such a reaction.

31. Calculate the activation energy of a reaction whose rate constant is tripled by 10°C rise in temperature in the vicinity of 27°C , what will be the activation energy of the reaction ?

32. In general , the rate of a chemical reaction doubles with every 10°C rise in temperature. If the reaction is carried out in the vicinity of 27°C , what will be the activation energy of the reaction ?

33. Can the activation energy of a reaction be zero or negative ?

34. What value of the rate constant is predicted by Arrhenius equation if T®∞ ? Is this value physically reasonable ?

35. The half life for radioactivity of ¹⁴C is 5730 y. An archaeological artefact contained wood had only 80% of the ¹⁴C found in a living tree. Estimate the age of the sample.

36. One of the hazards of nuclear explosion is the generation of ⁹⁰Sr and its subsequent incorporation in bones. This nucleide has a half life of 28.1 years. Suppose 1 microgram was absorbed by a new born child, how much ⁹⁰Sr will remain in his bones after 20 years ?

56. If the rate constant for a reaction is 1.6 x 10^-⁵ and 6.36 x 10^-³ s^-¹ at 600 K and 700K respectively, calculate energy of activation for the reaction.

57. The slope of the line in the graph of log₁₀k (k = rate constant) versus 1/T for a reaction is - 5841 . Calculate the energy of activation for this reaction.

58. The rate constant of a first order reaction becomes 5 times when temperature is raised from 350 K to 400 K. Calculate activation energy for the reaction.

MOLECULARITY OF A REACTION

According to the collision theory, a reaction takes place when the molecules of the reactants undergo collisions with one another. The number of molecules which undergo simultaneous collisions decides the molecularity of a reaction. The number of molecules which undergo simultaneous collisions decide the molecularity of a reaction. It may be defined as follows:

The number of reacting species (atoms, ions, molecules) which must collide simultaneously in order to bring about a chemical reaction is called molecularity of a reaction.

The molecularity of a reaction is a whole number and may have values 1, 2, 3, … etc. When only one molecule of the reactant is involved in the reaction, the molecularity of the reaction is 1 and the reaction is said to be a unimolecular reaction. When two molecules of reactants collide together to bring about a chemical change, the molecularity of the reaction is taken to be equal to 2 and the reaction is said to be a bimolecular reaction. In case three molecules of the reactants collide simultaneously to bring about a chemical change, the molecularity of the reaction is 3 and the reaction is said to be a trimolecular or termolecular reaction. Reactions having molecularity more than three are very rare because the chances of simultaneous collisions of more than three molecules are remote.

Molecularity and Elementary reactions

The simple chemical reactions which occur only in one step are called elementary reactions. The molecularity of such a reaction is decided by the number of molecules taking part in it. The molecularity of an elementary reaction is equal to the number of reacting species (atoms, ions or molecules) as represented by the balanced chemical equation of the reaction.

For example,

H₂O₂ ® H₂O + ½ O₂

( molecularity = 1, unimolecular reaction)

2 HI(g) ® H₂ + I₂

( molecularity = 2, bimolecular reaction)

2 NO(g) + O₂ ® 2 NO₂(g)

( molecularity = 3, trimolecular reaction)

Molecularity of complex reactions

The reactions which occur in two or steps are complex reactions. The molecularity of complex reactions cannot be decided by the stiochiometry of their balanced equations. This is because their balanced equations may possess a large number of reactant molecules. For example,

2 NO + 2 H₂ ® N₂+ 2 H₂O

2 FeCl₃ + 6 KI ® 2 FeI₂ + 6 KCl + I₂

2 KMnO₄ + 16 HCl ® 2 KCl + 2 MnCl₂ + 5 Cl₂ + 8 H₂O

If these reactions occur in a single step, a large number of molecules (as indicated by their stoichiometry) of reactants are required to collide simultaneously. For example, if the reaction between KMnO₄ and HCl occurs in a single step, the reaction can occur only when 2 molecules of KMnO₄ and 16 molecules of HCl i.e., a total of 18 molecules collide simultaneously. This is not feasible. The possibility of simultaneous collisions of three or more molecules is very small. In fact, the chances of simultaneous collisions of reacting molecules go on decreasing with increase in the number of molecules i.e., with increase in molecularity. This is why only a few trimolecular reactions are known. The reactions with higher molecularity are quite rare.

Therefore , in order to determine the molecularity of complex reactions, it is suggested that the complex reactions proceed through a series of steps, each involving one, two, or at the most three molecules. Each step in a complex reaction is an elementary reaction. Thus a complex reaction is supposed to proceed through several elementary reactions. Each elementary step in a complex reaction has its own rate. Some elementary steps may be slow while others may be fast. The overall rate of a complex reaction is governed by the rate of the slowest elementary step. The slowest elementary step is therefore termed as rate determining step. The number of reacting species taking part in the lowest elementary step decides the molecularity of the reaction. Thus, the molecularity of a complex reaction may be defined as follows.

The number of reacting species (atoms, ions or molecules) taking part in the slowest elementary step (i.e., rate determining step) of a complex reaction is called the molecularity of the complex reaction.

For example, the reaction between nitric oxide and hydrogen is a complex reaction.

2 NO + H₂ ® N₂ + 2 H₂O

This reaction is supposed to take place in the following two steps.

Step 1 : 2 NO + H₂ ® N₂ + H₂O₂ (slow)

Step 2 : H₂O₂ + H₂ ® 2 H₂O (fast)

Since step 1 is slow, it is the rate determining step. The chemical reaction corresponding to step 1 involves three molecules of reactants. Hence, the molecularity of the slowest step is 3 and is actually taken as the molecularity of the overall reaction.

It is to be noted that although the molecularity and order of a reaction may be similar in some cases, yet the two are quite different from each other. The main points of difference between the molecularity and the order of a reaction are summarized in TABLE.

Difference between order and molecularity

*Molecularity*	*Order*
1. It is the number of reacting species undergoing simultaneous collision in the reaction	1. It is the sum of the powers of the concentration terms in the rate law expression.
2. It is a theoretical concept.	2. It is determined experimentally.
3. It can have integral values only.	3. It can have fractional values also.
4. It cannot be zero.	4. It can be zero.
5. It does not tell us anything about the mechanism of a reaction.	5. It tells us about the slowest step in the mechanism and hence it gives some clue about mechanism of the reaction.

Pseudo Molecular reactions

For elementary reactions , the molecularity and order are usually the same. However, there are several first order reactions in which moleculaity differs from the order.

The first order reactions having moleculaity greater than one are called pseudo-unimolecular reactions.

A pseudo-unimolecular reaction is obtained when one of the reactants is presents is present in large excess. The reactant present in large excess do not contribute to the rate of reaction. Its concentration remains almost constant during the course of reaction and therefore the rate of reaction does not depend upon its concentration. For example,

The moleculaity of each reaction is 2. Both the reactions have found to be of first order. The rate of the first reaction is found to vary with the concentration of ethylacetate while that of the second one is found to vary with the concentration of sucrose as is clear from their rate law. This is because , in both the reactions, water is present is in large excess and its concentration does not change appreciably with the progresss of the reaction in each case. Therefore, the rate of reaction depends only upon the concentration of ester in the first case and upon that of sucrose in the second case. Hence both are pseudo unimolecular reactions.

A reaction of higher order follows the kinetics of first under special conditions. Such reactions are called pseudo first order reactions. For example, consider a hypothetical reaction:

A + B ® AB

Now if concentration of B is large enough so that it does not change appreciably during the reaction, then its rate will depend on the single concentration term and it will follow the kinetics of the first order. Acidic hydrolysis of ethyl acetate is a common example of this type.

CH₃COOC₂H₅ + H₂O ® CH₃COOH + C₂H₅OH

Here the rate of reaction is given by the expression :

Rate = k [CH₃COOC₂H₅]

Since the concentration of H₂O is quite large and does not change appreciably, therefore it does not appear in the rate law.

MECHANISM OF REACTION

The first step for determination of the mechanism of a complex reaction requires the determination of the stoichiometry of the reaction i.e., we require the information about the number of moles of each reactant consumed to the number of moles of each of final products. In some reactions, intermediates are produced which accumulate during the early period of the reaction, reach to a maximum concentration and then react to give the final products. The determination of mechanism of a reaction is not an easy task and requires the experience and ingenuity of the scientist. The necessary condition for a mechanism to fulfil is that it must lead to correct rate law. However , this does not guarantee that the mechanism is true. The kinetic study has to be supplimented by various techniques to verify the desirability of including certain elementary steps in the reaction mechanism. We shall discuss the mechanism of some simple reactions.

1. Reactions involving two first order consecutive steps

In such reactions, a reaction takes place in two steps both of which are first order.

Now I is produced by step (i) and consumed in step (ii) . The intermediate I accumulates and reaches a maximum after which it decays to zero concentration and is converted into the final product as shown in Fig.

Concentration profiles of R (reactants) , I intermediate and P (products) as a function of time.

2. Reactions involving slow step

If a reaction takes by a sequence of steps and one of the steps is slow, then the rate determining step is the slow step. The rate of this step may be slow either due to low value of the rate constant or very low concentration of one or more of the reacting species in the elementary reaction. For example, in the reaction,

if k₁ << k₂ then I is converted into the product as soon as it is formed and we can say

Example

In the reaction,

3 ClO^- ® ClO₃^- + 2 Cl^-

Various steps are :

ClO^- + ClO^- ® ClO₂^- + Cl^- (slow step)

(rate determining step)

ClO₂^- + ClO^-® ClO₃^- + Cl^- ( fast step)

\ rate = k₁[ClO^-]²

3. Reactions for which steady state hypothesis is valid

A reaction may take place in a number of steps and may have several intermediates. In the steady state hypothesis , we assume that the intermediates are so reactive that after a brief initial period (called induction period) their concentrations rise from zero to a small value and remain constant for most of the duration of the reaction, i.e., we assume that change in concentration with time for these reactive intermediates is zero. This assumption is very helpful for deriving the rate expression for complex reactions.

4. Reactions involving intermediates in Equilibrium with the Reactants

In some reactions involving H⁺ and OH^- , there is an equilibrium established between the reactants as both forward and reverse reactions have very large rate constants. The intermediate thus formed reacts so slowly that equilibrium concentration of the intermediate is not disturbed much. For example, the displacement of C₂H₅O^- form o-hydroxy amino ethyl benzoate is catalysed by OH^-. The following mechanism has been suggested :

k₃ is much smaller than k₁ and k₂ and therefore the equilibrium concentration of the intermediate is given by :

This is a second order reaction but the overall rate constant involves all the three rate constants, or we have

QUESTIONS

1. What is meant by reaction rate ? Give its symbolic expression and units for the reaction :

CO(g) + NO(g) ® CO₂(g) + NO₂(g)

2. Distinguish between reaction rate and specific reaction rate of a reaction.

3. Define specific rate of a chemical reaction.

4. How can you determine the velocity of a reaction at a given time ?

5. How can you express the velocity of a reaction ?

6. Define and explain the term 'instantaneous rate' of a reaction.

7. For a reaction : X ® Y , the rate of a reaction can be denoted by -dX/dt or +dY/dt. State the significance of the plus and minus signs in these expressions.

8. For the reaction,

2 N₂O₅ ® 4 NO₂ + O₂

the rate of reaction can be expressed in three ways:

-d[N₂O₅]/dt = k [N₂O₅]

+d[NO₂]/dt = k'[N₂O₅]

+d[O₂] /dt = k''[N₂O₅]

Establish the relationship between k and k' and between k and k''.

9. Explain the average rate of a reaction.

10. How can you measure the rate of the following reaction ?

CH₃COOH(aq) + CH₃OH ® CH₃COOCH₃(aq) + H₂O(l)

11. For a chemical reaction , A ® B the initial concentration is 0.30 mol L^- ¹ and its concentration at the end of 30 minutes is 0.18 mol L^-¹. What is the average rate of this reaction ?

12. How can you express symbolically the rate of the following reactions ?

i) 2 N₂O₅ ® 2 N₂O₄ + O₂

ii) 4 NH₃ + 5 O₂ ® 4 NO + 6 H₂O

13. Write, in brief, on molecularity and order of a reaction.

14. What do you understand by the order and molecularity ?

15. Distinguish between order and molecularity as applied to reactions.

16. Calculate the units of :

i) first order rate constant.

ii) second order rate constant.

iii) third order rate constant.

iv) zero order rate constant.

17. With the help of the following rate expressions of the reactions, find out the overall order of the reactions and order with respect to each reactant :

i) 2 NO(g) + O₂(g) ® 2 N O₂(g)

Rate = k[NO]²[O₂]

ii) 2 N₂O(g) ® 2 N₂(g) + O₂(g)

Rate = k[N₂O]

iii) 2 NO₂(g) ® 2 NO(g) + O₂(g)

Rate = k[N O₂]²

iv) 2 NO(g) + 2 H₂(g) ® N₂(g) + 2 H₂O(g)

Rate = k[NO]²[H₂]²

v) SO₂Cl₂(g) ® S O₂(g) +Cl₂(g)

Rate = [SO₂Cl₂]

18. What is the rate of reaction and the order of the reaction, if the mechanism is :

2 NO +H₂® N₂ + H₂O₂ (slow)

followed by :

H₂O₂ + H₂® 2H₂O (fast)

19. Explain, with suitable examples, how the molecularity of a reaction is different from the order of a reaction ?

20. With the help of suitable examples, explain the terms molecularity and order of reactions.

21. What is the order of a chemical reaction ? Write the rate equation for a zero order reaction.

22. The kinetics of a reaction :

A + 2 B ®C + D

obeys the rate equation :

Rate = k [A]^x [B]^y

For it, find out :

i) order of the reaction.

ii) order of the reaction, when B is present in large

excess.

23. Give one example each of unimolecular and bimolecular reactions.

24. Write a note on pseudo-unimolecular reactions. Give examples to illustrate it.

25. Derive the equation for the rate constant for a first order reaction.

26. What would be the unit of the first order rate constant, if the concentration is expressed in moles per litre and time in seconds.

27. Show that for a first order reaction, the time required for half the change (half-life of a reaction) is independent of initial concentration.

28. Explain half-life of a reaction.

29. Derive an equation for rate constant of a first order reaction.

30. What are the characteristics of a first order reaction ?

31. How will you prove that a chemical reaction is of first order ?

32. The rate constant for a first order reaction is 0.0005 min^-1. Calculate its half-life.

33. A first order reaction takes 30 minutes for 60% of the reaction to be completed. What is the velocity constant for the reaction ?

34. Show that in case of a first order reaction, the time required for 99.9% of the reaction to take place is 10 times that required for half of the reaction.

35. Show that the time required for the completion of three-fourth of a first order reaction is twice the time required for completion of the half reaction.

36. The rate constant of first order decomposition of N₂O₅ at 320 K is 3 x 10^-² min^-¹. Calculate half-life of the reaction.

37. What is Arrhenius theory of reaction rates ?

38. Define and explain activation energy.

39. Explain the term activated complex.

40. Discuss the factors which influence rates of chemical reactions.

41. How is rate of a reaction related to the concentration of the reactants ?

42. What is the function of a catalyst in a chemical reaction ? How does it affect the rate of a reaction ?

43. Define the term activation energy. How does a catalyst affect it ? Explain with a diagram.

44. What is the significance of activation energy ? Using this concept, explain the influence of a catalyst on the rate of a reaction.

45. Sketch a labelled potential energy diagram for endothermic reaction in the presence and in absence of a catalyst.

46. For a chemical reaction, what is the effect of catalyst on the following :

i) Equilibrium constant of a reaction.

ii) Free energy change of a reaction.

iii) Activation energy of a reaction.

47. An exothermic reaction : A ® B has an activation energy of 13 kJ mol^-1 of A and energy of the reaction is 44 kJ. Find out the activation energy of the reverse reaction : B ® A.

48. How does change in temperature affect the rate of a chemical reaction ? Give the mathematical expression describing this effect.

49. Sketch a potential energy diagram which represents an endothermic reaction. Mark on the diagram activated complex, activation energy and net energy change for the reaction.

50. What is energy of activation? How is the rate constant of a reaction related to its activation energy ?

51. Describe the influence of temperature on reaction rates.

52. Why increasing the temperature increase reaction rate ?

53. Explain why rise in temperature increases the rate of a reaction.

54. Define activation energy and explain the temperature dependence of the reaction rate on its basis.

55. For 10 K rise in temperature, the rate constant is nearly doubled. Explain.

56. How can you determine experimentally the activation energy of a reaction ?

57. What is meant by 'rate controlling step' in a reaction ?

58. If a reaction : A + 2 B ® C + D proceeds in two steps :

A + B ® AB (slow)

AB + B ® C + D (fast)

which one of the two steps determines the rate of

formation of products ?

59. What are photochemical reactions ? Give examples.

60. Write notes on :

i) photosensitization.

ii) Chain reaction.

61. Write a note on the mechanism of a reaction.

62. Give some instances of the role played by photochemistry in our daily life.

63. Explain the following :

i) A lump of coal burns at moderate rate in air ; while coal dust burns explosively.

ii) Liquid bromine reacts slowly as compared to bromine vapour.

64. Can activation energy for reactions be zero ?

65. What is the main difference between a photosensitizer and a catalyst ?

+2 UNIT 6 PAGE- 3

PLUS TWO

QUESTIONS