Chemistry - A Molecular Science

atoms. A similar argument holds for the OCl

molecules, but OCl molecules are formed by

the reaction and then consumed, so they are called

intermediates

. Catalysts appear first

as reactants and must be added to the reaction only to be released later, while intermediates are formed in th

e reaction but consumed later.

OO^3

O--O

Cl

O

+O

+{Cl}

22

OC

l+O

2

O

--C

l

3

O

+O

+{Cl}

3

Energy

E

=^17

kJ/mol

a

E’a

DE

=-392

kJ/mol

Figure 9.8 Energy diagram for O

(g) + O(g) 3

→

2O

(g) 2

Figure 9.8 shows the reaction energy diag

ram for the reaction in the absence and

presence of Cl atoms. In the absence of Cl, the transition state of the final reaction involves an O

species, and the activation energy of the reaction is 17 kJ/mol. Such a high 4

activation energy means that this reaction occu

rs very rarely in the atmosphere. However,

the activation energies for the two reactions

involving chlorine are both very small (~2

kJ/mol), so the overall reaction occurs much

more readily in the presence of chorine

atoms. Thus, the chlorine atoms catalyze the re

action because they increase the rate of the

reaction, but they are unchanged by it. A worldwide ban on the production of CFC’s has resulted as a consequence of the discovery of these reactions in the atmosphere.

Diagrams are in the absence (blue) and in the presence (red) of catalytic chlorine atoms. The activation energy E’

is 2.1 kJ/mol. a

Note that OCl lies in a shallow well between the reactants and the products, which is typical of intermediates, while the transition states O

-O, O 3

-Cl, and O-OCl all lie at peak maxima. 3

In summary, there are three ways to

increase the rate of a reaction:

1.^

increase the concentrations of the reacta

nts to increase the frequency of collisions;

2.^

increase the temperature to increase the co

llision frequency and the fraction of collisions

with sufficient energy to form the transition state; and

3.^

add a catalyst

to decrease the activation energy.

9.11

EQUILIBRIUM AND THE EQ

UILIBRIUM CONSTANT

In Section 9.8, we discussed the thermodynamic definition of equilibrium (

G = 0), and Δ

we now turn to the kinetic definition. Consider the reaction

CH

I + OH 3

1-^

→

CH

OH + I 3

1-.

Once some product molecules have formed, they

can also collide to produce the reverse

reaction, CH

OH + I 3

1-

→

CH

I + OH 3

1-, and as the forward reaction proceeds, the

concentrations of I

1- and CH

OH increase causing the rate of the reverse reaction (R 3

= r

k[CHr

OH][I 3

1-]) to increase. Simultaneously, the concentrations of OH

1- and CH

I 3

decrease, reducing the rate of the forward

reaction. Eventually, the two rates become

equal; that is, the transition state is reached at

the same rate from both sides. At this point,

all species are being formed and consumed at

the same rate, and there is no longer any

change in concentration; the system has re

ached equilibrium.* The reaction continues at

equilibrium, but it does so in both directions

with equal rates. Equilibria in which

competing processes continue at equal rates ar

e called dynamic equilibria, which is quite

different from a static equilibrium in whic

h the competing processes stop. Double arrows

are used to represent an equilibrium in order

to show that the reaction continues in both

* The thermodynamic view is that the reaction causes concentration

changes, which change the driving forces for the forward and reverse reactions. At equilibrium the driving forces are equal. In kinetics, the concentration changes cause rate changes, and equilibrium occurs when the rates equal.

Chapter 9 Reaction Energetics

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State

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