Reaction rates by wanghonghx

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• pg 1
```									The Rate of Reaction
   Reaction rates are the rates at which reactants
disappear or products appear.
   This movie is an illustration of a reaction rate.

1
Collision Theory of
Reaction Rates
        Three basic events must occur for a
reaction to occur the atoms, molecules or
ions must:
1.    Collide.
2.    Collide with enough energy to break and form
bonds.
3.    Collide with the proper orientation for a reaction
to occur.

2
Collision Theory of
Reaction Rates
   One method to increase the number of collisions and
the energy necessary to break and reform bonds is to
heat the molecules.
   As an example, look at the reaction of methane and
oxygen:
CH 4(g)  O 2(g)  CO 2(g)  H 2 O (g)  891 kJ
   We must start the reaction with a match.
   This provides the initial energy necessary to break the
first few bonds.
   Afterwards the reaction is self-sustaining.
3
Collision Theory of
Reaction Rates
   Illustrate the proper orientation of molecules
that is necessary for this reaction.
X2(g) + Y2(g) 2 XY(g)
   Some possible ineffective collisions are :

X                    Y
Y Y               X   X           X X   Y Y
X                    Y

4
Collision Theory of
Reaction Rates
   An example of an effective collision is:

X    Y        X    Y         X       Y
+

X    Y        X    Y         X       Y

5
Collision Theory
Activation energy
enthalpy

Ea

reactants

∆Hrxn

products

time
6
Transition State
   Transition state theory postulates that
reactants form a high energy intermediate,
the transition state, which then falls apart into
the products.
   For a reaction to occur, the reactants must
acquire sufficient energy to form the transition
state.
   This energy is called the activation energy or Ea.
   Look at a mechanical analog for activation
energy                                                   7
Transition State

8
Transition State
   The distribution of molecules possessing different
energies at a given temperature is represented in
this figure.

9
Factors That Affect Reaction
Rates
    There are several factors that can influence
the rate of a reaction:
1.    The surface area of the reactants.
2.    The concentration of the reactants.
3.    The temperature of the reaction.
4.    The presence of a catalyst.
     We will look at each factor individually.

10
Surface Area
   This movie illustrates how changing the surface area
of reactants affects the rate.

11
Concentrations of Reactants
   This is a simplified representation of the
effect of different numbers of molecules in the
same volume.
   The increase in the molecule numbers is
indicative of an increase in concentration.
A(g) + B (g)  Products

A   B                  A B                   A B
B                   A B
A B                    A B                   A B
4 different possible    6 different possible   9 different possible
12
A-B collisions          A-B collisions         A-B collisions
Catalysts
   Catalysts change reaction rates by providing an
alternative reaction pathway with a different
activation energy.

13
Catalysts
   Homogeneous catalysts exist in same phase as the
reactants.
   Heterogeneous catalysts exist in different phases
than the reactants.
   Catalysts are often solids.

14
Catalysts
Number of reactants particles

Fraction of particles with E > Ea
for the uncatalized reaction

Fraction of particles
with E > Eacat.

Eacat   Ea
Kinetic energy

15
Catalysts
   Examples of commercial catalyst systems include:

 and

C8 H18g  + 25 O 2g  NiOPt16CO2g   18 H 2 O g 

2 COg  + O 2g     2 CO2g 
NiO and Pt


2 NOg     N 2g   O 2g 
NiO and Pt

Automobile catalytic converter system

16
Catalysts

17
Catalysts
   This movie shows catalytic converter chemistry
on the Molecular Scale

18
Catalysts
   A second example of a catalytic system is:


2 SO 2g   O 2g    2 SO3g 
V2 O5 or NiO/Pt

Sulfuric acid preparatio n

19
Catalysts
   A third examples of a catalytic system is:

N 2g   3 H 2g    2 NH3g 
Fe or Fe 2 O 3

Haber Process

20
Catalysts
   Look at the catalytic oxidation of CO to CO2
   Overall reaction
2 CO(g)+ O2(g) 2CO2(g)
   Absorption
CO(g) CO(surface) + O2(g)
O2(g) O2(surface)
   Activation
O2(surface)  O(surface)
   Reaction
CO(surface) +O(surface)  CO2(surface)
   Desorption
CO2(surface)  CO2(g)            21

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