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1. Temperature and Catalyst
Dr Omed Ameen Charmo University

2.  Concentration
 Collision number
Temperature
Surface area
Catalyst

3. Collision Theory
A model called collision theory is used to relate the properties of
particles to the rates of chemical reactions.
•According to collision theory, atoms, ions, and molecules can react to
form products when they collide if the particles have enough kinetic
energy.
•Particles that do not have enough energy to react bounce apart
unchanged when they
collide.

4. Collision Theory
Collision
theory
Effective collision Ineffective collision
Successful Collison between the
particles
An effective collision of oxygen and
hydrogen molecules produces water
molecules.
Unsuccessful Collison between the
particles
An ineffective collision of oxygen and
hydrogen molecules produces no reaction;
the reactants bounce apart unchanged

5. Collision theory states there are three things that must Happen
before a chemical reaction will take place.
1) Molecules must collide
The above is know as a “2 body collision” since only
Two reactant molecules need collide to produce
Products.

6. Collision Theory states that MOST chemical reactions
Involve 2 body collisions
So… If most reactions are only two body collisons then
How do we explain reactions with more than 2 Reactants?
For example: 2NO(g) + F2  2NOF(g)
i.e. NO + NO + F2  NOF + NOF
According to this eqn, 2 nitrogen (II) oxide molecules and a fluorine
molecule must collide at the same time to yield 2 NOF. But this looks like a
3 body collision, right?

7. The answer is that a 3 body collision DOES NOT actually
Take place. The reaction proceeds in 2 separate 2 body
collisions:
Step 1: NO(g) + F2(g)  NOF2 Intermediate Product
2 separate two body collisions
Step 2: NOF2 + NO  2NOF Product
Notice that the sum of the two separate steps equals
The chemical equation as initially written

8. 2) Be in the right orientation at collision
Orientation
When two molecules react, bonds need to be broken
And bonds need to be created
For this to happen molecules need to be in the correct Orientation
when collision occurs
For example: H2 + Cl2  2HCL
H H Cl
Cl
Poor orientation for
Formation of new bonds
H H
Cl
Cl
Bond breaks
Bond breaks
Bonds begin
To form
Good orientation
Collision
Occurs
No Rxn Takes Place

9. Example

10. Kinetic Energy
3) Sufficient kinetic energy when they collide for the reaction to occur
– in other words they must collide with sufficient velocity
Even though the orientation may be right, without enough KE there is insufficient energy
to break bonds that need to be broken before new bonds can form

11. Kinetic Energy
At a given temperature some molecules have more
energy than others (The Maxwell-Boltzman Distribution).
Even though the temperature increases (inc. KE) not
all molecules will have sufficient energy to break bonds
even if the orientation is correct.
At 100 deg there
Is a greater
Fraction of
Molecules with
Sufficient KE for
Rxn to occur
Point above
Which there is
Sufficient KE
For Rxn to
occur
Inc. T
V ( )

12. The Maxwell-Boltzmann apparatus
 Maxwell and Boltzmann performed an experiment to determine the
kinetic energy distribution of atoms
 Because all atoms of an element have roughly the same mass, the
kinetic energy of identical atoms is determined by velocity (KE= ½mv2)

13. The Maxwell-Boltzmann distribution
 The resulting disk looks like this:
Basically, if we plot the
intensity of the dots on a
graph we get a graph of
fraction of atoms/molecules
vs. kinetic energy:
Fraction of
molecules
Kinetic energy 
Molecules
hit disk first
Molecules
hit disk last

14. Why is the graph skewed?
 This curve is characteristic of all molecules
 The curve is elongated due to how atoms collide, and to the units of the
graph
 Recall all particles are in motion. An average speed will be reached.
 The graph is skewed because 0 is the lower limit, but theoretically there
is no upper limit
Same data, different
axes. E.g. v=1, KE=1
v=2, KE=4
v=3, KE=9
velocity KE
• More than that the graph is skewed because
the x-axis has units of energy not velocity

15. All three required properties
The minimum energy that
colliding particles must
have in order to react is
called the activation
energy.
You can think of the
activation energy for a
reaction as a barrier that
reactants must cross
before products can form.

16. Collision Theory
•When two reactant
particles collide, they
may form an activated
complex.
•An activated complex
is an unstable
arrangement of atoms
that forms for a
moment at the peak
of the activation-
energy barrier and is
usually short lived.

17. Transition State Theory
Addresses the limitations of collision theory in explaining
chemical reactivity
Key principle: During the transformation of reactant into
product, one or more very short-lived chemical species
form that resemble (but are different from) reactant or product;
these transitional species contain partial bonds; they are called
transition states (TS) or activated complexes.
The activation energy is used to stretch/deform specific bonds
in the reactant(s) in order to reach the transition state.
Example reaction:
CH3Br + OH- CH3OH + Br-
What might the TS look like for this substitution reaction?

18. The TS is trigonal bipyramidal; note the elongated C-Br
and C-O bonds
The proposed transition state in the reaction between CH3Br and OH-

19. Collision Theory
The activation-energy
barrier must be crossed
before reactants are
converted to products.
• The lifetime of an activated
complex is typically about 10-13
seconds.
• Its brief existence ends with the
reformation of the reactants or with
the formation of products.
• Thus, the activated complex is
sometimes called the transition
state.

20. Collision Theory
Remember: An
endothermic
reaction absorbs
heat, and an
exothermic reaction
releases heat.

21. What factor determines whether a
molecular collision results in a
reaction?
Quiz

22. Effect of Temperature
According to kinetic theory (do you remember this?) as the
temperature increases the particles in a substance move about more
quickly.
Reaction at 300C Reaction at 500C
As the temperature increases the number of collisions increases as well
as the energy of the collisions. So temperature has a big effect on the rate
of reaction. For every 100C increase the rate approximately doubles.
Higher temperature = faster reaction

23. Temperature and reaction rate
 By understanding the Maxwell-Boltzmann distribution, we can begin to
understand the two reasons why an increase in temperature causes an
increase in reaction rate
where R is the gas constant and T is the Kelvin temperature.
f = e
-Ea
RT
This fraction of molecules can be found through the
expression

24. Endo and Exothermic Reactions

25. rate = k[A]
Consider the rate law for a first order reaction:
Where is the temperature dependence?
Answer: It is embodied in the rate constant, k, that is,
k depends on the temperature at which the
reaction is conducted.
What does the T-dependence of k look like?
R-COOR’ + H2O R-COOH + R’OH
ester acid alcohol
organic ester hydrolysis
test
reaction

26. Dependence of k on temperature for the hydrolysis of an organic
ester
k increases
exponentially!
Note that both
reactant
concentrations are
held constant

27. /
a
E RT
k Ae

 k = rate constant at temperature T
 A = frequency factor
 Ea = activation energy
 R = Gas constant, 8.31451 J/K·mol

28. 1
ln( ) ln( )
a
E
k A
R T
 
 
 
 
 Simplifies solving for Ea
 -Ea / R is the slope when (1/T) is plotted
against ln(k)
 ln(A) is the y-intercept
 Linear regression analysis of a table of
(1/T) vs. ln(k) can quickly yield a slope
 Ea = -R(slope)

29. © 2009, Prentice-
Hall, Inc.
Arrhenius Equation
Taking the natural
logarithm of both sides,
the equation becomes
ln k = - ( ) + ln A
1
T
y = m x + b
Therefore, if k is determined
experimentally at several
temperatures, Ea can be calculated
from the slope of a plot of ln k vs.
.
Ea
R
1
T
Taking the natural
logarithm of both sides,
the equation becomes
ln k = - ( ) + ln A
1
T
ln
k2
k1
=
Ea
R
-
1
T2
1
T1
-
Or

30. PLAN:
SOLUTION:
Determining the energy of activation
PROBLEM: The decomposition reaction of hydrogen iodide,
2HI(g) H2(g) + I2(g)
has rate constants of 9.51 x 10-9 L/mol.s at 500. K and 1.10
x 10-5 L/mol.s at 600. K. Find Ea.
Use a modification of the Arrhenius equation to find Ea.
ln
k2
k1
=
Ea
-
R
1
T2
1
T1
- Ea = - R ln
k2
k1
1
T2
1
T1
-
-1
ln
1.10 x 10-5 L/mol.s
9.51 x 10-9 L/mol.s
1
600. K
1
500. K
-
Ea = - (8.314 J/mol . K)
Ea = 1.76 x 105 J/mol = 176 kJ/mol
-1

31. Higher Surface Area = greater rate of reaction
•Particles more likely to collide
• The smaller the particle size, the greater the surface area is for a
given mass of particles.

32. Surface Area
• The smaller the particle size, the greater the surface area is for a
given mass of particles.
• The result of an increase in surface area is an increase in the
frequency of collisions and the reaction rate.
Only atoms at the surface of the metal
are available for reaction.
Dividing the metal into smaller pieces
increases the surface area and the
number of collisions.
Mg(s) + 2H+(aq)  Mg2+(aq) + H2(g)
When a piece of
magnesium is placed in
dilute acid, hydrogen
ions can collide with
magnesium atoms.

33. Catalyst
A catalyst is a substance that
increases the rate of a reaction
without being used up during
the reaction.
Catalysts permit reactions to
proceed along a lower energy
path.
The activation-energy barrier for
the catalyzed reaction is lower
than that of the uncatalyzed
reaction.

34. Mechanism of Catalyst
Collisions only result in a reaction if the
particles collide with a certain minimum
energy called the activation energy for
the reaction. The position of activation
energy can be determined from a on a
Maxwell-Boltzmann distribution:
To increase the rate of a reaction, the number of
successful collisions must be increased. One possible
way of doing this is to provide an alternative way for
the reaction to happen which has a lower activation
energy. In other words, to move the activation
energy to the left on the graph:

35. Adding a catalyst has this effect on
activation energy. A catalyst
provides an alternative route for
the reaction with a lower activation
energy.

36. Catalyst
Homogeneous
Catalyst
Heterogeneous
Catalyst
Homogeneous
catalysts are those
which exist in the
same phase (gas or
liquid) as the
reactants
In a heterogeneous
reaction, the
catalyst is in a
different phase from
the reactants.
Enzyme
Catalyst

38. Heterogeneous catalysis
How the heterogeneous catalyst works (in general terms)
A) One or more of the reactants
are adsorbed on to the surface of the
catalyst at active sites.
B) At this stage, both of the reactant
molecules might be attached to the
surface, or one might be attached and
hit by the other one moving freely in the
gas or liquid.

39. The double bond between the
carbon atoms breaks and the
electrons are used to bond it to
the nickel surface.
Hydrogen molecules are also adsorbed on
to the surface of the nickel. When this
happens, the hydrogen molecules are
broken into atoms. These can move
around on the surface of the nickel
The bond between the carbon and the nickel is
replaced by one between the carbon and
hydrogen.
That end of the original ethene now breaks free of the
surface, and eventually the same thing will happen at the
other end.

40. As before, one of the hydrogen atoms forms a bond with the carbon, and that
end also breaks free. There is now space on the surface of the nickel for new
reactant molecules to go through the whole process again.
The product molecules are desorbed.
Desorption simply means that the product molecules break away.
This leaves the active site available for a new set of molecules to
attach to and react.
C) Desorption

42. Will be more on the catalyst ……………………….

43. Other Remaining factors will be
next…………….