Collision Theory
and Activation
Energy
Unit 3: Chemical Kinetics and Equilibrium
Review: Chemical Reaction
Generic Format of reaction
A + B
Reactants
C + D
Products
5 different types of chemical reactions
1) Synthesis
4) Double Displacement
2) Decomposition
5) Combustion
3) Single Displacement
Review: Kinetic Molecular Theory
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Used to explain many observations and chemical events!
All matter is made up of microscopic-sized particles
(atoms, ions, molecules)
These particles are in constant motion (possess
kinetic energy)
There is space between the particles (speed and
spacing determine the physical state of matter)
Adding energy increases the speed of the moving
particles (thus inc kinetic energy)
How do reactions occur?
Collision Theory
 In order for a chemical reaction to take place, the
reactants must come in contact and collide!!!!!
 The collision transfers kinetic energy needed to
break the necessary bonds so that new bonds can be
formed.
Turns out……
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By calculating how many collisions are taking place per
second and how quickly product is being produced…
chemists learned that most collisions are not successful
(no product formed)
To think that reactant particles collide and products are
automatically produced is over simplified.
There must be other requirements for a collision to be
successful!
The Collision Theory
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is an explanation of what is necessary for a
chemical reaction to occur.
When a chemical reaction takes place, the
reactant particles must meet two conditions (or
requirements) during collision for the collisions
to be successful:
1.
2.
Proper orientation
Particles must collide with a certain minimum amount
of energy, called activation energy.
1. Orientation
o
Particles must
collide with the
proper geometry
or orientation for
atoms to come in
direct contact
and form the
chemical bonds
of the products.
2. Activation Energy
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Particles must collide with a certain
minimum amount of energy, called
activation energy (Ea).
This energy is required to break chemical
bonds in the reactants.
Note: The energy of each particle is not
important, it is the energy of the collision.
2. Activation Energy
Effect of Temperature:
Potential Energy Diagram:
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We can represent
the increase in
potential energy
during a chemical
reaction using a
potential energy
diagram
The kinetic energy of
reactants is
transferred to
potential energy as
the reactants collide
“The hill”
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The hill represents the activation energy that
must be overcome for the reaction to occur
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Top of hill is called “change over point”
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“high hill”  slow rate of reaction
“low hill”  fast rate of reaction
There is a chemical species that exists here is
referred to as “activated complex”
Activated complex
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Neither reactant or product
Partial bonds, highly unstable
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Potential Energy Diagram Handout
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If both of these conditions are not met,
particles will merely collide and bounce off
one another without forming products.
Some collisions are successful….
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Although, the percentage of successful collisions
is extremely small, chemical reactions still take
place at a reasonable rate because there are so
many collisions per second between reactant
particles!
Reaction Mechanism:
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Converting reactants to products often
involves more then one step…as you know!
Each step is called an elementary reaction
Molecules formed during elementary
reactions are called reaction intermediates
(neither reactants nor products)
Example: 2NO + O2  2NO2
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Step 1: NO + O2 NO3
Step 2: NO3 +NO 2NO2
The Rate Determining Step:
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In multi-step reactions there can be 2 or more
elementary reactions….
 There is always one reaction that is slowest …this
determines the overall rate of reaction
 Thus, the slowest elementary reactions is called the
rate determining reaction.
Catalysts:
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Works by lowering the activation energy of a
reaction so that a larger fraction of reactants
have sufficient energy to react
They do this by providing an alternative
reaction mechanism
For example:
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A + B  AB (no catalyst)
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With catalyst:
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Step 1:
Step 2:
Overall:
A + catalyst  A-cat
A-cat + B  AB + catalyst
A + B  AB
Both steps are faster then the original,
uncatalyzed reaction
Catalyst remains unchanged in the end
Catalysts cont’d
Real life catalyst example:
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Manganese dioxide (a black powder) will catalyze the
breakdown of hydrogen peroxide.
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2H2O2 (aq) → 2H2O(l) + O2 (g) (uncatalyzed)
With catalyst:
Car exhaust pipes use catalytic converters to get rid of
some of the nasty gases from the engines.
Practice:
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Read Chapter 12!!!!!
Pg 484 # 9, 12a,b,d
Pg 486 # 1-4, 12