The Collision Theory and
Activation Energy
Explaining how and why factors
affect reaction rates
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)
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
Molecules
hit disk last
Molecules
hit disk first
Kinetic energy 
COLLISION THEORY
• Rate proportional to fraction of effective
collisions x collision frequency
• Effective collisions with enough KE =
ACTIVATION ENERGY
Factors that affect number of
effective collisions
related to activation energy
chemical nature of reactants
Temperature (faster moving molecules means
more collisions per unit of time).
Catalyst
Factors that affect collision
frequency
•
•
•
•
Related to number of collisions
Concentration
Surface area
temperature
Temperature and reaction rate
• By increasing the temperature, a small number
of molecules reach Ea. The reaction is
exothermic, further increasing temperature and
causing more molecules to reach Ea, etc.
Shift due to higher
temperature
Fraction of
molecules
Kinetic energy 
Ea
kinetic vs. potential energy diagrams
• Recall the Maxwell-Boltzman distribution (i.e.
kinetic energy diagram)
Ea
Kinetic energy 
Path of reaction 
• The Ea is a critical point. To examine it more
closely we can use a potential energy graph
• The axes are not the same, thus the Ep graph
is not a blow up of the Ek graph; however it
does correspond to the part of the Ek graph
that is circled
potential energy graph: a closer look
Activated complex /
transition state
Reactants
Products
Ea
H
Collision
begins
Collision ends
A2 & B2 molecules molecules
2AB
rush
slow down speed up molecules
together Ep , Ek  Ep , Ek  float apart
Path of reaction
Ep + Ek =constant Overall Ep(reactants)>Ep(products)
throughout
Ek(reactants)<Ek(products)
Ep graph: Important points
Exothermic
Ea
forward
Endothermic
Ea
reverse
 H is
positive
• Forward and reverse reactions are possible
• Ea is the difference between Ep at transition
state and initial or final Ep
• The graph depicts an exothermic reaction.
Endothermic reactions are also possible
• H is the difference between initial and final
Ep. It is -ve for exothermic,+ve for endothermic
The collision theory
• Related to the Ep graph is the “collision theory”
- the idea that for molecules to react they must
meet with sufficient force
• Factors that affect reaction rate can be
explained via the collision theory:
• Increased temperature causes molecules to
move faster (increased number of collisions
per unit time and greater kinetic energy)
• Increased concentration means more collisions
• Homogenous reactions occur faster because
reacting molecules collide more frequently
• Catalysts decrease Ea, decreasing the amount
of kinetic energy needed to overcome Ea
Catalysts
• Recall, catalysts speed a reaction
• This can be explained by the Ek or Ep graphs
• In both, the catalyst works by lowering the Ea:
Fraction of
molecules
potential
energy
Kinetic energy 
Path of reaction
• Catalysts speed forward and reverse reactions
• However, most reactions favour the side that
has the lowest potential energy (most stable)
• Catalysts are heterogenous or homogenous
• They provide a substrate (p. 768) for a reaction
or they can bond temporarily to a molecule,
increasing the odds of a favourable meeting
Transition state lab: purpose
Purpose: 1) to visualize an activated complex, 2)
to observe the influence of a catalyst
We will be examining the following reaction:
NaKC4H4O6(aq) + H2O2(aq)  CO2(g) + …
Procedure:
1. Turn hot plates immediately to medium heat
2. Get a 10 mL graduated cylinder, a 100 mL
beaker, a test tube, and a rubber stopper.
3. Weigh 1.7 g NaKC4H4O6. Add to beaker along
with 10 mL distilled H2O. Swirl to dissolve.
4. Add 4.5 mL of 10% H2O2 to beaker. Heat.
5. Get 5 mL of CoCl2 but don’t add it yet.
Transition state lab: procedure
Procedure:
6. As soon as tiny bubbles start to form and
rise, remove the beaker from the hot plate.
Add the CoCl2 at this point.
7. Record your observations (in order to answer
the questions). Clean up – wash everything
down the drain, wipe off your lab bench.
Questions: answer on a separate sheet of paper
1. Look at the chemical equation that
represents the reaction. What physical sign
will there be when a reaction is occurring?
2. The products of the reaction are colourless.
What colour are the reactants?
Transition state lab: conclusions
Questions: read 18.11 (pg. 767 – 769)
3. What was the catalyst in the lab? What colour
was it? Is it homogenous or heterogeneous?
4. At the beginning of step 5, both reactants
were present; why was there no reaction?
(Illustrate with a Ek diagram).
5. Why is the reaction still slow after heat is
added? (illustrate using the Ek diagram)
6. Was the catalyst a different colour at the end
of the experiment than at the beginning?
7. What colour was the activated complex?
8. Illustrate the affect the catalyst had on the
reaction (using both Ek and Ep diagrams)
Answers
1. The production of CO2 (bubbling) is a
physical sign that the reaction is occurring
2. The reactants are colourless
3. CoCl2 was the catalyst in the lab (pink,
homogenous)
4. There was no reaction because the Ea was
not reached (Illustrate with a Ek diagram).
5. The reaction still slow after heat is added
because very few molecules exceed Ea.
Fraction of
molecules
Kinetic energy 
Answers
6. The catalyst was the same colour at the end
of the experiment (catalysts don’t change).
7. The activated complex was green
8. Illustrate the affect the catalyst had on the
reaction (using both Ek and Ep diagrams)
Fraction of
molecules
potential
energy
Kinetic energy 
Path of reaction
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