Collision Theory and Activation Energy

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Collision Theory and Activation Energy

• Unit 3 Chemical Kinetics and Equilibrium

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Review Chemical Reaction

• Generic Format of reaction

A B C D
Reactants
Products
5 different types of chemical reactions
1) Synthesis
4) Double Displacement
2) Decomposition
5) Combustion
3) Single Displacement
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Review Kinetic Molecular Theory

• 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)

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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.

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Turns out

• 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!

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The Collision Theory

• 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
• Proper orientation
• Particles must collide with a certain minimum
  amount of energy, called activation energy.

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1. Orientation

• Particles must collide with the proper geometry
  or orientation for atoms to come in direct
  contact and form the chemical bonds of the
  products.

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2. Activation Energy

• 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.

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2. Activation Energy
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Effect of Temperature
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Potential Energy Diagram

• 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

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The hill

• The hill represents the activation energy that
  must be overcome for the reaction to occur
• high hill ? slow rate of reaction
• low hill ? fast rate of reaction
• Top of hill is called change over point
• There is a chemical species that exists here is
  referred to as activated complex
• Activated complex
• 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.

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Some collisions are successful.

• 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!

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Reaction Mechanism

• Converting reactants to products often involves
  more then one stepas 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
• Step 1 NO O2 ?NO3
• Step 2 NO3 NO ?2NO2

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The Rate Determining Step

• 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.

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Catalysts

• 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

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For example

• A B ? AB (no catalyst)
• With catalyst
• Step 1 A catalyst ? A-cat
• Step 2 A-cat B ? AB catalyst
• Overall A B ? AB
• Both steps are faster then the original,
  uncatalyzed reaction
• Catalyst remains unchanged in the end

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Catalysts contd
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Real life catalyst example

• Manganese dioxide (a black powder) will catalyze
  the breakdown of hydrogen peroxide.
• 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.

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Practice

• Read Chapter 12!!!!!
• Pg 484 9, 12a,b,d
• Pg 486 1-4, 12