Rate Theories

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Rate Theories

• Two theories explaining why reactions occur at
  different rates
• Transition State Theory
• Collision
• These two theories are basically the same!
• Transition State includes more information

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Energy Relationships

• Activation energy (Ea) minimum amount of energy
  needed for reactions to occur
• Like an energy barrier that separates the
  reactants from products
• The greater the activation energy, the more
  difficult it will be for the reaction to occur
  ? Slower reaction

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Activation Energy

• Reactants do not immediately change into the
  products
• As reactants come together, an activated complex
  is formed
• Activated complex is higher in energy than the
  reactants
• Activation energy difference in energy between
  reactants and activated complex

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Energy Graph

• activated complex
• Energy Ea
• reactants
• DE
• products
• Reaction Pathway

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Change in Energy

• DE for a reaction
• energy of products energy of reactants
• DE for a reaction
• Activation energy reverse direction
  activation energy forward direction

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CH3NC ? CH3CN
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Overcoming Activation Energy

• Molecules are constantly moving
• This movement is kinetic energy
• When molecules collide, some of their energy is
  transferred
• Eventually, two molecules collide with sufficient
  energy to overcome activation energy

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Making a Reaction Occur

• Three things are needed for a reaction to occur
• Molecules must collide
• Collision must have sufficient energy
• Molecules must have proper orientation

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Making a Reaction Occur

• Effective Collision Ineffective Collision
• Sufficient
• Speed ?
• Molecules are changed Molecules
  remain unchanged
• Reaction occurs No
  reaction occurs

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Reaction Mechanisms

• In many reactions, it is too difficult for all of
  the reactants to collide properly in one step.
• Most reactions occur in a series of steps
• Reaction Mechanism step by step description of
  how reaction occurs
• Deduced from Rate Law and Activation Energy

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Reaction Mechanisms

• Made up of a series of single step reactions
• Reactions that CAN occur in one step
• Single step reaction also known as Elementary
  reaction
• Each reaction in the mechanism elementary
  reaction

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Reaction Mechanisms

• Three types of elementary reactions
• Unimolecular ? One reacting molecule
• A ? C D
• Bimolecular ? Two reacting molecules
• 2A ? C D
• A B ? C D
• Termolecular ? Three reacting molecules
• 2A B ? C D

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Reaction Mechanisms

• Each step in the mechanism is labeled as fast or
  slow
• The slow step is the rate determining step
• The rate law for the reaction can be predicted
  from the slow step
• This rate law must be compared to EXPERIMENTALLY
  determined rate law

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Reaction Mechanisms

• A mechanism is valid IF
• The steps add up to the original equation
• The rate law based on the mechanism matches the
  experimental rate law

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Reaction Mechanisms

• Coefficient from balanced equation is the
  reaction order

Reaction Rate Law Rxn order (reactants) Rxn Order (overall)
A ? C D Rate kA 1st 1st
A B ? C D Rate kAB 1st 2nd
2A ? C D Rate kA2 2nd 2nd
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Reaction Mechanisms

• Consider the reaction
• NO2 CO ? NO CO2
• Rate law written from reaction
• Rate k NO2 CO
• Rate law determined experimentally
• Rate k NO22

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Reaction Mechanisms

• Reaction NO2 CO ? NO CO2
• Proposed mechanism
• 2NO2 ? NO NO3 slow
• NO3 CO ? NO2 CO2 fast
• Rate law based on slow step
• Rate k NO22

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Reaction Mechanisms

• Mechanisms can include Intermediates ? compounds
  that do not appear in the overall equation
• Intermediates can NEVER appear in the rate law.
• 2NO2 ? NO NO3
• NO3 CO ? NO2 CO2

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Reaction Mechanisms

• If an intermediate is part of the slow step, it
  must be replaced in the rate law with its
  compliment
• If the first step is fast AND equilibrium, be on
  the lookout for intermediates.

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Reaction Mechanisms

• Reaction Cl2 CHCl3 ? HCl CCl4
• Rate Law Rate k Cl21/2 CHCl3
• Mechanism
• Cl2 ??2Cl fast
• Cl CHCl3 ? HCl CCl3 slow
• Cl CCl3 ? CCl4 fast

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Factors Affecting Rate

• Catalyst substance that changes the rate but is
  not part of the reaction
• In general, catalysts INCREASE rate
• Catalysts increase the rate of reaction by
  creating a different reaction pathway
• The new pathway has a lower activation energy
• More molecules have sufficient energy

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Catalysts

• Energy Diagram
• uncatalyzed
• Ea
• Energy Ea catalyzed
• reactants
• DE
• products
• Reaction Pathway

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Catalysts

• Two types of catalysts
• Homogeneous ? catalyst in same phase as reactants
• Heterogeneous ? catalyst in different phase than
  reactants
• Usually finely divided metals

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Rate vs. Temperature

• Reaction rates INCREASE when temperature
  increases. WHY??
• Increasing the temperature increases the average
  kinetic energy of the molecules.
• More molecules will have enough energy to
  overcome activation energy

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Rate vs. Temperature

• Arrhenius equation summarizes temp rate
  relationship
• ln k ln A Ea OR k Ae(-Ea/RT)
• RT
• A frequency factor Ea activation energy
• R gas constant k rate constant
• T temperature (K)

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Rate vs. Temperature

• ln k ln A (Ea / RT)
• Frequency factor constant that describes number
  of effective collisions
• As activation energy increases, rate constant
  decreases
• As temperature increases, rate constant increases

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Rate vs. Temperature

• We can modify the Arrenhius equation to compare
  the rate constant at different temperatures
• ln k1 Ea 1 - 1
• k2 R T2 T1

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Rate vs. Temperature

• C2H5Br ? C2H4 HBr
• Rate constant 2.0 x 10-5 / sec at 650 K
• Frequency factor 9.58 x 1012 / sec
• Find the rate constant at 670 K.