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Reaction rates

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					The Rate of Reaction
   Reaction rates are the rates at which reactants
    disappear or products appear.
   This movie is an illustration of a reaction rate.




                                                        1
Collision Theory of
Reaction Rates
        Three basic events must occur for a
         reaction to occur the atoms, molecules or
         ions must:
    1.    Collide.
    2.    Collide with enough energy to break and form
          bonds.
    3.    Collide with the proper orientation for a reaction
          to occur.


                                                               2
Collision Theory of
Reaction Rates
   One method to increase the number of collisions and
    the energy necessary to break and reform bonds is to
    heat the molecules.
   As an example, look at the reaction of methane and
    oxygen:
    CH 4(g)  O 2(g)  CO 2(g)  H 2 O (g)  891 kJ
   We must start the reaction with a match.
       This provides the initial energy necessary to break the
        first few bonds.
       Afterwards the reaction is self-sustaining.
                                                                  3
Collision Theory of
Reaction Rates
   Illustrate the proper orientation of molecules
    that is necessary for this reaction.
                  X2(g) + Y2(g) 2 XY(g)
   Some possible ineffective collisions are :


X                    Y
         Y Y               X   X           X X   Y Y
X                    Y

                                                       4
Collision Theory of
Reaction Rates
   An example of an effective collision is:


      X    Y        X    Y         X       Y
                                       +

      X    Y        X    Y         X       Y




                                               5
Collision Theory
                                    Activation energy
 enthalpy




                         Ea

            reactants



                        ∆Hrxn

                                products



                                           time
                                                        6
Transition State
   Transition state theory postulates that
    reactants form a high energy intermediate,
    the transition state, which then falls apart into
    the products.
   For a reaction to occur, the reactants must
    acquire sufficient energy to form the transition
    state.
       This energy is called the activation energy or Ea.
   Look at a mechanical analog for activation
    energy                                                   7
Transition State




                   8
Transition State
   The distribution of molecules possessing different
    energies at a given temperature is represented in
    this figure.




                                                         9
Factors That Affect Reaction
Rates
    There are several factors that can influence
     the rate of a reaction:
1.    The surface area of the reactants.
2.    The concentration of the reactants.
3.    The temperature of the reaction.
4.    The presence of a catalyst.
     We will look at each factor individually.

                                                    10
Surface Area
   This movie illustrates how changing the surface area
    of reactants affects the rate.




                                                       11
 Concentrations of Reactants
     This is a simplified representation of the
      effect of different numbers of molecules in the
      same volume.
         The increase in the molecule numbers is
          indicative of an increase in concentration.
                       A(g) + B (g)  Products

      A   B                  A B                   A B
                               B                   A B
      A B                    A B                   A B
4 different possible    6 different possible   9 different possible
                                                                    12
A-B collisions          A-B collisions         A-B collisions
Catalysts
   Catalysts change reaction rates by providing an
    alternative reaction pathway with a different
    activation energy.




                                                      13
Catalysts
   Homogeneous catalysts exist in same phase as the
    reactants.
   Heterogeneous catalysts exist in different phases
    than the reactants.
       Catalysts are often solids.




                                                    14
Catalysts
  Number of reactants particles




                                                   Fraction of particles with E > Ea
                                                   for the uncatalized reaction

                                    Fraction of particles
                                    with E > Eacat.




                                     Eacat   Ea
                                  Kinetic energy

                                                                                       15
Catalysts
   Examples of commercial catalyst systems include:

                          and
                             
C8 H18g  + 25 O 2g  NiOPt16CO2g   18 H 2 O g 
                                
         2 COg  + O 2g     2 CO2g 
                              NiO and Pt


                        
           2 NOg     N 2g   O 2g 
                      NiO and Pt


        Automobile catalytic converter system


                                                            16
Catalysts




            17
Catalysts
   This movie shows catalytic converter chemistry
    on the Molecular Scale




                                                     18
Catalysts
   A second example of a catalytic system is:

                             
    2 SO 2g   O 2g    2 SO3g 
                         V2 O5 or NiO/Pt


            Sulfuric acid preparatio n




                                                 19
Catalysts
   A third examples of a catalytic system is:

     N 2g   3 H 2g    2 NH3g 
                           Fe or Fe 2 O 3


                   Haber Process




                                                 20
Catalysts
   Look at the catalytic oxidation of CO to CO2
   Overall reaction
                       2 CO(g)+ O2(g) 2CO2(g)
   Absorption
                       CO(g) CO(surface) + O2(g)
                           O2(g) O2(surface)
   Activation
                        O2(surface)  O(surface)
   Reaction
                 CO(surface) +O(surface)  CO2(surface)
   Desorption
                         CO2(surface)  CO2(g)            21

				
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posted:11/27/2011
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