Chapter 12: Chemical Kinetics

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Chapter 12 Chemical Kinetics

• Reaction Rates
• Rate Laws Differential and Integrated
• Experimental Determination of the Rate Law
• Integrated Rate Laws Concentration/Time
  Problems
• Collision Theory of Reactions
• Reaction Mechanisms
• Predicting Rate Laws
• Temperature Dependence of the Rate Constant
  (Arrhenius)
• Catalysis

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As the reaction proceeds, it gets slower because
the rate depends on NO2.
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Figure 12.1
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Example 1

• Determine the rate law and the value of the rate
• constant for the following reaction.
• 2 NO Br2 ? 2 NOBr

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Example 2

• Determine the rate law and the value of the
• rate constant for the following reaction in which
• OH- is a catalyst.
• OCl- I- ? OI- Cl-

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Example 3
The concentration of H2O2 was monitored over time
for the following reaction at 25C H2O2 (aq) ?
H2 (g) O2 (g) Find the rate law and the
value of the rate constant for this reaction at
25C.
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Example 4

• The decomposition of N2O5 to NO2 and O2 is
• first order with a rate constant of
• 4.80 x 10-4 /s at 45C.
• If the initial concentration of N2O5 is 1.65 x
  10-2 mole/L, what is the concentration after 825
  s?
• How long would it take for the concentration of
  N2O5 to decrease to 1.00 x 10-2 mole/L if its
  initial concentration wa that given in (a)?

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Example 5

• For the following decomposition reaction,
• AB ? A B
• Rate kAB2
• k 0.20 L/mole s
• How long will it take for AB to reach one
• third of its original value of 1.50 M? What
• is AB after 10.0 seconds?

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Zero Order Half Life

• t1/2 A0 / 2k
• Note that this half life
• decreases as the reaction
• proceeds.
• There is less reactant to
• Consume and it runs at the
• same rate so it takes less
• time.

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First Order Half Life

• t1/2 0.693 / k
• The first order half life is constant
• for a particular reaction and
• temperature.
• As the initial concentration
• decreases, the decrease in rate
• is compensated for by less
• reactant to consume.

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Second Order Half Life

• t1/2 1 / A0 k
• Note that this half life increases
• as the reaction proceeds.
• The rate decreases so strongly
• with concentration that it still
• takes longer to use up half of
• the reactant even though the
• amount of reactant consumed is
• less for each half life.

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Example 6

• A decomposition reaction has a rate
• constant of 0.0012 yr-1.
• What is the half life of this reaction?
• How long does it take for the concentration of
  the reactant to reach 12.5 of its original value?

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Example 7

• It took 143 s for 50.0 of a particular
• substance to decompose. If the initial
• concentration was 0.060 M and the
• decomposition reaction follows second
• order kinetics, what is the value for the
• rate constant?

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Example 8

• For the reaction A ? products,
• successive half-lives are observed to be
• 10.0, 20.0, and 40.0 min for an
• experiment in which A00.10M.
• Calculate the concentration of
• A at the following times.
• 80.0 min
• 30.0 min

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Example 9

• (Book 34)O NO2 ? NO O2
• NO2 in large excess,
• 1.0x1013molecules/cm3
• Find the order of the reaction with respect to
  O.
• Reaction is known to be first order in NO2.
  Find the value of k.

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Example 10

• Consider the following mechanism for the
  decomposition
• of hydrogen peroxide catalyzed by I-.
• Step 1 H2O2 I- H2O IO-
  fast equilibrium
• Step 2 IO- H2O2 ? H2O O2 I-
  slow
• a) Write a balanced equation for the overall
  reaction.
• b) List any catalysts.
• c) List any intermediates.
• d) Write the rate law for the overall reaction.

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Example 11

• Consider the following mechanism for the
  production of
• phosgene
• Step 1 Cl2 (g) 2 Cl (g) fast
  equilibrium
• Step 2 Cl (g) CO (g) COCl (g)
  fast equilibrium
• Step 3 COCl (g) Cl2 (g) ? COCl2 (g)
  Cl (g) slow
• Step 4 2 Cl (g) ? Cl2 (g) fast
• a) Write a balanced equation for the overall
  reaction.
• b) List any catalysts.
• c) List any intermediates.
• d) Write the rate law for the overall reaction.

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Figure 12.11 2BrNO ? 2 NO Br2
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Figure 12.13
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Example 12

• The activation energy of the following
• reaction is 3.5 kJ/mole and the change in
• enthalpy is -66.6 kJ/mole. Calculate
• the activation energy of the reverse
• reaction.
• OH HCl ? H2O Cl
• Hint Sketch the reaction energy diagram.

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Example 13

• For the reaction A2 B2 ? 2 AB, the
• activation energy of the forward reaction is
• 125 kJ/mole and that of the reverse
• reaction is 85 kJ/mole. Assuming the
• reaction occurs in one step
• Draw a reaction energy diagram.
• Calculate DH for the reaction.
• Sketch a possible transition state.

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Figure 12.2 Collisions with E gt Ea at two
temperatures.
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Example 14

• The isomerization reaction
• CH3NC ? CH3CN
• has an activation energy of 161 kJ/mole.
• If the rate constant at 600K is 0.41/s,
• what is it at 1000K?

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Example 15

• The rate constant of a reaction is
• 4.50x10-5 L/mols at 195C and
• 3.20 x 10-3 L/mols at 258C.
• What is the activation energy of this
• reaction?

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Catalysts Work by Lowering EaThey provide a
different mechanism!
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More collisions result in reaction when Ea is
lower.
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Homogeneous Catalysis

• In the stratosphere, NO catalyzes the
• destruction of ozone, O3.
• NO (g) O3 (g) ? NO2 (g) O2 (g)
• O (g) NO2 (g) ? NO (g) O2 (g)
• Overall
• O3 (g) O (g) ? 2 O2 (g)

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C2H4 H2 ? C2H6with a metal catalysta.
reactantsb. adsorptionc. migration/reactiond.
desorption
Heterogeneous Catalysis
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Enzymes are Catalysts
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Enzymes are Catalysts
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Enzymes are Catalysts