Answer the following questions:

1
Answer the following questions

• Define the term spontaneous reaction.
• Which of the following reactions do you think are
  spontaneous?
• 1) A nail rusting.
• 2) The decomposition of carbon in diamond to
  graphite.
• 3) The combination of hydrogen and oxygen to
  form water.
• 4) The decomposition of hydrogen peroxide.
• 5) The combustion of butane.
• What does the spontaneity of a reaction indicate
  about the speed at which the reaction occurs?

2
Chemical Kinetics
3
Kinetics

• The study of reaction rates.
• Spontaneous reactions are reactions that will
  happen - but we cant tell how fast. (Spontaneity
  implies nothing about speed)
• For example Diamond will spontaneously turn to
  graphite so slow it is not detectable.
• Reaction mechanism- the steps by which a reaction
  takes place.

4
Factors that Affect Reaction Rates

• Nature of the reactants
• Concentration of the reactants
• Temperature
• Presence of a catalyst

5
Collision Theory of Reaction Rates

• For a reaction to occur, particles must collide.
• Increasing the concentration or reactants results
  in a greater number of collisions and therefore,
  a faster rate.
• Not all collisions are effective collisions.

6
Effective Collisions

• For a collision to be effective
• 1) particles must possess a necessary minimum
  energy to break bonds and form new ones
  (activation energy)
• 2) particles must have proper orientation at
  the time of the collision. (see page 554)

7
Reaction Rate

• Defined as the change in the concentration of a
  reactant or product per unit time.
• RateConc.of A at t2-Conc. of A at t1
• t2 t1
• Rate DA Dt

8
For this reaction N2 3 H2 ? 2 NH3
As the reaction progresses the concentration H2
goes down
CONCENTRATION
H2
TIME
9

• As the reaction progresses the concentration N2
  goes down 1/3 as fast

CONCENTRATION
N2
H2
H2
TIME
10

• As the reaction progresses the concentration NH3
  goes up.

N2
H2
NH3
11
Calculating Rates

• Average rates are taken over long intervals
• Instantaneous rates are determined by finding the
  slope of a line tangent to the curve at any given
  point because the rate can change over time

12
Defining Rate

• We can define rate in terms of the disappearance
  of the reactant or in terms of the rate of
  appearance of the product.
• In our example N2 3H2 ? 2NH3 -DN2
  -3DH2 2DNH3
  Dt Dt Dt

13
Rate Laws

• Reactions are reversible.
• As products accumulate they can begin to turn
  back into reactants.
• Initially, the rate will depend on only the
  amount of reactants present.
• Therefore, reactions must be studied at a point
  soon after the reactants are mixed.
• This is called the Initial rate method.

14
Rate Laws

• The concentration of the products do not appear
  in the rate law because this is an initial rate.
• The order must be determined experimentally,
  and cant be obtained from the equation
• The order indicates to what degree the
  concentration affects the rate.

15
Types of Rate Laws

• Differential Rate law - describes how rate
  depends on concentration (often simply called
  rate law)
• Integrated Rate Law - describes how concentration
  depends on time.
• For each type of differential rate law there is
  an integrated rate law and vice versa.
• Rate laws can help us better understand reaction
  mechanisms.

16
Reaction Mechanism

• The step by step pathway by which a reaction
  occurs is called its mechanism.
• The individual steps are called elementary steps.
• The rate law for an elementary step can be
  written from its molecularity.
• Molecularity is defined as the number of species
  that must collide to produce the reaction
  indicated by that step.

17
Examples of Elementary Steps
Elementary Step Molecularity Rate Law
A ? products Unimolecular Rate kA
A A ? products Bimolecular Rate kA2
A B? products Bimolecular Rate kAB
A A B? products Termolecular RatekA2B
A B C?products Termolecular RatekABC
18
2 NO2 ? 2 NO O2

• The rate will only depend on the concentration of
  the reactants.
• Rate kNO2n
• This is called a rate law expression or
  differential rate law.
• k is called the rate constant.
• n is the order of the reactant -usually a
  positive integer and must be determined
  experimentally
• Overall reaction order is the sum of the orders
  for each individual reactant.

19
Integrated Rate Laws

• Integrated Rate Laws express the reactant
  concentration as a function of time (instead of
  rate as a function of the reactant
  concentration).
• We will only be considering reactions involving a
  single reactant when determining integrated rate
  laws.

20
First-Order Rate Laws

• Consider the following reaction
• 2N2O5 ? 4NO2 O2
• The rate law (differential) is determined to be
• Rate kN2O5
• This means that if the concentration of N2O5 is
  doubled, the rate of production of the products
  is also doubled.

21
Integrated First-Order Rate Law

• The previous differential rate law can be
  integrated and as a result, written in the
  following form
• lnN2O5 -kt lnN2O50
• Where
• ln indicates the natural log
• t is the time
• N2O5 is the concentration at time t
• N2O5o is the initial concentration

22
Integrated First-Order Rate Laws

• All integrated first order rate laws take the
  following general form
• lnA -kt lnAo
• Items to note
• 1) The equation shows how the concentration of
  A depends on time

23
Integrated First-Order Rate LawslnA -kt
lnAo

• 2) The above equation is of the form ymx b,
  where a plot of y vs. x is a straight line with
  slope m and intercept b.
• In the above rate law
• ylnA xt m-k blnAo
• As a result, plotting lnA vs. time always
  gives a straight line. (This fact is often used
  to test whether a reaction is 1st order or not)
• 3) The integrated rate law for a 1st order
  reaction can also be written as
• ln(Ao/A) kt

24
Half-Life

• The time required for a reactant to reach half of
  its original concentration is called the
  half-life of a reactant and is designated with
  the symbol t1/2.
• The general equation for the half-life of a first
  order reaction is
• t1/2 0.693/k
• (Half-life does not depend on concentration)

25
2nd Order Rate Laws

• Consider the following reaction
• 2C2H6 (g) ? C8H12(g)
• The rate law (differential) is determined to be
• Rate kC2H62
• This means that if the concentration of C2H6 is
  doubled, the rate of production of the products
  is quadrupled.

26
Integrated Second-Order Rate Laws

• The previous differential rate law can be
  integrated and as a result, written in the
  following form
• 1/C4H6 kt 1/C4H6o
• All integrated second order rate laws have the
  form
• 1/A kt 1/Ao
• Note
• 1) A plot of 1/A vs. time produces a straight
  line with a slopek.
• 2) The half-life equation for a 2nd order
  reaction is t1/2 1/kAo

27
Zero-Order Rate Laws

• The rate law for a zero order reaction is rate
  k
• This means that changing the concentration has no
  effect on the rate of the reaction.
• Example N2O5 ? 2NO2 1/2O2
• The integrated rate law for a zero order reaction
  is
• A -kt Ao
• Note
• 1) A plot of A vs. time results is a straight
  line
• 2) the half life equation is
• t1/2 A0/2k

28
Interpreting Differential Rate Laws
29
A B C ? products

• The rate law is determined to be
• rate kAB2
• What happens to the reaction rate
• when we make the following changes
• to the concentrations?

30
The concentration of A is doubled, while B and C
remain unchanged.

1. The rate remains the same.
2. The rate doubles.
3. The rate is cut in half.
4. The rate is quadrupled.
5. The rate is increased by a factor of 8.

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31
The concentration of B is doubled while A and C
remain unchanged.

1. The rate remains the same.
2. The rate doubles.
3. The rate is cut in half.
4. The rate quadruples.
5. The rate increases by a factor of 8.

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32
The concentration of C is doubled while A and B
remain unchanged.

1. The rate remains the same.
2. The rate is doubled.
3. The rate is cut in half.
4. The rate is quadrupled.
5. The rate is increased by a factor of 8.

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33
The concentration of all three reactants are
doubled simultaneously.

1. The rate remains the same.
2. The rate is doubled.
3. The rate is cut in half.
4. The rate is quadrupled.
5. The rate increases by a factor of 8.

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34
The concentration of A is cut in half while B and
C are doubled.

1. The rate remains the same.
2. The rate is doubled.
3. The rate is cut in half.
4. The rate is quadrupled.
5. The rate increases by a factor of 8.

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35
Which of the following is the rate law for a
first order reaction?

1. Rate kA
2. Rate kA2
3. Rate kAB
4. Rate kA2B
5. Rate kA2B2

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36
Which is a rate law that is second order with
respect to B?

1. Rate kA
2. Rate kA2
3. Rate kAB
4. Rate kA2B
5. Rate kA2B2

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37
Which is the rate law for a third order reaction?

1. Rate kA
2. Rate kA2
3. Rate kAB
4. Rate kA2B
5. Rate kA2B2

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38
Which rate law represents a reaction in which the
initial rate will increase by a factor of eight
when A and B are doubled?

1. Rate kA
2. Rate kB2
3. Rate kAB
4. Rate kA2B
5. Rate kA2B2

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39
Which rate law represents a reaction in which the
rate will increase by a factor of two when A
and B are doubled?

1. Rate kA
2. Rate kB2
3. Rate kAB
4. Rate kA2B
5. Rate kA2B2

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40
Which rate law represents a reaction in which the
rate will not change when A is doubled and B
is held constant?

1. Rate kA
2. Rate kB2
3. Rate kAB
4. Rate kA2B
5. Rate kA2B2

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41
A reaction occurs according to the following rate
law rate kA. If the temperature of the
reaction chamber were increased, which of the
following would be true?

1. The rate of reaction and the rate constant will
   increase.
2. The rate of reaction and the rate constant will
   not change.
3. The rate of reaction will increase and the rate
   constant will decrease.
4. The rate of reaction will increase and the rate
   constant will not change.
5. The rate of reaction will not change and the rate
   constant will increase.

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42
Determining the form of a rate law(Method of
Initial Rates)

• The initial rate is the instantaneous rate right
  after the reaction begins (t0)
• Eliminates the effect of the reverse reaction.
• The initial concentrations of the reactants are
  varied.
• The reaction rate is determined for each trial.
• See the example problem on p. 536

43
A 2B ? C
Trial Initial A Initial B Initial Rate of formation of C
1 0.010 M 0.010 M 1.5 x 10-6 M/s
2 0.010 M 0.020 M 3.0 x 10-6 M/s
3 0.020 M 0.010 M 6.0 x 10-6 M/s
44
The rate law for the reaction is

1. Rate k AB
2. Rate k A2B
3. Rate k AB2
4. Rate k A2B2

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45
Reaction Mechanism (continued)

• In most mechanisms, one step is slower than the
  others.
• A reaction can never occur faster than its
  slowest step.
• The slow step is the rate-determining step.

46
Determining Possible Reaction Mechanisms

• The balanced equation for the overall reaction is
  equal to the sum of all of the individual steps.
• The rate law expression matches the coefficients
  of the rate determining step.

47
Example

• NO2 CO ? NO CO2
• rate kNO22
• One proposed mechanism
• NO2 NO2 ? N2O4 (slow)
• N2O4 CO ? NO CO2 NO2 (fast)
• NO2 CO ? NO CO2
• N2O4 is an intermediate (forms in one step and is
  consumed in a later step)
• This proposed mechanism is consistent with the
  experimentally determined rate law.

48
Another possible mechanism

• NO2 NO2 ? NO3 NO (slow)
• NO3 CO ? NO2 CO2 (fast)
• NO2 CO ? NO CO2
• In order to be a possible mechanism the following
  two criteria must be met
• 1) The individual steps must add up to the
  overall reaction
• 2) The mechanism is consistent with the
  experimentally determined rate law expression

49
Equilibrium Step

• An equilibrium step is a step in which a product
  can rapidly re-form the reactants to reach
  equilibrium.
• The concentration of products is equal to the
  concentration of the reactants at equilibrium.
• An equilibrium step is indicated by a double
  arrow. (lt -- gt )
• Substitutions can be made in the rate law
  expression between the reactants and the products.

50
Example

• 2NO Br2 ? 2NOBr
• rate kNO2Br2
• Proposed mechanism
• NO Br2 lt --gt NOBr2 (fast)
• NOBr2 NO ? 2NOBr (slow)
• 2NO Br2 ? 2NOBr
• Rate kNOBr2NO can be substituted with
• Rate kNOBr2NO or RatekNO2Br2

51
A multi step reaction takes place by the
following mechanismA B ? C DA C ? D
EWhich of the species shown is an intermediate
in the reaction?

1. A
2. B
3. C
4. D
5. E

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52
Which of the following is the rate law for the
predicted mechanism for the production of NO2?NO
O2 lt-gt NO3 (fast)NO3 NO ? 2NO2 (slow)

1. Rate kNOO2
2. Rate kNO3NO
3. Rate kNO2O2
4. Rate kNO22

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53
Homework problems to be completed by tomorrow

• Chapter 12 problems 51-53, 55, 65, 75, and 82

54
Example Problem p. 569-571

• Complete homework problems 32,34,39-42, 44, 46,
  48, and 89.