Title: Chemical Kinetics
1Chapter 14
2Review Section of Chapter 14 Test
- Net Ionic Equations Update to different
material. Maybe Chapter 11 stuff
3Reaction Rate
- The rate of a chemical reaction is measured as
the decrease in the concentration of a reactant
or the increase in the concentration of a product
in a unit of time.
4Reaction Rate
- The rate of a chemical reaction is measured as
the decrease in the concentration of a reactant
or the increase in the concentration of a product
in a unit of time.
?
Rate
?time
5Reaction Rate
- The rate of a chemical reaction is measured as
the decrease in the concentration of a reactant
or the increase in the concentration of a product
in a unit of time.
?
What units would we use for rate?
Rate
?time
6Reaction Rate
- The rate of a chemical reaction is measured as
the decrease in the concentration of a reactant
or the increase in the concentration of a product
in a unit of time.
?
Rate
?time
2H2O2(aq) ? 2H2O(l) O2(g)
7Reaction Rate
- The rate of a chemical reaction is measured as
the decrease in the concentration of a reactant
or the increase in the concentration of a product
in a unit of time.
?
Rate
?time
2H2O2(aq) ? 2H2O(l) O2(g)
How could the rate be expressed for this reaction
in terms of H2O2?
82H2O2(aq) ? 2H2O(l) O2(g)
92H2O2(aq) ? 2H2O(l) O2(g)
102H2O2(aq) ? 2H2O(l) O2(g)
What is the rate of the reaction from 0s to 2.16
x 104s?
112H2O2(aq) ? 2H2O(l) O2(g)
What is the average rate of appearance of O2 from
0s to 2.16 x 104s? 1.16 x 10-5 mol O2 L-1 s-1
12General Rate of Reaction
a A b B ? c C d D
Rate of reaction rate of disappearance of
reactants
or
Rate of reaction rate of appearance (formation)
of products
We can use the coefficients in the equation to
compare the reaction rates for all the substances
in the reaction.
1315-1 The Rate of a Chemical Reaction
- Rate is change of concentration with time.
2 Fe3(aq) Sn2(aq) ? 2 Fe2(aq) Sn4(aq)
t 38.5 s Fe2 0.0010 M
?t 38.5 s ?Fe2 (0.0010 0) M
14Rates of Chemical Reaction
2 Fe3(aq) Sn2(aq) ? 2 Fe2(aq) Sn4(aq)
Rate of formation of Fe2 2.6 x 10-5 mol L-1 s-1
What is the rate of formation of Sn4?
1.3 x 10-5 mol Sn4 L-1 s-1
What is the rate of disappearance of Fe3?
2.6 x 10-5 mol Fe3 L-1 s-1
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16What does the slope of the line represent?
17What is the concentration at 100s for the
reaction 2H2O2(aq) ? 2H2O(l) O2(g)?Given
H2O2i 2.32 M
2.15 M
18What does it mean when the rate of a reaction
reaches zero?
- For a normal reaction it means that one or more
of the reactants are used up and the reaction has
stopped. - For a reversible reaction it means that the
reaction has reached equilibrium.
19Factors Affecting Reaction Rates
- The nature of the reacting substances.
20Factors Affecting Reaction Rates
- 2. The state of subdivision of the reacting
substances (surface area).
21Lycopodium Powder
22Factors Affecting Reaction Rates
- 3. The temperature of the reacting substances.
23Factors Affecting Reaction Rates
- 4. The concentration of the reacting substances.
- (Except in zero order reactions)
Air (21 oxygen) 100 oxygen
24Factors Affecting Reaction Rates
- The presence of a catalyst.
- Catalysts speed up reactions but are left
unchanged by the reaction.
25The Rate Law
a A b B . ? g G h H .
Rate k AmBn .
Rate constant k (k is constant for a
particular reaction at a specific
temperature)
Order of A m Order of B n Overall order of
reaction m n .
26Temperature and Rate
- Generally, as temperature increases, so does the
reaction rate. - This is because k is temperature dependent.
- Therefore the temperature dependence of reaction
rates is contained in the temperature dependence
of the rate constant.
27Temperature dependence of k
.
.
.
.
.
28Concentration and Rate Summary
- After finding the trials to compare
- A reactant is zero order if the change in
concentration of that reactant produces no effect
on the rate. - A reaction is first order if doubling the
concentration of that reactant causes the rate to
double. - A reactant is nth order if doubling the
concentration of that reactant causes an 2n
increase in rate. - Note that the rate constant does not depend on
concentration.
29Use the data provided to write the rate law and
indicate the order of the reaction with respect
to HgCl2 and C2O42- and also the overall order of
the reaction.
30First determine the order of HgCl2
31Next determine the order of C2O42-
32Now write the rate law and determine the order of
the reaction.
33Calculate the rate constant k and its units.
Initial rate of disappearance HgCl2
mol L-1 min-1
34What is the average rate of disappearance of
C2O42- in trial 1?
Initial rate of disappearance HgCl2
mol L-1 min-1
35Use the data provided to write the rate law and
indicate the order of the reaction with respect
to NO2 and CO (support your answers). Also give
the overall order of the reaction.
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37Calculate the rate constant k and its units.
38What is the average rate of disappearance of CO
in trial 2?
39How do we make these charts?
Initial rate of disappearance HgCl2
mol L-1 min-1
- Rates can be measured experimentally using a
variety techniques - moniter pH changes
- Titrations
- Change in volume or mass (gas production)
- Basically we can use any method to follow a
reaction that produces a measurable change.
40How do we make these charts?
Initial rate of disappearance HgCl2
mol L-1 min-1
One important method involves the spectroscopic
determination of concentration through Beers
Law.
41Using Beers Law to Determine vs. time.
- For each trial, the reactants are mixed and the
reaction mixture is transferred into a test tube
or cuvette. - Without any delay, the reaction vessel is placed
into a spectrophotometer. The absorbance data is
then collected at the wavelength of maximum
absorbance as a function of time. - This absorbance data is then converted to
concentration data using Beers Law A ? l c
42Fe(s)CuSO4(aq)?Fe2SO4(aq)Cu(s)
- The solution gradually gets paler as the
concentration of copper sulfate decreases and the
concentration of iron sulfate increases.
Concentration of copper sulfate solution
1M 0.8M 0.6M 0.4M 0.2M
0s 30s 90s 200s 500s
43Using Beers Law to Determine vs. time.
- A graph of concentration vs. time can be prepared
and then used to experimentally determine the
rate.
44- What does this tangent allow us to measure?
45- We would want to use a tangent to measure the
initial rate
46Half Life of a First Order Reaction
- Half-life is the time required to convert one
half of a reactant to product. - For first-order reactions, half-life is often
used as a representation for the rate constant. - This is because the half-life of a first-order
reaction and the rate constant are inversely
proportional, and the half-life is independent of
concentration.
47Radioactivity
- Radioactive decay is the spontaneous breakdown of
unstable atoms into more stable atoms with the
simultaneous emission of particles and rays. - Radioactive decay occurs at a constant rate that
is a first order process.
48Radioactivity and Half - Life
- The half-life of carbon-14 is 5730 years.
- How old is a bone that has about 12.5 of the
carbon-14 that a living organism would have in it?
49Carbon Dating
50Big Question
- How can we experimentally determine the order of
a reaction?
51Make 3 Graphs
- In order to determine order of reactant, A. We
must collect data consisting of concentration
versus time. - One common way to determine concentration vs.
time data is through the use of a
spectrophotometer.
52Make 3 Graphs
- We then use the data to make three graphs.
- A versus t
- ln A versus t
- 1 / A versus t
- By examining these graphs we can determine the
order of the reaction with respect to a
particular reactant and determine the rate
constant.
53A versus t (linear for a zero order reaction)
k must be a positive number.
54ln A versus t (linear for a 1st order reaction)
551 / A versus t (linear for a 2nd order reaction)
56Collision Model
- Key Idea Molecules must collide to react.
- However, only a small fraction of collisions
produces a reaction. Why?
57Two Factors
- Collisions must have enough energy to produce the
reaction (must equal or exceed the activation
energy). - Orientation of reactants must allow formation of
new bonds.
582HI ? H2 2I
59Concentration and Collision Theory
- Why does an increase in concentration cause an
increase in reaction rate?
60Concentration and Collision Theory
- Why does an increase in concentration cause an
increase in reaction rate? - Increasing the concentration increases the number
of collisions and therefore there are more
collisions leading to product.
61Temperature and Collision Theory
- Why does a temperature increase cause the
reaction rate to increase?
62Temperature and Collision Theory
- Why does a temperature increase cause the
reaction rate to increase? - At higher temperatures there are more collisions
and a greater percentage of the collisions have
the energy necessary to create a successful
collision.
63Activation Energy
- The activation energy is the minimum amount of
energy necessary for a reaction to occur.
64Temperature and Activation Energy (Ea)
65Activation Energy
- The activation energy can also be thought of as
the energy necessary to form an activated complex
during a collision between reactants.
66Transition State Theory
- The activated complex is a hypothetical species
lying between reactants and products at a point
on the reaction profile called the transition
state.
67The activated complex is a transition state
between reactants and products where old bonds
have begun to break and new bonds have started to
form. It cannot be isolated.
68Determining the Activation EnergyThe Arrhenius
Equation
- Collisions must have enough energy to produce the
reaction (must equal or exceed the activation
energy). - Orientation of reactants must allow formation of
new bonds.
69Arrhenius Equation
A
- k rate constant
- A frequency factor
- Ea activation energy
- T temperature
- R ideal gas constant
frequency factor a value in the Arrhenius
equation indicating how many collisions have the
correct orientation to lead to products.
70Arrhenius EquationDetermination of Activation
Energy
- Graphical determination of activation energy
(Ea). - plot the ln k on the y-axis.
- Plot 1/T (use Kelvin temperature) on the x-axis.
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72Arrhenius EquationDetermination of Activation
Energy
- A plot of ln k versus 1/T (using Kelvin) will
have - slope of Ea/R
- y-intercept of the graph is ln A.
73ln A
Ea
Slope -
R
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75x-axis
y-axis
76x1,y1 1.25 x 10-3, -2.593
x2,y2 1.78 x 10-3, -14.447
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78For two reactions at the same temperature, the
reaction with the higher activation energy has
the lower rate constant (k) and the slower rate.
79For two reactions at the same temperature, the
reaction with the higher activation energy has
the lower rate constant (k) and the slower rate.
Disregard R, T and A and focus on k and Ea
ln k is proportional to ? Ea
Because Ea is negative a higher activation energy
results in a lower rate constant (k).
Hidden Teacher Only Slide
802O3 ? 3O2
- A chemical equation like the one above does not
tell us how reactants become products - it is
simply a summary of the overall reaction.
81- The reaction 2O3 ? 3O2
- Is proposed to occur through the two step process
given below - O3 ? O2 O
- O3 O ? 2O2
This two step process is an example of a reaction
mechanism
82Reaction Mechanisms
- A reaction mechanism is a step-by-step
description of a chemical reaction. - Each step is called an elementary reaction.
83Often Used Terms
- Intermediate formed in one step and used up in
a subsequent step and so is never seen as a
product. - Molecularity the number of species that must
collide to produce the reaction indicated by that
step. - Elementary Step A step within a reaction
mechanism whose rate law can be written from its
molecularity.
84Reaction Mechanisms
- Elementary Steps
- Molecularity the number of molecules present in
an elementary step. - Unimolecular one molecule in the elementary
step. - Bimolecular two molecules in the elementary
step. - Termolecular three molecules in the elementary
step.
85Reaction Mechanisms
- Elementary Steps
- It is not common to see termolecular processes
(statistically improbable). - Unimolecular reactions occur because collisions
with other molecules provide the activation
energy for the molecule to react. - Bimolecular reactions involve the collision of
two particles with sufficient energy and proper
orientation. - Termolecular reactions involve the simultaneous
collision of three particles with sufficient
energy and proper orientation.
86Reaction Mechanisms
- Rate Laws for Elementary Steps
- The rate law of an elementary step is determined
by its molecularity - Unimolecular processes are first order,
- Bimolecular processes are second order, and
- Termolecular processes are third order.
87Reaction Mechanisms
Rate Laws for Elementary Steps
88The Rate Determining Step
89Rate-Determining Step
- In a reaction mechanism, the rate determining
step is the slowest step. It therefore
determines the rate of reaction.
90Reaction Mechanisms
- Reaction mechanisms must be consistent with
- Stoichiometry for the overall reaction.
- The experimentally determined rate law.
91NO2(g) CO(g) ? NO(g) CO2(g)
- Reaction mechanism must be consistent with the
stoichiometry of the overall reaction. - Is the mechanism below consistent with the
overall reaction above? - NO2(g) NO2(g) ? NO3(g) NO(g)
- NO3(g) CO(g) ? NO2(g) CO2(g)
92Determining the stoichiometry of a reaction
mechanism.
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93Reaction Mechanisms
- The reaction mechanism must also support the rate
law.
94Reaction Mechanisms
- Rate Laws for Multistep Mechanisms
- with an initial fast step.
- Consider the reaction
- 2NO(g) Br2(g) ? 2NOBr(g)
95Reaction Mechanisms
- Mechanisms with an Initial Fast Step
- 2NO(g) Br2(g) ? 2NOBr(g)
- The experimentally determined rate law is
- Rate kNO2Br2
- Consider the following mechanism
96- The rate law is (based on Step 2)
- Rate k2NOBr2NO
- The rate law should not depend on the
concentration of an intermediate (intermediates
are usually unstable). - NOBr2 is an unstable intermediate, so we express
the concentration of NOBr2 in terms of NO and Br2
Since there is an equilibrium in step 1 we have
97- By definition of equilibrium
- Therefore, the overall rate law becomes
- Note the final rate law is consistent with the
experimentally observed rate law.
98Student ExampleDetermine the rate law for the
reaction and the balanced equation given the
mechanism below
- 2NO ? N2O2 fast
- N2O2 O2 ? 2NO2 slow
99- Use the slide that follows to show the students
the method for determining the slow step for each
mechanism. - It can only be used for students practicing how
to relate rate law to the mechanism if you treat
the first step of Mechanism B and C as an
equilibrium and H2O is treated as a liquid and
is therefore 1.
100Assume the rate law is Rate kH2O2H3OI- W
hich step would be the rate determining step?
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101This diagram shows a two-step mechanism for a
reaction with the first step being rate
determining.
102What is the mechanism for the reaction?
Overall Reaction
103Mechanism for Previous Reaction
- NO H2 ? NOH2 slow
- NO NOH2 ? N2O H2O fast
-
104Catalysts
- A catalyst is a substance that increases the rate
of a chemical reaction by reducing the activation
energy, but which is left unchanged by the
reaction.
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106What is the overall reaction?
107What is the overall reaction?
108Identify the intermediates.
109Identify the intermediates.
NO is a catalyst
A homogeneous catalyst is of the same phase as
the reacting substances. It lowers the
activation energy by forming intermediates which
allow the reaction to proceed by a different
pathway.
110Heterogeneous Catalysts
- A heterogeneous catalyst is of a different phase
than the reacting substances. - It provides a surface on which the transition
state is stabilized thus lowering the activation
energy and increasing the reaction rate.
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112Catalytic Converter A Heterogeneous Catalyst
- In a catalytic converter, the catalyst (in the
form of platinum and palladium) is coated onto a
ceramic honeycomb that are housed in a
muffler-like package attached to the exhaust
pipe. The catalyst helps to convert harmful
exhaust gases into safer ones.
113Catalysis
- Catalysis is the process of using a catalyst to
speed up a reaction.
114Heterogeneous and Homogeneous Catalysts can
Catalysis different Types of Reactions
- Acid Base Catalysis
- Enzyme Catalysis
- Surface Catalysis
115Acid Base Catalysis
- A chemical reaction is catalyzed by an acid or
a base. - A reactant either gains or loses a proton (H)
which causes an increase in the rate of the
reaction. - In acid catalysis all species capable of donating
protons contribute to reaction rate acceleration
with the strongest acids being most effective.
116Enzyme Catalysis
- Some enzymes accelerate reactions by binding to
the reactants in a way that lowers the activation
energy. - Other enzymes react with a reactant species to
form a new intermediate. - Enzyme catalysis essentially occurs when
substances catalyze reactions within a living
organism.
117Enzyme Catalysis
- There is a coulombic attraction between the
substrate and the enzyme. - After the reaction occurs the products do not
exhibit the coulombic attraction with the enzyme
due to changes in their structure and are
therefore are released by the enzyme.
118Substrate Interactions with the Active Sites in
Enzyme Catalysis
119Surface Catalysis
- In surface catalysis, either a new reaction
intermediate is formed, or the probability of a
successful collision is modified. - Most catalysts fall into this category.
120Inhibitor
- An inhibitor decreases the rate of a reaction.
It often does this by rendering a catalyst
ineffective.
Catalyst poisoning occurs when a catalytic
converter is exposed to exhaust containing
substances that coat the working surfaces,
encapsulating the catalyst so that it cannot
contact and treat the exhaust. The most notable
contaminant is lead, so vehicles equipped with
catalytic converters can only be run on unleaded
gasoline.
121Inhibitor
- An inhibitor decreases the rate of a reaction.
122Reaction Rate Lab
123Reaction Rate Lab Part A
- Use different containers for Reaction Mixtures I
and II. - Dont forget the starch.
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127Reaction Rate Lab Part B
- In part A you will perform five different trials
with various concentrations.
128Reaction Rate Lab Part C
- In part B you will perform trial 1 using a
catalyst.
129We will not be performing part C of the lab.