Chemical Kinetics

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Chapter 14

• Chemical Kinetics

2
Review Section of Chapter 14 Test

• Net Ionic Equations Update to different
  material. Maybe Chapter 11 stuff

3
Reaction 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.

4
Reaction 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
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Reaction 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
6
Reaction 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)
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Reaction 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?
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2H2O2(aq) ? 2H2O(l) O2(g)
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2H2O2(aq) ? 2H2O(l) O2(g)
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2H2O2(aq) ? 2H2O(l) O2(g)
What is the rate of the reaction from 0s to 2.16
x 104s?
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2H2O2(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
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General 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.
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15-1 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
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Rates 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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What does the slope of the line represent?
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What is the concentration at 100s for the
reaction 2H2O2(aq) ? 2H2O(l) O2(g)?Given
H2O2i 2.32 M
2.15 M
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What 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.

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Factors Affecting Reaction Rates

1. The nature of the reacting substances.

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Factors Affecting Reaction Rates

• 2. The state of subdivision of the reacting
  substances (surface area).

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Lycopodium Powder
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Factors Affecting Reaction Rates

• 3. The temperature of the reacting substances.

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Factors Affecting Reaction Rates

• 4. The concentration of the reacting substances.
  
• (Except in zero order reactions)

Air (21 oxygen) 100 oxygen
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Factors Affecting Reaction Rates

• The presence of a catalyst.
• Catalysts speed up reactions but are left
  unchanged by the reaction.

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The 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 .
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Temperature 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.

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Temperature dependence of k
.
.
.
.
.
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Concentration 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.

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Use 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.
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First determine the order of HgCl2
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Next determine the order of C2O42-
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Now write the rate law and determine the order of
the reaction.
33
Calculate the rate constant k and its units.
Initial rate of disappearance HgCl2
mol L-1 min-1
34
What is the average rate of disappearance of
C2O42- in trial 1?
Initial rate of disappearance HgCl2
mol L-1 min-1
35
Use 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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Calculate the rate constant k and its units.
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What is the average rate of disappearance of CO
in trial 2?
39
How 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.

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How 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.
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Using 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

42
Fe(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
43
Using Beers Law to Determine vs. time.

• A graph of concentration vs. time can be prepared
  and then used to experimentally determine the
  rate.

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• What does this tangent allow us to measure?

45

• We would want to use a tangent to measure the
  initial rate

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Half 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.

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Radioactivity

• 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.

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Radioactivity 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?

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Carbon Dating
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Big Question

• How can we experimentally determine the order of
  a reaction?

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Make 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.

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Make 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.

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A versus t (linear for a zero order reaction)
k must be a positive number.
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ln A versus t (linear for a 1st order reaction)
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1 / A versus t (linear for a 2nd order reaction)
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Collision Model

• Key Idea Molecules must collide to react.
• However, only a small fraction of collisions
  produces a reaction. Why?

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Two 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.

58
2HI ? H2 2I
59
Concentration and Collision Theory

• Why does an increase in concentration cause an
  increase in reaction rate?

60
Concentration 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.

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Temperature and Collision Theory

• Why does a temperature increase cause the
  reaction rate to increase?

62
Temperature 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.

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Activation Energy

• The activation energy is the minimum amount of
  energy necessary for a reaction to occur.

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Temperature and Activation Energy (Ea)
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Activation Energy

• The activation energy can also be thought of as
  the energy necessary to form an activated complex
  during a collision between reactants.

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Transition 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.

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The 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.
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Determining 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.

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Arrhenius 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.
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Arrhenius 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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Arrhenius 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.

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ln A
Ea
Slope -
R
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x-axis
y-axis
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x1,y1 1.25 x 10-3, -2.593
x2,y2 1.78 x 10-3, -14.447
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For two reactions at the same temperature, the
reaction with the higher activation energy has
the lower rate constant (k) and the slower rate.
79
For 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
80
2O3 ? 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
82
Reaction Mechanisms

• A reaction mechanism is a step-by-step
  description of a chemical reaction.
• Each step is called an elementary reaction.

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Often 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.

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Reaction 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.

85
Reaction 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.

86
Reaction 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.

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Reaction Mechanisms
Rate Laws for Elementary Steps
88
The Rate Determining Step
89
Rate-Determining Step

• In a reaction mechanism, the rate determining
  step is the slowest step. It therefore
  determines the rate of reaction.

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Reaction Mechanisms

• Reaction mechanisms must be consistent with
• Stoichiometry for the overall reaction.
• The experimentally determined rate law.

91
NO2(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)

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Determining the stoichiometry of a reaction
mechanism.
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93
Reaction Mechanisms

• The reaction mechanism must also support the rate
  law.

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Reaction Mechanisms

• Rate Laws for Multistep Mechanisms
• with an initial fast step.
• Consider the reaction
• 2NO(g) Br2(g) ? 2NOBr(g)

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Reaction 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.

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Student 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.

100
Assume the rate law is Rate kH2O2H3OI- W
hich step would be the rate determining step?
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101
This diagram shows a two-step mechanism for a
reaction with the first step being rate
determining. 
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What is the mechanism for the reaction? 
Overall Reaction
103
Mechanism for Previous Reaction

• NO H2 ? NOH2 slow
• NO NOH2 ? N2O H2O fast

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Catalysts

• 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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What is the overall reaction?

• O3 ? O2 O
• O3 O ? 2O2

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What is the overall reaction?
108
Identify the intermediates.
109
Identify 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.
110
Heterogeneous 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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Catalytic 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.

113
Catalysis

• Catalysis is the process of using a catalyst to
  speed up a reaction.

114
Heterogeneous and Homogeneous Catalysts can
Catalysis different Types of Reactions

• Acid Base Catalysis
• Enzyme Catalysis
• Surface Catalysis

115
Acid 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. 

116
Enzyme 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.

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Enzyme 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.

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Substrate Interactions with the Active Sites in
Enzyme Catalysis
119
Surface 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.

120
Inhibitor

• 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.
121
Inhibitor

• An inhibitor decreases the rate of a reaction.

122
Reaction Rate Lab
123
Reaction Rate Lab Part A

• Use different containers for Reaction Mixtures I
  and II.
• Dont forget the starch.

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Reaction Rate Lab Part B

• In part A you will perform five different trials
  with various concentrations.

128
Reaction Rate Lab Part C

• In part B you will perform trial 1 using a
  catalyst.

129
We will not be performing part C of the lab.