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Chapter 14. Chemical Kinetics
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 14.4 Concentration of Butyl Chloride 14.2 Reaction Rates
(C4H9Cl) as a Function of Time
Figure 14.6 Components of a Spectrometer 14.2 Reaction Rates
Figure 14.8 Kinetic Data for Conversion of Methyl 14.4 The Change of Concentration with Time
Isonitrile into Acetonitrile
Figure 14.9 Kinetic Data for Decomposition of NO2 14.4 The Change of Concentration with Time
Activities: Section:
Progress of Reaction 14.2 Reaction Rates
Decomposition of N2O5 14.3 Concentration and Rate Laws
Rates of Reaction 14.3 Concentration and Rate Laws
Arrhenius Model 14.5 Temperature and Rate
Animations: Section:
First-Order Process 14.4 The Change of Concentration with Time
Movies: Section:
CFCs and Stratospheric Ozone 14.4 The Change of Concentration with Time
Catalysis 14.7 Catalysis
3-D Models: Section:
Chlorine 14.3 Concentration and Rate Laws
Chemical Kinetics
197
Hydrogen Peroxide 14.7 Catalysis
Ethylene 14.7 Catalysis
Ethane 14.7 Catalysis
FeMo-cofactor of Nitrogenase 14.7 Catalysis
Nitrogen Dioxide 14.7 Catalysis
Other Resources
Further Readings: Section:
The Fizz Keeper, a Case Study in Chemical 14.1 Factors that Affect Reaction Rates
Education, Equilibrium, and Kinetics
Inflation Rates, Car Devaluation, and Chemical 14.3 Concentration and Rate Laws
The Arrhenius Law and Storage of Food 14.5 Temperature and Rate
Visualizing the Transition State: A Hands-On 14.5 Temperature and Rate
Approach to the Arrhenius Equation
Pictorial Analogies XIII: Kinetics and Mechanisms 14.6 Reaction Mechanisms
Doing the Dishes: An Analogy for Use in Teaching 14.6 Reaction Mechanisms
Reaction Kinetics
Auto Analogies 14.6 Reaction Mechanisms
Another Auto Analogy: Rate-Determining Steps 14.6 Reaction Mechanisms
Catalysis 14.7 Catalysis
Catalysis: New Reaction Pathways, Not Just a 14.7 Catalysis
Live Demonstrations: Section:
Appearing Red 14.1 Factors that Affect Reaction Rates
A New Twist on the Iodine Clock Reaction: 14.1 Factors that Affect Reaction Rates
Chapter 14
198
Determining the Order of a Reaction
Hydrogen Peroxide Iodine Clock: Oxidation of 14.1 Factors that Affect Reaction Rates
Potassium Iodide by Hydrogen Peroxide
Chemical Kinetics
199
Chapter 14. Chemical Kinetics
Common Student Misconceptions
• Students often assume that reaction orders may be determined from stoichiometric coefficients
regardless of the reaction mechanism.
Teaching Tips
• It is possible for mathematics to get in the way of some students’ understanding of the chemistry of
this chapter.
Lecture Outline
14.1 Factors that Affect Reaction Rates
1
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2
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4
• The speed at which a chemical reaction occurs is the reaction rate.
• Chemical kinetics is the study of how fast chemical reactions occur.
• There are several important factors that affect rates of reactions:
14.2 Reaction Rates
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6
• The speed of a reaction is defined as the change that occurs per unit time.
1
“Appearing Red” from Live Demonstrations
2
“A New Twist on the Iodine Clock Reaction: Determining the Order of a Reaction” from Live
Chapter 14
200
s
M
023.0
s 0s 20
M 0.000.46M
Rate Avg =
−
−
=
• It is often determined by measuring the change in concentration of a reactant or product with
time.
• For a reaction A → B
• For the reaction A → B there are two ways of measuring rate:
• the rate of appearance of product B (i.e., change in moles of B per unit time) as in the preceding
example, and
• the rate of disappearance of reactant A (i.e., the change in moles of A per unit time).
Change of Rate with Time
7
• In most chemical reactions we will determine the reaction rate by monitoring a change in
concentration (of a reactant or product).
• The most useful unit to use for rate is molarity.
Instantaneous Rate
• We can plot [C4H9Cl] versus time.
• The rate at any instant in time is called the instantaneous rate.
• It is the slope of the straight line tangent to the curve at that instant.
• Instantaneous rate is different from average rate.
Chemical Kinetics
201
Reaction Rates and Stoichiometry
8
• For the reaction: C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq)
• The rate of appearance of C4H9OH must equal the rate of disappearance of C4H9Cl.
• What if the stoichiometric relationships are not one-to-one?
• For the reaction: 2HI(g) → H2(g) + I2(g)
14.3 Concentration and Rate Laws
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• In general, rates:
• increase when reactant concentration is increased.
• decrease as the concentration of reactants is reduced.
• We often examine the effect of concentration on reaction rate by measuring the way in which reaction
rate at the beginning of a reaction depends on starting conditions.
• Consider the reaction: NH4+(aq) + NO2– (aq) → N2(g) + 2H2O(l)
8
Figure 14.6 from Transparency Pack
9
“Decomposition of N2O5” Activity from Instructor’s Resource CD/DVD
tt
=
−= OHHCClHC
Rate 9494
Chapter 14
202
• Once we have determined the rate law and the rate constant, we can use them to calculate initial
reaction rates under any set of initial concentrations.
Reaction Orders: The Exponents in the Rate Law
12
• For a general reaction with rate law:
Rate = k[reactant 1]m[reactant 2]n
• The exponents m and n are called reaction orders.
Magnitudes and Units of Rate Constants
• In comparing reactions to evaluate which ones are relatively fast and which are relative slow, the rate
Using Initial Rates to Determine Rate Laws
13
• To determine the rate law, we observe the effect of changing initial concentrations.
FORWARD REFERENCES
• The importance of pH in determining the rates of proton transfer reactions in biological
systems will be mentioned in Chapter 16 (section 16.4).
14.4 The Change of Concentration with Time
14
• Goal: Convert the rate law into a convenient equation that gives concentration as a function of time.
First-Order Reactions
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• For a first-order reaction, the rate doubles as the concentration of a reactant doubles.
• Therefore, we can write the differential rate law:
• A plot of ln[A]t versus t is a straight line with slope −k and intercept ln[A]0.
• Note that in this equation we use the natural logarithm, ln (log to the base e).
Second-Order Reactions
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• A second-order reaction is one whose rate depends on the reactant concentration to the second
power or on the concentration of two reactants, each raised to the first power.
• For a second-order reaction with just one reactant, we write the differential rate law:
• Integrating, we get the integrated form of the rate law:
Zero-Order Reactions
21
• A zero-order reaction is one whose rate is independent of the reactant concentration.
Chapter 14
204
Half-life
• Half-life, t½ , is the time required for the concentration of a reactant to decrease to half its original
value.
• That is, half life, t½, is the time taken for [A]0 to reach ½ [A]0.
• Mathematically, the half life of a first-order reaction is:
• Note that the half-life of a first-order reaction is independent of the initial concentration of the
reactant.
• We can show that the half-life of a second order reaction is:
• Note that, unlike for the first-order reaction, the half-life of a second-order reaction is dependent
on the initial concentration of the reactant.
FORWARD REFERENCES
• Rates of radioactive decay processes, half-lives of radioactive isotopes, and radiocarbon
dating will be further discussed in Chapter 21 (section 21.4).
14.5 Temperature and Rate
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• Most reactions speed up as temperature increases.
• We can illustrate this with chemiluminescent Cyalume® light sticks.
Chemical Kinetics
205
The Collision Model
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• Rates of reactions are affected by concentration and temperature.
• We need to develop a model that explains this observation.
• An explanation is provided by the collision model, based on kinetic-molecular theory.
• In order for molecules to react, they must collide.
The Orientation Factor
29
• The orientation of a molecule during collision can have a profound effect on whether or not a reaction
occurs.
• Consider the reaction between Cl and NOCl:
Activation Energy
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• Arrhenius: Molecules must posses a minimum amount of energy to react. Why?
• In order to form products, bonds must be broken in the reactants.
• Bond breakage requires energy.
• Molecules moving too slowly, with too little kinetic energy, don’t react when they collide.
• Activation energy, Ea, is the minimum energy required to initiate a chemical reaction.
• Ea will vary with the reaction.
• Consider the rearrangement of methyl isonitrile to form acetonitrile:
Chapter 14
206
• The energy associated with the molecule drops.
• The energy barrier between the starting molecule and the highest energy state found along the
reaction pathway is the activation energy.
• The species at the top of the barrier is called the activated complex or transition state.
• The change in energy for the reaction is the difference in energy between CH3NC and CH3CN.
• ∆Erxn has no effect on reaction rate.
• The activation energy is the difference in energy between reactants, (CH3NC) and the transition
state.
The Arrhenius Equation
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• Arrhenius discovered that most reaction-rate data obeyed an equation based on three factors:
• The number of collisions per unit time.
• The fraction of collisions that occur with the correct orientation.
• The fraction of the colliding molecules that have an energy equal to or greater than Ea.
Determining the Activation Energy
• Ea may be determined experimentally.
• We need to take the natural log of both sides of the Arrhenius equation:
A
RT
E
kalnln +−=
Chemical Kinetics
207
FORWARD REFERENCES
• The role of temperature in affecting the position of equilibrium will be discussed in Chapter
15 (section 15.7).
14.6 Reaction Mechanisms
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• The balanced chemical equation provides information about substances present at the beginning and
end of the reaction.
• The reaction mechanism is the process by which the reaction occurs.
• Mechanisms provide a picture of which bonds are broken and formed during the course of a reaction.
Elementary Reactions
37
• Elementary reactions or elementary processes are any processes that occur in a single step.
Multistep Mechanisms
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• A multistep mechanism consists of a sequence of elementary steps.
• The elementary steps must add to give the balanced chemical equation.
• Some multistep mechanisms will include intermediates.
Rate Laws of Elementary Reactions
• The rate laws of the elementary steps determine the overall rate law of the reaction.
• The rate law of an elementary step is determined by its molecularly.
The Rate-Determining Step for a Multistep Mechanism
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• Most reactions occur by mechanisms with more than one elementary step.
• Often one step is much slower than the others.
• The slow step limits the overall reaction rate.
• This is called the rate-determining step (rate-limiting step) of the reaction.
“Another Auto Analogy: Rate Determining Steps” from Further Readings
Chapter 14
208
• This step governs the overall rate law for the overall reaction.
Mechanisms with a Slow Initial Step
• Consider the reaction: NO2(g) + CO(g) → NO(g) + CO2(g)
• The experimentally derived rate law is: Rate = k[NO2]2
• We propose a mechanism for the reaction:
Mechanisms with a Fast Initial Step
• Consider the reaction:
2NO(g) + Br2(g) → 2NOBr(g)
• The experimentally determined rate law is:
Rate = k[NO]2[Br2]
• Consider the following proposed mechanism:
Rate = k2[NOBr2][NO]
• Problem: This rate law depends on the concentration of an intermediate species.
• Intermediates are usually unstable and have low/unknown concentrations.
• We need to find a way to remove this term from our rate law.
• We can express the concentration of [NOBr2] in terms of NOBr and Br2 by assuming that
there is an equilibrium in step 1.
• In a dynamic equilibrium, the forward rate equals the reverse rate.
• Therefore, by definition of equilibrium we get:
k1[NO][Br2] = k–1[NOBr2]
Chemical Kinetics
209
FORWARD REFERENCES
• The relationship between k1 and k−1 will be exploited in Chapter 15 (section 15.1).
• Mechanism of organic addition reactions will be discussed in Chapter 24 (section 24.4).
14.7 Catalysis
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• A catalyst is a substance that changes the rate of a chemical reaction without itself undergoing a
permanent chemical change in the process.
• There are two types of catalysts:
• homogeneous and
• heterogeneous.
• Catalysts are common in the body, in the environment, and in the chemistry lab!
Homogeneous Catalysis
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• A homogeneous catalyst is one that is present in the same phase as the reacting molecules.
• For example, hydrogen peroxide decomposes very slowly in the absence of a catalyst:
2H2O2(aq) → 2H2O(l) + O2(g)
41
“Catalysis” from Further Readings
42
“Catalysis” Movie from Instructor’s Resource CD/DVD
43
“Catalytic Decomposition of Hydrogen Peroxide: Foam Production” from Live Demonstrations
Chapter 14
210
Heterogeneous Catalysis
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• A heterogeneous catalyst exists in a different phase than the reactants.
• Often we encounter a situation involving a solid catalyst in contact with gaseous reactants and
gaseous products (example: catalytic converters in cars) or with reactants in a liquid.
• Many industrial catalysts are heterogeneous.
• How do they do their job?
• The first step is adsorption (the binding of reactant molecules to the catalyst surface).
• For example, consider the hydrogenation of ethylene to form ethane:
C2H4(g) + H2(g) → C2H6(g) ∆H° = −137 kJ/mol
• The reaction is slow in the absence of a catalyst.
• In the presence of a finely divided metal catalyst (Ni, Pt, or Pd) the reaction occurs quickly at
room temperature.
• First, the ethylene and hydrogen molecules are adsorbed onto active sites on the metal
surface.
Chemical Kinetics
211
Enzymes
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• Enzymes are biological catalysts.
• Most enzymes are large protein molecules.
• Molar masses are in the range of 104 to 106 amu.
• Enzymes are capable of catalyzing very specific reactions.
• Here, a substrate is pictured as fitting into the active site of an enzyme in a manner similar to a
specific key fitting into a lock. This forms an enzyme-substrate (ES) complex.
• Only substrates that fit into the enzyme lock can be involved in the reaction.
• The enzyme’s active site and the substrate thus have complementary shapes.
• However, there may be a significant amount of flexibility at the active site.
• Enzymes are extremely efficient catalysts.
• The number of individual catalytic events occurring at an active site per unit time is called the
turnover number.
• Large turnover numbers correspond to very low Ea values.
• For enzymes, turnover numbers are very large (typically 103 – 107 per second).
Nitrogen Fixation and Nitrogenase
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• Nitrogen gas cannot be used in the soil for plants or animals.
• Nitrogen compounds, NH3, NO2–, and NO3– are used in the soil.
Chapter 14
212
• The conversion between N2 and NH3 is a process with a high activation energy (the N2 triple bond
FORWARD REFERENCES
• The effect of catalysts on equilibrium and catalytic converters will be discussed in Chapter 15
(section 15.7).
• Metal ions will be mentioned as integral parts of many enzymes in Chapter 23 (section 23.4).
Chemical Kinetics
213
Further Readings:
1. Reed A. Howald, “The Fizz Keeper, a Case Study in Chemical Education, Equilibrium, and Kinetics,”
J. Chem. Educ., Vol. 76, 1999, 208–209.
6. Elizabeth Wilson, “Light Sticks,” Chemical and Engineering News, January 18, 1999, 65. A brief
article on chemiluminescent light sticks.
7. Kent W. Piepgrass, “Audience-Appropriate Analogies: Collision Theory,” J. Chem. Educ., Vol. 75,
1998, 72.
8. Lee A. Krug, “The Collision Theory and an American Tradition,” J. Chem. Educ., Vol. 64, 1987,
1000.
13. Arthur M. Last, “Doing the Dishes: An Analogy for Use in Teaching Reaction Kinetics,” J. Chem.
Educ., Vol. 62, 1985, 1015–1016.
14. Richard A. Potts, “Auto Analogies,” J. Chem. Educ., Vol. 62, 1985, 579. This brief article includes
analogies for reaction mechanisms and rate-determining steps of a reaction.
Chapter 14
214
18. Cynthia M. Friend, “Catalysis on Surfaces,” Scientific American, April 1993, 74–79.
19. Sir John Meurig Thomas, “Solid Acid Catalysts,” Scientific American, April 1992, 112–118.
23. Robert J. Farrauto, Ronald M. Heck, and Barry K. Speronello, “Environmental Catalysts,” Chemical
and Engineering News, September 7, 1992, 34–44.
24. H. Alan Rowe and Morris Brown, “Practical Enzyme Kinetics,” J. Chem. Educ., Vol. 65, 1988, 548–
549.
25. Todd P. Silverstein, “Breaking Bonds versus Chopping Heads: The Enzyme as Butcher,” J. Chem.
Educ., Vol. 72, 1995, 645–646.
Live Demonstrations:
1. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Appearing Red,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical Society, 1988),
pp. 145–146. An introductory kinetics experiment.
Chemical Kinetics
215
4. Lee. R. Summerlin, and James. L. Ealy, Jr., “The Starch-Iodine Clock Reaction,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 1 (Washington: American Chemical Society, 1988),
pp.107–108. The classic iodine clock experiment.
5. Bassam Z. Shakhashiri, “Lightsticks,” Chemical Demonstrations: A Handbook for Teachers of
Chemistry, Volume 1 (Madison: The University of Wisconsin Press, 1983), pp. 146–152.
