Chapter 14 Homework The Activation Energy The Difference Energy Between

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

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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

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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

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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

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• 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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• The speed of a reaction is defined as the change that occurs per unit time.

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“Appearing Red” from Live Demonstrations

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“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

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• 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.

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Reaction Rates and Stoichiometry

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• 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)

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Figure 14.6 from Transparency Pack

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“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

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• 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

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• 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

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• 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

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• 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.

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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

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• 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:

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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 +−=

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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

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• 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

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• 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]

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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)

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“Catalysis” from Further Readings

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“Catalysis” Movie from Instructor’s Resource CD/DVD

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“Catalytic Decomposition of Hydrogen Peroxide: Foam Production” from Live Demonstrations

Chapter 14

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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.

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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.

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• 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).

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