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Chapter 14Chemical Kinetics

Kinetics

- Studies the rate at which a chemical process

occurs. - Besides information about the speed at which

reactions occur, kinetics also sheds light on the

reaction mechanism (exactly how the reaction

occurs).

Factors That Affect Reaction Rates

- Physical State of the Reactants
- In order to react, molecules must come in contact

with each other. - The more homogeneous the mixture of reactants,

the faster the molecules can react.

Factors That Affect Reaction Rates

- Concentration of Reactants
- As the concentration of reactants increases, so

does the likelihood that reactant molecules will

collide.

Factors That Affect Reaction Rates

- Temperature
- At higher temperatures, reactant molecules have

more kinetic energy, move faster, and collide

more often and with greater energy.

Factors That Affect Reaction Rates

- Presence of a Catalyst Catalysts speed up

reactions by changing the mechanism of the

reaction. - Catalysts are not consumed during the course of

the reaction.

Reaction Rates

- Rates of reactions can be determined by

monitoring the change in concentration of either

reactants or products as a function of time.

Reaction Rates

C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)

- In this reaction, the concentration of butyl

chloride, C4H9Cl, was measured at various times.

Reaction Rates

C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)

- The average rate of the reaction over each

interval is the change in concentration divided

by the change in time

Reaction Rates

C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)

- Note that the average rate decreases as the

reaction proceeds. - This is because as the reaction goes forward,

there are fewer collisions between reactant

molecules.

Reaction Rates

C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)

- A plot of concentration vs. time for this

reaction yields a curve like this. - The slope of a line tangent to the curve at any

point is the instantaneous rate at that time.

Reaction Rates

C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)

- All reactions slow down over time.
- Therefore, the best indicator of the rate of a

reaction is the instantaneous rate near the

beginning.

Reaction Rates and Stoichiometry

C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)

- In this reaction, the ratio of C4H9Cl to C4H9OH

is 11. - Thus, the rate of disappearance of C4H9Cl is the

same as the rate of appearance of C4H9OH.

Reaction Rates and Stoichiometry

- What if the ratio is not 11?

2 HI(g) ??? H2(g) I2(g)

- Therefore,

Reaction Rates and Stoichiometry

- To generalize, then, for the reaction

Fig 14.3

Figure 14.3 Progress of a hypothetical reaction

A ??B. Each red sphere represents 0.01 mol A,

each blue sphere represents 0.01 mol B, and the

vessel has a volume of 1.00 L. (a) At time zero

the vessel contains 1.00 mol A (100 red spheres)

and 0 mol B (no blue spheres). (b) After 20 s the

vessel contains 0.54 mol A and 0.46 mol B. (c)

After 40 s the vessel contains 0.30 mol A and

0.70 mol B.

From the data given in the caption of Figure

14.3, calculate the average rate at which A

disappears over the time interval from 20 s to 40

s.

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PRACTICE EXERCISE For the reaction pictured in

Figure 14.3, calculate the average rate of

appearance of B over the time interval from 0 to

40 s. (The necessary data are given in the figure

caption.)

Answer 1.8 ? 10 2 M/s

Fig 14.4

Figure 14.4 Concentration of butyl chloride

(C4H9Cl) as a function of time. The dots

represent the experimental data from the first

two columns of Table 14.1, and the red curve is

drawn to connect the data points smoothly. Lines

are drawn that are tangent to the curve at t 9

and t 600 s. The slope of each tangent is

defined as the vertical change divided by the

horizontal change ?C4H9Cl/?t. The reaction

rate at any time is related to the slope of the

tangent to the curve at that time. Because C4H9Cl

is disappearing, the rate is equal to the

negative of the slope.

Using Figure 14.4, calculate the instantaneous

rate of disappearance of C4H9Cl at t 0 (the

initial rate).

SAMPLE EXERCISE 14.2 Calculating an Instantaneous

Rate of Reaction

PRACTICE EXERCISE Using Figure 14.4, determine

the instantaneous rate of disappearance of C4H9Cl

at t 300 s.

Answer 1.1 ? 10 4 M/s

(b) If the rate at which O2 appears, ?O2/?t, is

6.0 ? 105 M/s at a particular instant, at what

rate is O3 disappearing at this same time,

?O3/?t?

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Concentration and Rate

- One can gain information about the rate of a

reaction by seeing how the rate changes with

changes in concentration.

Concentration and Rate

- Comparing Experiments 1 and 2, when NH4

doubles, the initial rate doubles.

Concentration and Rate

- Likewise, comparing Experiments 5 and 6, when

NO2- doubles, the initial rate doubles.

Concentration and Rate

- This means
- Rate ? NH4
- Rate ? NO2-
- Rate ? NH NO2-
- or
- Rate k NH4 NO2-
- This equation is called the rate law, and k is

the rate constant.

Rate Laws

NH4(aq) NO2-(aq)

N2(g) 2 H2O(l)

- A rate law shows the relationship between the

reaction rate and the concentrations of

reactants. - The exponents tell the order of the reaction with

respect to each reactant. - This reaction is
- First-order in NH4
- First-order in NO2-

Rate Laws

- The overall reaction order can be found by adding

the exponents on the reactants in the rate law. - This reaction is second-order overall.

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PRACTICE EXERCISE Assuming that rate kAB,

rank the mixtures represented in this Sample

Exercise in order of increasing rate.

14.9 2N2O5 ? 4 NO2 O2 Rate k

N2O5 14.10 CHCl3 Cl2 ? CCl4 HCl Rate

kCHCl3 Cl2.5

PRACTICE EXERCISE 14.11 H2 I2 ? 2 HI Rate k

H2 I2 (a) What is the reaction order of the

reactant H2 in Equation 14.11? (b) What are the

units of the rate constant for Equation 14.11?

SAMPLE EXERCISE 14.6 Determining a Rate Law from

Initial Rate Data

Using these data, determine (a) the rate law for

the reaction, (b) the magnitude of the rate

constant, (c) the rate of the reaction when A

0.050 M and B 0.100 M.

(a) Determine the rate law for this reaction. (b)

Calculate the rate constant. (c) Calculate the

rate when NO 0.050 M and H2

0.150 M.

First-Order Processes

ln At -kt ln A0

- Therefore, if a reaction is first-order, a plot

of ln A vs. t will yield a straight line, and

the slope of the line will be -k.

First-Order Processes

- Consider the process in which methyl isonitrile

is converted to acetonitrile.

First-Order Processes

- This data was collected for this reaction at

198.9C.

First-Order Processes

- When ln P is plotted as a function of time, a

straight line results. - Therefore,
- The process is first-order.
- k is the negative slope 5.1 ? 10-5 s-1.

Second-Order Processes

- Similarly, integrating the rate law for a

process that is second-order in reactant A, we get

also in the form

y mx b

Second-Order Processes

- So if a process is second-order in A, a plot of

1/A vs. t will yield a straight line, and the

slope of that line is k.

Second-Order Processes

The decomposition of NO2 at 300C is described by

the equation

and yields data comparable to this

Time (s) NO2, M

0.0 0.01000

50.0 0.00787

100.0 0.00649

200.0 0.00481

300.0 0.00380

Second-Order Processes

- Graphing ln NO2 vs. t yields

- The plot is not a straight line, so the process

is not first-order in A.

Time (s) NO2, M ln NO2

0.0 0.01000 -4.610

50.0 0.00787 -4.845

100.0 0.00649 -5.038

200.0 0.00481 -5.337

300.0 0.00380 -5.573

Second-Order Processes

- Graphing ln 1/NO2 vs. t, however, gives this

plot.

- Because this is a straight line, the process is

second-order in A.

Time (s) NO2, M 1/NO2

0.0 0.01000 100

50.0 0.00787 127

100.0 0.00649 154

200.0 0.00481 208

300.0 0.00380 263

Fig 14.4

Figure 14.4 Concentration of butyl chloride

(C4H9Cl) as a function of time. The dots

represent the experimental data from the first

two columns of Table 14.1, and the red curve is

drawn to connect the data points smoothly. Lines

are drawn that are tangent to the curve at t 9

and t 600 s. The slope of each tangent is

defined as the vertical change divided by the

horizontal change ?C4H9Cl/?t. The reaction

rate at any time is related to the slope of the

tangent to the curve at that time. Because C4H9Cl

is disappearing, the rate is equal to the

negative of the slope.

Using Figure 14.4, calculate the instantaneous

rate of disappearance of C4H9Cl at t 0 (the

initial rate).

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Is the reaction first or second order in NO2?

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SAMPLE EXERCISE 14.8 continued

As Figure 14.8 shows, only the plot of 1/NO2

versus time is linear. Thus, the reaction obeys a

second-order rate law Rate kNO22. From the

slope of this straight-line graph, we determine

that k 0.543 M1 s1 for the disappearance of

NO2.

PRACTICE EXERCISE Consider again the

decomposition of NO2 discussed in the Sample

Exercise. The reaction is second order in NO2

with k 0.543 M1s1. If the initial

concentration of NO2 in a closed vessel is 0.0500

M, what is the remaining concentration after

0.500 h?

Answer Using Equation 14.14, we find NO2

1.00 ? 103 M

Half-Life

- Half-life is defined as the time required for

one-half of a reactant to react. - Because A at t1/2 is one-half of the original

A, - At 0.5 A0.

Half-Life

- For a first-order process, this becomes

ln 0.5 -kt1/2

-0.693 -kt1/2

NOTE For a first-order process, the half-life

does not depend on A0.

Half-Life

- For a second-order process,

PRACTICE EXERCISE (a) Using Equation 14.15,

calculate t1/2 for the decomposition of the

insecticide described in Sample Exercise 14.7.

(b) How long does it take for the concentration

of the insecticide to reach one-quarter of the

initial value?

Temperature and Rate

- Generally, as temperature increases, so does the

reaction rate. - This is because k is temperature dependent.

The Collision Model

- In a chemical reaction, bonds are broken and new

bonds are formed. - Molecules can only react if they collide with

each other.

The Collision Model

- Furthermore, molecules must collide with the

correct orientation and with enough energy to

cause bond breakage and formation.

Activation Energy

- In other words, there is a minimum amount of

energy required for reaction the activation

energy, Ea. - Just as a ball cannot get over a hill if it does

not roll up the hill with enough energy, a

reaction cannot occur unless the molecules

possess sufficient energy to get over the

activation energy barrier.

Reaction Coordinate Diagrams

- It is helpful to visualize energy changes

throughout a process on a reaction coordinate

diagram like this one for the rearrangement of

methyl isonitrile.

Reaction Coordinate Diagrams

- It shows the energy of the reactants and products

(and, therefore, ?E). - The high point on the diagram is the transition

state.

- The species present at the transition state is

called the activated complex. - The energy gap between the reactants and the

activated complex is the activation energy

barrier.

MaxwellBoltzmann Distributions

- Temperature is defined as a measure of the

average kinetic energy of the molecules in a

sample.

- At any temperature there is a wide distribution

of kinetic energies.

MaxwellBoltzmann Distributions

- As the temperature increases, the curve flattens

and broadens. - Thus at higher temperatures, a larger population

of molecules has higher energy.

MaxwellBoltzmann Distributions

- If the dotted line represents the activation

energy, as the temperature increases, so does the

fraction of molecules that can overcome the

activation energy barrier.

- As a result, the reaction rate increases.

MaxwellBoltzmann Distributions

- This fraction of molecules can be found through

the expression - where R is the gas constant and T is the Kelvin

temperature.

f e-Ea/RT

SAMPLE EXERCISE 14.10 Relating Energy Profiles to

Activation Energies and Speeds of Reaction

Consider a series of reactions having the

following energy profiles

Assuming that all three reactions have nearly the

same frequency factors, rank the reactions from

slowest to fastest.

Solution The lower the activation energy, the

faster the reaction. The value of ?? does not

affect the rate. Hence the order is (2) lt (3) lt

(1).

PRACTICE EXERCISE Imagine that these reactions

are reversed. Rank these reverse reactions from

slowest to fastest.

Answer (2) lt (1) lt (3) because Ea values are

40, 25, and 15 kJ/mol, respectively

Arrhenius Equation

- Svante Arrhenius developed a mathematical

relationship between k and Ea - k A e-Ea/RT
- where A is the frequency factor, a number that

represents the likelihood that collisions would

occur with the proper orientation for reaction.

Arrhenius Equation

- Taking the natural logarithm of both sides, the

equation becomes - ln k -Ea ( ) ln A

y mx b

Therefore, if k is determined experimentally at

several temperatures, Ea can be calculated from

the slope of a plot of ln k vs. 1/T.

SAMPLE EXERCISE 14.11 Determining the Energy of

Activation

The following table shows the rate constants for

the rearrangement of methyl isonitrile at various

temperatures (these are the data in Figure 14.12)

(a) From these data, calculate the activation

energy for the reaction. (b) What is the value of

the rate constant at 430.0 K?

Solution Analyze We are given rate constants, k,

measured at several temperatures and asked to

determine the activation energy, Ea, and the

rate constant, k, at a particular

temperature. Plan We can obtain Ea from the

slope of a graph of ln k versus 1/T. Once we know

Ea, we can use Equation 4.21 together with the

given rate data to calculate the rate constant at

430.0 K.

SAMPLE EXERCISE 14.11 continued

Solve (a) We must first convert the temperatures

from degrees Celsius to kelvins. We then take the

inverse of each temperature, 1/T, and the natural

log of each rate constant, ln k. This gives us

the table shown at the right

A graph of ln k versus 1/T results in a straight

line, as shown in Figure 14.17.

We report the activation energy to only two

significant figures because we are limited by the

precision with which we can read the graph in

Figure 14.17.

SAMPLE EXERCISE 14.11 continued

Thus,

Note that the units of k1 are the same as those

of k2.

PRACTICE EXERCISE Using the data in Sample

Exercise 14.11,(slide 68) calculate the rate

constant for the rearrangement of methyl

isonitrile at 280C.

Reaction Mechanisms

- The sequence of events that describes the actual

process by which reactants become products is

called the reaction mechanism.

Reaction Mechanisms

- Reactions may occur all at once or through

several discrete steps. - Each of these processes is known as an elementary

reaction or elementary process.

Reaction Mechanisms

- The molecularity of a process tells how many

molecules are involved in the process.

Multistep Mechanisms

- In a multistep process, one of the steps will be

slower than all others. - The overall reaction cannot occur faster than

this slowest, rate-determining step.

Slow Initial Step

NO2 (g) CO (g) ??? NO (g) CO2 (g)

- The rate law for this reaction is found

experimentally to be - Rate k NO22
- CO is necessary for this reaction to occur, but

the rate of the reaction does not depend on its

concentration. - This suggests the reaction occurs in two steps.

Slow Initial Step

- A proposed mechanism for this reaction is
- Step 1 NO2 NO2 ??? NO3 NO (slow)
- Step 2 NO3 CO ??? NO2 CO2 (fast)
- The NO3 intermediate is consumed in the second

step. - As CO is not involved in the slow,

rate-determining step, it does not appear in the

rate law.

Fast Initial Step

2 NO (g) Br2 (g) ??? 2 NOBr (g)

- The rate law for this reaction is found to be
- Rate k NO2 Br2
- Because termolecular processes are rare, this

rate law suggests a two-step mechanism.

Fast Initial Step

- A proposed mechanism is

Step 2 NOBr2 NO ??? 2 NOBr (slow)

Step 1 includes the forward and reverse reactions.

Fast Initial Step

- The rate of the overall reaction depends upon the

rate of the slow step. - The rate law for that step would be
- Rate k2 NOBr2 NO
- But how can we find NOBr2?

Fast Initial Step

- NOBr2 can react two ways
- With NO to form NOBr
- By decomposition to reform NO and Br2
- The reactants and products of the first step are

in equilibrium with each other. - Therefore,
- Ratef Rater

Fast Initial Step

- Because Ratef Rater ,
- k1 NO Br2 k-1 NOBr2
- Solving for NOBr2 gives us

Fast Initial Step

- Substituting this expression for NOBr2 in the

rate law for the rate-determining step gives

k NO2 Br2

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(a) Is the proposed mechanism consistent with the

equation for the overall reaction? (b) What is

the molecularity of each step of the mechanism?

(c) Identify the intermediate(s).

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(a) Write the equation for the overall reaction.

(b) Write the rate law for the overall reaction.

The experimental rate law is rate kO3NO2.

What can you say about the relative rates of the

two steps of the mechanism?

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What is the expression relating the concentration

of Br(g) to that of Br2(g)?

Catalysts

- Catalysts increase the rate of a reaction by

decreasing the activation energy of the reaction. - Catalysts change the mechanism by which the

process occurs.

Catalysts

- One way a catalyst can speed up a reaction is by

holding the reactants together and helping bonds

to break.

Enzymes

- Enzymes are catalysts in biological systems.
- The substrate fits into the active site of the

enzyme much like a key fits into a lock.

The decomposition reaction is determined to be

first order. A graph of the partial pressure of

HCOOH versus time for decomposition at 838 K is

shown as the red curve in Figure 14.28. When a

small amount of solid ZnO is added to the

reaction chamber, the partial pressure of acid

versus time varies as shown by the blue curve in

Figure 14.28.

- (a) Estimate the half-life and first-order rate

constant for formic acid decomposition. - (b) What can you conclude from the effect of

added ZnO on the decomposition of formic acid?

SAMPLE INTEGRATIVE EXERCISE continued

Solution (a) The initial pressure of HCOOH is

3.00 ? 102 torr. On the graph we move to the

level at which the partial pressure of HCOOH is

150 torr, half the initial value. This

corresponds to a time of about 6.60 x 102s, which

is therefore the half-life. The first-order rate

constant is given by Equation 14.15 k

0.693/t1/2 0.693/660 s 1.05 ? 103 s1.

(b) The reaction proceeds much more rapidly in

the presence of solid ZnO, so the surface of the

oxide must be acting as a catalyst for the

decomposition of the acid. This is an example of

heterogeneous catalysis.

(c) If we had graphed the concentration of

formic acid in units of moles per liter, we would

still have determined that the half-life for

decomposition is 660 seconds, and we would have

computed the same value for k. Because the units

for k are s1, the value for k is independent of

the units used for concentration.

SAMPLE INTEGRATIVE EXERCISE continued