Transcript CHEMICAL KINETICS CHAPTER 13

CHEMICAL KINETICS
CHAPTER 13
I. Introduction
A. Definition of Chemical Kinetics
“The study of the speed or rate of reactions and the
nanoscale pathways or processes by which reactants
are transformed into products.
B. Examples of Reactions and Rates
Rusting of Iron
Combustion Reaction
C. Significance of Studying Kinetics
D. Factors Affecting Reaction Rate
1. Concentration of Reactants
2. Temperature
3. Presence of a Catalyst
4. Surface Area of a Solid Reactant or
Catalyst
5. Properties of Reactants and Products
II. Understanding Reaction Rates
A. Kinetic Molecular Theory
Matter composed of particles in constant motion.
Increase in temperature increases particle’s kinetic
energy.
B. Collision Theory
For a reaction to occur , reactant molecules must
collide with the proper orientation and with an
energy greater than some minimum value.
Activation Energy (Ea) – minimum energy
required for reaction to occur.
Importance of Orientation
One hydrogen atom
can approach another
from any direction …
Effective collision; the
I atom can bond to the
C atom to form CH3I
… and reaction will still occur; the
spherical symmetry of the atoms means
that orientation does not matter.
Ineffective collision;
orientation is important
in this reaction.
Distribution of Kinetic Energies
At higher temperature
(red), more molecules
have the necessary
activation energy.
C. Transition State Theory
1. Note Reaction Profile:
CO(g) + NO2(g)

CO2(g) + NO(g)
2. Note Ea for either forward or reverse reaction.
“At a given temperature, the higher the energy
barrier, the slower the reaction.”
3. Transition State or “Activated Complex”
“Transition structure (between reactants and
products) which is always found on top of the
energy hill (energy of activation).
4. Is reaction, exothermic or endothermic as
written from left to right?
Reconsideration of Factors Affecting
Reaction Rate!!
1. Concentration
2. Temperature
3. Catalyst
III. Rates of Reactions
A. Definition
1. Reaction rate: expresses how much
product is appearing or how much reactant in
disappearing per unit time.
2. Units for reaction rates:
(Examples)
Ms-1 or M/s or mol L-1s-1
For Reaction: A → P
Rate of Disappearance of “A” = - ΔA / Δt

A
A

tt
Rate of Formation of “P” = + ΔP / Δt
B. Example Average Rate Determination
For Rxn of Cisplatin and Water:
H2O + Pt(NH3)2Cl2  Pt(NH3)2Cl(H2O) + ClCisplatin
Time (min)
0
200
400
[Cisplatin] (mol/L)
0.01000
0.00747
0.00558
What is average rate of disappearance of cisplatin
(in mol L-1 min-1) for the first 200 min?
What is average rate of disappearance of cisplatin
(in mol L-1 min-1) for the next 200 min?
C. Instantaneous Rate Determination
Significance of Measuring Instantaneous Rates!!
D. Reaction Rates and Stoichiometry
Given the reaction:
2 N2O5(g)  4 NO2(g) + O2 (g)
If the rate of NO2 formation is 0.060 mol L-1 s-1:
1. What is the rate of disappearance of N2O5?
2. What is the rate of formation of O2?
IV. Concentration and Rxn Rate
A. Rate Law Equation:
1. An equation that relates the rate of a
reaction to the concentrations of reactants
(and catalyst) raised to various powers.
2. Must be experimentally determined!!
3. For reaction:
A + B  C
Rate is proportional to reactant concentrations
Rate = k[A]m [B]n
k = rate constant
(exponents “m” and “n” must be
experimentally determined).
4. For reaction:
2 NO2(g) + F2(g)  2 NO2F(g)
(experimentally determined Rate Law is:)
Rate = k [NO2]1[F2]1 = k [NO2][F2]
Exponents not necessarily same as rxn coefficients!!
4. For hypothetical reaction:
2 A(g) + B2(g)  2 AB(g)
(experimentally determined Rate Law is:)
Rate = k [A]2
Not all reactants necessarily show up in the
rate law equation!!
B. Reaction Order
For the general equation:
aA + bB  pP
The rate equation is:
Rate = k [A]m [B]n
m and n are experimentally determined
and are usually integers (0, 1, 2, 3, …).
They may be fractions.
This reaction is said to m th order with
respect to A and n th order with respect
to B.
The overall reaction order is the sum of
the individual orders, or
Overall Reaction Order = m + n
Example:
For the following reaction:
2 NO(g) + Cl2(g)  2 NOCl(g)
The observed rate law is:
Rate = k [NO]2 [Cl2]
What is the reaction order with respect to NO?
What is the reaction order with respect to Cl2?
What is the overall reaction order?
How would the rate of the reaction be affected by
doubling the concentration of both NO and Cl2?
C. Determination of Rate Law Exponents
Done experimentally by measuring initial
rates for several different known
concentrations of reactants.
Consider the reaction:
2 NO(g) + 2 H2(g)  N2(g) + 2 H2O(g)
Given the information on the next slide:
1. Determine the rate law.
2. What is the order of the reaction?
3. What is the value of the rate constant? “units”
Experiment
Initial Concentration (M) Rate
[NO]
[H2]
mol / L.s
1
0.100
0.100
1.23 x 10-3
2
0.100
0.200
2.45 x 10-3
3
0.200
0.100
4.93 x 10-3
4
0.300
0.100
1.11 x 10-2
1. Determine the rate law.
a. Determine general form of rate law.
rate = k [NO]m[H2]n
b. Determine exponents.
For each reactant, compare two
experiments or trials where its
concentration is changing and all other
reactant concentrations are held
constant.
Methods For Determining Rate Law Exponents
Method 1 – solve analytically
Substitute data into rate law and compare
Divide equation with larger rate by eq. with
smaller rate. Cancel terms and solve.
or simplifying
Method 2 - solve by inspection
How does changing conc. affect rate?
2. Determine order of the reaction.
3. Determine rate constant (include units).
Use rate law and either set of data.
4. What is the rate of the reaction when
[NO] = [H2] = 0.200 M ?
V. Integrated Rate Law
Equations derived from rate law (by using
calculus) which are convenient for solving
concentration versus time problems.
For:
1. First Order Rxns – only one covered
2. Second Order Rxns
3. Zero Order Rxns
A. Integrated First Order Rate Law
For reaction:
aA  Product
rate = - Δ[A] / Δt = k[A]
and using calculus:
[ A ]t
ln
  kt
[ A ]o
where [A]o is the concentration of A at
time zero (t = 0) and [A]t is the
concentration at time t.
*** A and A0 may be replaced by quantities
that are proportional to concentration !!!!
B. Problem
The sugar, sucrose, will undergo the following
(first order) hydrolysis reaction
C12H22O11 + H2O  C6H12O6 + C6H12O6
sucrose
glucose
fructose
with a rate constant of 6.2 x 10-5 s-1 at 35oC.
A sample of 0.20 mol of sucrose was initially
dissolved in a total volume of 500 mL.
1. What is the sucrose conc. after 2 hours?
2. What will be the glucose concentration
after the 2 hours have elapsed?
3. How many minutes will it take for the
sucrose concentration to drop to 0.30M?
C. Half-Life and First Order Reactions
1. Definition (t1/2) – the time required for the
concentration of a reactant to fall to one half
its initial value.
2. Significance:
Useful in describing radioactive decay rates
Useful in describing rates of 1st order reactions
3. Equation derived from Integrated Rate Law
t1 
2
.6 9 3
k
4. Problem - Radioactive Iodine-131 has t1/2 of 8
days. If you had a sample of 10,000 I-131
atoms initially, how many I-131 atoms would
remain after 32 days?
5. Problem – The first order hydrolysis reaction of
sucrose
C12H22O11 + H2O  C6H12O6 + C6H12O6
sucrose
glucose
fructose
has a rate constant of 6.17 x 10-4 s-1 under
experimental conditions.
a. What is the half-life for the hydrolysis of
sucrose?
b. How many minutes are required for 75% of the
initial sucrose to react?
VI. Temperature and Rxn Rate
A. Nanoscale Explanation as to why
increasing temperature increase reaction
rate.
B. Mathematical Relationship
Arrhenius Equation
k  Ae
 Ea / RT
(not responsible for problem solving)
VII. Reaction Mechanisms
A. Introduction
Reaction mechanism is a series of elementary
reactions or simple steps whose overall effect is
given by the net chemical reaction (equation).
1. It is a theory of how the reaction occurs
which is based on experimental data.
2. Cannot be absolutely proven.
3. Steps must be elementary reactions.
B. Elementary Reaction
1. Definition- the simplest step in what is
often a multi-step mechanism for an
observed chemical reaction.
a. The equation for an elementary reaction
shows exactly which molecules, atoms,
or ions take part in the elementary
reaction.
b. For an elementary reaction, the rate
law is directly determined from the
elementary reaction.
2. Elementary Reactions in Mechanisms
Types
a. Unimolecular Reaction – structure of a
single particle (atom, molecule, or ion)
rearranges to produce a different
particle or particles.
b. Bimolecular Reaction - two particles
(atoms, ions, or molecules) collide and
rearrange into products.
c. Termolecular Reaction – (less likely)
3. Problems
Identify the type of elementary reaction and
give the rate law for the following elementary
reactions:
a.
Cl + Cl  Cl2
b.
N2O5  NO2 + NO3
C. Properties of Valid Mechanisms
1. Must consist of only unimolecular,
bimolecular, or termolecular elementary
reactions. (True for any mechanisms given to
you.)
***2. Sum of the elementary reactions should be
equal to the overall reaction equation.
***3. Should predict the experimentally observed
rate law.
The overall rate of the reaction is
dependent on the slowest step in the
mechanism - the rate-limiting step.
D. Example Mechanism Problems
1. For the overall reaction:
2 NO2Cl  2 NO2 + Cl2
the following mechanism is proposed.
NO2Cl  NO2 + Cl (slow)
NO2Cl + Cl  NO2 + Cl2 (fast)
a. Does the sum of the elementary processes
equal the overall reaction?
b. What is the rate law for the overall rxn?
c. Identify any reaction intermediates.
2. For the overall reaction:
(CH3)3CCl + OH-  (CH3)3COH + Clthere are two proposed mechanisms.
1) Concerted mechanism
(CH3)3CCl + OH-  (CH3)3COH + Cl2) Two Step Mechanism
(CH3)3CCl  (CH3)3C + + Cl (CH3)3C+ + OH -  (CH3)3COH
(slow)
(fast)
From kinetic data, the correct rate law for the
overall reaction is:
rate = k[(CH3)3CCl]
Questions:
1. Which is the correct mechanism? Why?
2. What is the order of the overall reaction?
3. Identify reaction intermediates in each
proposed mechanism.
VIII. Catalysts
A. Definition
Catalyst – a species that increases the rate of
an overall reaction but is not consumed in the
reaction.
Not shown in overall reaction.
Will show up in rate law for catalyzed
reaction.
B. How Do They Increase Reaction Rate?
1. Catalysts alter / participate in rxn mechanism.
2. Lower activation energy. Speeds reaction.
3. See Fig 13.17 (next slide)
4. Catalyst changes kinetics, but not
thermodynamics of reaction.
Increases speed of reaction.
Does not change net energy of reaction, type
of product produced, or direction of reaction.
C. Homogeneous vs. Heterogeneous Catalysts
Heterogeneous Catalyst- catalyst in different
phase from reacting substance.
(hydrogenation of vegetable oils – next slide)
(catalytic converters in autos)
Homogeneous Catalyst – catalyst in same phase
as reacting substance.
(enzymes)
Heterogeneous Catalysis
Hydrogen is adsorbed
onto the surface of a
nickel catalyst. A C=C
approaches …
… and is adsorbed.
Hydrogen atoms
attach to the carbon
atoms, and the
molecule is desorbed.
D. Enzymes (Homogeneous Catalysts)
1. Protein that catalyzes reaction.
2. Most efficient catalysts known to man.
3. Specifically binds to reactants (substrates),
holding them in correct position for reaction
to occur.
4. Lower activation energy by stabilizing
transition state or altering mechanism.
5. Examples: Lysozyme Next slide