Transcript Document

Chemical Kinetics
Chapter 14
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Chemical Kinetics
____________________ – does a reaction take place?
____________________– how fast does a reaction proceed?
____________ ____________ is the change in the
concentration of a reactant or a product with time (M/s).
A
B
D[A]
___ = Dt
D[A] = change in concentration of A over
time period Dt
D[B]
Dt
D[B] = change in concentration of B over
time period Dt
___ =
Because [A] decreases with time, D[A] is __________ .
14.1
A
B
time
D[A]
rate = Dt
D[B]
rate =
Dt
14.1
2Br- (aq) + 2H+ (aq) + CO2 (g)
Br2 (aq) + HCOOH (aq)
time
393 nm
light
Detector
393 nm
Br2 (aq)
D[Br2] a DAbsorption
14.1
Br2 (aq) + HCOOH (aq)
2Br- (aq) + 2H+ (aq) + CO2 (g)
slope of
tangent
slope of
tangent
slope of
tangent
[Br2]final – [Br2]initial
D[Br2]
____________ _____ = =Dt
tfinal - tinitial
____________ _____ = rate for specific instance in time
14.1
Rates of the Reaction Between Molecular Bromine and Formic Acid at 25ºC
rate a [Br2]
rate = k [Br2]
k=
rate
= ____ _________
[Br2]
= 3.50 x 10-3 s-1
14.1
2H2O2 (aq)
2H2O (l) + O2 (g)
PV = nRT
n
P=
RT = [O2]RT
V
1
[O2] =
P
RT
D[O2]
1 DP
rate =
=
RT Dt
Dt
measure DP over time
14.1
2H2O2 (aq)
2H2O (l) + O2 (g)
14.1
Reaction Rates and Stoichiometry
2A
B
Two moles of A disappear for each mole of B that is formed.
1 D[A]
rate = 2 Dt
aA + bB
D[B]
rate =
Dt
cC + dD
1 D[A]
1 D[B]
1 D[C]
1 D[D]
rate = ==
=
a Dt
b Dt
c Dt
d Dt
14.1
Write the rate expression for the following reaction:
CH4 (g) + 2O2 (g)
D[CH4]
rate = =
Dt
CO2 (g) + 2H2O (g)
D[O2]
=
Dt
D[CO2]
=
Dt
D[H2O]
Dt
14.1
The ______ ______ expresses the relationship of the rate of
a reaction to the rate constant and the concentrations of the
reactants raised to some powers.
aA + bB
cC + dD
Rate =
reaction is _______ order in A
reaction is _______ order in B
reaction is _______________ order overall
14.2
F2 (g) + 2ClO2 (g)
2FClO2 (g)
Rate Data for the Reaction between F2 and ClO2
rate =
Double [F2] with [ClO2] constant
Rate _____________
x=1
Quadruple [ClO2] with [F2] constant
Rate ______________
y=1
rate =
14.2
Rate Laws
•
Rate laws are always determined _______________.
•
Reaction order is always defined in terms of ________
(not _____________) concentrations.
•
The order of a reactant is not related to the
___________________ ____________________
of the reactants in the balanced chemical equation.
F2 (g) + 2ClO2 (g)
2FClO2 (g)
rate = k [F2][ClO2] 1
14.2
Determine the rate law and calculate the rate constant for
the following reaction from the following data:
S2O82- (aq) + 3I- (aq)
2SO42- (aq) + I3- (aq)
Experiment
[S2O82-]
[I-]
Initial Rate
(M/s)
1
0.08
0.034
2.2 x 10-4
2
0.08
0.017
1.1 x 10-4
3
0.16
0.017
2.2 x 10-4
rate =
y=
x=
rate =
Double [I-], rate ___________ (experiment 1 & 2)
Double [S2O82-], rate ___________ (experiment 2 & 3)
rate
k=
=
2[S2O8 ][I ] (
M)(
M/s
M)
= ______ /M•s
14.2
-Order Reactions
A
k=
product
D[A]
rate = Dt
rate
M/s
=
= 1/s or s-1
M
[A]
[A] = [A]0exp(-kt)
rate = k [A]
D[A]
= k [A]
Dt
[A] is the concentration of A at any time t
[A]0 is the concentration of A at time t=0
ln[A] = ln[A]0 - kt
14.3
The reaction 2A
B is first order in A with a rate
constant of 2.8 x 10-2 s-1 at 800C. How long will it
take for A to decrease from 0.88 M to 0.14 M ?
[A]0 = 0.88 M
ln[A] = ln[A]0 - kt
[A] = 0.14 M
kt = ln[A]0 – ln[A]
ln[A]0 – ln[A]
=
t=
k
ln
[A]0
[A]
k
ln
=
0.88 M
0.14 M
2.8 x
10-2
s-1
= 66 s
14.3
-Order Reactions
The half-life, t½, is the time required for the concentration of a
reactant to decrease to half of its initial concentration.
t½ = t when [A] = [A]0/2
ln
t½ =
[A]0
[A]0/2
k
ln2
0.693
=
=
k
k
What is the half-life of N2O5 if it decomposes with a rate
constant of 5.7 x 10-4 s-1?
0.693
t½ = ln2 =
= 1216 s = 20 minutes
-4
-1
k
5.7 x 10 s
How do you know decomposition is first order?
units of k (s-1)
14.3
First-order reaction
A
product
# of
half-lives
1
[A] = [A]0/n
2
4
3
8
4
16
2
14.3
-Order Reactions
A
product
D[A]
rate = Dt
rate
M/s
=
= 1/M•s
k=
2
2
M
[A]
1
1
=
+ kt
[A]
[A]0
rate = k [A]2
D[A]
= k [A]2
Dt
[A] is the concentration of A at any time t
[A]0 is the concentration of A at time t = 0
t½ = t when [A] = [A]0/2
1
t½ =
k[A]0
14.3
-Order Reactions
A
k=
product
rate
= M/s
0
[A]
[A] = [A]0 - kt
D[A]
rate = Dt
D[A]
=k
Dt
rate = k [A]0 = k
Reminder: what is the
value of any number
raised to the zero power?
x0= ??
[A] is the concentration of A at any time t
[A]0 is the concentration of A at time t=0
t½ = t when [A] = [A]0/2
[A]0
t½ =
2k
14.3
Summary of the Kinetics of Zero-Order,
First-Order and Second-Order Reactions
Order
0
Rate Law
rate = k
1
rate = k [A]
2
[A]2
rate = k
Concentration-Time
Equation
[A] = [A]0 - kt
ln[A] = ln[A]0 - kt
1
1
=
+ kt
[A]
[A]0
Half-Life
t½ =
[A]0
2k
t½ = ln2
k
1
t½ =
k[A]0
14.3
Implications of the Collision Theory of Chemical Kinetics
Rate a
The ____________ ____________ (Ea ) is the minimum
amount of energy required to initiate a chemical reaction.
p.448
A+B
__________ Reaction
C+D
__________ Reaction
The __________ __________ (___)is the minimum amount
of energy required to initiate a chemical reaction.
14.4
The Rate Constant Depends on Temperature
What happens to the rate of a chemical
reaction as the temperature increases?
Why does this happen?
Explain this observation in terms of the
Collision Theory of Kinetics.
We express this dependence of rate on
temperature in the ___________ equation.
14.4
The Rate Constant Depends on Temperature
Arrhenius equation
k=
Ea = the activation energy (J/mol)
R = the gas constant (8.314 J/K•mol)
T = the absolute temperature
A = frequency factor/collision frequency
e = the base of the natural log scale
Why do we write it in this form?
Ea 1
lnk = + lnA
R T
14.4
Why do we rewrite the equation in this form?
Ea 1
lnk = + lnA
R T
y =
m x +
b
slope
Now a plot of lnk versus 1/T is a __________ line,
Ea/R is the _________, and the
_____________________is lnA.
p.450
Ex.14.6, p.451
Ea 1
lnk = + lnA
R T
14.4
A balanced chemical equation does not tell us
how the reaction actually takes place.
Usually it just represents the sum of a series
of simpler elementary steps or elementary
reactions.
We try to deduce these steps and propose a
_______________ _______________ .
14.5
Reaction Mechanisms
The overall progress of a chemical reaction can be
represented at the molecular level by a series of simple
elementary steps or elementary reactions.
The sequence of elementary steps that leads to
product formation is the __________ _____________.
2NO (g) + O2 (g)
2NO2 (g)
N2O2 is detected during the reaction!
Elementary step:
NO + NO
N 2O 2
+ Elementary step:
N2O2 + O2
2NO2
Overall reaction:
2NO + O2
2NO2
14.5
Species that appear in a reaction mechanism but not in
the overall balanced equation are called ___________.
An ___________ is always formed in an early elementary
step and consumed in a later elementary step.
Elementary step:
NO + NO
N 2O 2
+ Elementary step:
N2O2 + O2
2NO2
Overall reaction:
2NO + O2
2NO2
The _______________ of a reaction is
the number of molecules reacting in an elementary step.
•
____________ reaction – elementary step with 1 molecule
•
___________ reaction – elementary step with 2 molecules
•
___________ reaction – elementary step with 3 molecules
14.5
Rate Laws and Elementary Steps
Unimolecular reaction
A
products
rate = k [A]
Bimolecular reaction
A+B
products
rate = k [A][B]
Bimolecular reaction
A+A
products
rate = k [A]2
Writing plausible reaction mechanisms:
•
The sum of the elementary steps must give the overall
balanced equation for the reaction.
•
The rate-determining step should predict the same rate
law that is determined experimentally.
The rate-determining step is the __________ step in
the sequence of steps leading to product formation
14.5
The experimental rate law for the reaction between
NO2 and CO to produce NO and CO2 is rate = k[NO2]2.
The reaction is believed to occur via two steps:
Step 1:
NO2 + NO2
NO + NO3
Step 2:
NO3 + CO
NO2 + CO2
What is the equation for the overall reaction?
What is the intermediate?
What can you say about the relative rates of steps 1 and 2?
rate =
is the rate law for step 1 so
step 1 must be ___________ than step 2
14.5
A _____________ is a substance that increases the rate of
a chemical reaction without itself being consumed.
k = A • exp( -Ea/RT )
Ea
uncatalyzed
k
catalyzed
ratecatalyzed > rateuncatalyzed
Ea' < Ea
14.6
In ____________________________________ catalysis,
the reactants and the catalysts are in different phases.
•
Haber synthesis of ammonia
•
Ostwald process for the production of nitric acid
•
Catalytic converters
In ________________________ catalysis, the reactants and
the catalysts are dispersed in a single phase, usually liquid.
•
Acid catalyses
•
Base catalyses
14.6
Haber Process
N2 (g) + 3H2 (g)
Fe/Al2O3/K2O
catalyst
2NH3 (g)
14.6
Ostwald Process
4NH3 (g) + 5O2 (g)
Pt catalyst
2NO (g) + O2 (g)
2NO2 (g) + H2O (l)
4NO (g) + 6H2O (g)
2NO2 (g)
HNO2 (aq) + HNO3 (aq)
Pt-Rh catalysts used
in Ostwald process
Hot Pt wire
over NH3 solution
14.6
Catalytic Converters
CO + Unburned Hydrocarbons + O2
NO + NO2
catalytic
converter
catalytic
converter
CO2 + H2O
N2 + O2
14.6
______________ Catalysis
14.6
uncatalyzed
enzyme
catalyzed
14.6