Chapter 12. Saturated Hydrocarbons
Sections
CHEM 102, Fall 2009, LA TECH
1-1
Chemistry 102(01) Fall 2009
Instructor: Dr. Upali Siriwardane
e-mail: upali@chem.latech.edu
Office: CTH 311 Phone 257-4941
Office Hours: M,W, 8:00-9:00 & 11:00-12:00 a.m Tu,Th,F 9:00 10:00 a.m.
Test Dates: March 25, April 26, and May 18; Comprehensive Fina
Exam: 9:30-10:45 am, CTH 328.
September 24, 2009 (Test 1): Chapter 13
October 22, 2009 (Test 2): Chapters 14 & 15
November 17, 2009 (Test 3): Chapters 16, 17 & 18
Comprehensive Final Exam: November 19, 2009 :
Chapters 13, 14, 15, 16, 17 and 18
CHEM 102, Fall 2009, LA TECH
1-2
Chapte
13.1
13.2
13.3
13.4
13.5
Reaction Rate
Effect of Concentration on Reaction Rate
Rate Law and Order of Reaction
A Nanoscale View: Elementary Reactions
Temperature and Reaction Rate: The Arrhenius
Equation
13.6 Rate Laws for Elementary Reactions
13.7 Reaction Mechanisms
13.8 Catalysts and Reaction Rate
13.9 Enzymes: Biological Catalysts
13-10Catalysis in Industry
CHEM 102, Fall 2009, LA TECH
1-3
How do you measure rates?
Rates are related to the time it required to decay
The rate reaction = change in concentration of re
Average rate
ction = – D[reactant]/Dt
Instantaneous rate
ction = – d[reactant]/dt
CHEM 102, Fall 2009, LA TECH
1-4
Rate of Appearance & Disappearance
NO2 (g) + O2 (g)
based on reactants
D[N2O5]/ D t
ased on products
[NO2]/ D t
D[O2]/ D t
s of Appearance.
= - 4/2 D[N2O5]/ D t
- 1/2 D[N2O5]/ D t
CHEM 102, Fall 2009, LA TECH
1-5
Measuring Rate
CHEM 102, Fall 2009, LA TECH
1-6
Reaction of cis-platin with Water
CHEM 102, Fall 2009, LA TECH
1-7
Disappearance of Color
CHEM 102, Fall 2009, LA TECH
1-8
An example reaction where gas is p
Gas
buret
Constant temperature bath
CHEM 102, Fall 2009, LA TECH
1-9
Time vs. volume of gas
Time (s)
0
300
600
900
1200
1800
2400
3000
4200
5400
6600
7800
CHEM 102, Fall 2009, LA TECH
Volume STP O2, mL
0
1.15
2.18
3.11
3.95
5.36
6.50
7.42
8.75
9.62
10.17
10.53
Here are the results
for our experiment.
1-10
2
5
CHEM 102, Fall 2009, LA TECH
2
2
1-11
Graph of 2 N2O5(g) ---> 4 NO2 (g) + O2 (g)
CHEM 102, Fall 2009, LA TECH
1-12
Graph
CHEM 102, Fall 2009, LA TECH
1-13
Factors that affect rates of chemical rea
a) Temperature
b) Concentration
c) Catalysts
d) Particle size of solid reactants
CHEM 102, Fall 2009, LA TECH
1-14
Effect of Particle Size on Rate
CHEM 102, Fall 2009, LA TECH
1-15
Chemical Kinetics Definitions and Co
a) rate law
b) rate constant
c) order
d) differential rate law
c) integral rate law
CHEM 102, Fall 2009, LA TECH
1-16
Rate Law
Every chemical reaction has a Rate La
The rate law is an expression that rela
The power of a concentration is called
CHEM 102, Fall 2009, LA TECH
1-17
Rate Law
E.g.
A + B -----> C
rate a [A]l[B]m
rate = k [A]l[B]m; k = rate constant
[A] = concentration of A
[B] = concentration of B
l = order with respect to A
m = order with respect to B
l & m have nothing to do with stoichiometric coefficien
CHEM 102, Fall 2009, LA TECH
1-18
Rate Constant
E.g.
A + B -----> C
rate a [A]l[B]m
rate = k [A]l[B]m;
k = rate constant
proportionality constant of the rate law
Larger the k faster the reaction
It is related inversely to t½
CHEM 102, Fall 2009, LA TECH
1-19
Decomposition Reaction
CHEM 102, Fall 2009, LA TECH
1-20
Rate Law
E.g.
2 N2O5(g) -----> 4 NO2 (g) + O2 (g)
rate a [N2O5]1
rate = k [N2O5] 1;k = rate constant
[N2O5] = concentration of N2O5
1 = order with respect to N2O5
Rate and the order are obtained by experiments
CHEM 102, Fall 2009, LA TECH
1-21
Order
The power of the concentrations is the order with resp
E.g.
A + B -----> C
If rate law: rate = k [A]1[B]2
The order of the reaction with respect to A is one (1).
The order of the reaction with respect to B is two (2).
Overall order of a chemical reaction is equal to the
sum of all orders (3).
CHEM 102, Fall 2009, LA TECH
1-22
Finding rate laws
Method of initial rates
The order for each reactant is found by:
•
•
•
Changing the initial concentration of that reactant.
Holding all other initial concentrations and conditions constant.
Measuring the initial rates of reaction
The change in rate is used to determine the order for that specific reactant. The
process is repeated for each reactant.
CHEM 102, Fall 2009, LA TECH
1-23
Initial rate
CHEM 102, Fall 2009, LA TECH
1-24
How do you find order?
A + B -----> C
rate = k [A]l[B]m;
Hold concentration of other reactants constant
If [A] doubled, rate doubled
-1st order, [2A]1 = 2 1 x [A]1 , 2 1 = 2
b) If [A] doubled, rate quadrupled
-2nd order, [2A]2 = 2 2 x [A]2 , 2 2 = 4
c) If [A] doubled, rate increased 8 times -3rd order, [2A
CHEM 102, Fall 2009, LA TECH
1-25
Rate data
CHEM 102, Fall 2009, LA TECH
1-26
Determining order
CHEM 102, Fall 2009, LA TECH
1-27
Determining K, Rate Constant
CHEM 102, Fall 2009, LA TECH
1-28
Overall order
CHEM 102, Fall 2009, LA TECH
1-29
Units of the Rate Constant (k)
1
first order: k = ─── =
s-1
s
L
second order k =
───
mol s
L2
third order k =
CHEM 102, Fall 2009, LA TECH
───
mol2 s
1-30
First order reactions
CHEM 102, Fall 2009, LA TECH
1-31
Differential and Integral Rate Law
Rate Law
Differential Rate Law
rate = k [A]0 -D [A]/Dt =k ; ([A]0=1)
rate = k [A]1
Integral Rate
[A]f-[A]i = -kt
-D [A]/Dt = k [A]
ln [A]o/[A]t = kt
rate = k [A]2 = -D [A]/Dt = k [A]2
1/ [A]f = kt + 1/[A]i
CHEM 102, Fall 2009, LA TECH
1-32
Integrated Rate Laws
CHEM 102, Fall 2009, LA TECH
1-33
Graphical method
Rate
Integrated Rate
Law
Graph
X vs.
time
Slope
[A]t
-k
Order
Law
0
rate = k
[A]t = -kt + [A]0
1
rate = k[A]
ln[A]t = -kt + ln[A]0 ln[A]t
2
rate=k[A]2
CHEM 102, Fall 2009, LA TECH
1
[A]t
= kt +
1
[A]0
1
[A]t
-k
k
1-34
Graphical Ways to get Order
CHEM 102, Fall 2009, LA TECH
1-35
First-order, Second-order,
and Zeroth-order Plots
CHEM 102, Fall 2009, LA TECH
1-36
Finding rate laws
0.2
100
0 order plot
2nd order plot
80
1/[N2O5]
[N2O5]
0.15
0.1
0.05
60
40
20
0
0
0
2000
4000
6000
8000
Time (s)
0
1000
2000
3000
4000
5000
6000
7000
8000
Time (s)
Time (s)
-1.5
0
2000
4000
-2
8000
As you can see from these
plots of the N2O5 data,
only a first order plot
results in a straight line.
-2.5
ln[N2O5]
6000
-3
-3.5
-4
-4.5
CHEM 102, Fall 2009, LA TECH
1st order plot
1-37
Comparing graphs
This plot of ln[cis-platin] vs.
time produces a straight line,
suggesting that the reaction
is first-order.
CHEM 102, Fall 2009, LA TECH
1-38
Reactions
A ----> B
CHEM 102, Fall 2009, LA TECH
1-39
t1/2 equation
0.693 =
t1/2 =
CHEM 102, Fall 2009, LA TECH
k t1/2
0.693
---k
1-40
Half-life
The half-life and the rate constant are related.
t1/2 =
0.693
k
Half-life can be used
to calculate the first order rate constant.
For our N2O5 example, the reaction took 1900 seconds to react half way so:
k =
=
0.693
t1/2
CHEM 102, Fall 2009, LA TECH
-4 -1
= 3.65 x 10 s
0.693
1900 s
1-41
A Nanoscale View:
Elementary Reactions
Most reactions occur through a series of simple st
Elementary reactions could be
unimolecular - rearrangement of a molecule
bimolecular - reaction involving the collision of tw
termolecular - reaction involving the collision of th
CHEM 102, Fall 2009, LA TECH
1-42
Elementary Reactions and Mechanism
2NO2 (g) + F2 (g)
2NO2F (g)
If the reaction
took place in a single step the rate law would be: rate = k
2
[NO2] [F2]
Observed: rate = k1 [NO2] [ F2]
If the observed rate law is not the same as if the reaction took place in a single
step that more than one step must be involved
CHEM 102, Fall 2009, LA TECH
1-43
Elementary Reactions
A possible reaction mechanism might be:
Step one NO2 + F2
NO2F + F (slow)
Step two NO2 + F
NO2F
(fast)
Overall
2NO2 + F2
2NO2F
Rate Determining Step
slowest step in a multi-step mechanism
the step which determines the overall rate of the re
rate = k1 [NO2] [ F2]
CHEM 102, Fall 2009, LA TECH
1-44
Reaction profile of rate determining s
Potential
Energy
This type of plot
shows the energy
changes during
a reaction.
DH
activation
energy
Reaction coordinate
CHEM 102, Fall 2009, LA TECH
1-45
What Potential Energy Curves Show
Exothermic Reactions
Endothermic Reactions
Activation Energy (Ea) of reactant or the minimum en
Effect of catalysts
Effect of temperature
CHEM 102, Fall 2009, LA TECH
1-46
Examples of reaction profiles
Exothermic reaction
Endothermic reaction
CHEM 102, Fall 2009, LA TECH
1-47
Examples of reaction profiles
High activation energy (kinetic)
Low heat of reaction (thermodynamic)
Low activation energy (kinetic)
High heat of reaction (thermodynamic)
CHEM 102, Fall 2009, LA TECH
1-48
Unimolecular Reaction
cis-2-butene
CHEM 102, Fall 2009, LA TECH
trans-2-butrne
1-49
Bimolecular Reaction
-
I + CH3Br
CHEM 102, Fall 2009, LA TECH
ICH3 + Br
-
1-50
Orientation Probability: Some Unsucce
I + CH3Br
CHEM 102, Fall 2009, LA TECH
ICH3 + Br
-
1-51
Arrhenius Equation: Dependence of R
Rate constant (k)
k = A e-Ea/RT
A = frequency factor: A = p x z
Ea = Activation energy
R = gas constant
T = Kelvin temperature
p = collision factor
z = Orientation factor
CHEM 102, Fall 2009, LA TECH
1-52
Energy Distribution Curves:Activati
CHEM 102, Fall 2009, LA TECH
1-53
Arrhenius Equation: ln form
An alternate form of the Arrhenius equation:
-Ea/RT
k=Ae
ln k =
+ ln A
Ea a straight
1 line of slope -Ea/RT is obtained.
If ln k is plotted against 1/T,
- R
T
Activation energy - Ea
The energy that molecules must have in order to react.
( ) ( )
CHEM 102, Fall 2009, LA TECH
1-54
Calculation of Ea
CHEM 102, Fall 2009, LA TECH
1-55
Rate vs Temperature plot
•
Reaction rates are temperature dependent.
7
Here are rate constants
for N2O5 decomposition
at various temperatures.
o
4 -1
T, C
k x 10 , s
6
4 -1
k x 10 (s )
5
20
25
30
35
40
45
4
3
2
0.235
0.469
0.933
1.82
3.62
6.29
1
0
20
25
30
35
o
CHEM 102, Fall 2009, LA TECH
Temperature ( C)
40
45
50
1-56
Calculation of Ea from N2O5 data
3
y = - 1 2 3 9 2 x + 4 0 .8 0 9
2
S lo p e = - 1 2 3 9 2
= 8 . 3 5 J /m o l K
Ea
= 1 03 kJ / m ol
ln k
1
R
0
-1
-2
0.003 1
CHEM 102, Fall 2009, LA TECH
0.003 2
0.003 3
-1
T
0.003 4
0.003 5
1-57
Collision Model
Three conditions must be met at the nano-scale lev
the molecules must collide;
they must be positioned so that the reacting groups
and the collision must have enough energy to form t
CHEM 102, Fall 2009, LA TECH
1-58
Effect of Concentration
on Frequency of
Bimolecular Collisions
CHEM 102, Fall 2009, LA TECH
1-59
Activated Complex or
Reaction Intermediates
an unstable arrangement of atoms that has the hig
CHEM 102, Fall 2009, LA TECH
1-60
Catalyst
A susbstance which speeds up the rate of a reacti
Homogeneous Catalysis - a catalyst which is in th
Heterogeneous Catalysis - a catalyst which is in th
catalytic converter
• solid catalyst working on gaseous materials
CHEM 102, Fall 2009, LA TECH
1-61
Catalysts Lowers Ea
CHEM 102, Fall 2009, LA TECH
1-62
Catalyzed & Uncatalyzed Reactions
CHEM 102, Fall 2009, LA TECH
1-63
Conversion of NO to N2 + O2
CHEM 102, Fall 2009, LA TECH
1-64
Catalytic Converter
H2O(g) + HCs
catalyst
 CO(g) + H2(g)
(unbalanced)
catalyst
2 H2(g) + 2 NO(g)  N2(g) + 2 H2O(g)
catalyst
HCs + O2(g)  CO2(g) + H2O(g) (unbalanced)
catalyst
CO(g) + O2(g)  CO2(g)
(unbalanced)
yst = Pt-NiO
ned hydrocarbons
CHEM 102, Fall 2009, LA TECH
1-65
Enzymes: Biological catalysts
Biological catalysts
Typically are very large proteins.
Permit reactions to ‘go’ at conditions that the body can tol
Can process millions of molecules every second.
Are very specific - react with one or only a few types of mo
CHEM 102, Fall 2009, LA TECH
1-66
The active site
Enzymes are typically HUGE proteins, yet only a sm
The active site has two
basic components.
catalytic site
binding site
Model of
trios-phosphate-isomerase
CHEM 102, Fall 2009, LA TECH
1-67
Relationship of Enzyme to Substrate
CHEM 102, Fall 2009, LA TECH
1-68
Enzyme Catalyzed Reaction
CHEM 102, Fall 2009, LA TECH
1-69
Maximum Velocity for an Enzyme Catalyze
CHEM 102, Fall 2009, LA TECH
1-70
Enzyme Activity Destroyed by Heat
CHEM 102, Fall 2009, LA TECH
1-71
Reaction Mechanism
A set of elementary reactions which represent the
CHEM 102, Fall 2009, LA TECH
1-72
Mechanism Oxidation of
Iodide Ion by Hydrogen Peroxide
CHEM 102, Fall 2009, LA TECH
1-73
Rate Law of Oxidation of
Iodide Ion by Hydrogen Peroxide
Step 1.
+ I-  HOI + OHslow step - rate determining step, suggests that th
= k[HOOH][I-]
CHEM 102, Fall 2009, LA TECH
1-74
Mechanisms with a Fast Initial Step
(g)
 2NOBr(g)
= k[NO]2[Br2]
CHEM 102, Fall 2009, LA TECH
1-75
Mechanism of NO + Br2
Rate = k[NOBr2][NO]
CHEM 102, Fall 2009, LA TECH
1-76
Rate Constants for NO + Br2
, rate constant k1
), rate constant k-1
constant k2
Step-1
+ rateStep2
NOBr2] - k2[NOBr2]
CHEM 102, Fall 2009, LA TECH
1-77
Relationships of Rate Constants
Br2] ~ k-1[NOBr2]
thus
= (k1/k-1)[NO][Br2]
stituting into
k2[NOBr2][NO]
1/k-1)[NO][Br2])[NO]
2[Br ]
k
/k
)[NO]
2 1 -1
2
CHEM 102, Fall 2009, LA TECH
1-78
Chain Mechanisms
chain initiating step - the step of a mechanism which s
chain propagating step(s) - the step or steps which kee
chain terminating step(s) - the step or steps which bre
CHEM 102, Fall 2009, LA TECH
1-79
Chain Mechanisms
combustion of gasoline in an internal combustion e
chain initiating step - additives which generate free
chain propagating step(s) - steps which generate n
chain terminating step(s) - steps which do not gene
CHEM 102, Fall 2009, LA TECH
1-80