This chapter discusses the study of chemical reactions through mechanisms, kinetics, and thermodynamics. It introduces tools for determining reaction mechanisms like equilibrium constants, free energy change, enthalpy, entropy, bond dissociation energy and activation energy. Free radical chain reactions are covered as well as initiation, propagation, and termination steps. Factors that determine the rate and favorability of reactions like temperature, reaction order, and transition state stability are also examined.

1. Chapter 4
The Study of
Chemical Reactions
Jo Blackburn
Richland College, Dallas, TX
Dallas County Community College District
© 2003, Prentice Hall
Organic Chemistry, 5th
Edition
L. G. Wade, Jr.

2. Chapter 4 2
Tools for Study
• To determine a reaction’s
mechanism, look at:
Equilibrium constant
Free energy change
Enthalpy
Entropy
Bond dissociation energy
Kinetics
Activation energy =>

3. Chapter 4 3
Chlorination of Methane
• Requires heat or light for initiation.
• The most effective wavelength is blue, which
is absorbed by chlorine gas.
• Lots of product formed from absorption of
only one photon of light (chain reaction).
=>
C
H
H
H
H + Cl2
heat or light
C
H
H
H
Cl + HCl

4. Chapter 4 4
Free-Radical Chain Reaction
• Initiation generates a reactive intermediate.
• Propagation: the intermediate reacts with a
stable molecule to produce another reactive
intermediate (and a product molecule).
• Termination: side reactions that destroy the
reactive intermediate.
=>

5. Chapter 4 5
Initiation Step
A chlorine molecule splits homolytically
into chlorine atoms (free radicals)
=>
Cl Cl + photon (hν) Cl + Cl

6. Chapter 4 6
Propagation Step (1)
The chlorine atom collides with a
methane molecule and abstracts
(removes) a H, forming another free
radical and one of the products (HCl).
C
H
H
H
H Cl+ C
H
H
H
+ H Cl
=>

7. Chapter 4 7
Propagation Step (2)
The methyl free radical collides with
another chlorine molecule, producing
the other product (methyl chloride) and
regenerating the chlorine radical.
C
H
H
H
+ Cl Cl C
H
H
H
Cl + Cl
=>

8. Chapter 4 8
Overall Reaction
C
H
H
H
H Cl+ C
H
H
H
+ H Cl
C
H
H
H
+ Cl Cl C
H
H
H
Cl + Cl
C
H
H
H
H + Cl Cl C
H
H
H
Cl + H Cl =>
Cl Cl + photon (hν) Cl + Cl

9. Chapter 4 9
Termination Steps
• Collision of any two free radicals
• Combination of free radical with
contaminant or collision with wall.
C
H
H
H
Cl+ C
H
H
H
Cl
Can you suggest others?
=>

10. Chapter 4 10
Equilibrium constant
• Keq = [products]
[reactants]
• For chlorination Keq = 1.1 x 1019
• Large value indicates reaction “goes to
completion.”
=>

11. Chapter 4 11
Free Energy Change
• ∆G = free energy of (products -
reactants), amount of energy available to
do work.
• Negative values indicate spontaneity.
• ∆Go
= -RT(lnKeq)
where R = 1.987 cal/K-mol
and T = temperature in kelvins
• Since chlorination has a large Keq, the free
energy change is large and negative.
=>

12. Chapter 4 12
Problem
• Given that -X is -OH, the energy difference for
the following reaction is -1.0 kcal/mol.
• What percentage of cyclohexanol molecules
will be in the equatorial conformer at
equilibrium at 25°C?
=>

13. Chapter 4 13
Factors Determining ∆G°
• Free energy change depends on
enthalpy
entropy
∀∆H° = (enthalpy of products) - (enthalpy
of reactants)
∀∆S° = (entropy of products) - (entropy of
reactants)
∀∆G° = ∆H° - T∆S° =>

14. Chapter 4 14
Enthalpy
• ∆Ho
= heat released or absorbed during
a chemical reaction at standard
conditions.
• Exothermic, (-∆H), heat is released.
• Endothermic, (+∆H), heat is absorbed.
• Reactions favor products with lowest
enthalpy (strongest bonds).
=>

15. Chapter 4 15
Entropy
• ∆So
= change in randomness, disorder,
freedom of movement.
• Increasing heat, volume, or number of
particles increases entropy.
• Spontaneous reactions maximize
disorder and minimize enthalpy.
• In the equation ∆Go
= ∆Ho
- T∆So
the
entropy value is often small.
=>

16. Chapter 4 16
Bond Dissociation Energy
• Bond breaking requires energy (+BDE)
• Bond formation releases energy (-BDE)
• Table 4.2 gives BDE for homolytic
cleavage of bonds in a gaseous molecule.
A B A + B
We can use BDE to estimate ∆H for a reaction.
=>

17. Chapter 4 17
Which is more likely?
Estimate ∆H for each step using BDE.
CH4 HCl+ +Cl CH3
CH3 + Cl2 CH3Cl + Cl
or
Cl+CH4 CH3Cl + H
H Cl2+ HCl Cl+
104 103
58 84
=>
104 84
58 103

18. Chapter 4 18
Kinetics
• Answers question, “How fast?”
• Rate is proportional to the concentration
of reactants raised to a power.
• Rate law is experimentally determined.
=>

19. Chapter 4 19
Reaction Order
• For A + B → C + D, rate = k[A]a
[B]b
a is the order with respect to A
a + b is the overall order
• Order is the number of molecules of that
reactant which is present in the rate-
determining step of the mechanism.
• The value of k depends on temperature as
given by Arrhenius: ln k = -Ea + lnA
RT
=>

20. Chapter 4 20
Activation Energy
• Minimum energy required to reach
the transition state.
• At higher temperatures, more
molecules have the required energy.
=>
C
H
H
H
H Cl

21. Chapter 4 21
Reaction-Energy Diagrams
• For a one-step reaction:
reactants → transition state → products
• A catalyst lowers the energy of the
transition state.
=>

22. Chapter 4 22
Energy Diagram for a
Two-Step Reaction
• Reactants → transition state → intermediate
• Intermediate → transition state → product
=>

23. Chapter 4 23
Rate-Determining Step
• Reaction intermediates are stable as long
as they don’t collide with another molecule
or atom, but they are very reactive.
• Transition states are at energy maximums.
• Intermediates are at energy minimums.
• The reaction step with highest Ea will be the
slowest, therefore rate-determining for the
entire reaction. =>

24. Chapter 4 24
Rate, Ea, and Temperature
X + CH4 HX + CH3
X E a Rate @ 300K Rate @ 500K
F 1.2 kcal 140,000 300,000
Cl 4 kcal 1300 18,000
Br 18 kcal 9 x 10-8
0.015
I 34 kcal 2 x 10-19
2 x 10-9
=>

25. Chapter 4 25
Conclusions
• With increasing Ea, rate decreases.
• With increasing temperature, rate
increases.
• Fluorine reacts explosively.
• Chlorine reacts at a moderate rate.
• Bromine must be heated to react.
• Iodine does not react (detectably).
=>

26. Chapter 4 26
Chlorination of Propane
• There are six 1° H’s and two 2° Η’s. We
expect 3:1 product mix, or 75% 1-
chloropropane and 25% 2-chloropropane.
• Typical product mix: 40% 1-chloropropane
and 60% 2-chloropropane.
• Therefore, not all H’s are equally reactive.
=>
1° C
2° C
CH3 CH2 CH3 + Cl2
hν CH2
Cl
CH2 CH3 + CH3 CH
Cl
CH3

27. Chapter 4 27
Reactivity of Hydrogens
• To compare hydrogen reactivity, find
amount of product formed per hydrogen:
40% 1-chloropropane from 6 hydrogens
and 60% 2-chloropropane from 2
hydrogens.
• 40% ÷ 6 = 6.67% per primary H and
60% ÷ 2 = 30% per secondary H
• Secondary H’s are 30% ÷ 6.67% = 4.5
times more reactive toward chlorination
than primary H’s. =>

28. Chapter 4 28
Predict the Product Mix
Given that secondary H’s are 4.5 times as
reactive as primary H’s, predict the
percentage of each monochlorinated
product of n-butane + chlorine.
=>

29. Chapter 4 29
Free Radical Stabilities
• Energy required to break a C-H bond
decreases as substitution on the carbon
increases.
• Stability: 3° > 2° > 1° > methyl
∆H(kcal) 91, 95, 98, 104
=>

30. Chapter 4 30
Chlorination Energy Diagram
Lower Ea, faster rate, so more stable
intermediate is formed faster.
=>

31. Chapter 4 31
• There are six 1° H’s and two 2° Η’s. We
expect 3:1 product mix, or 75% 1-
bromopropane and 25% 2-bromopropane.
• Typical product mix: 3% 1-bromopropane
and 97% 2-bromopropane !!!
• Bromination is more selective than
chlorination. =>
1° C
2° C
CH3 CH2 CH3 + CH2
Br
CH2 CH3 +Br2
heat
CH3 CH
Br
CH3
Bromination of Propane

32. Chapter 4 32
• To compare hydrogen reactivity, find amount of
product formed per hydrogen: 3% 1-
bromopropane from 6 hydrogens and 97% 2-
bromopropane from 2 hydrogens.
• 3% ÷ 6 = 0.5% per primary H and
97% ÷ 2 = 48.5% per secondary H
• Secondary H’s are 48.5% ÷ 0.5% = 97 times
more reactive toward bromination than primary
H’s.
=>
Reactivity of Hydrogens

33. Chapter 4 33
Bromination Energy Diagram
• Note larger difference in Ea
• Why endothermic?
=>

34. Chapter 4 34
Bromination vs. Chlorination
=>

35. Chapter 4 35
Endothermic and
Exothermic Diagrams
=>

36. Chapter 4 36
Hammond Postulate
• Related species that are similar in energy are
also similar in structure. The structure of a
transition state resembles the structure of the
closest stable species.
• Transition state structure for endothermic
reactions resemble the product.
• Transition state structure for exothermic
reactions resemble the reactants.
=>

37. Chapter 4 37
Radical Inhibitors
• Often added to food to retard spoilage.
• Without an inhibitor, each initiation step
will cause a chain reaction so that many
molecules will react.
• An inhibitor combines with the free
radical to form a stable molecule.
• Vitamin E and vitamin C are thought to
protect living cells from free radicals.
=>

38. Chapter 4 38
Reactive Intermediates
• Carbocations (or carbonium ions)
• Free radicals
• Carbanions
• Carbene
=>

39. Chapter 4 39
Carbocation Structure
• Carbon has 6 electrons,
positive charge.
• Carbon is sp2
hybridized
with vacant p orbital.
=>

40. Chapter 4 40
Carbocation Stability
• Stabilized by alkyl
substituents 2 ways:
• (1) Inductive effect:
donation of electron
density along the
sigma bonds.
• (2) Hyperconjugation:
overlap of sigma
bonding orbitals with
empty p orbital.
=>

41. Chapter 4 41
Free Radicals
• Also electron-
deficient
• Stabilized by alkyl
substituents
• Order of stability:
3° > 2° > 1° > methyl
=>

42. Chapter 4 42
Carbanions
• Eight electrons on C:
6 bonding + lone pair
• Carbon has a negative
charge.
• Destabilized by alkyl
substituents.
• Methyl >1° > 2 ° > 3 °
=>

43. Chapter 4 43
Carbenes
• Carbon is neutral.
• Vacant p orbital, so
can be electrophilic.
• Lone pair of
electrons, so can be
nucleophilic.
=>

44. Chapter 4 44
End of Chapter 4