Chapter 10
Radical Reactions
Created by
Professor William Tam & Dr. Phillis Chang
Ch. 10 - 1
About The Authors
These Powerpoint Lecture Slides were created and prepared by Professor
William Tam and his wife Dr. Phillis Chang.
Professor William Tam received his B.Sc. at the University of Hong Kong in
1990 and his Ph.D. at the University of Toronto (Canada) in 1995. He was an
NSERC postdoctoral fellow at the Imperial College (UK) and at Harvard
University (USA). He joined the Department of Chemistry at the University of
Guelph (Ontario, Canada) in 1998 and is currently a Full Professor and
Associate Chair in the department. Professor Tam has received several awards
in research and teaching, and according to Essential Science Indicators, he is
currently ranked as the Top 1% most cited Chemists worldwide. He has
published four books and over 80 scientific papers in top international journals
such as J. Am. Chem. Soc., Angew. Chem., Org. Lett., and J. Org. Chem.
Dr. Phillis Chang received her B.Sc. at New York University (USA) in 1994, her
M.Sc. and Ph.D. in 1997 and 2001 at the University of Guelph (Canada). She
lives in Guelph with her husband, William, and their son, Matthew.
Ch. 10 - 2
1.
Introduction: How Radicals Form
and How They React
 Heterolysis
A B

heterolytic
bond
cleavage
A + B
ions
Homolysis
A B
homolytic
bond
cleavage
A + B
radicals
Ch. 10 - 3
1A. Production of Radicals

Homolysis of covalent bonds
● Need heat or light (hn)
R
O
Cl
O
Cl
R
heat
hn
(alkoxyl radical)
2R
O
(chlorine radical)
2 Cl
Ch. 10 - 4
1B. Reactions of Radicals

Almost all small radicals are short-lived,
highly reactive species
Cl + H
CH3
H + CH3
Cl
R
R +
C
C
C
C
Ch. 10 - 5
2. Homolytic Bond Dissociation
Energies (DH°)
H + H
H H
Ho = 436 kJ/mol
Cl + Cl
Cl
Ho = 243 kJ/mol
Cl
H
H
H + H
Ho = 436 kJ/mol
Cl
Cl
Cl + Cl
Ho = 243 kJ/mol
Bond formation is
an exothermic
process.
Reactions in which
only bond breaking
occurs are always
endothermic.
Ch. 10 - 6

The energies required to break
covalent bonds homolytically are called
homolytic bond dissociation
energies, and they are usually
abbreviated by the symbol DH °
Ch. 10 - 7

Single-Bond Homolytic Dissociation
Energies (DH°) at 25°C
Bond Broken
kJ/mol
H–H
436
F–F
159
Cl–Cl
243
Br–Br
193
I–I
151
Ch. 10 - 8

Single-Bond Homolytic Dissociation
Energies (DH°) at 25°C
Bond Broken
kJ/mol
H–F
570
H–Cl
432
H–Br
366
H–I
298
Ch. 10 - 9

Single-Bond Homolytic Dissociation
Energies (DH°) at 25°C
Bond Broken
kJ/mol
Bond Broken
kJ/mol
H3C–H
440
H3C–F
461
H3C–Cl
352
H3C–Br
293
H3C–OH
387
H3C–I
240
H3C–OCH3
348
Ch. 10 - 10

Single-Bond Homolytic Dissociation
Energies (DH°) at 25°C
Bond Broken
Cl
Cl
Cl
kJ/mol
Bond Broken
Br
354
355
349
Br
Br
kJ/mol
294
298
292
Ch. 10 - 11

Single-Bond Homolytic Dissociation
Energies (DH°) at 25°C
Bond Broken
H
H
Ph
kJ/mol
Bond Broken
H
423
413
H
400
H
375
Ph
HC
C
kJ/mol
369
H
465
H
474
H
547
Ch. 10 - 12
2A. Use Homolytic Bond Dissociation
Energies to Calculate Heats of Reaction
o
o
(DH = 432 kJ/mol) ☓ 2
(DH = 436 kJ/mol)
o
(DH = 243 kJ/mol)
H
H
+
Cl
+679 kJ is required
to cleave 1 mol of
H2 bonds and
1 mol of Cl2 bonds
Cl
2H
Cl
-864 kJ is evolved
in formation of
bonds in
2 mol of HCl
Ch. 10 - 13
o
H

= 2 (432 kJ/mol) + (436 kJ/mol + 243 kJ/mol)
= 864 kJ/mol + 679 kJ/mol
= 185 kJ/mol
Overall, the reaction of 1 mol of H2 and
1 mol of Cl2 to form 2 mol of HCl is
exothermic
Ch. 10 - 14
2B. Use Homolytic Bond Dissociation
Energies to Determine the Relative
Stabilities of Radicals
Ho = +423 kJ/mol
H
+ H
Propyl radical
(a 1o radical)
Ho = +413 kJ/mol
+
H
H
Isopropyl radical
(a 2o radical)
Ch. 10 - 15
+
tert-Butyl radical
H
(a 3o radical)
H
Ho = +400 kJ/mol
H
+
Isobutyl radical
(a 1o radical)
H
Ho = +422 kJ/mol
Ch. 10 - 16

Relative Stability
● Carbon radicals are considered to
be electron deficient (similar to
carbocations), thus electron
donating groups stabilize radicals

o
C
o
3 >2 >1
CH3
CH3
o
H
CH3 > CH3
C
H
CH3 > CH3
C
H
H>H
C
H
(positive inductive effect of alkyl groups stabilize radical)
Ch. 10 - 17
3. The Reactions of Alkanes with
Halogens

Alkanes have no functional group and
are inert to many reagents and do not
undergo many reactions

Halogenation of alkanes is one of the
most typical free radical reactions
Ch. 10 - 18

Alkanes react with molecular halogens
to produce alkyl halides by a
substitution reaction called radical
halogenation
R
H + X2
heat
or
light (hn)
R
X + H
X
Ch. 10 - 19
3A. Multiple Halogen Substitution
H
H
C
H
H + Cl2
heat
or
light
H
H
C
Cl
Cl + H
H
C
H
+H
Cl
H
Cl
+ Cl
C
Cl
Cl + Cl
C
Cl
Cl
Cl
Ch. 10 - 20
3B. Lack of Chlorine Selectivity

Chlorination of most higher alkanes
gives a mixture of isomeric monochloro
products as well as more highly
halogenated compounds
● Chlorine is relatively unselective;
it does not discriminate greatly
among the different types of
hydrogen atoms (primary,
secondary, and tertiary) in an
alkane
Ch. 10 - 21
Cl2
Cl +
light
Isobutane
Isobutyl
chloride
(48%)
Cl
tert-Butyl
chloride
(29%)
Polichlorinated
+
+ HCl
products
(23%)
● Because alkane chlorinations
usually yield a complex mixture of
products, they are not useful as
synthetic methods when the goal is
preparation of a specific alkyl
chloride
Ch. 10 - 22
● An exception is the halogenation of
an alkane (or cycloalkane) whose
hydrogen atoms are all equivalent.
[Equivalent hydrogen atoms are
defined as those which on
replacement by some other group
(e.g., chlorine) yield the same
compound.]
Ch. 10 - 23
+ Cl2
Neopentane
(excess)
heat
or
light
Cl
+ H
Cl
Neopentyl
chloride
● Bromine is generally less reactive
toward alkanes than chlorine, and
bromine is more selective in the
site of attack when it does react
Ch. 10 - 24
4. Chlorination of Methane:
Mechanism of Reaction

Most radical reactions include 3 stages
(steps)
(1) chain initiation
(2) chain propagation
(3) chain termination
Ch. 10 - 25

Mechanism of Free Radical Chlorination
of CH4
(1) Chain initiation
Cl
Cl
hn
(homolytic
cleavage)
2 Cl
● Radicals are created in this step
Ch. 10 - 26
(2) Chain propagation
H
(i) Cl
+
H
C
H
H
Cl +
CH3
H
(ii) CH3
+
Cl
Cl
CH3Cl
+
Cl
● Repeating (i) and (ii) in a chain
reaction provides the product CH3Cl
● In chain propagation, one radical
generates another and the process
goes on
Ch. 10 - 27
(2) Chain propagation
● Other than CH3Cl, other
chlorination products can be formed
in the chain propagation step
Cl
(ia) Cl
+
H
C
H
H
Cl +
CH2Cl
H
(iia)
CH2Cl +
Cl
Cl
CH2Cl2 +
Cl
Ch. 10 - 28
(2) Chain propagation
Cl
(ib) Cl
+
H
C
Cl
H
Cl +
CHCl2
H
(iib)
CHCl2 +
Cl
Cl
CHCl 3
+
Cl
Cl
(ic) Cl
+
H
C
Cl
H
Cl +
CCl3
Cl
(iic)
CCl3
+
Cl
Cl
CCl4
+
Cl
Ch. 10 - 29
(3) Chain termination
Cl
+
CH3 +
Cl
+
CH3
Cl
CH3
H3C
CH2Cl
CH2Cl2
CHCl2 + CCl3
CH3
Cl2HC
CH3
CCl3
Ch. 10 - 30
(3) Chain termination
● Free radical reactions cannot be
completed without chain
termination
● All radicals are quenched in this
step
● Radical reactions usually provide
mixture of many different products
● Synthesis of CH3Cl or CCl4 is
possible using different amounts of
reactants (CH4 and Cl2)
Ch. 10 - 31
e.g.:
CH4 (large excess) + Cl2
hn
CH3Cl (mainly)
CH4 + Cl2 (large excess)
hn
CCl4 (mainly)
Ch. 10 - 32
5. Chlorination of Methane:
Energy Changes

Chain initiation
Step 1
Cl
Cl
(DHo = 243)
2 Cl
Ho = +243 kJ/mol
Ch. 10 - 33

Chain propagation
Step 2
H3C
CH3 + H
H + Cl
(DHo = 440)
Cl
(DHo = 432)
Ho = +8 kJ/mol
Step 3
CH3 + Cl
Cl
(DHo = 243)
H3C
Cl + Cl
(DHo = 352)
Ho = 109 kJ/mol
Ch. 10 - 34

Chain termination
CH3 + Cl
H3C
Ho = 352 kJ/mol
Cl
(DHo = 352)
CH3 + CH3
H3C
o
CH3 Ho = 378 kJ/mol
(DH = 378)
Cl
+ Cl
Cl
Cl
Ho = 243 kJ/mol
o
(DH = 243)
Ch. 10 - 35

The addition of the chain-propagation
steps yields the overall equation for the
chlorination of methane
H3C
CH3 + H
H + Cl
Cl
Ho = +8 kJ/mol
CH3 + Cl
Cl
H3C
Cl + Cl
o
H = 109 kJ/mol
H3C
H + Cl
Cl
H3C
Cl + H
Cl
Ho = 101 kJ/mol
Ch. 10 - 36
5A. The Overall Free-Energy Change
o
o
G =  H – T S

o
For many reactions the entropy change
o
is so small that the term T S in the
above expression is almost zero, and
o
o
G is approximately equal to H
Ch. 10 - 37
CH4 + Cl2  CH3Cl + HCl


2 mol of the products are formed from the
same number of moles of the reactants
● Thus the number of translational
degrees of freedom available to
products and reactants is the same
CH3Cl is a tetrahedral molecule like CH4, and
HCl is a diatomic molecule like Cl2
● This means that vibrational and
rotational degrees of freedom available
to products and reactants should also be
approximately the same
Ch. 10 - 38
CH4 + Cl2  CH3Cl + HCl




o
S = +2.8 J K-1 mol-1
At room temperature (298 K) the TS
term is 0.8 kJ mol-1
o
o
H = 101 kJ mol-1
o
G = 102 kJ mol-1
Ch. 10 - 39
5B. Activation Energies

A low energy of activation means a reaction
will take place rapidly; a high energy of
activation means that a reaction will take
place slowly
Chain initiation
Step 1
Cl2  2 Cl •
Eact = +243 kJ/mol
Chain propagation
Step 2
Cl • + CH4  HCl + CH3 •
Step 3
Eact = +16 kJ/mol
Cl • + Cl2  CH3Cl + Cl •
Eact = ~8 kJ/mol
Ch. 10 - 40

Estimates of energies of activation
(1) Any reaction in which bonds are
broken will have an energy of activation
greater than zero. This will be true even
if a stronger bond is formed and the
reaction is exothermic. The reason:
Bond formation and bond breaking do
not occur simultaneously in the
transition state. Bond formation lags
behind, and its energy is not all
available for bond breaking
Ch. 10 - 41

Estimates of energies of activation
(2) Activation energies of endothermic
reactions that involve both bond
formation and bond rupture will be
o
greater than the heat of reaction, H
H + Cl
CH3 + H Cl
o
o
(DH = 440) H = +8 kJ/mol (DHo = 432)
Eact = +16 kJ/mol
H3C
H + Br
CH3 + H Br
o
o
(DH = 440) H = +74 kJ/mol (DHo = 366)
Eact = +78 kJ/mol
H3C
Ch. 10 - 42
Ch. 10 - 43

Estimates of energies of activation
(3) The energy of activation of a gasphase reaction where bonds are broken
homolytically but no bonds are formed
o
is equal to H
Cl
Cl
(DHo = 243)
2 Cl
o
H = +243 kJ/mol
Eact = +243 kJ/mol
Ch. 10 - 44

Estimates of energies of activation
(4) The energy of activation for a gasphase reaction in which small radicals
combine to form molecules is usually
zero
2 CH3
H3C
CH3
(DHo = 378)
o
H = 378 kJ/mol
Eact = 0
Ch. 10 - 45
5C. Reaction of Methane with Other
Halogens
FLUORINATION
o
H
(kJ/mol)
(kJ/mol)
+159
+159
F • + CH4  HF + • CH3
130
+5.0
CH3 + F2  CH3F + F •
302
small
Eact
Chain initiation
F2  2 F •
Chain propagation
•
o
Overall H =
432
Ch. 10 - 46
CHLORINATION
o
H
(kJ/mol)
(kJ/mol)
+243
+243
+8
+16
109
small
Eact
Chain initiation
Cl2  2 Cl •
Chain propagation
Cl • + CH4  HCl + • CH3
•
CH3 + Cl2  CH3Cl + Cl •
o
Overall H =
101
Ch. 10 - 47
BROMINATION
o
H
(kJ/mol)
(kJ/mol)
+193
+193
+74
+78
100
small
Eact
Chain initiation
Br2  2 Br •
Chain propagation
Br • + CH4  HBr + • CH3
•
CH3 + Br2  CH3Br + Br •
o
Overall H =
26
Ch. 10 - 48
IODINATION
o
H
(kJ/mol)
(kJ/mol)
+151
+151
I • + CH4  HI + • CH3
+142
+140
CH3 + I2  CH3I + I •
89
small
Eact
Chain initiation
I2  2 I •
Chain propagation
•
o
Overall H =
+53
Ch. 10 - 49
6.
Halogenation of Higher Alkanes

Mechanism for radical halogenation of
ethane
Chain initiation
Step 1
Cl2
light
or heat
2 Cl
Chain propagation
Step 2
CH3CH2 H + Cl
CH3CH2 + H Cl
Step 3
CH3CH2 + Cl Cl
CH3CH2 Cl + Cl
Ch. 10 - 50
Chain termination
CH3CH2
+
Cl
CH3CH2
+
CH3CH2
Cl
+ Cl
CH3CH2 Cl
CH3CH2 CH2CH3
Cl Cl
Ch. 10 - 51
Cl
Cl2
Cl +
light
25oC
Cl2
Cl +
light
Cl
25oC
Cl2
Cl
+
300oC
Cl
+
+
Cl
Cl
Ch. 10 - 52
6A. Selectivity of Bromine

Bromination is slower than chlorination
because the 1st propagation step is
more endothermic (overall still
exothermic). As a result, bromination
is more selective than chlorination
H + Br2
hn
Br
(99%)
+
H
Br
(< 1%)
Ch. 10 - 53

Br
Mechanism
hn
Br
H
2 Br
Br
+ H
Br
(major; 3o radical more stable)
H
H
Br
H
+
H
Br
(minor; 1o radical less stable)
Ch. 10 - 54

Mechanism
Br
(major)
Br
H
H
Br
Br
(minor)
Ch. 10 - 55
Br2
Br +
hn
127oC
(trace)
Cl2
Cl +
hn
25oC
Br
(> 99%)
(63%)
Cl
(37%)
Ch. 10 - 56
7.
The Geometry of Alkyl Radicals
p-orbital
R
C
R
R
sp2 hybridized

Planar, similar to carbocation
Ch. 10 - 57
8. Reactions That Generate
Tetrahedral Chirality Centers
Cl
Cl2
Pentane
(achiral)
Cl
achiral
1-Chloropentane
(achiral)
+
*
( )-2-Chloropentane
(a racemic form)
+
Cl
3-Chloropentane
(achiral)
Ch. 10 - 58

The Stereochemistry of chlorination at
C2 of pentane C2
CH3CH2CH2CH2CH3
Cl
+ Cl
CH3
Cl
+ Cl
CH3
Cl2
H
(a)
CH2CH2CH3
(S)-2-Chloropentane
(50%)
Cl2
C
H
H3C
H
(b)
H3CH2CH2C
Cl
CH2CH2CH3
trigonal planar
radical (achiral)
enantiomers
(R)-2-Chloropentane
(50%)
Ch. 10 - 59
8A. Generation of a Second Chirality
Center in a Radical Halogenation
Cl
Cl
Cl2
3
2
diastereomers
H
hn
H
(S)-2Chloropentane
(chiral)
2
3
H
2
3
Cl
Cl
(2S,3S)Dichloropentane
(chiral)
(2S,3R)Dichloropentane
(chiral)
from bottom Cl2
face
Cl
+
Cl
trigonal planar
from top Cl2
face
Ch. 10 - 60

Note that other products are formed,
of course, by chlorination at other
carbon atoms
Ch. 10 - 61
9.
Radical Addition to Alkenes: The
Anti-Markovnikov Addition of
Hydrogen Bromide
 Anti-Markovnikov addition of HBr to
alkenes – peroxide effect
● Addition of HBr to alkenes usually
follows Markovnikov’s rule
HBr
Br H
not
H
Br
Ch. 10 - 62
● In the presence of peroxides (RO–
OR), anti-Markovnikov addition is
observed
HBr
H
Br
not
Br H
RO OR
heat
Ch. 10 - 63

Mechanism
● Via a radical mechanism
RO
OR
RO
+
heat
(homolytic cleavage)
(chain initiation)
H
Br
2 RO
ROH + Br
Ch. 10 - 64
Br
+
Br
o
(3 radical, more stable)
Br
not
Br
+
o
(1 radical, less stable)
H
+ H
Br
Br
+ Br
Br
Ch. 10 - 65

Synthetic application
Br
HBr
(via more stable
2o carbocation)
HBr
RO-OR
heat
Br
(via more stable
2o radical)
Ch. 10 - 66

Hydrogen bromide is the only hydrogen
halide that gives anti-Markovnikov
addition when peroxides are present

Hydrogen fluoride, hydrogen chloride,
and hydrogen iodide do not give antiMarkovnikov addition even when
peroxides are present
Ch. 10 - 67
10. Radical Polymerization of
Alkenes: Chain-Growth Polymers
n
CH2
CH2
(monomer)
peroxide
heat
CH2CH2
n
(polymer)
Ch. 10 - 68

Via radical mechanism
(i)
RO
OR
(ii)
RO
+
(iii)
RO
(iv)
RO

H2C
2 RO
CH2
CH2CH2
CH2CH2
+
H2C
RO
CH2CH2CH2CH2
CH2CH2CH2CH2
RO
RO
CH2
+
CH2CH2
H2C
2
CH2
CH2CH2
Ch. 10 - 69
(v)
RO
CH2CH2
RO
(vi)
RO
CH2CH2
+
2
CH2CH2
+
CH2CH2
n
OR
CH2CH2OR
CH2CH2
n
RO
CH2CH2
RO
CH2CH2
n
CH2CH2
OR
Ch. x
10 - 70

Other common polymers
n CH2 CHCH 3
Polypropylene
n CH2 CHCl
ROOR

CH3 n
ROOR

PVC (plumbing polymer)
n CF2 CF2
CH2CH
ROOR

Polytetrafluroethene (Teflon)
CH2CH
Cl
CF2CF2
n
n
Ch. 10 - 71

Other common polymers
CH3
n
CH2
C
COOMe
ROOR

CH3
CH2C
COOMe
Polymethyl methacrylate
(windshield, contact lenses)
n
CH2
CHPh
ROOR

Polysterene (styrofoam,
coffee cup, etc.)
n
CH2CH
Ph
n
Ch. 10 - 72
 END OF CHAPTER 10 
Ch. 10 - 73