BOND DISSOCIATION ENERGIES IN SULPHUR ORGANIC COMPOUNDS
•·"
•
BOND DISSOCIATION ENERGIES
IN
SULPHUR ORGANIC
CO~OUNDS
by
Robert L. Lauchlan, B.Sc.
A thesis submitted to the Faculty of
Graduate Studies and Research in· partial
fulfilment of the requirements for the
degree of Doctor of Philosopny.
Department of Chemistry,
McGill University,
Montreal.
October 1965.
•
AOKNOWLEDGEMENTS
The author wishes to acknowledge the guidance
of Dr. A. Sehon in the supervision of this research
and in the preparation of the manuscript.
Mr. S. Meyerson of the American Oil Company,
Whiting, Indiana, contributed mass spectroscopi e
analyses,· providing the author invaluable assistance.
Grants from the Petroleum Research Fund of the
Ameriean Chemical Society and the National Research
Oouncil of Canada made to Dr. A. Sehon in support of
this research are gratefully acknowledged .
The author is deeply indebted to his wife for
the editing and typing of this manuscript.
TABLE OF CONTENTS
Introduction
1
Factors Affecting Bond Dissociation Energies 7
Methode of Measuring Bond Dissociation
Energies
18
Spectroscopie Methode
19
Electron Impact Methode
23
Thermal Methode
A.
Equilibrium Methode
28
B.
Kinetic Methode
34
Carbon-Sulpnur Bond Dissociation Energies
Review
47
Present Investigation
51
Experimental
54
Materials
54
Apparat us
54
Procedure and Analyses
61
Resulte and Discussion
A.
B.
67
Phenyl Ethyl Sulpnide
Products
67
Mechanism of Decomposition
72
Bond Energies
78
Phenyl n-Propyl Sulpnide
Products
88
Mechanism of Decomposition
89
TABLE OF CONTENTS (cont.)
Bond Energies
C.
94
Phenyl Benzyl Sulphide
Pr~oo~
1~
Mechanism of Decomposition
105
Bond Energies
106
General Discussion
Experimental Errors
111
111
Dissociation Energies and Heats of
Formation of Radicals
112
The Stability of the Pheny1sulphide
Radical
115
Appendix A
Pyrolysis of Diphenyl Disulphide
116
Appendix B
Data and Calculations for Experiment 20
121
Summary and Contributions to Knowledge
123
Bibliography
126
!NDEX TO F IGJ.Jl\ES
1.
2.
Graphie representation of bond dissociation
energies and ionie-covalent resonance
14
Appara tus
58
3A. Section of apparatus for the introduction of
liquid su1phides
59
3B. Section of apparatus for the introduction of
solid sulphides
59
4.
Reaction Vessel
60
5.
Titration curves for mercaptans
69
6.
Titration curves for su1phides
70
7.
Separation of ethane and ethylene on column of
high activity silioa gel by gas ohromatography
8.
Plot of log k vs. 1/T
(pheny1 ethyl sulphide)
9.
71
81
Plot of log k vs. 1/T
(phenyl n-propyl sulphide)
98
10. Plot of log k' vs. 1/T
(phenyl benzyl sulphide)
108
INDEX TO TABLES
1.
Bond dissociation energies and bond energy
terrns
6
2.
C-H bond dissociation energies in hydrocarbons
8
3.
Dissociation energies of RX bonds
12
4.
Dissociation energies of the metal alkyls
16
5.
Primary C-S bond dissociation energies
49
6.
Products of the decomposition of phenyl ethyl
sulphide
7.
82
The decomposition of phenyl ethyl sulphide
Variation of k with contact time
8.
83
The decomposition of phenyl ethyl sulphide
Variation of k with the partial pressure of
84
sulphide
9.
The decomposition of phenyl ethyl sulphide
Variation of k with toluene pressure
85
10. The decomposition of pnenyl ethyl sulphide
Effect of the packed reaction vessel on k
86
11. A comparison of abstraction and decomposition
reactions for C2H5
12. Products of the decomposition of phenyl
87
n-propyl sulphide
99
13. The decomposition of phenyl n-propyl sulphide
Variation of k with contact time
100
~~
~
--
~~~~~~~--~~~~---
INDEX TO TABLES (cont.)
14. The decomposition of phenyl n-propyl sulphide
Variation of k with the partial pressure of
sulphide
101
15. The decomposition of phenyl n-propyl sulphide
Variation of k with toluene pressure
102
16. The decomposition of phenyl n-propyl sulphide
Effect of the packed reaction vessel on k
103
17. Products of the decomposition of phenyl benzyl
sulphide
109
18. Products of the decomposition of phenyl benzyl
sulphide in toluene-J:-d
llO
disulphide
120
3
19. Products of the decomposition of diphenyl
INTROIUlCTION
In a discussion of the strengths of ohemioal bonds
it is important to distinguish between two concepts, the
'bond dissociation energy' and the 'bond energy term'.
Since chemical reactions involve the formation and
rupture of bonds, it is clear that bond dissociation
energies must be considered among the main determining
factors of chemical reactivity.
However, bond properties,
auch as the bond length, are more effectively correlated
with the bond energy term which representa that fraction
of the total molecular binding energy that can be ascribed
to a particular bond as it exista in the undisturbed
molecule.
The bond dissociation energy, as a pbysical concept,
has a great advantage over the bond energy term, since it
can be defined without ambiguity.
Szwarc (1) has defined
the bond dissociation energy, D(A-B), of a bond A-B as the
change· ih internai energy at absolute zero, .6E~, in the
ideal gas state for the reaction
A-B
~
A+B
(1)
the products existing in their ground states or specified
electronic states.
Dissociation energies derived from
spectral data refer to absolute zero, but those derived
- 2 -
from equilibrium or kinetic data often refer to the
temperature used in the particular experimenta, since
the necessary heat capacity data required to make the
correction to absolute zero are frequently lacking.
0
0
Cottrell (2) has shawn the difference, AH2gs - AE 0 to
be about 4RT, a value which is well within the experimental error involved in bond dissociation energy
measurements.
In contrast to bond dissociation energies, the bond
energy term cannat be defined with complete precision
except in the restricted case of a diatomic or symmetrical
polyatomic molecule.
The sole criterion of correctness in
apportioning the total energy of a molecule among the
various bonds is that the sum of all its bond energy terms
is equal to the heat of atomization, H:.
The simplest
scheme, as proposed by Fajans (3), assumes a theory of
local pair bonds without interaction between bonds.
Each
bond of a given type has a constant energy term, irrespective of the molecular environment, and all bond energy
terms are additive.
For a molecule ABn containing only
A-B bonds, the bond energy term is then defined as 1/n
the heat of atomization of the gaseous molecule at 298.15°K.
(2)
- 3 Applied to methane, this definition gives a value of
E(C-H)
= 99.29
kcal./mole (4).
However, the extension
of equation (2) to the evaluation of E(C-H) and E{C-C)
terms in the straight-chain paraffin hydrocarbons leads
to various values depending on which members of the series
are compared.
For example, the heats of atomization of
butane and pentane give a value of E(C-H)
= 98.70
kcal./
mole, which is slightly lower than the value calculated
for E(C-H) in methane.
In addition, this bond energy
scheme fails to account for the significant differences
in the heats of formation of isomers such as 1:4-pentadiene
and 1:3-pentadiene (5), and erroneously classifies all
gaseous redistribution reactions as thermoneutral (6).
The initial proposal by Zahn (?) that interaction
between localized bonds within the same molecule is
significant received theoretical support from studies by
Dewar and Pettit (8), who used the perturbation theory for
their ca1cu1ations, and by Brown ,(9), who emp1oyed the
1inear combination of bond orbitals approximation.
\
Severa1 authors have subsequently proposed schemes which
differ from that of Fajans in at 1east one of the initial
postu1ates.
Laid1er (10,11) has accounted for the heats
of formation of paraffin hydrocarbons in terms of a scheme
which assumes that the strength of the C-C bond is constant,
- 4 but that of the C-H bond varies according to whether
it is primary, secondary, or tertiary.
Tatevskii (12)
has extended this classification to include
c-c
bonds.
Similarly, Dewar and Schmeising (13) devised a scheme
for olefine in which the term values for different
types of C-H and
C~C
bonds are classified according to
the state of hybridization of the carbon atoms involved.
Bond energy terms in general have been criticized
(14) as being unrealistic since the atoms in molecules,
particularly in the transition state, seldom exhibit the
same valence as in their ground state.
The role of the
valence state cannat, however, be ignored in the
evaluation of dissociation energies.
This is the prin-
cipal reason not only for the difference in the energies
of the stepwise dissociation processes of a polyatomic
molecule but also for the difference between these values
and the corresponding bond energy terms as shawn in Table
I.
For example, the C-F bond dissociation energies in
tetrafluoromethane are neither constant nor equivalent to
the bond energy term.
Values of D(CF 3-F) : 123 kcal./mole
(15) and D(CF-F) • 120 kcal./mole (16) have been estab-
lished from electron impact and predissociation etudies,
respectively.
uncertainty exista in the evaluation of the
final dissociation energy, D(C-F), although a value of llO
- 5 -
kcal./mole is consistent with available data {16).
These values require that D(CF 2 -F) be 110 koal./mole,
determined as the difference between the sum of the
ether three bond dissociation energies and the heat of
atomization of the molecule.
Previous unoertainties
(17) in the stepwise dissociation energies in hydrogen
sulphide have been resolved and the values of D(HS-H) =
92 koal./mole (18) and D(H-S)
= 81.4
kcal./mole {19)
have been determined from thermoohemical and spectroscopie data.
In general, bond dissociation energies have not
been used as frequently as bond energy terms in estimating thermochemical quantities sinoe they are both
diffioult to measure and sensitive to their partioular
molecular environment.
- 6 -
Table I
Bond Dissociation Energies and Bopd Energy Terms
E(A-B)
References
Molecule
lh
ll2
HgC12
81
25
53
(20,21)
HgBr2
72
16
44
(20,21)
Hgi2
61
8
35
(20,22)
102
87
124
80
99
{23)
123
110
120
110
116
(15,16)
TiC1 4
80
101
106
124
103
{20, 24)
SH2
92
81.4
87
(18,19)
OH2
117.5
101.5
110
NH3
103
92
85
93
91
95
119
102
CH
CF
4
4
A1C13
ll3
(2)
(25 ,26)
(20)
- 7 -
Factors Affecting Bond Dissociation Energies
The bond dissociation energy is equivalent to the
difference between the heats of formation of the fragments
and the heat of formation of the parent molecule or
radical,
(3)
and is thus influenced directly by factors which affect
any one of these terms.
In order to avoid the ambiguities which result in the
evaluation of molecular resonance energies from bond energy
terms, Szwarc (27) defined the experimental resonance of a
radical in terms of dissociation energies.
resonance energy,
~'
The experimental
of the radical formed upon dissocia-
tion of the molecule, RH, was defined as the relative
lowering of the bond dissociation energy with respect to
the value of D(OH -H) when a methyl radical is formed by
3
the dissociation of methane
(4)
Table II shows the lowering of the OH bond dissociation
energies in a series of hydrocarbons.
- 8 -
Table II
C-H Bond Dissociation Energies in Hydrocarbons
Hydrocarbon
Bond (R-H)
Dissociation
Energy
AHr(R)
Rr
CH -H
3
102
32
0
(23)
C2Hs-H
97
25
5
(23)
n-c 3a,-H
99
22
3
(28)
s-C 3Ry-H
94
17
8
(28)
t-C 4H9-H
89
5
13
(5)
C 6 H C~-H
80
40
22
(29)
77
29
25
(30)
c 6H -H
5
102
70
(5)
H2C:CH-H
105
65
(31)
5
2
H C:CHC~-H
References
- 9 -
The resonance energies of the alkyl radicals were
attributed to
~
-hyperconjugation (32); the greater the
number of appropriate valence bond structures, the
greater will be the resonance energy of the radical.
analogy,
By
n -conjugation accounts for the large resonance
energies associated with the allyl and benzyl radicals
(33).
The parent molecule might be subject either to
stabilization from affects such as 1T-conjugation and ioniecovalent resonance, or to destabilization as a result of
the release of 'reorganization' energy associated with
reduction in steric strain or change in valence state.
The reduction in length of the central
c-c
bond in
0
butadiene to 1.47 A from the 'normal' C-C bond length of
0
1.54 A has been attributed to lT-conjugation (34), although
the extent of this affect has recently been questioned by
Dewar and Schmeising (13).
central
c-c
A parallel strengthening of the
bond is reflected in the bond dissociation
energy, D(C 2H3-c 2H3 ) = 104 kcal./mole, calculated by the
combination of AHf(C2H 3 ) = 65 kcal./mole (31) with the
appropriate thermochemical data.
about 20 kcal./mole than the
c-c
This value·is larger by
bond in ethane.
A calculation of the central C-C bond dissociation
energy in diphenyl from the beats of formation of the
phenyl radical (35) and diphenyl (36) gives the value
- 10 -
D(C6H5-c 6H5 ) = 99.5 kcal./mole. This strengthening of
the c-c bond by about 16 kcal./mole, with respect ta the
c-c
bond in ethane, can be attributed to TT-conjugation,
the total effect of which bas been reduced due to steric
interaction between four of the hydrogena which prevents
complete planarity in the molecule (37).
Combination of the heats of formation of the benzyl
radical (29) and dibenzyl leads ta a value of 49.5 kcal./
mole for the dissociation energy of the central C-C bond
in dibenzyl.
If the weakening of the C-C bond were
attributed solely ta the high resonance energy of the
benzyl radical, then the bond dissociation energy should
be lesa than that in ethane by 2 x 22 kcal./mole, giving
a value of 39 kcal./mole.
This strengthening of the bond
has been attributed ta the partial double bond character
resulting from hyperconjugation (38).
A parallel
0
shortening of the C-C bond length ta 1.48 A (39) provides
independant support for this conclusion.
Szwarc (38) proposed that the extra weakening of the
central
c-c
bond in hexaphenylethane ta the value of 11
kcal./mole was due not only ta resonance energy associated
with the tripnenylmethyl radicale produced upon dissociation, but also ta steric repulsion between the two radicale.
The activation energy required for the recombination of
- 11 -
triphenylmethyl radicale has been measured by Ziegler
(40) as 8 kcal./mole, and may be considered to give an
approximate measure of' the contribution of steric
hindrance to the bond weakening in hexaphenylethane.
The effect of the ionie character of a bond on the
bond dissociation energy was initially discussed by
Pauling (41) for diatomic molecules.
It was proposed
that the energy of a 'normal' covalent bond between
atoms A and B be represented by either the arithmetic
or geometrie mean of the experimental dissociation
energies of the bonds A-A and B-B.
The difference
between the observed bond dissociation energy and that
of a normal covalent bond between the same atoms was
defined as the ionie-covalent resonance energy of the
bond.
Rie
= D(A-B)
-
i( D(AA) + D(BB) )
(5)
Baughan, Evans and Polanyi (42) employed the valence
bond method in a consideration of ionie-covalent resonance in alkyl halides.
In Table III the dissociation
energies, given in kcal./mole, of various organic
halides are compared with the corresponding hydrocarbons.
- 12 Table III
D~ssooiation
Energies of RX Bonds
References
lm!:
CH3
102
81
68
54
(28)
C2H5
<J7
80
67
53
(28)
n-C 3H7
99
82
68
54
(43)
(CH3) 2CH
94
82.
68
53
(43)
(CH3) 3c
89
78
64
48
(43)
C6HsC~
80
68
51
39
(43)
This relative strengthening of the RX bonds in the
series of alkyl halides has been attributed to the increase
in ionie-covalent resonance energy in the molecule, and
thus, to the increased stability of the R+x- ion pair.
The
resonance stabilization of the carbonium ion, R+, increases
with the number of increasing valence bond structures and,
therefore, with the number of carbons in the alkyl radical.
There are thus two opposing factors which influence bond
- 13 dissociation energies in alkyl halides.
The radical
resonance energy, Rr, tends to deorease the bond energy,
whereas the ionie-covalent resonance energy of the molecule, Rio' tends to strengthen the bond.
The difference,
ARia' between the ionie-covalent resonance energies of
the molecules RX and CH3X is defined by
(6)
ARia • D(R-X) - D(CH3-X) + Rr
Figure 1, reproduoed from Baughan, Evans and Polanyi
(42), shows the potential energy of the configuration R-X
as a funotion of the interatomio distance.
The full
ourves Hii' Hoc' and E, represent the energies of the ionie,
covalent, and ground states, respectively.
The curve for
the ionie state refera to the R+x- ion pair since for
halides the curve representing the n-x+ ion pair lies
sufficiently above the others to render contribution to
the energy of the molecule negligible.
The broken curves,
Hii*, H00 *, and E* refer to configurations in whioh either
the io.n R+ or the radical Ris resonance stabilized.
The
value for D(CH3-x) is greater than that expected for a
purely covalent speoies by an amount, Rie' equivalent to
the ionie-covalent resonance energy in the CH x molecule.
3
If the radical R in the molecule RX is resonance stabilized,
then the aseymptotic energy value will be lower than that
for CH3X by the radical resonance energy Rr•
However, if
- 14 -
Figure 1
Graphie representation of bond dissociation
energies a!]d ionie-covalent resonance.
'
'
\
\
\
D(R-X)
\\~J~''-----------·-A~~~tc_:____l_
INTERNUCLEAR DISTANCE .
•.
~ :'u ·
'llo;*
· ,.,;, .,,,.:,,,_ •.. ·
~-
(
:. , . . . . . . . . .
u."'~q,,,n:.ntA<.t-eÎuœ.i;. .,~~~!i'*"'•Alillidliii'Jiii1Aïu<e"'"'w.lii"":+I«m,!'!!!f<u."- N
_;_ill'•...,..u ...
...
- 15 the ionie-covalent resonance is more effective in RX
than in
CH ~,
3
then the ground state for RX will be
lower than that for CH:;X by an a.mount, Ll.Ric' equivalent
to the difference in resonance energies in the two molecules.
The Rie values of 9, 5, and 7 kcal./mole as
calculated by use of the equation (6) for benzyl chloride,
bromide and iodide, respectively, demonstra.te the effect
of ionie-covalent resonance on the bond dissociation
energy.
The
relationsh~p
between the bond dissociation
energies of the metal alkyls and the valence state excitation energy of the metal atom has been discussed by
Skinner (44).
If the dissociation energies for the process
.
MR + R
M+ R
are represented as D1 and D2 , res pee ti vely, then reference
to Table IV shows that D1 is greater than n2 , and that the
differences
n1 -D 2
are not constant.
- 16 -
Table IV
DissociAtion Energies of the Metal Alkyle
lll-~
References
8
43
(45)
42
8
34
(46)
Hg (i-c3a, )2
27
15
12
(47)
Cd(CH3 ) 2
45
22
23
(48)
Zn (CH )
3 2
47
35
12
(49)
Molecule
Dl
~
Hg(CH ) 2
3
51
Hg(C2H5)2
The bond in the molecule MR, where M is metal, is beat
described by a wave function consisting of'terms for
covalent, ionie, and non-bonding structures.
The valence state energy of the atom, M, in a
.
li
covalent bond M -R is dependent upon the nature of the
hybrid orbitals, h 1 and h 2 , formed by the two valence
electrons. ~ese orbitals are not necessarily equivalent
hybrida, although the bonding orbital h 1 will assume that
admixture of s- and p-character which resulta in the
- 17 strongest bond relative to the excitation energy required.
However, due to the non-bonding term, the 'effective'
valence state energy will be somewhat less than that
required to promote the atom M from the zero to divalent
1
state.
The!bond dissociation energy,
n2 ,
is therefore
lowered by an amount comparable to this transition.
The
'effective' valence state energy of the atom M in MR2 will
be different from that in MR not only because the orbitais
are then equivalent sp hybrida but also because the contributions from ionie and no-bond structures will differ.
This may partly explain the variation in the differences
(Dl-D2) •
Pilcher and Skinner (50) have calculated the valence
state excitation energies for the compounds TiC1 ,· TiC1 ,
4
3
TiC1 2 , and TiCl by determining in each instance that
particular hybrid obtained by mixing d3s and sp3 which
resulta in the maximum orbital strength.
Bath the valence
state excitation energies and the ionization potentials of
the titanium atom were found to deorease in the sense
>TiC1 2 ) TiCl. The deoreasing ionization
3
potential signifies an inoreasing ionie character along the
TiC14 > TiC1
series, as indicated by the deorease of p-character from
9/8
8/27
p
in TiCl4 to p
in TiCl in the isovalent configuration.
Bath the diminishing valence state energy and the increasing
- 18 -
ionie-covalent resonance tend to increase the bond
dissociation energies in the same sense D1 < D2 <D3 (D
4
as observed experimentally (24).
Methode of
Me~suring
Bond Dissociation Energies
Bond dissociation energies may be estimated directly
by measuring the amount of energy involved in either the
bond fission or bond formation processes.
Measurement of the heat of recombination of atoms or
radicals provides the only feasible method of determining
the energy liberated in bond formation.
Bichowsky and
Copeland (51) obtained a value of D(H-H)
= 105
kcal./mole
from a measurement of the hea.t of recombina.tion of hydrogen
atoms by direct calorimetry.
The atoms were initially
generated by electrical discharge and subsequently recombined
catalytically in a platinum calorimeter.
was estimated by the effusion method.
Their concentration
However, an extension
of the method to the recombination of oxygen atoms resulted
in high values for the 0-0 bond dissociation energy (52),
perhaps due to the participation of metastable oxygen atoms
in the recombination reaction.
Studies of bond fission have proved to be more profitable in the evaluation of bond dissociation energies, the
methods employed being classified according to the mode in
- 19 which the energy of dissociation is supplied.
P.h.oto-
chemical msthods include those in which energy is
supplied in the form of radiation, such as absorption
spectra and predissociation Pbenomena.
Electron impact
studies utilize the kinetic energy of a beam of
electrons to dissociate and ionize molecules.
Thermal
and pyrolytic methode may be divided into equilibrium
and kinetic studies.
Whereas, the former allows the
calculation of the heat of dissociation from data based
on the equilibrium between the undissociated molecules
and the fragments derived from the bond fission, the
latter leads to the determination of the activation
energy for
~e
dissociation process based on the kinetics
1
of the bond!breaking reaction.
Spectroscopie Methods
The sp,ctroscopic
determi~ation
of the dissociation
energy of a'molecule requires a knowledge both of the
energy of a dissociation limit above the ground state and
of the state of excitation of the products of dissociation
at this limit.
When a band convergence with its adjoining continuum
is observed lin absorption, the dissociation limit of the
upper state can be accurately evaluated.
In the case of
- 20 -
iodine, it can be shawn (53) that at the convergence
limit dissociation takes place into one normal atom
2
2
in the P3; 2 state and one in the metastable P
112
state. Since dissociation energies refer to a~oms in
their grounà state, D(I2 ) will be lees than the
dissociation limit of the upper state by an amount
equivalent to the energy of excitation of the
state.
2
P1; 2
Similar spectra have been used to determine
dissociation energies of bromine (54) and oxygen (55).
For the numerous cases in which no band convergence
is observed, Birge and Sponer (56) proposed that an
approximate value for the dissociation limit can be
obtained by an extrapolation from observed bands.
The
dissociation energy in a given electronic state is equal
to the sum of all the vibrational quanta
D (A-B) =
L
v
AG(v + 1/2)
and can therefore be approximated by the area under the
h.G · curve in a plot of the vibrational quanta .â G against
the vibrational quantum number v.
extrapolation of the
~G
Although a linear
curve to eut the v axis gives
values that represent an upper limit for the dissociation
energies of non-polar molecules, more ac cura te extra-polation methods (57) can be employeà if sufficient vibra-
- 21 -
tional quanta are obeerved to detect the curvature of
the
~G
curve.
A Birge-Sponer extrapolation gave a
value of 81.4 kcal./mole (19) for the dissociation
.energy of the hydrosulphide radical, in reaeonable
agreement with that evaluated thermochemically.
When a continuoue spectrum is observed without the
accompanying band structure, the long wavelength limit
of the continuum gives an upper limiting value for the
dissociation limit under consideration.
The most accurate
value for hydrogen was determined in this way (58).
The diffuseness of certain band spectra and the
breaking off in the rotational structure of some emission
bands can often be attributed to predissociation.
This
effect is due to a radiationless transition from a discrete electronic state to a continuous state at the same
energy, with a resultant spontaneous decomposition of the
molecule.
The degree of diffuseness resulting from pre-
dissociation is determined by the Kronig selection rules
and the Franck-Condon principle.
The ons et of diffuse-
ness gives an upper limiting value for the dissociation
limit of the state causing the predissociation.
This
method was used to determine dissociation energies for
nitrogen (59), sulphur monoxide (60), and sulphur (25).
- 22 When a continuous or diffuse absorption spectrum
corresponds to dissociation into a normal and excited
atom, then the absorption process may be followed by
the emission of a corresponding atomic line in fluorescence.
The long wavelength limit for the appearance
potential of the atomic fluorescence gives an upper
limiting value for the dissociation limit.
Terenin (61)
has applied this method to the alkali halides.
The longest wavelength of incident light that will
cause pnotodissociation gives an upper value for the
dissociation limit.
In this way, Flory and Johnston (62)
attributed the photodissociation observed when irradiating
0
.
nitric oxide with the mercury line 1832 A to predissociation and thus established an upper limit of 6.77 ev. for
D(N-0).
Ambiguities that e:xist in dissociation energies
derived from spectr0scopic data arise chiefly from uncertainties in the state of excitation of the products.
However,
in cases where the Wigner-Witmer correlation rules fail to
differen-tiate between several sets of allowed atomic states
resulting from a particular molecular state, it is often
possible to select the true atomic state if an approximate
value for the dissociation energy is known from thermochemical evidence.
An exact value of the dissociation energy
- 23 can then be determined.
For example, from the observed
predissociations, three values for D(S ), 76.1, 83.0,
2
and 101.5 kcal./mole, are possible depending on whether
it is assumed that bath the atoms are in the normal 3P
state or one in the 3P state and one in the excited 1D
state.
The approximate value of 97 ± 5 kcal./mole
recently obtained by Colin et al. (63) from mass spectrometric measurements on the vapour pnase above calcium
sulphide supports the highest value for D(S 2 }.
Electron Impact Methode
The iDjteraction of mono-energetic electrons with
molecules often resulta in molecular dissociation and ionization.
The fragment ions produced are characterized by
their maas-charge ratios in a mass spectrometer.
The
minimum electron energy at which a fragment ion is just
produced, referred to as the appearance potential, is
measured by the accelerating potential of the electrons at
which the appropriate ion current first appears.
For the·
process in which a positive ion and a neutral fragment are
generated
R1~ + e ---+- R1 + + ~ + 2e
the appearance potential, A(R1 +), of the ion R1 + is related
to the dissociation energy and the ionization potential,
- 24 I(Rl), of radical R1 by the equation
. +
ACl\1 ) = D(R1 - R ) + I (R1 )
2
provided t~t no excess energy or activation energy is
i~
involved
In
the process.
gen~ral,
electron impact experimenta on poly-
atomic molecules have been interpreted on the assumption
that
radica~s
and radical ions produced are in their
respective ground states.
For example, the cyanide
radical formed in electron impact etudies on cyanogen
was considered to be in the excited (A21T i) state.
On
this basis, Stevenson (64) rejected the value D(NC-CN)
6.90 ev. in favour of the lower value 4.64 ev.
1
=
However,
a subsequent study of the reaction of grapnite with
nitrogen at·2500°K by Berkowitz (65) established a value
of D(NC-CN)
= 6.20 ev., which correlates with a cyanide
radical in the ground state in electron impact.
If R1 is an atom, its ionization potential is known
from spectroscopie measurements. However, when R is a
1
complex free radical, the ionization potential must be
determined by an electron impact method.
Stevenson
(~6)
Hipple and
combined the directly measured ionization
potentials of the methyl and ethyl radicals with the
appearance potentials of these ions from methane and ethane
to obtain values of 4.44 and 4.23 ev., for D(CH3-H) and
- 25 D (C 2Hs-H), :respectively.
These values correspond to
heats of formation of 32 and 25 kcal./mole, respectively,
for the
i
me~hyl
and ethyl radicals.
Since ldirect measurement of the ionization potentials
of free raqicals is difficult experimentally, it is often
easier to dbtain bond dissociation energies indirectly by
calculation from appearance potentials.
Stevenson (67)
and Hipple and Stevenson (66) have shown that, by measuring
the appeara~ce potential of R + from R R and R H, in
1 2
1
1
conjunction with appropriate thermochemical data, D(R2 -H)
1
can be calciulated.
Addition of A2 and
R2H
1
3 and subtraction of A1 gives
~ R +H
D(R2 -H)
2
~H
Using this method Stevenson measured A(0 2HS+) = 15.2 ev.
from ethane and 14.5 ev. from propane and, in combination
re~uired
heats of formation, evaluated D(OH -H)
3
4.38 ev., in excellent agreement with the value measured
with the
directly.
Electron impact studies on thiols, sulpnides and di-
=
- 26 sulPhides have provided much of the available information
on S-H,
c-s,
and S-S bond dissociation energies and the
heats of formation of sulphur organic radicale.
and
Lumpki~
Franklin
(68) studied the process
~H+e
R+ + SH + 2e
for the ethyl, n-propyl, and t-butyl mercaptans, deriving
a value of 38 kcal./mole for the heat of formation of the
hydrosulphide radical, based on measurements of th~ c2H +,
5
+
+
n-c 3H? , and t-c H ions from the corresponding mercaptans.
4 9
In the light of more recent information, Mackle (18)
reinterpretbd the data for the ethyl mercaptan system and
arrived at a va~ue of AHf(SH)
= 35.2
± 4.5 kcal./mole.
From a similar study of the process
SH+ + H + 2e
Palmer and Lessing (69) derived a value of
~Hf(SH)
:
33.7 + 3 kcal./mole based on direct measurement of the
ionization and appearance potentials of the hydrosulPhide
ion.
Both values agree, within the limita of experimental
error, with the value of 34.9
± 4.5 kcal./mole derived by
Sehon and Darwent from the study (70) of the pyrolysis of
benzyl mercaptan.
Palmer and Lessing (69) also determined
the heat of formation of the methyl sulPhide radical
indirectly from measurements of the appearance potentials
1
of the radical ion formed from methyl sulPhide and dimethyl
- 27 disulphide ·
.
!
CH3sqH3 + e
CH3SSCH + e
3
The value o1f .llHf(SCH3 )
the equation
CH + SCH + + 2e
3
3
CH s + SCH + + 2e
3
3
= 31.8
kcal./mole evaluated from
is in agreement with the value 30.5 ± 5 kcal./mole based
on the pyro}ysis of benzyl methyl sulphide by Braye et al.
(71).
Although improved techniques for producing monoi
energetic electrons have simplified the -problem
o:f
deter-
mining appearance potentials, the interpretation of these
appearance potentials in terms of bond energies requires
an exact knqwledge of the fragmentation process.
The va\ue for the ionization potential of the benzyl
radical as d~termined directly is at variance with the
value determined from appearance potentials.
Schissler and
Stevenson (72) derived a value of 8.51 ev. for the ionization potential of the
potentials
e~aluated
ethylbenzene~
c7H7
radical based on the appearance
:from the mass spectra of toluene,
and dibenzyl.
However, Lessing and his co-
workers (73) obtained the lower value of 7.76 ev. from
- 28 -
direct measurement of the ionization potential of
radicale produced from toluene.
c7H?
Combined with the
appearance iPOtential from toluene these ionization
values of D(C 6H5CH2-H) = 77 and 95 kcal./
mole, the former being in agreement with Szwarc's value
potentials
~ive
77.5 kcal./mole (74).
Howeve~, the appearance potential of the c7H?t ion
depends upo~ the structure of the parent ion before
fragmentation.
Rylander, Meyerson and Grubb (75) proposed
i
the rapid rearrangement of the parent ion from toluene to
form the symmetrical cycloheptatriene molecule-ion with
subsequent decomposition to yield a stable tropylium ion
plus a hydrogen atom.
The higher value for the C-H bond
dissociation energy of toluene as measured directly could
be explained if the appearance potential of the tropylium
ion were higher than that of the benzyl ion.
Thermal Methpds
A.
Equilibr!um Methode
i
Bond dissociation energies, as determined by the equilibrium method, are based on the measurement of the equilibrium
constant for the dissociation reaction
1
AB
...
.
A+ B
- 29 The equili~rium constant can be evaluated by either
the second or third law of thermodynamics.
1
1
of the
equ~librium
Knowledge
constant over a range of tempera-
tures allows calculation of the heat of dissociation
'
from the van't Hoff equation
d ln K
dT
=
AR~
from which .AH~ can be ca.lculated if the relevant heat
capacity data are known.
From a knowledge of the equili-
brium constant at a single temperature and of the
appropriate'free energy functions for the reactants and
products, the heat of dissociation can be evaluated by
the third law method
0
- R ln K = .6Ho t
T
0
" 0
.6(F -Ho)
T
The free energy function for gases may be calculated
conveniently by means of statistical mechanical methods
when data
a~e
availa.ble from spectroscopie and electron-
diffraction .studies.
Third law procedures give the more
accurate res.ults since the measurement of the variation
in dissociation pressure with temperature is less accurate
than the dis~ociation pressure itself.
- 30 i
To
de~ermine
bond dissociation energies from
equilibrium studies requires a knowledge both of the
'
dissociation products and their exact concentration.
1
1
Direct measurement of the total pressure of a
gaseous system allows calculation of the partial
pressure of, each species on the assum.ption tha.t the
gases are ideal.
The most precise work of this type was
carried out by Perlman and Rollefson (76) on the dissociation equ~librium I 2 ~ 2I, in the temperature range
The value D(I2 ) = 35.556 ± 0.023 kcal./
mole thus obtained is in excellent agreement with the
723° to 1274°K.
value 35.514 ± 0.050 kcal./mole (77) derived spectroscopically.
Indirect methoda of measuring partial pressures
generally involve the mea.surement of weight-loss from a
system;
include the Langmuir and Knudsen techniques.
thes~
The Langmuir'technique (78) requires the measurement of
the rate of evaporation from a filament of the vaporizing
material, assuming there is no activation energy opposing
condensation .i
The rate of evaporation into a vacuum from
unit area of 1a heated surface is given by
.!..
2
aP(M/2~RT) ,
'
where a is
th~
accomodation coefficient, M the molecular
weight,, and pl the vapour pressure.
1
For example, this
method was applied to the evaluation of the dissociation
- 31 energies of the oxides of calcium and strontium (79).
The Knudsen method (80) involves the effusion of a gas
from a heated cell through a hole of small diameter and
infinitely small thickness into an evacuated region.
The various, molecular species in the gas phase escape
at a rate proportional to their partial pressures.
Wise
(81) appliep the method to the dissociation of fluorine
in the temperature range 500° to 650°K and evaluated
D(F 2 )
= 37.6
± 0.8 kcal./mole from second law considerations.
A difficulty in the Knudsen and Langmuir methods is
that of unequivocal identification of the vapeur species
involved.
Inghram, Chupka et al. (82) and Honig (83)
solved this problem through development of a mass spectrometer capable of analyzing the high temperature beam.
This technique yields heats of vaporization or dissociation
from a plot of IT versus 1/T, where I is the intensity of
the particular mass signal at T°K, or from third law methods
where absolute vapeur pressures can be obtained by comparison of the peak intensity of the substance under study with
that for a calibrating substance for which the vapeur
pressure is well known.
For example, the vapeur in equilibrium with graPhite
at 2500°K (82) was shown to consist of C atoms,
c2
and
molecules, the heats of vaporization of these species
i
c3
- 32 being 171, 190, and 200 koal./mo1e respectively.
An
1
important study of the reaction of graphite with nitrogen at 250d°K by Berkowitz (65) enabled evaluation of
the equilibrium constant of the gaseous reaction
C + iN ~ CN, and thus, a heat of formation of the
2
cyanide radical of 109 kcal./mole. This is equivalent
to D(CN) = 7.5 ± 0.1 ev. and is in agreement with the
recent shock wave value of 7. 60 ± 0.13 ev. (84).
Goldfinger and Jeunehomme (63) studied the
s2
~.
Colin,
2S
reaction in,equilibrium with solid calcium sulphide over
the temperature range 1750° to 2300°K and derived a value
of D(S2)
= 97
± 5 kcal./mole, which supports the spectro-
scopie value of 101.5 kcal./mole.
The dissociation energies
for a series of metallic sulphides were determined by
Berkowitz and Marquart (85).
Recently, Drowart et al. (86)
have derived dissociation energies of the gaseoua monoxides
of magnesium, calcium, and strontium and oompared their
values with those previoualy obtained using different
techniques.
To prevent undesirable aide reactions from occurring,
the dissociation equilibrium can be established in a flow
system so that although a statio concentration of reaotants
and produots is set up, the produot molecules do not remain
in the reaction zone long enough for aide reactions to occur.
- 33 The concentration of the dissociation products is
usually determined spectroscopically.
Dwyer and Olden-
1
berg (87) s1tudied the equilibrium
deduced that D(H-OH)
= 118.5
= 100.4 ± 0.9 kcal./mole.
2~0
+ 02
~
40H and
± 0.7 kcal./mole and D(O-H)
The same principles can be
applied to flame spectrophotometry in which a solution
of a salt of the element studied is sprayed in known
concentration into a flame, the temperature of which is
measured by the sodium line reversai method.
The intensity
of a characteristic line of the element is measured and
the concentration calculated.
The decrease in the concen-
tration of the element is assumed to be solely due to the
formation
of
the gaseous diatomic oxide.
Uncertainties
can result from temperature measurements and the possible
1
formation of gaseous hydroxides.
Thus, the dissociation
energies represent upper limita.
For example, the value
D (Mg-0)
= 102
kcal./mole, as obtained by Sugden and Bule-
wicz (88) from flame spectrophotometry, is about 16 kcal./
mole higher than that obtained by Knudsen maas spectrometry (86).
Detonation waves have been used to attain the high
temperatures necessary to produce appreciable dissociation
in molecules of high stability.
1
A detonation wave is a
shock wave maintained by the energy released in a chemical
- 34 -
reaction.
.The wave veloci ty depends on this energy
release and therefore on the dissociation equilibrium
at high temperatures.
K:istiakowsky, Knight and Malin
(89) used the measurement of the detonation velocity in
the cyanogen-oxygen system to determine the dissociation
energies of nitrogen and carbon monoxide.
However, only
a decision between widely separated values derived by
other methods was possible, the most suitable being D(N2 )
= 9.76 ev., D(CO) = 11.11 ev., and D(CN): 7.6 ev.
B,
Kinet~c
Methods
Measurements of reaction rates may be used to deduce
heats of reaction, since the equilibrium constant is
equal to the ratio of the rate constants of the forward
and reverse reactions
A
+
B
C + D
where
If the temperature dependance of the rate constants k 1
and k 2 can be expressed by the Arrhenius equation
ln k
=A -
E/RT
where E is the activation energy for the reaction, then
- 35 the change in internai energy is
A knowledge of the temperature dependence of the rate
constants allows the calculation of
~E.
Kistiakowsky and coworkers (90,91) applied this
method to tbe determination of D(CH3-H), using the
reactions
(1)
(2)
The occurrence of these reactions in the photobromination
of methane was indicated by the measurement of the rate
of disappearance of bromine from which the kinetic mechanism was established.
The activation energy E was
1
found to be 17.8 kcal./mole, corrections being made for
the temperature dependence of the pre-exponential factor.
T.he activation energy E2 of the reverse reaction was
estimated by analogy to be 1.5 kcal./mole, a value whioh
was recently oonfirmed by Fettis and Trotman-Dickenson
(92).
A E for the reaction between bromine atoms and
methane is therefore 16.3 kcal./mole.
This value, in
oonjunction with the value for D(H-Br), yields D(CH -H)
3
= 103
kcal./mole at 450°K or 101 koal./mole. at 0°K.
- 36 Similar studies with various compounds gave bond
dissociatio'n energies in agreement with those derived
by other methods, with the exception of fluoroform and
toluene.
A value of 109.5 kcal./mole was obtained by
Whittle and' coworkers (93) for D(CF -H), whereas
3
kinetic stuclies of the reaction CF + CH ~ CF H·+ CH
3
4
3
3
indicated a value of 102 kcal./mole. In addition, the
1
heat of formation of the trifluoromethyl radical based
on Whittle's studies leads to dissociation energies
that are incompatible with the presently.accepted values.
For example, the calculated dissociation energies
D(CF 3-Br)
= 72
and D(CF -F)
3
= 133
kcal./mole are higher
than the values of 65 kcal./mole obtained by Sehon and
Szwarc (94) from a study of the pyrolysis of CF Br and
3
123 kcal./mole derived by Reed and Snedden (15) from
electron impact studies for CF • It is obvious that
4
further investigation is required to clarify these
inconsistencies.
A value of D(C 6H cH2-H) : 89.5 kcal./mole was
5
obtained from bromination studies, in contrast with the
value of 77.5 kcal./mole deduced by Szwarc (74) from
pyrolysis studies.
Anderson and coworkers (95) found
the activation energy for the reaction
Br + C6HsCH3
_.,..
C6HsC~
+ HBr
(1)
- 37 -
to be 7 kcal./mole and the difference in activation
energies for the reactions
HBr + c 6H5cH2
Br 2 + c 6H5CH2
to be 5 kcal./mole.
---
c 6H5CH3 + Br
(2)
c 6H5cH2Br + Br
(3)
As in the analogous reaction in the
bromination of methane, the activation energy of reaction
(3) was assumed to be zero.
change
~
On this basis the energy
for reaction (1) was 2 kcal./mole, from which
the dissociation energy was calculated.
However, as
pointed out by Sehon and Szwarc (96) some activation may
be associated with reaction (3) which, unlike the analogous reaction with methyl radicals, is only exothermic
by about 4 kcal./mole.
The derived value of the
dissociation energy could therefore be higher by an
amount equivalent to E3 and should be considered as an
upper limit.
The reaction of organic iodides with hydrogen
iodide was originally reported by Ogg (97) and subsequently reinterpreted by Sullivan (98) and Benson and
0'Neal (99).
From studies of the methyl iodide-
hydrogen iodide system, Flowers and Benson (100)
determined D(CH3-I), based on the reactions
- 38 CH + I
3
2
(1)
(2)
These reactions were considered to be rate determining
steps in the kinetic mechanism as estab1ished from
measurements of the rate of formation of iodine.
The
activation energy for reaction (1) was found to be 20.5
kca1./mole and that of the reverse reaction (2) was
estimated to be 0.6 kcal./mole from collision frequencies.
This led to a value of 19.9 kca1./mole for the energy
change AE, which after correction to 298°K and together
with the value for D(I2 ), yielded D(CH3-I) = 55.6 ± 1 kcal.
/mole. This agrees favourably with the value 54 ± 1.5
kcal./mole estimated from thermochemical data.
Similar
studies have yielded bond dissociation energies of 52.0,
53.5, and 40.6 kcal./mole for ethyl iodide (98), trifluoromethyl iodide and benzy1 iodide (101·), respectively.
The direct kinetic determination of dissociation
energies depends on the assumption that for the unimolecular dissociation reaction
the activation energy for the reverse reaction is zero.
The activation energy of the forward reaction is then
- 39 -
equal to the energy change AE of the reaction.
In
general, dissociation energies based on this assumption
are consistent with those derived by other methods.
Szwarc (1) derived the expression
ku =
.KI (1-e
-h 'Y0 /kT
-D/kT
) e
h
for the rate constant of a unimolecu1ar reaction in
terms of the absolute reaction rate theory.
extreme cases were considered:
Two
(1) h)l0 (( kT, and (2)
h>b)) kT, where )10 is the fundamental vibration frequency
of the bond.
For case (1) which applies to high
temperatures or weak bonds, the rate constant reduces to
where the experimental activation energy E of the unimo1ecu1ar dissociation reaction is equal to the dissociation
energy, D.
The pre-exponential factor corresponds with
the fundamental vibration frequency of the bond, which is
12
13
-1
about 10
- 10
sec.
(77).
For case (2) the rate constant becomes
ku=
U e
h
-D/kT
- 40 where the experimental activation energy E
=D
t
RT
and the pre-exponential factor is approximately lol3
sec.
-1
at T • 500°K.
Thua, this theoretical treatment
predicts values of the order of 1012 - 101 3 sec.- 1
for the pre-exponential factors of unimolecular decomposition reactions involving the rupture of one single
bond.
The direct measurement of the rate of bond fission
is often complicated by secondary reactions initiated
by the radicals produced in the primary process..
In an
ideal system these radicals are removed irreversibly
before reaction with the parent molecules occurs.
In studies of the pyrolysis of a series of organic
iodides, Butler and Polanyi (102) sought to reduce
secondary reactions through the use of a flow system in
which the reaction time was short and the total decomposition small.
A simple reaction mechanism was
proposed in which the cleavage of the C-I bond was
considered rate determining.
R-I
__,....
R + I
It was assumed that iodine atoms combine to give
molecular iodine rather than recombining with radicals,
the rate of iodine formation providing a measure of the
- 41 initial rate of bond fission.
For many iodides, the
activation energies calculated with assumed preexponential factors of 101 3 sec.-l yielded bond
dissociation energies in approximate agreement with
those obtained by other methods (103).
However, the mechanism of the pyrolysis is not
known with any certainty.
Although Szwarc (1)
accepted the dissociation energies derived, he criticized the mechanism on grounds that the combination
reaction of iodine atoms which is termolecular was
not sufficiently fast to prevent the bimolecular back
reaction.
Alternately, he proposed a chain reaction
involving the rapid attack of alkyl radicale on the
parent iodide with subsequent decomposition of the
iodide radical produced to give an olefin and iodine
R + RI
R' I
- - t..
~
RH + R' I
--1....
olefin + I
Benson (104) recently challenged the validity of the
basic assumption that cleavage of the C-I bond is the
rate determining step and proposed instead two
mechanisms considered to be rate limiting in the
decomposition of alkyl iodides.
- 42 -
.
RI
I + RI
HI + olefin
~
I + HI + olefin
HYdrogen iodide is maintained at a low stationary
state concentration by rapid reaction with the parent
iodide to form an alkane and molecular iodine.
Whereas
such considerations may be valid for a static system,
their applicability to a rapid flow system is questionable.
Szwarc extended the application of a fast flow
system to the pyrolysis of toluene over the temperature
range 680- 850°0.
At short contact times and low
percentage decomposition the experimental data were best
explained by the following mechanism.
C6H5CH;
H + c 6H5cH3
H + C6H5cH
3
1
CH3 + c 6~cH 3
2C6H C~
5
~
____._
.__.....
c 6H CH2 + H
5
c 6H5CH2 +~
c 6H6 + CH:?
CH4 + c 6H5 C~
..__.,_
(C6H5CH2)2
The scheme was simplified by the rapid removal of
hydrogen atoms and methyl radicals through reaction with
- 43 excess toluene and by the thermal stability of the
benzyl radical.
The initial C-H.bond fission was
considered to be rate determining, the speed of the
reaction being measured by the rate of formation of
hydrogen plus methane.
The decomposition was found
to be homogeneous and first order with a frequency
factor of 2 x 101 3 sec. -l and an activation .energy of
77.5
± 1.3
kcal./mole which Szwarc identified with
5
D(C 6 H c~-H).
However, the accuracy of this value was
questioned by Steacie and coworkers (105) who obtained
an activation energy of about 90 kcal./molè from
similar studies conducted over the higher temperature
range 860° to 945°0.
In an attempt to explain these
discordant values, Takahasi (106) reinvestigated the
pyrolysis over a temperature range sufficiently wide
to. cover both previous investigations.
The plot of
log k versus 1/T was found to be slightly curved, the
apparent activation energy increasing from about 74 to
104 kcal./mole with increasing temperature.
Estimates
of the heat of formation of the benzyl radical based on
recent thermochemical studies faveur a value of 40
kcal./mole (cf. page 47) which leads to a value of
D(C6H5CH2-H)
= 80 ± 2
kcal./mole.
One of the most successful techniques for the
±2
- 44 determination of bond dissociation energies by kinetic
methods is the toluene carrier gas technique developed
by Szwarc (1).
As evidenced by the thermal decomposi-
tion studies, toluene not only possesses a labile
hydrogen atom and would provide an efficient radical
scavenger, but also forms a thermally stable and unreactive benzyl radical.
Szwarc proposed that if a
molecule R1R2 were decomposed to form radicale R1 and
R2 in the presence of an excess of toluene, these
radicale would rapidly abstract a hydrogen atom from
toluene to form stable products, and that the benzyl
radicale subsequently formed would dimerize to dibenzyl
rather than initiate a chain reaction.
RlR2
Rl + C6HSCH3
R2 + C6HSCH3
2C 6H5CH
2
____._
~·~
R1H + c6H CH2
5
----.- R H + c H CH2
6 5
2
----.- (C6ffsCH2)2
----+-
According to this mechanism, the rate of formation of
R1R and R2H, which should equal the rate of production
of dibenzyl, provides a measure of the initial bond
fission.
This scheme was applied successfully tà the thermal
decomposition of benzyl mercaptan in the presence of an
- 45 excess of toluene (70).
The decomposition was found
to be a first order, homogeneous reaction and the
principal products, dibenzyl and hydrogen sulPhide,
were formed at equal rates.
These observations were
explained by the following series of reactions.
SH
c 61Isc~SH
c H CH2
6 5
+
C6H5cH3
2C 6H5CH2
c6H5CH2
+ H2S
+
---
SH
(C6H5C~)2
The frequency factor was 3 x 101 3 sec.- 1 , and the
observed activation energy of 53 ± 2 kcal./mole was
identified with the bond dissociation energy.
In
conjunction with the appropriate heats of formation,
this value yields a heat of formation for the hydrosulphide radical of 34.9
± 4.5 kcal./mole, in agree-
ment with the values previously derived from electron
impact etudies.
Decomposition studies in an excess of toluene are
carried out in a flow system with a short contact time
111 tn 1m ri!.e
to ·elimilutte possible secondary reactions. The radicals
produced should be sufficiently reactive to abstract a
hydrogen atom from toluene and be produced at a temperature high enough for this reaction to be rapid.
The
abstraction reaction is favoured by a high ratio of
- 46 -
toluene to reactant, preferably greater than 50:1.
If the activation energy of the dissociation is to
be identified wi th the bond dissociation energy, the
decomposition must be established as being a first
order unimolecular dissociation reaction.
The method
is limited to determining dissociation energies that
are less than that of toluene, in order to prevent
interference from the decomposition of toluene itself.
Studies have been made with a large number of
compounds including benzyl methyl sulphide (71), benzyl
methyl sulpnone (107), benzylamine (108), benzyl bromide
(109), ethyl benzene (110), and 1-butene (32).
In
general, the dissociation energies measured by the
toluene carrier technique are in agreement with those
derived by other methods within a few kcal./mole.
The toluene carrier technique is particularly
suited to compounds of the type benzyl-R, where R is a
reactive radical, since the large resonance stabilization
of the benzyl radical tends to make this the weakest bond
in the molecule.
However, in
ord~r
to evaluate the heat
of formation of radical R, an accurate value is required
for the heat of formation of the benzyl radical.
In a review of the thermochemical data available to
1958, Skinner (5) proposed a value of 40 ± 4 kcal./mole
----------'-'----~...._:..,_
__
...................
-·~
- 47 for the heat of formation of the benzyl radical.
Two new evaluations which support this proposai can
now be made in the light of the recent measurements
of the heats of formation of benzyl bromide and iodide
by Carson and coworkers (111).
A value of 40 ± 2 kcal.
/mole is derived from calculations based on the value
D(C6H5CH2-Br)
= 50.5
kcal./mole (109) as determined by
the toluene carrier technique.
A similar value of 40.5
±2 kcal./mole arises from calculations based on the
value D(C 6H5CH2-I) = 40.6 kcal./mole (101) as derived
from a kinetic study of the reaction
Carbop-Su1tihur Bond Dissociation Energies
Reyiew
Reliable information concerning the bond dissociation
energies of sulphur-containing bonds is important for the
elucidation of a number of chemical processes.
Among
these are the vulcanization of rubber (112) and protein
transformations (113) such as the permanent setting of
hair, the felting and shrinking of wool, and the gelation
of egg white and serum albumin.
These processes all
involve the cleavage or formation of H-S, S-S, or C-S
- 48 bonds.
Sulphur compounds are being used increasingly
in the manufacture of new polymers, such as the polysulphones (114) which are formed by the co-polymerization of olefins with sulphur dioxide.
The deter-
mination of the kinetic mechanism of this complex
polymerization process requires a knowledge of the
dissociation energies of the first and second
c-s
bonds in sulphqnes.
Comparatively little information is available on
the energy properties of sulphur-containing bonds.
The
only direct measurements of the dissociation energies
of carbon-sulphur bonds are thoae derived either by
Franklin and Lumpkin (68) and Palmer and Lessing (69)
from electron impact etudies or by Sehon and coworkers
(70,71,115) and Busfield and Ivin (107) using the
toluene carrier method.
In conjunction with other
relevant thermochemical data, auch studies yield values
for the heats of formation of sulphur-containing free
radicale.
In a recent review, Mackle (18) selected the
following mean values for the radical heats of formation:
âllt (SH)
: 34 • 6 ± 4 kcal. /mole
AHf(SCH3) : 30.5 ± 5 kcal./mole
LlHf(SC2~)
- 25.5 ± 3 kcal./mole
ÂHf (S02CH3) = -60 ± 2 kcal./mole
- 49 A series of C-S bond dissociation energies based on
these values is reproduced in Table V.
Uhless other-
wise stated, the uncertainty in the first two columns
±3
is ± 5 kcal./mole and in the last two columns is
kcal./mole.
The values in
pa-r-en +he ses
~Paekets
were determined
directly.
Table Y.
Primary c-s Bond Dissociation Energies
D(R-SH)
D(R-SCH~
(kcal./mole)
D(R-SC2~l
D(R-S0 2.QR~
CH3
73
72
72
62 (60.5!2)
c2~
71
70
71
64
n-c 3~
73
72
72
i-C3~
70
69
70
62
C6H5
78
77
77
48 (49. 5±2. 5)
c 6 ~cH2
(53±2)
(51. 5±2)
53
- 50 In a study of the thermal decomposition of phenyl
methyl sulpnide, Back and Sehon (115) proposed a value
of 60 kcal./mole for the dissociation energy,
D(C 6H5S-CH3), and concluded that the phenylsulpnide
radical was resonance stabilized. It was assumed that
the radical resonates between the several canonical
structures
in a manner analogous to that proposed for the benzyl
radical (116).
Together with the recently determined
heat of formation of pnenyl methyl sulphide (117), this
value of the carbon-sulphur dissociation energy yields a
heat of formation of 51 kcal./mole for the phenylsulphide
radical.
However, recent attempts to determine the carbonsulpnur bond dissociation energy of phenyl alkyl sulphides
from electron impact studies yielded higher values than
that proppsed by Back and Sehon for phenyl methyl
s ulphide.
Palmer and Lessing (69) found the appearance
potential of the phenylsulphide ion from phenyl methyl
sulphide to be 12.1
±
0.1 ev.
Combined with the ionization
- 51 potential of the phenylsulphide radical, this appearance
potential yielded a value of D{C6H5S-CH3)
mole.
=
83 kcal./
Gowenlock and coworkers (118) derived a
similarly high value of D(C6HsS-C 2HS) = 85 kcal./mole
based on the appearance potential of the phenylsulphide
ion from phenyl ethyl sulphide.
This discrepancy was attributed to the formation
of a phenylsulphide ion with considerable excess energy.
However, an alternate explanation might involve
expansion of the parent ion of the phenyl alkyl sulphide
to a seven-membered ring before fragmentation, as
proposed for the alkylbenzenes (75) and alkyl thiophenes
(119).
Dinneen and coworkers (120) were able to explain
some of the ions produced in the electron impact studies
of deuterated phenyl mercaptan by assuming ring expansion
in the parent ion.
Present Investigation
·The principal aim of the present investigation was
to establish the heat of formation of the phentlsulphide
radical
and
to evaluate the carbon-sulphur bond dissocia-
tion energies for a series of phenyl organic sulphides.
To accomplish this, the thermal decompositions of phenyl
ethyl sulphide, phenyl n-propyl sulphide, and phenyl
- 52 -
benzyl sulphide were studied using the toluene carrier
technique.
Due to the apparent resonance stabilization
of the phenylsulPhide radical, it was expected that the
weakest bond in the sulPhides would be the C-S bond
joining the alkyl or benzyl radical to the phenylsulphide radical.
It was further assumed that the
phenylsulphide radical would be formed at a sufficiently
high temperature in each of these pyrolyses to react
rapidly with toluene to form phenyl mercaptan.
In the decomposition of phenyl ethyl sulphide, it was
postulated that the ethyl radical produced might not only
react with toluene to form ethane, but also decompose to
form ethylene plus hydrogen.
Regardless of the fate of
the ethyl radical, Phenyl mercaptan and
c2
hydrocarbons
were expected to be formed in equimolar amounts, and
consequently, the rate of decomposition could be followed
by the rate of production of either product.
In the decomposition of Phenyl n-propyl sulPhide,
the propyl radical formed was expected to decompose
further to yield ethylene plus a methyl radical.
If the
methyl radicals were to react rapidly with toluene rather
than dimerize, then ethylene and Phenyl mercaptan were
expected to be formed in equimolar amounts.
The rate of
formation of either product should, therefore, have given
- 53 a measure of the rate of initial bond fission.
In the decomposition of phenyl benzyl sulphide,
it was expected that the benzyl radicals would dimerize
to form dibenzyl rather than recombine with the sulphide
radicals.
The rate of formation of either phenyl
mercaptan or dibenzyl should then have provided a measure
of the rate of decomposition.
Preliminary resulta of a study designed to determine
the sulphur-sulphur bond dissociation energy in diphenyl
disulphide are presented in Appendix A.
- 54 -
EXPERIMENTAL
Mate rials
P.henyl ethyl sulphide and phenyl benzyl sulphide
were kindly supplied by Dr. Hawkins of the Bell Telephone
Laboratories, Murray Hill, N.J., and used without further
purification.
Phenyl n-propyl sulphide was obtained from
the Wateree Chemical Co., Lugoff,
s.e.,
by distillation at low pressure.
Nitration grade toluene
and was purified
was supplied by the Gulf Petroleum Co., Pittsburgh, Pa.
The toluene was dried and distilled through an efficient
column.
studies.
The fraction boiling at 110.5°0 was used in these
Toluene-oG-d
3 was obtained from Merck, Sharp
and Dohme, Montreal, P.Q.
Appapatus
The pyrolysis etudies were conducted in a high
vacuum flow system shown in Figure 2.
The toluene was
introduced into the system from a 100 ml. round-bottom
flask, F1 , equipped with a short capillary, c1 • The
flask could be maintained at any desired temperature by
means of a thermostated water bath, B1 •
Weighed amounts of the phenyl etbyl sulphide and
phenyl n-propyl sulphide were introduced into the toluene
- 55 -
stream through a capillar.y tube,
c2 ,
from a detachable
vial, F 2 • In the case of the phenyl benzyl sulphide,
a weighed pellet of the sulpnide was sealed into a Utube in the toluene flow line.
These sections of the
apparatus are shown in Figures 3A and 3B, respectively.
In both instances, the vapour pressure of the sulpnide
was controlled by the heat provided by a removable bath
of di-butyl pnthalate, B2 •
The flow line leading to the furnace was heated
electrically with Nichrome wire to prevent condensation
of the sulpnide.
A small manometer, M, was connected
to this section of the apparatus to measure the flow
pressure during an experiment, the pressure being read
with a cathetometer to an accuracy of ± 0.001 cm.
The reaction vessel, R, shown in Figure 4 was made
entirely of silica and was joined to the apparatus with
graded silica to pyrex seals.
Tlie thermocouple well was
sufficiently long to allow measurement of the temperature
along the length of the reaction vessel.
Section B of
the reaction vessel was considered to be the effective
reaction volume.
The temperature gradient over this
section was less than four degrees and dropped sharply
across sections A and C, thus minimizing the decomposition
in these zones.
The temperature in the reaction vessel
- 56 was measured by means of a chromel-alumel thermocouple
and a Leeds and Northrup potentiometer.
The surface dependence of the reaction was examined
by using a similar reaction vessel in which section B
was packed with quartz woel as supplied by Micro Chemical
Specialties, Berkeley, Cal.
The average diameter of the
fibres used was 3 x lo-3 cm. and the average length 15
cm. The weight per fibre was about 3 x 10-4 g. and about
6.2 g. were used for the packing. The surface area of
one fibre was about 0.15 cm. 2 so that the total surface
area of the packing was about 3100 cm. 2
The radius of
the reaction vessel was 18 mm. and the total surface area
about 230 cm. 2
The surface/volume ratio of the unpacked
vessel was about 1.1 cm.
about 15.1 cm.
-1
-1
.
and that of the packed vessel
The surface/volume ratio was therefore
increased by a factor of about 14 in the packed reaction
vessel.
The reaction vessel was enclosed in a furnace, H,
which consisted of a refractory tube (Norton Co., Worcester,
Mass.) wound in two sections with Nichrome wire to allow
adjustment of the temperature gradient along the reaction
vessel.
The furnace was thoroughly insulated with rock-
woel contained in a metal drum 23 in. in diameter and 2
ft. in length.
An Inconel tube was inserted between the
- 57 -
wall of the refractory tube and the reaction vessel
to ensure equalization of temperature.
The temperature
was regulated by a Thermo Electric controller which
maintained the temperature within ± 2°C.
The sensing
element of the controller was an independant thermocouple inserted in the thermocouple well of the
furnace.
The exit flow line from the furnace, heated with
Nichrome wire, contained a length of capillary tubing,
c3 ,
which controlled the flow rate.
This led to a
series of three traps, T , T2 , and T . The first trap
1
3
was a removable U-tube of 12 mm. O.D. tubing and the
others were of conventional design.
The third trap was
connected to an analytical section of the apparatus
consisting of a gas burette containing a small trap, T4 ,
and a removable sample flask, F • Non-condensible gases
3
were removed by a high efficiency Edwards diffusion pump,
DP, through a small trap, into a series of three bulbs
connected by mercury cut-offs.
The pressure of non-
condensible gases in this collection system was measured
by use of a McLeod gauge, MG.
A Toepler pump, TP, led
to a gas burette that was connected through a mercury
eut-off valve to a small trap, T5 , and a copper oxide
combustion furnace, CF.
- 58 -
Figure 2
Apparat us
M
__~-
..__
TP
'
j
"
'i
ij
-~ _ ___..:..___..:..~___..:.._ _ _ _ _ ______;_.......;_____;:_,_;__
__:___ _ _ _ _ _ '1
- 59 -
Figure 3A
Section of apparatus for the introducti on of liquid
sulphides .
Figure 3B
Section of apparatus for the introducti on of solid
sulphides .
tToV
~REACTION
.
TOLU EN E
--
ES§ EL __ _
.
. i
~
7~
~
~
'--""'
. TOLUENE
...
_FIG3A
TO REACTION
t
VESSEL .
...
PELLET
OFSULPHIDE
,.
-FIG 38
•.
- 60 -
Figure 4
Reaction Vessel
___
........._"·-~--.
u·~
!"li/
-<'·tr&,;;,;,?"""Mk'
-·'""-·-·-·------
e
r
e
~
EOGE OF FURNACE
1
.
1
1
THERMOCOUPLE 1
_.. . . . .
WELL
..
!
~~T
••:
t
=
'\
~
t
A
_
·/ •;
"'\.
8
VOLUME CC
LENGTH CM
l
------~--·--
1
"'
OUTLET
C
?
,.
1
1
1
·,
A
68.5
24.0
. B
c
201.5
21.5
20.5
7.5.
- 61 The apparatus for gas-liquid chromatograpby was
similar ta that described by Ca.llear and Cvetanovic
(121) but redesigned ta function as a single pass unit
with a reference column.
A Gow-Mac thermal conductivity
cell, madel TR II B, with tungsten filaments, was used
as a detecter.
The signal was amplified by a Keithley
microvolt-ammeter, madel 150, and recorded by a pen
Speedomax Recorder, T,ype G.
The power input was
regulated with a Sorenson a.c. voltage regulator, madel
500s and a d.c. law power regulator built by the Applied
Chemistry Division of the National Research Council of
Canada.
The apparatus was equipped with sample and
reference columns, i in. in diameter and 16 ft. in length,
packed with Burrell high activity silica gel.
Erocedure and Analyses
In the pyrolysis etudies of pnenyl ethyl sulpnide and
phenyl n-propyl sulpnide, the samples of sulpnide were
weighed and the sample vial attached ta the appara.tus by
use of silicone grease on the standard taper joint.
In
the pyrolysis study of pnenyl benzyl sulphide, weighed
pellets of the sulphide were introduced into the inlet
U-tube and the inlet arm wa.s sealed.
In bath cases the
reactant was kept frozen with a dry-ica-acetone bath while
- 62 -
the apparatus was evacuated to 10 -5 mm. Hg .• pressure.
After weighing, the toluene flask was attached to the
apparatus through a standard taper joint.
It was
degassed by successive freezing with liquid air and
5
melting until the final pressure was about 10- mm. Hg.
The water bath was adjusted to the desired temperature
and placed around the flask.
Previous to the start of an experiment, the three
product traps, T1 , T , and T , were surrounded with brine
2
3
(-5°C), dry-ice-acetone (-78°C), and liquid air (-l88°C),
respectively.
Liquid air was also placed around the small
trap in the collection system.
Immediately prior to the
initiation of the toluene flow, the sulpnide sample was
warmed to room temperature.
The di-butyl phthalate bath
was placed around the sample container and the toluene
and sulphide were then admitted simultaneously into the
reaction system.
The toluene flow pressure was generally about one cm.,
and the partial pressure of sulphide less than one per cent
of the toluene flow pressure.
The temperature of the bath
surrounding the sulpnide was varied from 80° to 100°C with
phenyl ethyl sulphide, from 95° to 115°C with phenyl
n-propyl sulphide, and from 90° to 115°C with phenyl benzyl
sulphide.
- 63 The temperature of the reaction vessel and the
flow pressure were each recorded at frequent intervals
throughout an experiment. ·At the conclusion of the
experiment, the toluene flask and the vial of sulPhide
were weighed again.
In the case of phenyl benzyl
sulPhide, the residue remaining after pyrolysis was
determined by an amperometric titration (122).
In the study of the decomposition of phenyl ethyl
sulphide and phenyl n-propyl sulphide, the material
collected at -5°C in the first trap, T1 , was weighed
and titrated amperometrically for sulphide content with
O.lN potassium bromate-bromide solution (122).
Dibenzyl
was then determined as a difference in weight.
A series
of experimente were conducted to show that the disulphide
content of this trap was negligible.
The disulphide was
reduced quantitatively with sodium borohydride to mercaptan
and the mercaptan was subsequently titrated with O.lN
silver nitrate by use of a Radiometer pH Meter equipped
with silver and mercurous sulphate electrodes (123).
In the pyrolysis of phenyl benzyl sulphide, the
disulphide and sulphide content of the material in the
first trap were determined in each experiment.
The
disulphide was reduced to mercaptan by treatment with
zinc amalgam according to the procedure outlined by
- 64 -
Kolthoff et al. (124) and the mercaptan formed was then
titrated amperometrically with O.lN silver nitrate
solution (125).
After precipitation of any excess
silver nitrate with hydrochloric acid, the solution was
filtered and analyzed for sulphide content.
Dibenzyl
was estimated as a difference in weight.
The second trap, T2, which contained mostly toluene,
was removed and the contents were titrated for mercaptan
with O.lN silver nitrate solution by use of a rotating
platinum electrode and micrometer (125).
It was
necessary in a series of experimente ta measure the
amount of sulphide present in the toluene in arder ta
establish a mass balance.
This was accomplished by
extraction of the mercaptan from a fraction of the
toluene with a 10 per cent sodium hydroxide solution to
prevent interference in the sulphide titration.
The condensible gases trapped at -188°C were
initially distilled into the small trap, T4 , that formed
part of the gas burette. The product trap, T3 , was
maintained at -78°C during the transfer to eliminate
any contamination of the products by toluene that may
have been carried over from trap T2 • The mercury eut-off
valve was closed, trap T4 warmed and the gas measured by
a pressure-volume measurement.
Fractions of this gas
- 65 -
were then transferred to the flask, F , for analysis.
3
In the pyrolysis of phenyl ethyl sulPhide, the
condensible gases were analyzed for ethane and ethylene
content using a gas chromatograPhy column of high
activity silica gel.
In the pyrolysis of phenyl
n-propyl sulphide, the condensible gases were essentially
ethylene but contained small quantities of hydrogen
sulPhide (about 5 percent of the total).
The conden-
sible gases were distilled into the flask, F , which
3
contained 10 ml. of lN sodium hydroxide solution kept
frozen with liquid air.
After warming, the contents of
the flask were shaken to promote absorption of the
hydrogen sulPhide, acidified and titrated iodometrically
(126).
The pressure of the non-condensible gases in the
calibrated volumes was measured with the McLeod gauge.
A
sample was transferred by a Toepler pump into the gas
burette for combustion over copper oxide at 300oc.
The
initial pressure of the sample was measured, the sample
admitted to the combustion furnace through a mercury eutoff valve and the U-trap, T , surrounded with a dry-ice5
acetone mixture. Within two hours the hydrogen was
completely oxidized to water.
The residual gas, methane,
was returned to the gas burette and the pressure measured.
- 66 The composition of the non-condensible gases determined
from these measurements agreed favourably with mass
spectroscopie results.
The decompositions were studied over the temperature
range 843° to 935°A for phenyl ethyl sulphide, 832° to
951°A for phenyl n-propyl sulphide, and 759° to 880°A for
phenyl benzyl sulphide.
- 67 -
REiSUL'U?
A.
AND
DISCUSSION
Phenyl Ethvl Sulphide
Products
The main products of the decomposition, as shown
in Table VI, were mercaptan, ethane, ethylene, hydrogen,
methane and dibenzyl.
Mass spectroscopie analysis of the products in the
trap maintained at -78°C indicated that the mercaptan
was totally pnenyl mercaptan, and that no ethyl mercaptan
was present.
Large amounts of the undecomposed sulpnide
and small quantities of dibenzyl were also identified.
The mercaptan, in addition to the undecomposed sulphide,
accounted for 99.5 per cent of the sulphide pyrolyzed.
Samples of the amperometric titration curves for the
mercaptan are shown in Figure 5.
Both chemical and mass
spectroscopie analyses of the products in the trap
maintained at -5°0 indicated only trace quantities of
diphenyl disulphide ((2 percent of the mercaptan produced).
Similar trace amounts of phenyl benzyl sulphide
were detected by mass spectrometr,y.
Dibenzyl was identified
both by infrared analysis and mass spectroscopie analysis.
Dibenzyl and pnenyl ethyl sulpnide accounted for more than
- 68 -
97 per cent of the products of this trap.
Samples
of the amperometric titration curves for sulphide are
shawn in Figure 6.
The condensible gases were shawn both by gas
chromatographie and mass spectroscopie analyses to be
exclusively ethane and ethylene.
The amount of ethylene
varied from 52 to 79 per cent depending on the reaction
temperature.
A sample chromatogram is shawn in Figure 7.
flYdrogen accounted for 65 to 70 per cent of the noncondensible gases, the remainder being methane.
The ratio of ethane plus ethylene to mercaptan
varied from 0.85 to 1.0.
The ratio of hydrogen plus
methane to ethylene varied from 0.7 to 1.0, increasing
with increasing ratio of toluene to sulphide.
The number
of dibenzyl radicals produced, considered to be twice the
number of moles of dibenzyl formed, was approximately
80 per cent of the total amount of hydrogenated products
(methane, hydrogen, ethane and mercaptan).
- 69 -
Figure 5
Titration curves for mercaptans
••
•
-
.
28 ""
Cl)
w
0::
w
1
•
...
EXPT. 20
1-
<(
0
a::
-u
~~
12
8
1
•
1
•
CALIBRATION /
1
•
1.
1
•
•
4
1-
1
l.
3.5
..
•
1
.
•
•
1
1-
p.
1
•
0...
~ 16
•
1
2.4 20
1
'
l
•
4
.
1
4.5
•
5
VOLUME O.IN SILVER NITRATE SOLUTION (ML.)
- 70 -
Figure 6
Titration curves for sulphides
(.
•
14
(J)
12.
w
a:
w ro
a_
~·a
o·
'
/
EXPT. 20 • •
;'
l
0:::6
(.)
2.4
2.
.
/
./
./
'
_ CALIBRATION
•
1
1
. 1
1•
j.
/•
3
3.5
4.5
4
, VOLUME O.IN BROMATE-BROMIDE SOLUTION (ML)
....
~'--·-
...•.
'-"-
'ti"';':·· ......_1
'- 71 -
Figure 7
Separation of ethane and ethylene on oolumn of high
aotivity silioa gel by gas ohromatograpny.
4~
,ETHYLENE
~
V)
l-l
0
>
-_J
_J
~
3~
.........,
ETHANE
1-
:::r:
:x:
(!)
LLJ
2.
..
::x:
<(
LJJ
a.
1
'u'
17
_,.a
23
ELUTION TIME (MIN,)
- 72 -
Mechanism of Decomposition
The experimental results indicateà that the primary
step in the pyrolysis of phenyl ethyl sulphide was the
rupture of the C-S bond,
(1)
with the formation of the phenylsulpnide and ethyl
radicals.
The phenylsulphide radical reacted with the
toluene to form mercaptan,
(2)
whereas the ethyl radical, in addition to its reaction
with the toluene to form ethane, decomposed to give
ethylene and a hydrogen atom.
c2H5
+
c6H5CH3
C2H5
---
c2H6
+
c6H CH2
C2H4 + H
5
(3)
(4)
The decomposition step (4) proposed for the ethyl radical
is in accord with the results of Leigh and Szwarc (127)
and Bywater and Steacie (128).
Previous investigations
(74) have shown that hydrogen atoms react rapidly with
toluene in this temperature range to give hydrogen and
methane.
- 73 (5)
{6)
{7)
The dibenzyl formed was considered to have been produced
by the combination or the benzyl radicals beyond the
'hot zone' of the reaction vessel.
{8)
In an attempt to turther elucidate the reaction mechanism,
phenyl ethyl sulphide waa pyrolyzed at 627°C in a stream
of àeQt&Pate• toluene-oC-d ; isotopie purity or the sample
3
used was 98.8 per cent d3 and 1.2 per cent d2 • The
hydrogen and methane produced in the pyrolysis gave the
follow1ng isotopie analyses:
These resulta indicated that hydrogen atoms and
deuterated methyl radicals extracted hydrogen (or deuterium)
from both the ring and side chain or
toluene-~-d3,
action with the side chain being preterred.
the re-
Similar conclu-
sions were reached by Blades and Steacie (129) from a study or
- 74 the pyrolysis of toluene-oC-d 3 of low isotopie purity,
the major difference being that the extent of deuteration of the hydrogen and methane was less.
In addition,
an analysis of the benzene produced showed no deuteration, thus indicating reaction (6) to be more important
than reaction (9).
(9)
Cher (130) has shown that at 60°C methyl radicale, derived
from the photolysis of azomethane, extract hydrogen with
equal probability from the ring or side chain of toluenecC-d3.
This result, though surprising, may indicate a
difference in the relative reactivities of ring and side
chain hydrogena at 60°C and 627°C.
The mercaptan produced was found to be only 10 per
cent deuterated.
This indicated that either the phenyl-
sulphide radical reacted preferentially with the ring
hydrogena of toluene, or that an extensive exchange
reaction occurred between the deuterated thiophenol and
the other products in the toluene trap before analysis.
An attempt was made to determine the amount of deuteration in the condensible gas products, but contributions of
variously deuterated species of ethane and ethylene overlapped to such an extent as to make a confident analysis
- 75 -
not feasible.
The fact that only trace quantities of
diphenyl disulphide and phenyl benzyl sulphide were
detected indicated that reactions
2C 6H5S
c H ssc 6~
(10)
c6H5S + c 6 H 5 c~
C6:ffsSCH2c6H5
(11)
6 5
occurred to a negligible extent.
The absence of butane in the reaction products
showed that the combination reaction
(12)
did not occur.
In a recent review of the reactions of
alkyl radicals, Kerr and Trotman-Dickenson (131) have
selected a value of kd/kc
= 0.14
for the ratio of the
rate constants for disproportionation to combination
reactions.
Since this ratio appears to be temperature
independant, then reaction (13) may also be negligible.
(13)
The proposed mechanism for the pyrolysis of phenyl
ethyl sulphide requires that:
(a) one mole of
c2 hydrocarbons be produced for
every mole of phenyl mercaptan,
- 76 -
(b) one mole of hydrogen be produced for each mole
of ethylene,
(c) one-half a mole of dibenzyl be produced for
each mole of phenyl mercaptan or
c2
hydrocarbons.
The experimental evidence seems to support these
conclusions.
The fact that the amount of
c2
hydrocarbons
is generally slightly less than the amount of phenyl
mercaptan produced may indicate that some of the ethyl
radicals were removed by combination with benzyl radicals
rather than by reaction with toluene.
The extent to which dibenzyl is formed supports the
conclusion that the decomposition of phenyl ethyl sulphide
occurs via a radical split.
This eliminates the possibility
of an intramolecular rearrangement.
(14)
The first order rate constants for the decomposition
of phenyl ethyl sulphide were calculated from the rate of
formation of thiophenol, and are recorded in Table VI.
sample calculation is shown in Appendix B.
A
The effect of
changes in the contact time, partial pressures of phenyl
ethyl sulphide and toluene are shown in Tables V.II, VIII,
and IX, respectively.
The effect of the increased surface/
volume ratio in the packed reaction vessel is recorded in
- 77 -
Table X.
T.hese resulta suggest that the decomposition
is a first order homogeneous process.
As observed in
Table VIII, the k values increased slightly with
increasing toluene pressure.
This may indicate that
radical recombination reactions occurred to a minor
extent at lower toluene pressures.
However, it was not
experimentally feasible to operate at the higher pressures
necessary to obtain sufficient data to make an extrapolation of k values to infinite pressure.
The plot of log k versus 1/T is given in Figure 8,
the straight line corresponding to an activation energy of
59.5 kcal./mole and a frequency factor of 6 x 1014 sec.- 1
The temperature dependance of the ratio of ethane to
ethylene can be expressed approximately by the equation
as shown in Table XI.
Based on the rate equations for
reactions (3) and (4), the ratio of ethane to ethylene
should be
=
- 78 If
4
E ~40
kcal./mole as suggested by several
workers (128,132), then E is about 13 kcal./mole.
3
The activation energy for hydrogen abstraction reactions
involving ethyl radicals could be as much as 4 kcal./
mole higher than for those involving methyl radicals,
since D(C 2Bs-H) is about 4 kcal./mole less than D(CH -H).
3
In view of this fact, the calculated value of 13 kcal./
mole is compatible with E15 = 8.3 kcal./mole as determined
by Trotman-Dickenson and Steacie (133) for the reaction
-
(15)
and gives further support for the proposed mechanism.
Bond Energies
From the experimental results the activation energy
of the initial dissociation process
(1)
was derived as 59.5 kcal./mole.
If the recombination of
the phenylsulphide and ethyl radioals is assumed to take
place without any activation energy, as is usual for the
recombination of free radicals, then D(C 6H5s-c 2H ) may be
5
identified as 59.5 kcal./mole, with a probable error of
! 2 kcal./mole.
- 79 This is considerably lower than the value of 85
kcal./mole implicit in the electron impact studies of
Gowenlock and coworkers (118).
However, these authors
considered that the higher value might be due to a
rearrangement process, rather than to dissociation of
the c-s bond.
From the equation
in conjunction with the recently determined value of
~Hf (C6H 5sc2H5)
= 18.4 :t
0.7 kcal./mole (117) and the
) = 25.5 ± 2 kcal./mole (5), i t follows
5
that the heat of formation of the phenylsulphide radical,
value of
~Hf (C 2H
6Hf(C6H5S), is 52.4
:t
4.7 kcal./mole.
This compares
favourably wi th the value of 6Hf (C H S) = 51.0 ± 3. 7
6 5
kcal./mole calculated from the values D(C6H5S-CH ) = 60
3
±2 kcal./mole, as determined by Back and Sehon (115), and
6Hf(C6H5ScH3 )
= 23.5 ± 0.7
kcal./mole (117).
Based on the value D(C 6H5CH2 -scH 3 ) = 51.5 kcal./mole,
as determined by Braye, Sehon and Darwent (71), Mackle
(18) proposed a value of 30.5
± 5 kcal./mole for
Together wi th the appropria te values for 6Hf (C 2H5 ) and
ô. Hf CCH3SC2H5 ) (134), this value yielded D (CH 3s-c 2H ) =
5
70
±5
kcal./mole (as shown in Table V).
3
~Hf(SCH ).
- 80 -
The difference
would therefore represent the difference between the
resonance energies of the methyl sulpnide and phenylsulphide radicals.
Renee, the experimental resonance
energy of the phenylsulphide radical, Re(C 6H5S), is
probably equal or larger than 10.5 kcal./mole. This
compares favourably with the value of 12.0 kcal./mole
proposed by Back and Sehon.
- 81 -
Figure 8
Plot of log k vs. 1/T
Filled circles denote experiments done in packed vessel.
------------------~-·-·--~---····
.8
.6
.4 .
.z
1.09
1.10
1.12
1.14
1/T X 103
1.16
f.f8
- 82 -
Table VI
Products of the decomposition of phenyl ethyl sulPhide
e
e
i'1l
<1>
i'1l
i'1l
Cl>
+>
Pi
<
0
0
Cl>
i'1l
F-I
p.;M
r-1
os
M~
E-1
31
843
•726
11.74
33
843
•715
21
853
30
r-1
0
t
~
r-1
:;:$
i'1l
"Bm
0
omCl>
i'1l
~
~
Pi'll
0~
i'1l
<1>
Cl>
~..;t
0
~i'll
Cl>
'Or-1
PO
oa
os
~
\0
~
C\1
~
~..;t
~T
r-1
0
r-1
m
Ni'll
P<l>
Cl> r-i
,00
:>:.
Cf.)
a
AS
~
rS(
1
r
Pi
,0+
•
..;t
f=l~
0
!><+
(\J
~~~..;t
~r~
a
0
0
<1>
A
r-1
1
0
Cl>
i'1l
f:i!'it!:l
C\1
~
130
.092
65.7
34.3
.160
46.2
53.8
.204
.116
.79
.63
1.07
13.0
.190
11.56
384
.056
63.7
36.3
.066
47.2
52.8
.112
.164
.59 1.65
1.62
13.4
.196
.243
11.84
133
.101
64.0
36.0
.140
40.5
59.5
.194
-
1.21
6.5
.272
853
•726
11.65
102
.176
67.9
32.1
.340
39.3
60.7
.410
38
857
.604
17.84
253
.133
64.1
35.9
.215
48.8
51.2
41
857
.587
6.47
75
.112
70.2
29.8
.245
34.3
20
863
.238
11.75
131
.178
66.6
33.4
.299
29
863
.728
11.51
85
.187
68.5
31.5
34
863
.715
11.60
221
.092
67 .o
44P 863
.245
11.25
135
.140
37
867
.588
17.88
285
42
867
.600
6.48
13
871
.254
10.56
+>
e-ca
Cf.)
0
C\1
0
~
·r-i
C\1
C\1~
li1
,.!4
. 72
-
.211
.83
.59
.85
18.4
.276
.260
.152
.83
.61
1.21
18.4
.335
65.7
.285
.220
.86
.92
.70 15.0
.286
35.8
64.2
.343
.252
.87
.80
.93
9.3
.415
.388
40.0
60.0
.457
.352
.85
.88
.81
25.7
.405
33.0
.143
34.7
65.3
.186
.129
. 79
•79
.99
27.9
.456
65.7
34.3
.267
28.8
71.2
.300
.292
.89 1.13
.74
10.5
.430
.148
64.5
35.5
.250
32.9
67.1
.290
.264
.86 1.01
.88
27.9
.555
110
.104
-
-
.145
32.5
67.5
.205
.162
•71
.91
1.06
26.1
.503
228
.160
66.0
34 .o
.259
28.6
71.4
.300
.200
.87
•75
.90 15.1
.653
e
e
1
III
tf.)
r-1
Dl
Q)
m
m
Q)
ex:
+'
0
Pl
~0
r=<:~l2i
E-1
0
Q)
Dl
+'
H
~~
r-1
os
E-IS
r
•r-I
,.g
Pl
r-1
r-1
0
roS
J:l
0
s
oroQ)
J:lOl
oro
Dl
Dl
Q)
Q)
Dl
~
~
ro
t!:lm
III"<T
ror-1
J:lO
12it!:l
III
0
91
.095
80.0
20.0
C\1
~
Q)
os
os
.187
:;:::$
tf.)
~
\.0
III
III"<t-
0
0
C\1
C\1
r-1
0
~
1
III
tf.)
r-1
t>;,
NOl
J:l Q)
<l>r-1
,.co
·r-1 a
~
l
\.0
t::JIII
J:lC\1
<DO
,b +
·r-I
1=1
IIIItn
+
~
0
~+
Pl
C\1
IIII"<T
+III
s0
0
Q)
1=1
r-1
1
0
Q)
Dl
~~
~
.81
. 70
15.4
•705
.94
-
•70
16.4
.657
1.05
-
-
16.7
.675
.89
.72
1.01
ocr'
.140 .94
I=IS
C\1
C\1 III
.!<l
10
873
. 705
10.61
11
873
.280
12.90
94
.167
67.8
33.2
.342
30.0
70.0
.366
2
875
.273
11.50 182
.124
84.7
15.3
.515
-
-
.490
19
883
.240
10.93
176
.197
67.7
32.3
.314
38.0
62.0
.355
25
883
.246
5.56
96
.163
76.2
23.8
.352
25.3
74.7
.383
-
.92
-
.62
21.9
.990
28
883
. 730
11.50 125
.167
70.0
30.0
.338
34.0
66.0
.392
.288
.86
.86
. 75
50.0
.945
45P 883
.241
11.55
150
.173
66.0
34.0
.302
26.9
73.1
.325
.230
.93
. 79
.80
20.3
.950
36
889
.580
16.78
384
.148
62.5
37.5
.243
33.0
67 .o
.275
.223
.88
.89
.91
57.8 1.47
14
892
.242
11.61
410
.167
65.7
34.3
.246
29.5
70.5
.285
.86
-
.97
30.5 1.51
22
892
.243
11.06
122
.264
73.3
26.7
.509
27.0
73.0
.550
.340
.93
. 72
•71
26.7 1.39
35
898
.698
11.19
168
.180
69.9
30.1
.320
32.6
67.4
.365
.313
.88
.97
.81
73.0 1.88
18
899
.247
10.90 194
.187
71.8
28.2
.305
25.6
74.4
.350
.247
.87
.80
.83
37.0 1.87
27
904
.693
11.01
.181
67.4
32.6
.234
23.5
76.5
.374
.204
•79
.67
1.01
79.2 2.25
510
27.5
72.5
.200
.240
-
21.5 1.01
e
e
tf.l
(1)
tf.l
tf.l
(1)
<
+>
0
A
MO
J":i!IZi
8
0
(1)
tf.l
+>
~
P-1~
r-1
os
E-iS
r
!
:;:$
r-1
0
ros
§S
Otf.l
(1)
!=ltf.l
Oct!
tf.l
tf.l
tf.l
r-1
(1)
~
C\1
~
-.:;jIJ:1
~tf.l(1)
ror-1
!=lO
os
(1)
~
\.0
IJ:1
C\1
~
p:j-.:;jC\1
0
~
r-1
p:jr
~
C'lJ
Ntf.l
j:lCD
CD r-I
1
,co
ro
I=IS
IJ:1
•rf
s
~1
IJ:1Iro
o<)'
r·
•
Pt
s0
-.:;j-
f=~p:j
:${p:j-.::r
0
C\1~
i~
1=1
0
+
(1)
r-1
1
0
(1)
tf.l
0
oa
0
72.8
27.2
.274
22.2
77.8
.324
.220
.85
•79
.82
48.5 2.13
ro
Jzit!l
IJ:1
250
.177
0
~
.!.4
24
904
.310
6.10
26
904
.227
11.66
480
.264
67.3
32.7
.441
21.7
79.3
.482
.320
.92
•76
•76
40.0 2.23
46P 904
.235
11.64
235
.177
69.2
30.8
.288
25.0 75.0
.310
.226
.93
.81
.82
40.5 2.20
17
911
.240
11.28
395
.314
70.2
29.8
.486
21.0
79.0
.550
.4(]7
.88
.84
.82
51.8 3.04
23
915
.233
10.96
345
.187
68.5
31.5
.279
22.5
77.5
.330
.247
.85
.85
.87
55.8 3.50
8
916
.277
13.21
165
.118
95.3
4.7
.615
20.9
79.1
.625
.431
.98
.99
.24
57.0 3.03
16
923
.234
11.44
225
.300 73.2
26.8
.488
22.5
77.5
.550
.362
.89
•76
.80 68.0 4.85
47P 923
.228
11.35
316
.186
70.0
30.0
.270
20.7
79.3
.290
-
.93
-
.87
66.0 4.72
15
.315
11.62
242
.315
71.8
28.2
.510
20.4
79.6
•565
.90
.99
•78
79.1 6.85
935
.490
- 83 Table VII
Tqe Decomposition of P.henyl Ethy1 Su1tihide
variation of k with Contact Time
Expt.
No.
T
oA
Contact
Time
k
sec.
sec.- 1
21
853
0.243
0.272
30
853
0.726
0.276
20
863
0.238
0.415
29
863
0.728
0.405
11
873
0.280
0.657
10
873
o. 705
o. 705
19
883
0.240
1.01
28
883
0.730
0.945
18
899
0.247
1.87
35
898
0.698
1.88
26
904
0.227
2.23
27
904
0.693
2.25
- 84 -
l'able VIII
The Decomposition of P.henyl Ethyl Sulpnide
Variation of k nith the Partial Pressure of Sulpnide
Expt.
No.
T
OA
Sulphide
Pressure
k
mm.
sec.-1
Hg
31
843
0.090
0.190
33
843
0.030
0.196
29
863
0.135
0.405
34
863
0.052
0.456
11
873
0.137
0.657
2
875
0.063
0.675
22
892
0.091
1.39
14
892
0.028
1.51
8
916
0.080
3.03
23
915
0.032
3.50
- 85 -
Table IX
The Decomposition of Phenyl Ew;yl Sulpnide
variation of k with Toluene Pre§ sure
Expt.
No.
T
OA
Toluene
Pressure
k
mm.
sec. -1
Hg.
38
857
17.84
0.335
41
857
6.47
0.286
37
867
17.88
0.555
42
867
6.48
0.503
19
883
10.93
1.01
25
883
5.56
0.990
26
904
11.66
2.23
24
904
6.10
2.13
- 86 Table X
~e
Decompositio n of P.henyl Ethyl Su1tihide
Effect Qf the Packed Reaction Vesse1 on k
Expt.
No.
T
k
OA
sec.- 1
20
863
0.415
44P
863
0.430
19
883
1.01
45P
883
0.95
26
904
2.23
46P
904
2.20
16
923
4.85
47P
923
4.72
- 87 Table XI
A Comparison of Abstraction and Decompos1tion
Re~ctions
C 2 H~
C2H4
Expt.
No.
fgr
c2E5
: 3.5 x 10-4 e27000/RT(C6H5CH3)
T
Toluene
OA
moles_
litre 1
e27000/RT
C2H6
C2H4
Cale. Exper.
2.24 x 10- 4
2.20 x 10- 4
11.2 x 106
11.2 x 106
0.87
0.86
0.86
0.89
2.23 x 10-4
2.21 x lo- 4
8.9 x 10 6
8.9 x 10 6
0.69
0.68
0.68
0.65
2.19 x 10-4
2.16 x 10- 4
7.4 x 106
0.57
0.56
7.4 x 106
0.56
0.53
6.4 x 106
6.0 x 10 6
0.44
0.40
873
1.95 x 10-4
1.95 x 10- 4
0.41
0.38
11
873
2.38 x 10- 4
0.50
0.43
28
883
0.39
0.51
14
892
2.10 x 10-4
2.09 x 10- 4
6.0 x 106
5.2 x 10 6
0.33
0.42
22
892
4.5 :x: 106
4.5 x 10 6
0.31·
0.37
18
899
0.27
0.34
27
904
4.0 x 106
3.8 x 10 6
0.26
0.31
26
904
0.28
0.27
17
911
0.22
0.27
23
915
1.99 x 10- 4
1.92 x 10- 4
3.8 x 106
3.1 x 10 6
0.21
0.29
16
923
2.00 x 10- 4
3.1 x 10 6
2.7 x 10 6
0.19
0.29
15
935
1.95 x 1o-4
2.2 x 106
0.15
0.25
31
843
33
843
21
853
30
853
20
863
34
863
13
871
10
1.99 x 10- 4
1.95 :x: 10- 4
1.95 x 10- 4
2.07 :x: 10- 4
- 88 -
B.
Phenyl n-Pronvl Sulnhide
Products
The main products of the decomposition, as shawn in
Table XII, were mercaptan, ethylene, methane and dibenzyl.
Bath chemical and mass spectroscopie analyses of the
products collected in the trap maintained at -5°C
indicated only trace quantities of diphenyl disulpnide
(less than 2 percent of the mercaptan produced).
The
sulphide content was predominantly phenyl n-propyl
sulphide, although minor amounts of phenyl benzyl
sulphide (about 7 per cent) and pnenyl methyl sulpnide
(about one percent) were detected.
Dibenzyl and phenyl
n-propyl sulphide accounted for more than 96 per cent of
the products of this trap.
Mass spectroscopie analysis of the products
collected in the trap maintained at -78°C indicated that
phenyl mercaptan and phenyl n-propyl sulphide were the
major constituants.
No propyl mercaptan was present
although small quantities of phenyl methyl sulphide and
dibenzyl were detected.
The mercaptan, in addition to
the sulphides present in the two initial traps, accounted
for 98 percent of the sulphide pyrolyzed.
The condensible gases were essentially ethylene, the
principal impurities being propylene (about 2 per cent)
- 89 and hydrogen sulphide (about 4 percent).
The former
was detected by mass spectrometry and the latter, which
was initially detected both by mass spectrometrie and
gas chromatographie analyses, was subsequently determined iodometrically.
Trace amounts of ethane were also
observed.
Methane accounted for more than 95 per cent of the
non-condensible gases, the remainder being hydrogen.
In
general, the hydrogen was considered sufficiently small
to be neglected.
Both the ratio of mercaptan to ethylene, which varied
from 0.7 to 1.0, and the ratio of methane to ethylene,
which varied from 0.6 to 0.9,
ratios of toluene to sulphide.
increas~~
with the increasing
sum~+he
TheVnumber of moles of
mercaptan and methane formed is approximately twice the
number of moles of dibenzyl.
Meçhanism of Decomposition
The experimental results indicated that the primary
step in the pyrolysis of phenyl n-propyl sulphide was the
rupture of the
c-s
bond,
(1)
- 90 -
with the formation of the phenylsulpnide and n-propyl
radicals.
The phenylsulphide radical reacted with the
toluene to form mercaptan,
(2)
whereas the n-propyl radical decomposed to give ethylene
and a methyl radical
(3)
The small amounts of hydrogen and propylene formed
indicated that a second mode of decomposition of the
n-propyl radical occurred to a minor extent
(4)
These modes of decomposition of the n-propyl radical are
in accord with the studies of Eywater and Steacie (135)
and Kerr and Trotman-Dickenson (136).
The reaction was
more complex since the methyl radicals did not react
completely with toluene to form methane.
...
(5)
Szwarc and Taylor (137) have shown that at 780°C about
25 per cent of the methyl radicals produced in the
thermal decomposition of acetone were removed by combination with benzyl radicals
- 91 -
CH:3 + c 6rr crr2
5
.__.....
c 6rr5crr2cH
3
(6)
and about 1 to 2 per cent by dimerization
CH3 + CH
(7)
C2H6
3
The presence of pheny1 methyl su1phide in the products
of the trap maintained at -78°0 indicated that some of
the methy1 radica1s were a1so removed by combination
with pheny1su1phide radicals
(8)
A pyrolysis of pheny1 n-propyl sulphide at 900°K in a
stream of
toluene-~-d
3 produced methane of the fo1lowing
isotopie analysis:
Thus the attack of the metbyl radical was directed both
to the side chain and ring hydrogens, the former being
more reactive.
Similar conclusions were reached for the
reaction of deuterated methyl radicals on toluene-oe-a
3
in the pyrolysis of pheny1 ethyl sulphide.
Since on1y trace amounts of diphenyl disulphide and
phenyl benzyl sulphide were detected, it may be concluded
that reactions
_______
.........
------~---------------.....:......__
-~· -~-··
- 92 -
....
(9)
(10)
occurred only to a negligible extent.
The minor amounts of hydrogen sulpnide were
apparently produced in a side reaction, the exact nature
of which is open to speculation.
Since D(C 6H -SR) where
5
R = H, CH3, c2H5 , has a value of about 77 kcal./mole
(see Table V), it appeared improbable that hydrosulphide
or alkyl sulphide radicals could be produced in significant amounts as intermediates, from which hydrogen
sulphide could be derived.
However, it may be possible
that an isomerization similar to that occurring in the
+k ta.«.lka.nes
linear tftielkaEea (138) takes place in the phenyl n-propyl
sulphide to form benzyl ethyl sulphide.
Such an isomer
would readily decompose under the conditions of this
pyrolysis to give ethylsulphide radicals, a source of
hydrosulphide radicals.
(11)
(12)
(13)
(14)
- 93 Based on reactions (1) to (10), the simple mechanism
proposed for the pyrolysis of pnenyl n-propyl sulpnide
requires that:
(a) one mole of mercaptan be produced for each
mole of ethylene,
(b) one mole of methane be produced for each mole
of ethylene,
(c) one-half a mole of dibenzyl be produced for
each mole of methane and mercaptan.
In general, the experimental results support these
conclusions.
Deviations of the ratios from unity can be
attributed to side reactions of the methyl radicals to
form pnenyl me thyl sulpnide and ethylbenzene.
'!he amount
of dibenzyl measured together with the knowledge that
some dibenzyl has been lost due to inefficient trapping
indicates that the primary mode of decomposition of pnenyl
n-propyl sulpnide occurred via a free radical process and
did not involve an intramolecular rearrangement such as
reaction (15).
..
(15)
The first order rate constants for the decomposition
of pnenyl n-propyl sulpnide were calculated from the
formation of ethylene and are recorded in Table XII.
- 94 The effect of changes in the contact time, partial
pressures of phenyl n-propyl sulphide and toluene are
shown in Tables XIII, XIV, and XV, respectively.
The
effect of the packed reaction vessel on the first order
rate constant is indicated in Table XVI.
These results
are compatible with a first order homogeneous reaction.
The plot of log k versus 1/T, given in Figure 9,
is remarkably linear, thus indicating the validity of
this treatment of the experimental data.
The straight
line corresponds to an activation energy of 60.0 kcal./
mole and a frequency factor of 8 x 1014 sec.-l
Bond Energies
The experimental results indicate that the activation
energy of the initial dissociation process
is 60.0
~cal./mole.
If the recombination of the phenyl-
sulphide and n-propyl radicals occurs without any
H?)
activation energy, then D(C6H5s-n-c 3
is equal to 60.0
kcal./mole, with a probable error of± 2 kcal./mole.
The equation
- 95 can be employed to calculate the heat of formation of
the phenylsulpnide radical.
Although the heat of
formation of pnenyl n-propyl sulphide has not been
determined, it can be calculated from the known value
of b.Hf(c 6H5sc 2H5 ) = 18.4 kcal./mole (117) by use of
Franklin's group method (139).
5
b.Hf(C 6 H S-C3~)
- .6.Hf(C 6H sc 2H ) + .6.Hf(-CH2 -)
5
5
= 18.4 - 4.9 = 13.5 kcal./mole
In order to calculate the heat of formation of
....
the phenylsulphide radical, a reliable value is required
for the heat of formation of the n-propyl radical.
Skinner (5) ini tially proposed a value of ÂHf (n-c 3~) 22
±
3 kcal./mole based on the electron impact deter-
minations of Stevenson (140).
However, Calvert and
Sleppy (141) and Kerr and Calvert (142) recently found
34.9 and 34.5 kcal./mole as the activation energy for
the decomposition of the propyl radical.
In combination
with Brinton's (143) activation energy for the addition
of a methyl radical to ethylene, these values yield
D(n-c 3H?-H)
= 97.3
kcal./mole.
In the light of this, Pope
and Skinner (144) have proposed a new value of ÂHf(n-c 3H7 )
= 20.5 ± 2 kcal./mole which, in conjunction with
ÂHf(C6H5S-C3H?)
= 13.5
~ 1 kcal./mole, gives the heat of
- 96 -
formation of the phenylsulphide radical as
- 53.0
~
~Hf(C H S)
6 5
5 kcal./mole.
This value agrees, within experimental error, with
the previously determined values of 51 kcal./mole from
the pyrolysis of phenyl methyl sulphide and 52.4 kcal./
mole from the pyrolysis of phenyl ethyl sulphide.
As recorded in Table V, Mackle (18) calculated
D(n-c 3H?-SCH 3 ) - 72
± 5 kcal./mole,
based on the value
3H7 ) = 22 ± 3 kcal./mole. However, the actual
dissociation energy should be 1.5 kcal./mole lower in
~Hf (n-c
view of the more recent value of ~Hf (n-C H7 )
3
kcal./mole.
= 20.5 ± 2
The difference
may then be considered to represent the difference between
the resonance energies of the methyl sulphide and phenylsulphide radicals, which may be identified as a lower
limit for the resonance energy associated with the phenylsulphide radical
This value is identical to that obtained in the pyrolysis
of phenyl ethyl sulphide and 1.5 kcal./mole lower than
- 97 that obtained in the pyrolysis of phenyl methyl
sulphide.
The consistency of the values for both ÀHf(c H S)
6 5
and Re(C H S) further supports the validity of the
6 5
conclusions derived in these studies.
- 98 -
Figure .9.
Plot of log k vs. 1/T
Filled circles denote experimenta done in packed vessel.
1.0
.a
.6
.4
.2
~0
(.!)
0
..J-.2
-.4
-.6
-.a
0
1.06
1.08
1.14
1.12
1.10
1/T X 10 3
.
l
· .. ''·:.·'1
. :,_:.;i '·. ·. ~
·,
..·
'
..
'
1.16
1.18
1.20
- 99 -
Table XII
Products of the decomposition of Phenyl n-propyl sulPhide
-
e
(!}
M~
..j...)
Pl
o
<
0
8
~
~ 8~
0
<D
Q)
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p., _tlO
(!}
IJ:l
(1)
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(!}
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œ
t!:s
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<D
r-l
o
Jro
11:1
om
12.24
lOO
.209
.232
r-l
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r-l r-l
::1
J:!
o a
oa
§
+->
~ ~ ~a
o:::t
11:1
rn
0
IJ:l
C\1
. Ç\1
Pl
r-l
>.,
(!}
t'J
(!}
0 r-l
~ o
<D
Q)
.0
r-l
«1 a>
a
J:~ <D
·r-i
o
a
RS
~0 ~
1
o:::t
o:::t
C\1
C\1
11:1
~ 11:1•
rn
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.:f
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a
r-l
r
~ ~
Pl
0
o
0
<D
Q)
(!}
1=1
.230
.195
1.01
.85
.89
10.8
.112
0
.250
.204
.95
.80
.86
15.5
.175
3.7
.280
.248
-73
.60
.97
17.8
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0
.165
.110 1.14
1.07
.69
17.2
.226
95.2
4.8
.173
.127
.87
.75
.79
5.9
.237
.300
95.0
5.0
.265
.207
-93
.75
.85
22.4
.262
.351
.491
97.4
2.6
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.362
.go
.74
.93
22.0
.326
63
.218
.406
98.1
1.9
.302
.190
.76
.55
.73
27.9
.310
139
.123
.168
94.1
5.9
.150
.95
.78
-
19.9
.303
.925
12.63 183
.209
.248
99.0
1.0
.255
.103
1.04
.85
.45
31.0
.423
859
.262
10.15
130
.183
.275
96.4
3.6
.215
.204
.81
.69
1.02
11.0
.440
66P
859
.248
10.39
147
.136
.182
97.3
2.7
.150
.105
.85
.77
.73
8.9
.400
20
868
.289
12.91 108
.245
.444
94.4
5.6
.320
.182
.76
.58
.64
17.4
.669
1.01
35
832
36
839
.975
12.17
120
.210
.264
37
839
.944
12.12
43
.231
.399
33
846
.955
12.39
317
.155
.145
49
847
.252
10.70 136
.148
.209
34
848
.965
12.39
130
.214
42
854
.746
20.37
15s
43
854
5.72
44
855
.730
5.24
32
859
47
1.05
99.0 1.0
lOO
96.3
lOO
-
e
tl
m
m
Q)
c::t:
+>
0
Pl
~0
0
Q)
m
H
P-t~
r-1
~~
li
m
Q)
m
m
~m
Q)
r-1
Q)
'Or-1
os
g~
39
.264
.623
fl:l~
Ho
ri:!~
E-1
+>
21
870
.295
12.63
22
870
.298
11.48
448
.085
.110
23
875
•740
20.78
160
.240
.350
24
875
11.76
91
.266
.523
38
880
.926
12.28
281
.238
.300
60
883
.294
5.70
76
.148
.334
65P
883
.240
10.77
189
.131
7
883
.273
12.55
185
31
883
.935
12.33
41
891
.705
18
894
29
30
H
ai
.....,
fl:l~
ca(\J
(\J
0
fJ:l
~
~
96.8
P!m
«JQ)
Or-1
HO
a>S
::l!EiS
NI
H
Q)
r-1
~
,.0
•ri
ts~m
HCI.>
Q)r-1
Pl
IJ:I
+
R ca
s0
0
Q)
r-1
1
0
Q)
m
s
·~::
.220
.50
.44
.78
19.4
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.89
.77
-
18.8
.702
,.00
•ri
AS
~
fl:l~
~~
J4 ~ 0
0
(\J
R
~
,!4
3.2
.300
0
.098
97.1
2.9
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.150
.84
•71
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44.8
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96.6
3.4
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.220
.69
.53
•72
57.6
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0
.295
.220
.98
•79
.82
56.2
.882
94.0
6.0
.200
.200
.64
.47
1.15
27.9
1.13
.186
96.8
3.2
.146
.110
.81
•73
.80
21.8
1.04
.201
.340
94.1
5.9
.250
•78
.63
-
27.0
1.14
170
.238
.357
98.0
2.0
.311
.220
.89
.68
.80
67.4
1.19
19.75
615
.300
.403
97.5
2.5
.374
.275
.95
•76
.82
63.6
1.46
.276
12.72
180
.320
.514
99.0
1.0
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.230
.84
.63
.60
40.0
1.84
894
.900
12.50
130
.247
.435
90.8
9.2
.324
.200
.82
.63
.70
74.0
1.50
895
.912
12.60
382
.165
.190
97.4
2.6
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.200
1.03
.89
1.13
72.5
1.49
1.04
100
100
-
-
e
e
fJ2
fJ2
fJ2
Q)
<
+>
0
A
0
Q)
fJ2
H
P-i~
rl
Q)
r
•r-I
fJ2
fJ2
~
rorl
P=l ~
0
os
os
81
.354
.850
.490 1.22
ii
rl
'<:;J-0
Ml2:i
E-1
+>
os
E-fS
19
897
.269
12.90
28
897
.890
12.60
13
39
902
.894
12.31
205
.376
58
902
.226
10.65
235
59
902
.292
5.76
63P
902
.239
53
906
6
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ct!
tbro
::$
tf..!
Q)
PO
tf..!
C\1
(\(
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P<l>
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Q)
rl
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~~~(\(
.225
.57
.43
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42.5
2.03
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.42
-
83.5
2.03
HO
.00
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~a
AS
<I>S
a0
P=l
'<:;j-
A
1
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•r-I
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P=l
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l=l
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.,-~
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p::'<:;j-
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97.1
2.9
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95.7
4.3
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98.3
1.7
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.91
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84.5
2.07
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.328
94.5
5.5
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.192
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39.1
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122
.177
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96.5
3.5
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1.00 47.0
2.17
10.85
276
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0
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-
.94
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-
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11.35
204
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.435
92.0 8.0
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.84
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913
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12.47
440
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96.5
3.5
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•79
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1.00 56.2
3.75
50
913
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9.57
160
.240
.451
96.7
3.3
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.318
.77
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1.10
53.9
3.36
1
913
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11.71
200
.290
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0
.390
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• 65
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60.0
3.52
52
915
.368
13.25
830
.173
.228
96.5
3.5
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.93
•79
-
69.8
3.47
14
927
.267
12.00 230
.400
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96.5
3.5
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.395
.82
.58
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77.1
5.50
15
927
.242
12.00
.288
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95.1
4.9
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.234
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76.0
5.87
42
lOO
100
-
-
42.3
2.29
62.0 2.60
-
r-1
m
m
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r-1
ror-1
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56
927
.082
8.32
165
.177
.311
95.2
4.8
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-
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61
927
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5.84
156
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97.5
2.5
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•79
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78.0
5.14
62
927
.221
11.16
307
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95.3
4.7
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72.3
5.75
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927
.226
10.66
252
.189
.288
95.5
4.5
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.184
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75.4
5.67
11
935
.265
12.16
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.290
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96.8
3.2
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89.0 7.45
55
941
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7.53
488
.161
.323
93.8
6.2
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•73
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. 90
50.8
12
951
.276
12.25
159
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95.7
4.3
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+'
Pl
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p:::J!2i
L
m
~
~-
(!)
i
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m
(!)
m
ct!
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+'
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P-1~
r-1
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(1)
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os
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IJ:I
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r-1
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ct!(!)
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Or-1
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C\l
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roo
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C\JO
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0
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r-1
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33.8
5.10
8.20
- lOO -
Table XIII
tBe Decomposition of P.henyl n-Propyl Sulrihide
vari§tion of k With Contact Time
E:x:pt.
No.
T
oA
Contact
Time
k
sec.
sec.-1
49
847
0.252
0.237
34
848
0.965
0.262
47
859
0.262
0.440
32
859
0.925
0.423
7
883
0.273
1.14
31
883
0.935
1.19
18
894
0.276
1.84
29
894
0.900
1.50
58
902
0.226
2.18
39
902
0.894
2.07
56
927
0.082
5.10
61
927
0.295
5.14
- 101 Table XIV
The Decomposition of Phenyl n-Propyl Sultihide
variation of k with the Partial Pressure of Sultihide
Sulphide
Pressure
k
OA
mm. Hg.
sec.-1
36
839
0.105
0.175
37
839
0.295
0.205
33
846
0.039
0.226
34
848
0.095
0.262
20
868
0.118
0.669
21
870
0.315
0.718
30
895
0.033
1.49
29
894
0.096
1.50
6
913
0.027
3.75
1
913
0.058
3.52
14
927
0.052
5.50
15
927
0.285
5.87
Expt.
No.
T
- 102 Table XV
The Decomposition of Phenyl n-Propyl Su1phide
variation of k with To1uene Pressure
Expt.
No.
T
oA
To1uene
Pressure
k
mm.
sec.-1
Hg,
42
854
20.37
0.326
44
855
5.24
0.303
23
871
20.78
0.802
24
871
11.76
0.809
7
883
12.55
1.14
60
883
5.70
1.13
59
902
10.65
2.18
58
902
5.76
2.17
62
927
11.16
5.75
61
927
5.84
5.14
- 103 Table XVI
The Decomposition of Phenyl n-Propyl Sultihide
Effect of the Packed Reaction Vassel on k
Expt.
No.
T
k
oA
sec.-l
47
859
0.440
66P
859
0.400
7
883
1.14
65P
883
1.04
58
902
2.18
63P
902
2.29
14
927
5.50
64P
927
5.67
- 104 -
C1
Phenyl Benzyl Sultihide
Eroducts
The main products of the decomposition, as shawn
in Table XVII, were mercaptan, disulphide and dibenzyl.
Mass spectroscopie analyses of the products
collected in the trap maintained at -78°C showed that
the mercaptan was exclusively phenyl mercaptan.
Bath chemical and mass spectroscopie analyses
indicated that dibenzyl, diphenyl disulphide and phenyl
benzyl sulphide were the main products collected in the
trap maintained at -5°C.
At high reaction temperatures
trace amounts of thiobenzophenone were identified by
ultraviolet analysis (less than one per cent of the
mercaptan and disulphide formed).
Diphenyl disulphide, phenyl mercaptan and phenyl
benzyl sulphide accounted for 99 per cent of .the sulphide
pyrolyzed.
The ratio of diphenyl disulphide to mercaptan
varied from 0.8 to 0.08 over the temperature range from
759° to 880°K, decreasing with increasing temperature.
The ratio of dibenzyl to mercaptan and disulphide varied
from 0.5 to 1.0.
- 105 -
M§chanism of Decomposition
The experimental results indicated that the primary
step was the rupture of the C-S bond
(1)
with the formation of the phenylsulphide and benzyl
radicals.
The phenylsulphide radical not only reacted
with toluene to form mercaptan
(2)
but also dimerized to form diphenyl disulphide
(3)
The benzyl radicals combined to give dibenzyl
(4)
At high reaction temperatures, small quantities of thiobenzophenone were produced, along with small amounts of
non-condensible gases, presumably hydrogen.
(5)
To determine the extent of the recombination reaction (6)
(6)
- 106 two different approaches were employed.
Initially,
the recombination reaction was assumed to be negligible
and first order rate constants were calculated from the
amounts of mercaptan and disulphide formed.
A plot of
log k versus 1/T is shown in Figure 10, the straight line
corresponding to an activation energy of 37.5 kcal./mole
and a pre-exponential factor of 3 x 1010 sec.-1 The fact
that this pre-exponential factor was much lower than that
for the other C-S bond ruptures indicated that reaction
(6) was indeed not negligible.
Finally, to determine the
extent of reaction (6) a series of pyrolyses were studied
using toluene-oC-d
as the carrier gas. Thus, deuterated
3
benzyl radicals were introduced into the system. An
analysis of the pyrolyzed phenyl benzyl sulPhide, as
shown in Table XVIII, revealed considerable amounts of
the d2 component. Analysis of the dibenzyl produced gave
a measure of the ratio of benzyl radicals to deuterated
benzyl radicals in the system.
This suggested that the
recombination reaction was predominant, and that kinetic
results based on mercaptan and disulphide were misleading.
Bond Energies
Although the kinetic study does not lead to a value
for the bond energy, this may be calculated from thermo-
- 107 chemical data.
By use of Franklin's group method (139)
a value can be derived for AHf(C 6H5ScH2C6H ).
5
AHf(C 6H 5 S-C~C 6 H 5 ) = lOAHf(~CH) + 2AHf()C-) +.ôHf(-S-) +AHf(CH2 )
= 10(3.30)
+ 2(5.57) + 11.6 - 4.9
= 50.9 kcal./mole
Combination of this value with an average value of
L:!.Hf(C6H5s)
= 52.1
kcal./mole (see general discussion) and
with the currently accepted value of
6 5 2 = 40
~Hf(C H cH )
kcal./mole (5) yields a value of 41.1 kcal./mole for
D(C 6H5S-CH2C6H5 ).
- 108 -
Figure 10
Plot of log k' vs. 1/T
1.0
0.5
....
~
(.!)
oo
....J
~5
1.15
.e
' 1.20
1/T X 103
1.25
1.30
- 109 -
Table XVII
Products of the decomposition of Phenyl benzyl sulphide
e
e
(!)
ra
ro
(!)
+l
p.
0
ex:
()
(!)
m
F-i
P-t~
'0
§
r
•ri
..t;l
p.
.,.;
:>;,
o..ro
«!Q)
1îro
r-I ID
Nl'f.l
F-i
.roo
,.; s
,00
·ri
+l
Or-1
os
8S
r-lr-1
:3
E-irn
~~
r-1
r-1
0
:3.-1
~(!)
ID r-I
s
~ro
rn
+
~
p.
s
0
()
(L)
A
r-1
1
()
(!)
m
...
~~
32
759
.17
6.85
262
.026
.016
-
-
7.4
.45
21
766
.28
10.35
180
.055
.035
.031
• 35
18.7
.61
31
775
.15
6.95
270
.029
.024
-
-
16
779
.29
10.20
156
.072
.046
.090
.77
24.5
.97
15
781
.30
8.95
116
.076
.048
-
-
24.7
.95
30
798
.16
7.25
230
.047
.047
-
-
18.9
1.33
20
807
.32
10.25
127
.258
.079
.325
.96
44.3
1.95
39
823
.21
5.80
133
.074
.045
-
37.2
2.20
38
824
.15
5.30
89
.165
.068
45.5
3.90
34
824
.13
4.40
90
.070
.056
-
-
37.2
3.51
35
825
.14
5.00
322
.050
.019
-
-
34.2
2.89
37
825
.16
4.80
159
.110
.107
-
-
34.6
2.71
33
828
.14
6.60
263
.142
.063
.53
34.2
2.87
8
+l
AS
AS
rn
~
.108
9.45
~
.66
•
e
-!-)
Pt
~0
Ml2i
ct:
0
E-f
0
Q)
rn
-!-)
rn
rn
Q)
H
P-i~
Q)
§
os
r
ri
E-fS
•ri
,.g
Pt
ri ri
:3
ro
ra
-a rn
ri
~~
mo
s
RS
Q)ri
·ri
,.00
•ri
1
-!-)
Pt li!
ctiQ)
Ori
HO
·ri
ria>
:3ri
~
Nlil
~Q)
s
RS
ro
~ro
;oi
·ll:l
ro
Pt
s
0
0
Q)
R
~
ri
1
0
Q)
l1.l
-
,!.:1
23
831
.27
9.80
127
.470
.075
.530
.97
61.0
3.50
25
833
.11
7.10
237
.158
.054
.128
.60
37.8
4.13
24
835
.13
5.50
120
.220
.145
.398
1.06
40.9
4.00
46
839
.45
5.60
66
.275
.050
-
-
85.4
4.25
26
869
.10
7.40
322
.490
.046
.395
.74
67.9
10.8
28
875
.13
6.65
183
.565
.061
.474
. 76
80.0
12.2
29
880
.13
6.70
340
.332
.025
.280
. 79
80.6
12.5
- llO -
Table XVIII
Products of the decomposition of phenyl benzyl sulphide
in toluene-cC -d3
.ç::..
j-I
.ç::..
.ç::..
.ç::..
0
.ç::..
\.N
())
())
())
-.J
1\)
1\)
j-I
.ç::..
.ç::..
1..0
•j-I
\.N
•
•1\)
0'\
0
\.N
())
-.J
())
•
IJ1
1\)
jExpt
No
-.J
IT OA
.ç::..
())
j-I
1..0
. . . . .
IJ1
())
0'\
Q)
IJ1
IJ1
IJ1
IJ1
j-I
j-I
j-I
-.J
0
IJ1
j-I
j-I
1\)
1\)
.ç::..
j-I
1\)
0
0
0
0'\
\.N
\.N
1..0
j-I
-.J
j-I
\.N
-.J
\.N
. .
1\)
0
0
Q)
1\)
1\)
IJ1
\.N
0
0
•
. .
0
0
\.N
0
\.N
~
1\)
\.N
1\)
.p.
1\)
1\)
0
•
. . . .
0'\
IJ1
1\)
0
0
0
0
\.N
())
-.J
-..J
0
0
j-I
lt
sec
Tol Press
mm Hg
1\)
. . . . .
.ç::..
•
1..0
1\)
Phenyl Benzyl
Sulphide Pyrolyzed
mmoles
Mercaptan
mmoles
Disulphide
mmoles
. . . . .
0
0'\
0
0
1..0
0
j-I
j-I
j-I
\.N
.p.
IJ1
\.N
.p.
IJ1
1\)
.p.
0
. . . .
0
0
1'\)
1..0
j-I
j-I
V1
1'\)
0
0
0
-.J
IJ1
-..J
\.N
-..J
())
CD
1..0
1'\)
CD
CD
1\)
())
1'\)
j-I
j-I
1\)
1'\)
1\)
j-I
1..0
j-I
V1
1..0
. . . . .
1..0
Phenyl
Benzyl
Sulphide
Recovered
After Pyrolys is
mmoles
ldo
.
•Q)
•
V1
j-I
j-I
\.N
1'\)
1\)
1'\)
0
-.J
Q)
j-I
1'\)
•
0
•Q)
-.J
. . . . .
d2
1o d 0 Dibenzyl
l%
d2 Dibenzyl
1% a4
Dibenzyl
e
- 111 -
GENERAL DISCUSSION
Experimental Errors
The primary source of error in the determination
of the rate constants was considered to be the measurement of the temperature of the reaction vessel during
an experiment.
of
± 1°
The estimated fluctuation in temperature
resulted in a considerable error in the rate
constant.
For example, as evaluated from the grapn in
Figure 9, the rate constant for the decomposition of
phenyl n-propyl sulphide at 900°A and 902°A was 2.09
sec. -1 and 2 •.24 sec. -1 , respectively.
The error in the
rate constant resulting from an uncertainty of 2° in the
temperature is therefore about 7 per cent.
The observed
scatter in the rate constants may be explained for the
most part by these fluctuations in temperature.
For
example, the differences between values of k for the two
pairs of experimenta (42, 43) and (39, 58) done at 854°A
and 902 0 A, respectively, are both 5 per cent.
An error of
7 per cent in k would result in an error of less than 1
kcal./mole in the experimental activation energy as
calculated from the slope of the plot of log k versus 1/T.
The least favourable slope for the Arrhenius plot based on
values of k with maximum scatter at both high and low
- 112 temperature values would result in an error of less
than 2
kcal~/mole
in the experimental activation energy.
In general, a factor of ± 2 kcal./mole was assigned to
the
activat~on
energies determined for the thermal
decompositions of phenyl ethyl sulphide and phenyl
n-propyl suiphide.
Dissociation Energies and
He~ts
of Formation of Radicale
This study indicated the applicability of the toluene
carrier technique for the determination of C-S bond
dissociation energies in the phenyl alkyl sulphides.
An
essential requirement of this technique is that all radicale
produced in the initial bond rupture be both sufficiently
energetic and formed at a high enough temperature to react
rapidly with toluene, thus preventing radical recombination
reactions.
In the etudies of the decomposition of both
phenyl ethyl sulphide and phenyl n-propyl sulphide, the
reaction of the phenylsulphide radical with toluene was
found to be only moderately fast in the temperature range
·o
850°A to 950
A.
To ensure complete removal of the phenyl-
sulphide radicale from the system it was necessary to
operate at toluene/sulphide ratios greater than 100:1.
In the study of the decomposition of phenyl ethyl sulphide,
the rate of C-S bond fission was measured by the rate of
- 113 production of phenyl mercaptan rather than by the rate
of productifn of
c2
hydrocarbons, since ethyl radicals
'
were apparently also removed from the system by
combinationlwith benzyl radicals.
This was not an
entirely ideal procedure since some uncertainty existed
as to whether all phenylsulphide radicale were converted
to phenyl m'rcaptan, particularly under conditions
favouring a high percentage of decomposition.
In the
decomposition of phenyl n-propyl sulphide, the n-pro.pyl
radical decomposed completely to give ethylene plus a
methyl radiJal, the rate of production of ethylene
providing a more satisfactory measure of the rate of
C-S bond fission.
However, the consistency of the
thermochemical values derived from both studies indicated
the validity of either treatment.
The fatlure of the toluene carrier technique in the
study of the decomposition of phenyl benzyl sulphide was
attributed to the comparative slowness of the reaction
between the phenylsulphide radical and toluene at temperatures belo~ 850°A.
The fact that the recombination
rather than the abstraction process was predominant was
shown by the formation of deuterated phenyl benzyl
sulphide when benzyl radicals were produced in the system
through the use of
toluene-~-d
3 as a carrier gas.
- 114 The frequency factors of 6 x 1014 and 8 x 1014
sec.- 1 determined for the rate constants of the C-S
bond fissions in pnenyl ethyl sulphide and phenyl
n-propyl sulphide, respectively, are similar and lie
in the upper end of the 'normal' range for unimolecular reactions involving breakdown into two free
radicals (145). Rate constants with frequency factors
of 3 x 1014 and 3.3 x 1014 sec.-1 have previously been
determined for C-S bond ruptures in phenyl methyl
sulphide (115) and benzyl methyl sulphone (107).
As proposed by Szwarc (1), the experimental activation energy can be identified with the bond dissociation energy within an uncertainty of RT.
As previously
discussed, the dissociation energy is not significantly
different from the change in enthalpy at 25°C.
Therefore,
the experimental activation energies derived in this
study were employed without correction in conjunction
with the corresponding standard heats of formation at
25°0 to calculate the heat of formation of the phenylsulphide radical.
An average value of 52.1
±5
kcal./mole was
proposed for the heat of formation of the phenylsulphide
radical based on the values 52.4 and 53.0 kcal./mole
obtained in this study and the value of 51.0 kcal./mole
- 115 derived from the previous study of the decomposition
of phenyl methyl sulphide (115).
The Stability of the Ehenylsu1Phide Radical
In both studies of the phenyl alkyl sulphide
decompositions, a lower limit of 10.5 kcal./mole was
determined for the resonance energy of the phenylsulphide
radical from the expression
The resonance stabilization is further supported by the
strengthening of the c-s bond in the phenylsulphide
radical with respect to that in the methylsulphide radical.
The average value LlHf(c 6H s) = 52.1 kcal./mole together
5
with .6Hf(C 6H5 ) = 70 kcal./mole (5) and AHf(S) = 65.9
kcal./mole (146) leads to a value of D(C 6H5-s)
mole as calculated from the expression
= 84
kcal./
This value for the C-S bond dissociation energy in the
phenylsulpnide radical is 16 kcal./mole higher than that
proposed by Mackle (18) for the methylsulphide radical.
- 116 -
APPENDIX A
Pyrolysis of Dirihenyl Disulrihide
Introduction
Presently accepted values for the S-S bond dissociation energies in disulpbides have been derived from
the heats of formation of the alkylsulphide radicals and
the gas phase heats of formation of the relevant disulphides.
Sehon et al. (71) attempted to determine
D(CH3S-SCH3 ) in a kinetic study of the thermal decomposition of dimethyl disulphide.
However, the reaction was
found to proceed by a molecular mechanism rather than by
a free radical mechanism.
A
similar attempt was made in
this investigation to determine D(C 6H S-SC 6H5 ) through a
5
kinetic study of the thermal decomposition of diphenyl
d isulphide.
Experimental
The diphenyl disulphide was generously provided by
Dr. Hawkins of Bell Telephone Laboratories, New Jersey,
and was run without further purification.
Pyrolysis in a toluene flow system at 765°K, 832°K,
and 901°K indicated that the major products were mercaptan
and dibenzyl.
Small amounts of sulphide were identified
•
- 117 at the lowest reaction temperature.
0
The products collected in the trap at -5 C were
dibenzyl, disulphide and sulphide, presumably benzyl
phenyl sulPhide.
The disulphide was determined by
reduction to mercaptan with sodium borohydride followed
by subsequent titration with O.lN silver nitrate using
a platinum and mercurous sulphate electrode system (123).
The sulphide was detected by titration with 0.05M potassium
iodate in 90 per cent acetic acid using platinum and
calomel electrodes (147).
The method allowed an estimate
of small amounts of sulphide to be made in large amounts of
disulphide.
The dibenzyl was not measured.
The products
collected in the trap at -78°C were mainly mercaptan.
This was determined in the usual manner (125).
Discussion
The experimental results indicated that the primary
step was the rupture of the S-S bond to give PhenylsulPhide
radicals.
(1)
The phenylsulphide radicals reacted in part to give
mercaptan, but the reaction was slow.
(2)
•
- 118 Detection of a sulPhide not only indicated that the
combination reaction
(3)
occurred to some extent, but also that some phenylsulphide radicals may have dimerized.
(4)
Thus, any kinetics based on the amount of mercaptan
as an indication of the extent of bond breakage will be
in error.
However, if it is assumed that the bond
breakage is a first order homogeneous process with a
'normal' pre-exponential factor of 101 3 sec.- 1 , an upper
limit can be estimated for the bond energy by use of the
Arrhenius equation.
The average of the values so deriyed
is D(c 6n5ssc 6H5 ) =· 48 kcal./mole (see Table XIX).
The thermochemical values of ~Hf(c n s) = 52.1
= 58.4
D(c n s-sc 6n5 )
6 5
6 5
kcal./mole and ~Hf(c 6 n5 ssc 6 ES)
kcal./mole (148)
allow the evaluation of
since
D(C 6H5ssc 6n5 ) = 2 liHf(C 6H5S) - LlHf(c 6n ssc 6H5 )
5
• 2(52.1) - 58.4
• 45.6 kcal./mole
- 119 -
Although the experimental values are unreliable, they
are at least of the correct arder of magnitude, indicating
that the suggested mechanism is probable.
- 120 -
Table XIX
Products of the decomposition of diphenyl disulphide
e
e
Expt
No
T
0
t
A sec
To1uene
Pressure
To1uene
Su1phide
mm Hg
Disu1phide
mmo1es
Pyro1yzed Recovered
Mercaptan
~
Recovery
(SS+SH)
mmo1es
k
E
sec-1
kca1/mo1e
1
901 .272
10.70
213
0.407
0.009
o. 729
91.7
8.25
49.5
2
832 .299
11.15
216
0.432
0.045
0.455
63.0
2.52
48.0
3
765 .375
11.65
160
0.505
0.360
0.130
84.0
0.37
47.0
- 121 -
APPENDIX B
Data and Calcu1ations for Experiment 20 (at 863°A)
1.
Data
A.
P.hysical Measurements
time of experiment (t)
= 43.8
min.
= 0.482
mmoles of sulphide = 3.675
moles of toluene = 131
moles of toluene
moles of sulphide
toluene pressure (p)
= 1.175
reaction temperature (T)
total number moles gases
~.
Qhemical
cm.
= 863°A
(n) = 0.486
MeAs~ements
1.
Analysis of the contents of trap T1 :
mmoles of sulphide = 1/2 x 50/25 x 0.102 x 2.85
= 0.290
(see Figure 6)
weight of contents of trap Tl
weight of sulphide
weight of dibenzyl
mmoles of dibenzyl
2.
= 0.0910
= Q.0440
-- 0.0470
g.
g.
g.
= 0.252
Analysis of the contents of trap T2:
mmoles of mercaptan
= 0.10 x 3.43
= 0.343 (see Figure
5)
- 122 3.
Analysis of the contents of trap T :
3
mmoles of condensible gases = 0.299
ratio of peak areas
c2H4/C2H6 = 1.80
(measured from the gas chromatogram shown in Figure 7)
'fo
C2H4
'fo
C2H6
4.
= 64.2
= 35.8
Analysis of the non-condensible gases:
mmoles of non-condensible gases
combustion analysis:
P1
'fo
'fo
2.
= 0.178
= 11.1 cm.; P2 = 3.7
CH4 = P2/P1 = 33.4
H2 = P1 - P2/Pl = 66.6
Calculations
contact time (tc)
= Y*(cc.)
x p(cm.) x t(sec.)
RxTxn
= 201.5
x 1.175 x 2628
6234 x 863 x 0.486
= 0.238 sec.
*
V
= volume
of the reaction vessel
rate constant (k)
=~
tc
ln
_â_
a-x
1
ln 3.675
0.238
3.332
- 0.415 sec. -1
cm.
- 123 -
SUMMARY AND CONTRIBUTIONS TO KNOWLEDGE
The thermal decompositions of phenyl ethyl sulPhide,
phenyl n-propyl sulphide, and Phenyl benzyl sulPhide were
studied using the toluene carrier technique.
The products of the decomposition of phenyl ethyl
sulphide were phenyl mercaptan, hydrogen, methane, ethane,
ethylene and dibenzyl.
The following mechanism was
postulated to account for these products:
c 6H5sc 2H5
...
...
C2HS
C2H6 + c 6H5CH2
C2H4 + H
C2H5
H + c 6H5cH3
H + C6H5CH3
t
c 6H SH + c 6H5cH2
5
c 6H5S + c 6H5cH3
C2H5 + c 6H5cH3
c 6H5S
5
C6H c~
...
+ H2
c 6H6 + CH3
C6IfsCH3
c 6H5cH2 + CH4
2C H C~
c ~cH CH c 6H
CH3 t.
6
5
6
2
2
5
The rate constant for the first order, homogeneous
decomposition of phenyl ethyl sulphide
- 124 -
as measured by the rate of production of phenyl mercaptan
over the temperature range 843°A to 935°A was expressed
by k = 6 x lo14 e- 59 , 500/RTsec.-l This activation energy
was identified with D(c 6n s-c 2n ) and the heat of forma5
5
tion of the pbenylsulpbide radical was thus calculated
to be 52.4 kcal./mole.
The resonance energy of the phenyl-
sulphide radical was estimated to be .). 10.5 kcal./mole.
The products of the decomposition of phenyl n-propyl
sulphide were phenyl mercaptan, ethylene, methane and
dibenzyl.
The mechanism was similar to that postulated
for pbenyl ethyl sulphide with the exception that the
n-propyl radical decomposed rapidly to form ethylene· anâ
hyargg9I~.
a metk~f
rd tettl.
..
The rate constant? as measured by the rate of production
0
0
of ethylene over the temperature range 832 A to 951 A was
14 -60,000/RT
-1
expressed by k = 8 x 10 e
sec.
The activation
energy was identified with D(c 6n s-n-c 3H?), and the heat
5
of formation of the phenylsulphide radical was calculated
to be 53.0 kcal./mole.
The resonance energy was again
found to be .)- 10.5 kcal./mole.
- 125 The products of the decomposition of phenyl benzyl
sulphide were phenyl mercaptan, diprrenyl disulphide and
dibenzyl.
The following mechanism was postulated to
account for these products.
c H ScH2 C6~
6 5
c 6H S + c 6H5CH 3
5
2C 6H5S
2C6H5CH2
-
C H S + CH2C H
6 5
6 5
c H SH +
6 5
c
H CH
6 5 2
c6H ssc 6H
5
5
c 6H5CH2CH2C6H5
However, appreciable decomposition occurred at low temperatures and under these conditions, the abstraction
reaction with toluene was not sufficiently fast to prevent
radical recombina ti on.
Kinetic data based on the rate of
production of phenyl mercaptan and diphenyl disulprride
were misleading, and no dissociation energy could be
d etermined.
A similar mechanism was proposed for the decomposition
of diphenyl disulphide based on preliminary experimente.
Although mercaptan and dibenzyl were the chief products of
the decomposition, no assurance could be given that the
abstraction reaction proceeded rapidly enough in this
temperature range to prevent radical recombination.
- 126 BIBLIOGRAPHY
il, 75 (1950).
1.
Szwarc, M., Chem. Rev.
2.
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