239 22 13MB
English Pages 271
222 37 2MB Read more
An authoritative guide to the theory and practice of static and dynamic structures analysis Static and Dynamic Analysis
528 123 9MB Read more
This book presents a collection of studies on state-of-art techniques for converting biomass to chemical products by mea
199 99 34MB Read more
BOND—DISSOCIATION ENERGIES ALKYL
THESIS submitted by
STANLEY ROBERTS M.Sc.
for the degree of
in the University of Manchester
ProQuest N um ber: 13871063
All rights reserved INFORMATION TO ALL USERS The quality of this reproduction is d e p e n d e n t upon the quality of the copy subm itted. In the unlikely e v e n t that the a u thor did not send a c o m p le te m anuscript and there are missing pages, these will be noted. Also, if m aterial had to be rem oved, a n o te will ind ica te the deletion.
uest. ProQuest 13871063 Published by ProQuest LLC(2019). C opyright of the Dissertation is held by the Author. All rights reserved. This work is protected against unauthorized copying under Title 17, United States C o d e M icroform Edition © ProQuest LLC. ProQuest LLC. 789 East Eisenhower Parkway P.O. Box 1346 Ann Arbor, Ml 4 8 1 0 6 - 1346
The basis of the method of pyrolysis is examined.
The development of the flow method is reviewed, with particular reference to the 'toluene carrier* technique. In the present work, some of the principles of these previous methods have been applied to a static pyrolysis system. 2.
Previous methods of pyrolysis of some alkyl iodides are
described and the inherent difficulties are discussed. It is found that some of the complications are removed when excess toluene is added in the decompositions of ethyl iodide and sec-propyl iodide. 3.
Energy relationships involving the acetyl, formyl and
benzoyl radicals are discussed.
The decomposition of
diacetyl has been studied, but the bond-dissociation energy is too high to allow a complete investigation in this syzicc, 4.
The decompositions of di-t-butyl peroxide and diethyl peroxide have been studied.
Di-t-butyl peroxide gives
clean kinetics which allow an assessment of the (0-0) dissociation energy.
Reasons are suggested for the
greater complexity of the pyrolysis of diethyl peroxide. 5.
An account Is given of the computation (using the method
of molecular orbitals) of conjugation and hyperconjugation energies in hydrocarbon molecules and radicals.
scheme can describe the general relationships between dissociation energies with reasonable success. inconsistencies are examined in detail.
The nature of
the conjugation effects in carbonyl compounds is discussed.
My sincere thanks are due to PROFESSOR M. POLANYI F.R.S. and
PROFESSOR M.G. EVANS F.R.S., for the opportunities and encouragement to begin and continue this research, DR. M. SZWARC,
for his Invaluable help and advice during his guidance of this work,
DR. H.A. SKINNER,
for the benefit of many stimulating discussions and his generous direction of the theoretical section;
my colleagues in the department, for their help in so many ways, and the technical staff, for countless services.
This research was done during the tenure of a grant from the DEPARTMENT OF SCIENTIFIC AND INDUSTRIAL RESEARCH.
CONTENTS page INTRODUCTION............................. .
SECTION I. THE DEVELOPMENT OF METHODS OF PYROLYSIS (a) comparison with the equilibrium method (b) the basis of the method of pyrolysis... (c) static methods of pyrolysis ••••••••••• (d) flow methods of pyrolysis ••••••••••••• the mirror method ................. further developments of the flow method (l)9(ii) the pyrolysis of iodides., (ill) extension to the pyrolysis of toluene.• (iv) the •toluene-carrier*technique (e) the present (static) method
3 5 12
15 15 17 21 26
SECTION II. The Static Pyrolysis of ETHYL IODIDE and sec-PROPYL IODIDE (a) mechanisms previously suggested for these decompositions (i) ethyl iodide............ ....... (ii) sec-propyl iodide •••••••••••••••• (b) previous static pyrolyses •••••...... (c) review of the photochemical decomposition of seme alkyl iodides (1) photolysis of ethyl iodide . (ii) polymerisation of ethylene induced by alkyl iodides ......
31 33 35
THE PRESENT INVESTIOATION apparatus .... 43 analysis of products ........... 48 52 results and discussion ••••••••...... • further discussion •••••••••••••••••••• 60 independent estimates of C-I dissociation energies ••••• 64
SECTION III. BOND-DISSOCIATION ENERGIES IN DIACETYL AND OTHER CARBONYL COMPOUNDS (a) (b) (c) (d)
introduction •••••••••••••••••••••••••• the stability of the acetyl radical ... previous decompositions of diacetyl ... THE PRESENT INVESTIGATION
67 69 76
apparatus ••••••••••••••••••••••••••••• analysis of products ..... results of the pyrolysis of diacetyl •• discussion ........ (e) (1) the pyrolysis of benzyl-oothyl ketone (ii) the pyrolysis of dlbenzoyl ..... (f) general discussion ••••••••••••••.•••••
78 80 81 84 87 89 90
page SECTION IV. BOIID—DISSOCIATION ENERGIES IN ALKYL PEROXIDES IVA. The static pyrolysis of di-t-BUTYL PEROXIDE (1) introduction........ (ii) THE PRESENT INVESTIGATION experimental............... ..... (ill) results.......... ................... (iv) discussion (v) comparison vith flow experiments ... (vl) conclusion................. 117
103 105 110 114
(vii) the activation energy of thereaction, CH^ ♦ C6H5CH3 — » CH4 -*■ C6HrCH2. (a) previous estimates....... 118 (b) evidence from pyrolysis of di-t-butyl peroxide ... 122 (c) evidence from pyrolysis of other compounds ....... 127 (d) conclusions •••••••••••••••••••• 130
IVB. The static pyrolysis of DIETHYL PEROXIDE (1) introduction..... .......... •••••• (ii) preparation of diethyl peroxide ..... (Hi) experimental .................... (lv) results; products.......... ••••••• the over-all reaction ••••••••....... (v) discussion................ conclusion............................. (vi) relative stabilities of alkoxy radicals
132 134 135 136 140 150
A review of bond-dissociation energies in other peroxides (i) dipropyl peroxide .••••••••••••••• 157 (11) u^^her alkyl peroxides ••••••••••• 157 158 (ill) I acetyl peroxide ••••••••........ (iv) benzoyl peroxide ........ 161 162 (v); hydroperoxides ••••...... ••••••••
THEORETICAL DISCUSSION of bonddissociation energies (a) general formulation in terms of conjugation and hyperconjugation .. (b) types of conjugation......... (c) standard *bond-energy terms* ...... (d) empirical resonance energies...... (e) calculated resonance energies (1) hydrocarbon molecules ••••...... (2) hydrocarbon radicals........ conclusion •••••••...... (3) extension to molecules containing oxygen .... types of conjugation in carbonyl compounds (i) second-order conjugation across the C=0 bond .... (ii) third-order conjugation across the C=0 bond •••• (iii)first-order conjugation ...... carbonyl radicals •••••••••••••••••
164 165 169 173 174 177 184 186
188 190 191 193
appendix a fuller compilation of bond-dissociation energies. 195 SUMMARY
All energies are expressed in kcal.mole*1* (1 kcal. = 4184 abs. joules)
Qf (= - AHj) denotes the heat of formation of a substance, from elements In their standard states, at 25°C,
INTRODUCTION The most general problem which will emerge during this account may be stated as follows* We know that the dissociation energy of a bond between two given atoms can be very sensitive to the detailed molecular context of that bond.
energy, we shall, throughout, understand the quantity defined ^ ^ 2 ^ unambiguously as the endothermicity of the reaction* which breaks the given bond in a molecule or radical to produce two fragments). In principle, this sensitivity might be so great, and the variations might be so complicated, that we should have to abandon all hope of prediction and attempt to find ways of direct measurement of a great number of dissociation energies before we could build any comprehensive scheme.
Our hopes of some degree of order
arise from our knowledge that regularities hold over a wide range of compounds as far as certain other borndproperties are concerned (e.g. the relationships between bond-length and force-constant).
We shall try
A precise definition^ would specify the molecule (or radical) and the fragments to be in the gas phase at zero pressure, and at 0°K (but see p.9).
to formulate some Idea of where the balance lies, and, in the final section, enquire whether we have yet sufficient theoretical background to enable us to understand (or even predict) any of the energy relationships* The particular bonds with which we shall be concerned experimentally are *(a) carbon-iodine bonds in alkyl iodides (section 2) (b) carbon-carbon bonds in some carbonyl compounds (section 3) (c) oxygen-oxygen bonds in di-alkyl peroxides (section 4) The direct objective of the method of pyrolysis is the estimation of bond-dissociation energies in the molecules (whose dissociations produce radicals or atoms). Indirectly, however, we shall often be concerned with other energies, concerning the radicals produced from the original molecules.
In (b) and (c) we shall be
particularly concerned with bond-dissociation energies in certain radicals and with the activation energies involved in the attack of simple radicals upon molecules. This leads us to the verge of another realm of energy relationships which are, however, almost as important as dissociation energies themselves for the understanding of the mechanisms of simple decompositions.
I. THE DEVELOPMENT OF METHODS OF PYROLYSIS* (a) Comparison with the equilibrium method The kinetic method of pyrolysis, In which we study the rates of unlmolecular decompositions, has provided much more scope for the derivation of bond-dissociation energies than has the other possible method, the equilibrium method.
In the latter, the endothermicity
of the bond-breaking process R-LR2—
♦ R2 is derived
from the teraperature-dependence of the equilibrium constant of the reaction R1R2 ^
R1 * R2
(In this case, of course, the heat of reaction is obtained, unambiguously*) Unfortunately, this method is of general application only in the case of the dissociation of diatomic
2X where only the two species X2 and X can be found in the system.
In the general case of radical production,
* The methods of
have been revie
s type of equilibrium
particularly in a static system, the method is beset with many difficulties caused by decomposition or further reactions of the radicals*
exception is the equilibrium (C6H5)3C-C(C6H5)3 ^
C H gsC H -C H g. ♦ CH^
by methane production* The method has been applied to benzyl and allyl bromides, above*
7° ^ estimating the HBr produced by reaction (3) In principle, comparison of the molar yields of
HBr and dlbenzyl would always make it possible to discriminate between HBr produced from R-Br— > R + Br, followed by (3) and that which might be produced from R-Br — ^ define + HBr, This is upset in the case of allyl bromide by the fact that the allyl radical is not •reactive* in the present sense, and does not only dlmerise, but probably combines with a benzyl radical instead of producing one
A similar effect was notice el in the
pyrolysis of butene-1 (above)* (27 ) The decomposition of benzylamine' ' in the presence of toluene gives ammonia and dibenzyl, ^ C^H^CHg* ^ HHg* ♦ NH2 — > C^H^CHg. ♦ NH^ A particularly good example of the use of the toluene carrier technique is provided by the pyrolysis of h y d r a z i n e , w h e r e the correlation of the gaseous products with the dibenzyl produced made it possible to distinguish between the homogeneous, radical decomposition — > 2N V and the heterogeneous decompositions which occur at the same time, 3R2**4 — * W2 + 4NH3 2N2H4
■■) Hg ♦ Ng + 2HH^
The bond-dissociation energies established by the toluene carrier technique are summarised in table 3«
Bond-dissociation energies established by the *toluene carrier* technique
' S S S S i
temp* range °C.
-3 Hg, N2 dibenzyl
C6H5CH2”Bp benzyl bromide
(29) ( n
CHrjjCHCH0—Br allyl bromide
products from which rates were calcS.
^ , sec”
heterogeneous decomp, occurs together with HgH-HHg— > 2HH2
(e) The present method. In the experiments described in the following sections, we have used the principle of the toluene carrier technique in a static system. In a static system, it is essential to be able to work with as little decomposition as possible.
the temperature must be low enough for no decomposition of dibenzyl to occur in the system. The conditions we have used compare with those of the normal high-temperature flow method as follows «-
100 - 350°C
500 - 800°C
time of reaction
15 mins to
0.1 sec. to 1 sec.
several hours partial pressure of reactant
10 - 50 mm Hg.
0.05 - 0.5 mm Hg.
partial pressure of toluene
50 - 300 mm Hg
5 - 10 mm Hg.
For most of the compounds pyrolysed, it has been possible to work with less than 5# decomposition.
peroxides of section IV the decomposition was generally higher than this.
ll. The Static Pyrolysis of ETHYL IODIDE and sec-PROPYL IODIDE (a) Mechanisms previously suggested for these decompositions. (i) ethyl iodide As outlined in section I, in the previous work on ethyl iodide in these laboratories a flow method has been used.
We now discuss some details of the results of
Gowenlock^D and of Szwarc.^2) It was shown in this later
work that the
decomposition, measured by theiodine production was a homogeneous, first-order reaction giving an activation energy of 54 to 56 kcal. from the temperature-dependence of k* .
The actual values of the activation energy and
the frequency factor found in these two investigations were iGowenlock
E = 55-56 kcal. 'O =
E = 54 kcal.
5 x 101^sec“1
'O « 12 x 1 0 ^ sec
As Butler and Polanyi^ had shown, however, besides the simple decomposition C2 H? ♦ I
there is the alternative possibility CgH^I — ^ CgB^ ♦ HI...................
(The first-order velocity constants obtained from HI production are approximately equal to those obtained from iodine production). Further, the removal of the ethyl radicals produced in (1) must be accounted for.
Since the decomposition
was very low, the most probable fate of the ethyl radicals was regarded as the reaction CgH^ ♦ CgH^I — > C2H6 ♦ followed by the decomposition of CgB^I — » C ^
the C^H^I radical
(probably outside the reaction vessel). The activation energy of reaction (3a) is probably of the order of 10-15 kcal.
(cf. section IVA).
is no doubt that D(C-I) in the radical CgH^I
will be very
much less than D(C-I) In ethyl iodide itself,* and so the rate-determining step in this mechanism would be the primary dissociation (1 ). In the results of Szwarc, quoted above, the frequency factor is rather high, and the activation energy corresponding to a more "normal" frequency factor of 10*3 would be reduced to about 50 kcals. Thus from these flow pyrolyses the limits can reasonably be set at 50-54 kcals. (ii) sec-propyl iodide Butler, Mandel and Polanyi^4 ^ studied the flow pyrolysis of sec-propyl iodide.
From the iodine
m We can make a rough estimate of the strength of the
C-I bond in CgH^I on the assumption that D(C-H) in ethyl iodide is equal to D(C-H) in ethane. The thermochemical data then give Q^CCgH^)^ -44 kcal. and D(C 2H4-I)- -6 kcal. i.e. on this assumption the C-I bond energy is a small negative quantity. Allowing for the uncertainty in this assumption, it seems safe to predict that the activation energy of reaction (3b) is less than 10-20 kcal,
production at 400°C, assuming ^ = lO1^, the activation energy was found to be 46 kcals.
over the temperature range was not good enough for the activation energy from temperature-dependence to have any real significance.
(The actual figure derived in
this way was about 29 kcals.). In this work the hydrocarbon products were analysed and found to consist entirely of propane and propylene. HI was again produced.
After allowing for the amount
of propylene which would be formed together with the HI from the reaction sec C^HrpI *— ) C^H^ ♦ HI the ratio of iodine (I) s propane s propylene was found to be 2 « 1 s 1. This was accounted for by assuming disproportionation of the sec-propyl radicals 2
— » 03Hg + C3H6
after the original decomposition ♦ I
The observed proportions of products could equally veil be accounted for by a mechanism analagous to reactions (3a) and (3b) proposed for ethyl iodide.
(b) previous static pyrolyses The static pyrolysis of alkyl iodides has generally given rather complex kinetics, without yielding any reliable activation energies for the primary process of dissociation. Ogg and Jones^^ made several attempts to follow the pyrolysis of alkyl iodides by a static method in the region of 300°C. With n-propyl iodide, for example, the amount of iodine formed on complete decomposition was checked both by direct titration and by measurement of the residual iodine pressure after removal of gases such as propane and propylene.
The quantities agreed with the over-all
equation, (1) A constant value for the ratio of final to Initial pressures was obtained, and hence the concentrations of iodine and iodide at any time could be deduced from the pressure measurements.
The kinetics were not simple,
the pate being best given by ■
(C3 V ) = k f C ^ D d j j i
It was thought that the most likely mechanism involved the equilibrium, n - C^I
— > nCjHy ♦ £l2
............... (3 )
with the rate-determining step n - OjHp ♦ I2 — * iso-C^Hyl ♦ I
followed by rapid decomposition of iso-C^H^I into the final products. The over-all velocity-constant in (2) thus involves the equilibrium constant of (3 ) and the velocity constant of (4). Previously, Ogg^®^ had studied the kinetics of the thermal reactions of gaseous alkyl iodides with hydrogen iodide, and concluded that the process was almost exclusively RI ♦ HI — » HH + I2
There vas no change in pressure during the reaction and after complete reaction, when all other (more volatile)
products had been removed, the residual pressure of iodine agreed very closely with the original pressure of the iodide. The reaction was followed by colorimetric estimation of the iodine concentration. The kinetics were more complex in these reactions. It seemed that for methyl, ethyl and n-propyl iodides, the rates were best represented by an equation of the form
Ogg Interpreted this as evidence for two concurrent mechanisms, the rate-controlling steps being (a) the blmolecular reaction between the alkyl iodide and HI and (b) the unlmolecular dissociation of the alkyl iodide into the alkyl radical and an iodine atom. The analysis of the kinetics suggested that, in all three cases, this unlmolecular dissociation had an activation energy of about 43 kcals.
It Is now clear
that this value is much too low for any of these
In the present work we have attempted to avoid some of the complexities associated with a static pyrolysis (a) by working vith as little decomposition as possible, (b) by adding mercury vapour to take up the iodine atoms, following the reaction by the gaseous products, and (c) by adding an excess of toluene to capture the alkyl radicals produced. First, we review briefly some photochemical work concerned with the influence of mercury on the decomposition of alkyl iodides.
(a) methyl iodide DCCH^-I) is well-established at ~ 54 kcals (see below). (b) ethyl iodide DCC^^-I)^ 50 kcals (see below). (c) n-propyl iodide
D(C^H^-I) probably rather less than D(C0Hcr-I) - Butler and Polanyi 5 give ^ 50 kcals.
(c) photochemical decomposition of some alkyl Iodides, (1)
The ordinary photolysis of ethyl Iodide at room
temperature was studied In detail by West and Schlessinger and Ginsburg The main products were ethane and ethylene in equal amounts*
In the presence of a silver surface
the rate of the reaction vas greatly Increasedf between
10- and 30- fold, without any great change in the ratio of the products* The mechanism proposed was s-
(1 ) (2 ) (3) (4) (4a) (5)
(6) I♦ I♦ X
i2 ♦ X
The equal amounts of ethane and ethylene were taken as evidence that the main secondary reactions occurring were (3), (4) or (4a), and (6 ).* The absence of butane (I.e. the absence of the recombination
was taken as an Indication
of the low stationary concentration of radicals during the photolysis.
(11) polymerisation of ethylene Induced by alkyl iodides. The photochemical work at higher temperatures (In the region of the present Investigation) has been concerned with the effect of alkyl iodides on the polymerisation of ethylene. Joris and Jungers^42^, and later, Jungers and Yeddanapalli^4^ region of 200°C.
studied this polymerisation in the Without the addition of mercury, the
quantum yield for polymerisation was almost nil.
the experiments of lungers and Yeddanapalll mercury was added to take up the iodine produced and the reaction was followed by the decrease in pressure caused by
* although the possibility of a chain reaction, with (3) followed by (5) is not ruled out by the observation of equal amounts of ethane and ethylene.
polymerisation of the ethylene. For any one iodide, the "polymerisation yield" (the decrease in pressure expressed as a fraction of the initial iodide pressure) increased with increasing ethylene pressure and with decreasing iodide pressure. The polymer was regarded as resulting from successive additions of ethylene molecules to radicals through blmolecular reactions e.g.
CH^ + ^2^4 -- ^ C3H7 °3B7 + °2H4 — > C5H11
Increasing the iodide pressure would increase the radical concentration) so that reactions of radicals with each other would become more important compared with the additions to ethylene, which would explain the decrease in polymerisation yield. Over the series of radicals the yield increased down the series, Me, Et, n-Pr, sec-Pr and this was explained in terms of the decreasing steric factor (for dimerisation) down this series. The time of half reaction decreased along the series, and to explain this it was postulated that there was, in fact, still some back reaction, R + I — > RI.
Thus the mercury was regarded as not completely removing the iodine atoms, but merely competing for them with the radicals R.
The gradation in reaction time was
then also ascribed to the change in sterlc factor in this back reaction. The ultimate fate of the radicals was considered to be, e.g.
CgHj ♦ C2H5I — > C^ 6 ♦ C ^ 4I
C2H4I — * C2H4 * 1
at low radical concentrations, or dlmerlsatlon (at higher radical concentrations).
(d) THE PRESENT INVESTIGATION materials lethyl Iodide (Ward and Blenkinsop) was distilled through an efficient fractionating column. Boiling-point 71.9 - 72.2°C. sec-propyl Iodide (B.D.H.) was washed with sodium bicarbonate and thiosulphate solutions, then with water, dried over CaCl2 and distilled through the column.
Boiling-point 88.2 - 88.5°C. toluene was a product which had been
purified by partial pyrolysis In the flow system at a high temperature, followed by distillation.
The Apparatus. The apparatus was exactly as shown in Fig.l. except that the section between taps Tg and T^ was not used in the experiments with iodides (the mercury pump P and the trap
being added for the experiments described in
sections III and IV). A 5-litre Pyrex flask, V, was used as reaction vessel. This was placed Inside a heavy aluminium block which was heated electrically and thermally Insulated by a layer of kleselgur.
The space above the neck of the flask was
packed with asbestos wool.
The aluminium block was
wound with nichrome tape, the circuit being divided into four independent sections and the current in each regulated separately.
The temperature-variation within V could be
made as little as - 1°C.
The temperature was measured
by a mercury-in-quartz thermometer (the calibration of which was checked at 100°C and 132°C), the thermometer pocket extending below the centre of V. On the vertical tube leading from V were two groundJoints Sj,, S2 to which were connected the side-arms used for the introduction of the reactants.
The outlet tube,
including tap T^ was heated by nichrome wiring, and this heating was continued as far as the trap
arms were enclosed by auxiliary heaters during an experiment, and the heating adjusted so that the whole of the 'dead-space* up to tap T^ was maintained at 150°C. (This ensured a constant pressure of mercury of about 3 mm.).
Tap T^ and the Joints of S were lubricated
by silicone "high-vacuum" grease. to regrease tap trap
(It was necessary
after every experiment).
the tap T 2 led to the "gas receiving section",
comprising the trap H2 ,the bulb W and a McLeod gauge (beyond T^) reading pressures up to 3 nm. Hg.
volume of this section was accurately determined. tap T^ led to the main vacuum line, evacuated by a
mercury vapour pump, backed by a rapid oil-pump.
T^ led to the analytical section* The reaction vessel was previously •baked-out* under high vacuum, and immediately before the start of an experiment, air was in V, and the rest of the apparatus, beyond T^, was under high vacuum. The ethyl iodide and mercury (and, in the appropriate experiments, the toluene) were weighed into the side-arms and these were fitted to
iodide was frozen in a liquid-air bath and the toluene was surrounded by a bath at -78°C*
Tap T^ was then
gradually opened, and the reaction vessel evacuated. The iodide was •degassed* by the usual method (evacuating while frozen, closing T^, thawing at -78°C). At the end of the *degassing*, T^ was closed (with a vacuum of about 10*^ mm. in the whole system).
iodide was thawed at -78°C, and the baths surrounding both side-arms removed.
The side-arms were then quickly
rotated about their ground-joints, allowing first the toluene and then the mercury and the iodide to enter the reaction vessel.
The side-arms were then in the upright
position, and the auxiliary heaters were quickly placed in position surrounding them.
The time of the start of
the reaction was taken as the time when these heaters were
In position, which was only some 30 secs, after the inversion of the iodide side-arm. The trap
was surrounded by a bath at -78°C and
trap H2 by liquid air.
At the end of an experiment,
was opened (with T^ closed) and the pressure in the whole of the system up to
the pressure of •permanent gases* in the system, and provided a check against leakages during the experiment. This pressure was usually of the order of 5 x 10
(corresponding to some 5# of the products later measured) and the irregular small variations in the exact amount indicated that this was a trace of air, remaining from the freezing-out process, and not a genuine product of the reaction. After checking that there was no leakage, all the taps bounding the receiving section were dosed* and the liquid-air bath on trap C replaced by a bath at -78°C.
m If the pressure recorded was higher than about 2 x 10
mm. the whole system was evacuated for a few minutes (retaining liquid air on Hp) until the pressure fell to this order. It was checked that no loss of *02* products occurred during this period (these products have a vapour pressure of ^ 5 x 10*3 nos. at - l8o°C).
The pressure of the products volatile at this temperature was then measured, sufficient time being allowed for the true final pressure to be reached.
volatile, at -78°C, but not at -180°C, will, for convenience be referred to as ,tC 2 hydrocarbons".
products may, of course, contain higher hydrocarbons, or hydrogen iodide). The nature of the system is such that the amounts of C2 products left in V after the end of the experiment can be considered as negligibly small.
This is because
during each experiment there is a total pressure in V of up to 300 mm Hg* and when tap T^ is opened, both the unused reactants and the C2 products are blown quickly through, the reactants condensing in trap H^.
products and reactants are blown through together, even , if there were a small residual pressure in V this would comprise products and reactants in the proportion they were in Just before the end of the experiment9 e.g.
3 mm products, with 300 mm reactants - 1% of products in the residual gas in V.
Further confirmation of the
quantitative transfer of products from V into the receiving section is provided by observations made with the modified apparatus described in sections III and IV.
All experiments after No, 20 were carried out exactly as described above.
In the earlier experiments
the iodide was cooled at -78°C instead of by liquid air during the evacuation.
This may have caused the loss
of a small fraction of the iodide in those cases where long periods of evacuation were necessary, as the vapour pressure of ethyl iodide is not negligible at -78°C. (Extrapolation of the I.C.T. vapour pressure values from 0°C to -78°C gives 0.13 mn for the vapour pressure of ethyl iodide at the latter temperature*). As sec-propyl iodide is less volatile than ethyl iodide, no such precautions were needed in this case and this iodide could be kept at -78°C during evacuation.
Analysis of products. The gaseous products from several experiments (see table 5(a) were tested for hydrogen Iodide by condensing them in the side-tube, beyond T^, containing about 1 ml.
x 'Blank' experiments (very short times of reaction at low temperature) proved that the amount of ethyl iodide recorded as 'C2 products* in the receiving section was, in fact, negligibly small.
TABLE gas Facpt.No.
(Br2 method) 30
excess Hg and toluene
TABLE gas analyses
of aqueous KOH (previously "degassed" as usual) and allowing the contents of the tube to warn up to room temperature, with
With the side-tube at -78°C once more, the "C2 products" were then returned to the normal receiving section and the pressure measured again.
was repeated several times. No decrease in pressure was observed in any of these experiments.
(A qualitative test of the solution
in the side-arm confirmed the absence of HI). The first method of testing for unsaturated hydrocarbons was by a similar absorption in the side-arm by bromine solution (4 drops Br2 in 1 ml KBr solution). The results (table 5(a) ) are reasonably reproducible but the analysis requires a long time and the presence of bromine vapour in the vacuum system has unpleasant effects. Consequently, the elegant method described by Robb and Melville^44^ was used in the later analyses.
method depends upon the photosensitised hydrogenation of the unsaturated hydrocarbons.
The apparatus (Fig. 2.)
consists essentially of a large, modified McLeod-type gauge. The volume of bulb B is about 200 ml. and Q is a clear
quartz tube (about 8 ml*), connected to the capillary tubing aboUte B by a B.14 ground joint.
The taps T11
and T12 were used as a •doser* for the introduction of hydrogen.
The hydrogen was purified by passing it slowly
through heated platinised asbestos and was stored in bulb C at a pressure of about 600 mm (the pressure could be checked *
on a manometer beyond T^).
led to the main
vacuum line and T14 to air. The mixture for analysis was admitted to the gauge through T^, the mercury level raised to the fixed mark X on the capillary above B, and the initial pressure recorded (in arbitrary units - no calibration is needed since all measurements are made for the constant volume defined by X). This initial pressure was usually of the order of 40 mm Hg. Hydrogen was then admitted to the •doser*, I^2 closed and T-^ opened, with the mercury level lowered to Y (Just below the point where the side capillary from T ^ enters the vertical tube). measured again.
T ^ was then closed and the pressure
The dimensions of the •doser* were chosen
so that on expansion into B and compression into Q, the pressure of hydrogen was about 100 mm.
The mercury level was raised Just above the groundjoint of Q and the mixture was Eradiated with a lowpressure quartz mercury lamp.
This brought about the
hydrogenation of the define by means of the photochemlcally-produced hydrogen atoms.
In pressure, measured with the mercury level at X, gave a direct measure of the proportion of olefine In the mixture.*
The final pressure was reached after about
10 minutes itradiation.
With the pressures used, this
method was accurate to within - 3% for about 50% olefine.
It was not possible to separate and estimate the dibenzyl from the reactions In the presence of toluene, as the dibenzyl would be in the complex mixture blown through Into trap H^.
» This is true whether the final product of hydrogenation
is ethane or butane (cf. Jungers and Taylor' In the first case, we have, 0 ^ 4 ♦ H — > C^Hj-
+ H 1 mole
* 1 mole H2 — >1 mole
In the second case, 2 + 2H — > 20^5 °4BI0 i.e.
* 1 mole H2 — > 1 mole C4H1Q
RESULTS The decomposition of ethyl iodide vas studied between 270 and 330°C.
We consider first the results
of the experiments without toluene#
The rate constants,
kg, are calculated assuming that two moles of hydrocarbon products result from the primary decomposition of one mole of iodide.
For ethyl iodide, for example, this is
based on the mechanism i-
C2H5* — > ®2^5 * *
(2 ) (3) This type of mechanism requires, of course, equal amounts of saturated and unsaturated hydrocarbon products. Analysis both by the bromine method and by the hydrogenation method supports this in the case of secpropyl iodide.
With ethyl iodide, however, the
percentage of ethylene estimated is never as high as 50#, and ranges from 15/6 to 40#.
There seems to be no
relation between this percentage and the temperature or the fraction decomposed. The absence of HI is an important feature.
is contrary to what was found in all the previous flow
ethyl Iodide with mercury
k, X 105 sec“l
ethyl Iodide with mercury - packed reaction vessel k0 x 10* sec-1
ethyl iodide pyrolysis without Hg. expt.
*2 x 10^ sec*l
without Hg •- packed reaction vessel 24
without Hg - I2 added expt.
I2 ( & 0
koxl0'? % dec 2 sec*1
0.35 • 2.2
ethyl iodide with mercury and toluene.
x 10* time j (mins) sec"1 1 1
ratio toluenesEtI (moles)
x not Included In graph 2.
CH^ ♦ CO Steacie and Darwent^3) foundthat the nature of the surface did not affect . This has beentaken as evidence against wall-recombination, but their pressures certainly seem high enough (50 to 100 mm) for neither of the alternatives to be important.
In any ease, the pressures In Gorin’s work were so high that his estimate of E for the decomposition of CH^CO must certainly apply to the true unlmolecular decomposition, and not to reaction (c).
In the other
two cases quoted above (Benson and Forbes, Herr and Noyes) It may be that (c) was occurring to sane extent* Our conclusion, as the result of this digression on the photolysis of acetone, Is, then, that the higher ’stability* of CH^CO, as favoured by Gorin, Is substantially correct. There is a serious discrepancy In the literature on the photolysis of acetaldehyde.
This Is definitely a
chain reaction at higher temperatures, presumably caused by the reaction CH^ ♦ CH^CHO — » CH4 ♦ CH^CO ch3co — > ch3 «■ CO There is excellent agreement about the value of the apparent activation energy of the chain reaction, several Independent authors^64V 123“126) giving values between
8 and 10 kcal*
Blacet and Loeffler^4 \ gave 9.6 kcal*
which they identified with the reaction, CH^CO ♦ M — > CH^ ♦ CO ♦ M.
It is quite dear, however, that the previous authors identified the activation energy with the methane producing step, CH3 + CH^CHO
CH4 ♦ CH^CO
and this is Justified by a kinetic analysis. Using the value B -15-18 kcal. selected during the above discussion, and remembering that some small activation energy is probably required for the back reaction, ve choose the value 15 kcals. for the C-C bond-dlssociatlon energy in the acetyl radical."
We must remember, however, in
our future use of this value, that the uncertainty must be placed as high as - 5 kcal.
■ The activation energy for the analogous decomposition of the formyl radical, HCO — ^ H ♦ CO has been estimated at-^26 kcal. by Gorin^56). The evidence for this is (a) the photolysis of CH2O gives 78 kcal. as the upper limit for *D(H-CHO) which corresponds to D(H-CO)— 27kcal< (b) the effect of temperature on HI production in the photolysis of CH3CHO, I2 mixtures gives an activation energy of about 26 kcal. for the decomposition of the formyl radical. The latter piececqf.evidence has been criticised by Blatfet et.al.,'°°'^b'' in whose opinion the system is not adequately represented by the reactions CH3CHO — > CH3 ♦ CHO CH3 ♦ I2 — » CH3I ♦ I CHO — * CO ♦ H H + I2 — » HI ♦ I
(c) Previous decompositions of diacetyl. Rice and Walters
studied the decomposition of
diacetyl between 420 and 470°C, following the reaction by the change in pressure.
It seemed that the main over-all
reaction was ch3cococh3 — » ch4 + CO ♦ ch2co This over-all reaction was first-order, with an activation energy of 66.5 kcal. and a frequency factor ^ = 4 i 101^sec“1. The formation of ketene was accounted for by the following chain mechanism, CH^COCOCH^— * 2CH3C0 CH^CO — ^ CH^ ♦ CO
CH^ +CH^COCOCH^ — > CH4 + •CHgCOCOCH^ •CH2C0C0CH3 — > CH2C0 + CH^CO
CH^ +.CHgCOCOCH^ — » CgH^COCOCH^ (termination) Although acetyl radicals are undoubtedly involved in the chain process it cannot be proved that the primary
step is reaction (1)*
The sequence of reactions,
CH^COGOCH^ — > CH^COCO* ♦ CH3 CH^COCO — > CH3CO ♦ CO CH^CO —
CH3 ♦ CO
could not be distinguished from the pair of reactions (1 ) and (2). This work was extended by Walters^ 2\ who checked the validity of the pressure method by analysis of the undecomposed diacetyl* ^ = 10^
Walters found E = 63 kcal* with
for the over-all reaction, and left little doubt
that the mechanism is a free-radical one, of the type outlined above* In the photolysis of diacetyl^
CgH^ and CO
are the main products at low temperatures, while at higher temperatures the products are similar to those of the pyrolysis*
From observation of the threshold in the
photolysis, Anderson and Rollefson^-^ have estimated 64 kcal* as the upper energy limit for the primary dissociation*
Dlaeetyl (Light and Co.)vas purified by distillation, using the column, and the fraction boiling at 89°C. was used in the experiments. As before, the toluene used vas a product purified by flow pyrolysis.
Blank experiments at the (higher)
temperatures used for dlaeetyl again shoved no decomposition of the toluene alone. Apparatus. The apparatus vas exactly as shown in Fig.l.
as described for the iodide experiments in the previous section, with the addition of the section between taps Tg and T^).
The mercury pump P is needed now, as the
products measured in this case are CH4 and Co, and these must be pumped directly into the receiving section. As with sec-propyl iodide it is not necessary to cool the diacetyl In liquid air during evacuation - a bath at -78°C. can be used. An experiment was ended exactly as before, with trap
at -78°C and
cooled by liquid air.
The pressure of permanent gases produced was measured first in the whole system, without pumping, with taps T2t T CgH^ ♦ (C H ^ C O O C C C H ^
(c h 3 )2c o o c (c h 3 )3 — * (c h ^
♦ (c h 3 )3c o
This explanation is not necessary when k values are used as the basis of comparison. Our absolute k values are in fair agreement with those of Vaughan et al.
We find that at 140°C. our rates
are about 20# lower than those of Vaughan, while at 160°C. our rates are about 40# lower.
This demonstrates how
sensitive is the activation energy over the temperature range of Vaughan, since even this difference is sufficient
to lead to an activation energy 5 kcal. greater than ours. Our greater temperature range, together with the more normal frequency factor, adds confidence in our lower estimate of the activation energy.
(v) Comparison with flow experiments. Further confirmation of the mechanism and kinetics of the decomposition has come from the flow experiments of Murawski^122^.
In this work, between 200°C and 270°C,
the pressure of toluene was kept constant throughout (at 10 mm Hg.) the toluenesperoxide ratio being changed by changing the partial pressure of peroxide.
lower partial pressures (about 0.01 mm Hg.) about 10# of methane was observed.
With the higher pressures (0.1 mm)
negligible amounts of methane were found, the only gaseous product being ethane. Using benzene as carrier in place of toluene^even for the lower partial pressure of peroxide the methane was now negligible.
(This confirms that the 10$ of
methane found above was a genuine product of the reaction of methyl radicals with toluene). Since so little methane was produced In these flow /
experiments, the difference between k and k negligible.
It was found that k was independent of the
ten-fold variation In partial pressure of peroxide. No change in k was found when benzene was used as carrier.
It was checked that the reaction could be
followed to complete decomposition (actually estimated as 103# decomposition). Packing the reaction vessel did not affect the rate or the products. The temperature-dependence over the range 200°C. to 270°C. gave an excellent straight-line plot, leading to an activation ehergy of 34 - 1.5 kcal., with a frequency factor of lO^sec”1. This is in excellent agreement with the result from our static pyrolysis *static system, E = 33.5 kcal.^*= 2 x lO^sec*1 flow system,
E = 34
kcal. v * 1 x 10
Comparing absolute rates, we find on extrapolation that they differ by a factor of about three,* the flow results
■ It seems that the true difference may be even greater than this. Direct comparison has been made between our thermometer and the thermocouple used in Hurawskl's flow experiments. In the vapour of bdLing aniline (184.5°C) our thermometer recorded. 185°C whereas the thermocouple recorded 195°C. If the difference Is the same at higher temperatures, the absolute values of Murawski's rates should be even higher, making the ratio,k(flow) to k(static) perhaps as high as 8 or 10.
Part of the difference might be explained
by the uncertainty in the definition of the reaction temperatures.
This applies particularly to the flow
system, where the volume of the reaction vessel is needed in the calculation of k, and the actual extent of the reaction zone is not known precisely.
It seems hard to
account for the whole of the discrepancy in this way, however,* The only other explanation would be by invoking some extent of chain reaction in the flow system.
observation of equal rates with benzene and with toluene as carriers is, admittedly, only negative evidence against a chain reaction.
Equal rates would still be
observed if there were a chain reaction with a very low activation, so that toluene and benzene were equally inefficient as chain-breakers.
The evidence against a
chain reaction in the static system is more convincing since, here, the toluene causes a much more drastic change in the fate of the methyl radicals, •fixing* a lare proportion of them as methane.
x e.g. it would need a difference of almost 10°C. (in the right direction) to account for a factor of three.
(vl) Conclusion From the results of the experiments described above. In both the static and flow systems, we conclude that the best value for the activation energy of the decomposition is 34 kcal.
We identify this with the 0-0 bond
dissociation energy.* D ft.BuO - Ot.Bu)
In conjunction with the heats of formation given by V a u g h a n ^ \ this corresponds to Q^(tBuO) = 25 kcals. and D(tBuO-H) = 104 kcal.
■ Yaughan quotes a ’calculated* value of 39 kcal. for this bond dissociation energy. It Is Impossible, however, to cpmpute this bond energy from the data he uses, unless it is assumed that D(tBu-OH) is Identical with D(tBuO-H). This value has, therefore, no thermochemical support.
(vli) The activation energy of the reaction. CH^ + C^H^CH^
CH^ ♦ C^H^CH2»
The appearance of ethane In our products from the pyrolysis of di-t-butyl peroxide was an unexpected featuret since the previous estimate of the activation energy of the reaction, CH4 * C6H5CH2# was low enough for us to expect methane formation to occur under our conditions, to the entire exclusion of the dlmerisatlon, CH3 ♦ CH. — > C ^ 6
(a) The previous estimate was by Taylor and Smith, ^114 who measured the rates of reactions of the type, CH^ + RH — * CH4 ♦ R
by the methane production in the photolysis of mixtures of mercury dimethyl with a number of hydrocarbons. proposed mechanism of these reactions was, .
Hg + 2CH^
This work added valuable evidence supporting the idea that the C-H bond dissociation energy in the hydrocarbons RH depends upon the nature of the radical R.
shown that the variation in the activation energy of reaction (1) was parallel to the usually-assumed variation in the C-H energy in the hydrocarbons)• With toluene as the hydrocarbon, the activation energy was estimated as
Although the gradation established by Taylor and Smith is undoubtedly right, it seems likely that the numerical values of the activation energies may require some corrections^11^M120>(121)^ Steacle et al.
pointed out that some difficulties
arise from the assumptions made by Taylor and Smith in deriving the activation energies of reactions (1) from methane production.
Steacle concluded that the estimates
were not valid unless very small steric factors (of the -4 —1 order of 10 to 10 could be assumed for these reactions.
This idea of low steric factors has been
opposed by Evans and Szuarc^11^
on the basis of other
m The temperature-range was wide (140°C. to 290°C.) but
the number of results was small.
evidence pointing to ‘normal1 steric factors* in radical reactions. A possible explanation would be that these activation energies are, in fact, higher than originally estimated.
There is now distinct confirmation that the
values of Taylor and Smith were, in general, at least 2 or 3 kcal. low. Gomer^121^ has re-investigated the photolysis of mercury dimethyl with n-butane, and concludes that the activation energy for, CH3 + CA o is 8.5 kcal.
- > CH^ ♦ C4H9
The corresponding estimate by Taylor and
Smith was 5»5 kcal.
Gamer advances reasons for this
difference and, indeed, states that "these arguments probably apply to other cases where the activation energy of hydrogen abstraction was calculated on the basis of the increase in CH^ formation as a function of temperature".
■ i.e. P - l in the expression, k = pze~E/RT
Steacle and Trotman-Dickenson^120^ have estimated activation energies for a series of paraffins from the photolysis of mixtures of acetone and these paraffins. The competing reactions here are, CH^ + RH — » CH4 ♦
CH ♦ CH3 — > C ^ 6 ........................
+ c h ^c o c h ^ — >
♦ c h 3 c o c h 2»
Results were obtained over a wide range of temperature (100-30(ft3), pressure and light intensity. There was no indication that (4) was a third-body or a wall-reaction. Comparison of ethane and methane production led to estimates of E^ - ^E4 (depending upon E^) for the series of paraffins.
Table 16 shows the values of E^, making
the usual assumption that E4 = 0.
In the four cases
where comparison is possible with Taylor and Smith*s values it is seen that the latter are consistently the lower. The excellent agreement between Gomer and Steacle for the case of n-C^H^Q may be noted. The corresponding steric factor, P^, was given as 10~3 to 10"4.*
IE Assuming P
The estimate actually gave. . Px/ p = 10-3 to 107*
r&(C^l6 )' [tolueney
values of A
dt (packed reaction vessel) A (with t
TABLE 15(b) values of A
- unpacked reaction vessel
A (with t In mlnsi (toluene] in grains)
estimated activation energies for the reactions »■ CH^ + RH — > CH^ ♦ R.
Taylor and Smith*114 ^ 1;L^ E (kcal)
Steacle and Trotman-Dickenson*120* E (kcal)
Gomer*121^ gives E * 8,4 kcal.
Trotman-Dickenson^2^ has Investigated the activation energy of the reaction ch^ + c6hjCHj — > cn4 +
c 6h jcn2
by photolysis of mixtures of toluene and CD^COCD^. ratio
gave the relative rates of the reactions 4
CI>3 + C^H^CH^ — ^ CD^H + C6H^CH2 and
CD^ + CD^COCD^ — * CD4 ♦ CD3COCD2
This led to an estimate of E2 = 8*3 kcal. with a steric factor of 7 x 10*4.
The experimental error in
E2 was quoted as - 0.3 kcal.*
(b) evidence from the pyrolysis of dl-t-butyl peroxide In our static pyrolysis of di-t-butyl peroxide, we have the competing reactions,
CH3 + C6HjGH^ and
CH3 + CE^
CH4 + C^HjCHg*
This estimate of E2 again depends upon the estimated activation energy for the corresponding reaction involving acetone.
Here y too, we have evidence that (4) is a homogeneous, gas phase reaction. We can compare the rates of formation of CH^ and as follows t-
d fC2%] dt
dt whence 4
Assuming that (4) requires no activation energy, we may compute the activation energy of (2) by plotting l°g kg2 against 1/T.
The values of log
o kp are given in tables 15(a) ^4
(for the unpacked and packed reaction vessels,
respectively). evaluated with in minutes.
In these tables the quantity has been jexpressed in grams and the time
124 2 Although the absolute values of kg are consistently ^4 lover for the unpacked reaction vessel, the tenperaturedependence (plotted In figs, 5(a) and 5(b) ) gives activation energies in good agreement tunpacked vessel, Eg = 13 - 2 kcal, packed vessel, E2 * 12 • 3 kcal. (Although there is some scatter of the points on the graphs, the slopes give twice Eg, so that the uncertainty is halved correspondingly in the estimates of Eg). In the flow pyrolysis by Hurawski^122^ we have the main reactions, (ch 3)3c o -o c (ch 3)3 — » 2(ch 3)3co (ch 3)3co — > (ch 3)2c = 0 ♦ CH3 + CH3 - ^ CgH^
In this case only about 5jf of the methyl radicals reacted according to reaction (2), so that the stationary concentration of methyl radicals may be calculatedt11^) from the equation, 2 ky(peroxide] lCH 3i
k4 icn3 ^
\2 k?[peroxideJ A 4|^
The 0 T
. 1 ratio is obtained from the equation,
(2k^)^ [peroxide]^ k4^ k2^toluene]
For generality, ve express the rate constants of reactions (2) and (4) as follows, k2 = p2 x 1012 e'V R I k4
p4 x 1012
using Collision numbers* of 1012 cc/mole.sec, denoting the sterlc factors by p2 and p4, and assuming E4 = 0, We substitute these rate constants in the above expression for lC2H6l■ together with the observed value LCH4~1 of k^ and the known values of [peroxide] and [toluene] Solving for Eg we find, Eg^ 12.0 + 1.15 log(p22/ *>4 In our static pyrolysis, where ve have comparable amounts of CgH^ and CH4, we may calculate the approximate
stationary concentration of methyl radicals in the same way as for the flow pyrolysis." We use the results in the region of 130°C., where our observed
With a toluenesperoxide
ratio of about 10 il we found L C 2H ^ ^ | C H j •
concentrations were,[peroxide] * 7x10“^ mole/cc. and [toluene] = 7 x 10*^ mole/cc. Substituting,as before,9 —in--the for' --- o , — ------------ expression -- r — ;--- -----
2 x 10
13.3 - log p /
E2^12.2 ♦ 0.92 log P22 p4
■ Strictly, we should solve the quadratic equation, [CH^l2 ♦ k2 [toluene] [CH^i c 2k^[peroxide j but this requires assumptions about k2 and k4 , which k4
would destroy the generality of the present method.
We may note how insensitive is this estimate of Eg to errors in the observed ratio of C2H6 and to the other 4 quantities used in the calculation* Even an error as great as a factor of ten would only affect Eg by 1 kcal* We now have the following estimates of Bg i(1)
Ep = 13 - 2 kcal* (from temperature-coefficient of
in static pyrolysis)
p 12 ♦ 1.15 log p2 —*
(from ratio of products in flow pyrolysis)
(3) Eg - 12 + 0*92 log pg2
(from ratio of products in static pyrolysis)
All the evidence from the pyrolysis of di-t-butyl peroxide would, then, be consistent with the value E ?y=12 kcal*, the corresponding conditions, imposed by (2) and (3) being that log
Pg2-a0, i.e* pg2 - p4* P4
(c) evidence from the pyrolysis of other compounds Szwarc and Boberts^^ • have shown that a value of Eg as high as 12 kcal* would not be inconsistent with the results of S z w a r c f o r the pyrolysis of ethylbenzene with
At the higher temperatures ^700°C.) which
were needed for this decomposition the gas produced was almost entirely CH4, from reaction (2). Small amounts of C2 hydrocarbons were produced, however, (about l-2£ of the amount of CH4 ). assumed^11^
If it is
that these were ethane, produced by reaction
(4), then we can obtain at least an upper limit for E2 from the ratio •*
In this case, where we have mainly CH4, we can calculate the stationary concentration of methyl radicals from the equation, k3
■ step in
2 The expression for
[toluene] is derived in the same way as
before, and substitution of the known quantities'2^ E2 ^ 14 ♦ 2.3 log
Finally, ve have the evidence from the pyrolysis of mercury dimethyl.
It was shown^11®^ that the first step
here was Hg(CH3 )2 — » HgCH3 ♦ CH3
followed by the rapid reaction, HgCH3 — ^ Hg ♦ When the reaction was carried out In a stream of toluene (at
500°C) it was found* that for every methyl
radical which was removed by reaction (4) four methyl radicals reacted with toluene according to reaction (2), A treatment Identical with that for ethylbenzene leads to i[C^] [CH4^
2 k 4 k^[HgMe2] k 2 2 Ltoluene"(2
Substitution leads to, E2 - 13.0 ♦ 1.8 log
> Szvarc, unpublished results, quoted by Covenlock^11®^
(d) Conclusions Reviewing the evidence from the pyrolysis of these compounds, Szvarc and Roberts^11^
concluded that the
best consistency was obtained for
log Iz ~ P4 The estimates of E2 would then become (to the nearest 0.5 heal.) (1)
E g ^ l l . O heal.
(static pyrolysis of di-t-butyl peroxide)
E g ^ l l . O heal.
(flow pyrolysis of di-t-butyl peroxide)
E2 ^ U . 5 kcal." (pyrolysis of ethylbenzene*) Ep^-ll.O heal.
(pyrolysis of mercury dimethyl)
These values are In fair agreement with the estimate from the temperature-dependence of
In the static pyrolysis
k4 of the peroxide, and It was concluded that the best estimate
1 2 - 2 heal.
■ It must be emphasised again that the numerical value In this case is much less reliable than In the other casesf since the amount of CpH^ was so small, and its origin not certain.
The condition log
p^2 ~ -1 means that
P4 P4 Assuming, as usual, that p2 4 p4 and since both are £ 1, this condition leads to 0.12 moles permanent gases
basis of 1 mole peroxide— »2 moles permanent gases* Table 19 shows that, between 150°C. and 170°C.t the highest apparent decomposition which can be attained Is about "33%n
(even after times of reaction up to twenty
times those needed for 15% apparent decomposition), i.e. complete decomposition of 1 mole of peroxide does, In fact, produce only two-thirds of a mole of (CO ♦ CH4 ). This limiting value of the apparent decomposition Is practically constant over the range 150°C to 170°C. (The toluenesperoxide ratio is between 15 and 30 In all these experiments).
As noted previously, the ratio
C2II6 is CH4 ♦ CO cannot be admitted as possible at such low temperatures*
decomposition of CH^CHO at these temperatures would almost certainly have to go through some kind of radical chain process* Horiya had found evidence of some chain reaction (see (i) above) but his proposed mechanism* for the slow decomposition does not seem very attractive*
Our present interpretation of the mechanism is as follows* (a)
It is considered very probable that the reaction CgH^O — > CH3 + CH20
is still occurring to an appreciable extent*
reaction is by far the most likely source of the CH20 produced.
Further, the presence of CgH^
confirms the production of methyl radicals at some stage*
■ based on the decomposition of radicals, thus tCpll^O ^ H + CH^CIIo followed by CH^CHO — * CH4 ♦ CO and, which seems particularly unlikely, C2H5O ® ^ CgH^OII
The production of 55-60# of CO in the permanent
gases suggests that acetyl radicals may be involved at some stage (perhaps arising from the decomposition of any CH^CHO produced), these radicals decomposing in the usual way iCHgCO — » CH^ ♦ CO
The Important condition is that whatever reaction produces CH^CO radicals, it must not produce CH^ at the same time.
If one CH4 were produced for each CH^CO,
there would be at least 50# of CH4 in the (CH^ ♦ CO) mixture to begin with, even before allowing for the extra proportion of CH^ which must arise when the methyl radicals from reaction (7) undergo further reaction such as that with toluene (4). This means that the reaction CH^ + CH^CHO — » CH4 + CH^CO
is not sufficient to account for the observed proportion of CO.
The most attractive remaining possibility is CgH^O ♦ CH3CH0 — -* CgH^OH ♦ CH3C0
The main objection is that it is difficult to understand how CH3CH0 can compete for radicals with the toluene which is present in far greater concentration. .
To account for the Intermediate production of
CH^CHO or CH^CO the following chain mechanism Is proposed
as the basis for the decomposition, CgH^O-OC^ — > 2C2H^0
♦ C^O-OC^ —>C^^OII + Ci^CHO-OCgH^ c h 3 c h o - o c 2h 5
c h 3c h o
the chain-ending being by reaction (2) or (9) above. (Alternatively, the decomposition of the radical CI^CHO-OCgH^ might be formulated as the rearrangement, CE^CHO-OC^ — > CHjCO + CgH^OH
This would remove the difficulties involved in postulating radical reactions of CI^CHO as the source of CH3C0 radicals).
To complete the formal scheme, ve add the
complication that if methyl radicals are produced in the system, they need not simply toluene.
dimerise or react with
In fact, most reaction*postulated for CgH^O
radicals can equally well be written for methyl radicals too.
A possible fate of the CgH^O radicals which has not been included in this scheme is the disproportionation CgH^O ♦ CgH^O — > CgH^OH «• CH^CHO
a non-chain reaction giving the same products as the successive reactions (10) and (11).
The only argument
against reaction (12) is that the concentration of undecomposed peroxide will be much greater than that of CgH^O radicals, at least in the early stages of the reaction.
Reaction (12) is favoured, energetically, by
its formation of two molecules, but the abstraction of a hydrogen atom is involved, so that an appreciable activation energy will still be required.* If we consider the radical CH^CHO-OCgH^, suggested as forming part of the chain reaction above, we might expect (by analogy with the acetyl and formyl radicals) to find a marked resonance stabilisation.
have the system iH H R-C-O-O-C-R H
The energy of the secondary C-H bond in the C2H^0 radical is calculated as 30 kcals. (assuming that D(0-0) in the peroxide ^ 33 kcals, and using the heat of formation given by Stathis and Egerton(°°)).
in acetyl and formyl.
Because of the resulting weakness of its secondary C-H bond, it seems very reasonable that the peroxide itself should be able to compete with toluene for radicals. It is interesting to note that, to explain the proportion of gases in the explosive decomposition, Harris and Egerton proposed the over-all reaction CgH^O ♦ CgH^O-OCgH^ — > CH3CO ♦ 2 0 ^ OH which is qualitatively the same a3 our proposed chain mechanism (reactions 10 and 11).
Conclusion, The decomposition of diethyl peroxide seems to be so complicated that we cannot claim any great significance for our activation energy based on permanent gas production.
The fair agreement between our first-order
rate constants and those of the previous authors does, however, add confidence in the estimates of an over-all activation energy in the region of 30-33 kcal,
activation energy will be provisionally identified with the 0-0 bond dissociation energy, but it must be remembered how uncertain is this estimate compared with that for di-t-butyl peroxide. The essential causes of the complexity of the decomposition of diethyl peroxide appear to be (1) the presence of the secondary hydrogen atoms adjacent to the oxygen atoms in diethyl peroxide,
profitable position for the attack of radicals upon the peroxide itself, (ii) the relative stability of the radical C^i^O with respect to the decomposition
(vi) the relative stabilities of the alkoxy radicals Rust, Seubold and Vaughan prepared the unsymmetrical alkyl-t-butyl peroxides, t.BuO-OR.
By analysis of the products from the
decomposition of these peroxides, in cyclohexene solution at 195°C., these authors have deduced the gradation in 'stability' along the series of alkoxy radicals.
•stability' is defined as Lalsotoll_________ [alcohol] + [carbonyl product] in the decomposition products.
The sequence deduced was i-
This gradation is in harmony with the general
experience that the decomposition of di-t-butyl peroxide is by far the cleanest among the di-alkyl peroxides, which virtue is partly due to the relative Instability of the t-butoxy radical. It should be emphasised, however, that the term 'stability' as used by Vaughan et al. is a quantity defined under given experimental conditions.
absolute scale of stability would be that based upon the relative activation energies of the unlmolecular decompositions of the alkoxy radicals.
The results of
Vaughan et al. cannot be taken as conclusive evidence of this sequence of activation energies, for two reasons. (1)
Comparison in the above way assumes that the alkoxy
group of each peroxide behaves exactly analogously to the t-butoxy group, i.e. it is assumed that there are no chain reactions, or other complications caused by the lower alkoxy groups.
Such complications might
misleading ratio of alcoholtcarbonyl compound. (ii) Even if there were no complications in the mechanism, the above 'stabilities' would, strictly, give the differences in activation energy between the unlmolecular decomposition of the alkoxy radical and its blmolecular reaction with the solvent (cf. page i s )•
It cannot be
assumed that the latter activation energy is exactly constant throughout the series.
It does seem, however, that the complicating factors In (1) and (11) are not sufficient to upset the overfall gradation established by Rust, Seubold and Vaughan. We recall our previous estimate (page ho ) of (CH^CO
> CH^ ♦ CH^COCH^ - 6 kcal.
If 33 kcal. can be taken as a fair approximation to D(0-0) In diethyl peroxide, ve may calculate in a similar way t— CgH^O — ^ CH^ + CHgO - 22 kcal. (approx.) (The uncertainty arising from D(0-0) Is halved In this calculation).
Further evidence concerning this latter
decomposition comes from the work of Rice and Rodowskas^102^ on the decomposition of ethyl nitrite.
It seems clear
that the primary step here Is CgH^OHO — * C ^ O ♦ NO
and that there is a competition between the reactions, CgH^O + CgH^ONO — * CgH^OH ♦ CH^CHOHO
C2H5° --- > CH2° * CH3
................. (cf. page 13 )
Rice and Radowskas found that, at low pressures (~1 mm) reaction (3) predominated.
(Steacie and Shaw^103^ had
found reaction (2) at their higher pressures*) and were able to make a rough estimate of E^-E2 ~ l 6 kcal. The activation energies assigned to the individual reactions were Eg
- 3 0 kcal.
This estimate agrees quite well with our above estimate of D ^ ~ 2 2 kcal, particularly when it i3 remembered that in this case
^ E^, but
= E3 -[activation energy
of reverse reaction]( thus we would estimate
Our corresponding estimate for the t-butoxy radical is more difficult to reconcile with the other evidence unless it is assumed that there is a very considerable activation energy for the reverse reaction,
CH^ «• CH^COCH^ — > (CH3 )3CO. Thus, on the usual assumption of not more than about 3 kcal. for a radical addition reaction, we would estimate,
C gH ^C H O
The apparent activation energy was 36.5 kcal., the first-order rate-constants being within 20$ of those found b y Harris and E g e r to m^ 1 ^ for diethyl peroxide. It will be obvious, however, that the mechanisms of these decompositions are too complex for us to have any confidence in conclusions about the difference in bonddlssociatlon energies based on comparison of these rates. (ii) higher alkyl peroxides Milas and Surgenor^10^
shoved that the mechanism
of the pyrolysis of di-t-amyl peroxide was exactly analogous to that which they had established for di-t-buty] peroxide
Ph C2H5 - C - O - O - C '
ch3 — >2 0 ^ - C - 0 CH.
CH3 C2H5 - ?
Raley, Rust and Vaughan^ 1 ^ studied the kinetics of this decomposition under the conditions they had used for di-t-butyl peroxide.
The kinetics were not quite so
accurately first-order, but the activation energy was determined as 39 - 2 kcal. (again over a rather small temperature range).
The first-order rate constants for
the two peroxides were almost identical. In this case, with the cleaner decompositions, we feel confident in concluding from the similarity of the rates, that the 0-0 bond-dissociation energies in the two peroxides are equal, within 1 or 2 kcal.
i.e. D(0-0) ^
34 kcal, using the rather lower value which we favour for t-butyl peroxide. Milas and Perry(108) haye Shovn the generality of the above type of mechanism for higher tertiary peroxides.
(iii) acetyl peroxide Walker and Wild^®) studied the decomposition of
acetyl peroxide9 both in toluene solution and In the gas phase. In toluene solution at 80°C•, the products were almost entirely methane and carbon dioxide (In almost equivalent proportions).
Rate constants were detennlned by
titration of the unused peroxide, and the reaction was found to be first-order.
The estimated activation energy
over a very small range of temperature (less than 10°C) was 31 kcal.
The corresponding frequency factor would be
These observations would be consistent with the mechanism, (CH3C00)2 — > 2CH3COO CH^COO — > CH^ + co2 ch3 ♦ c6h5ch3 — > ch4 ♦ c6h5ch2. If this mechanism were correct then
28 kcal. would seem
to be a fair estimate of the 0-0 bond-dissociation energy In this peroxide (taking a more normal value for
In the gas phase at 100°C. the products were mainly ethane and carbon dioxide in a ratio of about It2.
amounts of methane were formed (about 20% of the amount of ethane).
This pointed to the main over-all reaction (CH^COOjg — * CgH^ ♦ 2C02 which would be accounted for most simply by the mechanism, (CH3C00)2 — » 2CH.C00
ch3coo — * ch3 ♦ co2 2CH3 — * C2H6 Edvards and Mayo^10^ have recently studied the llquld-phase decomposition of acetyl peroxide in mixtures of CC14 and hydrocarbon solvents (Including toluene). The ratio of CH^Cl to CH4 was measured for several solvents, in order to compare the reactivities towards the methyl radicals expected from the decomposition of the peroxide. These reactivities were compared with the established reactivities of these solvents towards higher alkyl radicals In polymerisation systems.
Anomalies ver^lfound which
could only be understood (1) if the methyl radicals behaved very differently from the higher hydrocarbon radicals. or (11) If the ultimate products from the acetyl peroxide decomposition arise not from methyl radicals but from CH^COO radicals which decarboxylate as they react, CH^COO ♦ RH — > CH4 «■ C02 ♦ R. Edvards and Mayo prefer this second explanation.
(lv) benzoyl peroxide The liquid-phase decomposition of benzoyl peroxide has been the subject of many investigations*
conoluslon is that the over-all reaction is approximately of the first-order (at least In dilute solution) and that the activation energy is about 30 kcal*
There is also
general agreement that the choice of solvent can affect the rate considerably, without, hovever9 causing any significant change in the apparent activation energy. According to Hozakl and Bartlett(109)f the over-all reaction comprises two simultaneous processes - a unlmolecular decomposition, (c 6h ^c o o )2 — * 2C6HjC00
c 6h ^c o o
— * C6H^C00C6H£ + co2
and a chain reaction involving, C^H^COO ♦ (C6H^C00)2 -- * C6H^COOC6H^ ♦ C02 ♦ CgHjCOO. Rates were compared in over thirty different solvents. In benzene, the activation energy over the range 60 to 8o°C. was estimated as
33 kcal., while in acetic anhydride the
estimate was 31 kcal. was
(the corresponding frequency factor
lO^sec -1 in eaCb case).
Data quoted by Brown'
, for a wider temperature
range lead to an estimate E ^ 27.5 kcal. with ^ * 3 x 1012
.0.-1 (2 )
of the rate upon the solvent and estlrated that for the first-order reaction in infinitely dilute solution, E^31 kcal. Instead of estimating the residual peroxide, Bawn and Mellish^2®^ have developed a novel method of following the decomposition with the aid of the radical,
This stable, coloured radical is rapidly removed from solution by combination with the benzoyl radicals produced from the peroxide, and this disappearance can be followed colorimetrically.
This method gives an estimate
of about 35 kcal. for the activation energy, with to lO^sec-1, the rate again depending on the solvent.
Frank^4^ has reviewed the decomposition of hydroperoxides, substances of great importance in theories
of hydrocarbon oxidation*
are| in general, outside the scope of the present account, but it is worth noting that several tertiary hydroperoxides have been shown to decompose by a mechanism similar to that for tertiary peroxides*
Milas and Surgenor^®^ found that
t-butyl hydroperoxide could decompose by the splitting-off of an OH radical, (ch3 )3cooh — » oh «■ (ch3 )3co. It is of particular interest that this applies to the explosive reaction at 250^C*
The smooth decomposition at
100°C* produces t-butyl alcohol and oxygen*
us of diethyl peroxide, where Harris and Egerton^®1 ^ found the explosive decomposition to have the cleaner over-all reaction (page
THEORETICAL DISCUSSION OF BOHD DISSOCIATION ENERGIES
(a) general formulation In terms of conjugation and hyperconjugation In the theory of Baughan, Evans and Polanyi^1 ^ are the energies of compression or extension of the standard C-C and C=C bonds from equilibrium length to length r.
was taken as 64 kcal. (cf. Baughan and Polanyi^1^
) cs and
Cd were calculated from Morse functions. The equilibrium bond-lengths assumed in these calculations refer to the lengths of the standard bonds and so are slightly greater than the observed bond-lengths in ethane and ethylene. The magnitude of
chosen is largely determined by
the value of ~ 5 kcal. preferred for xCCgH^).
Using p CH
= 125 kcal.., the value of the third-order conjugation energy H H in the system -C-C- varies from 2.5 kcal. at the ethane (C-C)distance to 7 kcal. at the ethylene distance, i.e. x