Heats of dissociation of some hexaarylethanes

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U n p u b l i s h e d t h e s e s s u b m i t t e d f or the M a s t e r * s and D o c t o r ' s d e g r e e s a n d d e p o s i t e d in t h e N o r t h w e s t e r n U n i v e r s i t y L i b r a r y a r e o p e n f o r i n s p e c t i o n , b u t a r e to be u s e d o n l y w i t h d u e r e g a r d to t h e r i g h t s o f t h e a u t h o r s . Bibliographical r e f e r e n c e s m a y be n o t e d , b u t p a s s a g e s m a y b e c o p i e d o n l y w i t h t h e p e r m i s s i o n o f t h e a u t h o r , a n d p r o p e r c r e d i t m u 3 t be g i v e n in s u b s e q u e n t w r i t t e n o r p u b l i s h e d w o r k . E x t e n s i v e c o p y i n g or p u b l i c a t i o n of t h e t h e s e s in w h o l e or in p a r t r e q u i r e s a l s o t h e c o n s e n t o f t h e D e a n o f t h e G r a d u a t e S c h o o l of N o r t h w e s t e r n _ University. T h i s t h e s i s by h a s b e e n u s e d by t h e f o l l o w i n g a t t e s t t h e i r a c c e p t a n c e of the




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ACKNOWLEDGMENT To Professor P. W* Selwood for the direction, support and encouragement necessary to the successful completion of these researches*

TABLE OF CONTENTS page INTRODUCTION.............................................




1* Heats of Dissociation............................... 9 Preparation of Materials ......................... 9 Filling the Sample Tubes ......................... 10 Density Measurements . . . . . . . . ............ 12 Analytical Measurements............................. 13 Magnetic Measurements...................... . . # # 14 2# Heat of Activation of the Disproportionation R e a c t i o n ...................... . . ...............• * 19 Treatment of Data................................... 20 3# The Spectrographic Investigation .................. 21 4, The Ebulliometric Investigation.................. . . 23 Treatment of Data................................... 24 RESULTS . # # .......................................... #


1* Heats of Dissociation.................................26 The Disproportionation Reaction..................... 26 The Color Effect # • ............................... 27 Equilibrium Constants# ....................... 30 Free Energies, Heats, and Entropies of Dissociation .......................... 31 Discussion of Results............................... 31 2. Heat of Activation of the Disproportionation ....................................... 36 Reaction Discussion of Results............................... 36 3* The Spectrographic Investigation.................... 39 4. The Ebulliometric Investigation.................. 41 Discussion of Results............................... 42 SUMMARY

............................................. 46

VITA....................................................... 48




Since Gomberg's famous discovery of triphenylmethyl, radi­ cal chemistry has provided a fertile field for research to both organic and physical chemists.

Much of the interest to the

physical chemist has centred around the measurement of dissocia­ tion constants of the hexaarylethanes from which then can be calculated such quantities as the heats, free energies, and en­ tropies of dissociation.

The effect of aryl substitution in the

parent ethane in reducing the carbon-carbon bond energy to the i point where dissociation occurs at and below room temperature has aroused considerable interest, as also the possibility of determining the heat of activation of the dissociation process. Measurements on solutions of radicals in indifferent solvents give a convenient means for checking the conclusions obtained in theories advanced to explain the stability of radicals and i

the extent to which dissociation will occur in a given ethane.

| One of the distinguishing characteristics of free radicals I Iis their color, so colorimetry has been applied to the problem, !

and Ziegler


has used the photometer in determining extinction

coefficients, which give a means of determining the dissociation. Cryoscopic and ebullioscopic measurements of molecular weights |of ethanes in solution have been used to determine their dissocia tion.

These measurements, however, are of limited applicability,

since with a given solvent only two temperatures are available


(1) See Preckel "M.S. Thesis", Northwestern University (1941). (2) Ziegler and Ewald, Ann. 473 y 163 (1929).


for investigation, the freezing and boiling points, respectively. Colorimetry and spectrophotometry provide more versatile modes of attack, in that the temperature may be chosen at will, but difficulties may easily arise due to uncertainties in measure­ ment of the temperatures, invalidity of Beer's law, and the strong probability that there is always present, in a given ethane solution, a certain amount of highly colored impurity3 resulting from some side reaction occurring with the preparation of the ethane from triarylmethyl chloride and "molecular11 silver. The uncompensated spin of the "odd" electron of a radical makes available another physical chemical tool in the measure­ ment of its concentration, its magnetic susceptibility.


knowledge of the magnetic susceptibility of the ethane in solu­ tion will permit the calculation of the extent of dissociation, !if there is no disturbing factor such as an orbital contribution. However, in molecules like the triaryl methyl radicals, all orbi4s

jtal moment can be assumed to be damped out and all paramagnetic i !contribution to the ethane susceptibility will then arise from the electron spin.

The molar paramagnetic contribution of an

unpaired electron is well known, having a value of about |+1270 x 10"*6 e.s.u. at 20°C.

Thus, knowing the diamagnetic por­

tion of the susceptibility of the molecules, which can be cal­ culated from Pascal's constants, and having measxired the resultant of the para- and diamagnetism, we may easily calculate the |

! (3) Gomberg and Sullivan, J.A.C.S. 1810 (1922). ! (4) Mttller and Mftller-Rodloff, Ann. 521, 89 (1935),

3. i

proportion of radicals present.

In recent years the magnetic

method has heen rather well established as a very useful one for measurements on radicals, though it is generally necessary that the dissociation to

radicals exceed about one or two percent ,4

Hexaphenylethane is

the only compound of thistype to have



investigated by several different methods by different investiga­ tors,

Mtlller4 obtained a value of A H for this substance using

the Gouy balance magnetic method, reporting

A H = 11,6 + 1,7 kcal,,

which corresponds with Ziegler’s2 photometric determination and j 0 |fairly well with Wooster’s colorimetric value. The heat of t Idissociation is obtained from the change of equilibrium constant with temperature: inK^ K2

_AH Ta - T ± R Ti x Ts ’

Integrated from van't Hoff's d In K AH # I dT RT 2 j Bachmann 9 has discussed the theories of radical stability. Most important of these theories are the steric theory, the re­ sonance theory, presented chiefly by Pauling and coworkers, and Conant’s semiempirical attempt at predicting the heats of dissocia­ tion of a number of ethanes. The steric theory, for which Bent 1o and coworkers present (5) (6 ) (7) (8 ) (9)

Roy and Marvel, J,A,C,S, 6 g, 2622 (1937), et seq, Michaelis and Fetcher, ibid, 5g, 2460 (1937), et seq, Selwood, ibid. £1, 3168 (1939)7 et seq. Wooster, ibid. 5g, 2166 (1936). j Bachmann "Treatise on Organic Chemistry”, ed. Gilman, Vol. I, j p. 489, Wiley, New York (1938). j (10) Bent, J.A.C.S. 5g, 1786 (1931), et seq.

their measurements of electron affinities of the radicals and the calculations of the ethane bond energy from heats of hydro­ genation and oxidation of trlphenylmethyl as evidence, assumes that the weakening of the ethane bond is due to steric hindrance. Thus it is purely a space effect, and the more space that a group takes up near its connection to the radical carbon, the more it should promote dissociation.

Bent concluded that the ethane

bond in hexaphenylethane has been weakened by about 30 kcal., which he attributes to the steric hindrance effect, the remainder of the weakening being assumed to be resonance energy. This leaves approximately the same amount of bond energy loss to be accounted for— the bond energy in ethane being about i

80 kcal. and the measured bond energy of hexaphenylethane being about



The proponents of resonance




as the promoter of dissocia­

tion reason that the large number of resonance structures avail­ able in the substituted methyl radicals simply compensates for the i

large amount of energy required to break what is considered to be a substantially normal ethane bond.

Thus we expect a biphenylyl

radical to be more efficient in producing dissociation than a simple phenyl, and also that a-naphthyl should be more active than fi -naphthyl, since the former has one more canonical form than the latter.

The beginnings of an extension of the resonance

theory to include saturated structures have also been made, but this aspect of the theory requires much further work before it (11) Pauling and Wheland, J. Chem. Phys. 1, 362 (1933). (12) Wheland, ibid. 2, 474 (1934).

becomes as powerful a tool as the resonance involving unsatura­ tion* The system worked out by Conant


was, in essence, the

assignment of a value to each given substituent group which re­ presented the amount it reduced the heat of dissociation of ethane.

Thus it required the available data on dissociated

| ethanes and, from the degree of dissociation determined colorimetrically, or from a rough estimate of the temperature at which the dissociation is one per cent, the free energy of dissocia­ tion was calculated*


The assumption was made that the entropy

j of dissociation of all ethanes of similar structure is the same, |and constant with temperature at the value measured for hexaIphenylethane, and the heats of dissociation were then calculated* ! ;Further calculation gave values to be assigned to each substitu­ !

ent, such that, supposedly, any heat of dissociation of a sub|stituted ethane could be calculated arithmetically* i | An objective of this research, then, was to test the various |theories by means of experiments designed to confirm or deny

j the conclusions arrived at by each. redetermine the

Thus it was projected to

A H for hexaphenylethane in order that, knowing

its equilibrium constant at 20°C., the

A F of the dissociation

reaction might be calculated, and, thence, the AS*



was then to be compared with the values obtained for the other ethanes investigated in order to test the validity of Conant*s |basic assumption of equal entropy changes for the dissociation i

| | (13) Conant, J. Chem. Phys. 1, 427 (1933).


of different ethanes*

Another test of the Conant system would

result from the measurement of ethane and comparison with


H for di-a-naphthyltetraphenyl­

A H for he xaphenyle thane*

The pre­

diction is that the former should be 4 kcal. less than the latter* The relative merits of the resonance and steric theories may be sought into by comparison of di-o-tolyltetraphenylethane with di-a-naphthyltetraphenylethane, also tetra-o-tolyldiphenyl­ ethane (seeTab.V).

The steric theory predicts that the di-o-

tolyl compound should be dissociated to about the same extent as the di- oC-naphthyl while the tetra-o-tolyl compound should be more highly dissociated than either of the others*

The resonance

considerations on the other hand indicate that the ol -naphthyl group should be more effective in enhancing dissociation than an o-tolyl group by virtue of the fact that it has seven posi­ itions

available for resonance of the odd electron while the | o-tolyl has only three, the same number as the phenyl group it­ self, barring possible resonance with the o-methyl group, about which no quantitative calculations are at present available* In determinations of the susceptibility of a radical soluItion against temperature, it has been found that, above a cer­ tain temperature, the radical concentration diminishes with time, and the higher the temperature, the more rapidly the radical disappears*

This critical temperature is known to vary from

radical to radical, being about 0°C* for di-o-tolylphenylmethyl, j60°C* for o-tolyldiphenylmethyl and well over 90° for triphenylj methyl* In view of the importance of radicals as postulated


intermediates in many reactions, it becomes of interest to inves­ tigate this disappearance reaction more thoroughly* By a simple extension in technique of the measurements necessary for

A H determinations, it is possible to determine

the heat of activation of the reaction by which the radical disappears.

Thus, if a solution is heated in the Gouy balance

setup in the conventional manner, and kept at a temperature such that the radical disappears with measurable rapidity, a reaction rate constant for the disappearance of the radical may be calcu­ lated.

Then repetition of this experiment at a different tempera­

ture will enable a calculation of the heat of activation of the reaction to be made.

Thus, it was projected to measure a

heat of activation value for one of the radicals studied, in the hope that this might lead to some suggestion of the nature of the slow step in the reaction of the radicals. It has been observed that the color of a radical solution is not necessarily an indication of the extent of dissociation, and in fact it has been found that such a solution heated over a long period until all radical has disappeared is still colored. It appeared to be desirable to determine roughly what happens to the absorption spectrum of a radical solution upon prolonged heating, at, say, the boiling point in benzene solution.


this way it is possible to decide whether or not colorimetry is applicable to radical solutions, which seems unlikely, except in the ease that the radical is formed by a perfectly clean cut

i i

reaction of the methyl chloride with the molecular silver and is very stable after formation, so that colored byproducts or


decomposition products would not be present.

It seems probable,

also, that, by an extension of the technique to spectrophotometry with the determination of extinction coefficients, at least part of the problem of what happens to the radical upon its irreversible disappearance may eventually be solved. Certain determinations of the dissociation of radicals have been made ebulliometrically, especially in the field of organometallics.

Since this method provides a convenient check on the

effective overall molecular weight as heating continues over a long period of time at the boiling point, it seemed that some interesting data might result from following the apparent molecu­ lar weight of one of the radicals as it was destroyed by the heating.



has developed ebulliometry to a very

useful state of accuracy so that it appeared that the measure­ ments obtained might be meaningful in relation to those given by the magnetic kinetic study and the determination of the heat of activation. I | It will be seen, then, that the objective of this research iwas fourfold: magnetic measurements of the heats of dissociai ition of four he xaarylethanes, determination of the heat of activa­ tion of the radical disproportionation reaction, spectrographic investigation of the color of a radical solution undergoing dis­ proportionation, and ebulliometric determination of the change in apparent molecular weight with disproportionation. i I (14) Swietoslawski, "Ebulliometry" Chemical Pub. Co. of New York (1937).




1* Heats of Dissociation, Preparation of Materials.

The quantities of triarylmethyl

chlorides required for this investigation were obtained, for the most part, through the courtesy of Professor C. S. Marvel and Mr* C. M* Himel of the University of Illinois.

The sample of

triphenylmethyl chloride was prepared by Dr. J. D. Malkemus, formerly of this University.

All of these materials were supplied

in a high state of purity, as attested by sharp melting points and their appearance as nearly white powders or crystalline solids.

They were obtained vacuum sealed into glass ampoules

which, upon being opened, were either immediately evacuated and resealed or stored in a vacuum desiccator. i

"Molecular11 silver was prepared by internal electrolysis, substantially as described by Gomberg. Toluene was prepared from ,,C.P.tt toluene by repeated extrac­ tion with pure concentrated sulfuric acid until the extract no

jlonger showed color. Subsequently this extracted toluene was | Ishaken with successive portions of dilute sodium hydroxide solu­ tion and distilled water, dried over phosphorus pentoxide, shaken with C.P. mercury and distilled from sodium, the middle half being collected. The radical solutions were prepared by shaking a solution of triarylmethyl chloride with a tenfold excess of molecular silver under an atmosphere of toluene vapor at the vapor pressure

of the solution.

The shaking was continued for a longer or

shorter period depending on the chloride used, though in no case was the shaking period less than two hours.

The solution was

then filtered through a sintered glass disk into the bulb from which the expansion pycnometer and magnetic sample tube were ; filled.

In order to wash the silver chloride precipitate and

excess silver, as well as the preparation bulb, some of the sol­ vent toluene was distilled back into the preparation bulb.


ter the third washing further washings were never colored, so a threefold washing was considered sufficient.


Filling the Sample Tubes.

The apparatus evolved for use

in filling the expansion pycnometer and the magnetic sample tube | is shown in Fig. I.

It was sealed at A to an evacuation system

| consisting of a mercury vapor pump backed by a Cenco Megavac and was mounted on a board to facilitate shaking.

The first

step was to seal the apparatus, without the sample tubes attached, under high vacuum for


hours with freshly melted sodium in B

in order to destroy water vapor. ; silver was placed in bulb C.

The required amount of molecular

Then the sample tubes were sealed

into their respective positions with 2-3 cc. portions of anhydrous


I toluene in them.

Subsequently, this toluene was distilled back

through the apparatus and out at A into a dry ice-toluene trap. Toluene was introduced into B over the sodium, the toluene being heated to the point of reflux several times to melt the sodium, i thus ensuring removal of traces of water.

The quantities of

molecular silver and triarylmethyl chloride required were introduced


through the addition tube at D, which was later sealed off at E. The apparatus was then vigorously evacuated, the toluene in B being protected against extensive evaporation by cooling in a dry ice— toluene bath*

After evacuation, the apparatus was

sealed off at A and the toluene in B distilled into C by allow­ ing it to warm to room temperature and by sharply cooling C with solid dry ice*

When the distillation was complete, the

bulb B was sealed off from the remainder of the apparatus at F. The resulting triarylmethyl chloride solution was shaken with the tenfold excess of molecular silver for several hours, the solution becoming colored immediately owing to formation of free radical*

The solution, containing the ethane, free radical and

unreacted chloride, was filtered through the sintered glass disk into bulb H*

Bulb C and the molecular silver were washed

by distilling 2-3 cc. portions of the solvent using the same technique as previously.

Then the preparation of the solution

was complete and proper tipping and tilting of the apparatus served to fill the sample tubes, which were washed with the solu­ tion before the final filling.

After the proper amount of solu­

tion had been introduced into the magnetic sample tube, a seal was made at I by melting some Wood*s metal on the side of the tube and allowing it to flow into a constriction*

This was

done so that the solution in the sample tube could be manipulated satisfactorily without danger of changing its concentration^ thus, solvent was distilled up and down to clean the walls of the filling tube, and the solution was cooled with dry ice before

the tube was sealed off at J*

Similarly, after the pycnometer

had been filled a metal seal was made at K and the pycnometer subsequently sealed off at L.

Finally, in order to make the

apparatus less unwieldy for weighing operations, etc., the fill­ ing arm was sealed off at M. Density Measurements.

Densities of the solutions over the

temperature ranges employed were determined by means of an ex­ pansion pycnometer of about


ml. capacity with calibrated ex­

pansion neck some 10 cm. long.

The neck was semi-capillary tub­

ing of about 3 mm. internal diameter with a steel scale securely attached alongside so that it could be read with a eathetometer telescope.

The scale was marked in sixtyfourths of an inch or

approximately 0.4 mm. divisions.

With the cathe tome ter tele-

I scope, readings could be made to half a scale division or + 0 . 0 2 cm.

It will be seen that this corresponds to less than 0.0014

ml. error, which, in a total volume of in density of less than +0.0002.


ml. gives a tolerance

Calibration of the pycnometer

was performed with freshly redistilled water, and no measurable | hysteresis was observable$ the reading at a given temperature was the same whether the temperature was approached from above or below.

The pycnometer was immersed in a tall cylindrical oil

bath equipped with an efficient variable speed electrical stirrer and heated with a 250 watt knife heater whose heating rate was controlled with a Variac.

The temperature was held con­

stant to +0.1°C. throughout a period of about ten minutes for each reading.

Temperature indication was by means of a calibrated

eopper-constantan thermocouple used in connection with a Rubicon

potentiometer and a Leeds and Northrup HS galvanometer.

For the

measurements below room temperature, it was found quite satis­ factory to make measurements on the solution at 30°, 40°, etc. and at


°, and to obtain intermediate values and that for -1 0 °

by interpolation and extrapolation respectively. Analytical Measurements.

The pycnometer solution was

analyzed for total solute by evaporation.

The pycnometer was

cleaned by dipping it in successive baths of methyl alcohol and benzene until a standardized procedure gave weighings reproducible to + 0 . 1 mg.

The sealed tip was then broken open, care being

i taken to collect all pieces resulting from the break, and the toluene distilled into a dry ice— toluene trap under vacuum,


using a Cenco Megavac.

Thus air was excluded, preventing forma­

tion of peroxides, and evaporation was ensured. i

As the evapora­

tion proceded, heat was gradually applied to the pycnometer until a temperature of about 100°C. was attained.

Then it was allowed

to cool, still under vacuum, and weighed along with the pieces from the break of the seal.

Then the pycnometer was cleaned of

| ethane and chloride and weighed again.

The aggregate error in

jweight of the solute was not more than +0 . 2 mg., which, in a total solute weight of 100-300 mg., would not exceed 0.2$.


similar weight error in a total of about 8.5 g. of solution would obviously contribute a negligible error in the density calculation. In order to obtain the concentration of ethane in the solu­ tion, the preparation part of the apparatus (Figure I), sealed at points E, F, and M, was weighed to 0.01 gm.

The solution was

then removed after breaking the seal between bulbs C and H, care being taken to account for a 1 1 pieces of glass from the break. After the apparatus was washed with toluene, the weight of solu­ tion was obtained by reweighing the dismembered apparatus, again to 0.01 gm.

The concentration of ethane was obtained from the

weight of silver chloride obtained by the reaction producing the ethane.

Thus the silver chloride— molecular silver mixture re­

maining in bulb C was extracted repeatedly with 3 N ammonium hy­ droxide solution, giving a solution which was filtered and from which the

silver chloride was reprecipitated by additionof a

slight excess of cipitated

conc. nitric


After being cooled,the pre­

silver chloride was collected in a sintered glass Gooch

crucible,washed with slightly acidified distilled water little pure acetone, dried in an oven at +0.05 mg.

The error here is aggregated to

120° 0.1

and a

and weighed to mg. in 200-400 mg.,

which is less than 0.05$ or less than the error in weighing the solution, which was +0 . 0 2 gm. in 20-60 gm. or less than


.1 $.

Thus, the tolerance in concentration of ethane may be reported as less than

0 .2 $,

the concentration of the triarylmethyl chlor­

ide, obtained by difference, may be in somewhat larger percentage error, since it was generally present in smaller concentration than the ethane.

However, little significance attaches to this

fact, since the diamagnetic susceptibility of the chloride is not very different from that of the solvent. Magnetic Measurements.

The magnetic measurements were per15, 1 formed with the Gouy balance previously described. A new (15) Haller and Selwood, J.A.C.S. 61, 85 (1939).

modification was the heating coil of fine nichrome wire wound about the lead heat reservoir in the Dewar system.

This, in con­

junction with a Variac, was used for control above room tempera­ ture in an entirely satisfactory manner.

Control to +0*1° was

quite simple and could easily be refined to +0 *0 1 ° if necessary. Below room temperature, the previously described technique of cooling the heat reservoir by dropping oxygen liquified in a liquid nitrogen condenser into a copper tube imbedded in the reservoir was again used.

However, at the relatively high tempera­

tures here employed, -10°C. or above, it was found advantageous to maintain the temperature, once it had been reached, by passing a stream of tank nitrogen through the cooling system.

At the

temperatures used this procedure permitted much better control than did the use of liquid oxygen^ here, also, the control to + 0 .1 ° was very easy and doubtless susceptible to further refine­ ment. The magnetic sample tube was of a modified Freed and Kaspar

type which permitted use of rising or falling temperature during the course of a series of measurements.

The modification em­

bodied the use of a thin capillary tube communicating between the lower half of the sample tube and the bottom of the partly empty toluene reservoir.

In the original Freed and Kaspar tube, eola­ tion munication to the liquid reservoir was through a e o n & n e ­ cessitating the use of constant or rising temperature. Treatment of Data.

The measured quantity in the Gouy method

(16) Freed and Kaspar, Phys. Rev. 36, 1002 (1930).


Aw, the apparent change in weight of the sample on applica­

tion of the magnetic field* may be calculated*

From this quantity the susceptibility

In this research, using the double ended

|sample tube modified from Freed and Kasp a r ^ original design, it was necessary first to measure the "tube constant** over the range of temperatures used*

These data were obtained by filling

|both ends of the tube with toluene and measuring A w at 12*5 |jamperes, the standard field coil current producing 13,100 oersteds, j'at the various temperatures.

These values correspond to the

!magnetic disymmetry of the two ends of the sample tube which |was of constant internal diameter (about 5 mm. +0*01 mm., Pyrex). |

The relative susceptibilities of the toluene at various

||temperatures were then measured by emptying and evacuating one i|

j!end of the tube and again measuring iture*

A w as a function of tempera-

These measurements supplemented by an absolute value of


ithe susceptibility of toluene obtained at 25° gave the absolute j; insusceptibilities of toluene over the desired temperature range. Table I Temp. °C. 25° -10 0

+10 20 30 40 50 60 70 80 90 100

0.7060 x 10~ 6 e.s.u* .7036 .7044 .7051 .7058 .7065 .7076 .7081 .7088 .7095 .7101 .7110 .7116


In order to obtain » XI P XI Pi


1 t*.


*H r0


X Pi ♦H iH rH 0 p 1 0 1 0 iH ! P 0 P


Table V Best; Probable Values for the Degree, Free Energy, Heat, and Entropy Change of Dissociation for Some Hexaarylethanes Free Radical


for 0*03 molal solution at 2 0 °



Fo keal?

HSo kcal.

s20/ cal./°C




cP o












C jX c H s ch 3


.6 ?



confirms the observations of Wooster

J3s "b

and of Bent ,2 3 and shows

that Dr* Conant*s interesting attempt to predict degrees of dissociation rests upon a faulty assumption*

His theory would

also suggest that the heat of dissociation of d i - a -naphthyltetraphenylethane would be 4 kcal* less than that of hexaphenyl— ethane*

Such is not the case*

In connection with the resonance theory of free radical stability it has often been pointed out that & —naphthyldiphenylmethyl with thirteen


structures is

much more stable than the fb -naphthyldiphenylmethyl with twelve* The present work shows that this difference in stability may or may not exist*

It also shows that di-o-tolyl tetraphenyle thane

with only nine contributing resonance structures has the same heat of dissociation as di- (X-naphthyltetraphenylethane with thirteen*

The degree of dissociation of hexaphenylethane is

increased by an o-tolyl group almost as much as by an C(-naphthyl group*

The alkyl group may, of course, also contribute resonance

energy, but the extension of the resonance theory to include alkyl groups awaits elaboration*

It is quite evident from this

work that the addition of an ortho-alkyl group, and especially of two such groups (as in tetra-o-tolyIdiphenylethane) is far more effective in promoting dissociation than a large increase in resonance energy.

It should be pointed out that the a — x^ factory results as shown in the Table VII. The second order constants show maximum variations of always less than 30$ and the greatest deviation from the mean is about 15$. activation obtained is


The heat of

kcal. which is probably correct to

within a kilocalorie or so. Measurements by numerous investigators on hexaphenylethane solutions using iodine or an odd molecule such as NO to destroy the equilibrium radical concentration and observation of the rate of restoration of the equilibrium in a large variety of

Table VI Time



0.75 hr. 1.25 1.75 2.25 2.75 3.25 3.75 4.50 5.50 6.50 7.50 8.50 9.50 10.50 15.00 17.00 19.00 21.00

23.00 25.00 27.00 30.00 33.00 46.25


0.809 .802 .740 .755 .720 .691 .669 .680 .629 .602 .577 .545 .536 .516 .496 .421 .393 .365 .343 .317 .317 .302 .283 .263 .206



0.0836 .0819 .0766 .0780 .0744 .0714 .0691 .0704 .0651 .0623 .0596 .0563 .0554 .0534 .0513 .0435 .0406 .0377 .0355 .0328 .0328 .0313 .0293 .0271 .0213


0.58 hr. 2.16 3.00 4.10 5.10 6.40 8.00

12.50 23.75



0.761 .704 .584 .495 .433 .396 .342 .307 .218 .160

0.1082 .1102

.0831 .0704 .0616 .0563 .0486 .0436 .0310 .0227

Table VII



6 10 12 20

25 40 Mean

,2 nd 80°

0.062 .063 .062

0.697 .714


.743 .727

- 1 st *95°

v2nd *95°

v3rd **95°


0.17 .16 .16

1.47 1.54 1.65

9.5 15.3 18.0




.047 .033

.3rd 80°


3 5

„r1 st T 80°

o .


.787 .756 0.732


29.0 1. 85 1.49

14.0 1.60




solvents have led to the value 19 kcal. for the heat of activa­ tion of the reaction

This is shown in the graph Fig. V*,

The graph also shows, then,

that the heat of activation of the reverse of the above reaction is about


kcal. and that the measured heat of activation of

the disproportionation reaction is about 13 kcal., or only 5 kcal. greater, and considerably less than that for the original dis­ sociation reaction.

This suggests that the heat of activation

of the dissociation process as reported may be somewhat too high. Possible application of the magnetic method to this problem would give an interesting answer to the question. A bimolecular mechanism for the reaction may be proposed as fo


Here the normal radical with unpaired electron on the methyl carbon forms a bond with the carbon of the phenyl holding the methyl radical of the o-tolyl, with subsequent decomposition to give a molecule of o-tolyldiphenylmethane and a highly unsaturated o-quinoid structure which would undoubtedly polymerize rapidly.





.0 4 95

20 F IG . 6


4 0 HRS.


This mechanism and others will be further discussed when the ebulliometric data have been presented* Postulation of other reaction mechanisms presents the diffi­ culty that hexaphenylethane, for example, is much more stable in solution, and a para substituted radical having is much less stable*

Hence we assume that

ft hydrogens

(X hydrogens are

unusually vulnerable to reaction and that the lower reactivity shown by the o-tolyl as compared to p-tolyl is due to a steric effect*

Meta substituted radicals having

ft hydrogens are quite

stable in solution, presumably because the meta quinone structure cannot form*

Thus, the mechanism evidently must involve the

substituent group and the carbon to which it is bound, with the necessity of an

ft hydrogen for one mechanism in the case of

ortho and para substituents and probably another mechanism for the meta substituted radicals and ortho and para substituted varieties lacking

ft hydrogens.

The fact that di-o-tolylphenyl-

methyl disproportionates much more rapidly and at a lower tempera­ ture (10°C.) than any other radical studies supports the view expressed in the mechanism postulated* 3* The Sneetrographic Investigation* Table VIII shows the results of the micropho tome trie measure­ ments on the spectrograms of the solutions of o—tolyldiphenyl— methyl.

Since these values are proportional to the transmission

through the darkened negative, it will be seen that the recipro­ cals will give values proportional to the absorption at the various points.

The composite graphs of Fig. VII show the

Table VIII Microphotometer readings for the spectrogram of the solution 0*0007 m* in o-tolyldiphenylmethyl at 80° after 4*5 hours heating. Heading 3980 K 4065 4145 4225 4310 4395 4480 4565 4650 4735 4825 4910 5000 5090 5180 5270 5365 5455 5550 5645 5740 5840 5935 6035 6135 6235 6335 6435 6540 6640 6745

2i5 cm* 2*5 2*7 3.7 5.2 6*8

7.4 7.8 Reference line 10.7 12.4 14.8 15.3 31.3 36.0 37.0 44.6 47.6 47.8 47.9 47.9 48.0 48.0 48.1 48.1 48.1 48.0 45.5 13.4 5.6 4.0

100/Reading 40.0 40.0 37.0 27.0 19.2 14.7 13.6 12.8

9.4 8.1 6.8

6.5 3.2 2.8

2.7 2.2 2.1 2.1 2.1 2.1 2.1 2.1 2.1 2.1 2.1 2.1 2.2

7.5 17.9 25. 0

Absorption S p e c t r a of o-tolyldlphenylmethyl Undergoing Disproportion*tion. Cone. 0.0007 ■ ethane











Wave-length In A*U.

A 'jr o r y t lo n

d p o c tru n

o f o - to ly ld ly n e iv - lr a e th y l

O le p r o r > o r t lo n n . t l o : n ,

C one.

U n d e r c o lo g

0.03™ . e t h a n e

25 0

BO1 SO1 801 25' 5200


5300 W a v o - le n y t h



SlOO I n A .U .




absorption traces for two concentrations•

The more dilute solu­

tion shows structure toward the blue which is entirely blotted out with the more concentrated solution.

It originally showed

a transmission band near 4500 & which, upon heating, shifted with the reversible dissociation toward the red.

Prolonged heat­

ing caused this band to progressively lose its identity until, at the end of about 40 hours, at which point magnetic measure­ ments indicated that not more than


of the radical could re­

main, the band lost its identity entirely and cooling the solu­ tion to room temperature caused no substantial change in the transmission, showing that no reversible change was occurring. The more concentrated solution shows only a broad, more or less continuous band in the red region at room temperature before heating.

TJpon heating, the band contracts reversibly

toward the red and gradually shrinks in that direction as the heating progresses.

Cooling to room temperature after more

than 40 hours heating causes no significant change in the spectrum,

showing again that a reversible heat effect is no

longer obtained. These results seem to indicate that, while the colorimeter cannot be used as a measuring tool for radical solutions, it may yet be possible to work out a spectrophotometric mode of attack which will yield an easier, simpler procedure than the magnetic method. 4. The Ebulliometric investigation. Two solutions of nearly the same concentration were studied


with the ebulliometer already described*

The boiling point

differences (in terms of the measured potential, P) are recorded along with the time of observation in the Table IX*


each potential by 360 will give the actual temperature differ­ ence in degrees C*

The graph Fig* VTII shows a plot of the

temperature difference against time of boiling$ the apparent molecular weight is also shown* Discussion of Results*

It is evident from the graph that

the apparent molecular weight in solution at ebullition in­ creases rapidly, and this much more rapidly than the radical concentration decreases according to the magnetic measurements. Thus, at the end of about 1*6 hours in either case the apparent molecular weight has increased to something between two and four times its original value*

(In the case of the solution

0*0629 m. in ethane, there was also present a concentration of 0*0314 m* o-tolyldiphenylmethyl chloride, and the solution 0.0567 m. in ethane was only 0.0113 m. methyl chloride.)


magnetic data, on the other hand, show that over 80$ of the radical was still present in a solution originally 0.0724 m* in ethane (and 0*0209 m. in methyl chloride) after



at 95°C., or 15° higher than the temperature of the ebulliometer experiment. Thus the mechanism previously proposed for the reaction to explain the decrease in radical concentration comes into serious question

when the attempt is made to apply it to the

marked and rapid increase in apparent molecular weight at the

Table IX [ethane] = 0.0629 m. Time


hr. 0.30 .50 .70 .75

(137 yH V . 67 41 25



1.25 1.5 2.0

3.00 5.00 6.00

7.00 9.00 12.00

16.00 21.00

27.00 35.00 50.00


34 34 35 36 32 30 30 32 31 30 28 24 23 21 20

App. Mol. Wt. )

(279) 593 970 1530 1820 1120 1120

1090 1060 1190 1270 1270 1190 1230 1270 1360 1590 1660 1820 1910

[ethanej = 0.0557 m. Time hr. 0.25 .33 .50 .75 1.08 1.58 2.33 3.00 4.25 4.92 17.50 18.00 0


App. Mol,

(107 At 86

76 67 64 61 57 56 56 54 53 52 51

(287) 357 405 459 480 504 539 549 569 580 591 603



-4 0 0




HRS. 0.0629


20 F IG .8



beginning of the boiling point experiments.

It might be suggested

that the molecular weight increase is due to the fact that the residual methyl chloride reacts rapidly with the proposed in­ termediate to give a high polymer (m.w. ca 1300).

In that case,

with the 0.0629 m. solution, the expected P might decrease from 137^6.v. to about 95^a v .$ for the other the expected value would be a little higher due to the fact that a much smaller proportion of chloride was present.

Thus, in the one instance our mechanism

must explain a P only one-third as great, and in the other onehalf as great as is to be expected.

It is obvious that the

mechanism already proposed, occurring as rapidly as necessary to decrease P as observed, would reduce the susceptibility of the solute ethane to a point far lower than magnetic measure­ ments allow. A second mechanism which may be offered involves singleelectron bonds, which at present enjoy little prestige except in the case of the hydrogen molecule-ion:

HC— C H2 | CL L_


C — ur2 (I ^ cr

(p-C- (f C || JC


r i

c^ c / c

* 9