摘要:

JOHN G. ASTON 79 RESTRICTED ROTATION IN ETHANE BY L. J. OOSTERHOFF Received 2nd February, 1951 The electrostatic interaction between the methyl groups in ethane is in- vestigated. The basic idea is the same as in the investigation of Lassettre and Dean, but the calculations have been performed along different lines. The charge distributions of the C-H bonds are calculated with wave functions of the same type as those used by Coulson when discussing the dipole moment of this bond. By expanding the reciprocal distance, which occurs in the Coulomb interactions, in a Fourier series it proves possible to calculate directly the cos 34 term in the interaction energy, The cos 64 term proves to be negligible in comparison. The results of the calculations differ in detail from those of Lassettre and Dean, but the general trend is the same.With a purely covalent bond the electrostatic energy of the methyl groups depends only to a negligible extent on the angle of rotation 4. The right order of magnitude of the potential barrier is obtained if an ionic contribution of the type C-H+ is added, leading to a bond dipole moment of 0.7 D with the sense C-H+. Since the relation between the calculated dipole moment and the empirical values as derived, for example, by means of spectroscopy, is still vague, we consider i t premature to conclude that the electrostatic intei actions of bonds with fixed charge distributions are the origin of the potential barrier in ethane. Information about the magnitude of the potential barriers hindering the free rotation about single bonds has resulted from measurements of thermodynamic quantities such as specific heat, entropy, and chemical equilibrium constants.For ethane the specific heat measured at low temperatures, the entropy’ determined with the Fernst theorem and the equilibrium constant of the hydrogenation of ethylene indicate a potential80 RESTRICTED ROTATION barrier of about 2750 cal./mole, as was first suggested by Kemp and Pitzer-l where 4 is the angle of rotation about the C-C bond with 4 = o in the eclipsed or in the staggered configuration. The other terms in the Fourier expansion of the potential energy are usually considered to be negligible. This is in accord with the zig-zag configuration of paraffin molecules in the crystal- line state, with the preponderant occurrence of the rigid isomer of cyclo- hexane, the non-planar structure of cyclopentane,8 etc.Spectroscopic evidence also supports the assumption of a potential barrier of the said order of magnitude with the staggered position as the stable 0ne.4 Various theoretical explanations of the experimentally determined barrier have been suggested starting from a quantum mechanical analysis of the electronic structure of the ethane molecule. The standard methods of treating the electronic structure of saturated molecules like ethane, the H.L.S.P. method or the molecular orbital method, lead to a very small potential barrier.6, Refinements of the H.L.S.P. procedure were introduced by Eyring et aZ.7 Next to the 2s and z p orbitals he considers the influence of 3d and 4forbitals, which proves to be small.More important is the reson- ance of the usual structure with excited structures as, for example, those containing a bond between two hydrogen atoms belonging to different methyl groups, and a double bond between the carbon atoms. According to Eyring et al. these effects may lead to a barrier of the right magnitude, but with the opposed configuration as the stable one. Since now all the experimental evidence is in favour of the staggved form as the more stable, one has to conclude that either a more exact calculation will change the results of Eyring et al., or that still other effects may be large enough to give results in the opposite direction. Another way of attack has been followed by Lassettre and Dean8 and by Oosterhoff.s They calculate the Coulomb interactions of the charge distributions of the two methyl groups.The charge distributions are determined by using for each C-H bond a molecular orbital or valence bond wave function with a certain degree of ionicity. The present paper mainly deals with this method. It differs from the usual H.L.S.P. method by including ionic structures and taking into account the electrostatic interaction of overlap charges in different C-H bonds, which amounts to the same thing as including some of the higher order permutations. A quite different suggestion of Pitzer lo is based on the idea that the length of, for instance, the carbon-carbon bond in ethane is determined by the equilibrium of the interaction of the bonding electrons tending to shorten the bond and the repulsion of the other valence electrons of the carbon atoms.If the charge distributions on the carbon atoms deviate from the cylindrical symmetry, for example, because of a contribution of the d-orbitals to the C-H bonds, the resulting trigonal distributions will a Hassel and Viervoll, Arch. Math. Naturvidenskab., 1944, 47, im. 13. Hassel and Viervoll, Acta Chem. Scand., 1947, I, 149. 3Aston, Schumann, Fink and Doty, J . Amer. Chem. SOC., 1941, 63, 2029. Aston, Fink and Schumann, J . Amer. Chem. SOC., 1943, 65, 341 ; cf. Kilpatrick, Pitzer and Spitzer, J . Amer. Chem. Soc., 1947, 69, 2483. One is led to this value assuming a barrier of the shape v = v, cos 34, * (1) Probably the staggered configuration is the stable one. Kemp and Pitzer, J .Chem. Physics, 1936, 4, 749. * Smith, J . Chem. Physics, 1949, 17, 139. 5 Penney, Proc. Roy. SOC. A , 1934, 144, 166. 7 Gorin, Walter and Eyring, J . Amer. Chem. SOC., 1939, 61, 1876. 8 Lassettre and Dean, J. Chem. Physics, 1948, 16, 151, 553 ; 1949, 17, 317. IOPitzer, J . Amer. Chem. SOC., 1948, 70, 2140. Eyring, J . Amer. Chem. SOC., 1932, 54, 319. Oosterhoff, Thesis (Leyden, 1949).L. J. OOSTERHOFF 81 always give a larger repulsion if the hydrogens of one methyl group line up with those of the other methyl group than in the staggered position. The decreased repulsion in this configuration will allow the C-C bond to shorten a little, thereby increasing its binding energy and stabilizing the staggered form. Without an explicit calculation it is difficult to see whether this effect may account for the magnitude of the experimentally determined barriers.Method of Calculation.-In relating the experimental barrier in ethane to the electrostatic interaction of the charge distributions of the methyl groups, we assume that those parts of the charge distributions which contribute to the barrier do not overlap appreciably. In the calculation of the charge distribution in a C-H bond we make use of the wa.ve €unctions which were used by Coulson 11 in the discussion of the dipole moment of this bond. These functions are either localized two-centre molecular orbitals or electron pair wave functions of the H.L.S.P.-type including ionic contributions. They are built up of a carbon function #t and a hydrogen IS function #h: The carbon function, one of a set of four equivalent wave functions suited to form four tetrahedrally arranged bonds, is where that particular #22p function has been chosen which has rotational symmetry around the line joining the two bonded atoms.h? = 4#28 + * 1 / 3 $ z p The electron pair wave function is s = l$t#hdT* Coulson does not introduce a term with #h(~)#h(z) since the energy of a negative hydrogen ion and a positive carbon ion at the C-H distance of the normal C-H bond is supposed to be very high. Although we do not think this argument very convincing, since the situation indicated by #h(I)#h(2) is rather different from a negative hydrogen ion, we will likewise leave this function out of consideration, mainly because a really satisfactory investigation of the problem along the present lines would anyhow require a much more refined discussion of the individual bonds.The charge distribution of each electron can be considered as a linear combination of the three normalized charge clouds $t8, #&/s and h 2 , and is given by where a#t2 f h!%#hIs + Y#h2, . . ‘ (3) 2 S2 2 AS I I y=y2-* Using localized two-centre molecular orbitals it is possible, by adjusting the coefficients, to make the charge distributions nearly the same. This does not necessarily imply that the energies corresponding to these wave functions are also nearly equal, since the spatial correlations of the two electrons are always different. l1 Coulson, Trans. Faruduy Soc., 1942, 38, 433.82 RESTRICTED ROTATION We are only interested in that part of the Coulomb interaction of the methyl groups which changes on rotation about the C-C bond.Now the sum of the t,htz distributions of the three C-H bonds of one methyl group is a distribution with cylindrical symmetry about the C-C direc- tion. The interaction with these distributions can therefore be left out of the calculation. A t,hh2 charge distribution has spherical symmetry around the hydrogen nucleus. So in a calculation of the magnitude of the Coulomb interaction with this distribution we may reckon as if the total charge is concentrated at the hydrogen nucleus. More complicated are the calculations in which the overlap charges are involved. To a first approximation it is allowable to replace a charge cloud distributed according to $&/S by an equal charge placed in the centre of charge.For a further approximation it is necessary to consider also the higher moments of this distribution. Electrostatic Interaction of Two Charge Clouds .-The Coulomb interaction of two charge clouds with densities u, and u, is given by the integral L is the distance between a point in the first cloud and a point in the second one. i’ FIG. I. d -Co-ordinate system for the interaction of C-H bonds. The charge distributions which enter into the prcsent calculations always have cylindrical symmetry around a C-H bond. Therefore the centre of charge lies on the line through the carbon and the hydrogen nucleus. For the evaluation of integral (4) it is convenient to consider the reciprocal distance L-l as a function of the distance between the charge centres and the distances between the two above-mentioned points and the charge centres.In Fig. I the two carbon nuclei are shown at a distance d apart. The charge centres are lying at distances Y, and r2 from the carbon nuclei. Each charge centre serves as the origin of a co-ordinate system. The z-axis lies along the C-H line in the sense C-H. The x-axis lies in the plane through the C-C bond and the C-H bond, the positive sense pointing to the other methyl group. The y-axis is fixed by a right- handed co-ordinate system. The two co-ordinate systems can be dis- tinguished by the indices I or 2. + is the angle of rotation around the C-C bond ; + = o if the two C-H bonds lie in one plane. We call L the distance between the points (xl, yl, 2,) and (x2, y2, +), Lo being the distance between the points (0, 0, 0) and (0, 0, 0).Assumrng tetrahedral bond angles we get We will proceed in the following way. 16 9 3 3 9 Lo2 = R2 - -ylv2 cos + ; R2 = y12 + Y~~ + Zr,d + %,d + s l y , + d2. ( 5 )L. J. OOSTERHOFF 83 We will now develop L-l into a power series in x,, yl, 2, and x,, ye, z,. This may conveniently be written as The Greek letters, which have the same meaning as the Latin letters, have been used to bring out explicitly that the differential quotients refer only to the co-ordinates in L. The index o indicates that the differential quotients have to be taken at the origin. The series (6) may be used in (4) on conditions of non-overlapping, etc., which we will assume to be fulfilled. Inserting this series in ( 4 ) and introducing the abbreviations (I, 2) means that the same term as the foregoing has to be added but with the indices I and z interchanged.Ql(Jl2 - +p12) is the only non-zero component of the quadrupole moment of the charge distribution I. Q1 (5,3 - a<,p12) is the only non-zero component of the octupole moment. The dipole moment Qltl and the other components of the quadrupole moment and the octupole moment are zero in consequence of the choice of the origin in the centre of charge and of the cylindrical symmetry of the charge clouds. Fourier Expansion of Reciprocal Distance .-It has already been said that we are only interested in that part of the Coulomb interaction which changes on rotation of the methyl groups around the C-C bond.In consequence of the trigonal symmetry of the methyl groups, the inter- action energy may. be expected to depend on the angle of rotation $ according to provided that the zero of $ has been suitably chosen. This formula suggests that it will be advantageous to insert in the energy expression (7) a Fourier development of the reciprocal distance, for in the resulting development only’ the coefficients of cos 34, cos 64, etc., need to be con- sidered. The other terms cancel when adding together all the interactions between the charge distributions of the two methyl groups. Expanding the reciprocal distance V = V o + QV, cos 34 + QV, cos 64 + . . ., . - (8) 16 = (R8 ---y1y2 cos 4 LO 9 in a Fourier series yields the result : € 0 = I, €, = 2, m = 1 , 2 , . . . .84 RESTRICTED ROTATION The coefficient of cos m$ is closely related to the Legendre functions of the second kind.12 In fact, this coefficient may also be written as In the present calculations the series for the coefficients of cos 3+ and cos 64 proved to converge rapidly, and even the term with cos 64 is very small with respect to the cos 34 term.Inserting the expression (9) in (7) yields the desired development of the Coulomb interaction energy of two charge clouds. The interaction of point charges can be calculated easily. For the computation of the interaction energy of the higher moments the derivatives of Lo1 are re- quired which involves a somewhat lengthy but quite feasible numerical calculation. Charge Distributions .-The charge densities u have been calculated using Slater functions for the carbon $2s and $21zp functions which figure in $t.Coulson and Duncanson l3 have shown, in a discussion of the momentum distribution in CH,, that there are reasons for increasing Slater's screening constants (I for hydrogen, 1.625 for carbon 2s and z p ) by a factor w = 1.1. Following this suggestion we get the set of functions (in atomic units) $5 is a IS hydrogen function. 0 = 1'1 The length of a carbon-hydrogen bond is taken as 1.093 A. of the overlap integral S = 0,636. I,ht$h,/s with respect to the centre of charge are The value The moments of the overlap charge The distance of the charge centre from the carbon nucleus is 0-794 A. The moments of the +t2 distribution are - - } (13) 5 2 = 0-213 A2, 53 =- 0.082 A3, p2 = 0-307 A2, 5p2 = 0.021 A3, 5 2 - 1 2 p 2 = 0.059 Hi2, 5 3 - $ 5 ~ 2 = - 0.114 A 3 .- - The distance of the centre of the +t2 distribution from the carbon nucleus is 0-370 A. Re s w l t s The quantities 01, B, y measure the chance tor one electron t o be found in one of the three charge distributions I,hr2, $,I,hh/S or In Fig. 2 they are plotted as a function of A, which is a measure of the degree cjf ionicity of the C-H bond. The bond dipole moment p is plotted against h in Fig. 3. This dipole moment was calculated, according t o Coulson, for a charge distribution consisting of the positive charge of the hydrogen nucleus, an equal charge at the carbon nucleus and the charge cloud of the two electrons involved in the formation of the bond. In order t o facilitate comparison of our results with those of Lassettre and Dean the only non-zero components of the quadrupole moment P and the octupole 12 Whittaker and Watson, Modern Analysis (Cambridge, 1946), 4th edn., 13 Coulson and Duncanson, Proc.Camb. Phil. SOC., 1942, 38, 100. p. 316.L. J. OOSTERHOFF 85 moment 0 of the bond are also plotted in Fig. 3. by the formulae The three moments are defined 1 p = ZQ + 5ud7, s The summations refer to the two equal positive charges of the nuclei. integrations extend over the charge cloud of the two electrons. the co-ordinate system has been chosen midway between the two nuclei. units are based on a unit of charge equal to 10-l~ electrostatic unit and the as a unit of length. The The origin of The 1 2 3 4 , - A FIG. 2.-a, p, and y as a function FIG.3.-Bond dipole moment of the ionicity A. p, quadruple moment P and octupole moment 0 as a function of the ioni- city A. The results of the calculation of the electrostatic interaction of the methyl groups are plotted in Fig. 4. The dotted line a gives the value of the barrier height if the charge clouds had been concentrated as point charges a t their centres of charge ; b and c represent the additional interactions, caused by the quadru- pole moment and the octupole moment of the overlap charges. The total inter- action, including the three effects, is given by the curve marked V,. The co- efficient of 4 cos GI$, viz. V,, proved to be negligible as compared with V 3 (about one thousandth of V,). * It is gratifying that the main contribution to the cos 34 term in the inter- action energy results from the interaction of the point charges.The con- tribution of the terms involving the quadrupole moments is very low but the contribution of the octupole moments is rather high. So we do not feel sure that the hgher moments will give a negligible contribution. We do not expect, however, that the values of V , given in the Table will change very much if they are calculated exactly, using the same charge distributions. The values of V,, calculated according to the method of Lassettre and Dean,8 and indicated by V8’ are plotted as a dashed line in Fig. 4. In this method the charge distributions in the G - - H bonds are not split up into separate charge clouds. The interaction of the bonds is calculated using the dipole momcnt, quadrupole moment, etc., of the C-H bond, as defined by (14).From the calculation of the electrostatic interaction for a great number of values of 4, Lassettre and Dean derive the formula (converted to the units used in this investigation), With our method we derive the formula V3’ (cal./mole) = zg7.5p2 + 339.9 p P + ~ I I - ~ P ~ . . - (15) V,’ (cal./mole) = zg6-0p2 + 343-9 p P + 315.9 P2 - 736.7 PO, (16) * The limiting value of V , for A 3 00 is 4706 cal./mol.86 RESTRICTED ROTATION which agrees very well with (15) except that the dipole-octupole interaction has been included. This latter term is only of importance for high values of A corresponding to a high degree of ionicity of the bond. From Fig. 4 i t can be seen that the agreement between the values of V 3 and V3' is not very satisfactory, but the general trend is the same.Discussion From the graph of V , in Fig. 4 one might conclude that the electro- static interaction of the methyl groups can be used as an explanation of the potential barrier in ethane. In that case one has to assume a large ionicity of the C-H bond. In chemistry the C-H bond is usually con- sidered to be a typical covalent bond and, on the analogy of covalent diatomic molecules, one is inclined to attribute at most a small dipole moment to this bond. Coulson 11 pointed out, however, that the relation between bond character and dipole moment of the C-H bond is quite different (cf. Fig. 3). A small dipole moment is connected with a mixture of the covalent wave function and an ionic wave function of the type C-H+.FIG. +-Potential barrier in ethane as a function of ionicity A. a, contribution of charge clouds concentrated a t their charge centres ; b, con- tribution of quadrupole moments ; c. contribution of octupole moments ; V,, total height of potential barrier ; V,', height of barrier according to (16). The main arguments in favour of a small value of the dipole moment of the C-H bond, say smaller than 0.7 D, have, for the greater part, been derived either from an analysis of measured dipole moments of molecules or from the intensities in vibration spectra (cf. Gent 1 4 ) . A closer examination of the meaning of the bond dipole moments, derived from these sources, makes it doubtful whether they refer to the same quantities as the moments calculated according to Coulson.Therefore we do not think that a covalent character of the C-H bond can be excluded. Lassettre and Dean * assume the value of 0.4 D of the C-H bond dipole moment to be correct. With a small dipole moment the main contribution to the potential barrier calculated with (15) or (16) comes l4 Gent, Quart. Rev., 1948, 2, 383.L. J. OOSTERHOFF 87 from the quadrupole-quadrupole interaction. The calculated value of the potential barrier being too small, Lassettre and Dean have reversed their reasoning and derive an empirical value of the quadrupole moment from the known value of the potential barrier. The ratio of the empirical quadrupole moment t o the calculated one is about the same as the ratio between the accurate value of the quadrupole moment of the hydrogen molecule and that calculated with the crude molecular orbital method. Along these lines Lassettre and Dean come to the conclusion that the quadrupole-quadrupole interaction generally can be considered as the origin of the potential barriers hindering rotation. In view of the results represented in Fig. 4 we do not think that the method of calculation according to formula (15) or (16) is satisfactory. Therefore, since in our opinion it is at present not justifiable to use the bond dipole moment as a guide in selecting the most appropriate value of A, we prefer the conclusion that it may be useful to reckon with the possibility that the potential barrier in ethane is due to electrostatic interaction of C-H bonds, which have a rather high degree of ionicity. On the other hand, it should be borne in mind that the C-H bond may be a nearly covalent bond. In any case a more refined analysis of the electronic structure of ethane is necessary in order to arrive at a definite conclusion. The author wishes to express his thanks to Prof. H. A. Kramers for his interest and advice during the course of this research, to Mr. J. H. Kruizinga for his assistance in several mathematical derivations and for his indispensable help in carrying out many of the numerical cal- culations, and to Mr. J. A. van der Heiden fclr his help in preparing the figures. This paper is published by permission of the Management of the N.V. de Bataafsche Petroleum Maatschappij, The Hague. Koninklijke /Shell-Laboratoriuun, A m s tevd am.

ISSN:0366-9033    

DOI:10.1039/DF9511000079

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

L. J. OOSTERHOFF THE ISOMERS OF CYCLOHEXANE BY P. HAZEBROEK AND L. J. OOSTERHOFF Received 2nd February, I 95 I The equilibrium between the rigid and the flexible isomer of cyclohexane is calculated. The results indicate that a t ordinary temperatures the flexible isomer occurs a t most to a very slight extent. At higher temperatures it may occur in concentrations which are large enough to allow an experimental con- firmation. Cyclohexane resembles in many respects the saturated paraffinic hydrocarbons. This has led to the hypothesis that the valence state of the carbon atom in cyclohexane is similar to the valence state in saturated molecules where, according to Van't Hoff and Le Bel, one has to assume a tetrahedral arrangement of the four valencies. On the basis of this hypothesis, Sachse derived by geometrical considerations that two isomers of the cyclohexane molecule might be possible.The angles between two consecutive carbon-carbon bonds being less than IzoO-in our case 1 0 g O 28' 16"--it is possible to construct an infinite number of configurations which can be divided into two groups. In one group the six-membered ring can pass continuously through an infinite number of configurations. Sachse, Rer., 1890, 23, 1363 ; 2. physik. Chem., 1892, 10, 203.88 ISOMERS OF CYCLOHEXANE The other group consists of only one configuration which cannot change into the other configurations without a temporary alteration of the valence angles. The first group may be indicated as the flexible isomer, the second as the rigid isomer.Current names are boat and chair form, but these are less appropriate since the boat configuration is only one of all the possible configurations of the flexible isomer, whereas the pictures which are sup- posed to suggest a similarity with a chair do not emphasize the high sym- metry (D,d) of the rigid isomer. The ideas of Sachse, which were amplified by Mohr,2 were very fruitful in many chemical investigations of Boeseken and Hermans 4 into the properties of substituted cyclohexanes. A more definite proof of the structure of the cyclohexane molecule was based on the analysis of infra- red 4 and Raman spectra 6# and of the entropy and specific heat in the ideal gaseous state.'* These investigations proved the preponderant occurrence of the rigid isomer. The results of X-ray analysis of the crystal and electron diffraction of the gaseous molecules,D although less convincing, point in the same direction.The study of cyclohexane de- rivatives and other molecules containing a six-membered ring has also led to the conclusion that the ring preferably has the rigid s t r u c t ~ r e , ~ although exceptions are known (cf. Dallinga II). Cyclohexanedione-I : 4, for example, has a dipole moment of 1.3 D, which indicates that at least in part the flexible structure is present.lO, l1 At first sight, the higher free energy of the flexible isomer is some- what astonishing. The internal motion may be expected to add to the free energy the (negative) contribution of a fully excited degree of freedom, whereas the corresponding degree of freedom of the rigid form will be a partly excited vibration.Besides the rigid form has the higher sym- metry and therefore, also for this reason, would have a higher free energy. Thus one is led to the conclusion that there will be a difference in potential energy making the flexible isomer less stable. It is a plausible suggestion to hold the interactions between the C-H bonds responsible for this potential energy difference, as will be clear from a closer inspection of the geometry of both isomers. In Fig. I the rigid form and two configurations of the flexible form are pictured. The rigid form has a threefold axis together with the other symmetry ele- ments of the symmetry group D a d . The boat configuration has the sym- metry Czv, one twofold axis, and two planes of symmetry containing this axis.The stretched configuration of the flexible form has the symmetry V , three mutually perpendicular twofold axes. Tie other configurations of the flexible isomer have a twofold axis only. From the Fig. I it can be seen that in the rigid structure all pairs of consecutive CH, groups are in the staggered position, whereas in the boat configuration two pairs are in the eclipsed and four pairs are in the stag- gered position. I f , following Pitzer, we assume an energy difference of 2750 cal./mole between the staggered and the eclipsed position, the energy difference between the rigid and the boat configuration is 5500 cal./mole. The potential energy of the other configurations is more difficult to determine, but it will be shown that the potential energy is nearly a con- stant as far as the interactions of consecutive CH, groups are concerned. Mohr, J .prakt. Chem., 1.918 (z), 98, 322. Boeseken, Bull. SOC. Chzm., 1933, 53, I. Hermans, Rec. trav. chim., 1938, 57, 333. 6 Kohlrausch and Wittek, 2. physik. Chem. B, 1941, 48, 177. 6 Gerding, Smit and Westrik, Bec. trav. chim., 1942, 61, 561. 7 Aston, Schumann, Fink and Doty, J . Amer. Chem. Soc., 1941, 63, 2029. 8 Pitzer and Spitzer, J . Amer. Chem. Soc., 1947, 69, 2488. 9 Hassel and Vieroll, ArcJz. Math. Naturvidenskab., 1944, 47, no. 13. 10 Le Fkvre, Dipole Moments, (London, 1948)~ 2nd edn. *I Dallinga, Thesis (Leyden, 1951). Hassel and Viervoll, Acta Chem. Scand., 1947, I, 149.P. HAZEBROEK AND L. J. OOSTERHOFF 89 Pitzer,s in the interpretation of the experimentally determined entropy of cyclohexane reduced to the ideal gas state, only takes account of the boat configuration when calculating the entropy of the flexible isomer. Therefore the difference between the entropies of the two isomers is mainly due to the different symmetry.This difference, which is equal to R In 3, is large enough to conclude, on the basis of the experimental value of the entropy, that the rigid form occurs predominantly, at least at ordinary temperatures. If one takes account of the other configura- tions of the flexible isomer as well, the calculated value of the entropy of this isomer will be increased, which strengthens the argument of Pitzer. \/ 0 - 0 8 - y Sji-etched con fiy urabions. Boa1 cotrfigura~;on~ R/>l’d I5omer Flexi6ie homer FIG.I.-The isomers of cyclohexane. It would not be correct, however, to set the free energy difference equal to 55oo-RRT In 3 and to conclude that this value is in accord with the slight occurrence of the flexible isomer, since for a calculation of the equilibrium between both isomers one has to consider accurately the effect of the internal motion. In this paper we will indicate a method enabling the calculation of the free energy contribution of the internal motion of the flexible isomer, from which an estimate of the equilibrium ratio of the two isomers will z m , &Jq fQ9-5 be derived. Method Of CalCUlatiOn.-GEOMETRICAL 31 13 CONS ID ERAT I 0 N S .-h the description Of the possible coilfigurations of the cyclohexane molecule we will provisionally leave the hydrogen atoms out of consideration and fix our attention on the relative positions of the carbon atoms.Thus we have to treat FIG. 2. the problem of an appropriate description of the configurations of an equilateral and equiangular hexagon with angles of 1 0 g O 28’ 16”. The length of a side we put equal to I. The six sides are represented by six unit vectors, a,, u2, . . . a, which fulfil the equation 34 3 + Zai = 0. . * ( 1 ) + --+ Scalar multiplication with successively a, . . . a6 leads to the equations 6 1 Sij 7 0 J = I , * . . . 6, . * (2) i - 090 ISOMERS OF CYCLOHEXANE --++ with Sij = ( ~ i , aj). According to the suppositions Si; = I , Sii = 1/3, when i and j refer to consecutive vectors. In consequence of these equations and Sij = SF, there remain nine unknown Sij, which have to be determined by the six equations ( 2 ) and by additional equations which may be derived in the following way.Four vectors ai are always linearly dependent, so that we may write, e.g., --+ Sll Sl, Sl3 S2 1 ( 3 ) . . - - s44 - 1s41* ! I f and similar equations for all other four-rowed determinants contained in the matrix Sij. Of the equations ( 3 ) only three need be added to the six equations ( 2 ) as the other ones in the case of the flexible isomer do not restrict the consequences of the eight equations mentioned. We will now try to find a variable on which the Sij will depend in a symmetrical way. From equations ( 2 ) and the numerical values already mentioned it can be derived that s 1 3 = s46 ; S14 = - Q - S1, - S24 ; s ---6-s -s s24 = s51 ; s35 = s 2 S i and S25 =- Q - S24 - S35 ; a6 - 3 a5 46.With the substitutions 3S1, = 5 - I ; 3S,, = 7 - I ; 3S,, = 5 - I eqn. (3) can be written and two corresponding equations which can be obtained by cyclic per- mutations of [, 9 and <. Beside the trivial solution f = 9 = 5 = 0, which corresponds to the rigid structure, these equations possess a one-dimensional manifold of solutions, corresponding to the flexible isomer. Subtracting eqn. (4) in pairs leads, after some transformations, to the equations t 2 T 2 - 24(S2 + q2 - 677) - 32([ + 9) = 0, - * (4) and Eqn. (5) suggests the introduction of variables s and 0 by the substitutions 7 f q = ss cos 3e cos 8, or 5 = - s p - 2 cos2 8), 1 I I 5 = - s[J - 2 cos2 8 + - ( ;.,I*} - - = - s[: - 2 COSZ (8 - y ) 1.J The value of s follows from (6), which after substitution of (8) reads Sa - -cos2 38 + 36s - 32 = 0. . a (9) To a first approximation s = 8/9 and a more accurate value is 8 8 s =- 9 (I + T-COS2 39) . .P. HAZEBROEK AND L. J. OOSTERHOFF 91 The variable 6 is appropriate to define the possible configurations of the flexible isomer in a symmetrical way. The value 6 = o corresponds with a configuration of symmetry F' where the distance between carbon atoms I and 4 is as large as possible and which accordingly may be in- dicated as a stretched configuration (see Fig. I). 6 = 7~/2 corresponds w-ith a boat configuration, where the ca.rbon atoms I and 4 have the closest approach to one another, 8 = 7~ corresponds again with a stretched configuration which is the optical isomer of configuration 0 = 0, etc.POTENTIAL ENERGY.-For a calculation of the potential energy of the isomers of cyclohexane it is necessary to make an assumption about the interaction of the CH, groups. It looks plausible to assume an inter- action of the same kind as between the CH2 groups in %-butane and other straight chain paraffins for which Pitzer suggests a value of 3600 cal./mole for the potential barrier. On the other hand, Spitzer and Huffman,14 on account of heat of combustion data of cycloparaffins, consider even a value of 2750 cal./mole too high. This latter value has been used by Pitzer in considering cyclic molecules. Provisionally we will also assume a value of 2750 cal./mole. Besides these interactions between neighbouring CH, groups, there may also occur important repulsions between H atoms linked to opposite carbon atoms, since in a boat configuration the distance between two of these hydrogen atoms (1.84 A) is much lower than the sum of the van der Waals radii (2.4 A) according to Pauling.ls To begin with, however, we will restrict ourselves to the first type of interactions for which we will assume the formula v = I375 (cos 3 41 + cos 348 + cos 343).- * (11) + - - t 4 - P Defining the angle 43 as the angle between the vectors [a,, a,] and [%, as] for iqstance we find and similar formulae for the other angles. The total energy according to formulae (11) is The rigid isomer is characterized by [ = q = t = o . In this case we get the result For the flexible isomer we have to make use of (8) and (10) leading to Y (rigid) =- 3 x 2750.- (14) 95 7 T;.' (flexible) = - 81 x 2750 + x 2750 (r - cos 68). . (15) The energy of the flexible isomer is seen to be nearly constant, the term proportional to cos 6 8 not exceeding RT at room temperature. This picture may change if we include the repulsive interactions between the hydrogen atoms linked to opposite carbon atoms. These repulsions, which are most pronounced in the boat configurations and, to a less extent, in the neighbouring configurations, will increase the l2 Spitzer and Huffmann, J . Amer. Chem. SOL, 1947, 69, 211. 13 Pauling, The Nature of the Chemical Bond (1945), 2nd edn.92 ISOMERS OF CYCLOHEXANE potential energy of the flexible isomer at values of 8 equal to 30°, goo, 150°, etc.The magnitude of V , is difficult to assess. According to calculations of Barton l4 it may be as high as several kcal./mole. In the free energy calculations we will consider also the effect of a term like (16). MOMENT OF INERTIA.-For a calculation of the moment of inertia of the internal motion it is advantageous to introduce cylindrical co-ordinates to mark the position of the atoms. The twofold axis is taken as z-axis, Y measures the distance to this axis, and x the azimuth. The centre of gravity of the molecule is taken as origin. In order to study the internal motion separately, unperturbed by the overall rotation, the moment of momentum of the molecule has to be zero. This condition is expressed by So we may add to (15) a term of the type V,(I - cos 68) * (16) the summation extending over all atoms.calculated from The moment of inertia can be I = z m { ( - ) dz 2 + ( $ ) 2 + Y 2 ( g ) 2 } . d9 Using these formulae we find PARTITION FUNCTION.-In consequence of the nearly constant value of the moment of inertia the calculation of the partition function of the internal motion of the flexible isomer is particularly simple. In the case of a constant value of the potential energy the partition function of the completely free rotation is given by I g 39-6 x I O - ~ O + 0.4 x I O - ~ O cos 68 g.cm.2 . * (19) 277 (2rIkT)+ Qj=; 7 ( 2 0 ) where U; is the symmetry number of the internal motion. As all the distinguishable configurations are already represented by the &values in the range o < 9 Q r / 3 , ui = 6.Now if we add to the constant potential energy a term V,(I - cos 68) the value of the partition function Q can easily be deduced from the tables of Pitzer and Gwinn.lS In these tables R In (QlQj) is given as a function of Qj and of the ratio of the height of the potential barrier to RT. Results Equilibrium Concentrations of the Isomers of Cyclohexane .-For a calculation of the equilibrium concentrations of the isomers in gaseous cyclo- hexane we have to estimate the contribution to the partition function of the rigid isomer of the vibration, which we will consider as the counterpart of the internal motion of the flexible isomer. From a table of calculated vibration frequencies of the rigid isomer (Ramsay and Sutherlandls) we will select the vibration with the lowest frequency (w = 206 cm.-]) which will give an upper estimate for the partition function Qw and hence for the concentration of the rigid isomer as far as this vibrational contribution is concerned.Since the symmetry number of the rigid isomer is q = 6 and that of the overall rotation of the flexible isomer is U, = 2 the ratio of the number of mole- cules of the flexible isomer N , to the umber of molecules of the rigid isomer N is given by l4 Barton, J . Chem SOC., 1948, 340. 15 Pitzer and Gwinn, J . Chem. Physics, 1942, 10, 428. 16 Ramsay and Sutherland, Proc. Roy. SOC. A , 1947, 190, 235.P. HAZEBROEK AND L. J, OOSTERHOFF 93 where U is the potential energy of the flexible isomer, without the (I - cos 68) term, minus the potential energy of the rigid isomer.The percentage of flexible isomer of cyclohexane present in equilibrium with the rigid isomer is plotted in Fig. 3 against the absolute temperature for different values of v6. The dotted curve refers to a lower value of U (1900 X 148/81) which is taken into consideration following a suggestion of Spitzer and Huff mann.12 Calculations along similar lines have been performed for cyclohexanedione- I : 4. The results indicate that the stability of the flexible isomer and the magnitude of the dipole moment can be explained, assuming that the interaction between the 6 0 group and the CH, group is the same as that encountered in acetone. FIG. 3.-% flexible isomer in cyclohexane as a function of temperature for two U = 1900 x 148/81 cal./mole v 6 = 3600 cal./mole. values a t V , : and for Discussion In view of the uncertainties in the potential energy difference between the rigid and the flexible isomer of cyclohexane, the results, plotted in Fig. 3, should be considered as giving only a rough indication of the con- centrations of the flexible isomer which might occur. These results suggest the possibility of an experimental study of the flexible isomer of cyclohexane at temperatures above room temperature. A closer study of the isomers of cyclohexane and of cyclohexane derivatives will contribute not only to a better understanding of the behaviour of these molecules, but also to the solution of the problem of the forces hindering free rotation. The authors wish to express their thanks to Prof. H. A. Kramers for his interest and advice during the course of this research, to Mr. J. H. Kruizinga for his assistance in several mathematical derivations and for his indispensable help in carrying out many of the numerical calculations, and to Mr. J. A. van der Heiden for his help in preparing the figures. The paper is published by permission of the Management of the N.V. De Bataafsche Petroleum Maatschappij, The Hague. KoninklijkelShell-Laboratorium, N . V . De Bataafsche Petroleum A msterdam. Maatschappij, The Hague .

ISSN:0366-9033    

DOI:10.1039/DF9511000087

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

MOLECULAR CONFIGURATION AND HYDROCARBON REACTIVITY BY A. R. UBBELOHDE AND J. C. MCCOUBREY Received 21st February, 1951 Various modes of ‘ I action at a distance ” in saturated hydrocarbons are discussed in relation to the comparative reactivity of homologues and isomers. Electronic effects and the influence of vibration coupling on reactivity are summarized. Detailed consideration is given to the average degree of coiling of n-paraffins . Experimental results quoted include infra-red and Raman measurements on crystal and melt, molar volumes of the liquids at the boiling points and critical temperatures, entropies of vaporization of the liquids, viscosities and temperature coefficients of viscosities of the vapours, thermal conductivities of the vapours and the second virial coefficients of the vapours.These results show that in the vapour phase flexible polymethylene mole- cules are coiled till the space occupied is about the same as for the corresponding isoparaffins. A moderate degree of uncoiling takes place on forming the liquid, but on average the fully stretched molecules are only found in the crystal. Various consequences of this molecular coiling for the reactivity of paraffins incIude steric and energy transfer effects. In hydrocarbon chemistry, processes such as bond rupture are formally similar at various parts of the molecule, and are also similar in homologues and isomers. This gives particular importance to detailed examination of the ways in which remoter parts of a molecule can affect the reactivity of specific bonds. In polpethylene hydrocarbons, two modes of “ action at a distance ” have previously been discussed in some detail in the light of the quanti- tative evidence available at the time.l To summarize these two modes briefly : (i) With reference to electronic action at a distance, it has been noted that in a series of 35 isomeric nonanes the molecular refractivity shows a progressive drift from about 43-9 units for the least branched to about 43.0 units for the most branched nonanes.However, electronic influences on bond strength fall off quite rapidly as the bond becomes more remote. For example, there is a constant increment to the heat of formation for each CH, group added beyond n = 4, showing that the end CH, groups have little effect on the more central bonds beyond this order of remote- ness.Again, ionization potentials of n-paraffins tend to a limiting value beyond about n = 4.2 (ii) Owing to the similarity of fundamental vibration frequency of the C-C bonds, the skeletal vibrations of the carbon atoms in a polymethylene hydrocarbon are coupled.8 This vibration coupling has a number of con- sequences. In polymethylene hydrocarbons : (a) Skeletal vibrations of longer wavelength than the fundamental are appreciably excited thermally at substantially lower temperatures than the C-C vibration in ethane, for which the characteristic temperature is Ubbelohde, Proc. Roy. SOC. A , 1935, 152, 361 ; Rev. Inst. Franpis Pitrule Honig. J . Chew. Physics, 1948, 16, 110. For experimental evidence, cf. Brown, Sheppard and Simpson, Faraday Ann. Combust.Liquides, 1949, 4, 488. SOC. D ~ S C U S S ~ O ~ S , 1950, 10, 000. 94A. R. UBBELOHDE AND J. C. McCOUBREY 95 around 8 = 1300~ K. The average zero point energy per C-C bond de- creases and the average vibrational energy increases at ordinary reaction temperatures, with increasing n. (b) Coupled vibrations can act at a distance by transmitting energy t o any part of the coupled system. But the coupling is interrupted whenever there is a substantial break in the vibrational system, as at a double bond, or where C is joined to a group such as NHa or C1. Activation processes can only draw readily on the reservoir of vibrational energy if the bond forms part of a strongly coupled system. Thus the pyrolysis of the C-X bond in C,H,,+,X, where X is a halogen or NHa or OH, or of should show quite a different dependance on n from the pyrolysis of hydrocarbons.Evidence has recently been obtained of activation in- volving coupled vibrations in the pyrolysis of n-~araffins.~ Molecular Coiling and Hydrocarbon Reactivity.-A third mode of action at a distance is possible in flexible hydrocarbons. The idea is that if these coil up so as to bring remoter groups into the neighbourhood of the bond whose reactivity is under consideration, the van der Waals forces can : (a) smudge the quantum levels of the bond by providing a polariz- able medium with variable interaction according to distance ; (b) affect transition probabilities in vibrational and electronic excitation of the bond. At the time this suggestion was first put forward, only indirect estimates could be made of the degree of coiling of homologous normal paraffins. This gap has since been filled from a number of sources of evidence. (I) Infra-red and Raman measurements on crystal and melt of normal paraffins show that whereas the molecules are stretched planar zig-zags in the crystal, in the melt crumpling takes place to form one or more “ rota- tional isomers ”. The Raman spectra of liquid n-hydrocarbons C 4 4 , are found to con- tain more lines than can be interpreted on the basis of a single configur- tion .s AH for a change from one rotational isomer to the next on progressive coiling is estimated to be about 500 cal./mole.In the vapour the con- figuration is probably similar, from the similarity of the infra-red absorp- tion spectra, which do not, however, permit quantitative comparisons.It should be noted that second virial coefficients for the hydrocarbons indicate some temporary dimerization in the vapour phase.6 (2) Pyknometric measurements of molecular volumes have shown that the n-paraffins occupy practically the same volume in the liquids as do the corresponding isoparaffins at the respective boiling points.’ Illus- trative data are recorded in Table I for C, and Cs ; much the same con- clusions have been reached with the intervening paraffins. Ratios of the critical volumes show the same effect. From the similarity in volumes and the similarity in entropies of vaporization it seems likely that the n-paraffins are “bunched” or “ coiled ” in both liquid and gas phase, to an extent such that they occupy about the same average space as the corresponding isoparaffins.(3) Measurements have been made on the viscosities and the tem- perature coefficients of viscosities of vapours of paraffin homologues and isomers.’ The results again show that the molecules of the %-paraffins must, on average, be quite highly coiled in the gas phase. Some of these features are illustrated in summarized form in Table 11. C,.H,,CH = CHC,H,, These are further discussed below. Cf. Ingold, Stubbs and Hinshelwood, Proc. Roy. SOC. A , 1950, 203, 486. Sheppard and Szasz, J. Chem. Physics, 1948, 16, 704 ; 1949, 17, 86 ; Rank 6 Hirschfelder, McClure and Weeks, J . Chem. Physics, 1942, 10, 201. 7 McCoubrey, McCrea and Ubbelohde, J . Chem.Soc., (in course of publication). and Axford, J. Chem. Physics, 1949, 17, 430.96 CONFIGURATION AND REACTIVITY TABLE I.-RATIOS OF MOLAR VOLUMES AND ENTROPIES OF VAPORIZATION 13 Paraffin n-Butane : isobutane . n-Pentane : isopentane . n-Hexane : isohexane (di-iso- %-Octane : iso-octane (2 : z : 4- %-Octane : iso-octane (di-iso- ProPYl) - trimethyl pentane) . butane) . Molar Volumes Ratio at b.p. - 1.0035 - 1.0148 - Ratio of Entropies of Vaporization - 1.0262 1'0532 1.052 Molar Volumes Ratio at Critical Temperatures 1.004 1'01 I 1'028 1.017 TABLE II.-RATIOS OF VISCOSITIES OF ISOMERIC PARAFFINS Normal: isoButane . Normal : isoPentane Normal : isoHeptane : (z : z : 3-trimethyl butane) . Normal : iso-Octane : (2 : 2 : 4-trimethyl pentane) . . . TOK 291.1 373'1 393'1 298.1 373'1 380.1 398.1 416.9 436'5 398.1 4 16.9 436.5 Ratio 0.993 1'000 I '000 0.973 0.978 0.939 0.938 0'945 0.95 1 0.955 0.954 0.955 739 947 998 676 ,341 729 763 800 83 9 724 757 785 Facts of importance for the reactivity of hydrocarbons can be obtained from Table 11.In the simplest form of kinetic theory of momentum transfer, the molecular u diameter is related to the viscosity by the equation, Inspection of the ratios of viscosities shows that at any rate so far as momentum transfer is concerned, the n-paraffin molecules have an average ma with respect to molecular collisions, which is only from 0-5 % larger than that of the branched isomers. This must imply that the normal paraffins behave predominantly as coiled molecules in the gas phase. A similar conclusion has been put forward from more restricted data on diffusion of the polymethylene chains C,,-C,, in air.s A further fact from Table I1 is that no large temperature coefficient is observed in the ratio of viscosities. This indicates that an explanation of the observed coiling in terms of repulsion potentials between 8 Molar volumes, McCoubrey, McCrea and Ubbelohde, ref.(7). Critical volumes, quoted by Partington, An Advanced Treatise on Physical Chemistry, VoE. I (Longmans, 1949) p. 645. Entropies of vaporization, Nut. Bur. Stand., 1947, Project 44. =aa = $1/(K/N.r) . 1/(MT)/q = 5.694 x I O - ~ ~ 1/(MT)/q. (1) Bradley and Shellard, Proc. Roy. Soc. A, 1949, 198, 239.A. R. UBBELOHDE AND J. C. McCOUBREY 97 adjoining CH, groups lo cannot form a complete theory of the effect, since Taylor's theory requires the diameter of a flexible molecule to decrease as the temperature rises.The observed temperature coefficients of viscosities permit calcula- tion of apparent collision areas according to the Sutherland equation, 1-594 X I O - ' ~ ~ / M . Alternatively, the method K where the collision area, A = of Hirschfelder, Bird and Spotz can be used. Values as established from the new data and from previocs observations are summarized in Table 111.' TABLE I11 Hydrocarbon n-Butane . isoButane . n-Pentane . isoPentane . n-Hexane . n-Heptane . isoHeptane . n-Octane . iso-Octane . Hirschfelder Collision Area (AZ) I 9-6 22-4 26.2 - 27'4 - - 34'7 - Sutherland Collision Area A (A? I 6-6 17-1 19-6 I 8.9 25'4 26-7 31.1 2 7'9 32-6 From Table I11 the Sutherland collision diameters of the n-paraffins in the gas are rather smaller than for the corresponding isomers, in con- trast with Table I1 in which the diameters appear slightly larger.This is because the Sutherland collision areas do not refer to quite the same molecular quantity as the momentum transfer collision areas at any one temperature (eqn. (I)). From the way they are defined, the smaller Sutherland values indicate a greater " compressibility " of n-paraffins on collision, compared with isoparaffins of the same molecular weight. The Hirschfelder collision areas refer to zero velocity of approach of the mole- cules and again suggest that n-paraffins are somewhat more compressible than isoparaffins in a molecular collision. Although the information de- tailed in Tables I1 and I11 does not refer to collision diameters defined in quite the same way, the figures in all cases point to quite considerable crumpling of the n-paraffins.Evidence* for some further coiling on vaporization can be obtained from a comparison of the entropies of vaporization of isomers (cp. ref. (I) and Table V) and from the difference in attraction potentials between molecules in liquid and gas (Table IV). If the flexible mole- cules were fully extended the force constants Elk would increase linearly with n. It will be seen that at the critical temperature there are some indications of an increasing trend of E/k with n. But in the gas there is no definite trend. This agrees with the view that the flexible molecules are crumpled somewhat more in dilute gas phase than in the liquid.The observed increase of viscosity with temperature implies a smaller average collision diameter for momentum transfer and a smaller average collision frequency as the temperature rises. Unless account is taken of lo Taylor, J . Chem. Physics, 1948, 16, 258. * From ref. (6), p. 207. D98 Force Constant e l k ~ in Liquid (Critical 1 Temperatures) CONFIGURATION AND REACTIVITY TABLE IV.-COMPARISON OF FORCE CONSTANTS FOR MOLECULAR ATTRACTION BETWEEN LIKE MOLECULES Molecule Force Constant elk in Gas (Viscosity Data) TABLE V.-~OMPARISON OF ENTROPIES OF VAPORIZATION OF ISOMERIC OCTANES AT THE BOILING POINT Molecule n-Octane . 3-Ethylhexane . 2-Methylheptane . 2 : e-Dimethylhexane . 2-Methyl-3-ethylpentane .2 : z : 3-Trimethylpentane . 2 : 2 : 3 : 3-Tetramethylbutane . hIoIar Entropy of Vaporization (Cal./mole deg.) 20.96 20.9 I 20.5 5 20'34 20'2 I 20.08 19-92 this effect a spurious temperature coefficient may be introduced in hydro- carbon kinetics. For bimolecular collisions in butane the temperature coefficient of collision frequency from this cause would simulate an activation energy of about 600 cal., which is usually less than the experi- mental error. Some Consequences of the Coiling of Flexible Hydrocarbons for Reactivity.-The coiling of the n-paraffin molecules in the vapour phase can have a number of important effects on the reactivity of individual bonds. (a) STERIC EFFECTS.-on average, the end methyl groups in a poly- methylene chain will have not less than tw9 C-H bonds available for direct reactions such as the radical-forming process, I 1 I I I I I I 4 - H + X -+ XH + -C .. . where X is an atom or free radical, or for bond rearrangement reactions such as -C--H + C1, + HC1 + -C-Cl. In some of the crumpled configurations, intermediate CH, groups may have one or both -C-H bonds blocked owing to the coiling. It is true that the energy required in a collision in order to uncoil a molecule is not large. But the probability of uncoiling is small for geometrical reasons, except for end-onA. R. UBBELOHDE AND J. C. McCOUBREY 99 collisions. Furthermore, in the time required for uncoiling the collision energy' primarily transmitted to the paraffin will tend to be redistributed throughout the coupled vibrators.As a consequence, direct reactivity of the end groups immediately following a collision is sterically favoured in spite of the somewhat lower bonding energy of secondary and tertiary C-H bonds. This seems to be the explanation of " end attack " on n-paraffins, to which reference has been made in previous papers.12 It also appears to account for the higher steric factors which compensate for higher activation energies in relative reaction rates of primary', secondary and tertiary C-H with methyl radicals.13 Indirect reactivity which follows on the primary activating collision only after some redistribution of energy within the molecule does not necessarily favour the end groups, apart from a possible " organ-pipe " effect previously mentioned.1 It seems likely that uncoiling will in many cases precede sorption into capillaries of the molecular sieve type 14 and may influence the kinetics of sorption of flexible hydrocarbons.15 (b) POLARIZATION EFFECTS DUE TO CoILING.-Points previously dis- cussed need only be briefly summarized here. (i) Owing to coiling, the CHB groups form an environment akin to that of a solvent around certain bonds.This " smudges " quantum re- strictions on certain processes, such as the postulated peroxide radical formation in gaseous oxidation of hydrocarbons,l RCH2 + 0 2 + R-CH,-O-O-. gas gas When small, a single molecule formed from a binary collision would only have very short life unless it is stabilized by collision with a third body. But in the longer ut-paraffins crumpling and vibration coupling facilitates energy storage and the satisfying of quantum conditions by smudging the sharpness of the energy levels rather similarly to the smudging of quantum conditions in a liquid.(ii) Comparison of infra-red and Raman spectra of molecules in gas and in solution l 6 shows that a polarizable medium in the neighbourhood of the absorbing bond can markedly affect relative probabilities of excita- tion. Molecular coiling must to some extent produce this effect likewise, and can in consequence modify reactivity in certain hydrocarbon re- actions. (c) ENERGY TRANSFER AND MOLECULAR CoILING.-Though some of the effects discussed below refer to data which at present are rather frag- mentary, they are included on account of their experimental and theoretical interest.The transfer of a few quanta of rotational or vibrational energy does not necessarily follow the rules that apply for activation sufficient to produce reaction." It is nevertheless interesting to review informa tion on energy transfer as a function of the number of carbon atoms and flexibility of the molecule, and as a function of the structure of isomers. Two experi- mental techniques can give information. (a) ACOUSTIC MEASUREMENTS OF THE ADIABATIC C,/C, = y AS A FUNCTION OF FREQUENCY OF VIBRATIONS.-AS is well known, gases such as C02 and C2H, show relaxation effects when the frequency of the sound exceeds about 105 c./sec. This is attributed to restrictions on the transfer l2 Ref. (I) : cf. 2. Elektrochem., 1936, 42, 468. l3 Steacie, Darwent and Trost, Faraday SOC.Discussions, 1947, 2, 86, but see Trotman-Dickenson and Steacie, J.Chem. Physics, 1950, 18, 1097. l4 Barrer, Quart. Rev., 1949, 3, 293, Is Colloque sur E'Adsorptron et La Cinetique Heterogzne, Lyon, 1949. l6 E.g. Herzberg, I%fra-red and Raman Spectra of Polyatomic Molecules 1' CasteIIan and HuIburt, J . Chem. Physics, 1950, IS, 312. (Van Nostrand, 1945).I00 CONFIGURATION AND REACTIVITY of vibrational energy in molecular collisions. No evidence of any dis- persion regions of this kind have been found la for the molecules propane or n-hexane or cyclohexane up to 40 x I O ~ c./sec. Ultrasonic dispersion was found with ethane, benzene, and cyclopropane. Restrictions on the interconversion of translational and vibrational energy appear to be prominent for rigid molecules in which vibrations of low frequency are absent.At first sight the direct excitation of vibrational energy of a bond in a paraffin as a result of molecular impact should be subject to restrictions not very different in degree from those in ethane, where dispersion effects are observed. However, a flexible molecule can store some of the energy of collision in the form of torsional oscillations. These can subsequently be redistribhted within the spectrum of internal molecular vibrations, in accordance with the degree of coupling of these vibrations. As a result, the metrical conditions for the indirect activation of vibrational energy in molecular collisions can be much less rigorous than for non-flexible mole- cules. This may explain the distinction between ethane and hexane ob- served by' Lambert and Rowlinson, though further experimental data are desirable to test this suggestion.(b) MEASUREMENTS OF THERMAL CONDUCTIVITY K.-As has been pointed out previously la restrictions on the transfer of vibrational energy can be examined from thermal conductivity data. If the overall thermal conductivity of a gas is split into the terms referring respectively to the transport of translational, rotational and vibrational energy K = Ktram f Krot f Kvib ; then according to the Chapman-Enskog formula Ktram = 2.5 &trans and according to Eucken's hypothesis Krot = VCrot. If vibrational energy contributes without restrictions to heat transfer, Kvib = &vib. But in the extreme case, restrictions on the translational + vibrational process lead to the suppression of this term.The observed thermal con- ductivity may be expected to lie between the extreme values K(norestrictions) = 77(2*5ctram + Crot + Cvib) K(vibrationa1 restrictions) = 7)( 2'5ctrans + crot) - ce = ctrans -f- Crot + Cvib and Writing and the thermal conductivity will lie between the extreme values (referred to I mole in lieu of I g. where M is the mol. wt.), MKmxl~ = Ce + 1.5Ctrans = C, + 4.5, MKminlT = 2.5Ctrans + G o t = 10.5. Ctmm = G o t = 3R/2, Values of K and Cv are somewhat variable in quality but Table VI collects the information available. It illustrates the trend in normal paraffins with increasing number of carbon atoms. By comparing the last column with the values of M K / q i t will be seen that there is no evidence of any restricted transfer of vibrational energy in the thermal conductivity of the n-paraffins.There is in fact evidence, particularly for the higher paraffins, of a correlation such that the la Lambert and Rowlinson, Proc. Roy. Soc. A , 1950, 204, 424. l9 Ubbelohde, J . Ckem. Physics, 1935, 3, 219.A. R. UBBELOHDE AND J. C. McCOUBREY I 0 1 ' I hottest '' molecules translationally also contain more than the average vibrational energy since the experimental behaviour suggests . - where a > I. More reliable data are required to test this possibility, which could have important consequences for preferred energy transfer in the chain reactions of hydrocarbons.20 TABLE VI.-TEST OF RESTRICTIONS ON VIBRATIONAL ENERGY TRANSFER IN MOLECULAR COLLISIONS IN RELATION TO THERMAL CONDUCTIVITIES OF NORMAL PARAFFINS (GAS PHASE) Molecule CH, .C2H6 C,H8 - n-C4Hlo . n-C,H12 . n-C,H14 . ~so-C~H~O. M 16.032 30.048 58.080 58.080 72.096 86.112 44'044 (1) K x 1o6 7-2 I 3-60 3-22 3'32 3'12 2-96 4'36 ( 2 ) q x I@ 10'22 8.5 I 7'46 6.86 6-86 6-21 5.86 MKIrl) I 1-31 15'39 21.3 27'3 27'3 36.2 43'5 (6.34) 9'72 14-16 19-17 24'9 29'9 - -- co + 4'5 ( 10.84) I 3.2 18.6 24-2 29'4 34'4 - (I) At oo C : quoted by Partington, Advanced Treatise on Physical Chemistry. I (2) At oo C : from Schuil, Phil. Mag., 1939, 28, 679. (3) At oo C : calculated irom values of C, (Pitzer, Ind. Eng. Chem., 1944, 36, 829) and Berthelot's equation using values of p c and To quoted by Partington above). Other Collision Diameters in Relation to the Molecular Flexibility of Hydrocarbons.-Any molecular collision process can in principle be affected by molecular flexibility and the coiling of n-paraffins.According (Longmans, 1949)~ p. 893. TABLE V I I . 2 1 - E ~ ~ ~ ~ ~ ~ ~ ~ COLLISION DIAMETERS FOR QUENCHING Hg LINE (A2537 A) Molecule C3H8 - n-C,H,o ~ s o - C ~ H ~ ~ . n-C,H,, . iso-C6H1z . neo-C6H12 . 2-Methylpentane . 2 : 2-Dimethylbutane . n-C,H16 . n-C6H14 . u x 108 (cm.) I-- - 0.33 1-13 1-73 2-2 I 2-94 3'52 1-22 4.0 I 4'54 2'39 5'39 I to the process envisaged, the relative collision diameters of homologues and isomers may show a different sequence. Systematic studies are not at present very numerous. Optical studies give some information Zo Cf. Ubbelohde, ref. (I), p. 369. 21 Darwent, J. Chem. Physics, 1950, 18, 1532.I 0 2 CONFIGURATION AND REACTIVITY about ‘ I quenching diameters ” in fluorescence and about effective diameters for collisions of the second kind with excited atoms.Pressure broadening measurements on infra-red and micro-wave ab- sorption lines due to collision with paraffins does not at present appear *, to have been carried beyond C,. The effective diameters for the quenching emission from metastable molecules have been evaluated for a number of hydrocarbon homologues and isomers. Generally the effective diameter is larger for the isoparaffins, presumably’ owing to the greater proportion of secondary and especially tertiary C-H bonds, Bond rupture occurs in the quenching process, and collision diameters do not appear to depend primarily on the volume of the molecule but on a summation of the bonds present. Although no process appears to have been identified in general hydrocarbon reactivity which can be immedi- ately compared with this type of collision, the possibilities may arise in certain chemiluminescent reactions. DifPerence between Homo-molecular and Hetero -molecular Colli - sions.-Precise collision diameters depend on the degree of interpenetration of the molecules. Certain heteromolecular collisions show about the same collision diameter independent of the second molecule. Thus di-butyl- phthalate diffusing in air has u = 8-94 A, in H, CT = 9-4 A, and in f r e ~ n , ~ ~ u = 10.5 A. On the other hand when specific interaction is suspected, as for H, with C,H, distinctive effects are observed. TABLE VIII.-RATIO OF HOMO-MOLECULAR COLLISION DIAMETERS TO COLLISION DIAMETERS WITH H, -- I Molecule I Ratio -- I- (1’00) 1-34 1-57 1-70 1-83 This aspect of collision processes is under investigation for flexible molecules, in view of the special role of H, in certain hydrocarbon re- actions. 25 Department of Chemistry, Belfast. Queen’s University, a2 Coggeshall and Sair, -1. Chem. Physics, 1947, 15, 65 ; Bleaney and Penrose, 23 See also Darwent and Steacie, J . Chem. Physics, 1948, 16, 381. ,4 Birks and Bradley, Proc. Roy. SOC. A , 1949. 198, 226. 25 Cf. Small and Ubbelohde, J . Cham. Soc., 1950, 723. Proc. Physic. Soc., rg48, 60, 83.

ISSN:0366-9033    

DOI:10.1039/DF9511000094

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

GENERAL DISCUSSION + I, or or GENERAL DISCUSSION as one of a number of distinguishable alternatives ? Prof. A. R. Ubbelohde (Belfast) said In the ionization of a poly- methylene molecule how far is i t permissible to regard the charge a,s located a t a specific part of the molecule ? To select one out of a number of possible examples is it significant to write the process Hydrocarbon -f hydrocarbon + electron For example i- + CH3-CH2-CH2-CH2-CH~-CH2-CH2-CH3 + CH~-CH~-CH~-CH~-CH2-CH~-CH2-CH3 + I \ CH~-CH2-CH2-CH~-CH2-CH2-CH2-CH3 CI-I3-CH2-CH2-CH2-CH,-CI-I,-CH2-CH CH3-CH2-CH2-CHZ-CH2-CH2-CH-CH3 or + I + I, + in which the ionization energies I to I need not be the same. If these processes cannot be distinguished the molecule could be regarded as a linear conductor for the charge rather like the polyenes.But if the above states are distinct it becomes relevant to consider (i) The rate of transfer of an electron from one part of the molecule to another and how the time required for transfer compares with the time before disruption of the molecule in the mass spectrograph. (ii) How such transfer is affected by structural modifications to the polymethylene chain such as substitution of H by CH or other groups or the replacement of single bond C-C skeletons by unsaturated skeletons. (iii) M’hether any proton migration can occur in the gas phase in ion- ized unsaturated molecules The analogy between these problems and conductors semi-conductors and insulators in three-dimensional systems may be suggestive for certain problems of molecular reactivity of ionized molecules.Sir John Lennard-Jones (Cambridge) said It is not possible to give a brief answer to all the interesting questions put by Prof. Ubbelohde. The ionization of a molecule involves the removal of an electron from a molecular orbital spread over the whole molecule so that the electron distribution in the ionized molecules may be described as that of the neutral molecule minus the distribution of one electron appropriate to the vacated molecular orbital. Corresponding to this there is a proba- bility distribution of finding “ the hole ’’ or resultant positive charge in the various parts of the molecule just as the probability distribution of an electron is obtained from the orbital which i t occupies.It will not be correct to assign different ionization potentials (such as the I . . . I mentioned in the question) to the several parts of the molecule. Each of the possible ionization potentials (for example those listed in Table I1 of our second paper) corresponds to the removal of an electron from a particular molecular orbital and is associated with a characteristic proba- bility distribution of the net positive charge. The quantitative deter- mination of these distributions is a possible application of the theory but no examples have been calculated yet. Since the alternative structures mentioned by Prof. Ubbelohde are not stationary states and therefore not distinguishable the remaining questions need reformulation and cannot be answered without a dis- cussion of the energy zones of molecules similar to that of the Brillouin zones in solids.For a sufficiently large molecule i t may be possible to define a time of relaxation which governs the rate of approach to one of the stationary states and this may be related to the transfer of charge along the molecule. GENERAL DISCUSSION 1 04 Dr. P. Torkington (Brit. Rayon Res. Assoc.) (communicated) There are one or two points in Prof. Lennard-Jones’ papers which seem to have interesting even unexpected implications. Firstly if the factor determining configuration in general is repulsion between electrons de- scribable as in different localized bonding orbitals then in systems con- taining lone pairs presumably repulsions involving these will contribute.In particular in triatomic systems in which the central atom has two lone pairs the bond angle would be greater than or less than tetrahedral according as lone-pair repulsion was less than or greater than the inter- bond repulsion. One deduces that if Prof. Lennard- Jones’ hypothesis is correct the electron distributions in H,O and F,O are (apart from a scale factor) nearly identical and that in both molecules the lone-pair repulsion exceeds the inter-bond repulsion. Secondly ionization is said to be more closely related to the (equivalent) molecular orbital than the localized bond orbital description. Surely the first transition can only properly be described as from a localized orbital even if the first excited state is a combination of localized orbitals equivalent to a molecular orbital.The molecular orbital description would one supposes only apply to cases where there were several inter- vening excited states so that the ionizing model approximated closely in behaviour to the true molecule. X-H Regarding the description of the cheniical bond as a potential well between two nuclei i t might be mentioned that the only satisfactory description of the hydrogen bond (as opposed to “ system giving the correct calculated energy ”) involves the concept of electrons over-localized in the perturbed X-H bond leaving an abnormally-exposed proton weakly bound to an atom Y having an inert pair the lengthening of the bond being related to the inter-electronic repulsion within the X-H potential well.The perturbed X-H linkages in hydrogen-bonded systems are the only naturally-perturbed bonding potential wells that we have. It would seem that they are essentially similar to an excited U-T state in which the molecular orbital has remained localized ; (this in turn incidentally being presumably one of the conditions for chemical reaction to take place). Lastly in the double bond I cannot see how the two descriptions and bent-bond are equivalent. Surely in the bent-bond system there is negligible electron charge concentration along the line joining the two nuclei whereas in the (1-n structure there is a a-bond. I am assuming that the bent-bond can be looked at as arising from overlapping of two sp hybrids from each carbon atom; the initial angle with an angle of 1 2 0 O between the other two bonds on each atom being I O I O 32’.Actually this angle would not change at all without some separation of n-component. Is a better overall description one of resonance between the U-n and bent-bond models ? Dr. G. W. R. Bartindale (Manchestw) said I should like to ask Sir John Lennard-Jones what allowance has been made for the IS electrons of the Be atom and of the C atom of the CH molecule ? Sir John Lennard- Jones (Cambridge) said Although as Dr. Torkington points out the electron distributions of the separate orbitals are rather different in the u T and equivalent orbital descriptions the total electron density at any point is of course the same in both descriptions.Thus the electron density due to one equivalent orbital is + ( u + T ) ~ and that due to the other & ( u - T ) ~ . The sum of these is ( u ~ + T ~ ) . In reply to Dr. Bartindale the IS orbitals have much lower energies and are more localized than the others. Consequently no advantage is to be gained by including them in the transformation to equivalent orbitals. Dr. A. Burawoy (Manchsster University) (communicated) Sir John Lennard- Jones’s description of molecular structures by introducing equiv- alent orbitals must be considered as one of the outstanding theoretical GENERAL DISCUSSION IOj developments for many years since i t lays the foundation to a better understanding of the relative importance of the factors determining the stabilities and shapes of molecules.It has many important qualitative features. One is the emphasis of the contribution and the need for con- sideration of the electrostatic repulsions of electrons forces which have been rather neglected or the significance of which has not been fully recog- nized in older approximate theoretical and semi-empirical treatments. Another is the justification of the old views of organic chemists that the multiple linkages in ethylene and acetylene can be accepted to possess equivalent though strained 1inkages.l However the authors still maintain that the stability of conjugated hydrocarbons in particular the observed strengthening and shortening of the single C-C linkages is due to a partial delocalization of the orbitals of the multiple linkages.Attention map be therefore called to some recently published papers in which the writer has shown that the valency conception of non-localized bonds has no justification being in disagreement with numerous if not all relevant observations. As an example this view demands that the multiple linkages in con- jugated hydrocarbons are longer and weaker than in ethylene and acetylene. This has been clearly stated and supported by approximate quantitative computations in past years by Sir John Lennard- Jones,3 M~lliken,~ Coulson and others. However investigations by the X-ray and electron diffraction methods show unambiguously that the internuclear distances of the multiple linkages in conjugated hydrocarbons remain unchanged or are in many cases actually shortened.6 This is well illustrated by the bond distances observed 7 for diacetylene dicarboxylic acid HOOC-CrC-C=C-COOH.distance is found to be only 1-33 A i.e. as short as the double bond in The middle C-C ethylene. In spite of it the C=C bond distances are only 1.185 A i.e. not longer but shorter than the CEZC distance observed for acetylene (1.204 A). As recently shown this and numerous other difficulties disappear if the valency conception of non-localized bonds is abandoned. The (con- stitutive) changes of linkages and their properties in all polyatomic mole- cules are easily accounted for by changes of the effective nuclear charge (the screening) of atoms i.e. changes of the electrostatic repulsions of electrons (and nuclei).Thus the increased bond energy of the middle C-C linkages in con- jugated hydrocarbons is not due to a delocalization of the orbitals of the multiple linkages and the resultant exchange forces but to the reduced repulsion between the electrons of the multiple linkage and those of the single bond (and possibly of the second multiple linkage) as compared with ethane. Similarly the reduced internuclear distance (and increased bond energy) of the CZC bond is expected. The replacement of a hydrogen atom in acetylene by the much more strongly electron attracting acetylene group will withdraw the electron cloud from the C atom of the triple bond. Its effective nuclear charge will be increased i.e. the repulsion between the triple bond electrons and those of the single bond will be reduced and cause the shortening and strengthening of the CEC bond.For a discussion and explanation of the properties of multiple linkages cf. also Burawoy.2 Burawoy V . Henri Mem. VoZ. (Desoer LiBge 1948) ; cf. also Trans. Faraday Soc. 1944 40 537 ; Chem. and Ind. 1944 434. Lennard-Jones Proc. Roy. Soc. A 1937 158 280. Mulliken Rieke and Brown J . Amer. Chem. SOC. 1941 63 41. Coulson and Jacobs J . Chem. Soc. 1949 2805. 6 For a summary of the available data cf. Burawoya2 Dunitz and Robertson J . Chem. SOC. 1947 1145 ; cf. also Jeffrcy and .Rollett. hiatwe 1950 166 475. GENERAL DISCUSSION I 06 It is not a coincidence that the (writer’s) empirical analysis of the properties of simple and complicated polyatomic molecules (e.g.saturated and conjugated hydrocarbons) and the new theoretical analysis a t least of simple systems by the authors lead qualitatively to the same con- clusions the great importance of the electrostatic repulsions of electrons (and their variations) and the negligible contributions by exchange forces arising from delocalization. However the empirical analysis shows t h a t these conclusions are also valid for conjugated and other complicated systems. The acceptance of the equivalence of the bonds of a multiple linkage may have already removed the theoretical necessity if not justification for the hypothesis of non-localized bonds which a t any rate appears to be in disagreement with experience. The purpose of this comment is to draw attention to the need for an unbiased theoretical treatment of the structure of conjugated systems which would give information on the relative importance of the true factors responsible for their characteristic properties.Sir John Lennard- Jones (Cambridge) (communicated) We appreciate Dr. Burawoy’s favourable comments on our papers. On the other hand he gives the bond lengths of diacetylene dicarboxylic acid and asserts that the unusually short links which are observed cannot be explained by the usual methods. We ourselves do not know of any molecular orbital calculations on such a molecule and cannot say whether his statement is true. It is evident that the system is a complicated one and cannot readily be treated by current methods but it would be unsafe to predict that these methods would not be successful.Dr. Burawoy thinks that the observations can be explained in terms of the change of electrostatic repulsion between the electrons of neighbouring links but it is difficult to form a judgment on this until a quantitative treatment is available. Prof. E. C. Baughan (Shrivenhaw) (communicated) There is one point in the papers of Lennard-Jones and Pop!e to which I would like to draw attention. Their theoretical principle that lone pair electrons play a vital part in determining structures is well supported for several ele- ments by an inductive survey of experimental stereochemistry by Sidgwick and Powell who summarize the evidence thus “ Nearly (but not quite) all the structures can be even more simply related to the size of the valency group by assuming that the mean positions of the electron pairs in this group are the same whether they are shared or not.” Dr.C. A. McDowell (Liverpool University) (comnzunicated) Sir John Lennard-Jones and Dr. Pople have indicated that methylene may exist in a triplet state in which case it would probably have a linear structure or in a singlet state when i t would be expected to be angular. Very little is known about this compound but its ionization potential 9 is known to be 11.9 eV and I should like to ask these authors if i t would be possible to calculate the expected ionization potential for each of these two possible states. A comparison with the known experimental value might then help to decide whether methylene has a singlet or triplet ground state.Sir John Lennard-Jones (Cambridge) said I am grateful to Prof. Baughan for calling our attention to the paper by Sidgwick and Powell in which the structures of a large number of molecules are classified. These structures are readily underst;& if the lone pairs of electrons are assumed to be as important as the bond electrons and to have distribu- tions which fit in with the symmetry type. In reply to Dr. McDowell we have not calculated the ionization poten- tials of the methylene radical but the methods we have described may be applicable to such a structure. PYOC. ROY. SOC. A 1940 176 153. Langer and Hipple Physic. Rev. 1946 69 691. GENERAL DISCUSSION C-C 107 Dr .A. Burawoy (Munchester University) (communicated) The problem of the variations of bond energies in saturated hydrocarbons has a con- siderable bearing on the calculation of the bond energies and stabilization energies of most of the molecules discussed by Prof. Glackler. As recently l o pointed out the available experimental data allow two empirical inter- pretations of the changes of bond energies in saturated hydrocarbons. (i) The bond energies of the C-C and C-H linkages may be assumed to decrease in the order of participating CpTi. > C,,,. > Ctert. > Cquart. This generally accepted view leads to unlikely conclusions. The Table shows the calculated atomic heats of formation and the possible C-C and C-H bond energies for methane ethane propane isobutane and neopentane i.e.the bond energies are arranged in the order C y r i . to Cquart. Assuming the C-C bond energies remain unchanged the C-H respectively (column I). Any decrease of the C - C bond energies will be accompanied by’ a correspondingly’ greater decrease of the C-H energies (for an example cf. column 2). bond energies will decrease a t least by 1.1 1.0 and 0.9 kcal. bond These values are not arbitrarily chosen. In order to obtain unchanged bond energies the C-H bond energy X in ethane must satisfy the equation X = 97.0 - a where u is the difference between the C-H bond energies in the methyl groups of ethane and neopentane. This difference can be only small and may be ignored for the purpose of this discussion.It will be observed that the C-H bond energies decrease more strongly than the C-C bond energies. This is contrary to theoretical expectations and experience according to which the polarizability of C-C is greater than that of C-H. The smallest possible bond energy decrease from Cpri. to Cquart. (3 kcal.) required by this interpretation also appears to be unreasonably high. Moreover this assumption leads to the improbable conclusion that the stabilization energy of the C-C linkage in propene (89.2 - 84-5 = 4-7 kcal.) is only slightly higher than that of the C-H linkages in ethylene (101.6 - 98.1 = 3.5 kcal.) and this small increase is only apparent since it depends on the assumption that the bond energy of the -=C-H linkage adjacent to the methyl group in propene is only 99-7 kcal.as compared with the 101.6 kcal. of the other =C-H bond energies .I1 (ii) The bond energies of the C - C and C-H linkages may be assumed to increase in the order Cprf. to Cpuart. . If the C-H bond energies remain unchanged the C-C bond energies will increase by 1.2 1.1 and 1.0 kcal. respectively (column 3). Any increase in the C-H bond energies can be shown to be accompanied by a correspondingly greater increase of the C-C bond energies (for an example cf. column 4). bond energies vary now more strongly than the C-H The C-C bond energies as one should expect. Moreover the C-4 bond energy in propene will now be a t least 85.6 kcal. even if the C-H bond energy in the methyl group is taken as high as 99.0 kcal.i.e. the C - C stabiliza- tion energy will be greater than 8-3 (85-6 - 77-3) kcal. This empirical analysis is independent of the physical cause of the bond energy changes which has been discussed elsewhere.12 It indicates that the second interpretation must be the correct one. Thus the C - C bond energy in ethane will be 77-3 kcal. or smaller by an amount dependent on the increase of the C-H bond energy passing from methane to ethane. The C - C dissociation energy of ethane is of course independent of any possible individual bond energies. It is 83.3 kcal. if the atomic heats of formation of the methyl radical and of ethane are 291.4 and 665-9 kcal. respectively. l o Burawoy V . Henri Mem. V d . (Desoer LiBge 1948) p. 80. l1 Based on the values suggested by Glockler this Discussion.Based on the heats of formation values at oo K from Selected Values 1s and L(C) ;3 169.7 kcal. J3 Nut. Bur. Stand. (Cir. 461) 1947. GENERAL DISCUSSTON TABLE I I 2 4 3 C H C-C C-H C-C C H G H C-C C-C - 98.1 89.9 96-0 94.1 87.8 92-4 I 08 Substance Me& - 98.1 74'3 98.6 76.1 99.0 77'7 99'3 79'1 88.8 86.9 Q 2 392'4 HCH . MeCH . 665.9 Me,CH . 941.8 Me,CH . 121g*8 1499.6 - or C-H - because although not obtained in absorption in the . Dr. P. Torkington (Brit. Rayon Res. Assoc.) (communicated) Prof. Glockler interprets the bond length against bond energ41 data of the series C,H, C2H4 C,H as demonstrating the tendency of C-H bonds of high energy to be associated with C - C bonds of high energy in general.One would have thought that this type of association can only be accepted as probable on the above evidence in systems in which the C - C bonds can increase their energy by taking some multiple-bond character. Does Prof. Glockler interpret his results for the normal paraffin series as implying some double-bond character at the ends of the chains ? It seems possible that the stronger bonds at the ends of the paraffin chains arise from an overall molecular orbital component in the structure which favours the ends irrespective of whether they are C-C bonds-the data might almost be taken as proof of the reality of such orbitals. But it would seem difficult to relate this phenomenon to the changes in bond strength arising from hybridization changes associated with unsaturation.Dr. A. G. Gaydon (Imperial College) said Regarding the dissociation energy of CH my observations on the effect of pressure and other con- ditions on the spectrum leave me in no doubt that the predissociation is strong and real and not of the type which could be confused with a per- turbation. The only possible loop-hole which I see is that the ground state of CH may not be ,l7 but an unobserved quartet state ; the failure to observe CH bands in the laboratory would appear to support this view but as pointed out by ,z later speaker the occurrence of CH in inter- stellar absorption makes the possibility remote. Dr. George Porter (Cambridge) said Dr. Gaydon makes the suggestion that the 2 l 7 state of CH to which the dissociation energy of 80 kcal.refers may not be the ground state which is particularly interesting in view of the considerable discussion about the ground state of methylene. I think however that strong experimental evidence is available that the ground stat? is laboratory transitions from this level are observed in interstellar space. It is well known that owing to the rare occurrence of collisions in this region molecules exist in their ground electronic vibrational and rota- tional states a single line only being in general observed. The *Z - transition is not strictly forbidden even in the absence of collisions and therefore i t would appear that the 2 l 7 state must have the lowest energy.Dr. G. J. Minkoff (Imperial College) said Dr. Stevenson has men- tioned the paper of Voge in support of a value of ca. 80 kcal. for C-H in CH,. This result followed however from the original assumption by A and D(tert.-C,H,-H) given by Stevenson Voge of an angle of goo for H-C-H in CH (since this implies C-H bonds of p rather than sp character). Dr. B. G. Gowenlock (Swansea) (communicated) From the values of D(CH,-H) D(sec.-C,H,-H) together with the value of D(C,H,-H) used by Glockler we have values of 0 4 7 and Ir#g kcal. for the resonance energies of the methyl ethyl sec.-propyl and tert.-butyl radicals respectively. It can thus be seen that these experimentally observed figures agree in relative order of GENERAL DISCUSSION 109 magnitude with the calculations of Baughan Evans and Polanyi.l4 The calculations of these authors give resonance energies for the above radicals in the relative order of magnitude o I z 3 corresponding to the con- tribution of 0 3 6 and g equivalent hyperconjugation structures to the actual state of the radical.Thus the electron impact data result in a closer relative agreement with the simple valence-bond theory of alkyl radicals than the D(R-H) values derived from the D(R-I) pyrolysis data of Butler and P01anyi.l~ Dr. Alwyn G. Evans (Manchester) said With regard to the results of Dr. Stevenson on the ionization potentials of sec.-propyl and tert.-butyl radicals I should like to emphasize the importance of this work in relation to the interpretation of organic reactions involving carbonium ions.The ionization potentials 7-77 eV for sec.-propyl and 7-19 eV for tert.- butyl were estimated 16 from the earlier work of Stevenson and Hipple l7 using the dissociation energies given by Butler and Polanyi,l* and the marked decrease of ionization potential from methyl to tert.-butyl was shown to be a major factor in determining the changing tendency for S,I reaction from primary to tertiary alkyl halides.16 The ionization potentials of sec.-propyl (7-45 f 0-1 eV) and tart.-butyl (6.90 f 0-1 eV) obtained by Dr. Stevenson from this very detailed investigation are close to the earlier values and show the same difference in energy (0.55 eV). The marked decrease in ionization potential from methyl to tert.-butyl means that the resonance energy of the tert.-butyl ion is greater than the resonance energy of the tert.-butyl radical.This indicates that the energies of the ionic resonance states H+ H (CHJ ,C=C-H and (CH,) ,C-C-H H I H (9 of these) H. (CH3)zkC-H H I (9 of these) + + I lie closer together than do the energies of the radical resonance states H This may be due to the fact that in the excited resonance structures of the radical the free hydrogen atom is situated at the normal covalent distance from the carbon so that this structure involves a high repulsion energy between these two atoms while in the excited resonance structures for the ion this large repulsion will not occur because of the small size of the proton.16b If the observed decrease in ionization potential from methyl to tert.- butyl is to be explained in this way one might expect that the change from methyl to allyl or from methyl to benzyl would not involve such a decrease since in allyl and benzyl no unbonded particles are present in the resonance structures for the radical (e.g.CH2=CH-6H and or for the ion (e.g. CH,=CH-CH and CH,-CH-CH,). CH,-CH=CH,) methyl. l7 Hipple and Stevenson J . Amer. Chem. SOC. 1942 64 1590 2766 2769. + Thus the resonance energy of the radical may be similar t o that of the ion and hence the ionization potential of the radical similar to that of l4 Baughan Evans and Polanyi Trans. Faraday SOC. 1941 37 377. l5 Butler and Polanyi Trans. Faraday SOC. 1943 39 19. 16 Evans A.G. (a) Trans. Faraday SOC. 1946 42 719 ; (b) Reactions of Organic Halides in Solution (Manchester Univ. Press 1946). Butler and Polanyi Trans. Faraday SOC. 1943 39 19 GENERAL DISCUSSION I I 0 Evidence that this may be the case comes from a comparison of the S,I reactions of allyl and benzyl chlorides with that of tert.-butyl chloride. In spite of the fact that the carbon-halogen dissociation energies in ally1 and benzyl halides are less than that in tert.-butyl halide the S,I reactions of allyl and benzyl chlorides are much slower than that of tert.-butyl chloride. We believe this to indicate that the ionization potentials of allyl and benzyl radicals are greater than that of tert.-butyl.19 by McDowell and Warren for D(CH,-CN) - D(H-CN) viz. + 0.24 eV Dr.D. P. Stevenson (Euneryville California) said The value reported is not only of sign contrary to that expected by chemical intuition but is inconsistent with the heats of formation of CH,CN(g) (AH;, = + 19.4 kcal./mole) HCN(g) (AH,", = + 30.7 kcal./mole) and CH4(g) (AH,", = - 18.2 kcal./mole). Thus CH4 + HCN = CH,CN + H, AHzgl = + 19'4 - 30'7 + 18.2 = + 6.9 kcal./mole = + 0.30 eV/molecule 4-48 eV/molecule H = ZH 4-42 eV/molecule and hence CH4 + HCN = CH,CN + ZH AHl = 4-78 eV/molecule. While CH4 = CH + H CH3CN = CH + CN HCN = H + CN CH4 +HCN = CH,CN + ZH + 4-42 eV/molecule D = D = D(CH3-CN) AH = - D(CH,-CN) + D(HCN) D(H-CN) = 4-78 eV/molecule. Thus D(H-CN)-D(CH,-L") = + 0.36 eV/molecule not - 0.24 eV/molecule as given by McDowell and Warren.It may be noted that if Smith's value for A(CH:) = 15.8 eV/molecule in the methane mass spectrum is taken instead of that of McDowell and Warren 15.3 eVjmolecule then one may compute D(H-CN) - D(CH,-CN) = +0-28 eV/molecule by the method of McDowell and Warren in good agreement with the thermochemical value for this quantity' and consistent with the quantity D(H-CN)-D(NC-CN) = + 0.26 eV/mo1ecule.20 Dr. C . A. McDowell and Dr. J. W. Warren (Liverpool) (communi- cated) We are entirely convinced by Dr. Stevenson's thermal chemical data and also by his chemical intuition that D(HCN) > D(CH,CN). This is also substantiated by our measurements of the appearance potential of the CH,+ ion from methyl cyanide.We obtain a value of 14.6 eV for Y(CH,+) which with I(CH,) of 10.08 eV leads to a value of 4-52 eV for D(CH,CN). Taking D(HCN) as 4-81 21 we get D(HCN) - D(CH CN) = 0.3 eV in agreement with the value given by Stevenson. The contrary deduction given in our paper was based on measurements of the appearance potential of the CH,+ ion from methane and from methyl cyanide and emphasizes how careful one has to be in using appearance potential measurements. Both these measurements seem to be correct as far as we can determine. In each case the CH,+ ion beam is fairly large and the ionization curves are similar in shape and well defined. It may be that our method of determining appearance potentials of ions Evans A. G. and Hamann Trans. Faraday SOC.1951 47 28. 2o J . Chem. Physics 1950 18 1347. 21 Stevenson J . Chem. Physics 1950 18 1347. GENERAL DISCUSSION 1 1 1 produced in multiple dissociation processes tends to give low results when there is a large difference in the masses of the fragments. Another possible cause of the difference is that the CH,+ ion from methyl cyanide may be produced with an excess of kinetic energy of about 0.5 eV. This would be sufficient to bring the results on this ion into agreement with the data given above. If such an amount of kinetic energy were liberated in the process leading to the production of CH2+ ions from methyl cyanide one would have thought that our beam half-width method would have detected the broadening due to this for in this case the ion is only about one-harf of the mass of the neutral fragment and so i t should receive about two-thirds of the excess kinetic energy.It will be recalled that the CH,+ ion beam in methyl cyanide showed evidence of broadening. It should perhaps be pointed out that the discrepancy to which Dr. Stevenson has drawn attention while important in itself in no way requires any of the main conclusions of our paper to be altered. . - (4 Dr. A. F. Trotman-Dickenson (Manchester) said I should like to make two points concerning the C-H bond strengths in alkanes. First the original papers 22 on the determination of the C-H bond strengths in methane and ethane by photobromination gave D(CH,-H) = 102 kcal. and D(C,H,-H) = 99 kcal. a t 298" K. These figures have been generally accepted and are based on the activation energies of the following reactions (the activation energies are in kcal./mole). Br + C,H 13.3 C,H + HBr. and 0.8 The determinations of the activation energies of the forward reactions were very precise and it is unlikely that the value of 2 kcal. for the reverse reaction of (I) is greatly' in error. But it is very improbable that the reaction of an ethyl radical with hydrogen bromide has a lower activation energy' than the reaction of a methyl radical. It is contrary to our general information about radical reactivities and the evidence in favour of an activation energy of 0.8 kcal. is not compelling. Accordingly I suggest that the photobromination data should be interpreted as giving D(CH,-H) = 102 f I kcal.at 298" K D(CH,-H) - D(C,H,-H) 2 4-5 kcal. These values are in better agreement with those of Stevenson and of Butler and P~lanyi.~ Secondly Dr. Steacie and I have investigated the reactions of methyl radicals with 10 alkanes. In a series of reactions of the type CH + CH -3 CH4 + CH CH3 + C2H6 ,-f CH4 -1- C,H the bonds broken and formed are so similar that it might be expected that if a relation of the type AE = ccAH exists for this series then cc should be closely equal to 0.5. This presumption may be stated in terms of potential energy curves. It is necessary that the two curves should have the same slope at the point of intersection and that the contribution of resonance energy to the potential energy of the transition complex should not alter in passing from one member of the series to the next.One would expect the intuitive argument to be less applicable as the bonds become more different an extreme case being the reaction of methyl with z-methyl Z t Kistiakowsky and Van Artsdalen J . Chem. Physics 1944 12 469. Anderson and Van Artsdalen J . Chem. Physics 1944 12 479. 43 Butler and Polanyi Trans. Faraday SOC. 1943 39 19. GENERAL DISCUSSION AH . I I 2 propane. A consideration of the potential energy curves shows that it is probable that a is less than 0.5 when the C-H bond in the substrate is weak. Below is a Table of the results obtained with a = 0.5 ; D(CH 3-H) = 102 kcal. is taken as a fixed point. butane Compound Methane . Primary Ethane .2 2-Dimethylpropane 2 2 3 3-Tetramethyl Secondary n-Butane n-Pentane Cyclopentane . n-Hexane Cyclohexane . TABLE I1 All figures are in kcal./mole. E ca. 12.8 10.4 &o-4 10.0 fo.3 9'5 f o . 4 8.3 fo.2 8-1 j o - z 8-3 foe2 8.1 5 0 . 2 8.3 i 0 . 2 AE X = halogen atom) . Data from Trotman-Dickenson and Steacie J . Chem. Physics 1951 19 329. It appears from the results that all primary and all secondary C-H bonds have the same strength the experimental data on the compounds containing tertiary C-H bonds in which case allowance was made for the contribution of the primary hydrogen atoms to the overall rates are not so satisfactory. Dr. H. Steiner (Manchester) said Szwarc in his paper and also Trotman-Dickenson used the relation between changes in the heat of reaction and in the activation energies derived by Evans and Polanyi 2 * AE = aAH .This relation holds for series of reactions of the type (diabatic reactions) RX + Na = Na+X- + R (R = alkyl radical . . Tertiary 2-Methylpropane 2 3-Dimethylbutane (mean) 2 3 4-Trimethylpentane . Rev 1949 3 1. when changes are made in R. From theoretical considerations relation (I) can be expected to hold when resonance in the transition state is small. In the metathetical reactions involving homopolar bonds (adiabatic reaction) such as treated by Szwarc and Trotman-Dickenson resonance is a very. important factor in determining the magnitude of the activation energy as was actually mentioned by Szwarc.Nevertheless i t is an empirical fact that relation (I) holds also in a number of series of adiabatic metathetical reactions.25 24 Evans and Polanyi Trans. Faraduy SOC. 1936 32 I933 ; 1938 34 2 2 . 25 Steiner and Watson Trans. Faraday SOC. 1947 2 88. Bolland Quart. - (1) 1 13 GENERAL DISCUSSION However in my opinion Szwarc’s method of neglecting resonance altogether in discussing reaction series of this type is not justified on the basis of the limited experimental evidence available particularly because the effect of resonance on the absolute value of the activation energies involved is very large. Dr. D. P. Stevenson (Emeryville California) (communicated) In reply to Trotman-Dickenson hitherto unpublished data on the appearance potentials of C,H,+ in the mass spectra of n-butane and n-pentane and of C,H,+ in the mass spectra of n-pentane and n-hexane lead to D(CH,-H) -D(C,H,-H) equal to 0.19 and 0.25 eV/molecule respec- tively.When these values are taken with the earlier electron impact data 26 0.22 eV/molecule we have for this difference 0-22 f 0.03 eV/mole- cule or 5-1 5 0-7 kcal./mole. in good agreement with Dr. Trotman- Dickenson’s interpretation of the photobromination work of Van Artsdalen. It may be noted that data on the appearance potentials of C,H,+ in the mass spectrum of n-hexane and of C,H,+ in the mass spectra of n-heptane and n-octane when combined with the data referred to in the preceding paragraph result in D(CH,-H) - D(n-C,H,-H) = 0-1 eV/molecule D(CH,-H) - D(n-C,H,-H) = 0.0 eV/molecule and D(sec .-C,H,-H) + Iz(sec.-C,H,) - D(sec.-C,H,-H) - Iz(sec.-C,H,) = 0-3 eV/molecule.These latter results are not in complete accord with Trotman-Dickenson’s conclusion that primary and secondary carbon-hydrogen dissociation energies are invariant for propane and higher alkanes. However these latter electron impact results are not as unambiguous as those on methyl ethyl sec.-propyl and tert.-butyl. Dr. M. Magat (Paris) said The problem of the fragmentation of poly- atomic molecules by electron impact can be approached from two different sides. Either one considers one by one relatively simple molecules and tries to interpret the effects of say isotope substitution of the mass spectrographic pattern in the light of the Franck-Condon principle 27 gradually increasing the complexity of the problem.Or one considers large series of related molecules and attempts to establish empirical rules which may give some lead to the factors determining the frag- mentation of more complicated molecules where a detailed theoretical analysis seems for the moment hopeless. This is the approach we have chosen with Dr. Viallard in a paper now in press 28 and from which I would like to quote a few results. Our analysis is almost exclusively based on the compilation of mass spectra of hydrocarbon published by the Bureau of Standards and obtained with a Consolidated mass spectrograph. The discrimination of this apparatus seems to be small except possibly for hydrogen and hydrogen molecule ions as can be seen from the com- parison of the amount of ions obtained under identical conditions from rare gases 29 with the values of total ionization of Smith.30 We have hence made no corrections for discrimination.Here are our main findings some of which are not new. The amount of parent ions in straight chain paraffins decreases with increasing chain length. The amount of parent ions decreases somewhat with any branching and is practically reduced to zero if 2 methyl groups 26 J . Chem. Physics 1942 10 292. 27 E.g. Schaeffer and Hastings J . Chem. Physics 1950 18 1048. Schaeffer J . Chem. Physics 1950 18 1501. Turkevich et al. J . Amer. Chem. Soc. 1948 70 2638. 28 J . Chim. Phys. See also Compt. rend. 1949 228 1118.29 Diebeler Mohler and Reese J . Res. Nat. Bur. Stand. 1947 38 617. 30 Smith Physic. Rev. 1930 36 1293. GENERAL DISCUSSION bonds the amount of C,H,+ ions being abnormally high. 114 are substituted on the same carbon atom. The introduction of a double bond increases the amount of parent ions as compared to corresponding paraffins for chains containing less than six carbon atoms 31 but increases for longer chains. The particularly large amount of parent ions in arom- atics was already noticed and interpreted by Rurton.S2 The rupture of carbon-hydrogen bonds occur in the way to favour the formation of radical ions with odd numbers of hydrogen atoms ; this has already been mentioned by Delfosse and Bleaknay 33 for propane but is a general rule valid for all fragments from any paraffin and with any number of carbon atoms.It is also true for perfluorinated paraffins. It map be related to the fact that these radical ions are in fact saturated molecules no unpaired electrons being left. The breaking of carbon-carbon bonds shows some very marked regularities (a) In normal paraffins and olefins never more than I carbon-carbon bond is broken at once-the ratios CH,+/CH,+ and C,H,+/C,H,+ decrease first when the chain length increases and tend then to a constant value. (b) In branched paraffins " double scission " occurs in the neighbour- hood of the branching fragments recombining usually to give larger ions. For example neopeqtane yields 17 yo of C,H,+ among the fragments ; neononane (3 3-diethylpentane) yields about 40 yo of C,H,+.This recombination has to occur in the activated states before the fragmentation and sets a fairly high limit to the life-time of the activated ions. It seems to us reasonable to assume that several vibrations of these ions take place before the fragments separate. The energy can hence be redistributed and i t may be necessary to investigate carefully how far the Franck-Condon principle is sufficient to interpret the phenomenon. With neononane one observes also a simultaneous breaking of three C-C (G) In normal paraffins comprising m carbon atoms one can define for each bond n a fragmentation index which gives the ratio of the per- centage of scission occurring a t a given bord and giving rise to C$ and CL- ions to the a priori percentage one would expect if the scissions were at random.The result of this calculation is given in Fig. I. It seems to be a general rule that scission occurs preferentially between the 3rd and 4th carbon atoms. Scissions at the ends which would give CH sf and C m-lHz+ are extremely improbable. The scission probability decreases also continuously towards the centre of the molecule. (d) With branched paraffins olefins diolefins and aromatics a new phenomenon sets in-the capture of H atoms by the charged fragment. For instance in the butadiene spectrum one finds a fairly large number of CH3+ ions 0.61 yo as compared to 0.88 yo of CH,+ ions. More striking is perhaps the presence of 17.1 yo of C2H,+ besides 26.6 yo of C2H3+ which means that a hydrogen capture occurs with un- saturated hydrocarbons even if an even number hydrogen ions is produced.Sometimes the hydrogen is captured not only from the neigfibouring but also from a further carbon atom. For example among the fragments of 2 3-pentadiene H,C-C-C=CH-CH one finds 1-3 yo of C2H,+ and 4'5 yo C,H,+ ions. In branched paraffins the hydrogen capture (migra- tion) occurs only if i t leads to an ionic fragment with an odd number of hydrogen atoms (e.g. format on of C3H,+ from neopentane). This hydro- gen migration frustrates for the moment all attempts to establish frag- mentation indices for all except symmetrical compounds. With olefins i t induced us previously to make the erroneous statement that terminal or secondary double bonds are never split.As the matter stands now 31 Roberts and Johnson Anal. Chem. 1948 690. 32 J. P?ysic. Chem 1948 52 567. 33 Physzc. Rev. 1939 56 250. I I 5 GENERAL DISCUSSION we can say that (i) central double bonds in 3-hexene and 4-octene are split preferentially ; (ii) terminal double bonds either split preferentially or what seems less probable strongly favour the splitting of the single bond a t the opposite end of the cha n ; (iii) double bonds in the secondary position either do not change the sciss on probability of the bond as com- pared to the corresponding paraffins (2-hexene) or decrease i t or less prob- ably decrease the sciss on probability of symmetrical single bond. CH3- L H ~ - CHz- CHz-M2-CHz-CHz-CHz-CHz-CH~-CH3 FIG.I. Fragmentation indices of normal paraffins - Electron impact data - - - Pyrolysis data of Hinshelwood et al. (this Discussion) Dr. George Porter (Cambridge) said Although there is still little agreement about the energies of dissociation of the second and third hydrogen atoms from methane the first D(CH,-H) and the last D(C-H) have been widely accepted as IOI and 80 kcal. respectively and Stevenson gives a summary of the convincing evidence for the former. GlockIer now challenges the latter on the grounds that the predissociation in the spectrum of CH from which the value is derived may be a perturbation I 16 GENERAL DISCUSSION and this is an important revision if correct. The predissociation to which he refers however involves a complete breaking off of rotational structure in several bands and not merely one missing line and is in fact a very good example of this type of predissociation.It would nevertheless be useful to have an independent confirmation of this value and this would appear to have been given by McDowell and Warren by the method of appearance potentials. It is therefore necessary to point out that these authors in using the value of I(CH) which was derived by Douglas and Herzberg via a cycle involving D(CH) have argued in a circle so that other things being correct they would automatically obtain the same value for D(CH) as is used in the first place which in this case was 80 kcal. This is only important in so far as their claim to have “ derived ” a value for D(CH) is concerned and has no bearing on their other conclusions.I wonder whether they know of any electron impact data for I(CH) because D(CH+) being well known an independent value for D(CH) could then be obtained without further assumptions. Dr. C. A. McDowell and Dr. J. W. Warren (Liverpool) (communicated) Dr. Porter has raised an obvious point but the position is not quite so simple as he implies. In using Douglas and Herzberg’s value for the ionization potential of the CH molecule we were of course aware that it had been derived by a thermodynamical cyclic process in which I(C) D(CH+) and D(CH) were involved. I(C) is well known and the value used for D(CH+) namely 3-6 eV is in good agreement with that deduced from our values for the appearance potentials of CH+ and C+ i.e.3-8 eV. Incidentally this good agreement between these two values can be taken as evidence that the ionization process leading to tLe formation of C+ ions from methane must involve the production of a t least two hydrogen atoms ; and further i t shows that if any kinetic energy is produced in the formation of CH+ ions then a similar amount must be involved in the production of C+ ions. It will be recalled that we could find no evidence for kinetic energy‘ in the formation of either of these two ions and this is in agreement with the above statement. Herzberg’s value for D(CH) has been generallv accepted by spectroscopists as being reliable and we in our paper show that it is consistent with our data. We in no way intended that our results implied that D(CH) was being determined in- dependently A different presentation of the data would perhaps have shown this more clearly but we thought i t evident from the nature of the problem that a completely independent value for D(CH) cannot at present be obtained from mass spectrometric data.One great difficulty is to obtain satisfactory source of CH radicals. Most discharge tubes do not give a sufficiently intense beam of radicals and there are also considerable experimental difficulties in getting the radicals from the dis- charge into the ionization box of the mass spectrometer. One obvious method is to use molecular beam but in this case the low intensity of the CH radicals makes the detection of the resultant CH+ ions extremely difficult.A different method of detecting ion beams of low intensity would seem to be necessary and plans have been made to use a photo- multiplier at the collector end of the mass-spectrometer. Prof. E. A. Guggenheim (Reading) said Pitzer has shown us the remarkable agreement between theory and experiment for ethane with respect both to its entropy and to its heat capacity over the temperature range from IOOOK to 300OK. There is one experimental field where further tests of the theory might be possible. I refer to heat capacities a t still higher temperatures. It would be interesting to confirm that the contribution of the internal rotation to the heat capacity passes through a maximum as predicted by theory. I realize the difficulty of such a comparison owing to the anharmonicities of the remaining seventeen normal modes but if I am not mistaken these would be serious in the case of not more than three modes.I should be greatly interested to hear Prof. Pitzer’s opinion GENERAL DISCUSSION Mr. N. W. Luft (Waltham Abbey) (communicated) It seems worth- while to emphasize the value of the electrostatic theory of internal potential barriers put forward by Lasettre and Dean. Since internal molecular rotation is periodic within 2~ the barrier can be represented by the Fourier series v = ck cos (Re) - E,. . k This may be expanded in terms of inverse powers of the distances Lik between the charge clouds (i) ( k ) whose origins are at distances ri and ~ 1 1 respectively from the ends of the rotation axis d (cf.Oosterhoff’s diagram). Thus but cih’s may change with internal rotation. Normally ck+l < ck which where E = minimum energy. ck’s are constants for any given system decreases the number of terms with appreciable contribution considerably. When the symmetry properties of certain rotating groups are also taken into consideration the effective terms in eqn. ( I ) may be less than four. In eqn (2) contributions may be expected round n = 5 i.e. for n = 3 to 6 say due to the spatial extension of the phenomenon. Any theory which demonstrates the above properties and correlates the four major ck’s with the four major c,’s jn a unique way is acceptable. Lasettre and Dean’s theory does this elegantly. They consider three types of charge distributions viz.u-bonds n-bonds and lone p (or hybrid- ized) electron pairs and use experimental dipole moments. Higher moments for which data are lacking are defined in swh a way as to represent simply accurately known barriers. Recent spectroscopic results for Me. Me MeOH MeNH and MeCF are further evidence of the justification of this method. Apparent disagreement between ex- perimental and calculated results for H,O,S4 and N,H 35 can be largely attributed to an incorrect evaluation of spectroscopic data in the case of asymmetric barriers. Calculations for internal rotation of CH groups are particularly simple since with trigonal symmetry lone pairs make no contribution as correctly stated by Lasettre and Dean though overlooked by Aston.Thus the contributions from bonds (s) opposite to CH and a t angles 46s are where A represents the interactions from more distant bonds. The secondary influences B may well be compensated for by slight deviations from rigidity during vibration. To a first approximation c,(s) N 0 as long as c3(s) > 0. In calculating three-fold internal potential barriers I suggest the use of bond increments V(s) to be added if the bonds include tetrahedral angles and subtracted if they include an angle +8 = 180~. The following semi-empirical values are recommended for V,,) (kcal.) of bonds rotating at standard distances from the CH group C-H = 0.95 C-F = 1.1 C-Cl = 1.0 C-C = 1.05 0-H = 1.0 S-H = 1-35 N-H = 1.1 ~0-C = 2.0 C-0 = 1.2 N-C = 1-75 C-N = 1-25 C=C = 2.6 c=o = 1.9.Single-bond increments correspond to the experimental dipole moments and quadrupole moments P N 0.755~ - 1.0 L = bond length. A simple procedure can also be devised for estimating the contributions A in eqn. (3) 34 GiguCre J . Chem. Physics 1950 18 88 ; Can. J . Res. B 1950 28 485. 35 Scott et al. J . Amer. Chem. SOL 1949 71 2293. I 18 GENERAL DISCUSSION for the slight deviation of angles and from their standard values and for barriers with +8 = 180° i.e. with n = 6 minima. Thus internal potential barriers can be built up which agree closely with best available values obtained from spectroscopic and C measure- ments. This appears to present a simple method of predicting unknown barriers e.g. ethyl amine = 3-75 a-methyl pyrol = 1-5 cc-methyl furane = 1-5 for those who do not wish to be concerned with theoretical details.A full account of this work is to be published elsewhere. Dr. M. Magat (Paris) said Berthelot s6 in our laboratory has investig- ated recently the rotational isomerism of the lower aliphatic alcohols by' the Raman spectra method. As ought to be expected isomers due to rotation around the C-C bond were observed for n-propanol and n- butanol. The differences in energies for the two isomers are in good agreement with those quoted by Prof. Aston 0.82 + 0.18 kcal. for propanol and 0.67 & 0-19 kcal. for n-butanol as against 0.78 &- 0.16 kcal. for butane 3' 0.2 kcal. for n-pentane.s8 But rotational isomers or tautomers exist also in ethanol where they are due to rotations around the C-0 bond.The energy difference between the two isomers is slightly higher 1-0 f 0.2 kcal. It is remarkable that this energy is so low despite the existence of hydrogen bonds. and 0.62 Prof. E. A. Guggenheim (Reading) said The procedureused by Aston for computing the thermodynamic functions of butadiene by treating this as a mixture of trans and cis tautomers is valid only a t temperatures sufficiently low that states near the top of the barrier make no contribution. When the temperature is sufficiently low for this condition to be satisfied this procedure is unnecessarily complicated. In my opinion i t is simpler to construct a complete partition function for all states which make an appreciable contribution. This is quite easy because at these low tem- peratures the number of such states is small.Alternatively no appreci- able error would be made by using as partition function the sum of two partition functions for harmonic oscillators. Such a procedure avoids any need to consider the entropy of mixing. 38 Mizushima and Okazaki J . Amer. Chem. Soc. 1949 71 3411. Dr. N. Sheppard and Mr. J. K . Brown (Cambridge) (communicated) In his paper on hindered rotation in hydrocarbon molecules Prof. Aston has made the important point that at least in the case cf 2 3-dimethyl butane the spectroscopic data so far available indicate the possibility of solid solution between rotational isomeric configurations.39 40 This might imply a residual entropy a t the absolute zero a fact of great importance in the theoretical interpretation of the measured value of the entropy.The evidence for solid solution comes from the previous observation that although the liquid-state Raman and infra-red spectra are consistent with the existence of two rotational isomeric forms of this molecule with only a small energy difference,3Qj 40 4l no simplification occurs on transition to the solid state. Such a simplification occurs on crystallizing the straight chain hydrocarbons e.g. w-butane,40$ 41. corresponding to only one configuration being stable in the solid state. We have recently been able to show however that rapid cooling of this branched paraffin such as has been carried out in previous spectroscopic work leads to an amor- phous glassy solid but that on careful warming this glass undergoes a 36Berthelot Compt.rend. 1950 231 1481 ; J . Chim. Phys. (in press). A misprint in the Compt. rend. papers attributes to Aston et al. ( J . Chem. Physics 1944 12 336) the idea that the cis form is the stable one. In fact we agree with this author that tvaizs is the stable form throughout. 37 Szasz Sheppard and Rank J . Chem. Physics 1948 16 704. 30 Szasz and Sheppard J . Chem. Physics 1949 17 93. 40Axford and Rank J . Chem. Physics 1950. 18 51. Sheppard and Szasz J . Chem. Physics 1950 18 145. 42 Rank Sheppard and Szasz J . Chem. Physics 1949 17 83. GENERAL DISCUSSION I I 9 transition to a crystalline mass.43 This glass to crystal transition is accompanied by a very marked simplification in the infra-red spectrum indicating that only the centro-symmetric isomer is stable in the crystalline lattice.Hence as a result of this new evidence there now seems little possibility that a residual entropy due to the solid solution of different isomers can occur for this molecule under the usual experimental con- ditions of calorimetric work. The specboscopic simplification accompanying crystallization also leaves no doubt that two rotational isomeric species are present in con- siderable concentration in the liquid state corresponding to a low energy diff erence.39 Dr. K. S . Pitzer (Washington D.C.) said Although strictly outside the scope of a meeting on hydrocarbons the problem of internal rotation in methvl alcohol has received much attention and was mentioned by Prof.Aston. Consequently i t seems appropriate to mention that Dr. Weltner has measured the heat capacity of the vapour as a function of pressure and has found a striking non-linear pressure dependence below I atm. near the boiling point. We have interpreted this as arising from the presence at eqcilibrium of a small fraction of a tetramer presumably held together with hydrogen bonds. This is analogous to the situation with hydrogen fluoride. The reason this finding is important for the internal rotation problem is that the entropy of methyl alcohol in the ideal gas state has been obtained on the assumption of normal gas im- perfection behaviour and i t is such values that disagree with the spectro- scopic results. When our virial coefficients are used the values are in substantial agreement.The potential barrier will be obtained much more exactly from spectroscopic data than from ours but will be in the vicinity of 1000 cal./mole. Prof. J. G. Aston (Pennsylvania) (communicated) In connection with Pitzer’s comments on my paper recently E.M.F. determinations of the hydrogen chloride partial pressures above alcohol solutions of methyl- amine (partial pressure known) saturated with methylammonium chloride have yielded a value of the entropy of the reaction HC1 (g) + CH,NH (g) = CH,NH,CI (s) which at present agrees with experiment only if the spectroscopic entropy of methylamine be used and a zero-point entropy of R In 2 assumed in methylammonium chloride. The former may indicate residual entropy in methylamine at the absolute zero and the latter indicates lack of dis- crimination between the CH and NH ends in the methylammonium ion.The calculation of the spectroscopic entropy of methylamine made use of a line in the infra-red corresponding to a barrier of 1520 cal./mole and yielded a value 0-5 cal./mole deg. higher than the calorimetric value.44 Dr. J. W. Linnett (Oxford) said Stitt 45 studied the vibrations of ethane and hexadeuteroethane and concluded that “ one must introduce interaction terms between the two methyl groups” into the potential energy function to account for the observed vibration frequencies. This as Stitt pointed out is consistent with the existence of the barrier to free rotation which also requires an interaction between the methyl groups in ethane.On examination of Stitt’s potential energy function i t is found that if one methyl group is distorted symmetrically’ the bond lengths and three-fold axis of symmetry being maintained the configuration of the other methyl group which minimizes the potential energy is almost the same as if the first methyl group were not distorted. So for this particular distortion there appears to be very little interaction between the methyl groups. However i t is not my object to stress any result but to recall that in the measured vibration frequencies of ethane and 4 3 Brown and Sheppard J . Chem. Physics (in press). “Aston and Doty J . Chem. Physics 1940 8 743. 45 Stitt J . Chem. Physics 1939 7 297 I 2 0 GENERAL DISCUSSION hexadeutero-ethane there exist additional data which might be used for studying the interaction between the methyl groups and for testing theories about this interaction.Dr. J. Sheridan (Birmingham) said I should like to hear views on the possible extension of our understanding of restricted rotation by measure- ment of the splittings within each torsional vibration Such measure- ment is now becoming possible from the pure rotation spectra of certain molecules which lie in the microwave region. In the symme'ric top molecules H ,CCF 3,47 H ,CSiF 4* 49 and H ,CSiH 3,50 successive torsional levels are associated with increasing moments of inertia (I,) for the whole molecule and hence with successively lower frequencies for each pure rotational absorption.The lowering of frequency relative to the ground state spectrum appears roughly proportional to the energy of torsional vibration which can be estimated from the relative intensities of the spectra at known temperatures. For H ,CSiF rotational absorptions for the first and second excited torsional levels are found 49 to be doublets the first close the second considerably wider. This is interpreted as due to the splitting of each of these torsional states into two levels. Since roughly known torsional vibration quanta produce known frequency separations between the doublets and the ground state absorption the frequency separations of the doublets themselves can be used to estimate the extent of the splitting of the levels.The first and second excited torsional levels are about 140 cm.-l and 280 cm.-l respectively above the ground level the first being split into two levels about 1-4 cm.-l apart and the second into two about 12 cm.-l apart. Only doubling of the levels is observed here which is not unexpected since the SiF group contributes 96 yo of the moment of inertia of the molecule about the symmetry axis (Ib) ; the methyl group undergoes torsional vibration virtually against a rigid franie.51 In H,CSiH,,50 where the two groups contribirte nearly equally to I* the expected splitting of torsional levels into triplets (for K > 0) has been reported. Dr. J. J. A. Blekkingh (Rotterdam Holland) (communicated) A clear distinction has been made between isomers and so-called transitional fornzs of each of these isomers.52 IsomeIs are characterized by a certain distribction of the substituted atoms or groups over the available 12 non-ring valencies of cyclohexane.When the distribution is different then another isomer is obtained. The various possible forms which one and the same isomer can take on when the atoms move to a greater or lesser extent with respect to each other without however any change in the characteristic distribution of the substituted atoms or groups over the 12 valencies taking place are called transitional forms. Hassel 53 distinguishes two kinds of valencies in the immovable forms viz. 6 erect ( E ) and 6 lying ( K ) . This formclation does not lead to all the possible isomers. Nor is this the case when use is made of a flat six-membered ring as is still done in the most recent literature ; 54 here also two kinds of valencies are recognized viz.6 above and 6 beneath the plane of the ring This may be why as a rule only 8 different isomers of C,(HX) are distinguished whereas in reality g different isomers are theoretically possible which are completely equivalent and have exactly the same right to the name " isomer ". 46 Herzberg Infru-red and Raman Spectra (van Nostrand 1g45) p. 252. 47 Dailey Shulman and Minden Physic. Rev. 1949 75 1319. 48 Minden Mays and Dailey Physic. Rev. 1950 78 347 ; Minden and Dailey Abstr. Amer. Physic. Society Meeting (New York) Feb. 1951. 49 Gordy and Sheridan J . Chem. Physics (to be published). 5O Lide and Coles Physic.Rev. 1950 80 91 I. 51 See e.g. Koehler and Dennison Physic. Rev. 1940 57 1006. 52 Blekkingh Rec. trav. chim. 1949 68 345. 53 Hassel Tids. Kjemi Bergvesen Met. 1943 3 32. 54 Cristol Hause and Meek J . Amer. Chem. SOC. 1951 73 674. I 2 1 GENERAL DISCUSSION In accordance with the trigonal symmetry of the immovable forms of cyclohexane the IZ available valencies should be divided into four kinds viz. 3 straight up ( E ) 3 straight down ( - E ) 3 inclined upwards ( K ) and 3 inclined downwards ( - K ) . This formulation is correct but how- ever does not take into account all the elements of symmetry and there- fore gives rise to mistakes rather easily. To exclude any possible mis- understanding the simplest and most accurate formulation is that in which the 12 valencv directions are indicated by the figures I 2 3 4 5 6 and I’ z’ 3.’ 4i 5’ 6’ as in the accompanying diagram.The valencies of the immovable form I which are directed straight up are numbered anti-clockwise I 3 5. The valencies of this form which are directed straight down are numbered clockwise 2 4 6. The remaining (lying) valencies bear the same numbers but with accents as the straight valencies of the carbon atoms opposite in the ring. As is obvious each immovable form can be placed in six positions. The numbering is such that there are suc- 5 I cessive the valencies 3 4 2 transitions ) J 1 1 , 4 3 6 I from 6 5 2 I 5 6 2 1 of 4 6 3 ‘ 9 the 3 5 4 I 1 to 2 3 4 5 6 directions Rule A of The same with the numbers with 15 FIG 2.3 > 9 I J J 5 4 2 3 2 I accents. The six possible positions of one and the same form can consequently be deduced from each other with the aid of rule A i t is sufficient if only one form is known. By intramolecular movements involving slight deformation the im- movable form I can be changed via a movable form into the immovable form 11. The “ back ” of the chair becomes the “ foot end ” and vice versa the chair is turned round by the intramolecular change. By turning the whole chair i t returns to the same position as the immovable form I. Through this process of changing form I into form I1 by intra- molecular movements and afterwards adjusting the chair again to the original position of form I the following directions are interchanged ,> I J 4 2 3 I becoming 4’ 3‘ 2‘ I ’ and > 9 , 4’ 3’ 2‘ I’ becoming , 4 I \ \Rule B.7 9 5’ > 6’ , 6’ 5 6 5’ Y 3 3 5 6 , With the aid of rule €3 i t is thus possible to change an immovable form in which the valency I’ occurs into the other immovable form with the valency I. The formulation of an isomer is chosen from the twelve possible formula- tions (six positions of the one immovable form and six of the other) such that the formulation always begins with the valency numbered I followed by the lowest possible numbers. GENERAL DISCUSSION I22 Mention may be made of the special case that there is only owe im- movable form of the isomer concerned.Rule A indicates the possible positions and by’ application of rule 8 to them the turned chair positions are obtained. If the last-mentioned positions are identical to the normal ones the two immovable forms are identical. This is the case with the isomers II’ IZ‘ 16’ 11/22‘ 11’33’ 11/66’ 12’3’4 11/22/55’ and 11’33’55’. There are only six different movable transitional forms corresponding to the six positions of the one immovable form and the six positions obtained from them by applying rule B (therefore corresponding to the other immovable form). These movable forms all possess the well- known boat form. The distance between the two opposite carbon atoms with the valencies I (4’) and I’ (4) is slightly smaller than that between the other pairs of opposite carbon atoms.It is however not correct to distinguish two so-called stretched configurations and two boat con- figuration~.~~ There are in fact six boat forms in which the valencies according to rule A differ from each other e.g. 11’ 22’ 33 44’ 55’ and 66’. There are therefore a maximum of eight different transitional forms of each isomer two immovable and six movable and between the movable ones there is of course an infinite number of intermediate positions. Each isomer of a cyclohexane derivative is completely defined in the way de- scribed by indicating the valencies I 2 3 4 5 6 and 1’ 2’ 3’ 4’ 5’ 6’ to which the substituting atoms or groups of atoms are bound.Prof. L. J. Oosterhoff (Amsterdanz) (communicated) In reply to Dr. Blekkingh’s remark all the boat configuration stretched configurations and intermediate configurations have been considered as appears from the symmetry number used. Dr. Julian H. Gibbs (Princeton) said In connection with the paper of Hazebroek and Oosterhoff I should like to mention some results which Dr. Smyth and I obtained recently a t Princeton. We measured the molar polarization of I 4-dioxane over a range of temperatures in the vapour phase in order to obtain some information concerning the population of the molecules in the various possible con- figurations. For dioxane these possible configurations are (I) the sym- metrical cis form of symmetry Czv in which both oxygen atoms project out on the same side of the plane of the four carbons ; (2) the two un- symmetrical cis forms of symmetry C, in which two para methylene groups project out on the same side of the plane of the other two carbon atoms and the two oxygen atoms (these two forms being mirror images of each other) and (3) the trans form of symmetry Czh in which the two oxygen atoms project out on opposite sides of the plane of the four carbon atoms.56 The latter form is the analogue of the chair tautomer of cyclo- hexane and the symmetrical cis and unsymmetrical cis forms together form the analogue of the flexible or boat tautomer of cyclohexane.Our molar polarization values are plotted against reciprocal temper- ature along with some obtained from earlier data of Schwingel and Greene 57 and Kubo 56 on the appended figure.In making this plot it was necessary to lower all the values of Kubo by 6-5 yo in order to bring the mean height of the curve through his points down to that of Schwingel and Greene and our own. An inaccurate cell constant determination is probably responsible for this discrepancy between Kubo’s magnitudes and those of Schwingel and Greene and ours. At any rate we are primarily inter- ested in the shape of the curve rather than in its height ; therefore it is justifiable to make this alteration in his values in order that they may be more easily compared with our own and those of Schwingel and Greene. It can be seen from the plot that the polarizations drop slowly and more or less linearly with increasing temperature (decreasing I/T) at the 55 Oosterhoff Thesis (Leiden 1949) ; Hazebroek and Oosterhoff this Dis- cussion.Dallinga Thesis (Leiden 1951). 56 Kubo Sci. Papers Inst. Phys.-Chem. Res. Tokyo 1936 29 122. 67 Schwingel and Greene J . Amer. Chem. Soc. 1934 56 653. 123 GENERAL DISCUSSION lower temperatures and then rise sharply with temperature at the higher temperatures. It is interesting to note that the highest temper- ature polarization values obtained in both previous investigations in- dicated this rise a t higher temperatures. However in both cases the existence of a minimum and subsequent rise was not mentioned because the indications of this were only barely outside the estimated experi- mental error.This curve affords a means of determining the order of energies of the various forms of dioxane since the total molar polarization may be expressed as a function of temperature in terms of the dipole moments and energies of the various forms by means of the Boltzmann distribution law and since the moments of the various forms may be roughly evaluated (# Schwingel and Greene. 0 Kubo. Gibbs and Smyth. FIG. 3.-Molar polarization against reciprocal temperature for dioxane. by vector addition of the various group moments allowing for a small amount of inductive lowering of the moments of adjacent dipoles. That is if one assumes that the rotational and vibrational factors in the partition functions for the various forms are the same for all forms (they all have the same symmetry number) they may be cancelled and one obtains for the mean square moment in which the lowest electronic energies E of the various forms have all been referred to that of the trans form as a zero point.The meaning of the subscripts is obvious. No term representing a contribution due to the trans form appears in the equation since this form has no moment. This expression is t o be substituted into the Clausius-Mosotti-Debye equation written in the form P = A + BFZlT in which A and B are constants. One then has an equation for the experimental curve. Unfortunately i t is not possible to fit this equation to the experimental curve by adjusting the parameters E, and E,,. This is undoubtedly due to the approximations made in deriving the equations.However i t is possible to show that a dose approach to such a fit can only be obtained if the order of energies of the forms is E > E, > Et. That the trans form is the most stable is indicated by the small values of the polarizations. If E, and E, are not too close to each other in magnitude the equation exhibits two maxima each one being due to one of the polar forms. The interpretation is then that the experimentally observed portion of the curve is that lying between the two maxima. I f the energy of the highly polar symmetrical cis form were lower than that of the moderately polar unsymmetrical cis I 24 GENERAL DISCUSSION forms the lower temperature maximum would be due to the symmetrical cis form. That the situation must be the other way around is indicated by the small polarizat'on values and positive slope of the right-hand portion of the experimental curve.That is the small concentrations of symmetrical cis which would be required to give small enough polariza- tions would also give polarizations which were still rising with increasing temperature whereas the low temperature maximum has already been surpassed in the experimental region. Furthermore the fact that the high temperature maximum is apparently larger than the low temperature maximum also indicates that E > E,. The fact emphasized by Hazebroek and Oosterhoff that the boat or flexible form of cyclohexane is in reality a series of interconvertible forms corresponds for dioxane to easy interconvertibility of the sym- metrical cis and unsymmetrical cis forms.The activation energy for this conversion is not zero since there is a difference in energy between the forms themselves and since there is a slight strain associated with the intermediate configurations. However the lack of a sizable activation energy for this conversion is undoubtedly partly responsible for our failure to obtain an exact fit of our theoretical equation to our experimental curve since the starting point in the derivation of this equation is that the problem may be analyzed into contributions of distinct forms with each of which a partition function may be associated. This whole treatment was based on the curve of molar polarization against reciprocal temperature rather than dipole moment against tem- perature in order to avoid the necessity of making a guess at the mag- nitude of the atomic polarization which is involved in the calculation of a dipole moment for each temperature.It is true that in the unsymmetrical cis form a pair of para methylene groups are fairly close to one another so that one might at first sight expect this to be the form of highest energy. However in the symmetrical cis form the two oxygen atoms are placed in such a way that the repulsive interaction between the approximately tetrahedrally oriented lone-pair electron orbitals should be large thereby giving a high energy to this form. Dr. K. S. Pitzer (Washington D.C.) said The very thorough analysis of the electrostatic contribution to the internal rotation restricting barrier in ethane by Dr.Oosterhoff tends to reinforce my conclusion that more than one phenomenon contributes to the total barrier. In this regard I doubt if Eyring's calculations have proved that distortions of the spa orbitals on the carbon atoms have a negligible effect. I have made rough calculations which tend to indicate a very appreciable effect from this source although the calculations are too crude to have quantitative significance. Prof. A. R. Ubbelohde (Belfast) said In the freezing of cyclohexane i t has been observed that even after numerous recrystallizations the mother liquor freezes about 0 - 0 2 ~ below the crystals. This could be due to a temporary enrichment of the " boat " isomer in the last part of the liquid to freeze if the rate of the processes boat Z-f chair is comparable with the time of crystallization.The percentage enrichment required to explain the results does not exceed 5 8 about 0.008 yo. glutaric acid (CH,(CH . CH,COOH),) exists in two optically inactive Dr. C. N. Davies ( L o ~ d o n School of Hygiene) said ad-Dimethyl forms the meso or maleinoid (m.p. 128OC) which readily yields an an- hydride with acetyl chloride a t low temperatures59 and shows other evidence of facile cyclization,60 and the racemic or fumaroid form (m.p. 141 O C) which does not yield an anhydride except under extremely vigorous 58 Cf. Thompson and Ubbelohde Truws. Faraday SOC. 1950 46 349. 59 Auwers and Thorpe Annulen 1895 285 31 I. 60 Thorpe and Young J .Chem. Soc. 1903; 83 354. FIG. 46. FIG. 4a. FIG. 5b. FIG. 50. [To face page 125 GENERAL DISCUSSION 125 treatment when i t is converted into the anhydride of the maleinoid acid.61 Of these two acids only the racemic is of course resolvable into its enantiomorphs (m.p. 80" C) ; these optically acids differ from the inactive acid from which they are obtained in that they yield optically active anhydrides almost as readily as the malenoid acid.G2 The following interpretation of these facts is suggested. Inspection of models shows that the meso or maleinoid acid might exist in three forms each designated (+ -) with reference to the two asymmetric carbon atoms they contain. The first form which is believed to correspond to the known acid (Fig.4a) is internally compensated but two optically active enantiomorphs are also conceivable. The arrangement shown in Fig. 4a postulates a strong dipole inter- action between the carboxyl groups with resonance which confers a plane of symmetry on the molecule. This cis (+-) acid is thus a true vneso acid and is identified with the maleinoid acid (m.p. 128'). In view of the powerful interaction between the carboxyl groups this configuration should be stable and especially favourable for cyclization. In Fig. 4b two enantiomorphs are shown which can be formed from Fig. 4a by simple rotation about bonds. With the (+-) acids i t is only possible to achieve enantiomorphs in which the comparatively weak dipole attraction between a single methyl and a single carboxyl group is effective.This may be insufficient to prevent bond rotation taking place at ordinary temperatures and thus account for the fact that these acids are unknown. The racemic or fumaroid acid is a ( f +) or ( - -) acid and two pairs of enantiomorphs are theoretically possible. Fig. 5a shows models of the cis forms which arise from carboxyl group attraction as in the meso acid. These should also form anhydrides readily and we regard the optically active acids (m.p. 80") as having this structure. The optically inactive cis acid (&&) is not known but could presumably be obtained by mixing equal amounts of these isomers. Resonance between the carboxyls in this case does not confer a plane of symmetry as in the model of the meso acid.From these models simple bond rotation leads to the molecules indicated in Fig. 9. These are trans (+ +) and trans ( - -) acids. Each carboxyl group is favourably placed for association with a methyl group so that these forms are stabilized by h - o internal carboxyl-methyl links instead of one which as we saw above is apparently insufficient to pre- vent bond rotation. These acids would clearly be reluctant to form anhydrides. We regard the fumaroid acid (m.p. 141~) as the externally compensated mixture of these two that is the trans ( f -&) acid. The individual acids are unknown. Dr. Manfred Gordon (Royal Technical College Glasgow) said Ubbelohde and McCoubrey have drawn attention to the influence of molecular coiling on the reactivity of hydrocarbons in the gas phase.Cyclization reactions might well provide some quantitative information in this respect by cor- relation of rate measurements with theoretical calculations based on suitable models of kinked or coiled structures. As an example consider the cyclization of rubber. Though this has been studied in the liquid and not in the gas phase similar effects of coiling may be anticipated in solution. It has recently been demonstrated 63 that the reaction leads exclusively or very predominantly to ring formation between adjacent isoprene units of the same rubber chain. First an ethylenic carbon of one unit is converted to a carbonium ion. When this attacks another double bond to form a new C-C bond i t is thus in practice the double bond of the adjacent unit which is attacked most rapidly i.e.the double . 6a Moller Moller Ber. Lund's 1910 Univ. 43 %%. 1919 16 56 ; Ckem. Abstr. 1920 14 942. 63 Gordon PVOG. Roy. SOC. A 1951 204 569. GENERAL DISCUSSION available to the chain in which the two carbons to be linked come within bond- ing distance. Clearly f is a measure of the fraction of time geometrically favourable to reaction and should cor- I 26 bond in the neighbouring loop of the hydrocarbon coil (Fig. 7) reacts faster than more distant double bonds of the same or other chains. Quite analogous reactions can be carried out with low molecular weight terpenes and i t appears feasible (though this is a guess) that dihydromyrcene could be cyclized with BF in the gas phase.It would be most instructive if the rates of such cyclizations could be measured on different model compounds having the necessary two double bonds at different spacings. The rates r of such cyclizations should be compared with calculations based on suitable models of the fraction f of equi-energetic configurations ?f->Df 0 C + c - r ; FIG. 6. relate with Arrhenius frequency factors. Moreover f would vary widely from one model to another e.g. from a randomly kinked to a more regularly coiled model. Do Ubbelohde and McCoubrey agree that such rate measurements and cal- C culations would in principle promise to give information on molecular configuration complementary to the results of physical measurements ? Dr.M. Magat (Paris) said I would like to make three comments on the very interesting paper presented by Prof. Ubbelohde. I would like to call attention to a different way in which flexibility of hydrocarbon chains influences the reaction rate. I n certain cases the so-called steric hindrance in bimolecular reactions can be traced back to this property of hydrocarbons. Besides the cases analyzed by Ingold Hughes and their co-workers and by A. G. Evans where the steric hindrance increases the activation energy there are cases where the reaction rate decreases along a homologous series the activation energy remaining constant. Such cases have been recently compiled by Miss Ivanoff and myself 64 and we found that they all concern “ chain ” hydro- carbons and nucleophilic reactions.If one now considers the correspond- ing models i t appears that the number of possible chain configurations is reduced in the transition state as compared to the initial state. For instance i t is clear from the model that in the initial state acetone can have (disregarding the hydrogen positions) just one configuration while diethyl ketone has nine configurations (3 different positions for each -CH,. group). In the transition state for the reaction with semi- carbazine acetone and diethylketone can both have just one confi9.u ation (the “ pure trans ” one in the last case). Hence everything else being unchanged the entropy of activation is larger for the diethyl ketone than for acetone and one finds immediately that the rate ought to be g times slower for the diethylketone than for acetone.The experimental frequency factors are according to Price and Hammett 65 2-55 and 0.25 in good agreement with the prediction. It can be shown 6 6 that the changes in the translational and rotational terms of the activation entropy could account for a rate decrease of only 30-40 :/o. Prof. Ubbelohde assumes that relatively low normal paraffins are ‘‘ bunched ” or “ coiled ” in the liquid state. How is this to be reconciled with the findings of Kratky on the di-iodo-1 11-undecane in undecane solution for which X-ray diffraction favours almost entirely a stretched configuration ? Prof. Ubbelohde suggests that the probability of energy exchange from kinetic into vibrational on collision is favoured by the existence of low-frequency vibrations in non-rigid molecules.Although this idea 64 Ivanoff and Magat J . Chim. Physic. 1950 47 914. 65 Price and Hammett J . Amer. Chem. SOC. 1941 62 2387. 66 Bauer and Magat J . Claim. Phys. 1950 47 922. GENERAL DISCUSSION 127 appeals to me i t seems to be in contradiction with the formula for the probability of such an exchange derived by Jackson and Mott G 7 in which this probability is proportional to the frequency Y. Dr. K. S . Pitzer (Wmhington D.C.) said With respect to Prof. Ubbelohde’s paper i t seems to me that the indication is that the viscosity data are not simply interpretable in terms of the coiling of n-paraffins. The fully coiled configuration is 500-1000 cal./mole higher in energy for each C-C bond over the extended form as has been well established now by both statistical thermodynamic and spectroscopic methods.Thus a n-paraffin must become more coiled with increasing temperature whereas the viscosity ratios in Table I1 of Prof. Ubbelohde’s paper show no sig- nificant change with temperature. Prof. E. G . Cox (Leeds) said It should be pointed out that a comparison of the “ straight zig-zag ” and “ helix ” as representations of extremes of possible hydrocarbon chain configurations may be misleading unless i t is realized that there must be many zig-zags in one turn of the helix ; thus the helix is not a very realistic model to use for the discussion of the small molecules (up to octane) chiefly discussed in this paper.I should like to ask Prof. Ubbelohde whether he has been able to make any critical assessment of the accuracy of the various methods of deter- mining apparent molecular diameters ; can viscosity measurements for example give results which are in some way better or more accurate than those from diffusion methods ? Mr. R. S . Bradley (Lee&) (communicated) The point raised by Prof. E. G. Cox is mentioned in my paper with Dr. A. D. Shellard G8 where i t is stated “the hypothesis of coiling has previously been advanced by Mack and his co-workers to explain the comparatively low values of the collision radii of the lower hydrocarbons although up to C,H, the ‘ helix ’ on our model would occupy scarcely one t u r n ”. In view of the small amount of coiling of the lower hydrocarbons i t is somewhat doubtful whether any shielding of end groups occurs with hydrocarbons up to octane.In any case the term “ coiled molecule ” is misleading and might with advantage be replaced by “ crumpled mole- cule ” and the end groups of a long hydrocarbon are not necessarily near to the surface of the molecular “ drop ”. The work of Prof. Ubbelohde supports our viems on the normal paraffins Cl6-CI8 for which we found from diffusion measurements in air collision areas in agreement with the values based on statistical “ coiling ” and also the more recent work of the writer with Waghorn 6 9 in which we find nearly the same value for the collision radius for (n-C,H,,) ,CH (447 %i at 25OC) as for n-C18H38 (4-53 at 25O C).68 Proc. Roy. SOC. A 1949 198 239. 69 Proc. Roy. Soc. A 1951 206 65. Prof. A. R. Ubbelohde and Mr. J. C. McCoubrey (Belfast) (corn- municuted) With regard to the condition of n-paraffins in the liquid state present information suggests that it is important to distinguish between the degree of crumpling near the f.p. and the greater crumpling near the b.p. In general liquids behave as quasi-crystals near the f.p. and quasi-gases near the critical temperature. It is anticipated that at the b.p. which is approximately equal to two-thirds the critical tem- perature liquids should be quasi-gaseous. Our data on molar volumes and entropies of vaporization suggest moderate crumpling at the b.p. Near the f.p. considerably greater adlineation is suggested by existing evidence.”J The presence of dipoles could further favour this adlineation so that there is no present inconsistency between X-ray work on di-iodo- I II-undecane and our molar volumes.The exchange of energy referred to by Jackson and Mott deals with a 137 Jackson and Mott Proc. Boy. SOC. A 1932 137 703. 70 Cf. Quart. Rev. (in press). I28 GENERAL DISCUSSION much more restricted process than the various modes of transfer from kinetic to vibrational energy in flexible hydrocarbons. In reply to Prof. Cox the present consistency of collision diameters as obtained by’ various authors has been carefully reviewed.71 In the most favourable cases the agreement is within I yo of the measured viscosities. We do not necessarily expect collision diameters in homomolecular assemblies to agree with collision diameters in hetero-molecular assemblies.We agree with Mr. Bradley that “ crumpled ” is a better term than “ coiling ” for molecules of such chain length that they can still be readily vaporized. Dr. Peter Gray (Cambridge) (communicated) Ubbelohde and Mc- Coubrey state that in coiled flexible hydrocarbon molecules intermediate CH groups may have one and frequently both C-H bonds blocked owing to the coiling. But if a coil is formed will not these C-H bonds tend to be on the “outside ” of the coil and no less accessible than in an extended molecule ? Dr. P. Torkington (Brit. Rayon Res. Assoc.) (communicated) Prof. Ubbelohde implies in his paper that vibrational coupling is advantageous to a concentration of energy at a particular part of the molecule.One would have thought that such a concentration of energy would be analogous to a decrease in entropy and that the coupling of vibrations would deter not aid any drift towards an activated state. Prof. Ubbelohde (Belfast) (communicated) The coupling of vibrations in a molecule permits the transfer of vibrational energy to any particular bond where reaction is to occur. If there is a break in the coupling the vibrational energy in separate parts of the molecule is not generally available for activation beyond the break. . Dr. A. F. Trotman-Dickenson (Munchester) said Ubbelohde and McCoubrey have cited the supposed decrease in the steric factors of the reactions of methyl radicals with alkanes containing primary secondary and tertiary hydrogen atoms as evidence in favour of the theories out- lined in their paper.The evidence for the decrease comes from an inter- pretation 7 2 of early experimental work.73 However the figures cannot be relied upon because i t has been shown 74 that the experiments were in- adequate and that the interpretation is unsuitable. Dr. Steacie and I 75 have obtained values for the steric factors of these reactions and the parallel ones with the alkenes where the cc-methylenic hydrogen is attacked. These values are given in the annexed table. A statistical correction for TABLE III.-THE STERIC FACTORS FOR THE REACTIONS OF METHYL RADICALS WITH ALKANES AND ALKENES Alkanes 1 Alkenes 3 StericFactors x xd I 2 Type of H atom Primary Secondary .Tertiary 2 2 3 . 30 These figures do not appear to support the theories of coiling. the number of H atoms involved has been made ; 1-5 and 3-5 A were taken as the collision diameters of the hydrogen atoms in the hydrocarbons and the methyl radicals respectively. 71 J . Chem. SOC. (in press). 73 Steacie Darwent and Trost Faraday SOC. Discussions 1947 2 86. 73 Smith and Taylor J. Chew. Physics 1939 7 390. 74 Trotman-Dickenson and Steacie J. Physic. Chem. (in press). 76 Trotman-Dickenson and Steacie J . Chem. Physics 1951 19 329. GENERAL DISCUSSION GENERAL DISCUSSION Prof. A. R. Ubbelohde (Belfast) said In the ionization of a poly-methylene molecule how far is i t permissible to regard the charge a,s located a t a specific part of the molecule ? To select one out of a number of possible examples is it significant to write the process as one of a number of distinguishable alternatives ? Hydrocarbon -f hydrocarbon + electron For example, CI-I3-CH2-CH2-CH2-CH,-CI-I,-CH2-CH, \ CH~-CH2-CH2-CH~-CH2-CH2-CH2-CH3 + I,, or CH~-CH~-CH~-CH~-CH2-CH~-CH2-CH3 + I , or CH3-CH2-CH2-CH2-CH~-CH2-CH2-CH3 + I, or CH3-CH2-CH2-CHZ-CH2-CH2-CH-CH3 + I,, + + i-+ in which the ionization energies I to I need not be the same.If these processes cannot be distinguished the molecule could be regarded as a linear conductor for the charge rather like the polyenes. But if the above states are distinct it becomes relevant to consider : (i) The rate of transfer of an electron from one part of the molecule to another and how the time required for transfer compares with the time before disruption of the molecule in the mass spectrograph.(ii) How such transfer is affected by structural modifications to the polymethylene chain such as substitution of H by CH or other groups, or the replacement of single bond C-C skeletons by unsaturated skeletons. (iii) M’hether any proton migration can occur in the gas phase in ion-ized unsaturated molecules The analogy between these problems and conductors semi-conductors and insulators in three-dimensional systems may be suggestive for certain problems of molecular reactivity of ionized molecules. Sir John Lennard-Jones (Cambridge) said It is not possible to give a brief answer to all the interesting questions put by Prof.Ubbelohde. The ionization of a molecule involves the removal of an electron from a molecular orbital spread over the whole molecule so that the electron distribution in the ionized molecules may be described as that of the neutral molecule minus the distribution of one electron appropriate to the vacated molecular orbital. Corresponding to this there is a proba-bility distribution of finding “ the hole ’’ or resultant positive charge in the various parts of the molecule just as the probability distribution of an electron is obtained from the orbital which i t occupies. It will not be correct to assign different ionization potentials (such as the I . . . I, mentioned in the question) to the several parts of the molecule.Each of the possible ionization potentials (for example those listed in Table I1 of our second paper) corresponds to the removal of an electron from a particular molecular orbital and is associated with a characteristic proba-bility distribution of the net positive charge. The quantitative deter-mination of these distributions is a possible application of the theory but no examples have been calculated yet. Since the alternative structures mentioned by Prof. Ubbelohde are not stationary states and therefore not distinguishable the remaining questions need reformulation and cannot be answered without a dis-cussion of the energy zones of molecules similar to that of the Brillouin zones in solids. For a sufficiently large molecule i t may be possible to define a time of relaxation which governs the rate of approach to one of the stationary states and this may be related to the transfer of charge along the molecule 1 04 GENERAL DISCUSSION Dr.P. Torkington (Brit. Rayon Res. Assoc.) (communicated) There are one or two points in Prof. Lennard-Jones’ papers which seem to have interesting even unexpected implications. Firstly if the factor determining configuration in general is repulsion between electrons de-scribable as in different localized bonding orbitals then in systems con-taining lone pairs presumably repulsions involving these will contribute. In particular in triatomic systems in which the central atom has two lone pairs the bond angle would be greater than or less than tetrahedral according as lone-pair repulsion was less than or greater than the inter-bond repulsion.One deduces that if Prof. Lennard- Jones’ hypothesis is correct the electron distributions in H,O and F,O are (apart from a scale factor) nearly identical and that in both molecules the lone-pair repulsion exceeds the inter-bond repulsion. Secondly ionization is said to be more closely related to the (equivalent) molecular orbital than the localized bond orbital description. Surely the first transition can only properly be described as from a localized orbital, even if the first excited state is a combination of localized orbitals, equivalent to a molecular orbital. The molecular orbital description would one supposes only apply to cases where there were several inter-vening excited states so that the ionizing model approximated closely in behaviour to the true molecule.Regarding the description of the cheniical bond as a potential well between two nuclei i t might be mentioned that the only satisfactory description of the hydrogen bond (as opposed to “ system giving the correct calculated energy ”) involves the concept of electrons over-localized in the perturbed X-H bond leaving an abnormally-exposed proton weakly bound to an atom Y having an inert pair the lengthening of the X-H bond being related to the inter-electronic repulsion within the X-H potential well. The perturbed X-H linkages in hydrogen-bonded systems are the only naturally-perturbed bonding potential wells that we have. It would seem that they are essentially similar to an excited state in which the molecular orbital has remained localized ; (this in turn incidentally being presumably one of the conditions for chemical reaction to take place).Lastly in the double bond I cannot see how the two descriptions U-T and bent-bond are equivalent. Surely in the bent-bond system there is negligible electron charge concentration along the line joining the two nuclei whereas in the (1-n structure there is a a-bond. I am assuming that the bent-bond can be looked at as arising from overlapping of two sp hybrids from each carbon atom; the initial angle with an angle of 1 2 0 O between the other two bonds on each atom being I O I O 32’. Actually this angle would not change at all without some separation of n-component. Is a better overall description one of resonance between the U-n and bent-bond models ? Dr.G. W. R. Bartindale (Manchestw) said I should like to ask Sir John Lennard-Jones what allowance has been made for the IS electrons of the Be atom and of the C atom of the CH molecule ? Sir John Lennard- Jones (Cambridge) said Although as Dr. Torkington points out the electron distributions of the separate orbitals are rather different in the u T and equivalent orbital descriptions the total electron density at any point is of course the same in both descriptions. Thus the electron density due to one equivalent orbital is + ( u + T ) ~ and that due to the other & ( u - T ) ~ . In reply to Dr. Bartindale the IS orbitals have much lower energies and are more localized than the others. Consequently no advantage is to be gained by including them in the transformation to equivalent orbitals.Dr. A. Burawoy (Manchsster University) (communicated) Sir John Lennard- Jones’s description of molecular structures by introducing equiv-alent orbitals must be considered as one of the outstanding theoretical The sum of these is ( u ~ + T ~ ) GENERAL DISCUSSION IOj developments for many years since i t lays the foundation to a better understanding of the relative importance of the factors determining the stabilities and shapes of molecules. It has many important qualitative features. One is the emphasis of the contribution and the need for con-sideration of the electrostatic repulsions of electrons forces which have been rather neglected or the significance of which has not been fully recog-nized in older approximate theoretical and semi-empirical treatments.Another is the justification of the old views of organic chemists that the multiple linkages in ethylene and acetylene can be accepted to possess equivalent though strained 1inkages.l However the authors still maintain that the stability of conjugated hydrocarbons in particular the observed strengthening and shortening of the single C-C linkages is due to a partial delocalization of the orbitals of the multiple linkages. Attention map be therefore called to some recently published papers in which the writer has shown that the valency conception of non-localized bonds has no justification being in disagreement with numerous if not all relevant observations. As an example this view demands that the multiple linkages in con-jugated hydrocarbons are longer and weaker than in ethylene and acetylene.This has been clearly stated and supported by approximate quantitative computations in past years by Sir John Lennard- Jones,3 M~lliken,~ Coulson and others. However investigations by the X-ray and electron diffraction methods show unambiguously that the internuclear distances of the multiple linkages in conjugated hydrocarbons remain unchanged or are in many cases actually shortened.6 This is well illustrated by the bond distances observed 7 for diacetylene dicarboxylic acid HOOC-CrC-C=C-COOH. The middle C-C distance is found to be only 1-33 A i.e. as short as the double bond in ethylene. In spite of it the C=C bond distances are only 1.185 A i.e.not longer but shorter than the CEZC distance observed for acetylene (1.204 A). As recently shown this and numerous other difficulties disappear if the valency conception of non-localized bonds is abandoned. The (con-stitutive) changes of linkages and their properties in all polyatomic mole-cules are easily accounted for by changes of the effective nuclear charge (the screening) of atoms i.e. changes of the electrostatic repulsions of electrons (and nuclei). Thus the increased bond energy of the middle C-C linkages in con-jugated hydrocarbons is not due to a delocalization of the orbitals of the multiple linkages and the resultant exchange forces but to the reduced repulsion between the electrons of the multiple linkage and those of the single bond (and possibly of the second multiple linkage) as compared with ethane.Similarly the reduced internuclear distance (and increased bond energy) of the CZC bond is expected. The replacement of a hydrogen atom in acetylene by the much more strongly electron attracting acetylene group will withdraw the electron cloud from the C atom of the triple bond. Its effective nuclear charge will be increased i.e. the repulsion between the triple bond electrons and those of the single bond will be reduced, and cause the shortening and strengthening of the CEC bond. For a discussion and explanation of the properties of multiple linkages cf. Burawoy V . Henri Mem. VoZ. (Desoer LiBge 1948) ; cf. also Trans. Lennard-Jones Proc. Roy. Soc. A 1937 158 280. Mulliken Rieke and Brown J .Amer. Chem. SOC. 1941 63 41. Coulson and Jacobs J . Chem. Soc. 1949 2805. Dunitz and Robertson J . Chem. SOC. 1947 1145 ; cf. also Jeffrcy and also Burawoy.2 Faraday Soc. 1944 40 537 ; Chem. and Ind. 1944 434. 6 For a summary of the available data cf. Burawoya2 .Rollett. hiatwe 1950 166 475 I 06 GENERAL DISCUSSION It is not a coincidence that the (writer’s) empirical analysis of the properties of simple and complicated polyatomic molecules (e.g. saturated and conjugated hydrocarbons) and the new theoretical analysis a t least of simple systems by the authors lead qualitatively to the same con-clusions the great importance of the electrostatic repulsions of electrons (and their variations) and the negligible contributions by exchange forces arising from delocalization.However the empirical analysis shows t h a t these conclusions are also valid for conjugated and other complicated systems. The acceptance of the equivalence of the bonds of a multiple linkage may have already removed the theoretical necessity if not justification, for the hypothesis of non-localized bonds which a t any rate appears to be in disagreement with experience. The purpose of this comment is to draw attention to the need for an unbiased theoretical treatment of the structure of conjugated systems which would give information on the relative importance of the true factors responsible for their characteristic properties. Sir John Lennard- Jones (Cambridge) (communicated) We appreciate Dr. Burawoy’s favourable comments on our papers.On the other hand he gives the bond lengths of diacetylene dicarboxylic acid and asserts that the unusually short links which are observed cannot be explained by the usual methods. We ourselves do not know of any molecular orbital calculations on such a molecule and cannot say whether his statement is true. It is evident that the system is a complicated one and cannot readily be treated by current methods but it would be unsafe to predict that these methods would not be successful. Dr. Burawoy thinks that the observations can be explained in terms of the change of electrostatic repulsion between the electrons of neighbouring links but it is difficult to form a judgment on this until a quantitative treatment is available. Prof. E. C. Baughan (Shrivenhaw) (communicated) There is one point in the papers of Lennard-Jones and Pop!e to which I would like to draw attention.Their theoretical principle that lone pair electrons play a vital part in determining structures is well supported for several ele-ments by an inductive survey of experimental stereochemistry by Sidgwick and Powell who summarize the evidence thus “ Nearly (but not quite) all the structures can be even more simply related to the size of the valency group by assuming that the mean positions of the electron pairs in this group are the same whether they are shared or not.” Dr. C. A. McDowell (Liverpool University) (comnzunicated) Sir John Lennard-Jones and Dr. Pople have indicated that methylene may exist in a triplet state in which case it would probably have a linear structure, or in a singlet state when i t would be expected to be angular.Very little is known about this compound but its ionization potential 9 is known to be 11.9 eV and I should like to ask these authors if i t would be possible to calculate the expected ionization potential for each of these two possible states. A comparison with the known experimental value might then help to decide whether methylene has a singlet or triplet ground state. I am grateful to Prof. Baughan for calling our attention to the paper by Sidgwick and Powell in which the structures of a large number of molecules are classified. These structures are readily underst;& if the lone pairs of electrons are assumed to be as important as the bond electrons and to have distribu-tions which fit in with the symmetry type.In reply to Dr. McDowell we have not calculated the ionization poten-tials of the methylene radical but the methods we have described may be applicable to such a structure. Sir John Lennard-Jones (Cambridge) said : PYOC. ROY. SOC. A 1940 176 153. Langer and Hipple Physic. Rev. 1946 69 691 GENERAL DISCUSSION 107 Dr . A. Burawoy (Munchester University) (communicated) The problem of the variations of bond energies in saturated hydrocarbons has a con-siderable bearing on the calculation of the bond energies and stabilization energies of most of the molecules discussed by Prof. Glackler. As recently l o pointed out the available experimental data allow two empirical inter-pretations of the changes of bond energies in saturated hydrocarbons.(i) The bond energies of the C-C and C-H linkages may be assumed to decrease in the order of participating CpTi. > C,,,. > Ctert. > Cquart. This generally accepted view leads to unlikely conclusions. The Table shows the calculated atomic heats of formation and the possible C-C and C-H bond energies for methane ethane propane, isobutane and neopentane i.e. the bond energies are arranged in the order C y r i . to Cquart. Assuming the C-C bond energies remain unchanged, the C-H bond energies will decrease a t least by 1.1 1.0 and 0.9 kcal. respectively (column I). Any decrease of the C - C bond energies will be accompanied by’ a correspondingly’ greater decrease of the C-H bond energies (for an example cf.column 2). In order to obtain unchanged C-C bond energies the C-H bond energy X in ethane must satisfy the equation X = 97.0 - a where u is the difference between the C-H bond energies in the methyl groups of ethane and neopentane. This difference can be only small and may be ignored for the purpose of this discussion. It will be observed that the C-H bond energies decrease more strongly than the C-C bond energies. This is contrary to theoretical expectations and experience according to which the polarizability of C-C is greater than that of C-H. The smallest possible bond energy decrease from Cpri. to Cquart. (3 kcal.) required by this interpretation also appears to be unreasonably high. Moreover this assumption leads to the improbable conclusion that the stabilization energy of the C-C linkage in propene (89.2 - 84-5 = 4-7 kcal.) is only slightly higher than that of the C-H linkages in ethylene (101.6 - 98.1 = 3.5 kcal.) and this small increase is only apparent since it depends on the assumption that the bond energy of the -=C-H linkage adjacent to the methyl group in propene is only 99-7 kcal.as compared with the 101.6 kcal. of the other =C-H bond energies .I1 (ii) The bond energies of the C - C and C-H linkages may be assumed to increase in the order Cprf. to Cpuart. . If the C-H bond energies remain unchanged the C-C bond energies will increase by 1.2 1.1 and 1.0 kcal. respectively (column 3). Any increase in the C-H bond energies can be shown to be accompanied by a correspondingly greater increase of the C-C bond energies (for an example cf.column 4). The C-C bond energies vary now more strongly than the C-H bond energies as one should expect. Moreover the C-4 bond energy in propene will now be a t least 85.6 kcal. even if the C-H bond energy in the methyl group is taken as high as 99.0 kcal. i.e. the C - C stabiliza-tion energy will be greater than 8-3 (85-6 - 77-3) kcal. This empirical analysis is independent of the physical cause of the bond energy changes which has been discussed elsewhere.12 It indicates that the second interpretation must be the correct one. Thus the C - C bond energy in ethane will be 77-3 kcal. or smaller by an amount dependent on the increase of the C-H bond energy passing from methane to ethane. The C - C dissociation energy of ethane is of course independent of any possible individual bond energies.It is 83.3 kcal. if the atomic heats of formation of the methyl radical and of ethane are 291.4 and 665-9 kcal. respectively. These values are not arbitrarily chosen. l o Burawoy V . Henri Mem. V d . (Desoer LiBge 1948) p. 80. l1 Based on the values suggested by Glockler this Discussion. Based on the heats of formation values at oo K from Selected Values 1s and L(C) ;3 169.7 kcal. J3 Nut. Bur. Stand. (Cir. 461) 1947 I 08 GENERAL DISCUSSTON TABLE I Substance HCH . MeCH . Me,CH . Me,CH . Me& . Q 2 392'4 665.9 941.8 121g*8 1499.6 I C-C G H 2 C-C C H - 98.1 89.9 96-0 88.8 94.1 87.8 92-4 86.9 -3 C-C C-H 4 C-C C H - 98.1 74'3 98.6 76.1 99.0 77'7 99'3 79'1 -Dr.P. Torkington (Brit. Rayon Res. Assoc.) (communicated) Prof. Glockler interprets the bond length against bond energ41 data of the series C,H, C2H4 C,H as demonstrating the tendency of C-H bonds of high energy to be associated with C - C bonds of high energy in general. One would have thought that this type of association can only be accepted as probable on the above evidence in systems in which the C - C bonds can increase their energy by taking some multiple-bond character. Does Prof. Glockler interpret his results for the normal paraffin series as implying some double-bond character at the ends of the chains ? It seems possible that the stronger bonds at the ends of the paraffin chains arise from an overall molecular orbital component in the structure which favours the ends irrespective of whether they are C-C or C-H bonds-the data might almost be taken as proof of the reality of such orbitals.But it would seem difficult to relate this phenomenon to the changes in bond strength arising from hybridization changes associated with unsaturation. Dr. A. G. Gaydon (Imperial College) said Regarding the dissociation energy of CH my observations on the effect of pressure and other con-ditions on the spectrum leave me in no doubt that the predissociation is strong and real and not of the type which could be confused with a per-turbation. The only possible loop-hole which I see is that the ground state of CH may not be ,l7 but an unobserved quartet state ; the failure to observe CH bands in the laboratory would appear to support this view, but as pointed out by ,z later speaker the occurrence of CH in inter-stellar absorption makes the possibility remote.Dr. George Porter (Cambridge) said Dr. Gaydon makes the suggestion that the 2 l 7 state of CH to which the dissociation energy of 80 kcal. refers may not be the ground state which is particularly interesting in view of the considerable discussion about the ground state of methylene. I think however that strong experimental evidence is available that the ground stat? is because although not obtained in absorption in the laboratory transitions from this level are observed in interstellar space. It is well known that owing to the rare occurrence of collisions in this region molecules exist in their ground electronic vibrational and rota-tional states a single line only being in general observed.transition is not strictly forbidden even in the absence of collisions and therefore i t would appear that the 2 l 7 state must have the lowest energy. Dr. G. J. Minkoff (Imperial College) said Dr. Stevenson has men-tioned the paper of Voge in support of a value of ca. 80 kcal. for C-H in CH,. This result followed however from the original assumption by Voge of an angle of goo for H-C-H in CH (since this implies C-H bonds of p rather than sp character). Dr. B. G. Gowenlock (Swansea) (communicated) From the values of D(CH,-H) D(sec.-C,H,-H) and D(tert.-C,H,-H) given by Stevenson together with the value of D(C,H,-H) used by Glockler we have values of 0 4 7 and Ir#g kcal.for the resonance energies of the methyl ethyl, sec.-propyl and tert.-butyl radicals respectively. It can thus be seen that these experimentally observed figures agree in relative order of The *Z - GENERAL DISCUSSION 109 magnitude with the calculations of Baughan Evans and Polanyi.l4 The calculations of these authors give resonance energies for the above radicals in the relative order of magnitude o I z 3 corresponding to the con-tribution of 0 3 6 and g equivalent hyperconjugation structures to the actual state of the radical. Thus the electron impact data result in a closer relative agreement with the simple valence-bond theory of alkyl radicals than the D(R-H) values derived from the D(R-I) pyrolysis data of Butler and P01anyi.l~ Dr. Alwyn G. Evans (Manchester) said With regard to the results of Dr.Stevenson on the ionization potentials of sec.-propyl and tert.-butyl radicals I should like to emphasize the importance of this work in relation to the interpretation of organic reactions involving carbonium ions. The ionization potentials 7-77 eV for sec.-propyl and 7-19 eV for tert.-butyl were estimated 16 from the earlier work of Stevenson and Hipple l7 using the dissociation energies given by Butler and Polanyi,l* and the marked decrease of ionization potential from methyl to tert.-butyl was shown to be a major factor in determining the changing tendency for S,I reaction from primary to tertiary alkyl halides.16 The ionization potentials of sec.-propyl (7-45 f 0-1 eV) and tart.-butyl (6.90 f 0-1 eV) obtained by Dr.Stevenson from this very detailed investigation are close to the earlier values and show the same difference in energy (0.55 eV). The marked decrease in ionization potential from methyl to tert.-butyl means that the resonance energy of the tert.-butyl ion is greater than the resonance energy of the tert.-butyl radical. This indicates that the energies of the ionic resonance states H (CH,) ,C-C-H and H + I I lie closer together than do the energies H H+ (CHJ ,C=C-H H (9 of these) of the radical resonance states H. (CH3)zkC-H I H (9 of these) This may be due to the fact that in the excited resonance structures of the radical the free hydrogen atom is situated at the normal covalent distance from the carbon so that this structure involves a high repulsion energy between these two atoms while in the excited resonance structures for the ion this large repulsion will not occur because of the small size of the proton.16b If the observed decrease in ionization potential from methyl to tert.-butyl is to be explained in this way one might expect that the change from methyl to allyl or from methyl to benzyl would not involve such a decrease since in allyl and benzyl no unbonded particles are present in the resonance structures for the radical (e.g.CH2=CH-6H and CH,-CH=CH,) or for the ion (e.g. CH,=CH-CH and CH,-CH-CH,). Thus the resonance energy of the radical may be similar t o that of the ion and hence the ionization potential of the radical similar to that of methyl. + + l4 Baughan Evans and Polanyi Trans.Faraday SOC. 1941 37 377. l5 Butler and Polanyi Trans. Faraday SOC. 1943 39 19. 16 Evans A. G. (a) Trans. Faraday SOC. 1946 42 719 ; (b) Reactions of Organic Halides in Solution (Manchester Univ. Press 1946). l7 Hipple and Stevenson J . Amer. Chem. SOC. 1942 64 1590 2766 2769. Butler and Polanyi Trans. Faraday SOC. 1943 39 1 I I 0 GENERAL DISCUSSION Evidence that this may be the case comes from a comparison of the S,I reactions of allyl and benzyl chlorides with that of tert.-butyl chloride. In spite of the fact that the carbon-halogen dissociation energies in ally1 and benzyl halides are less than that in tert.-butyl halide the S,I reactions of allyl and benzyl chlorides are much slower than that of tert.-butyl chloride. We believe this to indicate that the ionization potentials of allyl and benzyl radicals are greater than that of tert.-butyl.19 Dr.D. P. Stevenson (Euneryville California) said The value reported by McDowell and Warren for D(CH,-CN) - D(H-CN) viz. + 0.24 eV, is not only of sign contrary to that expected by chemical intuition but is inconsistent with the heats of formation of CH,CN(g) (AH;, = + 19.4 kcal./mole) HCN(g) (AH,", = + 30.7 kcal./mole) and CH4(g) (AH,", = - 18.2 kcal./mole). Thus CH4 + HCN = CH,CN + H,, and H = ZH hence While CH4 + HCN = CH,CN + ZH CH4 = CH + H CH3CN = CH + CN HCN = H + CN CH4 +HCN = CH,CN + ZH AHzgl = + 19'4 - 30'7 + 18.2 = + 6.9 kcal./mole, = + 0.30 eV/molecule, D = 4-48 eV/molecule AHl = 4-78 eV/molecule. D = 4-42 eV/molecule D(CH3-CN) D(H-CN) AH = - D(CH,-CN) + D(HCN) + 4-42 eV/molecule = 4-78 eV/molecule.Thus D(H-CN)-D(CH,-L") = + 0.36 eV/molecule not - 0.24 eV/molecule as given by McDowell and Warren. It may be noted that if Smith's value for A(CH:) = 15.8 eV/molecule in the methane mass spectrum is taken instead of that of McDowell and Warren 15.3 eVjmolecule then one may compute, D(H-CN) - D(CH,-CN) = +0-28 eV/molecule by the method of McDowell and Warren in good agreement with the thermochemical value for this quantity' and consistent with the quantity D(H-CN)-D(NC-CN) = + 0.26 eV/mo1ecule.20 Dr. C . A. McDowell and Dr. J. W. Warren (Liverpool) (communi-cated) We are entirely convinced by Dr. Stevenson's thermal chemical data and also by his chemical intuition that D(HCN) > D(CH,CN).This is also substantiated by our measurements of the appearance potential of the CH,+ ion from methyl cyanide. We obtain a value of 14.6 eV for Y(CH,+) which with I(CH,) of 10.08 eV leads to a value of 4-52 eV for D(CH,CN). D(HCN) - D(CH CN) = 0.3 eV in agreement with the value given by Stevenson. The contrary deduction given in our paper was based on measurements of the appearance potential of the CH,+ ion from methane and from methyl cyanide and emphasizes how careful one has to be in using appearance potential measurements. Both these measurements seem to be correct as far as we can determine. In each case the CH,+ ion beam is fairly large and the ionization curves are similar in shape and well defined. It may be that our method of determining appearance potentials of ions Taking D(HCN) as 4-81 21 we get Evans A.G. and Hamann Trans. Faraday SOC. 1951 47 28. 2o J . Chem. Physics 1950 18 1347. 21 Stevenson J . Chem. Physics 1950 18 1347 GENERAL DISCUSSION 1 1 1 produced in multiple dissociation processes tends to give low results when there is a large difference in the masses of the fragments. Another possible cause of the difference is that the CH,+ ion from methyl cyanide may be produced with an excess of kinetic energy of about 0.5 eV. This would be sufficient to bring the results on this ion into agreement with the data given above. If such an amount of kinetic energy were liberated in the process leading to the production of CH2+ ions from methyl cyanide, one would have thought that our beam half-width method would have detected the broadening due to this for in this case the ion is only about one-harf of the mass of the neutral fragment and so i t should receive about two-thirds of the excess kinetic energy.It will be recalled that the CH,+ ion beam in methyl cyanide showed evidence of broadening. It should perhaps be pointed out that the discrepancy to which Dr. Stevenson has drawn attention while important in itself in no way requires any of the main conclusions of our paper to be altered. Dr. A. F. Trotman-Dickenson (Manchester) said I should like to make two points concerning the C-H bond strengths in alkanes. First the original papers 22 on the determination of the C-H bond strengths in methane and ethane by photobromination gave D(CH,-H) = 102 kcal.and D(C,H,-H) = 99 kcal. a t 298" K. These figures have been generally accepted and are based on the activation energies of the following reactions (the activation energies are in kcal. /mole). 13.3 Br + C,H C,H + HBr. . - (4 0.8 The determinations of the activation energies of the forward reactions were very precise and it is unlikely that the value of 2 kcal. for the reverse reaction of (I) is greatly' in error. But it is very improbable that the reaction of an ethyl radical with hydrogen bromide has a lower activation energy' than the reaction of a methyl radical. It is contrary to our general information about radical reactivities and the evidence in favour of an activation energy of 0.8 kcal. is not compelling. Accordingly I suggest that the photobromination data should be interpreted as giving D(CH,-H) = 102 f I kcal.at 298" K, and D(CH,-H) - D(C,H,-H) 2 4-5 kcal. These values are in better agreement with those of Stevenson and of Butler and P~lanyi.~, Secondly Dr. Steacie and I have investigated the reactions of methyl radicals with 10 alkanes. In a series of reactions of the type CH + CH -3 CH4 + CH, the bonds broken and formed are so similar that it might be expected that if a relation of the type AE = ccAH exists for this series then cc should be closely equal to 0.5. This presumption may be stated in terms of potential energy curves. It is necessary that the two curves should have the same slope at the point of intersection and that the contribution of resonance energy to the potential energy of the transition complex should not alter in passing from one member of the series to the next.One would expect the intuitive argument to be less applicable as the bonds become more different an extreme case being the reaction of methyl with z-methyl CH3 + C2H6 ,-f CH4 -1- C,H, Z t Kistiakowsky and Van Artsdalen J . Chem. Physics 1944 12 469. Anderson and Van Artsdalen J . Chem. Physics 1944 12 479. 43 Butler and Polanyi Trans. Faraday SOC. 1943 39 19 I I 2 GENERAL DISCUSSION propane. A consideration of the potential energy curves shows that it is probable that a is less than 0.5 when the C-H bond in the substrate is weak. Below is a Table of the results obtained with a = 0.5 ; D(CH 3-H) = 102 kcal. is taken as a fixed point.TABLE I1 All figures are in kcal./mole. Compound Methane . Primary Ethane . 2 2-Dimethylpropane . 2 2 3 3-Tetramethyl butane Secondary n-Butane n-Pentane . n-Hexane . Cyclopentane . Cyclohexane . Tertiary 2-Methylpropane . 2 3-Dimethylbutane (mean) 2 3 4-Trimethylpentane . E ca. 12.8 10.4 &o-4 10.0 fo.3 9'5 f o . 4 8.3 fo.2 8-1 j o - z 8.1 5 0 . 2 8-3 foe2 8.3 i 0 . 2 AE AH Data from Trotman-Dickenson and Steacie J . Chem. Physics 1951 19 329. It appears from the results that all primary and all secondary C-H bonds have the same strength the experimental data on the compounds containing tertiary C-H bonds in which case allowance was made for the contribution of the primary hydrogen atoms to the overall rates are not so satisfactory.Dr. H. Steiner (Manchester) said Szwarc in his paper and also Trotman-Dickenson used the relation between changes in the heat of reaction and in the activation energies derived by Evans and Polanyi 2 * AE = aAH . - (1) This relation holds for series of reactions of the type (diabatic reactions) : RX + Na = Na+X- + R (R = alkyl radical X = halogen atom) when changes are made in R. From theoretical considerations relation (I) can be expected to hold when resonance in the transition state is small. In the metathetical reactions involving homopolar bonds (adiabatic reaction) such as treated by Szwarc and Trotman-Dickenson resonance is a very. important factor in determining the magnitude of the activation energy as was actually mentioned by Szwarc.Nevertheless i t is an empirical fact that relation (I) holds also in a number of series of adiabatic metathetical reactions.25 24 Evans and Polanyi Trans. Faraduy SOC. 1936 32 I933 ; 1938 34 2 2 . 25 Steiner and Watson Trans. Faraday SOC. 1947 2 88. Bolland Quart. Rev 1949 3 1 GENERAL DISCUSSION 1 13 However in my opinion Szwarc’s method of neglecting resonance altogether in discussing reaction series of this type is not justified on the basis of the limited experimental evidence available particularly because the effect of resonance on the absolute value of the activation energies involved is very large. Dr. D. P. Stevenson (Emeryville California) (communicated) In reply to Trotman-Dickenson hitherto unpublished data on the appearance potentials of C,H,+ in the mass spectra of n-butane and n-pentane, and of C,H,+ in the mass spectra of n-pentane and n-hexane lead to D(CH,-H) -D(C,H,-H) equal to 0.19 and 0.25 eV/molecule respec-tively.When these values are taken with the earlier electron impact data 26 0.22 eV/molecule we have for this difference 0-22 f 0.03 eV/mole-cule or 5-1 5 0-7 kcal./mole. in good agreement with Dr. Trotman-Dickenson’s interpretation of the photobromination work of Van Artsdalen. It may be noted that data on the appearance potentials of C,H,+ in the mass spectrum of n-hexane and of C,H,+ in the mass spectra of n-heptane and n-octane when combined with the data referred to in the preceding paragraph result in D(CH,-H) - D(n-C,H,-H) = 0-1 eV/molecule, D(CH,-H) - D(n-C,H,-H) = 0.0 eV/molecule, and D(sec .-C,H,-H) + Iz(sec.-C,H,) - D(sec.-C,H,-H) - Iz(sec.-C,H,) These latter results are not in complete accord with Trotman-Dickenson’s conclusion that primary and secondary carbon-hydrogen dissociation energies are invariant for propane and higher alkanes.However these latter electron impact results are not as unambiguous as those on methyl, ethyl sec.-propyl and tert.-butyl. Dr. M. Magat (Paris) said The problem of the fragmentation of poly-atomic molecules by electron impact can be approached from two different sides. Either one considers one by one relatively simple molecules and tries to interpret the effects of say isotope substitution of the mass spectrographic pattern in the light of the Franck-Condon principle 27 gradually increasing the complexity of the problem.Or one considers large series of related molecules and attempts to establish empirical rules which may give some lead to the factors determining the frag-mentation of more complicated molecules where a detailed theoretical analysis seems for the moment hopeless. This is the approach we have chosen with Dr. Viallard in a paper now in press 28 and from which I would like to quote a few results. Our analysis is almost exclusively based on the compilation of mass spectra of hydrocarbon published by the Bureau of Standards and obtained with a Consolidated mass spectrograph. The discrimination of this apparatus seems to be small except possibly for hydrogen and hydrogen molecule ions as can be seen from the com-parison of the amount of ions obtained under identical conditions from rare gases 29 with the values of total ionization of Smith.30 We have hence made no corrections for discrimination.Here are our main findings some of which are not new. The amount of parent ions in straight chain paraffins decreases with increasing chain length. The amount of parent ions decreases somewhat with any branching and is practically reduced to zero if 2 methyl groups = 0-3 eV/molecule. 26 J . Chem. Physics 1942 10 292. 27 E.g. Schaeffer and Hastings J . Chem. Physics 1950 18 1048. Schaeffer, Turkevich et al. J . Amer. Chem. Soc. 1948 J . Chem. Physics 1950 18 1501. 70 2638. 28 J . Chim. Phys. 29 Diebeler Mohler and Reese J . Res. Nat. Bur. Stand. 1947 38 617. 30 Smith Physic.Rev. 1930 36 1293. See also Compt. rend. 1949 228 1118 114 GENERAL DISCUSSION are substituted on the same carbon atom. The introduction of a double bond increases the amount of parent ions as compared to corresponding paraffins for chains containing less than six carbon atoms 31 but increases for longer chains. The particularly large amount of parent ions in arom-atics was already noticed and interpreted by Rurton.S2 The rupture of carbon-hydrogen bonds occur in the way to favour the formation of radical ions with odd numbers of hydrogen atoms ; this has already been mentioned by Delfosse and Bleaknay 33 for propane but is a general rule valid for all fragments from any paraffin and with any number of carbon atoms. It is also true for perfluorinated paraffins.It map be related to the fact that these radical ions are in fact saturated molecules no unpaired electrons being left. The breaking of carbon-carbon bonds shows some very marked regularities : (a) In normal paraffins and olefins never more than I carbon-carbon bond is broken at once-the ratios CH,+/CH,+ and C,H,+/C,H,+ decrease first when the chain length increases and tend then to a constant value. (b) In branched paraffins " double scission " occurs in the neighbour-hood of the branching fragments recombining usually to give larger ions. For example neopeqtane yields 17 yo of C,H,+ among the fragments ; neononane (3 3-diethylpentane) yields about 40 yo of C,H,+. This recombination has to occur in the activated states before the fragmentation and sets a fairly high limit to the life-time of the activated ions.It seems to us reasonable to assume that several vibrations of these ions take place before the fragments separate. The energy can hence be redistributed and i t may be necessary to investigate carefully how far the Franck-Condon principle is sufficient to interpret the phenomenon. With neononane one observes also a simultaneous breaking of three C-C bonds the amount of C,H,+ ions being abnormally high. (G) In normal paraffins comprising m carbon atoms one can define for each bond n a fragmentation index which gives the ratio of the per-centage of scission occurring a t a given bord and giving rise to C$ and CL- ions to the a priori percentage one would expect if the scissions were at random. It seems to be a general rule that scission occurs preferentially between the 3rd and 4th carbon atoms.Scissions at the ends which would give CH sf and C m-lHz+ are extremely improbable. The scission probability decreases also continuously towards the centre of the molecule. (d) With branched paraffins olefins diolefins and aromatics a new phenomenon sets in-the capture of H atoms by the charged fragment. For instance in the butadiene spectrum one finds a fairly large number of CH3+ ions 0.61 yo as compared to 0.88 yo of CH,+ ions. More striking is perhaps the presence of 17.1 yo of C2H,+ besides 26.6 yo of C2H3+ which means that a hydrogen capture occurs with un-saturated hydrocarbons even if an even number hydrogen ions is produced. Sometimes the hydrogen is captured not only from the neigfibouring but also from a further carbon atom.For example among the fragments of 2 3-pentadiene H,C-C-C=CH-CH one finds 1-3 yo of C2H,+ and 4'5 yo C,H,+ ions. In branched paraffins the hydrogen capture (migra-tion) occurs only if i t leads to an ionic fragment with an odd number of hydrogen atoms (e.g. format on of C3H,+ from neopentane). This hydro-gen migration frustrates for the moment all attempts to establish frag-mentation indices for all except symmetrical compounds. With olefins i t induced us previously to make the erroneous statement that terminal or secondary double bonds are never split. As the matter stands now The result of this calculation is given in Fig. I. 31 Roberts and Johnson Anal. Chem. 1948 690. 32 J.P?ysic. Chem 1948 52 567. 33 Physzc. Rev. 1939 56 250 GENERAL DISCUSSION I I 5 we can say that (i) central double bonds in 3-hexene and 4-octene are split preferentially ; (ii) terminal double bonds either split preferentially or what seems less probable strongly favour the splitting of the single bond a t the opposite end of the cha n ; (iii) double bonds in the secondary position either do not change the sciss on probability of the bond as com-pared to the corresponding paraffins (2-hexene) or decrease i t or less prob-ably decrease the sciss on probability of symmetrical single bond. CH3- L H ~ - CHz- CHz-M2-CHz-CHz-CHz-CHz-CH~-CH3 FIG. I. Fragmentation indices of normal paraffins -Electron impact data - - -Pyrolysis data of Hinshelwood et al. (this Discussion) Dr.George Porter (Cambridge) said Although there is still little agreement about the energies of dissociation of the second and third hydrogen atoms from methane the first D(CH,-H) and the last D(C-H) have been widely accepted as IOI and 80 kcal. respectively and Stevenson gives a summary of the convincing evidence for the former. GlockIer now challenges the latter on the grounds that the predissociation in the spectrum of CH from which the value is derived may be a perturbatio I 16 GENERAL DISCUSSION and this is an important revision if correct. The predissociation to which he refers however involves a complete breaking off of rotational structure in several bands and not merely one missing line and is in fact a very good example of this type of predissociation.It would nevertheless be useful to have an independent confirmation of this value and this would appear to have been given by McDowell and Warren by the method of appearance potentials. It is therefore necessary to point out that these authors in using the value of I(CH) which was derived by Douglas and Herzberg via a cycle involving D(CH) have argued in a circle so that, other things being correct they would automatically obtain the same value for D(CH) as is used in the first place which in this case was 80 kcal. This is only important in so far as their claim to have “ derived ” a value for D(CH) is concerned and has no bearing on their other conclusions. I wonder whether they know of any electron impact data for I(CH) because D(CH+) being well known an independent value for D(CH) could then be obtained without further assumptions.Dr. C. A. McDowell and Dr. J. W. Warren (Liverpool) (communicated) : Dr. Porter has raised an obvious point but the position is not quite so simple as he implies. In using Douglas and Herzberg’s value for the ionization potential of the CH molecule we were of course aware that it had been derived by a thermodynamical cyclic process in which I(C), D(CH+) and D(CH) were involved. I(C) is well known and the value used for D(CH+) namely 3-6 eV is in good agreement with that deduced from our values for the appearance potentials of CH+ and C+ i.e. 3-8 eV. Incidentally this good agreement between these two values can be taken as evidence that the ionization process leading to tLe formation of C+ ions from methane must involve the production of a t least two hydrogen atoms ; and further i t shows that if any kinetic energy is produced in the formation of CH+ ions then a similar amount must be involved in the production of C+ ions.It will be recalled that we could find no evidence for kinetic energy‘ in the formation of either of these two ions and this is in agreement with the above statement. Herzberg’s value for D(CH) has been generallv accepted by spectroscopists as being reliable and we in our paper show that it is consistent with our data. We in no way intended that our results implied that D(CH) was being determined in-dependently A different presentation of the data would perhaps have shown this more clearly but we thought i t evident from the nature of the problem that a completely independent value for D(CH) cannot at present be obtained from mass spectrometric data.One great difficulty is to obtain satisfactory source of CH radicals. Most discharge tubes do not give a sufficiently intense beam of radicals and there are also considerable experimental difficulties in getting the radicals from the dis-charge into the ionization box of the mass spectrometer. One obvious method is to use molecular beam but in this case the low intensity of the CH radicals makes the detection of the resultant CH+ ions extremely difficult. A different method of detecting ion beams of low intensity would seem to be necessary and plans have been made to use a photo-multiplier at the collector end of the mass-spectrometer.Prof. E. A. Guggenheim (Reading) said Pitzer has shown us the remarkable agreement between theory and experiment for ethane with respect both to its entropy and to its heat capacity over the temperature range from IOOOK to 300OK. There is one experimental field where further tests of the theory might be possible. I refer to heat capacities a t still higher temperatures. It would be interesting to confirm that the contribution of the internal rotation to the heat capacity passes through a maximum as predicted by theory. I realize the difficulty of such a comparison owing to the anharmonicities of the remaining seventeen normal modes but if I am not mistaken these would be serious in the case of not more than three modes. I should be greatly interested to hear Prof.Pitzer’s opinion GENERAL DISCUSSION Mr. N. W. Luft (Waltham Abbey) (communicated) It seems worth-while to emphasize the value of the electrostatic theory of internal potential barriers put forward by Lasettre and Dean. Since internal molecular rotation is periodic within 2~ the barrier can be represented by the Fourier series, v = ck cos (Re) - E,. . k This may be expanded in terms of inverse powers of the distances Lik between the charge clouds (i) ( k ) whose origins are at distances ri and ~ 1 1 : respectively from the ends of the rotation axis d (cf. Oosterhoff’s diagram). Thus, where E = minimum energy. ck’s are constants for any given system but cih’s may change with internal rotation. Normally ck+l < ck which decreases the number of terms with appreciable contribution considerably.When the symmetry properties of certain rotating groups are also taken into consideration the effective terms in eqn. ( I ) may be less than four. In eqn (2) contributions may be expected round n = 5 i.e. for n = 3 to 6, say due to the spatial extension of the phenomenon. Any theory which demonstrates the above properties and correlates the four major ck’s with the four major c,’s jn a unique way is acceptable. Lasettre and Dean’s theory does this elegantly. They consider three types of charge distributions viz. u-bonds n-bonds and lone p (or hybrid-ized) electron pairs and use experimental dipole moments. Higher moments for which data are lacking are defined in swh a way as to represent simply accurately known barriers.Recent spectroscopic results for Me. Me MeOH MeNH and MeCF are further evidence of the justification of this method. Apparent disagreement between ex-perimental and calculated results for H,O,S4 and N,H 35 can be largely attributed to an incorrect evaluation of spectroscopic data in the case of asymmetric barriers. Calculations for internal rotation of CH groups are particularly simple since with trigonal symmetry lone pairs make no contribution, as correctly stated by Lasettre and Dean though overlooked by Aston. Thus the contributions from bonds (s) opposite to CH and a t angles 46s are where A represents the interactions from more distant bonds. The secondary influences B may well be compensated for by slight deviations from rigidity during vibration.To a first approximation c,(s) N 0 as long as c3(s) > 0. In calculating three-fold internal potential barriers I suggest the use of bond increments V(s) to be added if the bonds include tetrahedral angles and subtracted if they include an angle +8 = 180~. The following semi-empirical values are recommended for V,,) (kcal.) of bonds rotating at standard distances from the CH group C-H = 0.95 C-F = 1.1, C-Cl = 1.0 C-C = 1.05 0-H = 1.0 S-H = 1-35 N-H = 1.1, ~0-C = 2.0 C-0 = 1.2 N-C = 1-75 C-N = 1-25 C=C = 2.6, Single-bond increments correspond to the experimental dipole moments and quadrupole moments P N 0.755~ - 1.0 L = bond length. A simple procedure can also be devised for estimating the contributions A in eqn. (3), c=o = 1.9. 34 GiguCre J .Chem. Physics 1950 18 88 ; Can. J . Res. B 1950 28 485. 35 Scott et al. J . Amer. Chem. SOL 1949 71 2293 I 18 GENERAL DISCUSSION for the slight deviation of angles and from their standard values and for barriers with +8 = 180° i.e. with n = 6 minima. Thus internal potential barriers can be built up which agree closely with best available values obtained from spectroscopic and C measure-ments. This appears to present a simple method of predicting unknown barriers e.g. ethyl amine = 3-75 a-methyl pyrol = 1-5 cc-methyl furane = 1-5 for those who do not wish to be concerned with theoretical details. Dr. M. Magat (Paris) said Berthelot s6 in our laboratory has investig-ated recently the rotational isomerism of the lower aliphatic alcohols by' the Raman spectra method.As ought to be expected isomers due to rotation around the C-C bond were observed for n-propanol and n-butanol. The differences in energies for the two isomers are in good agreement with those quoted by Prof. Aston 0.82 + 0.18 kcal. for propanol and 0.67 & 0-19 kcal. for n-butanol as against 0.78 &- 0.16 kcal. for butane 3' and 0.62 0.2 kcal. for n-pentane.s8 But rotational isomers or tautomers exist also in ethanol where they are due to rotations around the C-0 bond. The energy difference between the two isomers is slightly higher 1-0 f 0.2 kcal. It is remarkable that this energy is so low despite the existence of hydrogen bonds. Prof. E. A. Guggenheim (Reading) said The procedureused by Aston for computing the thermodynamic functions of butadiene by treating this as a mixture of trans and cis tautomers is valid only a t temperatures sufficiently low that states near the top of the barrier make no contribution.When the temperature is sufficiently low for this condition to be satisfied, this procedure is unnecessarily complicated. In my opinion i t is simpler to construct a complete partition function for all states which make an appreciable contribution. This is quite easy because at these low tem-peratures the number of such states is small. Alternatively no appreci-able error would be made by using as partition function the sum of two partition functions for harmonic oscillators. Such a procedure avoids any need to consider the entropy of mixing. Dr. N. Sheppard and Mr. J. K . Brown (Cambridge) (communicated) : In his paper on hindered rotation in hydrocarbon molecules Prof.Aston has made the important point that at least in the case cf 2 3-dimethyl butane the spectroscopic data so far available indicate the possibility of solid solution between rotational isomeric configurations.39 40 This might imply a residual entropy a t the absolute zero a fact of great importance in the theoretical interpretation of the measured value of the entropy. The evidence for solid solution comes from the previous observation that although the liquid-state Raman and infra-red spectra are consistent with the existence of two rotational isomeric forms of this molecule with only a small energy difference,3Qj 40 4l no simplification occurs on transition to the solid state.Such a simplification occurs on crystallizing the straight chain hydrocarbons e.g. w-butane,40$ 41. corresponding to only one configuration being stable in the solid state. We have recently been able to show however that rapid cooling of this branched paraffin such as has been carried out in previous spectroscopic work leads to an amor-phous glassy solid but that on careful warming this glass undergoes a A full account of this work is to be published elsewhere. 36Berthelot Compt. rend. 1950 231 1481 ; J . Chim. Phys. (in press). A misprint in the Compt. rend. papers attributes to Aston et al. ( J . Chem. Physics, 1944 12 336) the idea that the cis form is the stable one. In fact we agree with this author that tvaizs is the stable form throughout. 37 Szasz Sheppard and Rank J .Chem. Physics 1948 16 704. 38 Mizushima and Okazaki J . Amer. Chem. Soc. 1949 71 3411. 30 Szasz and Sheppard J . Chem. Physics 1949 17 93. 40Axford and Rank J . Chem. Physics 1950. 18 51. 42 Rank Sheppard and Szasz J . Chem. Physics 1949 17 83. Sheppard and Szasz J . Chem. Physics 1950 18 145 GENERAL DISCUSSION I I 9 transition to a crystalline mass.43 This glass to crystal transition is accompanied by a very marked simplification in the infra-red spectrum, indicating that only the centro-symmetric isomer is stable in the crystalline lattice. Hence as a result of this new evidence there now seems little possibility that a residual entropy due to the solid solution of different isomers can occur for this molecule under the usual experimental con-ditions of calorimetric work.The specboscopic simplification accompanying crystallization also leaves no doubt that two rotational isomeric species are present in con-siderable concentration in the liquid state corresponding to a low energy diff erence.39 Dr. K. S . Pitzer (Washington D.C.) said Although strictly outside the scope of a meeting on hydrocarbons the problem of internal rotation in methvl alcohol has received much attention and was mentioned by Prof. Aston. Consequently i t seems appropriate to mention that Dr. Weltner has measured the heat capacity of the vapour as a function of pressure and has found a striking non-linear pressure dependence below I atm. near the boiling point. We have interpreted this as arising from the presence at eqcilibrium of a small fraction of a tetramer presumably held together with hydrogen bonds.This is analogous to the situation with hydrogen fluoride. The reason this finding is important for the internal rotation problem is that the entropy of methyl alcohol in the ideal gas state has been obtained on the assumption of normal gas im-perfection behaviour and i t is such values that disagree with the spectro-scopic results. When our virial coefficients are used the values are in substantial agreement. The potential barrier will be obtained much more exactly from spectroscopic data than from ours but will be in the vicinity of 1000 cal./mole. Prof. J. G. Aston (Pennsylvania) (communicated) In connection with Pitzer’s comments on my paper recently E.M.F.determinations of the hydrogen chloride partial pressures above alcohol solutions of methyl-amine (partial pressure known) saturated with methylammonium chloride have yielded a value of the entropy of the reaction, HC1 (g) + CH,NH (g) = CH,NH,CI (s) which at present agrees with experiment only if the spectroscopic entropy of methylamine be used and a zero-point entropy of R In 2 assumed in methylammonium chloride. The former may indicate residual entropy in methylamine at the absolute zero and the latter indicates lack of dis-crimination between the CH and NH ends in the methylammonium ion. The calculation of the spectroscopic entropy of methylamine made use of a line in the infra-red corresponding to a barrier of 1520 cal./mole and yielded a value 0-5 cal./mole deg.higher than the calorimetric value.44 Dr. J. W. Linnett (Oxford) said Stitt 45 studied the vibrations of ethane and hexadeuteroethane and concluded that “ one must introduce interaction terms between the two methyl groups” into the potential energy function to account for the observed vibration frequencies. This, as Stitt pointed out is consistent with the existence of the barrier to free rotation which also requires an interaction between the methyl groups in ethane. On examination of Stitt’s potential energy function i t is found that if one methyl group is distorted symmetrically’ the bond lengths and three-fold axis of symmetry being maintained the configuration of the other methyl group which minimizes the potential energy is almost the same as if the first methyl group were not distorted.So for this particular distortion there appears to be very little interaction between the methyl groups. However i t is not my object to stress any result but to recall that in the measured vibration frequencies of ethane and 4 3 Brown and Sheppard J . Chem. Physics (in press). “Aston and Doty J . Chem. Physics 1940 8 743. 45 Stitt J . Chem. Physics 1939 7 297 I 2 0 GENERAL DISCUSSION hexadeutero-ethane there exist additional data which might be used for studying the interaction between the methyl groups and for testing theories about this interaction. Dr. J. Sheridan (Birmingham) said I should like to hear views on the possible extension of our understanding of restricted rotation by measure-ment of the splittings within each torsional vibration Such measure-ment is now becoming possible from the pure rotation spectra of certain molecules which lie in the microwave region.In the symme'ric top molecules H ,CCF 3,47 H ,CSiF 4* 49 and H ,CSiH 3,50 successive torsional levels are associated with increasing moments of inertia (I,) for the whole molecule and hence with successively lower frequencies for each pure rotational absorption. The lowering of frequency relative to the ground state spectrum appears roughly proportional to the energy of torsional vibration which can be estimated from the relative intensities of the spectra at known temperatures. For H ,CSiF rotational absorptions for the first and second excited torsional levels are found 49 to be doublets, the first close the second considerably wider.This is interpreted as due to the splitting of each of these torsional states into two levels. Since roughly known torsional vibration quanta produce known frequency separations between the doublets and the ground state absorption the frequency separations of the doublets themselves can be used to estimate the extent of the splitting of the levels. The first and second excited torsional levels are about 140 cm.-l and 280 cm.-l respectively above the ground level the first being split into two levels about 1-4 cm.-l apart and the second into two about 12 cm.-l apart. Only doubling of the levels is observed here which is not unexpected since the SiF group contributes 96 yo of the moment of inertia of the molecule about the symmetry axis (Ib) ; the methyl group undergoes torsional vibration virtually against a rigid franie.51 In H,CSiH,,50 where the two groups contribirte nearly equally to I* the expected splitting of torsional levels into triplets (for K > 0) has been reported.Dr. J. J. A. Blekkingh (Rotterdam Holland) (communicated) A clear distinction has been made between isomers and so-called transitional fornzs of each of these isomers.52 IsomeIs are characterized by a certain distribction of the substituted atoms or groups over the available 12 non-ring valencies of cyclohexane. When the distribution is different, then another isomer is obtained. The various possible forms which one and the same isomer can take on when the atoms move to a greater or lesser extent with respect to each other without however any change in the characteristic distribution of the substituted atoms or groups over the 12 valencies taking place are called transitional forms.Hassel 53 distinguishes two kinds of valencies in the immovable forms, viz. 6 erect ( E ) and 6 lying ( K ) . This formclation does not lead to all the possible isomers. Nor is this the case when use is made of a flat six-membered ring as is still done in the most recent literature ; 54 here also two kinds of valencies are recognized viz. 6 above and 6 beneath the plane of the ring This may be why as a rule only 8 different isomers of C,(HX) are distinguished whereas in reality g different isomers are theoretically possible which are completely equivalent and have exactly the same right to the name " isomer ".46 Herzberg Infru-red and Raman Spectra (van Nostrand 1g45) p. 252. 47 Dailey Shulman and Minden Physic. Rev. 1949 75 1319. 48 Minden Mays and Dailey Physic. Rev. 1950 78 347 ; Minden and Dailey, 49 Gordy and Sheridan J . Chem. Physics (to be published). 5O Lide and Coles Physic. Rev. 1950 80 91 I. 51 See e.g. Koehler and Dennison Physic. Rev. 1940 57 1006. 52 Blekkingh Rec. trav. chim. 1949 68 345. 53 Hassel Tids. Kjemi Bergvesen Met. 1943 3 32. 54 Cristol Hause and Meek J . Amer. Chem. SOC. 1951 73 674. Abstr. Amer. Physic. Society Meeting (New York) Feb. 1951 GENERAL DISCUSSION I 2 1 In accordance with the trigonal symmetry of the immovable forms of cyclohexane the IZ available valencies should be divided into four kinds viz.3 straight up ( E ) 3 straight down ( - E ) 3 inclined upwards ( K ) and 3 inclined downwards ( - K ) . This formulation is correct but how-ever does not take into account all the elements of symmetry and there-fore gives rise to mistakes rather easily. To exclude any possible mis-understanding the simplest and most accurate formulation is that in which the 12 valencv directions are indicated by the figures I 2 3 4 5, 6 and I’ z’ 3.’ 4i 5’ 6’ as in the accompanying diagram. The valencies of the immovable form I which are directed straight up are numbered anti-clockwise I 3 5. The valencies of this form which are directed straight down are numbered clockwise 2 4 6. The remaining (lying) valencies bear the same numbers but with accents, as the straight valencies of the carbon atoms opposite in the ring.As is obvious each immovable form can be placed in six positions. The numbering is such that there are suc-cessive transitions of the directions of the valencies from 3 ) J 4 5 6 ‘ 9 Rule A I 1 to 2 3 4 5 6 2 , I 6 5 4 3 4 1 1 3 2 1 6 5 5 I 6 I 2 3 4 J J 5 4 2 3 2 I The same with the numbers with accents. The six possible positions of one and the same form can consequently be deduced from each other with the aid of rule A i t is sufficient if only one form is known. 15 FIG 2. By intramolecular movements involving slight deformation the im-movable form I can be changed via a movable form into the immovable form 11. The “ back ” of the chair becomes the “ foot end ” and vice versa By turning the whole chair i t returns to the same position as the immovable form I.Through this process of changing form I into form I1 by intra-molecular movements and afterwards adjusting the chair again to the original position of form I the following directions are interchanged : the chair is turned round by the intramolecular change. I becoming I ’ and I’ becoming I \ 2 ,> 2‘ , 2‘ ,, 5 7 9 5’ > 5’ Y 3 5 6 , 6’ , 6’ 3 6 4 3 I J 4’ 3‘ > 9 I 4’ 3’ 3 > 9 4 \Rule B. With the aid of rule €3 i t is thus possible to change an immovable form in which the valency I’ occurs into the other immovable form with the valency I. The formulation of an isomer is chosen from the twelve possible formula-tions (six positions of the one immovable form and six of the other) such that the formulation always begins with the valency numbered I followed by the lowest possible numbers I22 GENERAL DISCUSSION Mention may be made of the special case that there is only owe im-movable form of the isomer concerned.Rule A indicates the possible positions and by’ application of rule 8 to them the turned chair positions are obtained. If the last-mentioned positions are identical to the normal ones the two immovable forms are identical. This is the case with the isomers II’ IZ‘ 16’ 11/22‘ 11’33’ 11/66’ 12’3’4 11/22/55’ and 11’33’55’. There are only six different movable transitional forms corresponding to the six positions of the one immovable form and the six positions obtained from them by applying rule B (therefore corresponding to the other immovable form).These movable forms all possess the well-known boat form. The distance between the two opposite carbon atoms with the valencies I (4’) and I’ (4) is slightly smaller than that between the other pairs of opposite carbon atoms. It is however not correct to distinguish two so-called stretched configurations and two boat con-figuration~.~~ There are in fact six boat forms in which the valencies according to rule A differ from each other e.g. 11’ 22’ 33 44’ 55’ and 66’. There are therefore a maximum of eight different transitional forms of each isomer two immovable and six movable and between the movable ones there is of course an infinite number of intermediate positions. Each isomer of a cyclohexane derivative is completely defined in the way de-scribed by indicating the valencies I 2 3 4 5 6 and 1’ 2’ 3’ 4’ 5’ 6’, to which the substituting atoms or groups of atoms are bound.Prof. L. J. Oosterhoff (Amsterdanz) (communicated) In reply to Dr. Blekkingh’s remark all the boat configuration stretched configurations and intermediate configurations have been considered as appears from the symmetry number used. Dr. Julian H. Gibbs (Princeton) said In connection with the paper of Hazebroek and Oosterhoff I should like to mention some results which Dr. Smyth and I obtained recently a t Princeton. We measured the molar polarization of I 4-dioxane over a range of temperatures in the vapour phase in order to obtain some information concerning the population of the molecules in the various possible con-figurations.For dioxane these possible configurations are (I) the sym-metrical cis form of symmetry Czv in which both oxygen atoms project out on the same side of the plane of the four carbons ; (2) the two un-symmetrical cis forms of symmetry C, in which two para methylene groups project out on the same side of the plane of the other two carbon atoms and the two oxygen atoms (these two forms being mirror images of each other) and (3) the trans form of symmetry Czh in which the two oxygen atoms project out on opposite sides of the plane of the four carbon atoms.56 The latter form is the analogue of the chair tautomer of cyclo-hexane and the symmetrical cis and unsymmetrical cis forms together form the analogue of the flexible or boat tautomer of cyclohexane.Our molar polarization values are plotted against reciprocal temper-ature along with some obtained from earlier data of Schwingel and Greene 57 and Kubo 56 on the appended figure. In making this plot it was necessary to lower all the values of Kubo by 6-5 yo in order to bring the mean height of the curve through his points down to that of Schwingel and Greene and our own. An inaccurate cell constant determination is probably responsible for this discrepancy between Kubo’s magnitudes and those of Schwingel and Greene and ours. At any rate we are primarily inter-ested in the shape of the curve rather than in its height ; therefore it is justifiable to make this alteration in his values in order that they may be more easily compared with our own and those of Schwingel and Greene.It can be seen from the plot that the polarizations drop slowly and more or less linearly with increasing temperature (decreasing I/T) at the 55 Oosterhoff Thesis (Leiden 1949) ; Hazebroek and Oosterhoff this Dis-cussion. Dallinga Thesis (Leiden 1951). 56 Kubo Sci. Papers Inst. Phys.-Chem. Res. Tokyo 1936 29 122. 67 Schwingel and Greene J . Amer. Chem. Soc. 1934 56 653 GENERAL DISCUSSION 123 lower temperatures and then rise sharply with temperature at the higher temperatures. It is interesting to note that the highest temper-ature polarization values obtained in both previous investigations in-dicated this rise a t higher temperatures. However in both cases the existence of a minimum and subsequent rise was not mentioned because the indications of this were only barely outside the estimated experi-mental error.This curve affords a means of determining the order of energies of the various forms of dioxane since the total molar polarization may be expressed as a function of temperature in terms of the dipole moments and energies of the various forms by means of the Boltzmann distribution law and since the moments of the various forms may be roughly evaluated Gibbs and Smyth. (# Schwingel and Greene. 0 Kubo. FIG. 3.-Molar polarization against reciprocal temperature for dioxane. by vector addition of the various group moments allowing for a small amount of inductive lowering of the moments of adjacent dipoles. That is if one assumes that the rotational and vibrational factors in the partition functions for the various forms are the same for all forms (they all have the same symmetry number) they may be cancelled and one obtains for the mean square moment in which the lowest electronic energies E of the various forms have all been referred to that of the trans form as a zero point.The meaning of the subscripts is obvious. No term representing a contribution due to the trans form appears in the equation since this form has no moment. This expression is t o be substituted into the Clausius-Mosotti-Debye equation written in the form in which A and B are constants. One then has an equation for the experimental curve. Unfortunately i t is not possible to fit this equation to the experimental curve by adjusting the parameters E, and E,,. This is undoubtedly due to the approximations made in deriving the equations.However i t is possible to show that a dose approach to such a fit can only be obtained if the order of energies of the forms is E > E, > Et. That the trans form is the most stable is indicated by the small values of the polarizations. If E, and E, are not too close to each other in magnitude the equation exhibits two maxima each one being due to one of the polar forms. The interpretation is then that the experimentally observed portion of the curve is that lying between the two maxima. I f the energy of the highly polar symmetrical cis form were lower than that of the moderately polar unsymmetrical cis P = A + BFZl I 24 GENERAL DISCUSSION forms the lower temperature maximum would be due to the symmetrical cis form.That the situation must be the other way around is indicated by the small polarizat'on values and positive slope of the right-hand portion of the experimental curve. That is the small concentrations of symmetrical cis which would be required to give small enough polariza-tions would also give polarizations which were still rising with increasing temperature whereas the low temperature maximum has already been surpassed in the experimental region. Furthermore the fact that the high temperature maximum is apparently larger than the low temperature maximum also indicates that E > E,. The fact emphasized by Hazebroek and Oosterhoff that the boat or flexible form of cyclohexane is in reality a series of interconvertible forms corresponds for dioxane to easy interconvertibility of the sym-metrical cis and unsymmetrical cis forms.The activation energy for this conversion is not zero since there is a difference in energy between the forms themselves and since there is a slight strain associated with the intermediate configurations. However the lack of a sizable activation energy for this conversion is undoubtedly partly responsible for our failure to obtain an exact fit of our theoretical equation to our experimental curve since the starting point in the derivation of this equation is that the problem may be analyzed into contributions of distinct forms with each of which a partition function may be associated. This whole treatment was based on the curve of molar polarization against reciprocal temperature rather than dipole moment against tem-perature in order to avoid the necessity of making a guess at the mag-nitude of the atomic polarization which is involved in the calculation of a dipole moment for each temperature.It is true that in the unsymmetrical cis form a pair of para methylene groups are fairly close to one another so that one might at first sight, expect this to be the form of highest energy. However in the symmetrical cis form the two oxygen atoms are placed in such a way that the repulsive interaction between the approximately tetrahedrally oriented lone-pair electron orbitals should be large thereby giving a high energy to this form. Dr. K. S. Pitzer (Washington D.C.) said The very thorough analysis of the electrostatic contribution to the internal rotation restricting barrier in ethane by Dr.Oosterhoff tends to reinforce my conclusion that more than one phenomenon contributes to the total barrier. In this regard I doubt if Eyring's calculations have proved that distortions of the spa orbitals on the carbon atoms have a negligible effect. I have made rough calculations which tend to indicate a very appreciable effect from this source although the calculations are too crude to have quantitative significance. Prof. A. R. Ubbelohde (Belfast) said In the freezing of cyclohexane, i t has been observed that even after numerous recrystallizations the mother liquor freezes about 0 - 0 2 ~ below the crystals. This could be due to a temporary enrichment of the " boat " isomer in the last part of the liquid to freeze if the rate of the processes boat Z-f chair is comparable with the time of crystallization.The percentage enrichment required to explain the results does not exceed 5 8 about 0.008 yo. Dr. C. N. Davies ( L o ~ d o n School of Hygiene) said ad-Dimethyl glutaric acid (CH,(CH . CH,COOH),) exists in two optically inactive forms the meso or maleinoid (m.p. 128OC) which readily yields an an-hydride with acetyl chloride a t low temperatures59 and shows other evidence of facile cyclization,60 and the racemic or fumaroid form (m.p. 141 O C) which does not yield an anhydride except under extremely vigorous 58 Cf. Thompson and Ubbelohde Truws. Faraday SOC. 1950 46 349. 59 Auwers and Thorpe Annulen 1895 285 31 I. 60 Thorpe and Young J .Chem. Soc. 1903; 83 354 FIG. 4a. FIG. 50. FIG. 46. FIG. 5b. [To face page 12 GENERAL DISCUSSION 125 treatment when i t is converted into the anhydride of the maleinoid acid.61 Of these two acids only the racemic is of course resolvable into its enantiomorphs (m.p. 80" C) ; these optically acids differ from the inactive acid from which they are obtained in that they yield optically active anhydrides almost as readily as the malenoid acid.G2 Inspection of models shows that the meso or maleinoid acid might exist in three forms, each designated (+ -) with reference to the two asymmetric carbon atoms they contain. The first form which is believed to correspond to the known acid (Fig. 4a) is internally compensated but two optically active enantiomorphs are also conceivable.The arrangement shown in Fig. 4a postulates a strong dipole inter-action between the carboxyl groups with resonance which confers a plane of symmetry on the molecule. This cis (+-) acid is thus a true vneso acid and is identified with the maleinoid acid (m.p. 128'). In view of the powerful interaction between the carboxyl groups this configuration should be stable and especially favourable for cyclization. In Fig. 4b two enantiomorphs are shown which can be formed from Fig. 4a by simple rotation about bonds. With the (+-) acids i t is only possible to achieve enantiomorphs in which the comparatively weak dipole attraction between a single methyl and a single carboxyl group is effective. This may be insufficient to prevent bond rotation taking place at ordinary temperatures and thus account for the fact that these acids are unknown.The racemic or fumaroid acid is a ( f +) or ( - -) acid and two pairs of enantiomorphs are theoretically possible. Fig. 5a shows models of the cis forms which arise from carboxyl group attraction as in the meso acid. These should also form anhydrides readily and we regard the optically active acids (m.p. 80") as having this structure. The optically inactive cis acid (&&) is not known but could presumably be obtained by mixing equal amounts of these isomers. Resonance between the carboxyls in this case does not confer a plane of symmetry as in the model of the meso acid. From these models simple bond rotation leads to the molecules indicated in Fig. 9. Each carboxyl group is favourably placed for association with a methyl group so that these forms are stabilized by h - o internal carboxyl-methyl links, instead of one which as we saw above is apparently insufficient to pre-vent bond rotation.These acids would clearly be reluctant to form anhydrides. We regard the fumaroid acid (m.p. 141~) as the externally compensated mixture of these two that is the trans ( f -&) acid. The individual acids are unknown. Dr. Manfred Gordon (Royal Technical College Glasgow) said Ubbelohde and McCoubrey have drawn attention to the influence of molecular coiling on the reactivity of hydrocarbons in the gas phase. Cyclization reactions might well provide some quantitative information in this respect by cor-relation of rate measurements with theoretical calculations based on suitable models of kinked or coiled structures.As an example consider the cyclization of rubber. Though this has been studied in the liquid and not in the gas phase similar effects of coiling may be anticipated in solution. It has recently been demonstrated 63 that the reaction leads exclusively or very predominantly to ring formation between adjacent isoprene units of the same rubber chain. First an ethylenic carbon of one unit is converted to a carbonium ion. When this attacks another double bond to form a new C-C bond i t is thus in practice the double bond of the adjacent unit which is attacked most rapidly i.e. the double 6a Moller Lund's Univ. %%. 1919 16 56 ; Ckem. Abstr. 1920 14 942. 63 Gordon PVOG. Roy.SOC. A 1951 204 569. The following interpretation of these facts is suggested. These are trans (+ +) and trans ( - -) acids. Moller Ber. 1910 43 I 26 GENERAL DISCUSSION bond in the neighbouring loop of the hydrocarbon coil (Fig. 7) reacts faster than more distant double bonds of the same or other chains. Quite analogous reactions can be carried out with low molecular weight terpenes, and i t appears feasible (though this is a guess) that dihydromyrcene could be cyclized with BF in the gas phase. It would be most instructive if the rates of such cyclizations could be measured on different model compounds having the necessary two double bonds at different spacings. The rates r of such cyclizations should be compared with calculations, based on suitable models of the fraction f of equi-energetic configurations available to the chain in which the two carbons to be linked come within bond-ing distance.Clearly f is a measure of the fraction of time geometrically favourable to reaction and should cor-relate with Arrhenius frequency factors. model to another e.g. from a randomly kinked to a more regularly coiled model. FIG. 6. Do Ubbelohde and McCoubrey agree that such rate measurements and cal-culations would in principle promise to give information on molecular configuration complementary to the results of physical measurements ? Dr. M. Magat (Paris) said I would like to make three comments on the very interesting paper presented by Prof. Ubbelohde. I would like to call attention to a different way in which flexibility of hydrocarbon chains influences the reaction rate.I n certain cases the so-called steric hindrance in bimolecular reactions can be traced back to this property of hydrocarbons. Besides the cases analyzed by Ingold, Hughes and their co-workers and by A. G. Evans where the steric hindrance increases the activation energy there are cases where the reaction rate decreases along a homologous series the activation energy remaining constant. Such cases have been recently compiled by Miss Ivanoff and myself 64 and we found that they all concern “ chain ” hydro-carbons and nucleophilic reactions. If one now considers the correspond-ing models i t appears that the number of possible chain configurations is reduced in the transition state as compared to the initial state.For instance i t is clear from the model that in the initial state acetone can have (disregarding the hydrogen positions) just one configuration while diethyl ketone has nine configurations (3 different positions for each -CH,. group). In the transition state for the reaction with semi-carbazine acetone and diethylketone can both have just one confi9.u ation (the “ pure trans ” one in the last case). Hence everything else being unchanged the entropy of activation is larger for the diethyl ketone than for acetone and one finds immediately that the rate ought to be g times slower for the diethylketone than for acetone. The experimental frequency factors are according to Price and Hammett 65 2-55 and 0.25 in good agreement with the prediction. It can be shown 6 6 that the changes in the translational and rotational terms of the activation entropy could account for a rate decrease of only 30-40 :/o.Prof. Ubbelohde assumes that relatively low normal paraffins are ‘‘ bunched ” or “ coiled ” in the liquid state. How is this to be reconciled with the findings of Kratky on the di-iodo-1 11-undecane in undecane solution for which X-ray diffraction favours almost entirely a stretched configuration ? Prof. Ubbelohde suggests that the probability of energy exchange from kinetic into vibrational on collision is favoured by the existence of low-frequency vibrations in non-rigid molecules. Although this idea 64 Ivanoff and Magat J . Chim. Physic. 1950 47 914. 65 Price and Hammett J . Amer. Chem. SOC. 1941 62 2387.66 Bauer and Magat J . Claim. Phys. 1950 47 922. r ; 0 Moreover f would vary widely from one ?f->Df C + c - GENERAL DISCUSSION 127 appeals to me i t seems to be in contradiction with the formula for the probability of such an exchange derived by Jackson and Mott G 7 in which this probability is proportional to the frequency Y. Dr. K. S . Pitzer (Wmhington D.C.) said With respect to Prof. Ubbelohde’s paper i t seems to me that the indication is that the viscosity data are not simply interpretable in terms of the coiling of n-paraffins. The fully coiled configuration is 500-1000 cal./mole higher in energy for each C-C bond over the extended form as has been well established now by both statistical thermodynamic and spectroscopic methods. Thus a n-paraffin must become more coiled with increasing temperature whereas the viscosity ratios in Table I1 of Prof.Ubbelohde’s paper show no sig-nificant change with temperature. Prof. E. G . Cox (Leeds) said It should be pointed out that a comparison of the “ straight zig-zag ” and “ helix ” as representations of extremes of possible hydrocarbon chain configurations may be misleading unless i t is realized that there must be many zig-zags in one turn of the helix ; thus the helix is not a very realistic model to use for the discussion of the small molecules (up to octane) chiefly discussed in this paper. I should like to ask Prof. Ubbelohde whether he has been able to make any critical assessment of the accuracy of the various methods of deter-mining apparent molecular diameters ; can viscosity measurements for example give results which are in some way better or more accurate than those from diffusion methods ? Mr.R. S . Bradley (Lee&) (communicated) The point raised by Prof. E. G. Cox is mentioned in my paper with Dr. A. D. Shellard G8 where i t is stated “the hypothesis of coiling has previously been advanced by Mack and his co-workers to explain the comparatively low values of the collision radii of the lower hydrocarbons although up to C,H, the ‘ helix ’ on our model would occupy scarcely one t u r n ”. In view of the small amount of coiling of the lower hydrocarbons i t is somewhat doubtful whether any shielding of end groups occurs with hydrocarbons up to octane. In any case the term “ coiled molecule ” is misleading and might with advantage be replaced by “ crumpled mole-cule ” and the end groups of a long hydrocarbon are not necessarily near to the surface of the molecular “ drop ”.The work of Prof. Ubbelohde supports our viems on the normal paraffins Cl6-CI8 for which we found from diffusion measurements in air collision areas in agreement with the values based on statistical “ coiling ” and also the more recent work of the writer with Waghorn 6 9 in which we find nearly the same value for the collision radius for (n-C,H,,) ,CH (447 %i at 25OC) as for n-C18H38 (4-53 Prof. A. R. Ubbelohde and Mr. J. C. McCoubrey (Belfast) (corn-municuted) With regard to the condition of n-paraffins in the liquid state present information suggests that it is important to distinguish between the degree of crumpling near the f.p.and the greater crumpling near the b.p. In general liquids behave as quasi-crystals near the f.p. and quasi-gases near the critical temperature. It is anticipated that at the b.p. which is approximately equal to two-thirds the critical tem-perature liquids should be quasi-gaseous. Our data on molar volumes and entropies of vaporization suggest moderate crumpling at the b.p. Near the f.p. considerably greater adlineation is suggested by existing evidence.”J The presence of dipoles could further favour this adlineation so that there is no present inconsistency between X-ray work on di-iodo-I II-undecane and our molar volumes. The exchange of energy referred to by Jackson and Mott deals with a at 25O C). 137 Jackson and Mott Proc. Boy. SOC. A 1932 137 703. 68 Proc. Roy. SOC. A 1949 198 239. 69 Proc. Roy. Soc. A 1951 206 65. 70 Cf. Quart. Rev. (in press) I28 GENERAL DISCUSSION much more restricted process than the various modes of transfer from kinetic to vibrational energy in flexible hydrocarbons. In reply to Prof. Cox the present consistency of collision diameters as obtained by’ various authors has been carefully reviewed.71 In the most favourable cases the agreement is within I yo of the measured viscosities. We do not necessarily expect collision diameters in homomolecular assemblies to agree with collision diameters in hetero-molecular assemblies. We agree with Mr. Bradley that “ crumpled ” is a better term than “ coiling ” for molecules of such chain length that they can still be readily vaporized. Dr. Peter Gray (Cambridge) (communicated) Ubbelohde and Mc-Coubrey state that in coiled flexible hydrocarbon molecules intermediate CH groups may have one and frequently both C-H bonds blocked owing to the coiling. But if a coil is formed will not these C-H bonds tend to be on the “outside ” of the coil and no less accessible than in an extended molecule ? Dr. P. Torkington (Brit. Rayon Res. Assoc.) (communicated) Prof. Ubbelohde implies in his paper that vibrational coupling is advantageous to a concentration of energy at a particular part of the molecule. One would have thought that such a concentration of energy would be analogous to a decrease in entropy and that the coupling of vibrations would deter, not aid any drift towards an activated state. Prof. Ubbelohde (Belfast) (communicated) The coupling of vibrations in a molecule permits the transfer of vibrational energy to any particular bond where reaction is to occur. If there is a break in the coupling, the vibrational energy in separate parts of the molecule is not generally available for activation beyond the break. Dr. A. F. Trotman-Dickenson (Munchester) said Ubbelohde and McCoubrey have cited the supposed decrease in the steric factors of the reactions of methyl radicals with alkanes containing primary secondary and tertiary hydrogen atoms as evidence in favour of the theories out-lined in their paper. The evidence for the decrease comes from an inter-pretation 7 2 of early experimental work.73 However the figures cannot be relied upon because i t has been shown 74 that the experiments were in-adequate and that the interpretation is unsuitable. Dr. Steacie and I 75 have obtained values for the steric factors of these reactions and the parallel ones with the alkenes where the cc-methylenic hydrogen is attacked. These values are given in the annexed table. A statistical correction for TABLE III.-THE STERIC FACTORS FOR THE REACTIONS OF METHYL RADICALS WITH ALKANES AND ALKENES Alkanes 1 Alkenes StericFactors x xd Type of H atom Primary . Secondary . Tertiary . 2 2 3 3 30 I 2 These figures do not appear to support the theories of coiling. the number of H atoms involved has been made ; 1-5 and 3-5 A were taken as the collision diameters of the hydrogen atoms in the hydrocarbons and the methyl radicals respectively. 71 J . Chem. SOC. (in press). 73 Steacie Darwent and Trost Faraday SOC. Discussions 1947 2 86. 73 Smith and Taylor J. Chew. Physics 1939 7 390. 74 Trotman-Dickenson and Steacie J. Physic. Chem. (in press). 76 Trotman-Dickenson and Steacie J . Chem. Physics 1951 19 329

ISSN:0366-9033    

DOI:10.1039/DF9511000103

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

II. HYDROCARBON REACTIONS A. THERMAL REACTIONS THE THERMAL DECOMPOSITION OF HYDROCARBONS BY F. J. STUBBS AND C. N. HINSHELWOOD Received I 5th February, 195 I The following topics are discussed in the course of a general survey of the decomposition reactions of the normal paraffins : the nature of the radical- chain reactions, the probability that the residual reaction after nitric oxide inhibition of chains is a molecular reaction, the rate-pressure law which indicates that the mechanism is kinetically composite, the variation of activation energy with pressure, the relative probability of rupture of the carbon chain a t various points, and the interpretation of experiments on the composition of products. Further questions about energy-entropy relations in unimolecular reactions arise in connection with these various matters.The thermal decomposition of a paraffin follows the overall course : Higher paraffin = lower paraffin + olefine. The detailed occurrences can be formulated equally well as molecular reactions involving a relatively simple transfer of hydrogen accompanied by the splitting of a carbon-carbon bond, and, alternatively, as chain reactions initiated by the splitting of the paraffin into two radicals which then suffer further changes. The olefine is sometimes relatively stable, and sometimes suffers a rapid subsequent decomposition into a pair of molecules of lower olefines, or other more complex reactions. The general conclusion reached from a considerable amount of ex- perimental work is that the chain processes and the purely molecular processes occur simultaneously, and in the following section the evidence for this will be reviewed.The kinetic relations observed in both these types of process possess considerable interest in connection with the theory of reaction velocity, and in fact open up some new aspects of the theory of unimolecular re- actions generally. This bei-ng so, the first important matter is to examine with some care the status of the supposed molecular reaction, The Suppression of Radical Chains and the Question of the Nature of the Residual Reaction.-Nitric oxide, which suppresses radical chain reactions, reduces the decomposition rate of the paraffins to a constant limiting value. The residual reaction might be (a) the primary process of what in the absence of inhibitors would be a chain reaction, (b) the usual chain reaction imperfectly suppressed by nitric oxide, (c) a surface re- action, (d) a second type of chain reaction not influenced by nitric oxide, (e) a molecular reaction.(a) is supported by Steacie and Folkins,l who found that the products formed from butane have the same composition whether the rate is con- trolled by nitric oxide or not. They consider that a chain reaction and a competing molecular reaction would not yield the same products in the same proportions. The inhibited reactions usually account for a fifth to a third of the total normal change and if they are all primary processes this implies very short chains. Now while initially formed radicals can in fact easily go through cycles of reaction to give the same Steacie and Folkins, Can.J . Res. B , 1940, 18, I. E 129DECOhlPOSITION O F HYDROCARBONS products as the molecular processes would give, they could hardly do this if they all reacted with nitric oxide. On the whole, therefore, the evidence from the work of Steacie and Folkins might be taken to support the opposite view at least equally well. (b) has been put forward by Echols and Pease a on the basis of ob- servations with n-butane with and without nitric oxide. The Ap-time curves tend to become parallel as the reaction proceeds, giving the im- pression that the inhibition ceases (although hardly any nitric oxide is consumed). Echols and Pease assume an equilibrium in which as many radicals are regenerated from a combination with nitric oxide as are removed. This explanation is unsatisfactory because, firstly, propylene reduces the rate to the same limiting value as nitric oxide, and the forms of the A$-time curve of the inhibited reactions are the same ; and secondly, direct addition of extra nitric oxide during the reaction has no further effect.3 The apparent acceleration in the Ap-time curve is in fact an in- herent property of the reaction, and due to stoichiometric causes.The sigmoid nature of the curve for the fully inhibited reaction becomes more evident as the paraffin series is ascended. The higher paraffins on decom- position yield higher olefines which themselves subsequently decompose with further increase of pressure, whence the apparent acceleration. The Ap-time curves for the reactions with and without nitric oxide tend to become parallel not through the ceasing of nitric oxide inhibition but rather (as shown by direct experiment) because in the normal decom- position the olefine formed assumes the role of inhibitor.(c) has to be considered carefully as the evidence may easily be mis- leading. Surface has little or no effect on the uninhibited decomposi- tion, so that the chains are broken in the gas phase. The nitric oxide- inhibited reaction, however, in certain cases shows a slight apparent retardation by increased surface, and this becomes quite marked in a vessel whose internal surface is coated with potassium chloride. These effects prove, however, to be secondary.* In the early stages nitric oxide may catalyze a quite independent condensation reaction, and the re- sultant pressure decrease vitiates the measurement of the true decom- position rate.In uncoated vessels this effect is transitory, but in potassium chloride-coated vessels it is more serious. If, however, propyl- ene is used as inhibitor, the limiting rate is measurable without this dis- turbance, and is exactly the same as that observed in a vessel with a clean silica surface. The conclusion is that the residual reaction is essentially homogeneous. (d) A possibility for a second kind of chain reaction is one involving hydrogen atoms which might conceivably be immune to the action of nitric oxide. In presence of added hydrogen there would be replacement of alkyl radicals by hydrogen atoms generated in the reaction and since this would occur in competition with their removal by nitric oxide, more of the latter would be required to produce a given degree of inhibition.Now addition of hydrogen does in fact cause a marked increase in the rate of the uninhibited decomposition, so that the chains are evidently lengthened when hydrogen atoms replace methyl radicals according to the reaction shown above. But the limiting value to which the rate is lowered by sufficient nitric oxide is very little changed and moreover no greater amount of nitric oxide is now required for a pro- portional reduction of rate.4 It is highly unlikely, therefore, that the residual reaction depends simply upon hydrogen chains unsuppressible by nitric oxide. CH,. + HZ -+CHI + He, a Echols and Pease, J .Amer. Chem. SOL, 1939, 61, 1024. Stubbs and Hinshelwood, Proc. Roy. Soc. A , 1950,200, 458. Stubbs and Hinshelwood, Proc. Roy. Soc. A , 1950,201, 18 ; Ingold, Stubbs and Hinshelwood, Proc. Roy. SOC. A , 1950, 203, 486.F. J. STUBBS AND C. N. HINSHELWOOD ( e ) In the light of the foregoing facts and until evidence for somc alternative mechanism appaers, there is a prima facie case for the con- clusion that the reaction observable when the rate has been reduced to its limiting value by propylene or nitric oxide is that of a molecular de- composition of the hydrocarbon. It is of interest, therefore, to consider what properties this molecular reaction, if indeed it is such, possesses. This question will be discussed in a later section. General Characteristics of the Chain Reactions.-Firstly, however, we should deal with the chain reaction itself.For the thermal decomposition of a straight chain paraffin the types of process involved must be as follows : (I) Initiation steps : e.g. RH = R, + R,. ( 2 ) Reaction of radicals with hydrocarbons : e.g. ( 3 ) Further breakdown of radicals (4) Recombination of radicals, e.g. : 2R, = R,Rl. ( 5 ) Reaction of radicals with inhibitors, e.g. : (I) must clearly result in the production of alkyl radicals or hydrogen atoms. ( 3 ) gives lower alkyl radicals or hydrogen atoms. From the results referred to in the preceding section we may infer that no important kinetic differences in the overall character of the reaction arise if alkyl radicals are replaced by hydrogen atoms in ( 2 ) . ( 3 ) is necessary to account €or olefine production in the chain process. An important kinetic difference arises according as to whether an in- hibitor such as NO reacts predominantly with a radical which would otherwise have suffered a decomposition of type (3), or with one which would normally have reacted with more paraffin as in ( 2 ) .A decision as to which case prevails can be reached by studying the rate as a function of nitric oxide pres~ure.~ If Y , is the rate of the uninhibited reaction, Y that at a given nitric oxide concentration and vrn that at high nitric oxide concentrations, then where y is a function independent of [RH] when all the radicals affected by nitric oxide suffer alternative processes of type (3). This case is in fact found in the thermal decomposition of ether.6 Actually with the hydrocarbons y varies inversely as a power of [RH] which in various examples has been found to lie between 0.71 and 1.0.From this may be concluded that direct attack on the hydrocarbon occurs in competition with removal by the inhibitor. This suggests that any processes of type (3), that is, decompositions of longer alkyl radicals into olefine and shorter alkyls are usually rapid. The overall reaction (for the chain part of the decomposition) is usually of an order in the hydrocarbon which is between I and 2 . The rate of process (I) can vary as a power of [RH] between I and 2, that of z will be of the first order in [RH] itself, while (4) could either be inde- pendent of [RH] or involve it as a first power if ternary collisions played a part in the radical recombination.No particular problem arises, therefore, in accounting for the fractional order of the total reaction with respect to [RH]. The apparent chain length (which really indicates the relative pro- portion of chain reaction to molecular reaction) decreases as the initial pressure becomes greater, as the temperature rises, and as the number of carbon atoms in a normal paraffin increases. 5 Hobbs and Hinshelwood, Proc. Roy. Soc. A , 1938, 167, 439 ; Staveley, Proc. Roy. SOC. A , 1937. 162, 557. BHobbs, Proc. Roy. Soc. A , 1938, 167, 456. Rl + RH = R,H + R. R,CH,CH,- = R, + CH, : CH,. e.g. R, + NO = KINO. (Y - Yrn)/(Yo - Yrn) = {(r"O1)Z + I)+ - r"O1,1 32 DECOMPOSITION OF HYDROCARBONS The Rate-Pressure Law for the Molecular Reaction.-The rate of reaction has sometimes been represented as proportional to the 3 / 2 power of the pressure, but this is an oversimplification of the facts.In general, it varies as a power of the pressure between the first and the second, and with the higher paraffins (for the residual molecular reaction) is expressible over a considerable range by the formula : With ethane it conforms to the simple expression for a unimolecular reaction with collisional activation, showing a transition from second order to first as the pressure rises. With propane there is conformity with the expression rate = A p , + Bpo2. and there is, therefore, good reason to suppose that the general form should be the reaction mechanism being composite.4 On any other view it would be difficult to explain how in general over a considerable range of pressure a second order process gains predominance over a first order process, and then may subsequently return towards the first order.Two matters are relevant to the assessment of this result. In the first place, there is the question how far the pressure-time measurements constitute a valid criterion of reaction rate. They appear in fact to be quite satisfactory for comparative measurements on a given paraffin over a range of temperatures and pressures, since the curves of Ap against time for widely different conditions can be brought into complete super- position by simple changes of scale. The second question is how far the different kinetic mechanisms are associated with the formation of different sets of chemical products.This, remarkably enough, appears not to be so, as will be shown in another section. Variation of Activation Energy with Pressure.-One of the most interesting characteristics of the normal paraffin decompositions is that from n-butane upward the activation energy of the molecular reaction becomes a well-defined function of the pressure. The pressure variation, moreover, is closely correlated with the kinetic behaviour. With ethane, where E is constant, the reaction is of the first order except at lower pressures (where the normal transition to the second order occurs). In other words, there is no evidence of a kinetically composite character, and the results can be interpreted in terms of a single unimolecular reaction.A similar result has more recently been found by Mr. M. G. Peard with isobutane which satisfies the conditions of a single unimolec- ular reaction, and likewise shows no variation of E with pressure. On the other hand, with the higher normal paraffins there is a steep fall in the value of E as the pressure is increased, and this corresponds to an increasing predominance of the p 2 term in the expression for the reaction rate (see previous section). By extrapolation to low and high pressures respectively, the following values are found : %-Butane . n-Pentane . n-Heptane . E (kcal.) Lowest Pressures 69 93 88 High Pressures 9 63 59F. J. STUBBS AND C. N. HINSHELWOOD I33 A given paraffin chain can of course break in various ways yielding, for example, methane or ethane as the saturated product.These alter- natives do not, however, correspond to reactions with a different pressure or temperature dependence as shown by the analytical results to be referred to later. What we must probably conclude is that there are kinetically different modes of activation, corresponding to different relations between total energy in the molecule, energy required in a critical location, and the decomposition probability of the energized molecule. Relation to the Theory of Unimolecular Reactions.-The prima facie case statable on the basis of the experimental results and open to theoretical analysis is as follows. At low pressures large total energy accumulations in the molecule have time for redistribution, and provoke decomposition when an appropriate part is concentrated in the right bond, as in the classical theory of unimolecular reactions.' There is no difficulty with molecules of the size of butane and the higher paraffins in accounting for an energy accumulation rapid enough to provide the necessary activation rate.At higher pressures, however, an increasing part is played by a process in which a smaller amount of energy enters the molecule in a more specificially distributed manner, possibly even causing an immediate activation of one of the critical bonds. Unless decomposition super- venes rapidly this energy will flow into the rest of the molecule. This mode of activation is attended, therefore, with the necessity for rapid transformation of the molecule, and it remains of the second order up to considerably higher pressures.It involves a lower activation entropy as well as a lower activation energy, but, on balance, can compete with the other mode of reaction. Being of the second order over a large part of the experimental range, it gains on its competitor as the pressure rises. The picture thus disclosed widens in one rather interesting respect our usual conception of the energy-entropy relations in decomposition reactions. In so far as the total energy in a molecule may be thought of as the sum of amounts associated with individual bonds, the probability of a given distribution, with El in the first bond, E, in the second and so on, is proportional to e - WRT . e - EatRT . . . i.e,, to e - ~ Q I R T . A given total E can be made up in very many different ways, so that the probability of its occurrence is the sum of a great many terms of the above type which will come to Fe - E I R T , where F is large.The chance that this energy shall be localized (in the simplest model in a single bond) so as to give the transition state of the reaction, is a small fraction only of this, which, for the case of a localization of E , out of the E may be written f ( E , E,) Fe - EIRT = Ae - E P T . The temperature coefficient of the reaction is determined by E rather than E,, so that the activation energy is high. f ( E , E,) will, in general, be small but F is very large, so that A , which determines the entropy of activation, can be high. Roughly speaking, the reason why the entropy factor can be large while the activation energy is high is that although E is a large quantity, the amount of energy per degree of freedom to which it corresponds is not very many times the average when the molecule is fairly complex.Now although the probability of correct localization of this energy is small in any interval of time, dt, the whole long period between molecular collisions is available, and the reaction occurs 7 Rice and Ramsperger, J . Amer. Chem. SOC., 1927, 49, 1617; Kassel, J . Physic. Chem., 1928, 32, 225 ; Eley, Trans. Faraday Soc., 1943, 39, 168 : Evans and Rushbrooke, Trans. Faraday Soc., 1945, 41, 621 ; Barrer, Travs. Faradny sot., 1948, 44, 399.I 3 4 DECOMPOSITION OF HYDROCARBONS irrevocably should the correct distribution be reached for an infinitesimal fraction of this period.Let us now envisage the actual entry of energy into a molecule in a collision. This will in the very first instant be a quite localized affair, and the energy communicated may be thought of as entering one bond first and then being dissipated throughout the molecule in the course of a very brief relaxation time. If E,1 is communicated, E , may remain in the bond long enough for decomposition to occur. It is quite reason- able to suppose that EO1 - E , is very much smaller than E - E,. The activation energy will now be much lower than before, but the entropy of activation will be small also, since it contains no large factor F but corresponds to the requirement of a more or less precisely defined collision. Of course even with Eol in the whole molecule, E , would eventually collect into the required bond without a collision, but the time required for this would be longer than the interval during which the molecule is normally left undisturbed, and so the contribution to observable reaction would be negligible.Position of Rupture of the Carbon Chain.-When a straight chain paraffin decomposes it gives a lower paraffin and an olefine, the saturated fragment being always the shorter of the two. The probability of rupture is nearly always greatest at the link between the first and second carbon atoms, methane being the paraffin formed in highest proportion. The probability of rupture at the second link, that is, between the second and third carbon atoms, is in general rather lower, though the proportion of ethane is considerable.The formation of propane and higher paraffins is very much smaller. The rapid fall in the probability of rupture with distance from the end of the chain is shown in Table I. TABLE I.-RELATIVE PROBABILITIES OF BOND RUPTURES n-Butane . n-Pentane . n-Hexane . $2-Heptane . n-Octane . 1'0 1'0 1'0 1'0 1'0 1'0 0.78 0.5 0.64 1'0 L 0.25 0.35 0'10 In view of the fact that the activation energy for the higher paraffins varies with pressure, and that the kinetic relations indicate a composite mechanism, it is tempting at first sight to associate the various modes of rupture with processes of different activation energy. This supposition, however, would be incorrect. -4 detailed analytical investigation which Mr. K. U. Ingold 8 has recently made in this laboratory has shown that the relative proportions of the saturated hydrocarbons formed in the decomposition of n-hexane, n-pentane and n-heptane are independent of temperature and pressure over a range where marked shifts would certainly occur were their production conditioned by activation energy differences comparable with those actually observed.It appears, therefore, that the methanelethane ratio which measures the relative probability of rupture at the two different places, is deter- mined by processes of nearly equal activation energy. Inspection of the normal modes of vibration of a long carbon chain reveals no consistently applicable reason why the CIp2 and CZp3 ruptures should be so greatly favoured. On the other hand, thc relatively small differences between the methane and ethane fractions compared with the large drop in the frac- tions of any higher paraffins is suggestive of an electronic effect where 8 Ingold, Stubbs and Hinshelwood (in course of publication),F.J, STUBBS AND C. N. HINSHELWOOD I35 methyl and ethyl have a much greater facility than any higher alkyl groups for capturing an extra hydrogen atom. The delocalization of electrons in the bonds of methane is believed to confer extra stability upon it.O With ethane there is also a considerable measure of sym- metry which will favour delocalization, but with propane and any higher paraffin the carbon-carbon valency angles are such as to lower markedly the symmetry of the system of hydrogen atoms regarded as a whole. That the symmetry of the complete set of six hydrogens in ethane confers stability in a way analogous to the delocalization in methane is a possi- bility which might be susceptible of an interesting theorctical treatment.Comparison of Paraffins and 0lefins.-In comparable ranges of tem- perature a major difference between paraffins and olefins is that the reactions of the latter show no evidence of chain processes according to the usual tests. In view of the fact that olefines are themselves effective inhibitors of radical chain reactions this is not surprising. FIG. I. The general course of the decomposition is a little complex, and for the example of propylene in the range 570-650° has been examined in this laboratory.1° Primary fissioii of the propylene is followed by co- polymerization of activated intermediates with more propylene, and this reaction in turn by further breakdown and further polymerization. A generally similar course probably applies to the reactions of higher olefins, recently studied by Miss M.J. Mo1era.l' The rate under these circum- stances is mainly determined by the primary fission (activation energy 57 kcal. for propylene, 53 kcal. for hexene-I, considerably less than that found by Szwarc l2 for a more profound fission into radicals which inter- venes about 150' higher). Remarkably enough the reactions show none of the signs of composite kinetics, and are uniformly of the first order (with the usual tendency towards the second order at low pressures). The activation energies are generally lower than those for the corresponding paraffins. In Fig. I rates for paraffins and olefines are compared directly (the two branches of the curve for the latter corresponding to extreme 9 Coulson, Quart. Eev., 1947, t, 144. lo Ingold and Stubbs, J. Chem. SOC. (in press). l1 Molera (unpublished observations). l2 Szwarc, J . Chem. I'hyszcs, 1949, 17, 284.'I 36 REACTION OF METHYL RADICALS assumptions about the relation of pressure change and actual amount of primary decomposition). The higher olefines decompose more rapidly than the paraffins with an equal number of carbon atoms. This fact, taken in conjunction with the simpler order of the primary fission, suggests that the initial seat of reaction with the higher olefines is adjacent to the double bond in each case. It is perhaps tempting to attribute this to the circumstance that there is a greater relative gain in symmetry of structure accompanying the primary fission in changes such as 9 CHECH + CH3CHzCH3 CH, = CH . CH,CH,CH,< CH,=CH, + CH,=CHCH, than there is in CH,CH,CH2CHzCH3 -+ CH, -t CHz=CH - CHZCH,. With propane, however, there is a great gain in the process and this is in fact slightly faster than the .decomposition of propylene of which the first step is probably CH,CH,CH3 -+ CH, + CH,=CH, CH,CH=CH, -+ CH=CH + CH,. In the absence, however, of a quantitative measure of symmetry such considerations unfortunately lack precision, and will therefore not be enlarged upon. Physical Chemistry Laboratory, South Parks Road, Oxford

ISSN:0366-9033    

DOI:10.1039/DF9511000129

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

'I 36 REACTION OF METHYL RADICALS THE REACTION OF METHYL RADICALS WITH HYDROGEN BY R. D. ANDERSON, S. DAVISON, AND &I. BURTON Received 29th January, 1951 The reaction between methyl and molecular hydrogen follows the simple mechanism CH, +H,+CH, +H. Interpretations of the older literature are confused by use of low activation energy and low steric factor for this reaction. The effects of reaction order, hot radicals, surface and uncertainties regarding elementary reactions have been largely eliminated. The results lead t o a value of El? 13 kcal. and S1= I O - ~ . The former figure is in agreement with recent best values on activation energy of the reverse reaction and pertinent bond dis- sociation energies. The reaction CH, + H, +CHI + H, . * (1) CH, + RH +CHI + R, , * (2) has received much attention because it is the simplest of the important series where RH represents hydrogen, a simple aliphatic compound, or an oxygen derivative of the latter.However, the activation energy E, and the steric factor for reaction (I) have been disputed. 1 Contribution jointly from the Radiation Chemistry Project operated by the University of Notre Dame under U.S. Atomic Energy Commission Contract AT( I I-1)-38 and the Sinclair Research Project. Sinclair Fellow.R. D. ANDERSON, S. DAVISON AND hl. BURTON 137 Table I includes a list (in chronological order) of the various activation energy values proposed for reaction ( I ) with a brief description of the basis for each value. It will be shown that the new and higher value recom- mended herein of about 13 kcal.for El and a steric factor of about IO-, remove apparent anomalies which exist when the lower values for these quantities are assumed. TABLE ACTIVATION ENERGY VALUES Authors Hartel and Polan yi F. 0. Rice Paneth, Hofeditz and Wunsch Patat and Sachsse Taylor and Rosenblum Cunningham and Taylor Phibbs and Darwent Anderson and H. A. Taylor Davison and Burton Davison and Burton Basis Decomposition of methyl halides by sodium atoms in hydrogen as carrier gas Detection of methyl radicals from pyrolysis of Pb(CH,)* in the presence of hydrogen Recombination of methyl radicals from the pyrolysis of Pb(Cg,), in hydrogen as carrier gas Para-ortho hydrogen conversion in the presence of methyl radicals Photolysis of acetone plus hydrogen Photolysis of mercury dimethyl in the presence of hydrogen Photolysis of mercury dimethyl in the presence of hydrogen Photolysis of cadmium dimethyl in the presence of hydrogen High temperature photolysis of acetaldehyde with deuterium High temperature photolysis of acetone, with deuterium, and with a mixture of hydrogen and deuterium The value of E , may also be calculated from presently accepted bond dissociation energies and the activation energy of the reverse reaction H + CH, --f CH, + H, .where E , = 11 f 2 k ~ a l . , ~ , 4 CHI --f CH, + H - IOI 5 I kcal.5 . * (4) H, 3 2H - 103 kcaLg . - ( 5 ) Hence El should be equal to about 13 kcal. This value agrees excellently' with the new higher value for E, found by Anderson and Taylor and by Davison and Burton.8 In both cases E, has been evaluated by direct experimental methods free from several objectionable features of earlier work.Early Work.-Hartel and Polanyi investigated decomposition of the methyl halides (chloride, bromide and iodide) by sodium atoms in hydrogen as carrier gas. They assumed, for methane production, the mechanism Na + CH,X +NaX + CH, - (6) CH, + Ha +CH, + H. . * (1) Steacie, Darwent and Trost, Faraday SOC. Discussions, 1947, 2, 80. Evans and Szwarc, Trans. Faraday SOC., 1949, 45, 940. 5 Szwarc, Chem. Rev., 1950, 47, 75. 6 Gaydon, Dissociation Energies and Spectra of Diatomic Molecules (Chapman 7 Anderson and H. A. Taylor, forthcoming paper in J . Physic. Chem. 8 Davison and Burton, forthcoming publication. and Hall, London, 1947). Hartel and Polanyi, 2. 9hysik.Ckem. B, 1930, 11, 97.13s REACTION OF METHYL, RADICALS No mention was made of the fate of the hydrogen atoms produced by re- action (I). Furthermore, it was assumed that the methyl radical con- centration was temperature independent. The rates of methane pro- duction were inferred from the rates of hydrogen consumption, and the temperature dependence of the latter provided an estimated El value of 6-8 kcal. The experimental approach did not permit a very accurate estimate of the actual reaction temperature. The low El values pro- posed by Hartel and Polanyi should not, consequently, have received much credence. However, they actually had much in'fluence on later considerat ions. objected to the low value of about 8 kcal. since it led to the result that at 600° C only exp (- 8000/2 x 873) or about IOO collisions of CH, and Ha would be required for reaction.If this low number of col- lisions were correct, it should not be possible to detect methyl radicals in the presence of hydrogen at 600' C. Actually, Rice did detect methyl radicals in experiments carried out under such conditions and on the basis of collision theory, proposed a high El value of about 20 kcal. This estimate would be consistent with El= 13 kcal. and steric factor Paneth, Hofeditz and Wunsch l1 investigated reactions of methyl radicals from thermal decomposition of lead tetramethyl in hydrogen as carrier gas. Although the authors did not do so, an El value of about 15 kcal. was calculated from their data by Patat. Patat and Sachsse l2* la, l 4 published a series of papers concerning occurrence of chain reactions, measurement of methyl radical concentra- tions, and selection of activation energy value and steric factor for reaction (I).Their general procedure was to photolyze or pyrolyze a compound in the presence of para-hydrogen. A steady state hydrogen atom con- centration was assumed due to the reactions Rice SIC= 10-2. and CH, + H, +CH, + H . - (1) H + A + H , + R , . - ( 7 ) where A represents the parent organic compound, e.g., acetone, acetal- dehyde, ether, etc., and R represents a free radical. The usual kinetic treatment yields the expression Experimental procedure kept (A) and (H2) essentially constant. The (H) was determined from the observed rate of conversion of para-hydrogen to ortho-hydrogen by the method of Geib and Harteck.If a very accurate value for K,/K, could have been found, a method for the measurement of the (CH,) would have been available. In order to evaluate the ratio kl/k,, Patat and Sachsse assumed a particular value (- 8 kcal.) for El and used the appropriate E7 value available at the time. The (CH,) in the acetaldehyde case was evaluated on the basis of a detailed decomposition mechanism of Leermakers and was used to complete the evaluation of k1/K7. The selection of a steric factor of about 1 0 - 4 for (I) was a result of the use of the El value of 8 kcal. The entire scheme for measuring methyl radical concentration depends so completely upon assumptions that, in the words of Steacie,16 " it is lo Rice, J . Amer. Chem. SOC., 1934, 56, 488.11 Paneth, Hofeditz and Wunsch, J . Chem. Soc., 1935, 372. 12 Patat and Sachsse, 2. physik. Chem., 1935, 31, 105. Patat, 2. physik. Chem. B, 1936, 32, 274. 14Patat, 2. fihysik. Chem. B, 1936, 32, 294. 15 Steacie, Atomic and Free Radical Reactions (Reinhold, New Yorlr, 1946), P- 64-R. D. ANDERSON, S. DAVISON AND M. BURTON 139 difficult to obtain unequivocal results by this method ”. Nevertheless, Patat and Sachsse calculated methyl radical concentrations by the application of (I) to the decomposition of each substance studied. They then compared the results with the theoretical values obtained by application of a detailed Rice-Herzfeld chain mechanism to each decom- position. Where the latter was not available for a particular reaction, the general expression (CHd EI-E7 Expermental (cal.) (E1==g kcal.in Formula (I)) Substance CH30CH, 797 -2000 Io-11.8 CH3COCH3 820 0 10 -10.23 C2H5CH0 8 20 2000 10-10.04 C3H8 834 0 10-10.2 CH,CHO 820 I 800 10-9.91 was used. hexpt., and kchsin represented the rate constant for the reaction The observed experimental rate constant was represented by CH, + A -+CHI + R. . * (8) Table I1 contains some of the data reported by Patat. The fact that the Rice-Herzfeld theory leads to methyl radical concentrations which are larger in every case than those found by Patat has been used as an argument against the former theory. In column 6 we have recalculated the experimental (CH,) using El = 13 kcal. The methyl radical con- centrations so obtained are in better agreement with those obtained by the Rice-Herzfeld chain mechanism.However, there should be no expectation of close agreement, since the method of Patat l4 for measuring free methyl radical concentration is probably not accurate. In any case, it does not provide sufficient reason for abandoning the concept of chain mechanisms in reaction kinetics. (CH3). (CH3) Theoretical Experimental (Chain (El=13 kcal. Mechanism) in formula (I)) I 0 -9.39 10-10.7 10-9.66 1 0 -9.18 10-8.65 10-8.89 10-9.56 10-9.2 10-8.16 10-8.86 * All concentrations are in reality calculated. See discussion. Taylor and Rosenblum 16 photolyzed acetone in presence of hydrogen. They assumed that methane was produced by reaction (I) and proposed several methods for ethane production. Several mathematical treat- ments were employed with their results to give an average estimate of - II kcal.€or El. Cunningham and Taylor investigated the photolysis of mercury dimethyl in the presence of hydrogen and proposed the following mechan- ism for methane production Hg(CH3)2 + h v +- 2CH3 + Hg . - (9) CH, + H2 +- CH, + H . - (1) H + Hg(CH3)z -+ CHI + HgCH3 - (10) By the usual steady state treatment, they were able t o show that the rate of methane production was proportional to the product of concentra- tion of methyl radicals and concentration of hydrogen. They assumed both of the latter to be constant over the temperature range studied and l6 Taylor and Rosenblum, J . Chem. Physics, 1938. 6, 119. l7 Cunningham and Taylor, J . Chem. Physics, 1938, 6, 359.130 REACTION OF METHYL, RADICALS found El = 8.1 kcal.from the temperature dependence of the rate of methane formation. However, they referred to the work of Hartel and Polanyi, and of Taylor and Rosenblum, and then reported g f 2 kcal. as the “ most trustworthy ” value for El. In a further study of mercury dimethyl Phibbs and Darwent l8 showed that a “ h o t ” methyl effect was a plausible explanation for results of photolysis in the presence of hydrogen. Should this effect persist into the higher temperature range, the calculated value for El would be lower than the true one. Since methyl radicals have been shown to react with metal deposits, a back reaction between methyl radicals and mercury undoubtedly takes place. A complication arises from the fact that at lower temperatures any mercury set free during photolysis is deposited almost entirely on the walls of the reaction vessel, while in the high temperature range, the mercury would exist entirely as a vapour.Any back reaction of methyl radicals with mercury would be of a heterogeneous nature in the former cases, but of a homogeneous nature in the latter case. Recent Experimental Results .-Anderson and Taylor investigated the photolysis of cadmium dimethyl in the presence of hydrogen. Enough runs were made to permit calculation of rates of methane production at zero time. The quartz reaction vessel was a short cylinder 5 cm. long which was fitted at each end with plane quartz windows 10 cm. in dia- meter. The surface of each window was irradiated by its own 2537 A light source. The relatively large surface irradiated prevented serious light attenuation by the cadmium deposits produced during photolysis.It should be noted that the cadmium was deposited almost entirely on the walls of the reaction vessel over the complete range of temperature em- ployed (5ooC to 250’C). Hence, any back reaction between methyl radicals and cadmium would be of a heterogeneous nature at all times. Certain precautions were taken to ensure accurate temperature control. The furnace containing the reaction cell was controlled automatically so that its temperature was within 0.1’ C of the desired temperature. The light sources were in a thermostat whose temperature was within 0.2’ C of the constant temperature at which the latter was maintained. Suf- ficient pre-heating periods were used so that the reaction mixture would be at the temperature of the furnace during a run.Loading of the reaction vessel was carried out in such a way that no mercury vapour could enter the reaction vessel, and all analyses were carried out on a Consolidated Engineering mass spectrometer. The cadmium dimethyl used was made from the Grignard reaction involving methyl magnesium iodide and anhydrous cadmium chloride in ethyl ether. The last traces of ether were removed from the crude product by very efficient fractional distillation. The resulting product had a freezing point of - 2.4’ C over the entire range of solidification. The important part of the decomposition mechanism is as follows : Cd(CH,), + hv -+ 2CH3 + Cd * (11) 2CH3 + C,H,. (12) CH, + H, --f CH, + H .* (1) H + Cd(CH,), -+CH, + Cd + CH, . * (13) The reaction of methyl radicals with cadmium dimethyl to yield ethane was ruled out as insignificant since ethane production was independent of the concentration of cadmium dimethyl. A plot of log B,, against I / T gave a line whose slope increased constantly with increasing temperature. The best explanation of this was an increasing concentration of methyl radicals. From reaction (12), RlZ = h12 - (CH,I2 - . (111) :. (CH,) = (R12/K12)*. . . (IV) Phibbs and Danvent, Trans. Faraday Soc., 1949. 45, 541.R. TI. ANDERSON, S. DAVISON AND M. BURTON 141 However, the total rate of methane production, and by the usual steady state treatment Rm = 2K1. (CH,)(H,) . - (VI) But (H,) is essentially constant and may be set equal to B .When the proper substitutions are made for (CH,) and (H,) in (VI), there is obtained the expression The different (R,/Rl,*) values were substituted in pairs in the integrated form of the Arrhenius equation in such a way as to make equal use of all values. Temperature interval weighting was employed, and the final ( E l - +El,) value obtained was 13 f z kcal. I f the value of El, is taken as zero, then El = 13 f z kcal. A consideration of the energies involved in the reactions Cd(CH,), -+ 2CH, + Cd * (14) and Hg(CH3)2 --t 2CH3 f Hg * (15) indicates that the " hot '' methyl effect in the case of cadmium dimethyl should be much less than that attributed to mercury dimethyl. Actually no such effect was found with cadmium dimethyl over the range where methane analysis was accurate.The low El value (9 kcal.) and the low steric factor (10-4) sponsored by Steacie et aZ.,, 2 o for reaction (I) seem to be based upon the work with mercury dimethyl, whose defects have been noted above, and upon the conclusions of Patat and Sachsse in their attempt t o reconcile their experi- mental data with the low El value proposed by Hartel and Polanyi. The new high value for El removes the need for assuming a low steric factor to account for the slowness of reaction (I) at ordinary temperature. Furthermore, the work of Evans and Szwarc 4 shows that relatively high steric factors are to be expected for reactions between methyl radicals and the simple hydrocarbons. Reaction ( I ) is practically in this class, and Davison and Burton actually found a steric factor of -10-2 for ( I ) .Davison and Burton * studied photolpses of acetone and of acetal- dehyde in presence of deuterium, and in presence of equimolar mixtures of hydrogen and deuterium. One result of this work was elucidation of the mechanism of the reaction between methyl and molecular hydrogen. Eyring and his associates 21 calculated a potential energy surface for re- action (I). They proposed formation of a stable CH,-H, complex based upon a rather deep hole intercepting the path of reaction on the potential energy surface. Methane formation could then be ex- plained as reaction of this complex with another methyl radical. As a matter of fact, Taylor and Burton 2 2 had indicated that some of the diffi- culties presented by the Patat and Sachsse results were explicable on the basis of such a termolecular mechanism.However, the experimental results of Davison and Burton indicated that such a reaction was extremely unlikely since a value of - I was obtained by them for the ratio HD/CH,D during the photolysis of acetone plus deuterium. Such a ratio would be expected from the successive reactions CH, + D2 -+CH,D + D - (16) and D + C,H,CO -?. HD + CH,COCH,. * (17) 19 Long and Norrish, Phil. Trans. Roy. Soc. A , 1949, 241, 587. 2o Trotman-Dickenson and Steacie, J . Chem. Physics, 1950, IS, 1097. 22 Taylor and Burton, J . Chem. Physics, 1939, 7, 676. Gorin, Kauzmann, Walter and Eyring, J . Chem. Physics, 1939~ 7, 633.142 REACTION OF METHYL RADICALS The formation of HD and the value of the ratio HD/CH,De I provide strong evidence for the reality of reaction (I).When acetone is photolyzed in the presence of deuterium by the self- reversed radiation of a high pressure mercury arc, reaction (16) is ac- companied by The pressures of acetone vapour and deuterium are essentially constant so that it is possible to write CH, + (CH,),CO -+ CH, + CH,COCH,. . - (18) (VIII) In order to ensure the strict applicability of (VIII), and also to minimize the effect of side reactions of the products such as the experimental rates of methane production were corrected to zero time at each temperature. A plot of log (h16/i&) against I/T for different temperatures in the range from 1 5 0 O C to 4z5O C was then used to deter- mine the quantities (& - E18) and &6/&8, where S is used to represent steric factor.A similar procedure was used for the photolysis of acetal- dehyde in the presence of deuterium. Reactions (I) and (16) should have different rate constants because of the difference in zero point energies of H, and D,. To correct for this difference, a series of runs was made where acetone was photolyzed in the presence of an equimolar mixture of hydrogen and deuterium. Since some methane arises by reaction (18) it is necessary t o calculate this amount .by substituting the known (CH,D) in (VIII). This correction is subtracted from the total (CH,) resulting from the photolysis of acetone in the presence of deuterium and hydrogen. The (CHJcorr. and the measured (CH,D) are then used in the following expression D + CH, -+CH,D + H - (19) to evaluate the ratio &/k16 at different temperatures. A plot of log (kl/&) against I/T will permit the evaluation of (El6 - El) which was found to be 1-1 kcal.The difference in zero point energies between the H, and D2 molecules is 1.8 kcal. Since this method yields differences in activation energies and ratios of steric factors, it is necessary to select suitable reference points. Trotman-Dickenson and Steacie 2 o have studied the rates of production of methane and ethane in the photolysis of acetone and have arrived at the values El, - +El, = 9.7 kcal. Similar work by Saunders and Taylor 23 gave a value of 9.6 kcal. for The work of Grahame and Rollefson a4 on the high temperature and S,*/S,,B = 10-3. (El8 - *-w- photolysis of acetaldehyde where CH3 + CHSCHO + CH4 + CHaCO * (20) gave a value of 8.6 kcal.for (E20 - *El$). Since other experimental work indicates a small or zero value for El, we can set E12 = o in the above expressions. Recent work by Gomer 25 and by Szwarc and Roberts 26 shows that S,, is close to unity. From the assumption S,, = I 2s Saunders and Taylor, J . Citem. Physics, 1941, 9, 616. 24 Grahame and Rollefson, J . Chem. Physics, 1940, 8, 98. 25 Gomer, J . Chem. Physics, 1950, 18, 998. 26 Szwarc and Roberts, Trans. Faraday Soc., 1g50,46, 625.R. D. ANDERSON, S. DAVISON AND M. BURTON 143 and the quantities evaluated by Davison and Burton El6 - El = 1-1 kcal. El6 - El, = 4.6 kcal. El, - E,, = 6 2 kcal. and sl*/sl = 10-1, the data of Table I11 have been calculated. El kcal. from Reaction (CHdaCO CHsCHO TABLE III.-DAVISON-BURTON RESULTS Steric Factor CH, + H2 -+ CH, + H CH, + D, -+ CH,D + H In the case of acetaldehyde the hot radical effects are insignificant because the experiments were conducted in a temperature range where the chains were long. Thus, if the difference between the results from acetone and acetaldehyde is real, the upper values are to be preferred. The Davison-Burton method is completely independent of assumptions regarding the methyl radical concentration since methane production was always first order with regard to the (CH,) terin which cancelled out of the kinetic expressions used. Hence, it is not necessary to main- tain a constant value of (CH,) throughout the reaction vessel during the runs. Department of Chemistry, For the same reason, constant light intensity is not required. University of Notre Dame, U.S.A. I 3-2 10-2 14-3

ISSN:0366-9033    

DOI:10.1039/DF9511000136

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

R. D. ANDERSON, S. DAVISON AND M. BURTON 143 THE THERMAL STABILITY AND REACTIVITY OF HYDROCARBON RADICALS BY M. SZWARC Received 12th February, 1951 The bond dissociation energies of some hydrocarbon radicals are calculated and the vpriation in their magnitudes discussed in relation to radical structure. Various types of reactions involving radicals are treated, special attention being paid to the problem of repulsion forces in metathetic reactions of the type R + XR* --f RX + R*. Some conclusions are drawn regarding the shapes of various repulsion curves. The problem of the temperature independent factors of radical reactions is examined. It is shown that the experimental findings can be well accounted for by the transition state theory. The behaviour of radicals in chemical reactions is somewhat different from that of ordinary molecules.The essential difference is caused by their low thermal stability and by their high reactivity. The Thermal Stability of Radicals.-The low thermal stability of many hydrocarbon radicals when compared with the parent hydrocarbons is due to the considerably decreased dissociation energy of some of their bonds. CH, . CH,*H + CH, : CH, + 11. Let us clarify this point by considering the dissociation processI44 REACTIVITY OF RADICALS In this reaction the extension of the C-H bond coincides with a coupling process in which the free electron of the radical couples with one of the electrons participating in the C-H bond. Consequently, the rupture of the C-H bond occurs simultaneously with the formation of a new T bond (the “ second half of the double bond ”) and the energy 1ibera.ted in the formation of the latter bond is utilized in the process of fission of the former bond.Hence, the C-H bond dissociation energy in the ethyl radical is lower than the C-H bond dissociation energy in ethane by the energy of the “ second half ” of the double bond. Table I presents some C-H and C-C bond dissociation energies in a series of simple hydrocarbon radicals. The listed values have been calculated by applying the equation D(R-olefine) = AHf(R) + AH, (olefine) - AHf (R . olefme). R . olefine denoting the relevant radical, and R a hydrogen atom or a smaller radical produced in the dissociation process. The required heats of formation of radicals have been taken from a previous article.1 TABLE I D(*CH,.CH,-H) = 38 kcal./mole D(‘CH”)CH-H) =-- 39 kcal./mole D(CH, . &t. CH,-H) =: 45 kcal./mole CH, *CH; ( 3 3 3 \CH-H) = 40 kcal./mole \c. CH2-H) = 4G kcal./mole D(*CH,. CH2-CH,) = 25 kcal./mole D(*CH,. CH2-C2Hs) = 26 kcal./mole *CH YC-13) = 68 kcal./mole *CH D( CH,/ \C-CH,) = 57 kcal./mole Although the quoted values are not altogether certain and should be taken with reservation, they do show definite regularities and enable one to draw interesting conclusions.* (I} In the series, CH, . CH2-H, CHZ . CH-H, *CH, . CH-H t H 3 &S we find a small increase in the C-H bond dissociation energy. other hand, in the series On the CH3. CH2-H, CH3. CH-H, CH, . CH-H 6H , i‘.2HS there is a definitz and quite considerable decrease in the C-H bond dis- sociation energies, the relevant values in kcal./mole being 98, 89 and 86 -f respectively .Szwarc, Chem. Rev., 1950, 47, 75. *The results for D[(CH,),CH-HI and D[(CH3),C-H] of Stevenson in this 7 D( cH3>H-H) was estimated by subtracting from the value of \CH-H), 3 kcal./mole, which is the difference between D(CH,. CH,-H) Discussion invalidate some of our conclusions. CZHS CH, D and D(C,H, . CH,-H). (CH3/M, SZWARC I45 A useful interpretation of these results can be given by considering the ( I ) (2) following series of reactions : CH, . CH, . H +CH,. CH,. + H - D(X. CH,. CH,-H) kcal.jmole, and CH, . CH. H + CH, . CH- + H - D(CH3 . CH-H) kcal./mole, x x k k 2 X denoting an alkyl or a q l radical. energy is not affected by the presence of the X group, i.e.I n the second case this assumption cannot be made since the proximity of the X group to the reacting centre certainly changes the CH, . CH-H Let us assume that in the case (I) the X . CH, . CH,-H dissociation D(CH3. CHS-H) w D ( X . CH, . CHZ-H). dissociation energy making k D(CH3. CH2-H) > D(CH3. CH-H). ;ri. This decrease in the C-H bond dissocistion energy might be attributed to the inereas? in the resonance energy of the CH, . CH- radical as com- pared with that of the CH, . CH2* radical.8 In another way, one might say that the delocalization of the free electron is greater in the radic.al CH, . CH- than in the radical CH, . CH,.. 2 k It is plausible that the same gain in delocalization energy would be observed if one compares the hypothetical processes (3) CH, .CH . H 4 CH, . CH* + H k ;;r CH2. CH, . H -+ CH8 . CH,* + H However, the energy gained in process (3) as compared lost in process (5). CH, . CH* 3 CH, : CHX . x and (4) (3) * (4) with (4) would be and if the delocalization of 7~ electrons in CH, : CHX is about the same a s that in CH, : CH,, the results would be D(CH2. CH2-H) D(*CH,. CH-H). k The resultc quoted in Table I seem, therefore, to indicate th3.t the sub- stitution of a group X for a. hydrogen on a carbon atom having a free electron has much greater effect on its delocalization energy than it hds on the delocalization energy of the T electrons of the corresponding olefin. one finds a considerable increase in the C-H bond dissociation energy. This effect is in agreement with the preceding discussion.The substituent X affects only slightly’ the dissociation energy of the hypothetical process (2) Comparing D(.CH2 ~ CH,-H) with D(CH, . dH . CH,-H) CH . CH, . H -+ CH . CH,. + H - E , kcal./mole CH, . CH, . H -+ CH, . CH,. + H - E , kcal./mole. x 5i i.e. E, w E,. Szwarc, J . Citem. Physics, 1950, 18, 1160REACTIVITY OF RADICALS However, in the coupling processes CH . CH,. -+ CH : CH, + E,’ kcal./mole, ;;c k and CH, . CH,. -+ CH, : CH, + E,,’ kcal./mole, the difference E,‘ - E,’ is approximately equal to the difference between D(CH,. CH,-H) and D(CH,. CHX-H) (i.e. the “ delocalization ” energy in the CH, . CH, radical is lost in the coupling process). A We conclude, therefore, that D(CH,. 6H. CH,-H) - D(&I.,CH,-H) w D(CH,. CH,-H) CH, - D ( cH3>H-H).Inspection of the experimental values shows that the former difference is 7 kcal./mole, whilst the latter is g kcal./mole. ( 3 ) The loss of considerable resonance energy of the allyl radical is the reason for a great increase in its C-H bond dissociation energy as compared with that in n-propyl radical. (4) The difference in D(R-H) and D(R-CH,) is only slightly smaller for radicals than for molecules. For example, D(CH,. CH,-H) - D(-CH,, CH,-CH,) = 13 kcal./mole. D(CH,. CH,-H) - D(CH, . CH,--CH,) = 16 kcal./mole. As mentioned previously the decrease in the thermal stability of a radical is due to the decrease in the relevant bond dissociation energy. Whenever the radical is stabilized by resonance the relevant bond dis- sociation energy increases, and therefore its thermal stability increases too (see, e.g., n-propyl radical as compared with allyl radical).For the sake of accuracy one has to point out that the thermal stability is measured not by the dissociation energy but by the activation energy of the dis- sociation process. The latter is larger than the bond dissociation energy by the activation energy of the addition process olefine + R --f olefine . R. If R is a hydrogen atom or a methyl radical the relevant activation energy of the addition process seems to be small, and most likely it does not change the gradation in the stabilities of radicals as obtained from con- siderations of bond dissociation energies. A great increase in the thermal stability of a radical is found foi all cases where the decomposition cannot lead to the formation of a double bond. For example, no double bond can be formed in the decomposition of benzyl radical, and this makes the benzyl radical thermally more stable than the allyl radical, although the resonance stabilization of the latter is perhaps even higher than that of the former (estimates quoted from ref.(2) lead to ca. 24 kcal./mole for the resonance energy of benzyl radical and to ca. 25 kcal./mole for that of allyl radical). On similar grounds a high degree of thermal stability might be expected for such radicals as CH,. I It would be interesting to investigate the thermal stability of cyclopropyl radical, since the formation of a double bond would lead in this case to a highly strained structure. Of course, this does not exclude the possi- bility of an easy isomerization of the cyclopropyl radical to an allyl radicalM.SZWARC 147 and the subsequent decomposition of the latter to allene and a hydrogen at om. The Reactivity of Radicals.-In discussing the reactivity of radicals we shall consider two aspects of this problem : (a) reactions between two radicals and (b) reactions between radicals and molecules. The reactions between radicals may lead either to their dimerization or to their disproportionation. I t seems that the dimerization of many radicals does not require any activation energy. Evidence favouring this opinion has been obtained recently by independent studies of Gomer,3 of Rice and Lucas,* and of Dodd.6 All these workers investigated the rate of dimerization of methyl radicals CH, + CH3 + C&&, and their findings show that this reaction takes place nearly at every collision,* i.e.its activation energy is negligible. Saying that a re- action takes place at every collision implies that we can calculate in- dependently the frequency of encounters among the particles (or that we know the relevant cross-sections). However, it seems that the concept of a collision may be interpretzd in more than one way. Whether an encounter is considered as a collision depends on the type of interaction in which one is interested. It is common practice in chemical kinetics, especially when the collision theory is applied, to calculate the number of collisions from data obtained in measurements of the viscosity or the heat conductivity of a gas.This implies, of course, that we are interested in such collisions in which the translation energy of colliding particles can be exchanged. It is not certain, however, that this type of collision is relevant when a chemical reaction is considered (see, e.g., Glasstone, Eyring and Laidler ).6 Hence, it appears that the normal value for 2 (i.e. collision number) used in chemical kinetics is chosen quite arbitrarily. A more analytical approach to the problem of chemical kinetics is provided by the transition state theory.', * J * Unfortunately this approach requires a knowledge of the structure of the transition state (also referred to as the activated complex), which is a hypothetical species, not accessible to direct experimental measurements.Consequently, any calculations based on the latter method must remain vague, and the a priori computa- tion leading to numerical values is a matter of intelligent conjecture of the structure of the transition state. Nevertheless, this treatment helps in the understanding of the trends which are actually observed and makes it possible to interpret them in terms of some molecular models. Returning to the problem of recombination of methyl radicals, we have to point out that the high value for the rate constant of about 5 x 1013 cm.3 moles-1 sec.-l, reported in the recent investigations, might indicate that the collision diameter for the recombination process is large. Thi: would mean that the interaction between two methyl radicals which lead: to the dimerization might take place at a greater distance than that a1 which, say, two methane molecules can exchange their kinetic energy.On the other hand, one would expect that the methyl radical has to approach quite closely the molecule RH in order to react with it according to the equation CH, + RH -+ CH, + R. The last statement requires some clarification. Gomer, J . Chem. Physics, 1950, 18, 998 ; Gorner and Kistiakowsky, ibid., Rice and Lucas, J . Chem. Physics, 1950, 18, 993. 1951, 19, 85. 5 Dodd, Trans. Faraday SOC., 1931, 47, 56. * See, however, Miller and Steacie, J . Chem. Physics, 1951, 19, 73. 6 Glasstone, Eyring and Laidler, Themy of Rate Processes (McGraw-Hill, 1940). 7 Eyring and Polanyi, Z. physzk. Chem. B, 1931, 12, 279. Eyring, J . Chem. Physics, 1935, 3, 107.Evans and Polanyi, Trans. Faraday SOC., 1935, 31, 875.148 REACTIVITY OF RADICALS If this view is correct then one would ascribe to the methyl radical a large collision cross-section when a recombination process is considered and a small collision cross-section when it reacts with RH. The purpose of this discussion is to show that one can, in principle, ascribe different cross-sections to the same species according to the type of reaction which is investigated, a point which has been realized for some time. For instance, it is known that in experiments in which the scattering of particles by some medium is measured, the cross-section of the particle depends on its energy. An example is provided by the recent studies of the scattering of hydrogen atoms by hydrogen molecules.lo Although it is likely that on the whole the recombination of radicals does not involve any activation energy, some radicals have to surmount an activation barrier before recombining, e.g. in the dimerization of tri- phenyl methyl radicals.11 It seems that in this process the activation energy arises due to the necessity of overcoming the steric hindrance between the bulky phenyl groups.l* The recombination of polar radicals was discussed by Weiss 13 who suggested that these reactions might re- quire activation energies to overcome the repulsion between the approach- ing dipoles. A mathematical treatment of such a problem has been given by Heitler and Rummer.14 It seems that our present knowledge is not sufficiently well advanced to allow us to make any definite statement about the relative rates of disproportionation of radicals as compared with their dimerization. The experimental facts reported in the literature are highly confusing, l5 and it appears that experiments with as simple systems as possible are needed before one can hope for the elucidation of this problem.From general considerations, it would be expected that the disproportionation of radicals would involve only a small activation energy, and that the probability of occurrence of such a process, i.e. the PZ factor, would be smaller than for the dimerization reaction. Reactions between radicals and molecules will be considered under two headings : the activation energy and the temperature independent factor (the PZ factor). In the methathetic reactions of the type R + XR, +RX + R, the exothermicity of the process is given by D(R-X) - D(R,-X).It was suggested that for some processes of this type, namely, RX + Na + R + X- Na+ there is a simple relation between the activation energy of the process and its exothermicity, 16 AE and AH denoting the increase in the activation energy and in the exothermicity of the sodium flame reaction when two reacting species RX and R,X are compared. In deriving this relation one has to assume that the sodium atom can approach the reacting centre without overcoming any repulsion, and that the repulsion curve for the R-CI- interaction is independent ot the nature of the radical R. These assumptions seem to be justified for sodium flame reactions (see, however, ref.(IZ)), but they might not be valid for such reactions as, e.g., RH + CH, + R + CH,. AE = a. AH, lo Amdur, Kells and Davenport, J . Chem. Physics, 1950, 18, 1676. l1 Ziegler, Trans. Faraday Soc., 1934, 30, 10. l2 Szwarc, Faraday SOG. Discussions, 1947, 2, 39. l3 Weiss, Trans. Faraday SOL, 1940, 36, 856. l4 Heitler and Rummer, 2. Plzysik, 1931, 68, 12. I5 See, e.g., Steacie, Atomic and Radical Beactions (Reinhold, 1946). * The disproportionation process is highly exothermic, since the dissociation energy of the bond formed is greater than the dissociation energy of the bond broken. l6 Evans and Polanyi, Trans. Faraday Soc., 1936, 32, 1933 ; 1938, 34, 22.M. SZWARC 149 In considering the variation in the activation energies of such reactions one has to bear in mind the changes of the resonance energy in the transi- tion state 17 which is not negligible for this type of process.It seems profitable to compare a metathetic reaction of the type RA + Rl -+ R + AR, with the thermal decomposition of RA into two fragments produced by the rupture of one bond only. In the latter case the bond ruptured is always the weakest bond, i.e. the bond which has the lowest dissociation energy. However, this need not be the case for methathetic reactions. For example, let us consider two methathetic reactions between a mole- cule AMB and a radical R, namely, AMB + R +AM + BR and BMA + R +- BM + AR. It is not sufficient to consider the dissociation energies of AM-B and BM-A bonds in order to say which of the two reactions is more likely to occur.The result depends very much on the values of D(R-A) and D(R-B), since the energy liberated in the process of bond formation provides the driving force for the reaction. This problem was treated by Evans and Polanyi,l* who also showed how important is the consideration of the repulsion forces for understanding these processes. Let us consider, e.g. a series of thermoneutral reactions of the type R , X + R* + R + XR* R . Y + R* + R + YR*. Evans and Polanyi suggested that this reaction can be considered as taking place in three steps : (i) R* approaches the rigid molecule RX until the distance R . . . R* is that in the transition state complex, without affecting the R-X bond length, i.e. the configuration obtained is R-X . . . R*. - (1) R . . . X-R* * (11) (ii) Keeping the R .. R* distance constant, we move X from the position shown in (I) to the position shown in (11) where X-R* is the normal R*-X bond length in the molecule R*X. (iii) R is removed to infiGity. Fig. I and 2 represent schematically the overall process, the arrows denoted by number I, z and 3 representing the three stages of the reaction suggested. The total activation energy E , can be considered as composed from two parts, E,. resulting from the repulsion forces and Et resulting from the extension of the bond which is ultimately broken, i.e. It is plausible to expect that Et increases with increasing R-X dissociation energy, and hence the knowledge of bond dissociation energies enable us to understand the trend in Et. However, our knowledge of E, is scanty’ and its variation with molecular structure is not known.Some observa- tions of Kharasch may shed light on this problem. investigated the addition of chloroform and bromoform to double bonds and it follows from their observations that radical R reacts with chloroform and bromoform according to eqn. E, = Et + E,. Kharash and his colleagues R + CHBr, -+ RBr + CHBr,. . 17 Evans and Szwrarc, Trans. Faraday Soc., 1949, 45, 940. 18 Evans and Polanyi, Trans. Furuduy Soc., 1938,34, 11. 19 Kharash, Jensen and Urry, J . Amev. Chern. Soc., 1947, 60, 1100.150 REACTIVITY OF RADICALS Data of bond dissociation energies indicate that in general D (C-H) > D (C-CI) > D (C-Br) . Therefore we would expect that for reactions R + CHCI, -+ RH + CCl, - (1) R + CHCl, -+ R .C1 + CHCI, (2) . Et, >Eb. On the other hand, the results of Kharash seem to indicate that E,< E,. To explain these observations one would assume t h a t the repulsion forces increase along the series H, C1, Br, i.e. E, < Er? < E,, the results of Kharash being then interpreted in terms of the diagram (Fig. 3). ET, < Er2 < Er, Et, > Et, > Et, Ea, < Ea, < EQ- I R-X d / ; h c e --+ FIG. I. c 4 I I 6 h A - - ---- x FIG. 2. TIT Our knowledge of repulsion forces is very scanty, and it is gratifying that the above discussion might furnish information about the shape of various repulsion curves. Fig. 3, taken from the paper of Evans and Polanyi,I* shows the method for finding the position of the transition state and for calculating the activation energy of the process.Curves I and I’ represent the repulsion potentials which hinder the approach of a radical R to a rigid molecule XR*, and of a radical R* to a rigid molecule XR. Curves z and 2’ represent the potential energies corresponding to X-R* and X-R bond extension processes. The position of curve I relative to curve I’ is fixed by the properties of the system investigated. To find the position of the transition state (point T in Fig. 3 ) one has to slide curves 2 and 2’ along the curves I and I’ respectively until their crossing point is as low as possible ; this crossing point then represents the transition state. Its elevation from the 0 line gives the required activation energy of the process.151 M. SZWARC Simple geometrical considerations show that when curves 2 and 2’ attain the required positions, the slope of curve z a t the crossing point T must be equal to the slope of curve I a t point A.If we replace atom X in the molecule XR* by an atom Y, and let D(R*-Y) be greater than D (R*-X) . The potential energy curve representing the extension of R*-Y bond would be steeper than the one corresponding to the R*-X bond. Consequently, if one uses the same repulsion curve for both pro- cesses (i.e. involving R*X and R*Y) one finds the RY curve (the dotted curve in Fig. 3 ) to be placed higher than the corresponding RX curve. Since E, is measured by the level on which point A (or A’) is placed, we conclude that E, would increase with increasing dissociation energy of the bond in question, if the same repulsion curve is used for both reactions : and R + XR* -+ RX + R* R + YR* -+RY + R*.,e-- I / FIG. 3 We have seen, however, that the results of Kharash suggest a decrease in E, for a series R-Br, R-Cl, and R-H, i.e. for a series in which the relevant bond dissociation energy increases. To explain these results we assume that the repulsion curves are different for different atoms (X, Y , etc.) transferred in the reaction. For example, for R approaching HR* the repulsion curve should have the shape of curve I in Fig. 4, i.e. the repulsion force, being negligible for the first stages of reaction, increases suddenly when R is sufficiently close to HR*. In this case not much work is performed in overcoming the repulsion forces while the system arrives at point A corresponding to a relatively steep slope (i.e.E, is small). On the other hand, the repulsion curve representing the interaction of R and ClR* would have the shape of curve I1 in Fig. 4. The repulsion force appears at a greater distance, but i t increases gradually, and in consequence, a considerable amount of work has to be done before the system arrives a t point B which corresponds to a comparatively flat slope (i.e. here E, is greater than in the previous case). The difference in the shapes of these two repulsion curves is plausible. One can say that the approach to a small but “compact” hydrogen demands a repulsion curve represented by I, while the approach to a bigger but more “ compressible ” chlorine would call for a repulsion curve repre- sented by 11.Although this representation is crude and highly ap- proximate it seems to be reasonable. At any rate, it gives us a glimpse into the problem of repulsion forces, which are of great importance for the understanding of chemical kinetics. Let us turn now to the problem of “ temperature independent ” fac- tors of radical reactions. This problem has recently caused some dis- cussion, and it seems profitable at this juncture to survey the questionREACTIVITY OF RADICALS again. A paper published by Steacie, Danvent and Trost 20 initiated a discussion on the magnitude of steric factors in reactions involving radicals or atoms. These authors tentatively suggested that the steric factors for the metathetic reactions of the type RH + X + R + HX, where X denotes a hydrogen atom or a methyl radical, are very small, e.g.a steric factor of the order of I O - ~ was suggested for the reaction The idea of low steric factors was taken up by Bamford and Dewar who stated : ‘‘ It seems reasonable to infer that organic radical reactions with activation energy less than 20 kcal./moIe in geneVal have frequency factors of the order of 107 and not 10l2 * as is commonly assumed.” C,H,o + €3 --f CIH, + Ha. FIG. 4 It seems that the above papers created an impression that a new approach is required in dealing with the rates of reactions involving radicals or atoms. It might appear that these reactions cannot be treated by the usual transition state method, and that this theory should be abandoned, or at least modified, if it is to cope successfully with radical reactions.The paper by Evans and Szwarc l7 attempted to point out that there is no justification for such an extreme conclusion. The review of the existing experimental data, collected in Tables I and IA of the latter paper, indicates that the steric factors in reactions involving atonzs are of the order 0.1, to within a factor of 10. It was pointed out that the transition state theory leads to lower steric factors if a reaction involves radicals. For example, a steric factor of the order I O - ~ (within one power of 10) would be expected for a reaction involving the comparatively small CH, radical (see p. 945 of their paper). For reactions involving more complex radicals, still lower steric factors are to be expected, examples being given by the recombination of triphenyl methyl radicals, the growth of a polymeric chain, or chain transfer reactions.In all these cases three degrees of rotational freedom are lost in the formation of the transition 20 Steacie, Darwent and Trost, Faraday SOC. Discussions, 1947, 2, 80. Bamford and Dewar, Nature, 1949, 163, 256. * These numbers refer to cm., mole. -l sec. -l units.M. SZWARC I 5 3 state complex, and in terms of this loss of rotational freedom steric factors as low as I O - ~ have been described and understood. In consequence, the paper concludes that “ . . . there is no justification for assuming abnorwuZZy low steric factors for such (radical) reactions. Small steric factors will occur in reactions involving two complex radicals or mole- cules.This is not new, and has frequently b2en discussed in terms of the loss of rotational freedom of the reacting centres.” The paper by Szwarc and Roberts 22 re-emphasized the conclusion of Evans and Szwarc. The former authors discussed the reaction CH, + CIH, . CH, --f CH, + C,H, . CH, and deduced that the magnitude of the steric factor in this reaction does not contradict the expectations of the transition state theory, i.e. the frequency of the effective collisions (the temperature independent factor, or PZ factor) for this reaction is of the order 10ll-10~~ cm., sec.-l mole-1 (see summary of their paper, p. 625). In calculating the magnitude of stsic factors one has to assume the magnitude of collision encounters (usually denoted by the letter 2) which, in turn, depends on the collision diameter.Evans and Szwarc pointed out, however, that there are some uncertainties in estimating the latter entity, since it depends on the type of phenomenon which is in- vestigated and this point has been discussed further in the present paper. Obviously, if a large collision diameter is chosen, one must derive a low steric factor. For example, Gomer and Dorfman 28 used the data of Szwarc and Roberts 22 for the reaction CH, + toluene, and assuming a collision diameter as large as 10 derived a steric factor as low as 7 x I O - ~ . The magnitude which is measured directly and which is indeed essential for reaction kinetics is the PZ factor, i.e. the temperature independent factor (or the effective collision frequency). It would be better to con- centrate attention on this entity, and so avoid the apparent controversy due to usage of various values for the collision diameter.The transition state theory enables us to calculate approximately the PZ factors. For example, for a reaction of the type RH + CH, - + R + CH, the calculation would lead to PZ of cu. 10ll-10~~ (within a power of 10). It is remarkable that the experimental results of various workers give this order of magnitude (see Table 11). TABLE 11* I I Reaction Investigators A = P Z cm.S/moles Sec. I I CH, + CH,. CO . CH, CH, + CH,. CO. CH, CH, + CH, . C,H, CH, + CH,. C,H, Trotman-Dickenson z4 and Steacie Jaquiss, Roberts 24 and Szwarc Trotman-Dickenson 24 and Steacie Szwarc and Roberts 22 5’7 x I011 3’3 x I011 1.6 x 10ll 6.4 x 1o12 Szwarc and Roberts, Trans. Faraday SOC., 1950, 46, 625. 23 Gomer and Dorfman, J . Chem. Physics, 1951, 19, 136. *The values for A listed in this Table were computed from the formulae taking for AC~H,, the value 7 x 1013 ~m.~/rnoles sec. as reported by G ~ m e r , ~ who assumed E . c ~ H ~ = 0.I 5 3 TRICHLOROMETHYL RADICALS We conclude, therefore, that all the experimental evidence reported recently for reactions of the type CH, + RH -+ CH, + R seems to be in concord with the requirements of the transition state theory. No new hypothesis seems to be required to account for the rates of these radical reactions. Of course, the PZ factors need not be constant, and the work of Trotman-Dickenson and Steacie 24 is of the greatest value in showing how the change of the reacting species affects the numerical value of these factors. It is important, however, that their order of magnitude remains within the range predicted by the transition state theory. It is a pleasure to express my thanks to Prof. M. G. Evans, F.R.S., for very helpful and stimulating discussions. Chemistry Department, Manchester University, Manchester, 13. z4 Trotman-Diekenson and Steacie, J . Amer. Chem. Soc., 1950, 72, 2310 ; 45 Jaquiss, Roberts and Szware, unpublished data. J . Chem. Physics, 1950, 18, 1097; 1951, 19:

ISSN:0366-9033    

DOI:10.1039/DF9511000143

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

I 5 3 TRICHLOROMETHYL RADICALS THE KINETICS OF THE INTERACTION OF TRICHLOROMETHYL RADICALS WITH CYCLOHEXENE BY H, W, MELVILLE, J. C. ROBB AND R. C. TUTTON Received 2nd February, 1951 This paper outlines the technique adopted in obtaining kinetic data for the photochemically induced reaction between trichlorobromo methane and cyclo- hexene with the formation of I-trichloro-methyl 2-bromo-cyclohexene. The standard methods which have been previously developed in the investigations of high polymer reactions have been modified to suit this reaction. The reaction exhibits chain properties and will be specially useful in measuring the absolute reactivities of olefinic or unsaturated compounds in general towards a specific and unambiguous free radical, CCl,. It also gives valuable information concerning reactions of the type responsible for the phenomena of transfer in the polymerization of vinyl compounds in a way that is not possible by other means .Preliminary results indicate that the reaction CCl, + C6H,, proceeds with a velocity constant k , = 2-56 x loa 1. moles-l sec.-l a t 30' C and a termination velocity constant zk4 = 1.0 x 108 1. moles-lsec.-I, being that for z CC1, -+ C,Cl,. The overall energy of activation is 3-4 kcal./mole and it is indicated that the energy of activation for termination of the chains is also low. Quantitative radical chemistry has now progressed to such a stage that not only is the mechanism of certain reactions thoroughly known but the velocity coefficients of the constituent radical reactions have been determined with reasonable accuracy.In order to build up a general picture of radical reactivity both towards molecules and radicals it is essential to explore the possibilities in reactions other than those involving polymerization and oxidation. One type of reaction which appears to be suitable is the addition of simple organic halogen compounds to ethylenic molecules. Kharasch and his co-workers have investigatedH. W. MELVILLE, J. C. ROBB AND R. C. TUTTON 155 the chemistry of a number of reactions of this kind and have shown that they take place to form well-defined products, that they can be initiated easily photochemically and proceed by a radical chain mechanism.1 A typical example is the addition of trichlorobromomethane to cyclohexene. The terminal product has the structure CH2-CH2-CHCCI I I CH2-CH2-CHBr.It is suggested that initial fission of the halomethane is into Br + CCl,, that the trichloromethyl radicals attack cyclohexene, and the radical then formed abstracts a bromine atom from another molecule of bromo- trichloromethane. The CCl, radical starts the process and the cycle comes to an end by the interactions of the hydrocarbon and trichloromethyl radicals. It should be possible by so arranging the conditions of re- actions that the two rate determining steps would be CCl, + cyclohexene and CCl, + CCl,. A comprehensive investigation of these reactions is of importance for many reasons. Radical double bond interactions are the rate deter- mining step in addition polymerization. By experiments on copolymer- ization systems it is possible to measure such interactions in which the attacking radical is kept fixed in structure and its absolute reactivity towards a whole series of monomers investigated.Br . CCl, is further a good transfer agent in polymerization and the rate determining step is the debromination by a polymer radical. Nothing, however, is known of the rate of addition of the CCl, radical to start the growth of a new polymer radical. Recent experiments in the addition of atomic hydrogen to olefines have thrown some light on the factors that affect chemical reactivity in this system.2 It is important to explore other systems in which the nature of the attacking radical is varied. Radical-radical reactions are of great interest since although in general little energy of activation is needed in these encounters, quite highly specific factors enter into such processes.Further data are required to see how structure affects the interaction coefficients in this kind of reaction. The present paper is specifically concerned with the elucidation of the precise mechanism of the addition of trichlorobromomethane to cyclo- hexene initiated photochemically, together with the determination of the velocity coefficients of the constituent radical reactions involved. Towards the end of this investigation the authors became aware of a similar investigation being carried out on the interaction of CCl , radicals with methallyl acetate., Theoretical.-The most suitable method for producing CCl radicals in a controlled fashion is by the direct photochemical decomposition of trichlorobrornomethane CCI , .Br which absorbs light of wavelength around 3650 A, which has sufficient energy to rupture the C-Br bond, (I) CC1,Br + hv -+ CCI, + Br I Initiation The reaction mechanism is as follows : 1 Kharasch, et al., J . Amer. Chem. SOG., 1947. 69, 1105 ; J . Org. Chem., 1949, 14979s 239,537- Melville and Robb, Proc. Roy. SOC. A , 1949, 196, 445, 466, 479, 494 ; 1950, 202. 18T. 3 Matheson and Halter (private communication).TRICHLOROMETHYL RADICALS (4) 2 CCl, +- C,CI, 2 c> CCl, 3 \ / k5 I Termination k6 I By arranging the experimental conditions such that the unsaturate con- centration is low, reaction (2) can be made rate controlling, in which case a normal kinetic analysis leads to the expression If conditions are such that low CCI, .Br concentrations are employed and reaction ( 3 ) becomes rate determining, assuming long chains, By the application of standard techniques, developed to deal with polymerization and chain reactions, the kinetic constants in eqn. ( z ) , (3), (4) and ( 5 ) can be evaluated. The remaining constant for reaction (6) can be obtained by working in a region of intermediate concentration such that all three termination processes occur and, having previously determined the others, it is a relatively simple task to evaluate k6. The bromine atom produced in the initiation step is used up, presum- ably in addition to cyclohexene. If this radical is capable of propagating a chain, it means that two chains are started for each molecule of CCl, .Br decomposed. This does not affect the picture since the inhibitor tech- nique used later measures the number of chains starting, irrespective of the nature of the initiating radical atom. The technique adopted for following the progress of reaction was by dilatometry. This necessitates knowledge of the densities of initial and final components in the system, but in this case, direct measurement of the density of the product is somewhat difficult and liable to introduce errors of considerable magnitude. The method adopted was to make use of the fact that due to utilization of unsaturate, the reaction rate decreases with time. Let xco be the meniscus drop for roo yo reaction (at t = a), xt be the meniscus drop for a time t, Y o be the amount of unsaturate in the reaction vessel at time t = 0, Yt be the amount of unsaturate in the reaction vessel at time t.Then dxjdt = C’Yt, . . (3) where C’ is a constant. Now xm is proportional to Y o and xt is proportional to the amount of unsaturate utilized between t = o and t = t, therefore the amount of unsaturate remaining in the reaction system at time t is proportional to (xm - xt). where c is a proportionality constant. Substituting in eqn. ( 3 ) , dx/dt = G ( X ~ - ~ t ) , . * (4) - In (xa - xt) = ct + const. . - ( 5 ) Integrating eqn. (4), When t = 0, xt = o and so the constant of integration is - In xa, Hence orH. W. MELVILLE, J. C. ROBB AND R. C. TUTTON 157 By taking the values for xt at two times tl and t,, * (7) . xh - xtl = xin e-cti - xtle-&.The values of t,, t,, xtl, xh, are known and by plotting the function on the right-hand side of eqn. (7) against arbitrarily chosen values of C, a curve is obtained from which the value of c, which gives the right hand side function equal to (xtl - xt,) is obtained. This value for G is substituted back into eqn. (6) and a value for xco is obtained. This gives the relation- ship between extent of reaction and percentage contraction of the re- acting mixture. The kinetic constants are determined as follows. Measurement of the overall rate gives the value of R2(I/2k4)* and the temperature coefficient of the overall rate gives the overall activation energy, where Measurement of the rate of chain starting I is accomplished by means of an inhibitor technique, described more fully below, thus giving a value for k2(2k4)-*.By means of experiments done with a rotating sector producing intermittent illumination, the value of 2k4 is obtained, thus leading to the individual values of k2 and k4. This investigation has been restricted to conditions where reaction (2) is rate determining. Experimental Apparatus and ~ateria~S.-TRICHLOROBROMOMETHANE was prepared by allowing anhydrous aluminium bromide to stand in contact with carbon tetra- chloride a t room temperature for 3 days.4 The resulting mixture was filtered and the filtrate washed with aqueous potassium carbonate and with water, dried in the usual manner and fractionally distilled under reduced pressure, the fraction boiling a t 49'9O C a t I I G mm. pressure being collected and stored in a dark bottle.CYCLOHEXENE was obtained commercially and purified by distillation, the fraction boiling a t 33" C a t 80 mm. piessure being collected. It was found to be necessary to treat the cyclohexene before distillation by refluxing with copper stearate.5 This treatment removes the peroxide formed on standing in contact with air. Failure to remove this peroxide leads to the development of pure thermal reactions, which complicate the kinetic analysis. TETRACHLORO-0-BENZOQUINONE was prepared by passing chlorine into a hot solution of catechol in acetic acid.6 After two final crystallizations from anhydrous di-isopropyl ether, the dark red crystalline solid had a melting point of 130' C . The size of dilatometer used throughout these investigations had a volume of about 10 ml.and a capillary cross-section of about 0-02 sq. cm. Under the conditions of illumination employed, these dimensions made i t possible to follow the meniscus fall readily with the aid of a cathetometer and to measure ac- curately the extent of reaction. The dilatometer was mounted rigidly and reproducibly in a thermostat controlled by a conventional mercury-toluene regulator to better than f IO-, "C. The source of active radiation was a 125-W Osira lamp, the radiation being filtered through 5 mm. soft glass and about 15 cm. water. In the sector ex- periments, the radiation also passed through 5 mm. of Perspex and a soft glass lens system, The complete thermostat was enclosed in a light tight box, fitted with suitable apertures for irradiation of the reaction vessel and for observation of the meniscus level.The sector used in the intermittent illuminaticn experiments was constructed by blackening a 180' sector of a 16-in. Perspex wheel driven by a D.C. motor, either direct or via a 15/1 reduction gear box. The sector chopped the light The density of the product was d425'C = 2.005 g./ml. JVesper and Rollefson, J . Amer. Chem. Soc.. 1934, 56, 1456. George and Robertson, Trans. Faraday SOL, 1946, 42, 217. Zincke, Ber., 1887, 20, 1776.TRICHLOROMETHYL RADICALS beam a t the focus produced by a suitable lens and could be used to give flash times of from 1/20 sec. up to about 3 sec. For longer flash times, a manually operated shutter was empIoyed. All experiments were carried out in absence of air, since oxygen was found to retard the reaction markedly.The techniques for the vacuum manipulation of materials in this kind of investigation are well established and need no further explanation. The outgassing of the liquids used was achieved by repeated freezing and pumping until no more non-condensible gas was evolved. Results Extinction coefficient for CCl, . Br.-This has been experimentally determined and is shown as a function of wavelength in Fig. I. Since the light used was largely a t A = 3650 A, the molar extinction coefficient is I = 5-0 x 10'~. The amount of light absorbed by the reaction vessel containing pure CC1,. Br, of thickness about I cm., is therefore about 6 yo of that entering the system, so that the photochemical initiation takes place uniformly throughout the reaction vessel.0 -1 '% -2 3700 3{GOd 3400 3600 3800 FIG. I .-Extinction coefficient for CCl,. Br. I = I, IOUCd ; E = extension coefficient, c = concentration in moles/litre, d = thickness of sample in cm. Dependence on light intensity.-Table I shows the constancy of the intensity exponent n obtained by using an intensity screen having a transmission factor of 0.418. The constancy of the exponent indicates that termination throughout the range of experimental conditions is accurately bimolecular. TABLE I Temp. O C CCla . Br/CqH,O molar ratio Intensity I Exponent I O / I I O / I I5/I 2O/I I O / I * I O / I 0.48 0.53 0.50 0.52 0.49 0.52 * This exponent was measured after inhibition for IOO min. by tetrachloro-o-benzoquinone.H. W.MELVILLE, J. C. ROBB AND R. C. TUTTON 159 Rate dependence on cyclohexene concentration.-Table I1 shows how the rate is accurately proportional to the concentration of cyclohexene, thus establishing the proposed kinetic analysis. TABLE I1 Temp. (o c) Temp. (" C) 25 Rate of removd of cyclohexene moles 1.-1 Sec-1 30'7 40 *The standar Relative amount of cyclohexene* 1 0.76 0.52 0.76 0.55 I Relative rate 1 0.75 0-52 I 0.78 0.54 concentration emdoved was I , CC1,. Br/C,H,, = IO/I molar. Measurement of absolute rate .-A typical experimental run is shown in Fig. 2 in which the fall of the meniscus is plotted against time. It will be seen that the rate falls off with time as the cyclohexene is used up. By applying the mathematical treatment outlined above in order to obtain the rate of re- action in terms of the amount of cyclohexene used up per unit time, the curve FIG.2.-Fall of meniscus as a function of time. shown in Fig. 3 is obtained, being a plot of the function xt, . e-ell - xtl. e-cl, against c. The value of xtt - xtl chosen corresponded to 3-535 and for a value of c = 0.0053, the above function has this value. Substitution of this value for c in eqn. (6) enables a value for x, to be obtained from which the curve in Fig. 2 is drawn. The curve fits the experimental points well. In this way, the rate of removal of cyclohexene can be computed and for three temperatures, the rate is shown in Table 111. In the usual manner, by plotting log rate against r / T , the overall activation energy is obtained. This plot is shown in Fig.4, The overall energy of activation measured from the slope is 3-4 kcal. /mole. TABLE I11 I 1-66 x I O ~ 1-84 x 10-6 49'2 2'00 x 10-6I 60 TRICHLOROMETHYL RADICALS Measurement of rate of initiation.-It has been shown that the com- pound tetrachloro-o-benzoquinone is a particularly efficient inhibitor for re- actions involving trichloromethyl radicals. It has, unfortunately, a high extinction coefficient, the absorption spectrum being shown in Fig. 5. C -004 -006 - 008 -3-49 I I I t FIG. 3.-Curve for obtaining absolute rate of reaction. f = xb e-di - xtl e-clo. 31 32 33 I FIG. 4.-Dependance of overall reaction rate on temperature. In Fig. 6 are shown two typical inhibited runs with an uninhibited one f a comparison. It will be seen that there is a retardation of the subsequent rate by the products of the inhibited period of the reaction and i t is found that the degree of retardation does in fact increase with increasing amounts of inhibitor.On account of the fairly strong absorption by inhibitor, measurement of the overall duration of the inhibition period would not be a true measure of the rate of initiation. The technique adopted was to use low concentrations of inhibitor such that its absorption of 3650 A is weak, and t o measure the rate of inhibitor removal photometrically. This was done by removing the dilatometer from the thermostat at various intervals and placing i t in the cell-holder of a Unicam Diffraction Grating Visible Spectrometer. The wavelength used wasH. W. MELVILLE, J. C. ROBB AND R.C. TUTTON 161 5000 A and the instrument was calibrated by using solutions containing known amounts of tetrachloro-o-benzoquinone. Two such runs are shown in Fig. 7. Removal is quite linear over a considerable period near the end of the induction period and kom this slope, the rate of initiation can be calculated. In this way, the rate of initiation is found to be 1-17 x 10-6 mole 1.-l sec.-1 and is sensibly inde- pendent of temperature over the range of temperatures required. The overall rate of reaction is k I& given by 2 (C,,Hlo). Knowing the overall rate I , and (C,,Hlo), the ratio k,/(2k4)4 can be obtained and is given in Table IV. To obtain the absolute values of k2 and k, it is necessary t o measure another characteristic property of the reaction, namely the lifetime of the radical involved in the termination step.(2k4P TABLE IV 30 40 50 0.0256 0.0308 0.0365 20 I 40 I 3600 4400 5200 6000 6600 I 1 FIG. 5.-Extinction coefficient for tetrachloro- o-benzoquinone as a function of wavelength. 200 & FIG. 6.-Inhibition by tetrachloro-o-benzoquinone concentration of inhibition. A : 5.5 x I O - ~ moles/l. B : 8.0 x 10-8 moles/l. Measurement of lifetime of the CCl, radical.-This was done using the For these experiments, the rate of Fig. 8 shows the curve Burnett and now wellestablished sector technique.? initiation was reduced t o 7-16 x 10-9 mole 1.-1 sec.-l. Melville, Proc. Roy. SOC. A , 1947, 189, 456. Chapman, Briers and Walters, J. Chem. SOG., 1926, 562. RI 62 TRICHLOROMETHYL RADICALS obtained by plotting log I&t against R,/R,, where t is the duration of the flash, R, the sectored rate and R, the unsectored.The value of log (2k4)* is obtained by measuring the displacement of this experimental curve from the theoretical.* The curve drawn in Fig. 8 is that for a value of 2kq = 1.0 x 108 1. mole-l sec.-l. Furthermore, under these conditions, the lifetime of the radical is found to be t = 1-18 sec. It will be Seen that a t the two temperatures a t which sector experiments have been carried out, there is no significant change in the position of the curve. This means that only a low energy of activation can be ascribed to the termination process. Further work is proceeding in an attempt to obtain a value for the termination activation energy. FIG. 7.-Inhibition concentration as a function of time.FIG. &--Effect of intermittent illumination on the rate of reaction. Conclusion.-The kinetic values obtained are tabulated in Table V. TABLE V 2-56 x 102 1. mole-l sec.-l 3-08 x 102 ), ), 1'0 x I08 ) J J , 3.4 kcal. /mole low 6-96 x 105 -1.0 x 108- 3.65 x Io2 JB > J 8 Dickinson, Photochemistry of Gases, by W. A. Noyes and P. A. Leighton (Reinhold, New York, 1941)~ p. 207.H. W. MELVILLE, J. C. ROBB AND R . C. TUTTON 163 The low energy of activation for the addition of CCl, to cyclohexene is in general accord with the high reactivity of this radical as, for example, indicated by the fact that CC1, is an effective transfer agent in polymeriz- ation. The temperature independent factor is perhaps surprisingly low for a relatively simple addition of this kind.When, however, the attack- ing radical is a hydrogen atom the velocity coefficient increases to 6 x IO* 1.mole-l sec.-l at 1 5 O C.a The present measurements indicate a low energy of activation for termination but again the temperature in- dependent factor is considerably less than the normal value as happens in polymerization and in interactions in oxidation processes. The general conclusion from the results is that the velocity coefficients of these radical are controlled to a large extent by factors which affect the temperature independent factors rather than the energies of activation. More refined experiments are necessary to measure the termination constants with sufficient precision while it is hoped, by conducting similar experiments with IO/I molar excess of cyclohexene, to produce further constants for reactions (3) and (5). By working in the intermediate regions, the values for reaction (6) should be available. This paper serves to outline, very briefly, how study of systems such as that already described, yields information which can be obtained in no other way concerning the absolute reactivities of unsaturated com- pounds towards a specific radical. One of us (J. C . R.) is indebted to the Department of Scientific and Industrial Research for a Senior Award (1948-50), during the latter part of the tenure of which this work was carried out, and to the University of Birmingham for the award of an I.C.I. Fellowship (1950-51). Another of us (R. C. T.) is indebted to the Department of Scientific and Industrial Research for a maintenance allowance. We are also indebted to Messrs. Imperial Chemical Industries for the loan of the spectrophotometer. Chemistry Department, The University, Edgbaston, Birmingham, I 5.

ISSN:0366-9033    

DOI:10.1039/DF9511000154

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

H. W. MELVILLE, J. C. ROBB AND R . C. TUTTON 163 STRUCTURAL FACTORS GOVERNING THE REACTIVITIES OF a -METHYLENIC GROUPS TOWARDS ACTIVE FREE RADICALS BY E. C. KOOYMAN Received 5th February, 195 I Twenty-two hydrocarbons of varying structures are compared with regard to reactivity towards trichloromethyl radicals, using their retardation of carbon tetrachloride addition to cetene as a basis. Reactivities are expressed as a semiempirical constant, which shows correlation with other numerical con- stants bearing on hydrogen abstraction reactions at a-methylenic groups. Structural influences on a-methylenic reactivities are discussed on the basis of degree and type of substitution a t the reactive group. The higher reactivity of compounds in which the a-methylenic group forms part of a ring system as compared with analogous open-chain compounds is tentatively attributed to an entropy effect.The results tend to suggest that hydrogen abstraction by free radicals involving hydrocarbons is largely governed by the heats of reaction.REACTIVITIES OF a-METHYLENIC GROUPS a-Methylenic groups * are often preferential points of free-radical attack resulting in hydrogen abstraction, the removal of these atoms being facilitated by the low dissociation energies of the corresponding carbon-hydrogen bonds ; these are associated with the appreciable reson- ance energies of the allyl- or benzyl-type radicals formed. These reactions are of primary importance in the autoxidation of olefins, of rubber and the fatty oils, and of alkylated aromatics.Further typical examples are the bromination by means of N-bromosuccinimide, as discovered by Ziegler, as well as dehydrogenation reactions occurring under the influence of sulphur, selenium and quinones.1 Recently, a hydrogen abstraction reaction of the above type was found to be responsible for chain termination in the benzoyl peroxide initiated addition of carbon tetrachloride to olefins : X’ + R-CH,-CHzCH, __t XH + R-CH-CH=CH,. . (I) I In this equation X’ denotes an active radical, i.e. a radical capable of continuing the chain reaction, e.g. by abstracting a chlorine atom from carbon tetrachloride to produce a new CC1, radical. The stable allyl-type radicals formed in (I) fail to do so ; in experiments with cyclohexene as the olefin, they were detected in the form of their dimer 3 : 3’-dicyclo-hexenyl.I n line with this, foreign compounds capable of giving stable radicals by reactions analogous to (I) were found to exert a marked retarding influence on the carbon tetrachloride addition reaction. The present paper describes a simple method for the quantitative comparison of re- activities of this type ; it is based on the magnitude of the above retarding effects. the propagation sequence of the carbon tetrachloride addition to olefins may be represented in a manner analogous to that obtaining for the “ab- normal ’’ addition of hydrogen bromide to olefins : CCl,’ + R-CH,-CH=CH, --f R-CH,-CH-CH,-CCl, . (2) (trichloroalkyl radical) I Outline of method.-According to Kharasch, Urry and Jensen R-CH,-CH-CH,-CCI , + CCl, + R-CH,-CHC1-CH,-CC1 , + CCl ,’.(3) I It might be expected, therefore, that the chain-carrying trichloromethyl radical and the trichloroalkyl radical would play a predominant part in the termination step (I), forming chloroform and a trichloroalkane respectively. However, detailed analyses 2h 9 of the products formed when reacting a * a-Methylenic is employed here to indicate the positions adjacent to one or two double bonds or aromatic nuclei ; i t may apply to substituted or unsubstituted methyl or methylene groups (-CH,) having at least one hydrogen atom. 1 E.g. Waters in The Chemistry of Free Radicals, 2nd edn. (Clarendon Press, Oxford, 1948), pp. 168, 241-3, 255-7. (a) Kooyman, Rec. t m v . chim., 1950, 69,492 ; (b) Kooyman and Farenhorst, Rec. trav. chim., 1951 (in press).Kharasch et al., J . Amer. Chem. Soc., 1947, 69, 1100. p It should be noted that the above analyses were only possible because of the relatively short overall kinetic chain lengths of the carbon tetrachloride addition. With normal a-olefins, about z yo of peroxide (based on the olefin) is required 2 * to give complete conversion of the olefin. However, 3 : 3-dimethylbutene-I, CH, I H,G-C-CH==CH, LH3 which does not contain any or-methylenic hydrogen atoms, is completely con- verted under the influence of as little as 0.15 yo. Here, the termination step must be of a differentE. C. KOOYMAN 165 threefold excess of carbon tetrachloride with olefin in the presence of benzoyl peroxide revealed the formation of nearly two molecules of chloroform per molecule of peroxide decomposed, whereas the amounts of trichloroalkanes were relatively unimportant (a few per cent based on peroxide).The other hydrides XH expected on the basis of eqn. (I) such as benzene and benzoic acid, which are formed from the phenyl and benzoate radicals arising from the peroxide, were also present in minor amounts. From the above data it was concluded that the trichloromethyl radicals furnish the main contribution to chain termination. Apparently, re- action ( 3 ) proceeds quite readily under the reaction conditions employed, giving the trichloroalkyl radicals little opportunity for reaction according to (I). MILLMOLES OF A ao 6 0 - 5 0 - 4 0 - SLOPE 55‘4 (BELOW 75% CONVERSION) MEAN DEVIATION 0.5 MlLLlMOLES OF BENZOYL PmOXlOE DECOMPOSED FIG.ca carbon tetrachloride bound to cetene (A) as a function of benzoyl (Initial molar ratio CCl,!cetene 3-00 ; temp. g ~ + &+ O C.) The overall kinetics of the system carbon tetrachloride + cetene (molar ratio 3.0) + benzoyl peroxide were investigated a t 918 f *‘C.* A linear correlation was found between the amount A of carbon tetra- chloride bound to olefin and the amount of peroxide decomposed AP, irrespective of the intake of peroxide : This empirical relation holds up to olefin conversions as high as 80 yo ; it may be readily explained when assuming (2) to be irreversible. In that case, all trichloroalkyl radicals eventually give products containing one molecule of carbon tetrachloride; the very small amounts of tri- chloroalkanes (below 5 yo based on peroxide, or less than 0.1 yo based on A ) may be neglected.This argument is not affected by telomerization reactions (addition of trichloroalkyl radicals to olefin rather than reaction with CCl,) as the telomers also contain one molecule of CC1,. peroxide decomposed. A P = k A . . - (4 Combination of (I) and ( 2 ) gives In this equation T designates the trichloromethyl radicals eliminated by * This temperature was chosen because it is easily maintained in the course of experiments with molar ratio 3.0, the heat of reaction being carried off by the gzntly refluxing carbon tetrachloride.I 66 REACTIVITIES OF a-METHYLENIC GROUPS ( I ) , Kt and K, are the rate constants for (I) and (2), respectively ; 0 repre- sents the olefin. All quantities are expressed in moles or millimoles.When only the CC1, radicals are operative in (I), (b) may be integrated to give (u). However, K (eqn. ( a ) ) was found slightly to increase (from 0.0180 to 0.0196) when decreasing the initial molar ratio of halide and cetene to 2 . 0 ; it showed a decrease to 0.0168 when using the ratio 4-2. This suggests that K contains small contributions from other components besides K (eqn. (b)) ; for example, the trichloroalkyl radicals will give increasing contributions to termination as the excess of carbon tetrachloride decreases. Assuming only the trichloromethyl radicals to be operative in chain termination reactions with retarders, the influence of retarders may be expressed by the following equations : or AP' dA - = K , J (0). . R Here, T' denotes the trichloromethyl radicals eliminated by the retarder. The amount of retarder R was much larger than that of the peroxide decomposed and was therefore taken as a constant.AP' represents the amount of peroxide consumed by the retarder ; in connection with (b), it equals the difference between the amounts of peroxide decomposed and the amount of peroxide which would have sufficed to effect the same extent of addition A in the absence of retarder. K, constitutes our measure of or-methylenic reactivity, k , applying to the same olefin (cetene) in all experiments. 1% may be determined from the experimental plot of 1/(0) as a function of A ; K, constitutes the slope of the plot AP'/(R) against 18. Both the amount of peroxide and the amount of retarder were varied in order to obtain a number of points throughout the range of olefin conversions from 10 to 60 yo M.In using eqn. (d), the possible influence of the reaction products of the allyl-type radicals formed in (I), e.g. dimers or disproportionation products, had to be neglected. However, they are only forming a small proportion of the total olefin present, except at very high conversions of olefin. Experimental Materials .-Analytical grade carbon tetrachloride was employed in all experiments. Benzoyl peroxide was purified by dissolving in chloroform and precipitating with methanol ; the purity of the samples thus obtained varied between 97.5 yo and 99.0 yo. The preparations and properties of the hydro- carbons employed are indicated on opposite page. Procedure.*-(I) The determination of A may be illustrated by the following example.25 ml. of carbon tetrachloride (259 mmoles) and 25 ml. of cetene (87 mmoles, 19-50 8.) were pipetted into a flask containing 250 mg. (97-5 % pure) of benzoyl peroxide (I-01 mmoles) and a boiling chip. The mixture was rapidly heated to 91' C and then placed in an oil bath maintained between 93' and 9 4 O C. In this manner, the reaction temperature was kept a t g1-92OC for two hours. Solid carbon dioxide was periodically added in order to avoid the entrance of air. The mixture was then cooled to room temperature and made up to IOO ml. in a volumetric flask by means of the required amount of carbon tetrachloride. 5 ml. of this solution contained 4-50 mg. of peroxide, indicating that 1534 mg. (0.634 mmole) of peroxide had been decomposed.The remaining solution was then evaporated in a weighed flask at reduced pressure below 70°C ; the last traces of carbon tetrachloride were removed by increasing the temperature and * Cf. ref. (zb). The carbon tetrachloride refluxed very slowly.E. C. KOOYMAN 167 PROPERTIES AND PREPARATIONS OF HYDROCARBONS EMPLOYED Compound n-Decane . . . iso-Octane (2 : 2 : 4-tri- methylpentane . Decalin (cis) . . Decalin(truns) . . Benzene . . . Toluene . . . p-Xylene . . . m-Xylene . . . Ethylbenzene . . Cumene . . . Mesitylene . . tent.-Butyl benzene . Allylbenzene . . Tetralin . . . Diphenylmethane . Dibenzyl . . . meso-2 : 3-Diphenyl- Triphenylmethane . butane Cetene n-Heptene-3 . . 2 : 2 : 5 : 5-Tetramethyl- hexene-3 Cyclohexene .. Fluorene . . . g : lo-Dihydroanthra- cene . . . Physical Constants t 1.p. (" C/mm. Hg) Found 173 99 I93 185 80 I11 138 138 I35 152-3 x63-4 I 68 158-60 IOo/ZO II3lI - - - 170137 96.0 122 83 - - Lit. I74 99 I93 185 80 138 I39 136 152-3 165 169 I11 158-60 207 261-2 - - - 155115 96.0 125 83 I - Found 1.4121 1.3916 1.4828 1.4700 1.5010 1.4966 1.4956 1'4970 1.4959 1.4913 1.4990 2.4927 1.5136 1.5464 - - - - 1.4411 1'4050 1'4113 1'4734 - - Lit. 1'41203 1.3916 1.4828 1.4702 1~50110 1.49682 1.49580 1.49715 1-49580 1.4913 1.4990 1'4925 1.5143 1.5462 - - - - 1.4411 1.4049 1'4115 1.4727 - M.p. (" C) Founc Preparation Commercial sample * * 1, >, Fractionation of a commer- cial sample through a 30 plate column * Commercial sample * * * * * I , s, IS 22 I 1 1 , 25 ,, I , I , Org.Synth., 2, 341 * Standard procedures Commercial sample * Meyer and Wurster (Ber., Standard procedures $ Cf. Siside and NozakiJ J . Amer. Chem. SOC., 1948, Org. Synth., 1, 548 $ * 1, 2 ) 187336,964 $ 70,776 Standard procedures Howard et al., J . Res. Nal. Bur. Stand., 1947, 38, 1 2 2, 365 § Standard procedures 3 Commercial sample $ Wieland, Ber., xgm, 45,492 * Repeatedly shaken with 96 "I,, sulphuric acid until the acid layer remained nearly colourless, washed t Most of the literature data are from the Handbook of Ckmistry and Physics, 31st edn. ; those for the $Recrystallized from absolute ethanol. by further reduction of pressure. Finally, the residue was heated for 10 min. a t I I O - I I ~ O C a t 0-1-1 mm. Hg ; its weight then amounted t o 23.83 g.; this was not changed by another 5 min. heating. In calculating A a correction was used for volatile peroxide decomposition products amounting to 50 % of the peroxide converted.2b Thus the original reaction product contained (23.83 x 20/1g - 19-50 - 0.09 - 0.08 =) 5-42 g. of carbon tetrachloride bound to olefin, or 35-2 mmole. Both this procedure and ( 2 ) were checked by experiments without peroxide ; under the experimental conditions used, all the unreacted cetene remained behind. Chlorine deter- minations in the residues tallied very satisfactorily with the results obtained by direct weighing. ( 2 ) Determinations of A in the experiments with retarders were made in the same way, taking account of the volatility of the retarders. The latter were mostly either sufficiently volatile to be removed along with carbon tetrachloride, or remained behind.(Tetralin gave some difficulties owing to its intermediate with water, dried with calcium chloride and distilled over sodium. alkylated benzenes were taken from Francis, Chem. Rev., 1948, 42, 126. Distilled over sodium.168 REACT IV IT IES O F a-METHY LEN IC GROUPS volatility.) Methods (I) and (2) required some practice ; in order to ensure reproducible results all manipulations had to be carried out in exactly the same way. Duplicate experiments yielded A values hardly ever differing by more than 0.5, A ranging from 8 to 45 in procedure (2) and from 5 to 75 in &IN MILLMOLES-I 0.110 0.100 0.090 I 0.080 a070 0.050 4040 0.030 a020 0.010 QOOO A (IN MILLIMOLES) FIG. z.-1/(0) as a function of A (0, = 87 millimoles) .procedure (I). (3) Determination of plot of 1/(0) as a function of A was made by determining the uncon- verted olefin (bromine addition method according to McIlhiney), taking aliquot samples of the original reaction products. This plot (Fig. 2) was then graphically integrated ; for convenience, the results were plotted as Jg = f ( A ) . Results The mean K , values obtained J% by plotting AP’/(R) against have been recorded in Tables I, I1 and 111. Figures placed immediately after the name of the retarder refer to the numbers of experiments made ; figures after f indicate the largest deviations from the mean values. For the unretarded addi- tion to cetene, the mean deviation from the mean value amounted to not more than I % (Fig.I). basis of eqn. ( d ) . Fig. 3 shows a number of plots obtained on the lo’. A PyR (IN MILLIUOLES ) 6.50 r 600- 550 - WO. - 4 5 6 - do0 - 3.50 - 300- 2.50 - 2.00 - ,1.50 - la0 - f* FIG. 3.-K, values from equation (d).E. C. KOOYMAN TABLE I.--K, VALUES OF COMPOUNDS HAVING NO WMETHYLENIC HYDROGEN ATOMS Compound n-Decane (2) . * . iso-Octane (I) . (2 : z : q-trimethylpentane) . Decalin (trans) (2) . Benzene (I) . ted-Butylbenzene (I) . z : z : 5 : 5-Tetramethylhexene-3 (2) Decalin (cis) (2) . . . Kr x 102 0'0 0'0 0.08 f 0.02 0.17 & 0.05 0.03 0'1 0'0 TABLE II.-INFLUENCE OF TYPES AND DEGREES OF SUBSTITUTION AT THE Q-METHYLENIC GROUPS Compound Benzene (I) . Toluene (4) . Ethylbenzene (4) . Cumene (3) . tert.-Butylbenzene (I) . p-Xylene (4) . m-Xylene (3) .Mesitylene (4) . Toluene (4) Kt x 102 0.03 0.42 f 0.01 1-28 f 0-ob 1-75 f 0'1 0'1 0.95 f 0-04 0.84 f 0.04 1-29 f 0.08 0.42 f 0.02 Compound Diphenylmethane (6) . Triphenylmethane * (5) Fluorene (4) . Cetene (35) . n-Heptene-3 (6) . Allylbenzene (6) . . Dibenzyl (5) . meso-z : 3-Diphenyl- butane (2) . KT x 102 3'35 f 0.2 7 z t I 4'9 0.3 I 2 f I 1-3 f 0.1 1-13 f 0.05 * This compound failed to give a satisfactory straight line. i This value was estimated as follows. Allylbenzene and carbon tetra- chloride (molar ratio 1/3) were reacted with varying amounts of benzoyl peroxide (temp.. g ~ + " C). The plot AP = f ( A ) showed a nearly straight line, some deviation towards lower A values being observed a t increasing conversions of allylbenzene. Assuming the ratio of the rates of addition of the CC1, radical (K, in eqn.( b ) ) to allylbenzene and to n-octene to be 0.7 (as roughly determined by Kharasch and Sage) and the k , for n-octene and cetene to be equal, this gives 0.7 X 0.17 = 12 X I O - ~ for K, (allylbenzene). This value obviously constitutes merely a very crude approximation. The slope amounted to 0.17. TABLE III.-INFLUENCE OF CYCLIC STRUCTURES ~~ ~~ Cyclic Compound K, x 102 Analogous Acyclic Compound Cyclohexene (8) . Tetralin (9) . g : 10-Dihydro anthracene (7) . n-Heptene-3 (4) . Ethylbenzene (4). Diphenyl- methane (6) . KT x 13 I Discussion The significance of K,.-Inasmuch as the plots obtained satisfy eqn. ( d ) the largest deviation from the mean slopes being about 5-6 yo, whereas the mean deviations are about z yo, both the experiments with and without retarder may be interpreted in terms of a single radical operative ICharasch and Sage, J. Org.Chem., 1949, 14, 537.1 70 REACTIVITIES OF a-METHYLENIC GROUPS both in termination and addition. No analytical evidence is available, however, that the trichloromethyl radicals alsc predominate in the experi- ments with retarders. In the above sense, our K , values constitute a measure of a-methylenic reactivity ; it is realized that small contributions from other than the trichloromethyl radicals are incorporated in K, in a manner not indicated by the experimental data, As only retardations are measured, reactions of active radicals not leading to termination, i.e. to stable radicals incapable of continuing the chains, escape observation.For example, paraffins might be attacked to a small extent, giving alkyl radicals, which would in turn attack carbon tetrachloride to regenerate trichloromethyl radicals. Triphenylmethane appeared to be the only compound showing a systematic deviation, viz. towards decreasing K, values at increasing 1% values. This is tentatively attributed to the special properties of the triphenylmethyl radical, the dimerization of which is known to proceed towards an equilibrium ; because of the very high steady state equilibrium concentration (" high " compared with the concentrations of the other stable radicals) chain-starting reactions by these radicals might occur. Structural factors affecting &.-The values of K, found for the different compounds investigated have been arranged in Tables I, I1 and I11 in order to show that ( a ) (cf.Table I) : little or no retardation occurs by compounds not containing a-methylenic hydrogen atoms, 2 : 3-diphenylbutane (Table 11) forming the only exception ; (b) (cf. Table 11) : a-methylenic reactivity is increased by. substitution at the a-methylenic groups, methyl, phenyl and vinyl having " loosening " efficiencies increasing in the order mentioned, whereas the first substituent has more influence than the second ; ( G ) (cf. Table 111) ; a-methylenic groups forming part of a six-mem- bered ring are more reactive than the corresponding open-chain compounds. Conclusion ( a ) is limited to the classes of hydrocarbons investigated, polycyclic aromatics having marked retarding effect^,^ which are probably due to the capturing of free radicals to form new, stable radicals, rather than to hydrogen abstractions.Thus, naphthalene was found to act as a retarder, whereas Kharasch and Dannley6 have shown both the I- and 2-naphthyl radical formed by decomposition of the corresponding dinaphthoyl peroxides to be quite reactive, abstracting a chlorine atom from CCI, with formation of the two corresponding chloronaphthalenes. The formation of naphthyl radicals therefore cannot lead to chain termination. It is interesting to note that the tertiary hydrogen atom in iso-octane appears to react quite slowly ; otherwise the formation of tertiary radicals would lead to chain termination ; the latter type of radicals has been found to be incapable of abstracting a chlorine atom from carbon tetra- chloride except when the free valence occurs at a bridgehead.7 The small difference found between the K , values of the two decalins is, admittedly, at the limits of our experimental accuracy. However, cis-decaljn has a somewhat higher heat of combustion (2-3 kcal./mole) and is more readily oxidized than the trans isomer ; this would suggest that the radicals formed from both isomers, when removing the hydrogen 5 Kooyman et al.(to be published shortly). 6 Kharasch and Dannley, J. Org. Chem., 1945, 10, 406. 7 Kharasch ef al., J . Amer. Chem. SOC., 1942, 64, 1621 ; 1943, 65, 2428. Criegee, Ber., 1944, 77, 22.E. C. KOOYMAN atom at the g position, are identical and that their formation is more facile with cis-decalin than with the trans compound, in agreement with our result.The very low value found for z : 3-diphenylbutane seems to form an exception, as this compound contains two a-methy'lenic hydrogen atoms. Moreover, its value could be expected to be higher than that found for dibenzyl. However, model considerations showed that the bulky methyl groups probably cause steric hindrance ; in line with this, dibxzyl is readily dehydrogenated by chloranil to give stilbcne, whereas 2 : 3-diphenylbutane fails to react.* With 2 : 2 : 5 : 5-tetramethyl hexene-3, one might have expected some influence of addition of free radicals to the double bond; here too, steric hindrance will prevent the addition. Moreover, this olefin even failed to add on bromotrichloromethane, a halide which is known9 to react quite readily with olefins having the double bond in non-terminal positions. The very low values found with benzene and with tert.-butylbenzene, besides showing the effect of absence of a-methylenic hydrogen atoms, also indicate that-if any reaction occurs at all-the products formed with trichloromethyl radicals are sufficiently active to continue the chain.With regard to ( b ) , the influences of substitution on reactivity as expressed by this rule would seem to be analogous to the influence on bond strengths. Few data are, however, available as regards the values of the dissociation energies of the C-H bonds in most of the retarders used. Table IV gives a number of dissociation energies recently pub- lished by Roberts and Skinner 1 0 which illustrates the influences of sub- stitution on C-H bond energies in simple hydrocarbons.TABLE IV.-DISSOCIATION ENERGIES OF C-H BONDS Compound I l- Do-€I kcal . I02 97'5 90.8 86-5 77'5 78 I02 Difference with Respect to DC-= in Methane (kcal.) 0 4'5 15'5 24'5 24 11'2 0 Similarly, the dissociation constants of a number of extremely weak acids l1 show the same trends as the corresponding K , values (Table V). Apparently, it makes little difference that the pK values apply to the formation of a negative ion, whereas K, applies to the formation of a free radical. The methylated benzenes are seen to possess nearly equal reactivities when taking account of the number of reactive H atoms. Recent bond energy data obtained by Szwarc and by Szwarc and Sheon l2 are in line with this result (Table VI). These results suggest that little or no inter- action occurs between methyl groups in the meta isomers, whereas a * Private communication by Dr.N. Dost of this laboratory. lo Roberts and Skinner, Trans. Faraday Soc., 1949, 45, 339. l1 (a) Conant and Wheland, J . Amer. Chem. Soc., 1932, 54, 1212 ; l2 Szwarc, J . Chem. Physics, 1948, 16, 128; Kharasch, et al., J . Amer. Chem. SOC., 1947, 69, 1105. ( b ) Gilman, (c) McEwen, J . Amer. Chem. SOC., 1936, Szwarc and Shcon, J . Chem. J . Amer. Chem. SOC., 1938, 60, 2336; 58, 1124. Physics, 1950, 18, 237.172 REACTIVITIES OF a-METHYLENIC GROUPS loosening of hydrogen atoms in methyl groups occurs in the ortho and para isomers. Dibenzyl forms an interesting case, its K; being only 0.3, whereas ethylbenzene has 0.65.As pointed out by Szwarc lS both the shortening of the central C--C link (to 1-48 A) and thermal data indicate that this bond possesses some double bond character. The fact that the ct-methylenic hydrogen atoms are somewhat more tightly bound may be related to the higher Do-= in ethene as compared with that in ethane. The relatively small difference in K,.' observed between triphenyl- methane and dipheny'lmethane (7 and 1-68, respectively) suggests that the C-H bond energies are rather similar. This is in line with current conceptions that th2 high stability of the triphenylmethyl radical as compared with the diphenylmethyl and benzyl radical is largely due to steric factors, the resonance energy of the triphenylmethyl radical being not much higher than that of the others, as the three bulky phenyl groups cannot lie in one plane.14 TABLE V.-DISSOCIATION CONSTANTS AND ct-METHYLENIC REACTIVITIES (pK VALUES GIVEN BY CONANT AND WHELAND~~ COMPARED WITH K,) I I Fluorene .Triphenylmethane . Diphenylmethane . Cumene . 25 33 35 37 47 7 3'35 1-75 TABLE VI.-DISSOCIATION ENERGIES AND K, VALUES FOR METHYLATEII BENZENES Compound Toluene . m-Xylene . o-Xylene . p-Xy lene Mesitylene . Dissociation Energy of C H Bonds in Methyl Groups (kcal./mole) 77'3 77'1 74.8 76.2 - Kr ( x 102) 0.42 0.86 0.95 1'29 1.02 K,'(x 14) (= Kr per Reactive H Atom) 0.140 0.143 0.1 70 0.158 0.143 With regard to ( c ) , this higher reactivity of the cyclic hydrocarbons may be caused partly by a difference in the entropy of activation.Thus, Price and Hammett 16 interpreted the higher entropies of activation for the reaction between semicarbazide and cyclic ketones as compared with those for the acyclic ketones in the following manner. The cyclic ketones will lose fewer degrees of freedom upon entering the (rigid) transition state than the acyclic ketones ; consequently, the activation entropy will be more strongly positive for the former class of compounds. Whereas the analogous pairs tetralin + ethylbenzene and cyclo- hexene + n-heptene-3 have K,' values differing by a factor of about 24 (1.8 and 0.65; 2.8 and 1.2, resp.), the difference between dihydro- anthracene and diphenylmethane (81 and 1-68) is much larger. Dihydro- l3 Szwarc, Faraday SOC. Discussions, 1947, p. 39. l4 Faraday SOC.Discussions, 1947. pp. 70, 71. l6 Price and Hammett, J . Amer. Chem. SOC., 1941, 63, 2387 ; Hammett et al., 1943, 65, 1824.E. C. KOOYMAN I 7 3 anthracene is obviously an extremely “ rigid ” compound ; the removal of a hydrogen atom from the 9-position can hardly be expected to pro- duce a radical of much higher resonance energy than diphenylmethyl. The high reactivity of fluorene may be caused mainly by the low dis- sociation energy of the C-H bonds in the methylene group, the fluorenyl radical possessing an even higher resonance energy than the diphenyl- methyl radical. Probably, the rigidity of the fluorene molecule is of secondary importance. Correlations between the K, values and rate constants of other reactions involving a-methylenic groups .-CHAIN TRANSFER CONSTANTS IN THE POLYMERIZATION O F STYRENE.-Gregg and Mayo l6 recently described chain transfer experiments involving styrene ; several of the solvents used by these authors were also studied by us.Inasmuch as alkylated benzenes can also react by adding on a radical, which enters Kr x 10’ 105 x CHAIN TRANSFER CONSTANTS FIG. 4.--K, values as a function of chain transfer constants for styrene polymerization (temp. 914 C ) . into the chain transfer constant, but not into K,, we have corrected the values given by Gregg and Mayo for the small contributions arising from this type of reaction ; it was assumed that this contribution would be the same for all compounds containing a phenyl group. The resulting values were then recalculated for our reaction temperature 919°C.The values thus obtained were plotted against the corresponding K, values It is seen that both constants are very nearly on a straight line.* As the chain transfer constant represents the ratio between rate of transfer with solvent for a growing radical and rate of addition to styrene monomer, it is a measure of a-methylenic reactivity in those cases where other types of reactions are playing a minor part. (Fig. 4)- Gregg and Mayo, ref. (13), p. 328. * The values for fluorene (G. and M., about 1100 x I O - ~ , K, about 47 x I O - ~ ) do not fit into this picture, the K , value expected on the basis of a linear relation- ship being about 115 x I O - ~ .I 74 REACTIVITIES O F a-METHYLENIC GROUPS OXIDATION RATES.-An extensive series of oxidation experiments with The vital propaga- unsaturates was recently summarized by Bolland.’‘ tion step in these oxidations may be schematically represented by RO,’ + RH -+ R0,H + R’.* (4) The rate constant for these steps ( k , in Bolland’s nomenclature) was calculated for various types of olefinic compounds ; among other things, the results led the author t o the formulation of a number of rules con- cerning a-methylenic reactivity. These rules, as well as the general trends in the numerical values obtained, were quite similar to ours. Thus, a fair agreement is observed in the relative reactivities of the radicals CCl 3, substituted benzyl and RO, towards different a-methylenic groups. Whereas both CCl,’ and RO,’ are fairly reactive as well as rather electrophilic in character, the substituted benzyl radical is relatively stable ; its more rapid addition to monomers such as maleic anhydride as compared with styrene itself tends to suggest that it possesses a nucleo- philic character.On the other hand, heats of reaction have been shown in several in- stances to be closely related to the activation energies when comparing a series of related exothermic reactions. Thus, Butler and Polanyi suggested a linear relation between these two quantities ; in reactions involving hydrogen abstraction at a-methylenic groups, the differences in the heats of reaction are equal to the differences in the resonance energies of the stable radicals formed. From this point of view, the above relation- ships between the K,, values and the other rate constants are not sur- prising; they are the more significant inasmuch as they apply to con- stants of differing orders of magnitude. From the above, it may be inferred that polar properties are relatively unimportant in hydrogen abstraction reactions, the bond strengths of the C-H links involved being more decisive. This inference is admittedly based on a limited number of data ; no pertinent numerical data are available as regards the influence of strongly polar groups on a-methylenic reactivity. Anyway, it stands in contrast to the strong influence exerted by polarity factors on the addition of free radicals to double bonds. As the above-described retardation method is both a simple and a rapid one, it might serve as a yardstick for investigating a-methylenic reactivities even in mixtures and, probably, for compounds in which the “ loosening ” of the hydrogen atoms is caused by other groups, such as in ketones. An extension of the method to highly reactive retarders, which requires the more reactive system styrene + bromotrichloro- methane, will be reported on later.lS The author wishes to express his thanks to the Management of the N.V. de Bataafsche Petroleum Maatschappij for permission to publish this communication. Koninklijke /Shell-Laboratory, A msterdam. l7 Bolland, Quart. Rev., 1949, 3, I . Butler and Polanyi, Trans. Faraday SOC., 1943, 39, 19 ; cf. Evans and Walling and Mayo, ref. (13), p. 295 ; Mayo Polanyi, Trans. Faraday SOC., 1938, 34, T I . e l al., ref. (13), p. 285. 19 Cf. Price, ref. (13), p. 304.

ISSN:0366-9033    

DOI:10.1039/DF9511000163

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

ETHANE-ETHYLENE AND PROPANE-PROPYLENE EQUILIBRIA BY G. B. KISTIAKOWSKY AND A. GORDON NICKLE* Received 23rd January, 1951 Accurate measurements of the above equilibria are described, in which the equilibria were approached from both sides on a reduced chromic hydroxide catalyst and corrections were made for side reactions. For the reaction C,H, + C,H, + H, the equilibrium constants were found to be and 5-13 f 0.13 x I O - ~ atm. at 723.2' K . . . 4-04 f 0.17 x I O - ~ atm. a t 6 5 3 . ~ ~ K. For the reaction C3H8 -+ C,H, + H, the equilibrium constants were found to be 5-17 & 0.15 x ~ o - ~ a t m . a t 648.2' K and 3-67 f 0.17 x I O - ~ atm. a t 583.2O K. From these equilibrium constants the heats of reaction are calculated which are in satisfactory agreement with the calorimetric data.Using statistical methods on accepted molecular constants and calorimetric heats of reaction, equilibrium constants are calculated for the ethane system which are in excellent agreement with the present data. On the basis of these calculations and of other considerations, i t is suggested that the best value for the heat of hydro- genation of ethylene a t 298' K is AH = - 32,600 f 50 cal. The agreement of statistical calculations and the present data in the propane-propylene reaction is not quite as good, but the discrepancy may be due to the deviations of real molecules from the harmonic oscillator and rigid rotor models used in statistical calculations. An analysis of systematic errors in calorimetric work indicates that the best value for the heat of hydrogenation of propylene is AH = - 29,850 f 50 cal.a t 298" K. Quite some time ago a note appeared in print giving the results of our measurements of the above equilibria.1 Since they have some bearing on other thermodynamic properties of lower hydrocarbons, we present now a more detailed account of the measurements and their interpretation. and one of propane-propylene * have been described in the literature. The results are not internally consistent and are not fully in accord with the statistical calculations of equilibrium constants utilizing direct measurements of the heats of these reactions. A critical study of these measurements suggests two major sources of error. One arises from the methods chosen for the determination of the small amounts of olefins in the equilibrium mixtures, the other from the neglect of side reactions.We have attempted to eliminate both these errors in the present work. Four major investigations of the ethane-ethylene equilibrium 2s 4 e Experimental The equilibria were attained by repeated passes over a catalyst of gaseous mixtures of near-equilibrium composition a t controlled rates of flow and were measured by the determination of total pressure in the system and by analysis for hydrogen, total olefins and by-product olefins resulting from side reactions. * Standard Chemical Co. Ltd., Toronto, Canada. 1 Kistiakowsky and Nickle, J . Chem. Physics, 1942, 10, 78. Pease and Durgan, J . Amer. Chem. Soc., 1938, 50, 2715 Travers and Pearce, J . SOC. Chem. Ind., 1938, 53, 322. Frey and Huppke, Ind.Eng. Chem.. 1933, 25, 54. Pease and Byers, J . Amer. Chem. Soc., 1938, 60, 2489. Smith and Vaughan, J . Chem. Physics, 1935, 3, 341. I751 76 ETHANE-ETHYLENE, ETC., EQUILIBRIA Reaction System.-The reaction system consisted of a catalyst chamber with a thermocouple in a well inside the catalyst mass and a preheating chamber, placed in a metal furnace with manual temperature control. The temperature was measured by a chromel-alumel thermocouple, read on a Leeds and Northrup type K potentiometer with a HS galvanometer. The thermocouple was cali- brated a t several temperatures within the range used against a Bureau of Standards calibrated platinum resistance thermometer. The absolute tem- peratures given in the following are accurate to 0.5'.On a relative scale, the temperatures are accurate to 0.2'. The gas mixture was forced from a litre flask through a cold trap, the pre- heater chamber, and catalyst chamber into another litre flask by mercury, driven by a circulating pump patterned after the design of Urey.7 This pump permitted steady xates of flow of 4 to 250 ~m.~/min., but only Aows in the middle of this range were employed during equilibration runs. The all-glass system permitted multiple passes of the gases over the catalyst and was provided with the usual vacuum pumps and an accurate manometer. Catalyst.-Two catalysts were tried. The one described by Taylor and Joris,8 reduced copper on magnesia, failed to maintain high activity even a t 450° C. Good results were obtained with the catalyst described by Frey and Huppke,*~ and we followed their method of preparation except that the chromic oxide gel was washed by centrifugation rather than filtration.This catalyst retained its activity so well tbat one portion of the material was used in all the runs described. The precipitated and thoroughly washed chromic hydroxide was placed in an oven whose temperature was raised to 150' C during an hour and was dried a t that temperature for 5 hr. This material was ground and some of the -20 f 4 0 mesh material dried in vucuo a t 250°C for 18 hr. About 3 ~ m . ~ of the catalyst was now put in the catalyst chamber and hydrogen passed over i t a t 400' C for 12 hr. and then a t 490" C for 12 hr. Finally the catalyst was pumped out and flushed with ethylene, but still it retained much hydrogen which ap- peared in the gases of the first few equilibration runs.The same catalyst was used for propane after flushing i t with hydrogen a t 450' C and evacuation. There was always either a loss of olefin or an excessive production of hydrogen in the equilibration runs. To have the final gas mixture contain approximately equal amounts of olefin and hydrogen, it was necessary t o add olefin to the alkane on approaches from the low side, and to have olefin in excess of hydrogen on approaches from the high side. Materials .-HYDROGEN.-The hydrogen used in the catalyst treatment and in all equilibration experiments was Ohio Chemical Company's water-pumped hydrogenation grade gas. For use in equilibration runs i t was purified by passing over heated platinized asbestos and drying over Dehydrite.ETHANE.-Ethane was prepared by hydrogenation over copper of Ohio Chemical Company's 99-5 yo ethylene and was purified by repeated condensation, pumping off, and slow " bulb-to-bulb " distillations. One sample, which was used in Runs 36 to 44 inclusive, showed a difference of vapour pressures of 9.9 mm. in 490 when the last 5 yo was compared with the first 5 yo after a slow isckhermal distillation. 1 0 Most of the impurity was undoubtedly ethylene, since this material con- tained also the first fractions from the distillations of the second sample. This one was purified after hydrogenation by passing through 96 Yo sulphuric acid, potassium hydroxide solution and pellets, and a dry ice trap. The first 15 yo of this material, distilled off in vacuum, was combined with the first sample.On comparing the first and last 5 yo of the remainder, only a 2.2 mm. in 490 pressure difference was observed. Olefins were determined by bromination as described below, and found to the extent of 0.012 mole yo. This is not sufficient to account for the vapour pressure difference, and the impurity is most probably a paraffin formed by slow polymerization on the hydrogenation catalyst. Using the integrated form of the Rayleigh distillation equation,ll its concentration is computed as 0.02 mole Yo if the impurity is butane. Thode and Urey, J . Chem. Physics, 1939. 7, 34. Taylor and Joris, Bull. SOC. chim. BeZg., 1937, 46, 241. See also Lewis and Taylor, J . Amer. Chem. SOL, 1938, 60, 877.9 Frey and Huppke, U.S. Put. 2,098,959. loshepherd, J . Res. Nut. Bur. Stand., 1934, 12, 185. 11 Walker, Lewis, McAdams and Gilliland, Principles of Chemical Engineering (McGraw-Hill Book Company), 3rd edn., p. 532.G. B. KISTIAKOWSKY AND A. GORDON NICKLE 177 ETHYLENE.-Ethylene used in the control analyses and in equilibration runs was Ohio Chemical Company’s 99-5 yo material, purified somewhat by several slow bulb-to-bulb distillations with large rejections a t both beginning and end of each distillation. A difference in vapour pressure of 1-2 mm. in 490 was obtained when the first and last 5 yo of a slow isothermal distillation were com- pared. Assuming the impurity to be ethane, a similar calculation to that de- scribed above for ethane shows an impurity here of about 0.06 mole yo.PRoPANE.-Propane was Ohio Chemical Company’s 99.9 yo pure material purified by passing through 96 yo sulphuric acid , potassium hydroxide solution and pellets and collected in a dry ice trap. The product was further purified by filtration through glass wool at liquid nitrogen temperature to remove water, according to Roper’s method,12 then distilled bulb-to-bulb several times. A test for olefin showed none to the sensitivity of the method, viz., about 0.002 yo. The difference in vapour pressures of the first and last 5 yo was zero to the sensitivity of the method, viz., a few tenths of a millimetre at 600 mm. total pressure. Because of a crack in the apparatus, part of the sample later became mixed with a little air and was re-purified only by bulb-to-bulb distillation.Thus i t may have been contaminated with a trace of water, but no difference was noted in runs made with the slightly impure gas and those with the best material. PRoPYLENE.-Propylene used as a standard for analytical work was given us by Dr. F. D. Rossini, and was labelled by him as the remainder from their best sample used in the determination of the heat of combustion of propylene. A full discussion of the preparation and purity is given in the paper where the heat of combustion is reported.13 The propylene used for addition, when ap- proaching the equilibrium irom the high side, was from the Linde Company, labelled 99.9 % pure. The only purification was condensation in liquid air and pumping to remove any non-condensable material. Analyses.-Various precautions recommended in precision analytical work were applied in making, standardizing and preserving reagent solutions.Analytical weights and volumetric apparatus were recalibrated. The procedure €or the determination of the composition of equilibrated mixtures was: (i) combustion of hydrogen over copper oxide, (ii) photo- chemical bromination in the vapour phase of a portion of the hydrocarbon fraction with a subsequent determination of excess bromine to determine total olefins, and (iii) catalytic bromination of another sample of the hydrocarbons, followed by a determination of the freezing point of the dibromides to evaluate the amount of the pertinent olefin in the total olefins. HYDROGEN.-After the run, about 200 ~ m . ~ of the gases were transferred into an all-glass system containing hand-operated Toepler’pumps, a gas burette, a heated copper oxide tube, a phosphorus pentoxide tube, a cold trap and suitable connections with numerous stopcocks.The hydrocarbons were frozen out in the trap by liquid nitrogen, hydrogen was pumped off, the hydrocarbons melted and re-frozen, then pumped off again. The hydrogen-rich fraction was measured, then repeatedly passed over copper oxide at 250° C and water removed by phosphorus pentoxide. The volume of the residual gases was again measured and hydrogen determined by difference. The residual gases (largely methane, with a little ethane in the C , equilibria) were then combined with the material in the cold trap. This procedure was tested for the following sources of error : (u) oxidation of hydrocarbons on copper oxide, (b) incomplete combustion of hydrogen, (c) sorption or desorption of gases on the solids in the system.None was found to affect the results by more than 0.1 yo. A control mixture con- sisting of hydrogen, ethane and ethylene was analyzed and the result agreed within 0-1 yo with the actual composition. This is slightly better than the estimated probable error of 0.2 %. TOTAL OLEFINS.-About 500 ~ m . ~ of the hydrocarbons from the run were introduced into flask D shown in schematic manner in Fig. I. After the intro- duction the stopcock was closed and by means oi an external magnet the dry ground-glass plug B was pushed into place, protecting bromine vapour from interaction with stopcock grease.Then, with another magnet, the glass-enclosed iron piece E was lifted and dropped on the thin-walled, sealed glass capsule F, containing a weighed amount of bromine, breaking it. The breaker was moved up and down a, number of times to mix bromine vapour with the contents of the flask, and the latter was then exposed for 10 min. a t 15 cm. distance to the l2 Roper, Ind. Eng. Chem. (Anal.), 1937, 9, 414. 13 Rossini and Knowlton, J . Res. Nat. Bur. Stand., 1937, 19, 249.1 78 ETHANE-ETHYLENE, ETC., EQUILIBRIA light of a 500-W incandescent lamp. Neutral potassium iodide solution was poured into bulb 2, the ground glass plug lowered, and the solution was run into flask D without admitting air, and shaken well to scrub the bromine from the vapour phase. The upper connections and top part of the flask were washed down by admitting water through bulb A, then the ground glass joint C (held together with a low-melting thermoplastic cement applied only to the outer edge of the joint) was separated and the solution titrated with thiosulphate using starch as indicator.When 0.897 3'418 3'542 0.458 FIG. I . 0.900 0-3 3'452 3'4 3.586 4'4 0.496 3'8 the end point was reached, neutral potassium iodate solution was added and the liberated iodine titrated again. In a well-known manner the second of these titrations measures the quantity of acid present and so the extent of the sub- stitution reaction of bromine. The quantity of olefins present was evaluated from these titrations by assuming that the reactions occurling were : C,H,, + Br, + ~ , H , , B r , C,H,,+, + Br, -+ C,H,,+,Br -t HBr CnH2a + HBr -+ C,H2,+1Br C,H,,,Br, + Br, -+ CnH2,1+lBr3 + HBr, etc.Hence moles of biomine consumed in the reaction less moles ot acid formed equal moles of olefins. Since the amounts of acid involved are small, the flask had to be scrupulously free from traces of acid ; in particular, i t was found that cleaning solutions containing divalent acids could not be used because these acids were strongly adsorbed ; prolonged subsequent cleaning was required before the flask was sufficiently free of acid. to be serviceable. Much work was done on this procedure to obtain an exact equivalence between theoretical and experimental results, but the quantity of olefins found was always slightly less than that introduced.The magnitude of this loss, shown in Table I, was tound to depend on the quantity of bromine introduced and on the olefin being determined. No wav was found to prevent this loss and an empirical correction, obtained by plotting the data of Table I , was applied to a d runs. In the worst case (C, hydrocarbons at 450" C) the correction changed the equilibrium constants by about 1-5 %. The difficulty is probably due to a free radical induced polymerization of the olefins, and it is interesting to note that work on this subject l4 has shown that the rate is four rimes as fast with ethylene as with propylene, in qualitative agreement with the results obtained by this analytical method. TABLE 1.-Loss OF OLEFIN IN PHOTOCHEMICAL BROMINATION Olefin Ethylene Propylene I Bromine Added (Moles x 104) 1'427 5'582 5'957 6.133 1.849 5.952 6.116 Bromine Found (Moles x 104) ~~ 0.530 2.164 2.415 5'675 0'547 I 1.898 2.016 Olefin Total Olefin Olefin Found 1 Added 1 Loss (Moles x 104) (Moles x 104) (Moles x 10-6) I I PURITY OF 0LEFINS.--The extent of the substitution reactions taking place on illumination is such that the recovered bromides are not suitable for the determination of by-products.Therefore a portion of the equilibrium mixture was broniinated catalytically under conditions which minimize the substitution reaction. Slightly more than the required amount of bromine vapour was picked l4 Reeck and Rust, J , Chenz. Physics, 1941, 9, 480.G. 13. KISTTAKOWSKY AND A. GORDON NICKLE 179 up by a stream of nitrogen and mixed with a stream of the hydrocarbons just before the catalyst chamber.The catalyst was made by saturating chips of porous plate with a strong calcium bromide solution and drying, and was operated a t a temperature of about 50°C. Excess bromine was removed by passage over copper, the products were frozen out in a cold trap, lower-boiling constituents were removed by pumping and the high-boiling material was then used for melting point determinations which were carried out in a micro-adapta- tion of the melting curve technique.15. l6 Approximately 0.01 ml. of the dibromides was introduced into a thin-walled capillary, about the size of a large melting point tube, and a copper-constantan thermocouple of fine wires inserted into the caplllar?, well under the surface oi the liquid.By friction, the thermo- couple suspended the capillarv in the centre of an electrically heated cylindrical shield, closed at the top and bottom by corks. The whole was mounted in an enclosing glass shield and immersed into a constant temperature bath slightly colder than the melting point of the dibromides (melting ice for C,, melting chloroform for CJ. Because of the small size of the sample, the absence of stirring 400 MOLE % IMPURITY FIG. 2.-The effect of impurities upon the melting point of ethylene dibromide. 0 Propylene dibromide ; 1-butylene dibromide ; A 2-butylene di- caused no difficulty, but the magnitude of thermal leaks, as represented for instance by the size and length of the thermocouple wires, had to be carefully and reproducibly adjusted.This done, the device functioned very satisfactorily and permitted the determination of freezing points by observing melting curves almost as accurately as the larger devices. Fig. 2 shows the effect of several likely impurities on the freezing point of ethylene dibromide ; Fig. 3 provides similar data for propylene dibromide. In these figures, temperatures are plotted as microvolts, the reference junction being in an ice bath. The reliability of this technique may be. appreciated from the following. Brominations of mixtures of ethane and " pure " ethylene, and propane and " pure " propylene gave dibromides which analyzed 99.9 yo pure ; when I mole % of ethylene in the ethane-ethylene mixture was replaced by propylene, the dibromide was determined as 99.1 mole yo pure.Such tests indicate that i t is possible to determine the concentration of ethylene 01 propylene dibromides in the mixtures t o about 0.2 mole yo accuracy. ALKmES.-The partial pressure of ethane or propane in the equilibrium mixture was determined by difference. The quantity of methane found in bromide ; x 68 yo I-butylene dibromide, 32 yo propylene dibromide. 15Kistiakowsky. Ruhoff, Smith and Vaughan, J . Amer. Chem. SOC., 1935, 16Roper, J . Amer. Chem. SOC., 1938. 60, 1693. 579 876.I 80 ETHANE-ETHYLENE, ETC., EQUILIBRIA the gaseous residue after the combustion of hydrogen on copper oxide was always far too small to account for the olefinic impurity plus its alkane, if they were formed by disproportionation reactions such as and zC,H, -+ CH, + C3H8 2C3H8 -+ CH, $- C,H12.The methane is evidently formed by cracking reactions such as and and the olefinic impurities by independent polymerization reactions such as and No attempt was made to determine the composition of the impurities, as this information is not necessary for the interpretation of the freezing point data. > -'950r------- a t- z 0 - a -2000 c3 z t- -I W - -2050 MOLE % IMPURITY FIG. 3.-The effect of impurities upon the melting point of propylene dibromide. a Ethylene dibromide ; 0 2-methyl-2 : 3-dibromopentane. The alkane impurity in equilibrium with the olefin impuiity must be present in relatively very small amounts since one alkane is in equilibrium with several isomeric olefins. and of Kilpatrick, Frosen, Pitzer and Rossini 18 to calculate the concentrations of butanes and hexanes in equilibrium with the pertinent olefins, i t is found that the partial pressure of the alkane impurity would not be more than 0-2 yo of the total pressure under the conditions of this research.Rather than introduce a doubtful correction of such small magnitude, i t has been ignored. The difference between the total pressure prevailing during the last pass of the equilibration run and the sum of the partial pressures of the other constituents discussed above was assumed to be ethane or propane as the case might be. It was assumed that all the gases concerned act ideally a t the temperatures near 400° C a t which equilibration runs were made. The pressure and volume values obtained a t room temperature were converted to molal quantities using equation of state data from the following sources : hydrogen, ethane and ethylene from International Critical Tables ; *Q propane from Kemp and Egan ; 2 O propylene from Roper.21 The equilibrium constants were then calculated in the usual manner.The individual values were converted to stated temperatures by the van't Hoff isochore, using 34.0 kcal. per mole for the heat of reaction in the C, system and 30.8 kcal. per mole for the C , system. The maximum temperature correction was 1*5OC, and i t was usually only a few tenths of a degree. 17 Kilpatrick, Prosen, Pitzer and Rossini, J . Res. Nut. Bur. Stand., 1946, 1* Prosen, Pitzer and Rossini, J . Res. Nat. Bur. Stand., 1945, 34, 403. 1Q Int. Crit. Tables, 1st edn., 3, 3. 2 O Kemp and Egan, J .Amer. Chew SOL, 1938, 60, 1521. 21 Roper, J . Physic. Chem., 1940, 44, 835. Using the data of Frey and Huppke 36, 559-G. B. KISTIAKOWSKY AND A. GORDON NICKLE 181 L RUI No - 29 30 32 33 34 35 42 43 44 36 37 38 39 40 41 Results ETHANE.-Tabk I1 summarizes the results of all ethane runs after No. 28, with the exception of Run 31 in which the sample of dibromides was lost and hence no purity determination could be carried out. Runs preceding No. 29 were either of preliminary nature or were incomplete for one reason or another. The data a t 450' C excepting Run 30, are extremely self-consistent, the mean deviation being only 1.4 yo. These runs differed in direction of approach, contact time, total pressure and in composition. The reproducibility of the equilibrium constants is therefore the best proof that equilibrium was really attained and was correctly measured. The reason for the failure of Run 30 to agree with the rest remains unknown; this run was not included in the average since i t deviates by almost seven times the mean deviation of the others.The data a t 380'C are less extensive and accurate, the mean deviation being about 3 %. The data suggest that the equilibrium was not quite attained from either side, but against this are the data of Runs 36 and 37 which agree with the rest even though the contact time in these runs was only half as long. TABLE II.-SUMMARY OF DATA FOR ETHANE-ETHYLENE-HYDROGEN EQUILIBRIUM Approx. Initial Pressures - C2H4 (Atm X 10-2) nil 4.1 4' 2 5'0 4'1 5'2 1'5 4'9 1'1 0-6 1.1 0.3 nil 1'2 1'2 - HI 10-2) ( Atm X nil 3'1 nil 2.5 3'1 nil 3'8 nil 3'9 nil 0.9 nil nil 1'1 1'2 - %Ha Atm.0.80 0.71 0.80 0.84 0.97 0-71 0.63 0.76 1'00 0.86 0.78 1-02 1'01 0.80 0.78 - Con tacl Tim Per Pas: :Set, 2 2 2 2 I 2 2 2 2 2 2 2 2 2 2 - No. o Passe 5 5 4 3 5 5 7 8 6 5 5 I0 I1 I0 11 Im- purity % of Total Olefin - 3.1 6.3 9.6 9'0 7-7 7.3 6.0 8.6 8.3 6.0 5'0 6.6 6.3 6.2 3'1 C2H4 (Atm. 10-2) X 0.894 0.887 1.404 1'774 1.967 1.440 I -5 06 1.941 2.062 7.78 7'23 6.93 6-20 5'39 5'77 H4 (Atm. 104) X 4'495 3'97 2'994 2'557 2.762 2'591 2.364 2.728 2.088 4'35 4'72 5'79 6-55 5-91 543 - GH6 (Atm.] - 0.780 0.726 0.803 0.838 0'993 0.695 0.663 0.976 0.790 0.85 I 0.783 1.006 1.007 0.780 0.778 KT (Atm.) 5-15 X 10-4 4-85 5'23 5-41 5'47 5'37 5'37 5'43 5'45 3-98 x I O - ~ 4-36 3'99 4'03 4-08 4'33 Temp.("C) 451-2 451'3 451.1 451.5 451'4 451'3 451'4 451'3 451'3 380.5 380.6 380.6 380.6 380.6 380.5 Ko23.2 (Atm.1 4.95 x 10-4 4.63 5-04 5'15 5'23 5'15 5'13 5-20 5.22 K653*2 3.90 x 10-5 4-26 3-91 3'95 3-98 4-23 A careful consideration of several sources of systematic errors, as well as of accidental errors, the magnitude of which is inferred from Table 11, leads us to the conclusion that the standard error of the average value of the equilibrium constant a t 4jo' C, whch is K,,,., = 5-13 x I O - ~ , is less than 2-5 yo. The standard error of the value a t 380° C, K653.2 = 4-04 x I O - ~ , is less than 4 yo. Propane.-The results of all completed runs are shown in Table 111, and i t will be noted that the values obtained in the first few runs are quite low.The catalyst was pumped extensively before the runs previous to No. 52, and an examination of the table shows a distinct separation in values for the equilibrium constant between approaches from the high olefin side and approaches from the low side. In subsequent runs this pumping was dispensed with and the difference has disappeared. The reasons for this effect are not clear, but the important fact is that with the cessation of pumping, practically identical results were obtained on approaches from both sides in several runs. This seems a re asonable criterion of equilibrium. From the data of Table 111 we compute K,,,., = 5-17 x I O - ~ and. K683.2 3-67 x 10-6. Taking into account systematic errors as well as the accidental errors shown by the scatter of the results, we believe that the standard errorI 82 Vo. of 'asses 3 3 3 5 5 g 4 6 5 7 8 ETHANE-ETHYLENE, ETC., EQUILIBRIA Im- purity yo of Tot a1 -- 2-8 4'5 3'3 3'5 2-8 2.5 ' 2'1 2'1 2'0 2'0 1.5 of the average value of the equilibrium constant at 3 7 5 O C is less than 3 yo ; the value at 310' C has a standard error of less than 4-5 yo.Discussion Ethane.-Fig. 4 shows a plot of the equilibrium constants determined here, together with those given in the literature. Since a wide range of values of equilibrium constants is covered, it has been necessary to add a function of temperature (essentially a term of the type AHIT) to the logarithm of the equilibrium constant to get a sufficiently large ordinate scale to allow critical comparison of the results. Their agreement is not perfect, but in view of several sources of error in the earlier work it is surprisingly good.TABLE III.-SUMMARY OF DATA FOR PROPANE-PROPYLENE-HYDROGEN 0.828 0.982 0.984 0.647 0.664 0.814 0.838 0.884 0.918 0.920 04g1 EQUILIBRIUM -___ 5-06 5.01 5-32 5.09 5-24 5.28 5'30 5-21 5'29 3.66 3.66 Run No. - 45 47 48 49 50 51 54 55 56 52 53 Approx. Initial Pressures C3H6 X [Atrn. 10-2) 2'4 3'2 5'6 2'4 3'2 3'6 nil 4'4 0.6 nil 1'2 - H2 10-2) Atm. X nil nil 3'0 nil 3'0 3'1 nil 4'4 nil ni 1 1'2 - C3H8 htm.) 0.86 0.96 0' 66 0.65 0.80 0.87 0.86 0.94 I '00 0.93 0.89 - Con- tact rime per Pass [Sec,) 3'3 3'3 3'3 3' 3 3'3 3'3 3'3 3'3 3'3 3'3 3'3 CsH6 X Atm. xo-2) 2'033 1.942 2.693 2.107 1.649 1.899 1.996 1.962 2-3 2 2 0.653 0.641 H2 (Atrn. xo-2) X 2.061 2.533 x-944 1.563 2.108 2.263 2-225 2.349 2.093 0.516 0.509 Temp.("C) 375'4 375'2 375'6 375'0 375'0 375'2 375'0 375'0 375'0 309'9 310.0 Km-9 (Atm.) 4.98 x I O - ~ 4'97 5-21 5'09 5'24 5-23 5'30 5.21 5'29 K803.2 3.68 x IO-~ 3.66X 1o-I' In the work of Pease and Durgan, Travers and Hockin,22 and Travers and Pearce,s equilibrium was obtained by the heating of near-equilibrium mixtures in silica bulbs for appropriate times. In all this work there are two main potential sources of error. The first is the method of analysis for ethylene, which in no case is entirely above suspicion, but errors would be such as to yield high values for the equilibrium constant. The second is the extent of side reaction to produce (mainly) methane, and the dis- turbing effect this might have on the equilibrium. At 973'K, Pease and Durgan found more than 30 yo methane in the reaction mixture and comment that " the most serious source of error is indicated by the presence of considerable amounts of methane among the products of some of the experiments.The corresponding side reaction undoubtedly has displaced equilibrium to some extent." As low as 863" I<, Travers and Pearce frequently found more than 20 yo methane, but claim that " it is a fact very definitely established by experiment that the equili- brium relations between ethane, ethylene and hydrogen in a mixture initially in equilibrium, and undergoing pyrolysis, remain undisturbed ". They carried out experiments under widely different conditions, and indeed no trend is observable between the value for the equilibrium constant and any other measured quantity.However, at 883OK, the diffsrence between the equilibrium constant obtained from Fig. 4 and 22 Travers and Hockin, Proc. Boy. Soc. A , 1932, 136, I,G. F,. KISTIAKOWSKY AND A. GORDON NICKLE 183 that reported by Travers and Pearce is only about 7 yo. Detailed exam- ination of their results shows variations of f 5 yo in the individual values for the equilibrium constant ; these variations being of about the same magnitude as the effect in question, they could mask a trend between the extent of side reaction and the value of the equilibrium constant. In the work of Frey and Huppke,4 the high values obtained by extra- polation of equilibrium data to infinite contact time have already been recalculated.23 Since total olefins were determined by fuming sulphuric acid and were reported as ethylene, the agreement of the recalculated values with the present work is satisfactory. The work of Vvedenski and Vinnikova 2 4 is subject to many criticisms, their high results being attributable to the inclusion of higher olefins as ethylene. Thus, while the equilibrium constants determined here are not in perfect agreement with those given in the previous literature, we believe that the fault is not ours. 7.32 7.28- FIG. 4.-Comparison of equilibrium data on ethane -ethylene reaction from various sources. Curve I : statistical cal- culations based on AH of Kistiakowsky e f al. ; Curve 2 : statistical calculations based on AH of Prosen and Rossini ; A this re- search; 0 Frey and Huppke, original ; 0 Frey and Huppke, recalculated ; x Travers and Pearce ; @ Travers and Hockin; Pease and Durgan.- 7 24- 7.20 + y 716- - ; I 7.12 7.08 0 I I 7 - x 0 0 700: 1.0 1 1 1 2 1.3 1.4 1.5 T 6 By the use of the van’t Hoff isochore, the heat of hydrogenation of ethylene is calculated from the data of Table I1 as AH,,, = - 34-03 kcal. per mole with an uncertainty of 1-2 kcal. p-r mole. This value agrees extremely well with the value calculated for this temperature from the heat of reaction determined directly at 355°K and the heat capacity data discussed below, which gives AH,,, = - 33-95 kcal. per mole. A more sensitive test of the self-consistency of equilibrium data is obtained by comparing the best value of the equilibrium constant determined here with that calculated statistically from the directly deter- mined heat of reaction and an assortment of molecular constants.These calculations are by now so well known that we shall merely state the authorities whose molecular constants we have used and give the results. 23 Kistiakowsky, Romeyn, Ruhoff, Smith and Vaughan, J . Amer. Chem. SOL, 193.5, 57, 65- 24 Vvedenski and Vinnikova, J . Gen. Chem. (U.S.S.B.), 1934, 4, 120.184 ETHANE-ETHYLENE, ETC., EQUILIBRIA Over several years there have been published by the National Bureau of Standards extensive and self-consistent tables of thermodynamic functions of hydrocarbons. Of these the tables by Rossini, Pitzer, ct aLZ5, 2 6 ~ 27 and by Brickwedde, Moskow and Aston are appropriate for the calculation of the above-mentioned equilibrium constant. Both sets of data give essentially the same results and indeed are based on almost the same physical constants.A new assignment of vibrational frequencies of ethylene which was recently proposed by Arnett and Crawford 20 reduces all calculated equilibrium constants by 4 to 5 yo. In addition to the statistical data, the heat of hydrogenation of ethylene is required. Three values for this, based on modern data, should be considered : (i) - 32,575 cal./mole at 298.16~ K from the direct determination of the heat of hydrogenation of ethylene at 82' C 23 with the original data corrected for the new atomic weight of carbon. (ii) - 32,777 cal./mole at 298.16~ K from the heats of combustion of ethane,a0, 31 ethylene (iii) - 32,732 cal./mole at 298.16~ K estimated by Prosen and Rossini 33 to be the best value for the heat of hydrogenation.In Table IV are shown the values for the equilibrium constants using these heats of reaction and the tables published by Brickwedde, Moskow and Aston 2* for the statistical calculations. The first choice would ob- viously be the correct one, were it not that anharmonicity of vibrations TABLE IV.--COMPARISON OF EXPERIMENTAL AND CALCULATED EQUILIBRIUM CONSTANTS FOR THE REACTION C,H, --f C,H, + H, and hydrogen.27, *2 Heat of Reaction a t 298'i', A H (kcal. ) - 32'575 321777 32,732 Heat of Reaction: at 688'. A H (kcal.) 34'0 f 1'2 33'95 34'15 34-11 I Source Present data Kistiakowsky et al. Rossini et al. Prosen and Rossini 4-04 f 0.17 4-17 3'57 3-70 and the stretching of molecules by rotation were neglected in the statis- tical calculations for ethane and ethylene.The magnitude of the necessary corrections is unknown and hence the final selection among the three choices of Table IV must be postponed, although certain plausi- bility considerations, which we shall now give, indicate the best choice. When the simplifications of rigid rotors and harmonic oscillators are eliminated, the calculated h2a.t capacity of ethane should suffer a greater increase than that of ethylene because of looser structure of ethane and a larger number of vibrations. Thus the chemical potential of ethane will decrease more than the sum of the chemical potentials of ethylene 25 Kilpatrick and Pitzer, J. Res. Nal. Bur. Stand., 1946, 37, 163 ; 1947, 38, 191.26 Pitzer, Ind. Eng. Chem , 1944, 36, 829. 27 Wagman, Kilpatrick, Taylor, Pitzer and Rossini, J . Res. Nut. Bztr. Stand., 1945,34, 143. 28 Brickwedde, Moskow and Aston, J . Res. Nut. Bur. Stand., 1946, 37, 263. 29 Arnett and Crawford, J . Chem. Physics, 1950, 18, 118. 30 Prosen and Rossini, J . Res. Nut. Bur. Stand., 1945, 34, 263. 31 Rossini, J . lies. Nut. Bur. Stand., 1934, 12, 735. 32Rossini, J. Res. Nat. Bur. Stand., 1939, 22, 407. 33 Prosen and Rossini, J . Res. Nut. BUY. Stand., 1946, 36, 269.G. B. KISTIAKOWSKY AND A. GORDON NICKLE 185 and hydrogen and the net result should be an increase in the value of AGO, thus a smaller equilibrium constant for the reaction under discussion. The values calculated for the equilibrium constant using the heats of reaction proposed by Rossini et al.are already too low, so that a more refined treatment of the vibrations and rotations could not improve the agreement but would make it worse. The same reasoning applies a fortiori when the Arnett-Crawford assignment for ethylene is used, since it drops all calculated K's by a few per cent. On re-studying the papers 23, 3% 3% on the heat of hydrogenation of ethylene, we find ourselves unable to agree with Rossini's arguments against the lower value. The same catalyst as was used for ethylene was used for very extensive work with higher olefins, and in no case was any polymerization noted. The extent of polymerization of the higher olefins was certainly less than 0-5 yo and to suppose that ethylene did polymerize to the extent of 3 yo is quite unjustified.As regards the samples of ethylene studied by Kistiakowsky', Romeyn, Ruhoff, Smith and Vaughan,s6 the data on their own purified ethylene still seem to us the best because this ethylene was prepared by a method which excluded ethane, because it was repeatedly distilled, but mainly because the middle and last fractions gave heats of hydrogenation which were identical, within the very small average deviations of the two series of measurements, 44 cal./mole. If ethane was present in the material which was subjected to the final distillation, it would certainly be enriched in the " last fraction. In deference to Rossini's higher figure, we suggest, therefore, the rounded value of AH,,, = - 32,600 (f 50 ?) cal./mole, thus almost within Rossini's own estimated " precision uncertainty " which is in the nature of a probable error and hence leaves considerable probability for a slightly larger deviation than the precision indicated.It also places the Harvard figure, for the ethylene sample obtained from the Linde Air Products Company, as recalculated by Rossini allowing for 0.25 yo ethane, within the probable deviation range of the final value. Using this reaction heat, the calculated equilibrium constant at 723.2" K comes out to be 5-15, or 4.89 if Arnett and Crawford's assignment is used, in good agreement with the experimental value unless the anharmonicity correction is unexpectedly large. Propane .-Fig. 5 shows a plot of experimental and calculated equili- brium constants. As before, skew co-ordinates have been used.The work of Frey and Huppke 4 is obviously in disagreement with this research, and the same objections to their work are applicable here as for ethane, namely, uncertainty in the analytical method for olefin and no proof of the purity of the olefin found. In connection with some work primarily concerned with the rate of polymerization of propane, Travers 36 reported the value 0.085 for the C, equilibrium constant at 826°K. Careful examination of his data shows numerous discrepancies between the equilibrium constant which he gives and those which can be calculated from his reported analytical data. Presuming the analytical data to be correct, an average value of 0-077 is obtained for the equilibrium constant at 826' K and is the value plotted in Fig.5. The individual values scatter considerably, and the mean value could easily be in error by 5 yo. As in ths ethane work, it is possible that polymerization processes have had a disturbing effect on the equilibrium. By use of the van't Hoff isochore, the heat of hydrogenation of propylene is calculated from the data of Table I11 as AH,,, =- 30.5 f 1.2 kcal. per mole. This value agrees well with the value calculated from the heat of reaction determined directly at 355" K. 34Ro;sini, J . Res. Nat. Bur. Staf-zd., 1936, 17, 629. 35 Kistiakowsky, Ruhoff, Smith and Vaughan, J . Amcr. Chem. SOC., 1936, 58, 137. 38 Travers, Trans. Faraday SOL, 1937, 33, 751.I 86 ETHANE-ETHY LENE, ETC., EQU ILIBR IL4 For statistical calculations, adequate tables have been published by Pitzer et aLza, a3, a4 Three values are available for the heat of hydro- genation : (i) - 29,877 cal./mole at 298.16~ K from the direct determination of the heat of hydrogenation of propylene at 355037 with the original data corrected for the new atomic weight of carbon.(ii) - 29,532 cal./mole at 298.16~ K from the heats of combustion of propane,29, 31 propylene l3 and hydrogen.27, aa (iii) - 29,699 cal./mole at 298.16~ K estimated by Prosen and Rossini 33 to be the best value for the heat of hydrogenation. 321 7 28 1 y ! I 7 I 7 121 7 08 1 0 10 1 5 1.8 7 04 1x0 T FIG. 5.-Comparison of equilibrium dath on propane-propylene reaction from various sources. Curve I : statistical calculations based on AH of Prosen and Rossini ; Curve z : statistical calculations based on AH oi Kistiakowsky et al.; A this research; Travers; 0 Frey and Huppke, recalculated. In Table f a r e shown the values for the equilibrium constants cal- culated using these heats of reaction and the above-mentioned tables. None of the values gives as close a check with the experimental value as does ethane, but here the cal- culated values are higher than the experimental. As in the ethane reaction, this is the direction in which errors would be expected in the calculated values arising from the deviations of real molecules from the rigid rotor, harmonic oscillator models. Furthermore, th? greater uncertainty in the molecular models, arising from doubtful frequency assignments and inaccuracies in internal rotation barriers, makes the discrepancy between the cal- culations and the present measure- ments understandable. It prevents a positive identification of the best ” value for the heat of this reaction. The following considera- tions, however, suggest that the best value is close to - 29,877 cal. determined in hydrogenation ex- periments. Since Rossini’s value is lower, its difference from the hydrogenation experiments cannot be explained by such assumptions as the polymerization of olefin or the presence of saturated impurity in hydrogenation work, but can only be due to the presence of a more unsaturated impurity in the latter. In view of the two different methods of preparation, the extensive purification and the close agreement between values obtained for the heat of hydrogenation of the two samples of proyylene, this source of error does not seem reasonable. The clean- ness of the roaction and freedom from cracking during hydrogenation were amply demonstrated by appropriate methods. Systematic error of the calorimeter or calorimetric technique is ruled out by the close agree- ment between the values obtained by Rossini et al. and Kistiakowsky et U Z . ~ ~ for the formation of water from the elements. Thus, since we can find no adequate source of error in the hydrogenation value, and since 37 Kistiakowsky, Ruhoff, Smith and Vaughan, J . Amer. Cjiem, Sac., 1935, 58,876.G. B. KISTIAKOWSKY AND A. GORDON NICKLE 187 this value is within the estimated uncertainty of that determined by Rossini et al., we believe the former to be more nearly correct and recommend AH,,, =- 29,850 f 50 cal. as the final choice. KSP 9. I , Heat of Reytion Heat of Reaction at 298.1 at 616O, Source AH (kcal:) AH (kcal.) Atm. x 104 - 30.5 & 1-2 Present data 5-17 0.15 29377 30.85 Kistiakowsky 6.04 299532 30.50 Rossini ef al. 7-89 29,699 30'67 Prosen and 6.92 et al. Rossini TABLE V.-~OMPARISON OF EXPERIMENTAL AND CAIXU LATED EQUILIBRIUM CONSTANTS FOR THE REACTION C,H, -+ C,H, + H, Kba3.2, Atm. x 105 3-67 + 0-17 4-15 5-58 4-85 Gibbs Chemical Laboratory, Harvard University, Department of Chemistry, 12 Oxford Street, Cambridge 38, Mass., U.S.A.

ISSN:0366-9033    

DOI:10.1039/DF9511000175

出版商:RSC    

年代:1951

数据来源: RSC