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Proceedings of the Chemical Society. September 1962 |
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Proceedings of the Chemical Society ,
Volume 1,
Issue September,
1962,
Page 289-316
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PROCEEDINGS OF THE CHEMICAL SOCIETY SEPTEMBER 1962 CENTENARY LECTURE* The Mechanism of Oxidation of Acetylene at Very High Temperatures By G. B. KISTIAKOWSKY (HARVARD CAMBRIDGE, UNIVERSITY MASS.,U.S.A.) THIS lecture is a partial account of the work of a the front of the wave remains essentially constant. If number of my recent collaborators Drs. P. H. Kydd the usual thermodynamic functions of the gaseous W. C. Gardiner jun. J. N. Bradley L. W. Richards mixture are known a measurement of the speed of and C. W. Hand who have developed several tech- the wave can be used to calculate the temperature niques for studying very fast oxidations and other and pressure in the wave. When a wave is reflected reactions taking place under conditions of high from the end of the tube the temperature and pres- temperature in shock waves.Specifically I shall deal sure are greatly increased but still can be calculated with the oxidation of acetylene by oxygen for which from the velocity of the wave and the thermodynamic detailed information on the mechanism has been parameters of the gases. gained. A number of investigators among whom I would Shock waves can be generated in a very simple mention particularly N. Davidsonl and D. F. apparatus consisting of a long tube separated into Hornig,2 have used shock waves to investigate very two compartments by a frangible diaphragm. One simple chemical reactions such as dissociations of compartment is filled with the “driver” gas for diatomic molecules. instance hydrogen at a few hundred millimeters Of the techniques to observe the more complex pressure the other by the “driven” mixture for reactions which we have attempted to develop the instance argon with a few mole % of the experi- most informative is that of letting the gaseous mix- mental gases e.g.oxygen and acetylene at a total ture effuse continuously from the shock tube through pressure of a few millimeters of mercury. When the a small pin-hole in its end plate into the evacuated diaphragm is suddenly ruptured the expansion of space of a specially designed time-of-flight mass the high-pressure gas causes a flat-topped wave of spe~trometer.~ This apparatus a modification of the supersonic speed to travel down the tube into the commercial Bendix time-of-flight mass spectro-low-pressure mixture.If no chemical reaction takes meter permits the display of the mass spectrum at place in the shock wave the thermodynamic state of a repetition frequency of 20 kilocycles per second. the gas for a considerable distance and time behind While the records are very poor compared with * Delivered before the Society at Manchester College of Science and Technology on November 14th 1961 at Burlington House London W.l on November 16th 1961 and at Cambridge on November 17th 1961. Davidson and Schott Discuss. Faraday SOC.,1954 17 58. Palmer Britton and Hornig J. Chem.Phys. 1957 26 98. Kistiakowsky and Kydd J. Amer. Chem. SOC.,1957,79,4825. Wiley and McLaren Rev. Sci. Instr. 1955 26 1150. 289 those which can be obtained in a matter of minutes frcm other mass spectrometers the rapidity of dis- play permits us to observe chemical events taking place within the gases in a matter of milliseconds or less a time which is appropriate for shock-wave studies.A second technique developed in my laboratory scme years ago measures the density of the shocked gas as a function of time with a resolution time of about one microsecond by displaying the time- dependence of absorption of soft X-rays. To ensure adequate X-ray absorption the argon in the experi- mental mixture is replaced by xenon. The density of the gas within the shock wave decreases when an exothermic reaction takes place whereas an endo- thermic reaction causes an increase in density. The third technique consists in observing the emission of radiation.s In addition to radiation in the visible and near-ultraviolet region of the spectrum we found to our surprise that oxidation of hydro- carbons also causes emission of radiation in the far- ultraviolet region between 1800 and 1500 A.To detect this radiation we have used lithium fluoride optics and a specially designed photomultiplier which is sensitive only to radiation of wavelengths shorter than about 1800 A. Our last technique that was used in the work I am reporting involves measurement of ionisation caused by chemical reaction,' by introducing the so-called Langmuir probe into the shock tube. The observa- tions of changes in gas density of radiation and of gas ions have to be displayed by means of cathode- ray oscilloscopes to achieve the necessary high-speed response of the apparatus.But apart from the need to work with very short time intervals the technique presents no special difficulties. Typical mixtures in our experiments on the oxida- tion of acetylene contained 95 mole % of argon or xenon the remainder being various mixtures of acetylene and oxygen. On compression by the expanding hydrogen from the high-pressure compart- ment of the shock tube the total pressure of the re- active gas reaches a few hundred millimeters of mer-cury and the temperature ranges from 1200" to 3500"~.Measurements of these mixtures by the different methods revealed that the reaction has a well-defined induction period. The end of this period is shown in the mass-spectrometric data8 by the appearance of reaction products namely water and carbon monoxide as well as diacetylene the last PROCEEDINGS reaching a stationary concentration.The end of the induction period is also indicated by a decrease in gas density manifested by the lowering of X-ray abs~rption.~ This indicates that an exothermic re- action is taking place. At the same time radiation begins to be emitted in measurable intensity and gaseous ions begin to be formed. The intensity of radiation was found to increase strictly exponentially with time both in the visible and the far-ultraviolet region of the spectrum. After reaching its maximum intensity the radiation rapidly decreases virtually to zero although on the evidence of the mass spectro- meter substantial concentrations of diacetylene still exist in the gas phase.The ion concentration in the gas also increases exponentially the time-dependence being identical within experimental error with that found for different frequencies of the electromagnetic radia- tion. However when the maximum is passed the ion concentration decreases to zero less rapidly accord- ing to a bimolecular rate law. This means that prob- ably no new ions are formed when the emission of radiation ceases and that those already in the gas gradually recombine. It was observed that to a very good approxima- tion the product of (i) the time constants for the exponential functions describing the rise in radiation intensity and ionisation and (ii) the oxygen concen- tration is dependent on temperature only being approximately independent of acetylene concentra- tion and completely independent of argon concentra- tion.Over a wide temperature range from about 1500" to about 2500"~,the exponential time constant can be represented by the equation log,o(~[O,])(molesec.l.-l) = -11-5 + 17,100/4.587' where T is the temperature in This equation is OK. essentially identical with that obtained by Schott and KinseylO for the time constant of the exponential rise of the concentration of free hydroxyl radicals in the hydrogen-oxygen reaction occurring in shock waves at high temperatures. In fact at 2000"~,about the middle of our temperature range the two equa- tions are identical and the slight differences at other temperatures are well within the combined experi- mental errors.This observation as well as the forma- tion of diacetylene suggested to us that the oxidation of acetylene proceeds by a branching-chain mechan-ism similar formally to that proposed by Schott and Kinsey for the hydrogen-oxygen reaction and with the same reaction as the rate-determining process Kistiakowsky J. Chem. Phys. 1951 19 1611. 'Kistiakowsky and Richards J. Chem. Phys. 1962,36 1707. 'Hand and Kistiakowsky J. Chem. Phys. in print. Bradley an! Kistiakowsky J. Chem. Phys. 1961 35 254. * Gardiner j~n.,J. Chem. Phys. 1961 35,2252. lo Schott and Kinsey J. Chem. Phys. 1958 29 1177. SEPTEMBER 1962 kl (1) H + 0 -+ OH + 0 AH = 16 kcal.mole-' k2 (2) 0 + C2H -f OH + C,H AH I20 (?) k3 (3) OH + C,H -f OH + C,H AH 5 1 (?) k* (4) C2H + C,H + CIH + H AH 5 -23 (?) H + 0,+ 6C2H2 -+ 3CdH2 + 2H2O + 3H The net result of this branching reaction cycle is that one atom of hydrogen reacting first with oxygen eventually produces diacetylene water and three atoms of hydrogen thus ensuring branching. The solution of appropriate differential equations leads to the following expression for the time con- stant 7 of the exponential growth in the conceatra- tion of intermediates:' 7 = 1/2k,[02] if the assurr.ption is made that reaction (1) is completely rate-determining. A more refined calculation assum- ing that reaction (2) is of comparable speed to reaction (l) whereas reactions (3) and (4) both of which are thermoneutral or exothermic are much faster leads to a more complicated expression which indicates that in the second approximation the ex- ponential time constant 7 should have a slight dependence on acetylene concentration as observed.The reference to slower reactions should not be taken too literally. Under typical experimental con- ditions the entire time from the compression of the gas to the essential completion of the above reaction is so short that an average molecule of one reactant suffers only some ten-thousand collisions with mole- cules of the other reactant. The time constants 7are such that the concentrations of reaction intermediates double in tens of microseconds.The induction period that is the time between the shock compression of the gas and the first evidence of chemical reaction is to be thought of as the time required for a few chain carriers probably produced by random thermal activation to multiply by the mechanism defined above to such concentrations that the rate of reaction becomes measurable. Experimental observations show that the duration of the induction period in acetylene oxidations with a high degree of consistency is equal to ten times the exponential time constant 7.This means that the concentrations of the intermediates increase by about four powers of ten during the induction period. The experiments of course do not reveal the nature of the spontaneously formed intial chain carriers.There is a further rise of concentration of intermediates during the observable part of the reaction by about two orders of magnitude so that altogether the branching-chain reaction results in something like a million-fold rise of concentration of the free radicals involved in the chain. Callear and Norrish Proc. Roy. Soc. 1960 A 259 304. 291 A great deal of work has been described on the action of anti-knock compounds in internal com-bustion engines and in oxidations of hydrocarbon under laboratory conditions.ll At the comparatively low temperatures of these experiments tetraethyl-lead retards the reaction and reduces branching. We have also carried out experiments with traces of tetra-methyl-lead and acetone added to argon-acetylene- oxygen mixtures.Under experimental conditions tetramethyl-lead and acetone reduce very sub-stantially the duration of the induction periods without having a significant effect on the exponential time constants of the subsequent reaction. We inter- pret this as meaning that methyl radicals formed within very few micro-seconds by the thermal dis- sociation of these compounds act as chain initiators e.g. by CH + 0,= CH,O + 0. On the other hand the chain-breaking reactions at lower tempera- tures due probably to oxides of lead are ineffective at the high temperatures of the present experiments because they are followed by other reactions re-forming chain carriers. The emission of radiation at least in the ultra- violet region in the spectrum and the formation of gaseous ions-the negative ions being mostly free electrons-are undoubtedly in the nature of chemi- luminescence and chemi-ionisation since they exceed by many orders of magnitudes the intensity and the concentration calculated from thermal-equilibrium considerations.However they are still the result of only comparatively unimportant side reactions. Measurements of absolute intensities indicate that approximately a hundred thousand molecules of acetylene are oxidised per photon emitted in the vacuum-ultraviolet range and several million are oxidised per ion pair formed. The exact nature of this side reaction has not been demonstrated but it is made rather probable by the following consideration. The emission of photons and the formation of ions cease rather abruptly whereas according to mass-spectrometric measure- ments diacetylene persists for a longer time.Hence their production must be associated with the re- actions taking place during the branching-chain process outlined above and is not the result of the secondary oxidation of diacetylene. The exponential time constants of radiation growth and ion-con- centration growth are identical under a great variety of conditions. Moreover we have found that the temperature coefficient of the maximum intensity of radiation and of the maximum concentration of ions are identical being equivalent to 23 kcal. of activa- tion energy. This finding strongly suggests that the same rate-determining process leads to photon emis- sion and to ionisation and that therefore the mechanisms of the two must be very similar.PROCEEDINGS Searching for reactions which could provide the is almost certainly the fourth positive system of needed energy at least 230 kcal. in the case of carbon monoxide which further supports the ionisation and somewhat less for radiation we con- proposed reaction scheme. clude that the reaction sequence (5-10) is most likely. The experiments which I have described illustrate C,H + 0,+-CH* + CO,; C2H+ 0 + CH* + CO; CH* (A2dor B2Z) -+ CH (X27r)+ h; co + 03 co,+ hv CH*(A2dor Beg + 0 -f HCO++ e-; CH (X%) + 03 CO (Ah) + H; AH = -75 kcal. mole-1 AH = -70 kcal. mole-’ AH = -66-6 or -774.6 kcal. mole-’ AH = -40 to -50 kcal.mole-’ AH = +I0 kcal. mole-1 CH*(A2Aor B22) + 0-f CO (Ab) + H; AH = -56 to -65 kcal. mole-’ In favour of this scheme it should be pointed out that reactions (6)and (7) the emission of visible and near-ultraviolet radiation are we11 known from the study of flames. Moreover in experiments which Dr. Hand carried out after leaving my laboratory on acetylene-oxygen flames burning in a helium atmosphere and photographed with an ultraviolet spectrograph,12 he observed the long-wavelength portion of the Condon parabola of the fourth positive band system of carbon monoxide (Aln-+ Xlz). The emission of shorter wavelength was obscured in these experiments by absorption. The unobserved branch of the Condon parabola of the fourth positive system lies precisely in that region of the spectrum which is transmitted by lithium fluoride optics and to which our special photomultiplier was sensitive.Thus the emission of vacuum-ultraviolet radiation from the acetylene-oxygen reaction in shock waves l2 Hand J. Chem. Phys. in the press. the information that may be obtained by the special techniques developed for the study of chemical pro- cesses taking place in shock waves. It must of course be realised that the reactions studied occur under very special conditions the reactants are at very low partial pressures and are diluted with a great excess of inert gas. Moreover the temperatures are very high compared with those in studies of other authors. Therefore it is not surprising if the kinetics described here differ from those observed or proposed for re- actions taking place at much higher partial pressures or much lower temperatures.It seems to me how- ever that the conditions chosen for this work help greatly to simplify the kinetics so that the study of complicated reactions under our conditions may pro- vide kinetic information on the basis of which the more complex kinetics involved in other conditions can then be analysed more conclusively. SCIENTIFIC MEETING THE following papers were read and discussed at a Scientific Meeting of The Chemical Society held at __ -Burlington House on June 7th 1962. Reaction of Anthracene with Free Radicals derived from 2,2,4Trimethylpentane (Iso-octane). By A. L. J. BECKWKH.WHENdetermining values of “methyl affinity” for aromatic hydrocarbons Szwarc and his co-workers assumed that C8H,,. radicals produced by hydrogen- atom abstraction from 2,2,4-trimethylpentane (iso- octane) by methyl radicals do not react with the aromatic substrate. The validity of this assumption has now been investigated. The only volatile product obtained from thermal decomposition of di-t-butyl peroxide in an iso-octane solution of anthracene was t-butyl alcohol whose formation indicates that t-butoxy-radicals produced by homolysis of the peroxide react solely by hydro- gen-atom abstraction from the solvent. The absence of volatile olefins suggests that iso-octyl radicals do not readily disproportionate. Of the anthracene used one third was recovered the remainder having been converted into various products.By taking account of the reasonable assumptions that decomposition of one molecule of peroxide yields two iso-octyl radicals and that one or possibly two of the latter are consumed in forma- tion of product from one molecule of anthracene it was calculated that more than half of the iso-octyl radicals formed reacted directly with anthracene. The isolation of two crystalline compounds formu- lated as stereoisomers of 9,9’ 10,10’-tetrahydro-lO,lO’-di-(2,4,4-trimethylpenty1)-9,9’-bianthryl on the basis of their oxidation to anthraquinone and 2,4,4-trimethylpentanoicacid confirms the occur- SEPTEMBER1962 rence of a direct reaction between anthracene and iso-octyl radicals by a mechanism analogous to that proposed for other free-radical reactions of anthracene.A number of compounds were synthesised for comparison with reaction products. Formation of 9,9’,10,10’-tetrahydro-lO,lO’-di-(l,l,3,3-tetramethyl- butyl)-9,9’-bianthryl by treatment of sodium anthra- cenide with 2-chloro-2,4,4-trimethylpentaneaccords with the radical-ion structure of the metal adduct. Other compounds prepared include 2,2,4-and 2,4,4-trimethylpentanoicacid. Dielectric Studies of Carboxyiic Acid-Pyridine Com- plexes in Benzene. By MANSEL DAVIES and LUCJAN SOBCZYK. IN structural terms the systems studied may be represented as X -H ... Y + X-... [H -Y]+. When both forms play a significant rde the energy can be represented by a curve with two minima in each configuration there is a hydrogen bond but the donor-acceptor relationship is interchanged with the transfer of the proton.The equilibrium represents the essentials of the acid-base function presumably in its simplest form and its study in non-polar or aprotic solvents minimises solvent effects. One little- explored feature arises from the appreciable change in polarity with the displacement of the proton; its oscillation might give rise to a dielectric adsorption centred at a frequency characteristic of the process. In the examples studied the simple AB (acid-base) complex is the only species shown by conventional stoicheiometric assessment. The acids were acetic monochloroacetic and trichloroacetic.Dielectric measurements (E’ and E”) were made over the frequency range 5 c./sec. to 1700 Mc./sec. It was deduced that the very small electrolytic con- ductances (h to arise from minute con- centrations of complex ions. The effective dipole moments p (AB) were 2-93~ 4-67~ and 7.78~. These suffice to show that with acetic acid the equi- librium is essentially completely to the left with tri- chloroacetic it is largely (if not completely) to the right. The relaxation times for the dielectric polarisation confirm that the polar species are monomeric rigid- dipole complexes. Over the frequency range covered no evidence is found in these instances for a proton- jumping process. Thermodynamic and Electrochemical Characteristics of the Ionic Dissociation of Triethylmethylammonium Iodide in Methylene Dichioride.By J. H. BEARDand P. H. PLESCH. WE have measured the electrical conductivities of solutions of triethylmethylammonium iodide in methylene dichloride between 0” and -95” in a vacuum apparatus over the concentration range 0.26-4-5 x lo4 molefl. From the results we evaluated A and the equilibrium constant K,and hence AGO AH’ and AS” for the dissociation of ion-pairs to free ions. -log K = 3-63+ 5.20 X lW3 T AGO= 2.379 x low5T2 + 1.661 X T kcal. mole-l -AH” = 2.379 x T2kcal. mole-l -AS” = 4.758 x 10-~T + 1.661 x loL2 kcal. mole-l deg.-l According to current theory (Fuoss Ramsey) the relation between K and the interionic distance a in the ion-pairs is of the form log K = A + B/cTa where E is the dielectric constant of the solvent and A and B are constants.We found that our results are not compatible with this equation unless a increases with increasing ET (decreasing 7‘). The numerical value of the “activation energy” of (1 is about 6% greater than that of viscosity. The average diameter of the ions obtained from the Walden product and Stokes’s law also increases with decreasing tempera- ture. All this evidence indicates that the extent of solva- tion of both free ions and ion-pairs increases as T diminishes. We were unable to obtain reproducible results with solutions exposed to the atmosphere the effect on K being much stronger than that on A,,. We attribute this to traces of water sohating free ions preferentially and thus enhancing the dissociation.Professor M. C. R. Symons in the ensuing discus- sion held that there was evidence (Griffiths and Symons Mol. Phys. 1960 3 90) that when alkyl- ammonium ions are in contact with iodide ions an intense charge-transfer band characteristic of this contact ion pair is found in the 300 mp region. This band is only detected in non-polar solvents such as carbon tetrachloride and the addition of traces of polar solvents results in the loss of this band and the appearance of a band in the 220-250 mp region characteristic of solvated iodide ions. Since the spectrum of such solids in methylene chloride is characterised only by the solvated iodide band he had concluded that there were no contact ion pairs in these solutions.If this conclusion is correct he must assume that the major constituent of these solutions is not the contact ion pair but some form of solvent-separated ion pair. Mr. J. H. Beard and Dr. P. H. Plesch in reply PROCEEDINGS commented on the interest of Professor Symons' persal that the charge-transfer process which gives findings but did not agree that these necessarily im- rise to the absorption in the 300 mp region was plied the existence of solvent-separated ion-pairs. It suppressed. However it seemed likely that whatever was possible that the solvation of a contact ion pair the actual inter-ionic distance in the ion-pairs their had the effect of producing the spectrum of the own findings did indicate that it was not temperature- solvated iodide ion and also so large a charge dis-independent.COMMUNICATIONS The Gas-phase Reaction of Chlorine Monoxide with Hydrocarbons Chlorination by the C10* Radical By L. PHILLIPS and R. SHAW OF AVIATION RESEARCH ESTABLISHMENT, (MINISTRY EXPLOSIVES AND DEVELOPMENT WALTHAM ABBEY,ESSEX) IThas long been known' that chlorine monoxide can results therefore show that chlorine atoms are not react rapidly or even violently with organic materials the main chlorinating agent and the only logical but except for the studies of Goldschmidt and reaction scheme is that annexed here. Schussler,2 who found that chlorination was the main reaction with a variety of organic compounds ClOCl -+ CI.+ CIO. . . . (1) in solution in carbon tetrachloride no information CI. + CH,*CH2.CH -+ HCI + CH3*CHCH3 . . . (2) exists in the literature on these reactions. CI. + CH,CH2.CH3 -+ HCI + CH,.CH,-CH . . . (3) CIO. + CH,-CH2CH -+ By analogy with the well known reactions of HOCl + CH,-CH*CH, AH = -4 kcal. mole-' . . (4) chlorine atoms with hydrocarbons3 it might be ex- CIO. + CH,-CH,*CH -+ pected that similar long-chain chlorination would HOCl + CH,.CH,.CH,. AH = 1 kcal. mole-1 . . . (5) occur in the gas phase with chlorine monoxide and CH3*CH.CH3+ ClOCl -+ CH3*CHCI*CH3+ CIO., hydrocarbons the C10. radical rather than the AH = -46 kcal. mole-l . . . (6) chlorine atom being the chain carrier. Edgcombe CH,*CH,*CH2* + ClOCl -+ CH,.CH,-CH,.CI + CIO. Norrish and Thrush4 have pointed out that the AH = -46 kcal.mole-' . . . (7) C10. radical should be a good chain-carrier because of its high dissociation energy and slow bimolecular Steps (4) and (6) and (5) and (7) constitute non- decay. branching chain reactions in which the C10. radical Preliminary studies on these reactions with mix- is the chain carrier. Since it will be shown that the tures of 12 mm. of chlorine monoxide with 120 mm. chains are very long R(PriCl)/R(PrnC1) = k,/k = of propane in a 450-ml. reactor at 100"c show that 7.2 f0.5 at 100". The C10. radical is therefore (a) no oxygen which would be expected from more selective than the chlorine atom for hydrogen straight pyrolysis of the oxide is formed (b) the abstraction. oxide is consumed far more rapidly than would be The calculated equilibrium constant K for the expected from the results of Beaver and Stieger5 on system C1.+ .OH + HOCl is 5 x atm.-l at straight pyrolysis and (c) hypochlorous acid and 127" so no complications should arise from pyro- n-and iso-propyl chloride are the only detectable lysis of the hypochlorous acid formed in these re- products and the molar ratio of these chlorides is actions. The constancy of molar ratio of the mono- constant at 7.2 & 0.5 up to about 40% decomposi- chloropropanes suggests that reactions R + ClOH tion of the oxide. -RCl + .OH are unimportant in the early stages The results of Knox and Nelson6 on competitive of reaction. chlorination with photolytic chlorine atoms indicate It can be shown from the Slater theory that at the that with propane the molar ratio of the isomeric pressures used the initial step in the pyrolysis of monochloropropanes is unity at 100".The above chlorine monoxide is in the second-order region. Mellor "Comprehensive Treatise on Inorganic and Theoretical Chemistry," Longmans Green & Co. London Suppl. 11 1956 p. 520. Goldschmidt and Schussler Ber. 1925 58 566. Trotman-Dickenson "Gas Kinetics," Butterworths Scientific Publns. London 1955 p. 181. Edgcombe Norrish and Thrush Chem. SOC.Special Publ. No. 9 1957 p. 135. Beaver and Stieger 2.phys. Chem. 1931 B 12 93. Knox and Nelson Trans. Faraday SOC.,1959,55 937. SEPTEMBER 1962 295 From the calculated equilibrium constant Kc of 1-40 x lW9 (concentration in moles cm.g at 100" for the system ka CI + OCI + M + ClOCl + M kb and an assumed probable rate of 10l6 cm.6 mole-2 set.-' for the recombination rate k, the rate kbof pyrolysis of the Cl-OCl bond at 100"in the second- order zone with [MI equal to 112 mm.of propane can be shown to be 3.4 x sec.-l and the half- life 2 x los sec. Since the measured half-life in the presence of an excess of propane is 7 x 103 sec. the chain length is at least lo4. Iredale and Edwards,' in a note on the photolysis of chlorine monoxide with hydrogen reported that the rate of reaction of the monoxide was accelerated with formation of hydrogen chloride. They postu- lated the C10- radical as a chain carrier but preferred the reaction C10. + H23HCI + .OH. It is prob- able that the main reaction is the same as for propane i.e.CIO. + H 4 HOCl + H. He + ClOCl + HCI + CIO. but complications due to photolysis of hypochlorous acid probably occur. (Received June 6th 1962.) ' Iredale and Edwards J. Amer. Chem. SOC.,1937 59 761. The Reactions of Methoxyl Radicals with Methyl Formate; Comparisons between Alkoxyl and Alkyl By J. C. J. THYNNE and PETER GRAY (SCHOOL OF CHEMISTRY UNIVERSITY OF LEEDS) ALTHOUGH correlations of the reactivity of alkyl and alkoxyl radicals have been made and their thermo- chemical basis has been explored,l there have been few quantitative comparisons e.g. between the velo- city constants for such isoelectronic pairs as methoxyl and ethyl. The observations so far rep~rted~,~ have different implications.One investigation2 suggests that methoxyl is considerably more reactive than methyl in abstracting hydrogen atoms from methyl acetate and that the difference arises from the activa- tion energies (5 and 10 kcal. mole-l respectively2). The other,3 concerning hydrogen abstraction from alkanes suggests that methoxyl and methyl have rather similar reactivities. A contribution to these differences may arise from the fact that it is generally difficult when studying hydrogen abstraction by methoxyl to determine what proportion of the total methanol produced comes from reactions 1 (abstraction from a substrate RH) and 2 (autodisproportionation) MeO-+ RH 4 MeOH + R. . . . (1) 2Me0. -f MeOH + CH,O . . . (2) since the precise estimation of formaldehyde which otherwise could be used as the index of reaction (2) presents many problems.A system in which on the other hand the amount of Re produced can be estimated precisely is available when R is the methoxycarbonyl radical C02Me and RH is methyl formate since previous work4 on the methyl radical- induced decomposition of methyl formate has shown that the radical C02Me decomposes quantitatively to carbon dioxide and methyl. Accordingly the abstraction of hydrogen atoms from methyl formate by methoxyl has been studied between 120" and 185"c with thermal decomposi- tion of gaseous dimethyl peroxide as source of methoxyl radicals. The course of induced decom- position can be represented by the mechanism (3-8).Me,O -f 2Me0. . . . (3) MeO-+ HC0,Me + MeOH + -CO,Me . . . (4) CO,Me -f CO + Me-Me. + H-CO,Me -f CH + CO,Me MeO. + Me -+ CH + CH,O MeO. + Me. -f Me,O 2Me. 3 C2H . . . (40) . . . (5) . . . (6) . . . (7) . . . (8) The carbon dioxide produced by reaction (44 was used as a measure of reaction (4) and the methoxy- radical concentration was related to the rate of formation of dimethyl ether. Interpreting the results obtained in the light of this mechanism leads by two independent routes to values for the velocity constant and Arrhenius parameters of the hydrogen abstraction (4). If values5 for log, (k,; mole-l ~111.~ sec.-l) and log (ks; mole-l cm? sec.-l) of 14 and 13.36 respectively are used measurements of carbon dioxide yield log (k4;mole-l cm.3 sec.-l) = (13.01 f0.10) -(8.18 f0*48)/2*303RT and measurements based on methyl-radical forma- tion (as shown by methane ethane and dimethyl ether) yield log (k4;mole-l cm.3 sec.-l) = (13.13 f0.20) -(8.46 f0.49)/2*303RT.Gray and Williams Chem. Rev. 1959 59 239. Wijnen J. Chem. Phys. 1957 27 710. Shaw and Trotman-Dickenson J. 1960 310. Thynne Trans. Faraday SOC.,1962,58 676. Shepp,J. Chem. Phys. 1956,24,939. At 182"c these expressions lead to values for log k of 9-07 and 9.05 respectively. The corresponding value4 for hydrogen abstraction by methyl radicals (log k5)is 6.59. For ethyl radicals attacking ethyl formate log k is about 7-18 at 182". It is therefore clear that methoxyl radicals are considerably more reactive than methyl radicals in hydrogen abstraction PROCEEDINGS from methyl formate.Since however the activation energies E4 and E5 are 8.2 and 8.6 kcal. mole-l respectively it is apparent that the different re-activities arise from different A factors. we are to D.S*l*R.for a research grant* (Received July 6th 1962). A Direct Comparison of Proton Availability in Dual-acid Systems By D. P. N. SATCHELL and J. L. WARDELL OF CHEMISTRY LONDON, (DEPARTMENT KING'S COLLEGE W.C.2) THAT Lewis acids of the metal halide type will strengthen Brransted acids by co-ordination is known:' what is not known with certainty is which particular combinations provide the most acidic systems from the Brransted viewpoint.Hydrogen-exchange studies previously led us2 to the generalisation that in the absence of steric effects those Brransted acids whose corresponding anions carry the greatest electron density on the atom co- ordinated with the metal provide the most acidic systems and therefore make the best co-catalysts. Thus we found an order of co-catalytic efficiency CH3CO2H > CH,Cl.CO,H > CHCl,.CO,H > CFSC02H. However conclusions from cationic-polymerisa- tion studies3 are not always in accord with this generalisation and it appeared important to devise a direct measure of proton availability in dual-acid systems. We now find that for solutions of stannic chloride sorption measurements. Lewis acids alone also affect the absorption of indicators often in a manner similar to that of Brransted acids? Fortunately in the present case the effects produced by the concentra- tions of stannic chloride necessary for our purpose are sufficiently small to permit easy study of the appreciable further reductions in absorption found on addition of the different Brsnsted acids whose behaviour we wish to compare.Indeed for p-nitro- diphenylamine the effect of stannic chloride alone was almost negligible. Results with this indicator are in the Table. Similar behaviour was found at other (lower) chloride concentrations and also with 4-chloro-2- nitroaniline. In calculating [BH+] :[B] allowance has been made for the effect on the indicator of the Brernsted acid when present alone i.e.the observed enhancement of ionisation is recorded. However except for trifluoroacetic acid the .effect of the Brransted component alone was negligible. Effect of increasing concentrations of diflerent Brmsted acids on the ionisation of p-nitrodiphenylamine (B) in o-dichlorobenzene in the presence of 8.6 x 10-3M-stannic chloride at 21" I2". 102[Acid](M) 1.0 2.0 4.0 CH,-CO,H 0.19 0.25 0.46 CH,Cl-CO,H 0.18 0.21 0-26 CHCl2*CO2H 0.17 0.19 0.23 CF3*C02H 0.14 0.16 0.20 in o-dichlorobenzene-a solvent with a dielectric constant (-10) common in studies of co-catalysis -the ionisation of basic indicators may be used for this purpose. 4-Chloro-Znitroaniline and p-nitrodiphenylamine have absorption maxima at 410 and 385 mp respec- tively in o-dichlorobenzene.As in aqueous solutions the addition of sufficient of a suitable Brsnsted acid reduces this absorption to negligible proportions the conjugate acids formed absorbing at shorter wave- lengths. At intermediate concentrations the ionisa- tion ratios pH+] :PI,may be calculated from ab- E.g. Satchell J. 1958 3910. Satchell J. 1961 1453 3822. Pepper Quart. Rev.,1954 8 88. 6.0 8.0 0.68 0.78 i 0.31 0.36 0.28 0.30 ) [BH+l:IB1 0.22 0.23 J At least four conclusions follow from these results (i) All the Brransted acids are strengthened by the presence of stannic chloride. (ii) For Brsnsted acid- Lewis acid concentration ratios < 10:1 the most acidic system is provided by CH3C0,H-SnC14 and the least by CF,CO,H-SnCl mixtures.(iii) Our previous generalisation is thus certainly correct for this class of Brernsted acids with stannic chloride. (iv) Discrepancies with polymerisation studies must arise because the Brernsted component there plays other roles (e.g. in chain termination) besides that of co-catalyst. (Received,July 17th 1962.) E.g. Shuba and Zenchelsky J. Amer. Chem. SOC.,1960 82 4136. SEPTEMBER 1962 297 Rotational Analysis of the A 3n0+,u -X lCp+System of the Chlorine Molecule By W. G. RICHARDS and R. F. BARROW CHEMISTRY OXFORD (PHYSICAL LABORATORY UNIVERSITY) A STUDY of the published analysis’ of the visible 11111111l111111 absorption bands of chlorine reveals that some of the spectroscopic constants for this molecule are very t -1 imperfectly known.In particular the vibrational con-stants of the ground state are uncertain since only a rough estimate of the anharmonicity constant x&z is available and the numbering of the vibra-: tional levels in A 3n0+ is also in doubt. We have now photographed the spectrum in the -,5 18-- -region 4700-6000 A with a second order of a 6.5m. grating spectrograph using absorption paths from 1 to 8 m. at pressures from 20 to 60 cm. The rota-16-v’-l -tional analysis presented no unusual problem and is in general agreement with the previous analysis. The l1 I ’ I’4 IIIIII 111 )I 11 1 I I I I 554rr A is 5521 550 I I I I I I I I I I I !I 500 1000 J(J + I) FIG.1. dR(J) dP(J) = dG -aJ(J + 1) plotted against J(J + l) for vn = 0,l (curve A) and 1,2 (curve B)for the ground state of 35C12.derivation of the vibrational constants for the ground state is illustrated in Fig. 1 where R(J)d,v*-R(J)U*,Vn+l = P(J)v*,vn -P(J)tJ*,zI’+l = dG,*,,~+1-CW”J(J + 1) is plotted against J(J + l) for =Cl,. The slopes give 01’’ = 0.0015 cm.-l and the intercepts are AG;, = 554.36 dG;’,2 = 548-95 cm.-l. Thus G,” = 559*78(v+ 4) -2*70,(v + &)2 significantly different from the previously accepted value Gi = 564.9(~+ 4) -4*0(~ + &)2. The determination of the vibrational numbering in The expression for G and for B for A3&+ may be un-the upper state is illustrated in Fig. 2 in which the reliable outside the limits of values of v’ so far studied vibrational shifts between 35C135C1and 35C137C1are i.e.at 6 > vf < 22. compared with calculated shifts for different as-(Received,June 13th 1962.) Elliott Proc. Roy. Soc. 1929 A 123 629; 1930 A 127 638. a Herzberg “Molecular Spectra and Molecular Structure. Part I. Spectra of Diatomic Molecules,” van Nostrand, New York 1950. Rosen “Donnkes Spectroscopiques concernant les Molkules Diatomiques,” Hermann Paris 1951. PROCEEDINGS The Mechanism of Nucleoside Synthesis Pyrimidine 03 N-Glycosyl Rearrangement By T. L. V. ULBRICHT (TWYFORD Lm. LONDON,N.W. 10) LABORATORIES CONSIDERATION of the mechanism of dealkylation of alkoxypyrimidines and of 0-N-alkyl rearrangement in such compounds led to the suggestion1 that it might be possible to rearrange 0-glycosyl-to N-glycosyl-pyrimidines and with other evidence to a proposal2 that reaction of many pyrimidine mer- cury salts with halogeno-sugars initially gives an 0-glycoside which rearranges to an N-glycosyl compound in the presence of mercuric salts.Some recent experiments are consistent with this. N-AcetyIcytosine3 gives a silver salt which may be condensed with acetobromoglucose to give a pro- duct (I) which is formulated as an 0-glycoside on the basis of its ultraviolet spectrum (Amax. 230 273 Amin. 245) positive Fehling's test and liberation of N-acetylcytosine with dilute acid at room tempera- ture. With mercuric bromide in hot toluene this pro- duct rearranges in 55% yield to N3-/?-~-glucosyl- cytosine penta-acetate4 (IV) Amax.250 299 Amin. 277. If mercuric chloride is used in place of mercuric bromide the 0-glycoside (I) is recovered unchanged,. which suggests that failure to obtain an N-glycoside from a pyrimidine mercury salt may in some cases be due to the use of a glycosyl ~hloride.~ The glyco- side (I) is also unaffected by sodium iodide in acetonylacetone or lithium bromide in dimethyl- formamide indicating that the formation of a com- R = 2,3,4,6-Tet ra-0-acet y I-D-gI ucosy1 plex (11) is probably an essential step in the sequence (I-IV). The glycosyl halide may not be present as a free entity as indicated for simplicity in (111). From the known stereochemistry of nucleoside synthesis from metal salts and acylated glycosyl halides,2 compounds (I) and (IV) can be formulated as 19-nucleosides.(Received July 30rh 1962.) Ulbricht J. 1961 3345. Ulbricht Angew. Chem. (Internat. Edn.) 1962 1 in the press. a Brown Todd and Varadarajan J. 1956 2384. Kindly supplied by Dr. I. Wempen (Fox Yung Wempen and Doerr J. Arner. Chem. SOC.,1957,79 5060). Hoffer Chem. Ber. 1960,93 2777. The Reactions of Corymine an alkaloid isolated from Hunteria corymbosa (Apocynaceae) By A. K. KIANGand G. F. SMITH DEPARTMENTS OF SINGAPORE OF MANCHESTER) (CHEMISTRY UNIVERSITY AND UNIVERSITY THEleaves of Hunteria coryrnbosa contain 0.57 % of total basic material chromatography of which on alumina has yielded a new alkaloid corymine CZ~H~~N~O,, m.p. 189-192" $-27.3" (in CHCI,) (0.076%based on dry leaves).It is a mono- acidic base pKa 7.86 forming the following cry- stalline derivatives B,HNO, m.p. 204-205"; B,C6H3N30, m.p. 136-137"; B,CH,I m.p. 188-190";and a mono-0-acetate m.p. 159-1 61 *. It does not reduce Fehling's solution but reduces Tollens' reagent. The infrared spectrum reveals an ester car- bony1 group at 1725 cm.-l and a hydroxyl group at 3 150 cm.-l. The nuclear magnetic resonance spectrum in deuterochloroform shows the presence of the following groups methoxyl (6.18 r) N-methyl (7.28 T) and ethylidene (a doublet centred at 8.35 T and a quartet centred at 4.62 r,J = 6-5 c./sec.). Corymine methiodide is decomposed by hot 1% aqueous potassium hydroxide to deformylcorymine methine C22H28N203 m.p. 159-160". Under the same conditions corymine is converted into de- formykorymine C21H26N203, m.p.170-1 72" the methiodide of which is also decomposed by hot 1 % aqueous potassium hydroxide into deformylcory- mine methine. This methine is very weakly basic having PKa -3-0.Its infrared spectrum shows a new carbonyl absorption at 1690 cm.-l in addition to the ester-carbonyl absorption at 1730 cm.-' and absence SEPTEMBER 1962 of hydroxyl absorption. The methine is not reduced by sodium borohydride and does not react with methyl iodide at room temperature. Reduction of corymine with lithium aluminium hydride in ether gives a triol C21H2,N203 m.p. 241 -243 O which forms an isopropylidene derivative m.p. 230-234" that still contains one hydroxyl group having vmax.(in Nujol) 3550 cm.-l and is hydrolysed back to the triol by dilute hydrochloric acid. Reduction of corymine by borohydride gives a dihydrocorymine C22H,8Nn04 m.p. 220-222" which however still reduces Tollens' reagent. hemiacetal grouping for the corresponding carbonyl absorption is absent from the infrared spectrum. This is confirmed by the nuclear magnetic resonance spectrum in which the hemiacetal CH is seen as a band at 4.86 r. The persistent survival of the aryl-N-C-N system in so many derivatives of corymine suggests a struc- tural analogy with echitamine2 (VII). This leads to structure (I) which we propose for corymine. The analogy with echitamine allows a simple rationalisa- tion of the hemiacetal system and of the properties of deformylcorymine methine which can be given The ultaviolet spectrum of corymine Amax 258 structure (VI) and is thus simply N(a)-methylallo- This finds strong support from the 3 14 mp in EtOH (log E 4.99,4.50) Amax.245,299 mp e~hitamine.~~~ HCl and those of mass spectra of the two compounds which are very (log E 540,4.51) in 0*2~-ethanolic all the basic degradation products described above closely similar the only differences being due to the with the exception of the methine are characteristic extra N(a)-methyl group in (VI). The very weakly of the aryl-N-C-N chromophore.l basic N(b)and anomalous carbonyl group of (VI) are The very easy loss of the elements of carbon mon- the result of transannular interaction of the type first oxide speaks for the presence of a p-formyl ester observed in ~ryptopine.~ group which is supported by the conversion of the Since the specific rotations of alloechitamine lithium aluminium hydride reduction product (111) [a],-2Oo (in EtOH) and of deformylcorymine into the isopropylidene derivative (IV).The aldehydic methine [01],-21 (in EtOH) are almost identical O function in corymine must be largely involved in a and of the same sign and since the absolute con- C0,Me * Me HO Me HO (V) I it (a t2 CO Me C0,Me HO-/ Me Me K HO.H,C 50,Me --t *c,-3 HNH OH Reagents (1) Mel then aq. KOH. (2) NaBH,. (3) COMe,-H,O.+ Hodson and Smith J. 1957 1877. Hamilton Hamor Robertson and Sim Proc. Chem. Soc. 1961 63. Conroy Bernasconi Brook Ikan Kurtz and Robinson Tetrahedron Letters 1960 No.6 1. Birch Hodson Moore and Smith Proc. Chem. SOC.,1961 62. Anet Bailey and Robinson Chem. and Ind. 1953 944. 300 PROCEEDINGS -~ ~ figuration of echitamine has been determined by Research Institue Kepong Mr. E. L. Tan Mr. J. X-ray crystallography6 structure (I) represents the Goh Dr. J. N. Shoolery and Dr. W. I. Taylor; and absolute configuration of corymine. one of us (A.K.K.) thanks Professor A. J. Birch for the hospitality of his laboratory. We are indebted for various services to Professor K. Biemann the Chief Research Officer Forest (Received August lst 1962.) Manohar and Ramaseshan Tetrahedron Letters 1961 814. Metal-Metal Conjugation in Group IVB Organometallic Systems By D.N. HAGUE and R. H. PRINCE (UNIVERSITY LABORATORY, CHEMICAL CAMBRIDGE) THERE is a close similarity between the electronic spectra in cyclohexane solution of many compounds of the type Ph,M.X (M = Si Ge Sn or Pb; X in-cludes Ph,l OH F C1 and Br) in the region 230-270 mp. The benzenoid fine-structure typical of a &band2 is observed and for a given element M the values of Amax are approximately constant. It is found that molar extinction coefficients (which lie between about 500 and 1200) are approximately additive indicating3 that any interaction between the phenyl rings and the group X is insufficient to affect the transitions responsible for the ultraviolet spectra. When X is MPh, however a very intense K-band2 is observed (cf.Table and Figure) which effectively masks the fine structure typical of a phenyl group.4 When X is iodine and M is tin or lead a shoulder is observed of medium intensity. It appears that in the compounds Ph,M2 there is intense interaction between phenyl groups on different M atoms across the metal-metal bond Ph6Siz Ph6Ge2 Ph6Sn2 Amax 246.5 239 247-5 E 32,600 30,400 33,900 x 24 0 I I I Si GC Sn Pb P As Sb Bi / x-x Ph,M (M = Group IVB). -Ph,MJM = Group VB). Rao et a1.' that compounds of the type Ph4MCl Ph3MCI2 etc. have spectra of the same form as those of Ph3MX (M = Group IVB; X = Ph OH etc.) i.e. show typical benzenoid bands at about 260 mp. Compounds in the series Ph,M however Ph6Pb2 Ph,SnI Ph,PbI 245* 293 245* 278* 33,900* 35,000 6610* 5660* * = Shoulder.(there being good evidence5 against the formation of radicals of the type Ph,M.). It seems likely that this unusual type of conjugation arises through overlap of suitable vacant d-orbitals on the M atoms. With the iodides the interaction occurs either with the lone-pair electrons of the iodine atom or through the overlap of iodine d-orbitals with those of M. It is interesting to compare these systems with those of the Group VB organometallic compounds Ph3M and Ph,MCl. It has been found by Jaffe6 and have fairly strong bands with no fine structure. The absence of fine structure is attributed to interaction between the non-bondedp-electrons on M and the n-orbitals of the phenyl rings.Evidently this cannot be the explanation for the observed strong bands in the spectra of compounds Ph6M since there are no suitable non-bonded p-electrons on M. However it is interesting that the same general trend in Amax values is observed in both series. (Received July 26th 1962.) Cf. La Paglia J. Mol. Spectroscopy 1961 7 427 and refs. therein quoted. 2 Nomenclature as in Braude Ann. Reports 1945 42 105. Cf. Fehnel and Carmack J. Amer. Chem. SOC.,1949,71 88. Cf. spectrum of trans-stilbene Jerchel and Melloh Annalen 1959 622 53. ti Hague and Prince unpublished work. Jaffk J. Chem. Phys. 1954 22 1430. Rao Ramachandran and Balasubramanian Canad. J. Chem. 1961 39 171. SEPTEMBER 1962 301 Dalbergiones A New Group of Natural Products By W.B. EYTON, W. D. OLLIS,I. 0.SUTHERLAND (UNIVERSITY OF BRISTOL) L. M. JACKMAN (IMPERIAL LONDON, COLLEGE S.W.7) 0. R. GOTTLIEB, and M. T. MAGALH~ES DE QU~MICA MINIST~FUO (INSTITUTO AGR~COLA DA AGRICULTURA Rro DE JANEIRO,BRAZIL) examination of the substances J-1 and 5-2 bergiones and their quinol diacetates with summation FURTHER isolated,' in addition to caviunin,1,2 from the heart- curves corresponding to the individual chromophores wood of Dalbergia nigra Fr. Allem. has yielded three indicated in the formula showed that the two new natural products.* Compound J-1 is 4"-hydroxy- chromophores were not conjugated with each other m.p. 172-1 78 O [a] or with an olefinic double bond. dalbergione Cl,~1103~OMe (in dioxan) -51 ;5-2 is an equimolecular mixture Since these quinones and the quinol diacetates are of dalbergione CISHl,O,-OMe m.p.114-1 16" optically active the [C,H,]-Ar residue must be [a] (in CHC13) + 13" (in dioxan) -51" and 1 -arylallyl (cf. 111) or 2-arylcyclopropyl but the latter 4"-methoxydalbergione CISH,,O,(OMe), m.p. was excluded because the nuclear magnetic resonance 109~5-111",[a],(in CHC13) -138" (in dioxan) spectra of the three dalbergiones and their quinol -32". diacetates were completely compatible with the The colour and ultraviolet and infrared spectra of structures (111) and (IV) respectively. These spectra the dalbergiones indicated a quinonoid structure and also established the location of the substituent on comparison of the spectra with those of other ring B and the para-relation of the two protons on quinones4p5 showed the presence of a 2-alkyl-5-ring A.methoxybenzoquinone chromophore. With perman- Catalytic hydrogenation of the quinones gave ganate dalbergione gave benzoic acid and 4"-tetrahydro-derivatives which yielded optically active methoxydalbergione gave p-anisic acid in accord dihydro-derivatives (V) on oxidation by air in with the partial structures (Ia and b respectively). aqueous potassium carbonate. The ultraviolet spectra No corresponding oxidation product from 4"-of corresponding pairs (HI and V) confirmed the hydroxydalbergione was formed which coupled non-conjugation of the olefinic group with either with active hydrogen determination (1H) and its chromophore. infrared spectrum [v(OH) 3470 cm.-l] suggested the Ozonolysis of the dihydro-derivative (Va) gave biogenetically acceptable partial structure (Ic).(-)-a-ethylphenylacetic acid (negative plain optical rotatory dispersion curve in CHCl& of the estab- MeOfiOH lished configuration (Via) whereas the dihydro- derivative (Vb) similarly gave (+)-a-ethyl-p-meth-oxyphenylacetic acid the enantiomer of (VIb) (posi-tive plain curve). Throughout for (a) (b) and (c) X=H OMe and OH "'OW 0 respectively. AcO \ Mam The presence of quinone residues in the structures (Ia b and c) was confirmed by reduction (Na,S,O,) to the colourless quinols (IIa b and c) characterised Meom as diacetates and oxidised without racemisation to 0 the quinones by air in aqueous potassium carbonate.Comparison of the ultraviolet spectra of the dal- "Dr. Gonwlves de Lima has also reported the isolation of these compounds from Dalbergia nigra; he has studied their antibiotic activity and suggested that they be called dalbergi~nes.~ Gottlieb and Magalhiies J. Urg. Chenz. 1961 26 2449. Dyke Ollis and Sainsbury J. Urp Chem. 1961 26 2453. 0. Gonqalves de Lima and M. H. Dalia Maia Lecture at the 3rd Annual Meeting Associaqiio BrasiIeira de Quimica (Rio de Janeiro November 6-10th 1961). Braude J. 1945 490. Yates Ardao and Fieser J. Amer. Chem. SOC., 1956 78 650. Levene Mikesha and Passoth J. Biol. Chem. 1930 88 40; Mislow and Heffler J. Arner. Chem. Soc. 1952 74, 3668; Sjoberg Acta Chem. Scand. 1960 14 273. 4“-Me t hoxy- and 4 “-hy droxy-dalbergione (I11 b and c) had complex optical rotatory dispersion curves which showed Cotton effects that corresponded closely; they differed appropriately from the curve of dalbergione (IIIa).This established that dalbergione had the absolute (R)-configuration (VIIa) whereas 4”-methoxy- and 4”-hydroxy-dalbergione had the (S)-configuration (VIIb and c). The isolation from a single plant of several compounds with biogenetically equivalent structures but opposite absolute stereo- chemistry is very unusual among natural phenolic compounds. Compound 5-2 is thus a quasi-racemate of (VIIa) and (VIIb). The dalbergiones (VII) and caviunin (VIII),19z which have been isolated from Dalbergia nigra have some corresponding structural features and the rela- tionship of the dalbergiones to the 4-arylcoumarins dalbergin (IXa) and 0-methyldalbergin (IXb) which were both isolated’ from the generic relative Ahluwalia and Seshadri J 1957 970.Grisebach and Ollis Experientia 1961 11 4. PROCEEDINGS Dalbergia sissoo is particularly striking and could well be of biosynthetic significance? Ph (IXa :R=H b ;R=Me) Resorcinol monomethylether and cinnamyl bromide gave the corresponding ether. This ether was rear- ranged thermally to the phenol (X) which by oxida- tion with Fremy’s salt gave racemic dalbergione (IIIa). We thank Dr. S. F. Mason (University of Exeter) for generously providing us with the facilities for the optical rotatory dispersion studies. Two of us (O.R.G. and M.T.M.) are indebted to the Conselho Nacional de Pesquisas Brazil for financial aid.(Received July 4th 1962.) Hydrogen Bonding in Crystalline 4,4’-DihydroxythiobenzophenoneMonohydrate By LJ. M. MANOJLOVICH* and I. G. EDMUNDS (PHYSICS DEPARTMENT MANCHESTER, COLLEGE OF SCIENCE AND TECHNOLOGY 1) WHEREAS thiobenzophenone and most of its deriva- tives are blue to green crystalline 4-hydroxythio- benzophenone and 4,4’-dihydroxythiobenzophenone both first described by Brocklehurst and Burawoy? are red. Brocklehurst2 suggested that the red colours are probably due to comparatively strong hydrogen bonds between a phenolic hydroxyl group and the sulphur atom of a neighbouring molecule H0‘C6H4>= S*.’HO*C6H,.CS*C6H4X X.C6H4 where X = H or OH.Since we could find no re- corded experimental evidence for a strong hydrogen- bonding between sulphur and oxygcn we determined the crystal structure of 4,4‘-dihydroxythiobenzo-phenone monohydrate from two-dimensional X-ray diffraction data of the (100)and the (010) projection. The compound crystallises in the monoclinic space group Czh-P2,/c and has four molecules in a unit cell of dimensions a = 5-62,b = 10-95,c = 20.24 8, and 18 = 103.5”. The structure consists of con-tinuous chains of the organic molecules in which the molecules are linked by hydrogen bonds 0 -H-0 between phenolic hydroxyl groups but the expected hydrogen bond S * -H-0 is not present. Hydrogen bonds from a water molecule to a phenolic group in each of two organic molecules link the chains to- gether.The sulphur atom of a third organic molecule is distant 3.42 & 0.03 8,from the oxygen atom of the water molecule and lies in such a direction that a hydrogen bond between these atoms would conform to the angular requirements given by Donohue3 and F~ller.~ It is therefore likely that a weak hydrogen bond exists between these two atoms although our two-dimensional X-ray results cannot locate the hydrogen atom within this bond with any certainty. We are much indebted to the late Dr. A. Burawoy for drawing our attention to this problem and Professor H. Lipson for his interest. (Received July 23rd 1962.) * Present address Institute of Nuclear Sciences “Boris Kidrich,” Belgrade Yugoslavia. 1 Brocklehurst and Burawoy Tetrahedron 1960 10 118.Brocklehurst Ph.D. Thesis University of Manchester 1956. s Donohue J. Phys. Chem. 1952,56,502. Fuller J. Php. Chem. 1959,63 1705. SEPTEMBER 1962 303 Ease of Displacement of Thiol and Oxide Anions from Methyl Groups Carbon Basicities of Anions of Oxygen and Sulphur By BERNARD MILLER (CHEMICAL LABORATORIES DIVISION, RESEARCH AGRICULTURAL AMERICAN Co. PRINCETON, CYANAMID NEWJERSEY U.S.A.) IN contrast to the mass of data on the activities of and sulphur since the higher polarisability of the oxide and thiol anions as nucleophilic no sulphur atom should decrease the net charge separa- quantitative studies have been conducted on the re- tion in the transition state2 and thus make a thiol verse reactions to determine the relative ease of anion a better leaving group than an oxide anion.elimination of thiol and oxide anions in SN2 Although the reverse reactions of phosphate and reactions. thiophosphate anions with methyl phenyl sulphide The use of thiophosphate and phosphate anions as do not proceed to a measurable extent we attempted leaving groups seemed to us to have several virtues to obtain some idea of their relative rates by gas- (a) the basicities of thiophosphate and phosphate chromatographic analysis of the products of reaction anions are almost identical;3 (b)the very low p&’s of of the monothiophosphate anion with methyl iodide the conjugate acids3 suggest that the anions will be in ethanol dimethyl sulphate in ethanol and methyl good leaving groups; and (c) the use of isomeric iodide in benzene.Although as little as 0.25 % of the S-and 0-alkyl monothiophosphates offers an oppor- 0-alkyl derivative (11) would have been detected tunity to study the displacement of sulphur and only the S-alkyl derivative (111) was observed in all oxygen groups which give rise to identical anions. three reactions. Methylation of sulphur in the mono- Methyl esters (I-IV) were demethylated by thiophosphate anion therefore proceeds at least 250 sodium thiophenoxide in 95% ethanol. The disap- times as fast as methylation of oxygen. pearance of the thiophenoxide anion was followed by Combination of the rates of methylation and de- ultraviolet spectroscopy. The results are listed in the methylation shows that the equilibrium constant for Table.reaction (A) is at least lo4 times as large as that for Ester k (1. mole.-l sec.-l) Rel. rates at 58” pKa of acid3 (in 7% EtOH) (I) (PhO),PO.OMe (11) (PhO),PS.OMe (111) (PhO) ,PO-SMe (IV) (PhO),PS.SMe 2.10 x 9.30 X 8-40 x 3-4 x 10-1 618 lo- 277 10-4 2.5 10-4* 1 1.7 1.9 1.9 1-5 * Estimated from rate at 70”. It wilI be seen that changing from a P=S to a (A) RS-+ MeX + RSMe + X-P =0 structure increases the rate of demethylation (B) RO-+ MeX + ROMe + X-by a factor of roughly two while changing from an S-methyl to an 0-methyl ester increases the rate by reaction (B) when RO-and RS-have similar a factor of 200-300. basicities; that is the “carbon ba~icity”~ of a thiol anion is more than lo4times as great as its “hydrogen The much slower rate of cleavage of the S-Me ba~icity”.~ The very high nucleophilicities of thiol bond compared with the 0-Me bond cannot be ex- anions may be attributed in large part to their high plained as due to better solvation of oxide than of “carbon basicities,” with polarisation and solvation sulphide anions since the identical anion is produced factors of lesser importance.from esters (11) and (HI) nor can it be explained on the basis of polarisability differences between oxygen (Received,July 19th 1962.) Swain and Scott J. Amer. Chem. Soc. 1953 75 141. Miller J. Amer. Chem. SOC.,1962 84 403. Kabachnik Mastryukova Shipov and Melenty’eva Tetrahedron 1960,9 10. * We employ the terminology suggested by Parker (Proc. Chem. Sac. 1961 371) in which “carbon basicity” refers to equilibrium constants for displacements on carbon and “carbon nucleophilicity” to rate measurements.Bunnett and his co-workers (Proc. Chem. SOC.,1961 305) have suggested that “thermodynamic affinity for carbon parallels that for hydrogen.” Since they compared the carbon basicity of hydroxide (pKa for conjugate acid 16) to thiophenoxide ion (pKa 6-7),it seems difficult to draw conclusions from their work. PROCEEDINGS The Structure of the Tdtion State in the Disproportionation and Combination of Ethyl Radicals By M. MATSUOKA A. P. STEFANI, P. S. DIXON and M. SZWARC (DEPARTMENT STATE COLLEGE OF CHEMISTRY UNIVERSITY OF FORESTRY SYRACUSE SYRACUSE UNIVERSITY 10 NEW YORK) VARIOUS workers agree that combination and dis- since some side reactions were eliminated and the proportionation of ethyl radicals in the gas phase range of d(l/T) was increased to 2.4 x Thus have kdis/kcom in the range 0.124.15.1No tempera- the accuracy of the temperature-coefficients found ture-dependence of kdis/kcom was observed i.e.was improved. The ratio C2H4:C4HlOwas used for &is -Emm was claimed to be zero. calculating kdis/kcom and the ratio of { 1 -(C2H4.+ We recently investigated these reactions in the C4H10)/N2) :{ (C,H4 3-C4H10)/N21 for evaluating liquid cage and in the gas phase. The radicals were keec/(kdis + kcom),kesc being the rate constant of produced by photolysis of azoethane and the re- escape from the cage. Gas chromatography was used action was investigated in the temperature range in the analysis of the nitrogen ethane ethylene and -78" to +So.Iso-octane and propan-2-01 were butane mixtures.used as solvent and styrene or 1,1 -diphenylethylene The negative temperature-dependence observed in Disproportionation and combination of ethyl radicals. Cage reaction Solvent Temp. NO. Of kdislkcorn kesc/(kdis -f kcom) expts. Iso-octane -78" 6 0.208 5 0.007 0.171 f0.054 0 3 0.177 f0.005 0.497 f0.033 39 25 4 0.163 f0.005 0.613 & 0.029 9 45 3 0.152 f0.005 0.666 0.029 9 65 3 0.144 0.005 0.795 & 0-032 9 85 4 0.144 f0.005 1.01 f0.060 9 Propan-2-01 -65 4 0.265 5 0.005 0.177 0.020 0 4 0.204 & 0.005 0.220 f0.015 3 65 4 0.178 f0.005 0.550 & 0-020 9 Gas phase Diluent P (mnl.) Temp.No. of kdidkcom expts. 50 -65" 5 0.159 f0.005 co2 co2 700 -65 4 0.162 f0.005 None 0.2 0 2 0.129 f0-005 coz 50 0 4 0.131 f0.005 700 0 4 0-134f0-005 co2 H2O 10-50 40 4 0.119 fO*OO5 as scavenger. The gas-phase studies were carried out our studies is incompatible with the model of head- at Pa, = 0.24.3 mm. carbon dioxide or water to-head collision for combination and head-to-tail vapour (at 40")being used as a diluent. No effect of for disproportionation which was advocated by the diluent pressure upon the kdis:kcom ratio was several workers.2 It may be accounted for by the observed. The results of our studies are given in the model proposed by Bradley3 and independently by Table. Plots of log (kdis/kcom) against 1/Tare linear. Kerr and Trotman-Dickenson4 who suggested that The use of low temperatures was advantageous disproportionation results from the bendingvibrations Ausloos and Steacie Bull SOC.chim.belges 1954 63,87; Brinton and Steacie Canad. J. Chem. 1955 33 1840; James and Steacie Proc. Roy. Soc. 1958 A 244 289; Turner and Cvetanovic Canad. J. Chem. 1959 37 1079; Cerfontain and Kutschke ibid. 1958 36 344; Kerr and Trotman-Dickenson J. 1960 1611. Wijnen and Steacie Canad. J. Chem. 1951,29 1092; McNesky Drew and Gordon J. Phys. Chem. 1955,59,988. Bradley J. Chem. Phys. 1961,35,748. Kerr and Trotman-Dickenson Progr. Chem. Kinetics 1961 1 113. SEPTEMBER 1962 of the “hot” butane molecule formed on linking two ethyl radicals (see diagram). We propose that this process competes with the intramolecular dissipation of the energy of the incipient C-C bond which yields butane.This suggestion is based on the fact that kdig/kcom is independent of the pressure of the added diluent. Neither process requires any activation energy. We wish to elaborate Bradley’s model. The “hot” butane resulting from a successful encounter may be represented by a set of points in a suitable configura- tion space of all the internal co-ordinates. Such a set occupies some finite volume which represents all the permissible configurations of this species but only a part of this volume represents a “compressed” mole- cule of the “hot” butane which may undergo dispro- portionation. We suggest that the potential energy of the configurations forming this part of the volume is lowered by the interaction of the or-radical carbon with the hydrogen of the methyl group (see diagram).Thus the probability of these configurations in- creases with decreasing temperature and thus the observed temperature-dependence follows. More- over any external force compressing the “hot” mole- cule would be expected to exert a similar effect upon kdis/kmm.Such forces are encountered in liquids and therefore kdis/kmm should be higher for a cage than for a gas reaction. Also the higher the internal pres- sure of the liquid the larger should be the ratio. Our results confirm these predictions kd&/kcom is larger in propan-2-01 than in a hydrocarbon solvent and both ratios are greater than that found for the gas- phase reaction.Further the temperature-coefficient increases accordingly E = -0-27 kcal./mole in the gas phase -0.32 in iso-octane and -0.44 in propan-2-01. The solvation of radical ends may be proposed as an alternative hypothesis to account for some of these phenomena. Desolvation of only one end would be required for disproportionation while both ends would need to be desolvated for combination. However this hypothesis would not account for the results obtained in the gas phase. Moreover it would predict that the extent of cage reaction should be greater in propan-2-01 than in iso-octane since solva- tion hinders both combination and disproportiona- tion :the opposite behaviour was observed. Financial support by the Office of Ordnance Research and by the National Science Foundation is gratefully acknowledged.(Received June 29th 1962.) Dichloronitrosylnickel By C. C. ADDISON and B. F. G. JOHNSON (DEPARTMENT OF CHEMISTRY THE UNIVERSITY NOTIINGHAM) WHEN nickel carbonyl reacts with nitrosyl chloride in organic solvents nickel dichloride is the usual product. However if the reaction is carried out in the gas phase (the reactants being diluted with argon) a grey-green powder is deposited. Analysis and physical properties correspond with the formula Ni(NO)Cl, and the ratio of reactants may be varied in wide limits without changing this composition. The compound can be heated to 150” before nitric oxide is evolved. This is the first example of a new class of nitrosyl dihalides of nickel and its isolation emphasises the value of gas-phase reactions as a preparative technique.Its high thermal stability is surprising since the number and stability of metal nitrosyl halides decreases in the order iron cobalt and nickel. Again the thermal stability of the known’ nickel nitrosyl monohalides Ni(N0)X decreases in the order X = I > Br > C1. In such compounds the bonding of the NO group to the nickel atom by three electrons reduces the formal valency state of the Hieber and Nast 2.anorg. Chem. 1940,244 23. Seel 2.anorg. Chem. 1942 249 308. nickel to nickel(I) so that monohalides are the ex- pected compounds in this class and conform to Seel’s series of pseudo-atoms.2 In the dihalides the Ni-NO bonding must differ from that in the monohalides and the chemistry of the two types of compound is quite different.Dichloronitrosylnickel is paramagnetic with Xg = 4.120 x c.g.s. unit indicating two unpaired electrons per nickel atom. In reaction with triphenyl- phosphine in a sealed tube at loo” chlorine is dis- placed; deep-blue needles of the compound (Ph,P),Ni(NO)Cl are formed which show strong NO absorption at 1710 cm.-l. In contrast nitric oxide is evolved when dichloronitrosylnickel is added at room temperature to solvents which are able to act as ligands to the metal atom. The rates of decomposition are in the increasing order EtOH < Me,SO -MeCN < H20. The compound neither reacts with nor dissolves in benzene n-hexane di- ethyl ether or ethyl acetate.No solvent for the com- PROCEEDINGS pound has yet been found and it is probably poly- attempting the preparation of the parent compound meric; the infrared spectrum shows only two bands Ni(NO),Cl, analogous to Pd(NO),C1,.3 in the sodium chloride region at 1835s and 1870s assigned to the co-ordinated NO group in two We are indebted to the D.S.I.R. for a maintenance different environments in the solid. We are now grant to B.F.G.J* (Received August 16th 1962.) * Manchot and Waldmuller Ber. 1926 59 2363. Reactions of Cobalt(n1) Tetrammines with Chromium(I1) By K. D. KOPPLE* and R. R. MILLER OF CHEMISTRY OF CHICAGO, (DEPARTMENT UNIVERSITY CHICAGO 37 ILL. U.S.A.) WE have examined using recorded methods,l the probably accounted for by further stabilisation of the reduction of several carboxylatotetramrninecobalt(II1) acceptor antibonding orbital this stabilisation result- perchlorates by aquochromium(I1).The Table shows ing from a decrease in electron density at the co- that replacement of ammonia by a second oxygen ordinated oxygen on protonation of the carboxylate function increases the uncatalysed rate of electron group. transfer about a hundredfold ;also the diacido-com- The hydrolysis of (I) is also acid-catalysed but it plexes (I)and (11) in contrast to carboxylatopentam- is a negligible side reaction under the conditions of minecobalt(II1) derivatives undergo an acid-catalysed the redox reaction. By following the decrease in reaction. The resuIts [compare (I) and (V)] do not absorption of (I) at 550 mp and 25" we have ob- suggest that the enhanced uncatalysed reactivity of tained for its hydrolysis k = 04004 04011 and the tetrammines is the result of lowered charge on 0.0013 sec.-l at perchloric acid concentrations of 0.1 the oxidant.2 It is more likely that replacement of 0.5 and 1.0 M respectively.Reduction of cobalt(II1) ammines by chromiurn(r1) at 25" and p = 1.0. Complex trans-[Co( NH3)*( OAc) 2]+ cis- [Co(NH3),(02CCH,C02)]+ k(mo1e-' 1. sec.-l) 15 + 50[H30+] a 30 + 20[H30+] a AH^ (kcal./mole) 7.9 6.8 8.5 4-1 dSt (em) -26 -28 -24 -39 [CO(NH,),(OAC)]~+ cis- [CO(NH,),(H,O),]~+ cis- [Co(NH,),(H,O)(OAc) 12+ [Co(NH3)5(H@)13+ 0.18 7 + 8.4[H3O+]-l 47 + 2.8[H3O+]-' 0-5 + 1.57[H30+]-1a 3.5 2.9 14-2 -50 -52 -18 a Activation parameters for the acid-dependent term which include AHDand ASD are the lower set at right.Data from Sebera and Taube J. Amer. Chem. Soc. 1961,83 1785. Data of Zwickel and Taube J. Amer. Chem. SOC.,1959 81 1288. (20° p = 1.2.) ammonia by a ligand producing a weaker field in- The bridging group involved in electron transfer creases the probability of electron transfer by lower- to (V) was determined by ion-exchange chromato- ing the energy of the acceptor antibonding d orbital graphy of the products of reaction between (V) and of cobalt. less than the stoicheiometric quantity of chrom- In both 1 SOand 0.05 M perchloric acid solu- No pK for the conjugate acids of the diacido- ium(~~).~ complexes is available but a maximum value of -1 tions more than 80% of the chromium(II1) was re- is indicated by the linear dependence on acid of the covered as CrOAc2+.Were (V) the trans-isomer reaction rate and an estimate of -2 to -3 may be reaction at the lower acidity would be expected on interpolated from the known pK,'s of water car- the basis of the observed rate law to produce at least boxylic acids their conjugate acids and aquoam- as much Cr3+ as CrOAc2+. It is suggested therefore minocobalt(II1) complexes. The protonated diacido- that (V) is the cis-ion and that the intermediate tetrammines are therefore 1O-lOOO times more re- complex may be doubly bridged. active than their conjugate bases a phenomenon (Received,July 2nd 1962.) * Present address General Electric Co. Research Laboratory Schenectady New York U.S.A.l Svatos and Taube J. Amer. Chem. SOC.,1961 83 4172. In the exchange of aquochromium(n) and fluorochromium(m) ions; CrF2+is somewhat more reactive than cis or trans CrF,+ (Chia and King Discuss. Faraday SOC.,1960 29 109). Orgel "Report of the Xth Solvay Conference," Brussels 1956 p. 289. Linhard and Weigel 2.anorg. Chem. 1949 260 65. Taube J. Amer. Chem. Soc. 1955,77,4481. SEPTEMBER 1962 307 The Stereochemistry of the Transition State of a Displacement at a Phosphoryl Centre By M. GREENand R. F. HUDSON EUROPEAN INSTITUTE GENEVA, (CYANAMID RESEARCH COLOGNY SWITZERLAND) THE configurational change in the displacement re- Hence the observed rate constant k, gives a true action between methyl ethylphenylphosphinate and measure of the rate of exchange.methoxide ions in methanol has been established by The rate of racemisation k, of the optically active comparing the rates of racemisation of the optically ester1 (01% -3-49") was determined by means of an active ester and of isotopic exchange. E.T.L. photoelectric polarimeter (sensitivity ca. For the exchange experiments the [14C]methyl 0~0002") for the same solution of sodium meth- ester was prepared by the action of [14C]methanol on oxide and the same ester concentration (k = 12-46 the chloride b.p. 112"/0.6 mm. n2," 1.5495 prepared 12-35 x sec.-l at 25.0"). The change in con- as follows centration of sodium methoxide in the reaction mix- ture during the reaction was found to be negligible, EtOH PhP(OEt) -+ Etl showing the absence of side reactions (e.g.,attack on PhPCI -+ the methyl group).COCI PhEtPO-OEt-+ PhEtPOCl The rate of racem'sation is almost exactly twice the rate of exchange showing that each act of chemical The rate of exchange k, was measured by distilling reaction proceeds with inversion of configuration. portions of the reaction mixture at 0.1 mm. at known Previous investigations2 have shown that inversion is times and measuring the activity of the methanol in possible for reactions at the phosphoryl group but a liquid scintillation counter (k = 6.23 and 6.06 x in all cases products of very low optical purity were sec.-l at 25.0" [MeO-] = 0.42 mole l.-l obtained. The present results therefore provide the [ester] = 0.62 mole l.-l). The reaction is effectively first example of a stereospecific reaction at the phos- irreversible owing to the low 14C content in the initial phoryl centre.ester and to the removal of the released radioactive methoxide ions by methanol We thank Dr. R. Collet of the Cantonal Hospital Geneva for assistance with the radioactivity PhEtPO.OMe* + MeO-+ PhEtPO-OMe+ Me*O-Me*O-+ MeOH + Me*OH + MeO-measurements. (Received July 4th 1962.) Green and Hudson Proc. Chem. SOC.,1961 145. Green and Hudson Proc. Chem. SOC.,1959 227; Michalski and Ratajczak Chem. and Ind. 1960 1241; Aaron, Uyeda Frack and Miller J. Amer. Chem. SOC., 1962 84 617. The Absolute Configuration of t~~eo-olp-Dihydroxy-~-methylbutyric Acid By Bo W. CHRISTENSEN KJWR and ANDERS (ORGANIC CHEMISTRY DEPARTMENT VETERINARY AND AGRICULTURAL ROYAL COLLEGE COPENHAGEN, DENMARK) HYDROLYSIS of the Veratrum ester alkaloids germi- tetrinel (germitetrine B2s3) and neogermb~dine~ affords in addition to the alkamine germine and volatile acids a lavorotatory orp-dihydroxy-or-methylbutyi-ic acid (I)identical with an acid obtained by resolution of the racemic acid synthesised by per- manganate oxidation of tiglic acid.4 The synthetic acid was therefore formulated as the cis-adduct but named there5 as erythro instead of as here threo because a different convention was used.g We now report that ( +)-threo-orp-dihydroxy- or-methyl butyric acid has the (orR,pS)-configuration (11).Myers Glen Morozovitch Barber Papineau-Couture and Grant J. Amer. Chem. SOC.,1956,78 1621.Nash and Brooker J. Amer. Chem. SOC.,1953,75 1942. Kupchan and Deliwala J. Amer. Chem. SOC.,1953 75,4671. Myers Morozovitch Glen Barber Papineau-Couture and Grant J. Amer. Chem. SOC.,1955 77 3348. Kupchan and Gruenfeld J. Amer. Pharm. ASSOC., Sci. Edn. 1959 48 737. Kupchan personal communication. PROCEEDINGS On reduction with lithium aluminium hydride this acid4 {[a]~ + 5.7" (c 4.6 in H,O)) was converted into (+)-threo-2-methylbutane-1,2,3-triol(III; R = H),7b.p. 138"/10 mm. nt5 1.4693 d;l 1,136 [a]","+ 8.9" (neat). Tosylation of the trio1 in pyridine at -20" gave the primary tosylate (In; R = p-CH3C6H,.SO~,m.p. 83" [a]k3-4.5" (c 5.0 in CHCl,) which on hydrogenolysis with lithium alu- minium hydride in ether yielded (+)-2-methyl-butane-2,3-diol (IV) b.p.74"/10.5 mm. n:5 1.4363 {R]:~ + 4-6" (neat). The absolute configuration of (IV) was established upon synthesis of its enantiomer b.p. 71.5"/9 mm. nt5 1.4367 d;Oe50.9716 [a]t3 -5-2" (neat) from (R)-methyl lactate ([a]il+ 9") and methylmagnesium iodide. The antipodal diols ex- hibited identical infrared and proton resonance spectra. Hence the laevorotatory a/%dihydroxy- a-methylbutyric acid derived from germimetrine1s8 and neogermbudine4s5 possesses the (aS,PR)-configura-tion i.e. the mirror image of (II). (Received June 25th 1962.) Satisfactory analyses and consistent infrared and proton resonance spectra have been obtained for all new com- pounds reported. Kupchan and Ayres J. Amer. Pharm. ASSOC. Sci. Edn.1959 48 440. Novel Synthesis of a Trimeric P-N Ring System By E. W. ABELand G. WILLEY (DEPARTMENT OF PHYSICAL AND INORGANIC CHEMISTRY THE UNIVERSITY BRISTOL 8) As part of a study of the reactions between com- pounds containing silicon-nitrogen bonds and the metal and metalloid halides we have treated N-alkyl- bistrimethylsilylamines R.N(SiMe,), with phos-phorus halides. The reaction between N-ethyl bis- trimethylsilylamine (1 mol.) and phosphorus tri- chloride (1 mol.) proceeded as shown in the equation and (I) PCI + Et.N(SiMe,) -+ ZMe,SiCI + [.NEt.PCI.] The chlorotrimethylsilane was distilled from the reaction vessel in virtually quantitative yield and low-pressure distillation of the remaining oil gave the trimer (11; R = Et) in 50 % yield b.p.128-130"/ 0.2 mm. d3O 1.371 nz1.5732 [Found C 21.9; H 4.3; N 12.9%; M 329 (ebullioscopic in benzene). CGH,,C13N,P requires C 21.9; H 4.6; N 12.8% M 3291 as a colourless liquid. The compound is stable under dry nitrogen but in the presence of moisture and oxygen fumes and undergoes rapid hydrolysis and oxidation. In addition to the trimer (11) formed as the major product a small quantity of a lower-boiling fraction is believed to contain the dimer which is already known for the phenyl ana1ogue.l Also an involatile pale yellow wax remained that is being investigated for the possible presence of larger ring systems and linear polymers. Compounds (11) are of interest in that they are isomeric with the well-known trialkyl derivatives of I ? I (I I R"'CI (ID) phosphonitrilic halides (III).2 The strongest absorp- tion in the infrared spectrum of the latter (e.g.R = Me) in the region 4000-650 cm.-l occurs at about 1200 cm.-l; assignment2 of this peak as the P=N stretch of the ring vibration characteristic of the cyclotriphospha(v)azene ring system is widely sup- p~rted.~ Absence of this predominant band from the spectrum for our compound (I1 ;R = Et) excludes the presence of a ring system containing P=N bonds. We thank the Department of Scientific and In- dustrial Research for a Research Studentship (to G.W.) the Microanalytical Laboratory of Imperial College for analyses and molecular-weight deter- mination Midland Silicones for a gift of chlorotri- methylsilane and Dr.Peter Slota for a copy of the infrared spectrum of compound (111; R = Me). (Received May lst 1962.) Michaelis and Schroeter Ber. 1894 27 490. a Tesi and Slota Proc. Chem. SOC.,1960 404. Daasch J. Amer. Chem. SOC.,1954,76 3403; Shaw Chem. and Ind. 1959 54. SEPTEMBER 1962 309 A Convenient Procedure for the Decarboxylation of Acids By D. H. R. BARTON and E. P. SEREBRYAKOV (IMPERIAL LONDON, COLLEGE S.W.7) THEuse of lead tetra-acetate-iodine (with illumina- acetate and iodine react as follows Pb(OAc) + I tion) for the generation and photolysis of the hypo- -2MeI + 2C02 + P~(OAC)~. No doubt a com- iodites of alcohols was recently described.l We now pound of the type R.CO.O.Pb(OAc) is generated in report that these conditions also provide a con-the general case.venient procedure for the decarboxylation2 of pri-An intermediate acyl hypoiodite (R-CO.O.1) may mary and secondary monocarboxylic acids in high be postulated which collapses by an S,i process. In yield. Illustrative results are given in the Table. agreement 3/%acetoxy- 1 1 -oxobisnorallocholanic acid To a 5% w/v suspension of lead tetra-acetate gives only one iodide which is tentatively formulated (1 mol.) in refluxing carbon tetrachloride (tungsten as the 20a-isomer (retention of configuration). Reaction Products (yields) Carboxylic acid Nor-iodides Other products Primary acids 1 2-Acetoxystearic 11-acetoxy-1-iodoheptadecane - 6-Benzamidohexanoic Secondary acids Cyclohexanecarboxylic 3fLAcetoxy- 1 1 -oxobisnorallocholanic 5-benzamido- 1 -iodopentane3 iodocyclohexane (91 %)3p-acetoxy-20 cc-iodoallopregnan- (82%) (63 %I -- 11-one (85%) Tertiary acids Pivalic Aromatic acids t-butyl iodide (10 %) isobutene (20 %) recovered pivalic acid ( > 40%) Benzoic 2-Naphthoic iodobenzene (56 %)2-iodonaphthalene (40 %) 1,4-di-iodobenzene (1 7 %)- Dicarboxylic acids Glutaric Adipic 1,3-di-iodopropane (1 2 %) 1,4-di-iodobutane (33 %) recovered glutaric acid (55 %) recovered adipic acid (1 5 %) Cyclohexane-1 ,Zdicarboxylic acid (cis) cyclohexene* (34 %) recovered -starting acid (15 %) * The reaction product was treated with zinc dust before being worked up.lamp) is added the carboxylic acid (1 mol.) and then One of us (E.P.S.) thanks the Royal Society and iodine (1 mol.) in the same solvent until the iodine the British Council for financial assistance under the colour persists.The reaction proceeds more slowly Royal Society-Academy of Sciences of the U.S.S.R. and less cleanly (lower yields) in the dark. Scientific and Cultural Exchange. Without addition of the carboxylic acid lead tetra- (Received June 26th 1962.) Meystre Heusler Kalvoda Wieland Anner and Wettstein Experientia 1961 17 475; Helv. Chim,Acta 1962, 45 1317. Compare Cristoi and Firth J. Org. Chem. 1961 26 280. Braun and Steindorff Ber. 1905 38 169. 3 10 PROCEEDINGS NEWS AND ANNOUNCEMENTS Election of New Fellows.-56 Candidates whose names were published in Proceedings for July have been elected to the Fellowship. Deaths of Fellows.-We regret to announce the death of Dr.E. C. Britton (31.7.62) of the Dow Chemical Co. Michigan U.S.A. Symposium on the Evaluation of Hydrogen Bonding in Biological Molecules.-The Society is sponsoring this Symposium to be held at the South-West Essex Technical College Walthamstow London E.17 on Saturday February 16th 1963 from 9.45 a.m. to 6.15 p.m. Inter- and intra-hydrogen bonds are frequently invoked to explain simple hydration and ring closure as well as more complex interactions in protein and nucleic acids. The Symposium contributions will approach the evaluation of hydrogen bonding from various physical and chemical aspects; it is hoped that they will add significantly to our understanding of the uses and limitations of experimental tools in this field.All persons interested in attending the Symposium should communicate with Dr. S. Lewin South-West Essex Technical College Walthamstow London E. 17. Symposia etc.- A Symposium on Fungi and their Metabolites sponsored by the Plant Phenolics Group and the British Mycological Society will be held at the School of Pharmacy Brunswick Square London W.C. 1 on October 1 st-2nd 1962. Further enquiries should be addressed to Dr. J. G. Manners Depart-ment of Botany The University Southampton. A Conference on X-Ray Diffractometry spon- sored by the X-Ray Analysis Group of the Institute of Physics and The Physical Society will be held at the Institution of Civil Engineers London on November 16-1 7th 1962. Further enquiries should be addressed to the Honorary Conference Secretary Dr.U. W. Arndt The Royal Institution Albemarle Street London W. 1. A Conference of the International Waste Rubber and Plastics Federation will be held in Antwerp on November 22nd-23rd 1962. Further enquiries should be addressed to the Secretary R. G. Kirk-patrick Moorgate Hall Moorgate London E.C.2 England. A Conference on “Sorption Properties of Vacuum Deposited Metal Films” sponsored by the Institute of Physics and The Physical Society in collaboration with the Joint British Committee for Vacuum Science and Technology will be held at the Univer- sity of Liverpool on April 17-19th 1963. Further enquiries should be addressed to Dr. J. H. Leck Department of Electrical Engineering The Univer- sity of Liverpool Brownlow Street Liverpool 3.The 1st European Symposium on Vacuum Engi- neering will be held in Frankfurt-am-Main on June 5-6th 1963. Further enquiries should be addressed to the Secretary of the Symposium Rheingauallee 25 Frankfurt-am-Main Germany. A Symposium on the Chemistry and Biochemistry of Fungi and Yeasts will be held in Dublin on June 18-20th 1963. Further enquiries should be addres- sed to Professor T. S. Wheeler Department of Chemistry University College Science Buildings Upper Merrion Street Dublin Ireland. The 34th International Congress of Industrial Chemistry will be held in Belgrade on September 23rd-27th 1963. Further enquiries should be addressed to Professor S.Stankovie Savez Hemicara- Tehnologa Jugoslavije Kneza Milosa Br.7/ 1 1 1 Belgrade Yugoslavia. Personal.-Dr. J. W. ApSimon formerly a Canadian National Research Council Postdoctoral Fellow has been appointed Assistant Professor of Chemistry Carle- ton University Ottawa. Professor W. F. Barker has been appointed Visiting Professor of Chemistry at the Mary Washington College University of Virginia U.S.A. Dr. S. Basu at present at the University of Upp- sala will return to the University College of Science Calcutta at the end of the year. Dr. N. Beredjick has been appointed Principal Research Scientist Associate at the Scientific Labora- tory of Ford Motor Company Dearborn Michigan U.S.A. Dr. C. E. Berkofl formerly a Research Fellow University of Southampton is now with Biorex Laboratories Ltd.Dr. C.J. L. Booker has returned from his National Research Council of Canada Fellowship at Ottawa and has taken up a position with the Corrosion of Metals Group at the National Chemical Laboratory Teddington. Mr. D. H. Booth recently accepted an appointment as Senior Research Chemist Laporte Acids Ltd. Castleford. Mr. G. CampbeZZ has been appointed Managing Director of Magnesium Elektron Ltd. Manchester. The title of Professor of Pharmaceutics has been conferred on Dr. A. M. Cook in respect of his post at the School of Pharmacy London. Dr. R. R. Davies has been appointed Associate Research Manager of the Dyestuffs Group Imperial Chemical Industries Limited Dyestuffs Division. Dr.J. E. Dove has been appointed Assistant Professor in Chemistry at the University of Toronto Canada. SEPTEMBER 1962 Dr. J. E. Drake and Dr. R. M. Scrowston have been appointed to Assistant Lectureships in the University of Hull. Dr. T. R. Griffithshas taken up an appointment for one year as Visiting Scientist at the Oak Ridge National Laboratories Tennessee U.S.A. Dr. F. Hartley is resigning his position as Director of Scientific Services of The British Drug Houses Ltd. to take up an appointment as Dean of the School of Pharmacy University of London in succession to Pruf‘ssor W. H. Linnell as from November 1st next. Dr. P. W. Jefs has taken up an appointment as a Research Associate at Indiana University U.S.A. Dr. H.Kitchen has been appointed Senior Lectur- er in Food Technology at the Royal Melbourne Institute of Technology. Mr. G. Nonhebel has been appointed a member of the Clean Air Council for the period to April 30th 1965. Dr. W. D. Ollis has been appointed to a Reader- ship in Organic Chemistry at the University of Bristol and will be Visiting Professor at the Univer- 311 sity of California Los Angeles during the autumn semester 1962/63. Dr. A. felter of the National Institute of Medical Research has been appointed Lecturer at the Univer- sity of Manchester. Dr. V.Petrow has been appointed Manager of the Research and Development Division of British Drug Houses Ltd. Dr. W. D. Scott of BTR Industries Ltd. has been elected Chairman of the Council of the Rubber and Plastics Research Association of Great Britain.Mr. D. W. WiZsun has been appointed Head of the Department of Chemistry at Sir John Cass College in succession to Dr. A. J. Lindsey. Mr. H. N. Wilson has retired from his position as Analytical Group Manager Billingham Research Department Imperial Chemical Industries Limited. Dr. F. P. Wuodford has been appointed Guest Investigator at (successively) the Rockefeller Institute New York; the Department of Physiology Univer- sity of Tennessee Memphis; and the National Heart Institute Bethesda Maryland during the academic year 1962163. PROGRAMME OF MEETINGS* OCTOBER 1962 TO JANUARY 1963 London Thursday October 1 lth 1962 at 6 p.m. Pedler Lecture “Amino-acid Sequences in Certain Enzymes,” by Dr.F. Sanger F.R.S. To be given in the Lecture Theatre The Royal Institution Albe- marle Street W.l. Thursday October 25th 1962 at 1.45 p.m. One-day Symposium on “The Chemistry of the Early Transition Elements,” to be held in the Main Lecture Theatre Chemistry Department University College Gower Street W.C.1. Thursday November 15th 1962 at 6 p.m. Tilden Lecture “Nuclear Magnetic Resonance Spectroscopy,” by Dr. R. E. Richards M.A. F.R.S. To be given in the Anatomy Lecture Theatre King’s College Strand W.C.2. Thursday December 13th 1962 at 6 p.m. Meeting for the Reading of Original Papers. To be held in the Rooms of the Society Burlington House w.l. Thursday January 17th 1963 at 6 p.m. Centenary Lecture “The Mechanism of the Enz- ymic Decarboxylation of Acetoacetic Acid,” by Professor F.H. Westheimer M.A. Ph.D. To be given in the Large Chemistry Lecture Theatre Imperial College of Science and Technology South Kensington S.W.7. Aberdeen Thursday October 18th 1962 at 8 p.m. Lecture ‘‘ Quantitative and Qualitative Aspects of Radical Addition,” by Dr. J. I. G. Cadogan. Joint Meeting with the Royal Institute of Chemistry and the Society of Chemical Industry to be held in the Medical Physics Lecture Theatre Marischal College. Wednesday October 31st at 8 p.m. Tilden Lecture “Nuclear Magnetic Resonance Spectroscopy,” by Dr. R. E. Richards M.A. F.R.S. To be given in the Chemistry Department The University. Friday November 23rd at 8 p.m.Lecture “The Prospects for Leaf Protein as a Food in Various Parts of the World,” by N. W. Pirie F.R.S. Joint Meeting with the Royal Institute of Chemistry and the Society of Chemical Industry to be held in the Medical Physics Lecture Theatre Marischal College. Aberystwyth (Joint Meetings with the University College of Wales Chemical Society to be held in the Edward Davies Chemical Laboratory.) Thursday October llth 1962 at 5 p.m. Lecture “Scientific Careers in Industry,” by Dr. J. P. Parke. * Offprintsof this programme can be obtained from the General Secretary The Chemical Society Burlington House, London W.l. 3 12 Thursday October 25th at 5 p.m. Lecture “Some Aspects of Fluorine Chemistry,” by Dr. R.Stephens. Tuesday November 6th at 5 p.m. Lecture “Mechanisms of Inorganic Redox Re-actions,” by Professor F. S. Dainton Sc.D. F.R.S. Thursday November 22nd at 5 p.m. Lecture “Directive Effects in Addition Reactions,” by Professor H. B. Henbest D.Sc. F.R.I.C. Thursday December 6th at 5 p.m. Lecture “Champagne,” by Professor F. Mackenzie M.A. D. ks 1. Belfast Thursday December 6th 1962 at 7.45 p.m. Lecture “Catalytic Superactivity of Metal Wires,” by Dr. A. J. B. Robertson M.A. Joint Meeting with the Royal Institute of Chemistry and the Society of Chemical Industry to be held in the Department of Chemistry David Keir Building Queen’s University. Birmingham (Joint Meetings with the University Chemical Society to be held in the Chemistry Department The University.) Friday October 19th 1962 at 4.30 p.m.Lecture “The Active Centres of Enzymes,” by Professor H. N. Rydon D.Sc. F.R.I.C. Friday November 9th at 4.30 p.m. Lecture “The Simplest Charge Atom and Mole- cule,” by Professor F. s.Dainton Sc.D. F.R.S. Bristol (Joint Meetings with the Society of Chemical In- dustry and the Royal Institute of Chemistry to be held in the Department of Chemistry The University unless otherwise stated.) Thursday October 4th 1962 at 6.30 p.m. Lecture “Recent Developments in Polyurethanes,” by R. L. Stafford B.Sc. A.I.R.I. Thursday October llth at 6.30 p.m. Lecture “Science in Art and Archaeology,” by Dr. A. E. A. Werner A.R.I.C. To be given in the Grange Barn Cinema Street Somerset.Thursday October 18th at 6.30 p.m. (Joint Meeting with the Royal Institute of Chemistry the Society of Chemical Industry and the Plastics Institute to be held at the Gloucester Technical College.) Thursday November lst at 6.30 p.m. Pedler Lecture “Amino-acid Sequences in Certain Enzymes,” by Dr. F. Sanger F.R.S. Thursday November 15th at 5.15 p.m. Lecture “Some Problems in the Chemistry of Cell- wall Materials,” by Professor E. L. Hirst C.B.E. D.Sc. F.R.S. PROCEEDINGS Thursday November 15th at 6.30 p.m. Social Evening (Films and Talk) to be held at Cheltenham. Thursday November 22nd at 6.30 p.m. Ladies’ Night “Perfumes,” arranged by Dr. R. Favre of Proprietry Perfumes Ltd. To be held at the College of Science and Technology Ashley Down.Thursday December 6th at 6.30 p.m. Lecture “Modern Methods of Aluminium Produc- tion,” by A. R. Carr B.Sc. A.R.I.C. and Dr. C. E. Ransley F.I.M. Joint Meeting with the Royal Insti- tute of Chemistry and the Chemical Engineering Group of the Society of Chemical Industry. Thursday January loth 1963 at 6.30 p.m. Lecture “New Developments in Chelatometry,” by Dr. T. S. West F.R.I.C. Thursday January 31st at 6.30 p.m. Jubilee Memorial Lecture “Horizons in Disinfec- tions and Antisepsis,” by G. Sykes M.Sc. F.R.I.C. Cambridge (Joint Meetings with theuniversity Chemical Society to be held in the University Chemical Laboratory Lensfield Road.) Friday October 19th 1962 at 8.30 p.m. Lecture “Organic Semi-conductors,” by Professor D.D. Eley O.B.E. Ph.D. Friday November 2nd at 8.30 p.m. Lecture “Oxidative Cyclisation,” by Professor G. W. Kenner Ph.D. Sc.D. Friday November 16th at 8.30 p.m. Lecture “Euclid and the Chemist,” by Dr. A. F. Wells M.A. Cardiff (Meetings will be held in the Department of Chem- istry University College Cathays Park.) Monday October 29th 1962 at 5 p.m. Lecture “Directive Effects in Additions to Cyclo- alkenes,” by Professor H. B.Henbest D.Sc. F.R.I.C. Monday November 19th at 5 p.m. Lecture “Non-stoicheiometric Compounds,” by Dr. J. S. Anderson F.R.S. Monday January 21st 1963 at 5 p.m. Lecture “Tetraterpenes,” by Professor B. C. L. Weedon D.Sc. F.R.I.C. Dublin (Meetings will be held in the Department of Chem- istry Trinity College.) Friday October 26th 1962 at 7.30 p.m.Lecture “The Chemistry of the Tannins,” by Professor R. D. Haworth D.Sc. F.R.S. Joint Meet- ing with the Werner Society. SEPTEMBER 1962 Wednesday November 28th at 5.30 p.n Lecture “The Periodate Oxidation of Disac-charides,” by Dr. J. M. Clancy. Durham (Joint Meetings with the Durham Colleges Chemical Society to be held in the Science Laboratories The University.) Monday October 29th 1962 at 5 p.m. Lecture “Hydrocarbon Complexes of Transition Metals,” by Dr. J. Chatt M.A. F.R.S. Monday November 12th at 5 p.m. Lecture “The Benzidine Rearrangement,” by Sir Christopher Ingold D.Sc. F.R.S. Monday November 26th at 5 p.m. Lecture “The Electronic Structure of Molecules,” by Dr.J. W. Linnett M.A. F.R.S. Monday December 3rd at 5 p.m. Lecture “Unusual Co-ordination Numbers of the Transition Metals,” by Professor R. S. Nyholm D.Sc. F.R.S. Monday December loth at 5 p.m. Lecture “Emerging Patterns in the Chemistry of Gallium,” by Professor N. N. Greenwood Ph.D. F.R.I.C. Edinburgh Thursday October 4th 1962 at 7.30 p.m. Lecture “The Use of Urea as a Fertiliser,” by J. K. R. Gasser. Joint Meeting with the Royal Insti- tute of Chemistry and the Society of Chemical Industry to be held in the Heriot-Watt College. Tuesday November 20th at 4.30 p.m. Lecture “Some Theoretical and Practical Aspects of Conducting Flames,” by Dr. T. M. Sugden M.A. Joint Meeting with the University Chemical Society to be held in the Department of Chemistry The University.Thursday November 22nd at 7.30 p.m. Lecture “Some Chemotherapeutic Topics,” by Dr. F. L. Rose O.B.E. F.R.I.C. F.R.S. Joint Meeting with the Royal Institute of Chemistry and the Society of Chemical Industry to be held in the Heriot-Watt College. Thursday December 6th at 7.30 p.m. Lecture “Solid-state Polymerisation,” by Professor C. H. Bamford M.A. Sc.D. F.R.I.C. Joint Meeting with the Royal Institute of Chemistry and the Society of Chemical Industry to be held in the Heriot-Watt College. Thursday January 17th 1963 at 7.30 p.m. Lecture “The Extrusion of Sulphur from Organic Molecules,” by Dr. J. D. Loudon A.R.I.C. Joint Meeting with the Royal Institute of Chemistry and the Society of Chemical Industry to be held in the Herio t- Wat t College.Tuesday January 29th at 4.30 p.m. Lecture “The Bacterial Cell Wall,” by Professor J. Baddiley D.Sc. F.R.S. Joint Meeting with the Uni-versity Chemical Society to be held in the Depart- ment of Chemistry The University. Exeter (Meetings will be held in the Washington Singer Laboratories PMce of Wales Road.) Friday November 16th 1962 at 5.15 p.m Lecture “The Recombination of Atoms-The Simplest Chemical Reaction,” by Professor G. Porter Ph.D. F.R.S. Joint Meeting with the Exeter University Chemical Society. Friday November 30th at 5.15 p.m. Lecture “Oxidative Cyclisation,” by Professor G. W.Kenner Ph.D. Sc.D. Friday December 7th at 5.15 p.m. Lecture “The Benzidine Rearrangement,” by Sir Christopher Ingold D.Sc.F.R.S. Glasgow Thursday October ISth 1962 at 4 p.m. Lecture “Solid-state Polymerisation,” by Professor C. H. Bamford M.A. Sc.D. F.R.I.C. Joint Meeting with The Andersonian Society and the Alchemists’ Club to be held in the Chemistry Department The Royal College of Science and Technology. Thursday November 15th at 4 p.m. Lecture “Spectroscopic Aspects of Optical Rotatory Power,” by Dr. S. F. Mason M.A. Joint Meeting with the Alchemists’ Club to be held in the Chem- istry Department The University. Thursday January 17th 1963 at 4 p.m. Lecture “Carbanions to Carbenes,” by Professor R. N. Haszeldine D.Sc. F.R.I.C. Joint Meeting with the Andersonian Chemical Society to be held in the Chemistry Department The RoyaI College of Science and Technology.Hull Thursday October lSth 1962 at 7.30 p.m. Lecture “Molecular Biology,” by Dr. M. F. Perutz F.R.S. Joint Meeting with the Royal Institute of Chemistry to be held in the Physics Lecture Theatre The University. Thursday November 8th at 5 p.m. Pedler Lecture “Amino-acid Sequences in Certain Enzymes,” by Dr. F. Sanger F.R.S. To be given in the Physics Lecture Theatre The University. Tuesday November 20th at 5 p.m. Lecture “Electrochemical Methods of Studying Reactions,” by Professor W. F. K. Wynne-Jones D.Sc. Joint Meeting with University Students Chem- ical Society to be held in the Organic Lecture Theatre The University. beds Thursday October 25th 1962 at 5.45 p.m.Lecture “Stereochemistry of Some Enzymic Re- actions,” by Dr. J. W. Cornforth F.R.S. To be given in the Chemistry Lecture Theatre The University. Leicester Thursday November lst 1962 at 4.30 p.m. Lecture “The Electronic Structure of Molecules,” by Dr. J. W. Linnett M.A. F.R.S. Joint Meeting with the University Chemical Society to be held in the Department of Chemistry The University. Wednesday November 14th at 3.30 p.m. Lecture “Nuclear Magnetic Resonance in Organic Chemistry,” by Dr. R. A. Y.Jones M.A. M.S. Joint Meeting with the Colleges of Art and Tech- nology Chemical Society to be held in the Colleges of Art and Technology. Monday December 3rd at 4.30 p.m. Lecture “Aromatic Fluorine Compounds,” by Pro- fessor M.Stacey D.Sc. F.R.S. Joint Meeting with the University Chemical Society to be held at the Department of Chemistry The University. Liverpool (Joint Meetings with the University Chemical Society to be held in the Donnan Laboratories The Chemistry Department The University.) Thursday October 25th 1962 at 5 p.m. Tilden Lecture “Nuclear Magnetic Resonance Spectroscopy,” by Dr. R. E. Richards M.A. F.R.S. Thursday November 29th at 5 p.m. Lecture “Organic Reactions in Strong Alkalis,” by Professor B. C. L. Weedon D.Sc. F.R.I.C. Thursday January 31st 1963 at 5 p.m. Lecture “Carbonium-ion Rearrangements,” by Dr. C. A. Bunton. Manchester Thursday October 18th 1962 at 6.30 p.m. Lecture “Alkyl and Aryl Complexes of Transition Metals,” by Dr.J. Chatt M.A. F.R.S. To be given at the Manchester College of Science and Tech- nology. Thursday November lst at 6.30 p.m. Lecture “Some New Horizons in Reaction Kinetics,” by Professor F. S. Dainton Sc.D. F.R.S. To be given at the Manchester College of Science and Technology. Thursday November 29th at 5 p.m. Lecture “New Thoughts on Old Dyes,” by Dr. E. N. Abrahart. Joint Meeting with Students Union Chem- ical Society of the Royal College of Advanced Tech- nology to be held at the Royal College of Advanced Technology Salford. Newcastle-upon-T yne Friday October 19th 1962 at 5.30 p.m. PROCEEDINGS Bedson Club Lecture “The Properties of Glass,” by J. A. Frost to be given in the Chemistry Department King’s College.Nottingham (Joint Meetings with the University Chemical Society to be held in the Chemistry Department The University.) Monday October 22nd 1962 at 5 p.m Lecture “Sherry,” by D. Lloyd. Tuesday October 30th at 8 p.m. Lecture “The Chemistry of Cement Hydration,” by Dr. H. F. W. Taylor. Joint Meeting with the Royal Institute of Chemistry. Tuesday November 13th at 5 p.m. Lecture “Cationic Polymerisation,” by Professor D. C. Pepper M.A. Ph.D. Tuesday November 20th at 5 p.m. Lecture “The Study of Molecular Structure by Infrared Spectroscopy,” by Dr. N. Sheppard M.A. Tuesday November 27th at 5 p.m. Lecture by Dame Kathleen Lonsdale D.Sc. F.R.S. Tuesday January 15th 1963 at 5 p.m. Lecture “Radiocarbon Dating,” by H. Barker. Tuesday January 29th at 5 p.m.Lecture “Biosynthetic Pathways in the Amaryl- lidaceae,” by Professor A. R. Battersby Ph.D. North Wales Thursday October 25th 1962 at 5.45 p.m. Lecture “The Scope for Chemical Propellants,” by Dr. C. H. Johnson C.B.E. Joint Meeting with the .University College of North Wales Chemical Society to be held in the Chemistry Department University College Bangor. Wednesday October 31st at 7.30 p.m. Lecture “Some Reactions of Free Radicals with Aromatic Compounds,” by Dr. G. H. Williams. Joint Meeting with the Royal Institute of Chemistry to be held in the Chemistry Department Denbigh- shire Technical College Wrexham. Thursday January 31st 1963 at 5.45 p.m. Lecture “Chemical Aspects of Enzyme Specificity,” by Dr.W. J. Whelan F.R.I.C. Joint Meeting with the University College of North Wales Chemical Society to be held in the Chemistry Department University College Bangor. Oxford (Joint Meetings with the Alembic Club to be held in the Inorganic Chemistry Laboratory.) Monday October 22nd 1962 at 8.30 p.m. Lecture “Co-ordination Chemistry of Boron and Gallium,” by Professor N. N. Greenwood Ph.D. F.R.I.C. Monday October 29th at 8.30 p.m. Lecture “The Chemistry of the Bacterial Walls and SEPTEMBER 1962 Membranes,” by Professor J. Baddiley D.Sc. F.R.S. Monday November 19th at 8.30 p.m. Lecture “The Nature of the Intermediates formed by Hydrocarbons in Catalytic Reactions on Metals,” by Professor C. Kemball M.A. Ph.D. F.R.I.C. Reading Tuesday November 27th 1962 at 5.45 p.m.Lecture “Application of Nuclear Magnetic Reson- ance to Organic Chemistry,” by Dr. A. R. Katritzky M.A. Joint Meeting with the Royal Institute of Chemistry to be held in the Main Chemistry Lecture Theatre The University. St. Andrews and Dundee (Meetings will be held in the Chemistry Depart- ment Queen’s College Dundee.) Tuesday October 30th 1962 at 5 p.m. Tilden Lecture “Nuclear Magnetic Resonance Spectroscopy,” by Dr. R. E. Richards M.A. F.R.S. Tuesday November 13th at 5 p.m. Lecture “Electron-spin Resonance Studies of Simple Inorganic Oxy-radicals,” by Professor M. C. R. Symons D.Sc. F.R.I.C. Tuesday December 4th at 5 p.m. Lecture “A few Chemical Problems Connected with Cancer Chemotherapy,” by Professor F.Bergel D.Sc. F.R.S. Sheffield Thursday November 29th 1962 at 4.30 p.m. Lecture “Enol Elimination Reactions,” by Dr. J. Harley-Mason M.A. F.R.I.C. Joint Meeting with the Royal Institute of Chemistry and the University Chemical Society to be held in the Department of Chemistry The University. Southampton Tuesday October 16th 1962 at 7 p.m. Lecture “Explosives,” by Dr. B. D. Shaw M.M. T.D. To be given at the College of Technology Portsmouth. Friday October 19th at 5 p.m. Lecture “One-electron Bonds,” by Dr. J. W. Linnett M.A. F.R.S.To be given in the Chemistry Department The University. Friday November 16th at 5 p.m. Lecture “Biosynthetic Pathways in Amaryllidaceae,” by Professor A. R. Battersby Ph.D.To be given in the Chemistry Department The University. Friday November 23rd at 5 p.m. Lecture “Aliphatic Electrophilic Substitution,” by Sir Christopher Ingold D.S. F.R.S. Joint Meeting with the Royal Institute of Chemistry to be held in the Chemistry Department The University. Wednesday December Sth at 7 p.m. Lecture “Platinum Group Metals,” by E. C. Davies. To be given at the College of Technology Ports- mouth. Friday December 7th at 5 p.m. Lecture “The Reactivity of Solids,” by Dr. F. S. Stone. Joint Meeting with the Royal Institute of Chemistry to be held in the Chemistry Department The University. swansea (Joint Meetings with the Student Chemical Society to be held in the Chemistry Lecture Theatre University College.) Monday October 15th 1962 at 4.30 p.m.Lecture “Natural Polyacetylenes,” by Professor E. R. H. Jones D.Sc. F.R.S. Monday November 19th at 4.30 p.m. Lecture “Simple and Complex Metal Nitrates and Nitrites,’’ by Professor C. C. Addison D.Sc. F.R.I.C. Monday December loth at 4.30 p.m. Lecture “The Shapes and Spectra of Small Mole- cules,” by Professor A. D. Walsh Ph.D. F.R.I.C. Tees-side (Joint Meetings with the Royal Institute of Chem- istry the Society of Chemical Industry and the Society for Analytical Chemists.) Thursday October 4th 1962 at 8.15 p.m. Lecture “Some Thoughts on Ceylon and India as Potential Industrial Countries,” by Dr. I. J. Faulkner F.R.I.C. To be given at the William Newton School Norton. Wednesday October 17th at 7.45 p.m.Exhibition of Scientific Films to be held at the Synthonia Theatre Billingham. Wednesday November 7th at 7.30 p.m. Annual Dinner Dance to be held at the Billingham Arms Billingham. Monday November 12th at 8 p.m. Lecture “Solvent Extraction of Inorganic Com-pounds;Some Recent Developments,” by Professor H. M. N. H. Irving M.A. D.Sc. F.R.I.C. To be given at the Constantine Technical College Middles- brough. Wednesday December 5th at 8 p.m. Lecture “Organometallic Co-ordination Complexes of Some Group I1 and IDElements,” by Professor G. E. Coates M.A. D.Sc. To be given at the William Newton School Norton. Wednesday January 9th,1963 at 8 p.m. Lecture “Explosions and Explosion Research,” by Dr. J. H. Bergoin.To be given at the William Newton School Norton. APPLICATIONS FOR FELLOWSHIP (Fellows wishing to lodge objections to the election of these candidates should communicate with the Honorary Secretaries within ten days of the publication of this issue of Proceedings. Such objections will be treated as confidential. The forms of application are available in the Rooms of the Society for inspection by Fellows.) Abdy Colin George. “Hillside,” Hull Road Osgodby Selby Yorkshire. Anandavadivel Kumarasamy B. Sc. “Chandra Vasa,” Kopay Ceylon. Bacon James Edward B.A.Sc. 39 West Avenue South Hamilton Ontario Canada. Breisacher Peter B.A. Ph.D. 28212 Ella Drive Palos Verdes Peninsula Calif. U.S.A. Brennan Leo Charles Gerard BSc. 5 St. Lawrence Road Clontarf Dublin 3 Eire.Bullock Caroline Elinor B.Sc. “Enfield,” Lower Broomie- knowe Lasswade Midlothian. Burton James Michael B.Sc. Box 451 San Diego State College San Diego 15 Calif. U.S.A. Byrne Anthony Robert B.Sc. Londonderry Laboratory for Radiochemistry South Road Durham. Carrier John Paul. 45 Gregory Way Childwall Liver- pool 16. Chakraborty Jagadish M.Sc. Department of Bio-chemistry St. Mary’s Hospital Medical School, London W.2. Chapman John B.Sc. 89 Ladysmith Road Lipson Plymouth Devon. Clark Edmund Roy Ph.D. A.R.I.C. 15 Grasmere Close Tettenhall Wolverhampton Staffordshire. Cooper David John B.Sc. Keswick House 29 Sylvan Road Upper Norwood London S.E.19. Dabom Gordon Russell. Taychreggan Jenner Drive West End Woking Surrey.Dence Joseph B. 704 Richards Road Toledo 7 Ohio U.S.A. Dyson Ian Fraser B.Sc. Organic Chemistry Department The University Leeds 2. Farmer Geoffrey Charles Henry. 88 Chessington Hill Park Chessington Surrey. Feigl Dorothy Marie B.Sc. Department of Chemistry Stanford University Stanford Calif. U.S.A. Feldman Isaac Ph.D. Chemistry Department Univer- sity College Gower Street London W.C.1. Gadallah Abd El Aziz Ali M.Sc. 12 Borhan Street Manchiet El Bakry Cairo U.A.R. Hansford Geoffrey Spearing B.Sc. Afterglow 5 Crown Crescent Camps Bay Cape Town South Africa. John David B.Sc. Institut fur Anorganische und Allge- meine Chemie Technische Hochschule Wien VI, Getreidemarkt 9 Austria. Jones Barry David. 36 Molyneux Park Road Tunbridge Wells Kent.Kamlet Mortimer J. B.S. M.A. Ph.D. Embassy of the United States Office of Naval Research 429 Oxford Street London W.l. ADDITIONS TO Harrap’s standard French and English dictionary. Edited by J. E. Mansion. Part 1. French-English. Pp.912. Harrap. London. 1961. Industrial research in Britain. Edited by A. W. Haslett. 4th edn. Pp. 461. Harrap. London. 1962. Laboratory organization and administration. K. Guy. Pp. 386. Macmillan. London. 1962. (Presented by the publisher.) Inorganic chemistry for advanced and scholarship levels. J. H. White. Pp.496. University-of London Press. Lawrence Kenneth Gordon B.Sc. All Saints’ Vicarage Hampton Road Forest Gate London E.7. McKeown Robert Henry M.Sc. 80 Kellys Road St. Albans Christchurch New Zealand.Marks Robert Edward Ph.D. A.R.I.C. 1 Benhams Drive Horley Surrey. Miller David Ph.D. Pendle Cottage 61 Claygate Road Dorking Surrey. Miller Marcus William B.Sc. 17 Vale Road Claygate Esher Surrey. Mohammad Mi M.Sc. 54 Windsor Street N.W. Glas- gow Scotland. Morel Charles Joseph Ph.D. Bromhubelweg 9 Arles- heim,Switzerland. Morris Lindsay Johnston Ph.D. 39 Cherry Way Shepperton Middlesex. Pan. Yuh Kann. B.Sc.219 Bailey Street East Lansing -_ Michigan U:S.A. Paukstelis. JoseDh V.. B.Sc. Room 1 13A. Noyes Labora- tory University of sIllinois Urbana Illinoh U.S.A. Pike Malcolm Robert B.Sc. 41 Milton Road Chelten- ham Gloucestershire. Richards Kenneth Edward B.Sc. A.R.I.C. 28 Vicarage Crescent Caverswall Stoke-on-Trent Staffordshire.Riddell George Moir. 50 Kaystone Road Glasgow W.5, Scot land. Robinson Brian Harford M.Sc. 42 Randolph Street Woolston Christchurch New Zealand. Rossell John Barry B.Sc. A.R.C.S. 5 Swains Avenue Carlton Road Nottmgham. Roy Jyotibhushan M.Sc. Chemistry Department Uni- versity College Gower Street London W.C.1. Rutledge Peter Stewart Ph.D. Chemistry Department Auckland University Auckland New Zealand. Safford Joyce B.Sc. Alfred Deakin Hall Monash Uni- versity Clayton Victoria Australia. Spande Thomas Frederick M.A. Frick Chemistry Laboratories Princeton New Jersey U.S.A. Stuart Kenneth Lloyd B.Sc. Chemistry Department University of the West Indies Kingston 7 Jamaica W.I. Vassiliou Eustathios B.Sc.6 Wellington Crescent Manchester 16. Velandia Jose Alberto. Instituto Venezolano de Investi- gaciones Cientificas Apartado 1827 Caracas Vene zuela South America. Waite John. 16 Kimberley Road Harrow Middlesex. Walker David Alan PbD. Biochemistry Department Connecticut Agricultural Experiment Station P.O. Box 1106 New Haven 4 Connecticut U.S.A. Whish June Challis B.Sc. Department of Botany The Queen’s University of Belfast Belfast 7 Northern Treland. THE LIBRARY London. 1962. (Presented by the author.) Ion exchange. F. Helfferich. Pp. 624. McGraw-Hill. New York. 1962. Modem aspects of the vitreous state. Edited by J. D. Mackenzie. Vol. 2. Pp.260. Butterworth. London. 1962. NEW JOURNALS Chemical Processing (Chicago) from 1962 25.Fortschritte der Anneimittelforschung from 1959 1. Bulletin of the Royal Society International Scientsc Information Services from 1962 1.
ISSN:0369-8718
DOI:10.1039/PS9620000289
出版商:RSC
年代:1962
数据来源: RSC
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