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M# lavestigatioa of eom@ Group Throe Ilalldee with verloue Donor Moleoules including a Thermodynamic Study of Carhonyl Bromide m d some Aluminium Hal Idee
A thesis submitted by
MABTIN S&IO A&TB0B2I
In eandltatur© for the degree of Doctor of Philosophy of the University of Dondoa*
November 1971 Royal Holloway College# Englefleld Green#
X h ' T : T i b 4 A ry ' Surrey* CLASS _ÆÔB— Ho. _ B û L - ACC. No. io & m _ date acq ProQuest Number: 10096771
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ProQuest LLC 789 East Eisenhower Parkway P.O. Box 1346 Ann Arbor, Ml 48106-1346 The author would like to oxproso hla grateful appreciation of the Invaluable help given by Drs# P.J* Gardner and A* finch# and the many other people, too numerous to mention# who have contributed towards the completion of this thesis* Thanks are also due to the Science Research Council fo r providing a research grant.
To my parents# and to everybody who has helped me along the way# Amfmof
The Interaction between boron tribalid ©3 (bromide, chloride) and benzene derivatives (benzene, m»xylene, mesltylene, hexamethylbenzena) has been studied using the techniques of cryosoopy, vapour pressure measurement, proton magnetic resonance and Raman spectroscopy* ; No evidence was found to support previous aaeertions of weak complex formation in solution* The reactions between carbonyl halides (bromide, chloride) and Group Three halides (boron, alimimium; chloride, bromide, iodide) have been studied* Ho complexes could be Isolated, in contrast to the system C0 0 l 2/AlCl^ studied previously* No reaction was found using boron halides, but with aluminium halides rapid and extensive halogen exchange occurred at ambient temperatures. Aluminium halides, also catalysed the decomposition of carbonyl bromide *
COajPgCî-) « CO (g) ♦ Brg(g) L+ Srgtt)]
The enthalpy of hydrolysis of carbonyl bromide was measured by adiabatic solution calorimetry coargd) + nHjO — ^ a)g(g) ♦ t2E3r,(n-l)ri203( l)
hydrolysis « -49«06 — 0.15 koal mol“^ L-205.56 - 0,67 kJ mol“h
.H.C From this was derived
à ïû COâr-(l) » -34.70 - 0.21 koal mol"'’ t-145.lS Î 0.83 kJ m ol-h differing eignifioantly from previous values* Estimates have been made of thermodynamic fuactlene for carbonyl iodide. The standard enthalpies of formation of aluminium bromide and aluminium Iodide have been determined as AlBr^(o) « "11$.h " 0*6 heal 1-49$. 2 - 2.4 kJ mol^^]
&k| A1Ij (o) - -67.1 - 0.6 keal mol"’ L-230.9 - 2.4 kJ mol"’j by ieoperibol calorimetry of the reactions
AlXj|(o} ♦ i3Ha0X * ïillaÔHjàqLltaâl(OH)^ ♦ 3HaCl
AlGl.te) + L33aX + aBaOHjaq C#aAl(OK)^ + SBaCl ♦ 3HaX ♦ (n-4)HaOm]a% where X « Br o r I* In addition, êJî| KaAl(0H)^.15.000 HgO - -409.9 t 0.4 kcal mol"’
1-1713.0 - 1.6 keal mol"’j Infra-red studies on the complex fOGlj*3Cl,(o) have confirmed Its structure as dative covalent, with bonding from oxygen to boron. On running the Infra-red spectrum at room temperature, there Is considerable dissociation of the complex, and also reaction with the potassium bromide windows. Deter mination of the degree of dissociation has shown that there is only negligible g^sociatlon in the vapour* m t Fa%@ Interaction between boron trlfealldea and benzene derivatives Introduction 11 Phase Diagrams 12 Vapour Pressure Curves 27 Proton Magnetic Resonance 38 Raman Spectroscopy 43 Conclusions 45
Pan? T'm " Interaction of group 3 halides with carbonyl halides Imt roduet ion 48 Apparatus * Vacuum lin e 5 3 I*R, gas cell 5 3 Low temperature gas cell 56 Analysis procedure for Al, Cl, 8r. 60 COSrl preparation 61 AIX^ purification 65 BBry^COarg 67 Aiarg/GOBrg 8 2 AlCiycOBTg 83 Aisrycocig 2 2 AlIp/COGlg 94
Discussion 97 PAR? Page Thermochemioal properties of carbonyl bromlie Introduction 102 Adiabatic Calorimeter and procedure 104 &B* hydrolyala of CO&fg 112 AH® COBr^ 114 Thermodyn&mio functions fo r COIL 115 Carbonyl Iodide 116 Dissociation of GOXg 122 r&m? fmx. Thermochemleal properties of aluminium M l idee Introduction 124 C.T.E* calorimeter and procedure 131 Results 137 Enthalpies of formation 144
PAR? flT E The phoaphorjl, chlorlde-boron trichloride com.plex (l) 'Infra-Eed Spectrum and Interpretation Introduction 147 Practical procedure 1 5 2 Results 154 Interpretation of results 15 4 Page (2) Enthalpy of Dissociation Introduction 171 fresBure-measuri&g apparatus 173 C alibration 174 Procedure 180 Results and Discussion 182
Appendix One*- 12 7 Justification of the proportionality of the displacement from zero on the chart recorder against the change In thermistor resistance#
Appendix Two*- 189 Justification of the linear extrapolation of resistance-time fore- and after-periods on C.T.È# calorimeter.
Appendix Three *- 19 2 Justification of the extrapolation of fore- and after-periods to the midpoint of the heating period in G.T.E* calibration#
References 195 No. ra?e 1 Cryoacopio apparatus 14 2 Calibration of platinum resistance thermometer 15 3 CSbloroform/s-esitylene phase diagram 18 4 Boron trlbromlde/mesitylene phase diasrara 21 $ Boron trlbromlde/m-xylene phase diagram 24 6 Vacuum system used for vapour pressure 28 measurements 7 Boron trlbromide/benzene vapour pressure curve 32 8 Boron trt-iodide/beasene vapour pressure curve 35 9 Boron tribromide/mesitylene vapour pressure 37 curve 10 Vacuum system used for carbonyl halides/group 3 54 halides 11 Infra-red gas c e ll 55 12 Low temperature infra-red cell 57 13 1*IU spectrum of gaseous boron tribromide 69 14 I.B* Spectrum of gaseous carbonyl bromide 70 15 I.E. spectrum of gaseous boron trlbromide/ 72 carbonyl bromide 16 I.E. spectrum of solid boron trlbromide 76 17 I.H» spectrum of solid carbonyl bromide 78 18 I*t. spectrum of solid boron tribromlde/ 8 0 carbonyl bromide 19 I.E. spectrum of vapour from carbonyl bromide/ 83 aluminium bromide 20 I.E. spectrum of vapour from carbonyl bromide/ 87 aluminium chloride Am (continued) go* Pam 21 Î.S* spectrum of gaseous carbonyl chloride 90 22 I.E. spectrum of vapour from carbonyl 92 oaloride/aluminlum bromide 23 I.E . spectrum of vapour carbonyl clüoriâê/ 95 aluminium Iodide 24 Carbonyl bromide ampoule 106 23 Carbonyl bromide reaction trace 111 26 C.Î.I. calorimeter 132 27 Aluminium halide reaction end calibration 14 2 trace 29 I.E . spectrum of gaseous phosphoryl chloride 155 29 1,B* spectrum of gaseous boron trichloride 156 30 I.E. spectrum of gaseous boron trichloride/ 157 phosphoryl chloride 31 Low temperature I.E* spectrum of solid 15 8 fOClj.BCl. 32 I#R* spectrum Of POCl^.BCl^ (sublimed ©olid) 159 33 • I.E . spectrum of 32, after pmping cell 160 34 Î.B# spectrum of 2001^*801, (mujol mull) 161 33 I.E. spectrum of I0C1..BC1. (nujol mull on. 162 standing) 36 Pressure-measuring apparatus 175 37 Calibration of transducer a t zero pressure 17 6 39 Calibration of transducer with water 17 8 39 Vapour pressure of eyolohexane 17 9 40 Vapour pressure of propionic anhydride 181 I u
vmn mn
THS maJACMW. 9'-=TWSW 80B0S TRIHAUBES
ATÎ3 BEfiZESE OFJÎIVATITES lEgRQBüCTIO?» « 2 Previous work * has ©uggeated tha presenoa of weak eolttte-aolvent eomplaxlng In benzene solutions of boron tribromida and boron tri-iodide. In each case, the symmetric stretching mode, ï>.|, (278 B3r^, 196 cnT^ BI^) of the boron trihalide appeared as a strong band in the infra-red spectrum* Since the selection rules predict Raman activity only for this mode, it was suggested that deformation of ths boron trihalide was allowing the band to become active. nuclear magnetic resonance of the boron tribromide/ benzene system showed a residual high-field shift at infinite dilution of 0*48 p.p.m* (relative to pure boron tribromide) after correcting for the bulk diama^etio susceptibility of the system, again suggesting weak interaction* It was proposed that the principal orientation of boron was above or below the plane of the aromatic ring, as indicated in the following diagram;- The average distance between the aromatic ring and the boron o atom was estimated to be 4«7a * However, the same workers state that the phase diagrams for boron tribromlde ♦ benzene and boron tri-iodide + benzene showed no sign of complex formation* The object of the present work was to seek further evidence fo r complex formation, using the following teolmiquea r
(1) Cryoscopy, and determination of phase diagrams* (2) Vapour Pressure Measurement* (3) Proton Magnetic Resonance* (4) Raman Spectroscopy* Boron trich lo rid e and boron tri-io d id e were studied with the following aromatic compounds| benzene, m-xylene, meaitylene and hexamethylbenzene * the electron-donating methyl group should increase the Lewis basicity of the aromatic ring*
(t) ■ PHASE BIAGHAPB , The appai'atus was tested using the Mown chloroforn/ mesltylene system^, and then the mixtures
boron tribromide/mesitylene, boron tribromide/m-xylene, and boron tribromide/hexamethylbenzene were studied* Experimental Apparatus I- The orypmeter used {diagram 1) has been described la detail elsewhere*^ Temperature was measured using a miniature platinum resistance thermometer {Degussa lîanau) connected across one arm of a Wheatstone bridge, and calibrated using the freezing points of purified (see next paragraph) chlorobenzene (-45*2^0) and chloroform (-63*5^0)^ (diagram 2). Readings were accurate to - 0.1 ®C# Stirring was achieved by means of a solenoid-activated,reciprocating, chromium-plated* iron cylinder* Chemicals t-
Chlorobensens (May and Baker) was dried over size k t molecular sieves and d istille d at atmospheric pressure (b.p# 131.5®C, literature*^ 131 .V®0). d*lorofom (May and Baker) was shaken with water to remove added alcohol, dried over 4% molecular sieves, and dlatllled (b.p. 61,2°C, literature** 61.2®C). Mesityleae (B.D,H* Ltd.) was dried over 4A molecular sieves and distilled (b.p. 164*7^0, literature^ 164.7^0). lî-xylene (B.D.H. Ltd.) and hexamethylbeasens (Koch-Light) were used as supplied commercially, with stated purities of at least and respectively. A
DIAGRAM 1
CRYOSCOPIC APPARATUS
rubber .cap glass— coated iron
jacket
CELL PLATINUM STIRRER RESISTANCE THERMOMETER LOo
LU
cr h-
LU o
en
=)
o CJ oo
o o o CD oo u 33NViSIS3d Boron triBromlde (iCoch-Light) was purified by shaking with mercury to remove any free bromine♦ end then distilled at atmospheric pressure# the portion boiling at 91 ^ I^C being retained# (lite ra tu re b.p*^ 91#3*0)#
Procedure!- One component was placed In the cryometer# and successive aliquots of the second component added throu^sh the side-arm* Moisture sensitive boron trlbromide was always loaded In the cryometer In a dry-box# the side-arm covered with a self- sealing rubber cap# and the hydrocarbon aliquots added through the cap using a repeater syringe (Jencons ltd*)# weighed before and after each addition* fhe quantity of boron halide used was determined by post-experiment hydrolysis with water# separation of the hydrocarbon# end titration for bromide with 0*1B silver nitrate solution# using cosIn Indicator* The mixture was frozen using a tcluene/llquld nltroger. slush bath (ca-94*0)* The freezing point was taken when the temperature remained almost constant for approximately one minute# or longer# due to the latent heat of fusion evolved* there supercooling occurred# the temperature fell well below the freezing point# and then rose sharply as crystallisation commenced# falling again when equilibrium had been reached* The freezing point was taken as the maximum temperature reached after crystallisation^* 17
Results and Bieeusslon Ohlorofom ♦ #eeltylen@ The results of a single experiment are given In Table 1.
' 1 POIMTS 0? CHI^ROBOR!^ ^ rTlXTUREH
P.R.T. RSiUSTASCE m w . m a x'-oibî M012 rmcTioâ ^ (OBS?a t 0 ,1 o®5) (®0 * 0.1*0) 75.8 -63.5 100*0 73.1 -70.9 89*5 72.8 -71.5 87.3 72.2 -73.5 60.9 72.8 -71.5 77.4
75.4 —64*6 73*9 i 76.8 —60*8 69.6 I 79.1 -55.4 ! 63.2 ' ^ 1 60.1 -52,8 I 57.9 80.3 -50*4 53.4 1 81.0 -50.3 49.5 80.8 -50*4 46.2 73.6 —56*5 34*7 75.1 -65*7 26.1 74.9 —66*1 15.5 79.2 —54*6 0*0
The freezing points for pure chloroform and pure mesitylene agree very well with literature^ values of y
o -o > o
00 o u
l_L_ ■s g 2 u N LU § LU X cr u
•5 z o H- CJ < cr Lu ro LU o
o o LO
FREEZING POINT C aad *52*7^0 respectively* The phase diagram (diagram 3) ehowa a well-defined peak at .50 mole chlorofozim# Indicating a 111 complex with a 7 congruent melting point'# In effect# this corresponds to two. eutectic phase diagrams linked together# namely# for chloroform ♦ complex and complex ♦ mesltylene# The diagram agrees well with previous results^*
Boron Trlbromide ♦.Mesltylene On mixing boron tribromide and mesltylene# both of which were colourless# a very pale lemon-yellow colour appeared# which deepened considerably on freezing the mixture# In all# six experiments were carried out# the results of which are given in Table 2# ’ The results show very poor reproducibility below approx imately 50 mole ^ boron tribromide (diagram 4)# The reason for this is not entirely clear* One’ possible explanation Is that supercooling# which was particularly marked In this area# resulted in the true freezing point not always being reached# Upon commencement of crystallisation ♦ Crystals of mesitylene# being in excess# would tend to separate# so tliat by the time equilibrium was reached the liquid present would be richer in boron trlbromide# and the freezing point correspondingly lower# Hence the correct freezing point would never be attained# «lu
2 FREESias FQI3T3 m soRON THiqROwiag/i'îs-oim^iS igixtmsa
EXPZRlIdlT P.R#T* RZ3ISTiriC3 IgZSZIflO 20I2S? H0L3 rsACTion (OHM - 0.1 m m ) (*C 1 0.1*0) BonoN TRi3mr;iDT3 1 82,7 —46.0 100.0 1 81,0 , -50,2 90.8 1 79.3 -55.0 83.1 2 77.5 -59.3 74.5 6 77.0 —60.5 73.0 1 76,8 -61.0 71,1 t 74.7 —66,7 62.1 2 73.8 —63,8 59.3 6 73.7 -69.5 57.4 1 73.1 -70,9 55,2 6 71.9 -73.8 51.9 1 71.7 —74,4 49.6 5 69.6 -80.2 49.5 2 71.3 -75.5 49.3 3 70,5 -77.7 « 47.9 6 70.7 -77.1 47.4 3 70.4 -77.7 45,9 1 69.9 -79.2 45.1 3 70,5 -77.7 44.0 6 69.1 -31,0 43.5 2 68.9 - 81.6 42.2 3 70.4 -77.7 42.2 6 69.0 -81.3 40.3 5 69.5 —80.2 39.6 2 68.9 —8l .6 36.8 4 68.4 -82.2 36.8 » 5 69.4 —80,2 32.9 2 68,9 —81,6 32.7 3 70.5 -77.7 29.8 a.1
ro
LU o LU .o oo
LU o -O . CÇ
N
o o o FREEZING POINT °C a a. T4 3^,15 g (CCHÎlSUEa) zxPEiamsGT P$R.T* B2ÜISTAHCB P^2%IB9^gOIBT MQZIS FRACTIOm # KUN%2% (OHM • 0.1 OIM) 1 69.3 -80.5 29.4 5 71,6 -75,0 28,2 k 69.9 -79,1 23.0 2 70,7 -77,1 26,7 3 72.5 -72,5 24,7 2 72,4 -72,6 24.5 3 72,0 -73,7 23,0 k 71,4 -75.3 22.6 3 73,5 -69,8 21.9 3 74,0 —68,4 19,7 4 72,2 -73,3 13,9 3 74,3 -67,5 18,7 2 73,4 —70,0 18.3 5 74,7 —66,6 17,9 4 72,9 -71,3 16.3 3 75.4 —64,6 15.3 4 72,9 -71,3 14,3 3 79,9 -63,4 13.7 4 75,5 —64,5 12,7 9 75,6 —64,2 12.3 3 76,5 —62.0 12,0 3 76,6 —62 ,0 10,7 3 76,6 —62,0 9,7 4 76,0 -63,0 8,9 3 77,3 -59,6 8.9 3 77,3 -59,6 7.5 4 76,8 -61,1 6,8 79,2 -54,6 0.0 2.2 The true Treeaslng point la tbua taken aa the hlgheot value obtained for any given composition, when a reasonably smooth curve is obtained. A plateau is clearly discernible between approximately 33-43 mole ^ boron tribromide, but there is no evidence for a peak due to complex formation. The plateau is probably due to a m sciblllty gap'*7 , 8 containing, at equilibrium, two solids in equilibrium with the remaining liquid. iOn complete ©olidlflcatlon, a heterogeneous mixture of the two solids would presumably be obtained]. By applying the Phase.Buie to 7 Ê this equilibrium * , we obtain, for three phases and only two components, one degree of freedom. This la taken up by the pressure of the system, assumed constant. This means.that whatever the composition of the mixture within this range, the temperature of freezing must remain constant* A similar incongruency was found in the boron tri-iodlde/beasene phase diagram . Boron Trlbromide ♦ m-lylene Again, a yellow colour was observed, but only in the solid phase. The results of a single run are given In Table 3# The literature^ freezing point fOr m-xylene is -47.4*C, The freezing point-composition curve (diagram 5) ©hows a simple eutectic at -79^3, where the composition la 47/& boron trlbromide, 53# m-xylene. There la no mlsclbility gap as for 24- r o m CJ X o LU Q 2 . o CD S CQ cr M § cc o .O CD O I— LO W cn Ll LU—I o o o o oo FREEZING POINT C a. 5 f m M 3 fmmim of mmm TRiE ram m rs P.E.T. EEJISÏAKCS FBszaiMG yoiai KOLB ÎSACTIOH ;( (orm - 0.1 own) ( * ’0 1 0.1 OC) 0>KOÎ? TRI'înOMIDB 82.7 -46.0 100.0 74.6 -67.0 66.0 73.0 -71.1 60.8 72,8 -71.8 55.4 72.1 -73.5 52.6 71.0 -76.4 49.3 70.3 -73.2 46.3 72.7 -71.8 41.4 74.8 —66.2 35.7 82.0 -47.5 0.0 feorm trlbromld^meaityleme# and no evldeneo fo r complex formation# Boron frlbromide ■» Hexamethylbeneene A similar phase diagram could not be obtained, due to the very low solubility of solid hexasethylbeneena in boron tri- bromide i#e. the presence of a wide misclbllity gap# All solutions prepared at room temperature and cooled, gradually precipitated hexamethylbensene to give saturated solutions of approximate composition 1 mole ^ hexamethylbonzene, 99 mole ^ 'oi6 trlbrom ide, and freezing at On mixing hexamsthylbensQn® (a white so lid ) with boron tribromide (a colourless liquid), an orange colour appeared in the solutions obtained. 2.7 (2) VAPoim PHESBijHE The systems studied werei- Boron tribromide ♦ benzene Boron tri-io d id e ♦ benzene Boron tribromide ♦ m esityleae. Experimental Apparatusi- All the compounds were handled on a vacuum line capable of achieving pressures of 10**^ torr or better (diagram 6). Mixing of the compounds, and all vapour pressure measurements, took place in a section of the line completely free from grease, thereby avoiding any absorption of the grease or reaction with boron tribromide. Under these conditions, best vacuum seals wei*e obtained by using taps of the "Uni-form" type (Glass Precision Engineering Co. Ltd.,), and "0-ring" joints (y# Young (Scientific Glassware) Ltd.). Thus the compounds were exposed only to Teflon, glass and mercury vapour (from the manometer). Vapour pressures were determined using a mercury manometer, consisting of two parallel glass tubes, approximate diameter 2 cm, connected by a capillary tube. The mercury was carefully degassed before use by pumping for several hours while warming gently. The difference in height of the two mercury columns was measured using a cathetometer. Pressures were usually reproducible to -1 to rr. (D ID E o 3 CO h- Lf) I— z UJ :E UJ cr 3 if) < LÜ cr :e e cn UJ cr :E 3 ÜJ in I— in U1 UJ > cr if) Gl 2 cr 3 3 3 O O CL < < > > 0» CD 2 < cn o: :(/) o Chemicals I*** ristilled boron trlbromldc, b.p* 91^0 (lltem ture^ 91#9®C) was relist Iliad Into Its container on the vacuum line* where It was stored over mercury to remove any free bromine formed. I t was degassed by repeated freezing and pumping. Before each run* the sample used was distilled and trapped at ea, -80^0 (solid carbon dioxide - acetone slush bath)* at which tempera ture any hydrogen bromide Impurity would be pumped off (vapour pressure^ ^ $00 torr). It was then condensed Into the mixing vessel* and its purity checked by measurement of its vapour pressure at known temperatures «- Temp, (^c) Experimental V,p, (torr) Fit, Value^ (torr) 0 19 19 12,8 3? 35 20 52 54 A supply of pure boron tri-iodide was already available in the laboratory* with established purlty^*^ 93-»99f, Commercial samples of benzene and mesltylene (B,D.n, ltd ,) were distilled (b,p, 80,0®C and 164#7^C respectively* lite r ature values^ 80,1®G and 164* T^O* and dried and stored over grade 4Î molecular sieves ( ltd,), Each was degassed by repeated freezing and pumping* The p u rity was also checked by measurement of vapour pressure i- 30 Temp, (^0) Experimental V,p, (torr) lit. value^*^^(torr) Benzene 0.0 2$ 26.3 12,8 52 52,5 20,5 79 76,7 Keelty- 19 2 1,7 len© . î*rocedurei- A quantity of boron trlhalid© was either condensed into the mixing vessel, (3I?r^) or loaded in a dry-box and then connected to the vacuum lin e and degassed (31^)* Aliquots of benzene or mesitylen© were then distilled in, the amount added being deter mined by weighing the container before and after each addition, The mixture was now isolated from the pump, and allowed to come to equilibrium at the temperature selected for the pressure measurements. Readings were usually taken below room temper ature to avoid thermostating the manometer. To achieve this, either an ice bath (0 - I^C) or a p-xylena^llquid nitrogen slush bath (13 - 1^0) was used, la the case of boron tribroald© + mesltyleae, consistent readings were only obtained when the mixture was well stirred. The reason for this is act known. To achieve mixing, a magnetic stirrer and glass-coated follower were used# which meant however that the mixture could not b© thermostated# and all readings had to be taken a t room tem perature. The quantity of boron halide used was determined by post-experiment hydrolysis and titrim otrie analysis. At the end of each run# the mixture was frozen in liq u id nitrogen# and the mixing vessel cut off using a hot spot. The mixture was then analysed argentimetrlcally# as described previously (page 16)# for h alide. R esults and Diecussion Boron Tribromide ♦ Benzene A colourless solution resulted from mixing boron tri bromide and benzene# which remained colourless on freezing, in contrast to the other systems studied* four experiments were perforaed, and the results shown in Table 4# The results from experiment 1 were in very poor agreement with the rest, and were discarded in constructing the pressure composition curves {diagram 7), These show a positive deviation from Raoult's lawi- m Mole fra c tio n component A p^ m P a rtia l pressure component A P m Observed pressure, ' LJ o o o 33. CO o .O 00 LU Q LU N O cr cn cr .O CD o CO cr LO o LU CD o I— > < cr Ll. LU __i o IE 8 o o O CD CNl VAPOUR PRESSURE (TORR) 33 TA. y F, 1» ■ VAPOira FRERRimER Qg BORON AlIxggTÎBB EXPS3ir.n^:rr VAPOm rBE33UR3 (TORR * 1 TORR) Î’OIS PSACTIOÎÎ # Km mm x-FÆîî TaiBsoiîicg : 0^0 1 1^0 13^0 - 1®C 1 19 ## 100.0 2 19 37 100.0 3 22 38 97.9 3 24 43 94.0 3 26 45 66.5 3 21 49 61.6 1 41 53.6 j 2 29 49 48.8 3 29 52 45.4 1 2 30 51 33.3 3 23 52 35.4 2 23 52 34.6 2 21 51 31.7 k 21 51 23.2 k 23 51 23.7 2 27 51 ■ 22.9 1 31 #» 21.6 4 . 23 51 15.7 1 32 - , 11.4 1 30 #» - 8.3 3 23 52 0.0 This Implies that the major contribution to Interaction between boron trlbromlde and be name in the liquid phase comes from Van der Waal*s forces# An alternative repre— aentatlon* is that the attraction between two molecules of the same component Is greater than between two different molecules* If any significant complex formation occurred# a negative deviation would be expected* Boron Tri^lodlde 4- Benzene Great difficulty was experienced in dissolving boron trl-lodide# even with vigorous stirring* Consequently# measurements were only possible at 13^C#. and then only at concentrations below 30 mole ÿ boron trl-ioild©* The solutions were colourless# and remained so on freezing*, The results of the single experiment are given in Table 3* 3 VAPOUR 0? i30RGR TRI-IODIDE * BENZEDB VATCOH^raESSaSS tîOlE IIUCÎICÎI ^ {'i'CEE Î 1 .SCEE) n o m s 1 3 »C * 1 0 0 TSI-I 0 r iI 3 S 2 1 0 0 . 0 4 3 ...... 2 3 . 6 4 8 1 3 . 4 4 9 9 . 6 5 2 0 , 0 o k o •S N ÜJ Q Q O OC LO .O z O o (X o CÛ z .o o I— < oc ü_ co UJ -J o .o CN o o CD /<" VAPOUR PRESSURE (TORR) 36 Deaplta the email mumber of poiate, waà the fact that only one-third of the total oompoeition range ia covered# the preeeure-oompoeitlon curve (diagram S) ©hows a clea r p o sitive deviation from Eaoult*a Law, Again# there is no support for complex form ation. Boron Trjhromlde + Beeitylene AS in the cryoeoopio experiments# a yellow colour# particularly pronounced in the ©olid phase# appeared on mixing the two components. Because of the necessity for stirring to obtain consistent pressure measurements# readings were only taken at room temperature# 19®C. The results are shown in Table 6, fABLB 6 ^ VAPOim Of BOROTi TRIBRO^DE 4- K BITY Tm S VAfCmrPBESSOES IÊOI.E mCT103 f (TOp « 1 ÎOKH) BOKOK 19 i 1 ®c TRIBHOSIiîÊ 52 100.0 37 64.3 3S 61.7 33 49.7 26 35,0 2S 32.5 2 0.0 As with the previous experiments# there is a positive deviation from Eaoult^s Law# and no evidence of complexlng (diagram 9). .37 rO UJ o 00 CO UJ •s U 1 cn -8 O o o o CD OJ VAPOUR PRESSURE (TORR) 38 (3) i?;;0T03 sashctio rsjOkascs 4 E sh ifts were determined for mesltylene, and a 1 al mixture of mesltjlene and boron tribromide$ using tetra- methylsliana (T.M .5.) m internal standard. For comparison, the 1st mesitylene/chlorofom epeotrum was also run* Experimental Chemicals Mesltylene, chloroform and boron tribroslde were purified as described previously (p# 21 ), Tetramethylsilane was used as obtained commarolally (9.D.M. Ltd*). ProcedureI- All the spectra were run on a Varian HA CO^IL machine (60 IHz), locked on to the T.M.S* peak, with facilities for low temperature spectra using a liquid nitrogen/nitrogen gas flow for coaling. The R.g.B. tubes were filled with sample in a dry*bo% (including small quantities of as an internal standard), and then sealed* Spectra were obtained for all three samples at room temperature, and also at for meeitylene and meoitylene/ boron tribromide. 3 ^ Results Bnâ discussion 7 lirsi'i";,! IRTSESITY rHSWTO^If SHIfT S rilTTIE O RATIO Room teaperatura * 40*0 PATTERN Ez p.p.m. Hg p .p .a . OgHafcsg), 3 C113 128.7 2.145 127*^ 2 .1 2 7 At le a s t quartet* 1 H 336.1» 6.607 394A 6.573 At le a s t octet* 0.113(0153)3 3 SI3 1 2 9 .2 2.153 l 2w#42*140 I At le a s t 1 quartet* 1 H 396.S 6.613 393*0 6.583 1 At le a s t + 33r%J 1 octet* 3 CII3 130.0 2.167 mm 1 At le a s t 1 quartet* 4^ GhCl^ 1 n 400.6 6.677 At le a s t o c te t. Ail shifts are relative to T^S, and accurate to * 0*311^ (0*006 p.p.m.)* The separation between peaks due to splitting was 0*611^ for each system* 4 0 The predicted spectrum for mesltylene^V la two peaks with very complex splitting, elnea all the couplings are long range, em ail, and of the earns order. However, the chemical shifts and ohoervcd splitting in the present case are In good agreement with published speotra for pure meeltylene,^^*^^*^^' Before the shifts In the mixtures can be compared with those of pure mesltylene, allowance must be made for the variation in bulk magnetic susceptibility. This can be represented'^ as ''preferred orientations* of the molecules in the mixtures arising not from any chemical Interaction but from the inherent magnetic anisotropy of the molecules not averaging out to xero when all possible molecale-molecule orientations are considered. The correction is made using the eiuatio&5*15,l6 ^corrected “ ^observed J 7T ( ^ mixture - ^ mesltylene) where The values used were X mesltyleae « «0,652 c.g.s. unlts^ X 33r. a-0.900 * ^ X CMCl a -0.731 * ^ 41 Tm effect of the TiAB la this respect eaa be neglected, since it is only slii^htly affected by changea la bulk magnetic susceptibility, and It la present in only very low concentrât ions # After making the corrections, the shifts obtained (relative to pure mesltylene as sero) are*- T.ABta 5 CORRECTED P.7.R. SHIFTS « Q) mOTON COmlGAL SHIFT (F.P.Îâ#) ROOM -40*0 mz)iTfLm3E —0.24 —0.24 ♦ Biir^ H -0 .2 5 —0.24 -0 .0 6 « ♦ c n c i. H -0 .0 1 Due to the corrections made, shifts are now accurate to only — 0«0t p*p#m* All shifts are now up-fleld, showing more shielding, as expected. Aromatic protons or substituents are noiraially strongly deshielded by the fie ld produced by the ring ourren1?^*^f which opposes the applied field within the aromatic ring, but reinforces it outside the ring# Hence an up-fleld shift 4a suggests a weaker rin g cu rren t, would occur i f there was Interaction between the tt -electron eyetea and the Lewie acid* However, the s h ifts are sdwaya only very s lig h t, in d icating that any interaction must be extremely weak* Compare, fo r example, with the complexes of 1,3,5 trlnitrcbenzene with benzene and methylbenzenes*?, where shifts of 1.2 - 1.3 p.p.m. are obtained for the ring protons. But negligible shifts are also obtained for solutions of iodine in benzene^^, where a weak charge-transfer Interaction is also thought to occur^^. Previous studies^on the sh ifts of the chloroform proton 11 and B of boron tribromide, both in aesityleae so lu tio n , gave values for 1*1 mixtures of -1.1 p.p.m. and -0.60 p.p.m. res pectively, relative to pure chloroform and pure boron tribromide. The small values again suggest only weak Interaction, though stronger than indicated by the present work* This is probably due to the change in shielding of the Lewis acid atom, sited directly above the TT-electron cloud, being more pronounced than the indirect shielding of the rin g substituents. The identical shifts obtained a t room temperature and -40*0 for mesitylene/boron tribromide suggest l i t t l e or no temperature dependence fo r any interaction* This is ra th e r surprising, since the s h if t obtained would be a weighted average o f the sh ifts for "complexed* and *free** moleoules (assuming rapid exchange)* LSee, fo r example, reference 14, p.tCOj* One would expect an increasing proportion of molecules to be com plexe! as the temperature was lowered, with a corresponding increase in sh ift* 4-2 (4)^ RASAtî sj-Boramoos* l'uiiiii uii'iiiwiwn Mimiuiiw.m# The Eaman spectrum of an approximately 1*1 mixture of boron tribromide and benzene was run in the region 0-300 om"*\ while the ©ample was kept frozen, in order to detect the possible presence of lattice bands attributable to relative motion of molecules in the ©olid {©@e, for example, reference 20)# Because of the large masses of the molecule© Involved, such lattice modes would be expected at very low frequencies# Experimental Spectra were run using a Carey 8l instrument, with a 50 mW Helium-Eeoa laser, and 6328a exciting lino# The sample was placed in a flat-ended capillary, subseq uently sealed# This was surrounded by a wad of plasticine to act as an insulating jacket, frozen by immersion in liquid nitrogen, and placed in the laser beam# Periodic visual checks were made to ensure that the sample remained solid# Results and Diocusslon ■ Only two bands were obtained, at 113 and 230 both strong la Intensity# These are" attributable to the ( in-plane bend) and (symmetric stretch) modes of boron tribromide, which occur at 150 and 273 om"^ in liq u id boron tribromide^^*^^# Benzene has no absorption in this region^^# 44 The absence of any further bands Indleates that lattice bands must be so weak as to be Indlatlngulehable from back ground noise* 4-r CONCLUSIONS While the results obtained eliminate the existence of strongly-bonded complexes between boron trihalides and benzene derivatives, the possibility of weak contact charge- transfer bGndinap4,25 j^^î^ains open* The positive vapour pressure deviations from Eaoult*s Law, and the failure to find any discontinuities due to oomplexing in the phase diagrams suggest that the bonding energy must be extremely small, even by charge-trancfer stan- darda (for example, benzene/iodine is quoted as —1#32 k c a l/ mol t-5‘.S2 kJ/mole] and benzene/trinltrobensene as -1*71 kcal/ mol, C-7. IS* kd/mcl3)* A p a ra lle l can be dram with the benzene/ iodine system (see, for example, the reviews in references 19 and 27), where the bulk properties of the solutions show no deviations, but charge-transfer complexing is well established by other methods, such as infra-red end ultra-violet spectro- seopy, spectrophotometry, X-ray crystallography^etc* Also, as already mentioned, there is no measurable s h if t ta the benzene resonance on dissolving iodine*^# further support for loose eomplexing comes from the yellow colours observed when freezing many mixtures of boron tri bromide and methylbenzcnee. Absorption la the v isib le spectrum is a well-known phenomenon associated with charge-tranefer complexes19*24,2d^ caused by traaaitioaa between energy levels 46 associated viith the oharga-transfer bond* Indeed, the olosely related ayeteas aluminium bromide/benzene or methyl- benzenes have been shorn to form 1*1 and 2*1 solid complexes, all of which are various shades of yallow^^*^^* Two approaches which might provide more evidence o f' completing are the use of dipole moments and of ultra-violet spectral measurements of solutions of boron trihalides in benzenes* Both require further investigation* Many oharge- tranefer complexes, even when formed by molecular species each with very small individual dipoles, have been shown to exhibit ■appreciable dipole moment(e.g. iodine/benzene)* A deter- mlnation^ of the dipole moment of boron tribromide in benzene yielded a value of 0.194 e.s.u. Since very low values are unreliable i t was Inferred that the boron trlbromlde was planar However, In view of some of the other evidence, th is needs re-lnvestlgatlng, along with other systems. In case the low value represents a real deviation from zero. 47 M H T TWO ISSïMSIÎiSLÆ_£S2S2L±ï®SlJWSSM W I T H C U t B O m . HAtIBîS 4g IKTRODÜOTim Although reactioma of boron an! aluminium halides with organic carbonyls (and other organic ligands) have been extensively studied because of th e ir importance in Friedel- Crafts type syntheaes^^, with one exception very little work has been done on their reaction with carbonyl halides. The exception is the system aluminium chloride/carbonyl chloride (phosgene). The first reference to this is in a paper by Band“^^t who claimed the existence of several solid complexes varying from Al^Olg.^COdg to 2Al2Cl^,aOOl2# In the light of later work, this paper seems very unreliable, Qermmnr^ and co-workers carried mit extensive studies on the solutions, which contributed considerably to the formulation of acid-base theories for solvent systems^^*^^*^^. Aluminium chloride/phosgene solutions were believed to contain ionic species, such as COGl^ cr cations and hlC X ^ anions, and the solid complexes obtainable from these solutions were also thought to be ionic. Subsequently, Uuston^^ has shorn that the exchange rate of radioactive chlorine for very dilute (less than 2 equivalent fraction per cent) aluminium chloride solutions in phosgene is vary slow, exchange half-lives varying between 12 and 60 houra at O^C, This slow exchange would be extremely improbable i f ionic species were involved, Huston also showed the exchange to be homogeneous. In a later paper, however^^, he modified this view to allow for the appearance of ions a t 4-i higher oonoentratlme, the exact concentration at which ions first appear in the solution being dependent on temperature* (Approximately 10 equivalent fraction per cent a t -21^0, 2 at 0®C and 0.3 at 25®C), Mora racently, Jonea and Wood*^^ aad Claris hare shown that only the 111 complex e x ists a t ambient temperatures, this being a lËiite, easily dissociated solid, melting point 25®C. Both sets of workers overcame extreme experimental difficulties caused by the reactivity of the complex to obtain its infra-red spectrum* Although there are some wide discrepancies between the two seta of band assignments, both papers agree on the dative covalent nature of the complex, with co-ordination through the oxygen to aluminium* CL. .CL ^ C = 0 ------c r ^ C L Christo confirmed the 111 aluminium chloride/phosgene complex by measurement of vapour pressure — composition curves at 0®C and 25®G, when both a negative deviation from Raoult*s Law and a sharp break at a 1*1 mole ratio , were found# However, with boron trichloride/carbonyl chloride, he found a positiv e 5"o deviation from Eaoalt*s law# and no evidence for a complex. This l3 rather surprising, since boron trichloride Is generally accepted to be a stronger lewis acid than aiuainium cMoride^^*^, literature data on systems involving bromide is very sparse. A. von Bartal^^ claimed that aluminium bromide reacted with excess phosgene according to the equaticm A13r^ ♦ ^GOClg » AlOlg ♦ 30001Br. With the aluminium bromide in excess, however, at 100^0 aluminium dichlorobromide, AlSrClg# was supposed to be formed as a brick-red solid, melting at 142-143^0* Identification of this substance was based on analysis. The red colour seems rather surprising in view of the fact that both aluminium chloride and aluminium bromide are white. In a much later hG paper # Corbett and Gregory repeated the experiment, and found that the red colour was due to free bromine. Also, the crystals obtained resembled an annealed mixture of aluminium chloride and aluminium bromide rather than any mixed halide phase involving AlBr 0l 2 Î no analytical data were given. Aluminium dichlorobromide is known, however, from the reaction of hydrogen chloride on aluminium bromide^^*^^. The objects of the present work w ere:- (1) In view of the failure of Christe^^ to find a complex between boron trichloride and carbonyl chloride, to investigate the system boron trlbroaldc/carbonyl bromide. Boron tribrom ide SI la a better îi@wls acid than boron trichloride due to the smaller degree of 7V-boading involved in the molecule, leading to a smaller re organisation energy for the transition Bep^—~^3ep^ required on oomplexing^^*^^*^^# Carbonyl bromide, despite being thermodynamically much lees stable than phosgene with respect to the dissociation COXgd) # 00(g) ♦ X gd) o r (g) ISee P art 3] might b@ expected to behave as a better donor because of the lower electronegativity of bromine compared with chlorine^\ There is the added advantage that possible halogen exchange would be unimportant* Techniques employed were infra-red spectroscopy of gaseous mixtures and solid mixtures, and the ^^3 magnetic resonance of liquid mixtures* (ii) To re-investigate the reaction between aluminium bromide and phosgene, and also the other combinations# Aluminium ch lo rid ^o aib o ay l bromide Aluminium bromide/carbonyl bromide Aluminium iodide/carbonyl chloride with a view not only to establishing the products of excMnge reactions, but also determining any complex formation, both at room tem perature and below* s a Techniques employed were#- (à) Mixing the two components, both at room temperature and low temperatures, pumping off volatile products, then weighing and analysing the residues* (b) Infra-red spectroscopy of the vapours from the mixtures* In both (1) and (ii), infra-red and Eaman spectra of both solutions and residues would have been invaluable* However, sufficiently inert materials suitable for fabricating infra-red cells were not available ( compare reference 41, where the complex aluminium chloride/carbonyl chloride is stated to attack sodium chloride, potassium bromide, silver chloride and polythene). Infra-red spectra of the residues yielded only bands too broad and too weak to be assi^&d, Measurement of Bamaa spectra was attempted, but each time strong fluorescence obscured any bands* This was possibly due to bromine being present in very low concentrations, since it is well toowa^^ that minute quantities of coloured impurities can cause strong fluorescence. S3 Apparatus and Procedure All the compounds were handled on a modification of the grease-free vacuum line described in Part One {p# 3,7) (diagram 10). V o latile compounds were stored in the grease- free containers, A and B, from which samples could be distilled into the respective 0-tubes and then into the mixing tube, C* This mixing tube was fitted with a greaseleso tap and attached to the line via a greaseless joint, so that it could be read ily detached and weighed. Having obtained the required mixture, one of the storage containers could then be replaced by an infra-red cell or tube as required* The relatively iavolatile, solid, aluminium halides were loaded in to the mixing tube in a dry-bcx, and the mixing tube then attached to the line and evacuated. For infra-red gas phase spectra, a cell specially designed and constructed in the department was used (diagram 11). The cell, path-length 10 cm, was fitted with potassium bromide windows, held on and sealed by Viton *0-rings* and metal screw caps. It was also fitted with a *0nifgreaselesa tap (Glass Ireclsioa Engineering Co. l t d . ) , and with an *0-ring* joint for attachment to the vacuum line. It was capable of holding a vacuum of ca. 10*^^ torr. for filling with the vapour 54 LU Q _J U) < U) X on Si Qj LÜ in Q < cr o < Q S'S* niAGRAM 11 IN FRA-RED GAS CELL Polypropylene Cap PJ.RE. Plunger Metal Metal Screw S crew Cap Thread Plate o 0 Cork V i t o n W asher R u b b er from a re&etlom mixture# eell anà mixing vessel were both evacuatedf and then closed to the pump. The taps connecting the two were opened, end the eastern allowed to come to eq.uil^ ihrlua# After ©hutting the taps and pumping away excess vapour, the cell could usually he removed from the line simply by slid in g the jo in t a p a rt, d esp ite the vacuum inside* In a few instances, i t was necessary to shrink the #0*ring3* of the Jo in t, using a little liquid nitrogen, in orler to pull i t apart* To run the spectrim of a gaseous mixture of boron trihromide and carbonyl bromide, tW cell was attached to the line in place of the mixing vessel* Each of the components, stored in v essels A and B, was distilled into its respective 0 -tï^ p , where they were thermostated such th a t the two vapour pressures were nearly identical, and of a suitable value to obtain reasonable peak h e i s t s in the spectra. The boron tri^ bromide was kept at rocm temperature (approximately 20®C) and the carbonyl bromide at approximately *-$^0 (ice/salt bath), where both have vapour pressures in the region $0 torr^*^^» The vapours were then allowed to mix, and the cell filled# The ' spectrum of a solid boron tribromide/carbonyl bromide mixture ms obtained by using the cold c e ll shown in diagram 12, A l i t mixture of the vapours, prepared as described above, was sprayed through the fine nossle on to the caesium Iodide plate. DIAGRAM 12 LOW TEMPERATURE INFRA-RED CELL > Copper Holder Cs I KBr P late Plate 5 y Caeaiua Iodide was uaed due to Ito auparlor roslataaoe to ther^mllj-^lnduced strain compared witli moat common window materials# The plate was sited in a copper holder cooled liquid nitrogen# After spraying on the mixture» the plate could he rotated hy means of the greased joint at the top of the apparatus ©o m to align it with the potassium bromide windows# Care had to he taken that the greased joint did not ©les® due to the extreme cold# Seat mobility was obtained using Apiezon 2? grease (May and Baker)# Another pi^hlem encountered in humid weather waa that of water and carbon dioxide condensing on the outside of the potassium bromide p lates due to the cooling of the apparatus» ^d spurious bands were obtained in the spectra’’'it approximately 360Ô caT^ and 1600 csT^» (w ater) and 66$ caT^ (carbon dioxide)# The m e doubt about the principle of the method might fee that# for weak complexes where th ere is interaction in the solid but not the gas phase# spraying a gaseous mixture on to a cold plate gives not a mixed crystal» but a mixture of the individual crystals# However» using a sim ila r method fo r feromine/feenzene# rersoo showed that only the mixed crystal was formed, end complexing" was observed# A ll infra-red spectra described above were run on a Pemin" Timer 33? instrument, 'giving a resolution of approximately i 3 cm.*^# Infra-red spectra of the residues from aluminium h a lid e / carbonyl halide reactions were attempted ao nujoi mulls and dry powders# between potassium bromide plates# These spectra were run on a ?erkin*Elmer 32$ instrument#. However, the bands obtained were always too broad end weak to be of use in Identifying the residues# Beman spectra of solutions end so lid s were attempted on both Gary 8l (hellus/neon laser, 6328a exciting lin e ) and Bpex Bamalog (ionised argon laser, 3147A ex citin g lin e ) instruments, but a l l samples fluoresced so strongly as to completely obscure any bands# The samples were made up by filling thick-walled capillary tubes (with flattened bases of thinner glass) in a dry-box situ a te d in a fume cupboard, then freezing, and sealing off the tubes# nuclear magnetic resonance for boron tribromide/ carbonyl bromide and pure boron trlbromlde was carried out on a lerkin-Elmer 110 Instrument (60 fwo B#M#E# tubes, each ■Z ' containing a sealed capillary of boron trifluoride etherate, BfjtOBtgf ac external reference, and fitted with a.greaseless joint, were attached to the vacuum line and filled respectively with pure, boron, tribromide . and an approximately 1*1 mixture of boron tribromlde/oerbonyl brmide# They were then sealed, and the spectra run# Shifts were measured accurately to - 3 p#p.m# GO For studying th# reactions of aluminium halides m â carbonyl halides, the mixing vessel was evacuated, weighed, loaded with aluminium halide In a dry-box, then re-evacuated end re-^vsighed# Carbonyl halide was d istille d in, and the mixing vessel re-welshed# The system was then thermostated if required# , After pumping o ff carbonyl halide at low teiop- erature, usually -43 - 1^0, (chlorobenaene ©lush bath), the mixlns vessel was re-weighad to see i f there was any addition in weight corresponding to complex formation# Comparison with the method of preparation^^ and dissociation pressures^^ quoted for aluminium chloride/carbonyl chloride showed the unlikelihood of decomposing any complex formed by pumping away excess carbonyl M ild# a t -43^0# A fter pumping to dryness at room temperature the residue was weighed, hydrolysed, and anal ysed as follows I- Excess sodium hydroxide solution (concentration not critical) was placed in the mixing vessel above the tap, and the vessel etoppered# The tap was ©lowly opened to allow gradual hydrolysis of the residue# ■ After shaking to dissolve the residue completely, the hydrolysate was made up to a stan dard volume in © ll# tly acid solution (lüîO^)# The addition of acid was to prevent th e slow precipitation of various alumlnlUTi/ hydroxide species which occurs in alkaline solution on standing, and which are very difficult to rediseolve# The hydrolysis Itself had to be carried out in alkaline medium because aluminium bromide, unlike aluminium chloride, hydrolyses only extremely slowly at pB ^7 Total halide concentration was determined by Volhard*smethod^ , involving the precipitation of s ilv e r halide from nitric acid solution with excess 0.13 silver nitrate, filtering, and back-titrating the excess silver nitrate with 0,13 ammonium thlocyanate, using f e rric indicator. Chloride was determined^^ by first removing bromide by oxidation to bromine with potassium iodate in acid solution, followed by boiling. Excess iodate m s removed by reduction with phosphorous acid, and boiling off the iodine so formed. The chloride solution remaining ma then titrated by Volhard*a method as above, the bromide present being obtained by d ifferen ce, Aluminium was determined g rav i- metrically by precipitation as the 8-hydroxyquinolate (**oxinatô*), Al(0^1^03)j, in ammonium acetate buffer (reference 56, p, 383 and 316), The precipitate was collected in a sin tered glass crucible (porosity Ho. 4 ), washed, dried at 120^0, and weighed. Chemicals Cax^onyl bromide was prepared according to the method of Schumacher and lenher^^, by the reaction of carbon tetz^brmide with concentrated sulphuric acid (specific gravity 1,83) at 130-170^0. The carbon tetrabromlde was heated just to melting la a dry, nltrogea-fllled atmosphere, m â the eulphurlo add added dropwlse from a funnel * After all the sulphuric add had been added, the mixture was heated to 130-170^0, when a ll the volatile produots of the reaction distilled off, and were condensed using a cold finger filled with an Ice/salt mixture. The distillate was collected In a flask cooled in a so lid carbon dioxide/acetone slush bath# The dark-brown distillate contained, apart from carbonyl bromide, considerable quantities of free bromine, hydrogen bromide, and assorted sulphur compounds# The apparatus was open to the atmosphere via a drying bottle (concentrated sulphuric acid} to allow for the escape of carbon monoxide also formed in the reaction# The bromine was removed by the very slow addition of mercury while shaking and cooling the liquid* [This process must be carried out very carefully, since the large heat of reactio n between mercury and bromine can easily lead to decomposition of the carbonyl bromide COBr^d) ^ CO(g) * B r^ d ) ♦ (g) 3 The mercuric bromide formed • ta ile d off* to give a coating on the w all of the flask, while the liquid turned from dark brown to pal© yellow* The carbonyl bromide was distilled off at atmospheric pressure, the portion boiling between 61-65^0 being retained* 63 The la s t trac es of bromine were removed by the careful addition of email quantities of antimony powder, while cooling, and a further distillation at reduced pressure yielded colourless carbonyl bromide (boiling point at atmospheric pressure 54*5 - 0*5^0, literature value^^ 64-^5®G)* The y ield fo r the preparation varied between 20-40# of theoretical based on the equation C3r^ ♦ tOj —^ COBrg ♦ 8r^ , Carbonyl bromide is extremely poisonous, so all distill ations were carried out in a well-ventilated fume cupboard, and the apparatus was open to the atmosphere only via a con centrated sulphuric acid washbottlo (to exclude moisture) and a washbottle of concentrated alkali to decompose any carbonyl bromide vapour*- 23aOn ♦ COBTg —) 2Ba3r + 00% + H^O The carbonyl bromide was introduced into the vacuum lin e , degassed by repeated freezing and pumping, and stored in the dark at -196*0 (liquid nitrogen)* When a sample was required, it was fractionated through -45*0 (chlcroben&ene slush bath) to a trap a t -94*0 (toluene slu sh b ath ), a t «âalch temperature any hydrogen bromide impurity would be pumped away (vapour pressure^ approximately 500 to rr)# 64- The purity of the carbonyl bromide was tested by hydrolysing a sample with alkali, as described for aluminium halides above, and titrating by Volhard*s method, A ty p ic a l result, using a 0,339g sample made up to 200 ml solution was Br (found) 64,6# by mass Br (calculated) 65,1# by mass To test for the presence of free bromine, potassium iodide was added to a slightly acid poat-hydrolysis solution, followed by a drop of starch solution. The absence of any immediate blue colour due to starch-iodine complex showed the iodine concentration to be less than 2 x 10*% corresponding to a bromine content in carbonyl bromide of less than one half of one per cent. Commercial cazbonyl chloride (latheson Co, Ltd,), supplied in a cylinder under pressure, with stated purity 99.0# minimum, was further purified by fractionation as for carbonyl bromide. The cylinder was connected to the vacuum line by a length of reinforced polythene tubing, the joint either end being made vacuum tight by application of *forrseal* resin (Vsriaa Ltd), Boron tribromide was purified as described previously (p* )# OoEimerolaX aluminium chloride and aluminium bromide Ltd#), with minimum stated purities $8# and 99# respectively, were vacuum sublimed several times each, the f i r s t few times from a mixture w ith finely-divided aluminium powder to remove water and free halogen. k minimum of heating bad to be used for these sublimations, because cm. heating at low pressures, the reaction 2 A 1 ( s ) ♦ A Ilj(s) **—) 3Al%(g) tends to oceur^^. On condensing, the moaohalide dlspro- portionates to the trihalide and aluminium again, resulting in a transfer of aluminium metal to the final product# After sublimation, both trihalides were white, crystalline solids, though the trlbromide darkened slightly after several months even though stored in th e dark# The purities were tested by alkaline hydrolysis followed by halide and aluminium determ ination as described above# The re s u lts were AlGlj CI{found) 79#1# by mass Al(found) 20.3# by mass ClCcalo#) 79#7# by mass Al(calc#) 20.3# by mass AlBr, Br(found) 89#&#by mass Al(found) 10.2# by mass Br{calc#) 69#9# by mass Al(calc#) 10.1# bymass Attempts to purify commercial aluminium iodide {Alfa Inorganics $ no stated p u rity ) by sublimation resulted only in 66 further décomposition to aluminium and iodine# Likewise, preparation from the elements according to the method of ÛÙ Watt and H all resu lted only in diseoloured product# Probably this was. due to the difficulty in the method# which involved subliming iodine through heated aluminium granules# of packing the aluminium extremely ti^tly* Only one-third of the quoted weight of aluminium per unit volume of apparatus coaid be achieved. Therefore# tlia commercial product, supplied as a pink powder compared with the white colour of pure aluminium iodide, was used directly, since analysis showed the p u rity to be bettor than 99#, despite the d is colouration# AlXj 1(found) 93.2# by mass Al(fo%md)7#0# by mass I(calc# ) 93*4# by mass A l(oalc#) 6#6# bymass 67 Boron Tribromlde ■» Carbonyl Bromide On mixing boron trlbromide and carbonyl bromide, a eolourXeas solution resulted, which turned very slowly brown on standing in the light for several hours, due to decom position of carbonyl bromide. The rate of appeazmce of bromine colour was ekbout the same as for pure carbonyl bromide under the same conditions. The gaa phase spectra of boron tribromide, carbonyl bromide, and an approximate t i t gaseous mixture of the two are shorn in diagrams 13, 14 and 15 respectively, and the freq uencies listed in table 9. All frequencies are - § cm*\ The spectra of boron tribromide and carbonyl bromide agree very well with published speetra^^*^^, except for the presence of a PQB band at 1362 cm*^ In the carbonyl bromide opectrm. This is almost certainly due to the extremely strong fundamental (1352 ca"^) of sulphur dioxide^^, present as an impurity from the preparation. The Intensity of this band indicates that the sulphur dioxide was present in only small concentrations, and Indeed after the first few samples the band disappeared ( compare $ for example, with the carbonyl bromide/ boron tribromide speetxm, diagram 15). 62 9 R IA 5 E JMfM-KZO C? BOROH ?RI mOROVC \R K'NYZ; BRCHÎSS ft33I3K0SW CAS 3011 XL CAR3i;Hin. 30K0K T S l- ASSX3®E5T (fiE J . 61) 3 1 £ E Ig | Bsctaiii!/ SKCKISE cb“ ' 3CaOH S S I- -1 (ESF. 62) 3R0KIDR cm 2 »2 3633 m 3632 * ' -2 1822 8 1822 s 3 1730 ffi 1735 a 1715 w 1715 a 2 vj(^°a) 1633 a 1641 m 2 2 \ 1510 m 2 UgV ♦ X.*v 1480 a 1375 w 1373 a 2 v^(^’ a) s o , im parity 1362 a 4# 2 ♦ Vj 1214 w ## "4 ♦ ^1 1160 m 1132 a 1132 a 1093 a 1033 a v^+ Xj(1l3) #* 2 >^6 1013 w 1000 w ^3* 967 m 967 a 918 w 910 w 911 m j 2 X,* xgfllg) 855 r& 854 va : V (10a) ■ 8l8 va 615 va ; ( ti3 ) 786 v a . 704 va "l * *5 741 va 742 va " 4 575 w 574 m i -^3- x,(iOa) V 560 w «• 540 a 539 a j V j - V ^(^^3) \ ^6 512 a 512 w ■ \ ''1 427 a 423 m 60 CO < o LU O Q S s ÛÛ œ g o CÛ \ ü_ o § 5 3 q : cr (- LU CÛ ÜJ Z CL 3 LO LU Q LU CL O î 8 Ll Z co Z < Q CVi 100 60 20 TRANSMITTANCE »\' 70 LU O U ÜL. O in > o o o o o o o oo CD TRANSMITTANCE 7c o ■o LO ■o K“ Z o u z < cr o o < Q O O o LO 0 TRANSMITTANCE o LD O O cr o en O O o o LD TRANSMITTANCE o -o LO o o s “ CD LO > O o o o LO TRANSMITTANCE °A 74 The speotrum of the 111 mixture la clearly a aupezv im position of the two component spectra# whereas i f there was any association in the gas phase# shifts or changes in relative intensity would be expected in some of the bands# accompanied possibly by new bands attributable to a B ~ 0 bond* Compare# fo r example# with aluminium chloride/carbonyl chloride# where a band assi^ed to a shifted C«0 stretch appeared in the region 1600-1700 ki%k2^ Bliailar conclusions are reached from consideration of the spectra of the solids (diagram 16# 17# lB| table 10)# Although many of the bands show temperature shifts relative to the gas phase spectra# there are no shifts in the 111 mixture re la tiv e to the p ire components# fhis strongly suggests th a t no complex exists even at low temperatures {less than -lOO^C)# 11 The B nuclear magnetic resonance spectra again lend no support to the possibility of association# Samples of pure boron trlbromide and lit boron tribromide/carbonyl bromide both gave low-field shifts of 39 * 3 p#p#m# relative to boron t r i fluoride etherate# Ko corrections for changes in bulk magnetic susceptibility have been attempted# since these would be insignificant relative to the limits of accuracy of the measure ments# 7 ^ T4SI2 10 iïOlID »;ifw wm»-«MVT*w#eiwj mo:Uj3.V:ilWI3 TËimciMIDB 0AH30ÎÎYÎ. GAKaOHXIi 30H0H ÎK I- Ar.si3SMSnî SHOSIDS 320îai3f/ 330SIDS A3SI3S3SJT i om-1 30B0H THI- 8R0MI3S cm*^ 2 V2 3570 w *» ''2 1790 8 1792 s 1635 a 1683 a 2v,(^® 9) 1600 m 1602 a 2 v lO h ) 2 %4 1500 vw «a 2 »5 + »4 1430 m 1420 Bh 1403 m 1403 m 2 %,+ '^^0%) 2 Vj(1l3)1360 a 1365 a 2 Vj(1l3)1360 SOg impurity 1330 vw m 1 2 5 5 a 1 2 5 2 m ? 2 ♦ Vg 1210 w a» \ ♦ N 1140 a 1112 m 1103 m ^1* 1070 s 1057 8 2 ^ 6 1013 w 383 w 984 « 945 m 940 a v,+ i^i,('V3) 935 w 933 w 2 9 0 0 m 3 0 0 a 2 Vg(^^3) 890 w ■» 835 va 860-750 va y , (i° a ) 7 9 0 va . 3 ( ” b) "l ♦ ^ (8 0 0 - 7 0 0 T3 7 2 0 vs [ 700 va CO2 660 a 665 m 670 a OOg 5 0 6 m 505 a ^6 480 « 4 8 5 w V3 - ü,(1°3) 460 m 465 w v |- %,(lia) 420 w ^1 425 84 76 i TRANSMITTANCE "/( o o LO o o CJ I— z o CJ O GÛ CD CD < cr o < Q o o o o o LO TRANSMITTANCE % o Q .0 Li_ Zj 0 O o w LO 3 cr LU I— Q u LU T CL O 2 LO cr 0 CD Q LU >- cr cr z LU I o CÛ < 2: cr CD 3 Ll_ cr z z < LU o > 0 < .0 0 5 CO < cr o < 5 o o o LD TRANSMITTANCE Vo LO O o u LU O -O Œ) O •O O O O LO TRANSMITTANCE °/c Ll O 2 3 cr h- u ÜJ CL o • 8 csi Q >- LiJ cr < cr Ll CJ z LU m < LU X C l Q > o O U1 CO 00 2 < cr o < Q O o o LO o TRANSMITTANCE o o LO I— Z o o OD < c r o < Q O O o o o LO TRANSMITTANCE % Aluminium Bromide + Bromide On condensing a large exoeee (mole ra tio ca. 5*t) of carbonyl bromide on to aluminium bromide# a white solid mixture resulted whilst maintained at *196^0# indicating the absence of bromine* On allowing the mixture to warm# the carbonyl bromide melted# and the aluminium bromide floated to the surface of the liquid* fhe aluminium bromide slowly dissolved# but aa i t did so effervescence and the formation of a brown colour were both clearly visible at the solid/liquid interface* By the time room temperature had been reached# all the solid had dissolved# and the colour had spread to give a uniformly dark^brown solution* On pumping this solution at w#$^0# where pure carbonyl bromide has a vapour pressure of ca* 10 torr^^# a dark brown solid resulted# which obviously s till contained considerable quantities of free bromine* Therefore# the temperature was gradually raised# while still pumping# t i l l a white solid remained# This proved to be aluminium bromide# recovered almost quantitatively# two erperimsnts giving respectively $8*5# and 97*0# of the starting weight* This la in sharp eontmst to the preparation of aluminium chloride/carbonyl chloride co m p let\ where on pumping o ff excess carbonyl chloride at approximately 10 torr# the 111 complex remained# Although the aluminium bromide/carbonyl bromide mixture z o cr u. cr o O CL o < o > (N Ll o CJ CJ LlI CL if) Q ÜJ cr > ü_ Z CD 2 < cr o < Q o o o LO TRANSMITTANCE 0 g4 tam eà âarWbroim wltliln mlüut^a while miming to room temperature# a reference eample of pure esrhonyl bromide# kept a t room temperature under otherwise similar conditions# remained colourless# Evidently the aluminium brmide catalysed the decomposition of carbonyl bromide# aconclusion supported by the presence of a new# strong# band in the infra^ red spectrum of the vapour from the mixture (diagram 19)* Centred at 2150 ^ 10 cm^\ with well-defined ? and H branches and rotational fine structure# this band is attributable to carbon monoxide Even when maintained at -45^0# a freshly prepared mixture s t i l l turned brown, though much more slowly than a t room temperature# Aluminium Chloride + Carbonyl Bromide Several experiments were performed# w ith aüLumlnium chloride! carbonyl bromide TOtios ranging from 2 if to 1i30# On allowing some mixtures to slowly warm up from -19$®C# the reactio n appeared vistWLly identical to the aluminium bromido/ carbonyl bromide system# with so lid flo a tin g on the carbonyl bromide# and gradually dissolving# accompanied by effervescence and formation of bromine# to give a dark-brown solution# The appearance of bromine colour was extremely rapid# reaching its maximum intensity la well below five minutes whoa hand-heat was applied to bring the system to room temperature quickly. Even when maintained at -45^0# bromine waa e t i l l formed# but only very slowly# On pumping to dryness# also at -45^0# the residue ©bowed a weight increase f a r le ss than th at predicted fo r a l i t complex* Since the residue was brown# the we l ^ t increase was probably due# a t le a s t In part# to bromine# which at -45^0 is a solid with vapour pressure only c^» 1 torr^* This was confirmed by the appearance of brown vapour on warming the residue to room temperature* The re s u lt was t- S tartin g wt* AlCl^ » 0*0605 g f in a l wt* residue » 0*0654 g Hence# even a t these low temperatures and pressures, there is no evidence fo r an iso lab le complex* Two solutions# both containing a large excess of carbonyl bromide# were allowed to stand at room temperature for several hours# then pumped to dryness Uihm a white solid remained) and analysed* Besults were*- (a) (b) Ai # 0*33 X I0*^g*atom 0*73 x 10*“^ g# atom Cl a 0*40 X 10^^^ atom 0*71 % 10^^ g* atom Br « 0*46 X lO^^g-atom 1*62 x 10*^ g# atom Wt* residue (ealc)» 0*0662 g. 0*1757 g* Wt* residue(found)»0*0599 g* 0*1753 g* , S ta rtin g wt*AlClj«# 0*0469 g* 0*1244 g# S6 T».$î,E 11 Qg YAi-caa wm-' Kimumn Ciii/>aiDE/c^Ri»3Yr. s^mibb FREQUENCE enri ( # ^6 , 1 ) __ C0 % 2 3630 @ 2 ^ 2 I 2150 9 GO 1922 w ? 1 8 2 2 V3 . 1 0 0 %^ " 2 1780 a 3 ^5+ % G03r^ ; 1530 w COClg 1590 m 2 COClBr : ■ 1505 m 2 v^ 0 )8 r^ 1 1480 n 2^5+^4 008r^d I 1405 vw V,* V4 OOClg 1 1 3 1 8 m j V4 oociar 1 1213 m ' 2 V5 OOBr^ ! 1165 m coarg i V ' ' 4 1133 w Î 1 1094 w ccstj ! V ' ' 5 1085 w 2 V g ,. OOOlSr 1 1 0 2 2 coarg j m 2 ^ 6 960 w 1 Î i ' 920 m : 1 Vj*V 4 COSrg i 893 a { V5 cociar Î 651 m , ' I COClj J C0G13r■ i 8 0 8 v a % 785 v a ( a h ) ^ ' V,+ Vj costj ': 1 745 v @ ! ^4 : C03rg ! 1 610 m i COSfg 1 cociar j 548 m ^ 6 5 2 2 a COClBr i - ^ 1 512 8 Vg. COlÎTg 8 CGBTj 426 ' ’ 1 o o o LD ce Ll. cr 3 O Û_ Ll o O H- O LlJ CL. l/l CJ Q o LU “O o: ÇJ) < cr ü_ z o CSJ o 2 o < cr o < Q O o O LO TRANSMITTANCE %' Sotîi realduea correspond to AIX^ only, and also clearly demonstrate that halogen exchange has occurred, Further proof of exchange came from the in fra-red spectrum of the vapour from a room temperature so lu tio n , when hands for carhonyl bromide, carbonyl chlorobromide and carbonyl chloride were all clearly visible* The assignments are given in table 11, and part of a spectrum obtained at lower pressure (by expanding vapour into a larger volume) is shown in diagram 20 (compare with diagrams 14 and 21)# TJîhe much lower Intensity of carbonyl chloride bands compared to the others la presumably due to (a) The large excess of carbonyl bromide over aluminium chloride la the starting materials# (b) Oomplexing of carbonyl chloride with aluminium chloride# thus reducing i t s vapour pressure. Aluminium Bromide Carbonyl Chloride S im ilar results were obtained# with the solid dissolving, and formation of carbon monoxide and bromine# This occurred very rapidly at room temperature, but much more slowly a t -4^^G# On pumping two solutions (both with a large excess of carbonyl chloride) to dryness# one at •45^0# one at room temperature, the residues showed a decrease in weight compared to the starting aluminium bromide# ruling out any simple# tit complex formation# Results weret- (Brown residue) g# Initial wt# AlSr^ « 0#1250 Final wt# residue « 0*tlB3 Hoorn temperature I (White residue) Initial wt, AlBr^ » 0.2505 f in a l wt. residue «* 0*1453 Analysis of the room temperature residue corresponded to AlX^i A1 a 0.97 X 10*^ g# atom 01 w 2*72 X 10^^ g. atom Br « 0.25 % 10*^ g* atom Wt. residue (calc.) » 0,1428 g. Wt, residue (found) » 0*1453 g* In fra-red spectra of the vapour from a room temperature solution again showed the presence of carbonyl cMorlde# carbonyl ehlorobromlde# carbonyl bromide and carbon monoxide# all in appreciable quantities (diagram 22> table 12), The gas phase epeotnua of pure carbonyl chloride was run for comparison purposes (diagram 21), and gives good agreement with published spectra^^*^^. o ■8 Ll O S 3 cr h- o ÜJ LU CL U3 U Q LU cr < cr ü_ o z o U e n Z < cr o < Q O O LO TRANSMITTANCE 0 o -o u n O -o LU O -O CD > O -O O o o LD TRANSMITTANCE °/c o 'O LO 2 O cr Ll u ce 3 O >- CL Ll O :z 3 cr K u LU CL m O 00 o ^ Q Œ) ^ LU cr < cr 11. z OsJ (N 2 < cr o k—I< Q O O O LO TRANSMITTANCE 0 12 isPB\-gia gas.;,»g;cig3 op VAfom frou CAS ciü^hiee/ AiiîalBloa FKB,Ua4CÏ AS3IaSMiaT -1 ora ' (PSB. 61, . 64, J 5 I. 3G30 m 2 ^ 2 . COXg 2140 m 00 1825 s "^2 COX2 1680 a 2 0001.<5 1595 a 2 V4 • COGlBr 1405 a ^1* ^ 4 COClg 1310 m COGlSf V V4 1160 a 2 V,, 2 COClg 1080 w 2 Vg COGiar 1010 m COCl^ (r CJ z> o CL < o > -O Ll O :E %) cr I— CJ LU CL U1 Q LU cr > < CL Ll Z o o < CL o < Q O O o LO TRANSMITTANCE Vc 16 IPFRA-RE3 gHsggisîCig 01* yftpoim o.mnmxrj, cnt K îxnm jm loD im AssiGm&mr om«•t ' (Bsr. 64 aad 65) 3630 m 2 OOOlg 3 5 4 0 a %2 + 3 OOOlg 2 1 5 0 3 œ 2 0 7 0 m V ^ 4 - V ^ COClg 18 2 8 va V 2 GOGlg 1 6 8 0 3 2 COGlg 1422 m COGlg 1165 m .2 y,, 2 V g COClg 1015 s y ^ COClg 845 va y^ 0 0 0 1 ^ 745 w y^+y^ coGi^ 5 8 0 0 y 0 COCl^ 567 s y^ COClg 47 DiscanniON Seoaua# of the high olaotronegatlvltiea of the halogens, carbonyl halides might be predicted to aot m poor Lewis bases compared to other carbonyls, such aa organic kotones^^*^^# Such complexes as ex ist would be expected to be fairly unstable, a theory supported by the properties of the aluminium chloride/ carbonyl chloride complex# The present re su lts on boron tribromide/oarbonyl bromide give no evidence for complex formation, even at low temper^ atures, in acooi^ with Christens lindings^^ on the boron tr i* chloride/carbonyl clilorlde system# Ilozmially, boron 1mlides are considered b e tto r Lewis acids than aluminium halides (e#g# with phosphoryl chlorido^^ and ethylacetata^^)• Tet for carbonyl halides, while aluminium chloride forms an Iso lab le complex with phosgene, and there must be interaction la the other systems to account fo r halogen exoliangs and c a ta ly tic decomposition of carbonyl bromide, boron trichloride forms no complex, and boron trlbrcmide does not catalyse the same decomposition, (The boron tribromido/oarbonyl chloride and boron trichloride/carbonyl bromide systems remain to be studied), This discrepancy is probably due^^ to the need for boron to rehybridise from planar sp^ tetrahedral ep^ on coordination, involving the loss of considerable Tf-bond energy which has to be compensated by the formation of a new doaor-accoptor bond, fhua If the boron^onor bond energy is less than the rr-bond energy lo s s , no aâduct would be ezpectod* In the aluminium halides, however, the aluminium is already tetrahedrally coordinated, so the earns problem does not a ris e . As already stated, although no Isolable complexes were found, some Interaction in the systems alumimiuis bromide/ carbonyl bromide, aluminium chloride/carbonyl bromide and aluminium bromide/carbonyl chloride must be assumed to account for the halogen exchange and the catalytic decomposition of carbonyl bromide. The speed of halogen exchange, as \ exemplified by (a) • the quick appearance of strong infra-rad peaks for all three carbonyl halides, (b) the rapid production of bromine, is in contrast to the resu lts of Buston^^*^^, for solutions of similar concentrations to those used in the present work, he found half-lives for exchange between radioactive aluminium chloride and phosgene of 30 minutes or longer a t 25^0, In the present work, quantitative comparisons are difficult end very approximate, because of the need for cooling the systems initially to obtain the mixtures, which then slowly warm to room tem perature. On applying external heat, however, to 4 4 speed the attainment of room temperature, the bromine colour . seamed to reach maximum In te n sity in the order of 3 Blnutes# suggesting a faster exchange rate. The reason la not known, but the fast rate suggests that exchange probably occurs through an ionic rather than covalent mechaniam* for the concentration range involved, Huston postulated ionisation in the system aluminium chloride/carbonyl chloride. The presence of appreciable quantities of all three carbonyl halides, even after standing for several hours, suggests that the eschank'^e is an equilibrium process, Involving severe separate equilibria of the type AlCl, ♦ C03rgjf ^ CAlCly^A13r,] ♦ jr COCISr ^^AlBr. + COClg ca ♦ 3rg CAlGly'ilBrj] ♦ COClg ♦ OOSr^ * C0C13F CO ♦ Br^ In a closed apparatus, such an th a t u sed ,' equilibrium would be established, but la an open system the escape of carbon monoxide would p u ll the e q u ilib ria over, Even so , von Bartal*s claim^^ that only the reaction AlBr^ ^ 30001^ AlCl^ ♦ 300ClSr takes place seems a vast oversimplification, likewise, his claim, baaed on halide analysis, that with excess aluminium o o bromia#, aluminium chlorobromide is formed quantitatively also aaema dubious. Tiie nature of the aluminium halidsa in solutlouf and on pumplnj off carbonyl halldee, cannot ha jud#ad from tha résulta available. from the variation In analytical data, howavar, there must be a mixture of aluminium hvullden, whether the aim^le aluminium chlorldo/aluialnlum bromide mixture or Involving chamloally distinguishable mixed halides* In this ooaneotlon, the failure to obtain any useful or infra-red spectra Is unfortimate# Corbett and Gregory*^ assertlon^^ that the residue from aluminium bromide/phosgene Is an annealed mixture of aluminium chloride/ aluminium bromide Is merely qualitative, with no apparent experimental backing# The possibility romains that complexes may be formed a t temperaturos lower than at which temperatures halogen exchange and carbonyl bromide decomposition might be su ffic iently slow to enable complex Isolation# This Is most likely with carbonyl chloride, which Is liquid down to whereas carbonyl bromide solidifies at a higher temperature, estimated at -70 to *430^0, (although the exact freeal&g point Is not known) and there would be d iffic u lty In pumping o ff any excess* Sven i f complexes were Isolated at such temperatures, th e ir study would of eouz^e present great problems due to the need to maintain them at all times at the m m low temperatures to avoid decomposition* I tJ I TUBES OE GA&30RY& BBGMIS# ISee also M.S* Anthoney, Gardner and A. flnoh, J# Chem* TWrm* (1970) 1 697, a reprint cf which is included at the end of this thesiaj. o a. iKtaosuggi^a îhra® papers have been on the equili brium GO(g) ♦ Br^Cg) COlr^Câ) reporting equilibrium constants for temperatures ranging from to 456S:. Sines d (In Kn) _ %# AR^ -wwniw , d ( iA ) a the enthalpy of reaction and hence the standard enthalpy of , 71 formation for carbonyl bromide could be derived. tagman et.al# used îrautz*s (1915) resulte^^ to obtain ** -23.0 kcal moX**^ 1-96.2 kJ while Cox end lllcher^^ quota a *be@t* value of -20.1 kcal mol*^ 1-64*1 kJ mol“'^3 derived from Bclu^achcr and Bergmanswork (1931). Beerlnk'a reaults^^ (1S2S) lead to an Intermediate value of -20*9 kcal mol'^^ 1-67.4 k i mol'^^ j* In the absence of any further experimental data, and In view of the spread of values. It was decided to redetermine àn^ employing a different, calorlmetric, approach, The reactio n studied wia COBTgd) ♦ nl^o(l) — » COg(g) * UHBr.(n-1)ll203(l) a t 296*151 and one atmosphere pressure, for which the enthalpy change Is denoted as m hyd* By using water already saturated I o J w ith carbon dioxide for the hydrolyeIs, all the carbon dioxide produced In the reaction ©an he considered aa In I ts standard (gaseous) state. Since the enthalpies of formation of all the other reactan ts m â products are well known, measurement of AH hyd allows AllS 0 0 Br2 ( l ) to he calculated. Preliminary studies showed the reaction to he slow, 0.3339 g of carbonyl bromide la 200 ml of water taking ca forty minutes to dissolve completely* Since carbonyl bromide and water are Immiscible, and shaking speeded up the reactio n , the rate of hydrolysis la probably controlled by the surface area of the two liquids In contact. Analysis of the hydroly sa te gave 99.5^1 of the calculated bromide. Indicating that the reactio n goes to completion. In view of the slowness of re actio n , the use of an adiabatic solution calorimeter was considered easentlal* 04- Chemicals the carbonyl bromide was from the same batch used previously (p* 63} [Analysis# Br (found) 84.6# by mass, Br (calo ) 85,1# by mass]. Distilled water was saturated with carbon dioxide by adding a few lumps of dry Ice to a flask of water which was then placed in a thermostated bath at 298% to equilibrate. Calorimeter The calorim eter was basically similar to that described in Dart 4 (p, l2l)# but modified to operate Miabatioally* 75 Comprehensive detail® are already available The principle differences from the constant-temperature-envlronment c a lo ri meter described in the next chapter w ere;- (a) The calorimeter was immersed In a water jacket# the temperature o f which was automatically controlled such that the temperature difference between the inside of the calorimeter and the jacket was maintained at a constant# pre-set# level {m 0.06 - 0.002 degrees)# The design of the automatic adiabatic control was the same# with minor modifications# as that employed, in the Gallen& 74 A diabatic .Bomb Calorimeter- . /OS' (b) As part of the Miabatio control unit, the glass Dewar shown Im d lagr^ 26 (p./22j contained two additional thermistors# (c) A Soalamp galvanometer (1400 0 ) was used la place of the chart recorder as the eut-of-baXaace detector la the thermistor bridge for measuring the temperature Inside the calorimeter# Bead Inga of the thei^lstor resistance were taken every minute# The only m odifications In the present work of the apparatus and procedure given In reference Î3 were due to the fact that reproducible results were only obtained on eliminating all contact between carbonyl bromide and grease. The reason Is not kaom certainly, but one possible explanation Is absorption of carbonyl bromide vapour by the grease* ■ This was supported by ■ diserepaneles of up to 10# la the quantity of carbonyl bromide used as detci^laed by weighing and comparison with the results • of post-hydrolysis bromide analysis. fo overcome th is # - (1) The ampoule design was changed s lig h tly (diagram 24) to enable sealin g under vacuum. The narrowed, thickened portion facilitated sealing at a well-defined point, which meant th a t the ampoule volume could be determined beforehand by weighing first empty then filled to the constriction with water. A blank calorimeter run showed th a t the e ffe c t of heating the ampoule during sealing, and the enthalpy of ampoule breaking, were both negligible# , } 06 / P u DIAGRAM 24 COBr^ AMPOULE o / (11) No grease wm used la the calorimeter head. fee te showed that the two flangea of the glass Dewar were waterproof If held flî^ly together by a circular wire clip, and then lii^tly greased m the outside rim only* Nor was any grease used on the joint connecting the ampoule to the ampoule holder# Before and after using the calorimeter for carbonyl bromide, Its performance was checked by measuring the enthalpy of neutralisation of trle(hydroxymethyl)amlnomethane (TIIAM), { Aristar grade; 99.9# minimum purity), in excess aqueous 0*1% hydrochloric acid* , Eeeults were Before*- -7.11 kcal 1-29*75 kJ mol**^ j After*- -7.11 kcal mol*^ 1-29*75 kJ Previous THAW reeults^^ for this calorimeter gave a mean fo r ten runs of -7*11 kcal t-29.75 kJ mol^^ ] with a standard deviation of the mean of 0.005 kcal mol*"^ 10.02 kJ mol“"^J. The 1^0* TEAM mole ra tio varied from 1155 to 2227* These results compare with literature values of M W t l G S âa®kealaor%îi®kJ aoX”'* (BgO* TBAZiBATIO) 3-ann"^^ 373 • 1120 -7.107^.001 -.29,736^0.004 Sanaer aad Waiao*^® 1133-1462 -7.111^0.001 -29.753l0.004 Ojelwid ana ?'aaso^ 1320 - 1390 -7.112^.002 -29.757*0.008 « i l l S t. al,^® 1170- 1574 -7.109-0.001 -29.744^0.003 Procedure After Its volume had been determined, a dried mspoule was attached, via aa adaptor, to the greaseleas vacuum lin e described previously {p# 53 ), in place of the mixing vessel. The pressure was slowly reduced to minimise the danger of fracturing the fra g ile glass bulbs, and the ampoule tested fo r its ability to hold a pressure ^ 10'*'^ to r r , Carbonyl bromide was then condensed in by cooling the lower bulb in a solid carbon dlczlde/aoetone bath {-80^0), [The bulbs were found to crack at liquid nitrogen temperatures]. With the -80*C bath still in position, the ampoule was sealed at the con st riot ion, using as small a torch flame as possible. At -80^0, the low vapour pressure of acetone combined with the carbon dioxide gas evolved was sufficient to avoid ignition of the acetone* The carbonyl bromide was allowed to warm to room temperature, and the completion of the seal tested by the ability of the system above the constriction to hold pressures ^ 10*"^ to r r . The ampoule was removed from the vacuum lin e , attached to the holder by platinum w ire, and placed in position, together w ith the other inserts, in the glass Dewar. After pipetting in 200 ml of carbon dioxide-saturated water, the Dewar was lowered into the adiabatic jacket and allowed to equilibrate. The quantity of carbonyl bromide used was determined by analysis after the reaction and calibration* 10 ml samples from the 200 ml to ta l were titrated for bromide against 0*1N eliver nitrate solution, using eosla indicator* I I o ùa breaking ampoules of carbonyl bromide into water# the heavier carbonyl bromide sank to the bottom, where it was broken into small globules by the stirring* Effervescence was ' observed at the carbonyl bromide - water interfaces due to the production of carbon dioxide gas* " : The graph fo r m typical reaction (No, 3 in table: 14) is shorn in diagram 25, The calibration graph is similar except that the heating period corresponding to the reaction period is linear rather than.curved because the e le c tric heater supplies. heat a t a uniform rate# whereas in the reaction the ra te of heating decreases as the amount of carbonyl bromide decreases# Assuming the heat capacity of the ‘solution remains unchanged throughout (a v alid assumption fo r very d ilu te solutions such as those used) éïï reaction ^ ÀH calibration AT reaction AT calibration âU calibration « (Dower through heater) x (heating time) . whence êlî reaction = D xt _ 1 _ AT reaction (kcal mol"1 ) 41C4 ^ lï'"caliermticm M o LO Ll U U1 LlI h- ZD Z u Ljü :z .o CJ H- CD e n .O 3 .5 o CL fd rü _0 O ' O O O O O O OO CD LO CD THERMISTOR RESISTANCE ( O ) where ? m Power through beater (IxY joulea) t « beating time la « ho* moles COBTg usa& 6t s* Temperature rise 4 1 % 8» Oomrsrelon faetor# joules to koal* I t can be readily shorn (reference T3# p*22l) that* to m accuracy of 99*9^, for a thermistor B, èf...... reaction ...... ■BMW— # &lo% T? reaction lo% /B*> &T calibration âlog E calibration log where B. » Beaistanoe corresponding to temperature before reaction ^ 2 * ^ ^ ** ** after reaction. Bj w ® ** « « before ^ calibration. \ » •* « ^ « after Calibration. Since under adiabatic conditions the rate of heat loss can. be assumed constant (compare the similar slopes for fore^- and after^periods in diagram 25)# the resistances and carres- ponding to the températures after reaction and after calibration can be determined simply by extrapolating the relevant after- periods back to the times of ampoule breaking and switching on the heater respectively. lia. Table 14 gives the results for nine doterminâtioas of àîl h y d : n :igO(l) + Co^rgd) an co^(g) + L2mar. (m -i)iuo](i) fABlE !4 ENTHAt.Ff CP inrmOtYSIS 0? ÙARBÇnft M0 .MOLES —SH 0Î19 âfî oorr n —&E hyd , GOdrg usmi) {Jtcalmol~ ) (Steal sal* ) (KOtS RATIO (koal moT (% 10^) HgO 1 COlrg) 1 1,982 49.01 0.05 5537 43.95 2 1.028 49.06 0.12 10777 43,94 3 1.194 i 49.49 ; 0.11 9277 49.33 4 1.154 I 48.37 : 0.09 9594 46.28 3 1.197 Ï 43,56 ( 0.11 : 9255 43.45 6 1.359 ; 43.16 j 0.10 ! 8141 49.06 7 1 1.779 i 43.63 } 0.06 6224 49.62 8 1 2.056 j 49.61 ! 0.05 4384 ■ 49,56 9 ! 0.450 49.61 : 0.30 24615 43,31 The OQlusm headed &H eorr gives the oorreotlons applied to the observed enthalpies of reaction to allow for some carbonyl bromide being present as vapour. The la te n t heat of vaporisation#' &II vap* and the vapour pressure, I", of carbonyl bromide are both known a t 298%*^ £T#4 kcal and ISO torr respectively], and by measuring the volume# ? , ofthe ampoule (neglecting the small volume of COf3r2(l))# the correction becornea àU c e rr «» van % P % V « 7*56 % 10*^ V koal mol’*'* m X E X 293.15 ^ where V « Volume in ml of ampoule m m Ko# moles COBr^ used H M Gaa constant, 1.99 x 10**^ kcal K’*'*mol'~'* or 62.360 x 10 ^ ml torr In calculating &H hyd, thermal effects arising from the decreased solubility of carbon dioxide with the small temp erature rise during reactio n , and also from themixing of COgCa^) and Kdr(aq) were ignored# Kor was accounttaken of the increase in.pressure due to the carbon dioxide produced, from table 1%, a mean value cf âH hvâ ■» -49.06 î O.lS kcal «ol*^ 1-205.36 - 0.67 kJ is Obtained. The uncertainty interval is calculated as - one standard deviation of the mean# ^ cf n (n-1) 7Q following the recommendations of losslni * 11^ Standard Enthalpy of formation of Carbonyl Bromide From the equation for the hydrolysis of carbonyl bromide, it is apparent that the standard enthalpy of formation, âE^, of carbonyl bromide is given by c o a rg d ) m Êî^ cOgCg) ♦ z&n^ laar la | (a-llBgO] -sa| i^O - &a hyd. 7l Oslag accepted value# , âîl® COgCs) m -94.05-0.01 koal æorU-395.51^0.04 kJrasol"h Ah2 H,G(1 ) --68.315^0.001 kcal fflor’C-235.850i0.004 , la è(n-1)a20j--23.0li0.01 koalfflol"^ 1-121.3S-0.04Wmof (mean value between a » 5000 and a '» 25000) th is gives aa| C03fg(l) « -34.70 * 0.21 kcal »ol"^ 1-145.1& - 0.83 kJ ®ol“l -|tTi~r* I rill nuiT i—nflYr. i— r r "r * .**,##*#.# #i#iimm* imwa Using Echumacher and tenher*#^^ value of 7*4 kcal (no errors assignable) for the latent heat of vaporization, th is gives âH| C03r^(g) « -27.3 koal mol"'* L-114.2 kJ mol"^] ||W ^ 11 niiiji (rtwi ii.iiri^ differing markedly from the values derived by equilibrium 8tudiea^^*^^*^^L The reason for this is not knom definitely, but in view of the differences between equilibrium constants 1 I s fo r sim ilar temperatures quoted la the three papers# one possibility is that tru e equilibrium was not achieved in eome or all of the experiments, This would he in aooord with the fa c t that the decomposition of carbonyl bromide (and hence presumably the attainment of equilibrium) is known tobe very slow. Thermodynamic Functione for Carbonyl Halides Overend and Bvans^^ calculated the s ta tis tic a l thermo dynamic functions for carbonyl bromide from infra-red frequency measurements# using the harmonic c s o llla to r/rigid rotor ap.;roxl- Bù mations « This gave 8* COargLg), 293.15% - 73.82 oaX K"'s8X“’ UaXmg the following valuôo S® C (grE#hlte) « 1.3609 CïU. K”^aol”^ i 8® Og(g) # 24.502 eaX S® argd) « 36.4 cal S® 3rg(3) • 53.639 e a l «.''’mol*'’ the standard entropy of formation of carbonyl bromide Is derived âsj COar^tg) w 11#6 cal K"^mol"** [standard state Br^d)] - 10.7 cal K"^mol^\ [standard state Br^Cg)] from and &H^# âtfj 00^1r2(g) « -30*7ical mol - t These values are compared with for carbonyl fluoride and carbonyl chloride in table 15, TABLE 15 TESmiO.DfKAMIG FUI^CTIOKS FOR OOXg(g) SOF^ COClj COdTg -ûiij koal mol*^ 151.7 52.3 27.3 koal 143.0 40*9 30.7 -ASg eal 12.4 11.4 - 11,6 (10.7) The reason for deriving the entropy of formation of carbonyl bromide from both liquid and gaseous bromine can be seen from table 15# An apparently anomalous value for the standard entropy of formation is due to the condensed standard s ta te of bromine; as shown, a steady progression is obtained using the value for bromine gas. Carbonyl Iodide It is interesting at this stage to consider how table 15 might be extended to include carbonyl iodidt, OOIg, There are no reports in the literature of its preparation# and the results in Part Two (p, ^4-) suggest it la very unstable even at --45^C# Molecular models using * planar* carbon are easily made, end there la no reason to assume adverse structural features; hence considérât ion of its thermodynamic stability is appropriate. 117 The f i r s t need la fo r an estimate of fo r carbonyl iodide. Although many bond-eoergy echemee have been devised to assess none of them is appropriate la the present example because mo "correctiom parameters** are available to allow fo r two halogen atoms being bonded to the same carbonyl group, a situation unique to the carbonyl halides# As an example, the values for Aii|(g) calculated (without correction fa c to rs) tm m the schemes of laldier?^#^^, Allen^^*^^ and are compared with the experimental' values in table 16. TABLE 16 -AE^. CO%g(g), mTIMATED'TBOM BOîlD-mmif 80HEMI3 (kcal mol"^) GSPg COSlg eosr. Experimental 151.7 52.3 27.3 Laldler?2'82 119.3 38.2 16.0 A U eJ2.83 102.7 31.3 8.4 Cox®'* 115.5 37.5 11.5 85 In view of the email differences in C«0 bond lengths in going from COf^ to CO.Sr^, me shown in table 17, (compare with GOSg or C0 (CHj) 2 for example), a reasonable starting point would seem to be the assumption of a constant bond energy fo r the carbonyl group. I S 17 IE CO%.(g) £a#m.] QgsQ G-X C0?2 0.117 - 0.002 0.132 - 0.002 COCl^c; 0.1165 - 0.0002 0.1745 - 0.0004 COBTg 0.113 - 0.002 0.205 - 0.004 COEg 0.121 3 ) 2 0.124 From the cycle I: DISSOCIATION c(g) ♦ 0(g) ♦ 21(g) tli atomization C(graphite) * * X^iBtmô^avû s ta te ) 00X2 (3 ) •* &B stoaslzation - 25(0-1 > - B(0=0) where 2{C-X) is the mean bond d isso ciatio n energy fo r the C-X bond, and S(0«0) Is defined by the relation 2{C«0) w Dissociation Energies 2î(C-X) Using values from reference 71, AH atomization % # F 258.60 kcal mol"^ Cl 2 8 9 ,0 1 kcal mol"^ Br 2 8 4 . 3 2 kcal mol"^ The Values of D(C-Z) used were those fo r ©lace# despite the change tn. hybridisation of the carbon atom (which would be expected to have only a small e f f e c t ^ ) , of a ll the simple organic halides the bond lengtbs^^ in came closest to those in COJ^ (table 18) TABLE 1 8 CX^ aOKD LERaTia AED MEAE DISSOOIAIIDQ m m G f Bon length (mm) D(O-x) kcal mol"^ CF^(g) 0.1323 1 0.0003 117.0 CCl^(g) 0,1766 t 0,0003 73.1 CBr^(g) 0.1942 - 0.0003 64.7 Using the values for I(O-Z) in table lB# S(0aO) OOf^ # 1-86 kcal mcl"^ OCCl^a 183 kcal mol"? COBr^» 1 8 2 kcal mol"? la accord with the initial assumption. Despite the approximate and empirical nature of the method, the good agreement obtained justifies extending i t to carbonyl iodide, where AH atomization « 231,9 kcal mol"? D(CwO) «at 1 S3 kcal mol"? from above. Unfortunately# data fo r CX^ are not available# ao D(G-Z) has to be extrapolated from values for CH^Ï and by com parison with the other halides (ta b le 19), a o TA3DS 19 SÎA3 C-X B0Î5I! iJlîîSoOIâîIOa BiEH3ï FOR TSSAI, SIS30CIAÎI3Î! a(o-x) -1 kcal mol "SRSy kcal « 01—1 kQgil woX~ C3r^ 19 233.3 64.7 C'HSr, 4 299.6 67.7 CHgBFg #Mt CH,3r —3*4 352.7 63.1 —24*6 312.2 73.1 CCI4 CHGl, -24.7 335.3 79.6 Cl^Olg - 22.1 355.7 80.6 Ga,ci -13.3 376.0 81,4 c a ij 2?.0 299.6 5 2 . 6 eUji 3.1 350.0 55.4 [Eelevant values fo r D(G-H) taken from laldler^^ and Oo%**j.84 This gives D(O-i) oa 50 koal mol"? from wSiioh Ah| CQI^(g) « -3 - 5 kcal mol"? t —13 — 20 kJ mol ?3 ai Oalttg Kubascîiewski m û Evans* procedure fo r the estimation of entropies sfg3 COIgCg) «« 39.0 ♦ 0.34M - 6.2 X (c al where M *» Moleeul&r weight OOÎg# 28l,Sl leads to Sfgs OOIgCs) - 86 cal K’ ^aol’ ’ whence Ê s| 0 0 1 3 ( 3 ) » 32 cal K"Lo1"^ 1-2 eal Biol”^ i t star.âard State of lollne la taken as gaa] and &a| COIg(g) » -13 - ça 1 0 Steal aol"'* Clearly, however, carbonyl Iodide would be expected to exist as a condensed phase in its standard state* By com parison with the other carbonyl halides, the latent heat of vaporization is likely to be ca 8 koal mol"? [GO?* ea 4 kcal mol"? CQCl^ 6 kcal mol”?, CQBr^ 7.4 koal mol"?]^?* The entropy of vaporisation for molecular compounds is usually close •hr to tb@ Trouton figure of 22 cal mol" , giving COIg (1) ça -11 kcal mol"* às| celj (1) ça 10 cal K"*mol"* â a | COÏg (1) ça -14 kcal mol"* I a. a. Dissociation Equilibria In view of the great variation in stabilities for ths carbonyl halides, it la interesting to compare the equilibrium constant© at 293,15% for the dissociation 002^(g or X) # CO(g) + In (a-a) Th( nocsasary Xata ara glvea la tabls 20. ftiv E 20 ziiTHktrtt saeHoPY a ss gBisB msxQt ct nissoci^Tioî? TOR CAsgostî. H tsism AH dies* A3 disc* â3 dis©* kcal mot- 1 cal E"?mol"? kcal mol-1 COFgCg) 125,3 33.7 115.2 COClgfg) 25.9 32»7 16.1 COSrgCg) 0.9 9.8 —2.1 ( 1 ) 8,3 31.8 - 1,2 COIgCg) -23 -11 -20 (1 ) -8 11 -19 From these values, the equilibrium constants, %p, fo r dlssoclatim are derived as G0F2(g) 4,7 X 10,-95 cocigCg) 1,6 X 10-12 OOargCg) 34,4 ( 1 ) 7,6 14 COI^(g) 4 X 10 13 ( 1 ) 8 X 10 a s r&RT foim ihsssihisàl - z s â s s n s . Of fj_x!mnmi mxxnm I 2 4 INTRODUCTION Though the literature eoataine several determlaations of the standard enthalpy of formation of aluminium chloride, the values for aluminium bromide and aluminium Iodide have been much le ss thoroughly investigated# and ro ly mostly on f a ir ly old work* Table 21 lists the literature data. 21 lI?e%A?UR% VAtUgS ?0% TH5 0# OP (c) Real mol*'* 1 i AUTmOBS REFER*.! TEAR AlCl* Al.Br* A ll, ■ ÎS5SP. ENC23 1 m ^ m « •' K B erthelot 87, 71 1878 a 1170.6 126,8 77.6 .252 '' -e* 124,6 75.9 232 1-! klemm and 68, 9 0 1931) 167 121 71 2 9 3 Tanks ) “ I1 ,Klemm and 89 1932) \ Jacobi [ f ’ T 1 #& Slemonsen 91 1951 1 à !1167.5 * 2 9 s Dear and kley 96, 1954 ! 7 5 . 2 2 3 3 81, 99 1 Corbett 9 2 , 99 ! 1954 : Î • 73.7 293 and ' 1 Gregory i' ' .Coughlin l93 , 1953 i ^ i 163,97 2 9 3 . ! iGross and ^4 ‘ 1967 1 h 1 168*60 m mt 2 9 3 [iiaymaa ! * The nature of the crystal structures is irrelevant to the present work* «AIX3 " is used only as an abbreviation for "aluminium halide", and is not intended to suggest that the c ry sta ls consist of All* units* 7 .5 (a) Calculated using iertbelot^s values for the enthalpies of solution in water at 232 K, AlCljCe) -76,3 kcal aol"'' âlBr^(c) -85,3 kcal aol*'’ Allj (e) -89.0 kcal sol"' and 6 h2 Al^*(a%) and ^u2 x"(a^) from reference 71, aasusing B-0 , AU^(c) ^ ai^*(a%) + 3X'{a^) Lafter the JftHA? tables®^). (b) Berthelot also measured the enthalpies of solutioni A lSrj(o) * [JKOl.aSgOJfl) -86.9 kcal sol"') ( i) AlCljCo) ♦ tSkdr.aligO^d) -76.0 heal m cl"')(ll) AlIjCc) ♦ L3XCl.an203{l) -89,1 kcal mol"') (1) AlCl^fo) ♦ L3KI.nîÎ203{l) -76.6 kcalE o l“ ' ) (11) where m # ^ 1000* From these results, using the value fo r AXCl^(e) from Gross and Haymaa , and assuming (a fte r Berthelot) that the resultant solutions for each pair (i) and ( i i ) are id e n tic a l, AlBr^(c) » 6B (ii) * âH(i) ♦ 5 ÇA:yBr(a%) * êflf îCCl{aq)3 * £H| klClj (c) Klemm end Tanka œeasurad tha enthaiplsa of eolation; AH 80l“* AlCl^Co) ♦ pKCl.a lOOiÆHCaJ(l) -71,7 koal mol' AlCl^Cc) ■*■ tJKBr.ti 100i2iCl3Cl) -71,8 kcal mol' AlCl^Ce) ♦ L3KI.n lOOSiClJ(l) -72.3 kcal mol' A13rj{e) ■*• L3kCl.a lOOrfitClJd) -85.5 koal mol' A llj(o ) ♦ L3SC1.R 1OO£K0l3(l) -83.9 koal mol" where n represents a large, unspecified, eaoees# Using literature valuefor the enthalpy of réaction Al(c} ♦ U 100?® C lK l)------^ [AlGl,(n-3) 10OSIA] (1 H,(g) iHAl * they obtained AIX^ as m2&l%2(*)*Aa*0ln. + &3ai-d6B2BGl(8q) -d&Hg. %%(&%) + )&%% KC1(#%) (d) SieaoBsen used eolation calorimetry similar to (c)* (e ) Dear and Mey measured the enthalpy of eolation of aluminium iodide in water 1^1*4 * 1.9 kcal mol*^J, then followed the same procedure as in (a)# (f ) Corbett and Gregory measured the enthalpy of reaction AllgCe) + 3HCl(g) -----> AlClj(c) + 3Hl(g) Knowing th e other enthalpies of formation, Ah|â1Î^(c) was evaluated# 2.7 Coughlin measured the enthalpy of solution of aluminium chloride and aluminium in 41 hydrochloric acid , sim ilar to (o ). (h) Grose and Layman used combustion caloriastry to measure directly the enthalpy of formation of aluminium chloride Al(o) ♦ ^/aCIjts) ----- ) AlClj(e) Despite the seeming measure of agreement, there are several reasons why many of the values are open to doubt i (1) On dissolving aluminium halides in water or acid, hydrolysis results in the formation of considerable acid m ist above the solutions, this being more marked fo r the bromide and Iodide than the chloride (see references 87 and 83), This is bound to affect the accuracy of oalorimetrie re s u lts , especially early work where only one or two experiments were performed by each worker to determine enthalpies of solution* (2) The hydrolysis, particularly fo r the bromide and iodide, la elow^^ (see also Part Two, p, 61}, while all the calcul* at ions assume that the final state is attained almost instant* aneously# (3) The enthalpy of hydrolysis might be expected to depend cn the aluminium halide to water ratio# Indeed, for aluminium chloride, the enthalpy of dilution has been shown * to be appreciable and n egative over a large concentration range, due to hydrolysis. This is in contrast to moat common electrol* ytes, which have only small enthalpies of dilution, positive to ea 1M, and negative la more dilute solution, (4) The aeeumptioa ia made ia some methods that eolutloa yields simple, a.iuatad Al"'^ and 3% ions. In fa c t, the situation is very complicated, for the nature of aluminium halides in solution is very sensitive to both pH and concen tration changes, and the literature contains many papers postulating species such as lAl(OH)(îlpO)^3^'*‘ (98, 99, 100 , 106), AlgtOHlgOl (102), LAl(a20)^(CH)j2^* (1C1), lAlCOiUCl^Jg (103) e tc , in n eu tral c r acid so lu tio n . Only for very dilute solutions at high or low pH does the situation seem at all clear-cut? for dilute solutions in strong acid (pH 4 2) the cation is* predominant^while fo r dilute solutions in alkali (pH ^ 13) the anion AI(OH)^'^ is exclusively present^^^* The relevance of these points can be demonstrated by comparing the literature values for the enthalpies of solution of aluminium halides in water (table 22), For aluminium chloride, these values compare with an enthalpy of so lu tio n in 4.360M hydrochloric acid of - 72,31 - 0,03 kcal mol~^ (303K)^^, whereas for simple solution and complete dissociation to Al^^ ♦ 3 Cl"", the effect of the hydrochloric acid would be expected to be only slight. T A 3E 2 2 2 EZiTHAEPIgrX OF SOWTION OF HALIDr^ IN WATER AUTHORS itDFm- AlXjiBgO - m so w vio a ( k e a la o r ’) TMP ^01% RATIO AlCl^ A llj £ le rt he lo t 87 oa 1000 75.5 85.3 89.0 282 Thomsen 107 1230 75.5 291 Eoth and Buchner 108 2300-4730 73.1 293 Roth and 109 4 OGO-9 OGO 73.5 2 9 3 Bor go r Plotnikov 111, 2993 73.2 91.5 288 and I&kuboon 110 Dear and 96 ca 10,000 91,4 2 9 8 ZLey I t has also bean th a t the enthalpy of formation (and hence the enthalpy of solution) of aXomialura chloride in hydrochlorie acid is a linear function of the acid molality, with a spread of -7*6 kcal mol"*'* from m # 0 to m m 3# In view of the uncertainty, it was decided to redetermine the enthalpies of fom at ion of aluminium bromide and iodide by measuring the enthalpies of solaSioa in aqueous alkali, quan tities not previously determined, The method used was to measure the enthalpy of solution of aluminium bromide (or iodide) in 0,11 sodium hydroxide solution containing sodium 3 0 chloride, and the enthalpy of solution of aluminium chloride in alkali containing sodium bromide (or iodide) (1) AlX^(o)fWaOH4^3mGlj(aq) — 1) lBaAl(0H).43NaCl +3Na% + (n-4 )NaOH] (aq) (2) AlClj(o)fWiaOn*3NaX](aq) —^ BaAl(GR)^43NaCl (n-4)NaOR](aq) whence âîi|AlXj(e) » £^2 ~ * 3âli| HaX{à Besides the fa c t that aluminium Is present in only one form, th is method has the advantage over other solution methods that the acid mist formed above the solution is greatly reduced. I 3 EXPERIMENTAT! Chemicals same stocka of aluminium halides were used aa described in Part Two (p, 65"}. Analar grade (3#D*H. Ltd*) sodium chloride, ©odium bromide and sodium iodide (stated purities 99*9v^# 99# and 99*5# minimum respectively) were dried in an oven a t 120^0 before use* 0.1DON sodium hydroxide (B.d.H* volumetric eolation) was used directly. Apparatus Preliminary experiments showed the reactions to be extremely rapid, suited to conetant*temperature-envirozment (OTl) calorimetry# The calorimeter was an all-glass Dewar type (diagram 26), which has been used successfully in the Department for several y earst and is fully described e l s e w h e r e ^ Since a ty p ical experiment imvolvee long periods of waiting for equili b ra tio n , experiments were performed in overlapping pairs using two separate calorimeters. Minor modifications to the original design and the procedure for using the calorimeters are well described by Coley^^^* ' Temperature changea were monitored by an F53 thermistor (standard Telephone and Cable ltd*) connected as one arm of a 3 3- DIAGRAM 26 C.T.E. CALORIMETER (33 WbeatatoB* bridge# and with a resistance change of ca I 8 OQ per degree a t 2$*C. The only modification of Coley * 0 procedura^**^ was that the out-of-^balance bridge potential registered by a chart recorder (Bryans ltd ; model 27000) in place of a galvanometer. It can readily be shown (Appendix One, p#(27) that the displacement from sera shown by the recorder is linear with respect to changes in thermistor resistance# and by suitable adjustment of the recorder senait- ivity# resistances could be read off directly from a scale of 1 0 m l cm of chart paper. The Sample was contained in a thin^waXled glass ampoule# made with two small# very thin# glass bulbs, one above the level of sample in the ampoule and one below# such th a t on b rew ing the two bulbs calorimeter liquid could flow through the ampoule to ensure complete mixing. The ampoule was f ille d with aluminium halide In a dry-box# being weighed before and after f i l l in g . The ampoule of aluminium halide ms broken in to 200 ml of 0 . 1001$ sodium hydroxide containing sodium chloride o r sodium bromide (sodium io dide). Calibration was achieved by m% e le c tr ic heater of measured power (ca one watt)# consisting of a sp ira l of wire sited in a thin-*walled glass well f ille d with light oil of good thermal conductivity. The system was returned to the bath temperature (25.00 » O.Ot^O) between reaction and calibration by the addition of :?4- small pieces of solid carbon dioxide into the cooling tube. Procedure From Part Three, p. Ho , B &H reaction w f jt P % t ^ E- x 4184 w %/B ^ /c-ii-cn-T' c r ^ m w Molecular weight of sample# w « Weight of .sample used. The resistance-time plots for pro- and post-reaction, and pre- and post-calibration, periods can be regarded as linear for times less than cm ten minutes (Appendix Two, p . '2^). As the reaction was quick (complete in under 0.5 min, ), the values for the resistances before and after reaction, and Bg, could be obtained by extrapolating the straight fo re - . and'sfter-parioda to the time of ampoule breaking. In drawing the line for the afte r-p e rio d , the ty p ic a l **overshoot** which occurred for ^ one minute immediately following the reaction, was i.gnored. ' ; .During this time, the system was not at equili brium, the temperature of the calorimetric liq u id being higher than th a t of the glass, due to the former being stirred and having a .much'better theKsal capacity. however, electrical heating during calibration took ca four minutes, and allowance had to be made for the heat lost during this period when extrapolating the fo re - and a f te r - periods* Thlo wxb achieved by Slckinson*a. "equal areas" method**^*115 11â**’'* t *.l 7 la the present olrouæstances, this is simplified by the linear form of the resistance-time plot during the heating period ( ignoring the times needed for equilibration on switching the heater on and off, which are both ©mall and also tend to cancel each other), and are obtained by extrapolating the fore- and after-periods to the time corresponding to the halfway point between the in te r- ©eotions of the lin e a r heating period with the extended fo re - and after-periods (see Appendix Three and diagram 27)* The performance of the two calorimeters was checked periodically using the enthalpy of neutralisation of TEAM, as described in Part Three, p .f* ?. The results were 1st C«lorlm«ter 2r>4 Calorimeter Wt. ') -4H(feO »1 ) wt. ?5Am(%) -4H(kcal mol*^) 1 . 1 0 3 a 7.17 1.1S70 7,17 1 , 0 2 4 s 7,03 0.9289 7*10 Î.053S 8.05 1 .0 7 7 0 7*18 1.0757 7,03 1.0413 7.07 1.0567 7,11 0 . 3 3 2 4 7*10 1,0554 7.16 0.9954 7.11 0,9293 7,13 -1,1229 7.13 Ignoring run 3 for set 1, which lies outside three standard deviations from the mean, th is gives ( 76 1st Calorimeter 2nd Calorimeter Mean m -7*ti koal mol -1 Mean *» -7,12 koal mol^^ Standard deviation w 0,021 Standard deviation *» 0,015 of the mean koal mol*"^ of the mean koal mol-1 37 In a l l the reactions, the eolld aluminium halide dieaolved rapidly. In a l l casea, despite the alkali, some acid mist was formed above the solution in the ealorimeter, th is being most pronounced for aluminium bromide. However, the mist quickly dissolved, and bad no apparent effect on the reaction trace. Analyses of the end products for halide con firmed that very little, if any, halide was lost. AlBr. + maO^/NaOl Cl) Br (found) 1.46x10*^ g.atom (ii) Br(found) I.SOxlO^^ g. atom 3r (calo) 1.47x10^^ g.atom BrCcalc) 1,62x10*^^ g. atom 99.3# 98.6# All^ ^ EaOB/BaCl (1) 1( found) 1.42x10*^ g.atom (11) I (found) 1.22x10*^ g, atom I(calc) 1.43x10*3 g.atom l(cale) 1.23x10*3 g * atom 99.3# 99.2# AlOlj + BaClt/Eal (i) Cl(found) 2.63x10*3 g,atom (ii) 01 (found) 2.54x10*3 g.atom Cl(calo) 2.70x10*3 g.atom 01(calc) 2.56x10*3 g.atom 98.1# ^2# The re su lts are given in tables 23-26* 3 g 23 ERTHALPY 0? BEAOTirR FOR t,TA^ Al3r^(c) + W&OH 4^ 3aaCl](aq) —: 4^ 3Na3r ♦ 3#aCl + (a~4)B%G5](aq) 1st cxtonx'-szTsn 2ild CAiOHI:.îEïsa n , Al-3r, n —63g A13r, n , 1 Kcal raol” ' ...... Kcal raol ' 0.1113 47.7 93.03 0.0758 70.4 101.45 0.1336 33.5 101.47 0.1481 36.0 101.50 0.1756 30.4 100.85 0.1472 35.2 100.81 0.1728 30.9 102,73 0.1639 32.5 100.39 0.0307 53.3 101.43 0.1213 43.8 102.50 ■ 0.0886 60.2 102.33 0.1445 35,9 99.95 ■ 0.1308 40.8 101,57 ' Rua one ia the 1st Calorimeter is greater than three standard deviations from the mean, and was henoe ignored. Mean » -101.46 koal mol*^ t-424.50 kJ Standard deviation 0,22 kcal mol*\0.92 kJ mof^j of th e mean Standard deviation 3? „ » 0.77 koal mol*^D*22 kJ mof4-1 U frtm the mean 11 71 AlCl^(c) + UiBaOH ♦ 2#a3rj(aq) l#aAl(On)^ + 35aGl ♦ 3HaSr ♦ (n-4)EaOH3(aq) le t GALOlU^ETEa 2na CAlOHIiETSa wt, AlCl^ n . ah| n, A ici, B “ Ali? -1 ...... -.^ -...-..... koal mol . P I- . ^ ■ ...... |...|n kcal mol 0*1173 22,7 : 33,82 0.0478 55.8 85.80 ! 0.1090 24.5 85.35 0.0505 44.1 33.43 0,0916 } 29,1 83.70 0.1202 22.2 84.59 0,0602 ; 44,3 84.87 '0.0929 23.7 34.70 0,0831 32,1 83.54 0.0799 33.4 84.60 0.1103 24.1 33.97 Mean « -84*40 kcal mol*^ L-333*t4 kJ Standard deviation of the mean « « 0,23 koal mol*^ 10,96 kj mol*^j 11 X 10 Standard deviation m from the mean ■hâL. a 0,77 kcal BOl"'’i5.22 kJ 10 I 4-0 TASL8 25 ËNTHATIT OF FOE AH"F), WaAi(on)^+3:%WNaai f (ri-4)EaOHjaq 1 st CAlOBIiîEïEa 2nd OALoa: A llj a -&H® m . A ll, n ...... S__ ___ kcal mol*' __ A...... ___ kcal mol** : 0.0753 ; 108.3 106.30 0.3047 i 25.8 105.40 0,2527 32.3 106.14 0.3124 i 2 6 .l| 105.27 0.2735 2 9 . 8 1 105.43 0.1824 i 4 4 . 7 ] 105.13 0.2351 34.7 ! 105.67 0.2484 Î32.8! 105.57 : 0.2340 34.3 106.13 0.1950 141*8 105.76 0.1680 43.5 ! 105.33 Mean w -105,89 keal mol*^ 1-443.04 kJ Standard deviation 1.525 0.12 kcal mol*'* 10,50 kJ mcf *J of the mean, It X 10 Standard deviation - 0.39 keaX fflOl"*Ll.63kJraol"'*j from the mean 10 141 26 0? mAOfim was. AliS AlClj(o) ♦ LnSaGB*35al3ai [RaAl(oa)^*3KaI+3:aCl+(n-4) BaOllJaq 1st CA1031%:zrm 2nd CALÔE1 3 t, AlCl* n 1 S t., AlCl^ B I *^% kcal fflol ^ksal mol* I 0,0363 50,7: 84.53 0.0533 5 0 . 0 I 83.57 ! 0,0919 29.5 8 4 . 6 3 0.0743 35.9 1 85.25 I 0,0765 5 4 , 9 ; 82.53 0 . 0 9 5 3 27.8 1 83.04 = 0,1455 13,6 85.02 0.1164 22.9 j 84.23 1 0,1031 24,7 83.79 0.1235 20.# I 8 4 . 9 2 ; 0,1202 22,2 84.64 0.1133 23.4 1 84.95 Mean *» —84* 19 kû'î l mol** L- 3 5 2 . 2 5 kJ mol** J Standard deviation M fM » o;28 koal îac-l"*ll;l7 kJmoM of the mean 12 X 11 Standard deviation 10,26 « 0,97 koal mol*^£4,04 kJ mof’j from the mean 11 4- S_ CD OJ OD CO 00 00 00 o 00 * r TD ÜJ 00 Ü < o % : y X LO h- Qj d> XZ cr CO u z o H Ü < LU X o X -O CO LU _J JO < u > h- LU -LD < CD N N I < cr o < 3 T T T CNI o 00 THERMISTOR RESISTANCE ( Q ) CD 00 00 (\i /4-- OI ooi LO \hf X O C\i X in LO GDI CNjj CD I o oo LO CNJ (O C\J '''HERMISTOR r esistan ce ( Q ) 4-^ m s o u s s i m Krithalpy of Formation of Alumlnlua Halide@ from the Imtrodmotlom# m âîl^(Z) - + 3àI^i't»X(&q) -3Mi| PaCl(at) ♦ éh| AlCljCc) This Igaores an/ contrlhutlona from (i) Enthalpies of âllntitm arising from slig h t difference® in conceatration batmen tm final products of reaction® (l) and (2). the random spread of experimental results with regard to the weight of aluminium halide used shows these to be negligible# (ii) Sathalpiee of mixing of ♦ 8aOiI(aq) and MGl(aq) 4" ÈaOE(aq). these too might be expected to be rery small# and would tend to cancel each other. The a n c illa r / data used in the calculât ions were from reference 8l unless otherwise stated; HaCl.TSSlHgO « -97.265 kcal mol"^ IJaBr.SOTSHgO « -86.140 keaX mol"? Bal.sasaagO - -70.610 seal mol '1 &b| AlClj(e) « -168.80 î 0.20 keal ®ol"^ (from reference 94)# From which AXBr^(c) *» * 1 1 8 , 4 * 0.6 keal L-495.2 * 2.4 kJ mol"?] Clij 4 1 1 3 ( 0 ) » -67.1 - 0 . 6 kcal mol“^ ■■at- tumnrncmn w n . ■nnKMff Mimw L-280.37 - 2.4 kJ æol~’3 Enthalpy of formation of Aguecue Sodium Alumlnate From tha equation mZ(2) AlCl3(o)+[nSaOH+3HaX3(a'i) — " LKaAl(0H)j^+3BaCl +3K&X ♦ (n-4)gaOa](a%) the standard enthalpy of formation of aqueous sodium alamlnate can be obtained &a ma&l(CB)^(a%) - &Kg(2) + êl^AlCljCc) - 3AH^mCl(@q) + 44«pJaOH(aq) Using the values given above la conjunction with SaGH.SS^IigO « -112.12 kcal mol“' (reference 81) leads to a value of EaAl(Oiï)^. 1 5 0 0 0 HgO = -409.9 - 0.4 koal mol” ^ 1-1715.0 - 1.6 kJ mol'B This compares with the value obtained from references Î1 and 81 (o rig in not given) by adding âll| Ma*(aq) v êll|Al(Ou)|^'^{aq) both at unit molality# of AH| KaAl(OH)^.55.6«2® “ -413.5 kcal ®ol*^ 1-1730.0 kJ bo1“’ 3 /46 PAR? FIVE 8GË08 TRlCHüORlDE CO^fl.&X 1 ^ 7 iMTBODUOTIOg Âlthoui^ ea sily prepared and handled# the 1t1 complex between phoaphoryl chloride and boron trichloride has been the source of much controversy in recent years# on two mala fro n ts I (1) The nature and In terp retatio n of I ts lafra*red spectrum (2) The value for the enthalpy of dissociation# poci^.3ci^(c)# poci^(g) + ( l ) ir;rRA*RED SPEOrmi AI?D four possible structures could he conceived for the adduct* (a) Coordination through oxygen to boron# Cl^PO ^ BGlg (h) Coordination through chlorine to boron# Cl^fCl 301% ^11 0 (c) AS the extreme case of (b)# en Ionic structure involving the tetrachloroboi^te ion, IWOI^J^ IBCl^J^* (d) A combination of ionic and dative covalent bonding# Leocig.BCigJfci" Ixamples of both and ( 0 )^2 2 , 1 2 3 ^re described la the literature# while the ion Is also well documented* -*7 4WÆ CcBductimetric ^ and 'Ol^exohaage experiments have shown th a t there Is negligible ionisation for solutions of the complex in boron trichloride, while in phosphoryl chloride solution there is appreciable ionisation* the nature of which is not yet certain# Although ionization to (c) is more likely* (d) cannot he completely ruled out^^^# Certain predictions can be about the s h ifts expected in the in fra-red spectrum of (a)* (h ), (c) and (d) relative to the frequencies in pure boron trichloride and pure phosphoryl chloride. In (a), (b) and (o)* but not (d ), 2 3 the boron goes from sp to sp on complexing* which has the consequence of lowering the 3*01 stretching frequencies {993 and 9§6 in ^^3 and boron triclilorid® respectively for the asym étrie stretching mode). In ease (e)* the characteristic 301.^ frequenciesshould be observed, With coordination through oxygen, as in (a) and possibly (d), a considerable lowering of the stretching •#1 mode is expected (1290 om In phoaphopyl chloride), accompanied by a raising of the F*C1 stretching modes 1393 (gas) or $$1 (liquid) for the asymmetric mode, 490 cm*^ (gas) or 4Û5 (liquid) for the symmetric mode, in phosphoryl ch lo rid e), fo r (b) and (©) the opposite should occur, with a very slight increase in the stretching mode, and a marked decrease in the F-Cl frequencies, sim ilar to the situation occurring in the complexes of nitrosyl and acetyl halides’ 12S ♦ With (a ) and (b ), vibrations associated with the newly formed I 4-^9 K)-B or F01-B bond# are expected* However# these should be only weak i f the pattern of other boron-K>xygea complexes is fallowed* Three aain groups of workers have published the In fra red spectrum of fOGl^.BOl^# waldington e t.a l.^ ^ ^ used a Bujol mall between sodium chloride or potassium bromide p lates. They found a strong band at 1290 esT^ which they assigned to an unchanged l’«0 stretching frequency, Respite this, they postulated structure (a) on the basis of a weak band at 1190 cm # which they assigned to a B-0 stre tc h in g mode# and a large decrease in B-01 stretching frequencies {700 and €67 caT^ fo r and ^^3), In a second psper^^^# an additional weak band a t 1150 cm**^ was assigned to an stretching mode# and the 1190 cmT^ band to The bands assigned to F-Cl stretching modes showed only very s lig h t frequency shifts, Gerrard et.al.^^^*^^^ also concluded from a nujol mull {potassium chloride plates) that the P*0 stretching frequency was unchanged on complexing. This# together with the presence of what was termed a **3C1. envelope* in the TOO cm ’ region led them to the conclusion that the Ionic structure (e) was correct* They also claimed t!iat a slig h t Increase la the f-01 stretching frequencies (583 and 490 cmT^ ) was compatible with a change la hybridisation from tetrahedral fOCl^ to planar fOGlg*# Bo trace was found of the 1190 or 1150 bands assigned by Waddlngton to B-0 stretching modes, Ooubeaa et.al,127 obtained the low-temporature Infra-red spectrum and the room temperature Raman spectrum @3 part of a systematic Investigation of complexes (X s halogen, 11 a B, â l, Oa or In ), At low tem peratures, mo 1290 band was found, but a strong band at 11 ST was attributed to a stretching mode of 1%0 in the adduct, Strong bands at T13 —1 and 382 cm ' wore assigned to B-0 stretching and 3-0-P bending modes respectively. The asymmetric B-Cl stretch in g modes were lowered to the region 727 - 795 cm~% while the asymmetric and symmetric f-C l stretching modes showed a marked increase to 642 cmT^ and 505 resp ectiv ely . The room temperature Raman spectrum gave eubotantlally analogous re s u lts , Goubeau concluded on tae b asis of these assignments that structure (a). was correct* Ha also stated, giving no details, that the conflicting results of Waddington and Garrard were due to dissociation, hydrolysis and reaction with the cell windows. In addition to the three main works mentioned, SenGupta**^^ obtained the infra-red spectrum of the edduct la carbon disulphide and cyclohexane solutions, when the spectra compared in th e ir salient features with those of faldington, but not with those of Boubeau i ,e . B»0stretching mode a t 1300 cmT\ 3-0 band a t 1200 cm*^. The object of the present work was to try to resolve the r discrepancies between the various resu lts, both by repeating some of the published work and by obtaining spectra using different experimental techniques* The conditions employed were«- (1) Qm phase infra-red, ( i i ) low tempe%%turt)ÿ ^olid phase infra-red# ( i i i ) Room temperature so lid phase infra-red, obtained by subliming the adduct on to the windows of a gas cell, (iv) Hoorn temperature nujol m ill, (v) loom temperature so lid phase Ha^an spectrum* Chemicals Commercial boron trichloride Ltdf minimum stated p u rity and phosphoryl chloride ltd ; minimum stated purity 98#) were both further pirified by fractional vacuum distillation through -45®C (ehlorcbenaene slush bath) to -94*0 (toluene slush b ath ). The p u rity of samples of each were tested by chloride analysis (Volhard*s method) following alkaline hydrolysis. Results were POCl. aci* Cl (found) 68,9# by mass Cl (found) 90,8# by mass Cl (cale) 69,3# by mass Cl ( c a lc ) ’90,8# by mass « fim 1*1 sdduot was prepared by condensing m. excess of boron trichloride on to phosphoryl chloride# allowing the two to mix a t room temperature# thenpumping o ff eaoeee boron trichloride at -45*0# A white solid rezmined, A sample of the adduct was analysed for chloride*- Cl (found) 78.7# by mass Cl (calc) 78.6# by mass. Apparatus mid Procedure Gas phase spectra ware obtained using the cell described previously (p. )# fitted with potassium bromide windows. 1*1 mixtures of gaseous phosphoryl chloride and boron tri chloride were obtained by u tiliz in g the vapour above the so lid complex. Using the grease-free vacuum line described earlier (p . ^3 $ diagram 10)# a sample of the solid complex was placed in the mixing vessel# the line evacuated, and then isolated from the pump* with the remainder of the line, apart from the gas call, saturated with vapour from the complex# the tap on the mixing vessel was shut to iso la te the so lid . Vapour was then allowed to f i l l the gas ceil. This procedure ensured that the vapour in the cell was below saturation pressure, to avoid any sudden, alight, lo c a l cooling leading to condensation on the c e ll windows, Hoorn temperature spectra of the sublimed solid complex were obtained by making intentional use of the condensation 5*3 problem mentioned above. The gas cell was opened to vapour from the so lid complex, which was maintalaed by hand-heat at a slightly higher temperature. #lt& the cell saturated with vapour, the tap was shut, and the cell windows subjected to a flow of slightly-cooled air, when a thin film of solid condensed on the plates. With practice, the thickness of the iilm oGuld be controlled by the degree of heating or oooling applied. A fter the spectrum had Wen run, the c e ll was re-connected to the vacuum lin e , a ll v o la tile products pumped o ff, and a further spectrum obtained. Eoota temperature nujol mull spectra wore obtained by grinding a sazmple of the adduct with a l i t t l e nujol (dried over sodium wire) in a dry-bcx, and then spreading a thin, film of the mull between potassium bromide plates. The spectrum was run as quickly ms possible after making the mull ( eg two minutes) and again periodically a fter leaving the p lates ♦ mull open to the atmosphere. • The low temperature (ca -100*0) spectra of the solid were obtained using the apparatus described previously (p. g'6 ), by spraying vapour from the complex on to a liq u id nitrogen- cooled caesium iodide p la te . All the infra-red spsotra were run on a ferkin Elmer 337 machine. The Raman spectrum of a sample of the solid, contained in a ca p illa ry tuba, was obtained on a Carey 81 instrument (6326a exciting line). ARP Tabla 27 compares the bmade found in the various infra- red spectra, including for reference some literature sp ectra. In diagrama 23-35, the sp ectra are reproduced, in the region -1 1300-400 cm , for the following systernsi- 28 Phosphoryl chloride gas. 29 Boron tric h lo rid e gas. 30 1*1 gaseous mixture. 31 Solid adduct at low temperatures {ca -100*0). 32 Sublimed solid adduct at room temperature. 33 The c e ll windows a f te r the so lid in 32 had been pumped o ff. 34 Bujol mull# 35 The nujol mull in 34 after standing for 45 minutes. The spectra of gaseous phosphoryl chloride and gaseous boron trichloride agreed very well with published data (references 139 and 140 respectively). The spectrum of the gaseous mixture showed no deviations from a simple super- Imposition of the Individual spectra, with no evidence for association. In contrast, the spectrum of a 1*1 solid mixture at -100*0 showed almost total association. The T*0 stretching mode of free phosphoryl chloride (1320 cmT^ In the gas, 1290 omT* in the liq u id ) was absent or very weak. lA weak band was present at 1280 ca*% but since this coincides with an overtone (2 x 633 cmT^), i t is impossible to say whether O -O LO en O m -O u o ÛL ü- 0 o z D cr h - LiJ u O LU -o û_ e n LO Q > LU CC 1 < cr Li_ o o 00 OvJ < cr o < Q o o LO TRANSMITTANCE o o - o LO CJ> u CÛ O o Ll O Z) cr f— u CJ ÜJ CL LO ÜJ o Q o LU e n CL I < cr Ll Z OD CNJ o o < cr o < Q o o o LO TRANSMITTANCE °/o Ç1 o • o LO CJ> en O CÛ O O CL Li_ O o 2 3 cr o I— ■ o u e n lü C l lO LU Q ÜJ q : I < cr Ll o ZI o oî co 21 < cr OI < Q o .o ro O o LO o TRANSMITTANCE o - o LO 2 D cr K u UJ CL lO o - o Q ÜJ cr I < cr Ll CJ z LU cr LU D O I— ■ O < CD cr LU CL LU 2 > LU I— o O . o ro I < cr o < 3 o o o LD TRANSMITTANCE O -O LT) O -O Q Li_ O O LO z Q D LiJ CJ cr t- cr u CÛ ÜJ ZD O UJ CL LO ■ O in CD 3 z Q LU LlI O 5 cr CÛ I < m cr O Ll Z o CL O O C\J ro Z < cr o < Q O .O co O O O LO TRANSMITTANCE °/o o ‘O LO LU O O Z K. z =) ÛL (T LU CJ h - Ll_ < (N o -o :E < cr > o < Q cr o ü_ o •R o O o LO TRANSMITTANCE % o ' o LO _J _J ID _ i ro o —) O 3 - o z Nu CJ o - o CD > o o o o ro O O o LO TRANSMITTANCE I 6 ( ? c O •O LO «n LO ÜJ O -o LO cr e o z LU Q O CO Z O 2^ < en h - LO cr ÜJ H- ü_ < o .o o .o CO o o o LO TRANSMITTANCE 7< T&9&3 27. IKVRA-MED SFËeîP.A ao ij » S l j + fOOl^.aOl^ souamu, S30- SU3— ■ TEMP 10# TEMP U ÎS3 BUJOl (g) (g) acxj ERATURE ERATURE s o ils $0113 BULL (s) (a) (8) AFTER ruapiso 2636 w 2636* 2635V* 2240 vw 1930W 1990* 11990* 1920m' 1920 b 1920 V* 1580 V# ! ■ ■ . . 1468#;1465 w 1465 vw 1430® 1422 m 1422 w 1385® 1390m i1375 V# 1 I I320vs 1320 0 1 3 2 0 a à 1290 s 1305® ; i il2S0 vw 1 Î 1260 eh 1260® i 1205 a i llBOvw) 1180 m IlSOm 1 1 1180* j 1l60vw) j }l130e 1167s 1140m i 1140 w ! 1073 mI 1070 m i 1070 vw 1 i I 1 1 Î067* I ! 1020 V* i 1 5’! 1 995V8 995 8 (995 8 980* i j i i 1 1 193578 955 8 ? ! 955 a i 940 0 1 920# i 1i I 870 vw ! i 1 870 vw 1 i Î |815* i ! I i 800 ©h j?95 eh I i 1 {730w i i i ! } i OP THS gO al,.3C a. COMH.e s . ( e a ~ ^ ) E«JOi sA3i3iFiami SEÎSAaS STUTX (îiüjroi (îrjjoi ASSiaBBgBT OS K Ü i)* MUtl)* sTAsm m 3200 B 0-E str* (byiroXyâis) 2 ^2 (2 *0 8 tr .) 2(P*0 str*) adduct 2 (aa/* gtr*) 2 Vj (aay, atr*) V^BCl 2(a-Cl aay. a tr*) adduét V, ♦ V3 ’^aci^Cg) , ' ♦ ^ 3 11 ’ '3 C l,(g ) 3 ^’ bcI j Cs ) VgCfaO etr.) fcci^(g) 1305 a 1290 s u VgCrso str.) roci.(i) 2 (f-01 asy. str.) aâduct 3 0 1 4 “ ^3 *■ % L00l3 1190a 1190 SI B-0 ®tr* CfeydroX/ala) 2^i^(P-Cl a a j, B tr. ) POOl^(g) 1140 Yw 1150 w PaO str* adduet + \ POCl^(g) 7 2(P-C1 ey* str#) adduot 980 w 1010 m v^(âsy# 6tr.)^^aClj(g) ùr (1) 940 m 975 m Vj(aay# str#)^^BCI^(g) or (1) 936 e \ + *5 ro c i^ ta ) 8 8 0 w 7 ♦ «5 KîCljCg) adduct \ ♦ 1% ÏO Cljtg) TABLE 27 POC1 5 3Clj POCI3 + POCL3 .BCI3 G0U3EAU, SUB SUB (s) (g) 3 CI3 LOW TEMP LOW TEMP- LIMED LIMED NUJOL (g) ERATURE ERATÜRE SOLID SOLID MULL* (s) (s) AFTER PUMPING 765 sh 765 sh 7 ) 750 sh 700 ) 700 1 725 s 727 s s(br)) s(br)| 7 1 3 sh 7 2 8 w 725vw 7 2 0 w 700 1 7 0 0 m 700s (br) ; s(br)) 670 m 670 a 670 w 635 8 620 w 595 vs 595 s 5 9 0 m 595 w 5 9 0 8 585 s 5 0 5 8 5 0 5 8 5 0 5 8h 490 8 4 8 5 m ) 4 9 0 m 480m) 475 s j 4 8 5 m ) 475 m 480ml 455 8 |455 m 455 m s a Strong intensity V » v e r y m » medium intensity br « broad w * w eak i n t e n s i t y sh shoulder CC. (coaTinuEDj aUJGL 3&Bal53T0a asmiAia 'M3M, (saj-oL {sajoL A3%I3#33wT i cs ra«D * ■ w 'i.)* j i STAKDIKO" 1 7 7 3 s (,3-Cl a t r . 1 743 9 j(+ B -0 str* 7) 715 8 [fadduct I ;3 V4 iiaoigtg) 1 f Î ■ 7 I j V, ♦ V4 I i 700 a Î j iBcx- ' .1; 670 w i 670 a 1s(b r) SOI4 - , . ^ i Î 2 roca^ ( 1 V2 1^301 (g) or (1) i ^ y ' m only one band observed aay. *» asymmetric m * bands due to nujol s j . m symmetric omitted* str.» stretching mode t « not assignable \L 1 ttm free iW) atretchlag mode was totally absent ; Its intensity was certainly greatly reduced]. In I ts place, a new, strong, band appeared a t 1130 cm^\ attributable to the stretching frequency of a P»0 group actin g as Lewis base. In addition, the B-CI asymmetric stretching modes (995 and 955 omT^ In free and ^^301^) were sh ifted down to the 800-700 cm"*^ regloa, though the broad, ill-defined, nature of the bands made precise assignments impossible# The P-Cl asynmetrio stretching mode increased from 595 cm~"^ to 635 while the symmetrio stretching mode also Increased, from 490 cm ^ (gas) or 485 (liq u id ) to 505 These results agree well with those of Wartenberg and Goubea#^^?, the only discrepancies being that the completed IW) mode occurred ca 40 cmT^ lower (1130 than found by Goubeau (1167 cra*^^), and there was no sign o f the weak 1067 em“^ band found by Goubeau* la the present work, the bands were not sharp enough to soaign a iM> stretching mode (c.f. Goubeau 713 c!3T^). The shifts observed in the low-temperature spectrum are compatible only with an ozygen-ooordinated species, structure (a). The room, temperature spectra of the adduct, both as sublimed so lid and aujol mull# were v astly more eomplicated than those obtained at low temperatures# The so lid film spectrum In particular contained many more bands, some of which were attributable to dissociated vapour, present in the cell as 6 g a recuit of the technique used* the oomplexed P«0 stretching mode (1140 ) waa s t i l l p resen t, though with a much reduced in te n s ity compared to the low temperature spectrum* In ad d ition , however, strong &ards for free phosphoryl chloride were also observed (1320 gas, m â 1 2 9 0 liq u id ), suggesting a large degree of dissociation* This was confirmed by the presence of both completed and free B-01 stretching frequencies ( ca 750 csaT** and 995/955 respectively)* The complexe# stretch in g modes slowly moved, with time, to lower frequencies ^ 700 On attempting to pump all of the adduct out of the c e ll, a l l bands disappeared except fo r strong peaks a t 700 and 670 cmT\ and a weaker peak a t 1260 ca”L attrlautaala to ths 201^" loa In TMa Clearly demonstrates that reaction has occurred with the potassium bromide windows, a p h e n o m o m u which probably explains the ”301^** envelope" of Gerrard et*&l#^^^*^^^* The 1190 assignment of a "3-0 stretch " by Waddington et*al*^^*^^^also could not be substantiated* Only a very weak band was found in th is area, th is being attributable to an overtone of free phosphoryl chloride (2 v^, F-Cl asymmetric stretching mode, 390 coT^)* P-Cl symmetric stretching frequencies fo r both complexe# and free phosphoryl chloride were present (505 cm“^ and 490 cmT^), but fo r the asymmetric stretching modes, only the fre e band could be detected, (590 ), the complexe# —1 band (ca 640 cm ) being hidden* The only prominent feature 6^ in the spectrum which could not be assigned was a 3-typa band of moderate in ten sity at 830 cm~\ This had not been reported previoualy, and could not be aseigned to any conceivable reaction products, such as free phosphoryl chloride, free boron trichloride# phosphoryl bromide, boron tribromide, tetraohloroborate, tetrabromoborate, phosphorus acids etc* The nujol mull spectrum confirmed the findings of the so lid film spectrum, the main difference lying in the much weaker relative intensities of the adduct bands and of the 8B0 band# In addition, on standing, the vary weak 1180 band increased markedly in Intensity, and shifted «►i to 1190 cm , accompanied by the appearance of new bands a t 3200 cmT^ and 800 ciïT**, in the O-H stretching and bending regions. These observâtions together surest hydrolysis, and the spectrum of boric acid, indeed contains strong bands at 3150 0sT% 1190 cmT^ m â 800 cmT^# The other main band, a t 1428 crn^^, would be obscured by nujol peaks. This leads to the possibility that some, a t le a st, of the intensity of —1 the 1190 cm * band of Waddington e t.a l# may have been due to a hydrolysis product most of the other hydrolysis product bands being obaeured by stronger peaks* Due to fluorescence, the room-temperature IWmn spectrux of the solid adduct could not be obtained above cm 500 cm# # The bands that were observed agreed with the Raman epeotrum of Goubeau^^^, as shown in Table 23* I 7 ^ TABLE 28 a BFLCTau: 0? FOClj, soit scent #Ork Goubeau ^0012(1)127*139 aci^(i)127 Assignment I6la 170s 1933 I78w - 223s 2258 2553 3-Cl i n - p l a n e d e f . 265% 265* 267m 2-Clt deform. 306a 306% 337a 376a 382s 443s 445s 4728 S-Gl ey%* str. 505m 505a 43&S I-Cl sym. str. 1 642s 581* F-Cl asya. str. i 713m B-0 adduct 727@h 958* ^Td-Cl aaym, str. 7508& 7658h 996vw 1®a-Cl aaya. | Str* 7358% 1167* 1300m . . P—0 atr. I t appears that there la far less diaaociation under the conditions required for the room-tcmperature Mrnim spectrum than for the corresponding infra-red spectra. : In conclusion, i t may be sta te d that these resu lts are only compatible w ith an adduct involving coordination through oxygen, but which is readily dissociated at ambient temperatures^ I (2) Dir;$OG%AIlOM OP Values for the enthalpy of dissociâtloa POOl^.BCl^Ca) =^POül^(g) ♦ iJCl^(g) (âlî d lss^ cm be derived from three different sources# (1) Burg and Roea^^^ determined the vapour preasure-temperature curve fo r the complex in the region 10-35^0 from which, assuming complete dissociation in the gas, they obtained log,nlU P " 15.803 ~ 5714 — Since ^ d In Kp as th is gives AH disa «* 26.1 kcal mol*^ {109.2 kJ mol^^ ) From vapour density measurements. It was calculated that the Vapour was at least 90^ dissociated at a temperature (not quoted) higher than fo r the pressure measurements# ( i l 3 Outmann et.al# ^ ^^ measured the enthalpies of solution in phosphoryl ch lo rid e, of boron trichloride and of the adduct K)Cl^#$Clj, whence 301^(1) ♦ foei^(l) -4 rOGl^.BCl^(a) m = -7.9 kcal Eol'l £-33.1 kJ mol'1] Using values from references 71 and Si for the la te n t beats of vaporization of boron trichloride and phosphoryl chloride (5*6 and 9.2 kcal mol^^ resp ectiv ely ), 7 a èildiss» 22.7 kcal mor^ 195.0 W (ill) Finch, Gardner and Sea GuptaT^^* measured the enthalpy of hydrolyeis of the adduct in water FOOl^.mCl^te) + (a*6)HgO(l) m * AMgO](l} 6H hyd* ÈU bÿû m -123.2 koal bo1~1 i-515.6 kj sel*h This leads to âli® fOGlj.3Glj(e) « -271.2 kcal aiol*1 L-1134,7 kJ saol'L anâ, «slag values"^1*®1 of âiij rooij(g) - -134.3 koal raol'l èb| BGlgtg) - -96.7 kcal aol’1 th is gives Audios. 40.2 kcal mol'1 1-163,2 kJ mol'L Flack et.al.coaflraed tkl$ value bj repeating their measurements with a fresh sample of adduct. Burg m â Hass* values for the dissociation pressures were also confirmed, suggesting that the different values for the enthalpy of dissociation could not be attributed to differences between the samples of adduct used* It was postulated that m e explan ation for the discrepancies might he association in the gas phase at 29B1C, contrary to Burg*s assumption i.e. fOeiyBGl^(s) ^?001j*3Ca^(g) **FBCl^(g) e BGl^(g) It was the object of the present work to test this postulate by measuring the association in mixtures of the two 72 Ghealoal@ The same sample of POai^,BCl% waa used as fo r the spectroscopy (p* The water used for calibrating the pressure gauge was distilled from acid permanganate, and degassed by four trap* to-*trap vacuum distillations# The eyolohexaae for testing the pressure gauge was a high pu rity sample (l^hillips Petroleum Co* ltd ; stated minimum purity 99.99#), used directly after drying (grade kk molecular siev es) and degassing. Pure, distilled propionic anhydride (minimum purity 98#) was already available in the Department # Apparatus and Procedure Pressures were measured using a strain gauge (Bell and Howell ltd; type 4*366)* Pressures experienced by a thin, sta in less s te e l, diaphragm were converted to a controlled strain in a Wheatstone bridge whose four arma consisted of strain gauge windings. A rod welded to the centre of the diaphragm transmitted its displacement to a spring element, the movement of which tilte d two of the windings so as to increase their strain, and hence resistance, and the other two windings so as to decrease their s tra in and resistance* The system was arranged such that the output voltage of the out*of* balance Wheatstone bridge was a linear function of the applied 174" pressure* The nominal 40a? fall range output (0- ea §00 torr) was amplified to ^ 10?, giving 0*01§ v o lts per torr* The voltage was registered on a digital voltmeter (krooni ltd; TP 2660) reading correct to * 1m?. The pressure transducer mm incorporated in the epparatua shown in diagram 36. The metal take-off from the transducer sealed into glass tubing of slightly wider diameter using Torrseal resin (?ari@a Ltd.). The apparatus was fitted with a greaseleae tap and joint for connection to the vacuum lin e . The volume of the apparatus up to the seating of the tap was determined aa §61*2 and * 0.1 ml by filling with water from a b u rette and p ip e tte . In use, the apparatus was immersedè to the level of the top of the transducer, in a water jacket which could be ther&oetatted to - 0*1^0. Readings could be taken in the range 0*60^0. When taking a series of measurements over a wide range of temperatures, care had to be exercised that the tap maintained a good seal a t i t s lower seating. Since the coefficient of thermal expansion is much greater for Teflon than glass, on cooling there was a tendency fo r leaking, while at higher temperatures there was à danger that undue force might crack the tap* The thermal characteristics of the apparatus were d eter mined in the for?B of a voltage-temperature graph, a f te r evacuating the vessel to 10*"^ torr (i.e. effectively sero DIAGRAM 36 PRESSURE-MEASURING APPARATUS To Vacuum Line Greaseless V / Tap Level Of Immersion Transd ucer DIAGRAM 37 CALIBRATION OF TRANSDUCER AT ZERO PRESSURE ~2 A ■ ^ -26- 28- ERATÜRE 77 pressure), Diagram 3T shows th is to approximate to a straight lin e la the range with a slope of 0*00307 per degree# A mishap to the apparatus at a later stage aeoessl* tate# repeating this determination# when readings taken wer a wider temperature range (0-60*0) showed the graph to be an extremely gentle curve, approximating to a stmight lin e fo r smaller temperature ranges. These graphs were used to com pensate all measurements fo r thermal dependence of the voltage by arbitrarily selecting the 2$#0*G valueas a standard, and correcting voltage readings a t other temperatures to 25*0# The corrected voltage output from the transducer was calibrated in terms of pressure by using water as a standard^* The corrected voltage-pressure calibration graph is shown in diagrsKi 38 fo r the region 0-100 torr# This line cam be f itte d to the equation P - i70.373V ♦ 21.733 - 0.07 where P « Pressure in toir - 0*0? torr# ? « Corrected voltage in volte# To test the accuracy of the system, the vapour pressare- tamperature curve fo r oyolohexan* was determined in the range 25-40*0 (100-200 torr)# Cyclohexane has often been suggested as a standard for pressure measurements ’^# and its vapour pressure has been carefully studied*^?# The experimental points are plotted on the same graph as the literature curve in diagram 39# 7g -O en o cr cr o LiJ LO cr Z) ui LD LU cr CL O ro O O oo CD O Ô CD O O Ô CORRECTED VOLTAGE (VOLTS) DIAGRAM 39 VAPOUR PRESSURE OF CYCLOHEXANE 180 - 160- 120- Literature Curve Experi mental Values 100 - TE? RATURE go However# many of the l-OGl^/BOl^ pressures to be studied lay in a much lower range (0-20 torr) than the vapour pressures of cyclohexane# So, as a further check on the sem itlvity of the apparatus at low pressures (1-6 torr), the vapour pressure* temperature curve «as determined for a sample of pure propionic anhydride, already available la the laboratory* The results are plotted against the literature curvein diagram 40# The literature values are from very old papers, so it is possible that their accuracy may be questioned to account for the dlQcrepanoy from the present values* Two experiments were performed for phosphoryl chloride-» boron trichloride gaseous mixtures, one In the range 25*55^0, the second from 10-40^0* Below tO^O# maximum permissible pressures 3 torr) were so small that changes in pressure (ca 0.01 torr per degree or lower) could not be registered accurately* The vapour from the solid adduct was again used as the source of a 111 gaseous mixture of boron trichloride and phosphoryl chloride# The pres sure-measuring apparatus was attached to the grease-free vacuum line (p*f3 diagram 10)* The mixing vessel was filled with POvl%#BGl%, evacuated, weighed, and re-attacned to the vacuum line* With the line evacuated and isolated from the pump, the adduot was opened to the pressure apparatus while being thermostatted slightly below the lowest temperature at which readings were to be taken (l*e* ca 20^0 for the first experiment, ca 6^0 for the 1^1 \ O ' u _ LU o Q cr LU Q cr > - ID X CJ LO z LO < LU cr LU CL u q : ID z h- cr o < ID CL cr o LU CL o o CL < cr CO > CL LU o < (T O < Q o —1— — I— — I— C\J LO CO VAPOUR PRESSURE (TORR) eecoad)* fills was Beowaazy to achieve as high a gas pressure as possible while avoiding the problem of possible condensation during the experiment* After allowing a few minutes for equilibration» the tap on the pressure apparatus was shut#' and the remaining vapour condensed back into the mixing v essel, which could then be removed from the vacuum lin e and rewelghed to determine the welglit of vapour in the pressure vessel* Aa the re s u lt of aa accidental pressure shock between experiments one and two, the characteristics of the transducer were altered, necessitating recallbratloa# The sensitivity remained almost unchanged, but all voltage readings were decreased by a near-constant amount* hence the calibration graphs In diagrams 37 and 38 apply only to the f i r s t experiment# The calibrations for the second experiment gave lines nearly parallel to the originals, but displaced by ca -0#149V# Am The resu lts of the two experiments are given In table 29# 93 TABT,E 29 pamismm v @ASEoo$ Ex?mma$T oif^ # o.o%o g I'ooi^.aoi^ used* Voltage ( V) OompiBm at e d Volt age ( V ) Pressure (torr^ f empsjmture *o#oos -0.005 21.43 24.6 0.020 0.003 21.99 30.6 0.033 0.009 22.41 34.8 0.050 0.015 22.84 40.0 0.080 0.021 23.26 44.7 0.102 0.027 23.68 50.2 0.113 0.031 23.96 54.4 0.102 0.027 23.63 50.0 0.035 0.022 23.33 46.0 0.060 0.015 22.84 39.3 0.044 0,011 22.54 36.0 -0.003 -0.003 21.57 25.1 ^sHises^'Î T # 1 0.0355g 2001 j.m31. used. —0.584 -0,342 3.13 10.1 -0.375 -0.343 8.06 13.3 -0.364 -0.341 8.20 16.7 -0.353 -0.339 8.34 20.2 -0.335 -0.337 8.48 25.6 -0.317 -0,334 8,65 30.7 —0.302 -0.333 3.76 35.4 -0.235 -0.331 8.30 40.4 Assuming an equilibrium A(g) ^ B(g) ♦ G(g) then p observed «pA-#- p34- pC 94 and p calculated a p A + § ( p 3 + pO) where p observed a measured praosura p calculated » preaauro calculated from the simple gaa equation assuming■total as#0elation p 1,S,0 as partial pressure A$9*G. TIils leads to Degree of, « p ebs * p calc Slssociatioa ^ 2 Equilibrium, Km (p 003 ** p calc) constant 2 p calc - p ©33 The Values for the degree of dissociation, thus calculated are given in table 30# These values for the degree of dissociation neglect any deviations from Eaoult*© law due to fan der Waal©* repulsive fo rces. However, has shown th a t such int& raeticns can be neglected at pressures less than ea 20 torr# Calculated values for the equilibrium constant, K ^ , are not given because, aa (2 p calc p obs) is so small in every case (often less than the limits of error attached to p obs), the wide limite of error which must be assigned to % p render it almost meaningleea. These limits .range from - upwards, corresponding to limits of - cr more on In Kp ; Hence a graph of IhK. against the reciprocal of the temperature to determine âil is not sensible# ) s r t \ n i s 50 jisaass CF pissociATioa vooiyB ORB JUSTIFICATION 0^ Tim PSOPORTIOI^Atl^Y OF TH^ DISPI.40æ%E?4T PROM ?8&0 0# CHART RTCOÜDSR 43AI%ST GaA%33 IK RSniSTAUem OF TU2RMI%T0a ¥ v o lts Q CR -i- 0ATTERY (VOLTASS ¥ ) CR # Chart recorder I » Matched pair o f resistances 1#) 4,3 kfl Ry « Variable resistance ( = f m Thermistor resistan ce Displacement from zero & Applied voltage (Matois specifications) Applied voltage « ? -• ¥^ ¥ 4 - ^ 0 &R (Bg 2 T at balance) -f &B V _ ? Applied voltage ■> «0 «0 ♦ /IB + X 8g ♦ X 5 «aV "" f\'o ♦ t (Sg ♦ A3 ♦ X)^ (Sg + AB + %) (Sg + aa + x)^ ♦ 63 •*• x)^ m VI (fig ♦ ÉS ♦ 2)' vx a (8„a + x)^ ♦ asaCfig ♦ Since # » ca 5 Q « OB 4900 O X « 4.3 kO d (Applied voltage) jrnm m m Am t #MWTkiu, W m ww^wwWWMmWmw miiPH i a constant to a very good dm approx imat Ion# C < o.i VO two OS' LINEAR BX?R«K;i.a?10S Cl* ltS1IS?A 50S-TI^3 ?OR!î~ ATO ArgSg-ggHlGSS dT — ** ♦ w liewton^s law of cooling d t 4 where t » Temperature (E) t # Time (minutes) Tj » Jackettemperature (K.) K# w m constants* ♦% In IE(T*T :) R t + e by integration # K where o » integration constant. But f w B Thermistor relationship In % where A# fl » Thermistor constants H # Them is tor resistance . . i. - - m a m Differentiating# m ( - ) Ê LlA @K(t+c) d t a (Because of the modulus la la I I # the sign of ^V dt must he determined separately)# Differentiating again# A . 4 U ln B 2 (in ] 0 dt^ Jt ^ 9Ki/s$m0Marr>t>nn* _ m (la 3 For a straight fo re - mià after-period# » e o n s t a n t dt Clearly# this Is not rigorously true. However# typical values are a * 3 % lO^K H j « 3000 i l A m 0 .0 3 i l CJ m 0.003 K mla"'* Ki » —0*01 min.-1 a » 4950 a Bow# consider the after-period# % n CJ la m la where R « Resistance at start of after-period Rj w Thermistor resistance corresponding to jacket temperature c ♦ 380 min fo r a f te r period A M- (at t . 0 ) « -3 X 10^ (11.5)2 (-0. 003) dt T T lO ^ a 0 .7 il afm (at t m 0) # 5 % 10* 117500 ^ 50403 (5 % 10"^) dir 9 % 10^ - 10*2 % 5 a tgS (11.5)2 (3 g, io~3) 5 % 10^ # -0.006 il Bcmally# extrapolation is from t 5» 0 min to t ^ -2 mln fo r the after-period# during which time the slope of the realetanoe-time plot would change from# say# 0.700 s i min"^ to 0.712 S i mln~^ giving a maximum error la the extrapolated resistance of - 0.024 s i # Since resistances can only be read accurately to - 0 .1 n # the possible error will be in sign ifican t. ■Although the after-period only has been considered# sim ilar results (with a smaller ^ ^dt^ in fa c t) can be obtained for the fore-period* In conclusion# Itcan be stated that the extrapolation of fo re - and after-periods as straight lines is justified provided ( 1) The lin es are drawn using as short a section of the fo re - and after-periods as practicable (ignoring the "overshoot" which occurs at the s ta r t of the after-period) ( 11) the extrapolation periods are less than five minutes (ill) The apparatus design should be such as to minimise the values for ^Vdt (1* 8 * ^ and w both very small)# 1 a. ACTgiaix l’irnss Of Tag StgAJOtAflOB 0? FORS- M3 AFïEa-pün-iioa?; To gns Miapsiat oi? nm HSAfias m sim m TIME for the beatias period ^ ~ » K (M J ♦ w Newton*e law a t t ^ ® (T-fj) at ♦ 6J ( u ^ ^ ^ d- k«.^tr jr^ .D A '*3 where (î - 2.) at ^ M@on temperature 6iff#rene@# VJ (f * T.) over the MmtWg period# Kow, to ta l correct@4* temperature change on heating, â f la given by / Aî « Tg - 2^ * dT é t (*3 - t / where /dT \ Mean rate of temperature w change during th is period. o t m - T. 4^ % (f - 2j) at + ^ (tg - t^) iî â. t* - 2g - 2^-WC(& ♦ 3 4 S ♦ £)♦ w ( t ^ - If tjg la a time such that A « C, then âf «S f^ ^k(3 * 0 4" I) -*AD -* ^ (t^ t^) - 2 3 - 2 ,,^K (2g- îp ctg -tp * K(2^ - 2j}(t% - -r u (tg - t^) " ” 3 - ^A * 2j)(ta - tjj) 4 o (tg - t%)] A - (tj - t&)] 62 - 2.3 - 2 , * ( 2 ,. - 2 .,) 4 ( 2 » - 2 ,) as Burning the elope of the after^perlod is constant and equal to K - 2j) * w, and similarly for the fore-period # « * The resletanoea and corresponding to T^ and T^ are obtained by extrapolating the fore- and after-periods of the resistance-time graph to the time t^ i.e# To the mid-point of the heating period. k recent paper by draws similar conclusions to Appendices Two and Three on the use of thermistor## but from a different approach. 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JOB (i960) 1772. 142. E.H. Servose and H*B. Clark. J. Chem. Phys. (1957) 2ê 1175, 143, B.C. Brown, M.D. Taylor and M. Gersteia. J aCS (1944) 66 431. 144 H.C. arowi and M. Geretein. JAC3 (1950) 72 2923. 145. B.H. Gisnn. J » Chem. Therm, (1971) ^ 19, J. Chem. Thermodynamics 1970, 2, 697-700 The enthalpy of hydrolysis and thermodynamic properties of carbonyl bromide M. E. ANTHONEY, ARTHUR FINCH, and P. J. GARDNER Moore Laboratory, Royal Holloway College {University o f London), Engle field Green, Surrey, U.K. {Received 7 March 1970) By use of an adiabatic solution-réaction calorimeter, the standard enthalpy of formation of carbonyl bromide, A/7f(COBr2,1, 298 K), was determined as — (34.7 ± 0.2) kcal m ol"\ 1. Introduction Evidence for the existence and nature of weak interactions between carbonyl halides and Lewis acids, e.g. aluminium halides, is difficult to interpret. Thus COCI 2 and Al^Clg are known^'^^ to form a solid complex, with no halogen exchange, as shown by radio-chlorine measurements.^^) Vibrational spectroscopic work in this labora tory,^^) however, indicates that halogen exchange between COCI 2 and Al 2 Brg occurs, and further that there is (apparently catalytic) decomposition of C 0 Br2 by aluminium tribromide, following very weak complex formation. Clearly there are considerable differences between C0Br2 and COCI 2 , and this has prompted the present deter mination of the standard enthalpy of formation of carbonyl bromide. Despite its easy synthesis and thorough spectroscopic characterization^^) thermodynamic data are limited to the results of an early^^) investigation of the equilibrium constant of the reaction: C0Br2(g) = CO(g)4-Br2(g), at 346.4 and 454.8 K, whence AH[{C0B t2, g) = —23 kcal moP^.f Vapour pressure studies^®) corrected to 298 K, give AH„ = 7.4 kcal mop)^, and hence^^) A7f^(COBr2,1) = —30.4 kcalmoP^. Limits of error are impossible to assign from the literature data. In the present work, a solution calorimetric method at 25 °C was employed for the measurement of the standard enthalpy of hydrolysis. 2. Experimental MATERIALS C 0 Br2 was prepared from carbon tetrabromide and concentrated sulphuric acid according to established procedures,^^) distilled, allowed to stand over mercury, re-distilled (normal boiling temperature 64.5 °C; literature^®) 64 to 65 °C), and t Throughout this paper cal = 4.184 J, M = mol dm~®, and Torr = (101,325/760) kN 698 • M. E, ANTHONEY, A. FINCH, AND P. J. GARDNER Stored in a grease-free vacuum line at liquid nitrogen temperature. For sampling, fractionation through —30 to — 96°C was employed, the product being slightly discoloured. Samples were periodically checked by use of standard silver nitrate titrimetry on aqueous hydrolysates. It was found necessary to shake the liquid vigorously in stoppered flasks for not less than 40 min to ensure complete reaction. A typical result, using a 0.3339 g sample, was: Br(found), 84.6 per cent by mass; Br(calc.), 85.1 per cent by mass. Starch-iodide tests on the aqueous hydrolysate showed the free bromine content of C0Br2 to be less than 0.5 per cent by mass. CALORIMETER AND PROCEDURE The calorimeter was an all-glass Dewar type, containing 200 cnP of liquid, and similar to that previously reported,^^) but modified to operate adiabatically. The design of the automatic adiabatic control is the same (with minor modifications) as that employed in the Gallenkamp Adiabatic Bomb Calorimeter; full details are available elsewhere,^®) and will be forwarded on application. Before (i) and after (ii) runs using carbonyl bromide, the system was checked using the enthalpy of neutralization of tris(hydroxy- methyl)aminomethane (th a m ) in excess aqueous 0.1 M hydrochloric acid. Results were: (i), —7.11 and (ii), —7.11 kcalm oP ^; literature:^)®) —(7.109 + 0.001) kcal mol“). For this calorimetric system the standard deviation of the mean for this reaction has been shown^^) to be +0.02 kcal mol“ ). The reaction studied was : C0Br2(l)+«H20(l) = C02(g) + (2HBr + (n-l)H20}(l) for which we denote the enthalpy change by AT/hyd- Thus Ai7f°(COBr2,1) = AATf°(C02, g)+ 2A77^(HBr in \{n - 1)H20} - A77^(H20,1) - The water in the calorimeter was pre-saturated with carbon dioxide to ensure that the evolved gas would be in its standard state. Thermal effects arising from the decreased solubility of carbon dioxide with the small temperature increase during reaction and also from the mixing of C02(aq) and HBr(aq) were ignored. The rate of hydrolysis of samples of about 0.3 g in 200 cm^ of water was relatively slow » 10 min) and initial results, using a constant-temperature-environment calori meter, were abandoned in favour of adiabatic measurements. Reproducible calori metric results were obtained only after all contact between carbonyl bromide and grease was eliminated. Hence all-glass ampoules, sealed directly from a grease-free high-vacuum line (about 10“® Torr) were employed to contain the carbonyl bromide, and no grease was used in the calorimetric head. Blank runs showed that the effect of heating the ampoule during sealing, and the enthalpy of ampoule breaking, were negligible. The amount of carbonyl bromide used in each run was measured by post-hydrolysis analysis. Corrections AiT^orr for the carbonyl bromide present in the gas phase in the ampoule were applied. 3. Results The results are given in table 1 from which we obtain AH^ye = —(49.06 + 0.16) kcal mol“), the uncertainty interval being calculated as ôlln{n—\)Y'^. Using THERMODYNAMIC PROPERTIES OF CARBONYL BROMIDE 699 TABLE 1. Enthalpies of reaction, A^obs, and enthalpies of hydrolysis, A/fhya, o f COBra at (25.0 ± 0.1) °C m(COBr2) AJToba AH'2orr AHjiyd Run n g kcal mol"^ kcal mol“^ kcal mol“^ 1 0.3723 49.01 0.05 5587 48.96 2 0.1930 49.06 0.12 10777 48.94 3 0.2242 49.49 0.11 9277 49.38 4 0.2168 48.37 0.09 9594 48.28 5 0.2248 48.56 0.11 9253 48.45 6 0.2553 49.16 0.10 8141 49.06 7 0.3342 49.68 0.06 6224 49.62 8 0.3862 49.61 0.06 4386 49.56 9 0.0845 49.61 0.30 24615 49.31 accepted values,*^’') A i/f(C02, g) = —(94.05 ± 0.01) kcal mol ), A //f (H 2 0 , 1) = -(68.315 + 0.001) kcal m ol"), and AHf°(HBr in i(« -1 )H 2 0 } = -(29.01 + 0.01) kcal mol") (mean value betweenn = 5000 and n = 25000), we obtain AH^(COBr2 , 1) = -(34.70 ± 0.21) kcal m ol"). 4. Discussion Statistical thermodynamic functions, based on the usual harmonic oscillator rigid rotor approximation, have been evaluated^"*) for C0Br2(g). Using these results in conjunction with the enthalpies presented above, the standard Gibbs energy and entropy of formation of COB^ (assumed ideal gas) have been calculated, and are compared with the corresponding functions for the analogous halides in table 2. TABLE 2. Thermodynamic functions for COXz(g), (X = F, Cl, Br), at 298.15 K COF2 COCI2 COBra —AHf/kcal mol"^ 151.7“ 52.3“ 27.3o —A(?[/kcal moI“^ 148.0“ 48.9“ 30.7s -A,Sf7cal K-i mol-i 12.4 11.4 -11.56" (10.79 “ From reference 7. " Using data from references 4 and 7. ® Calculated taking the standard state of bromine as gas. The apparently anomalous value of the standard entropy of formation of COBrj, is due to the condensed standard state of bromine; as shown, this disappears if calculation using the appropriate value for bromine gas is adopted. It is a pleasure to thank Mr. S. 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