oxamic acid reaction, and this corrwponds also I o the order of increasing nucleophilicity. This coincidence is good verification for the hypothesis that the rate-determining step of the reaction is the formation of a transition complex between solute and solvent. The AH* values for the oxalic acid reaction likewise decrease progressively on going from aniline t o quinoline and from quinoline to triethyl phosphate-thus far in accord with theoretical predictions assuming a mechanism similar to that of the oxamic acid reaction. I n each of these solvents, furthermore, the AH* is higher for the oxamic acid reaction than it is for that of oxalic acid, a result which is consistent with the fact that the polarized carbonyl carbon atom of oxamic acid has a lower effective positive charge than does that of oxalic acid.2 I n each of these three solvents it will be observed, also, that the AS* values for the oxamic acid reaction are higher than they are for that of oxalic acid in spite of the nearly equal sizes of the two acids. This circumstance has been attributed, previously, t o the greater tendency of the dicarboxylic acid to associat,e through hydrogen-bonding to form a I ( supermolecule” cluster, I n the solvents dimethyl sulfoxide and 8methylquinoline the A H * as well as the AS* 1-aluesfor the oxalic acid reaction deviate abruptly from the regular order observed for the oxamic acid rea,ction (lines 4 and 5 of Table 11). Evidence has been presented previously that in these tn-o solvents primary ionization of the oxalic acid t!akes place, the acid oxalate ion, rather t,han the un -ionized di-acid, then undergoing decarhoxylation.2 The fact that, in these two solvents, the H* values for the acid oxalate ion reaction are slightly higher than they are for that of oxamic acid is ronsistent with the difference in the relative acidities of these two ~ p e c i e s . ~It will be noted, also, that the S* values of the acid oxalate ion reaction in these two solvents are slightly larger than those of oxamic acid-a result which is consistent with the diflerence in the relative sizes of the two species. In quinoline, the enthalpy of activation for the decarboxylation of oxalic acid is about 8 licnl./ mole higher than that of oxamic acid, whereas, in aniline, it is more than 19 kcal./mole higher (see Table 11, lines 1 and 2 ) . I n view of the fact that these two amines do not dif’fer appreciably in basicity ( p K for aniline is 9.42 for quinoline 9.2),6 this abnormally high value of A H * in the case of oxamic acid in aiiiline poses a problem. Inductive and steric effects noted in studies on the decarboxylation of oxamic acid in the two primary amines aniline and o-toluidine, as well as in the tertiary amines quinoline and S-methylq ~ i n o l i n e were , ~ indicative of the formation of a transition complex involving coordination between acid and solvent. Apparently, however, in the case of primary amines, the orientation of solute molecules with respect, to solvent must differ from that in the case of other nucleophilic sol(6) N A . Lange, “Elandbook of Chemistry,” 9 t h ed., IIandbook Publisliers, Inc., Sandusky, Ohio, 1956, p. 1204.
vent,s. This circuinst ance may bc connected with the structural relatioilship of oxamic acid to a-keto acids, for it has been ohserved that the decarboxylation of such acids is specifically catalyzed by primary anlines.’ The fact that oxalic acid, itself a type of a-keto acid, does not show tjhis same behavior may be at’tributed to its t’endency t’o associat8e through hydrogen-bonding to form a “supermolecule” cluster. Further work on this problem is conternplatcd. Acknowledgrnent.-‘l?he support of this research by the National Science Foundation, Washington, D. C., is gratefully acknowledged. (7) .J. Hine, “Physical Organic Chemistry,” MoGraw-Hill Book Co.. Inc., Now York, N. Y., 1956, p. 288.
TIIISRNOICIIE~\1[ISTRYOF ZPIKONPU3l HALIDES 1 3 A. ~ G. TCRNBULL Division o j M i n e r a l Chemistry. Chemical Research Laboratories, C.S.I.R.O. Australia
Received ApriE I , 1961
‘I’hermodynamic data for zirconium compoiititls are accurately known for the oxide and tetrachloridc on1y.132 Development of a consistent set of heat of formation and free energy data would assist in understanding the reactions of ore extraction, metal reduction and Zr-Hf separation, and also promises to reinforce the still sketchy ideas on zirconium solution chemistry. For the measurement, of AHoZga of ZrI3r4 and Zrl4, t’he direct union of cleinent,s in a bomb cxlorimcter is difficult t o init,iate and mmpletc. lIon-c\vrrr values relative .to those of ZrOz and ZrC14 may bc obtained by suitable soiution reactioiis. Thcsc two independent schemcs were used
+ 4NaOIJ,,, + + 4Na(:I,,, -i- 2F120 ( I ) ZrX((c) + 4NaOII,,, + + 4NaX,,. $- 2 H z 0 ( 2 j ZrCld(c) + 2 H z 0 --+ ZrOOII”,,. + 4C1kSG,,+ ;HIf ( 3 ) %rXa(c) -/- 21120 -+%rOOII+,,, -I- 4x-,, + 3T1+ ( 3 )
A. ZrCld(c)
Zr02hyd.
%iro?hyd.
I3.
Such reactions we:it to completion in a few minutes a t 25’, giving reproducible final states, and tlwir validity vas based 011 the recent systematic zirconium chemistry proposed by Blumen thal. Experimental Heats of reaction were measured in a suitably thermostated D e n w vessel containing the g l a ~ ssample bulb in 100 ml.of solution, glass stirrer, glass covered constantan heater and copper-constantan thermocouple. Temperature rise was amplified and recorded on a Leeds and Sorthrup Speedomax recorder at 4 in./hr. The maximum sensitivity 7va.s 0.1 ca,l., while heats of reaction were 100-150 cal. Electrical ca1ibrat)ions before and after each reaction and for different fillings agreed wit’hin 0.570, with power measured t o 0.27,. To verify the absence of systematic errors, the heat of solution of _ _ _ I _ -
(1) F. D. Rossini, el aE., National Bureau of Standards Circular 500, 1952.
(2) P. Xubaohewski and E. Evans, “1R.letallurgical Tliermoriic.riiistry,” 3rd Ed. Pergamon Prees, N e w York, N. Y . , 1958. (3) W. €3. Bluinenthal, “The Chemical Behavior of Zirconiririi.” 1). Van Kostrand Co., Kea. York, N. Y . , 1959.
Sept., 1961
1653
xs’03(c) in 500 moles of water was measured to be 8.55 rf 0.1 kcal./mole in agreement with the best literature value of 8.56 kcal./mole.’ The heat of breaking the evacuated sample bulbs was measured t o be 0.3 cal., representing only 0.2% of the total heat evolution. Samples of ZrBr,, Zr14 and HfC14 were made by halogenation of the pure metals in a silica tube a t 500-700”. The ZrClr was made by chlorination of zirconium carbide made from zircon with a Hf/Zr weight ratio of 0.017. A11 halides were vacuum sublimed into thin Pyrex bulbs and the sample weight found by difference. The chlorides and bromides were pure white, crystalline solids and it has been amply shown that such a preparation leads to the formula MX,.* The iodide was pale orange-brouvn of formula ZrI, within analytical accuracy (calcd. Zr 15.23%, found 15.15%). Spectrographic analysis of the Hf metal showed only 0.06% total impurity (apart from 2.4% Zr) and the “reactor grade” Zr metal had the same order of purity, so that no correction of reaction heats for impurities was warranted.
Results A. Reaction with Na0H.-The solid halides were treated with a small excess of aqueous NaOH, using a halide to water molar ratio of 1:1500. No dependence on sample weight was observed, as expected a t such high dilution. The product was pulpy white hydrated Zr02 in all cases. It was verified by separate measurements that the effect of the small excess of KaOH on AH298 of the NaX product was negligible. Also the solubility of ZrOz in the excess S a O H was extremely small.4 Since the Zr contained 1.7% Hf and the Hf contained 2.4% Zr, it was possible to solve simultaneous equations and find the heats of reaction of the pure tetrachlorides. The correction was negligible for ZrCl4 and only 0.3 kcal./mole for HfCL. The elements are so similar that any heat of mixing should be quite negligible. Values of A l l for reactions 1 and 2 are given in Table I. Beck5 treated ZrCl4 with NH40H(aq),but, his single value of --54.3 kcal./mole apparently is in error. Combination of reaction heats with values of AHoZ98 for NaOH(aq), NaCl(aq), NaBr(aq) and NaI(aq) * and the latest value for ZrCl,(c) of 234.7 rf 0.4 kcal./mole6 led to values of AHa,,, for ZrBr4(c), -182.2 +. 0.7 kcal./mole and ZrIe(c) -115.6 rf 0.8 kcal./mole.
1
2 3 4 5 6 Mole NaOH/mole ZrC14. Fig. 1.-Thermometric titration of 0.01 M ZrCL with 0
4 M XaOH.
and HI(aq),I and ZrC14(c)‘jled to values of AoH298 for ZrBr4(c), -181.0 f 0.5 kcal./mole and ZrL(c), -116.3 0.8 kcal./mole. For the thermometric titration of N/40 ZrCl4 with N/40 ISaOH, Chauvenet‘ found -46 kcal./ mole, which compares well with -43.7 kcal./mole found here from the difference of reactions 1 and 3 for a similar order of concentration. This titration was repeated to provide an independent check on the above results. A ZrCla solution was titrated with 1 cc. of 4 M SaOH using a syringe-type microburet inside the calorimeter. As shown in Fig. 1 the total heat evolved up to the equivalence point was -44.5 kcal./mole, in good agreement with the above value. The initial slope of the titration curve, - 13.6 kcal./mole, corresponded within experimental error to the heat of the reaction
*
4- OH, ,+f X ~ O H----f HzO After two moles of SaOH per mole of ZrCl4 had H’
been added, neutralization of the zirconium cation evidently commenced as the rate of heat evolution fell.
TABLE I
Discussion There appears t o be no significant difference in aq. HtO values of ALNL for ZrBr,(c) and Zr14(c)obtained by Material - 4 HNaOH (koa1 /mole) - 4 H (kcal./mole) the two methods. The final average values are Zr( 1.7% Hf)CL(c) 103.2,102.1,102.8 59.1,59.3,55.5 -181.6 f 0.6kcal./mole and -115.9 f 0.8 kcal./ Z r Cla(c) 102 7 59.0 mole8,respectively,whichare considerably lower than ZrBr,(c) 110.4,110.7,111.3 68.6,69.0,68.6 previous estimated values, - 192 ek 10 kcal./mole 114.6, 115.2, 115.8 70.5, 71.3, 70.9, ZII,(C) and -130 f 5 kcal./mole,2 respectively. En71.1 tropies S Z ~ofS ZrBr4(c) and Zr14(c)have been estiHf(2.4% Zr)CL(c) 108.7 65.7 mated by the method of Latimer, to be 52.1 j= 5 HfCli(c) 109.0 66.0 and 67.5 5 cal./omole, respectively, leading to B. Reaction with Water.-The solid halides values of A P L of -172.5 * 2 and -116.6 +. 2 were treat,ed with mater to give clear, colorless kcal./mole, respectively. For ZrC14(c) the value solutions. As above, the molar ratio was about of A F k is -213.4 0.6 kcal./mole, calculated 1:1500 and no dependence on sample weight was from the best reported values of AH!gg and S 2 9 8 ( 2 ) . observed. The correction for the Hf content of When values of for the tetrahalides of Ti, the ZrClr was again negligible. Values of AH for Zr and T h are plotted against values for the correactions 3 and 4 are given in Table I, and combined responding oxides, the regular shape of the curves uith values of AH0298 for 1520, I-ICl(aq), €IBr(aq) confirms the new values for ZrBr4 and ZrI4 and HEATSOF REACTION
*
*
F. K. XIcTagpart, Rev. Pure A p p l Chem., 1, 152 (1951). ( 5 ) G. Beck, Z. anorg. allgem. Chem., 174,31 (1928). (6) P. Gross, C. 1 L ) r n a n and D. L. Levi, Trans. Faraday Soc., 63, 1285, 1601 (1957). (4)
(7) E. Cbauvenet, Compt. rend., 164, 630 (1917). (8) Standard statesassumed to be Zr(c), Clz(g), Brz(l), Iz(c). (9) W. M. Latimer, “Oxidation Potentlals,” 2nd Ed., Prentice-Hall Inc., New York, N. P., 1952.
enables prediction of -160 =t 5 kcal./mole for AH898 of Thl,(c). The value of Af& for hydrous ZrOz, -260.3 f. 0.6 kcal./mole, found frdm reaction 1, is not significantly different from that reported for monoclinic ZrOz, -261.5 =k 0.2 kcal./mole. This would indicate that the m-ater in ZrOz hyd. must be only loosely held by physical trapping in the micelles of the solid. There is slight X-ray evidence reported of a tetragonal arrangement of ZrOz in the hydrous form on heating to below 6OOol3 but the difference in AH betmen monoclinic and tetragonal is only about 1 kcal./mole.* Latimerg already has cast doubt on the heat of hydration of ZrOz of - 25.3 kcal./mole previously reported.’ From reaction 3 the value of AHh8 for the ZrOOH+(aq) ion is -270.7 =k 0.5 kcal./mole. Acknowledgments.-The author wishes to thank Mr. N. Gye for chemical analyses and Mr. I. E. Newnham for continued support and advice. T H E I S F L U E S C E OF PRESSURE ON THE C I T I O N C POLYMERIZATION OF ISOAJIYL T’IXYL E T H E R BY S.11 H.of.raxh AKD D. R. TEPLITZKY~ Division of Physical Chemistry, A u s t r a l i a n C o m m o n w e a l t h Scientific a n d I n d u s t m a 1 Research Organization, Fishermen‘s B e n d , .Melbourne, Australza Received A p r i l 13, 1961
Although it is now well established that high pressures have a profound influence on t’he rates and products of free-radical polymerizations in the liquid p h a ~ e ,it~appears ,~ t’hat very little work has been done on ionic polymerizations under pressure. Kilroe and TVeale4have examined the cationic polymerizat’ioii of a-methylstyrene catalyzed by trichloroacetic acid, and found it to be accelerated by an increase in pressure in much the same way as the free-radical polymerization of the same monomer. But the kinet’ic form of this reaction is unknown, even at atmospheric pressure,j and it is difficult to draw any useful coiiclusions from the results. Here we report some measurements of the effect of pressure on the rate of polymerization of isoamyl vinyl ether, cat’alyzed by iodine. Eley and Richards’ and Eley and Saunders8 have made a detailed study of this type of polymerization a t atmospheric pressure and concluded that it occurs by a carbonium ion mechanism, initiated by I + ions. This supposition is supported by more re(1) Now a t t h e School of Biological Sciences, University of New South Wales, Sydney, Australia. ( 2 ) S.D. Hamann, “Physico-Chemical Effects of Pressure,” Butterworths, London, 1957, pp. 189-192. (3) M . C:. Gonikberp, “The Rates and Equilibria of Chemical Reactions a t High Pressures,” Acad. Sci. U.S.S.R., Moscow, 1953, pp. 156-167. (4) J. G. Kilroe a n d K. E. Weale, J . C h e m . Soc., 3849 (1960).
( 5 ) T h e kinetics m a y be t h e same as those for t h e cationic polymerization of 1 : 1 diphenylethylene ( L e . , 0-phenylstyrene) under t h e same conditions,%b u t this has not been proved. ( 6 ) A . G. Evans, N. Jones a n d J. H. Thomas, J . C h e m . Soc., 1824 (1955). (7) D. D. Eley a n d 9.W. Richards, T r a n s . F a r a d a y Soc., 45, 425 (1949). (82 D. D . Eley a n d J. Saundera, J. Chem. Sac,, 41G7 (1952).
cent evidenceg that although vinyl ethers are resistant to attack by free radicals, they are readily attacked by the stable tropylium carbonium ion. In the prePentJwork we have examined the polymerization both in diethyl ether solution and in the undiluted monomer, at pressures up to 10,000 atm. Experimental The reactions were carried out in glass hypodermic tubes, sealed with sliding glass plugs, and immersed in petroleum ether in a pressure vessel similar to the one described by David and Hamann.’O The progress of each polymerization was followed by extracting samples of the reaction mixture, quenching the reaction with an aqueous solution of sodium thiosulfate, and estimating the amount of polymer by removing the unchanged monomer and solvent by prolonged evaporation at 75” and 0.04 mm. The average molecular weight of the polymer was measured cryoscopically in benzene solution. The sample of isoamyl vinyl ether was kindly given to us by the Distillers Co., England, and was purified by fractionation under vacuum. The diethyl ether was dried over sodium. The reaction mixtures in diethyl ether were prepared by first dissolving iodine in the solvent and then adding the monomer. In the absence of solvent, the addition of solid iodine to the pure monomer caused a violent polymerization before the iodine could fully dissolve.” For the reactions in the “undiluted” monomer we were therefore forced to add the iodine in the form of a concentrated solution in diethyl ether. It is unlikely that the small amount of diethyl ether introduced in this way (1% of volume) would have affected the kinetics.
Results and Discussion Reactions in Diethyl Ether Solution.-In the ether solutions we confirmed Eley’s7v8 kinetics (where [MI = concentration of monomer, [Iz] = concentration of iodine) at atmospheric pressure, but we found discrepancies at higher pressures. It was obvious that iodine was consumed during the high pressure reactions and in some instances no free iodine was left after 40 minutes (starting from a concentration [IzIt= = 0.001 mole l.-l). It appeared that the iodine had formed a strong complex or an addition compound with the vinyl ether. This could be hydrolyzed, with the liberation of iodine, by adding a little water to the mixture. T o avoid this complication we were compelled to work at pressures below 1500 atm. and to estimatc the rate constants from the initial slopes of the reaction curves over fairly short times. Our results are given in Table I. Although the rate constants may not be very accurate, there is no doubt that the reaction waq accelerated at high pressures. On the other hand there was no significant change in the degree of polymerization between 1 and 1000 atm. This iq not surprising since the degree of polymerization is determined by the ratio of the rate constants (ka and kg in the notation of Eley and Saunderss) for two closely analogous reactions between growing polymer and monomer. It is understandablc (9) D. N. Kursanov, M. E. Volpin a n d I. S. Akhrem, D o k l a d v A k a d .Vauk S S.S R , 120, 531 (1958). (10) H. G. David and S. D. Hamann, TTan8. F a r a d a y SOL.5 0 , 1188 (1954).
(11)
Chalmers, C u n . J. Res., ET,464 (1932).
