Sept., 1961 the electrode reaction cannot be prepared conveniently. This work was supported in part by funds received from the United States Atomic Energy Commission, under Contract S o . ilT (11-1)-64, Project No. 17.
iB5i
XOTGS
40 0 mi. of GO2 at ATP on rompktc rraction) n x intro~ duced in the usual manner into the dried reaction f l d containing B weighed sample of solvent saturated with dry GOz gas.
Results The rate of decarboxylation of oxamic acid was measured in triethyl phosphate and in dimethyl sulfoxide over a temperature range of about 20". A t a fixed temperature no appreciable difference KINETIC STUDIES Or\' THE in the specific reaction velocity constant could be DECARBOXYLATION OF OXh3lIC ACID IN detected when the quantity of solvent was varied from 50 to 140 g. Generally about 130 g. of solDIMETHYL SULFOXIDE ,4SD IN vent was used in each experiment. Duplicate TRIETHYL PHOSPHATE experiments were performed a t each temperature. In every experiment 9501, or more of the theoretiBYLOUISTTATTS CLARK cal volume of C 0 2 was collected. There was Depaiinzent of Chemistry, W e s t e r n Carolina College, Culloiuhee, IVorth Carolina very little tendency for the reverse reaction to Receiued M a r c h 2?Y, 1961 take place in either of these solvents as shown by the fact that, on allowing the system to reniaiii The results of kinet,ic studies on the decarboxylaintact overnight after decarboxylation was comtion of oxalic acid in various polar solvents have plete, a very few milliliters, if any, of the evolved indicated that, in weakly basic solvents, unCOZ was resorbed. This was a marked contrast ionized oxalic acid decomposes by a tra,nsitionconiplex mechanism.1 I n more strongly basic with the behavior of the system using aromatic solvents primary ionization of the oxalic acid must amines as solvents. The amines acted as very occur, the acid oxalat,e anion then undergoing strong catalysts for the reverse reaction.3f4 The plot of log ( V , - V,) us. t was linear over decarboxyiation.2 I n the case of the un-ionized nearly the entire experiment, indicating that the di-acid, a "supermolecule" cluster composed of reaction is first order. The average rate constants several molecules of oxalic acid apparently takes part in the &e-determining st'ep, whereas, in calculated in the usual manner from the slopes of the reaction in ionizing solvents, the acid oxalat'e the logarithmic plots are shown in Table I. The parameters of the Eyring equation, based upon ion appears to be involved singly.4 Kinetic studies have been reported also on the the data in Table I, are shown in Table 11, along decarboxyiation of oxamic acid in several non- with those for the reaction in seyeral other solaqueous ~olvents.~,4Since oxamic acid is a vents previously investigated. Shown also for weaker acid than oxalic acid it does not ionize as comparison arc corresponding data for oxalic acid. readily. Also, having but oiie carboxyl group, it TABLE I apparently does not form an association clust'er, APPAREUTFIRST-ORDER RATE C O N S T A N T S FOR THE ])Esingle molecules only participating in the rate- CARBOXYLATIOU O F OX4fvfIC ACID IU DIMETHYL SULFOXIIIE determining step. The reaction has been found 4 h D I N TRIETHYL PHOSPH4TE to be reversible in the presence of aromatic amines. Temp IC x 10% >ol\ r n t ("C rot ) (we I n order to gain more information 011 this re140 1 4 1 87 f O 01 action, and especially in order to ascertain whether IJimeth) 1 sulfoxide 150 72 6 10 f 01 or not nou-nitrogenous solvents are able to cata160 60 17 12 & 02 lyze t'he reverse reaction, kinetic studies have been 131 23 1 65 f 01 carried out in this Laborat'ory on the decarboxyla- Triethyl phosphate 138 20 1 01 It 02 tion of oxamic acid in the solvents triethyl phos149.61 15 79 It 02 phatc and dimethyl sulfoxide. Results of this investigation are reported herein. TABLE I1 Experimental KINETICI-)ATA FOR THE n E C A R B O X Y L 4 T I O N OH O X - i l l I L -1)
Reagents.-(1) The oxamic acid used in this invc,stig;ltion was analytical reagent, grade, 100.Oyo ~ s s a ~( 2. ) Reagent grade triethyl phosphate was used. Further purification was effecbed by distillation of each sample a t atmoepheric pressure directly into the reaction flask immediately before the beginning of each experiment. (3) Tho dimethyl sulfoxide used in this research was 99.9% j)ure, the remainder being water. Distillation of this solvent was not feasible due to its instability a t the boiling point. Apparatus and Technique.-The kinetic experiments were conducted in a constant-temperature oil-bath by measuring the volume of COZ evolved a t constant pressure, as described in a previous paper.6 In each experiment, a 0.1585g . sample of oxamic acid (the amount required to produce __ (1) L. W.Clark, J . Phya. Chem., 61, ti99 (1957). (2) 1,. W.Clilrk, i b i d . , 62, 633 (1958). (3) L. W.Clark, ibid., 66, 180 (1961). (4) L. W. Clark. ibid., 66, 659 (1961). ( 5 ) L. '8. Clark, ibid., 60, 1150 (1956).
ACID B %l\
~ OXAIIL D ACID IN
rnt
SEIERAI.POLAR SOLIENTS'
--Oxaiiiic
acid-AS*
mole)
( e 11 j inole)
AH* (Lcal 1
(IC\
--Oxah AH* (kcal / mole)
aciil-AS* (E 11 j mole)
Aniline3 1 59 7 1-68 0 40 3 $16 2 Quinoline4 47 0 4-37 5 38 9 $15 8 Triethylphosphate' 40 9 +24 7 28 9 - 5 8 Dimethyl sulfoxide' 37 7 +14 9 40 6 +20 7 8 - M e t h y l q u i n o l i ~ e ~ ~36 ~ 0 4-12 2 37 7 $13 7 a The first superxript after the name of the solvent refers to the source of the oxamic acid data, the second to that of the oxalic acid data. For the solvents triethyl phosphate and dimethyl sulfoxide the single superscript refers to the murre of the oxalic acid data.
Discussion of Results 'l'hc solvents listed in Table I1 are arranged in the order of decreasing activation energy for thc
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-alues for 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’f er 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) -/- 21&0 -+%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.
