crl stallization of other sulfides of low solubility product (OH) - in quite high concentrations may be efficacious in increasing the solubility without causing oxide formation. It is interesting to note that LinaresI6 has observed that in molten NaOH at temperatures betxveen the point and 700" ZnS is converted to ZnO. Acknowledgments.-The authors wish to thank S. Geller and C. G. B. Garrett for discussions concerning the structure of ZnS, S. Geller and M. Reid for X-ray work and J. TT. Nielsen for ZnO secds. -1. ,I. Caporaso performed some of the experimeiitnl ivork. (10) R c'. I i n m r ,
I I I I ~i t < I O ~ I I I I I I I I I I ( ~ ~ I ~ J ~ .
STUDIES OS THE C'HEJIISTRY OF
HALOGESS A S I ) OF POLTHALIDES. SS. FOHSI-ITIOS COXSTAKT OF
lIIOSLOE-I:&I6 (19.58); W. R. Person, R . E. HumFhrey a n d A. T. T'opov, ibid.. 81, 2 i 3 (1959).
to check the formation constant of dioxane-IC1 complex. Since preliminary results disagree with those of Lilich arid Presnikova, a more thorough study of this constant was indicated. Experimental Part The source and purity of iodine monochloride, of carbon tetrachloride. and dioxane as well as cxnerimental details of spectral measurement,s have been des;ribed in previous publications . 6 # $ I n a typical run a series of eight solut.ions was prepared cont,aining small and constant concentrat,ions of iodine monochloride and varying large excesses of dioxane. Concentration of iodinr chloride was determined by an iodometric tit'ration. Dioxane solutions were prepared by the addition of a given amount of dioxanc from a weight buret to n given volume of carbon tetrachloride and diluting to the dcsired conrentration. Solutions used for infrared measur('ments had to be quite concentrated, and bocansc~of thv difficulties involved in morlting with concentrated iodin0 monochloride solut,ions Lrhirh wore described earlirr ,9 t.hrcmrcntrations of thrse solutions irere knoLvn otil>- with a n :muracy of itoyo. absorption spectra of ICl-dioxane niixturcs were determined in the visible and ultraviolet regions; since tht: uncomplexed IC1 showed some absorption in this region, tlie niethod of Ketelaar, et u L . , ~ ~ ~ was O used for computing tho formation constant. The data were treated b y the niethod of least squares and are summarized in Table I . TABLE 1 ABSORBANCE DATAAKD FORMATION COKYTISTOF DIOXAXE-IC1COMPLEX Ci"&
7.848 X 7.848 X 7.848 X 7.848 X 7.848 X 7.848 X 7.848 X 7.848 X ac I e Kf KW7.
C2L. ... 0.051 .I02 . I 3 ,204 ,255 ,357 ,510
355
nip
;ibsorbanrr 315 nip ~ 5 . in* j
:3rij
nip
0.022 0.025 0.048 0.120 ,370 ,500 .til3 ,695 ,460 ,635 ,755 ,865 ,520 .io5 ,Xti5 ,950 ,535 .71% ,!)I0 ,995 ,555 ,770 .!I38 1.025 ,590 ,805 !ti5 1.005 ,620 ,845 1.015 1.110 83.0 114.6 139.1 150.!) 23.86 23.12 23 07 23.03 23.27 i0 , :3-L
For thv infrared studies, solutions of dios:ine and IC1 in CC1, also were used. In one scriw of cjxpc'rinients :t solution of 0.05 31 IC1 \vas stndird with concc3ntr:~tionsof dioxane ranging from 1.0 to 0.1 A!. I n anothvr series, the IC1 concentration was 0.1 ,If, while the dioxanc concentration was varied from 0.1 to 0.0125 M . Thew conccntrations were not the most desirable values, hiit, m - w ~ r e limited by the reactivity of the halogen to fairly dilute solutions. However, tjhe concentration of the complex must be high enough so that its absorbance can he measured. This imposes rather narrov limits on the range of concentr:ttion which can be studied. As seen from Table I , the average valw for tlie form:rtion constant measured in t,he ultraviolet rrgiori is 23.3, n-hich is much lower than the value reported by Lilirli and Presnikova.6 However, a personal romminiic:~tio~ifrom Prof m o r Lilich disclosed that they also oht:iinec for Ki in a series of solutions i n whirh N vc' of dioxane was used, but :tt~t~rihiitedi; to n r i c~spcrirnciit:il error. Thr results of 1,ilic.h m t l Priwiikova, :is w~ll:is t h o w of Iltissel and Hvoslofs point to n distincxt possi1)iIit). t h a t in solutions wit,h only a small excess of tiioxanr-, :L tiiox.2IC'I complex might be forming, to sonic cxtcmt at, It.ci,st,. This \vould derrea.se t>hccoriwntration of frcv iodine nionorliloridc and w-ould account for t h c higher v:rlue of thcl ca1cul:ttttl formation constant. On the other hand, ivith :i,large vxress of dioxane only the 1:l complex has a chance to form. T\-hile it might be possible to prove the existence of tho (10) A. 1. Popov, C. Castellani-Bisi and (1958).
11. Craft. ibid.. 80, 6573
second complex spcct,rophotonict~ric~ally :ml cvon t'o determine it,s formation constant, such measurements would have to be carried out in solutions having an excess of iodine monochloride which has a considerable absorption in the near ultraviolet region and the results would not be too reliable. Also, concentrated IC1 solutions have a t.endcnc,y to halogenate organic compounds and such a side reaction with dioxane would render the measurement,s valueless. Thc procedure used to analyze the infrared data was that described earlicr.6 However, in using the Benesi-Hildebrand--Scott equation, we made two modifications. Instead o f using the initial concentration of dioxane as the independent variable, we used the actual concentration (obtained by assuming a value for K f ) . Also, we used the integrated absorbance, Bnl, instead of the absorbance, A , . The rrsulting value of k'fwas 35 I./mole. However, the cxperirnental scatter of the data, was sizable, and lines rould he drawn t~hrorigh the data giving values of Kf as high as 55 :and as low its 25. It should be noted that the infrarcil mt~muremenis were made on solutions with rclativdy high coric*ent,r:ttionsof ICL. Thus, it is possible that, thr higher valiic of Kf is significant, and duo t,o some contribiit.ion to oiir data from t.he 2: I complex. 1 3 ~riw 1 of elit! corrc'la tions fount1 c::trlierQ~ll in sttidiw of t,he iiifr:trccl iipcitr:~of charge-transfer comp1esc.s i t is of some iritercst. to determine the values of the frequency, halfintc.nsity width, and intenbit,y for thc I-C1 absorption band in thr rompl(,s. These were, respectively, Y = 352 cm.-l, A U ~ =: / ~ 1?1 c ~ i i . - ~and , B = 4000 cm.? millipole-l From this, A k / k = 0.12 and sa = 2 . 6 D / A . (see wnces cited for definitions). This value of sa is quitc: high for t,he corresponding value of A k / k , and t,hc resulting point f:ills quite high on the correlation plot (Fig. 7 of ref. 11) and outside the average deviation of the values shown i n that plot. This may also be due to the presence of a tw~~-to-onc complex.
Acknowledgment.-The authors are grateful to Professor L. Lilich for the discussion of this problem. (11) '7.B. Person, R . E. Erioksonand R. E. Buckles, J . A m . Chena. SOC.83, 29 (1960).
THE EFFECT OF HIGHER FATTY ACIDS O S THE D EC~~RBOSYL;ITIOS OF JLYI,OSIC ,lCID BI 1,cirlis \VATIS ( ' L ~ R R Coriliibu'ioii
110n (ha D e p n r f n e n t
o/ Chemisfry. Sarnl H n r y of the I'lnzias College, Dodoe Czty, Kansas
ReceiLed J a n u a r y 8 , 1900
Kintitic studies have been reported on the decomposition of malonic acid in forty-seven nonaqueous solvents comprising representatives of fourtern homologous groups.' The rate-determining step of the rraction in every case appears to he the form,ition of a transition complex, the nucleophdic c:trbo113d carbon atom of un-ionized malonic acid coordinating n ith an unshared pair of electrons +xian electrophilic atom of the solvent molecule, facilitating cleavage. The delicate affinity of un-ionized malonic acid for urishared electrons, mide manifest by the different rates of evolution of carbon dioxide, makes possible its use as n tool or technique for studying the electron and stclric strncturc.; of all kind5 of polar molcculcs. IntereLtiiig rr~11Jti ha^^ t)wn olitnincd prcviou4y from : study of the reaction i i i sm-era1 nf the lomrr fatty wids and their For the sake of cvmpleteness it was helieved n.orthwhile t o extend thr investigation into some of the higher membc rs of the series. ( I ) L W. Clark, Tirre J O [ J R \ % L , 64, 508 (1900) ( 2 ) L W. C l u h ibid , 64, 4 1 (19bO).
The presciit paper describes the results of kinetic studies carried out in this Laboratory on the decarboxylation of malonic acid in four additional monocarboxylic acids, namely, 2-methylbutanoic acid, n-valeric acid, hexanoic acid and decanoic acid, making nom a total of 51 solventb in which the reaction has been investigated. Experimental Reagents.-(1) Reagent grade mziloiiic nc id, 100 0% :may, was used in this investigation. ( 2 ) Solvmts: The fatty arids used in this research were either rmgent grade or highcst purity chemicals. Each sumple of eac 11 liquid M as fractionally distilled at atmospheric presbure into the reaction flask immediately before the beginning of earh experiment. Apparatus and Technique -Thc dct,iil. of t :q)p:wAis .tnd technique h a w h n desciibcd prc\ ioii4y I n thew csxpwimt.nts a sample of mitlonic atid nc~iglilng0.1857 g. ((*orrespondingto 40 ml. of COzat STP on c*omph%crcwtion) h ( b
calibrated Iiy the I;S. Bureau of St:tndaitlB.
Results Decarbosylatioii experiments were carried out in each solvent a t three or four different temperatures over approximately a 20" temperature range. Two or three experiments were performed a t each temperature in each solvent. First-order kinetics were observed for the first 50-75uo of the reaction, after which the reaction rate conqtant decreased slightly with time, due, undoubtedly, to slower side reactions. About 50 ml. of solvent was the amount generally used in the experiments HOWever, a wide variation in the ratlo of solvt~ntto wlute did not appear t o have any effect upon the rate of reaction. The average values of the appnrcnt first-order rate constants for the reaction in the four acids a t the various temperatures studied. obtnincd from the slopes of the experiment:d logarithmic plots, are listed in Table I. The parametrrs of the Eyriiig equation are shown in Table 11. Corresponding data for the reaction in propionic arid and butyric acid, as well as for the dec~nrboxylntlo~i of molten malonic acid, are included for co1np:trison. Discussion of Results In the c a v of the dccarhoxylation of nialoiiic acid iu the l(~\vcrnioiiocarboxylic acids it has been shown that the transition complex is probably formed by the coordination of the elcetrophilic carbonyl carbon atom of the malonir acid with one of the unshared pairs of electrons on the hydroxyl oxygen atom of the solrent inolccule.* Since the AIZ* of the reaction decreases n. thp effective ncgati\-e chargc on thc nurleophilic atom of thc v ) l \ ~ ~ir~crrnscs~ it thc tlecrcnsr in AII* on going from propionic ncitl to n-l)utyrit acid (line.: 5 and 6 of Tablc 11) is coiisisttwt ~ i t thc h fnct that the ethyl group exerts a greater positiJr. inductive effect than does the methyl group6 The slight (7) I, W Clark zbzd 6 0 , l l i 0 (1056) K J Laidler "Chemical Kinetics lnc., New York, N. Y., 1950, p. 138 (2)
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J J c ( ~ r ~ s - I l dBook 1 Co
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