NOTEB
VOl. 64
results of the solution measurements (Table 11) indicate moments near 1.2 D. Considerable reFrom this equation, the minimum average elec- striction to rotation of the alkoxy groups is, tric moment, possible, assuming free rotation and therefore, indicated. Acknowledgment.-We are indebted to the Office irrespective of the values of m2 and m3 is 1.86 Debyes. This is evident since a minimum calcu- of Ordnance Research, U. S. Army, for supporting lated p will occur n-hen mz - m3 = +0.38. The this work under Contract D-4-04-200-ORD-037. I/$,
this equation simplifies to ji' = (WL,
- ma)' - 0 . 7 6 ( ~ ~-2 m ~ + ) 3.611
NOTES THE DECL4RBOXYLATIOK OF THE TRICHLOROACETATE ION I N n-BUTYL ALCOHOL, %-HEXYL ALCOHOL AND n-CAPROIC ACID BY LOUISWATTSCLARK* Department of Chemistry, Saint M a r y of the Plains College, Dodge City, Kanaas Receined April 8, 1960
A large variety of organic reactions takes place as the result of an initial interaction between an electrophilic, polarized, carbonyl carbon atom (in aldehydes, ketones, esters, acids, amides and related compounds) and a nucleophilic agent. The extremely general Claisen condensation may be cited as an outstanding example of this princip1e.l It appears highly probable, also, that a great many reactions of biochemical interest, e.g., enzymic and immunization reactions, are of this same general type.2 Two closely related examples of this type of reaction which have been t'he subject of a great many kinet'ic studies over a relatively long period of time are the decarboxylations of malonic acid and the trichloroacetate ion in polar solvents. It has been demonstrated convincingly that the rate-determining st,ep for both reactions is the formation, prior to cleavage, of a transition complex involving coordinat>ion between the elect.rophilic carbonyl carbon atom of un-ionized malonic acid on the one hand, and that of the trichloroacetate ion on Dhe other, with an unshared pair of electrons on the nucleophilic atom of the solvent m ~ l e c u l e . ~The decomposition of malonic acid has been studied, to date, in approximately 53 non-aqueous s o l v e n t ~ ,that ~ of the trichloroacetat'e ion in more than a d ~ i i e n . ~ In those solvents containing mobile electrons in addition to the fixed electron pairs on the nucleophilic atom alii interest'ing difference in electron influence has been observed between these two species. It has been observed that, in the att'empted coordination between malonic acid and a polar molecule, the effective positive charge on the * Western Carolina College, Cullowhee. N . C. (1) R. Q . Brew-ster, "Organic Chemistry," 2nd Ed., Prentice-Hall. Inc., New York, N. I-., 1953,pp. 400 et seq. (2) J. H. Turnhull, Esperimtia, XV/8, 304 (1959). (3) (a) 0.Fraenkol, R. I,. Belford and P. E. Yankwich, J . A m . Chern. Soc., 76, 15 (1854); (h) L. W. Clark, Txrs J O U R K ~ L 68, , 99 (1959). (4) L. F.Clark, ibid., 64, 692 (1960). (5) L. W. Clark, ibid., 64, 917 (1960).
carbonyl carbon atom of the malonic acid tends to attract mobile electrons present in the solvent molecule, resulting in a decrease in AH*. On t h e other hand, the resonating negative ionic charge on the trichloroacetate ion tends to repel mobile electrons present in the solvent molecule, thus decreasing the electron density on the nucleophilic center a t the moment of reaction and resulting in an increase in AH*. The solvent molecule responds to the advances of the malonic acid, but resists the advances of the trichloroacetate ion. This phenomenon is not observed in polar solvents containing no mobile electrons.3b Although this principle has been tested already on a considerable number of compounds of various classes it was deemed worthwhile to carry out additional tests to see whether or not any exception could be found to it. For this purpose kinetic studies have been carried out in this Laboratory on the decarboxylation of the trichloroacetate ion in three additional polar solvents, namely, nbutyl alcohol, n-hexyl alcohol and n-caproic acid. The results of this investigation are reported herein. Experimental Reagents.-( 1) The potassium trichloroacetate used in this research was reagent grade, 100.0% assay. (2) The solvents were reagent grade chemicals which were freshly distilled a t atmospheric pressure directly into the reaction flask immediately before the beginning of each decarboxylation experiment. Apparatus and Technique.-The details of the apparatus and technique have been described previously.6 Temperatures were controlled to within fO.OlO, and were determined by means of a thermometer calibrated by the U. S. Bureau of Standards. The reaction flask was 100ml. capacity. Approximately 60 g . of solvent was used in each experiment, along with 330 mg. of potassium trichloroacetate (the amount required to furnish 40.0 ml. of COz a t STP on complete reaction).
Results In the decarboxylation of the trichloroacetate ion in n-butyl alcohol and in n-hexyl alcohol the log ( V , - V,) was a linear function of time over only approximately the first 75% of the reaction. In any one experiment no more than 85% of the theoretical volume of COz was collected. This probably is attributable to a partial absorption of the evolved COz by the alkoxide ions concomitantly formed. I n the decarboxylation of the trichloroacetate ion in n-caproic acid, on the other hand, (6)
L. W. Clark, zbzd., 60, 1150 (1956).
Nov., 1960
NOTES
1759
The increase in AS* of approximately 10 e.u. regular first-order kinetics were observed over the greater part of the experiment, and each experi- in the case of the trichloroacetate ion reaction ment went practically to completion. Duplicate on going from butanoic acid to hexanoic acid experiments were performed in each solvent a t (lines 1 and 2 of Table 11) points to a decrease four different temperatures over about a 20" in the degree of dimerization of the six-carbon interval. The average rate constants calculated in acid. the usual manner from the slopes of the experiIt is highly probable that decanoic acid (line mental logarithmic plots are brought together in 3) does not exist a t all in the dimeric form.5 The Table I. The parameters of the Eyring E q ~ a t i o n , ~fact that AS* decreases on going from hexanoic based upon the data in Table I, are shown in acid to decanoic acid in the malonic acid reaction Table 11. Included in the table for purposes and increases in the trichloroacetate ion reaction of comparison are corresponding data for the de- probably is connected with the fact that the composition of the trichloroacetate ion in several trichloroacetate ion, because of its electron reother solvents previously studied, as well as data pelling effect, is able to keep the long hydrocarbon for the malonic acid reaction where available. chain of decanoic acid a t a distance from the carboxyl group, whereas malonic acid is unable TABLE I APPARENTFIRST-ORDER RATE CONSTANTSFOR THE DE- to do this. CARBOXYLATION OF POTASSIUM TRICHLOROACETATE IN Hydrogen bonding (in the case of both acids SEVERAL LIQUIDS and alcohols) serves to immobilize or restrain some Temp., k X 106, Average of the mobile electrons present in a molecule. Solvent 'C. cor. see. - 1 deviation With the partial or complete destruction of hyn-Butyl aleohol 96.83 0.080 0.002 drogen bonding these bound electrons are re103.04 ,166 ,002 leased, and are thereby enabled to retreat to more 107.70 .357 .003 remote positions within the molecule under the 110.94 .399 .003 repelling influence of the trichloroacetate ion, or, n-Hexyl alcohol 77.03 .553 ,004 on the other hand, to respond freely to the at85.44 1.137 -01 tractive influence of the polarized carbonyl carbon 92.04 5.89 .02 atom of malonic acid. The result will be an in96.54 11.62 .02 crease in AH* in the first case, a decrease in n-Caproic acid 118.71 0.532 .005 AH* in the second. This principle is strikingly 124.54 1.537 .Ol illustrated by the data in lines 2 and 3, and again 134.23 5.09 .01 in lines 5 and 6 of Table 11. 138.86 8.96 .01 The large difference in AS* for the trichloroacetate ion reaction on going from ethanol to TABLEI1 KINETIC DATAFOR THE DECARBOXYLATION OF THE TRI- 1-butanol (lines 4 and 5 of Table 11) is commenCHLOROACETATE ION AND OF MALONIC ACID IN VARIOUS surate with the relative sizes of the associated complexes of the two alcohols which probably LIQUIDS consist of three or four molecules each.12 Trichloracetate ion Malonic acid AH* AH* It is interesting to note in Table I that, at about (kcal./ AS* (kcal./ AS* Solventa mole) (e.u.) mole) (e.u.) 9 6 O , the trichloroacetate ion decomposes 145 Butanoic acid*$9 35.1 $12.7 32.3 +2.5 times faster in 1-hexanol than in 1-butanol. The Hexanoic acid4 39.3 $22.1 32.5 $3.2 reason for this unexpected result can be deduced Decanoic acid",4 41.4 $27.7 26.6 -11.0 readily by studying the data in lines 6 and 6 of Ethanol'O 31.1 $15.3 Table 11. On going from 1-butanol to 1-hexanol l-Butanolll 31.0 $ 1.6 27.2 -4.4 the AH* for the trichloroacetate ion reaction 1-Hexanol'' 39.6 $33.7 26.0 -7.6 increases by about 9 kcal. and the AS* increases a The first superscript after the name of the solvent refers by more than 33 e.u. Apparently the association to the source of the trichloroacetate ion data, the second cluster or "supermolecule" of 1-hexanol (but not t o that of malonic acid. In the case of hexanoic acid, 1butanol and 1-hexanol, the superscript refers to the source of 1-butanol) has been broken up by the negatively charged trichloroacetate ion, releasing some of the of the malonic acid data. mobile electrons involved in hydrogen bonding, Discussion which then retreat as far as possible from the It is int,eresting to note that, in the case of all nucleophilic center. (The malonic acid reaction the aliphatic acids and alcohols listed in Table 11, does not ,show this kind of behavior.) The great both AH* and AS* for the decomposition of the increase in the rate of reaction in 1-hexanol is due trichloroacetate ion are generally much higher largely to the improl-ed entropy factor whirh than they are for that of malonic acid, also, in more than compensates for the retarding effect of nearly all cases where the AH* for the trichloro- the higher activation energy. acetate ion reaction increases with changes of The fact that the AH* for the malonic acid solvent! tha,t for malonic acid decreases. reaction is lower in both alcohols than is that of (7) S.Gladstone. K. J. Laidler and H. Eyring, "The Theory of Rate the trichloroacetate ion indicates that in the Processes," MoGrau;-tIill Rook Co., Inc.. New 'Pork, N. Y . , 1st ed., malonic acid reaction the alkyl groups of both 1911, p. 196, equation 158. ulcohols produce a +I effect, but in the reaction ( 8 ) L. W. Clark, THISJOCRNIL, 63, 1760 (l!J59). (9) L. W.Clark, ibid., 64, 41 (1960). (10) F. H. Verhoek, J . Am. Chcm. Soc.. 56, 571 (1934). (11) L. W.Clark. THISJOURNAL, 64, 508 (1960).
(12) W Huckel, "Theoretical Principles of Organic Chemistry," Vol. 11, Elsebier Pub. Co., New York, N. Y . , 1958,p. 341.
1760
VOl. 64
NOTES
with the trichloroacetate ion the potential +I effect is prevented from taking place. If decanoic acid and 1-hexanol do not associate through hydrogen bonding the difference of 1.8 kcal. in the AH* of the trichloroacetate ion reaction in these two solvents (lines 3 and 6 of Table 11) may be attributed to the additional electron withdrawal due to a -E effect on the carbonyl carbon atom of decanoic acid evoked by the negative charge on the attacking ion. The decrease of 7 e.u. in the AS* of the reaction on going from 1hexanol to decanoic acid is commensurate with the difference in the number of carbon atoms between the monomers of the two compounds plus the steric effect (screening effect) of the carbonyl oxygen atom in close proximity to the nucleophilic hydroxyl group. The results of this investigation strengthen the hypothesis that the electron repelling influence of the negative charge on the trichloroacetate ion tends to evoke unfavorable inductive and electromeric effects, and that the electron attracting influence of the electrophilic, carbonyl carbon atom of the un-ionized malonic acid tends to evoke favorable inductive and electromeric effects. Further work on this problem is contemplated. Acknowledgments.-(1) The support of this research by the National Science Foundation, Washington, D. C., is gratefully acknowledged. (2) The potassium trichloroacetate used in this research was supplied by Dr. James D. Solomon, Blackburn Laboratories, St. Elizabeth Hospital, Washington, D. C.
bands centered a t about 2000, 2150, 2300, 2400 and 2550 pi., with wave length uncertainties which range from 1 2 5 t o A40 A. His observed wave lengths are listed in Table I and have been claimed to be independent of cation. Audubert suggested in most of the papers cited that the bands could be explained in terms of the known energy levels of Nz, and in one paperI3 specifically compares four of his band wave lengths with those of one band each of four different Nz band systems. It is the purpose of this note to suggest that the broad features probably arise from groups of bands of the Nz Vegard-Kaplan (A% -+ X1&) system which originate from the v’ = 0, 1, 2, 3 levels of the A32: state. The band-head vacuum wave lengths from 1900-2600 k. of the gas-phase Vegard-Kaplan bands are also given in Table I together with their Franck-Condon factors Q ~ which ~ ~ give J ~ an indication of the relative band intensities for equal populations in the upper levels. It will be noticed in Table I that strong Vegard-Kaplan bands ( p v ~ v ~ ~0.1 or greater) occur in the region of each of the bands reported by Audubert. Positive identifications are, of course, difficult with such large uncertainties on wave lengths but the following paints lend strong support t o identification of the Vegard-Kaplan system as the most plausible emitter. (A) It is physically more reasonable to attribute all of the bands to one system than to rely on individual wave length coincidences between the bands and those of difference Nz systems. (B) In considering the mechanisms of dissociation of simple molecules, Henri16 compared lengths (re) and characteristic frequencies (we) of bonds ON T H E NATURE OF ULTRAVIOLET LIGHT within the molecules with those of the electronic WHICH ACCOMPANIES THE states of diatomic species into which the molecule may decompose. In particular he quotes the DECOMPOSITION OF SOME AZIDES’ values re = 1.13 k., we = 2060 cm.-l for the BY R. W. NICHOLLS* terminal NN bond in NaN3 t o be compared with Department of Phyaice, Univsrsit of Western Ontario, London, Onlar%o, the values re = 1.144 A.; we = 2046 cm.-l for 8ando free Nz in the C311ustate. From the close agreeReceived April $Q8 1960 ment between these data he went on to propose Some years ago Audubert and his colleagues that Nz is formed in the CSII, state in the slow reported that a very feeble luminosity accompanied thermal decompositio? of NaN3. The more recent the slow thermal decomposition of NaN3, 3,4.6.6 values of re = 1.148 A.; we = 2035.1 cm.-l for the ICN3,3AgNa,s*7 P ~ ( N s )C~B, (~N ~ )Ba(N3)z,3 ~,~ T1N8.8 C3IIU state,17 and the fact that the characteristic It also occurs near the anode during electrolysis frequency we = 2141 cm.-l is ncw associated with of NaN3 and The spectrum of the lumi- the antisymmetric stretching of the Na group as a nosity was studied over the wave length range 1800- whole18 while re for the terminal NN bond is 2600 k. a t very low dispersion using a quartz 1.128 k.l9 does not detract from the usefulness of monochromator and a copper iodide photon Henri’s suggestion which is really based (insofar counter,l0-lZand was found to be limited to broad as rough equality of re is concerned) upon the Franck-Condon principle. (1) This work ha8 been supported in part by contracts with the Air Force Office of Scientific Research, The Office of Naval .Research. The If it therefore be supposed that Nz is formed Air Force Cambridge Research Center and the Department of Defence during slow thermal decomposition of NaNa in the Production of Canada. 0’ = 0 level of C311u,three radiative transitions in (2) Temporarily on leave of absence at the National Bureau of Standards, Washington, D. C. cascade will be.expected to occur, viz., C311u -+
-
~
(3) R. .4udubert and H. Muraow, Compt. rend., 204, 431 (1937). (4) R . Audubert, ibdd., 204, 1192 (1937). ( 5 ) R. Audubert, ibid., 206, 748 (1938). (6) R. Audubert and J. Mattler, {bid., 206, 1639 (1938).
(7) R. Audubert, ibid., 206, 133 (1937). (8) R. Audubert. and C. R a w , ibid., 208, 1810 (1939). (9) R. Audubert,, ibid., 208, 1984 (1939). (IO) R. Audubert, J . Phys. et Rod.. 6, 452 (1935). (11) R. Audubert, Compt. rend., 200, 389 (1935). (12) R. Audubert, ibid., BOO, 918 (1935).
(13) R. Audubert. Trans. Faraday Soc., 86, 197 (1939). (14) R. W. Nicholls, Ann. Geophys.. 14, 208 (1957). (15) W. R. Jarmain, P. .4. Fraser and R. W. Nicholls, Astrophy. J . , 118, 228 (1953). (16) V. Henri, Compl. rsnd., 203, 67 (1936). (17) R. S. Mulliken, “The Threshold of Space,” Pergamon Preaa, Inc., New York. N. Y., 1957, p. 169. (18) E. H. Eyster and R. H. Gillette. J . Chem. Phye., 8, 369 (1940). (19) E. H. Eyster, i b d . , 8, 135 (1940).
Nov., 1960
1761
NOTES
TABLE I BAND-HEAD VACUUM WAVELENGTHS AND FRANCK-CONDON FACTORS FOR VEQARD-KAPLAN BANDS IN THE SPECTRAL REGION 1900-2600 A., AND AUDUBERT’S BANDSARISINQFROM TEIERMAL DECOMPOSITION AND ELECTROLYSIS OF SOME AZIDES Y V e g a r d - K a p l a n bandsA
(h.)
qv’d’b
1901 1936 1954 1990 2010 2027 2046
0.009 .07 ,003 .043
2085 2108 2125 2148 2189
.096 .004 ,044 .06 .lo2
2258 2302 2334 2347
.111 .Oil .053 ,044
2379 2425 2463 2474
.124 000 .lo6 .07
NaNa (Thermal)
AgNi (Thermal)
Audubert’s features ‘ I 1N: (Thermal)
(H.)
NaN;
(Elect)
-
HNa (Elect)
1950 f 25
1975 f 25 1990 f 52
.001
.099 .019
.
2040 f 30
2140 f 30 2150 f 25
2150 f 25
2300 f 35
2300 & 35
2400 f 40
2400 f 40
2140 f 25 2280 f 35
2511 .076 2561 .041 2605 ,158 2614 .019 * The bands predicted as occurring strongly are italicized. bert cites this band as doubtful.*
B311, (Second Positive Bands), BaIIg - A3Z: (First Positive Bands), AaZ: + X’Z’, (VegardKaplan Bands). Tables of Franck-Condon Factors which are available for these transitions16J0 may then be used to predict which bands will be most strongly radiated. They are u” = 0, 1, 2 u” 0, 1, 2 Y” 0, 2, 3, 4 v“ = 0, 1, 2, 5, 6 v” = 4, 5, 6, 7, 8 v” = 3, 4, 5, 9, 10, 11 v” = 2, 3, 6, 7, 8, 11, 12, 13 v‘ = 3 v” = 1, 2, 3, 5, 6 , 7, 13, 14, 15, 10
Nz Second positive bands. v’ = 0 N 1First positive bands. v‘ = 0 v’ = 1 v’ = 2 NzVegard-Kaplan bands: v’ = 0 v’ = 1 (From 1900-2600 A.) v’ = 2
The strong Vegard-Kaplan bands predicted by this scheme, which lie in the wave length range of Table I have been italicized there and it is noted that such strong bands appear in each of the five groups into which the table is divided. Additional Vegard-Kaplan bands can be predicted in a similar fashion if it be assumed that the v’ = 1, 2, etc., levels of C311, are also populated by thermal dissociation of azides. (C) It is possible that N, is formed during the decomposition of the azides directly into the AaZ: state from which the Vegard-Kaplan system is radiated directly. The primary Franck-Condon parabola for the system passes (in the usual v’, V” array) in the vicinity of the (3, 2), (2, 3), (1, 3), (1, 4) and (0,5 ) bands one of each of which is associated in Table I with each of Audubert’s spectral features. (20) W. Jarmain and R. W. Niaholls. Con. J . Phys., 31,201 (1954).
2130 f 25
2270 f 35 2300 f 35
2450 f 35
2390 f 40 2425 f 35
2500 f 40 2550 f 40
q,‘,”
2600 i 40 (2650)” is the Franck-Condon factor of the band.“ Audu-
The intrinsic weakness of the intercombination Vegard-Kaplan system may account in some part for the very feeble intensities observed by Audubert. The system nevertheless has been excited in the gas phase by electrical discharges of various kinds by Kaplan,21 Wulf and Melvin,22 Janin2a and Herman and Herman,24 by electron beams by Bernard,26and in solid N2by VegardZ6and by Peyron and B r ~ i d a . ~ ~ The bands observed during electrolysis of azides probably are excited in electrical discharges inside the small bubbles of N P liberated at the anode. Ozonizer-like electrical discharges caused by surface changes in such bubbles are known to give rise to light2* and as Wulf and Melvin point outz2 such discharges are favorable for the excit,ation of the Vegard-Kaplan system. One test for the operation of the cascade mechanism discussed in B above is to observe the predicted strong bands of the N2 second positive and first positive bands (which lay outside Audubert’s range of observation and which are also outside the visible spectrum). Experiments to search for these bands are in preparation. The absence of them will not weaken the case for identification of (21) J. Kaplan, Phys. Rev., 46, 675 (1934). (22) 0. R. Wulf and E. H. Melvin, Phys. Rev., 86, 687 (1939). (23) J. Janin, Ann. Phye., [121 1, 538 (1946). (24) R. Herman and L. Herman, J . Phya. et Rad., 7, 203 (1946). (25) R. Bernard, Compt. rend., 300, 2074 (1935). (26) L. Vegard. Skriftar Norske Videnekaps. Akad. Oslo, No. 9 (1938). (27) M. Peyron and H. P. Broida, J. Chem. Phys.. 30, 139 (1950). (28) P. Jarman, Proc. Phys. Sac., 73, 628 (1959).
Vol. 64
1762
Audubert's features with Vegard-Kaplan bands, but will make the cascade mechanism less likely.
THE GROWTH OF LARGE SINGLE CRYSTALS OF ZINC OXIDE' BY J. IT. NIELSEX~ AND E. F. DEARBORN Bell Telephone Laboratories, Inc., Murray H i l l , New Jersey Received M a y 3 , 1050
Single crystals of hexagonal zinc oxide can be grown from the vapor phase by the method of Scharowsky.3 Lander4 has improved upon Scharowsky's method and recently Laudise and Ballmanj have shown that zinc oxide can be grown hydrothermally. Both of these methods produce crystals which have a needle-like habit with the needle axis and the crystallographic c-axis coinciding. Growth perpendicular to the c-direction is usually quite slow and seldom exceeds a millimeter. Growth of single crystals from molten salt solution was attempted in the hope crystals large in the a-dimension could be obtained which could be used as seeds for the hydrdhermal method. As is usually t'he case with molten salt systems, insufficient phase equilibrium dat,a were available t'o choose a solvent for ZnO, and the selection had to he made by trial and error with the help of previous experience. It was quickly discovered that LiCl arid KF were very poor solvents for ZnO. PbF2, however, dissolved appreciable quantities of ZnO, and this solvent was selected. PbFz previously had been used by Jona, Shirane and Pepinsky6to grow single crystals of lead zirconate, and it had been used by the authors? to grow cryst.als of Mn203 and Lao.,Pb~.3Mn03. ZnO crystals grown from molten PbFz between 1050 and 1150" exhibited the desired platey habit, the plane of the plate coinciding with the (001) crystallographic plane. Experimental All crystal growth experiments were done in resistance furnaces which have been described previously.* The samples were contained in 100 ml. platinum crucibles fitted with lids. They were prepared from reagent grade chemicals or materials prepared from reagent grade salts. Two hundred g. of PbF2 and 22 g. of ZnO were convenient amounts to use. Temperatures were measured with Pt10% Rh-Pt thermocouples. It was found that 1150" was a convenient temperature from which to begin crystal growth. At this temperature the volatility of PbF2, though high, was not excessive. The samples usually were held a t 1150" for from 2 to 4 hours and then cooled a t a constant rate of from 1 to IO" per hour. The system reaches a steady state quickly and it is possible a holding time of less than two hours would bc satisfactory. The samples usually were withdrawn near 800" and allowed t,o cool in air. Both air and oxygen at____(1) Presented at, the 135th Meeting of the American Chemical Society, Boston, Massachusetts, April 1959. (2) Now at Airtron, Division of Litton Industries, ;LIorris Plains, N. J. (3) E. Scharowsky, Z. physik., 156, 318 (1953). (4) J. J. Lander (to be published). ( 5 ) R. A. Laudise and A. A. Ballman (to be published). (6) F. Jona, G. Shirane and R. Pepinsky, Phys. Rev., 97, 1584 (1955). (7)J. W. Nielsen and E. F. Dearborn (unpublished work). ( 8 ) J. n . Nielsen and E . F . Dearborn, Phys. Chem. of Solids,6, 202 (1958).
PbF,
5
13 15 20 25 M O L E P E R CENT ZflD.
Fig. 1.-Solubility
33
35
curve of ZnO in PbF2.
mospheres were used in the growing furnace and no change in growth characteristics resulted. Since no method of chemically separating PbF? from ZnO has been found, the crystals must be broken away from the melt. It is still possible, however, to separate crystals which are 1 to 5 cm. in their largest dimension (the adimension) by tapping the bottom of the crucible containing the sample. It is found that the PbFz does not cling very tenaciously to the ZnO crystals. Solubility Data.-A rather crude solubility curve of ZnO in molten PbF2 was determined. The system has only a simple eutectic between 0 and 35 mole % ZnO. The curve, which is shown in Fig. 1, was obtained by preparing melts with a large excess of ZnO present, holding them for a t least two hours a t the desired temperature and quenching them in water as quickly as possible. The excess ZnO which had floated on the surface of the melt was analyzed for zinc and lead. The percentage zinc oxide present was calculated assuming all the zinc was present as ZnO and the lead as PbF2. The eutectic composition m-as determined from a crystal run which was allowed to cool to solidification. The eutectic was found to be a t 8.8 mole % ZnO and 733". Two of the points near 850" were determined by approaching the equilibrium from the other side, that is, slowly cooling the melts to that temperature from 1150°, then quenching them. Since these points fell on the curve nicely, i t was assumed that two hours is ample time to allow for a steady state condition to be reached by samples held a t constant temperature. The average error in the ZnO ronrentration determined by this method is 5273 except a t the eutectic where the ZnO concentration is known more accurately. I t should be pointed out that the quench technique is not as reliable as one might wish when melts are highly fluid, but the great volatility of PbFz makes any other technique extremely difficult. Another possible source of error lies in side reactions. X Ray powder photographs of quenched melts of ZnO and PbFz revealed that only ZnO and P-PbF2 with a small amount of a-PbF2, were present. Thus side reactions, if they occur, do not yield products in amounts detectablr by X-ray diffraction . Spectrochemical analyses of the crystals indicated :t total impurity rontcnt of less than O.lTo.
Discussion Although the habit of the ZnO crystals 011 the melt surface was platey, it was discovered that crystals nucleated a t temperatures below 1050" tended to be more drum shaped. Thus the ratio of the rate in the a-direction to the rate in the cdirection varied from about 50 to 1150" to near unity a t 1050". It was also observed that crystals
NOTES
Nov., 1960 grown by evaporation of PbF2at 1150' were always very thin plates and exhibited the largest ratio of rate in the a-direction to rate in the c-direction. The reason for this change in habit is not known, but it probably arises from a rather subtle change in the structure of the PbF2-ZnO melt. The possibility that a buildup of PbO was responsible in some way for the habit change was eliminated by growing crystals in melts to which PbO had been added. Yo change in growth characteristics was observed. At present, no experimental studies of such melts which could lead to an explanation of the habit change have been made. I n conclusion, large single crystals of ZnO which possess an unusual and convenient habit have been grown from molten lead fluoride. Such crystals already have been used successfully as seeds in hydrothermal crystal growth experiment^.^ The combination of the two methods should lead to growth of the first very large and pure crystals of ZnO. Acknowledgments.-The authors wish to thank W. Hartmann and E. Bloom for the excellent chemical analyses required for this work.
fer were carried out either i n vacuo or in a carbon dioxide-free dry box. The heat content of the capsule was determined by separate experiments over the temperature range investigated. A platinum vs. platinumrhodium thermocouple, previously calibrated against a National Bureau of Standards therniocouple, was used for furnace temperature measurements. The measured heat contents of NazO above 298' K. are given in Table I. The results are based upon a molecular weight of 61.98' for Xa20 and are expressed in defined calories per mole (1 cal. = 4.1840 abs. joules). Corrections for the sodium carbonate impurity were made using the integrated heat capacity data of Popov and Ginzbergs and for sodium peroxide using the data of Chandrasekharaiah, Grimley and M a r g r a ~ e . ~It was necessary to extrapolate the data for Na2C03 and NazO2 to higher temperatures in order to correct the sodium oxide data. Measurements at higher temperatures were prevented by the attack on the gold solder used in fabricating the capsule by either the sodium oxide or the impurities. TABLE I JIEASURED HEATCONTENT OF SODIUM OXIDEABOVE 298.15'' K. (NazO, mol. wt. = 61.98)
THE HIGH TEMPERATURE HEAT CONTENT OF SODIUM OXIDE' BY ROBERT T. GRIMLEY~ AND JOHN L.
T,OK.
MARQRAVE
380.1 477.2 577.2 673.5 776.4
Department of Chemistry, University oJ Wisconsin, Madieon, Wisconsin Received M a y 6 , 1960
A review of available literature indicates that the only high temperature heat content data reported for the alkali oxides are those of Shomate and (>ohen3 for lithium oxide. In addition, Naylor4 has estimated the high temperature heat content of sodium oxide by employing the NeumannKopp rule in conjunction with experimental data for NazTi03, Na2Ti20s,Na2Ti30, and TiOz. Experimental values for the heat content of sodium oxide are presented in this paper. Measurements were made using a copper block drop-type calorimeter which has been described previou~ly.~The sample of sodium oxide was prepared from commercial grade du Pont sodium oxide by purification according to the method of Klemenc, Ofner and Wirth6 and X-ray examination of the product showed only the lines of sodium oxide. Analyses were made for NazO, NazOz and h-azCOs and the sample mas found to consist of 96.76% Na20, 2.33% Na2C03 and 0.91% Na20z. The presence of KaOH was ruled out on the basis of indirect analysis. The sample (9.2728 g. in vacu,o) was t'ransferred to x gold capsule, the neck of the capsule squeezed shut and then soldered with gold. All operations of preparation or t,rsns(1) Presented before the 134th Meeting of the American Chemical Society, Chicago, Illinois, September 8, 1958. (2) Department of Physics, University of Chicago, Chicago, Illinois. (3) C . H. Shomate and A . J. Cohen, J . A m . Chem. Soc., 7 7 , 285 (1955). (4) B. F. Naylor, ibid., 67, 2120 (1948). (5) J. L. Margrave and R. T. Grimley, THISJOURNAL,62, 1436 (1958). (6) A. Klemenc, G. Ofner and IT. Kirth. %. anorg. Chem., 265, 221 (1951).
1763
cal./mole
T, OK.
H T - Hsoe.ls, cal./mole
1418 3191 5257 7190 9276
876.0 980.4 1078.3 1174.6
11,447 14,045 16,539 19,409
HT
- H298.15,
The high temperature heat content data and the heat capacity may be represented by the equations HT
- H t ~ d =. ~14.497' ~ + 4.94 X
10-W' -4799 cal./mole (29&1170°K.; 1 2 % ) CP = 14.49 9.88 X 1O-T cal./deg./mole
+
A plot of the Shomate functionlo using the low temperature heat capacity data of Furukawa" gave some evidence of a low energy transition in A study of this range using the range 500-800'K. adiabatic methods on a sample of higher purity would appear to be desirable. Destruction of the capsule during measurements at higher temperatures prevented further study by the drop method. From the above equations and a value for S2g8 determined by Furukawal' to be 17.99 e.u., the heat contents, entropies and the free energy functions were calculated at 100' intervals and the results are listed in Table 11. I t is interesting to note that the sodium oxide heat content which Naylor4 estimated by subtracting the a-Ti02 contribution to the XazTiOS heat content agrees remarkably well with the experimental values. The agreement with Naylor's mean values obtained from data on additional compounds is somewhat (7) E. Wichers, J . A m . Chem. Soc., 80,4121 (1958). (8) A I . M. Popov and D. M. Ginrberg, J . Gen. Chem. (U.S.S.R.), 26, 1103 (1956). (9) M. S. Chandrasekharaiah, R. T. Grimley and J. L. Xargrave, Tms JOURNAL,63, 1505 (1959). (10) C. H. Shornate, rbid., 58, 36s (1954). (11) G. E'urukawa, private communication.
NOTES
1764
VOl. 64
poorer, but, in general, does not differ by more than lOql,.
Preparation of Cadmium Chloride.-The anhydrous cadmium Chloride used waa prepared from Baker Chemical Company analytical reagent cadmium chloride containing 2.5 molecules of water. The anhydrous salt was prepared TABLE I1 by recrystallizing CdC12.2.5H20from an aqueous solution, SMOOTHED HEATCONTENTS, ENTROPIES A N D FREEENERGYslightly acidified with hydrochloric acid. These crystals were dehydrated in a tube oven, heated to 150°, through FUNCTIONS FOR SOLIDSODIUJ~ OXIDE which was passed a stream of dry hydrogen chloride. (NazO, mol. wt. = 61.98) The excess hydrogen chloride was removed by means of a current of dry nitrogen. The resulting salt then was Hi,- H n e s . ~ , ST - s ~ s . 1 6 , (FT -THnga"') T ,OK. csl./mole cal./deg./mole cal./deg./mole evacuated until successive samples, taken a t hour intervals, gave a constant pH reading. 400 1 ,789 5.27 18.79 Measurement of pH.-The pH wag measured with a 500 3 ,683 9.49 20.11 Beckman Laboratory Model G pH meter using a Beckman 600 5,676 13.12 21.65 general purpose shielded glaas electrode and a fiber type sealed calomel Beckman electrode. The hydrolysis cell 700 23.23 7,768 16.31 contained openings for two electrodes, a thermometer, a 800 9 ,959 19.27 24.81 conductivity water inlet and a nitrogen inlet and. outlet. 12,247 900 26.34 21.96 The cell containing the wei hed cadmium chlon.de was 14,636 1000 24.48 27.83 flushed with nitrogen gas, before and durlng the time the 17,122 1100 29.27 purified water was added. A magnetic stirring bar agitated 26.85 the solution for 15 minutes before the pH readings were Acknowledgments.-The authors wish to ac- taken. When cadmium chloride is placed in water, these dissociknowledge the support of this research by the United States Navy through the Callery Chemical ation and association reactions may occur
Company under a contract with the Bureau of Aeronautics.
+
1 3 KARL ~ H. GAYERA N D RUDYM. HAAS Department o f Chemtstrv, Wayne Stale University, Delrozt, Mzchigan Received Y a y 0, 1060
The study of the hydrolysis reactions and the evaluation of equilibrium constants for cadmium chloride has been undertaken by other investigators. If a true hydrolysis constant can be obtained for any species of cadmium chloride over a definite concentration range, it can be related t o an over-all reaction. The hydrolysis constants (Kh) and the reactions predicted by others are: at 100" for Kullgren' reported a Kh of 3.3 X the reaction CdClf €320 CdClOH H+; Chaberek, Courtney and Martel12 calculated R Kh of 2.5 X lo-'* at 30" in a 0.1 M potassium chloride medium for the reaction Cd++ H20 Cd(0H) + H+. According t o this last reaction it is assumed that the chloride ion from the potassium chloride has no effect on the hydrolysis reaction of the Cd++ ion. Doubt is cast on this assumption because it is further assumed that the concentration of the Cd++ ion is equal to the total concentration of cadmium chloride used. Our results seem to indicate that the major hydrolysis product at 25" is CtlC1OI.Z.
+
+
+
+
Experimental Purification of Water.-Water was prepared in a manner similar to that described by Gayer and Woontner.3 Ordinary distilled water was distilled in a Pyrex distillation apparatus from an oxidizing solution of sodium hydroxide and potassium permanganate. The resulting water was then redistilled by adding one drop of orthophosphoric acid per liter of water. Last traces of carbon dioxide were removed by boiling the h a 1 distilled water for about 15 minutes. Purification of Nitrogen.-Purified tank nitrogen was passed through several gas washing bottles of dilute sodium hydroxide, sulfuric acid and water. (1) C. Kullgren, Z. physik. Chem. (Leiprig), 86, 466 (1913). (2) 9. Chaberek, Jr.. R. C. Courtney and A. E. Martell, J . Am. Chem. Soc., 74, SO57 (1952). (3) K . H. Gayer and L. Woontner, J . Chum. Ed., 93,296 (1956).
(1) (2) (3)
+
HYDROLYSIS OF C14DMIU&1CHLORIDE AT 25'
+ +
CdC12 CdCl+ C1C d C l + s Cd++ C1CdClz C1CdCLCdCli 2C1- If CdCla"
(4)
In order to obtain the activities of these species, it is necessary to know the dissociation constants of the reactions involved. These are not available as such. As an approximation, however, it was assumed by Harned and Fitzgerald that the first reaction goes to completion, the third and fourth do not occur, and the second is the only one in which an equilibrium is involved.* By means of electromotive force measurements, they calculated the dissociation constant ( K ) for this reaction to be 1.1 X lo-*. Other Kt values given in the literature, assuming only the presence of Cd++, CdC1+ and C1-ions are 0.0106 and 0.0101.6
- aCd"
2--
aCl-
aCdCIC
- CCd* CClCCdCl+
.fCdHfCIfCdCIC
Here a, C and f refer to the activity, concentration and activity coefficient, respectively. The various concentrations of CdCl+, C d + + and C1- ions can be evaluated by setting X . = Cd, X M = C1- and M - X = CdCl+, where M IS the molar concentration of the cadmium chloride used. The activity coefficients can be obtained by solving the regular form of the Debye-Huckel equation
+
where Zi is the charge on the ion, a, is its effective diameter and u is the ionic strength of the solution. The constants A and B were taken as 0.509 and 0.330, respectively, and.the effective diameters7 of the Cd++, C1-, CdC1+ and H + ions were taken as 5 X 10-8, 3 X 10-8, 4 x 10-8 and 9 X 10-8 cm., respectively. By the method of successive approximations, the following ionization equation can be solved by successively approximating X, the concentration of the Cd++ion. The activity coefficient of the Cd++ion, fx is given in the regular form of the Debye-Huckel equation
logf, =
-2 . 0 4 d 4 r M
1
+ 1.65 X 10-8-
(4) H. 9. Harned and M. Fitrgerald, J . A m . Chem. SOC.,68, 264 (1936). ( 5 ) H. L. Riley and V. Gallafant, J . Chem. SOC.,514 (1932). 2 6 , 592 (6) E. C. Righellato and C. W. Dsvies, Trana. Faradoar SOC., (1930). (7) J. Kiellsnd, J . Am. Ckem. Soc., 59, 1 6 i 5 (1937).
NOTES
Nov., 1960
The hydrolysis constant for cadmium chloride is then calculated from a value of the concentrations of CdCl+, Cd++, C1- and H + (from pH measurements) ions and the activity coefficients of these ions. All the following hydrolysis reactions will probably occur, to some extent; i t can however he assumed that one of them predominates over the rest. C d + + H20 _J CJ(OH)+ H + (1) C d + + 2H20 I_Cd(OH)2 2H+ (11) CdCl+ H2O )r CdClOH H+ (111) The K h equations for the above are given in consecutive order. CCd(OH)+CH+ fCd(OH)+ fH+ = Khi CCdu fedtt -CUT+- Ho+)CH+ fCd(OH)+ fE+
+ +
+ + +
+
cia*- (H+- a091f;cd*-
(E+- Ho+)l
1765
due t o the hydrolysis of the cadmium chloride. Again assuming that only one of the above reactions occurs, the quantity of CdClOH or Cd(OH)+ formed is equal to the increase in the hydrogen ion concentration, and the quantity of Cd(0H)Z formed is equal to one-half the increase of the hydrogen ion concentration. The concentrations Ccd++ and CCdCl+ and the activity coefficients fed" and fCdCI+ are the values of these species after hydrolysis has occurred whereas Ca* and CGCI+ and fCd* and faE,+are the values of these ions, respectively, before h y d r o e hrts orcurred. The terms [Cd++ l/z(H+ - Ho+)],CdC1+ ( H + - Ho+), etc., that is, those containing the bar symbols, simply illustrate how the measured values were corrected to give the equilibrium values used in the calculations.
-
Approximations .C(H'- Hof) < < ca* (71/2(H+- Ho") < < CG* C(E+- Eo+)
98% Cr(I1) oxidized), be obeyed throughout the specified concentration range. and then was exposed to the atmosphere, heated before the potassium hydroxide treatment failed to reveal with potassium hydroxide to free ligands attached solution of free glyoxal. to the Cr(III), and filtered while hot and again after the(7)presence This solution was buffered with HSOd- and SO'-5 Similar cooling in an ice-bath.* Tests made on portions of tailing off was observed with unbuffered solutions. (8) Although some oxidation of Cr(I1) by ethylenediaminetetraacet,he filtrate to determine the nature of the reduced tate was observed, the amount of oxidation in a few minutes was negspecies gave no indication of chloride, formate, ligible. formaldehyde, glyoxalate or glyoxal, but gave posi(9) H. Taube and H. Myers, J . Am. Chem. Soc., 1 6 , 2103 (1954). tive evidence for the presence of glycolate by the deep red-violet color of the 2,7-dihydroxynaphthalene test.5 Comparison of' the absorbance a t T H E HEAT OF COMBUSTION OF 540 mp of the red-violet product with the absorDICYANOBCETYLENE bancies given by solutions of known glycolate conBYGEORGE T. ARMSTRONG AND SIDNEY MARANTZ centrations indicated that the original filtrate conThe National Bureau of Standards, Washington, D . C. tained 95% of the amount to be expected for quanReceived June 6 , 1060 titative reduction of oxalate to glycolate by the Cr(I1) .'j Cyanogen, dicyanoacetylene and the dicyano(1) Supported by the A.E.C. under Contract AT(ll-1)-378. polyacetylenes are highly unstable thermodynami(2) The observations reported were made in the course of a study of cally. When they are burned with the proper the catalysis b y C r + + of the formation of Cr(II1) complexes. Alproportions of oxygen, carbon monoxide and though the research failed in its purpose with oxalate as ligand, a nitrogen may be the principal products of comstriking effect was demonstrated with pyrophosphate as ligand. (3) H. Moissan, Compt. rend., 92, 1051 (1881); Ann. chim. et p h y ~ . , bustion. These facts lead to the possibility of [5] !Xi, 401 (1882). obtaining very high temperatures of combustion, a (4) Much of the perchlorate was removed as KClO4; large amounts possibility which has been borne out by experiments of perchlorate interfered with some of the tests made. of Conway, Smith, Liddell and Grossel on the ( 5 ) On acidification with nitric acid and addition of silver nitrate, no cloudiness from silver chloride was observed. The qualitative combustion of cyanogen. Kirshenbaum and tests for the organic compounds were taken from F. Feigl, "Spot Testa Grosse2have calculated that even higher temperain Organic Analysis,"' D. Van Nostrand Co., Inc., Princeton, N. .I., QR,, =
6.3 f 0.6
where the limits are those of the derived value of
5th edition, 1954. (6) The potasaium hydroxide treatment may have resulted in at least partial conversion of any glyoxal to glycolate. Tests on the
(1) J. B. Conway, W. F. R. Smith, W. J. Liddell and A. V. Grosses J . Am. Chem. Soc.. 17, 2026 (1955). (2) A. D. Kirshenbaum and A. V, Grosse, ibid., 78, 2020 (195D).
Nov., 1960 tures, above 5000°K., should result from the combustion of dicyanoacetylene, C4N2. The heats of formation of C4Nzand related cyanoacetylenes are therefore of great interest. The heat of formation of cyanogen gas has been established to be 73.84 f 0.43 kcal./mole.a The only measurement of the heat of formation of any of the cyanoacetylenes is that by Moureu and B ~ n g r a n d . ~The lack of further thermochemical study in spite of their interesting linear conjugated structure probably is the result of the scarcity of materials. Moureu and Bongrand, using a bomb calorimeter, burned a C4N2sample contained in a glass bulb, which they ignited with a platinum fuse and a naphthalene promoter. Reaction was incomplete, some carbon being formed in all experiments. Carbon deposits were not recovered or their amounts determined in four experiments, in which the heats of combustion ranged from 6426 to I n a fifth experiment the carbon 6660 cal. g.-'. was collected and burned for analysis, a correction of about 1% was applied for the incompletely burned material and the heat of combustion at constant volume, 6790 cal. g.-l (solid), obtained from this experiment mas reported. From this they derived the heat of combustion (solid) a t constant pressure to be 515.5 kcal. mole-'. We have redetermined the heat of combustion of C4N2 in order to test the value reported by Moureu and Bongrand. Experimental Materials.-Two samples of dicyanoacetylene were SUPplied by Professor A. V. Grosse, of the Research Institute of Temple University. As received, these samples consisted of a 6.0-g. sample (KO. 1) and an 8.5-g. stmple (No. 2 ) reportedly melting a t 18-20" and 20.0-20.5 , respectively. The infrared spectra had shown no impurities containing C-H or C-OH groupings. Impurities, if present, were believed to be other materials having similar structures. The samples were refrigerated with Dry Ice to prevent thermal decomposition. During storage some brown coloration developed in the samples. Shortly before use each sample was purified by sublimation into ampoules refrigerated with liquid nitrogen in an evacuated glass apparatus. The material was divided at this step into portions suitable for single combustion determinations and sealed until ready for rise. The portions thus obtained were white crystals, quite free of any brown color, which formed a colorless liquid on melting. Mass spectrometer andysis5 showed no substances identifiable as impurities. Samples for the bomb calorimeter were prepared by filling small thin-walled glass sample bulbs of the type commonly used for volatile liquids in this Laboratory.6 Transfer from ampoule to sample bulb was made by hypodermic syringe in a dry-box filled with nitrogen. The filled bulbs were removed from the dry-box and sealed by torch to permit them to be weighed. Visible charring invariably occurred a t the point of the seal, and several samples were lost because the seal could not be made. Where successful seals were obtained, the charred material was considered to be a negligible fraction of the total weight of sample. In the time necessary for weighing, between the time of closure of the bulb and its introduction into the calorimeter, a pale straw color appeared in the liquid.
-
(3) J. W. Knowlton and E. J. Proaen, J . Research Natl. Bur. Standards, 46, 489 (1951). (4) C. Moureu and J. C. Bongrand, Ann. chin., 14, 5 (1920). (5) M a s s spectrometer analysis waa performed by V. H. Dibeler of the Mass Spectrometry Section a t the National Bureau of Standards. ( 6 ) R. 8. Jessup, "Precise meaaurement of heat of combustion with a bomb calorimeter," Natl. Bur. Standard6 Monograph No. 7, (U.S. Government Printing Office, Waahington 2.5, D. C., 1960).
1777
NOTES
At no time was there evidence of violent decomposition, even in the near presence of a flame. A deliberate effort to initiate explosion was made by turning a torch directly on a filled bulb which had failed to seal properly. No explosion occurred, but after a brief time of smoky flame, a very vigorous, non-smoking bright blue jet appeared, which, however, was not self sustaining under the conditions of the test. Combustion Experiments.-Samples were burned in a Dickinson type calorimeter using an Illium bomb, and techniques similar in most respects to our customary practices.6 The calorimeter was calibrated with benzoic acid, National Bureau of Standards standard sample 39 g, using the certified heat of combustion 26,433.8 j.g.+ (weight in vacuo). Absence of hydrogen in the sample required some modification of the ordinary technique. Although no water was to be formed by the combustion, one ml. of water waa placed in the bomb as usual in order to dissolve oxides of nitrogen. It waa found, however, that practically none of the oxides dissolved in the water when i t waa on the bottom of the bomb. The difficulty was attributed to the small surface area exposed for solution of the gases. The water was thereafter placed in a shallow platinum pan supported by one electrode and located directly over the flame. Evaporation of the water by the flame during combustion caused a deposit of fine droplets on all walls of the bomb. The oxides were found to be dissolved in the water aa nitric acid under these conditions, and were determined by titration. The carbon dioxide formed in each experiment was measured by absorption in Ascarite. Some carbon was formed during each combustion experiment. The amount was about 1% of the total carbon in experiment 4, the most pronounced case, and was much less in the other experiments. Gases from the bomb were tested for carbon monoxide, but none was found. Samples used were smaller than necessary to produce a three degree temperature rise of the calorimeter because of (a) the small amount of material available for combustion, (b) the expectation that a very fast reaction would cause an exceptionally high pressure in the bomb, and (c) the experience of Moureu and Bongrand that larger samples tended to produce more carbon. Heat Measurements .-Of eight experimental measurements made after the technique had been developed for dissolving oxides of nitrogen, one was discarded because of B leak in the bomb, one because of a very large inconsistency in the weights of carbon dioxide and C ~ N Zand , one because an unidentifiable black deposit collected in the carbon dioxide absorption tube and because a small leak apparently remained unsealed in the glass sample bulb. The results of the remaining five experiments are shown in Table I. Two experiments were made on sample No. 1 and three on sample No. 2. In Table I, msis the mass of CINl in grams; ma' is the mass of C ~ Ncalculated Z from carbon dioxide collected; -A& is the observed energy of the bomb process in joules; --pl is the correction in joules for nitric acid formed in the reaction; --QZ is the firing energy in joules; -q3 is a correction in joules for the amount of carbon formed; and -q4 is the Waahburn correction. The correction -ga was calculated on the assumption that the difference between m. and ma' waa due to the conversion of some CINz to carbon instead of carbon dioxide. AEOTis the internal energy change in kj. mole-' a t the final bomb temperature (300.2 t o 301.4"K.) for reaction 1; and A H O Z W 18 the enthalpy change at 298°K. for reaction 1
+
+
CdNz(1) 402 = 4coz Nz (1) The mean value for the enthalpy change of reaction 1 in the standard state is A"hJ(1)
= -2078.5 f 0.7 kj. mole-'
=
-496.8 f 0.2 kcal. mole-'
The uncertainty given is the standard deviation of the mean. The value found here is about 4% smaller than was found by Moureu and Bongrand.
Discussion The heat of combustion given above can be used to calculate the standard heat of formation of C I N t liquid. Using AHor288[C02]= -94.05 kcal.
1778
Vol. 64
MOTES
TABLE I COMBUSTION DATAAND HEATOF REACTION 1 Sample
1 1 2 2 2
me.
me',
g.
g.
0.65711 0.65667 .37582 ,37421 ,44996 .44915 ,38954 ,38607 ,23974 .23829 Mean Stand. dev. of the mean
-AE'b, 1.
18019.1 10282.2 12344.0 10582.5 6549.2
-,?I,
i.
73.6 45.9 47.8 48.0 24.4
m01e-l~~we obtain AHof~s8[C,Nz(1)]= 120.6 kcal. mole-'. The heat of vaporization is given by Saggiomo8 to be AHvap = 6.875 kcal. mole-l. If this is taken as the best approximate value available for the enthalpy difference between gas and liquid in their standard states, the heat of formation of the ideal gas is AHof298[C4N2(g)]= 127.5 kcal. mole-'. Using AH0f298[C] = 171.299 kcal. mole-' and AH0rz98[N]= 112.96 kcal. r n ~ l e - l , ' ~we * ~ obtain for the total binding energy of the C4Nzmolecule 783.6 kcal. mole-1 a t 298°K. (7) F. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine and I. Jaffe, "Selected values of chemical thermodynamic properties," Natl. Bur. Standards Circ. 500 (U. S. Government Printing Office, Washington, D. C., 1952)). (8) A. J. Saggiomo, J . Org. Chem., 2 2 , 1171 (1957). (9) W. H. Evans, private communication. (10) L. Brewer and A. W. Searcy, Ann. Revs. Phys. Chem., 7 , 259 (1956).
NON-RESONANT MICROWAVE ABSORPTION AND RELAXATION FREQUENCY ,4T ELEVATED PRESSURES BY A. A. MARYOTT~ A N D G. BIRNBAUM~ b
a National Bureau of Standards, Washzngton, D . C . , and The Research Laboratories, Hughes Aarcraft Company, Culver Ctty, Calzfornaa Receaved June 10, 1960
The present note is part of a systematic investigation of the non-resonant spectra of symmetrictop molecules in the vapor state. Previous work has dealt with the detailed shape of these spectra,l the dependence of relaxation frequency on temperature,2 and the use of these spectra for precise determinations of values of the electric dipole moment. 3 , 4 Previous measurements were confined, for the most, part, to moderately low pressures (several atmospheres or less). In this region, the relaxation frequency increases in proportion to the pressure, or density, in accord with a mechanism of binary collisions. In the present investigation, data on trifluorochloromethane are extended to inuch higher densities in order to examine ihe dependence of relaxation frequency on density in greater detail. Trifluorochloromethane has a comparatively small dipole moment, 0.50 debye. It was selected because, unlike symmetric tops with large dipole moments, the non-resonant spectrum is characterized b j a single relaxation frequency parameter (1) G. Ihrnbaurn, J Chem P h y s , 27, 360 (1957). (2) A. A. Maryott and G. Birnbaum, abid., 82, 1501 (1960). (3) A. A. Maryott and G. Birnbaum, %bid.,24, 1022 (1956). (4) A. A . Maryott and S. J. Kryder, abid.. 27, 1221 (1957).
-93,
j.
j.
i.
- AE",
- AHOms'
-92,
kj. mole-'
mole-'
22 .o 22.0 22.0 22.0 22.0
2.9 10.6 5.4 23 .O 9.6
-15.4 - 8.6 -10.4 - 9.4 - 5.6
2074.6 2076.5 2077.7 2073.7 2076.9
2077.3 2079.1 2080.3 2076.3 2079.5 2078.5 rt0.7
-91.
kj.
-
and because reliable compressibility data are available. I n addition, certain factors which complicate interpretation, such as dipole association and the problem of the local field, should be minimized. Experimental The method of measuring the dielectric loss in compressed gases with tunable, transmission-type cavities has been described.*J Values of the loss tangent were measured up to 30 atm. (near saturation) with a hybrid-mode cavity (loaded Q = 4000) at a frequency of 1190 Mc. (0.0397 cm.-l). Values of thc dielectric constant e' were measured only with the accuracy needed to convert loss tangent to loss factor. The CClFa was part of the sample used previously..' The purity indicated by a mass spectrographic analysis was better than 99%. All measurements were made a t 30".
Results and Analysis The experiment,al data are given in Table I. The compressibility factor, 2, (2 = pV/RT) was obtained by interpolation of data reported by Albright and Values of the dielectric loss factor were derived from the measured values of the loss tangent according to the relation E" = E' tan 6, where E' is the real part of the dielectric constant. TABLE I DIELECTRIC Loss FACTOR e" AND DIELECTRIC CONSTANT e' OF CClF, AT 1190 Mc. AND 30" P (atm.)
Z
e'
e"
0.355 ,463 ,569 ,833 1.028 .1233 1.82 2.90 4.88 7.61. 10.89 14.29 17.66
0.997 ,996 .995 ,994 ,992 990 ,985 ,975 ,956 ,931 .899 .863 ,825 ,784 .730 ,684 642
1 ,0088 1.0010 1,0013 1.0019 I ,0023 1.002T 1,0041 1 ,0068 1.0118 1.019 1.029 1.n a 1.052
17.6 X 1027.8 38.6 64.1 81.6 94.8 124.0 147.4 158.4 159.5 164.B 168.5 176.4
1.065
185.5
1.080 1.097
196.1 215.7 231.8
21.03
24.45 27.89 30.88
1.116
The theoretical relations and procedures used to evaluate the parameters characterizing the nonresonant absorption in symmet,ric-top molecules (5) H. E. Bussey and G. Birnbaum, Reu. Sci. Inatr., 80, 800 (1959). (6) L. F. Albright and J. J. Martin. Ind. Eng. Chem., 44, 188 (1953.
NOTES
sov., 1960
1779
have been described.'~~Some modifications, however, are needed in the present case, where the data extend well above atmospheric pressure. I n particular, allowance must be made for substantial departures from the ideal gas density and for changes in the local, or polarizing, field when the dielectric constant can no longer be considered as unity. There is also some resonant absorption associated with the pure rotational transitions. The frequencies of the rotational lines are all considerably higher than the applied frequency. However, only a t the lower densities, where the lines are comparatively narrow, will their effect on the measured loss be negligible, so that, in general t"
=
E"
+
(non-resonant)
E"
(resonant)
(1)
I
The non-resonant loss factor is given by E''
(non-resonant) =
3kT
where N o ( p / Z ) is the molecular density ( N o being the ideal gas density a t 1 atmosphere and p the pressure in atmospheres), p the total dipole moment, F the local field factor, v the applied frequency, Av the relaxation frequency, and q = XJK ~ J K K 2 / ( J 2 + J ) . The function ~ J is K the fractional number of molecules in the rotational state designated by the quantum numbers J and K . With Av expressed in units of wave number the relaxation time 7 = 1/2?rcAv. Since the dipole moment is small, and F should not differ greatly from unity even a t the highest density, an adequate approximation for the local field factor can be obtained from the ClausiusMossotti relation. If €0 - €k is the non-resonant contribution to the state dielectric constant, then F
(EO
+ 2)(€k
2)/9
(e'
/
I .
/
'
I
- 2
+ 2)'/9
We shall consider first the data below two atmospheres, where the loss from the rotational lines is quite negligible, to obtain the value of the dipole moment and the relaxation frequency parameter, 6 = A v / ( p / Z ) , applicable a t lower densities. Equation 2 can be rearranged to give (p/Z)'F/t" = u
Fig. 1.-Non-resonant loss below 2 atm. plotted according ? eq. 3, (p/Z)ZF/s'' a8 a function of the square of the density, p / Z .
+ b(p/Z)'
(3)
where
Figure 1 shows a plot of the data according to eq. 3. The data fit a straight line within the experimental uncertainties of about 2yo and give 6 = 0.0392 cm.-l atm.-I. With the rotational constant A = 1.92 cm.-' and I3 = C = 1.10 cm.-ll7 the summation q = 0.263. The dipole moment is then 0.496 debye. This value is about 1y0lower than previously reported, the difference being due partly to the allowance made here for non-ideal gas behavior. If Av continues to increase in proportion to the density, it is evident from eq. 2 that, when Av2>> v2, E"/F will approach a limiting value provided the absorption from the rotational lines remains negligible. The loss factor computed on the assumption that Av is proportional to the density over the entire range is shown by curve 3 of Fig. 2. ( 7 ) U. I
