LIQUID AMMONIA RESEARCH in 1933-A REVIEW GEORGE W. WATT The Ohio State University, Columbus, Ohio
D
URING the years which have elapsed since the report to the d e c t that there is no change in the publication of the early work of Franklin, solubility of lithium in going from -80" to 0" might be =an% and Cady which pointed out the pos- in error due to the nature of the method which was sibiities of liquid ammonia as a medium for the study used. Accordingly, they have measured the vapor of chemical phenomena, the attention of many chemists pressures of liquid ammonia solutions of high lithium has been attracted by this somewhat unusual field of concentrations and, by extrapolation of the pressureresearch. This widespread interest has not been con- composition curves, have obtained the composition of fined to those who have furthered the work of Franklin the saturated solutions a t various temperatures. in establishing an ammonia system of compounds, but They have reported the following values in terms of workers in other fields have often had occasion to grams of Li/100 g. NH8: 10.698, 10.866, and 11.319 a t take advantage of some of the unusual properties of -63.5", -33.2', and 0°, respectively, or an increase liquid ammonia and liquid ammonia solutions. For of approximately six per cent. over the temperature example, the Dains method for the quantitative deter- range investigated. mination of the halogens in organic compounds is now Solubility values for twenty-six inorganic compounds widely used. Further, it is to be remembered that and for three sets of mixed-salt pairs have been rethis solvent has played an important role in Kraus's ported by Hunt and Boncyk (4). They claim to have fruitful researches on the nature of the metallic state checked their method against the vapor pressure and on free radicals. method and to have found perfect agreement. I t is An annual review of liquid ammonia research should difficultto reconcile Hunt's statement to the effect that be of interest not only to those actively engaged in this sulfur appears to be insoluble in liquid ammonia a t 25' type of work, but also to others who wish to keep with repeated observations made in our laboratories informed as to the progress being made in this particular and particularly with the work of Ruff and Hecht (5) field. who have reported solubility values of 25.65 g. S/100 g. In the past, little attention has been given to the solution at 16.4' and 21.00 g. S/IOOg. solution at 30'. problem of accurate determinations of solubilities in Linhard and Stephan (6,7)have pointed out probable liquid ammonia. However, this difficulty is being errors involved in Hunt's method and offer values which rapidly corrected. During the past year, Wiebe and they believe to be more nearly correct. These values, Tremeame (1) have developed a satisfactory method in terms of grams of salt/100 g. of solution, a t 0°, are for measuring the solubility of gases in liquids having reproduced in the following table: high vapor pressures and have applied this method to the determination of the solubility of nitrogen in TABLE I liquid ammonia. At 25', the solubility of nitrogen CI E~ I NO2 was found to increase from 2.22 cc. Nz/gram of NHa a t ~i 1.43 N~ 11.37 39.00 60.88 68.05 25 atmospheres to 54.83 cc. Nz/gmm of NHa a t 1000 K 0.132 21.18 04.81 9.52 atmospheres. Rb 0.289 18.23 08.15 Ipat'ev and Teodorovich (2) have reported that co 0.381 4.38 00.28 Ag 0.280 2.35 84.15 57.90 70.99 the solubility of hydrogen a t 25' varies between NH, 39.91 0.004 0.158 97.3 cc. H2/100 cc. NH3 a t 15 atmospheres and 1575 ca MX 0.00~ 3.85 45.13 cc. H2/100 cc. NH3 a t 250 atmospheres. The solusr 0.008 0.308 28.77 Bn 0.017 0.231 17.88 bility of helium was determined over a more limited pressure range at 20'" and was found to vary between 11.56 cc. He/100 cc. NHa at 5.2 atmospheres and 75.22 Using the above values, Linhard and Stephan have cc. He/100 cc. NHa a t 38.3 atmospheres. ~ u r t h & , studied the influence of ionic radii, solvation of ions, it was found that the solubility of both hydrogen and and ammonate formation upon the solubility of inorhelium increases with rising temperature. ganic salts in liquid ammonia. Some very interesting Johnson and Piskur (3) have determined the solu- relationships have been found. Patscheke (8) has made a thorough study of the bility of metallic lithium in liquid ammonia. Since the solubility of sodium is known to decrease and the solubility of sodium chloride between -76.6' and 43". solubility of potassium to increase with increase in A solubility maximum, 15.37 g. NaC1/100 g. of solution, temperature, the above authors believed that an earlier was found a t -9.5". Above this temperature, the 339
solubility decreases rapidly to a value of 1.76 g. NaCl/- tion constant, and this method has been applied to the above conductance measurements in those cases for 100 g. of solution a t 43". Between -76.6' and -9.5", sodium chloride exists in liquid ammonia as the am- which sutliciently accurate data were obtained. Some monate, NaC1.5NHa. Using thermal data supplied very interesting theoretical conclusions have been by earlier workers, Patscheke has calculated the heat drawn from the results obtained in these studies. As a result of greatly improved methods, older of solution of this ammonate at - 10' and has reported a value of -7.75 kg.-cal./mol. I t is interesting to values for the conductances of the amides of sodium and note that Patscheke obtained a value of 11.52 for the potassium in liquid ammonia have been corrected by solubility of sodium chloride at 0°, as compared with the work of Hawes (19) and the concentration range has been extended. The data obtained for sodium the value of 11.37 obtained by Linbard and Stephau. Similar studies on the solubility of sodium chloride amide do not agree with the theory of Fuoss and Kraus have been carried out by Abe and Hara ( 9 ) . They have or with the simple mass action theory. Somewhat determined the solubility of sodium chloride in liquid better agreement was found in the case of potassium ammonia between -23- and 30' and have reported amide. The extent of anodic dissolution which occurs during values which are in good agreement with those reported by Patscheke for the same temperature range. They the electrolysis of certain salts in liquid ammouia has have found that the transition between sodium chloride been investigated by del Boca (20). I t was found that and the ammonate, NaC1.5NH3, occurs a t -9.7' and the loss in weight of the anode was, in general, greater have reported a value of 2.767 atmospheres for the de- than would be expected on the basis of Faraday's law. composition pressure of the ammonate a t the transition Magnesium was not dissolved, and the behavior of copper was unusual in that it gave both cuprous and point. Johnson and Krumboltz (9a) have described a cupric ions. Del Boca prefers to account for the method for the direct determination of solubilities in deviations from Faraday's law on the basis of cataliquid ammonia a t -33.g0 and have reported solubility phoretic transfers, but also suggests that they might values for ten inorganic salts. These values, in terms be due to the formation of complex cations. The photolysis of pure liquid ammonia and of liquid of grams salt/100 g. NHBare summarized in the followammonia solutions of sodium, potassium, and cesium ing table: TABLE I1 has been studied by Ogg, Leightou, and Bergstrom (21). In ultra-violet light of short wave-length, no photoCi Br Acetate CN decomposition of pure ammonia was observed, while Li 0.538 Na 3.325 17.62 with the solutions of alkali metals, amide formation K 0.213 40.32 1.028 4.55 with evolution of hydrogen resulted. They found that NH4 14.75 90.75 AE 0.215 pure liquid ammonia absorbs light to a longer wavelength than does gaseous ammpnia and that, solutiqns Further, the above authors have given a comprehensive of the alkali metals show continuous absorption and critical discussion of earlier work on the deter- throughout the ultra-violet. Howard and Browne (22) have found that hydrazine mination of solubilities in liquid ammonia. is formed in small quantities by the thermal action of Liquid ammonia, with its unusual solvent properties and low dielectric constant, affords a valuable medium incandescent filaments on liquid ammonia. The maxifor the study of conductances, since in solvents of low mum yield, approximately 03.8 g./gross kw.-hr., was dielectric constant ordinary electrolytes are measur- obtained with a straight tungsten filament a t a temably undissociated. In fact, no electrolyte is known to perature of 3000°. The influence of various factors be completely dissociated in liquid ammonia solution upon the yield of hydrazine obtained has been investigated. They believe that the action of the heated (19). filament is a purely thermal one and suggest possible' Although other solvents were used in this same connection, Kraus and co-workers (10-18) have made ex- mechanisms for the formation of hydrazine under tensive use of liquid ammonia in their studies on: the these conditions. In another paper (23) they have properties of electrolytic solutions. ' The.conductance* t.eported that the yield of bydrazine varies inversely of the following electrolytes have been determined with'the pressure and that the yield is not increased a t or near the boilmg point of liquid ammonia: sodimi by the .presence of solutes such as ammonium chloride, phenyl amide, sodium andpotassium diphenyl amsdes, gelati?: etc. sodium and potassium triphenylboron ammides, w - ..:Sliafenshtein and Monoszon (24) have measured the dium triphenyl stannide,. spdium triphenyl germanide; vapor pressures of liquid ammonia solutions and have sodium triphenyl methida; dispdium tetfaphenyl distan- used the .resulting data to calculate the molecular nide, dipotassium diphenylliydrazide, sodiuni beuz- weights of the solutes. Fairly accurate values were hydrolate, mono- and disodium diphenyl ketyls, and obtained for non-associated substances. Some qualithe sodium salts of phenol, a-and B-naphthols, ethyl tative data relative to the degree of association of mercaptol, thiophenol, and' trimethyltin. Fuoss .and various solutes are given. Ipatlev, Teodorovich, and Druschina-Artemovich Kraus (11) have developed a method for evaluating the limkidg equivalent conductance and the dissocia- (25), have develooed a method for the determination of
.
the diffusion coefficients of gases in liquids under pres- in liquid ammonia a t room temperature. No evidence sure. They have found that the diffusion coefficient of the formation of the salt, KHN.NO, was found, but for hydrogen in liquid ammonia a t 25' and 30' and a t diphenyl and di-p-tolyl nitrosamines were shown to pressures varying between 35 and 100 atmospheres is react quantitatively with the metal amides in acunusuallv . hizh. It has been shown that the diiusion cordance with the eauation : coefficient is independent of pressure over the pressure R%N,NO 2MNHz -+RsNM MOH NS NHI range investigated, but that it is more dependent upon A number of interesting relationships between the temperature than is the diffusion coefficient for hydrogen in water. Values for gases in solvents, other than water and ammonia systems have been pointed out by Wood and Bergstrom (30) in connection with their liquid ammonia, have also been determined. A patent bas been issued to Booth, Torrey, and Mer- unsuccessful attempts to prepare nitrogen derivatives lub-Sobel (26) covering a method for the electro- of divalent carbon. They have reported the preparadeposition of beryllium from a liquid ammonia solution tion of potassium salts of formic acid and of ethyl formate by treatment of these substances with potasof a beryllium salt such as the iodide or chloride. sium amide in liquid ammonia. Metal salts of benFranklin (27) has pointed out that since cyanamide bears the same relationship to ammonia that carbonic zimidazole and of diphenylmethylformamidine,CeH6M:acid bears to water, compounds of the type RzNCN, CH.N(CHa)C6Hs, were prepared with solutions of where R is an organic group, would be esters of an metals in liquid ammonia a t room temperature. The ammonocarbonic acid. In support of this view, he same authors have reoorted the meoaration of a diand his students have shown that these esters are am- potassium salt of tetrazine (31),HCP N % ~ by , treat\ w-.= d -. monolyzed by potassium amide, an ammono base, in ingitwith an excess of potassium amidein liquidammonia accordance with the equation: a t -40". By treating the ammonia solution of the RzNCN 2KNHz --+ G N H KINCN f N& salt with ammonium bromide, they were able to rethis reaction being strictly analogous to the hydrolysis cover the tetrazine. It is interesting to note that this of an aquocarbonate: compound could not be recovered unchanged from solvents other than liquid ammonia. OC(0GHJr 2KOH +KzCOa 2CIHsOH Baldinger and Nieuwland (32) have oreoared the The diethyl, di-n-propyl, and di-isobutyl esters were wphenyl-derivatives of n-propionitrile, n-bu&onitrile, studied. To show that guanidine may also be thought n-valeronitrile, n-capronitrile, and n-heptonitrile by of as an ammonocarbonic acid, they have shown that treating phenylacetonitrile with sodium in liquid amits N,N', N"-triphenylester may be ammonolyzed by monia and adding the appropriate alkyl halide to the ammonium chloride, an acid in ammonia, a t ZOO0, in solution of the sodium salt. The following equation accordance with the equation: probably describes the course of the principal reaction:
-
+
+
+ +
+
+
+
+
. .
With potassium amide a t ordinary temperatures, this ester is not ammonolyzed, but it is to be remembered that it is not hydrolyzed by aqueous hydrochloric acid unless heated to 250'. Franklin and co-workers have also prepared a number of alkali. metal salts of acid guanidine esters and describe the ammonolysis of acetic acid, acetamide, diphenylacetamide, diphenylbenzamide, acetanilide, benzotrichloride, and some nitrogen esters of urea. Franklin (28) has also reported the preparation of alkali metal salts of a number of acid amides, which supports the view that these compounds may be thought of as aquo-ammono acids, that is, acids which are a t the same time derivatives of water and of ammonia. These salts were prepared by treating the acid amides with potassium amide or sodium amide in liquid ammonia. Between 250° and 300' these salts were found to undergo pyrogenetic decomposition with the formation of saturated hydrocarbons and potassium cyanate (potassium aquoammono carbonate). Fernelius and Watt (29) have attempted to prepare a salt of an aquo-ammono nitrous acid by the saponification of esters of this acid (the diarylnitrosamines) with the amides of lithium, sodium, potassium, and calcium
-
Physical constants for the products have been reported. Vaughn and Nieuwland (33, 34) have studied the iodiuation of monosubstituted acetylenes. Vmyl, ethyl, butyl, amyl, heptyl, phenyl, and tolyl acetylenes were iodimated with a liquid ammonia solution of iodine. The yield obtained and the rate of iodination were found to depend upon the nature of the substituent group. Vaughn (35) has reported that dialkiuyl mercury compounds react with iodine in liquid ammonia to give iodoacetylenes. Thus, dihexinyl mercury gives a 54y0yield of butyliodoacetylene. During the past year, a number of interesting studies have been made on the organic derivatives of elements of the thud and fourth groups of the periodic arrangement of the elements. Kraus and Toonder (36, 37) have prepared trimethylgallium ammine by treating trimethylgallium etherate with liquid ammonia and have studied its properties and reactions. They believe that the following equations probably describe the reactions of this ammine with sodium in liquid ammonia : 2(CHJrGaNHa
+ N a + [(CHs)rGallNH2Na+ '/aH2 + NHa + 2Na + [(CH&Ga].Na~ f 2NH3
2(CH&GaNHa
Trimethyl gallium may be regenerated from the disodium compound by treating it with ammonium bromide in liquid ammonia. With one equivalent of sodium, dimethyl gallium chloride is reduced to dimethyl gallium which first reads with the solvent to form a monammine, then this ammiue loses hydrogen to form an amide:
With two equivalents of sodium, the coordinately linked compound, (CH&Ga.NaNH2, is formed and it will yield dimethylgallium ammine upon treatment with ammonium bromide. Johnson and Sidwell(38) have found a method for the of very pure germanic imide, HN: G ~NH, : consists in the ammonolysis of germaniumtetraiodide, either in liquid ammonia or carbon tetrachloride, in accordance with the equation:
ammonolyzing agent than ammonia), Kraus and Eatough (41) have prepared the lithium salt, (C6H& SiCi, which when treated with trimethyltin chloride in liquid ammonia yielded the compound, (CsHs)3Si-Sn(CH& The silicon-tin bond is broken by sodium in liquid ammonia with the formation of sodium salts of the two constituent groups. With bromine, the corresponding bromides are formed. It is interesting to note that with one equivalent of lithium in ethylamine, they obtained a compound, triphenylsilicyl ethylamine, (CeH6)sSiGHsNHz. This compound is unusual in that it contams an odd number of electrons and yet is remarkably stable, as shown by the fact that it can be distilled . . at 150' in a high vacuum without decomposition. Kraus and Eatough (42) have also studied the reactions of sodium triphenylstannide with polyhalogenated methanes. With methylene chloride and with chloroform, the following reactions occurred: 2(CsHa)rSnNa 3(CIHd . . SnNa
They have also reported work on the reactions of germanium tetraiodide with primary, secondary, and tertiary ethylamines. Flood (39) has reported that ethylgermanium triiodide is completely ammonolyzed by liquid ammonia to form a nitride which hydrolyzes to give monoethylgermanic oxide:
Ethylgermanium tribromide has been shown to undergo identical reactions. After allowing triphenylgermanium to react with sodium in liquid ammonia to form the sodium salt, (C8H&GeNa, Kraus and Sherman (4Q) have prepared the n-propyl, n-butyl, n-amyl, and benzyl derivatives of triphenylgermanium by adding the corresponding alkyl or substituted alkyl halide to the liquid ammonia solution of the sodium salt. By treating the sodium salt with triethylgermanium bromide in benzene, they have prepared triphenylgermanyltriethylgermanium, (GHn)3Ge-Ge(GHs)3. This reaction could not be carried out in liquid ammonia due to the fact that triethylgermanium bromide ammonolyzesreadily in that solvent. As a result of the work of Kraus and co-workers on the organic derivatives of the elements of the fourth group, it is seen that these elements form compounds of the type (RJA)?, where R is an organic group or hydrogen, and that these compounds may.be stable in liquid ammonia solution or may dissociate to give the free group, RJA. Ordinarily, the free group may be formed by the reduction of the corresponding halide with sodium in liquid ammonia. With excess sodium, the free group combines to form the sodium salt. The behavior of derivatives of tin, germanium, and carbon has been studied, but idue to ammonolysis the corresponding derivatives of silicon could not be obtained. However, by using lithium in ethylamine (a weaker
++ CHElr 4I(C&)tSnllCH* f 2NaCI CHClr 4I(CsHJsSnI8CH+ 3NaCI ..
~
~
A somewhat more complex reaction occurred with carbon tetrachloride. Gmner (43) has found that liquid ammonia may be used to remove from minerals water which exists in the mineral in the form of [Ca(HOH)J. He has applied this method to the partial dehydration of a number of alkali aluminum silicates. It was found that water combined with Si02 as H2SiOaor H4Si04and zeolithic water in the anion of the zeolite molecule could not be removed by this method. Formulas for a number of mineral zeolites have been proposed. The metathetic reaction, NaNOI
+ HnNCOONH,
+NI~NOI,
4H I N C O O N ~
has been studied by Kameyama and co-workers (44). The reaction was effected by passing carbon dioxide into a liquid ammonia solution of sodium nitrate. The utilization efficiency of sodium in the production of sodium carbamate.was found to be 96-98% and to be independent of concentration. Zintl and Kaiser (45) have made extensive use of liquid ammonia in their work on intermetallic compounds. They have investigated the ability of metals such as germanium, tin, lead, gallium, indium, and thallium to form complex anions. Three procedures were followed. 1. The metals were used as cathodes in the electrolysis of liquid ammonia solutions of alkali halides. 2. The metals were treated with potassium amide in liquid ammonia. 3. Alkali metal alloys were extracted with liquid ammonia. Information secured through the use of the latter method showed that germanium as well as tin and lead may form complex polyanions. By passing nitric oxide into a liquid ammonia solution of sodium, Zintl and Harder (46) have prepared nitrosyl sodium, (NaNO).. DebyeScherrer photographs show that this product is not identical with sodium hyponitrite. Biltz and LeBoucher (47) have
prepared anhydrous triammonium phosphate by treat- (52) has given a brief discussion of the ammonia system, ing diammonium hydrogen phosphate with liquid and Wolthorn (53) has presented a brief discussion of ammonia under pressure. liquid ammonia as a solvent and reaction medium. Jarrey (48) has separated phenols from tat. oils by From the foregoing discussion, i t is evident that extraction with liquid ammonia. Hydrocarbons were liquid ammonia is proving to be a useful tool in the removed from the liquid ammonia extract by successive solution of a wide variety of chemical problems and washmgs with butane. Phenols thus obtained were that a considerable volume of new facts is being sufficiently pure for use in the manufacture of resins. brought to light by those workers who are providing In connection with his studies on the viscosities of additional information in support of Franklin's amvarious refrigerants, Stakelbeck (49) has measured the monia system of compounds. In those cases where a given paper forms one of a absolute viscosity of ammonia over a temperature range of from -20' to 80°, and from 1 to 25 at- series, the references to earlier work have been omitted, mospheres pressure, by the Lawaczek dropping method. since they are invariably to be found in the paper under A number of papers of a more general nature have ap- consideration. Any suggestions or criticisms which peared during the past year. An interesting account will result in the improvement of a review of another of the beginning and early development of liquid am- year's work are welcomed. monia work in this country has been given by Taft The author takes this opportunity to acknowl(50). Much of interest to those interested in ammonia edge his indebtedness to Dr. W. C. Fernelius of this chemistry is to be found in the review of the chemistry University and to Mr. W. R. Stemen of the editorial of the alkali amides by Bergstrom and Fernelius (51). staff of Chemical Abstracts, for their many helpful In considering solvo-systems of compounds, Audrieth suggestions during the preparation of this paper. LITERATURE CITED
( 1 ) R. WIEBEAND T. H. TREMEARNE, J. Am. Chem. Soc., 55, 975-8 (Mar., 1933). AND W. P. TEODOROYICH, J. Gen. Chem. ( 2 ) W. W. IPAT'EV (U. S. S. R.), 2,305-10 (1932). ( 3 ) W. C. JOHNSON AND M. M. PISKUR, I.Phys. Chem.. 37, 93-9(Jan., 1933). AND L. BONCYK, 1.Am. Chem. Sac., 55,352&30 ( 4 ) H . HUNT (Sept., 1933). 2. anwg. Chem., 70, 6 1 4 (Feb., ( 5 ) 0. R m a AND L. HECHT, 1011\
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( 8 ) G. PATSC~KR, ibid., Al63,340-50 (Mar., 1933). ( 9 ) S. ABEAND R. HARA.I.Sac. C h . Ind. (Japan).36, Suppl. binding, 557-9 (Oct., 1933). (9,) W . C. JOHNSON AND 0. F. K R ~ O L 2. T physik. ~, Chem., A167.249-59 (Dec., 1933). (10) C. A. KRAUS AND R . M. Fuoss. J. Am. C h . Soc.. 55,2146 (Jan., 1933). (11) R. M . Fuoss AND C. A. Rnnns, ibid.. 55, 476-88 (Feb., 10x?\
(12) R . M. Fuoss AND C. A. KRAWS, ibid.. 55, 1019-27 (Mar.. 1933). (13) R. M . moss AND C. A. KRAUS,ibid., 55, 2387-99 (June,
A"-",. 1(177\
(14) C. A. Knnns AND W. W. HAWES, ibid., 55, 2776-85 (July. 1933). (15) C. A. KRAUSAND W. H. KARLEX,ibid., 55,353742 (Sept.,
.""",. 1077\
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-- -.
7 .41-1
(35) T. H. VAUGHN, J. Am. Chem. Sac., 55,3457 (Aug:, 1933). Proc. Nat. Acad. Sci., (36) C. A. KRAUSAND F. E. TOONDER, 19,292-8 (Mar., 1933). J. Am. C h . Soc., 55, (37) C. A. KUUS AND F. E. TOONDER, 3547-54 (Sept., 1933). (38) W,-C. JOHNSON AND A. E. SIDWELL. {bid.. 55, 1884-9 (May, 1Y33). (39) E. A. FLOOD, ibid., 55,493Fr8 (Dec., 1933). (40) C. A. KRAUSAND C. S. S ~ R M Aibid., N , 55,4694-7 (Nav., 1933). H , 55, 50084 (Dec., (41) C.-?i,-Guus AND H. E A ~ U Gibid., lY33). (42) C. A. KRAuS AND H. EAmucH, ibid., 55, 5014-6 (1933). (43) E. GRWNER.2. enorg. allgem. C h . . 211, 38547 (Apr..
(16) C. A. a n u s AND E. G. JOHNSON, ibid., 55,3542-7 (Sept., 1933). (17) C. A. KBnns AND P. B. BIBN,ibid., 55, 3609-14 (Sept., 1933). (18) R . M . Fuoss AND C. A. KRAUS, ibid., 55, 3614-20 (Sept., 1933). ibid., 55,442240 (Nov.. 1933). (19) W. W. HAWES, '"LL., "00,. 20) M. C. DELBOCA, Heh. C h m . Aclo, 16,565-71 (May, 1933). (46) E. ZI&L AND A. HARDER. Ber.. 66B. 760-1 (May. 1933). AND F. W. B E R G S ~ O Y ,(47) W. BILTZ 21) R. A. OGG,JR.. P. A. LEIGHTON, AND L. LEBOUCHBR, Anales sac. esf& fis. quim., I.Am. Chem. Sac., 55,1754-66 (May. 1933). 31,427-33 (June, 1933). (22) D. H. HOWARD, JR. AND A. W. BROWNE, ibid., 55,1968-74 48) R. JARREY, Compt. rend.. 196,16754 (May, 1933). 2. gm. Kiillc-Ind., 40,3340 (1933). (May, 1933). 49) H. STAKELBECK. (50) R. TAET,J . CEEM.EDUC., (23) D. H . HOWARD, JR.AND A. W. BROWNE, ibid.. 55, 3211-4 10.34-9 (Jan.. 1933). (Aug., 1933). (51) F. W. BERGSTROM AND W. C. FERNELNS. Chem. RN., 12, (24) A. I. SHATENSHTEIN AND A. M. MONOSZON, 2. physik. 43-179 (Feb.. 1933). Chem., A165,147-53 (June,1933). (52) L. F. AWDRIETH. Z. physik. Chem.. A165,324-7 (July, 1933). W. P. TEODOROMCA, AND S. I. DRUSCHINA-(53) H. WOLTHORN, (25) W. W. IPAT'EY, S C ~Sci. . Math., 33,288-92 (Mar.. 1933).
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