508
VOl. ti0
NOTES
cipal species. Metcalf and his eo-workers, studying more dilute solutions, found slightly different pH boundaries, but they are essentially in agreement with Chapin. From the experimental data of these investigators, we estimate K = lo6 for the reaction H+
+ PNH2Cl
+ NHi+
(3)
+ 3NHC12 = 2NC13 + NH,+
(4)
NHClz
c1-
I c1-
c1-
Free Energy Calculations.-Metcalf and his coworkersls determined the equilibrium constant for the following reaction a t 15” NHzCl
+ OHK16
= NH,
= 1.6 X
+ OC1-
lo-*
Clz
-1.60
- 1.36
Clr
- 1.42
Cl2
- 1.37
NHClz
I
- 1.39 - 1.36
NHzCl
I
- 1.48
I
and K = lo4for H+
-1.36
- 1.38
NC13
I
J
The C1--NH2C1 potential for alkaline solution (1 M OH- and 1M NH,) is -0.81 v., and the C1-NHC1- potential is approximately the same. Oxidation Potentials for Liquid Ammonia Solutions.-Using methods of calculation outlined elsewhere,20one obtains the potential diagram for acid liquid ammonia solutions.
By estimating the entropy of NHzCl as 30 e.u. and by using the knownl0 entropy arid free energy values for hydroxide, ammonia and hypochlorite, one calculates, for 25”, 18.6 kcal./mole for the free energy of formation of aqueous chloramine. Using -1.9 - 1.0 this datum and the equilibrium constants for Cl -(am) cm N€IzCl(am) reactions (3) and (4), one calculates 48 and 79 - 1.4 kcal./mole for the free energies of formation, respectively, of aqueous dichloramine and aqueous nitrogen trichloride. This potential diagram is significant in the anodic It is interesting to compare the above value for oxidation of chloride in liquid ammonia. The the free energy of formation of nitrogen trichloride NH4+-Nz potential is about zero volts and so with that calculated from the known thermal data. nitrogen evolution is the thermodynamically The heat of formation of NCl3 (in CCIS is 55 ked./ favored anode reaction. However, a t an anode m01e.l~ Estimating the entropy as 45 e.u., we with a high nitrogen overvoltage, it might be calculate 72 kcal./mole for the free energy of possible to form chloramine (and therefore hyformation in CC14. Nitrogen trichloride is very drazine). much more soluble in CCll than in water.18 In Acknowledgments.-The present work was carorder that the free energy of formation of NC13 ried out as one phase of a study of the synthesis of be 79 kcal./mole, we must assume a partition hydrazine sponsored by the Office of Ordnance coefficient of about lo6, which is not an unreason- Research, Contract DA-11-022-ORD-828. The able value. author wishes to thank Professor L. F. Audrieth Oxidation Potentials for Aqueous Solutions.of the University of Illinois for his aid and enThere are two interesting points which arise when couragement . considering the oxidation potentials of chloramine (20) W. L. Jolly, THIS JOURNAL, 68, 250 (1954). and its derivatives: First, one observes from equations (3) and (4) that not only the pH, but also the ammonium ion (or ammonia) concentration is important in determining whether NH2C1, NHCl2 T H E QUESTION O F A PHASE TRANSITION I N SILICON or NC13 is the important species in solution. (It is unlikely that either NH3Cl+ or NHC1- would BY ELIZABETH A. WOOD ever be encountered as the principal chlorine Be22 Telephone Laboratories, Inc., Murrau Hill,N. J . species in aqueous solution; these exist only in Received November 4, 1966 strongly acid19 or strongly basic solutions, respectively.) Second, there can be ambiguity as to the The existence of a non-cubic (probably hexagooxidation states of nitrogen and chlorine in chlor- nal) form of silicon was reported by Heyd, Khol amine and its derivatives. It seems plausible to and Kochanovsk4 in 1947.’ A report of this form the author t o assign chlorine to the +1 oxidation appears in the Structure Reports2 and in Wyckoff’s state and nitrogen to the - 3 oxidation state in these “Crystal Structures.”3 The original report was of compounds, The potential diagrams for solutions a preliminary nature, but no subsequent report was 1 mold in H+ and 1 molal in NH4+ follow published. It therefore seemed desirable to determine whether the results could be repeated. (16) R. E. Corbett. W. S. Metcalf and F. G. Soper, J . Chcm. Soc.. 1927 (1953). Heyd, et al., looked for a non-cubic form of sili(17) F. R. Bichowsky and F. D. Rossini, “The Thermochemistry con because of the existence of the non-cubic forms of the Chemical Substances,” Reinhold Publ. Corp., New York, N. Y.,
1
1936. (18) “Gmelins Handbuch der Anorganischen Chemie,” “Chlor,” Nr. 6, Verlag Chemie G.M.B.H., Berlin, 1927, p. 414. (19) The basic dissociation conatant of chloramine haa been estimated as 10-u by I. Weil and J. C. Morris, J . Am. Chem. Sac., ‘71, 8123 (1949).
I
(1) F. Heyd, F. Khol and A. Kochanovsk4, Collection CaeChOJlOV. Chem. Communa., 13, 502 (1947). (2) Wilson, Barrett, Bijvoet and Robertson. “Structure Reporls,” Vol. 11, N.V.A. Ooathoek‘e Uitgevera. Utrecht, 1951.
(3) R. W. G. Wyckoff, “Crystal Struatures,” Vol. I, Interscience Publishere, New York, N. Y., 1951.
NOTES
April, 1956 of carbon and tin. They tried to produce it in three different kinds of experiments: the first involved the addition of fluorides; the second, vacuum-evaporated silicon. The results of each of these were analyzed by X-ray diffraction a t room temperature. I n the third, Debye photographs were taken of pure silicon a t 700". The first two experiments were, in the opinion of this writer, open to alternative explanation. The following attempts were made to repeat the last experiment. Experiments The first attempt to repeat this experiment was performed with powdered silicon ("Hyper pure" from du Pont) in a fused silica capillary at 700". The Debye-Scherrer hotograph, taken in a camera designed by W. L. Bond, ofthese laboratories, showed none of the lines listed by Heyd, Khol and KochanovsU. A second attempt to repeat the experiment was made with silicon furnished by H. C. Theuerer of these laboratories which was prepared by decomposition of Sic14 with hydrogen. The resulting silicon was deposited from the vapor state onto a tantalum tape. A rod of this polycrystalline silicon was cut and fitted to the specimen holder of a commercially available high-temperature camera built by Central Research Laboratories, Inc., Red Wing, Minnesota. Debye-Scherrer photographs of this sample were taken at room temperature and a t 700, 800 and 900". No extra lines appeared at the higher temperatures, although the film was over-exposed to the point of halation. Additional evidence for the non-existence of the hightemperature form was obtained by P. D. Garn of these laboratories who ran a differential thermal analysis of the "hyper-pure" silicon and found no evidence of a phase transition between room temperature and 1O0Oo. This experiment waa repeated three times.
Conclusions X-Ray diffraction photographs taken a t 700,800 and 900" coupled with differential thermal analysis indicate no phase transition in pure silicon between room temperature and 1000". This evidence is contrary to the report by Heyd, Khol and Kochanovsk6 that a non-cubic modification of silicon exists a t 700". Acknowledgments.-The writer wishes to thank S. Geller for fruitful discussion, F. Barbieri for skillfully and carefully preparing the crystalline rod for the second experiment and V. Bala for taking the high temperature photographs. THE STABILITY OF METAL CHELATES OF SUBSTITUTED ANTHRANILIC ACIDS BY WILLIAM F. HARRIS^
AND
THOMAS R. SWEET
Department of Chsmietry, the Ohio State Uniuersitg, Co2umbu8, Ohio Received October $9, 1966
The effect of a number of substituents on the chelating properties of anthranilic acid was studied. This was done by determination of the apparent formation constants of copper and cadmium with 3 methylanthranilic acid, N - methylanthranilic acid, anthranilic acid, 5-sulfoanthranilic acid, Nphenylanthranilic acid and 3,5-diiodoanthranilic acid in 50% dioxane solutions. The Bjerrum2
-
(1) Abstracted from the doctoral dissertation of W. F. Harris presented to the Graduate School of the Ohio State University, August, 1955. (2) J. Bjerrum, "Metal Amine Formation in Aqueous Solution," P.Haase and Son,Copenhagen, 1941.
509
5
/
O
4
> Y 4 0 _I
2 PKa.
Fig. 1.-Correlation
of log ICav. with pKa: 0,Cu chelabe; , Cd chelates.
titration method as modified by Calvin and Wilson' was adapted to the present work. Experimental Materials .-The metal ion solutions were prepared and standardized by the method described in an earlier publication.' The dioxane waa urified by the method EU gested by Calvin and Wilson.8 &he anthranilic acid was oftained from Coleman and Bell Chemical Go. and was prepared for use by crystallizing several times from 50% acetic acid. The N-methyl, N-phenyl, and 3,5-diiodo derivatives were Esstman Kodak Go. white label reagents. The 3-methylanthranilic acid was obtained from Dr. H. Shechter of Ohio State University. It was prepared by reaction of 3-methylphthalic anhydride and hydrazoic acid in sulfuric acid.' 5-Sulfoanthranilic acid was prepared as described in an earlier publication.4 Procedure.-The weighed rea ent waa added as a solid to 50 ml. of purified dioxane. d a t e r , nitric acid and metal ion were added in this order. The final volume was 100 ml. Nitric acid was not used with the 5-sulfoanthranilic acid. The titration procedure was the same as that described previously.' Calculations.-The constants for the 5-sulfoanthranilic acid complexes were calculated by using the equations that the authors previously derived.'
TABLEI pKa VALUESOF ANTHRANILIC ACID AND SUBSTITUTED ANTHRANILIC ACIDSIN 50% DIOXANE Reagent
3,5-Diiodoanthranilic acid N-Phenylanthranilic acid 5-Sulfoanthranilic acid Anthranilic acid N-Methylanthranilic acid 3-Methylanthranilic acid
P K ~
5.59 6.09 6.24 6.53 6.58 6.64
(3) M. Calvin and K. W. Wilson, J . Am. Chum. Soc., 61, 2003 (1945). (4) W. F. Harris and T. R. Sweet, ibid., 77, 2893 (1955). (5) H. Barkemeyer. Master's Thesis, The Ohio State University, 1952.
