April, 1956
NOTES
plasma albuming in 0.2 M NaC1. The change in si5 is possibly due to a decrease in the hydration of the polyelectrolyte at higher temperatures. With salt present, hydration would be less at all temperatures, and a difference in temperature might not produce a perceptible effect.
The Acidities of Chloramine and Dichloramine. -For the following discussion, it will be convenient to define an aquo acid molecule m a water molecule in which one of the protons has been replaced by another group, R. The ionization of an aquo acid is thus represented by
507
HOR = H +
THE THERMODYNAMIC PROPERTIES OF CHLORAMINE, DICHLORAMINE AND NITROGEN TRICHLORIDE BY WILLIAM L. JOLLY Department of Chemistry, University of California, Berkeley, CaZ. Received November 10. 1066
I n both aqueous solutions’ and in liquid ammonia solutions,2 chloramine reacts with ammonia to form hydrazine. Therefore the thermodynamic properties of chloramine are of importance in discussing equilibria involved in the synthesis of hydrazine. I n this note, the free energies of aqueous chloramine and related species are estimated from literature data,. It is emphasized that some of the results are only rough approximations and are to be considered tentative until better experimental data are obtained.
+ OR-
(1)
The corresponding ammono acid molecule is an ammonia molecule in which one of the protons has been replaced by the group R HzNR = H f
+ HNR-
(2)
The pK values for several ammono acids and the corresponding aquo acids are listed in Table I. It appears that most of the aquo acids tabulated are roughly lo7 times stronger than the corresponding ammono acids. Chloramine, NH2CI, may be looked upon as the ammono analog of hypochlorous acid, HOCI. Since the pK value for hypochlorous acid is about 73, one may estimate pK = 14 2 for chloramine. Hence in very alkaline solutions one should expect an appreciable concentration of chloramide ion, NHCI-. The acidity of dichloramine, NHC12, may be estimated from Table I and the acidities of several
*
TABLE I pK VALUESFOR AMMONO AND AQUOACIDS Ammono acid
PK
Ammonia, HNHna -33 Acetamide, CHsCONH24 15 14.5 Benzamide, C~HICON Hz4 o-Nitroaniline, ( N O Z ) C ~ H ~ N H ~ 14 p-Nitroaniline, ( N02)C6H4NHz6 12 Sulfamide, HZNSOZNHZ~ -11 Cyanamide, HzNCN* 10.5 Benzenesulfonamide, C ~ H ~ S O Z N H Z ~ -10 Sulfanilamide, ( NHz)CBH~OZNH+ 10 Nitramide, HzNNOZ’ 7
Aquo acid
PK
Water, HOHa Acetic acid, CHaCOOHa Benzoic acid, C~H6COOH10 +Nitrophenol, (NOz)c6H4OH6 p-Nitrophenol, (N02)CeH40H6 Sulfamic acid, H2NSOZ0H1l Cyanic acid, HOCNla Benzenesulfonic acid, C6HsSOzOH13 Sulfanilic acid, (NHz)C6H4S0z0H14 Nitric acid, HONOz
The methods of calculation are straightforward except for the estimation of the acid ionization constants of chloramine and dichloramine. A method for estimating the acidity of an ammono acid from the known acidity of the corresponding aquo acid (or vice versa) is outlined. (1) L. F. Audrieth and B. A. Ogg. “The Chemistry of Hydrazine,” John Wiley and Sons, Inc., New York. N. Y., 1951. (2) H. H. Bider, F. T. Neth and F. R. Hurley, J . A m . Chem. Soc., 16, 3909 (1954). (3) G. E. K. Branch and M. Calvin, “The Theory of Organic Chemistry,” Prentice-Hall Inc., New York, N. Y., 1941. (4) G. E. K. Branch and J. 0. Clayton, J . A m . Chem. SOC.,60, 1680 (1928).
(5) A. I. Schattenstein, Acta Phueicochim. U.S.S.R., 10, 121 (1939). ( 6 ) Personal communication from Professor L. F. Audrieth. (7) Determined by the author by pH titration. (8) A. Albert and R. Goldacre, Nature, 149, 245 (1942). (9) Branch and Calvin8 have estimated A p K = 6 for the reaonance energy of the carboxylate ion. (10) W. M. Latimer, “Oxidation Potentials.” 2nd ed., PrenticeHall, Inc., New York. N. Y., 1952. (11) Estimated from the data of M. E. Cupery, I n d . Eno. Chem.. 80, 627 (1938). (12) “Iuternational Critical Tablea,“ Vol. VI, McGraw-Hill Book Co., Inc.. New York, N . Y., 1928. (13) R. S. Airs and M. P. Balfe, T r a m . Faraday SOC., 89, 102 (1 943). (14) W. Carr and W. J. Bchutt, ibid., 86, 579 (1939).
16 5 (5 4 (4 7 7 1 4
