1084
Vol. 65
NOTES
tributions to the design and fabrication of the samplers and Donald L. Guernsey for his careful analytical work.
SOME PHYSICAL PROPERTIES OF AQUEOUS l>ICOLINIC ACID SOLUTIONS
2.5 re
In q/qo = -1-
QFc
where P = 0.0917 1. mole-' is an "effective molal volume" and Q = 0.873 is an interaction parameter. (4) V. Vand, J . Phya. Chem., 6 2 , 277 (1948). (5) A. Einstein, Ann. Phys., 19, 289 (1906).
BY R. A. ROBINSON' AND R. W. GREEN University of New England, Armidale, N . S . W . , Australia and University of Sydney, N.S. W., Australia Received January dS, 1961
We have measured the density, refractive index and viscosity of some aqueous picolinic acid solutions a t 25'. Picolinic acid was purified by sublimation a t low pressure. Vacuum corrections were applied to the weights in making up the solutions. Densities were measured by the usual pyknometer method, refractive indices with an Abbe instrument and viscosities in an Ostwald viscometer. Kinetic energy corrections were made to the viscosity data. The results are given in Table I.
THE ISOTOPE EXCHANGE REACTION BETWEEN LABELLED NITRIC OXIDE, 16N0 AND NITROSYL CHLORIDE, 1 4 ~ 0 ~ 1 1 BY LESTER P. KUHNAND CHESTER BUTKIEWICZ Ballistic Research Laboratories, Aberdeen Promng Ground, M d . Received January 81, 1061
In a previous paper2 it was shown that a bimolecular exchange reaction occurs between 14N0 and RO15K0which is first order in nitric oxide and first order in nitrite ester. The reaction between nitric oxide and nitrate esters to yield nitrogen dioxide and nitrite ester, on the other hand, was found to be first order in nitrate ester but indeTABLEI pendent of nitric oxide pressure. It was conDENSITY,REFRACTIVE INDEXAND VISCOSITY OF AQUEOUS cluded that the isotope exchange reaction between PICOLINIC ACIDSOLUTIONS AT 25'" nitric oxide and nitrite esters occurs through the d, formation of an unstable intermediate which % m C z g. mL-1 n s/so 5.093 8.711 9.730 10.77 15.34 20.41 25.01 28.96 35.69 40.69 47.71 a x
index.
0.4359 .7751 .8755 .9804 1.472 2.083 2.709 3.311 4.508 5.573 7.411
0.4193 0.00779 1.01359 1.3431 1.103 .7258 ,01377 1.02570 1.3520 1.193 .8135 ,01553 1.02924 1.3538 1.214 ,9035 .01736 1.03273 1.3561 1.248 ,02583 1.04835 1.3669 1.394 1.306 1.768 .03617 1.06634 1.3788 1.604 .04653 1.08334 1.3889 1.829 2.201 2.582 ,05629 1.09769 1.3997 2.109 3.256 .07511 1.12329 1.4165 2.744 3.777 .09124 1.14277 1.4297 3.459 .I178 1.17092 1.4491 5.127 4.538
= mole fraction of picolinic acid; n = refractive
The partial molal volume at infinite dilution is 83.8 ml. mole-'; that of pyridine is 77.5 ml. mole-' calculated from the density of its aqueous solutions* or 80.9 ml. mole-' from the density of the pure liquid. The Traube increment for the carboxylic group (15.8 ml. mole-') added to the molal volume of pyridine therefore gives a t least 93 ml. mole-' for picolinic acid. The refractive index measurements probably were accurate within =k0.0002: the refractivity a t infinite dilu.tion therefore cannot be calculated with accuracy but it is approximately 31.9 ml. mole-'. That of pyridine is 23.97 ml. mole-' and the usual increment for the carboxyl group predicts 30.1 for picolinic acid. Thus the observed molal volume is low ttnd the refractivity high. This may well be due to the fact that picolinic acid exists mainly in the zwitterion form8 with increased resonance. The high viacosities are notable: a 48% solution is about as viscous as a 38% sucrose solution a t the same temperature. The relative viscosities can be represented with an average deviation of 0.005 by Vand's modification* of the Einstein equation6 (1) Correspondence to R. A. Robinson, National Bureau of Standards, Washington 2d, D. C. (2) "International Critiaal Tables," Vol. 111, McGraw-Hill Book Co.,Ino., New Yorlr, N. Y., 1928,p. 112. (3) R. W. Green trnd H. K. Tong, J . Am. C h m . Sac., 78,4896:(1956).
RO-'6N=O 1
I4
~
~
0
rearranges to give the exchanged products and does not occur by way of a displacement on alkyl oxygen. The object of the present work was to make a similar study of the reaction between labelled nitric oxide and nitrosyl chloride. The reaction was run in an infrared cell having CaFz windows. The reaction mixture was analyzed for 14NOCland WOC1 by means of infrared spectroscopy. The respective bands in the region of 5.5 I.C are nicely resolved by the CaFz prism as shown in Fig. 1. It was found that a t 25' with initial pressures of 14NOClof 5 to 10 mm., and a t pressures of 'WO such that the ratio KO/NOCl varied from 0.3 to 3, equilibration was complete in three minutes, which was the time required to get an infrared spectrum after the gases were mixed in the cell. The half-life of this reaction under our experimental conditions must therefore be of the order of seconds or less. Because of the speed of the reaction it was not possible to obtain kinetic data or to determine the order of the reaction with the technique we employed. The thermal decomposition of nitrosyl chloride has been studied thoroughly. At temperatures below 250' the reaction is second order in NOCl and the rate constant3 is exp(-22000/RT) mole-' cc. set.-'. The half-life of the bimolecular decomposition under our experimental conditions according to this rate constant would be about 1O1O seconds. At 300' the decomposition proceeds by a two-step mechanism4 (1) This work was supported by the Office of Ordnance Research, Durham, N. C. (2) L. Kuhn and H. Gunthard, Helu. Chim. Acta, I S , 607 (1960). (3) I. Welinsky and H. A. Taylor, J . Chrm. Phue.. 6,466 (1938). (4) P. G. Aahmore and J. Chanmugan, Trona. Faraday SOC.,49, 270 (1953).
NOTES
June, 1961 C1
+
1085
NOCl +NO C1 NOCl +NO Clz
+
+
The rate constants for the two steps are 10l2 exp(-38000/RT) set.-' and 1013 exp(-llOO/RT) mole-’ cc. sec. -I, respectively. According to this mechanism the half-life for the decomposition of NOCl under our conditions would be about 10I6 seconds. It is clear that the rapid isotope exchange reaction which we have observed cannot take place by a rate-determining decomposition of l4NOCI followed by a rapid reaction of 15N0with C1 or Clz, since the rate of decomposition of KOCl by either the bimolecular or unimolecular is much too slow. It seems highly likely that the isotope exchange reaction involves the attack of W O upon l4NOCI. We can picture this reaction as taking place in either of two ways: The first possibility is a displacement reaction on chlorine in which the transition state may be pictured as 0 “N--Cl-
14N0
0.8
0.6
0.2
The second possibility is the formation of a shortlived intermediate which rearranges to give the O~~N-CI
I
OW
exchanged products. The new bond that is formed in the intermediate involves the unshared pair of electrons of the nitrogen of NOCl being donated to the NO. If the first possibility is correct then one should find a similar bimolecular reaction between NO and NOG1 (first-order in NO and firstorder in N02C1) to give NOz and NOC1. If, on the other hand, the second possibility is correct one would not expect to find a reaction between NO and NO&l whose rate was dependent upon the pressure of KO since NOzCl does not have an unshared pair of electrons and hence is incapable of forming the postulated intermediate. A rapid reaction6 does in fact, occur between NO and NOzCl to give NO2 and NOCl which is first order in both NO and NO2CI, whose rate constant is 0.8 X 1Ol2 exp(-6900/RT) mole-l cc. sec.-l. This reaction has a half-life of several seconds under our conditions of temperature and pressure. Thus we conclude that the isotope exchange reaction between 15N0 and 14NOCl and the reaction between NO and N02C1 proceed by a displacement on chlorine and not via the formation of the intermediate pictured above. It is interesting to note that whereas this reaction takes place via a displacement on chlorine, the analogous reaction involving nitrite esters does not take place via a displacement on oxygen. Thus we have another example of the frequently observed phenomenon that free radical reactions occur readily by displacement a t a univalent atom (halogen or hydrogen) but not a t a polyvalent atom. Experimental was made in the manner previously described.* l4NOCI was purchased from the Matheson Com any and purified by distillation.B The infrared spectrum o?the puri16NO
( 5 ) E. Freiling, H. Johnston and R. Ogg, J . Chem. Phys., 8 0 , 327
(1952).
( 6 ) “Inorganic Syntheses,” McGraw Hill Book Co., New York, N. Y.. 1939, p . 55.
1
5.6 5.7 1,P. Fig. 1.-(1) 14NOCl; (2) 1sNOC1, pressure 10 mm., cell length 20 mm. fied material showed that it was essentially free of the two major impurities, HC1 and NOz. The mea were introduced into a 5-cm. infrared cell with Calf windows at the desired pressure using a conventional vacuum line. The mercury of the manometer was rotected from the NOCl by a layer of perfluorokerosene. &he spectra were obtained with a Perkin-Elmer Model 21 Spectrometer equipped with a CaF2prism. 5.5
ACIDITY COXSTANT OF A PROTEIN CONJUGATE I N DzO BY W.-Y. WEN AND I. M. KLOTZ Department of Chemistry, Northwestern Univermty, Evanslon, Illinois Received October 98, 1960
The relative strengths in solution of hydrogen bonds involving protons and deuterium atoms, respectively, have been the subject of numerous investigations. 1-6 For isolated molecules in the gaseous state, heats of dimerization’ show clearly that bonding by deuterium is stronger, but only slightly so. For solutes in solution, however, there are competing effects between substituents of the solute and the solvent, and the net effect on physicochemical behavior of the solute is not easy to predict. We have r e ~ e n t l y ~interpreted .~ the shifts in (1) S. Korman and V. K. LaMer,
J. A m . Chem. SOC., 68, 1396
(1936). (2) F. C. Nachod, 2. physik. Chem., A182, 193 (1938). (3) A. H. Cockett and A. Ferguson, Phil. Mag., 88, 693 (1939). (4) F. A. Long and D. Watson, J. Chem. Soc., 2019 (5) M. Calvin, J. Hermans. Jr., and H. A. Scheraga, J. A m . Chem. Soc., 81, 5048 (6) G. Dahlgren, Jr., and F. A. Long, ibid., 88, 1303 (1960). (7) A. E. Potter, Jr., P. Bender and H. L. Ritter, J. Phys. Chem., 69. 250 (1955). (8) I. M. Klotz, Science, 198, 815 (9) I. M. Klotr and H. A. Fiess, Biochim. st Biophya. Acto, 8 8 , 57
(1958).
(1959).
(1958).
(1960).
