NOTES
Dec., 1958 SOLVENT EFFECTS I N THE NUCLEAR MAGNETIC RESONANCE SPECTRA OF FLUORO-ORGANICS BY R.E. GLICKAND S.J. EHRENSON Contribution jrom the Department of Chemistry, Whitmore Laboratory, The Pennsuluania State Universitv, University Park, Pennsyluania. Received June 1968
The effect of solvent on the position of a fluorine high resolution nuclear magnetic resonance line has been reported2 recently to deviate substantially from that predicted by the magnetostatic model.a In this connection we have examined the behavior of the nuclear magnetic resonance frequency of four fluorine containing materials each extrapolated to infinite dilution in approximately forty solvents. The measurements were made on a Varian High Resolution nuclear magnetic resonance spectrometer a t 40 mc. with spectral shifts obtained by the usual frequency measuring techniques. *J The results for one of these materials, 1,2-dibromotetrafluoroethane (Fl), for a typical series of solvents are recorded in Table I, column 1. Column 2 contains a diamagnetic susceptibility correction as determined from the behavior of a proton containing material in the same solvent. (The latter, therefore, contains the experimental defect in Column 3 is the sum of columns 1 and 2 and is thus the deviation of F1 from equation 1. Another compound examined was benzotrifluoride in order t o compare the solvent behavior when both hydrogen and fluorine were present in the same molecule. I n this compound the hydrogen resonance frequency varied, essentially, according to equation 1 , 6 while the fluorine shifts paralleled thoseQf F1. The deviations, AH", listed in column 3 are found to be proportional to the molecular polarization ( P ) of the solvent molecule,6 as listed in column 4, according to equation 2 A H + AH' = AH" = -3.8P 62 (2) Thus, the experimental value of AH for Fl in the gas phase as compared to that in infinite dilution in carbon tetrachloride is found to be 167 f 5 cycles while the value determined from equation 2 is 176 f 8 cycles. A similar extrapolation through the
+
(1) This work was supported in part by the Office of Naval Research, Project NR055-328. Reproduction in whole or in part is permitted for any purpose of the United States Government. (2) D. F. Evans, Proc. Chem. SOC.,115 (1958). (3) The magnetostatic model is given by equation 1.
H i
=
siHo(1
- LYK)
(1)
For ( l ) , Ha is a reference resonant field (for a proton, Ho is the resonant field for the bare nucleus i n vacuo), si is a specific shielding factor and is a function of the electron environment of the ith nucleus, (I is a shape factor, x is the magnetic susceptibility of the medium, and H i is the observed field for the ith nucleus. Using the Lorena-Lorenta approximation for the molecular cavity and a cylindrical sample cell oriented transversely to the magnetic field, (I has the theoretical value of - 2 ~ / 3 (-2.09). The variations of proton resonance frequency with solvent have been found to follow equation 1 if a i s given the value of -2.60.' (4) A. A. Bothner-By and R. E. Glick, J . Chem. Phys., 26, 1647 (1957). (5) A. A. Bothner-By and R. E. Glick, ibid., 1651 (1957). (8) Y . K. Syrkin and M. E. Dyatkina, "Structure of Molecules and the Chemical Bond," Interscience Publishers, Inc., New York, N. Y., 1950, p. 201.
1599
three tetrahalo-methanes listed in Table I was made for carbon tetrafluoride. I n this instance AH for the gas phase and that for infinite dilution in carbon tetrachloride had a theoretical value of 369 f 15 cycles while that determined experimentally by Evans2was 369 f 3 cycles and in this study was 370 f 3 cycles. TABLE I F1 ext. to m dilution in
1
2
3
AHa,b cycles
AH',C cycles
AH"-d cycles
CHzClz CHC13
+ +
+
4 Polariaability of solvent (P) cc.6
0 2.1 2.1 16.4 -25.0 4.2 -20.8 21.3 cc4 -32.0 - 1.2 -33.2 26.1 CHzCIBr -27.0 +12.3 -14.7 19.3 CHzBrz -58.0 +22.3 -35.7 22.1 CHClzBr -45.0 +12.9 -32.1 24.1 CHClBrz -67.0 +21.3 -45.7 27.0 CHBrs -77.0 +27.1 -49.9 29.8 -48.0 6.2 CC13Br -41.8 28.9 CClzBrz -68.0 $12.9 -55.6 31.8 A H values are given in frequency equivalents referred to a fixed fre uency of 40 mc. b A H is the frequency shift from that oypure F1. AH'is the frequency shift for a pure proton containing material of the same diamagnetic susceptibility as F1 to infinite dilution in the given solvent. AH"is the deviation of F1 from equation 1. Calculated from ref. 6.
+
Since solvent polarization (including self-polarization in the pure liquid) contributes substantially to the fluorine resonant frequency, intermolecular influences may be evaluated exactly by an extrapolation to zero polarization. This procedure contrasts markedly with that for standard hydrogen chemical shift assignments where measurements in magnetically isotropic media such as carbon tetrachloride are ~ u i t a b l e . ~Nevertheless, proton resonance, as evidenced by the deviation noted in a, equation 1, may be influenced by solvent polarization. These results are presently being extended, both theoretically and experimentally, to include intramolecular chemical shifts as well as solution 'behavior in general. We wish to thank Professor L. A. Currie for assistance in preparing the gas eamples. DISPLACEMENT REACTIONS AT THE SULFUR ATOM. 11.' THE REACTION OF CYANIDE WITH TETRATHIONATE BYROBERT EARLDAVIS~ Converse MemoricE Laboratory of HarvGrd Universitu. Cambridge, M a s s Received June 18, 1968
Cyanide ion reacts quantitatively with tetrathionate in aqueous solution producing sulfate, thiosulfate and thiocyanate . CN-
+
5406-2
k2 + HzO + + SzO3-2 + SCN- + 2H+ Sod-2
(1)
Indeed all of the higher polythionates react quantitatively and the reaction has been used as an (1) Part I. R. E. Davis, J . A m . Chem. Soc., 80, 3565 (1958). (2) Public Health Service Research Fellow of the National Cancer Institute. Massachusetts Institute of Technology, Cambridge, Mass.
NOTES
1600
0.20t 0.30 I
0.10 4
8
I4
0
-0.10
-0.20
1
i.
t 0
0.4
0.2 p'/%
Fig. 1.-Salt effects at 25" in water: open circles, reactants only; full circles, with added sodium sulfate; vertical bar with added potassium thiocyanate; horivontal bar, with added potassium chloride while the reactants are at pl/a = 0.05.
analytical method, measuring the thiocyanate produced as the ferric c ~ m p l e x . ~Cyanide ion reacts further with thiosulfate3 forming thiocyanate and sulfite but this reaction4 is 1/104 as fast as reaction 1. The rate of reaction 1 has been briefly investigated in water. Ishikawa6 rapidly quenched the reaction in dilute sulfuric acid and titrated with iodine. The value of kz obtained a t 0" in 0.01 M sodium hydroxide was 0.205 1. mole-' see.-'. Forestie studied the reaction in the pH range of 7 to 8. The present study has extended the temperature range from 0 to 50" and the effects of salts and methanol have been measured. The ultraviolet absorption of the tetrathionate' was used to follow the reaction.
Vol. 62
Discussion Reaction 1 is described adequately by the rate expression when the cyanide ion concentration has been varied from 1.0 X low4to 5 X M and the tetrathionate concentration varied from2 X to 3 X M. The second-order rate constant k, is a function of the ionic strength. The addition of inert salts (potassium chloride, potassium thiocyanate or sodium sulfate) increases the rate (Fig. 1). The addition of varying amounts of methanol decreases the dielectric constant and decreases the rate. Thus a solvent of 90% water and 10% methanol by volume reduces the rate from 0.757 1. mole-' see.-' to 0.671 1. mole-' set.-' at 25' ( p = 0.012). The data are in accordance with the predictions of the Bronsted-Christiansen-Scatchard equation. It is unusual that marked deviation does not occur until rather high ionic strengths are reached. A mechanism consistent with the kinetic data has been postulated by FosslO involving a nucleophilic displacement of thiosulfate by the cyanide ion CN-
+ -0aS-S-S-SOs-
+ S-SOa-'
(3)
+ 20H- --ic Sod-' + SCN- + H20
(4)
-03s-SCN
L
followed by hydrolysis -0aS-SCN
Experimental Mallinckrodt AR sodium cyanide was used from freshly opened bottles. The solutions were standardized against silver nitrate.4 Sodium tetrathionate dihydrate was prepared from sodium thiosulfate and iodine3 and recrystallized several times from aqueous ethanol. The material was analyzed by the method of Jay.s The salts were AR grade and were to constant weight. Triply distilled conductivity water and pure methanol4 were used to prepare the solutions. All volumetric ware was calibrated and the temperatures were measured with NBS thermometers. Stock solutions of the reactants were prepared, thermostated and then rapidly mixed and transferred to a thermostated 1.00 cm. quartz cell in a cell compartment4 of a Beckman DU spectrophotometer. The optical density (generally at 260 mp) then was recorded as a function of time. The analysis of the data was done by the usual methods.Q
The ultraviolet absorption of thiosulfate" with only one S-S bond is much weaker (E = 105 at 260 mp) than that of tetrathionate which contains three S-S bonds (E = 1000 a t 260 mp in 8ater). The anion, -OaS-SCN, has not been prepared; its ultraviolet absorption would be predicted to be low on the basis of having only one S-S bond. Therefore, there is no evidence from the previous kinetic studies6J or from the present investigation as to the rate of hydrolysis of this postulated intermediate. Reactions 3 and 4 are consistent with the radiochemical data with labelled tetrathionate12 (dithio %) and cyanide ion. The sulfate formed is inactive, the thiocyanate is active and the outer thiosulfur of the thiosulfate is active. The activation parameters of reaction 1are: AH* = 11.0 kcal./mole and AS* = -22 cal./ mole degree at 25.00 f O.0lo, p = 0.0177. The activation energy is somewhat lower than the values generally observed in displacement reactions a t the carbon atom.la In fact the majority of displacement reactions a t the sulfur atom have low activations energies. It appears that the large negative A s + value is the dominant factor determining the rate. This has been commented upon on the basis of the transition-state theory.14a15
(3) 0.A. Nietzel and M. A. Desesa, Anal. Chsm., 27, 1839 (1955). (4) P. D. Bartlett and R. E. Davis, J. A m . Chem. Soc., 80, 2513 (1958). (5) F.Ishikawa, 2. physik. Chem., 130, 73 (1927). . 217, 33 (1934). (6) B. Foresti, 2.a n o ~ gChen., (7)A. D. Awtrey and R. E. Connick, J . A m . Chen. Soc., 73, 1842 (1951). ( 8 ) R. R. Jay, Anal. Chem., 2 6 , 288 (1953). (9) A. A. Frost and R. G. Pearson. "Kinetics and Mechanisms," John Wiley and Sons, Inc.. New York, N. Y.. 1953,pp. 27-53.
(10) 0.Foss, Acta Chsm. Scand., 4, 404 (1950). (11) D. P. Ames and J. E. Willard, J . A m . Chem. SOC.,7 6 , 3267 (1953). (12) A. I. Brodskii and R. K. Eremenko, J . Gen. Chsm. U.S.S.R., 26, 1189 (1955); cf. D. R.Stranks and R. G. Wilkins, Chem. Reus., 67, 841 (1957). (13) A. Streitwieser, Jr., ibid., 66, 571 (1956). (14) Ref. 9,p. 133. (15) W. F. K. Wynne-Jones and H. Eyring, J . Chem. Phys., 3, 492 (1935).
t
NOTES
Dee., 1958
1601
The author wishes to acknowledge discussions with Prof. P. D. Bartlett.
greater absorption of energy, Ar attains but a low maximum yield before its trend is reversed by its impurities. Inhibition by NZl4if not due to initial impurities, may be explained by the formation NOTE ON RADIATION-INDUCED of NH,6 which has i.p. of only 10.23 e.v. and could EXCHANGE O F HYDROGEN readily inhibit by charge transfer. Presumably, that small portion of the reaction BY S. C. LIND chain remaining when inhibition is complete (Fig. Oak Ridge National Laboratory, 1 Oak Ridge, Tennrssee 4)2 is the part propagated by free atoms? ApReceived July 81, 1968 parently the very long chain reaction between H2 I n their recent paper on the influence of ions of and Dz escaped earlier detection on account of the inert gases on the radiation-induced reaction presence of Hg vapor which, like Xe, inhibits the H2 D2 = 2HD, Thompson and Schaeffer2 ion propagation. Almost unattainable purity is dealt with a reactant having ionization potential necessary to avoid some suppression of sensitive (i.p.) midway among the five noble gases. Xenon long-chain reactions. Particularly substancea with and krypton, with i.p. below that of Hz (or Dz), lower ionization potential are to be avoided. suppress by charge exchange a hitherto unknown (4) The authors also have informed me that nitrogen depressed long-chain reaction between H2 and Dz, while He, instead of promoting as its higher i.p. would predict. Ne and Ar, having higher i.p. than Hz, do not do (5) J. C. Jungera, Bull. sac. chim. Belg., 41,393 (1932). so. This the authors interpret, correctly I believe, as indicating that the long-chain reaction is propagated by ions rather than by free atoms or SPECIFIC EFFECTS OF CATIONS ON radical^.^ Without entering here into the question RHODAMINE B EQUILIBRIA1 of which hydrogen ion is the propagator, I wish BY R. W. RAMETTE AND T. R. BLACKBURN to call attention to further conclusions t o be drawn from their results. Contribution from Leighton Hall of Chemistry, Carleton College, Northfield, Minnesota Comparison of their Fig. 22 for the influence of Received July fi6, 1068 additions of He, Ne or Ar with Figs. 3 and 4, which show inhibition by Xe and Kr, seems convincing I n both kinetics and equilibrium research in evidence that Ne, Ar and He, having i.p. above aqueous solutions there is frequent need for the that of H2, make small contributions in the op- variation of hydrogen ion concentration in the posite (positive) direction, also by charge transfer. range from 0.1 to 1.0 M . To minimize the simulThe authors do not stress this, apparently because taneous variation of activity coefficients it has been the accelerating effect is almost negligible in com- commoii practice to add an inert electrolyte which parison with inhibition. This is so because, in- maintains constant ionic strength throughout a stead of breaking short a very long chain as Xe series of solutions. Although the ionic strength and Kr do, only a little ionization is being added continues to be a useful generalization a t low values, by exchange with He+, Ne+ or Ar+, t o the direct there is ample evidence2that inert electrolytes show ionization of H2 by a-irradiation. Nevertheless, increasingly individualistic effects as the ionic the complete series presents an outstanding example strength becomes greater than about 0.05. Thereof ion exchange of H2+ in either direction ac- fore, when studies are made in a series of solutions cording to the relative i.p.'s of varying pH a t constant ionic strength the observed effects are not entirely due to changing He (24.58); Ne (21.58); Ar (15.77); HP(15.43); Kr (14.0); Xe (12.13). acidity. formation ---t + inhibition A common way of varying acidity a t constant Explanation of the subsequent drop following the ionic strength is to use mixtures of strong acid, rise with Ne, Ar and He (Fig. 2)2 requires some usually hydI ochloric or perchloric, with the corconsideration. (The authors have kindly in- responding alkali metal salt. The present work formed me that He also begins a sharp and con- was undertaken to compare the effects upon a tinuing drop at higher concentrations (20-38%) not chosen acid-base equilibrium of various 1:1 inert shown in their Fig. 2.) Impurities from some chlorides a t an ionic strength of unity. The source apparently inhibit the reaction and over- choice of rhodamine B for the equilibrium system come the initial rise. If the impurity were present was based upon several considerations. Rhodain the initial gases, He would appear purest, Ne mine B equilibria have been well characterizeda next and Ar least. The contribution of He is so that unsuspected behavior is minimized, The small due to its low stopping power. That of Ne orange-violet step in the equilibrium system can be reaches a maximum a t 3% Ne, which accounts studied nearly independently of other equilibria almost fully for its observed contribution of lo%, and the acid dissociation constant is about 0.2, if one takes into account that its specific ionization which makes it suitable for study in the acidity is more than threefold that of Hz. Despite its range mentioned above. The spectral characteristics of the dye are such that the equilibrium
+
-
(1) Operated for the United States Atomic Energy Commission by the Union Carbide Corporation. (2) 8. 0.Thompson and 0. A. Schaeffer, J . A m . Chem. Soc., 80.553 (1958). (3) H.Eyring, J. 0. Hirschfelder and H. 8. Taylor, J . Cham. Phys.. 4,479 (1936); P. C. Capron, Bull. 8 0 ~cha'm. . Belo., I S , 222 (1935); W. Mund and M. Van Meersche. ibid.. 67, 88 (1948).
(1) Based upon a thesis submitted by T. R. B. for the degree of B.A. with honors, 1958. (2) See, for example, A. A. Noyes, J . A m . Chem. SOC.,46, I098 (1924); F. A. Long and W. F. McDevit, Chem. Revs.,51, 119 (1952). (3) R. W. Ramette and E. B. Sandell, J . A m . Chem. SOC..1 8 , 4872 (1956).
.
