Polarographic Behavior of 1-Hydroxypyridine-2-Thione Alan F. Krivis Department of Chemistry, The University of Akron, Akron, Ohio 44304
Eugene S . Gazda Olin Mathieson Chemical Corp., New Haven, Conn.
A recent polarographic study of some metal ion derivatives of 1-hydroxypyridine-2-thione,which was made for analytical purposes, disclosed unusual phenomena concerning the electrochemical behavior of the thione functional group (1). Therefore, a more detailed study was undertaken in an attempt to discover the basis for some of the experimental results previously obtained. In particular, the effect of pH on the polarographic behavior of 1-hydroxypyridine-2-thionewas studied in aqueous and ethanolic media. The tautomeric equilibrium between l-hydroxypyridine-2thione (I) and 2-pyridinethiol-1-oxide (11) was studied by Jones and Katritzky (2), who found that the compound existed predominantly in the thione form (I). The pK values for
-
o = s
both addition (Equation 2) and abstraction (Equation 3) of a proton also were determined; the values obtained were - 1.95 and 4.67, respectively. Therefore, above pH .- 5, the anionic
OH 1 (1)
I
OH (Ill)
on I
EXPERIMENTAL
Apparatus. Sargent Polarographs Models XXI and XV and a Metrohm Polarecord E-261R were used in conjunction with a thermostated H-cell (25 f 0.1 “C)containing a saturated calomel reference electrode (10). No damping was used with the Sargent Polarographs and a damping setting of 5 was used with the Metrohm Polarecord. Several capillaries with constants varying from 1.975 to 2.281 mg2k116 (distilled water, open circuit) were used. A potentiostat and current integrator manufactured by Analytical Instruments, Inc., was used for controlled potential electrolysis. The electrolysis cell consisted of a waterjacketed (25 “C) 500-ml three-neck flask. A mercury pool served as the working electrode with an isolated carbon rod as the auxiliary electrode. A large calomel electrode (Beckman No. 39970) was used as a reference electrode. Chemicals. The sodium salt of l-hydroxypyridine-2thione ( O h Mathieson Chemical Corp.) was purified by acidification, several recrystallizations of the free acid, and reformation of the sodium salt from the purified material. Assay by a silver titration indicated the final product to be 98.0 pure. No decomposition product-Le., the disulfidecould be detected polarographically or from infrared spectra. Aqueous buffer solutions were prepared with double the concentration of material desired at final dilution and were diluted 1 :1 with either water or ethanol. The final concentrations and measured pH values for both media were as follows:
(IV)
PH
I
OH (1)
forms V and VI would predominate in solution. The more probable structure is VI, based on NMR spectra which indicated that changes in the sulfur had occurred in the basic pH range (3). Polarographic studies of thione compounds have been reported by several workers (4-7), as have the thione complexes of some metals (8). Zuman summarized the results obtained for a number of mercapto compounds and metal mercaptides (9). (1) A. F. Krivis, E. S . Gazda, G. R. Supp, and M. A. Robinson, ANAL.CHEM., 35,966 (1963). (2) R. A. Jones and A. R. Katritzky, J . Chem. SOC.,1960,2937. (3) H. Agahigian, Olin Mathieson Chem. Corp., private com-
munication, 1966. (4) R. A. F. Bullerwell, J. Polarog. SOC.,8,2 (1962). (5) M. Fedoroiiko, 0. ManouSek, and P. Zuman, CON. Czech. Chem.Communs., 21,672 (1956). (6) L. JenSovskf, ibid., 21, 459 (1956). (7) C. J. Nyman and E. P. Parry, ANAL.CHEM., 30, 1255 (1958). (8) T. J. Lane, Adv. Polarography, 2, 797 (1960). (9) M. Brezina and P. Zurnan, “Polarography in Medicine, Biochemistry, and Pharmacy,” Rev. English Ed., Interscience, New York, 1958. 212
ANALYTICAL CHEMISTRY
Buffer 1. 0.5MHCl 2. 0.1MHC1 3. 0.1M KCI-O.01M HCl 4. Britton-Robinson buffer 5 . 0.1M Acetic acid-0.01M sodium acetate 6. 0.1M Acetic acid-0.1 M sodium acetate 7. 0.05M Acetic acid-0.05M sodium acetate 8. 0.01M Acetic acid4.1M sodium acetate 9. Britton-Robinson buffer 10. Britton-Robinson buffer 11. 0.05M Borax
Aqueous 0.4 1.1
Ethanolic 0.4 1.2
2.0 2.9
2.2
3.7
4.5
4.9
...
...
5.3
5.5 6.8 7.9 9.2
6.3 7.4 8.5
3.8
...
Stock solutions of the sodium salt of l-hydroxypyridine2-thione were prepared with either oxygen-free water or ethanol and the flasks were covered with aluminum foil to minimize decomposition. Fresh stock solutions were prepared before each series of measurements and were not used if more than 4 hours old. Gelatin (Fisher Scientific Co.) was used as a maximum suppressor, and purified and equilibrated nitrogen was used to remove oxygen from all solutions. Procedure. (A) Five milliliters of buffer solution were transferred to a 10-ml volumetric flask and 0.1 ml of 0.5% (10) J. C. Komyathy, F. Malloy, and P. J. Elving, ANAL.CHEM., 24 541 (1952).
0.3
1
/
.
5
B I
.'
wI I 2
I 4
I
I
6
8
//
PH
Figure 1. Ell2 values for 1-hydroxypyridine-2-thioneas a function of pH in aqueous ( I ) and ethanol-aqueous (2) solvents
gelatin solution was added. The necessary volume of 1hydroxypyridine-2-thionestock solution was added and the solution was made up to volume with the proper solvent; the final solution containing ca. 1-2 m M electroactive species. After mixing, the solution was transferred to the H-cell, deoxygenated for 10 minutes and the polarogram was run. The cells were covered with aluminum foil to minimize decomposition during measurement. The polarograms were evaluated graphically with the average of the recorder traces, (B) Controlled potential electrolyses were carried out at pH 2.9 and 8.1 with the potentials maintained at -0.04 V and -0.1 V (us. SCE), respectively. Accurately weighed samples of ca. 0.3 gram of the sodium salt of l-hydroxypyridine-2-thione were quantitatively transferred to the deoxygenated solutions, which had been pre-electrolyzed to remove impurities, and the electrolysis was carried out. Initial currents of about 0.2 A decayed to about 10 mA at which point the electrolyses were essentially complete. The time required for a complete electrolysis was about 2 hours.
RESULTS AND DISCUSSION Two solvent systems were used for this study, a completely aqueous solvent and a 1 : 1 ethanol-water solvent. The latter medium was utilized to permit correlation with the data previously obtained for the metal chelate systems ( I ) ; the chelates are insoluble in water and ethanol is needed to enhance solubility. In terms of the data reported here, there were no unexpected differences in the results obtained in the two solvents. The presence of ethanol produced the anticipated increase in measured pH values and a diminution in diffusion current because of increased viscosity. Therefore, the evaluation of the polarographic results applies, interchangeably, to either solvent. Only one anodic wave was found in all of the polarograms run at all pH levels. The half-wave potential shifted linearly to more negative potentials as a function of pH, up to the pK us. value, and then became independent of pH. A plot of pH for both solvents is shown in Figure 1 ; the change in slope occurred at pH 4.6 in water and 5.4 in ethanol-water. The slope of the portion of the curve below the pK corresponded to 53.1 mV/pH unit in water and 56.7 mV/pH unit in ethanolwater. These values are in agreement with the theoretical shift of 59 mV/pH unit expected for an electrode reaction involving one proton. The mean value for the diffusion current constant, I , for the range of pH levels studied was 1.76 in water and 1.35 in
I
0 log i/(& -i)
-1
Figure 2. Plots of E us. log i/(id
+I
I
- i ) at pH 1.1 and 6.3
ethanol-water. By comparison, the disulfide, 2,2 '-dipyridyl1,l '-dihydroxydisulfide, gave an I value of 2.72 in water and 2.19 in ethanol-water. Because the disulfide reduction is known to be a two-electron reaction, the anodic reaction of 1-hydroxypyridine-2-thione appears to involve only one electron. On the basis of these data, the anodic reaction over the range of pH 0.4-5 can be represented as: HRS
+ Hg
RSHg
-+
+ H+ + e-
(4)
+ e-
(5)
and above pH 5 as: RS-
+ Hg
+
RSHg
This one-electron, one-proton reaction has more of the character attributed to a thiol anodic wave rather than a thione. The alternate thione mechanisms, such as those proposed for thiourea and its derivatives ( 5 ) and for 2-mercaptoimidazole (4), 2HRS
+ 2Hg
or 2HRS
-
RSHgHgSR
+ Hg
RSHgSR
+ 2Hf + 2e
+ 2H+ + 2e
(6)
(7)
d o not seem likely in the present case. For example, both Reactions 6 and 7 require that the equation for the polarographic wave be of the form
In that case, a plot of E us. log i/(id - i) is nonlinear, while a plot of E US. log i / ( i d - i)' is linear, with a slope of RT/nF. Thiourea ( 5 ) and 2-mercaptoimidazole ( 4 ) exhibit behavior consistent with Equations 6.4. However, plots of E us. log i/(id - i) for 1-hydroxypyridine-2-thionewere linear. Examples of the plots at pH 1.1 and 6.3, in either solvent, are shown in Figure 2. The slopes were 52 mV and 5 5 mV, respectively. VOL. 41, NO. 1 , JANUARY 1969
213
Deviations of the slopes from ideal values probably are caused by adsorption phenomena such as that found with 2-mercaptoimidazole (4). The log plots, therefore, support a one-electron reaction similar to a thiol, such as ergothionine ( I I ) , rather than the expected two-electron thione reaction. Further support for the one-electron mechanism was obtained from the controlled-potential electrolysis experiments. At pH 2.9, the reaction was found to involve 0.92 electrons/ molecule and, at pH 8.1, the value was 0.93. These data corroborate a one-electron reaction. The electrolysis product was isolated and analyzed, with the results establishing its formula as the mercuric dimercaptide, Hg(SR)z. This corresponds to similar findings by Stricks, Kolthoff, and Heyn-
RECEIVED for review May 20, 1968. Accepted September 20, 1968. Manuscript abstracted in part from a dissertation submitted by E. S. G . to St. Joseph College, Hartford, Conn., in partial fulfillment of the requirements for a Master of Arts degree.
(11) V. Preininger, M. Ctirnoch, and A. Santav);, Ado. Polarography, 3, 1087 (1960).
(12) W. Stricks, I. M. Kolthoff, and A. Heyndrickx, J . Amer. Chem. SOC.,76, 1515 (1954).
drickx (12) that compounds of the type RSHg are unstable in aqueous solution and decompose to (RS)2Hg and Hg. No cathodic wave attributable to an N-oxide reduction was found at any pH level. By contrast, pyridine N-oxide and 2-chloropyridine-N-oxideboth give cathodic waves. The presence of sulfur in 1-hydroxypyridine-2-thionemust stabilize the N-oxide linkage to reduction.
Determination of Hydrogen in Partially Deuterated Water Maurice M. Kreevoy and Thomas S . Straub
School of Chemistry, University of Minnesota, Minneapolis, Minn. THEDETERMINATION of the isotopic composition of partially deuterated water has been the subject of a number of investigations. Earlier methods, generally based on density, are described and discussed by Kirshenbaum ( I ) . Later methods have been based on a variety of physical and chemical properties, and numerous references are given by Leyden and Reilley (2). A previous method based on the spectrum in the 1900 nm region (3) lacks the precision, and, for many purposes, the convenience of the present method, although it has the advantage of using much smaller samples. The present note describes a method based on the spectrum of water in the 800-1300 nm region (4). Liquid HzO has peaks at 970 and 1192 nm and a valley at 1065 nm (4). A semiempirical relation is developed among AIl92,A1065, and xH. (Absorbance at wavelength, A, is A x . ) In this relation A1065 is, effectively, used in the place of a suitable reference material, which is unavailable. The method offers an accuracy of -BO2 in XH, the atom fraction of ordinary hydrogen. Since the analysis is carried out in 1-5 cm quartz spectrophotometer cells, it is particularly convenient for applications where other spectrophotometric measurements have been made on the same samples, as in many measurements of rates. EXPERIMENTAL
All spectra were made on a Beckman DK-2 spectrophotometer. Its wavelength scale was calibrated by means of the major water vapor lines at 1362, 1375, 1382, and 1400 nm (5). These were observed by scanning the transmitted energy in one beam with the cell compartment empty and the chopper turned off. Only air was used in the reference beam. A suitable reference substance would have to have no absorbance in the region of interest and a refractive index identical with water.
Densities used in the determination of isotopic composition were determined by a standard pycnometric technique (6) using a 50-ml pycnometer. Deuterium oxide was obtained from a number of commercial suppliers and was redistilled before density determinations. RESULTS
Figure 1 suggests that, for DzO, A1065 and AlI92are equal. Since the two wavelengths are so close together, it is reasonable to assume that scattering at the windows and by dust is the same at both of them. With this observation and assumption the quantity AA, defined as A1192 minus A1065 should be due entirely to absorption by HOD and HzO. Assuming that AeHoD( e 1 ~ 9 2-~ elOejHoD) ~ ~ is exactly half of AeHzO,then A A at constant pathlength, b, should be proportional to xR. Molar absorbtivities are designated enM where M i s the absorbing substance and n is the wavelength. Figure 2 shows that this relation is not obeyed with useful accuracy. If AeHoais given by rAeHzo,r being a constant independent of isotopic composition, A A is given by AA =
xH30
214
ANALYTICAL CHEMISTRY
+
X H ~ DUAeH’Ob.
(1)
The mole fractions of HzO and HOD, XH,O and X H O D , can be evaluated in terms of X H and K, the equilibrium constant governing the redistribution reaction, HzO
K
+ Dz0 e 2 HOD
(2)
If K is 4.00 as suggested by the rule of the geometric mean (7), they are XHiO
and xHOD
(1) I. Kirshenbaum, “Physical Properties of Heavy Water,” McGraw-Hill, New York, 1951. (2) D. E. Leyden and C. N. Reilley, ANAL.CHEM., 37, 1333 (1965). (3) W. E. Keder and D. R. Kalwarp, ibid., 38, 1288 (1966). (4) M. R. Thomas, H. A. Scheraga, and E. E. Schrier, J. Phys. Chem., 69, 3722 (1965). (5) E. K. Plyler, Phys. Reo. (2), 39 77 (1932).
AeHlob
=
=
(3)
XH2
2 XA(I-
XH).
(4)
Combination of Equations 1, 3, and 4 leads to (6) D. P. Shoemaker and C. W. Garland, “Experiments in Physical Chemistry,” McGraw-Hill, New York, 1962, p 126. (7) J. Bigeleisen, J. Chem. Phys., 23, 2264 (1955).
