Jail., 1954
POLAROGRAPHIC STUDIESIN
PORMAMIDE-ACETAMIDE MIXTURES
”ON-AQUEOUS
81
POLAROGRAPHIC STUDIES IN NON-AQUEOUS MEDIA. 11. FORMAMIDE-ACETAMIDE BY JOHNH. HOOP,^ HARRY LETAW,J R . , AND ~ ARMINH. GROPP Chemistry Department, University of Florida, Gainesville, Florida Received June 9,1068
A polarographic study of several inorganic and organic substances was carried out in formamide-acetamide mixtures. It was found that the Ell2 values obtained in these mixtures are in good agreement with those obtained in pure formamide. Exceptional behavior is described in which extreme departure of (id/C)V+/* values from their predicted linearity was observed
Introduction Virtually ideal polarographic behavior of a number of inorganic ions and organic compounds in formamide solutions has been reportedqs A further extension of polarographic investigations of nonaqueous media was carried out by studies of solutions in formamide-acetamide mixtures. The question of the effect of viscosity on the polarographic wave has been experimentally investigated several times in recent years.’ I n all the papers cited, save that of Vavruch, it has been reported that i d is inversely proportional to the square root of the viscosity of the solvent. Vavruch, in studies of saccharose solutions with viscosities as high as 16.0 cp., found that as the viscosities of solutions containing inorganic ions were increased, the i d values became systematically greater than those predicted by the relationship mentioned above. Although Brasher and Jones’ had studied NaOH and HzS04 solutions whose viscosities were in the neighborhood of 5.5 cp., they encountered no such anomalous behavior. It is believed that the increase in viscosity provided by the addition of acetamide to its next lowest homolog, formamide, yields a more nearly ideal system than those previously mentioned. Heyrovsky,S in the classical derivation of the expression for Ey,, related this quantity to the viscosity of the solution in which it is determined. Although differences in E,/, values have been reported for materials investigated polarographically in various solvent systems, there has been no systematic study of this problem. I n view of the fact that no such differences have been mentioned in previous discussions of the effects of changing viscosity, it must be presumed that the functional relationship between E l / ,and viscosity predicted by Heyrovsky has not been observed. The choice of systems available in the present research lead us to believe that this functionality, if demonstrable by (1) Adapted in part from a thesis presented by John H. Hook t o the Graduate Council of the University of Florida in partial fulfillment of the requirements for the M.S. degree, February, 1951. (2) Supported in part by a grant from the Research Corporation. (3) Presented in part at the Meeting-In-Miniature of the Florida Section of the American Chemical Society, May 5, 1951, at Orlando, Florida. (4) Research Corporation Fellow. (5) A.E.C. Predoctoral Fellow, 1950-1952. (6) H. Letaw, Jr., and A. H. Gropp, THISJOURNAL, 67, 964 (1953). (7) E. 8. Peracchio and V. W. Meloche, J . Am. Chem. Xoc., 60, 1770 (1938); D. M. Brasher and F. R. Jones. Trans. Faraday Xoc., 42, 775 (1946); C. H.R. Gentry, Nature, 157, 479 (1946): 0. Collenberg and A. Soholander, ibid., 158, 449 (1946); I. Vavruch, CoZlection Czecho-’ dov. Chem. Communs., 12,429 (1947); H. A. McKenzie, J . Council Sci. Ind. Research, 21, 210 (1948); P. Rutherford and L. A. Cha, A n d . Chem., 23,1714 (1951). (8) J. Heyrovsky, “PoIarographie,” Springer-Verlag, Vienna, 1941, P. 151 ff.
conventional polarographic instrumentation, could be determined. Experimental Materials.-Baker reagent grade Zn( C~H30&2H20 and Baker and Adamson reagent grade Pb(N0a)z were dried and used without further purification. Dr. F. H. Heath of this Laboratory prepared and recrystallized Tl?SO4 for this research. Recrystallized fluorenone was provided by Dr. F. E. Ray, Director of the Cancer Research Laboratories of the University of Florida. Recrystallized benzalacetone and benzophenone were given to us by Dr. C. B. Pollard of this Laboratory. Dried reagent grade KCl and KNOs were used as supporting electrolytes. Metal Salts Corporation triple-distilled mercury was used without further purification. Eimer and Amend chemically pure grade formamide, after being placed under a pressure of less than 1 mm. for two hours, was found to be polarographically pure in the range studied. Eimer and Amend chemically pure grade acetamide was found to be polarographically pure. All solutions exce t those of TlzS04 were prepared in stock solution of 0.1 M $Cl in the formamide-acetamide mixture of desired composition. A 0.1 M KNOa supporting electrolyte was used for the TlzSOl solutions. The solvent system studied consisted of formamideacetamide mixtures which were 20.3, 30.9 and 43.2 mole per cent. acetamide. Mixtures of higher concentrations were solids at 25”. Apparatus.-Polarograms were obtained using a SargentHeyrovsky Model XI1 Polarograph. Potentials were measured on a Gray “Queen” Potentiometer calibrated with a Weston Standard Cell. A 0.1 N calomel reference electrode equipped with an agar-KC1 bridge was used in order that transference effects across the interface be minimized. All potentials are referred to the saturated calomel electrode. The mercury reservoir was 45 om. above the level of the solutions. It was attached by neoprene tubing to a glass capillary of Corning “marine barometer” tubing 8 em. long with an internal diametep of 0.08 mm. The capillary constant was 2.98 mg.2/r/sec.-1/! at -0.40, and 2.94 mg.¶/I sec.-’/s at -0.80 applied volt in 0.1 1cf KCl in the 30.9 mole per cent. acetamide mixture. Tank nitrogen was bubbled through a tower of alkaline pyrogallol and then through a tower of formamide-acetamide mixture of the proper composition before being introduced into the Heyrovsky-type electrolysis cell for the purpose of sweeping out atmospheric oxygen. The solvent tower eliminated errors in concentrations caused by the removal of solvent from the cell by the sweeping gas. During all determinations, the polarographic cell was immersed in a waterbath thermostatically controlled at 25 & 0.1”.
Results I n Table I, the densities, viscosities and molar volumes of the formamide-acetamide mixtures used are given. The viscosities and the molar volume of the solutions are found to be linearly related t o the mole per cent. acetamide. TABLE I PROPERTIES O F FOR~IAMIDE-ACETah.rIuE hfIXTURES Acetamide d%, 111 iiiole per ceA. g./cin.$ rllp. ctn. a/lklole ?I
20.3 30.9 13.2
1.103 1.091 1.090
v,
46.43
44.0
54.93 65.19
45.9 47.7
JOHNH . HOOK,HARRYLXTAW,JR., A N D ARMINH. GHOPP
82
VOl. 58
TABLEI1 HALF-WAVE POTENTIALS Comppund or ion
Pormaniide (ref. ( )
Bansslacetone
-1.08 *0.03
Half-wqve potential referred to S C E (volt@) Formaniide-acetamide mixtures, niole per cent. acetamide 20.3
-1.28& .03 -1.26 i .01
Benzophenone Fluorenone Lea i(I1)
-0.87 & .02
- .38 k .02 - ,35i .01
Thrtllium(1) Zinc(I1)
30.9
-1.12 k 0 01 -1.26& .01 -1.22 =t .01 -0.89 f .01 . 3 6 k .O1 - .34 f .01 . 9 9 f .91
-1.13 f 0 01 -1.27 d~ .03 -1.22 f .01 -0.86 =t .01 - .35 z t .01 - .36 f .01
- .92f
- .97 j, .01
The half-wave potentiah previously reported in formamidea and those determined in the formamide-acetamide mixtures are found in Table 11. 1.6
-1.08 -1.29 -1.21 -0.89
-
.01
43.2
f 0.01
.03 .01 .02
&
& &
- .40k - .35 f - .97 f
.01
.01 .01
Average ('fd/C).?7i/z values for each of the compounds investigated are tabulated for 'each solvent system in Table 111. Graphs of id us. c curves are plotted for TI+, P b + f and fluorenone in Figs. 1, 2 and 3, respectively.
-
1.2
k
ii
4
0.8 0.8-
2
2 0.4
-
0, 0
Fig. le-id
I
I
I
I
I
1
8 16 24 MillIxpoles/liter X IO2. us. c curves for TI+:0,20.3; A, 30.9; 3, 4'&2 mole per cent, acetamide.
01.
,,
,!
I
I
I
J
2 4 6 Millimoles/liter Fig, ?.-&Ivs. c curves fpr fluorenone: 0,20.3; A, 30.9; O, 43.2 m o b per cent. acetamide. 0
ot 0
Fig. % - - i d
Discussion It can be seen from the data in Table I that the solutions dealt with were virtually ideal in the range studied. The linear nature of the inorease in viscosity with increasing mole per cent, of acetamide would seem to indicate thst the informatian obtained relative to the effects of viscosity is valid. The data collected in Table I1 indicate that the El/, values are not functionally related to the viscosity of the solvent in which they are determined. It must be remembered, however, that a functional relationship of this type may be so subtle as to be undetectable by the instrumental methods used, .Because the potential measurements involve an I I , I I unknown and variable junction e.m.f. between acet0.8 I .6 2.4 amide-formamide solutions and water extruded Millimoles/liter. vs. c curves for Pb++: 0, 20.3; A, 30.9; 0 , 43.2 from &p agar-KC1 bridge, it is not posBible to extend the precision of these determinations. Good mole per cent. aeetamide.
ELECTRICAL POTENTIAL
Jan., 1954
O F ATYL0n’ FIBERS IN
agreement of the E l / , values determined in formamide with those evaluated in formamide-acetamide mixtures was obtained. The data presented in Figs. 1, 2 and 3 are apparently anomalous. It is evident that id/C is not, in general, inversely proportional to the square root of the viscosity. The data in Table 111 indicate that Zn++ is the only ion studied which behaved approximately according to this relationship. TABLE 111 AVERAGE(i,JC)y1’2 VALUES,MA. MP.’/~MOLE-^ Compound or ion
Benzalacc tone Benzophenone Fluorenone Lead(I1) Thallium( I) Zinc( I1j
Formamide-acetamide mixtures, mole per cent. acetamide 30.9 43.2
20.3
34.0 25.5 27.2 27.2 44.3
30.2
32.2 28.5 34.2 57.4 40.8 30.7
38.0 31.2 38.4 54.0 33.1 31.6
The results of Vavruch’ have shown that exalted
id values result when T1+,Pb++ and Zn++ are studied in saccharose solutions of viscosity higher than 20 mp. In contrast to these results, we have found that (id/C)’771/2 is not necessarily a monotonically increasing function with increasing viscosity. There seems to be no clear-cut pattern in the data. The id vs. C curves for the organic compounds investigated show complete inversion, with respect to behavior predicted on increase in viscosity, for fluorenone and benzophenone, but inversion only of the 30.9 mole per cent. acetamide solution of benealacetone. Virtually ideal behavior was found for Zn++, but Pb++ was completely inverted with the 30.9 and 43.2 mole per cent. acetamide
AQUEOGSM E D I A
83
solutions reinverted. T1+ presents no inversion, but it is seen from Table I11 that i d / c decreases at a greater rate with increasing viscosity than is predicted by the formula stated. It is reasonable to assume that the inversions found for fluorenone and benzophenone are attributable to a greater solvation tendency exerted by formamide than by acetamide with respect to these two compounds. A reverse explanation would seem to pertain in the case of Tl+. Though this ion ordinarily is not solvated or complexed in aqueous solutions, it is entirely possible that it is solvated by acetamide. The description of the phenomena observed on the basis of solvation becomes somewhat tenuous in the light of the data obtained for Pb++ and benzalacetone. It would be necessary to state that a miiiimuni and a maximum, respectively, occur in selective solvation of these materials. This is not unreasonable, but the absence of data concerning solvation and complex formation prevents the unambiguous resolution of this question. Mixtures of formamide and acetamide are suitable solvents for polarographic determinations. Although extensive solubility work has not been carried out in this system, it offers advantages in that a wider choice of supporting electrolytes is available for use than is offered by the commonly used alcohols. It has been shown that the E+ values of several organic compounds and inorganic ions remain constant over a rather wide range of viscosities. Anomalous behavior of the diffusion currents of benealacetone, benzophenone, fluorenone, Pb++and T1+has been observed. These latter phenomena indicate that the viscosity relationship cited above must be applied with caution to other than aqueous systems.
ELECTRTOAL POTENTIAL OF NYLOhT FIBERS I N ,4QUEOUS MEDIA B Y FREDERICK T. W A L L AND
PATRICIA
M. SAXTON
Noyes Cheinichl Laboratory, Universify of Illinois, Ui’bana, Illinois Received Mau 86, 1956
The electrical potential of a nylon fiber interacting with solutions of bases is calculated from theoretical Considerations and previously obtained experimental data. It is shown that the fiber potential, which i R in the neighborhood of -0.3 volt, is relatively insensitive to temperature changes. The number of end groups on the fiber that must be charged to produce this potential is shown t o be negligibly small compared to the number of end groups interacting with the basic solutions. This calculation verifies the assumption, usually made for the theoretical development of the absorption equations, that equivalent numbers of cations and anions are absorbed by the fiber.
Introduction Several investigatorsl-6 have studied the interactions of polymeric fibers such as ~voolor nylon with acids and bases. Gilbert and RideaP developed the first theoretical equations for the special case of absorption of acids on materials having (1) G . A. Gilbert, Proc. Rou. SOC.(London), 1838, 167 (1944). (2) R. H.Peters, J . SOC.Dyers Colourists, 61,95 (1945). (3) 0. R. Lernin and T. Vickerstaff, ibid.. 63, 405 (1947). (4) F.C . MoGrew and A. K. Schneider, J . Am. Chenb. Soc., 72,2547 (1950). (5) W.R. Remington and E. K.Gladding, ibid., 72, 2553 (1950). (6) B. Oloffssen, J . Soc. Dyers C o l o u ~ i s t s67, , 57 (1951). (7) G. A. Gilberl and E. K. Rideal, P r o c . Rov. Soc. ( L o n d o n ) , 1828, 335 (1844).
equal numbers of carboxyl and amine groups. Wall and Swoboda8 subsequently extended these equations to cover the more general problem of absorption of acids or bases on fibers having either equal or unequal numbers of those groups. The latter theory was experimentally confirmed by measurements of the absorptioii of sodium hydroxide on commercial nylon fibers, and extension of this experimental work to different temperatures has enabled the present authorss to calculate the heats of reaction associated with the absorption processes. ( 8 ) F. T. Wall and T. J. Swoboda, THISJOURNAL, 56,50 (1852). (9) F. T. Wall and P. M. Saxton, ibid., 57,370 (1953).
