June, 1962
EFFECT 0 F SOLUTION CO&IPOSITIOX0‘s THE ~EPTUNUR.I.(TT)-(T71) C O U P L E
Peard and Pflaum6 prepared a number of solid complexes of si1,ver with aromatic heterocyclic amines. Only 4,8-dimethylquinoline and 4,4’bipyridine gave 1 : l complexes instead of [Ag(aminelz]”. The 1 : 2 complexes all melted without decomposition a t temperatures bclow 160’. Thc bipyridine complex, like our pyrazine complex, had a very high melting point (5!50°) which led the authors to suggest a polymeric structure. I n solid [Ag(Pa)]S03 there may be polymeric chains in which pyrazine acts as a bidentate bridging group. A similar structure has been suggested for solid nickel pyrazine complexes.2b A value of K2 less than K1 is very unusual for monodentate ligands. Of the 17 complexes of silver with aromatic heterocyclic amines listed in “Stability Constants”‘ log K1 and log K Zvalues are given for 10 (Table nos. 113, 117-9, 1,6643, 202-3, 246), of which only 4-methoxypyridine forms complexes with KZ< K,. The ratio log Kz/log K1 ranges from 0.95 to 1.32 for these complexes. For pyrazinesilver the value is 0.62/1.51 or 0.41. Weimer and Fernelius’sBdata on the silver complex of 6-methyl-2picolylamine give log K&og K1 = 0.87. It is (6) W. J. Peard a n d R. T . Pflaum, J. Am. Chem. SOC.,8 0 , (1958). (7) J. Bjerrum, G. Schwarnenbach, a n d L. G. Sill&, “Stability stants of Metal-Ion Cxnplexes,” P a r t I , The Chemical Society, don, 1967. (8) H. R. Weimer and R‘. C. Fernelius, J . Phys. Chem., 64, (1960).
1593 ConLon-
1065
thought that the coordination is through the primary amine rather than the ring IT so it cannot be cornparcd with the aromatic heterocyclic amine complexes considered above. Pyraeinesilver complex is atypical in yet another way. Thc ratio log Kl/pKa is approximately constant for silver complexes with a given type of amine.9 For aliphatic amines the ratio varies from 0.25 for primary amines to 0.32 for tertiary amines. For aromatic heterocyclic amines the value ranges from 0.31 to 0.43 with an average of 0.36. The fractional value is to be expected since the proton has a tremendously greater ionic potential than a metal ion. For pyrazinesilver, however, log K1/ pK, is 2.3. The stability is some hundred-fold greater than we would expect from the basicity of pyrazine alone. We do not know the source of this extra stability. If pi bonding were involved we would expect extra stability to appear with other aromatic heterocyclic amine complexes. Furthermore, there is evidence that pi bonding does not occur in solid cobalt, nickel, or manganese complexes with pyrazine.7b Acknowledgments.-This work was supported by the Graduate Council of Southern Illinois University. We also are indebted to the Wyandot)te Chemicals Corporation for a generous sample of pyrazine.
1951 (9) G. Schwarzenbach, Helv. Chim. Acta, 3 6 , 23 (1953).
EFFECTS OF SOLUTIOK CORIPOSITION ON THE POTESTIALS OF THE COUPLES: Np(V-VI), Fe(II-III), Hg(O-I), AND Hg(1-11)’ BY A. J. ZIELEN AND J. C. SULLIVAN Argonne A’ational Laboratory, Argonne, 1llinois Received December 4, 1061
A linear variation with hydrogen ion concentration was found for the formal potential of the Mp(V-VI) couple a t 25” in perchloric acid-sodium or lithium perchlorate solutions of ionic strength 2. The major portion and very probably all of this acid dependence is shown to be caused by activity coefficient changes, reaffirming ?r’pOs+and N p 0 2 + as the only Np(V) and Np(V1) species of consequence in acid solutione. Similar measurements in sodium perchlorate solutions were made of the Fe(I1-III), Hg(0-I), and Hg(1-11) couples. Equations for the formal potentials as a function of hydrogen ion concentration are given, and activity coefficient effects produced by the substitution of K a + or Li+ for H + a t constant ionic strength are discussed.
Introduction involved, especially if only minor amounts of other The assignment of the general formulas MOz+ forms are present. In the only reported study of the Np(V-VI) and MOZ2+as the dominant species in acid solution for the (V) and (VI) oxidation states of U, Kp, Pu, couple as a function of acidity, a definite though and Am is based on a variety of evidence.2 Earlier minor hydrogen ion dependencc was observed. this assignment was largely based on the reversi- Magnusson, Hindman, and LaChapelle4 studied the cell bility and acid independence of the actinide (V-VI) couples, but at present the infrared measurements of Jones and Penneman3 probably represent the single, most compelling proof. However, the hy- Holding the nitrate concentration constant a t 1.02 drogen ion dependence of these oxidation-reduction M , the hydrogen ion was varied from 0.12 to 1.02 couples remains the most direct test of the species ill. This resulted in a 20 mv. variation in the ob( 1 ) Based on work performed under the auspices of the U.S. btomio served e.m,f. compared to a 55 mb?. predicted Energy Commission. change if the (V-’VI) couple possessed a first power (2) J. J. Katz and C:. T. Seaborg, “The Chemistry of the Bctinide Elements,” Methuen a n d Co., Ltd., London, 1957, pp. 176-179. 220223, 292-297. 366-370. (3) I,. 13. Jones and R. A. Penneman, J . Chem. Phys., 21, 542 (1953)
( 4 ) L. B. Magnusson, J. C . Hindman. a n d T. J. LaChapelle, Paper 15.4, “The Transuranium Elements,” National Nuclear Energy Series, IV, RIcGraw-Hill Book Co., Inc., New York, N. Y., 1949, Vol. 140, P a r t 11, p. 1059.
.2. J. ZIELEK -,yn J . C . STJIAIVAS
1066
hydrogen ion dependence. However, independent experiments indicated that at least, 10 my. and perhaps all of this changc was due to the junction potcritial of t h r qaturatpd potasrjiimi nitrntc bridge. Thus, although a small hydrogm ion \ w y probable, tho aut hors cwicludr couplc to hc mniiily acid indcpcndcrit. In a recent) communication we h a w descri procedures for precise mcwurcmcnts of the Np VI) couple without liquid junction potential, using the glass electrode as an intermediatc for the standard hydrogen Determination of formal potentials by this method allowed a critical test of any hydrogen ion dependence of t,he Kp(V-VI) couple. I n order to investigate possible changes in activity coefficients, similar measurements also were made of the formal polentials of the couples: Fe (11-111), Hg(0-I), and I-Ig(1-11). Experimental Reagents and Analytical Procedures .-An approximately 1: 1 ferrous-ferric perchlorate stock solution was prepared by dissolving the reagent grade sulfates in water followed by titration with barium perchlorate. The barium sulfate precipitate was removed by centrifugation and pssage through a fine sintered glass filter. After standing several days, the filtration was repeated and the solution diluted to final volume with a known amount of dilute perchloric arid. It was estimated that the slight excess of barium perchlorate &! in the final cell solutions. used amounted to ca. 3 X The equilibrium ferrous roncentrations were determined by direct micro titrations of cell solution samples a t the end of each potential run, using standard ceric sulfate and ferroin indicator. A4comparison check of the micro titration procedure with macro scale results agreed within 0.1%. Ferric content was taken as the difference from the calculated total iron concentration obtained by a Zimmermann-Reinhardt process standardization of the ferrous-ferric stock. A mercurouB perchlorate and a mixed mercurous-merruric perchlorate stock were prepared by dissolution of a weighed amount of EIgO (yellow, Mallinckrodt, analytical reagent) in a known excess of perchloric acid. A portion of this solution was stored over metallic mercury with frequent and vigorous shaking. After several days equal aliquots of the resulting Hg( I ) solution and the original Hg( 11) solution were combined to form the mixed Hg(I-11) stock. All metallic mercury was triple distilled and filtered before use. The mercury(1) and (11) concentrations in the cell solutions all were calculated quantities based on the original weight of HgO used plus appropriate minor corrections for the equilibrium6 Hg'f IIg = Ilg,'+ (2) Primarily this assumed stability of the Hg(1-11) stock. This was checked by analysis for Hg(1) in the mixed stock a t the end of the cell measurements (a three-meek period) by the method of Willard and Young.7 The calculated and found concentration agreed within 0.2%. Details of the neptunium(\*) and (VI) preparation and all other reagents employed have been summarized previously.8 T u o independent stock solutions, A and B, of mixed neptunium(V-VI) perchlorates were used. In the final cell solutions stocks A and B, respectively, gave total neptunium concentrations of 2.167 and 1.005 X M and approximate Np(V)/Np(VI) ratios of 1.04 and 1.22. As before5 the actual Np(V)/Np(VI) ratio was determined directly by multiple spectrophotometric analyses of the cell yolutions a t the end of each potential run. The hydrogen ion concentrations of the cell solutions from stock A were determined by direct micro titration with standard base after reduction of all neptunium to the (V) state with exccss sodium nitrite.
+
T'ol. GG
This is a satisfactory procedure only a t relatively high acid and low neptunium concentrations. To extend our measurements below 0.2 M hydrogen ion, stock B was preand pared with a low free acid concentration (ca. 10-2 A!); the pII of the stock solution was nicasured after a tenfold dilut,ion with a Heclcnian Model G p H Meter. Conversion to moles/l. of hjdrogen ion was done by calibration with solutions of known acidity and similar ionic strength. The [H+] of the cell solutions then was calculated from make-up with a minimum of YiPJ, of the total acid obtained from the addit,ion of a pure perchloric acid stock. The hydrogen ion concentrations of the iron and mercury cells also were obtained by calculation. However, the free acid contents of t,hese stocks were accurately known, and their maximum contribut'ion to t.he total hydrogen ion \vas only about 10%. Potentiometric Measurements.-Illustrating with the Np(V--T'I) couplc, potentials as a function of hydrogen ion concentration or C h were made on the cell pairs Pt, Hz IC& HClO,, ( 2 - Ch - 3cs).!\fNaC104, Co&!i' &(C104)2Mass (3) glass jChn/I HClOa, ( 2 - C h - C, - 3C6)M NaC104, Cshf KpOnC10.1,C d f NpO*(C104)2jPt (4) With C,j and C6-the pu'p(V)and Kp(V1) concentrationsrelatively small (ca. 5 X 10-3 or the hydrogen ion responso of the glass electrode is assumed the same in the similar solutions of cells 3 and 4. Addition of the two potentials gives the junction-free e.m.f. of the reaction
+
+
'/zH? Sp(V1) Sp(V) H f (5) Details of the apparatus and procedure have been described.j However, it should be noted that in cells of type 3 there is a ca.ncellation of hydrogen ion activity and hence no need to match precisely the acidities of cells 3 and 4. At one atmosphere of hydrogen the e.m.f. of ( 3 ) supplies the "standard potential" of the glass clectrode with possible changes in the as1 mmctry potential of the glass as the only variable. Thus our bstsic assumption of identical glass electrode response in two different but similar solutions requires only matching of the asymmetry potentials. This has been verified under much more drastic conditions than encountered here and will be discussed in a subsequent manuscript. Cell pairs similar t o ( 3 ) and (4) were used for the I-Tg(I-11) and Fe(I1-111) couples, using Ys" as an inert substitute for Fe3+. The shiny platinum electrode of (4) was replaced by a mercury pool contacted by a platinum needle for the Hg(0-1) couple. Two types of Beckman glass electrodes were employed: the now obsolete 290-7 and the current, 89177. The st,ability performance of the latter type was much superior, probably because of the improved shielding which extends well into the body of the electrode. The 39177 electrode was used only in cells prepared from mixed neptunium stock B. All measurements were carried out a t 25.00 f 0.02" in solutions of ionic strength 2.00 with substitution of sodium or lithium perchlorate for perchloric acid. The moles/l. concentration scale was used throughout and the potentials are given in absolute volts. The signs of cells and relative electrode potentials conform t.o the Stockholm C o n v e n t i ~ n . ~ Hydrogen electrode potentials were converted t,o the standard one atmosphere by conventional barometric, hydrostatic, and water vapor pressure corrections. Cell solution vapor pressures mere calculated by linear interpolation from osmotic coefficient data of pure 2 'VI perchloric acid and sodium or lithium perchlorate.1°
Soc., 83,
Results The results for each couplo are present'ed in terms of the formal potential E'. The following equations summarize first, the method of calculating E' from t8heobserved e.m.f. E and second the relationship of E' to the standard potent,ial Eo and the pertinent activity coefficient terms. Concentrations in moles/l. are indicated by brackets and mean molar activit,y coefficients by y. The 2~ sub-
(1960).
(9) J. A. Christiansen, ibid., 82, 5517 (1960). (10) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," Academic Press, Inc., New York, h-.Y., 1959, p. 483.
( 5 ) J. C. Sullivan, J. C. Hindman, and A. J. Zielen, J . Am. Chem.
3373 (1061). (6) S. Hietanen and I,. G. SiIIEn, Arkiv K e m i , 10, 103 (1956). (7) H. H. W~llardand P. Young, J . Am. Chem. Soc., 52, 557 (1930). (8) J. C. Sullivan, A. J Zielen, and J. C. Hindman, tbzd., 82, 8288
scripts h. 5, 6, 2 , 3 , I , and 2' refer, respectivrly, to ~, I/& of HC104, SpOsClO4, ? ; P O ~ ( U O ~ )I"e(C104)2, l"r(C10e)3,IIg,(C104)2,and Hg(ClO&. Throughout rvc hllvc abbreviated the familiar 2.30238 RTIF term by IC, and at 2.5' the value of k used was 0.059156 abq. v. XI)(T')-Sp(VI) E' = E k log [E1+1[Np(V)I/~Np(~'I~1 E' = Eoh,cv-vr)- h log 1/1,~g?/1/6~
+
FC2+LFC3+
(en) (GI,)
+ +
+
E' = E k log [H'][Fc2 ']/[F(,(III)] E'tnlir (it$) (71)) E' = l ~ i l ~ ~ j ~h log l r y- 1~, ~l 1~/ ~)~ / v , ~ I
