mass spectrometry, spectrophotometry, and various electroanalytical measurements. Technological Purifications. The separations technologist should find profitable applications for this method in t h e recovery and purification of americium isotopes produced in various transplutonium processes. T h e relatively high radiation stability of this simple inorgaric system would appear t o offer definite advantages over the currently-used orgmic ion exchange and solvent extraction systems. The extremely high separiition factors described above for Am/Cm, Bk, Cf, lanthanides suggest, immediate application both b y the radioisotope processor and the heavy element chemist. ACKNOWLEDGMENT
The author is indebted to J. H. Cooper and C. J. Coley for capable assistance in some of the tests.
LITERATURE CITED
(1) Asprey, L. B., Stephanou, S. E., Penneman, R. A., J. Am. Chem. SOC. 72, 1425 (1950). (2) Bubernak, J., Lew, M. S., Matlack, G. M., ANAL.CHEM.30,1759 (1958). (3) Bubernak, J., Lew, M. S., Matlack, G. M., Ta2anta 6 , 167 (1960). (4)Caro, H., 2.Angew. Chem., 845 (1898). (5) Elbs, K., Schonhen, O., 2. Elektrochem. 1,468 (1895). (6) Hyde, E. K., Proc. Intern. Conf. Peaceful Uses At. Energy, Geneva, 1955, 7, Paper 728, p. 281, United Nations, New York (1956). (7) Kolthoff, I. M Miller, I. IC, J . Am. Chem. SOC.73, 3&5 (1951). (8) Koshland, D. E., U. S. Atomic Energy Commission Declassified Report, CN2041 (1945). (9) . . Moore. F. L., ANAL.CHEM.28, 997 (1956). ’ (10) Moore, F. L., Ibid., 33, 748 (1961). (11) Moore, F. L., Hudgens, J. E., Jr., Ibid., 29,1767 (1957). (12) Moore, F. L., Lyon, W. S., U. S . Atomic Energy Commission Declassified Report, ORNL-1101 (1951). (13) Orr, P. B., U. S. Atomic Energy Commission Unclassified Reaort. ORNL-3271 (1962). A
,
(14) Palme, H., 2. Anorg. Chem., 112, 97 11920). (15) Peppird, D. F., Mason, G. W., in
NAS-NS-3006, 34 (ref. 16). (16) Penneman, R . A., Keenan,.T. K., “The Radiochemistry of hmericium and Curium,” NAS-XS-3006 (1960). Available from the Office of Technical Services, Department of Commerce, Washington 25, D. C. 7) Pressly, R. S., U. S. Atomic Energy Commission Declassified Report, ORNL-2202 (1957). 8) Stevenson, P. C., Nervick, W. E., “The Radiochemistrv of the Rare Earths, Scandium, Y t k u m , and Actinium,” NAS-NS-3020 (1961); Available from the Office of Technical Services, Department of Commerce, Washington 25, D. C. (19) Thompson, S. G., Ghiorso, A., Seaborg, G., Phys. Rev. 80,781 (1950). (20) Thompson, S. G., Morgan, L. O., James, R. A., Perlman, I., “The Transuranium Elements,” NNES IV-14B, pp. 1339-61, McGraw-Hill, New York, 1949. (21) Ward, M., Welch, G. A., J . Chem. SOC.1954,4038. RECEIVED for review Sovember 28, 1962. Accepted February 21, 1963.
A S pec t rQ photo met ric a nd Pote nt io rnet ric St udy Of
Certain Metal Chelates of Pyridine-2,6-dialdoxime
SASWATI
P. BAG, GUINTUS FERNANDOf
and HENRY FREISER
Department o f Chemisfry, University of Arizona, Tucson, Ark.
,This investigation was carried out to evaluate the metal chelating properties of pyridine-2,6-~dialdoximewhich forms highly colored metal complexes that may be used for the spectrophotometric determination of a number of metal ions. The acid dissociation constants of pyridine-2,6-diaIdoxime and its chelate formation constants with Cu(ll), Ni(ll), Zn(II), and Mn(ll) have been determined. A study of the composition of the chelates in aqueous solution showed that Cu(ll) formed a metal chelate in which the meta1:ligand ratio WIYS 1 :1 whereas 1 :2 chelates were formed with Ni(ll), Zn(ll), Mn(ll), and Co1:Il). In the chelates of pyridine-2,6-dialdoxime, one of the oxime groups in each ligand molecule is intact and therefore the influence of the metal ion on proton release from this oxime group could b e studied. The metal ion had an acid-strengthening effect on the oxime group and this effect was directly proportional to the stability of the metal chelate.
A
a-dioximes, a most useful family of chelating agents, have received a great deal of attention ( I , 6,16) it is only recently t h a t LTHOUGH
metal chelation reactions of other reagents containing the oxime group have been studied (IS). For example, methyl-2-pyridyl ketoxime has been proposed as an analytical reagent for the spectrophotometric determination of Cu(1) and Fe(I1) (a), and phenyl-2 pyridyl ketoxime for the determination of iron and palladium (14, 15); pyridine-2-aldoxime forms intensely colored chelates with many transition metal ions and the properties of these complexes have been investigated recently ( 4 , 11). Pyridine-2,B-dialdoxime could act as a tridentate ligand and a study of its metal chelates should prove useful in determining the utility of this compound as an analytical reagent.
2,6-dialdoxime were determined potentiometrically at 25 + 0.1’ C. A weighed quantity of t h e compound was transferred t o a water-jacketed vessel a n d dissolved i n a measured volume of water. A known a m o u n t of HClOd was added t o t h e solution for t h e determination of pK,,. Nitrogen gas, saturated with water, was passed through the solution throughout the course of the titration with carbonate-free NaOH. All p H measurements were made with a Beckman Model G pH meter equipped with a glass-saturated calomel electrode pair and standardized with buffer solutions at pH 4.00 and 9.00. The first acid dissociation constant K,, was calculated from the experimental lsoints in the first buffer region in which only K,, was involved.
EXPERIMENTAL
where T R represents the analytical concentration of pyridine-2,B-dialdoxime. The second and third dissociation constants K,, and K,, were calculated as follows since the two buffer regions in which the second and third protons were released, showed some overlap. Values of [ ~ were ~plotted ~ against S [H+l (S- 1)
Compounds. Pyridine-2,6-dialdoxime was obtained from K and K Laboratories, New York. T h e compound was decolorized with charcoal a n d recrystallized from water t o give colorless needles (m.p. 209’ C., lit. m.p. 211.5’ C.) (12). Acid Dissociation Constants. T h e acid dissociation constants of pyridine-
-
(S- 2) VOL. 35, NO. 6, MAY 1963
719
where
In titrations in which HC104 was not added [C104-] = 0. The slope of the straight line obtained is equal to K,, and the intercept to K,,K,,. The potentiometric titration data in the overlapping buffer regions were also treated as follows: the function, p , which ia equal to the average number of protons bound to the most highly charged anionic species of pyridine-2,fi-dialdoxime, was plotted against pH. P=
When p = 0.5, 1.0, and 1.5, the corresponding p H values rvere equal to PK,, l / d ~ ~ K a , pK,J and pH,,, respectively. The averaged values of pK,,, pK,, ??d pK,, are given in Table I. Composition of t h e Metal Complexes. The composition of the Cu(II), S i ( I I ) , Co(II), Zn(TI), and hfn(I1) chelates of pyridine-2,6-dialdoxime \vas determined spectrophotometrically by Job's method of continuous variations (10). A series of solutions mi: prepared containing metal ion and ligand in ,varying quantities and the sum of their concentrations was kept constant a t 5.65 X 10-4X. The ionic strengths of all the solutions were maintained constant a t 0.1 with KaC104. The absorbances of the solutions were meawred a t fixed wavelength and pH, and plotted against the mole fraction of the metal ion. The results obtained are summarized in Table 11. These results
+
Table I.
HL- + M + 2 a M(HL)+ HLSI(HL)" e I\I(HL)z
(3) The stepwise formation constants kl and k2 are given by:
PKSI 5.1s 3.55 2.34 ( p ...
pK,,
pKas ...
=
0,005)
iO:i7 9.7 ( p 10.7
=
0.01)
10:7 ( p ...
=
0.01)
Composition of Metal Chelates of Pyridine-2,6-dialdoxime Determined by Method of Continuous Variation and by Mole Ratio Method
Metal ion
Color of
chelate
Cu(I1)
Green
Co(I1)
Yellow
Mn(I1)
Blue violet'
Ni(I1)
Pale yellow
Zn(I1)
Colorless
Wavelength of absorbance measurements ( m p ) (1% e ) 600(2.26)400(3.33)360(3.80) 630 (2.30) 400 (3.22) 360 (3.75)
610 (2.6)' 400 (3.58) 330 (3.72) 330 (3.77) 330 (4.01) 590 (2.48) 590 (3.05) 590 (3.13) 330 (3.47) 330 (3.55)
330 ( 2 . 8 1 ) 340 (3.17) 340 (3.30)
340 (3.42)
720
(2)
+
Acid Dissociation Constants of Pyridine-2,6-dialdoxime and Related Compounds at 25' h 0.1" C. in Water
Compound Pyridine (9) Pyridine-%aldoxime (8) Pyridine-2,6-dialdoxime Benzaldoxime (6)
Table II.
were verified by the mole ratio method (17). The absorbances of a series of solutions containing a fixed amount of metal ion (5.65 X 10-4JI) and a varying amount of ligand (5.65 X 10-5 to 5.65 X 10-3M) were measured a t fixed wavelength and pH, and plotted against the number of moles of ligand added. The results, identical for both methods, showed that the chelates of Zn(II), Mn(II), Co(II), and Si(I1) have a meta1:ligand ratio of 1:2 and the Cu(I1) chelate a meta1:ligand ratio of 1:l (Table 11). The spectra of all solutions were obtained with a Cary Model 11 or Beckman Model DB recording spectrophotometer employing 1-em. stoppered cells. Potentiometric Titrations in the Presence of Metal Ions. -4 weighed quantity of pyridine-2,6-dialdoxime together with 5 ml. of a standard metal perchlorate solution and 100 ml. of mater was introduced into a water-jacketed titration vessel maintained a t 25' + 0.1' C. The metal: ligand ratios were 1:2, l:5, and 1 : l O . Presaturated nitrogen gas was passed through the solution throughout the course of the titration. Carbonate-free SaOH m s added from a buret and pH measurements were made with a Beckman Model G p H meter. If the pyridine-2,6-dialdoxime is represented by H2L, the chelate formation reactions may be represented as follows:
ANALYTICAL CHEMISTRY
pH 3.8
5.5 7.5 3.0
4.0
6.7 3.8 7.8 9.0 3.0 4.0
Met a1: ligand ratio 1:l 1:l 1: 1
1:2 1:2 1:2
1:2 1:2 1:2 1:2 1:2
6.7 3.0
1:2
7.8
1:2
4.0
1:2 1:2
'
=
[lI(HL)-] [HL-] [nl+Z] and
The potentiometric titration data from the first buffer region were used to evaluate log kl and log k2 for the ZntII), l I n ( I I ) , and Ki(I1) chelates from plots of 5,the formation function us. p(HL-).
[Xa+] -
[H+]
+ [OH-])(/
W-1
+ [OH-]).
T M and Tx represent analytical concentrations of the total metal and ligand, respectively. Formation constants could not be calculated from the 130tentiometric titration data for the Cu(II) and Co(I1) chelates since proton release in the first buffer region was not solely due to the chelate formation process. Values of log L1and log k2are collected in Table 111. The reliability of these values is approximately 1 0 . 1 . The potentiometric titration data in the second buffer region, in the titration curves of the Zn(I1) and 1\In(II) chelates \vere used to calculate the acid dissociation constants K,' and K," of the unchelated oxime groups on the assumption that no chelate hydrolysis occurred in this region. Since the two buffer regions in the titration curve of the Xi(I1) chelate overlapped, only approximate values of the acid dissociation constants of the unchelated oxime groups could be calculated in this case. Corresponding values for the Cu(I1) and Co(I1) chelates could not be obtained for reasons mentioned below. Potentiometric titrations n-ere carried out on solutions containing a 1:2 meta1:ligand ratio. The titration curve of the Zn(I1) chelate had two well defined buffer regions. I n the first buffer region two protons per metal ion were released on chelate formation; in the second buffer region, further proton release from the metal chelate occurred. The titration curve of the lIn(1I) chelate resembled that of the Zn(I1) chelate and although the two buffer regions were less distinct, it was found that in the first buffer region two protons per metal ion were released on chelation. The two buffer regions overlapped in the titration curve of the Ki(I1) chelate and i t was not possible to determine the number of protons released on chelate formation. I n the titration curve of the Cu(1I) chelate, two protons per metal ion were released in the first buffer region; a second poorly defined buffer region n-as also present in which further proton release (greater than one proton per metal ion) from the metal chelate occurred. The titration curve for the Co(I1) chelate was unique since i t showed a single buffer region in which three protons per metal ion were released.
The potentiometric titration curves were repeated with solutions containing a meta1:ligand ratio of 1:5 and 1 : l O . The Zn(II), Mn(II), IJi(II), and Cu(I1) chelates showed the same behavior as they did when the rneta1:ligand ratio was 1:2. However, in the titration of the Co(I1) chelate, four protons were released in the first buffer region. Spectrophotometric: Determination of Formation Constmts. Values of log kl for t h e chelates of Cu(I1) a n d
S i ( I 1 ) were determir ed spectrophotometrically in soluticins containing a constant concentraticn of the ligand (1.13 x 10-6M) and a 100-fold excess of metal ion (1.14
