Absorption Spectra and Formation Constants of Some Metal Chelates of Quinoline-8-Selenol EllCHl SEKIDO,' QUINTUS FERNANDO, and HENRY FREISER Department o f Chemistry, University of Arizona, Tucson, Ariz.
b The formation constants of 1 : 1 lead(ll), cadmium(ll), and zinc(l1) complexes of quinoline-8-selenol (selenoxine) have been determined spectrophotometrically at 25" C. in a 50 v./v. dioxane-water medium a t 0.1 ionic strength. Although the values of these constants are smaller than those of the corresponding 8-mercaptoquinoline complexes, formation occurs from 0.6 to 1.0 pH units lower because of the higher proton displacement constants of selenoxine chelates. The spread in metal chelate formation constant values, of interest in reagent selectivity considerations, is about half that found for these metals with oxine and thiooxine. Both the cadmium and the lead complexes can b e transformed to unusual, protonated, MHL+2, species.
70
T
deals with the determination of a series of metal complex formation constants of the recently developed chelating agent, quinoline-8selenol (selenoxine), (4,5). Such data are of interest not only in characterizing the analytical behavior of a new reagent but also in a study of the role of donor atom properties on metal complex stability. HIS REPORT
EXPERIMENTAL
Stock Solution of Selenoxine. Pure 8,8'-diquinolyl diselenide (20.17 mg.) was dissolved in 1.0 ml. of 50Tc hypophosphorous acid and 1.0 ml. of 2M perchloric acid and then added to 20 ml. of water. The solution was warmed on a water bath for about 10 minutes t o reduce the diselenide t o selenoxine. The solution was diluted t o 25.00 ml. with water deaerated with nitrogen and then nitrogen gas was bubbled through the solution for 10 minutes, The stock solution of the reagent prepared by this method and stored in a glass-stoppered bottle was stable for a t least a week in the dark. The concentration of this stock solution was4.0 X 10-3M. Other Reagents. The 1,4-dioxane was purified as described previously (2). Stock solutions (about 0.8M) of metal perchlorates were prepared by dissolving them in deionized water.
Present address, Department of Chemistry, Kobe University, Kobe, Japan. 1
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ANALYTICAL CHEMISTRY
Their concentrations were determined by a gravimetric method using oxine or by a titrimetric method using EDTA. Apparatus. Spectrophotometric measurements were made with a Beckman Model D B spectrophotometer. All pH measurements were made with a Beckman Model G pH meter equipped with a glass-saturated calomel electrode pair and calibrated at suitable intervals with Beckman buffers. Measurement of Absorption Spectra of M e t a l Chelates of Selenoxine in 50 v./v. % Dioxane-Water Media. I n view of the very low pH region in which these chelates formed (4),their formation constants could not be determined potentiometrically and a spectrophotometric method was used to determine the formation constants. Since the 1 : 2 chelates of selenoxine are very insoluble even in 50% v./v. dioxane-water solutions, the metal ion concentration was maintained in very large excess of that of the reagent (100: 1). The equilibrium data obtained under such conditions give information on the formation constants of the 1: 1 metal chelates. Dioxane (12.50 ml.) was placed in each of a series of 25-ml. volumetric flasks and then aliquots of 11.5, 6, or 2Jf perchloric acid were added to make the hydrogen ion concentrations of the solutions increase from 0.01-1f t o 4.551. Next, the stock solution of the metal perchlorate \\as added to make the metal ion concentration 3.2 X 10-2M11. Then 0.93 ml. of 1.0031 sodium hydroxide solution was added to each flask because it \vas found expedient to compensate for the acid present in the reagent solution. Each solution was bubbled with nitrogen gas for 10 minutes. Finally, 2.00 ml. of the reagent solution was added into each flask, and then each solution was diluted to the mark with water deaerated with nitrogen. After a few hours standing in the dark, each solution was placed in a 1em. quartz cell (whose stopper was coated with silicone grease) under nitrogen gas and the absorption spectrum was recorded immediately using as reference solution, one which contained all components except the reagent. The measurements were carried out a t 25" f '1 C. Determination of Formation Constant of Zinc Selenoxine Chelate. The absorption spectra of the zinc selenoxine chelate in 50Tc v./v. dioxane-water media are shown in
Figure 1. The absorption maximum of the zinc selenoxine chelate is found at 387 mp, which is near the isosbestic point, 397 mp, exhibited by the reagent in acid solutions between pH -1.0 and 3.0 (Figure 2). The absorbance values of the metal chelate solutions were taken a t this isosbestic point a t various hydrogen ion concentrations and plottkd against pH (curve B , Figure 3). The formation constant for the zinc chelate is calculated as shown below. The acid dissociations of the ligand [the ligand exists almost entirely as zwitterion H+L- (5)]are HzL+ e H+L-
+ H+ [H+L-] [H +] [HzL+l
K,, =
H+L-
(1)
+ H+
e L-
and the equilibria for the formation of the metal chelate
JI+'
+ HzL+ e ?vIL++ 2H+ K1
N+z
+ H+L- F? ML+ + H + K,'
AI+'
[hlL+][H+Iz [M+'] [HZL+] (3)
=
=
[?vlL+][H+] [M+2][H+L-] (4)
+ L- e ML+ K
[ML+]
'l
- [M+*][L-] (5)
The absorbance a t 397 mp, A3g7Jis expressed by : A397
=
+
[hIL+] [H+L-] ~ 3
€397XL+
+
€397H+L-
x
[HzL+] (6)
9 7 ~ ~ ~ '
where ~ ~ ~ ~, h Ef~ L~ , H+ + L -and E ~ ~ ~ Hare ~ L the molar absorptivities at 397 mp of the 1: 1 metal chelate, ligand (zwitterion form), and protonated ligand, respectively. At 397 mfi, the isosbestic point, eSg7H+L-
=
E
~
~
~ =H E397 + R~
thus A397
=
€397ML+
[ML']
+ + [HzL+l)
([H+L-l
~ 3 9 7 ~
(7)
+
Figure 2. Absorption spectra of selenoxine in 50% aqueous dioxane media at various molarities of HClOi [Reagent] = 3.2 X 10-4M
Figure 1. Absorption spectra of 1 : 1 zinc chelate of selenoxine in 50% v./v. aqueous dioxane media containing varying molarities of HClO4 [Znfa] = 3.2 X [Reagent] 3.2
X 10-4M
With TM and TI, representing the concentrations of total metal ion and total ligand, respectively, T M = [M+’]
TL = [NIL+]
+ [NIL’]
Kf,= Kit - or log K,,
(8)
+ [H+L-] +
K1’ T L T M ( E ~ # ~-+E d ) K.,[H+Iz IH+l KI * TY an' * TL (10)
+
+
d
(11)
+
When the values of K O ,= 0.759, TL = 3.2 x 10-‘M, Tar = 3.2 x 10-2AM, eag7R = 484, c*#L+ = 2590 and observed Asg7and [H+]are substituted in Equation 10, the value of K1’ is obtained, and log K,, is derived from the relationship:
I
I
I
-04
0
0.4
I 0.8
I
u
I
I
Id
10
PH
The values of these constants for the zinc chelate are shown in Table I.
Figure 3. Plots of absorbances at 397 mp of 1 : 1 metal chelates of selenoxine as a function of pH
Determination of Formation Constant of Cadmium Selenoxine Chelate.
When T M : T L= 100:1, Equation 8 can be approximated by TM = [M+’]. From Equations 1, 4, 7, 8, and 9, the following Equation 10 can be derived. =
+ log K1’
KO,
[ H A + ] (9)
A391
= pK,,
The absorption spectra of the cadmium selenoxine chelate in various hydrogen ion concentrations in 50% dioxane-water solution are different from those of the zinc selenoxine chelate (Figure 4). That is, the absorption curves in the range of 0.010.05M hydrogen ion concentration, have an absorption maximum a t 380 mp (E = 2.6 x l O 3 ) , and indicate complete chelate formation. At higher hydrogen ion concentrations, however, the absorption bands shift to longer wavelengths and their intensities become lower. The isosbestic points occur a t 330 mp and 412 mp, In 1M acid solution the absorption maximum occurs at 390 mp (E = 1.7 X 103). Further
[Metal ion] = 3.2 X 10-2M [Reagent] = 3.2 X 10-4M A. l e a d chelate E. Zinc chelate C. Cadmium chelate
increase in acidity results in a shift of the absorption band to shorter wavelengths with the ultimate elimination of this band. An isosbestic point occurs in this acidity range a t about 365 mp. These spectral changes in high acidity cannot be attributed to dissociation of the chelate because the spectra of the ligand in the absence of metal ion under these conditions are
Formation Constants and Proton Displacements Constants for 1 : 1 Metal Chelates of Selenoxine, Thiooxine, and Oxine in 50% v./v. Aqueous Dioxane
Table 1.
Pelenoxine chelate log Kj,5 log K1‘ log K1 10.42 1.92 1.80
Thiooxine chelate Oxine chelate log Kf log KI’ log Ki log K/ log K1’ log K1 Pb 11.52b 2.32 0.58 10.61d -0.93 -4.90 11.850 2.64 0.90 Zn 10.17 1.67 1.55 ... ... 10.896 1.62 -0.12 9.96d -1.58 -5.55 11.050 1.85 0.11 Cd 10.49 1.99 1.87 2.10 -0.23 1O.7ga 1.59 -0.15 9.43d -2.11 -6.08 a Reliability of log Kj,, log KI’, log K1,log KI” and log K1”‘ is estimated to be within fO.l. Reference ( 1 ) . Reference ( 6 ) . d Reflog Ki” 2.20
log Ki”’ -0.40
0
erence (3).
~~
~~
VOL. 37, NO. 12, NOVEMBER 1965
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I
*
D
365
$12 -
,,‘,* *,
e ,,
0-
,,’,,/
A// /‘,*
,./ .../,..*- __--.--/---
,/; __,--
, ,.;:;:,,‘.”..’’. ... _..-, I
,e
2f::/..---*I
E-
l
I
I
1 thus lH+I
I
I
I
=
- €ML+) - _ 1_ . (A - € M L + TL) K1”’ (16)
T~(€HYL+*
K1”’
Plots of 1/(A - € X L + TL) US. I / [ H + ] as shown in Figure 5 should be linear and the intercept with the ordinate corresponds to -l/K1”‘. From measurements a t a series of appropriate wavelengths, one can obtain a more reliable value of K1”‘. The value of K1”, the equilibrium constant in Equation 17 is obtained as follows:
Figure 4. Absorption spectra of 1 :1 cadmium chelate of selenoxine in soy0 v./v. aqueous dioxane media containing varying molarities of HClOi [Cd+z] = 3.2 X 10-W [Reagent]
= 3.2 X lO-‘M
entirely different. For example, in liM perchloric acid, if complete chelate dissociation occurs, the absorbance a t 450 mp should be below 0.085. The absorbance observed a t this wavelength, however, is 0.22. The absorption spectra of the solutions which are greater than 4 X in perchloric acid become similar to that of the free ligand in strong acid solution. The relationship between the absorbance at 397 mp and pH is shown in Figure 3. Obviously the curve for the cadmium-chelate is different from that for the zinc. The curve seems to be composed of two steps which suggests two reaction processes. Furthermore, when the same calculation used for the zinc chelate is employed, the values of K,’ do not remain constant. These results seem to show that another species (probably HML+Z) exists in 1M-2M perchloric acid solution as shown below :
If it is assumed that the absorption maxima a t 380 mp and 390 mp correspond to the species ML+ and HML+2, respectively, equilibrium constants K1” and Kl”’ may be calculated. I n such calculations, the contribution of ML+ as well as H2L+ and H M L f 2 to the absorbance a t 390 mp must be taken into account. This is particularly true in the pH range 1.5-0.0 because below this pH range the chelate is completely protonated. On the basis of these considerations, the equilibrium constant, Kl”’, is obtained, using the following equations: HML+2 S M L +
+ H+
H2L+
+ hI+’ E HAIL+’ + H + Kl“ = [HML+z][H+] (17) [H*L+][ M + 2 1
Thus log K1” = log
[HML+2] [HzL+] -t pM pH (18)
~
-
The pH is read from the curve which shows the relationship between the absorbance and pH, when the absorbance is (A397HbfL+2 A397H2L+), in other words, whenlog [H,L+]/ [HML+2] is equal to zero. A397HML+a and ABg7%L+ are the absorbances a t 397 mp of the reagent which exist as HML+Z and H2L+. These values are calculated from the concentration of reagent and and tB97H2:+, respectively. The value of K,’ is derived from
+
and KI, is obtained from Equation 11. The results for the cadmium selenoxine chelate are shown in Table I. A represents the absorbance a t the appropriate wavelength and and e M L + represent the molar absorptivities at this wavelength for HML+2 and ML+ species, respectively. Equations 13 and 15 are substituted in Equation 12 to give
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ANALYTICAL CHEMISTRY
Determination of Formation Constant of Lead Selenoxine Chelate. The absorption spectra of the lead selenoxine chelate in various dioxanewater solutions are shown in Figure 6 and the relationship between the absorbance a t 397 mp and pH is shown in Figure 3. The lead chelate has an absorption maximum a t 383 mp above pH 2.0. When the hydrogen ion concentration increases from 0.01M to 0.15M this absorption band shifts to
slightly longer wavelengths with little change in the absorbance a t 397 mp. The increase in the absorbance in the range of 400 550 mp is large, similar to that observed with the cadmium chelate. Isosbestic points occur a t about 369 mp and 400 mp. Calculations similar to that carried out for the cadmium chelate give results which are shown in Table I. N
DISCUSSION AND RESULTS
+
The values of (pK,, pKa2) of the three reagents decrease in this order, oxine (15,51), thiooxine (10.94), and selenoxine (8.62) ( 5 ) . On the basis of proton affinity, the stability of their chelates would be expected to decrease in this order. The formation constants ( K f 1,) however, of the zinc(I1) , cadmium(II), and lead(I1) thiooxine chelates are greater than those of the corresponding oxine chelates. This fact may reflect the greater degree of covalent character in the metal-sulfur bond ( 1 ) . The formation constants of the cadmium (11) and zinc(I1) selenoxine chelates occur between that of corresponding oxine and thiooxine chelates and the formation constant of the lead(I1) selenoxine chelate is smaller than that of the thiooxine and oxine chelates. The covalent character of the metalselenium bond is probably the same or slightly greater than that of the metalsulfur bond inasmuch as the electronegativities of selenium and sulfur are almost the same. The proton affinity pK, = 8.62) is of selenoxine (pK., considerably smaller than that of thiooxine (10.94). Hence this latter factor would be of greater importance in comparing the K,, value of metal chelates of selenoxine and thiooxine. The spread of formation constant values, of interest in reagent selectivity considerations, is about half that observed for the same metals with oxine and thiooxine. A comparison of the pertinent proton displacement constants of the metal complexes of selenoxine and thiooxine gives the information of analytical interest, that those in the former series will form a t pH values from 0.6 to 1.0 unit lower. This pH advantage is not as significant as that resulting in the change from oxine to t hiooxine. It is interesting to compare the value of pK, for the cadmium(I1) and lead(I1) selenoxine chelates with the pK, of the neutral form of selenoxine and the pK, of its zwitterion form.
Figure 6.
Absorption spectra of 1 :1 lead chelate v./v. aqueous dioxane media containing varying molarities of HCIO4
of selenoxine in 50%
[Pbf2] = 3.2 X lO-'M [Reagent] = 3.2 X 10-'M
+
c
Se-
Se-
pK,- 8.50
SeR
pK,-
Cd (Pb' 1
1.90(R:H or CHS)
Pb*) pK1=.O.23 (for Cd) 0.40 (for Pb)
protonated metal-selenoxine, is seen to result in a further but smaller pK reduction. It must be kept in mind that this is not a direct measure of the further change in basicity of the nitrogen atom since the N-H bond is replaced If the formation by N-Metal. constant corresponding to the NMetal bonds in these complexes is similar to that formed in the corresponding pyridine-metal complexes, then the behavior of both zinc and lead seem inexplicable. Because zinc forms about as stable a pyridine complex as does cadmium, one would have expected the zinc to form a protonated selenoxine complex; lead, on the other hand, would not be expected to.
It can be seen that the value of pK corresponding to the dissociation of
\ / N is I
quite sensitive to the charge
H+ on the selenium atom. That is, neutralization of the selenium anion (--Sew) by a proton or methyl group results in a remarkable decrease in the pK value, from 8.50 to 1.90, reflecting a great decrease in the basicity of the nitrogen atom. The effect of the change in the charge from the anion to the ether species must contribute significantly in the pK drop. Further increase in the charge of the species, as in the
LITERATURE CITED
(1) Corsini, A,, Fernando, Q., Freiser, H., ANAL.CHEM.35,1424(1963). (2) Freiser, H.,Charles, R. G., Johnston, W. D., J. Am. Chem. SOC.74, 1383 (1952). (3) Johnston, W. D.,Freiser, Ibid., p. 5239. (4) Sekido, E., Fernando, Q., Freiser, H., ANAL.CHEM.35, 1550 (1963). (5)Zbid., 36, 1768 (1964). (6) Sekido, E.,Ueda, M., Preprint, 1965
Annual Symposium, Analytical Chemistry in Japan;
RECEIVED for review July 9, 1965. Accepted August 24, 1965. Work supported by the U. S. Atomic Energy Commission. VOL. 37, N O . 12, NOVEMBER 1965
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