An Investigation of the Dissociation Phenomena of Quinoline-8-seIenoI in Water and Aqueous Dioxane EllCHl SEKIDO,' QUINTUS FERNANDO, and HENRY FREISER Department o f Chemistry, University o f Arizona, Tucson, Ariz.
b The synthesis of a new chelating agent, quinoline-8-selenol (selenoxine), has been described. The acid dissociation constants of selenoxine have been determined spectrophotometrically both in water and in 50% v./v. dioxane-water. In addition, the acid dissociation constants of the Se-methyl derivative of selenoxine and of the S-methyl derivative of quinoline-8thiol (thiooxine) were determined. These data were used to elucidate the dissociation phenomena of oxine, thiooxine, and selenoxine in aqueous media. The extent of zwitter ion formation in solutions of these reagents has been quantitatively evaluated.
0
PARTICULAR IXTEREST in the investigation of the analytical reagent 8-mercaptoquinoline (thiooxine) is the study of the effect of the change of the Group VI bondingatom (from oxygen to sulfur) on the physical and chemical properties of the reagent (4). For example, the acid dissociation constants of thiooxine are significantly greater than those of oxine. It would be of interest to learn whether a similar increase in acid strength would be observed when sulfur is replaced by selenium. Furthermore, as the acid strength of one of the functional groups increases, it might be expected that the possibility of zwitter ion formation would become more significant. We have therefore determined spectrophotometrically, the acid dissociation constants of quinoline-8-selenol (selenoxine). These results have prompted a re-examination of the role of znitter ion formation in the dissociation phenomena of oxine and thiooxine as well.
F
EXPERIMENTAL
Synthesis of Selenoxine. The following reaction sequence was developed for the synthesis of selenoxine and 8,8'-diquinolyl diselenide.
I
I
N=N@
NHz
1 Present address, Department of Chemistry, Kobe University, Kobe, Japan.
1768
0
ANALYTICAL CHEMISTRY
air. When a small quantity of hydrogen peroxide is added to the filtrate with stirring, a white precipitate of 8,8'-diquinolyl diselenide forms rapidly, and is filtered, washed, and dried at 110" C. The product is best purified by dissolving it in HCl and hypophosphorous acid and repeating the abovementioned procedure. Yield, 12%; m.p., 205 " to 206 " C. Calcd. for C18HI2N2Se2 : C, 52.2; H, 2.9; Y, 6.8; Se, 38.1. Found: C, 51.5; H , 3.0; K, 7.2; Se, 38.2. Quinoline-8-selenol
I n view of the sensitivity of selenoxine
to atmospheric oxidation, the easily produced diselenide was made in quantity. This compound, from which the selenoxine can be conveniently obtained, is stable indefinitely. 8 - Selenocyanatoquinoline. Three grams (0.021 mole) of 8-aminoquinoline are dissolved in 100 ml. of 0.l.U H2S04. T h e orange solution is cooled below 5' C. in an ire bath and diazotized with a 5 M sodium nitrite solution to the starch-iodide end point. The diazotized solution is carefully neutralized with 30y0 NaOH while the solution is maintained below 5" C. Potassium selenocyanate (IO) (3 grams or 0.021 mole) and NaHCO3 (5 grams) are dissolved in 50 ml. of water and cooled below 5" C. in an ice bath. The cold, filtered diazonium solution is added dropwise from a dropping funnel into the potassium selenocyanate solution through nhich a stream of nitrogen gas is maintained. With the addition of the diazonium solution, a yellow-orange turbidity occurs which gradually turns into a brown-black tarry product upon further addition of the diazonium solution. This product, though impure, may be used in the next step of the synthesis. 8,8'-Diquinolyl Diselenide. Impure 8-selenocyanatoquinoline (0.20 gram) is dissolved in 6M HC1 and 2 ml. of hypophosphorous acid is added. The solution is heated on a water bath for about 10 minutes. The color of the solution changes from yellow to orange-red. After cooling, the solution is filtered, and then neutralized with 5M XaOH in an ice bath while nitrogen gas is bubbled through the solution. During neutralization, the color of the solution changes to redviolet and then to yellow in alkaline solution. To remove the impurities, the alkaline solution is filtered rapidly under nitrogen gas. The filtrate develops a m hite turbidity on contact with
(Selenoxine).
Purified 8,8'-diquinolyl diselenide (0.1 gram) is dissolved in 3.0 ml. of 6 M HC1 and 1.0 ml. of 50% hypophosphorous acid. I n order to complete the reduction of 8,8'-diquinolyl diselenide, the solution is heated under nitrogen on a water bath for about 10 minutes. After cooling, the solution is filtered j f necessary and kept under nitrogen, while 5M XaOH is added dropwise. The orange-red solution becomes deep red and then violet-brown needles gradually precipitate. The precipitate is rapidly filtered under nitrogen gas, and is washed with oxygen-free cold water. Selenoxine is dried and kept in a desiccator filled with nitrogen gas. Zinc Selenoxinate. .1 slight excess of freshly prepared selenoxine is added to 50 ml. of zinc perchlorate solution a t 60' C. under nitrogen gas. The mixture whose p H is adjusted to approximately 2 is allowed to digest for 15 minutes on a water bath. T h e yellow precipitate obtained is filtered, washed thoroughly with hot wat'er, and dried at 110" C. wit,hout any special precautions. Calculated for ClsH12S2ZnSe-C, 45.1; H, 2.5; N, 5.8. FoundC, 45.1; H, 2.7; N, 6.1. 8-M ethylselenoquinoline. 8,8'Diquinolyl diselenide (83 mg. 0.2 mmole) is dissolved in a mixture of 20 ml. of 1Jf HCI and 2.0 ml. of 50% hypophosphorous acid and t h e resulting solution is heated under nitrogen for a short time on a water bath. After cooling, sufficient 5M NaOH is added t o the solution under nitrogen t'o make the resulting solution I M in NaOH. This solution is treated, with 60 mg. (0.43 mmole) of methyl iodide in a flask fitted with a reflux condenser and stirred a t room temperature for 2 hours. The yellow solution becomes turbid and faint yellow crystals form. After filtration the precipitate is washed thoroughly with water, recrystallized from 1Jf HC1, reprecipitated, and dried in air. Yield, 70%; m.p., 59.5" to 60" C.
1
Selenoxine
0
Selenoxine
Thiooxine 0.A
I Oxine
I
I
I
I
I
58
54
50
I
I
I
Frequency x
d3
Figure 2. Relationship between absorption frequency and dielectric constant of the solvent 1.
2. 3. 4.
5.
Figure 1 . Absorption spectra of 2 X 10-4M solutions of selenoxine, thiooxine, and oxine A. 6. C.
8.
9. 10.
CHxOH C~HF,OH isa-C3H,OH n-CdH90H f-C4HgOH
Absorption Maxima in Visible Region of Selenoxine and Thiooxine in Organic Solvents
Selenoxine Ui-
Solvent Carbon tetrachloride Benzene Chloroform Methvlene chlGride &Butyl alcohol n-Butyl alcohol Isopropyl alcohol Acetone Ethanol Methanol Nitromethane Water a Reference ( 3 ) . Reference (8). Table II.
maximum
it was decided to use the 400-mu band for absorbance measurements in aqueous solution and the 498-mp band for 507, aqueous dioxane media. The values of pK,, and pK,, (Table 11) were determined from a plot of absorbance us. pH. Because pK,, is so low, it was not possible to maintain the ionic strength constant a t 0.1.
precautions to prevent contact of the solutions with air. From an examination of the spectra (Figure 1, Table I)
Table I.
6. 7.
Hz0 CHaNOz CH3COCHa CHzCIz CHCls
Strong acid solutions N e a r isoelectric point Alkaline solutions
Reagents. 1,-l-Dioxane was purified as described previously (6). A11 organic solvents used in spectrophotometric measurements were Spectrograde reagents wherever available. Other solvent,s were suitably purified. Buffer components were Analyzed Grade. Apparatus. Absorbance measurements were made with a Beckman Model DU spectrophotometer. All p H measurements were made with a 13eckman Model G p H meter equipped with a glass-saturated calomel electrode and standardized with a Beckman buffer solution a t p H 4.00. p H values in the vicinity of zero were calculated from total acid concentration. Spectrophotometric Determination of Acid Dissociation Constants. h l l acid dissociation constants reported here for selenoxine and its Se-methyl derivative were determined spectrophotometrically a t 25' i 1' C. The macroscopic (dissociation constants for selenoxine were determined as follows : 8,8'-diquinolyl diselenide (20.17 mg.) was dissoslved in 1 ml. of 507, HJ'02 and 1.0 ml. of 2.0011.1 HC1 and then diluted to exactly 25 ml. with water. Exactly :2 ml. of this solution were added to 20 ml. of solutions having various pH values and where appropriate, sufficient dioxane to make the final solution 50Tj v./v. in dioxane, as well as XaC104 to maint>ainthe ionic strength a t 0.1. Each of the mixtures was made up to the mark in a 25-ml. voluniet'ric flask and deaerated with nitrogen for about 10 minutes. ;\liquots of the solutions were transferred to the spectrophotometric cells with suitable
I
I
66
62
electric constant 2.2 2.28 4.81 9.1 10.9 17.1 18.3 20.7 24.0 32.6 35.9 78.5
Thiooxine em,,
vmSx X
X,
Color
mp
10-13
, , ,
...
Colorless Colorless Blue
616
48.7
Blue
596
50 3
Violet-red Red-violet Blue-violet Violet-red Red Violet-red Red-orange
...
...
...
523 522 575 514 503 535 460
l/mol. cm.
.., ...
500
450 ...
57 3 57 2 52 2 58 4 59 6 56 0 65 2
...
... 800 1100 1200
... 1500
emax
,,x, mp
l/mol cm.
589" 593s 555. ... 528b 509
17b 27
...
...
570a 503
, . .
490
43 * 1335
...
448 *
2032b
Acid Dissociation Constants of Selenoxine, Thiooxine, and Oxine in Water and 50% v./v. Aqueous Dioxane
PKO, Com50% v./v. pound Water dioxane-water Selenoxine -0.08a 0.12a Thiooxine 2.0 1.74b Oxine 5.13c 3.97d Reliability of both pK,, and pK,, is & 0 , 1 . Reference ( 4 ) . Reference (e). Reference ( 7 ) .
PK+ Water 8.7P 8.40 9 , 8gC
500c v./v. dioxane-water 8,50a 9.205 11.54d
VOL. 36, N O . 9, AUGUST 1 9 6 4
1769
Table 111. Acid Dissociation Constants of Certain Methyl-Substituted Derivatives of Selenoxine, Thiooxine, and Oxine in Water and 50% v./v. Aqueous Dioxane
hcid dissoriation constant ___~_________ 507; v./v. Compound
dioxanewater
Water
!!J = 0
3.90
SeCH3
(p =
1.90 1)
(Ai
=
...
0.1)
0.5)
band a t 460 mp is observed a t intermediate p H values. Selenoxine exhibits a corresponding absorption band in other solvents a t positions which vary with the dielectric constant and hydrogen bonding ability of the solvent (Figure 2). In general this absorption band undergoes a blue shift as well as an increase in intensity with increasing dielectric constant of the solvent. The change in intensity of this band seems to indicate the e\istence of a tautomeric equilibrium in selenoxine which could involve its zwitterion form.
3.53 (p =
scH3
w,
0.1)
3 . 50a
6.8b
...
OH CH3' 0
Reference (1).
* Reference ( 11)
The acid dissociation constant of the Se-methyl derivative was determined as follows: appropriate quantities of the compound and NaC1O4 were dissolved in water or 50Oj, v./v. dioxanea a t e r to give solutions that were 2.02 x 10-4-11 and 0.1 in ionic strength. Solutions of varying pH, obtained by adding S a O H , were examined spectrophotometrically at 345 mp and the pK, value was determined (Table 111). RESULTS A N D DISCUSSION
Selenoxine crystallizes as violet brown needles which are more soluble than thioosine in water. Selenosine is readily soluble in chloroform and ethanol and gives violet to purple solutions, but in contrast with thiooxine it is only slightly soluble in benzene and carbon tetrachloride. Such solubility behavior might well indicate the existence of selenoxine in a zwitterion form. Although solutions of selenoxine in the alkaline or strongly acid region are nearly colorless, a strong absorption
A similar eyplanation was advanced for the behavior of thiooxine (9) where, however, the change in the band intensity is much more pronounced. The decrease in band intensity from this point of view would result from a decreaqe in the relative concentration of the zwitter ion form. Thus in nonpolar solvents such as benzene or carbon tetrachloride where the molar absorbance is lorn, the nonpolar tautomer predominates. On the other hand, the intense absorption observed in solvents of increasing polarity reflects the growing importance of the zwitter ion form. I n the case of thiooxine, whose molar absorbance in water is 2032 and in methanol 133, the nonpolar tautomer predominates in solvents whose dielectric constants are under 30. I n the case of selenoxine the molar absorbance in solvents of such lo!$ polarity as CHC13 and CH2C12is sufficiently high (ca. 500) as to make it likely that the zwitter ion form is of great importance even here. Certainly the solubility characteristics of selenoxine in nonpolar solvents strengthen this conclusion. Osine also exhibits an absorption band in the same general region (430 mp) whose very low intensity would indicate the small contribution of the zwitter ion form in this reagent. In addition to the general blue shift observed in the position of this absorption band ith the dielectric constant there is, particularly at low dielectric
Table IV. Tautomeric Constants and Acid Dissociation Constants of the Cationic, Anionic, Neutral, and Zwitter Ionic Forms of Selenoxine, Thiooxine, and Oxine in Water and 50% v./v. Aqueous Dioxane
Compound Selenoxine Selenoxine
Xkdiurn Water
pKaa
pK.8
pKw
pKao
-0.08
3.26
8.75
5.41
1.90 3.50 5.14*
8.50 8.27 8.426
6.72 6.84 9.98*
507, v . / v .
aqueous dioxane Water Water
Thiooxine Oxine a Reference (1). * Reference ( 8 ) . 1770
Kt 2187
ANALYTICAL CHEMISTRY
59 27" 0.04b
0.13 2.07 6.60b
constants, a blue shift in going from a nonhydrogen-bonding to a hydrogenbonding solvent. This finding is in keeping with the stabilization of the ground state of the zwitter ion form that would be expected by its solvation with a hydrogen-bonding solvent. There might be a corresponding effect on the band intensity which would show the effect of the hydroxylic solvent in moving the tautomeric equilibrium toward the dipolar form. This same general e~planationcan be advanced for thiooxine whose behavior is seen to closely parallel that of selenosine (Figure 2). I n view of the foregoing discussion, acid dissociation phenomena of selenoxine as well as oxine and thiooxine require the following explanatory scheme :
XH
neutral
cation
X
= 0, S, or Se. The relationships between the various equilibrium constants in this scheme and composite or macroscopic acid dissociation constants, called K,, and K,,, are as follows:
Ka,
=
KoA
+
Kog
(1)
From the above scheme it is obvious that potentiometrically determined pK values must be composite-Le., macroscopic constants. In the case of oxine, the zwitter ion form represents such a small fraction of the total neutral species that the potentiometrically determined pK,, and pK,, correspond closely to P K ~ Band pKaD. In the case of selenoxine the pK,, and pKa2values were determined spectrophotometrically by using a band corresponding to the zwitter ion form alone. Even so, these pK, values are composite values since K,, characterizing the interconversion of neutral and zwitter ion forms is pH independent. Values of pK,, of both the thiooxine and selenoxine can be estimated by
measuring the dissociation constants of the S-methyl or Se-methyl derivatives (Table 111) since in these methyl derivatives participation of the zwvjtter ion forms is excluded ( b ) . From the values of Kogas well as from K,, and K,, from use of Equa,tions 1 to 3 the other K, values are readily obtained. With this information (Table IV:l it now becomes possible to evaluate the effect of structural changes on the dissociation phenomena in detail. esaniination of the dissociation scheme shown above leads to the following predictions. The values of pKaD, expected to be sensitive to the electronegativity of X , should decrease considerably from 0 to Se. Similarly, values of pK,, which correspond t'o the loss of a proton from i;hp group-XH in the cationic species, !should be lower than those of pK,, by virtue of the presence of the positive charge, but should change in the same manner as pK,,. Aiscan be seen from Table IV values cf pK,, drop 4.51units from oxine to sele;.oxine in the same order as observed in the acid strengths of H20, H,S, and H2Se. The magnitude of these values seem to be reasonable since that for osine is very close to the pK, of a-naphthol (9.85) and l.hat for thiooxine is close to the pK, of thiophenol (6.5). I t is of interest to note the similarity of the value of pK,, of oxine with that
(6.8) of the 'V-methyl analog of the cationic oxine species (12). The relationship between pK,, and pK,, can be expressed as: pK,, = 1.5 pKaD - 8.4. The relationship shows two consequences of the presence of the positive charge. First, its presence causes a general lowering of the pK,, and second, this lowering varies linearly with X, being greatest with selenoxine. This might be due to the greater sensitivity of the more highly polarizable selenium to the presence of the neighboring positively charged proton. The effect of changing X on the values of pKaBand pKac would not be expected to be appreciable since the X atom is not directly involved in the dissociation processes nor is there any appreciable conjugative interaction between the X and S atoms. This is the case in the pK,, values observed and, were it not for oxine, in the pK,, values as well. The deviation of the value of pK,B for oxine may reflect the stabilization of the neutral species through hydrogen bonding. K , would be eypected to increase with increasing strength of -XH and with increasing dielectric constant of the solvent. Both selenoxine and thiooxine exist almost entirely in their zwitter ion forms in water. I n 50% dioxane-water (dielectric constant = 32) selenoxine is still present entirely as the zwitter ion
form but thiooxine, with its significantly lower K 1 even in water, is probably largely in its neutral form in 50% aqueous dioxane. For osine, even in aqueous solution, the concentration of the awitter ion form is negligible. LITERATURE CITED
( 1 ) Albert, A , , Barlin, G. B., J . Chem. SOC. 1959, p. 2384. ( 2 ) Albert, A., Hampton, A . , Ibad., 1954, p. 505. (3) Bankovsky, Y. A , , Chera, L. XI., Ievinish, A. F., J . Anal. Chem., C'.S.S.R. 18, 668 (1963). ( 4 ) Corsini, A., Fernando, Q., Freiser, H., ANAL.CHEM.35, 1424 (1963). ( 5 ) Ebert, L., Z . Phys. Chem. 121, 385 (1920). ( 6 ) Freiser, H., Charles, R. G., Johnston, W. I)., J . Am. ('hem. SOC. 74, 1383 (1952). 1~, 7 ) Johnston. W. D.. Freiser. H.. J . Am. Chem. S O C . '5239 ~ ~ ,(19521. ' (8) Lee, H. S., Freiser, H., J . Org. Chem. 25, 1277 (1960). (9) IIason, S . F., J . Chem. SOC. 1958, p. 678. (10) Muthmann, W.,Schroeder, E., Rer. 33. 1766 11900). ( 1 1 ) 'Phillips, J. P., Keown, R. W., J . Am. Chem. SOC.73,5483 (1951). (12) Sekido, E., Fernando, Q., Freise:r, H., ANAL.CHEM.35, 1550 ( i963).
RECEIVEDfor review March 6, 1964. Accepted May 1, 1964. Work supported, by the U. S. Atomic Energy Commission.
ion Exchange Separation of Microgram Quantities of Osmium from Large Amounts of Base Metals J. C. VAN LOON and F. E. BEAMISH Department o f Chemisfry, University o f Toronto, Toronto
b Microgram amounts of osmium can be separated quantitatively from decigram amounts of the associated base metals, copper, iron, and nickel by cation exchange. Significant losses of osmium will result from an evaporation of hydrochloric acid solutions of hexachloroosmates. These losses may be controlled by a prior treatment of the osmium solution with suifur dioxide.
C
rnethods for quantitatively separating iron, copper, and nickel from trace amounts of platinum and palladium ( 6 ) , rhodium and iridium ( b ) , and ruthenium (8) have been recorded. Early efforts to separate osmium in a similar manner resulted in persistently low values. It was as,umed that the difficulties incident to the separation of ruthenium applied also in the case of osmium ( 2 ) , as the elements behave alike analytiATJOX EXCHANGE
5, Ontario, Canada
cally. With ruthenium, low recoveries from the cation exchange separation were associated with aging solutions caused, presumably, either by the slow production of dissolved cationic species, and/or by hydrolysis to hydrated oxides. An additional source of loss in the case of osmium, relatively inapplicable in the case of ruthenium, is the readily formed volatile octavalent osmium oxide. Various investigators have reported losses of osmium during the evaporation of the hydrochloric acid solutions used to collect the distilled octavalent oxide. T o avoid this loss the collecting liquid is usually treated with a reducing reagent ( 1 ) . 'Cnfortunately, while some pertinent data have been recorded, dealing with the solution composition of osmium-hydrochloric acid solutions, none allows a final interpretation of the identities of dissolved constituents in the various receiving solutions containing the reducing constit-
uents. Obviously, identification of the osmium constituents would allow more acceptable conclusions as to the source and amount of losses in experiments designed to provide safe methods of dissolution and separation. However, an examination by absorption spectrophotometry of osmium-hydrochloric acid solutions prepared by the evaporation of various collecting liquids used to receive the octavalent osmium oxide and by the direct addition of chloroosmate to hydrochloric acid solution indicates a common pattern. Thus irrespective of the exact identity of the dissolved species, and in the absence of more satisfactory data, the chloroosmate may be used to determine the degree to which osmium is lost during evaporation. The use of this constituent is particularly advantageous because the comparable ruthenium salt was used for a similar investigation with ruthenium whose separations frequently involve a similar technique ( 2 ) . VOL. 36, NO. 9, AUGUST 1964
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