kcal. per mole (4%). Therefore, the variability of the heat of precipitation is sufficient to account for errors of a few per cent. It is possible that the effects of changing concentrations could be overcome by the addition of a large excess of a n inert electrolyte. This was observed in the oxidation of iron(I1) and hexacyanoferrate(I1) by cerium(1V) (6). Here the reactions were carried out in 0.5M sulfuric acid which was apparently effective in swamping out the effects of small changes in reactant concentrations.
The effect of changing ionic strength on the value of reaction enthalpies should be investigated prior to using an enthalpimetric approach in analysis. LITERATURE CITED
(1) Dean, P. M., Newcomer, E., J. Am. Chem. SOC.47, 64 (1925). (2) Greathouse, L. H., Janssen, H. J., Haydel, C. H., ANAL. CHEM.28, 357 (1965). (3) Harries, R. J. N., Neath Technical College, Wales, private communication, Sept. 8, 1965. (4) Jordan, J., Alleman, T. G., ANAL. CHEM.29, 9 (1957). (5) Jordan, J., ,Ewing, G. J., “Thermometric Titrations,” in Meites, “Hand-
.
book of Analytical Chemistry,” p. 8-3, McGraw-Hill, New York (1963). (6) Jordan, J., Ewing, G. J., unpublished work, The Pennsylvania State UniverUniversit Park, Pa., 1960. inde, H. Rogers, L. B., Hume, D. N., ANAL.CHEM.25, 404 (1953). (8) McClure, J. H., Roder, T. M., Kinsey, R. H., Ibid., 27, 1599 (1955). (9) Ross1ni, F. C., et ul., “Selected Values of Chemical Thermodynamic Properties,” National Bureau of Standards Circular 500, Series I, U: S.Government Printing Office, Washington, D. C., 1962. (10) Wasilewski, J. C., Pei, P. T-S., Jordan, J., ANAL. CHEM. 36, 2131 (1964). RECEIVEDfor review July 12, 1966. Accepted August 8, 1966.
(7yiF
d,
I.
Substituted 8-Mercaptoquino~inesas Analytical Reagents Dissociation and Metal Chelate Formation Constants of 2-Methyl-8-Mercaptoquinoline DAVID KEALEY’ and HENRY FREISER Deparfmenf of Chemistry, Universify of Arizona, Tucson, Ariz. The macroscopic acid dissociation constants of 8-mercaptoquinoline and 2-methyl-8-mercaptoquinoline have been determined in water and 50% v./v. aqueous dioxane at 25’ C. These, together with the acid dissociation constants of the corresponding S-methyl derivatives, have been employed to elucidate the dissociation schemes of the two reagents. The first stepwise formation constants of the Cu(ll), Ni(ll), Co(ll), and Zn(ll) chelates of both reagents have been determined in soy0 v./v. aqueous dioxane. As a result of both ligand atom electronegativity and steric effects, the relative stability of Zn(ll) to Ni(ll) increases significantly indicating the usefulness of these structural design parameters in increasing selectivity of organic analytical reagents.
T
that steric hindrance arises by the interference of the oxygen atom of one ligand molecule with the methyl group of the other, when the two ligand molecules are co-planar. I n chelates in which only one ligand molecule is coordinated to the metal ion, any steric effect that is observed must be caused by the proximity of the alkyl group to water molecules or other donor groups co-ordinated to the metal. The present work describes the p r o p erties of 2-methyl-8-mercaptoquinoline and its reactions with metal ions. It was considered that an even greater steric effect than that observed with %methyl-8-quinolinol might occur due to the increased size of sulfur relative to oFgen. A comparison of the acid dissociation, tautomeric and chelateformation equilibria with those of 8mercaptoquinoline is presented.
HE INTRODUCTION OF SUBSTITUENTS
into the molecule of an organic chelating agent often results in changes in its chelating ability in addition to effects on other chemical and physical properties. The presence of methyl or other alkyl substituents close to the donor atoms has resulted in marked changes in the reactivities of these reagents toward metal ions. 8-Quinolinols substituted in the 2-position not only fail to form precipitates with aluminum(II1) (6, fO), but give chelates of significantly lower stability with the metal ions with which they do react (8). Molecular models of 2methyl-8-quinolinol chelates indicate
EXPERIMENTAL
Synthesis of 8-Mercaptoquinoline
and
2-Methyl-8-Mercaptoquinoline. The .reagents were prepared from 8 - aminoquinoline and 2 - methyl - 8aminoquinoline as described previously (9) and stored as their hydrated sodium salts. The sodium salt of 8-mercaptoquinoline is stable indefinitely, but that of the 2-methyl derivative is very slowly oxidized to 2,2’-dimethyl-8,Btdiquinolyl disulfide which can, however, be removed before use by treatment with hot benzene. Both salts decompose above 300” C. without melting. Synthesis of S-Methyl-Z-Methyl-8Mercaptoquinoline. A solution of
sodium salt of 2-methyl-8-mercaptoquinoline (2.3 grams, 0.01 mole) in 0.1M aqueous sodium hydroxide (100 ml.) is cooled to 0’ C. and treated with several 0.25-ml. portions of methyl iodide, with vigorous shaking between additions, until the mixture is decolorized. The excess of methyl iodide is removed by evaporation, and the white precipitate of S-methyl-2methyl-8-mercaptoquinoline is filtered, and washed with water. Recrystallization from absolute ethanol gives white crystals, m.p. 103’ C. The yield is almost quantitative. Anal. calcd. for CllHllNS: C, 69.80; H, 5.86; N, 7.40 Found: C, 70.1; H, 6.1; N, 7.4. Other Reagents. The lJ4-dioxane was Baker analyzed reagent and was further purified as described previously (4). Stock solutions of the metal perchlorates in deionized water were standardized titrimetrically using standard EDTA procedures. All other reagents were Analyzed Grade. Apparatus. All p H measurements were made with a Becknian Model G pH meter equipped with a glass and a saturated calomel electrode. T h e method of standardizing the electrode has been described previously (8). The spectrophotometric measurements were made with a Cary-14 recording spectrophotometer. Thermogravimetric Analysis of 2-
- -
Methyl 8 Mercaptoquinoline H y drate. Samples of the orange hydrate
of 2-methyl-8-mercaptoquinoline, recrystallized from de-aerated water Present address, Mobil Oil Corp., Paulsboro, N. J.
VOL. 38, NO. 11, OCTOBER 1 9 6 6
1577
and dried between filter papers, were weighed a t room temperature in a current of argon. Continuous passage of argon for several hours resulted in a weight loss which corresponded to 3 k 0.5 moles of water per mole of the mercaptan. The loss of water was accompanied by a gradual color change from orange to white. On raising the temperature to 100" C., no further weight loss was observed, but at 40' C., a rapid and irreversible solid to liquid phase change occurred, accompanied by a further color change to dark violet blue. These observations are discussed later.
quinoline or 2-methyl-&mercaptoquinoline was dissolved in 100 ml. of deaerated 50% v./v. aqueous dioxane containing sufficient perchloric acid to leave the solution slightly acid. The ionic strength was adjusted to 0.1 with solid sodium perchlorate, and the titration, with 0.3M sodium hydroxide, performed in the usual way (8) with nitrogen passing continuously through and above the solution. The second acid dissociation constant, Kaz,defined by the expression
Qualitative Reactions with M e t a l
where [HL], [H+L-1, and [L-] r e p resent the concentration of neutral, zwitterionic, and anionic forms of the reagent, was evaluated from the titration data with the help of the following expression derived by consideration of charge and mass balance relations:
Ions. A 0.5% aqueous solution of the reagent was prepared by dissolving a n appropriate amount of the sodium salt in de-aerated dilute perchloric acid. The solution was added dropwise to 0.1% solutions of 25 metal salts, perchlorates or nitrates where possible, at several acidities. Al(III), Cr(III), Mn(II), Mg(II), and U(V1) gave no observable reaction with the reagent at pH 5. Sb(II1) and Sn(I1) gave yellow precipitates which may have included hydrolysis products in 2M HC104. The results are shown in Table I. Acid
Dissociation
The pK.,, which is too low for potentiometric determination, was obtained spectrophotometrically. Solutions of the sodium salts, about lO-SM, were prepared in de-aerated 50% v./v. aqueous dioxane at pH values ranging from 0 to 5.2. The ionic strength was adjusted to 0.1 with solid sodium perchlorate for solutions of pH > 1. The absorbance of the 481 mp band, in the case of 8-mercaptoquinoline, and of the 464 mp band, in the case of 2-methyl-8mercaptoquinoline, which increased from 0 to a maximum with increasing pH, was measured at 25' C. using 2cm. quartz cells. The bands, which are attributed to the zwitterionic form of the neutral molecule ( I ) , are free of interference from other absorbing species. Values of the measured absorbances are given by
A
=
2EHtL-{ [HLI
+
[H'L-]) 2eR,L+[ffZL+] (4)
and the total reagent concentration by where T R is the analytical concentration of the reagent and
S
=
{ [H+]
+ [Na+] -
[C101-1 - [OH-]) (3) A plot of pH us. the logarithm term gives a line whose intercept on the pH axis is pK,,.
Constants.
Values of pK,, were determined potentiometrically at 25' C. About 200 mg. of the sodium salt of 8-mercapto-
TR = [ H L ]
+ [H+L-] + [HzL+]
(5)
assuming the species [L-] to be absent. Combination of Equations 4 and 5 and substitution for [H2L+] and [H+L-] in the expression for the acid dissociation constant gives
~~
Table 1.
M +n 2 : 3
Bi Cd+2 co+2
Cu+:' Fe Fe+8 Ga+a Hg+:' In
Ni+2
Pb+2 Pd+2
Qualitative Reactions of 2-Methyl-8-Mercaptoquinoline with Metal Ion
2M
orange ppt. d. brown ppt. ... yellow ppt.
...
brown ppt.
...
... ...
yellow ppt.
...
orange'ppt. red-brown ppt. brown ppt. reagent oxidized
P t +2 Pt +4 T1+ ... Zn +a ... , . , indicates not tested
HC104 1M orange ppt. d. brown ppt. no ppt. yellow ppt. no ppt. brown ppt.
0.1M
orange ppt. d. brown ppt. orange ppt. ellow ppt. rown ppt. brown ppt. no ppt. reagent oxidized pale yellow ppt. pale yellow ppt. pale yellow ppt. no ppt. orange ppt. orange-red ppt. brown ppt. reagent oxidized yellow ppt. yellow ppt.
i:
no ppt. yellow ppt. no ppt. ... orange ppt. red-brown ppt. brown ppt. reagent oxidized no ppt. no ppt.
PH 5 Acetate buffer orange ppt. d. brown ppt. ... ellow ppt. Town ppt. brown ppt. d. green ppt. d. brown ppt. pale yellow ppt. pale yellow ppt. pale ellow ppt. grey-Lown ppt. orange ppt. orange-red ppt. brown ppt. reagent oxidized yellow ppt. yellow ppt.
g
Table II. Acid Dissociation Constants of 8-Mercaptoquinoline, 2-Methyl-8Mercaptoquinoline, and Their S-Methyl Derivatives in Water and 50% v./v. Aqueous Dioxane at 25' C. and Ionic Strength 0.1
~Kai
P K ~
50% v./v.
Compound 8-Mercaptoquinoline
%Methyl-8-mercaptoquinoline &Methyl-8-mercaptoquinoline %Methyl-Zmeth .1 8-
50% v./v.
Water
dioxanewater
Water
dioxanewater
2.16 2.28 3 . 5 0 (11)
1.79 1.96 2.23
8.38 9.06
9.22 9.76
... ...
... ...
rnercaptoqmnoLe (4.08)" 2.78 Based on assumption that the differences in pK measured in the two media is the same as observed with %methyl-8-mercaptoquinoline. 1578
ANALYTICAL CHEMISTRY
whence pH
=
{
log A2,+::;:
t}:
+ PKa,
A graphical solution of the equation was obtained as described previously for pK,,. The pK,, of S-methyl-Zmethyl-8mercaptoquinoline, which is higher than that of the mercaptan, was determined potentiometrically. The pK,, and pK,, values obtained by both methods are shown in Table 11, together with the corresponding data for aqueous solutions. Spectrophotometric Determination
of Co, Ni, Cu, and Zn 1: 1 Chelate The 1 : 2 chelates are very insoluble, even in 50% aqueous dioxane, and the equilibrium data obtained are thus limited to the formation of the 1:l chelates. This was carried out by keeping the ratio of metal to reagent very high-Le. 100: 1. Solutions of the 1:1 metal chelates in 50y0 aqueous dioxane were prepared by mixing the following: 5 ml. of metal perchlorate solution, 10 ml. of perchloric acid-sodium perchlorate solution to adjust to the required pH; solid sodium perchlorate to adjust ionic strength to 0.1, 20 ml. of freshly distilled 1,4dioxane, 5 ml. of reagent solution in 0.01M perchloric acid. The total volume was 40 ml. and no further dilution was made. The final concentration of total metal ion was at least 1.25 X 1OW2Mand of total reagent
Formation Constants.
1.25 X lO-*M. The pH range covered was 1.1 to 6 and matched 2-cm. quartz cells were used throughout. The acid for dissolution of the reagent was deaerated with nitrogen before use. The absorbance of the 1:1 chelates decreased slowly as a function of time, and a linear extrapolation to zero time was made in each case. To obtain the absorbance of the 1:1 chelates, all components of the metal chelate solutions, except the reagent solution, were mixed. The reagent was then added using a 5 m l . pipet, and an electric timer started when the pipet had half emptied. The contents of the flask were mixed thoroughly and the cell was rinsed and filled. The absorbance a t A., . was recorded for a time equal to the time which had already elapsed since adding the reagent. The absorbance at zero-time was then obtained by adding the observed decrease to the initial reading. The procedure was repeated with all the solutions, which were measured against a blank from which the reagent had been omitted. Corrections due to reagent absorbance were incorporated into the calculations as shown below, where H + L - represents the zwitterionic form of the neutral molecule.
+ (H'L- + H L ) Kaz (H+L- + H L ) H + + L-
Table 111. Tautomeric Constants and Acid Dissociation Constants of the Cationic, Anionic, and Zwitterionic Forms of 8-Mercaptoquinoline and Its 2-Methyl Analog in Water and 50% v./v. Aqueous Dioxane Compound hledium KT p K e ~ p K 0 ~ pK.c PKOD 8-Mercaptoquinoline 50% v./v. 1.74 1.99 2.23 9.02 8.78
8-hlercaptoquinoline 2-Methyl-8-mercaptcquinoline %Methyl-S-mercaptoquinoline
aqueous dioxane Water ( 1 ) 50% v./v. aqueous dioxane Water
27 5.62
2.07 2.03
3.50
2.78
8.27 9.69
6.84 8.94
60
2.29
(4.08)
9.05
7.28
From Equations 7, 10, and 11 and putting [ M + 2 ] = Tu,
and substituting for [ M L + ]from Equation 13 gives
K.1
H * L + eH +
(6)
(7)
The equilibria for the formation of the metal chelates are given by M + 2
+ HzL++KIM L + + 2 H +
(8)
Kl'
M+2
+ ( H L + H+L-) e ML+ M+2
+ H+
+ L - KII eM L +
(9) (10)
Also, the total reagent is given by TR
=
+
+ +
Graphical solution of the final equation gave a line of unit slope confirming the absence of other complexes, such as M L 2 , which would require a different [ H + ] dependence. A value for K I , is obtained from the intercept on the p H axis when the left-hand logarithmic term is equal to zero. It should be noted that for absorbance measurements made where either H L , H+L-, or H2L+ do not absorb significantly, the final equation reduces to a simpler form. The values of K J : are shown in Table 111.
[ML+l ([HLI [If+L-l) [HzL+l (11)
If A represents the total absorbance of the solutions in a 2-cm. cell, then, A =
log Kf1
+
+
[ M L + ] 2ex + L - ( [ H L ] [H+L-I) 4- ~ C H , L + [ H Z L (12) +]
~ E M +L
where C H + L - is the molar absorptivity of the tautomeric mixture of H L and H+L-. Combinations of Equations 6 and 11 give expressions for ( [ H + L - ] [HL]) and [HZL+] which can be substituted in Equation 12. Subsequent rearrangement gives
+
[ML+]=
DISCUSSION
Structure and Roperties of 2Methyl 8 Mercaptoquinoline. T h e properties of 2 methyl 8 mercaptoquinoline closely parallel those of the parent compound. Thus, t h e anhydrous mercaptan is a dark redblue liquid with a pungent b u t not unpleasant odor. It is soluble in polar and nonpolar solvents and is slowly oxidized in air to the corresponding disulfide. The compound is amphiprotic, acid and alkaline solutions being yellow due to the presence of protonated and anionic species, respectively. A bright orange crystalline hydrate
-
-
-
- -
KO,
Tu
(16)
separates from neutral solutions, and thermogravimetric data indicate the presence of 3 0.5 moles of water per mole of mercaptan. The uncertainty in this figure results from the extreme ease of dehydration, which occurs more readily than with 8-mercaptoquinoline hydrate. For this reason, attempts to determine the melting point of the orange hydrate were unsuccessful, but a value within the range 30-36' C. is probable. The bright orange color of the hydrated mercaptan contrask with the deep red of 8-mercaptoquinoline dihydrate, and this is reflected in the visible absorption spectrum which shows a hypsochromic shift of the zwitterion band at 446 mp in the parent compound to 433 m r in the 2-methyl derivative. This effect is attributable to the influence of the methyl group on electron transitions occurring in the zwitterionic form. Steric hindrance of the methyl group with the bridging water molecules is discounted on the basis of examination of models. The presence of the methyl group must also be responsible for the white metastable anhydrous form, which has not been observed in the case of 8mercaptoquinoline, although a similar phenomenon has been reported (8) for some 5-halo-8-mercaptoquinolines. It is possible that anhydrous zwitterion is the initial product formed on dehydration, and the diffe'rence in behavior of
*
VOL 38, NO. 1 1 , OCTOBER 1966
1579
Table IV. 1 :1 Metal Chelate Formation Constants of 8-Mercaptoquinoline and 2-Methyl-8-Mercaptoquinoline in 50% v./v. Aqueous Dioxane at 25" C. and Ionic Strength 0.1
8-Mercaptoquinoline log Kfi 7.9
Metal ion COW) Ni(I1) CU(I1)
11.0
12.7
Zn(1I)
11.0
2-Methyl-82-Methyl-8merca toquinoline 8-Quinolinol ('4) quinolinol (8) Kfi log Kfi log Kfl 9.6 10.55 9.63 9.2 11.44 9.41 11.7 13.49 12.48 11.1 9.96 9.82
the substituted and unsubstituted mercaptans could be explained by consideration of the relative stabilities of the zwitterion forms of the ligands. The value of the tautomeric equilibrium constant, K,, decreases with decreasing water (polar solvent) content (11). The dipolar form of the 2-methyl compound, being more stable than that of the parent compound, can exist in the anhydrous form. Infrared data show medium to strong SH stretching bands at 2515 cm.-l and 2510 cm.-1 for anhydrous 8-mercaptoquinoline and liquid anhydrous 2methyl - 8 - mercaptoquinoline, respectively. Comparisons with thiophenol (2575 cm.-l) and 2-thionaphthalene (2590 cm.-l) indicate that the mercaptoquinolines are hydrogen bonded, and the absence of detectable shifts between the spectra of the neat mercaptans and those of carbon tetrachloride solutions a t a number of concentrations shows that this is intramolecular rather than intermolecular. Meaningful data could not be obtained from the solid anhydrous form of 2methyl-8-mercaptoquinoline due to its instability towards pressure, solvents, and Nujol. The qualitative reactions with metal ions (Table I) are similar to these of 8-mercaptoquinoline. A notable exception is Xit2 to which forms a grey-brown precipitate a t pH 5 but none a t pH 1. This contrasts with the bright crimson precipitate formed with 8-mercaptoquinoline at the lower pH, and provided a preliminary indication of the lower formation constant expected from steric hindrance. It was also noted that some larger ions-e.g,, Cd'2, Hg+2, Pb+*-gave precipitates in 2M HClOd under which conditions no such reaction with 8-mercaptoquinoline is observed. Acid Dissociation Phenomena. Since 2-methyl-8-mercaptoquinoline exists as a tautomeric mixture of neutral and zwitterion forms in aqueous dioxane, the measured acid dissociation constants, K., and K,,, are composites of those which describe the reactions of the following scheme: K.*
HzL+ e H' 1580
+ H'L-
ANALYTICAL CHEMISTRY
HL
KT + H+L-
It can be shown that
By assuming that the value of PK,, is the same as the pK. of the S-methyl derivative, the values of all of the dissociation constants in the scheme can be calculated. The values thus obtained are shown in Table 111. The value of pK,, is increased in the 2methyl derivative over 8-mercaptoquinoline (A pK,, 0.5) to the same extent as observed in the corresponding 8-quinolinol system. This is in keeping with the expected electronreleasing effect of the 2-methyl group on the basicity of the quinoline nitrogen. The effect of the methyl group on the basicity of the 8-mercaptide species 0.2) is less pronounced (ApK,, which is also parallel to the behavior of 8-quinolinol and its 2-methyl analog. The general increase in basic character resulting from the 2-methyl group produced the expected increase in stability of the zwitterion tautomer. I t is interesting to note that in the case of pK,,, involving a cationic acid, an acid-strengthening is observed ( ApK = 1.3) in changing from water to the less polar dioxane-water medium whereas for pK,,, for a neutral acid, an acid-weakening effect was found (ApK = 1.8 j~ 0.1). These effects are generally observed with acids of similar charge type. Chelate Formation Equilibria. Despite the lower proton affinity of 8-mercaptoquinoline (for which pK., pK,, is lower than that of 8-quinolinol by 4.5 units) its relatively high metal complexing ability which was reported
-
-
+
earlier (3) was confirmed in this study (Table IV). Similarly, although the pK,, for 2-methyl-8value of pK,, mercaptoquinoline is 4.4 units less than that of its oxygen analog, the metal complex formation constants of the two ligands are about the same. The introduction of the 2-methyl group results in about the same changes in the metal complex formation constants of both mercapto- and hydroxyquinoline reagents, at least in the case of Ni, Cu, and Zn. The results for Co are anomalous; the K f , for 8-mercaptoquinoline seems unusually small. It was somewhat unexpected to find the steric effect of the 2-methyl group to be so close in the mercapto- and hydroxyquinolines. Since sulfur is larger than oxygen i t might be expected to be more affected by steric hindrance, Perhaps this factor might be compensated by an increase in covalent character expected in the sulfur-metal bond. The decrease in stability arising from the 2-methyl group in 8-mercaptoquinoline indicates the operation of steric interference of the methyl group with the water molecules in the hydrated 1 : l metaI complexes. This effect becomes larger with decreasing metal ion radius (6-8). This serves to enlarge the stability difference between zinc and nickel already noted in the replacement of oxygen by sulfur (S), and brings the zinc complex almost to the level of stability of the copper complex. Thus, the combination of steric and ligand atom electronegativity factors is seen to be especially effective in changing the usual metal stability sequence. Since the complex formation constants of 2-methyl-8-mercaptoquinoline do not increase above those of the parent compound as do the acid dissociation constants, except for cobalt, the corresponding proton displacement constants, K P D .(defined as Kf, K,, K,,), are smaller. As a result, complex formation for the 2-methyl compound does not take place until higher pH values are reached. Thus, with nickel, formation of the 2-methyl-8-mercaptoquinoline complex occurs at 2.5 pH units higher than that with 8-mercaptoquinoline. In the case of cobalt, complex formation starts at one pH unit lower.
+
ACKNOWLEDGMENT
The authors gratefully acknowledge David Kingston's help with the thermogravimetric measurements. LITERATURE CITED
(1) Albert, A,, Barlin, G. B., J . Chem. SOC.1959, p. 2384. (2) Bankovskis, J., Chera, L..M., Ievin'sh, A. F., Zh. Analit. Kham. 18, 668 (1963).
(3) Corsini, A., Fernando, Q., Freiser, H.,ANAL.CHEM.35, 1424 (1963). (4) Freiser, H.1 Charles, R. G.1 Johnston1 w. D., J.Am. Chem. SOc. 74,1383 (1952). (5) Irving, H., Butler, E . J., Ring, hf. F., J. Chem. SOC.1949,p. 1489. (6) Irving, H., Mellor, D. H., Zbid., 1962, p. 5237.
(7) Irving, H., Rossotti, H. S., Zbid., 1954, p. 2910. (8) Johnston, W. D., Freiser, H., Anal. Chim. A , . ~11, 201 (1954). (9) K e a b , D*, Freiser~ H . ~Talanta 13, in Press (1966). (10) Merritt, L. L., Walker, J. K., IND.
ENG.CHEM.,ANAL.ED. 16, 387 (1944). (11) Sekido, E., Fernando, Q., Freiser, H., ANAL.CHEM.36, 1768 (1964). RECEIVEDfor review June 20, 1966. Accepted July 29, 1966, The financial assistance of the U. S. Atomic Energy Commission is acknowledged.
Solution Equilibria and Structures of MoIy bde num(VI) Che Iates (E th y le ned initri lo)tetra acetic Acid RICHARD J. KULA Department o f Chemistry, The University o f Wisconsin, Madison, Wis.
b Several of the equilibria in the Mo(VI) (ethylenedinitri1o)tetraacetic acid system have been investigated using proton nuclear magnetic resonance techniques and pH titrations. Between pH 9 and 5, two complexes are formed, one with a one-to-one metal-ligand ratio and the other with a two-to-one metal-ligand ratio. The formation constants of both chelates have been determined and compared with the formation constant of the one-to-one Mo(VI)-methyliminodiacetote chelate. In acid solutions the two-to-one chelate polymerizes by
0 forming Mo/ \Mo linkages between individual chelate species. From the nuclear magnetic resonance spectra, structures of the complexes are proposed; and from the aqueous solution equilibria, the possibility of a direct titrimetric determination of Mo(VI) using (ethylenedinitri1o)tetraacetic acid is considered.
A
the existence of an (ethylenedinitri1o)tetraacetic acid (EDTA) chelate having two Mo(V1) ions has been proved, no quantitative information concerning the stability of this species is available (,9,8, 11-13, 16). One report suggests that a one-to-one Mo(V1)-EDTA complex also forms, but no work seems to have been initiated to characterize this complex (5). I n the present work, the aqueous solution equilibria of the Mo(V1)EDTA system are examined in detail, and structural and bonding features of the solution species are considered. As in the preceding paper concerning the complexes of Mo(V1) and methyliminodiacetic acid (MIDA), the data were obtained mainly by proton nuclear magnetic resonance (NMR) and by pH titrations (6). I n addition, certain features of the aqueous infrared spectra are considered. LTHOUGH
12
I
10
t:.
I
two equivalents of acid are consumed for each equivalent of Mo between pH 11 and 5, and one equivalent of acid is consumed for each Mo below pH 5. By analogy with the RIIDA system, the overall stoichiometry between pH 11 and 5 can be viewed as the reaction 2Hf MOO^-^ e MOO, H20, followed by coordination of MOO, with EDTA. With two equivalents of M o and four equivalents of H + per EDTA, the resultant product is the two-to-one ( M O O J ~ E D T A - chelate. ~ In the MIDA system, the consumption of one equivalent of H + per Mo below pH 5 leads to a dimeric chelate. A similar reaction can be envisaged for the EDTA system, but on the basis of structures proposed for the EDTA-4 chelate (2, 8 ) , the product would be polymeric rather than dimeric. Consistent with this interpretation is the observation that solutions of pH less than 4 become cloudy after standing for about an hour. I n the vicinity of pH 2, a solution which is 0.5M in (MoO&EDTA-~ completely solidifies. This solid dissolves only very slowly if a NaOH solution is added. NMR. Spectra of solutions with varying metal-ligand ratios indicate that ligand exchange is relatively slow on the N M R time scale, and t h a t the predominant Mo(V1) chelate contains two Mo ions per EDTA. Most of the remaining work was carried out on solutions with this metal-ligand ratio. The spectra between pH 10 and 5 consist of resonances other than those of free EDTA (7) and of the two-to-one Mo-EDTA complex ( 2 ) , as illustrated in Figure 2. The main features of these additional resonances are a four-line A B multiplet whose chemical shift is independent of pH; a single sharp resonance whose chemical shift varies with pH and whose intensity is equal to the total intensity of the A B multiplet;
+
moles H' moles EDTA
Figure 1.
pH titration curves
- - - 0.2M NarEDTA4
-25'C; 0.4M Na2Ma04:
0.2M NarEDTA
titrated with 5 M H N 0 3
EXPERIMENTAL
The experimental techniques were identical to those outlined previously (6), a Varian A-60 N M R spectrometer being employed using water as a solvent and tetramethylammonium chloride as an internal reference. Crystalline Na4(MOO~)ZEDTA.~HZ was O synthesized according to the method of Pecsok and Sawyer (8). Solution infrared spectra were obtained on a Perkin-Elmer Model 421 instrument, equipped with sample cells made of BaFz and having 0.025-mm. spacing. The solutions for the infrared studies were 0.05M in the Mo-EDTA chelate, and 99.8% D 2 0 was the solvent. RESULTS
pH Titrations. The p H titration curves for a solution of NalEDTA and for a two-to-one solution of Ner MoOrNalEDTA are shown in Figure 1. The corresponding titration curve for a solution of NazMoO4 was presented previously and showed acid consump tion only below pH 6.2 (6). The titration curve for the t w o - t w n e solution of Mo and EDTA is similar to that for the one-to-one Mo-MIDA system where
+
VOL 38, NO. 11, OCTOBER 1966
1581
