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Equilibrium studies of uranyl complexes—I
J. Inorg.Nucl.Chem.. 1964,Vol.26, pp. 789 to 798. PergamonPressLtd. Printedin NorthernIreland
EQUILIBRIUM STUDIES OF URANYL COMPLEXES--I INTERACTION OF URANYL ION WITH SOME HYDROXYCARBOXYLIC AND AMINOCARBOXYLIC ACIDS*t K. S. RAJAN a n d A. E. MARTELL Department of Chemistry, Illinois Institute of Technology, Chicago, Illinois (Received 29 August 1963; in revised form 25 October 1963)
Abstract--A quantitative study of the interaction of uranyl ion with 5-sulphosalicylic acid (SSA), iminodiacetic acid (IMDA) and N-hydroxyethyliminodiacetic acid (HIMDA) has been carried out potentiometrically. Stability constants of the uranyl chelates containing a 1 : 1 ratio of metal to ligand have been determined at 25 ° in media maintained at 0"1 and 1.0 ionic strengths with potassium. nitrate. The hydrolysis and dimerization constants of the 1 : 1 uranyl HIMDA chelate are reported. Possible structures for the various species present in solution are indicated on the basis of an analysis of the potentiometric data. ALTHOUGH m a n y u r a n y l complexes have been r e p o r t e d , i n f o r m a t i o n on the equilibria involved in their f o r m a t i o n is generally lacking. Similarly, relatively little research has been carried o u t on the h y d r o l y t i c reactions o f uranyl complexes, a l t h o u g h the hydrolysis a n d p o l y m e r i z a t i o n o f the a q u o u r a n y l ion in a q u e o u s solutions has been studied b y m a n y investigators. (1,~) O f the u r a n y l complexes with oxygen d o n o r s , the tartratet3, 4) citrate, ('s) lactate, t5,6) malatetS,6, 7> a n d salicylate is) have been described a n d f o r m a t i o n constants for the glycolate tg) a n d ascorbate(l°, ~1) have been r e p o r t e d . FELDMAN a n d NEWMANt12) a n d FELDMAN and HAVmL iS) have d e t e r m i n e d the stoichiometry o f the possible complexes o f u r a n y l ion with aliphatic h y d r o x y acids by spectrop h o t o m e t r i c m e t h o d . O n the basis o f p o l a r o g r a p h i c a n d p o t e n t i o m e t r i c d e t e r m i n a t i o n s , F:ELDMANOa,~41 a n d c o w o r k e r s r e p o r t e d the f o r m a t i o n o f b i n u c l e a r u r a n y l citrate, tartrate, m a l a t e a n d lactate chelates below p H 3, a n d they have d e t e r m i n e d equilibrium constants for the f o r m a t i o n o f d i m e r in all these systems. T h e y considered * This work was supported by the U.S. Atomic Energy Commission under contract no. AT(1 I-1)1020. Abstract in part from material submitted to the faculty of the Illinois Institute of Technology in partial fulfillment of the requirements of the Degree of Doctor of Philosophy. (1~ j. SUTTON,J. Chem. Soc. Supp., No. 2, 275 (1949). (~ S. AHRLAND,S. HIETAtqENand L. G. SILLEN,Acta. Chem. Scand. 8, 1907 (1954). "~ E. PELIGOT,J. Prakt. Chem. 35, 153 (1945). (~) H. ITZIG, Bet'. Chem. Dtsch. Ges. 34, 3822 (1901). ~:')I. FELDMANand J. R. HAVlLL,J. Amer. Chem. Soc. 76, 2114 (1954). ~) I. FELDMAN,I. R. HAVILLand W. U. NEUMAN,J. Amer. Chem. Soc. 76, 4726 (1954). ~7) S. HAKAMORI,Science Reports Tohako Imp. University I, 20, 756 (1931); Chem. Abstr. 26, 2366 (1932). (") G. COURTOIS,Btdl. Soc. Chim. France (4), 33, 1761 (1923). ~9) N. C. LI, W. M. WESTFALL,A. LINDENBAUM,J. M. WHITEand J. SCHUBERT,J. Amer. Chem. Soc. 79, 5864 (1957). (~0~ 1. J. GALE,Proceedings of the International Conference on the Peaceful Uses of Atomic Ener~,), at Geneva, 1956, Vol. 8, p. 358. United Nations 1956. ~ A. GREGORCZYK,Acta Polon. Pharm. 15, 129 (1958); Chem. Abstr. 52, 19655 (1958). (~.ol1. FELDMANand W. F. NEWMAN,J. Amer. Chem. Soc. 73, 2312 (1951). ila) W. F. NEWMAN,J. R. HAVILLand I. FELDMAN,J. Amer. Chem. Soc. 73, 3593 (1951). t~4~ I. FEI DMAN,C. A, NORTH and H. B. HUNTER, J. Phys. Chem, 64, 1224 (1960). 789
790
K.S. RAJAN and A. E. MARTELL
further interaction o f the binuclear species with hydroxyl ion to result in the formation o f a ternuclear chelate. LI and coworkers ~15~ calculated stability constants for the m o n o n u c l e a r u r a n y l citrate chelate having a ligand to metal ratio o f 1:1. F r o m spectrophotometric studies FOLEY and ANDERSON~16) reported the existence o f only the 1 : i uranyl-sulphosalicylate complex in aqueous solution. However, BANKS and SINGH t17~have reported the formation o f 1 : 1 and 1 : 2 complexes o f UO~ 2+ with sulphosalicylic acid at p H values 4.5 and 7,5 respectively. RICHARD et al. tla) investigated the chelation, hydrolysis and olation o f uranyl ion with 8-hydroxyquinoline-5-sulphonate by potentiometric and spectrophotometric methods, and reported the hydrolysis and olation constants for the bis-(5-sulpho-8-quinolinolo) dioxouranium (VI) complex. GUSTAFSON et al. ~19~ studied the interaction o f uranyl ion with Tiron (disodium-l,2dihydroxybenzene-3,5-disulphonate) by potentiometric and spectrophotometric methods in the p H range 2-11 and reported equilibrium constants pertaining to the formation ofternuclear u r a n y l - T i r o n chelates having a 1 : 1 ratio o f l i g a n d to metal ion. The present investigation, the first o f a series o f equilibrium studies on uranyl complexes, consists o f a quantitative potentiometric study o f the interactions o f uranyl ion with the ligands 5-sulphosalicylic acid (SSA), iminodiacetic acid ( I M D A ) , and N-hydroxyethyl iminodiacetic acid ( H I M D A ) as well as any subsequent hydrolysis and olation reactions. These complexes were studied as a function o f p H over a wide range o f concentration in order to determine the nature o f the possible interactions which take place in aqueous solutions. EXPERIMENTAL The experimental method consisted of potentiometric titrations of the ligand in the absence of and in the presence of uranyl ion. The ionic strengths of the media were maintained constant at 0-1 and 1.0 by using potassium nitrate. Purified nitrogen was passed through the solution throughout the course of a titration in order to exclude carbon dioxide. The temperature was maintained constant at 25.0 °. A Radiometer pH meter (type PHM 4) fitted with glass and calomel extension electrodes was used to determine the hydrogen ion concentration. The electrode system was calibrated by direct titration of acetic acid in such a way that the observed pH meter reading was compared with the actual hydrogen ion concentration determined from the data tabulated by HARNEDand OwEr,r.~°~ In the pH regions below 3.5 and above 10.5 the meter was calibrated by measurement of solutions of known concentrations of HC1 and NaOH, respectively. Reagents. Aqueous uranyl nitrate solutions were prepared from Baker and Adamson's analysed reagent and were standardized gravimetrically by ignition of suitable aliquots to U3Os. The ligand 5-sulphosalicylic acid (SSA) was obtained from Distillation Products Industries, Rochester 3, New York. Disodium iminodiacetate obtained from the same source was acidified with hydrochloric acid in order to crystallize out the pure acid. A single reerystallization produced a product of 99.7 per cent purity. Hydroxyethyliminodiacetic acid (HIMDA) was donated by Dow Chemical Company Midland, Michigan. Aqueous solutions of the various chelating agents were standardized potentiometrically with standard carbonate-free sodium hydroxide. Carbonate-free sodium hydroxide was prepared from a nearly saturated aqueous solution of sodium hydroxide, and was standardized against potassium acid phthalate. ~1~ N. C. El, A. LINDENBAUMand J. M. WHITE,J. lnorg. Nucl. Chem. 12, 122 (1959). ~1~ R. T. FOLEYand R. C. ANDERSON,J. Amer. Chem. Soc. 71, 909 (1949). {17~C. V. BANKSand R. S. SINGH,J. lnorg. Nucl. Chem. 15, 125 (1960). txs) C. F. RICHARD,R. L. GUSTAFSONand A. E. MARTELL,J. Amer. Chem. Soc. 81, 1033 (1959). o91 R. L. GUSTAFSON,C. F. RICHARDand A. E. MARTELL,J. Amer. Chem. Soc. 82, 1526 (1960). t29} H. S. HARNEDand B. B. OWEN, The Physical Chemistry of Electrolytic Solutions (2rid Ed.), p. 523. Reinhold, New York (1950).
Equilibrium studies of uranyl complexes--I
791
Procedure. A measured quantity of standardized uranyl nitrate solution was pipetted into the titration cell and the desired amount of the acid form of the ligand was then added slowly while the mixture was agitated with a magnetic stirrer. The solution was made up to 50 ml with carbonate-free water and an appropriate amount of 1-0 M KNOB. When the mixture reached thermal equilibrium, the hydrogen ion concentration was determined by a number of successive readings of the pH meter. Increments of base were then added from a microburette and measurements were made until the pH reached 10.5-11.5.
7"'" s /
/
/ /
/
/C
i I l !
i /
/
/ / I
5
3
0
I
I
2
4
6
m
FIG. 1 .--Potentiometric curves of uranyl 5-sulphosalicylicacid chelates : A, free ligand; B, 1:1 molar ratio of UO22+ to ligand; C, 1:2 molar ratio of UO~2÷ to ligand; m --
moles of sodium hydroxide added per gram-ion of uranium present; 10 zM; t = 2 5 ~ z 0 . 0 5 ° ; /~-0"I(KNO3).
T~t = 3.9
RESULTS
Uranyl-sulphosalicylic acid complex.
P o t e n t i o m e t r i c titrations o f the free ligand as well as o f solutions c o n t a i n i n g 1 : 1 a n d 1 : 2 m o l a r ratios o f u r a n y l ion to sulphosalicylic acid are shown in Fig. 1. T h e t i t r a t i o n curve for the ligand shows a strong inflexion at two moles o f base p e r mole o f ligand. H 2 L - is considered to be the l i g a n d in the m a t h e m a t i c a l t r e a t m e n t , since the s u l p h o n i c acid g r o u p is c o m p l e t e l y dissociated u n d e r the reaction c o n d i t i o n s e m p l o y e d . T h e inflexion in the t i t r a t i o n curve at m = 3 for the 1 : 1 system indicates the f o r m a t i o n o f the simple chelate in which one m o l e o f l i g a n d combines with an equivalent a m o u n t o f the metal ion, with the d i s p l a c e m e n t o f two equivalents o f h y d r o g e n ion from the ligand : UO2 ~+ + H 2 L - ~ U O , , L - i - 2H + where H 2 L - represents the ligand.
792
K . S . RAJAN and A. E. MARTELL
The formation constant for this chelate compound was calculated from the usual relationships: KML =
(T~-
[L~-]X)
[L3-]2X
(1)
where
[L3-] =
2TL To~ -- [H+] 2- [- H + ] q- ~ [H+] -
-
[H+] ~
[H+]
X = K1----~ + - - ~ , + 1
(2)
(3)
and TM = total molar concentration of metal species, TT, = total molar concentration of ligand species, To~ = total molar concentration of hydroxyl ion added to the solution and K1 and K2 are the acid dissociation constants of the ligand. Potentiometric titrations of the UO2 s+ ion in the presence of an equivalent concentration of sulphosalicyclic acid were carried out over a twenty-fold metal ion concentration in order to determine the presence or absence of polynuclear complexes. The results obtained from the experiments carried out in 0.1 and 1.0 molar potassium nitrate media are presented in Table 1. Iminodiacetatodioxouranium (VI) complex. Potentiometric titrations of solutions of iminodiacetic acid as well as of 1 : 1 and 1 : 2 molar ratios of uranyl nitrate to iminodiacetic acid are illustrated in Fig. 2. The inflexion at two moles of base per grammeion of metal (m = 2) for the 1 : 1 complex is in accord with the reaction: UOs s+ + H2L ~ UO~L + 2H + The formation constant for this complex was calculated with Equation (1). Potentiometric measurements were made over a twenty-fold concentration range of the metal chelate in 0.1 and 1.0 molar potassium nitrate media. The results are presented in Table 1. N-Hydroxyethyliminodiacetato-dioxouranium (IV) complexes. Potentiometric titration curves obtained from solutions of the ligand (HIMDA), as well as 1 : 1 and 1 : 2 molar ratios of uranyl salt to the ligand are shown in Fig. 3. The 1 : 1 titration has a barely perceptible inflexion at m = 2 and a steep inflexion at m -----3, preceded by the appearance of a solid phase at m = 2.8 (pH 6.0). The initial reaction is probably metal chelate formation according to the following equation : UO2 s+ + HsL ~-- UOsL + 2H + where H2L represents the neutral form of the ligand. The concentration dependence of the above reaction was investigated over a twenty-fold range of metal chelate concentration in 0.1 and 1.0 molar potassium nitrate media and the data obtained are presented in Table 1. The equilibrium constant for the hydrolysis reaction: UO2L + H s O . K~. UO2(OH)L + H+ was determined from the potentiometric data between m = 2 and m = 2.6. The concentration dependence of the hydrolysis constant was studied over a twenty-fold
H ~¸
UO ~
H÷
U022+
H +
U O 22~
H"
H ~ + U O 2 ( O H ) L - ~ UO._,L
H ÷ + L 2- ~ H L H + ÷ H L - ~- HzL UO_~+ L_ L 2- ~ UOoL
H + 4" L 2- ~ H L H + + H L - ~ H2L UO22~ + L 2- ~ UO2L
H i ? L a- ~_ H L ~H + 4" H L ~- ~ H~LUO.fl + -!- L ~- ~ UO2L-
Reaction
25 °
5'22 -~ 0"05 5'51 ± 0"09 5"81 + 0"14
0"01 0.03 0.01 0.01 0.01
0'01 0"01 0'02 0"02 0'02
1'49 × 10 -2 3"59 ;~ 10 -3 7"18 ~: 10 -~
~ ~ 4. 4, 4-
4~ 4" ± 4"
4" 0"02 4" 0"02 4" 0"03 4" 0"03 4" 0"03
1.49 :~, 10 2 3.59 × 10 -a 7"18 × 10 -4
9"40 2"50 8"86 8'96 8"97
1l'40 2"27 10"56 10'62 10'67
~ ~: 4. 4. 4.
~ 4" 444"
0.01 0.01 0.01 0.01 0.04
0'01 0"01 0"02 0'02 0"04
4- 0"02 & 0"02 4" 0"02 4" 0"01 4" 0"05
5"08 ± 0"09 5"38 ± 0"12 5'66 ± 0"15
8"67 2"22 7"99 8.01 7'98
9"38 2'55 8-71 8'78 8"70
11"32 2"33 10"45 10"45 10'42
Log o f constant tt = 0"1 It = 1'0
8.72 2.20 8.24 8-34 8.37
1"59 × 10 -z 3"83 × 10 -3 7'67 × 10 -4
1"66 × 10 ~ 3"98 × 10 -3 7"97 × 10 -~
T~ or T~ (M)
F o r m a t i o n constant
DISSOCIATION CONSTANTS AND CHELATE FORMATION CONSTANTS t =
~-'" D. D. PERRIN, Nature Lond. 182, 741 (1958). (~2~ C. V. BANKS and R. S. SINGH, J. Inorg. Nuel. Chem. 15, 125 (1960). (~3~ S. CHABEREK and A. E. MARTELL, J. Amer. Chem. Soc. 74, 5052 (1952). c~ai S. CHABEREK, R. C. COURTNEY and A. E. MARTELL, J. Amer. Chem. Soc. 74, 5057 (1952).
N-Hydroxyethyliminodiacetic acid, H2L
Iminodiacetic acid, H~L
5-Sulphosalicylic acid, H~L
Ligand
1.--ACID
Positive ion
TABLE
1.961-
8.78/~,~,
a t / t ==0.I
9"12]c~3, 2'54! at # = 0'1
K, = I1"14t1~ K2 8-06 J a t / t ":~-0-1
11 "95 t ~2,~ 2"62J at i t ~ 0.1~)-15
Literature value
8
e-
¢e~
e-
O"
=_..
794
K . S . RAJAN and A. E. MARTELL
7 I' s
6
.,,'"
•!
A
t..__./
FIG. 2.--Potentiometric curves o f uranyl-iminodiacetic acid chelates: A, free ligand; B, 1:1 molar ratio o f uranyl ion to ligand; C, 1:2 molar ratio of uranyl ion to ligand; m = moles of base added per gram-ion of" uranium present; T~ = 4"20 x 10 -3. M; t -- 25°; /~ ~- 0-10 (KNOa).
I
4
5
2
©
I
J
I
I
2
5
m
12
IC I"
FIG. 3.--Potentiometric curves of uranyl-hydroxyethyliminodiacetic acid chelates: A, free ligand; B, 1 : 1 molar ratio of uranyl ion to ligand; C, 1 : 2 molar ratio of uranyl ion to ligand; T~ = 3.78 ×10-3 M; t = 25°;/~ ~ 0.1 (KNO3).
i "°
+ Z ¢,D 8
l "s
,•/, o f
I
4
2
0
I
!
2
4
ITI
I
6
Equilibrium studies of uranyl complexes--I
FIo. 4.--Demonstration of binuclear complex formation in the hydrolysis of the 1 : 1 uranyl-hydroxyethyliminodiacetic acid chelates: Ton = total molar concentration of base added; molar concentration of the metal chelate, [ML] = O , 1.49 x 10-3; ~,3.59 x 10-3; O, 7.18 x 10-4; t = 25 °, ionic medium = 0.1 M KNO3.
795
,o, 4 ,-~ ~ 3 ~' .~'-~
*_o 2
0
0 10-2 x [M I_] / [i-I+]
IC
8 FIG. 5.--Demonstration of binuclear complex formation in the hydrolysis of the 1 : 1 uranyl-hydroxyethyliminodiacetic acid chelate: Toa = total molar concentration of base added per mole of metal complex present; molar concentrations of the metal chelate, [ML] = ~, 1"49 x 10 2; ~, 3.59 x 10-~; O, 7'18 × 10-4; t = 2 5 °, ionic medium = 1'0 M KNO3.
q---, I I
6
÷
r~
4
Z qD
C
0
4
10-2x [M~ /[H~
796
K.S. RAJANand A. E. MARTELL
TABLE
2 . - - C O M P L E X FORMATION CONSTANTS FOR HYDROLYSIS AND OLATION OF THE 1 : I HYDROXYETHYLIMINODIACETIC
Reaction
URANYL--
ACID CHELATE AT 25 °
Ionic strength (KN03 medium)
Log of equilibrium constant
0"1 1"0 0'1 1"0 0"1
5"92 5'87 8.34 8'08 3"50
1"0
3"65
H+ + UOa(OH)L- ~ UO=L 2H+ + (UO=(OH)L)=~- ~ 2UO=L 2UO=(OI-t-)L- ~ (UO2(OH)L)~2L
range of metal chelate concentration in 0-1 and 1.0 molar potassium nitrate media and the results are presented in Table I. Since the increasing pK's presented in Table 1 indicate the possibility of a polymerization reaction, the standard slope intercept relationship (Equation 6) was employed to determine both the hydrolysis and dimerization constants: K~ -----
KI) =
[UO~(OH)L-][H +] [UO~L]
[(UO2(OH) L)2z-] [H+] ~ [UO2L_]2
[H+](Torr + [H +] -- [OH-]) = 2[UO2L ] - KD+KH [UO2L] [H +]
(4)
(5) (6)
where ToN is the total moles of base added per mole of uranyl complex present, and the other terms have their usual meaning. The values of KD and Krr determined from the plot of this equation illustrated in Figs. 4 and 5 are shown in Table 2. DISCUSSION The potentiometric curve (Fig. 1) obtained for a solution containing a 1:2 molar ratio of the uranyl salt to sulphosalicylic acid (SSA) corresponds to that which would be predicted for a mixture of the 1 : 1 chelate and the free ligand. This observation is in agreement with the continuous variation work of FOLEY and ANDERSON~19) who showed the existence of only a I : 1 UO2-SSA chelate. BANKSand SINGHtls) reported that 1 : 1 and 1:2 complexes of UO2-SSA are formed at pH values 4.5 and 7.5 respectively, based on their spectrophotometric measurements of solutions containing 1:2, 1 : 6 and 1 : 10 molar ratios of metal ion to ligand, and calculated stability constants from potentiometrie data. Although the constant reported by them for a 1 : 1 complex is very close to the one obtained in this research, no evidence is apparent in Fig. I for the presence of a 1 : 2 complex. Also, precipitation was observed above pH 5 in both 1 : i and 1:2 systems. A possible structure of the diaquo 5-sulpho-salicylato chelate compatible with experimental evidence is the following:
d/ I
Equilibrium studies of uranyl complexes--I
797
Treatment of the hydrolytic reactions taking place beyond m = 3 (Fig. 1) could not be carried out in view of the precipitations observed beyond m = 3.3. It is evident from the data presented in Table 1 for the UO2-IMDA chelate compound that the log KuL values obtained from the 0.1 M KNO3 medium increase slightly with decreasing concentration of the metal chelate. This observed trend cannot be due to polymer formation, since the formation of polynuclear complexes would give a trend in the opposite direction. It may be due to small changes in the ionic atmosphere with increasing contribution of uranyl salt. An average Log K:vi, value of 8.76 -k 0.03 was obtained in 1.0 M KNOz medium. It is therefore concluded that the UO2-IMDA chelate is present in aqueous solution only as a monomer. A possible structure for this complex is indicated by I1. /co--cH~.. 0~- - - 0 - - - 7 N H I\ /'~. \ II / / CH z
- -
-o-
-
-:6
j
II
Analysis of the 1 : 2 UO=-IMDA titration curves indicate the presence of only the 1 : 1 complex. From the observed trend in the log KMr, values obtained in 0.1 and 1-0 M KNOz media (Table 1) for the 1 : 1 UOz-HIMDA chelate, it is concluded that no polynuclear complexes are formed in this system in the presence of up to two equivalents of base per mole of the metal salt. A possible structure for the 1:1 chelate in this buffer region is the following: _/CO~CH2 _
/' /
~ /
/ " Itf1 \ \
/
/CHzCH20H
CO
~o . . . . . o . . . . oj III
On the basis of the mathematical treatment of potentiometric data in the buffer region between m = 2 and m = 3 (Figs. 4 and 5 and Table 1) it is concluded that a hydrolysed and polymerized species (IV) is present in solution under these conditions:
"°°'.;7 , coo/O
\c.
. IV
However, on the basis of potentiometric data alone it is not possible to determine if hydrolysis occurs by dissociation of a proton from a coordinated water molecule or from a hydroxyethyl group of the ligand. A comparison of the 1 : 1 chelate stability constants of various metal salts with nitrilotriacetic acid (NTA), N-hydroxyethyliminodiacetic acid (HIMDA) and iminodiacetic acid (IMDA) presented in Table 3 suggest
798
K.S.
RAJAN a n d A. E. MARTELL
TABLE 3.--COMPARISON OF CHELATION TENDENCIES OF N T A , H I M D A AND I M D A Log K Metal i o n
C u ~+ N i z+ Co 2+ M n ~+ Z n 2+ C d ~+ P b ~+ M g 2+ C a 2+ UO2 ~+
N T A t~j
H I M D A c2~
IMDA
12"7 11"3 10'6 7"4 10"4 9"5 11"8 5"4 6'4
10 9"5 8"3 5"6 8"6 7"1 9.5 3"5 4"8 8"32
10'6 c27~ 8-3 (27~ 7"0 taT~ 7"0 t~7~ 5"3 c ~ 3.6 ~s~ 3"4 c28~ 8"93
that the hydroxyethyl group of HIMDA is not coordinated to the uranyl ion in the 1 : 1 chelate below pH 4. It is seen that for the transition metal ions and many other metal ions the HIMDA chelate has a larger formation constant than does the corresponding IMDA chelate. Since HIMDA is the less basic ligand, this correlation has been taken to indicate involvement of the aliphatic hydroxyl groupin the co-ordination of the metal ion. In the case of the uranyl complexes, the lack of involvement of the hydroxyethyl group in co-ordination indicates that it also may not be involved as a bridging group in the polynuclear complex. ~,5~ G. SCHWAILZENBACH,Helv. Chim. Acta, 32, 1175 (1949); 34, 1492 (1949). t2e~ S. CHABEREK, R. C. COUgTNEY a n d A. E. MARTELL,. L A m e r . Chem. Soc. 74, 5057 (1952). t27~ S. CaABEREK a n d A. E. MARTELL,.]'. Amer. Chem. Soc. 74, 5052 (1952). (g8) G. SCHWARZENBACH et ak, Heir. Chim. Acta, 28, 1133 (1945).
EQUILIBRIUM STUDIES OF URANYL COMPLEXES--I INTERACTION OF URANYL ION WITH SOME HYDROXYCARBOXYLIC AND AMINOCARBOXYLIC ACIDS*t K. S. RAJAN a n d A. E. MARTELL Department of Chemistry, Illinois Institute of Technology, Chicago, Illinois (Received 29 August 1963; in revised form 25 October 1963)
Abstract--A quantitative study of the interaction of uranyl ion with 5-sulphosalicylic acid (SSA), iminodiacetic acid (IMDA) and N-hydroxyethyliminodiacetic acid (HIMDA) has been carried out potentiometrically. Stability constants of the uranyl chelates containing a 1 : 1 ratio of metal to ligand have been determined at 25 ° in media maintained at 0"1 and 1.0 ionic strengths with potassium. nitrate. The hydrolysis and dimerization constants of the 1 : 1 uranyl HIMDA chelate are reported. Possible structures for the various species present in solution are indicated on the basis of an analysis of the potentiometric data. ALTHOUGH m a n y u r a n y l complexes have been r e p o r t e d , i n f o r m a t i o n on the equilibria involved in their f o r m a t i o n is generally lacking. Similarly, relatively little research has been carried o u t on the h y d r o l y t i c reactions o f uranyl complexes, a l t h o u g h the hydrolysis a n d p o l y m e r i z a t i o n o f the a q u o u r a n y l ion in a q u e o u s solutions has been studied b y m a n y investigators. (1,~) O f the u r a n y l complexes with oxygen d o n o r s , the tartratet3, 4) citrate, ('s) lactate, t5,6) malatetS,6, 7> a n d salicylate is) have been described a n d f o r m a t i o n constants for the glycolate tg) a n d ascorbate(l°, ~1) have been r e p o r t e d . FELDMAN a n d NEWMANt12) a n d FELDMAN and HAVmL iS) have d e t e r m i n e d the stoichiometry o f the possible complexes o f u r a n y l ion with aliphatic h y d r o x y acids by spectrop h o t o m e t r i c m e t h o d . O n the basis o f p o l a r o g r a p h i c a n d p o t e n t i o m e t r i c d e t e r m i n a t i o n s , F:ELDMANOa,~41 a n d c o w o r k e r s r e p o r t e d the f o r m a t i o n o f b i n u c l e a r u r a n y l citrate, tartrate, m a l a t e a n d lactate chelates below p H 3, a n d they have d e t e r m i n e d equilibrium constants for the f o r m a t i o n o f d i m e r in all these systems. T h e y considered * This work was supported by the U.S. Atomic Energy Commission under contract no. AT(1 I-1)1020. Abstract in part from material submitted to the faculty of the Illinois Institute of Technology in partial fulfillment of the requirements of the Degree of Doctor of Philosophy. (1~ j. SUTTON,J. Chem. Soc. Supp., No. 2, 275 (1949). (~ S. AHRLAND,S. HIETAtqENand L. G. SILLEN,Acta. Chem. Scand. 8, 1907 (1954). "~ E. PELIGOT,J. Prakt. Chem. 35, 153 (1945). (~) H. ITZIG, Bet'. Chem. Dtsch. Ges. 34, 3822 (1901). ~:')I. FELDMANand J. R. HAVlLL,J. Amer. Chem. Soc. 76, 2114 (1954). ~) I. FELDMAN,I. R. HAVILLand W. U. NEUMAN,J. Amer. Chem. Soc. 76, 4726 (1954). ~7) S. HAKAMORI,Science Reports Tohako Imp. University I, 20, 756 (1931); Chem. Abstr. 26, 2366 (1932). (") G. COURTOIS,Btdl. Soc. Chim. France (4), 33, 1761 (1923). ~9) N. C. LI, W. M. WESTFALL,A. LINDENBAUM,J. M. WHITEand J. SCHUBERT,J. Amer. Chem. Soc. 79, 5864 (1957). (~0~ 1. J. GALE,Proceedings of the International Conference on the Peaceful Uses of Atomic Ener~,), at Geneva, 1956, Vol. 8, p. 358. United Nations 1956. ~ A. GREGORCZYK,Acta Polon. Pharm. 15, 129 (1958); Chem. Abstr. 52, 19655 (1958). (~.ol1. FELDMANand W. F. NEWMAN,J. Amer. Chem. Soc. 73, 2312 (1951). ila) W. F. NEWMAN,J. R. HAVILLand I. FELDMAN,J. Amer. Chem. Soc. 73, 3593 (1951). t~4~ I. FEI DMAN,C. A, NORTH and H. B. HUNTER, J. Phys. Chem, 64, 1224 (1960). 789
790
K.S. RAJAN and A. E. MARTELL
further interaction o f the binuclear species with hydroxyl ion to result in the formation o f a ternuclear chelate. LI and coworkers ~15~ calculated stability constants for the m o n o n u c l e a r u r a n y l citrate chelate having a ligand to metal ratio o f 1:1. F r o m spectrophotometric studies FOLEY and ANDERSON~16) reported the existence o f only the 1 : i uranyl-sulphosalicylate complex in aqueous solution. However, BANKS and SINGH t17~have reported the formation o f 1 : 1 and 1 : 2 complexes o f UO~ 2+ with sulphosalicylic acid at p H values 4.5 and 7,5 respectively. RICHARD et al. tla) investigated the chelation, hydrolysis and olation o f uranyl ion with 8-hydroxyquinoline-5-sulphonate by potentiometric and spectrophotometric methods, and reported the hydrolysis and olation constants for the bis-(5-sulpho-8-quinolinolo) dioxouranium (VI) complex. GUSTAFSON et al. ~19~ studied the interaction o f uranyl ion with Tiron (disodium-l,2dihydroxybenzene-3,5-disulphonate) by potentiometric and spectrophotometric methods in the p H range 2-11 and reported equilibrium constants pertaining to the formation ofternuclear u r a n y l - T i r o n chelates having a 1 : 1 ratio o f l i g a n d to metal ion. The present investigation, the first o f a series o f equilibrium studies on uranyl complexes, consists o f a quantitative potentiometric study o f the interactions o f uranyl ion with the ligands 5-sulphosalicylic acid (SSA), iminodiacetic acid ( I M D A ) , and N-hydroxyethyl iminodiacetic acid ( H I M D A ) as well as any subsequent hydrolysis and olation reactions. These complexes were studied as a function o f p H over a wide range o f concentration in order to determine the nature o f the possible interactions which take place in aqueous solutions. EXPERIMENTAL The experimental method consisted of potentiometric titrations of the ligand in the absence of and in the presence of uranyl ion. The ionic strengths of the media were maintained constant at 0-1 and 1.0 by using potassium nitrate. Purified nitrogen was passed through the solution throughout the course of a titration in order to exclude carbon dioxide. The temperature was maintained constant at 25.0 °. A Radiometer pH meter (type PHM 4) fitted with glass and calomel extension electrodes was used to determine the hydrogen ion concentration. The electrode system was calibrated by direct titration of acetic acid in such a way that the observed pH meter reading was compared with the actual hydrogen ion concentration determined from the data tabulated by HARNEDand OwEr,r.~°~ In the pH regions below 3.5 and above 10.5 the meter was calibrated by measurement of solutions of known concentrations of HC1 and NaOH, respectively. Reagents. Aqueous uranyl nitrate solutions were prepared from Baker and Adamson's analysed reagent and were standardized gravimetrically by ignition of suitable aliquots to U3Os. The ligand 5-sulphosalicylic acid (SSA) was obtained from Distillation Products Industries, Rochester 3, New York. Disodium iminodiacetate obtained from the same source was acidified with hydrochloric acid in order to crystallize out the pure acid. A single reerystallization produced a product of 99.7 per cent purity. Hydroxyethyliminodiacetic acid (HIMDA) was donated by Dow Chemical Company Midland, Michigan. Aqueous solutions of the various chelating agents were standardized potentiometrically with standard carbonate-free sodium hydroxide. Carbonate-free sodium hydroxide was prepared from a nearly saturated aqueous solution of sodium hydroxide, and was standardized against potassium acid phthalate. ~1~ N. C. El, A. LINDENBAUMand J. M. WHITE,J. lnorg. Nucl. Chem. 12, 122 (1959). ~1~ R. T. FOLEYand R. C. ANDERSON,J. Amer. Chem. Soc. 71, 909 (1949). {17~C. V. BANKSand R. S. SINGH,J. lnorg. Nucl. Chem. 15, 125 (1960). txs) C. F. RICHARD,R. L. GUSTAFSONand A. E. MARTELL,J. Amer. Chem. Soc. 81, 1033 (1959). o91 R. L. GUSTAFSON,C. F. RICHARDand A. E. MARTELL,J. Amer. Chem. Soc. 82, 1526 (1960). t29} H. S. HARNEDand B. B. OWEN, The Physical Chemistry of Electrolytic Solutions (2rid Ed.), p. 523. Reinhold, New York (1950).
Equilibrium studies of uranyl complexes--I
791
Procedure. A measured quantity of standardized uranyl nitrate solution was pipetted into the titration cell and the desired amount of the acid form of the ligand was then added slowly while the mixture was agitated with a magnetic stirrer. The solution was made up to 50 ml with carbonate-free water and an appropriate amount of 1-0 M KNOB. When the mixture reached thermal equilibrium, the hydrogen ion concentration was determined by a number of successive readings of the pH meter. Increments of base were then added from a microburette and measurements were made until the pH reached 10.5-11.5.
7"'" s /
/
/ /
/
/C
i I l !
i /
/
/ / I
5
3
0
I
I
2
4
6
m
FIG. 1 .--Potentiometric curves of uranyl 5-sulphosalicylicacid chelates : A, free ligand; B, 1:1 molar ratio of UO22+ to ligand; C, 1:2 molar ratio of UO~2÷ to ligand; m --
moles of sodium hydroxide added per gram-ion of uranium present; 10 zM; t = 2 5 ~ z 0 . 0 5 ° ; /~-0"I(KNO3).
T~t = 3.9
RESULTS
Uranyl-sulphosalicylic acid complex.
P o t e n t i o m e t r i c titrations o f the free ligand as well as o f solutions c o n t a i n i n g 1 : 1 a n d 1 : 2 m o l a r ratios o f u r a n y l ion to sulphosalicylic acid are shown in Fig. 1. T h e t i t r a t i o n curve for the ligand shows a strong inflexion at two moles o f base p e r mole o f ligand. H 2 L - is considered to be the l i g a n d in the m a t h e m a t i c a l t r e a t m e n t , since the s u l p h o n i c acid g r o u p is c o m p l e t e l y dissociated u n d e r the reaction c o n d i t i o n s e m p l o y e d . T h e inflexion in the t i t r a t i o n curve at m = 3 for the 1 : 1 system indicates the f o r m a t i o n o f the simple chelate in which one m o l e o f l i g a n d combines with an equivalent a m o u n t o f the metal ion, with the d i s p l a c e m e n t o f two equivalents o f h y d r o g e n ion from the ligand : UO2 ~+ + H 2 L - ~ U O , , L - i - 2H + where H 2 L - represents the ligand.
792
K . S . RAJAN and A. E. MARTELL
The formation constant for this chelate compound was calculated from the usual relationships: KML =
(T~-
[L~-]X)
[L3-]2X
(1)
where
[L3-] =
2TL To~ -- [H+] 2- [- H + ] q- ~ [H+] -
-
[H+] ~
[H+]
X = K1----~ + - - ~ , + 1
(2)
(3)
and TM = total molar concentration of metal species, TT, = total molar concentration of ligand species, To~ = total molar concentration of hydroxyl ion added to the solution and K1 and K2 are the acid dissociation constants of the ligand. Potentiometric titrations of the UO2 s+ ion in the presence of an equivalent concentration of sulphosalicyclic acid were carried out over a twenty-fold metal ion concentration in order to determine the presence or absence of polynuclear complexes. The results obtained from the experiments carried out in 0.1 and 1.0 molar potassium nitrate media are presented in Table 1. Iminodiacetatodioxouranium (VI) complex. Potentiometric titrations of solutions of iminodiacetic acid as well as of 1 : 1 and 1 : 2 molar ratios of uranyl nitrate to iminodiacetic acid are illustrated in Fig. 2. The inflexion at two moles of base per grammeion of metal (m = 2) for the 1 : 1 complex is in accord with the reaction: UOs s+ + H2L ~ UO~L + 2H + The formation constant for this complex was calculated with Equation (1). Potentiometric measurements were made over a twenty-fold concentration range of the metal chelate in 0.1 and 1.0 molar potassium nitrate media. The results are presented in Table 1. N-Hydroxyethyliminodiacetato-dioxouranium (IV) complexes. Potentiometric titration curves obtained from solutions of the ligand (HIMDA), as well as 1 : 1 and 1 : 2 molar ratios of uranyl salt to the ligand are shown in Fig. 3. The 1 : 1 titration has a barely perceptible inflexion at m = 2 and a steep inflexion at m -----3, preceded by the appearance of a solid phase at m = 2.8 (pH 6.0). The initial reaction is probably metal chelate formation according to the following equation : UO2 s+ + HsL ~-- UOsL + 2H + where H2L represents the neutral form of the ligand. The concentration dependence of the above reaction was investigated over a twenty-fold range of metal chelate concentration in 0.1 and 1.0 molar potassium nitrate media and the data obtained are presented in Table 1. The equilibrium constant for the hydrolysis reaction: UO2L + H s O . K~. UO2(OH)L + H+ was determined from the potentiometric data between m = 2 and m = 2.6. The concentration dependence of the hydrolysis constant was studied over a twenty-fold
H ~¸
UO ~
H÷
U022+
H +
U O 22~
H"
H ~ + U O 2 ( O H ) L - ~ UO._,L
H ÷ + L 2- ~ H L H + ÷ H L - ~- HzL UO_~+ L_ L 2- ~ UOoL
H + 4" L 2- ~ H L H + + H L - ~ H2L UO22~ + L 2- ~ UO2L
H i ? L a- ~_ H L ~H + 4" H L ~- ~ H~LUO.fl + -!- L ~- ~ UO2L-
Reaction
25 °
5'22 -~ 0"05 5'51 ± 0"09 5"81 + 0"14
0"01 0.03 0.01 0.01 0.01
0'01 0"01 0'02 0"02 0'02
1'49 × 10 -2 3"59 ;~ 10 -3 7"18 ~: 10 -~
~ ~ 4. 4, 4-
4~ 4" ± 4"
4" 0"02 4" 0"02 4" 0"03 4" 0"03 4" 0"03
1.49 :~, 10 2 3.59 × 10 -a 7"18 × 10 -4
9"40 2"50 8"86 8'96 8"97
1l'40 2"27 10"56 10'62 10'67
~ ~: 4. 4. 4.
~ 4" 444"
0.01 0.01 0.01 0.01 0.04
0'01 0"01 0"02 0'02 0"04
4- 0"02 & 0"02 4" 0"02 4" 0"01 4" 0"05
5"08 ± 0"09 5"38 ± 0"12 5'66 ± 0"15
8"67 2"22 7"99 8.01 7'98
9"38 2'55 8-71 8'78 8"70
11"32 2"33 10"45 10"45 10'42
Log o f constant tt = 0"1 It = 1'0
8.72 2.20 8.24 8-34 8.37
1"59 × 10 -z 3"83 × 10 -3 7'67 × 10 -4
1"66 × 10 ~ 3"98 × 10 -3 7"97 × 10 -~
T~ or T~ (M)
F o r m a t i o n constant
DISSOCIATION CONSTANTS AND CHELATE FORMATION CONSTANTS t =
~-'" D. D. PERRIN, Nature Lond. 182, 741 (1958). (~2~ C. V. BANKS and R. S. SINGH, J. Inorg. Nuel. Chem. 15, 125 (1960). (~3~ S. CHABEREK and A. E. MARTELL, J. Amer. Chem. Soc. 74, 5052 (1952). c~ai S. CHABEREK, R. C. COURTNEY and A. E. MARTELL, J. Amer. Chem. Soc. 74, 5057 (1952).
N-Hydroxyethyliminodiacetic acid, H2L
Iminodiacetic acid, H~L
5-Sulphosalicylic acid, H~L
Ligand
1.--ACID
Positive ion
TABLE
1.961-
8.78/~,~,
a t / t ==0.I
9"12]c~3, 2'54! at # = 0'1
K, = I1"14t1~ K2 8-06 J a t / t ":~-0-1
11 "95 t ~2,~ 2"62J at i t ~ 0.1~)-15
Literature value
8
e-
¢e~
e-
O"
=_..
794
K . S . RAJAN and A. E. MARTELL
7 I' s
6
.,,'"
•!
A
t..__./
FIG. 2.--Potentiometric curves o f uranyl-iminodiacetic acid chelates: A, free ligand; B, 1:1 molar ratio o f uranyl ion to ligand; C, 1:2 molar ratio of uranyl ion to ligand; m = moles of base added per gram-ion of" uranium present; T~ = 4"20 x 10 -3. M; t -- 25°; /~ ~- 0-10 (KNOa).
I
4
5
2
©
I
J
I
I
2
5
m
12
IC I"
FIG. 3.--Potentiometric curves of uranyl-hydroxyethyliminodiacetic acid chelates: A, free ligand; B, 1 : 1 molar ratio of uranyl ion to ligand; C, 1 : 2 molar ratio of uranyl ion to ligand; T~ = 3.78 ×10-3 M; t = 25°;/~ ~ 0.1 (KNO3).
i "°
+ Z ¢,D 8
l "s
,•/, o f
I
4
2
0
I
!
2
4
ITI
I
6
Equilibrium studies of uranyl complexes--I
FIo. 4.--Demonstration of binuclear complex formation in the hydrolysis of the 1 : 1 uranyl-hydroxyethyliminodiacetic acid chelates: Ton = total molar concentration of base added; molar concentration of the metal chelate, [ML] = O , 1.49 x 10-3; ~,3.59 x 10-3; O, 7.18 x 10-4; t = 25 °, ionic medium = 0.1 M KNO3.
795
,o, 4 ,-~ ~ 3 ~' .~'-~
*_o 2
0
0 10-2 x [M I_] / [i-I+]
IC
8 FIG. 5.--Demonstration of binuclear complex formation in the hydrolysis of the 1 : 1 uranyl-hydroxyethyliminodiacetic acid chelate: Toa = total molar concentration of base added per mole of metal complex present; molar concentrations of the metal chelate, [ML] = ~, 1"49 x 10 2; ~, 3.59 x 10-~; O, 7'18 × 10-4; t = 2 5 °, ionic medium = 1'0 M KNO3.
q---, I I
6
÷
r~
4
Z qD
C
0
4
10-2x [M~ /[H~
796
K.S. RAJANand A. E. MARTELL
TABLE
2 . - - C O M P L E X FORMATION CONSTANTS FOR HYDROLYSIS AND OLATION OF THE 1 : I HYDROXYETHYLIMINODIACETIC
Reaction
URANYL--
ACID CHELATE AT 25 °
Ionic strength (KN03 medium)
Log of equilibrium constant
0"1 1"0 0'1 1"0 0"1
5"92 5'87 8.34 8'08 3"50
1"0
3"65
H+ + UOa(OH)L- ~ UO=L 2H+ + (UO=(OH)L)=~- ~ 2UO=L 2UO=(OI-t-)L- ~ (UO2(OH)L)~2L
range of metal chelate concentration in 0-1 and 1.0 molar potassium nitrate media and the results are presented in Table I. Since the increasing pK's presented in Table 1 indicate the possibility of a polymerization reaction, the standard slope intercept relationship (Equation 6) was employed to determine both the hydrolysis and dimerization constants: K~ -----
KI) =
[UO~(OH)L-][H +] [UO~L]
[(UO2(OH) L)2z-] [H+] ~ [UO2L_]2
[H+](Torr + [H +] -- [OH-]) = 2[UO2L ] - KD+KH [UO2L] [H +]
(4)
(5) (6)
where ToN is the total moles of base added per mole of uranyl complex present, and the other terms have their usual meaning. The values of KD and Krr determined from the plot of this equation illustrated in Figs. 4 and 5 are shown in Table 2. DISCUSSION The potentiometric curve (Fig. 1) obtained for a solution containing a 1:2 molar ratio of the uranyl salt to sulphosalicylic acid (SSA) corresponds to that which would be predicted for a mixture of the 1 : 1 chelate and the free ligand. This observation is in agreement with the continuous variation work of FOLEY and ANDERSON~19) who showed the existence of only a I : 1 UO2-SSA chelate. BANKSand SINGHtls) reported that 1 : 1 and 1:2 complexes of UO2-SSA are formed at pH values 4.5 and 7.5 respectively, based on their spectrophotometric measurements of solutions containing 1:2, 1 : 6 and 1 : 10 molar ratios of metal ion to ligand, and calculated stability constants from potentiometrie data. Although the constant reported by them for a 1 : 1 complex is very close to the one obtained in this research, no evidence is apparent in Fig. I for the presence of a 1 : 2 complex. Also, precipitation was observed above pH 5 in both 1 : i and 1:2 systems. A possible structure of the diaquo 5-sulpho-salicylato chelate compatible with experimental evidence is the following:
d/ I
Equilibrium studies of uranyl complexes--I
797
Treatment of the hydrolytic reactions taking place beyond m = 3 (Fig. 1) could not be carried out in view of the precipitations observed beyond m = 3.3. It is evident from the data presented in Table 1 for the UO2-IMDA chelate compound that the log KuL values obtained from the 0.1 M KNO3 medium increase slightly with decreasing concentration of the metal chelate. This observed trend cannot be due to polymer formation, since the formation of polynuclear complexes would give a trend in the opposite direction. It may be due to small changes in the ionic atmosphere with increasing contribution of uranyl salt. An average Log K:vi, value of 8.76 -k 0.03 was obtained in 1.0 M KNOz medium. It is therefore concluded that the UO2-IMDA chelate is present in aqueous solution only as a monomer. A possible structure for this complex is indicated by I1. /co--cH~.. 0~- - - 0 - - - 7 N H I\ /'~. \ II / / CH z
- -
-o-
-
-:6
j
II
Analysis of the 1 : 2 UO=-IMDA titration curves indicate the presence of only the 1 : 1 complex. From the observed trend in the log KMr, values obtained in 0.1 and 1-0 M KNOz media (Table 1) for the 1 : 1 UOz-HIMDA chelate, it is concluded that no polynuclear complexes are formed in this system in the presence of up to two equivalents of base per mole of the metal salt. A possible structure for the 1:1 chelate in this buffer region is the following: _/CO~CH2 _
/' /
~ /
/ " Itf1 \ \
/
/CHzCH20H
CO
~o . . . . . o . . . . oj III
On the basis of the mathematical treatment of potentiometric data in the buffer region between m = 2 and m = 3 (Figs. 4 and 5 and Table 1) it is concluded that a hydrolysed and polymerized species (IV) is present in solution under these conditions:
"°°'.;7 , coo/O
\c.
. IV
However, on the basis of potentiometric data alone it is not possible to determine if hydrolysis occurs by dissociation of a proton from a coordinated water molecule or from a hydroxyethyl group of the ligand. A comparison of the 1 : 1 chelate stability constants of various metal salts with nitrilotriacetic acid (NTA), N-hydroxyethyliminodiacetic acid (HIMDA) and iminodiacetic acid (IMDA) presented in Table 3 suggest
798
K.S.
RAJAN a n d A. E. MARTELL
TABLE 3.--COMPARISON OF CHELATION TENDENCIES OF N T A , H I M D A AND I M D A Log K Metal i o n
C u ~+ N i z+ Co 2+ M n ~+ Z n 2+ C d ~+ P b ~+ M g 2+ C a 2+ UO2 ~+
N T A t~j
H I M D A c2~
IMDA
12"7 11"3 10'6 7"4 10"4 9"5 11"8 5"4 6'4
10 9"5 8"3 5"6 8"6 7"1 9.5 3"5 4"8 8"32
10'6 c27~ 8-3 (27~ 7"0 taT~ 7"0 t~7~ 5"3 c ~ 3.6 ~s~ 3"4 c28~ 8"93
that the hydroxyethyl group of HIMDA is not coordinated to the uranyl ion in the 1 : 1 chelate below pH 4. It is seen that for the transition metal ions and many other metal ions the HIMDA chelate has a larger formation constant than does the corresponding IMDA chelate. Since HIMDA is the less basic ligand, this correlation has been taken to indicate involvement of the aliphatic hydroxyl groupin the co-ordination of the metal ion. In the case of the uranyl complexes, the lack of involvement of the hydroxyethyl group in co-ordination indicates that it also may not be involved as a bridging group in the polynuclear complex. ~,5~ G. SCHWAILZENBACH,Helv. Chim. Acta, 32, 1175 (1949); 34, 1492 (1949). t2e~ S. CHABEREK, R. C. COUgTNEY a n d A. E. MARTELL,. L A m e r . Chem. Soc. 74, 5057 (1952). t27~ S. CaABEREK a n d A. E. MARTELL,.]'. Amer. Chem. Soc. 74, 5052 (1952). (g8) G. SCHWARZENBACH et ak, Heir. Chim. Acta, 28, 1133 (1945).