634
THOMAS L. JAMES AND RICHARD J. KULA
the following phase boundary equilibria
+ Fe2+V + 2Fea+V + 8eNull = 202-V + 4e-V + OZ
Doz- = constant X (PHz/PHz0)1/3
202 = Fe304
(9)
(10)
where the symbol V indicates a vacancy, e- represents an electron, and Null denotes a stoichiometric crystal without lattice deffects. These reactions can be expressed in terms of their equilibrium constants as P202 =
constant CFezfV X
X CSe-
constant = C2~2-vX C4e-v X Po2 but C
(11) (12)
F = ~~C F ~, ~ +ceV ~ ; =~~ C F , ~ +C,-v V ; = 2coz-v
:.
C F ~ Z + V= C O Z - ~=
constant
PO^^/^^
constant Po2-l’e
(13) (14)
Equations 13 and 14 for binary oxides can be compared directly with eq 4 and 5 which refer to the ternary oxide. Combining eq 13 and 14 with eq 6 and again assuming a linear dependence of the self-diffusion coefficient on the defect concentration, we obtain
D F ~ =+ constant X ( P c o / P c o J ~ / ~ ~ (15)
(16)
The exponents are in better agreement with the measured values than those predicted for ternary oxides. This supports the probability that the structure of magnetite is randomized within the temperature range of the studies (>500°) : it is not therefore a useful oxide with which to test the Wagner-Schmalzried model. The response of the oxygen diffusion coefficient to the oxygen partial pressure is particularly satisfactory since it shows that the oxygen sublattice remains in thermodynamic equilibrium with the atmosphere, even in an oxide in which the bulk of the stoichiometric deviations are accommodated in the cation sublattice. The change in the value of the self-diffusion coefficient over the complete composition range of magnetite is only one order of magnitude. This result, therefore, does not alter our earlier conclusion that the solid-state diffusion of oxygen is unlikely to be important in the oxidation of iron by steam or water.
Acknowledgments. The authors are grateful to the Central Electricity Generating Board for permission to publish this work, and also to Professor J. S. Anderson for valuable discussions.
A Comparison of the Proton Affinities of Neutral Oxygen and Sulfur in Chelating Ligands1 by Thomas L. James and Richard J. Kula Department of Chemistry, University of Wisconsin, Madison, Wisconsin 53706
(Received August #?9, 2 9 6 8 )
The nuclear magnetic resonance chemical shift behaviors of the iionexchangiiig carbon-bonded proton resonances of various 8-alkylthioacetic acids and alkoxyacetic acids were studied in aqueous solutioii as a function of solvent acidity to determine the relative proton affinities of the donor atoms in such chelating ligands. The relative electron-donating abilities, ascertained from the sites and sequences of protonatioa, were found to be in the order 00 0 8 This order of proton affinity seem to reflect the relative contributions of these donor atoms t o the stability of metal chelates for those metal ions with little tendency for the formation of dative s bonds. The suitability of internal references in the strongly acidic solutioiis used in this investigation was also studied. Preliminary proton nuclear magnetic resonance (nmr) studies of sonic metal ion complexes of various 8-alkylthioacetic acids (R-S-CH2COOH) indicated little, if any, metal-sulfur interaction. Differences in metal-sdfur bond strengths for different ligands are noverned bv the Lewis basicity of the sulfur atom and the teadel1cy for back-doiiatiol1 p bollding of sulfur. For transition inetal ions, the strength of the metal-
-
The Journal OSPhgsical Chemistry
thioether sulfur bonding is apparently determined primarily by the latter mechanism-formation of dative .R bonds with the empty d orbitals of sulfur.2 (1)This work was supported by a grant (GP4423) from t h o National Science Foundation and by a xational Institutes of Health doctoral Fellowship (T.L. JJ. (2) (a) H. H. Jaffe, J. Phys. Chem., 58, 185 (1954); (b) D. P. Craig, A. Maccoll, R. S. Piyholni, L. E. Orgel, and L. E. Sutton, J. Chem. sot., 332 (1954).
PROTON AFFINITIES OF NEUTRALOXYGENAND SULFURIN CHELATING LIGAKDS In the absence of back-donation, the Lewis basicity of the sulfur is presumably the predominant factor determining the metal-sulfur bond strength. Studies have been made of metal ion coordination with several S-alkylthioacetic acids and also with the corresponding alkoxyacetic acid (R-O-CH2COOH) ligand^.^ While these studies indicate that the extent of metal thioether sulfur bonding and metal-ether oxygen bonding depends on the structure of the ligand as well as the metal ion, no definitive data are available comparing the relative basicities of thioether sulfur and ether oxygen. One method of determining the relative Lewis basicities of oxygen and sulfur donor atoms is to compare the proton affinities of these atoms in a variety of carboxylic acid ligands containing thioether sulfur or ether oxygen. If there are no steric factors which affect the bonding, the proton affinities of the donor atoms in the ligands should in turn reflect the strength of the metal-donor bond for those metal ions which are incapable of forming dative ?r bonds with the donor atom. Rather than relying on comparisons of niacroscopic protonation constants for analogous ligands in order to determine the relative basicities of the donor atoms, the sites and sequences of protonation have been ascertained using proton nmr. The protonation schemes of ethylenebisthioglycolic acid (ETGA) , thiodiglycolic acid (TDG), (ethylthio) acetic acid (ETAA) , ethoxyacetic acid (EOAA) , acetic acid (HAC), dimethyl sulfide (DMS), and diethyl ether (DEE) were determined in aqueous solution.
Experimental Section Proton nuclear magnetic resonance spectra were obtained with a Varian A-60A high-resolution spectrometer operated at a probe temperature of 25 f 2'. Tetramethylammonium chloride, at a concentration of about 0.02 M , was used as an internal reference. The chemical shifts, v, are reported in cycles per second from the central resonance of the tetramethylammonium (TMA) triplet , negative shifts corresponding to a resonance downfield from TMA. The spectruni of each sample was obtained five tinies and the reported chemical shift is an average of these five individual measurements. Solution pH measurements were made at 2.5" using a Sargent Model DR line-operated pH meter equipped with a wide-range glass electrode. Conventional NBS buffers were employed for standardization of the pH meter. Reagents were of the highest available commercial purity and were used without purification with the exception of ethylenebisthioglycolic acid which was recrystallized froni water. Aqueous solutions of the reagents were prepared by weighing the requisite amounts of the chemicals in the case of solid reagents or by pipetting in the case of liquid reagents, and
635
diluting to the prescribed volume. The reagent concentration was normally between 0.1 and 0.2 M , and in this limited range no concentration dependence of the nmr spectra was observed. Solutions were adjusted to the desired pH using Concentrated potassium hydroxide or perchloric acid. Strongly acidic solutions were prepared with perchloric acid dihydrate (determined to be 73.5% by weight). Because the S-alkylthioacetic acids were found to be oxidized in perchloric acid at this concentration, the upper concentration limit for perchloric acid was about 60%. In some cases, additional data were obtained using concentrated hydrochloric acid for preparing the acid solutions. To check for possible chemical degradation or alterations of the organic ligands in the concentrated acids, solutions of the ligand were neutralized and the nmr spectra were obtained. If the spectra of the neutralized solutions were identical with those of the ligands which had not been acidified, it was presumed that no alteration of the ligand had occurred. The use of concentrated sulfuric acid as a proton source was discontinued after finding that ETGA decomposed in such a solvent. Solutions of the organic reagents ranged in acidity from pH 12 to 10 M perchloric acid in this study, but atj present no adequate scale exists which will give a quantitative measure of acidity over such a range. In principle a scale based on the Hanimett acidity function should be applicable. However, the Haniniett acidity function may be utilized with confidence only in those cases where the compounds under observation have structural features which are similar to those of tlic indicator bases used in establishing the acidity scale.4 For lack of a better scale, the acidity of the systems investigated here will be described by pH. For solutions of pH greater than 0.5, the pH signifies an activity function determined from the pH meter; for solutions of greater acidity, pH will..signify the negative logarit hili of thc hydrogen ioii molarity. This approach is justified because no absolute quantitative measurements are made in thc strongly acidic region. Good correlation of the chemical shift data mas obtained using hot11 types of pH measurement in the 0 t o 1 pH range. The acid dissociation contaiils for the loss of a proton from the neutral carboxylic acids were calculated from nmr data. From the chemical shifts of the carbonbonded ligand proton resonances at each pH, the coiiceiitrations of the protonated and iionprotoiiated forms of the compound were determined.6 The cal(3) (a) M. Yasuda, K . Yamasaki, and H. Ohtaki, BuZl. Chem. SOC. Japan, 3 3 , 1067 (1960): (b) K . Sueuki and K . Yamasaki, J . Inorg. Nucl. Chem., 24, 1093 (1962); ( c ) A . Sandell, Acta Chem. Scand., 15, (4) (5) 27,
190 (1961). K . Yates and H . Wai, J . Amer. Chem. Soc., 86, 5408 (1964). E. Grunwald. A. Loewenstein, and 9. Meiboom, J. Chem. Phys. 641 (1957).
Volume Yd, Number d
March 1060
636
THOMAS L. JAMES AND RICHARD J. KULA
culated dissociation constants contain the conjugate acid-base pair terms as concentrations and the hydrogen ion term as an activity function; the hydrogen ion activity was not converted to concentration because the error incurred in measuring the concentrations of the other species from the nmr spectra is greater than the magnitude of the activity correction. Also, no attempts were made to maintain a constant ionic strength by adding excess inert electrolyte, because these nmr studies necessitated the use of high concentrations of the species being studied. The experirrientally determined acid dissociation constants in
'
v)
a.
2
0-
is
Table I: Acid Dissociation Constants for Neutral Carboxylic Acids Calculated from Nmr Data Acid
Acid dissociation con stan t
ETGA TDG ETAA EOAA
5
4 X 2.4 x i o - 4 = 2 X 10-4a 7 . 1 x 10-4a 1.8 x 1 0 - 4 2.43 x 10-4 2.7 x 10-4 3 . 5 x 10-4
Reference This work 3b (at 25" a t 1 = 0.1 with NaClO,) This work 3a (at 25" a t 1= 0.1 with NaC104) This work 3c (at 20" in 1.00 M NaCIOa) This work 3c (at 20" in 1.00 M NaClOJ
(KIK2)"2.
Table I are compared with previously determined values of these constants.
Results Solvent Efect on Tetramethplammonium Ion. The tetramethylammonium ion (TMA) is a useful internal reference for measuring chemical shifts in aqueous solution because of its solubility, its closely spaced triplet spectrum which enables adjustment of the magnetic field homogeneity, and its chemical shift (3.20 ppm downfield from tetramethylsilane, TMS) which is in the region of many aliphatic ligand resonances. Further, in aqueous solutions the TMA chemical shift is generally unaffected by the presence of other ions or molecules. The first indication that TMA might be subject to solution effects was the anomalous upfield chemical shifts of the ETGA resonances (Figure 1) in solutions made progressively more acidic with HC1. This behavior was not observed when the solutions were acidified with liClOc and suggested an interaction of TMA with C1-. In order to determine the extent of specific ion or solvent interaction with TlIA, measurements were made of the chemical shift difference between the methyl resonances of TMA and sodium 3-(trimethylsilyl) -1-propanesulfonate (TMS*) for a variety of solution conditions. In neutral aqueous solutions the methyl resonance of 'l'MS* is 190.2 cps upfield from The Journal of PhvslsicaZ Chemistry
-30' -
t
-I
I
I
I
I
'
3
5
I
' 7
PH Figure 1. Chemical shifts of ethylenebisthioglycolic acid (ETGA) proton resonances as a function of pH.
TMA. The results of these studies are given in Table 11. These results can be interpreted (assuming that the methyl proton resonance of TMS* is not affected by C1- and that the TMA resonance is not affected by H+) in terms of an ion-pair interaction of TMA with C1- and a protonation of TMS". The ion pairing of TMA with C1- results in a perturbation of the TMA electronic structure, a decreased shielding o f the methyl protons, and consequently a downfield chemical shift of the TMA re~onance,~" manifested as a larger chemical shift difference between TMA and TMS". In the strongly acidic solutions, TMS" is Table 11: Chemical Shift Differences between the Methyl Resonances of Tetramethylammonium Ion and 3- (Trimethylsilyl) -1-Propanesulfonate under Various Solution Conditions Solution HzO
4 M NaC104 4 M NaCl 4 M HCl 8 M HCl 4 M HC104 6 M HClOi 10 M HClOd
AV = WM8v
- VTMA
190,2Q 189 1" 192.3 192.2 192,6 189.3 188.8 187.6 I
The difference in chemical shifts for these two solutions may be due t o association of the sulfonate group of TMS* with Na+, ion association of TMA with clod-, or some combination of the two effects in the concentrated NaClOd solution.
(6) A. D. Buckingham, Can. J. Chenz., 38, 300 (1960). (7) A. D.Buckingham, T. Schaefer, and W. G . Schneider, J. Chem. Phys., 32, 1227 (1960).
PROTON AFFINITIES OF NEUTRAL OXYGENAND SULFURIN CHELATING LIGANDS
637
Table 111: Sequence of Protonation and Chemical Shifts of the Variously Protonated Species of Chelating Ligands Containing Neutral Oxygen and Sulfur -Site
for addition of nth proton----
C
---
Chemical shiftsQ
I
0-
va
ETGA
0 1,2b
TDG
0 1,2b
ETAA
0 1
EOAA
0 1
HAC
0 1
DMS
0
DEE
0
-3.6 -14.8 -27.2c -4.1d -18.3d -22.5dsc -1.8 -13.2 -26.5c -43.2 -60.5 -76.70 75, 6d 64. gd 59.2dva 64.3 17.6~ -23.3' -57 * 1 c J
Symbol
3, 40
3,46
2=
20
26
10 10
Vb
VO
22.2 16.4 12.20
35.2f 31 .Of 25.3C8.f -23.3f -28.01 - 43,209,
117.2f 116.3' 112,O"J 118.9f 117.8f 110.2cJ
120.Of 104.9cgf
Value in a Chemical shifts in cps from internal tetramethylammonium ion (TMA) a t 25". b Both carboxylate groups are protonated. 9.7 M HClOa; protonation not complete. d At 38". 8 Value in 12.1 M HC1; protonation not complete. f Denotes center of multiplet.
protonated to form the sulfonic acid resulting in a small downfield shift of the TMS* methyl resonance, manifested as a smaller chemical shift differencc between TMA and ThIS*. This interpretation is contrary to the conclusions of Abraham and Thomas who failed to recognize the possibility of interaction between hydrochloric acid and tetramethylammonium ion and the shift of the TMS* resonance to lower fields upon protonation.*
a
1) ( C . P . 8 . ) b
35
c 120
11:
I IC
30 25
-10
-20
I
-3 0
-2
C H3C H@-CH2CO; c b a
0
2 .
4
6
PH
Figure 2. Chemical shifts of (ethylthio) acetic acid (ETAA) proton resonances as a function of pH.
8
Tetramethylsilane (TAIS) and TAIS* are unsuitable refereiicerj for these studies. T A I S is insoluble in aqueous solutions and TMS*, like a number of other cominon references, is itself being protonated in the region of interest. Although it could hardly be expected that TMA would remain completely adamant to the solvent change, especially in the concentrated solutions of perchloric acid, it is felt on the basis of the results observed that the use of TMA as an internal reference in perchloric acid solutions is probably superior to the use of other internal references or of an external reference, which would require corrections for bulk diamagnetic susceptibility differences for all solutions encountered because of the wide range of solvent composition : from water to concentrated perchloric acid. Observed Spectra. The nmr parameter of interest in this investigation is the chemical shift of resonances for the nonexchanging carbon-bonded ligand protons. The chemical shifts of these protons are pH dependent as seen in Figures 1-5. The breaks in the chemical shift us. pH curves correspond to protonation at one or more basic sites in the ligands. The structures and resonance assignments for the ligands are shown in the figures and in Table 111. The proton nmr spectra for ETGA have two singlets of equal intensity; the low-field resonance has been assigned t o the acetate protons (labeled a in Table 111) and the upfield resonance has been assigned to the b ( 8 ) R. J.
Abraham and W. A. Thomas, J. Chem. ~ o c . 3739 , (1964). Volume YB, Number B March 1060
638
THOMAS L. JAMES AND RICHARD J. KULA v
a
b
-40
was slowly oxidized. The downfield chemical shift change between pII 1 and pH -1 is 12.4 cps for the a resonance and 4.2 cps for the b resonance. The nmr spectrum of T D G was characterized by only a sharp singlet whose chemical shift had approximat'ely the same pH dependence as the a proton resonance of ETGA.
(C.P.83
c
-20
115
-25
110
-30 -50
-35 -4 0
-6 0
Y
-4 5
a
EOAA -7 0
I20
C H j C H 2 - 0 - C H2C 02 c b a
-
-80
-2
0
4
2
(C.P.S.)
6
8
PH
Figure 3. Chemical shifts of ethoxyacetic acid (EOAA) proton resonances as a function of pH. -40
protons. Potentiometric titrations revealed that two equivalents of acid were consumed by the ETGA dianion between pH 6 and 2. In this pH range the a resonance was shifted downfield by 11.2 cps and the b resonance was shifted downfieId by 5.8 cps. In stronger perchloric acid solutions, the resonances are shifted still further downfield but the chemical shifts have not reached a constant value even iii 9.7 f?! perchloric acid. In stronger solutions of perchloric acid, ETGX
60-
50' 7 u)
& 40L) w
h
30 -
DMS CH3-S-CH3
20 -
-5 0
-6 0
ti11
DE E CH3CHfO-CH2CH3 b
o
t : -I
I
3 PH
5
7
Figure 5. Chemical shifts of diethyl ether proton resonances as a function of pH.
The nnir spectra of ETAA are distinguished by three groups of resonances with integrated intensity ratios a:b:c of 2:2:3. The a proton resonance is a sharp singlet domiifield from the TJIA resonance. Spinspin coupling between the b and c protons produces the typical ethyl group triplet at higher fields for the c resonances and quartet at lower fields for the b resonances. In solutioiis of positive pH, a small amount of second-order coupling was evident and the triplet and quartet were further split into an A2B3pattern. The extent of the second-order coupliiig mas sufficiently sinall that for most purposes, the center of each multiplet could be taken as its chemical shift. In strongly acidic solutions, the pattern simplified to an A2X3. In the positive pH region, simple protonation results in an 11.4-cps downfield shift for the a resonance, a 4.2-cps downfield shift for the b resonance, and a 0.9-cps downfield shift for the c resonance. The a resonance line shift compares with the 11.2-cps shift for the a resonance of ETGA. The b resonance shift plus the c resonance shift gives approximately the same value as the b resonance shift of ETGA. Further acidification results in downfield shifts similar to those of the ETGA resonances.
; -I
I
3
5
PH Figure 4. Chemical shift of dimethyl sulfide proton resonance as a function of pH. The Journal of Physical Chemistry
7
PROTON AFFINITIESOF NEUTRALOXYGENAND SGLFURIN CHELATING LIGANDS The nmr spectra of EOAA are similar to those of ETAA except that the a and b resonances of EOAA are further downfield than the corresponding resonances of ETAA. The large chemical shift between the b quartet and c triplet is maintained over the entire pH range, obviating any second-order coupling phenomena. In contrast to ETAA, the a and b resonances are shifted downfield to approximately the same extent below pH 1. Dimethyl sulfide (DAIS) gives a single resonance upfield from TMA. The spectra for diethyl ether (DEE) approximate AzXa patterns. Both DMS and D E E show only one break in the chemical shift curves and in neither case has the chemical shift reached a constant value, even in the most concentrated acid solutions studied.
Discussion The chemical shifts of the various ligand resonances depend upon the solution pH, and in all cases the resonances are shifted downfield as the solutions are made more acidic. When a proton associates with a basic site of the ligand, a deshielding effect is produced on the ligand protons which results in a downfield shift of their resonances. Because the deshielding effect attenuates with distance from the site of perturbation, the magnitudes of the downfield shifts depend on which donor site within the ligand associates with the acidic proton, The resonances for those ligand protons which are further from the site of acidic proton association will be shifted less than the resonances for the ligand protons which are closer t o the protonated donor site.s In the ETGA system, the downfield shifts between pH 6 and 2 must correspond to -OOCCH&CH&HzSCHzCOO-
+ 2H+ +HOOCCHzSCH&HzSCHzCOOH The greater chemical shift change for the a than for the b resonance indicates that the carboxylate oxygens are the sites of protonation in this region which is in agreement with the consumption of two H+/EGTA in this region as was determined potentiometrically. The upfield shift of the ETGA a and b resonances below pH 1 in HC1 can be attributed t o ion pairing of TMA with C1- as discussed previously. The ETGA resonances are shifted downfield only slightly between pH 1 and 0 (cf. Figure 1 in HC104), but ion pairing of the reference ion, TMA, with C1- shifts the TMA resonance downfield, resulting in an apparent upfield shift for the ETGA resonances, Further proton interaction is indicated in solutions of negative pH by additional shifts of the resonances to lower fields. The greater downfield shift of the a resonance relative to the b resonance is not commensurate with protonation of the thioether sulfur. A more reasonable explanation would involve proton
639
interaction at the carbonyl oxygen
R-C
/OH
\
ZR-C
OH
+
/ \
tjH ZR-C+ OH
/OH
\
OH
The structures depicted may be formed via an intermediate in which the carbonyl oxygen is hydrogen bonded to the acidic solvent.'O For a comparison chemical shifts of the TDG resonance and the methyl proton resonance of acetic acid were found to have the same pH behavior as the a resonance of ETGA. Previous studies of acetic acid in concentrated sulfuric acid and in fluorosulfuric acid have provided evidence that the carbonyl oxygen is the site of protonation."'l2 Deno, et al.," obtained a change in chemical shift of 34 cps for the complete protonation of the carbonyl oxygen in sulfuric acid, which indicates that the protonation is no more than one-third complete in any of the cases considered here. The pH dependence of the chemical shifts for the various resonances of ETAA further substantiate the carbonyl oxygen as the site of the second protonation for S-alkylthioacetic acids. Thus, for ETAA below pH 1 the a resonance is shifted more than twice as much as the b resonance which is shifted further than the c resonance. The pH dependencies of the resonances of EOAA, the oxygen analog of ETAA, are depicted in Figure 3 . The behavior upon protonation of the carboxylate group (pH 6-2) is analogous to that of the S-alkylthioacetate ligands. In more acidic solutions (from pH 1 to - l ) , the a resonance of EOAA shifts 16.4 cps downfield, the b resonance shifts l5,3 cps downfield, and the c resonance shifts 7.7 cps downfield. The magnitude of the shifts of these resonances is not in accord with those of the corresponding shifts of the S-alkylthioacetic acids, which have been postulated to undergo protonation of the carbonyl oxygen. Rather, these results suggest that the predominant proton interaction occurs at the ether oxygen because the b resonance shifts approximately the same amount as the a resonance. Edward, Leane, and Wang investigated the chemical shift difference between the methyl proton resonance and the methylenic proton resonance (which they designated the "internal shift") of diethyl ether as a function of acidity in sulfuric acid.1° They interpreted the observed increase in (9) See, for example: (a) E . Grunwald, A . Loewenstein, and 9. Meihoom, J. Chem. P h y s . , 27, 641 (1957); (h) A. Loewenstein and J. D. Roberts, J . A m e r . Chem. Soc., 8 2 , 2705 (1960);(c) J. L. Sudmeier and 0 . N. Reilley, A n a l . Chem., 36, 1968 (1964). (10) J. T. Edward, J. B. Leane, and I. 0 . Wang, Can. J. Chem., 40, 1521 (1962). (11) N. C. Deno, C. I?.Pittman, Jr.,and M. J. Wisotsky, J. A m e r . Chem. Soc., 86, 4370 (1964). (12) T.Birchall and R. J. Giillespie, Can. J. Chem., 43, 1045 (1965).
Volume YS, Number S March 1969
THOMAS L. JAMESAND RICHARD J. KULA
640 “internal shift” with increasing acidity in terms of a mechanism involving hydrogen bonding of the solvent H 2 0 to the ether oxygen, followed by formation of the protonated ether species. Whether the shifts observed in the present study, in which the acidity is not so great as in the study of Edward, Leane, and Wang, are attributable to a hydrogen-bonding phenomenon or to a direct protonation is not determinable. However, whichever phenomenon predominates, it is clear from the relative chemical shift changes that the basic site within the ligand, which is associated with the acidic proton, is the ether oxygen. Therefore, in the alkoxyacetic acids the ether oxygen has a greater proton affinity than the carbonyl oxygen. These results reveal, on a qualitative basis, the differences in intrinsic Lewis basicities of neutral oxygen and sulfur. In the S-alkylthioacetic acids and the alkoxyacetic acids studied here, the donor tendencies are in the order
On the basis of this ordering, the ether oxygen is expected to contribute significantly more to the stability of chelates of metal ions having little capacity for back-bonding than the thioether sulfur in corresponding ligands. An investigation of the relative stability constants of some S-aryl and S-alkylthioacetic acids, including ETGA, and their oxygen analogs with the divalent ions of Zn, Cd, and P b coiicurs with the conclusions reached lierc; i.e., those ions with filled d orbitals showed greater affinity toward oxygen tliaii sulfur.3 The observed basicity of the carboiiyl oxygen has
The Journal of Physical Chemistry
interesting implications concerning the -C
// \
0
0-
group as a possible bidentate chelating group. Bidentate carboxylate bonding has been detected from X-ray studies of Na[UOz(CH&OO)3],13 Zn(CHsCOO)2.2H2014 and Cu[H21T-CH (CHzCHzCOO)-COO] -2H20L5 The carbonyl oxygen may also be utilized in polymeric chelates in which -COO- is an asymmetric bridging ligand; this type of structure has been detected by X-ray studies of Zn (HzNCH2C00)2-Hz0 and Cd (H2NCH2COO)2.H2O.l6 More recently, infrared studies of several crystalline metal chelates of methionine, CHsSCH2CH2CHNHzC0OH, have indicated that metal-carbonyl oxygen bonding is favored over metal-thioether sulfur bonding.” The results of the present work suggest that perhaps the carbonyl oxygen may be a factor to be considered in chelate formation in solution as well. (13) W. H. Zachariasen and H. A. Peltinger, Acta Cryst,, 12, 526 (1959). (14) J. H. Talbot, ibid,, 6 , 720 (1953), (15) H. 0.Freeman in “The Biochemistry of Copper,” J . Peisach. P. Aisen, and W. E . Bluinberg, Ed., Academic Press, Kew Yorlt, N. Y., 1966, p 84. (16) B . M. Low, I?. L. Hirshfeld, a n d F . &I. Richards, J. Amer. Chcm. SOC.,81, 4412 (1959). (17) C. A. RlcAuliffe, J. V. Quagliano, and L. R.5. Vallarino, Inorg. Chem., 5 , 1996 (1966).
