Solution Equilibria and Structures of Molybdenum(V1) Chelates N-Methy Iiminodia cet ic Acid RICHARD J. KULA Department of Chemistry, University of Wisconsin, Madison, Wis.
Nuclear magnetic resonance techniques have been employed to study the aqueous solution behavior of the methyliminodiacetate complexes of molybdenum(VI), all of which form with metal-ligand ratios of one. This stoichiometry is governed by the strong association of molybdenum(Vl) with oxygen, the central metal coordinating unit being MoOa. In alkaline solutions there are competing equilibria between molybdenum-methyliminodiacetate formation and free ligand protonation, whereas in acidic solutions there are competing equilibria between protonation of the chelate, the formation of dimeric chelates, and dissociation. The equilibrium constants for these processes have been determined by studying the concentrations of the various ligand species as a function of pH. Structural and bonding features of the complexes are discussed from considerations of their nuclear magnetic resonance spectro.
S
TUDIES OF METAGAMINOPOLYCARBOXYLIC acid chelates being carried
out in this laboratory have led to the investigation of several molybdenum systems. Molybdenum is capable of coordinating with an enormous variety of ligands as pointed out in the recent monograph by Busev on the analytical determination of this element (8). Interest in this laboratory has been focused on the +6 oxidation state, in which the element is diamagnetic, thereby enabling normal high resolution proton nuclear magnetic resonance (NMR) techniques to be employed for studying its complexes. NMR techniques are particularly advantageous for studying Mo(V1) complexes because the usual techniques for equilibrium studies are not always applicable. The absence of absorption bands above 250 mp for the metal makes spectrophotometric methods generally inappropriate unless the ligand itself absorbs at longer wavelengths. Also electrochemical methods have been found to be less suitable than for other metal ion systems because of the complicated equilibria 1382
0
ANALYTICAL CHEMISTRY
and electrode processes involved, as recently illustrated by studies of Mo(V1) in acid media (80). Although Mo(V1) forms complexes with the same chelating agents as most other metal ions, its chemistry is differentiated from other transition ions by its strong association with oxygen. Thus, in most Mo(V1) complexes, Mo02+* or MOO* is the central coordinating unit, which with octahedral geometry, severely limits the number of coordination sites available to the ligands. An interesting consequence of this behavior is illustrated by the complex formed between Mo(V1) and ethylenediamine-tetraacetic acid (EDTA) in which two Mo ions can coordinate with EDTA (14). This is in distinct contrast to the normal metal-EDTA chelates in which only one metal ion coordinates with each ligand. The structures which were proposed for the MorEDTA chelate led to speculation concerning the possibility of forming one to one Mo(V1) chelates with methyliminodiacetic acid (MIDA) and nitrilotriscetic acid (NTA), and proton NMR studies of these chelates confirmed that the predominate EDTA chelate does indeed contain two Mo ions and that stable, one to one molybdenum chelates are formed with MIDA and NTA (4). In this paper detailed investigations of the aqueous solution equilibria of the molybdenum chelates of MIDA and some further examinations of the structural features of the chelate are reported. EXPERIMENTAL
Proton NMR spectra were obtained on a Varian A-60 spectrometer. The 100 Mc. spectra were obtained by Ross Pitcher of Varian Associates on an HA-100 spectrometer. The sweep rates were normally between 0.2 and 1.0 c.p.s./second, and r.f. fields of less than 0.2 milligauss were employed. All of the spectra were recorded at a probe temperature of 35 f 2’ C . , and were obtained in either aqueous solutions or in solutions prepared using D?O. Tetramethylammonium chloride at a concentration of about 0.03M served as an internal reference. All chemical shifts, Y, we reported in C.P.S. from the central
resonance of the tetramethylammonium ion (TMA) triplet, which appears about 3.20 p.p.m. down-field from tetramethylsilane. Its resonant position is almost completely independent of the other ions in solution. In addition to its use as an internal reference, the TMA gives an indication of the field homogeneity which is adjusted until complete resolution of the three components (separated by about 0.5 c.P.s.) is achieved. The solution pH was measured with a Leeds and Northrup line-operated pH meter equipped with a high-range glass electrode. For measurements in D20 solutions, micro electrodes were employed and the “pH” was corrected for deuterium effects using the equa. ~ as ~ ~ ~ tion, PH = ( P H ) ~ ~ ~0.40, described by Glasoe and Long (8). For both electrode systems the meter was standardized using conventional N.B.S. buffer solutions, and all measurements were made at 25’ C. All the chemicals were analytical reagent grade and were used without any purification. The MIDA was obtained from hldrich Chemical Co. in the form of the diprotonated acid. Molybdenum was introduced either as anhydrous molybdic oxide or as sodium molybdate dihydrate. Crystalline Mo-MIDA was prepared by making up a solution 1M in NazMoO,, and 1M in HzMIDA, from which the chelate was precipitated as white, grainy crystals by adding about three volumes of 95% ethanol for each volume of aqueous solution. The crystals were then filtered, washed with ethanol, and dried over magnesium perchlorate in a vacuum desiccator for several days. An elemental analysis showed the material to have the composition NazMoOaMIDA.4HzO. Solutions of the chelate were prepared either hy weighing the requisite amounts of reagents and diluting to the desired volume or by weighing out the crystalline chelate and diluting. The chelate concentration was normally between 0.2 and 1.OM, and no dependence was noted on the ICMR spectra in this concentration range. The solution pH was monitored aa standard sodium hydroxide on nitric acid was added, these reagents being about 5M to avoid dilution effects, and 0.5-ml. samples were removed for the NMR studies. For investigations in acidic solutions where the ligand proton resonances are close to the water
+
resonance, signal distortions and extraneous lines from spinning sidebands interfere with the measurements. In such cases 99.8% DzO, obtained from Bio-Rad Laboratories, was utilized as a soivent. NMR signal intensities have been measured by either electrical integration of the resonance areas with the A 4 0 integrator when there is no overlap of signals, or by peak-height comparisons for overlapping signals. The former method is superior, no corrections being required for “cross-talk” from overlapping resonances. Also with peak-height comparisons one BSsumes that the line widths of the resonances being compared are equal, an assumption which is not always valid for metal-chelate systems. Generally, signal intensities can be determined within about 10% or within 2570 in the extreme cases. Such errors should not detract seriously from the qualitative interpretations, although the accuracy of the equilibrium constant calculations are lower than one would desire. In calculating constants for the various equilibria described in the Results section the following procedure was employed. The NMR spectrum was obtained at a measured pH. From the chemical shifts of the resonances and/or their relative intensities, the concentration of each component involved in the appropriate equilibrium was calculated. Thus, all the species except H + in the equilibria are given as concentrations rather than activities. The Concentration of H + in these equilibria is taken from the measured pH which is, of course, an activity function. However, it was not deemed necessary to make a conversion from hydrogen ion activity to concentration because the magnitude of the correction would be much less than the error in determining the concentrations of the other species from the NMR spectra. Although an equilibrium constant calculation could be made a t a single pH, a more reliable value can be obtained by calculating the constant a t a number of different pH values. The constants given in Table I are each the average of at least five independently calculated values at diiTerent pH’s, and the estimated error represents the range of the observed values. Also it should be noted that because of the high concentrations which must be employed for the NMR studies (minimum of 0.1M in metal and ligand), no excess “inert electrolyte” has been added and the ionic strength of the solutions is continuously changing as the pH changes. This behavior is an unfortunate and unavoidable c o m e quence of using NMR for equilibrium studies but is compensated somewhat by the additional information which is gained about the structural, bonding, and kinetic aspects of the complexes.
Of 0
I
I
I
I
I
2
I
I
3
I
I
4
moles H + moles Mo or MIDA Figure 1.
pH curves for the titrations of NasMoOG of MIDA-*, and of 1 : 1 Mo04-2-MIDA-* solutions MOO,-’
MQ+ 0.20M M I D A 3 0 n a = 0.20M Tihated with 6.22M HNOi T = 25O C.
ficiently slow that separate resonances can be observed for the ligands coordinated to Mo and for the free ligands (4). By varying the Mo to ligand ratio it was shown that Mo(V1) coordinates in a one to one mole ratio with MIDA; the results presented here substantiate this conclusion and indicate that this is the only combining ratio of metal and ligand at all pH’s. When studies are performed at MIDA to Mo ratios greater than one, the resonances for the free and for the complexed species are no broader than when each is studied separately. This indicates that the lifetime of Mo coordinated to a given MIDA ligand is probably greater than 3 seconds, so that intermolecular exchange processes will not complicate the interpretations. As a result of thew considerations the remaining work has
Table 1.
.quihbnum 1
2 3 4 5 6 7
RESULTS
Earlier NMR studies of the MoMIDA system have demonstrated that the exchange of MIDA ligands is suf-
I
log K
57.7 8 9
2.8 f 0.2 7 f 0.5
I
Cakulated Equilibrium Constants
9.6 f 0 . 2 9.65 f 0.05 8.5 f 0.1 18.2 -7.8 10.4
-8.9 50.5 52
been carried out with Mo and MIDA in equal concentrations. In Figure 1 the results of the pH titrations of a one to one Mo-MIDA solution, of a NssMoO, solution, and of an MIDA solution are shown. Although the titration curve for the MoMIDA solution is drawn-out, there are definite breaks corresponding to the consumption of 2H+/Mo and 3H+/Mo. The titration curve for NapMoO, merely indicates that no acid is consumed down to pH 6.2. Below this pH, acid consumption presumably leads to the formation of paramolybdate, M d r - 6 , or at much lower pH to the formation of octamolybdate, Mo&-‘, (16) The NMR spectra of one to one MoMIDA solutions are characterized by distinct resonaces for the free and for
Method of determination NMR Potentiometric NMR NMR Potentiometric Calculated from Equilibria 3 and 4 NMR NMR Potentiometric (in 2 M NaC1) Potenbometnc (in 3M NaClO.) NMR NMR
Reference This work This work This work This work (18)
This work
This work (‘3)
(16)
This work This work
VOL 38, NO. 10, SEPTEMBER 1966
1383
the complexed ligand, whose relative intensities depend upon the pH. For free MIDA two resonances are observed which have an intensity ratio of four to three, corresponding to the acetate and methyl protons, respectively, with the acetate resonance always appearing at lower fields. The chemical shifts of these resonances are pH dependent, shifting to lower fields as protonation occurs. The chemical shifts of the resonances for the one to one Mo-MIDA complex, which are first observed below pH 9, are pH independent down to pH 4. The NMR spectrum of this complex shows a single line for the methyl protons but a complex multiplet for the acetate protons. The results of these studies are illustrated in Figure 2. I n the upper portion of the figure the chemical shifts of the various resonances are plotted as a function of pH. For protons which give a complex multiplet pattern, the chemical shift corresponds to the center of the multiplet. The curves are discontinued at pH values where the resonances are no longer observable. In the bottom portion of Figure 2 are plotted the relative concentrations of the various ligand species as a function of pH. These concentrations were found from the relative intensities of the resonances corresponding to the species indicated. In addition the H-MIDAand MIDA+ concentrations may be determined from the chemical shift of the free MIDA resonance and from the total free MIDA concentration determined above. Studies of pure MIDA solutions verified that such a procedure is valid because the chemical shift of the MIDA resonance, relative to the shifts for H-MIDA- and MIDA-2, is a linear function of the fraction of free MIDA which exists as H-MIDA-. A material balance then gives the concentration by difference. Thus, the concentration of each species in solution is known and the equilibria and associated constants may be evaluated. Similar considerations were used in evaluating the equilibria in acid solutions. Referring to Figures 1 and 2, it can be seen that as two equivalents of acid are added to the one to one solution the free MIDA concentration decreases and the Mo-MIDA complex concentration increases so that at pH 5 the complex formation is virtually complete. However, even at pH 5 a finite quantity of the ligand (approximately 4% of the total MIDA) remains in the free form. Below pH 4 another equivalent of H + is consumed with two results: the resonances for the complex Mo-RIIDA are shifted downfield (the acetate multiplet resonances being shifted twice as much as the methyl resonance) while decreasing in intensity. Simultaneously, several new, broad resonances (line widths of 2 to 3 c.p.5. compared with 1384
ANALYTICAL CHEMISTRY
J J......... L...........
/
AC
PH
Figure 2. Upper: MIDA species
pH dependence of chemical shifts for various CHa denotes methyl protons Ac denotes acetate protons
LoweT: pH dependence of relative concentrationsof MIDA species
1 to 1.5 C.P.S. for the other resonances) begin to appear. Because these resonances are broad and are overlapped, exact spectral assignments and intensity measurements are difficult. However, in the chemical shift region normally associated with the acetate proton resonances two groups of four-line multiplets have been located. One of these groups is two or three times as intense as the other and appears at lower fields; the chemical shifts of both are pH independent. Associated with these multiplets is a single broad resonance, whose chemical shift is also pH independent. This resonance appears in the same region as the methyl proton resonances. Upon standing for a day or more, solutions with p H less than 2 are found to contain a white precipitate of molybdic oxide. The NMR spectra of the supernates show a diminishing intensity of the broad resonances with a concurrent increase in the intensities of the free MIDA resonances. The relative concentration curves below pH 4 in Figure 2 were for solutions whose NMR spectra
were obtained within hour of their preparation and do not represent an equilibrium situation. Previously, the multiplet resonances observed for the acetate protons of the MoOsMIDA-* complex were assigned as a four-line A B pattern (4). With sufficiently high resolution, however, the multiplet is found to consist of eight lines rather than four, as illustrated by the 60-Mc. spectrum in Figure 3A and the 100-Mc. spectrum in Figure 4A. While these spectra seem to result from a superposition of two individual AB patterns with coincident centers of gravity, a simple calculation indicates that this is not the case. Rather, the observed spectra can be explained as resulting from spin-coupling between protons in one acetate ann with those in the other acetate arm
protons producing one equivalent of HzO and neutral Moo3 to which MIDA-2 is coordinated. This stoichiometry was checked by NaOH titration of a solution prepared from the crystalline chelate. To within 1% the results indicated that two equivalents of base are consumed so that the one to one complex must be monomeric. Despite the apparently simple stoichiometry suggested for MoOJMIDA-2, its formation does not occur in one step as acid is added. When the first onehalf equivalent of acid is added to the basic solution the free ligand resonances shift downfield nearly linearly and no Mo-MIDA complex resonances are observed. This indicates that the only reaction is protonation of the free ligand. As more acid is added, complex formation competes with free ligand protonation. That these two competing reactions predominate, and that Mo polymerization does not occur, can be deduced from the pH titration curve for NslMoOc which shows no consumption
60 Mc
A
Figure 3. A. B.
MOO~MIDA-~ acetate resonances Observed 60-Mc. spectrum Calculated AzBz spectrum at
giving an A & (or an AA'BB') pattern. The theoretical 60- and 100-Mc. A& spectra are given in Figures 3B and 4B, respectively. These spectra were determined by the standard procedure (19) using a CDC 1604 computer. The parameters which give the best fit of theoretical to experimental spectra are indicated on the Figures. Spectra obtained a t 70" C. show slightly sharper lines, but otherwise they are identical to the room temperature spectra. In Table I1 the results of the NMR studies are summarized. The chemical shifts in the first two columns are the limiting values found from Figure 2 for the resonance of the MIDA species and for the various Mo complexes of MIDA. The coupling constants and chemical shift differences between geminal protons for the acetate multiplets are listed in the third and fourth columns, respectively.
60 Mc.
,
A
100 Mc IOcm
,
B
Figure 4. A. B.
MOO&IDA-~ acetate resonances Observed 1 00-Mc. spectrum Calculated A& spectrum at 100 Mc.
DISCUSSION
Equilibria. At the highest pH's both the p H titration data and the N M R data are consistent with the lack of any significant complex formation, the molybdenum undoubtedly existing as free and MIDA as the di-anion, MIDA+. When two equivalents of acid are added, there is an almost complete conversion of MOO,+ and ligand to the one to one Mo-MIDA complex. The most likely stoichiometry for complex formation, consistent with the consumption of two equivalents of acid, is the removal of one molybdenum oxygen atom by two
Table 11.
Observed NMR Parameters
Species methyl5 VAoststea JCierninalc MIDA-2 55 +5 ... H-M I DA 13 -36 ... Ha-MIDA+ +4 -64, ... MoOiMIDA-' +9 -21 16.5 H-MoOsMIDA,+4 -31, -16 (Mo-MIDAbdimer +5 -29, -16.1 (Mo-MIDAbdimer +5 -39' 17.6 a All chemical shifts in C.P.S. from internal TMA. b Denotes center of multiplet. Spin-coupling constant in C.P.S. between geminal protons. Chemical shift difference in C.P.S. between A and B components.
++
Avaeminald
...
...
...
16.2 ~ 1 6 -8.2 23.2
VOL 38, NO. 10, SEPTEMBER 1966
1385
of acid above pH 6.2. These processes can be represented by the equilibria:
+ MIDA-' e H-MIDAH-MIDA- + Mood-' + H + 5 MoOsMIDA-' + HZ0 H+
(1)
(2)
The equilibrium constant, K 1 , for Equation 1 is simply the reciprocal of the acid dissociation constant for the monoprotonated ligand. From studies of pure MIDA solutions at concentrations comparable to those which also contained Mo-i.e., about 0.2M-Kl was determined both potentiometrically and by NMRtechniques(9). The product K I .K z is equated to Ka, the equilibrium constant for:
+ MIDA-' + 2H+ *
MOO4-'
MoOaMIDA-'
+ Hz0
(3)
This represents the formation of the 1:1 complex, but the stability expression is complicated by the H + dependence. In order to compare the relative stability of the Mo-MIDA chelate with other metal-MIDA chelates, it would be desirable to eliminate the H + dependence. Schwarsenbach and Meier investigated the protonation of Moo4-' leading to HMo04- and H&fo04, which rapidly polymerize to paramolybdate (16). From their data they were able to determine the acid dissociation constant, K4, of the equilibrium: H&fOO4 s 2H+
+ MOO,-'
(4)
If H a o O , may justly be considered as a hydrated form of molybdic oxide-
i.e., H&foO, = MoOa.HtO-as has been proposed (IS), then the formation constant for: MoOa
+ MIDA-'
is K s = Ka.K4 = 2.4 x 10'0. Thus the MOO8 complex is among the most stable of $he one to one metal complexes formed with MIDA (17). The presence of small quantities of free MIDA between pH 6 and pH 4 is somewhat disconcerting in view of the large formation constant of the complex. Above pH 6 the free MIDA results from incomplete complex formation according to Equation 2. Below pH 6, however, an i n s i m c a n t quantity of free MIDA should be present according to the same equation. The existence of free MIDA below pH 6 seems to be due to a competing equilibrium such as the formation of paramolybdate,
+ 3HzO + 7 MoOaMIDA-' 7H-MIDA-
G
+ M G O M ~(6)
as evidenced by the experiments carried out a t pH 6.5 and 5.8 in which exce88 Moo4-' was added to one to one Mo-
1386
7Mo04-*
+ 8H+ S
+
M @ O Z ~ - ~mzo (7) K 6 is evaluated directly from the data and K7 is calculated from K7 = & . K Z 7 . The value obtained for K7 is consistent with values obtained by other workers using direct potentiometric titrations of MOO^-^ (6, 16). The fairly flat concentration profile for free MIDA between pH 6 and pH 4 in Figure 2 results from the dependence of the H-MIDAconcentration on only the '/, power of H + in Equation 6. These results suggest that unless one is willing to make an empirical correction for the slight di4'sociation of the MoO&IIDA-* chelate into H-MIDA- and polymeric Mo species, the use of MIDA as a titrant for the direct determination of Mo(V1) is not particularly attractive. Below pH 4 the NMR data again suggest that several equilibria are occumng simultaneously. The downfield shifts of the MoOslMIDA-' resonances correspond to protonation of this complex, whereas the decrease in their intensities and the appearance of new resonances is indicative of a conversion of the original chelate to some new complex. The protonation of the chelate is represented by Equation 8 :
MoO,MIDA-'
+ H+
$
H-MoOaMIDA-
S
MoOiMIDA-* (5)
H+
MIDA solutions. At the higher pH excess Moo4-* causes the intensities of the free MIDA resonances to diminish by forcing equilibrium 2 to the right. At the lower pH excess MOO^-^ has virtually no effect on the free MIDA resonances implying that the concentration of MOO4-' is not increasing, the excess presumably forming polymeric species according to Equation 7:
ANALYTICAL CHEMISTRY
(8)
Because no distinct resonances are observed for the H-MoONIDA- species the H + must be exchanging rapidly between M O O N I D A - ~ complexes. Thus, only the ligand resonances averaged over MoO&fIDA-* and HMoONIDA- can be seen. However, from the pH dependence of the chemically-shifted resonances & can be estimated in a manner exactly analogous to that employed for determining K I . The equilibrium which is in competition with 8, that is, the formation of a dilTerent complex, also must result in the consumption of one equivalent of acid per Mo as deduced from the pH titration data and the fact that all of the MoONIDA-' and H-MoOr MIDA- have been converted a t the equivalence point (pH 2). A reaction which can account for these observations and which seems likely from consideration of the catenation tendencies of molybdates is the dimerization of the one to one complex :
+
~ M O O ~ M I D A - 2H+ ~ e 0
A
(MIDA-MoO~
Mo02MIDA)-' HzO (9)
+
Further evidence for such a dimeric species is the pH independence of its chemical shifts, because additional polymerization and protonations are blocked. Because the spectra were obtained under non-equilibrium conditions, the constant, Ks, is only a crude approximation. The values of the equilibrium constante, & and Kg, given in Table I were those obtained for measurements in H20. It is interesting to note that for similar measurements in D20 the values of & and K g are increased to 2.5 X lo3 and 3 X lo7,respectively. Such changes in the constants are attributed to deuterium isotope effects, and the changes are of the same order of magnitude as have been observed for several other weak acid systems (8). Structural Considerations. Beginning with alkaline solutions the ions and complexes are discussed in the order of their formation as the solution pH decreases. Above pH 9 with no complex formation the ligand exists as the di-anion, MIDA-', and as the monoprotonated ion, H-MIDAEvidence has previously been presented that this protonation occurs almost exclusively at the nitrogen atom (9). The nature of the molybdenum ion in solution can not be determined from the data obtained here, but X-ray studies of crystalline Moo4-* salts (18) and Raman studies of both crystalline N e Mo04 and aqueous solutions Nazlllo04 (3) indicate that the Mo(VI) ion is tetrahedrally coordinated to the four oxygen atoms. In acidic solutions, however, MOO-* polymerizes to para- and octamolybdates, and X-ray studies of crystalline substances show that the Mo(V1) ion exists in an octahedral environment of oxygen atoms (11). In most Mo(V1) chelates the geometry about the Mo is probably also octahedral. A recent X-ray determination of the structure of MoOAiethylenetriamine has demonstrated that the Mo(V1) is six-coordmate in a distorted octahedron of ligands (the N-Mo-N bond angles are less than 90' and the 0-Mo-O bond angles are greater than 9 0 O ) and has suggested that the Mo-O bonds may be considered as b e i i double bonds (6). It is assumed that all the MIDA complexes in this study are such that Mo is also six-coordmate. Invoking such an assumption does not lead to any inconsistencies with regard to the experimental results. In the MoONIDA chelate with octahedral geometry three of the six Mo coordination sites must be occupied by doubly-bonded oxygen atoms, leaving
three sites available for coordination with the three donor groups of MIDA. The question of lability of the IMOnitrogen and Mo-carboxylate bonding, however, cannot be ascertained directly from the data. Because the acetate protons give rise to a complex multiplet, the two protons on a given acetate group must be magnetically nonequivalent. I n the previous work (4, it was assumed that this nonequivalence arose from the asymmetry introduced in the chelate by having all three MIDA donor groups rigidly coordinated to Mo. However, the protons on a given acetate group may also be magnetically nonequivalent even if the molybdenumcarboxylate bonds are labile, provided that the molybdenum-nitrogen bond is sufficiently long-lived. Such nonequivalence has been observed for various metal-EI)TA complexes when the metal-nitrogen bonds have lifetimes longer than a few milliseconds, and when the poliulations of the rotational conformers associated with the free carboxylate groups are different (7). In contrast to the one to one metal-EDTA chelates, which by necessity of the molecular geometry must have more labile metal carboxylate bonding than metal-nitrogen bonding, the metal complexes formed with MIDA may break a metal-nitrogen bond without necessarily disiu1)ting the metal-carboxylate bonds. In such a situation the geminal aretate protons would still retain their magnetic nonequivalencc and their ”esonances would be indistinguishable from thosc obtained for nonlabile ;netal-nitrogen and metal-carbo,xylate !)onding and for nonlabile metal-nitrogen with labile metal-carbosylate bonding. One feature of the MoO&fIDA-2 complcs which can be resolved is that the eight resonances of the acetate protons are not superimposed ABmultiplets and are therefore not assignable to two geoinetrical isomers of the coml)lcx. This conclusion is consistent with the observations by others (7, IO) that the centers-of-gravity of isomeric multiplets do not normally coincide. The complexity of the acetate proton resonances has been shown to arisc from internal coupling. That is, the protons in one -CHzC02arm, in addition to coupling with one another, are coupled to the protons in the other -CH2C02- arm giving an .4& pattern. Coupling with the methyl protons is ruled out because a very similar acetate multiplet pattern is obtained with the Mo(V1) complex of EDTA in which a CHI group has replaced the CH3 of MIIIA. In addition, the CH3resonance of the hIoO&IIID.4-* chelate shows no splitting and no significant broadening which it should if such coupling esisted. Why similar AZB2patterns have not been observed
in other metal-aminopolycarboxylate systems is difficult to understand. Undoubtedly the nature of the Monitrogen bonding plays a role here because such long range spin-spin coupling must be transmitted through the nitrogen. From examinations of molecular models one concludes that in MOOT MIDA-* the two MIDA carboxylate groups must be coordinated at cites which are cis to the nitrogen coordination site. The carboxylates, however, may be bonded at sites which are either cis or trans to one another, the statistical probabilities of these two configurations being 2/3 and 1/3, respectively. Because only one complex multiplet is observed for the acetate protons either there exists only one form of the complex, (that is, carboxylate groups coordinated either cis or trans to one another) or there is a rapid interconversion between the cis and trans configurations. Preponderance of one isomer over the other might result from strain introduced in the chelate rings in one form relative to the other form. This possibility has been suggested by Hoard from X-ray data on sexadentate metal-EDTA chelates (I@, and would favor a cis carboxylate configuration. If there were an interconversion between cis and trans isomers, in other words a rearrangement between a planar MOO, configuration, corresponding to trans carboxylate coordination, and one in which two oxygen atoms and the ion are planar with one oxygen out of the plane, corresponding to cis carboxylate coordination, it would have to be sufficiently rapid that only an averaged multiplet is observed. Such “tunneling” interconversions have been proposed for certain five and seven-coordinate metal-ligand complexes (1). Unfortunately both mechanisms lead to identical NMR results and the controlling process cannot be distinguished in this investigation. Below pH 5 two species are formed from the 3f00~M1DA-2 complex, the protonated form of the chelate, H-MoOr MIDA- and the dimeric chelate. Protonation could conceivably occur at the nitrogen atom. at the carboxylate groups, or at the doubly-bonded oxygen on Mo. The relative magnitudes of the chemical shifts of the ligand resonances when the chelate is protonated (-10 C.P.S. for the center of gravity of the acetate resonances and - 5 C.P.S. for the methyl resonances) seem to rule out nitrogen and carboxylate protonstion. For the former case, the chemical shifts of acetate and methyl resonances should be almost equal whereas in the latter case the acetate resonances should be shifted considerably more than twice as much as the methyl resonances (c.i. carboxylate protonation of MIDA in Table 11). For protonation of the
oxygen atoms coordinated to molybdenum the acetate shifts would be expected to be greater than the methyl shifts because perturbations can be effectively transmitted through both the nitrogen atom and the carboxylate linkages for the former, but only through nitrogen linkages for the latter. Moreover, the relatively small shifts for both types of protons are consistent with the large distances between the site of protonation and the protons in question, because such perturbations attenuate rapidly with distance (c.f. the 42 c.p.s. downfield shifts for acetate and methyl protons on nitrogen protonation and the 28 C.P.S. shift for carboxylate protonation in MIDA). The formation of a dimeric MoMIDA complex was suggested on the basis of the observed stoichiometry of its formation as well as the observation that the chemical shifts on the assigned resonances are independent of pH. Because the dimeric species must be formed by Mo-0-Mo linkages, one concludes that bridging might occur through any of the three oxygen atoms on each M o ion. Isased on assumption that the carboxylates are coordinated in positions cis to one another, several combinations are possible for dimerieation considering the two oxygens which are cis to the nitrogen atom and the single oxygen which is trans to the nitrogen atom. The appearance of the multiplet patterns for the acetate resonances in the dimer is not surprising since the structural and bonding features of the groups making up the dimer have probably not been altered significantly from those of the monomer. Furthermore, two or more multiplets might have been expected because of the possibility of forming geometrically distinguishable species depending upon which molybdenum oxygens provide the bridging. Just as observed for the acetate resonances, the chemical shift of the methyl resonance for the dimer is distinctly different from the monomer; its extreme breadth undoubtedly results from superpositions of two or more resonances due to methyl groups in the various geometrical isomers. LITERATURE CITED
( 1 ) Bartlett,, N., Beaton, S., Reeves, L. W., Wells, E. J., Can. J. Chem. 42, 2531 (1964). (2) Busev, A. I., “Analytical Chemistry of Molybdenum,” D. Davey and Co., New York, 1964.
(3) Busey, R. H., Keller, 0. L., J . Chem. Phys. 41,215 (1964). ( 4 ) Chan, S. I., Kula, R. J., Sawyer, D. T., J . Am. Chem. Soe. 86,377 (1964). ( 5 ) Cotton, F. A., Elder, R. C . , Inorg. Chem. 3, 397 (1964). ( 6 ) Daniele, G., Gasselta 90, 1371 (1960). (7) Day, R. J., Reilley, C. N., ANAL. &EM. 36, 1073 (1964). (8) Glasoe, P. K., Long, F. A., J. Phy8. Chem. 64, 188 (1960). VOL 38, NO. 10, SEPTEMBER 1966
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(9) Kula R. J., Sawyer, D. T., Chan, S. I., €hey, C. M., J. Am. Chem. Soc. 85, 2930 (1963). (10) Legg, J. I., Cooke, D. W., Znorg. Chem. 4, 1,576 (1965). ( 1 1 ) Lindqvist, I., Acta Cryst. 3 , 159 (1950). (12) Lindqvist, I., Nova Acta Regiae SOC. Sci. Upsaliemis 15, 22 (1950). (13) Maricic, S., Smith, J. A. S., J. Chem.. SOC.1958, p. 886. (14) Pecsok, R. L., Sawyer, D. T., J. Am. Chem. SOC.78, 5496 (1956).
(15) Sasaki, Y., Lindqvist, I., Sillen, L. G., J. Znorg. Nucl. Chem., 9, 93 (1959). (16) Schwarsenbach, G., Meier, J., J . Inorg. Nucl. Chem. 8 , 302 (1958). (17) Sillen, L. G., Martell, A. E., “Stability Constants of Metal-Ion Complexes,” Special Publication No. 17, The Chemical Society, London, 1964. (18) Weakliem, H. A., Hoard, J. L., J . Am. Chem. SOC.81, 549 (1959). (19) Wiberg, K. B., Nist, B. J., “The Interpretation of NMR Spectra,” W.
Pyridoxa I-GI yci nate Co mp lexes with Metal Ions
A. Benjamin, Inc., New York, 1962. (20) Wittick, J. J., Rechnits, G. A., ANAL.CHEM.37, 816 (1965).
RECEIVEDfor review July 6, 1965. Accepted May 4, 1966. This work was supported in part by the Research Committee of the Graduate School of the University of Wisconsin, with funds made available by the Wisconsin Alumni Research Foundation, and by a Grant (GP-4423) from the National Science Foundation.
Some Divalent
D. 1. LEUSSING and NURJAHAN HUQ Department of Chemistry, Ohio State University, Columbus, Ohio
b Titrimetric-pH investigations of pyridoxal-glycine systems have revealed the formation and coexistence of stable mono- and bis-Schiff base complexes with the metal ions Ni(ll) and Zn(ll). The pyridoxal pyridine nitrogen in these species is readily protonated to form pyridinium complexes. The formation constants of the protonated and unprotonated complexes have been evaluated and the implications toward earlier spectrophotometric studies reported in the literature are discussed. Under the acidic conditions encountered in these titrimetric studies Cu(ll)-pyridoxal-glycine solutions appear to undergo reactions other than complex formation. N o complex formation was observed with Ca(ll).
binary metal ion complexes as well as the ternary Schiff base complexes, similarly absorbing species) attempts were made to arrange conditions so as to favor the formation of only the lowest Schiff base complex having a pyridoxalamino acid metal ion ratio of 1 : l : l . Bis-Schiff base complexes however, have been isolated as solids with the ions Mn(II), Fe(II), Fe(III), Ni(II), and Zn(I1) (I). Cu(II), on the other hand, yields only the 1 :1 : 1 solid ( I ) . A continuous variation study on pyridoxalalanine-Ni(I1) solutions shows the formation of a stable 2:2:1 complex An interesting but further complicating aspect of these pyridoxal systems is the protonation of the pyridine nitrogen in the metal ion complexes. For the reaction R
S
basemetal ion complexes undergo a variety of interesting reactions (11) some of which duplicate those observed in enzymic systems involving amino acids and the cofactor pyridoxal (IS). Pyridoxal activity appears to arise via Schiff base intermediates and because of this much interest has been directed toward the investigation of simpler model systems of Schiff bases with the hope of improving our understanding of the mechanism of enzyme activity (4,9, IS). Bearing upon the present investigation previous spectrophotometric studies have been concerned with the solution stabilities of pyridoxal-valine complexes with the metal ions Mg(II), Mn(II), Ni(II), Zn(I1) (3) and Cu(I1) (8, 3 ) . Other amino acids were also cursorily investigated (3). Because of the complicated mature of the aqueous solutions of these species (dissociation of the Schiff bases, formation of simple ANALYTICAL CHEMISTRY
solids Christensen ( I ) reports somewhat higher values for Ni (pyridoxal-valine)? pKI. = 7.3, pKZD= 8.1 but pK. = 5.6 for the 1:1:l Cu complex. The basicity of this nitrogen appears to lie between that of pyridine (pK. = 5.45) and the dipolar form of pyridoxal,
I
(4).
I
CHIFF
1388
432 IO
I
H+ R I
Davis, Roddy, and Metzler (3) report pK.valuesof 6.5-6.7 [M+’ = Zn+a, Ni+*, R = (CHJpCH-1, 5.6[M’* = CU+’, R = (CH$zCH-], 6.05(M+*=C~+’, R = H ) . From the direct acidimetric titration of solutions of the dissolved
H+ pK.=8.5 (IO). Spectrophotometric studies on solutions in the absence of metal ions have demonstrated the formation of the species PV-2, PVH-, and PVHZ (P-= pyridoxal, V-=valinate) and the weaker Schiff bases PG-2 and PGH(G- = glycinate) (6, 8). The proton most likely occupies a hydrogen bonded position between the imine nitrogen and the phenolic oxygen. Titrimetric pH studies have been shown to have particular value in determining the equilibrium properties of solutions of Schiff base complexes especially where the uncomplexed Schiff base is highly dissociated (7). A large number of data points are easily obtained and subsequently processed using a high speed digital computer. Earlier studies in these laboratories have shown that the purely aliphatic Schiff bases, at least, form in solution a series of complexes in which the extensive coexistence of both mono and bis forme is possible. Since this overlap ww not considered in the previous pyridoxal imine studies, the present investigation
