24. SOLUTION CHEMISTRY of the LANTHANIDE ELEMENTS Gregory

.24.

SOLUTION CHEMISTRY OF THE LANTHANIDE ELEMENTS

Gregory R. Choppin Department of Chemistry Florida State University Tallahassee, Florida 32306,U.S.A.

ABSTRACT Complexes of the lanthanide elements provide insight into the factors which are important in ionically bonded systems. The enthalpy and entropy terms of complexation are a measure of the role of hydration effects while the free energy measures the metal-ligand interaction. The role of ligand basicity in the nature of the complex as an inner sphere or a solvent- separated outer sphere species is reviewed. An equation which has been used successfully to calculate stability constants which agree with experimental values is described and shown to be useful in establishing models for lanthanide complexing. The stability of complexes as a function of canonical structure of organic ligands and of ligand ring size is discussed. A study of charge polarization in conjugated organic ligands complexed by lanthanide cations is described. .25. SOLUTION CHEMISTRY OF THE LANTHANIDE ELEMENTS

I. INTRODUCTION

The development of separation methods using ion exchange resins has provided chemists with large quantities of very pure individual lanthanide elements. As a result, in the last 30 years research in lanthanide chemistry has been relatively intensive and has led to a rather satisfactory understanding of the chemical behavior of these elements as well as to the growing utilization of them in a variety of technological applications. A central feature of lanthanide chemistry is the strongly ionic character of the bonding between lanthanide cations and other atoms. As a result of this ionicity, they can be classified as "hard" (strongly acidic) cations. This hard acid nature is evident whether we consider the bonds between lanthanide cations with oxygen/ nitrogen, or halogen donors in complexes in aqueous solution (1) or the metal-car Don bonds of organometallic compounds (2). Due to the ionic nature of the bonding, lanthanide cations do not display the restricted stereochemistry typical of the d transition elements. The coordination number and the geometry of the coordination sphere in the d elements are essentially determined by the overlap of the metal and ligand oroitals. However, in the ionic complexes of the lanthanides, the number and arrangement of the ligand donor atoms are the result of the interplay between steric and electrostatic factors. This leads to a range of observed coordination numbers from 5 to 12 for lanthanides in solid compounds. Knowledge of the coordination number of lanthanide cations in aqueous solution is very limited. Primary sphere hydration numbers of 8 and 9 have been reported from x-ray (3) and neutron (4) scattering. It is likely that even in complexes with polydentate ligands such as EDTA, the lanthanide cation remains .26. partially hydrated with a total coordination number of 8 to 10. However, direct experimental verification of such a model is lacking and only lower limits of the coordination number can ba

established - e.g., 9 in the Ln(DP)3~ (DP » dipicolinate) complexes (5). In aqueous solution the hard nature of the lanthanide citicTiS excludes reaction with soft bases and restricts interaction to F, O, and N donors. Moreover, the interactions usually cause minimal changes in the optical or magnetic spectra involving the f orbital electrons. This insensitivity deprives the chemist of techniques which have been of major importance in the study of transition element complexes. To offset this disadvantage is the fact that the absence of significant covalent effects allows simpler interpretation of lanthanide complexation data in many systems. For example, the predominance of pseudo- contacts shifts in the NMR spectra of lanthanide compounds is the basis of their use as shift reagents. Moreover, lanthanide complexation can be used to test models of electrostatic interaction.

II. COMPLEXATION THERMODYNAMICS

An extensive body of data exists on the thermodynamic changes (AG, AH, and AS) associated with lanthanide complexation. Plots of these functions vs. lanthanide atomic number (or, vs. the reciprocal of the radii of the cations) often show a relatively linear change in the free energy but definitely non- linear curves for the enthalpy and entropy changes. This typo of pattern is shown in Figure 1 for the reaction:

The AH and AS curves are different in shape from the curve for AG but since both AH and AS are positive, in the expression &G « AH - TAS, the non-linearity largely cancels to produce an almost .27. linear AG function. This "compensation" in* the aH and AS values has a particular significance in the interpretation of the thermodynamic parameters of complexation. Complexation causes* a decrease in the hydration of the cations and anions, resulting in a positive entropy change. Since energy is required to break the cation-water and water-water interactions in the hydration sphere, such dehydration also results in a positive enthalpy change. The combination of the dehydrated cation and anion reduces the degrees of freedom (a negative contribution to the entropy) and forms an ionic bond (a negative enthalpy change). The experimental values of AH and AS reflect the sum of the opposed contributions of dehydration and cation-ligand combination. Since the observed values of AH and AS for most lanthanide complexes are positive, it seems that the dehyration effects are more significant than the combination of ions in these terns.

•800-

MS •60.0 -

•40.0 -

_ «200 -AG

000-

-20X)- o.

~°~-O -AM

-400 -

i I Pt I Pm I EM I Tb I to~l Tm I C« Hi Sm Gd L» Er Vb

Figure 1. Variation of the values of the free energies, enthalpies, and entropies of complexation of lanthanide cation» forming Lni1*^ . .28. The observed compensation of the enthalpy and entropy changes can be related to this dominance of the dehydration contribution to the net values. The stepwise reactions can be written as:

3 (a) Ln(H2O)* - (n-p)H2O h-AGD' AV AS (b) (m-q)H2O

3 (c) Ln(H2O)* AGR, AHR, ASR (3)

The net reaction is:

3 Ln(H2O)* + A(H2O)m + (n+m-p-q) HjO AG, AH, AS (4)

The net thermodynamic change is:

AG AG AHR) - ASR) (5) fl To explain the compensation effect between AH and AS, it is assumed that AG-'O. This has not been derived from first = principles but is a logical conclusion of AH_ TASD. It has been suggested that AG-=0 results from the fact that all hydrated species are always in equilibrium with the bulk solvent (6). We can go even further since the observation of positive values for AB and AS indicate that:

|AHH| > |AHR| and |AS |AS H

The consequence of this is that the observed entropy and enthalpy changes reflect primarily the sum of reactions 2(a) and 2(b) while the observed free energy change is related almost com* pletely to reaction 3. Based on these considerations, we can use the thermodynamic changes to describe both the hydrational .29. changes (from the AH and AS values) as well as the combination step (from the AG value) of the net complexation reaction. The principle involved in this model of the individual subreactions has been used also to identify innei and outer sphere complexation. In outer sphere formation, the cation retains its first sphere water and the cation and anion are thus separated by a water molecule. Since the dehydration in reaction 2(a) is minimal, the net enthalpy and entropy changes could be expected to be either negative or positive, but, in any case, small. The values for reaction 3 would be negative and small (due to the relatively large separation of the caticn and anion). Conversely, in inner sphere complexation, the cation hydration sphere is disrupted sufficiently to allow direct contact between the cation and the anion. As a result, the measured net AH and AS of the inner sphere complexation are expected to be positive and, also, larger in absolute value than for outer sphere formation. Based on this interpretation of the signs and absolute values of AH and AS of complexation, a predominantly outer sphere character has been assigned to the 1:1 lanthanide complexes with perchlorate, halides, nitrate, sulfonate, and trichloroacetate ligands, and an inner sphere nature to the iodate, fluoride, sulfate, and acetate complexes (7). A correlation was noted between the nature of the complex and the pK cl value of the acidic ligand (a measure of the basicity of the ligand): Inner Intermediate Outer

Ligand: Ac", F", ClAc", CljAc, CljAc", SCN", N03, Cl"

pKa: 4.5, 4, 2.8, 1.0, -0.5, -1.8, -2.5, -2.7

We return to a more quantitative evaluation of the inner vs. outer sphere character correlation with ligand pK in a later section of this paper. .30.

III. TEST OF AN IONIC MODEL

The correlation of the free energy of coroplexation, AG, with the metal-ligand step, Equation 4, suggests that an equation describing ionic interaction may be useful in estimating the AG values. The Born equation calculates the free energy of interaction of a cation of charge Z and an anion of charge Z. at an intemuclear distance d,- in a solvent of dielectric constant Deff

Z1Z2 AG

Munze (8) has proposed such an equation expanded to include a "cratic" term (for the change in the number of species when M + X combine to give MX) and a DeBye-Huckel term for activity coefficient correction when the ionic strength, \x , is not zero, The full equation is:

Ne - RT'vln 55.51 + RTILnf, . (u) (6) effdd1,2 where N » Avogadro's number; e - 4.80 x IO*"*"1100 esu; cation and anion ionic radii; v • -1.

2 ELnf(w) - -AZ ^ Dp with AZ2 ""CAZ MA" (ZM * Zl)^} a-4*3J B" °-33' c" °-75? D" °-15-

Munze used the bulk value of the dielectric constant with a tempera-

ture dependent term. We have chosen, instead to treat Deff as an .31.

adjustable parameter since it is very unlikely that the effective dielectric constant in the vicinity of ions is the same as the bulk solvent value. From the AG values for the formation of +2 +3 4 fluoride complexes with Ca Nd , and Th* , we set Dfiff equal +2 +3 to 79 for all M cations, 57 for all M cations, and 41 for all M +4 cations, independent of the anion, Figure 2 shows the agreement with the experimental values for the stability constants of a number of monofluoride complexes of di, tri, and tetravalent cations and the values calculated from Equation 6 (9). To use this equation with organic ligands, a value of Z^ must be used which reflects the true, not the formal, charge on the donor site. Since the pK value measures the attraction of

a proton to the organic anion, the Z2 value can be expected to be proportional to the pK . For monocarboxylie acids, the maximum value found for pK is ca. 5.1 while for dicarboxylic acids, the maximum found for the sum of pK^ + pK^ is ca. 10. From these observations, we set -1 as the charge for a carboxylate group with the maximum pK of 5 and, thus, establish the 80

70 60

E 30

M+4 •* 20

a 10

-10

0.25 0.30 035 0.40 045 050 055"

Figure 2. Comparison of experimental values of AG for formation of MF/aq) with the variation (solid lines) of AG with d12 calculated from Equation 6, .32.

calibration curve shown in Figure 3. The values of Z^ obtained from the ligand pK values with this curve are used in Equation 6. The value for r_ was 1.5S A while cationic radii values of Shannon (10) were used for coordination number 8. Figure 4 shows the correlation between calculated and experimental values using this procedure. The pK values of the nitrogen donors were included in the £pK to obtain the total Z, from Figure 3 for the polydentate Uganda. The excellent correlation between the experimental and calculated values in Figures 2 and 4 confirms the validity of the ionic model for the lanthanide complexes.

16 20 24 28

Figure 3. Relationship between Z2 values and the EpKa of organic acids. .33.

so too A6colc

Figure 4. Agreement between experimental values of AG for Sra(III) complexes with aminocarboxylate ligands and the calculated values from Equation 6.

IV. INNER VS. OUTER SPHERE COMPLEXATION

Equation 6 has been used to study the relation between outer and inner sphere complexation as a function of ligand pK using acetate and mono, di and trichloroacetate as ligands and lanthanides as cations (11). The experimental stability constant,

0T, includes the sum of the contribution from inner sphere, 8-, and outer sphere, 0Q, complexation: i.e., 0_ - 8- + 6Q. Values of flj were calculated with Equation 6 using Z- from Figure 3 for each ligand pK . The variation of 0_ and the calculated values a. I of 3j and &Q with pK& are shown in Figure 5. The relative extent of inner and outer sphere complexation was measured using La-139 NMR spectroscopy and agreed with the calculated values as shown in Figure 5 (12). This agreement further substan- tiates the utility of Equation 2 in analyzing the complexation of lanthanide cations. .34.

-05

Figure 5. Variation of log 8 (experimental), log 0O (outer sphere complexation) , log 3j_ (inner sphere complexation) with pKa of the ligand. Log 6^ was calculated from Equation 6.

V. DETERMINATION OF LIGAND ELECTRONIC STRUCTURE

In some systems, different canonical structures are possible for organic ligands. Such structures may involve different charges from the binding group. Equation 6 can be used to e-stiaiate the free energies of complexation to be expected for the anionic charges associated with the different electronic structures. Comparison of these calculated values with that from experiment provides evidence of the canonical structure for the complexed ligand. This approach has been used with croconic acid, a dihydric ligand (13). As shown in Figure 6, when croconic acid ionizes, the anion is electronically symmetric with a two electron, aromatic pi system extending over all 5 carbon and all 5 oxygen atoms. The resulting charge on each oxygen is -0.4. It may retain this structure when it bonds i->r the pi system may be more localized, extending over only four cl the carbon and oxy;óu atoira. This would lead to a charge of .55.

CROCONIC ACIO -.4

• Sm

I Z2- -0.8 -AGcdle

• 17.9 kj m"1 -AGMp»l7.7kJ m"

Figure 6. The possible canonical structures for complexation of croconic acid by Sm(III).

-0.5 on the four oxygens of the pi system. In the first case, EZ- • -0.8 which results in a value of -AG from Equation 6 for + -1 the 1:1 SmCroc complex of 12.6 kJ-m . In the more localized structure, EZ_ - -1.0 and -AG is calculated to be 17.9 kJ-m -1 Since the experimental value was found to be 17.7 kJ-ra , the canonical structure with the pi system encompassing only 4 carbons and 4 oxygens (leaving the remaining C-0 with its own carbonyl pi bonding) is apparently the one involved in the complexation. Other hydroxy acid ligands - kojate, maltolate, and tropolonate - were also analyzed (13). In the kojate and maltolate complexes, either a 6-electron aromatic pi structure or a 4- electron non-aromatic pi structure was possible. In both ligands the ggreement of calculated and experimental values of AG of complexation supported the existence of the 4-electron pi structure. Tropolonate, as expected, was found to possess a 6-electron pi structure in the 1:1 complexes with lanthanide cations.

VI, CHELATE RING SIZE AND STAIILITY

An interesting question in lanthanide complexing by organic ligands is the effect of increasing size of the chelate ring on the stability of the complex. How large can a chelate ring be .36. before simple monodentation replaces chelation? Again, Equation 6 can help to answer these questions as it allows calculation of the expected stability constants for bidentation and for monodentation. Comparison of these values with the experimental stability constants indicates which is the better model for the complexation. We have studied the complexation of lanthanides with a series of alkyldicarboxylie acids (14). The chelate structures possible for the systems studied are shown in Figure 7. The calculated and experimental values of AG are listed in Table I. We see a decreasing stability with ring size such that only monodentation seems to occur with adipate and even for the succinate complex (7 membered chelate ring), the chelation is rather weak. The oxalate (5 membered ring) is most stable while even the 6 membered chelate ring (malonate) is strained. However, tl:e CHDCA (cyclohexy ldicarboxy lie acid) complex, although a 9- meirx-eied ring, shows surprising stability to chelation. Fresumably, this ligand chelates with the cyclohexyl ring in the boat form, which gives a very favorable structure for chel^ ion (Figure 7). Comparison of the enthalpy and entropy data, when measured, could provide confirmation of these relationships of ring size and stability.

Ç —0.. S- — 0.. HjC —C —0.

c-a c-o- v-c-o'

OXALATE MALONATE SUCCINATE

c—c-o.. HJC—c— c—o. r/c—o. c > I > Í X—C—0* H2C—C—C—O' lAC—O' «2 \> H2 \ ^SQ 6LUTARATE AOlPATC CHOCA

Figure 7. Structures of dicarboxylic acid ligand» .tudied for chelation with lanthanide». .37. TABLE I

COMPARISON OF EXPERIMENTAL AND CALCULATED (WITH EQUATION 2) VALUES FOR THE REACTION; -2 +1 Ln Lnx -AG -AG -AG Ring Complex (exp.) (chel.) (monodent.) Size

EuOx 24.60 24.52 0 5 SmMal 25.86 35.48 18.85 6 SmSuc 20.09 37.96 18.61 7 SmGlu 18.50 40.44 17.86 8 SmAdi 17.70 40.44 17.74 9 SmCHDCA 25.23 42.43 18.85 9

Data from NMR spectra and thermodynamic measurements (15) and from calculations with Equation 6 have provided evidence that when the 5-membered chelate rin<;s of EDTA complexes with lanthanide cations are replaced by 6 and 7-membered rings - by

increasing the length of R, and R2 in (O2CCH2)2N-R1~N-(R2CO2")2~, the stability decreases. In these complexes, the lanthanide carboxylate bonding seems relatively unaffected as R. increases but the Ln-N bones are weakened. In fact, the sum of both Ln-M

bonds is roughly equivalent, when Rx = C3H6 or C4Hg, to about half the bonding interaction in the EDTA complexes (R. » C-HJ .

VII. INTRALIGAND CHARGE POLARIZATION

The croconate, kojate, etc., complexes have indicated that binding by lanthanide ions can influence the charge on the anionic bonding sites. The variation in pK values of sub- stituted carboxylic acids is also a reflection that the charge on the carboxylate group can be modified by inductive and resonance effects. These effects may cancel or reinforce each other and can be expected to be enhanced (relative to polarization of charge by a proton) by multiply charged cations. .38. Conjugated pi ligand systems should show such charge polarization ("intraligand charge transfer") more readily than saturated systems as resonance effects are usually stronger than inductive ones. Ligands studied recently in our laboratory (16) for such intraligand charge polarization are shown in Figure 8. The methoxybenzoates (MB) were chosen to avoid the possibility of chelation which also is not possible, due to steric constraints, for fumarate and isophthalate. For any particular lanthanide cation, the stability of the 1:1 complex, as measured by the experimental stability constant was found to be in the sequence:

p-MB > O-MB *" m-MB > B and p-MB - i-Ph - Fum

If the anionic charge in the complexes remained the same as in the acids, from the pK values we would expect the sequence of a stability to be:

iPh > p-MB > Fum 5 B > o-MB - m-MB.

,OCM, OCH

M3°°

BENZOATE O-METMQXYfjeNZOATE M-METHOXYBENZOATE P-M£TMOXYBCNZOATE (B) (O-MB) (m-MB) '.p-MBl

^C C M

H C '.O' ISOPHTHALATC fUMARATE

8. Ligande studied for intraligand charge polarization. .39.

The disagreement in the observed and predicted trends can be explained by assuming that resonance and inductive promoted charge polarization are both present in the p-methoxybenzoate and fumarate complexes. For example, in the p-methoxybenzoate complex, the resonance effect con be written as follows:

However, such a change in canonical structure cannot be written for the m-methoxybenzoate complexation so only the weaker inductive polarization is present. Although it is possible to write both canonical resonance structures for the o-methoxybenzoate complexes, resonance requires coplanarity of the carboxylate, the phenyl and the methoxy groups. This coplanarity is apparently sterically hindered by the large methoxy group when ortho to the carboxylate group, Consequently, only an inductive polarization can occur. By similar arguments, resonance cannot enhance the carboxylate group in the isophthalate complex since it is a meta isomer. We can attempt to estimate quantitatively these charge polarization effects. To do this, the experimental values of AG are used wich Equation 6 to calculate the corresponding Z2 values. From Figure 3, the Z_ values for the acid form can be obtained from the pK values. The results of these estimates cl are presented in Table II. Attention is called to the coiunon value of AZ- of 0.19 + 0.01 obtained for the o-methoxybenzoate, m-methoxybenzoate, and isophthalate complexation where only inductive polarization from the unbonded groups is proposed. The value of AZ- of 0.12 for benzoate complexation would be a measure of the polarization of charge from the pi system of the phenyl ring. If this is valid, the difference - +0.07 - can be assigned to the extra negative charge drawn from the unbonded methoxy and carboxylate groups to the bonded carboxylate site. .40. TABLE II

VALUES OF Z2 FOR Sin (III) COMPLEXATIQN

Ligand Z- (from pK ) 2_ (from ML) AZ2

Acetate -0.89 -0.B9 0 Benzoate -0.78 -0.90 0.12 o-Methoxybenzoate -0.73 -0.92 0.19 m-Methoxybenzoate -0.73 -0.93 0.20 p-Hethoxybenzoate -0.83 -1.06 0.23 Isophthalate -0.87 -1.06 0.19 Fumarate -0.79 -1.06 0.27

For fumarate and p-roethoxybenzoate, in which resonance polarization is expected to add to the inductive polarization, flZ_ has values of 0.23 and 0.27. This implies an extra charge polarization due to resonance of only 0.04 in p-raethoxybenzoate and 0.03 in fumarate. However, note in Table II that the anionic charge, Z~> on the complexed carboxylate group is -1.06 in both ligands. This charge corresponds (Figure 3) to a pK of ca. S.I, the maximum value found for monocarboxylic acids. The implication of such a value is that the bonded carboxylate group is saturated with its maximum negative charge. Therefore, even though additional electronic charge may be available for transfer to the bonding carboxylate via inductive and/or resonance effects, the carboxylate group cannot accept it. An important conclusion from this study is that, due to the limiting saturation of charge on a carboxylate site, we can define the maximum value of log 3 (or -LG) to be expected for Ln+ - carboxylate interaction. For Sin , this value is log &l - 2.8 or -áGl =16.0 kJ*m~ . In several carboxylic polyelectrolytea, the value of log & obtained for 1:1 lanthanide to carboxylate binding is 13.2 at full ionization (17). This could imply a large statistical effect since the lanthanide cation would have so many aites to choose from in a model of .41. specific site binding. Alternately, it may be evidence of the failure cf such a model of specific cation-anion site interaction in metal binding to polyelectrolytes. Manning (18) has proposed a "condensation" model which may satisfactorily interpret, such large binding constants. Further work on lanthanide binding to polyelectrolytes should be useful in establishing a valid model for metal-polyelectrolyte interaction which is important in many biological and environmental systems. The preparation of this paper was supported by a contract with the U.S.O.O.E. Division of Chemical Sciences.

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9. Choppin, G. R., Unrein, P. J., Transplutonium Elements, ed. W. Muller and R. Lindner, North-Holland, Amsterdam, 1976, p. 97. 10. Shannon, R. D., Acta Cryst3llog., A52 (1976),7S1. 11. Choppin, G. R., Lanthanide and Actinide Chemistry and Spectroscopy, ed. N. M. Edelstein, ÃÜS" Symposium Series 131, Am. Chem. Soc, Wasnington, 1980, p. 173. 12. Rinaldi, P. L., Khan, S. A., Choppin, G. R., and Levy, G. C, J. Am. Chem. Soc. 101 (1979), 1350. 13. Choppin, G. R., and Orebaugh, E., Inorg. Chem., r7 (1978), 2300. 14. Choppin, G. R., Dadgar, A., and Rizkalla, E. N., to be published. .42. 15. 3rock, J. L., Nuclear Magnetic Resonance and Thermodynamics of Aainocarboxylate Complexes, M. S. Thesis, Florida State University, 1983. 16. Choppin, G. R., and Liu, Q-L., J. Less Common Metals. 94 (1983), 406. 17. Koppold, F., and Choppin, G. R., to be published. 18. Manning, G. S.. J. Chen. Phys.. 5U (1969), 924.