A study of carboxylic and amino acid complexes of neodymium(III) by difference absorption spectroscopy

BIOINORGANIC

CHEMISTRY

3,15-25

(1973)

15

A Study of Carboxylic and Amino Acid Complexes of Neodymium(IIZ) by Difference Absorption Spectroscopy* EDWARD

R

BIRNBAUMt

and DENNIS

W. DARNALL$

Department of Chemistry, -New Mea&o State Univetsily,

Las Cruces, New Meztco 88003

ABSTRACT The interaction of neodymium(III) with acetate, alanine, hktidine, benzoate, and anthranilate has been %tudied using changes in the visible absorption spectrum of neodymium(II1) upon complexation. At pH’s below 6 only the carboxyl group of alanine coordinates to the metal ion, whereas at pH’s near 7 the P-smino group begins to coordinate. The carboxyl group alone of histidine coordinates below pH 4. Between pH 4 and 6.5 the imidazoIe group of histidine coordinates forming a bidentate complex, and above pH 6.5 the a-amino group also begins to coordinate forming a trident&e complex. From a comparison of benzoate and anthranilate cornplexes it is clear that the amino group of anthranilate does coordinate to form a bidentate complex at a pH very near the pI6 of the amino group of anthranilate.

INTRODUCTION

earth meta ions have recently been suggested as spectroscopic ion binding sites in proteins and enzymes [l-41. We have been interestedfor some time in using changes in t.he visible absorpt.ion spectrum of Nd3f as probes of metal binding sites in proteins [I, 51. Before an interpret.ationof the spectral changes of NdJf upon binding to proteins can be made, however, certain model ligands must be studied. In this article, we report the changes in the visible absorption spectrum of Nd3* upon complexationwith amino acids and carboxylic acids. It is well knovm that the visible-ultraviolet spectralabsorptionsresult.ing Rare

probes

of cakium

*This study was supported in part by Grants GEL15192 and GEL31374 from the National Science Founddi~n. i Address correspondenceto Edward R. Bimbanm at the above address. $ Career Development Awardee (K4GM-32014) from the National Institutes of He&h, United States Public Health Service. @ American EXsevierPublishing Company, Inc., 1973

16

E. R_ BIRKBATJN

AND 2). IV. DARKALL

from Laporte forbidden f-f transitions are rather insensitive to changes in the environment about the lanthanide ion. This is attributed to the fact that the radial e.xtensions of the 4f orbitals lie within the filled 5s and 5p subshells, causing the lanthanide ions to behave chemically in a manner similar to ions wit,h noble gas electronic cor&urations, such as the calcium(I1) ion cr the lanthanum(II1) ion itself. This pseudo-noble gas structure manifests itself in the limited number of isolable complexes formed by these ions in addition to the relative insensitivity of the f-f electronic transitions to environmental effects as compared to the dtransition metal ions (6). Several workers have attempted to use the absence of any differences in the absorption spectra of two complexes to conclude that the complexes are essentially the same, whereas the presence of diflerences has been used to conclude that t-hecompleues have different coordination geometries [7-9-j. The technique of di$immce absorption spectroscopy, in which the reference beam contains the same concentration of absorbing species a~ the sample beam, although in a different chemical environment, has several inherent advantages which make it especially useful for observing small changes in the absorption spectra of the lanthanide ions due to complex&ion. The difference spectrum observed by balancing the sample against the reference measures directly the difference in absorbance between the two soiutions, whiIe t.he absorption spectrum records the absorbance changes as only small perturbations of t,he basic spectrum. As a result, the difference spectrum is much more sensitive than the simple absorption spectrum measurement. Hales and Spedding have recently used this technique to calculate thermodynamic parameters for the complexation of sulfate ion nit,h Eu3+ [lo]. E ar1ier work, using difference spectra to identify the formation of species in soiution, has been reported for nitrate complexation [11-J_ We are interested in using this technique to determine under what conditions various functional groups of a given polyfunctional ligand will coordinate to a lanthanide ion in solution. This is particularly import,ant in studying the interaction of lanthanides with proteins since many donor groups are present, but only a few may actually coordinate. Similar conditions may prevail with simpler ligands. We have reported previously in a preliminary communication [127 that Herence absorption spectroscopy can be used to detect differences in the mode of coordination of polyfunctional amino acids such as alanine and i&Mine. Much of this work was done at high ligand-metal ion ratios, introducing the possibility of more than one complex existing in solution simuhaneously. To avoid this problem, in this study all ligand to metal ratios were kept at 1:l.

CARBOXYLIC

AKD

AMINO

ACID

CO%fl’LEU%

OF Nd(III)

17

EXPERIMENTAL Ma&&ls: Neodymium(II1) oxide with a purity of greater than 99.9% was provided by the Molybdenum Corporation of America and the KerrMcGee Corporation. L-Alanine (Nutrit.ional Biochemicals Corp.), E hi&dine (Sigma), sodium acetate (Baker Reagent Grade) and benzoic acid (Baker Primary Standard) were used without further purification. Anthrsnilic acid (Baker) was recrystallized once from aqueous hydrochloric acid. All other chemicals were reagent grade from Baker Chemical Company. Preparahn of Solutions: Stock solutions of XdCG (0.10 ilf) xvereprepared by digesting a slight stoichiometric excess of neodymium(III) oxide xvith ECI, filtering, and diluting to volume with distilled, deionized water. The pH of solutions prepared in this manner wss 6.20. The pH’s of the l&and solutions (0.10 M) were adjusted to VLU~OUS values with concentrated HCl or XaOH so that miuinxal changes in the ligand concentration occurred. Aliquots of equal volumes of the ligands (at various pII’s) and the neodymium solutions (sll at pH 6.2) n-ere than mixed, the fmal pH of these mixtures wss obtained and the spectra were recorded. Spectral ilfeaswwn~: Spectra were obtained on a Gary 14 spectrophotometer employing either l-cm or 5-cm cells and a O-0.1, 0.1-0.2 absorbance slidewire. The scanning rate was 10 &/sec for the most intense spectra at high pH and 20 A/set for the less inixnse spect.ra at low pH. The areas of the difference spectra were measured using an Allbrit compensating planimeter. The difference absorption spectrum of the neodymium(II1) ion was recorded over the range of 550-610 nm. This region includes a major band at 575 nm. This band is the “hypersensitive” 41gD+ 4Gsfl, ‘Grn transition which appears to be very sensitive to the environment about the neodymium ion [8]. The difference spectra n-ere obtained using an aqueous solution of the neodymium ion (pH 6.2) as a reference solution. The concentration of the neodymium ion was the same in both the sample cell containing the ligsnd and the reference cell so that the difference spectrum observed could only result from differences in the environment about the metal ion in the two wlls. To insure that dilution factors were not contributing to the difference spectra observed, solutions of the same neodymium ion coneentrstions were prepared separately for the sampIe and reference cells and no difference spectrum wss seen. Since the sample beam solutions were at pH’s other than 6.2, while the pH of the reference beam solutions wss held con&ant at 6.2 in these experiments, a check \yss made to see if different acidities alone could

E. R. BIRNBAUM

18

AXD

D. W. DARNALL

cause a difference spectrum. No difference spectrum was observed with the same concentration of neodymium ion in the sample and reference at a pH of 1.0 and 6.2, respectively (the pH was adjusted with HCl). The pB of the reference solution could not be raised signifkmtiy above pH 6.2 since hydroxide precipitation ensued, and thus no direct check of the effect of base on the difference spectrum could be made. As mentioned before, the chloride ion produces no difference spectrum [12]. The final concentrations of neodymium(II1) ion (and &and) used in the difference spectraI study were 0.05 31, with the exception of anthranilic acid and benzoic acid. Due to the low solubiity of these ligands the final concentration of neodymium ion (and ligands) was reduced to 0.01 M. To compensate for this reduction in concentrations, 5-cm cells instead of l-cm celis were used to record the spectra. RESULTS

AND DISCUSSION

The difference spectrum from 550-610 nm was measured for solutions of neodymium(III) chloride containing equal concentrations of each of the following ligands: acet.ate, benzoate, alanine, hi&dine, and anthraniIate over the pH range from 1 to 7. Figures 1 and 2 show representative clifkence spectra obtained for each of these ligands. The difference spectra

I

I

I

I

I

‘- Acetate pH=3.2

I

I

I

I

I

t-listidine

Alanine

pH=tiO

pH=65

001 I

I

I-

I

OOlA

.b

WAVELENGTHhd

Fig. 1. Difference absorption spectra of -03 M Xd+J solutions containing (a) -05 M acetate, pH 3.2; (b) -05 X bistidine, pH 6.0; and (c) .05 M alauine, pH 6.5. Spectra were takeu using l-cm pathlength celk

CARBOXYLIC AND AMINO ACID COMF’LEXES

T

I

I

1

OF xd(III)

19

I

Anthr3nilate

WAVELENGTH

Fig. 2. Difference absorption spectra of -01 M Nd*z soIutions containing (a) 0.01 &f benzoate, pH 3.62; and (b) -01 M anthraniIate, pH 3.67. Spectra were taken using 5-cm pathlength cells. (Figure 1) exhibit t.he same features, although the i-ten&es vary markedly. The spectra of the aromatic ligands ant.banilafe and benzoate (Figure 2) are very similar to the albhatic ligands but do have features different from those ligands in Figure 1. Both types of spectra indicate that the hypersensitivetransition is intenSed in the presence of a carboxylate ligand in comparison to water. The major Werences in the two types of differencespectra occur in the 568475 nm region. In this wavelength region, the intensity of components of the hypersensitivetransition is increased for the aromatic liga,ndsbut decreasedfor the aliphatie ligands (relative to the aquo ion). The intensit.iesof the differencespectra obtained nit.h all ligands varied markedly with PH. Figure 3 shows representative difference spectra obtained with Kd(II1) and hi&dine as a function of pH. For hi&dine and all Iigandsexamined,the intensitiesof the Merence spectra increased as the pH was increased,but the shapes of t.hespect.raremainedthe same. The area beneath the positive portion of t.he differencespectrum has been measured for all ligands as a function of pH. Figure 4 shows these data for acetate, alanine, and hi&line, whereas Fig. 5 shows the plots for anthranilic acid and benzoic acid. Earlier work [12] at high ligandmetal ratios suggestedthat only the carboxyl group of alanineand hi&dine

of acetate, alanine, and_ histidine

20

E. R. BIRNBAUM

AKl

I

D. W. DARNALL

I

I

‘7, % %

I

6,

WAVELENGTH Fig_ 3. Difierence absorption spectra of _05 _MXd+~solutions containing -05 M histidine at sever& pH’s_ Spectrs acre taken using l-cm cells_

i

bind neodymium ion near pH 4, and that additional coordinationof the imidazoIegroup of hi&dine occurs in the pH region of 6-7. The present differencespectral work at a 1:l rat.io of l&and to metal ion is consistent with the earlier study. The tikation curve of alanine (Fig. 4) shows a gradual increasein area as the pH increasesfrom 1 to 3, and then it levels off until pH 6 after which the area begins to increase again. Above pH 7.2 precipitation of hydroxy species occurs. A similar plot is observed for hi&dine (Fig. 4). A gradual rise in t,he measuredarea over the pH range 0.%!?.5 followed by a plateau from 2.5-4 after which a gradual increase in area occurs from pH 4-6.5 and finally a very rapid increasewhich continuesabove pH 7. The increasein area ok the alanine spectra over the.pH range l-3 can only be due to coordinationof the carboxyl group of alanine to the neodymium ion. Further increase in pH beyond that -A% not increase the binding of alanine to neodymium once the carboxyl group is completely

CARI3OXYIJC

AND

AMINO

ACID

COMPLEXES

OF Nd(III)

21

deprotonate& (the pK,, of the carboxy group of alaniue is 2.35). The increase in area above pH 6 can be due either to formation of hydroxy species OF coordination of t,he deprotonated SIX&O group (pI6 = 9.69). A similar explauation can be made for the his&line plot. The deprotonation of the carboxyl group of bistidine (prC, = 1.82) allows the formation of a neodymiumhistidiue complex over the pH range 0.25-2.5. Increasing the pH to 4 causes no additional complex to form. The increase in area occurs at a lower pH for histidine than for alanine because the pK of the l&&line carboxyl is lower than that of alanine. The smaller area measured for histidine, compared to alanine below pH 5, suggests that the protonatcd imidazole group is either interfering with the coordination of the carboxyl group causing a reduced stability constant (i.e., fewer complex species in solution) or else t.heincrease in the intensity of the hypersensitive trausition is smaller for histidine, The former explanation seems somen-hat better since at pH 6 and above the areas of the histidine spectra are greater than those of alanine. This change occurs near t.he pK, of the imidazole nitrogen of histidine (pI6 i S-O), hence the deprotonation of that group and

0

iiistidine

Fig- 4_ -4 plot of the difference spectral intensities (areas) of .05 Jf Xd+a solutions con-g (A) -05 iIf acetate, (x> _05 M alanineand (0) .Wi M histidine as a function of PH. Spectraweretakenusing l-cm cells.

22

E_ R BIRKBAUM

I

AND D. IV. DARNALL

I

I

I

I

I

8 p_+._ 0

o Anthranilate

7. _

/ 0

A 8enzoate

6Q-

I e

50-

-60

/ I

30-

-40

AMP

11 A

-30

d/

20-

.=.a /”

lo-

s/ /

ma Nd _-_____-______

-20

A _-_-__-

__-_----___

-,(-J

J

I 1

.A

2

I

3

I

I

4

5

I

6

’

PH

Fig. 5- -4 plot of the differencespectral inteiities of .05 M Sd+z solutions containing (4) -01 Ai benzoate and (01 -01 31 anthranilate 8s a function of PH. Spectra were takenusing5-cmceUs.

subsequent formation of a bidentate complex seem to be t.he cause of the Gwwse in area measured. The area of the difference spectrum of histidine over the pH range 5-6.5 increases, but beyond that point a much more dramatic increase in area occurs. This again suggests that either a hydroxy species is forming, or else the ol-amino group of histidine is coordinating. The difference spectral work done at high Iigand-metal ion moIe ratios [12 7 and the present study done at a 1: 1 ratio are both in accord with an nmr study of the Nd(III)-his&dine system reported earher [13]. In the rum work the variation in the isotropic shifts of the histidine protons in the neod_ymium complex were studied as a function of pH. It n-as clear that the isotropic shifts of the C, and Cd protons of t.he bistidine were sign&ant at Iom pH (<4), while the shifts of the imidaaole protons were very small_ This reflected coordination of t,he metal ion by the carboxyl group of hi&dine aIone. As the pH WLSraised above 4, t.he G and C4 protons of the imidazole ring underwent large isotropic shifts, while the

CARBOXYLIC

AND AMINO ACID COMPLEXES

OF Nd(III)

23

C, and Cp protons suffered only a smaU increase in the observed shift. This corresponded to coordination of the imidazole group to form a bidentate Nd-hi&line complex. As the pH was raised above 6, the observed isotropic shift of the C,, CD, and CL protons increases still more rapidly, indicative of additional coordination by the a-amino group or the formation of a hydroxy complex The similarity of the data presented in the earlier nmr work (Fig. 2, Reference 13) with the difference spectral area-pH plots shown here (Fig. 4) suggests that both techniques are suitable for detecting changes in the coordination sphere of a Iant.hanide ion. There are several pieces of evidence which suggest that coordination of the a-amino group is cauzing the rapid increase in area observed for both the alanine and histidine spectra near pH 7. First is the fact that the areas of the acetate spectra level off near pH 5.5 and become constant up to pH 7.1, whereas the areas of the alanine spectra begin to increase rapidly above pH 6.5. Second, the rapid increase in area for the hi&dine spectra occurs at a sigmfkantly lower pH than for the alanine spectra. This is in agreement both with the lower pK, of the a-amino group of bistidine and also with the fact that t.he hi&line complex containing a deprotonated amino group would be tridentate, whereas the analagous alanine complex would be bidentate. The formation of a tridentate complex should lower the apparent pK, of the a-amino group more than the formation of a bidentate complex. In order to help distinguish between amino group coordination and hydroxy group coordination, _the interaction of Nd3+ with benzoic acid and antbranilic acid was examined. The advantage of these aromatic systems is that the PIG’S are lower (benzoate pK, = 4.20; anthranilate pGs = 2.05, 4.95), and so the complete ionization of the amino group of anthranilate acid can be seen below pH 7 before hydroxide formation become important. The benzoic acid titration curve (Fig. 5) is essentially the same as that of acetate (Fig. 4), although the inflection point occurs at a lower pH for benzoate, as would be expected. The anthranilate acid curve has tso inflection points at pH 1.5 and at pH 4 corresponding to the ionization of the earboxyl group and amino group, respectively. The authranilate acid curve levels off above pH 5.5 and shows no increase in area up to pH 6.5. The area of the anthranilate spectrum above pH 5.5 is much greater than that of the benzoate or acetate spectra which might be expected if both the carboxyl and ammo groups of anthranilate are coordinating to the metal ion. This evidence, coupled with the fact that the acetate spectrum does not increase in intensity from pE 5.5-7, is strong evidence that the amino group of ala&e and histidine is beginning to coordinate near pH 7, rather than hydroxide ion causing the increase iu area. The possibility does remain however, that deprotonation of the

24

E. R. BIRXBAEM

AXD

D_ W_ DARNALL

a-amino group changes the character of the carboxyl group as a ligand enough to account for the observed spectral change. A comparison of intensities among the ligands is difficult. The absolute absorbance of a given solution is dependent upon both t.he extinction coefficient of the species in solution and the concentration of the species in solution_ The relationship of molecular st~ructure to intensity of rare earth ion transitions is in gene& not understood sufhciently to be able to predict what the intensity of a complex should be. In this work, the intensity of the difference spectra, and hence the intensity of the sohrtion containing the comples, appears to be a function primarily of the concent.ration of the species in solution rather than large changes in extinction coefficients_ Thus, for the acetate tit.ration curve, the gradual increase in area as a function of pH is due to t.he gradual ionization of the acetic acid molecules resulting in increasing amounts of 1 :l neodymium-acet.ate complex in solution. At these concent.rations, the neodymium ions are not completely complexed since the stability constant for Nd(CH&OO)*+ is only 83-93 (14). The Iolver overall areas recorded for alanine and histidine (beiow pH 5) probabIy reflected the smaller stability constants for the I rl complexes of those Iigands with t.he neodymium ion_ This is consistent with several stability constant,s reported for glycine, alanine, and histidine [14]. SimilarIy, the very small area recorded for anthranilate at low pH abo reffects a weak comples. The value of the log E, reported for ‘Xd+3 (3.23) is almost certainly for t,he bidentate comples with both the carbosyl and amino groups coordinated [lZJ. There are, however, differences in the extinction coefficients of the various complexes so that one cannot simply use the relative areas measured between two ligands to calculate equilibrium constants. This can be seen by using reported stability constants [14] for aianine (K = 6.5) and acetate (K = 83) to calculate the concentration of complex in solution in these experimental conditions. If the spectra of the complexes obey Beer’s Law then equation 1 should hoId;

RAE=

Ai/md-Ac]

= Az/Bd-ala],

(I)

n-here A, and Aa are the areas, i.e., the summation of the difference absorbances, of t.he I : 1 alanine and acetate complexes, respectively, pd-ala] and [h’d-Ac] are the concentrations of the 1: 1 alanine and acetate cornplexes, respectively, and Z Ae is the summation of the difference ext.inction coe6icient.s over the entire hypersensitive transition for both complexes. This calculation made at the plateau region of the titration curve (pH 5 for alanine, pH 6.5 for acetate) yields a value for Z AE of 900 for alanine and 1550 for acetate. Thus, although the areas are good indicators of the

CARBOXYLIC

AND

AMINO

relative magnitude of the stabiity ship is not quantitative.

ACID

COMPLEXES

OF Nd(II1)

constant for complex&on,

25

the rclation-

CONCLUSION The primary result of this n-ark has been the fact that difference absorption spectroscopy of neodymium complexes can be us&d to detect coordination of functional groups of polyfunctional ligands as a function of solution pH_ Specifically, the a-amino group of alanine is just beginning to coordinate with the neodymium ion at pH 7.2, forming a bidentate complex. Similarly, the imidazole group of histidine is coordinating to the neodymium ion over the pH range 5-6.5, while above pH 6.5 significant amounts of tridentate complex containing the o-amino group are formed. Hydroxide ion has no visible effect on the difference spectra observed for these complexes up to a pH of 7.2. The shapes of the difference spectra are very similar for the neodymium complexes which have single carboxyl ligands, carboxyl and amino ligands, carboxyl and imidazole ligands, or carboxyl, amino, and imidazole ligands. Thus it would appear that determination of the ammo acid side chains chains involved in neodymium ion binding in proteins will be di&ult from an analysis of the shape of the difference spectrum alone. However, the pH-dependence of the difference spectrum observed upon a neodymiumprotein interaction may be particularly sensitive in determining the pGs of t,he neodymium ion ligands. We wish to thank the MoL’ybdenunt Corporation of America and the Kerr McGee Chemical Corporation for providing samples of rare earth oxid&. REFERENCES 1. E. R. Bimbaum, .J. E. Gomez, and D. W_ Dam& J_ Amer. C/km. Sot. 92, 5287 (1970). 2. D. W_ Damall and E. R. Biibaum, J. Biol. C&m. 245,6484 (1970). 3. G. E. Smolka, E. R. Bimbaum, and D. \V_ Damall, Biochemistry10,4556 (1971)_ 4_ R. J_ P_ \Villiams, Quart. Rev. 24, 331 (1970). 5. D. W. Dar&I, E. R_ Btiba~~~, J. E. Gomez, and G. E. Smolka, Proc. Ar&!h fire Ear@ Res. Conf. 1,278 (1971). 6. T. MoeHer, D. F. M&in, L. C. Thompson, R. Fen-us, G. R. Feistel, aud W:. L. Randall, Chenr.Rev. 65,l (1965). 7. L. I. Katzin and al. L. Bamett, J. Phys. Chem. 68,3779 (1964). 8. D. G. I&x-raker,Inorg. C&m_ 6, 1863 (1967).

26

33.Et.BIRNBAUM

AND

D. IV_ DARNILL

9. I)_ G.-er,M 7,473 (1963)_ 10. C. F. Hale and F. H. Spedding, J. Phys. Chem. 76, 1337 (1972). 11. N. A. Cod and R. W. Kiss, J_ Phys. Chem. 70,213 (1966). 12. E. R. Bimbaum, C. Yoshida, J. E. Goxnez and D. W. Damall, Proc Ninth I&e Earth Res. Cmf_ 1,2M (1911). 13. A. D. Sherry, E. R. Bimbaum, and D. W_ Damall, J. Bid. Chem. 247, MS9 (1972). 14. A. D. Sherry, C. A. Yoshida, E. R. Bimbaum, and D. W. Damail, J- Amer. Chenz. Sot- 95,3011(1973)_ 15. M. Cefola, A. S. Tompa, A. V. C&no, and P. S. Gentile, Inorg_ Chem. 1,290 (1962). Ik-ceixedZS?February 1973