Spectroscopic study of lanthanide(III) complexes with aliphatic dicarboxylic acids

www.elsevier.nl/locate/ica Inorganica Chimica Acta 310 (2000) 248– 256

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Spectroscopic study of lanthanide(III) complexes with aliphatic dicarboxylic acids Z.-M. Wang *, L.J. van de Burgt, G.R. Choppin Department of Chemistry, Florida State Uni6ersity, Tallahassee, FL 32306 -4390, USA Received 17 March 2000; accepted 12 July 2000

Abstract The complexation of trivalent lanthanides with aliphatic dicarboxylic acids (malonic, succinic, glutaric and adipic) were studied at 25°C and 0.1 M (NaClO4) ionic strength by luminescence and absorption spectroscopy and luminescence lifetime measurements. The luminescence spectra and decay constants indicate that ML and ML2 complexes were formed. The stability constants of Eu(III) complexes with the dicarboxylic acids were calculated from the changes of the 5D0 ’ 7F0 excitation spectra of Eu(III). For the four dicarboxylic acids studied, both the stability constant and the number of water molecules released from the inner sphere of Eu(III) upon complexation decrease from malonate to adipate for both the ML and ML2 complexes. The results are interpreted as reflecting an increasing tendency from chelation to monodentation as the carbon chain length increases between carboxylate groups. The trend in the oscillator strength in the hypersensitive transition of the Nd(III)and Ho(III) complexes is the same as that in the ligand basicity. © 2000 Elsevier Science B.V. All rights reserved. Keywords: Lanthanide(III) complexes; Aliphatic dicarboxylic acids; Stability constants; Luminescence spectra; Complexation; Absorption spectroscopy

1. Introduction The interaction between metals and humic acids has attracted great attention in recent years because of the important role of humics in the speciation, transportation and migration of heavy radioactive nuclides in the environment [1–4]. Aliphatic and aromatic carboxylate are major metal binding sites in humic acids [5]. Therefore, metal ion complexation with aromatic and aliphatic carboxylic acids can serve as a basis for the interpretation of the interaction of metal cations with humic acids. In a previous paper, we reported the results of a study of lanthanide complexation with aromatic polycarboxylic acids [6]. In this work, lanthanide complexes with aliphatic dicarboxylic (malonic,

* Corresponding author. Present address: Mail Stop K8-96, Pacific Northwest National Laboratory, Richland, WA 99352, USA. E-mail address: [email protected] (Z.-M. Wang).

succinic, glutaric and adipic) acids ligands were studied by Eu(III) luminescence spectroscopy and absorption spectroscopy of hypersensitive bands of Nd(III) and Ho(III). The complexation of lanthanides with aliphatic dicarboxylic acids has been reported earlier from studies using potentiometry, calorimetry and luminescence method [7–10]. The high sensitivity and wide dynamic range of the laser-induced time-resolved luminescence spectroscopy makes it possible to systematically study the complexation of lanthanides with a series of ligands among which the solubilities and stabilities of both the ligand and the complexes may vary in a wide range. Luminescence lifetimes of Eu(III) can provide accurate measures of the hydration status of the metal ion as well as the complexes, giving useful information on the bonding and structure. The variations in the oscillator strength of hypersensitive bands of Ln(III) ions can also provide insight into the characteristics of the Ln(III) complexes with ligands.

0020-1693/00/$ - see front matter © 2000 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 0 - 1 6 9 3 ( 0 0 ) 0 0 2 5 9 - 0

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2. Experimental Oxides of neodymium (99.9%), holmium (99.9%), europium (99.99%), and terbium (99.9%) from Aldrich Chemical Co., were dissolved in perchloric acid to prepare aqueous solutions of the lanthanide perchlorates. The concentration of the lanthanide ions were determined by complexometric titration with ethylenediaminetetraacetic acid, EDTA, using 20% hexaethyltetramine as buffer and xylenol orange as indicator [11]. Acetic acid was purchased from Fisher Chemical Co. and the dicarboxylic acids (malonic acid, succinic acid, glutaric acid and adipic acid) from Aldrich Chemical Co. All of the acids were reagent grade and used without further purification. Distilled, deionized water (E-Pure, Barnstead) was used throughout experiments. Concentrated perchloric acid (60%), hydrochloric acid (36.5%), and nitric acid (71%) (reagent grade from Fisher) solutions were diluted to obtain stock solutions used in the experiments. Stock solutions of NaOH were prepared from reagent grade NaOH (Fisher). For most of the aliphatic carboxylic acids, complete dissolution only occurs under basic conditions. Hence, appropriate volumes of 1.0 M sodium hydroxide solutions were added to achieve dissolution. The final pH of the stock solutions was adjusted to pH 5 – 6 with solutions of sodium hydroxide or perchloric acid. Working solutions of the organic acids were prepared by dilution of the stock solution using distilled and deionized water. An appropriate amount of sodium perchlorate was added to maintain the ionic strength of the solutions at 0.1 M. EDTA solution was made by dissolving reagent grade disodium salt (Aldrich) in distilled water. Hexamethylenetetramine (20%) was prepared by dissolution of the reagent in distilled water. Xylenol orange was prepared as a 1:100 solid dispersion in NaCl. The pH measurements were performed with a Corning 130 pH meter equipped with a Corning Semi-Micro combination electrode filled with saturated sodium chloride solution. Standard buffer solutions at pH 7.0 and pH 4.0 (Fisher) were used for the calibration of the pH meter. Solutions of sodium hydroxide and perchloric acid of various concentrations were used for pH adjustment throughout the experiments. UV –Vis absorption spectra were recorded with a Cary-14 (OLIS modified [6]) scanning spectrophotometer. The laser-induced time-resolved spectrometry system for recording Eu(III) 5D0 ’ 7F0 selective excitation spectra and luminescence decay curves has been described elsewhere [12]. Briefly, the Eu(III) sample solutions were excited with the output of a dye laser pumped by a Nd –YAG laser with the dye laser wavelength scanned in the range of the 5D0 ’ 7F0 band of Eu(III). The light emitted by the sample is collected at 90°, detected by a Hamamatsu R928 photo-multiplier

249

tube. The signal was amplified in a Lecroy 6103 amplifier and then fed into a Lecroy TR8828C transient recorder connected with a GPIB interface to a computer where the signal was gated, summed, averaged and stored. The laser dye was a 50:50 (v/v) mixture of Rhodamine 590 and Rhodamine 610 (Exciton Chemical). The luminescence decay curve was obtained by monitoring the luminescence intensity at 616 nm while exciting at the maximum of the 7F0 ’ 5D0 band. Two series of solutions of Eu(III) complexes with carboxylic acids were prepared by titrating either an europium solution with the carboxylic acid solution at a fixed pH or titrating an Eu(III)-carboxylic acid solution mixture with sodium hydroxide or perchloric acid solution to different pH values between pH 2 and 5.5. The concentrations of Eu(III) and the ligand in the measured samples were calculated from the accumulative added concentrations of Eu(III) and of ligand, with correction for the dilution effect in the pH adjustment. The stability constants of Eu(III) complexes with carboxylic acids were calculated from the 5D0 ’ 7F0 selective excitation spectra of Eu(III) with the SQUAD program [13]. Absorption spectra of Nd(III) and Ho(III) in the hypersensitive transition ranges were measured for solutions of varying carboxylic acid to metal ion concentration ratios. The hypersensitive transitions 5G6, 5F1 ’ 5I8 for Nd(III) are located between 545.0 and 605.0 nm and 4 G5/2, 2G7/2 ’ 4I9/2 for Ho(III) between 435.0 and 463.7 nm. In a typical experiment, 500 ml of a metal ion solution was prepared by dilution of a calculated volume of the stock solution with 0.1 M NaClO4. Thirty milliliters of the solution was withdrawn to measure the absorption spectra of the metal ion with 0.1 M NaClO4 as reference and the rest of the solution was titrated by aliquots of solution of the carboxylic acid under study. After the addition of each aliquot of ligand, the solution pH was adjusted to a value between 5.0 and 5.5 with dilute NaOH or HClO4 solution. Thirty milliliters of the resulting solution was withdrawn to measure the absorption spectra of the complexes with the corresponding ligand solution as reference. Paired quartz cells of either 5.0 or 10.0 cm pathlength were used throughout the experiment. At each wavelength, the absorbance was taken as the average of five samplings and each spectrum was the average of three scans. The area under the absorption peaks was integrated by the trapezoidal method [14] to calculate the oscillator strength of the transitions. For Nd(III) and Ho(III) complexes of carboxylic acids, the mole fractions of the free metal ion, [M(III)], and the MLn (n=1–3) complexes were calculated from the reported protonation constants of the carboxylic acids (Table 1) and the stability constants of the com-

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Table 1 Protonation constants of the dicarboxylic acids (25°C, I= 0.1 M) Ligand Malonic acid Succinic acid Glutaric acid Adipic acid

pK1 2.65 4.00 4.13 4.26

pK2 5.28 5.24 5.03 5.03

Table 2 Solution conditions and luminescence decay constants of Eu(III)–succinic acid (SUC) system at 25°C and I= 0.1 M (NaClO4)

Ref. [15] [14] [15] [15]

plexes using the SPECIES program written in this laboratory [6]. The stability constants Nd(III) and Ho(III) complexes used in the calculations were estimated from the values obtained for Eu(III) complexes with these carboxylic acids from the luminescence excitation spectra based on the trends in variation of the stability constants of Nd(III) and Ho(III) with these carboxylic acids obtained from other methods among the lanthanide series.

No.

[Eu(III)] (M)

[SUC] (M)

pH

k (ms−1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26

9.8991E−5 9.7625E−5 9.7395E−5 9.7278E−5 9.7219E−5 9.7179E−5 9.7138E−5 9.7076E−5 9.7035E−5 9.6822E−5 9.6608E−5 9.6412E−5 9.6213E−5 9.6035E−5 9.5876E−5 9.4557E−5 9.3362E−5 9.1484E−5 8.9234E−5 8.6502E−5 8.4364E−5 8.2267E−5 7.9721E−5 7.7256E−5 7.4433E−5 4.0000E−4

0.0000E−0 9.9096E−4 9.8862E−4 9.8744E−4 9.8684E−4 9.8644E−4 9.8602E−4 9.8540E−4 9.8497E−4 9.8282E−4 9.8064E−4 9.7865E−4 9.7664E−4 9.7482E−4 9.7321E−4 2.1694E−3 3.3583E−3 5.1117E−3 7.3982E−3 1.0186E−2 1.2358E−2 1.4489E−2 1.7087E−2 1.9603E−2 2.2518E−2 0.0000E−0

3.660 2.505 2.821 3.085 3.322 3.554 3.768 3.976 4.201 4.388 4.590 4.777 4.973 5.157 5.335 5.209 5.221 5.201 5.222 5.239 5.230 5.246 5.253 5.253 5.283 2.500

a a a a a

9.488 9.700 9.327 9.120 8.655 7.889 8.122 7.961 7.772 7.648 7.108 6.920 6.641 6.435 6.315 6.197 6.130 6.066 5.995 5.923 9.200

a The signal is too weak to calculate the luminescence decay constants.

3. Results and discussion

Fig. 1. 5D0 ’ 7F0 selective excitation spectra of Eu(III) complexes with succinic acid(SUC). The solution conditions are: 1, [Eu3 + ]= 7.4433× 10 – 5 M, [SUC]= 2.2518×10 – 2 M, pH 5.283; 2, [Eu3 + ]= 8.2267× 10 – 5 M, [SUC]=1.4489×10 – 2 M, pH 5.246; 3, [Eu3 + ] =8.6502 × 10 – 5 M, [SUC]= 1.0186×10 – 2 M, pH 5.239; 4, [Eu3 + ] =8.9234 × 10 – 5 M, [SUC]= 7.3982×10 – 3 M, pH 5.222; 5, [Eu3 + ] =9.3362 × 10 – 5 M, [SUC]= 3.3583×10 – 3 M, pH 5.221; 6, [Eu3 + ] =9.4557 × 10 – 5 M, [SUC]= 2.1694×10 – 3 M, pH 5.209; 7, [Eu3 + ] =9.5876 × 10 – 5 M, [SUC]= 9.7321×10 – 4 M, pH 5.335; 8, [Eu3 + ] =9.6213 × 10 – 5 M, [SUC]= 9.7664×10 – 4 M, pH 4.973; 9, [Eu3 + ] =9.6608 × 10 – 5 M, [SUC]= 9.8064×10 – 4 M, pH= 4.590; 10, [Eu3 + ] =9.7035 × 10 – 5 M, [SUC]= 9.8497× 10 – 4 M, pH 4.021; 11, [Eu3 + ] =9.7179 × 10 – 5 M, [SUC]= 9.8644× 10 – 4 M, pH 3.554; 12, [Eu3 + ] =9.7625 × 10 – 5 M, [SUC]= 9.9096× 10 – 4 M, pH 2.505. For simplicity and clearness, only a part of the spectra are shown.

3.1. Stability constants and luminescence decay constants Fig. 1 shows spectra for the succinic acid –Eu(III) system as examples of the luminescence data. Addition of a ligand into a Eu(III) solution causes the appearance of a spectral peak on the 5D0 ’ 7F0 excitation spectra at a longer wavelength, 579.0 nm, than that of aqueous Eu(III), 578.84 nm, and the intensity of the peak increases with the amount of ligand added. Since the 5D0 ’ 7F0 transition is non-degenerate, each distinctive spectral peak corresponds to an unique complex. The appearance of the 590.0 nm band reflects the formation of inner sphere complex. As the ligand concentration further increases, a second peak appears at 579.2 nm, indicating the formation of higher complexes. Since luminescence quenching of Eu(III) in aqueous solution is caused primarily by the coupling of the excited states of Eu(III) with the vibrational levels of O–H of its hydration waters [15], the decrease of the luminescence decay constant indicates a decrease in the number of water molecules in the inner coordination sphere of Eu(III) due to complexation by the ligands (Table 2).

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Table 3 Stability constants of Eu(III) complexes with dicarboxylic acids a CA

Method

I (M)

t (°C)

log i1

log i2

MAL

lif gl gl gl gl gl

0.1 0.1 1.0 0.5 0.1 0.1

25 25 25 25 25 25

4.18 9 0.01 4.28 9 0.01 3.72 9 0.02 3.60 4.62 9 0.03 4.31

6.62 9 0.01

SUC

lif gl

0.1 0.1

25 25

2.99 90.01 3.46 9 0.03

4.90 90.02

p.w. [16]

GLU

lif gl

0.1 0.1

25 25

2.66 9 0.01 3.22 9 0.02

4.53 90.01

p.w. [16] b

ADI

lif gl

0.1 0.1

25 25

2.59 9 0.01 3.14 9 0.03

4.84 9 0.02

p.w. [16] b

a b

log i3

6.24 90.02 6.99

Ref. p.w. [7] [8] [9] [10] [16]

CA, carboxylic acid; lif, laser-induced fluorimetry; gl, glass electrode potentiometry; p.w., present work. These value are interpolated from stability constants of Sm(III) and Gd(III) complexes with the concerned carboxylic acid.

The stability constants of Eu(III) with carboxylic acids were calculated using the SQUAD program. Models including the ML, ML2, MHL and MH2L2 were used in the data analysis. The calculation failed when either or both of MHL and MH2L2 were included, indicating that, for the four carboxylic acid systems only ML and ML2 complexes were present under the conditions of these experiments. Table 3 compares the stability constants of the Eu(III) – aliphatic carboxylate complexes with the values reported by other methods. Potentiometric and calorimetric studies show that malonate [7–10] forms ML, ML2, ML3 and MHL with lanthanides depending on the experimental conditions. Succinate, glutarate and adipate [16] form both ML and ML2 complexes. Under the experimental conditions of the present work, only ML and ML2 complexes were observed for all four ligands. The data in Table 3 shows that the stability constants of Eu(III) complexes with malonate and succinate obtained by different methods are quite close. The stability constants of Eu(III) complexes with glutarate and adipate determined by luminescence are approximately 0.6 log unit smaller than the values by potentiometry. Potentiometric methods do not distinguish between inner-sphere and outersphere complexes. However, spectral measurements mostly reflect inner-sphere complex formation; so, it is probable that the differences between the stability constants obtained in this work and those from potentiometric methods are due to the exclusion of the contribution of outer-sphere complex formation in the spectral measurements. As hard acid, trivalent lanthanides favor complexation with hard bases such as carboxylates in which oxygens are the donor atoms. The stability constants of lanthanides with carboxylate ligands are expected to be proportional to the basicity of the ligands reflected by pKa values. Plots of log i1 of complexes of Ln(III) with

monobasic and dibasic ligands as a function of the pKa of the ligands have shown a linear relationship [17,18]. For the aliphatic dicarboxylates of this study, the stability constants would be expected to increase from malonate to adipate, based on the  pKa of the ligands. However, the opposite trend is observed. In these aliphatic dicarboxylate ligands, the length of the carbon chain between the two carboxylate groups increases from malonate to adipate. Consequently in chelation, malonate forms a six-membered ring while adipate forms a nine-membered ring. A plot of log i1 versus the size of the chelating ring is shown in Fig. 2 (data for oxalate [19] is included in the plot). Fig. 2 indicates that the stability constants decrease sharply as

Fig. 2. Correlation between logs i1 and the chelating ring size in Eu(III) complexes of aliphatic dicarboxylate.

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Table 4 Corrected stability constants, protonation constants and calculated effective charges of carboxylate ligands and Eu(III)–carboxylate complexes Ligand

M:L

log i1(eff)

Z2,HA a

Z2,HA(2) a Z2,MA b

Malonate

1:1 1:2

4.18 6.62

−1.04

−0.65

−1.27 −1.81

Succinate

1:1 1:2

2.99 4.90

−1.04

−0.85

−1.01 −1.43

Glutarate

1:1 1:2

2.66 4.53

−1.00

−0.86

−0.94 −1.35

Adipate

1:1 1:2

2.59 4.84

−1.00

−0.89

−0.93 −1.42

a Calculated anionic charges experienced by the proton for the coordinating carboxylates. In the calculation: B= 0.33; C= 0.75; D =0.15; a0 =4.3; Deff =15.5; d12 = 2.33 A, ; and Z1 = 1.0. b Calculated anionic charges experienced by the metal ion for the coordinating carboxylates. In the calculation: B= 0.33; C= 0.75; D =0.15; a0 =4.3; Deff =57; d12 = 2.64 A, ; and Z1 = 3.0.

Table 5 Luminescence decay constants, hydration number and dehydration number of Eu(III)–carboxylate complexes at 25°C and I = 0.1 M (NaClO4) Ligand

M:Ln

k (ms−1)

nh

nr a

Malonate

1:1 1:2

6.6 4.3

6.5 4.1

2.5 4.9

Succinate

1:1 1:2

6.9 4.8

6.8 4.6

2.2 4.4

Glutarate

1:1 1:2

7.2 5.0

7.1 4.8

1.9 4.2

Adipate

1:1 1:2

7.7 5.6

7.5 5.7

1.6 3.4

a

Assuming a hydration number of 9.0 for aqueous Eu(III).

the ring size increases from 5 to 7, then slowly from 7 to 9. However, it is also possible that chelation becomes less likely as the chain length increases. An earlier calorimetric study of the complexes of uranyl and some trivalent lanthanides with aliphatic dicarboxylates [16] showed that for a particular cation, the formation enthalpies were almost constant from oxalate to glutarate while the formation entropies decreased nearly 30%. Such an entropy decrease was attributed to an increasing loss of conformational entropy in the alkyl chain. The loss of conformational entropy may explain the decreased stability from Eu(III) complexes in the series from oxalate to adipate. Consequently, the ligands may bind Eu(III) in a unidentate nature and/or have increased outer sphere nature. In such cases, the stability constant would decrease even though the  pKa increases.

Based on the assumption that the metal –ligand interaction between a hard acid and a hard base is predominantly ionic, Munze [20] has developed an extended Born equation to calculate the free energy of complexation. This parameterized approach offered a way to derive the effective ligand charge in a metal-ligand complex from the thermodynamic stability constants and has been applied to the complexes of trivalent lanthanides with acetate and croconate and actinide complexes with fluoride ligands [21 –25]. A charge of +3 was used for Eu(III) with values from the literature for Deff, B, C, D and a0, as listed in Table 4 [21 –25], to calculate the effective charge on the aliphatic carboxylate ligands (Table 4). The data in Table 4 indicate that the effective negative charge on the dicarboxylate ligands decreases as the carbon chain increases which supports the hypothesis that as the length of the carbon chain increases, the degree of chelate binding with the ligand decreases. Complexation of ligands with Eu(III) involves the displacement of water molecules from the inner coordination sphere and, thus, the reduction of the luminescence decay constants. From the measured luminescence decay constant, the number of water molecules remaining in the inner coordination sphere of the complexed Eu(III) can be obtained. Since the luminescence decay curves were best fitted into a single exponential, we assume that the binding and the bulk ligand and water molecules in solutions of Eu(III) –carboxylate complexes are in rapid exchange on the time scale of the excited state [26] and the observed luminescence decay constant is a weighted average of the characteristic decay constants of all the species present in the solution. However, the relative luminescence intensity of the aqueous Eu(III) ion is almost negligible compared with that of the complexed Eu(III). Calculations using the weighted average method led to erroneous results. Therefore, the luminescence decay constants of Eu(III) complexes were obtained from the luminescence intensities vs. time obtained for sample solutions in which the molar fraction of both ML and ML2 are greater than 5%, which allows us to ignore any contribution to the overall fluorescence intensity from the uncomplexed aqueous Eu(III) cation. The hydration number in the complexes of Eu(III) with the four dicarboxylate ligands were obtained using the method of Horrocks and Sudnick [26]. The number of water molecules removed by complexation of ligands were obtained by subtracting the hydration number of the complex from the hydration number of the fully hydrated aqueous Eu(III) ion. In the literature, the hydration number of aqueous Eu(III) obtained by various techniques varies from 8.590.5 to 9.69 0.5 [27 – 33]. A study of Eu(III) perchlorate solutions over a range of Eu(III) concentrations resulted in hydration numbers, calculated by the method described elsewhere

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Table 6 Solution conditions and oscillator strengths in the system of Nd(III)–succinic acid at 25°C and I =0.1 M No.

[M]t (10−2 M)

[L]t (10−2 M)

pH

fM

fML

Pe (106)

Nd(III) 1 2 3 4 5 6 7 8 9 10

1.9472 1.9148 1.9043 1.8944 1.8839 1.8740 1.8630 1.8523 1.8410 1.8292

0.0000 0.1255 0.1692 0.2110 0.2554 0.2977 0.3432 0.3863 0.4335 0.4859

2.400 5.263 5.286 5.282 5.273 5.244 5.271 5.320 5.329 5.292

1.0000 0.9447 0.9251 0.9066 0.8870 0.8688 0.8480 0.8275 0.8060 0.7838

0.0000 0.0553 0.0749 0.0934 0.1130 0.1312 0.1520 0.1725 0.1940 0.2162

9.672 10.053 10.144 10.267 10.364 10.498 10.605 10.712 10.833 11.002

Ho(III) 1 2 3 4 5 6 7 8 9 10

2.0136 1.9849 1.9738 1.9620 1.9519 1.9419 1.9319 1.9201 1.9076 1.8950

0.0000 0.1229 0.1651 0.2096 0.2523 0.2980 0.3415 0.3883 0.4326 0.4814

5.008 5.040 5.186 5.328 5.290 5.203 5.183 5.224 5.335 5.385

1.0000 0.9497 0.9302 0.9093 0.8913 0.8733 0.8553 0.8341 0.8117 0.7888

0.0000 0.0503 0.0698 0.0907 0.1087 0.1267 0.1447 0.1659 0.1883 0.2112

6.110 6.390 6.563 6.671 6.778 6.920 7.054 7.196 7.296 7.473

[27], which varied with the concentration of Eu(III) perchlorate [34]. At 5× 10 − 5 M Eu(III) concentrations, the calculated hydration number was 9.8 compared to a hydration number of 8.7 at 0.01 M Eu(III) [34]. In this study we used a hydration number for Eu(III)(aq.) of 9.0. Table 5 lists the luminescence decay constant of each Eu(III) complex, kML, the hydration number, nh, and the number of replaced water molecules, nr. The method of calculating the hydration number carries an inherent uncertainty of 9 0.5 water molecule [27,28]; therefore, the hydration data in Table 6 have errors of 9 0.5. The data in Table 5 indicate that in complexation of Eu(III) with malonate to adipate, the number of water molecules replaced from the innersphere of Eu(III) decreases from 2.5 to 1.6 for the ML complex and 5.4 to 3.9 for the ML2 complex. Barthelemy and Choppin [27] studied this series of complexes at a higher Eu(III) concentration using the same method. By assuming a hydration number of 8.5 for aqueous Eu(III), they found that the dehydration number decreases from 2.3 (malonate) to 1.5 (adipate). If nonahydration of Eu(III) were assumed, these numbers would correspond to 2.8 and 2.0, respectively, which is similar to present results. Both studies indicate that as the carbon chain of the dicarboxylate increases, the number of water molecules released upon complexation decreases. This is in agreement with the interpretation of the trends in log i1 and the DS values which reflects decreased chelate contribution with increasing carbon chain length [16].

3.2. Absorption spectra and experimental oscillator strength Addition of carboxylic ligands to the solution of Nd(III) or Ho(III) caused an increase of the absorbance in the hypersensitive spectral bands studied. For Nd(III), the increase of the absorbance at the longer wavelengths was more significant than at the shorter

Fig. 3. Experimental oscillator strength versus the molar fraction of ML complex in the system of Nd(III) with aliphatic dicarboxylic acids.

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Fig. 4. Experimental oscillator strength versus the molar fraction of ML complex in the system of Ho(III) with aliphatic dicarboxylic acids.

wavelengths. Because of the large mole fraction of the free metal ion, the spectral characteristics of the free metal ion was dominant in all the spectra. The experimental oscillator strength of each solution was calculated from the total metal ion concentration and the area under the absorption peak. Typical data sets for the Nd(III)/Ho(III) – succinic acid systems are listed in Table 6. For each ligand system, the experimental oscillator strength plotted as a function of the mole fraction of the ML complexes are shown in Figs. 3 and 4. By linear regression, the oscillator strengths of the ML complexes were calculated with correlation coefficients better than 0.99. Table 7 lists the calculated oscillator strength, P, of the ML complexes along with the correlation coefficients, r, of the linear plots. After subtraction of the spectral intensity contribution from the free metal ion

Fig. 5. Absorption spectra of Nd(III) (a) and ML complexes of Nd(III) with aliphatic dicarboxylic acids, malonic acid (b), succinic acid (c), glutaric acid (d) and adipic acid (e). Offset to the molar absorptivity: (a) 0.0; (b) 3.0; (c) 6.0; (d) 9.0; (e) 12.0.

at every wavelength from the spectra of the solutions containing both free metal ion and the ML complex, absorption spectra of the ML complex of Nd(III) and Ho(III) with each carboxylate were obtained. Figs. 5 and 6 show the absorption spectra of ML complexes of Nd(III) and Ho(III) with carboxylic acids and Table 7 includes their spectral characteristics. A study of the hypersensitivity in the systems of Nd(III) and Ho(III) complexes with a series of monobasic and dibasic ligands [35,36], reported a linear correlation of the oscillator strengths of the hypersensitive transitions of the complexes with the total ligand basicity as measured by the  pKa values of the binding carboxylates although the slopes of the plots for monobasic ligands and dibasic ligands differed. How-

Table 7 Summary of oscillator strengths and correlation coefficients of the ML complexes and characteristics of absorption spectra of ML complexes of Nd(III) and Ho(III) with the carboxylic acids at 25°C and I= 0.1 M

Nd(III) P×106 R umax (nm) m (M−1 cm−1) Ho(III) P×106 r umax (nm) m (M−1 cm−1)

Ln(III)(aq.)

Malonate

Succinate

Glutarate

Adipate

9.6519 0.055

13.65690.059 0.998 577.4 9.808

15.627 90.093 0.998 576.6 12.944

16.389 90.111 0.998 576.6 15.416

16.586 9 0.096 0.998 576.7 18.156

10.497 9 0.082 0.997 450.1 6.232

12.616 9 0.107 0.998 450.1 7.571

13.840 90.170 0.997 449.8 8.420

14.073 90.148 0.998 449.6 8.827

575.0 7.189 6.0989 0.042 451.7 3.864

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ever, the relationship was not linear for the polyaminopolycarboxylate ligand complexation. Plots of oscillator strengths of ML complexes of Nd(III) and Ho(III) with carboxylate ligands in this work versus  pKaeff (Fig. 7) show that for aliphatic carboxylates the correlation could be linear but the limited data points only allow a conclusion of a trend of increasing oscillator strength with increasing  pKa. This correlation of oscillator strength and pKa reflects a dominantly ionic binding since it shows that the charge on the ligands increases in the same order as oscillator strength. However, non-linearity of the corre-

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lation suggests that besides ligand basicity, there are other factors that can affect the hypersensitivity of lanthanide complexes.

4. Conclusions In conclusion, using laser-induced time-resolved luminescence spectroscopy and lifetime measurement of Eu(III) and UV –Vis absorption spectroscopy of Nd(III) and Ho(III), the complexation of trivalent lanthanides with aliphatic dicarboxylic acids including malonic acid, succinic acid, glutaric acid and adipic acid, were studied. The 5D0 ’ 7F0 excitation spectra of Eu(III) clearly indicated the formation of both ML and ML2 complexes under the present experimental conditions. From the change of the 5D0 ’ 7F0 excitation spectra the stability constants of Eu(III) complexes with these four carboxylic acids were calculated. The resulting stability constants decrease as the carbon chain increases, indicating a decreased chelating effect. Consistent with the change of stability constant, the number of water molecules in the inner-sphere of Eu(III) removed upon ligand complexation also decreases from malonic acid to adipic acid. From the UV –Vis absorption spectra it was found that the oscillator strength seems to correlate with the trend in the basicities of the ligands.

Acknowledgements

Fig. 6. Absorption spectra of Ho(III) and ML complexes of Ho(III) with four aliphatic carboxylic acids, malonic acid (b), succinic acid (c), glutaric acid (d) and adipic acid (e). Offset to the molar absorptivity: (a) 0.0; (b) 1.5; (c) 3.0; (d) 4.5; (e) 6.0.

This research was assisted by a contract from the Division of Chemical Sciences, OBES-USDOE.

References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11]

Fig. 7. Correlation between oscillator strength and total ligand basicity in Nd(III) and Ho(III) complexes with aliphatic dicarboxylates. 1, malonate; 2, succinate; 3, glutarate; 4, adipate. .

[12] [13] [14]

R.A. Torres, G.R. Choppin, Radiochim. Acta 35 (1984) 143. G.R. Choppin, Eur. J. Solid State Inorg. Chem. 28 (1991) 319. G.R. Choppin, S.B. Clark, Marine Chem. 36 (1991) 27. J.V. Beitz, D.L. Bowers, M.M. Doxtader, V.A. Maroni, D.T. Reed, Radiochim. Acta 44/45 (1988) 87. M. Schnitzer, S.O. Khan, Humic Substances in The Environment, Marcel Dekker, New York, 1972. Z. Wang, L.J. van de Burgt, G.R. Choppin, Inorg. Chim Acta. 293 (1999) 167. S.M. Shanbhag, G.R Choppin, Inorg. Chem. 21 (1982) 1696. I. Dellien, I. Grenthe, Acta Chem. Scand. 25 (1971) 1387. J.E. Powell, J.L. Farrell, W.F.S. Neillie, R. Russell, J. Inorg. Nucl. Chem. 30 (1968) 2223. G. Degischer, G.R. Choppin, J. Inorg. Nucl. Chem. 34 (1972) 2823. A.I. Vogel, A Textbook in Quantitative Inorganic Analysis, Longmans, London, 1966. G.R. Choppin, Z.-M. Wang, Inorg. Chem. 36 (1997) 249. D.J. Leggettt, Anal. Chem. 49 (1977) 276. P.J. Davis, P. Rabinowitz, Methods of Numerical Integration, Academic Press, New York, 1975, p. 40.

256

Z.-M. Wang et al. / Inorganica Chimica Acta 310 (2000) 248–256

[15] W.DeW. Horrocks, Jr, D.R. Sudnick, Science 206 (1979) 1194. [16] G.R. Choppin, A. Dadgar, E.N. Rizkalla, Inorg. Chem. 25 (1986) 3581. [17] G.R. Choppin, A. Dadgar, R. Stampfli, J. Inorg. Nucl. Chem. 35 (1973) 875. [18] G.R. Choppin, in: J.-C.G. Bunzli, G.R. Choppin (Eds.), Lanthanide Probes in Life, Chemical and Earth Sciences: Theory and Practice, Elsevier, New York, 1989, p. 1. [19] A.S. Kereichuk, V.I. Paramonova, Radiokhim. 2 (1960) 549. [20] R. Munze, J. Inorg. Nucl. Chem. 34 (1972) 661. [21] G.R. Choppin, L.F. Rao, Radiochim. Acta 37 (1984) 143. [22] G.R. Choppin, E. Orebaugh, Inorg. Chem. 17 (1978) 2300. [23] G.R. Choppin, R.J. Unrein, in: W. Muller, R. Linder (Eds.), Transplutonium 1975, North-Holland, Amsterdam, 1976, p. 97. [24] G.R. Choppin, Radiochim. Acta 32 (1983) 43. [25] L.F. Rao, Ph.D. Dissertation, Florida State University, 1992, p. 9.

.

[26] W.DeW. Horrocks, Jr, V.K. Arkle, F.J. Tiotta, D.R. Sudnick, J. Am. Chem. Soc. 105 (1983) 3455. [27] P.P. Barthelemy, G.R Choppin, Inorg. Chem. 28 (1989) 3354. [28] W.DeW. Horrocks, Jr, D.R. Sudnick, J. Am. Chem. Soc. 101 (1979) 334. [29] G.H. Frost, F.A. Hart, C. Heath, M.B. Hursthouse, Chem. Commun. (1969) 1421. [30] M.J. Lochhead, P.R. Wamsley, K.L. Bray, Inorg. Chem. 33 (2000) 1994. [31] P.J. Breen, W.DeW. Horrocks, Jr, Inorg. Chem. 22 (1983) 536. [32] J.F. Giuliani, T. Donohue, Inorg. Chem. 17 (1978) 1090. [33] Z. Wang, Ph.D. Dissertation, Florida State University, 1994, p. 50. [34] S. Lis, G.R. Choppin, Mat. Chem. Phys. 31 (1992) 159. [35] R.L. Fellows, G.R. Choppin, J. Coord. Chem. 4 (1974) 79. [36] R.L. Fellows, G.R. Choppin, J. Coord. Chem. 3 (1973) 209.