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The thermodynamics of lanthanide complexing by fumarate and maleate
J. inorg, nucl. Chem., 1973, Vol. 35, pp. 875-880.
Pergamon Press.
Printed in Great Britain
THE THERMODYNAMICS OF LANTHANIDE COMPLEXING BY FUMARATE AND MALEATE G. R. C H O P P I N , A. D A D G A R * and R. S T A M P F L I Department of Chemistry, Florida State University, Tallahassee, Florida 32306 (Received 30 May 1972)
Abstract-The stability constants and enthalpies of formation of the lanthanide monofumarate and monomaleate complexes have been measured by pH and calorimetric titration in an ionic medium of 0.10 M (NaCIO4). The fumarate complexes are more stable than would be expected and it is proposed that this is due to the transmission of charge from the unbonded carboxylate group to the bonded one. INTRODUCTION
RELATIVELY little attention has been paid to the complexation of lanthanide cations by dicarboxylate ligands. Stability constants for the family of lanthanide ions have been reported for malonate, succinate and maleate as well as some substituted derivatives of the ligands [ 1-5]. We have reported the thermodynamic parameters for formation of the malonate complexes [1]. In this present work, we have extended these studies to the maleate and fumarate complexes. EXPERIMENTAL Reagents and apparatus Stock solutions of lanthanide perchlorates were prepared as described previously [6]. Final solutions were made by diluting the stock solution to 5 mM. Stock solutions of HCIO4 and NaC104 were used, respectively, to adjust the pH and the final ionic strength to 3.50and 0.10M. The solutions were analyzed for the concentration of the lanthanide ions by titrating with standard ethylenediaminetetraaeetate solution, using xylenol-orange as indicator and hexamethylenetetramine as buffer. Stock solutions of fumaric and maleic acids were prepared using analytical grade reagents (MathesonColman and Bell Co.). These solutions were standardized by potentiometric titration with a standard solution of sodium hydroxide. The measurements have been made with a calorimeter which was designed in this laboratory [1]. The pH-measuring system consisted of a Beckman Research pH meter, Model 1019, a glass electrode (Beckman) and a calomel reference electrode. Procedure and calculations
The pKI and pK2 values for fumaric acid reported by Christensen et al.[7] after correction to 0.10 M ionic strength, were used to calculate the stability constant of the lanthanide monofumarates. The log K1 for the lanthanide fumarates were determined potentiometricaUy by titrating the fumarate buffer *On sabbatical leave: Department of Chemistry, Aria-Mehr University of Technology, Tehran, Iran. 1. 2. 3. 4. 5. 6. 7.
G. Degischer and G. R. Choppin, J. inorg, nucl. Chem. 34, 2823 (1972). J.E. Powell, J. L. Farreil, W. F. S. Neillie and R. Russel, J. inorg, nucl. Chem. 30, 2223 (1968). J.E. Powell and D. K. Johnson, J. inorg, nucl. Chem. 33, 3586 (1971). R. Roulet, J. Feuz and T. VuDuc, Heir. Chim. Acta 53, 1876 (1970). J. M. Peacock and J. C. Jumes, J. chem. Soc. 2233 (1951). G. R. Choppin and J. A. Chopoorian, J. inorg, nucl. Chem. 22, 97 ( 1961). J.J. Christensen, R. M. Izatt and L. D. Hansen, J. Am. chem. Soc. 89, 213 (1967). 875
876
G . R . C H O P P I N , A. D A D G A R and R. S T A M P F L I
into the metal solution. The procedure for the calculation of log K1 has been described elsewhere [8]. The values of the enthalpy change, AH, for formation of the monofumarate and monomaleate complexes of the lanthanides were determined by titrating the metal solutions with the respective fumarate and maleate buffer solutions in the calorimeter. Heats of dilution for each buffer titrant were measured by titration of each respective titrant into a 0.10 M NaC104 solution. The heat data, corrected for the heat of dilutions and the corresponding number of moles of complex formed, were used to calculate the enthalpy change for each step in the complex formation [8]. A least-squares treatment of the data performed with the CDC-6400 computer of Florida State University was used to obtain the K~ and &Hi values and their standard deviations for the fumarate complexes and AH~ (and the standard deviation) for the maleate. The pK~ and pK2 values for maleic acid as well as the log K~ ~ = 0.1 M) values of the maleate complexes reported by Roulet, Feuz and Vu Duc [4] were used to obtain the concentration of the complexes in the enthalpy titrations. RESULTS
Tables 1 and 2 present sample sets of data for the determination of the stability constants and enthalpy changes of the formation of the monofumarate and monoTable 2. Pr(CIO4)a-maleic buffer; T = 25°C; tz = 0.10 M (NaC104) Table 1. Pr(C10~h-fumarate buffer; T = 25°C;/z = 0.10 M (NaC104) Volume (ml)
pH
Q (cad
0.50 1"00 1"50 2.00 2"50 3"00 3"50 4.00 4.50 5.00
3.466 3.540 3'590 3"628 3'656 3'680 3'701 3'718 3'733 3"745
0-009219 0.008740 0.009196 0.008338 0.008706 0.009629 0.007356 0.008024 0.009362 0.00
Initial volume = 50-00 ml; initial pH = 3.50; (metal conc) = 0.004868 M; [HL]--0.01244M; [L]=0.01312 M.
Volume (ml)
pH
0"25 0"50 0"75 1"00 1'25 1'50 1'75 2.00 2'25 2'50 2-75 3"00 3"25 3"50 3-75 4.00
3"801 4"011 4"182 4"302 4"389 4"455 4"508 4"553 4"593 4"629 4.663 4"694 4"725 4'756 4'785 4'815
Q (cal) 0.00959 0-01281 0.01490 0.01700 0.01903 0"01943 0.01912 0"01935 0"01920 0.01900 0-01864 0.01825 0.02071 0"02003 0"01782 0"01798
Initial volume = 50-00ml; initial pH = 3-607; [metal conc] = 0.00487 M; [HL] = 0.0267 M; [L] = 0.0301 M.
maleate complexes of the lanthanides as represented by the reaction 3+ -'~ L(aq) 2- - - LnL~aq) Ln~aq) -
-
where L~aq) 2- can be either the anion of fumaric acid or maleic acid. The thermodynamic parameters and their standard deviations are listed in Table 3 and 4. 8. A. Dadgar and (3. R. Choppin, J. coord. Chem. 1, 179 (1972).
Lanthanide complexing by fumarate and maleate
877
Table 3. Thermodynamic parameters for fumarate complexing; T = 25°C; /z = 0.10 M (NaCIO4)
Ion La Ce Pr Nd Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Log K~ 2.74_+0.01 2'80_+0-01 2.84_+0.01 2.74_+0.01 2-83_+0.01 2.86_+0.02 2.88_+0.02 2.77_+0.01 2.80_+0'02 2-80_+0.02 2.80_+0.02 2.81_+0-02 2.80_+0.02 2.81--+0-02
--AG ~ AH1 AS1 TAS~ (kcal/mole) (kcal/mole) (calldeg/mole) (kcal/mole)
3.73_+0.02 3.81_+0.02 3.86_+0.02 3.73_+0.02 3.85_+0.02 3.89_+0.02 3.91_+0.02 3.77_+0.02 3.81_+0.02 3.81_+0.02 3.81_+0.02 3.82_+0.02 3.81_+0.02 3.82_+0.02
2.69_+0.07 3.16_+0.03 3.24±0.06 3-67_+0.07 3.46_+0-04 3.44___0.04 3.84_+0.06 3.64_+0.05 3.81_+0.04 3'61_+0-08 3.88_+0.06 3.48_+0.03 3-80_+0.06 3.81___0.05
21.6±0-2 23.4_+0"2 23.8+0-2 24.8_+0.3 24-5±0.2 24.6_+0.1 25.9_+0-2 24-9__-0-2 25.5_+0.2 24-9_+0.3 25-8_+0"3 24.5_+0.2 25.5_+0'2 25.6_+0-2
6-42_+0.07 6.97___0.03 7.10_+0.06 7.40_+0.07 7.31_+0.04 7.33_+0.04 7.75___0.06 7.41_+0.05 7.62_+0.04 7-42_+0.08 7-69_+0.06 7.30_+0.03 7.61_+0.06 7.63_+0.05
Table 4. Thermodynamic parameters for maleate complexing; T = 25°C; /z = 0.10 M (NaC10~)
Ion La Pr Nd Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Log KI 3.44___0.01 3.63_+0.01 3-66+--0.01 3.82___0.01 3.83_+0.01 3.79_+0.01 3.74_+0.01 3.74_+0.01 3.67_+0.01 3-64_+0.01 3-62___0.01 3-64_+0.01 3.59_+0.01
--AG 1 (kcal/mole)
AHI AS1 TAS 1 (kcal/mole) (cal/deg/mole) (kcallmole)
4.69___0.01 4-95___0.01 4.99_+0.01 5.20_+0.01 5.23_+0"01 5.18_+0.01 5.10_+0'01 5-10_+0.01 5.00_+0.01 4.96___0.01 4-93_+0.01 4.96_+0.01 4-90_+0.01
3.06_+0.11 2.94_+0.14 2.89_+0-08 3.04_+0.09 3.40_+0.12 3.57_+0.11 4.15_+0.12 4.16_+0.11 4.36_+0.15 4-31_+0.13 4.38_+0.16 4-43_+0.15 4-41_+0.13
26.0__-0.4 26.5+0.5 26.4_+0.3 27.6_+0.3 28.9_+0.4 29.3_+0.4 31.0_+0-4 31.1_+0'4 31.4__-0.5 31'1 -+0.4 31.2_+0-5 31.5_+0.5 31.2_+0.4
7.75_+0.11 7.90_+0.14 7.88_+0.08 8.24_+0.09 8-63_+0.12 8-75___0.11 9-25_+0.12 9-27_+0.11 9-36_+0.15 9-27___.0.13 9.31_+0-16 9.39_+0-15 9.31_+0.13
DISCUSSION
In Fig. 1 we have plotted the values of log K1 for lutetium complexes of a variety of di- and monocarboxylate ligands as a function of pK2 (for the dicarboxylates) and pK1 (for the monocarboxylates). Ligands 1-7 are malonate and substituted malonates, forming 6 membered ring chelates. Log K~ values for ligands 2-7 were measured in/x = 0.10 M (KNO3) [2, 3]. They have been adjusted in Fig. 1 to values for/x = 0.10 M (NaCIO4) by using the difference (0.3 log K units) between the values of the malonate complex in KNO312] and NaCIO4[1]. The values for ligands 14 and 15 were similarly corrected for the difference in the acetate values in these two media [9, 10]. 9. R. S. Kolat and J. E. Powell, lnorg. Chem. 1, 293 (1962). 10. I. Grenthe, Acta chem. scand. 18, 283 (1964).
878
G.R.
CHOPPIN,
A. D A D G A R
and R. S T A M P F L I
5~
4~ I=M 2=EM 3=DMM 4=DEM 5=DPM 6=DBM 7=CPDA 8=MS
10 8
~0
/ 1~ 2.0
1.0" 2.0
I1---F 12=HM 13=A 14=CIA 15=CIP
/
~
, ~ l ' ~ tI 3.0
4.0
,
13 I
5.0 pK
,
I 6.0
,
I
ZO
,
I 8.0
Fig. 1. Relationship between the values of log/31 for the formation of L n X and the pKI (or the monobasic ligands) or pK2 (of the dibasic ligands): M = m a l o n a t e [ l ] ; EM = ethylmalonate[2]; D M M = dimethylmalonate[2]; D E M = diethylmalonate[2]; D P M = dipropylmalonate[2]; D B M = dibutylmalonate[2]; C P D A = l,l-cyclopentanedicarboxylate [7]; MS = methylsuccinate [4]; m a = malate [4]; m = maleate; F = fumarate; H M = bimalonate [11]; A = acetate [10]; C1A = chloroacetate [12]; CIP = chloropropionate [ 12].
The malonato (1) and ethylmalonato (2) complexes are somewhat more stable than their pK2 values would indicate when compared to the dialkyl substituted malonate (3-6) and the 1,1-cyclopentanedicarboxylate complexes (7). Powell and Johnson[3] have proposed that the steric effects of the alkyl substituents result in a diminished - O O C - C - C O O - bond angle which in turn affects the stability constants. This steric effect is reflected in the pK2 since hydrogen bonding results in a ring involving the - O O C - C - C O O - entity. However, we cannot expect that the stability of the chelate ring involving the large, trivalent lanthanide cations would be affected in exactly the same fashion as the protonated ring. The enhanced stability of the malonato and ethylmalonato complex may reflect the greater attraction associated with the larger - O O C - C - C O O - bond angles in chelate rings with the lanthanide cations. Comparison in Fig. 1 of the stability of the complexes of the malonate type ligands (1-7), which form six membered rings, with that of the maleate (10) and methylsuccinate (8) complexes which have seven membered rings indicates the effect of ring size. Malate (9) is considerably more stable than expected from the pK2 value, no doubt reflecting its behavior as a tridentate ligand. The fumarate (11) complex is more than an order of magnitude more stable than would be expected from the trend in monocarboxylate ligands, including bimalonate (12). However, the conjugated double bond system of fumarate may allow transmission of negative charge from the unbonded carboxylate group to the bonded one. That this occurs is suggested by the fact that the log/31 of lutetium 11. C. H. Ke, P. C. Kong, H. S. C h e n g a n d N. C. Li, J. inorg, nucl. Chem. 30, 961 (1968). 12. G . R . Choppin, G. D e g i s c h e r and E. G. Orebaugh. T o be published.
Lanthanide complexing by fumarate and maleate
879
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The variation of the thermodynamic parameters for the formation of the m o n o complexes with lanthanide atomic number: M = m a l o n a t e ; m = m a l e a t e ; F = fumarate.
monofumarate falls on the monocarboxylate line when PK1 + pK2 is plotted as shown by point (11) in Fig. 1. In Fig. 2 the thermodynamic parameters of the formation of the monomalonate (M), monomaleate (m) and monofumarate (F) complexes of the lanthanides are presented. We see that the enthalpy changes are similar and the major contribution to the differences in stability among these three ligands is attributable to the entropy differences. In a previous publication [1] we have shown the similarity in values between the thermodynamic parameters for the formation of the chelated lanthanide monomalonates, LnM +, and the monodentate lanthanide diacetates, LnA2 +. It might seem surprising that the monodentate lanthanide monofumarate complexes should have values so close to those of the chelates. However, if our conclusion from the position in Fig. 1 of log/31 for the fumarate complexation (that the bonded carboxylate draws extra negative charge to itself from the trans, unbonded carboxylate group) is correct, this similarity is understandable. We can test this explanation of the fumarate values further by comparing them with the analogous values for the complexation with a dinegative, monodentate ligand such as sulfate. The thermodynamic values of the formation of LnSO4 + have been reported for zero[13] and 2.0 M[14] ionic strength. We can estimate from these data values for formation of LnSO4 ÷ of AH1 - 3.5 kcal 13. D . P.
Fay and N. Purdie, J. phys. Chem. 73, 3 4 6 2 ( 1 9 6 9 ) . inorg, nucl. Chem.
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29, 7 3 7 ( 1 9 6 7 ) .
880
G . R . CHOPPIN, A. D A D G A R and R. S T A M P F L I
mole -1 and T A S 1 ~ 7.2 kcal mole -1 which agree well with the values for the formation of the lanthanide monofumarate. By contrast, analogous values for the formation of LnX +2 with uninegative monodentate ligands are: for acetate (estimated from/z = 2.0 M data[9, 10]), AH1 -- 1"4 kcal mole -1, T A S 1 ~ 5.4 kcal mole-l; for iodate (/z = 0.1M)[15], AH~ = 2.6 kcal mole -x, T A S = 4.3kcal mole -~.
Acknowledgements-This research was supported by the U.S.A.E.C. under Contract AT-(40-1)-1797. The computer calculations were partially subsidized by an N S F grant to the F S U computer center. R.S. wishes to acknowledge the assistance of a Swiss grant Stiftung fur Stipendien aiif dem Gebiete der Chemic. 15. G. R. Choppin and S. Bertha. To be published.
Pergamon Press.
Printed in Great Britain
THE THERMODYNAMICS OF LANTHANIDE COMPLEXING BY FUMARATE AND MALEATE G. R. C H O P P I N , A. D A D G A R * and R. S T A M P F L I Department of Chemistry, Florida State University, Tallahassee, Florida 32306 (Received 30 May 1972)
Abstract-The stability constants and enthalpies of formation of the lanthanide monofumarate and monomaleate complexes have been measured by pH and calorimetric titration in an ionic medium of 0.10 M (NaCIO4). The fumarate complexes are more stable than would be expected and it is proposed that this is due to the transmission of charge from the unbonded carboxylate group to the bonded one. INTRODUCTION
RELATIVELY little attention has been paid to the complexation of lanthanide cations by dicarboxylate ligands. Stability constants for the family of lanthanide ions have been reported for malonate, succinate and maleate as well as some substituted derivatives of the ligands [ 1-5]. We have reported the thermodynamic parameters for formation of the malonate complexes [1]. In this present work, we have extended these studies to the maleate and fumarate complexes. EXPERIMENTAL Reagents and apparatus Stock solutions of lanthanide perchlorates were prepared as described previously [6]. Final solutions were made by diluting the stock solution to 5 mM. Stock solutions of HCIO4 and NaC104 were used, respectively, to adjust the pH and the final ionic strength to 3.50and 0.10M. The solutions were analyzed for the concentration of the lanthanide ions by titrating with standard ethylenediaminetetraaeetate solution, using xylenol-orange as indicator and hexamethylenetetramine as buffer. Stock solutions of fumaric and maleic acids were prepared using analytical grade reagents (MathesonColman and Bell Co.). These solutions were standardized by potentiometric titration with a standard solution of sodium hydroxide. The measurements have been made with a calorimeter which was designed in this laboratory [1]. The pH-measuring system consisted of a Beckman Research pH meter, Model 1019, a glass electrode (Beckman) and a calomel reference electrode. Procedure and calculations
The pKI and pK2 values for fumaric acid reported by Christensen et al.[7] after correction to 0.10 M ionic strength, were used to calculate the stability constant of the lanthanide monofumarates. The log K1 for the lanthanide fumarates were determined potentiometricaUy by titrating the fumarate buffer *On sabbatical leave: Department of Chemistry, Aria-Mehr University of Technology, Tehran, Iran. 1. 2. 3. 4. 5. 6. 7.
G. Degischer and G. R. Choppin, J. inorg, nucl. Chem. 34, 2823 (1972). J.E. Powell, J. L. Farreil, W. F. S. Neillie and R. Russel, J. inorg, nucl. Chem. 30, 2223 (1968). J.E. Powell and D. K. Johnson, J. inorg, nucl. Chem. 33, 3586 (1971). R. Roulet, J. Feuz and T. VuDuc, Heir. Chim. Acta 53, 1876 (1970). J. M. Peacock and J. C. Jumes, J. chem. Soc. 2233 (1951). G. R. Choppin and J. A. Chopoorian, J. inorg, nucl. Chem. 22, 97 ( 1961). J.J. Christensen, R. M. Izatt and L. D. Hansen, J. Am. chem. Soc. 89, 213 (1967). 875
876
G . R . C H O P P I N , A. D A D G A R and R. S T A M P F L I
into the metal solution. The procedure for the calculation of log K1 has been described elsewhere [8]. The values of the enthalpy change, AH, for formation of the monofumarate and monomaleate complexes of the lanthanides were determined by titrating the metal solutions with the respective fumarate and maleate buffer solutions in the calorimeter. Heats of dilution for each buffer titrant were measured by titration of each respective titrant into a 0.10 M NaC104 solution. The heat data, corrected for the heat of dilutions and the corresponding number of moles of complex formed, were used to calculate the enthalpy change for each step in the complex formation [8]. A least-squares treatment of the data performed with the CDC-6400 computer of Florida State University was used to obtain the K~ and &Hi values and their standard deviations for the fumarate complexes and AH~ (and the standard deviation) for the maleate. The pK~ and pK2 values for maleic acid as well as the log K~ ~ = 0.1 M) values of the maleate complexes reported by Roulet, Feuz and Vu Duc [4] were used to obtain the concentration of the complexes in the enthalpy titrations. RESULTS
Tables 1 and 2 present sample sets of data for the determination of the stability constants and enthalpy changes of the formation of the monofumarate and monoTable 2. Pr(CIO4)a-maleic buffer; T = 25°C; tz = 0.10 M (NaC104) Table 1. Pr(C10~h-fumarate buffer; T = 25°C;/z = 0.10 M (NaC104) Volume (ml)
pH
Q (cad
0.50 1"00 1"50 2.00 2"50 3"00 3"50 4.00 4.50 5.00
3.466 3.540 3'590 3"628 3'656 3'680 3'701 3'718 3'733 3"745
0-009219 0.008740 0.009196 0.008338 0.008706 0.009629 0.007356 0.008024 0.009362 0.00
Initial volume = 50-00 ml; initial pH = 3.50; (metal conc) = 0.004868 M; [HL]--0.01244M; [L]=0.01312 M.
Volume (ml)
pH
0"25 0"50 0"75 1"00 1'25 1'50 1'75 2.00 2'25 2'50 2-75 3"00 3"25 3"50 3-75 4.00
3"801 4"011 4"182 4"302 4"389 4"455 4"508 4"553 4"593 4"629 4.663 4"694 4"725 4'756 4'785 4'815
Q (cal) 0.00959 0-01281 0.01490 0.01700 0.01903 0"01943 0.01912 0"01935 0"01920 0.01900 0-01864 0.01825 0.02071 0"02003 0"01782 0"01798
Initial volume = 50-00ml; initial pH = 3-607; [metal conc] = 0.00487 M; [HL] = 0.0267 M; [L] = 0.0301 M.
maleate complexes of the lanthanides as represented by the reaction 3+ -'~ L(aq) 2- - - LnL~aq) Ln~aq) -
-
where L~aq) 2- can be either the anion of fumaric acid or maleic acid. The thermodynamic parameters and their standard deviations are listed in Table 3 and 4. 8. A. Dadgar and (3. R. Choppin, J. coord. Chem. 1, 179 (1972).
Lanthanide complexing by fumarate and maleate
877
Table 3. Thermodynamic parameters for fumarate complexing; T = 25°C; /z = 0.10 M (NaCIO4)
Ion La Ce Pr Nd Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Log K~ 2.74_+0.01 2'80_+0-01 2.84_+0.01 2.74_+0.01 2-83_+0.01 2.86_+0.02 2.88_+0.02 2.77_+0.01 2.80_+0'02 2-80_+0.02 2.80_+0.02 2.81_+0-02 2.80_+0.02 2.81--+0-02
--AG ~ AH1 AS1 TAS~ (kcal/mole) (kcal/mole) (calldeg/mole) (kcal/mole)
3.73_+0.02 3.81_+0.02 3.86_+0.02 3.73_+0.02 3.85_+0.02 3.89_+0.02 3.91_+0.02 3.77_+0.02 3.81_+0.02 3.81_+0.02 3.81_+0.02 3.82_+0.02 3.81_+0.02 3.82_+0.02
2.69_+0.07 3.16_+0.03 3.24±0.06 3-67_+0.07 3.46_+0-04 3.44___0.04 3.84_+0.06 3.64_+0.05 3.81_+0.04 3'61_+0-08 3.88_+0.06 3.48_+0.03 3-80_+0.06 3.81___0.05
21.6±0-2 23.4_+0"2 23.8+0-2 24.8_+0.3 24-5±0.2 24.6_+0.1 25.9_+0-2 24-9__-0-2 25.5_+0.2 24-9_+0.3 25-8_+0"3 24.5_+0.2 25.5_+0'2 25.6_+0-2
6-42_+0.07 6.97___0.03 7.10_+0.06 7.40_+0.07 7.31_+0.04 7.33_+0.04 7.75___0.06 7.41_+0.05 7.62_+0.04 7-42_+0.08 7-69_+0.06 7.30_+0.03 7.61_+0.06 7.63_+0.05
Table 4. Thermodynamic parameters for maleate complexing; T = 25°C; /z = 0.10 M (NaC10~)
Ion La Pr Nd Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Log KI 3.44___0.01 3.63_+0.01 3-66+--0.01 3.82___0.01 3.83_+0.01 3.79_+0.01 3.74_+0.01 3.74_+0.01 3.67_+0.01 3-64_+0.01 3-62___0.01 3-64_+0.01 3.59_+0.01
--AG 1 (kcal/mole)
AHI AS1 TAS 1 (kcal/mole) (cal/deg/mole) (kcallmole)
4.69___0.01 4-95___0.01 4.99_+0.01 5.20_+0.01 5.23_+0"01 5.18_+0.01 5.10_+0'01 5-10_+0.01 5.00_+0.01 4.96___0.01 4-93_+0.01 4.96_+0.01 4-90_+0.01
3.06_+0.11 2.94_+0.14 2.89_+0-08 3.04_+0.09 3.40_+0.12 3.57_+0.11 4.15_+0.12 4.16_+0.11 4.36_+0.15 4-31_+0.13 4.38_+0.16 4-43_+0.15 4-41_+0.13
26.0__-0.4 26.5+0.5 26.4_+0.3 27.6_+0.3 28.9_+0.4 29.3_+0.4 31.0_+0-4 31.1_+0'4 31.4__-0.5 31'1 -+0.4 31.2_+0-5 31.5_+0.5 31.2_+0.4
7.75_+0.11 7.90_+0.14 7.88_+0.08 8.24_+0.09 8-63_+0.12 8-75___0.11 9-25_+0.12 9-27_+0.11 9-36_+0.15 9-27___.0.13 9.31_+0-16 9.39_+0-15 9.31_+0.13
DISCUSSION
In Fig. 1 we have plotted the values of log K1 for lutetium complexes of a variety of di- and monocarboxylate ligands as a function of pK2 (for the dicarboxylates) and pK1 (for the monocarboxylates). Ligands 1-7 are malonate and substituted malonates, forming 6 membered ring chelates. Log K~ values for ligands 2-7 were measured in/x = 0.10 M (KNO3) [2, 3]. They have been adjusted in Fig. 1 to values for/x = 0.10 M (NaCIO4) by using the difference (0.3 log K units) between the values of the malonate complex in KNO312] and NaCIO4[1]. The values for ligands 14 and 15 were similarly corrected for the difference in the acetate values in these two media [9, 10]. 9. R. S. Kolat and J. E. Powell, lnorg. Chem. 1, 293 (1962). 10. I. Grenthe, Acta chem. scand. 18, 283 (1964).
878
G.R.
CHOPPIN,
A. D A D G A R
and R. S T A M P F L I
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4~ I=M 2=EM 3=DMM 4=DEM 5=DPM 6=DBM 7=CPDA 8=MS
10 8
~0
/ 1~ 2.0
1.0" 2.0
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/
~
, ~ l ' ~ tI 3.0
4.0
,
13 I
5.0 pK
,
I 6.0
,
I
ZO
,
I 8.0
Fig. 1. Relationship between the values of log/31 for the formation of L n X and the pKI (or the monobasic ligands) or pK2 (of the dibasic ligands): M = m a l o n a t e [ l ] ; EM = ethylmalonate[2]; D M M = dimethylmalonate[2]; D E M = diethylmalonate[2]; D P M = dipropylmalonate[2]; D B M = dibutylmalonate[2]; C P D A = l,l-cyclopentanedicarboxylate [7]; MS = methylsuccinate [4]; m a = malate [4]; m = maleate; F = fumarate; H M = bimalonate [11]; A = acetate [10]; C1A = chloroacetate [12]; CIP = chloropropionate [ 12].
The malonato (1) and ethylmalonato (2) complexes are somewhat more stable than their pK2 values would indicate when compared to the dialkyl substituted malonate (3-6) and the 1,1-cyclopentanedicarboxylate complexes (7). Powell and Johnson[3] have proposed that the steric effects of the alkyl substituents result in a diminished - O O C - C - C O O - bond angle which in turn affects the stability constants. This steric effect is reflected in the pK2 since hydrogen bonding results in a ring involving the - O O C - C - C O O - entity. However, we cannot expect that the stability of the chelate ring involving the large, trivalent lanthanide cations would be affected in exactly the same fashion as the protonated ring. The enhanced stability of the malonato and ethylmalonato complex may reflect the greater attraction associated with the larger - O O C - C - C O O - bond angles in chelate rings with the lanthanide cations. Comparison in Fig. 1 of the stability of the complexes of the malonate type ligands (1-7), which form six membered rings, with that of the maleate (10) and methylsuccinate (8) complexes which have seven membered rings indicates the effect of ring size. Malate (9) is considerably more stable than expected from the pK2 value, no doubt reflecting its behavior as a tridentate ligand. The fumarate (11) complex is more than an order of magnitude more stable than would be expected from the trend in monocarboxylate ligands, including bimalonate (12). However, the conjugated double bond system of fumarate may allow transmission of negative charge from the unbonded carboxylate group to the bonded one. That this occurs is suggested by the fact that the log/31 of lutetium 11. C. H. Ke, P. C. Kong, H. S. C h e n g a n d N. C. Li, J. inorg, nucl. Chem. 30, 961 (1968). 12. G . R . Choppin, G. D e g i s c h e r and E. G. Orebaugh. T o be published.
Lanthanide complexing by fumarate and maleate
879
15.0 13.(] 11£
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o.-.-.r " - ' ~ '
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5.C
:
•
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La C o F i g . 2. ligand
I
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I
I
i
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Lu
The variation of the thermodynamic parameters for the formation of the m o n o complexes with lanthanide atomic number: M = m a l o n a t e ; m = m a l e a t e ; F = fumarate.
monofumarate falls on the monocarboxylate line when PK1 + pK2 is plotted as shown by point (11) in Fig. 1. In Fig. 2 the thermodynamic parameters of the formation of the monomalonate (M), monomaleate (m) and monofumarate (F) complexes of the lanthanides are presented. We see that the enthalpy changes are similar and the major contribution to the differences in stability among these three ligands is attributable to the entropy differences. In a previous publication [1] we have shown the similarity in values between the thermodynamic parameters for the formation of the chelated lanthanide monomalonates, LnM +, and the monodentate lanthanide diacetates, LnA2 +. It might seem surprising that the monodentate lanthanide monofumarate complexes should have values so close to those of the chelates. However, if our conclusion from the position in Fig. 1 of log/31 for the fumarate complexation (that the bonded carboxylate draws extra negative charge to itself from the trans, unbonded carboxylate group) is correct, this similarity is understandable. We can test this explanation of the fumarate values further by comparing them with the analogous values for the complexation with a dinegative, monodentate ligand such as sulfate. The thermodynamic values of the formation of LnSO4 + have been reported for zero[13] and 2.0 M[14] ionic strength. We can estimate from these data values for formation of LnSO4 ÷ of AH1 - 3.5 kcal 13. D . P.
Fay and N. Purdie, J. phys. Chem. 73, 3 4 6 2 ( 1 9 6 9 ) . inorg, nucl. Chem.
14. R . G . d e C a r v a l h o a n d G . R . C h o p p i n , J .
29, 7 3 7 ( 1 9 6 7 ) .
880
G . R . CHOPPIN, A. D A D G A R and R. S T A M P F L I
mole -1 and T A S 1 ~ 7.2 kcal mole -1 which agree well with the values for the formation of the lanthanide monofumarate. By contrast, analogous values for the formation of LnX +2 with uninegative monodentate ligands are: for acetate (estimated from/z = 2.0 M data[9, 10]), AH1 -- 1"4 kcal mole -1, T A S 1 ~ 5.4 kcal mole-l; for iodate (/z = 0.1M)[15], AH~ = 2.6 kcal mole -x, T A S = 4.3kcal mole -~.
Acknowledgements-This research was supported by the U.S.A.E.C. under Contract AT-(40-1)-1797. The computer calculations were partially subsidized by an N S F grant to the F S U computer center. R.S. wishes to acknowledge the assistance of a Swiss grant Stiftung fur Stipendien aiif dem Gebiete der Chemic. 15. G. R. Choppin and S. Bertha. To be published.