Mercury(II) complexes with sulfhydryl containing chelating agents: Stability constant inconsistencies and their resolution

J. inorg, nucl. Chem. Vol. 42, pp. 99-102 Pergamon Press Ltd., 1980. Printed in Great Britain

MERCURY(II) COMPLEXES WITH SULFHYDRYL CONTAINING CHELATING AGENTS: STABILITY CONSTANT INCONSISTENCIES AND THEIR RESOLUTION J. S. CASAS and MARK M. JONES Department of Chemistry and Center in Environmental Toxicology, Vanderbilt University, Nashville, TN 37235, U.S.A.

(Received 31 May 1979) Abstract--An experimental study has been carried out to provide information on the nature of the inconsistencies in previously reported values of the stability constants of mercury(II) complexes with sulfhydryl containing chelating agents. Stability constants have been determined for the mercury(II) complexes with D,L-penicillamine (PA), N-Acetyl-D,L-penicillamine(NAPA). 2,3-dimercaptopropanesulfonate (DMPS), and 2,3-dimercaptopropanol1 (BAL) using a titration procedure involving simultaneous measurements with a glass and a mercury electrode. The stability constants obtained in this manner were of the order of 1035-1044.Chelating agents with two sulfhydryl groups give complexes with mercury(II) whose K~ values are greater by a factor of l 0 7 than those which contain but a single sulfhydryl group. Arguments are presented in favor of the validity of such large values. INTRODUCTION

EXPERIMENTAL

A casual survey of the reported values for the stability constants of Hg 2+ complexes with sulfhydryl bearing chelating agents reveals a number of anomalies. These become especially troublesome when an attempt is make to systematize our understanding of the relative effectiveness of such chelating agents in the treatment of acute or chronic mercury poisoning. The values reported for three chelates with very similar structures illustrate this clearly. Thus, log K, for 2,3-dimercaptopropanesulfonate is reported as 39.711], while the value for the 2,3-dimercaptosuccinic acid is given as 17.312], and the value for 2,3-dimercaptopropanol-I is given as 25.74[3]. Turning to another closely related series, one finds values of log/32 = 43.6[4], log/32[Hg(LH)2] = 39.414a] or log K~ = 14.215] for cysteine; log K~ = 17.5 and log K2 = 6 for D-peniciilamine[6] or log Kt = 16.1513] or 16.416a]; and for N-Acetyl-D,L-penicillamine a report of log K2 = 1017]. The enormous range of values reported here indicate some serious problems. Their extreme variability has made it impossible to carry out any systematic study of the relationship of the stabilities of these mercury complexes and the relative effectiveness of the chelating agents as therapeutic antidotes for acute or chronic mercury poisoning[8]. For this reason a set of consistently obtained values for such compounds as have already proven useful in mercury poisoning is needed to provide information on the relative changes in stability constants associated with various structural changes. We thus set out to collect such data to assist us in our search for more effective chelating agents for use in chronic and acute mercury poisoning. Because the wide variations in the values reported for compounds which are structurally very similar might well be due to differences in the procedures, we decided to examine several of these compounds using the apparatus described by Schmid and Reilly[9, 10]. This allows titrations of the HgZ+-chelate solution using a standard solution of sodium hydroxide and the simultaneous measurements of the potentials of a glass and a mercury electrode. We hoped in this way to achieve at least a set of values for the stability constants which were capable of consistent intercomparison.

The chelating agents used were obtained from the following sources: 2,3-dimercaptopropanol-I (BAL), Sigma Chem. Co., 2,3dimercaptopropane-sulfonate, sodium salt (DMPS): Heyl and Co., Berlin; D.L-penicillamine(PA); Aldrich Chem. Co. puriss; N-AcetyI-D.L-penicillamine(NAPA); Sigma Chem. Co. All were obtained in the highest state of purity commercially available and were used without further purification. Their solutions were prepared on the day in which they were used and standardized by potentiometric titration prior to use in the stability constant studies. The solutions of mercuric chloride were prepared from analytical grade mercuric chloride (Mallinckrodt) and were standardized via complexometric titration [11].

Potentiometric titrations The titrations were carried out in a 400-ml jacketed titration cell through whose jacket water thermostated at 25.0-+0. I°C, was circulated. The cell was capped with a rubber stopper through which were inserted glass and j-tube mercury electrodes, two calomel electrodes, a microburet delivery tube and nitrogen (freed of CO2) inlet and outlet tubes. The solutions were stirred by an air-propelled magnetic stirrer. The pH and potential of the solutions were monitored by two Beckman Research Model pH-meters. Glass and calomel electrodes were standardized before each titration with pH 7.00 buffer solution. The operation of the J-tube mercury electrode was checked by the determination of the formation constant of the mercuric ion-EDTA complex. All potentiometric measurements were made at an ionic strength of approximately 0.1 M in sodium perchlorate. The concentration of the metal in the cell was 10-~M in all the experiments and the concentration of the ligand was approx. 6 × 10-3 M. The titrant used was approximately 0.1 M NaOH. In calculating the free-hydrogen and hydroxyl-ion concentrations, the mean activity coefficient /'~ was taken as 0.78 and K~ = 1.007 × 10 14112]. RESULTS The operation of the mercury electrode was standardized with several determinations of the stability constant of the HgEDTA complex. We obtained values of log K, for this as shown in Table I. The value reported by Schmid and Reilley[9] was 22.1. The variability in the individual log K, values in Table 1 is of the order of -+0.1 in log units. The results obtained for the constants involved in the present studies are listed in 99

100

J. S. CASASand M. M. JONES Table 1. Log K: for HgEDTA Mole r a t i o log K

(EDTA/Hg)

l.S 22.0

1

2

3

4

6

22.1

22.4

22.1

22.1

Table 2. Dissociationconstants of ligands and stabUityconstants of mercury(II)complexes i

PK1H

C h e l a t i n g Agent D,L-pencillamine

7.88 (3)

PK2H

log t 1

log t 2

10.43 (5)

38.5±0.1

6.1±0.2

N-Acetyl D , L - p e n c i l l a m i n e 3.60±0.05 (a) 10.0l±O.10 (a) 35.41±0.1

6.2±0.2

2 , 3 - D i m e r e a p t o p r o p a n e s u l - 8.655 (13) fonate 2 , 3 - d e m e r e a p t o p r o p a n o l - 1 8.618 (14) a.

11.939 (13}

42.2±0.1

10.9±0.2

10.567 (14)

44.8±0.1

7.11±0.2

This work

Table 2. These values, it should be noted, were all obtained under conditions which were, as nearly as possible, strictly comparable. The potential-pH diagram obtained with o£-penicillamine is shown in Fig. 1. The calculation of the log K~ and log K2 values were made between pH 5 and pH 7 for 2,3-dimercaptropropanesulfonate-I and between a pH of 6 and 8 for each of the other ligands. The reason for this was that the solutions were clear over these ranges and the measurements were more reproducible. For the 2,3dimercaptopropanesulfonate there was evidence of decomposition or oxidation in media with a pH greater than 7.

"tOO.

-3OO"

lilllm~ ro|io J~6

~

.

"

'SO0-

#s

Fig. 1. Plot of experimentalvalues obtained for measurementsof the potential and pH values in the titration of a mercuric chloride--penicillaminesystem with 0.1 M NaOH.

DISCUSSION The general reluctance to give credence to such large stability constants requires that we examine the process by which they have been obtained as well as other evidence bearing on the relative stabilities of such complexes. Because of the many orders of magnitude differences among these our attention must center on the way in which these have been obtained. The acid dissociation constants of the chelating agents were obtained from the literature except for the case of N-Acetyl D,L-penicillamine where our own data was used with the calculational procedure of Lenz and Martell[5]. Our results for these constants are in good accord with those obtained on this same compound at 20° by Doornbos and Faber[15]. The procedure we used in our calculations was based upon the assumption that we could make accurate simultaneous measurements of both [H +] and [Hg2÷] in the systems examined. In our case the chief problem in such an assumption was the assurance that the mercury electrode was operating properly in the media we used. The chelating species we used can also participate in the formation of polynuclear species[16, 17]. These polynuclear species can be sparingly soluble and may interfere with the proper functioning of the mercury electrode. To avoid this we used a chelating agent to mercury(II) ratio of 6:1 in our titrations. No constant is reported in the table for 2,3-dimercapto-succinic acid because the low solubility of the ligand and its complexes made it impossible to attain this molar ratio while the mercury concentration was 10-3 M. The experiments carried out at a molar ratio of 2:1 showed poor reproducibility, but seemed to suggest a stability constant for this system of the same order of magnitude as those reported in Table 2. In calculating the stability constants we made the following further assumptions: (i) that the mercury species in solution were: Hg2+, HgL and HgL22-; (ii) that given the large excess of chelating agent present we could provisionally ignore the chloride ion species in solution. In these systems it proved to he very advantageous to work with HgCI2 rather than mercuric perchlorate or nitrate because these latter salts gave solutions which were turbid at lower pH values. The significant equilibria involving the complexes were thus

Mercury(II)complexeswith sulfhydrylcontainingchelatingagents Hg2+ + L2-~HgL

K, =

[HgL] [Hg2+] [L 2-]

(1)

and HgL + L 2 ~HgL2 i -

K2= [Hgl-a2-] [HgL] [L2-]

(2)

The term [L 2-] was calculated from the usual expression: [L 2 ] = 2[H2L]o - [NaOH] - [H+]][OH -] (3) where +

o_[H I. P

-

+2

2[H ]

(4)

T ~ - / T " ~ ";-

and K1n and K2n are the first and second ionization constants of H2L. [Hg2+] was calculated from the measured potential differences between the mercury and the calomel electrodes using the expression log [Hg2+] = (E - E°)/(0.0296).

(5)

Here E ° is 0.614V vs a saturated calomel electrode. We may also write [H2L]o = IHgL] + 2[HgL22-] + $[L 2-]

[61

[Hg2+]o = [Hg2÷] + [HgL] + [HgL2],

(7)

and

where [H2L]o and [Hg2+]o are the initial total concentrations of the chelating agent and of mercury; [Hg2÷] is the free mercuric ion concentration, and IH+I

lH+l2

(8)

By combining these expressions we can solve for [HgL] as

[HgL] = 2[Hg2+]o+ [L2-]~b- [H2L]o- 2[Hg2*],

(9)

and [HgL2] can be solved for analogously. In order to ascertain the validity of the assumption about chloride ion, the stability constants for N-Acetyl D,L-penicillamine were determined in a solution of mercury(II) and perchloric acid, with no chloride present whatsoever. Under these conditions it was found that log K~ = 35.3 and log K2 = 6.5. These values do not differ significantly from those reported in Table 2 which were determined using mercuric chloride. It is now necessary to inquire into the basis for the great discordancy among the values available for the stability constants. The obvious places to look are in the assumptions underlying the two different procedures. For the studies using the mercury electrode the most obvious point is the reversibility and reliability of the mercury electrode under the conditions used. It is certainly true that there are conditions under which it is not reliable [9, I0]. The fundamental reason for the high values obtained

101

for these stability constants lies in the very low values obtained for [Hg 2÷] free. It is evident that the mercury electrode indicates that the value of this parameter is much less in these solutions than one estimates from other procedures by means of more indirect evaluation. It was for this reason that the reproducibility of the electrode was repeatedly tested under the experimental conditions utilized. In addition, even admitting a value of [Hg2÷] similar to that which exists in the system Hg2÷ :EDTA, the resulting constants would be of the order of 1024. Furthermore toxicological tests, in which the ligands compete with biological binding sites for bound Hg2* demonstrate that the stabilities of the complexes of these S bearing chelates with mercury are more stable than the Hg2÷-EDTA complex [8]. On this basis it is probable that the equilibrium constants must be greater than 1024 and this renders improbable the lower values reported previously. Furthermore, it should be noted that our values fall in the rough order of magnitude range of the value reported previously for 2,3-dimercaptopropanesulfonate[2]. The greater stability of the complexes with BAL and 2,3-dimercaptopropanesulfonate is what one would expect from the chemical structures of these ligands, because they have two -SH groups and it is generally accepted that the binding of the mercury(II) to -SH groups is firmer than its binding to N or O donor groups[18], n.L-penicillamine, and N-Acetyl D,L-penicillamine, each bearing but a single -SH group binds the mercury(II) less firmly by a significant factor. The reduced stability of the N-Acetyl D,L-penicillamine complex in comparison with the complex of O,L-penicillamine is what one would expect from the interference of the acetyl group with the normal donor properties of the amine group. Regardless of other considerations, the values are comparable among themselves and allow some conclusions to be drawn concerning their relative ability to bind mercury in animals and anticipated utility in cases of acute mercury toxicity. If the only parameter of importance were the stability constant of the Hg 2÷ chelate complex one would expect that the order of their effectiveness as antidotes in acute mercury poisoning would fall in the sequence: BAL > DMPS > DPA > NAPA. In fact this is only part of the picture. The appreciably greater toxicity of BAL and the frequent occurrence of side effects[19] as well as its demonstrated ability to carry methyl mercury to the brain [20], make it less attractive than the three remaining species. One may finally ask why the value for BAL reported here is so much greater than the one reported earlier from this laboratory[7]. This presumably lies in the assumptions made about the species present in the initial part of the titration curve[21]. In conclusion one may say that the evidence is overwhelmingly against these chelating agents giving 1 : 1 mercury chelates with stability constants less than 102o. Whether they are in fact as large as we report them here can only be determined when some independent method can be developed to check the functioning of the mercury electrode at the limits of its concentration range. Acknowledgements--This work was carried out under the auspices of the Center in EnvironmentalToxicology,Departmentof Biochemistry,School of Medicine,VanderbiltUniversity School of Medicine,Nashville,Tenn. 37235 and supported by Grant ES 01018-04A1-TOX of the National Institute of Environmental

102

J.S. CASAS and M. M. JONES

Health Sciences. We also wish to thank the firm of E. Heyl and Co., Goerza/lee 253, 1000 Berlin 37, West Germany for their sodium 2,3-dimercaptosulfonate [DMPS, Dimaval (R)] used in this work. REFERENCES 1. A. T. Pilipenko and O. P. Ryabushko, Ukr. Khim. Zhur. 32, 622 (1966). 2. I. E. Okonoshnikova, L. G. Egorova, V. I. Nirenberg and I. Ya. Postovskii, Khimiko-Farmasevticheskii Zharnal 4(1), 21 (1970). 3. R. L. Coates and M. M. Jones, J. Inorg. Nacl. Chem. 39, 677 (1978). 4. W. Stricks and I. M. Kolthoff, J. Am. Chem. Soc. 75, 5673 (1953); (a) W. F. Van der Linden, Anal. Chim. Acta 68, 143 (1973). 5. G. R. Lenz and A. E. Martell, Biochemistry 3, 745 (1964). 6. E. J. Kuchinskas and Y. Rosen, Arch. Biochem. Biophys. 97, 370 (1962); (a) Y. Sugiura, A. Yokoyama and H. Tanaka, Chem. Pharm. Ball. 15, 693 (1970). 7. D. A. Doornbos and J. S. Fabcr, Pharm. Weeblad 99(12), 289 (1964). 8. V. Nigrovic, Arzneimittelforschang 13, 787 (1963).

9. R. W. Schmid and C. N. Reilley,J. Am. Chem. Soc. 78, 5513 (1956). 10. C. N. Reilley and R. W. Schmid, Anal. Chem. 30, 947 (1956). I I. T. S. West, Complexometry with EDTA and Related Reagents, p. 195-196. BDH Chemical Ltd., Poole, Dorset (1969). 12. A.AibertandE.P.Serjeant, lonizationConstantsof Acidsand Bases, pp. 16, 42. Methuen, London 0962). 13. P. J. Antikainen and V. M. K. Rosi, Soumen Kemistilehti B36, 132 0963). 14. P. J. Antikainen and K. Tevanen, Saomen Kemistilehti B35, 224 (1962). 15. D. A. Doornbos and J. S. Faber, Pharm. Weekblad 99, 289 (1~). 16. K. H. Schreder, Acta Chem. Scand. 20, 881 (1966). 17. D. D. Perrin and I. G. Sayce, J. Chem. Soc. A. 53 (1968). 18. S. Ahrland, J. Chatt and N. R. Davies, Quarterly Reviews 12, 265 (1958). 19. w. T. Longcope and J. A. Luetscher, Ann. Intern. Med. 31, 546-547 (1949). 20. M. Berlin and T. Lewander, Acta Pharmac. Tox. 22, I-7 (1965). 21. N. Weinstein, privatecommunication (1977).