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Subscriber access provided by RICE UNIV Mercury(II) Complexes of Imidazole and Histidine Philip Brooks, and Norman Davidson J. Am. Chem. Soc., 1960, 82 (9), 2118-2123 • DOI: 10.1021/ja01494a008 Downloaded from http://pubs.acs.org on December 9, 2008 More About This Article The permalink http://dx.doi.org/10.1021/ja01494a008 provides access to: • • Links to articles and content related to this article Copyright permission to reproduce figures and/or text from this article Journal of the American Chemical Society is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 PHILIPBROOKSAND NORMAN DAVIDSON 211s The absorption band of medium intensity a t 1308 cm.-' in the reagent, which is also probably due to C-0 vibrations, is affected to a small extent in the case of Pb(I1) and Cd(I1) chelates but is shifted to a marked degree in the chelates of the transition metals. This implies that the C-0 group vibrations in the metal chelates of 4-hydroxybenzothiazole are influenced to a slight extent by the mass of the metal in the chelate ring. The interaction of the C-0 group with the n-bonds in the chelate ring as well as with the delectrons of the metal atoms play a much more important part in determining the extent of the shifts that are observed in C-0 group vibrations as we go from the reagent to the metal [COh"IR I U U r I O N S O . 2$YG F R O M THE Vol. 82 chelates. It was found that if the frequency of the band that appears between 1310 and 13% an.-' for the transition metal chelates of 4-hydroxybenzothiazole was plotted against the stability constant (log &) of the chelate, a linear relationship was obtained. This type of relationship has been found in a number of metal c ~ m p l e x e s ' ~and ~ ' ~confirms the assignments made for the C-0 vibrations in the 4-hydroxybenzothiazole chelates. Acknowledgment.-The authors are grateful for a grant from the U. S. Atomic Energy Commission in support of this work. (14) H. F. Holtzclaw, Jr., and J. P. Collman, THISJOURNAL,79, 3318 (1957). PITTSBURGH 13, PENNWLVANIA GATESAXD CRELLIN LABORATORIES O F CHEMISTRY, PASADENA, CALIFORSIA] Mercury(I1) Complexes of Imidazole and Histidine BY PHILIP BROOKSAND NORMAN DAVIDSON RECEIVED AUGUST 31, 1959 The association constants of Hg*I with imidazole and L-histidine have been calculated From potential measurements of a I-Ig, Hg" electrode in solutions containing the complexing agents. For imidazole, the main result is H g + + 2C3114N: Hg(C3H&il)zf+ = 1016.7 A{-2 + solid substance, H g ( C 3 1 i a S ~ ) C 1 0 ~ ~ Hprecipitates :0, from solutions a t p H (C3H3SzHg-)n+". For histidine, the main results are: H g + + 2LHgL2 H g + + L HL HgL(HL)+ K = 1021.2 A I - 2 K = 1018.4 AJ-Z I\ + + > 4. + I t probably contains an infinite ch:iin: Hg++ + 2HL A' Hg(HL)z++ 1015.0 A[-2 where L- = (C3H3N2)CHzCH( SH?)CO2- and I-IL = (C3H3N2)CHCH(NHa+)CO:-. Because of the tendency of H g + + to form two bonds in a linear configcration, chelate formation contributes only slightly to the ligand binding by Hg"+. There is also evidence for the complex HgL+, with a formation constant of M-' (from L-). I niidazole and histidine complexes with metal ions are of intrinsic interest as a part of the chemis- try of complex ions, as well as being of possible biological significance. Complex ion formation with a nuniber of metal ions has been studied.' There is evidence that the imidazole ring is an important binding site for metal ions in serum : ~ l b u r n i n . ~In~ ~ general, mercury(I1) has the greatest binding power for nitrogen ligands of any of the metal ions; the present investigation is concerned with the mercury(I1) complexes of imidazole and histidine. Experimental Imidazole, from the Eastman Rodak Co. (m.p. 88-90', Eastman grade) and the Aldrich Chemical Co., and L-histidine, cfp. grade (99.98y0 pure by nitrogen analysis) from the California Corporation for Biochemical Research were dissolved in doubly distilled water. Concentrations calculated from PH titrations and from the weight of solute taken agreed within 1yo. Imidazolium perchlorate solutions were prepared by adding standard perchloric acid t o imidazole solutions. Mercuric perchlorate solutions were prepared by dissolving reagent grade mercuric oxide in a known excess of 9 F perchloric acid and diluting. In all experiments sodium perchlorate was used to keep the ionic strength constant a t 0.15 M . (1) For a general review, see J. Bjerrum. G. Schwarzenbach, L. S i l k , "Stability Constants, P a r t I." Special Publication No. 6, T h e Chemical Society, T.ondon, 1957, Tables 26 and 207; see also t h e speciiic references in Table 111. (2) F. R. N. Gurd and D. S. Goodman, THISJ O U R N A L , 74, 070 (1932). ( 3 ) C . ' I a i i f ~ ~ ridb,d . , 7 4 , '21 I ( I X d ) . The principal experimental data were the potentials of a mercury, mercury(I1) electrode as a function of ligand concentration and pH. The cell, in a water bath a t 27.0 & 0.1', contained a J-type mercury electrode in which the mcrcury surface could be renewed by overflow, a 1.5 F S a S O a salt bridge connecting t o a saturated calomel electrode and provision for titrating in reagents and for maintaining a nitrogen atmosphere. Potential differences were measured with a Leeds and Northrup IC-2 potentiometer and a 0.01 p amp ./mm. galvanometer; p H measurements were made with a Beckman model GS p H meter. In every case ample time was allowed for the system t o attain equilibrium before final measurements were made. Results Imidazole Complexes.-As anticipated, the binding of mercury by imidazole is so great that, under practical conditions, the equilibrium Hg++ + nIniH+ eHg(Im),++ + %€I+ (1) lies far to the right except for rather acid solutions with p H 5 2 . Under these circumstances, the pH titration method for the study of complex ion formation is difficult. We have therefore used a potentiometric method. Insoluble mercury-imidazole compounds precipitate a t pH's > 4, so measurements were made in the pH 2-4 range. The predominant imidazole species is the imidazolium ion, ImH+. Equation 1 then represents the over-all chemical rexction. Tlic Nernst law e m bc writteti MERCURY (11) COMPLEXES OF IMIDAZOLE AND HISTIDINE May 5, 1960 For Eo', the potential of the hypothetical Hg, 1.00 M Hg++ electrode a t p = 0.15 M vs. the see. (saturated calomal electrode), we take - 0.607 v. It is inconvenient to measure this quantity because of the Hg, Hg++, Hgz++ equilibrium. Our choice is based on the values Eo (Hg++ vs. Hz) = -0.841 in 1 F HC104,4 Eo (Hg++ DS. H2) = - 0.854 v. ( p = O).6 It is t o be noted that the quantities E that we quote are oxidation-reduction potentials using the American convention; a Hg, Hg++ electrode is electrically positive with respect to s.c.e. For the equilibria Hg++ + ImH+ HgIm++ +H+ = Hg++ [ H + ][HgIm++] [Hg++][ImH+] + 2ImH+ J_ HgImp++ + 2 H + P2' we have Equation 2 can then be written as E = Eo' RT - -In 2T [ZHg"] + Plots of the quantity (25/RT)E + In (I: Hg") os. In [ImH+]/[H+J were linear with a slope of 1.95-2.00. This shows that the dominant term Kl'[ImH+]/[H+] B?'. in the power series 1 ( [ImH+]/[H+])2 is the quadratic term and the principal complex formed is Hg(Im)z++. Representative data analyzed in terms of the simple equation + + E = Eo' - 0.0298log [Z Hg"] TABLEI POTENTIAL MEASUREMENTS OF Hg", I m H + SOLUTIOSS 10-4 10-4 10-4 10-4 1.05 5.49 1.82 3.80 6.02 [Z ImH+] [ImH+] F M X 10- 0.0963 0.0055 X 10-9 .0960 ,0952 X 10-10 ,0954 .0946 X 10-1' ,0954 ,0946 X 10-1' ,0954 ,0946 2.53 2.46 2.54 2.58 2.50 X X X X 10-4 10-4 10-4 10-4 2.52 6.74 1.81 7.71 X X X X 10-7 10' 10-8 10- ,0160 ,0317 ,0642 .0972 2.55 2.54 2.53 2.51 X X X X 10-4 10-4 10-8 10-8 7.76 2.75 5.49 4.46 X X X X 10-8 10-8 1010- ,0147 .0696 .a692 .Of384 log 81' = [X Hgf*]b Fc 2 . 0 3 3 . 8 5 X 10-4 -E,a v. 0.369 ,361 ,317 ,267 .243 2.21 2.91 3.77 4.17 3.84 3.82 3.82 3.82 X X X X ,411 ,394 ,377 .366 2.12 2.12 2.10 2.10 3.87 4.08 4.10 3.88 2.58 5.11 2.60 5.26 2.59 1.01 ,388 2 . 6 0 6 . 4 4 [H$++] PH ,395 .352 ,361 Av. [ H+]2[HgIrn2+'] = [Hg++][ImH+]* 2119 .0153 .0309 ,0634 ,0965 log 021 ,0138 2 . 3 8 .0686 2 . 4 0 .0674 2 . 4 2 ,0567 2 . 4 6 2.50 f 0.06 As already emphasized, these are conventional oxidation-reduction potentials vs. s.c.e.; the electrode was electrically positive with respect to s.c.e. [ZHgII] is the total (formal) concentration of Hg", corrected for Hg2++. F and M arr formula weights per liter and moles per liter, respectively. a tion and imidazole concentration shows that the principal species is Hg(Im)2++. On changing concentrations of reagents, the potentials usually were established to f 0.5 mv. within a few minutes, although occasionally the potentials drifted for 10-30 minutes before becoming constant. In general a good way to investigate the possible influence of the first complex, HgIm++ and the term K,'[ImH+]/[Hf] is to start from the equation + Plots of the function on the left of equation 9 vs. [ImH+]/ [H+]gave fairly good straight lines; how[ImH+] (7) ever, it was not possible to obtain points close are presented in Table I. In all cases from the enough to the origin to result in a reliable extrapoactual potential, we know that [ZHg"]/ [Hg++] 2 lation for K1'. The difficulty is essentially this. 400. The equilibrium constant for the reaction To emphasize the contribution to the binding by HgTm++, one should work a t low [ImH+]/[H+] Hg Hg++ Hgz++ (8) ratios. This decreases the ratio [I: Hg"]/[Hg++]; has been measured as 130.s We have assumed the amount of Hgsff according to equation 8 then that the system was a t equilibrium with respect becomes large relative to [I: Hg"]. Under these to this reaction, although we do not know that such circunlstances, the corrections involved in analyzing is the case. With this assumption, [HgIm2++! = the data are too great to permit reliable interpreta[ 2 Hg"] = [Z Hq] - [Hg,++] where [I: Hg] is tion. There was no evidence for the formation the total amount of mercury added as Hg(I1) in of Hg(Im)3++or Hg(Im)4++ even a t [TmH+] = the solution and [I: Hg"] is the total amount of O.lOMandpH4. Hg", !.e., after correction for the mercurous ion The pKa of imidazolium a t 27" and p = 0.15 14 formed. Then [ImH+] = TI: IrnH+] - 2[Hg- was measured by a pH titration to be 7.12. This Imz++]. The corrections for the amount of Hgz++ can be compared with Edsall's value7 of 7.11 a t present and for the amount of imidazole tied up as 23O, ,u = 0.16 M , and Tanford and Wagner's HgTml++ are small in all cases and negligible for valuesof 7.12 a t 25", p = 0.15 M . Since high ratios [ImH+]/[Z Hg], so errors in the corHg++ 2IrnH+ = HgImz++ + 2H+, rections are not very important. PZ' = 102." (10) The constancy of the calculated values of p2' in Table I with varying PH, mercury(I1) concentra- H g + + + 2Im = HgImp++, 0.0596 log {B2' I + + (4) E. H. Swift, "Introductory Quantitative Analysis," PrenticeHall, Inc., New York, N. Y.,1950, p. 509. (5) W. bf. Latimer, "Oxidation Potentials," Prentice-Hall, Inc., Englewood Cliffs. N . J., 1956, p. 179. (6) L. G. Sillen, A. Jonsson and I. Qvarfort, Acta Chem. Scand., 1, 461 (1947). (7) J. T. Edsall, G. Felsenfeld, D. S. Goodman and F. R . N . Gurd, THISJOURNAL, 76, 3054 (1954). (8) C. Tanford and M. L. Wagner, ibid., 76, 434 (1953). PHILIPBROOKSAND NORMAN DAVIDSON 2120 1701. 82 This corresponds approximately to the formula Hg(C104)(C3H3N2)*H20a L-Histidine.-The ionization equilibria of histidine are HpL+ ~: -100O HL K, + H+ --f fi,.- 1,- + 2H' (12) where I t > w -200 -400 2.0 ' I 3.0 H I ' ' 4.0 ' ' ' I 5.0 PH 6.0 " 7.0 ' 8.0 I ' 3.0 Fig. 1.-Potential it$. PH for fixed ZHg" and histidine concentrations. These concentrations, [2HgTr] = 3.80 X F and [ZL] = 0.095 F, actually decreased slightly during the titration due to dilution; suitable corrections were made. A similar straight line, displaced as expected by the Nernst law, was obtained for [ZHgII] = 3.4 X F, [ZL] = 0.011 F. The solid line is calculated from equation 15, and the values log pzl) = 2.98, log 8 2 1 = 9.18, log /3zr = 14.99. Insoluble Hg (C3H3N2) (C104).H20.-Both the imidazole and histidine complexes with mercury become insoluble as the fiH is raised. The imidazole mercury complex precipitates copiously a t pH's greater than 4.5, but the histidine complex is more soluble, solutions a t very alkaline pH's (ca. 12-13) being just slightly cloudy. To ca. 150 ml. of solution containing 10 mmoles of Ha", 45 mmoles of Im, 35 mmoles of H f , was added about 13 mmoles of NaOH. About half the mercury precipitated as a white precipitate and t.he p H was 6.1. The precipitate was dried in a vacuum desiccator over Cas04 for several days and then analyzed. The analysis for mercury was carried out in 6 N HCW4 with NaSCN by a Volhard titration. About 60% of the NaSCN was added prior to the titration to dissolve the complex. The analyses for carbon, hydrogen, nitrogen and chlorine were done commercially by Dr. A. Elek. Element yo Present Hg 51.12 c1 9.41 N 7.63 H 1.21 C 9.75 0 (by difference) 20.88 Relative molar amouut 1.00 1.04 2.13 4.71 3.18 5.08 H Our measurements give pKa = 6.08, PKa' = 8.12; in good agreement with other results (AlbertJg 6.08, 9.20, 20°, I.L = 0.01 M ; Li and Manning.l0 6.05, 9.12,25',p = 0.15M). In a solution with constant ZHgl' concentration and a constant, rather large [ZL] (where [ZL] = [HzL+] [HL] [L-1) the potential was observed to be a linear function of pH over the p H range 2.5 to 9, with a slope of 0.060 v. per p H unit, as displayed in Fig. I. The linear relation is rather surprising since the transformation H2L+ F? H L H + has pKa = 6 0. I t is necessary to suppose therefore that the complexes also can be protonated and we propose the following equilibria. It is convenient for the analysis to write all equilibria terms of the species HL. + + + + H L If HgL+ + 1% H g + + + HL HgHL-+ H g + + + 2HL I_ HgLz + 2H+ H g f C + 2HL I_ HgL(HL)+ + H + Hg++ {9,0 (1';) $ 1 0.0 1391 4-2HL Hg(HL)s+ B.2 The measurements used for qudntitative interpretation were in the p H range 9-3, so that it is necessary to consider the three histidine species, L-, H L and H2Lf (We can neglect HaL++,which has a PK, of 1 8).g lye write 21, for the total uncomplexed histidine Hgf+ [ZL] = [L-1 + IHLl + [H*L+] understanding that any significant corrections for the amount complexed have been made (if necessary, by a successive approximation procedure). Then, from equation 12 (9) A. hlbert, Bioclrem. J . , S O , GSO (19.52). (IO) N. C . Li and R. A . l l a n n i n g . Tms J O U K X A L , 77, 5225 (1025). MERCURY (11) COMPLEXES OF IMIDAZOLE AND HISTIDIKE May 5, 1960 2121 TABLE I1 SLOPE AND INTERCEPT OF EQUATION 15b FOR EXPERIMENTS AT FIXED pH PH Intercept' Slope* 5.00 6.08 5.20 7.00 6.20 4.93(& 1)X lo9 1.30(* 1.3) X 1Olo 6.90(&2) X 10" 4.87(&3) X 1011 2.42(& 1 . 5 ) X 1013 5.46(*0.1)X IO1* 1 . 8 4 ( f 0 . 1 ) X 1 0 1 * 1.42(1O.2)X1O16 2.46(&OO.2)X10'5 9.75(&0.1)X 10l6 Intercept = ' q ( p l o IH 1 + Pll[H+]). bSlope = + PzI[H+]4-P22[H+l2). where the function f(H+) is defined by the second ments in the p H range 8-3; the data a t 9 and 2.5 equality. are also predicted fairly satisfactorily.) The total mercury(I1) concentration is given by the relation: [ZHgT1]= [Hg++] [HgL+] [HgHL++] [HgLzl [HgL(HL)+] [Hg4.0 k (HL)z++]. [HLl = [ZL]f(H+) from equation 14. The equilibria 13 then permit us to express the quantities [HgL+], [Hg(HL)z++], etc., in terms of [Hg++], [H+] and [ZL]; for example, o/o' [HgL(HL)+1 = Pzi [&?+I [~LIz(j"(H+))z/[H+l. By substitution of these relations into the equation for [ZHg"] we obtain + + + + + c where and 6 2 are implicitly defined by comparison of (15a) and (15b). The term on the left of (15) was plotted os. [ZL] for experiments a t several fixed values of the PH with varying [EL]. One of the best plots is shown in Fig. 2. The values of the intercepts and slopes are collected in Table 11. Examination of the plots, such as Fig. 2, strongly suggssts that there is a significant contribution by the 81 term, but the intercept cannot be estimated accurately, as we have tried to suggest by the estimated uncertainties in the table. The two best intercepts fo_rp H 7 and 5 give 81 = 2.6 X lo6 a t pH 7.00 and j31 = 6.5 X l o e a t pH 5.00. According to our formulation, 81 should increase with decreasing pH. It is quite possible that there is another complex such as HgL(OH), but we believe the data are not sufficiently good to warrant such a conclusion. We provisionally conclude only that there is a complex with one histidine per mercury, HgL+, that probably Plo = lo6 and that in the 5 to 7 p H range there is not much of a contribution_by Hg(HL) ++. A plot of 8 2 vs. pH is shown in Fig. 3, I n addition to the results in Table 11, we have used the data of Fig. 1 where the fiH was varied a t a constant high Z L concentration; a t this high concentration, only the 8 2 term is important. The solid lines in Figs. 1 and 3 are calculated from the values p20 = 2.98 log Pzi(M-') 9.18 log P Z Z ( M - ~ ) 14.99 log The good agreement justifies the interpretation given. Incidentally, since different EHg" concentrations were used in different experiments, the fact that the complexes are monomeric as regards mercury has been demonstrated. (The constants &, ~ ? Z Z were actually derived from the measure- 1.0 c 1 I v d o i 1 I 1 2 3 4 5 6 7 [PL] x I03 Fig. 2.-Plot pH = of eq. l5b for the determination of P1 and pz; 5.00 & 0.01, [ZHg"] = 1.4 X lo-* F. As already remarked, the fact that the variation of potential with p H is quite linear with a slope of RT/S in the p H 3-8 range (Fig. 1) was not to be expected in general. Reference to equation 15a shows that this linearity requires that the function log {[fW+)I2(Pzo Pzi[H+I Pzi[H+I')J should be constant. For pH<8 where the Ka' term can be neglected, this is equivalent to the constancy of + + which requires that (PzoIPzI) = K J 2 ; and ( P 2 1 , l P ~ z ) = 2Ka. With the chosen values of the 0's we have Discussion The binding constants analogous to P2 of several nitrogen bases for various metal ions are displayed in Table 111. Several significant trends are evident in Table 111. In the first place, as has been previously n 0 t e d ~ ~ 8 Jthe ~ ~ ' binding ~ Constants of imidazole and ammonia for complexing metals are approximately equal but the pKa of imidazolium is less (12) N. C.Li, T.L.Chu, C. T.Fuji and J. bf. White, THISJOURNAL, 77. 859 (1955). (13) R. B. Martin and J. T. Edsall, ibid., 80, 5033 (1958). PHILIPBROOKS AND NORMAN DAVIDSON 2122 Vol. 82 The increase in binding from imidazole to the negative histidine anion (L-) is very marked for the metals Zn++, Cd++, Ni++ and Cu++, whereas it is rather small for Hg++. In fact the ratios p&(L-)/pp2(Im) for these five metals are 2.8, 2.3, 2.7, 2.6 and 1.3, respectively. For the first four metals, chelate formation with the probable structuresQ H undoubtedly contributes very significantly to the binding. Because of the tendency t o form only two strong linear bonds, chelate formation contributes only slightly to the ligand binding for Hg++. This point can be illustrated further by these equilibria which can be calculated from data already given. H g + + -t 2L- - 0 9 8 7 6 5 4 3 2 PH Fig. 3.-Log PZZJS. PH; the circles are calculated directly from the experiments using l5b; the solid line curve is calculated from 15a and the values log ,920 = 2.98, log 1921 = 9.18, log /j22 = 14.99. than that of ammonia. Thus, relative to ammonia imidazole shows a preference for binding transition metal ions rather than protons. TABLE I11 FORMATION CONSTANTS OF COMPLEX IONS" Imidazole (CaHdN:) li.letal/ligand Histidine (L-) "11 H+ 14.28 18.48 18.24 Zn++ 4.17' 4.43 11.78l 4.901 4.47 Cd++ 11.101 Hg++ 16.74 17.5 21.22 4.79 15.8411 5.951 Xi++ cu++ 7.15* 7.33 18.601' 0 The values listed are Be's; ;.e., equilibrium constants for Culmz++. To compare with the reaction Cu++ 21m the 82 values, we list twice the pK, of the ligand, for example, the equilibrium constant of the reaction 2Hf 2NH1 2XH4+, etc. + + e As is well known, Hg++ ion typically binds two ligands very strongly in a linear configuration and shows only a small tendency to form a third and a fourth coordinate bond. The other metals in Table 111 generally have coordination numbers of 4 or 6, with, typically, only a small gradual decrease in the intrinsic binding constants MLn-I +L MLn K - WLnl [MLn-II [LI (where L here is any monodentate ligand and not histidine) as n increases from 1 t o 4 or 6. (11) N. C. Li, Br. E. Doody and J. M. White, 5859 (1957). K HgL + + HL H g + + + 2HL J _ Hg(HL)*++ H g + + LHgL(HL)+ = 1021.22 M - 2 (16a) K = 10l8.m Al-'(16b) K = 1014.9@ M-' ( 1 6 ~ ) The binding constant of HL by Hg++, presumably via the imidazole group, is less than the binding constant of imidazole for Hg++ by about lo2, which is just about what would be predicted from the differences in proton binding. The small increase in binding constants in the series of reactions 16c, 16b and 16a may be partly due t o chelate formation; it is also partly due to charge effects. It is instructive t o write out the acid ionization equilibria K = HL/H++LHg(HL)zC+ Hg(HL)L+ Hg(HL)L+ HgLz H+ + + H+ K 10-8.12 = 10-b.81 I( = 10-6.20 In each case, the protonated amino group is ionizing; coordination of Hg++ to the imidazole group increases the ionization constant by a factor of only lo3; a much larger factor would be expected if chelation by L- were very significant. In addition to its basic properties, iniidazole is ;&o a weak acid. The pK for this reaction has been estimated as 14.2.14 We presume that the insoluble mercury cornpound, Hg(CsI13N~)C104~H~0, contains the linear polymer T H I S JOORNAL, 79, (14) H. Walba and R. W, Isensee, ;bid. 77, 5488 (1955). H 2123 MERCURY(I AND 11) COhIPLEXES May 5, lOGO H H with one positive charge per polymer unit. A similar subrkance Hg(C3H3N2)C1 has been reported. 15 The structure proposed above is similar to that which occurs in H ~ ( N H ~ ) B and ~ ~HgB (NHp)Cl.l7 (15) K . Birttcher. Chcm. Zcnrralblolf, 102, 11, 2757 (1931). (18) L. Nijsseri and W. N . Lipscomb, Acla Cryst., 6, 604 (1952). (17) W.N. Lipscomb, ibid., 4, 266 (1951). [COSTRIBUTION No. 2522 FROM THE H H \’ H’\H Acknowledgments*-Mr* Tetsuo designed and developed the cell and potentiometric techniques which we have used. We have profited also from his advice and criticism regarding other features of this research. We are grateful to the AEC for support of this work via t h i contract AT(11-1)-188* PASADENA, CAI.IFORNIA GATESA N D CRELLINLABORATORIES OF CHEMISTRY, CALIFORNIA INSTITUTE OF TECHNOLOGY, PASADENA, CALIFORNIA] Complexes of Mercury(1) with Polyphosphate and Dicarboxylate Anions and Mercury (11) Pyrophosphate Complexes1 BY TETSUO YAMANE AND NORMAN DAVIDSON RECEIVED OCTOBER 30, 1959 It has been discovered that mercurous mercury forms complexes with pyrophosphate, tripolyphosphate, oxalate, adimethylmalonate and succinate. These complexes are stable toward disproportionation to mercury( 11) complexes and mercury, If L-9 is the anion, the principal complexes are Hg,L2-2q+2 and Hg?L(OH)-q+‘. The formation constants determined from the potential of a mercury-mercurous electrode in ligand solutions are: Hg2(P~O~)r-’,(2.4 f 0.6)101*M - * ; Hg,(OH)P201-3, (4.4 f 0.6)1015 M-2; Hgz(PsOiO)z-*, (1.7 f 0.3)1011. Hg,(OH)P3010-‘, (1.0 f 0.2)1015; H ~ J ( C ~ O ~ ) Z - : , (9.5 f 0.2)106; Hg2(OH)CzOa-l, (1.1 f 0.2)1013. H ~ [ ( C H ~ ) Z C ( C O L ) ~ (3.3 ] Z -f ~ , 0.6)10’; Hgz(OH)[(CH!)zC(COz)21 - , (3.8 f 0.5)lOl3; H g ~ [ ( C H 2 ) 2 ( C 0 ~ ) ~ ](1.9 ~ - 2 ,i 0.5)107; Hgz(0H) [(CH2)2(C02)2]-l, (2.8 f 0.6)1013. (The ionic strength The mercurous compounds have a was 0.75 M (NaNO?), except for oxalate and succinate, where I t was 2.5 M ( Nal’;Oa).) characteristic ultravlolet spectrum. Theory and experiment agree that mercurous complexes of ligands (such as NH3 and CN-) which form strong covalent bonds to mercury are unstable toward disproportionation to give mercuric complexes but “ionic” chelating ligands can form stable mercurous complexes. The mercury( 11) pyrophosphate complex was studied from the potential of a Pt electrode in Hg2, Hg”, pyrophosphate solutions a t PH 7-10. The principal species is Hg(0H)(P201)-3,with a formation constant of (2.8 f 0.6)1017M - 2 . Sillen and co-workers have suggested, on the basis of potentiometric evidence, that there are weak complexes of Hgz++ formed by nitrate, sulfate and perchlorate anions, with formation constants of: 2.5 M-’ (Hg,N03+), 0.5 (Hgz(N03)2), 20 M-’ (HgzS04), 250 (Hgz(S0&-2) and 0.9 The equilibrium is readily reversible. When a M-’ (HgzC104+)..2~5This presumably is mainly complexing ligand is added to a mercurous solution, ion-pair association. It is also possible that the the usual reaction that occurs is disproportionation assumption of constant activity coefficients a t of the mercurous ion to give elementary mercury constant ionic strength is not sufficiently reliable and a complexed mercuric ion. This occurs, for to enable one to identify such weak complexes example, with the complexing ligands, CN- with certainty by potentiometric experiments. I t is due to the relatively greater staand ”3. However some time ago, Stromeyers and then bility of the mercuric complexes. The same situa- Brand7 reported that when sodium pyrophosphate tion occurs for many insoluble compounds. Thus, solution is added to a mercurous solution, a white mercurous ion is unstable in basic solutions and precipitate forms and then redissolves in excess of in the presence of sulfide ion. Compounds such the reagent, which suggests the formation of a as “mercurous sulfide” or “mercurous oxide” strong, stable complex. reported in the past have been shown to be a mixWe have confirmed and extended these observature of mercury and the corresponding mercuric tions and have now found that mercurous ion forms comp~und.~,~ The general impression conveyed by textbooks stable complexes with pyrophosphate (Py-”, and by the chemical literature is that there are tripolyphosphate (TP-~), oxalate OX-^), a-dino known stable complexes of mercurous ion. (5) G.Infeldt and L. G. SillCn. Suensk kern. Tidskr., 6 8 , 104 (1946). The equilibrium constant €or the formation of mercurous ion from elementary mercury and mercuric ion is 130 in 0.5 M NaC104.2 (1) Presented at the 136th National Meeting of the American Chemical Society, Atlantic City, N . J., September, 1959. (2) S. Hietanen and L. G . Sill&, Arkiu Kcmi. 10, 103 (1956). (3) R. Fricke and P. Ackermann, 2. anorg. Chcm., a l l , 233 (1933). (4) W.N . Lipscomb, THISJOURNAL, 1 3 , 1015 (1951). (6) F. Stromeyer, Schwciggcr’s Journal, 68, 130 (1830); as reported in J. W. Mellor “A Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. IV, Longman, Greens & Co., Londoii, 1923, p. 1003. (7) Brand, 2. enol. Chcm., 28, 592 (1889).