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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
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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).