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Electrochemical generation of mono- and dinuclear mercuric cryptates
J. Electromul Chem., 202 (1986) 191-201 Elsevter Sequoia S.A., Lausanne Printed
191 m The Netherlands
ELECTROCHEMICAL GENERATION MERCURIC CRYPTATES
C. BOUDON
l, J.P. GISSELBRECHT
and J.M. LEHN
l
OF MONO-
AND DINUCLEAR
l, M. GROSS l. F. KOTZYBA-HIBERT
l
*
*
DtQxwtement de Chrmre, LJmc~ers~t~LAWS Pasteur. 4, Rue Blarse Pa.ccul, 67000 Strushourg (Fruncr) (Received
26th September
1985. m revised form 23rd December
1985)
ABSTRACT
The electrochemical oxidation of mercury electrodes was studted m orgamc solvents m the presence of macrocychc ligands. It resulted m the generation of very stable mono- and dinuclear mercuric complexes. The redox couple based on the dimercuric complex with the hgand T(DPO), exhibits the unprecedented property of transferring almost stmultaneously four electrons at + 0.2 V (~5 SCE).
INTRODUCTION
Fifty years ago, Ravenda [l] reported that the electrochemical oxidation of mercury was facilitated in the presence of complexing anions. After the pioneering work of Kolthoff and Miller [2]. several authors studied extensively the electrochemtcal oxidation of mercury, thermodynamically assisted by complexing ligands, and, in some instances, this specific reaction was applied to indirect titrations of cations [3-91. More recently, it was observed that polyazamacrocycles, and among them the diazamacrobicycles, also facilitate the anodic oxidation of mercury [lO.ll]. allowing either direct titration of the macrocycles in the solution, indirect titration of metal ions [ll], or indirect titration of other mono- and dications [12]. Thus, considering the prominent interest in a better understanding of molecular receptor-substrate interactions and in identifying the critical parameters in these interactions, analysis of the characteristic parameters involved in the electrogeneration of mercuric complexes with the mono- and ditopic ligands shown in Fig. 1 was attempted. The corresponding results are reported and discussed in the present paper. These results reveal an actual control of the cooperativity between the two
* Laboratoire ** Laboratoue
d’Electrochimte et Chimie Physique du Corps Sohde. U.A. au C.N.R.S. de Chimie Orgamque Physique. U A au C.N.R.S No. 422.
0012-0728/86/$03.50
‘c’ 1986 Elsevter Sequoia
S.A
No. 405.
192
cN204
2 Me
3
[T(P-‘d]
Fig. 1. The ligands
[Tb-P’&]
used in the present
stud).
mercuric ions through the length ditopic ligands (Fig. 1).
and the flexibility
of the lateral
chains
in the
EXPERIMENTAL
The electrochemical measurements were carried out in dimethylformamide (DMF) or in a mixed solvent (l/3 1,2-dichloroethane + 2/3 propylene carbonate (PC)). The solvents were purified before use as described previously [13-151. The ligands studied (Fig. 1) are poorly soluble in PC and very soluble in 1,2-C>H,Cl,. The supporting electrolyte was either tetra-n-hexylammonium perchlorate (THAP. from Fluka) or tetra-n-butylammonium perchlorate (TBAP, from Eastman Kodak). and both were recrystallized before use [16]. Electrochemical measurements were carried out with a multipurpose PRG4 device (SOLEA, Tacussel), in a classical three-electrode cell. on a dropping mercury electrode as the working electrode (m = 0.178 mg/s. h = 40 cm Hg). The auxiliary electrode was a platinum wire. The reference electrode was a calomel electrode in an aqueous KCl-saturated solution, electrically connected to the electrolysis cell by a bridge filled with the same organic solution as that introduced into the cell (solvent + supporting electrolyte). It was repeatedly checked that, with this electrolytic junction, the junction potential remained constant during the series of experi-
193
ments. The followmg electrochemical methods were used in the present study: normal pulse polarography, dc polarography with controlled drop time, cyclic voltammetry and controlled potential coulometry. In the latter experiments (coulometry), the working electrode was a mercury pool electrode (area = 15 cm’) in a cell containing about 10 cm3 of solution (the exact volume was measured in each experiment). The syntheses of the ligands used (Fig. 1) have been described previously: [N,O, 2Me] and [N,O, 3Me] [17]. [T(o-Ph),] [18]. [T(p-Ph),] [19], [T(BPh)?] [18] and
[-UDPO)~l [181. RESULTS
Anodic oxidatron of mercury electrodes in the presence of macroqvclic Iigands When no complexing ligand was present in the solution, the electrochemical oxidation of the mercury electrode occurred from +0.80 V (vs. SCE) in the solvent (2/3 PC + l/3 C,H,C12) containing also 0.1 M THAP or TBAP. At this potential. +0.80 V (vs. SCE), i was 0.18 + 0.02 PA cm-‘. In the solvent DMF + 0.1 M THAP or TBAP, mercury oxidation was observed from + 0.45 V (vs. SCE) (I = 0.26 IO.03 I_IA cm-’ at this potential). In the presence of the complexing agents shown in Fig. 1. the anodic oxidation of mercury exhibited the characteristics gathered in Table 1. In normal pulse polarography, the limiting current of the observed wave was controlled by the diffusion of the ligand to the oxidizing electrode, as reported earlier with other macrocyclic ligands [lO,ll]. The method used to ascertain the diffusion control has been described previously [20] and, for a 2 s drop time and a pulse duration of 100 ms, /S In tP], for 10 < t,/ms ,< 90 (i.e. from 10 to the measured slopes [ 6 In I,,, cnrrccred 90% of the pulse) are given in Table 1. With the ligand [T( p-Ph),], however. two distinct oxidation steps were obtained. at variance with the single step observed with all the other ligands of Fig. 1. Controlled potential coulometry indicated that two or four electrons (Table 1) were involved in each of these oxidation steps. Cyclic voltammetric experiments performed between -0.300 and +0.700 V (vs. SCE) on the mercury electrode consistently indicated [21] ( Zpc,/r,‘~” constant. a positive shift of E,, which is non-linear, i.e. AE, more than proportional to log ~1. AE, = E,, - E,, increasing and non-linear with log I’, and Z,,/I,,, decreasing from unity with increasing cl) that the anodic oxidation of mercury, in the presence of all involved a quasi-reversible electrochemical the ligands of Fig. 1 but [T( p-Ph),]. oxidation of mercury. This quasi-reversible oxidation is followed by a reversible chemical reaction. as indicated by the decrease of I,,/Z,, on increasing I’ in cyclic voltammetry and by the negative shift of the oxidation potential of mercury when complex-forming ligands are present in the solution. Thus for each mercury atom the corresponding oxidation scheme is probably a two-electron oxidation followed by complexation of the resulting dication.
c
Mercury oxidation parameters
’
0.49
2.3 ?r 0.5
2 2.1
43+3
-
2.240.5
2 2.0
40&S
-
in the presence
lrgands
1.2kO.2
4 2.2
32+2
-0.55
1.3 * 0.4
2+2
3S&2 5S&5
- 0.54
+0.38 i (+o.ol) + 0.270 d (+0.015)
1.4+0.2
4 3.9
60+5
- 0.56
4
4.5&S
- 0.52
1.010.2 pfrom 80 to 90% of the pulse).
d
0.29 ’ (zkO.01)
FW'h),l
given in Fig. 1 a
not detectable
+O.26 b
(SOW
i-O.26 b
(+_O.Ol)
+0.27 E
I+W'W,I
of the maerocycltc
( + 0.01)
['Uo-Ph),]
electrodes
a Method used: normal pulse polarography (drop time = 2.000 s, delay = 1.950 s, pulse width = 0.040 s. current samphng h 2/3 PC+ l/3 C,H,C!a +O.l M THAI’. ’ 2/3 PC+ l/3 C,H,C12 CO.1 M TBAP. ’ DMF+O.l M THAP. fn this solvent, only the ftrst oxrdatton IS observable for the hgand [T(p-Ph),]. e in the determlnat~o~ of S In I NPPiorrecced/S In ~a the following expertmental parameters were used in normal pulse polarography (drop time = 2.000 s. delay = 1 890 s, pulse width = 0.100 5, current samplmg over 2% of the pulse (1.e. 2 ms) shtfted from the first 10 ms of the pulse to the end of the pulse duration). ’ In the coulometric determmations of n, the followrng exact results were obtained, respectively total charge (Q,,,,). background charge (Qh). amount of matertal ( VC). leading to the value n reported tn the table: Q,,, , Qh. VC m the presence: [N,O, ZMej (0.98 C: 1.4~10^~ C: 4.8 pmol), INsO, 3Mej (095 C; 1.4x lo-* C: 4.9 pmol). [T(o-Ph),) (1.26 C.1.4x 10e2 C; 5.8 firnoB. [T(DPO)J (1.73 C: 1.4~10~’ C: 4.6 pmol).
lo6 D/cm2 s-’
NPP Coufometry
n electrons per molecuie of complexed ligand:
fog[I/fI, - I)1 =f(E) /mV dec-”
Slope of
6 In I NPP corrected s In t
f 0.09 h (rtO.01) + 0.09 c (SO.01) + 0.085 d (+0.005)
+0.19 b
( z!z0.01) +0.16 = (i:O.Ol) i-o.19 d (t0.01)
4/2/V
3W
of mercury
WE)
oxtdatton
[N,O,
Ligand
the electrochemrcal
IN&‘., 2Mel
L201 c
characterizing
1
parameters
TABLE
195
In the presence of the ligand [T(p-Ph),] (Fig. l), the same overall oxidation-complexation scheme remains valid for the oxidation of the mercury electrode, but a specific feature here is the occurrence of two distinct two-electron steps (Table l), corresponding, respectively, to the oxidation-complexation of a first mercury at El,* = +0.25 V (vs. SCE) followed at E,,, = +0.39 V (vs. SCE) by the same reaction for a second mercury to be accommodated in the ligand [T( p-Ph)2], thus totalling four electrons exchanged in both oxidation steps (Table 1). If the size parameters of the ditopic ligand [T(p-Ph),] are taken into account, it is possible to calculate that the distance left between the two mercuric cations in this ligand will be only slightly more than 0.3 nm. Thus, it is easy to rationalize the above-reported behaviour, which indicated that the complexation of a second mercuric cation in the moiety [T( p-Ph),, Hg’+]‘+ was made slightly more difficult by the presence of the first complexed cation. This slightly negative cooperativity between the two mercuric cations can be ascribed to the limited size of the macrocyclic ligand [T( p-Ph),]. This explanation is also supported by the occurrence of a single four-electron process observed in the presence of [T(BPh),] and of [T(DPO),], in identical experimental conditions: in the ligand [T(DPO),], the distance between two complexed mercuric cations is almost twice the distance prevailing in the ligand [T( p-Ph)*]. Stoichiometry
of the resultmg mercuric complexes
In the present study, the oxidation potential of the mercury electrodes was found to be independent of the concentration of the ligand in solution. This means that each molecule of the electrogenerated mercuric complexes involves only one ligand. Thus, the stoichiometry of the electrogenerated mercuric complexes may not be higher than one ligand per mercuric cation [3]. It remained therefore to be determined which of the (Hg’+/ligand) stoichiometries, l/l or 2/l, prevailed. The stoichiometry of the mercuric complexes obtained was established in two different ways. The first was controlled potential coulometry applied to the global oxidation reaction: x Hg-2xe-+L+(Hg,L)“+
(x=1
or2)
For each of the ligands studied, controlled potential coulometry carried out at +0.450 V (vs. SCE) revealed that two electrons were consumed for each of the ligands [N,O, 2Me], [N,O, 3MeJ and [T(o-Ph),] (Fig. 1) whereas four electrons were consumed for each of the ligands [T( p-Ph),], [T(BPh),] and [T(DPO),] (Fig. 1) introduced into the solution. After consideration of the very great stability expected for the generated complexes from the large shift of the oxidation potential of mercury (i.e. generation of inclusion [22] complexes), considering also the available space inside the ligands (CPK space filling models), and in agreement with all previous experimental results which indicated that oxidation of mercury in the presence of complex-forming agents strongly stabilizes mercuric vs. mercurous complexes [3], the above results
196
indicate l/l stoichiometry for the complexes between the mercuric cations and the ligands [N204 2Me], [N,O, 3Me] and [T(o-Ph),]. They also reveal that two [Hg’+] are complexed in each of the ditopic ligands [T( p-Ph),], [T(BPh),] and [T(DPO),]. In a second set of experiments based on polarography, the stoichiometry 2 Hg’+/l [T(DP0)2] was confirmed independently as follows: when 0.15 ml of a lo-’ A4 Hg(NO,), solution in 2/3 PC + l/3 C,H,Cl, was added to a solution of [T(DPO),] (10 ml, c = 1.46 X 10d4 M), the original anodic wave changed to a composite anodic-cathodic wave. It was observed that, in these conditions corresponding to equal amounts of the ligand [T(DPO),] and of the cation [Hg”] introduced into the solution, the limiting current corresponding to the anodic oxidation of the mercury electrode (this current being proportional to the concentration of free ligand in the solution [11,12]) was half of the limiting current measured before any addition of [Hg”] cations to the solution of [T(DPO),]. Also, this anodic limiting current became zero as soon as the amount of [Hg”] cations introduced into the solution was twice that of the ligand. At this point, it might be worth noting that the conclusions drawn from this set of experiments are validated by the very high stability of the resulting mercuric complexes, which will be discussed later in the present paper. The results obtained for the anodic oxidation of mercury electrodes in the presence of the ligand [T(o-Ph),] deserve specific attention because exhaustive controlled potential coulometry indicated that two electrons were involved per mole of the resulting mercuric complex whereas experiments carried out on a shorter time-scale, by stationary and cyclic voltammetry, revealed (by intercomparisons with the limiting or peak currents in solutions of other ligands given in Fig. 1) that four electrons were exchanged, in the single oxidation observed, per mole of the resulting complex. However, the stable form of the mercuric complex of [T(o-PhJ has l/l stoichiometry, as revealed by increasing additions of Hg(CF,SO,), to a solution of [T(o-Ph),] and subsequent analysis of the decrease of the limiting anodic current according to the procedure described above in studying the stoichiometry of the other mercuric complexes. The following oxidation scheme provides a consistent interpretation of the body of the results observed with the ligand [T( o-Ph),]: 2 Hg + [T( o-Ph),]
{ Hg’+,
- 4 e- s { Hg2+, [T( o-Ph),] , Hg’+}4+ External complex 2/l
[T( o-Ph),] , Hg2+}4+ 2
Hg2+ + 2 e- + Hg”
(Hg*+ c [T( o-Ph),])2+ + Hg*+ Stable inclusion complex l/l
(1)
(2)
(3)
On the time-scale of the voltammetry, only the first step was observed, whereas the whole set of steps (l)-(3) occurred on the large time-scale of controlled potential coulometry at +0.4 or +0.5 V (vs. SCE), thus generating the stable mononuclear l/l complex.
797
Electrochemical
reduction of the dimercuric complex ((Hg’ ‘)?,
T(DPO), / ’ +
A solution of the dimercuric complex of [T(DPO),] was prepared by exhaustive controlled potential oxidation of mercury at +0.450 V (vs. SCE) in 2/3 PC + l/3 C,H,C12 containing 0.1 M TBAP and 4.6 X 10e4 M [T(DPO)I]. In the resulting solution, the electrochemical reduction of the complex [(Hg”)*. T(DP0),14+ was studied by cyclic voltammetry between +0.7 and - 0.5 V (vs. SCE). The results obtained (Fig. 2a) revealed that the reduction of this dimercuric complex was achieved in a single step on the time-scale of the method.
Fig. 2. Reduction of the dlmercuric complex [(Hg’+ J2, T(DP0)2]4+ (4.6X 10m4 M) m the solvent (2/3 PC + l/3 C, H,Cl z ) contammg 0.1 M TBAP, on a Hg electrode. (- - -_) Current-potential curve for the supportmg electrolyte. (a) Cyclic voltammetry. L’= 5 V/s; the asterisk mdlcates the start of the scan. (b) Tast polarography (drop time = 2 00 s: pulse width = 2 00 s: current sampling from 90 to 100% of the pulse).
198
Analysis of the obtained cyclic voltammograms led to the conclusion that this reduction occurred via a fast electron transfer, followed by a reversible chemical reaction (negative shift of Epc with the scan rate u was non-linear, E,, = +0.155 V (vs. SCE) at u = 2 V/s and +0.030 V (vs. SCE) at u = 100 V/s. Zpc/u”’ almost constant = 3.1 PA s ‘I2 V-‘12 from 2 to 100 V/s, and I,,/Z,,, departing from unity when u increases: from 1.0 at 2 V/s to 0.63 at 100 V/s) [21]. Polarographic analysis (Fig. 2b) of this reduction reaction (I?,,, = 0.205 V (vs. SCE)) confirmed the overall reversibility of the reduction on the polarographic time-scale. Thus the following reduction scheme consistently accounts for the above results: electrochemical chemical
step :
step: { (Hg’+),. { (Hd,.
[T(DPO),]}~++ [‘W-O),]
}3
4 e-a
{ (Hg),,
[T(DP0)2]
}
Hg + [TWO),]
with k, + k, 2 10” ( kf/k,,)’ This scheme proposed for the electrochemical reduction of the dimercuric complex ) deserves a short specific comment of ]T(DPO) 2I OLerslble followed by Creverblble with respect to the reaction scheme proposed for the oxidation of the mercury electrode in a solution containing free macrocyclic ligand (Equasl__, followed by C re,,ers,b,e)in order to make clear that both schemes do not correspond to the same reaction step (direct and reverse reaction). Therefore, their electrochemical (E) and chemical (C) steps do not correspond at all to identical elementary events. This is reflected in the different values obtained, for instance, from the analysis of the corresponding polarographic curves (for the oxidation of the mercury electrode in a solution of free [T(DPO),] ligand, E,,, = +0.26 V (vs. SCE) with a slope of 60 mV/dec; whereas for the reduction of the complex [(Hg2)T(DP0)2J4+, E,,z = +0.20 V (vs. SCE) with a slope of 36 mV/dec). Thus, the overall oxido-reduction process involving mercury and [T(DPO),] will be described as follows: (i) Reduction of [(Hg)ZT(DPO),]4’ to Hg and [T(DPO),]: reduction of the complexed mercury (E) followed by its decoordination (C) from [T(DPO),]. The chemical step then involves zerovalent mercury and the ligand (plus the solvent). (ii) Electrogeneration of [(Hg)2T(DPO),]4’ from the Hg electrode and free [T(DPO),]: oxidation of the mercury electrode (E) followed by complexation (C) of the mercuric cations into [T(DPO),]. The chemical step then involves. besides the solvent, mercuric cations and the ligand. The voltammetric and polarographic curves also exhibit additional features which provide interesting insight into the reduction mechanism of [(Hg)2T(DPO),]4’. First, in cyclic voltammetry, the difference between the peak and the half-peak reduction potentials ( Ep - Ep,z = -53 mV), independent of ~1, corresponds to a sequence of two distinct redox steps (distinct values of the corresponding E’s) if the electron transfers are Nernstian and in the absence of any other rate-determining step. On the other hand, the log plot analysis (E vs. log[ Z/( Id - Z)] of the single stationary wave gives 36 mV/dec for the slope of this linear plot (and not 15
199
mV/dec which would correspond to n = 4). From this slope. the difference AE” = EP - E,” between the two successive two-electron transfers may then be calculated [23] equal to 27 mV, if the two steps are Nernstian and rate-determining. As 17.4 mV corresponds [23] to two non-interacting dielectronic redox centres, this experimental value of AE” indicates that the two redox centres are almost independent. The reduction scheme is thus: { (Hg2+),[T(DPO),]
}4++ 2 em g { (Hg’+)(Hg’)[T(DPO),]
{ (Hg2+)(Hgo)[T(DP0)2]2’}
+2 em g { (Hg”),[T(DPO)2]
}‘+
)
It is clear, however, that the above scheme is correct only if neither nor any conformation change parallels the two redox steps. Stabilrty of the mercutw
a chemical
step
complexes
It has been a general observation that the stability of the mercuric complexes involving polyazamacrocyclic ligands is high in water as solvent [24]. With the diaza ligands, this stability remains almost unaffected by substituting groups on the two nitrogens, and even by the cyclization of the ligand [25]: for instance, the formation constant K is about 10” in H,O (25°C and ionic strength of 0.1) for the l/l mercuric complexes of the diaza ligands H,_,N(CH,CH,OCH,CH,OCH, CH,),NH,_, with n = 1, 2 or 3 [25]. This high stability was confirmed by the important shift in the oxidation potential of mercury, in the presence of diazamacrobicyclic ligands [ 111. In the presence of the mono- and ditopic ligands studied in the present paper (Fig. l), the electrochemical oxidation of the mercury electrode is also much facilitated (Tables 1 and 2). However, as the corresponding redox reactions are not reversible, the difference between the observed oxidation potential and that of mercury in the absence of any ligand in the solution has no straightforward thermodynamic significance because it results both from a thermodynamic contribution-proportional to the logarithm of the stability constant of the generated mercuric complex-and from a kinetic term. In this global effect, the net effect of the kinetic term is to decrease the shift of the oxidation potential of the mercury which would be observed with the thermodynamic contribution alone. On the other hand, the selected value 0.8 V (vs. SCE) for the oxidation potential of mercury in the absence of a complex-forming agent is an experimental one. which is actually E” + (RT/2F) In [Hg2’]. The value +0.8 V is therefore less positive than E”(Hg’+/Hg’). Thus, for a given ligand (Fig. 1) in solution, dividing the actual shift in the oxidation potential of mercury by (0.060/n) V gives only an approximate, minimum value for the logarithm of the corresponding stability constant (Table 2). However, as the oxidation processes depart only slightly from reversibility, the values calculated in Table 2 are likely to be reasonable approximations of
700 TABLE
2
Relative
stabihtles
Llgands
L
of the electrogenerated
%/mV (SCE)
W204 2Mel
’
mercwc
AE/mV (800-E,
complexes
in 2/3
10
h
PC + l/3
bOm”
Cz H,CI z + 0.1 M THAP
R stab
2)
+ 190
610
2 x 102” (n=2)
2 2x10’“‘
+90
710
4x10” (ll=2)
> 4x10’3’
+ 260 + 380
540 420
lolX (n = 2) 10’4(n=2)/
/T(o-Ph),l
+ 270
530
2x10’5r (n=4)
> 2x10’5d.r
[T(BPh)ll
+ 290
510
lox4 (n = 4)
> 10’“”
IT(DPO),I
+ 260
540
1036(n=4)
> 1036d
[N,O,
[T(
3Mel
p-Ph)zl
\
2 10’2 J
a Half-wave oxldatlon potential of the mercury electrode (normal pulse polarography as m Table 1). h 800 mV (vs. SCE) IS the oxidation potential of the mercury electrode m the absence of any hgand m solutmn [ HgL’+ ] ‘ K\,& =
’ Kstah= [Hg’+ ]‘[L] for the external
e Calculated
the stability constants chemical processes.
complex
attached
(2 mercury/l
ligand).
to the mercuric
complexes
generated
in this electro-
CONCLUSION
The electrochemical oxidation of mercury is much facilitated in the presence of both monotopic and ditopic macrocyclic ligands (Fig. 1). The resulting mono- and dinuclear complexes are extremely stable, with a clear-cut effect of the length of the lateral chains in the ditopic ligands on the stoichrometries and stabilities of the complexes. The dinuclear complex {(Hg’+),, [T(DPO),]}4’ exhtbits the remarkable characteristic of being able to undergo an almost simultaneous four-electron transfer at +0.2 V (vs. SCE), which may be reversed on the time-scale of voltammetric methods, thus implying the generation of the reduced form {(Hg”)?, [T(DPO), } at the electrode. REFERENCES 1 J. Ravenda, Collect. Czech. Chem. Commun . 6 (1934) 453. 2 I.M. Kolthoff and C.S. Mdler. J. Am. Chem. Sot . 63 (1941) 1405. 2732
201 3 J. Heyrovsky and J. Kuta, Prmctples of Polarography, Academic Press, New York, 1966. pp. 172-178 and refs. ctted therem. 4 J. Goffart. G Michel and G. Duyckaerts. Anal. Chtm. Acta, 9 (1953) 184 5 G. Mtchel. Anal Chtm. Acta. 10 (1954) 87. 6 C.N. Retlley. W.G Scribner and C. Temple, Anal. Chem.. 28 (1956) 450. 7 M. Morattlle and B. Tremillon. Bull. Sot. Chum. Fr., (1961) 506. 8 B. Fleet. Soe Wm and T.S. West. Analyst. 94 (1969) 269. 9 L.L Jackson. J. Osteryoung and R.A. Osteryoung, Anal. Chem.. 52 (1980) 66. 10 F. Peter. M. Gross, L. Posptsd and J. Kuta. J. Electroanal. Chem.. 90 (1978) 239. 11 C Boudon, F. Peter and M. Gross. J Electroanal. Chem.. 135 (1982) 93. 12 C. Boudon, M. Gross, F Kotzyba-Hibert. J.M. Lehn and K. Satgo. J. Electroanal. Chem.. 191 (1985) 201. 13 A. Gtraudeau. H.J. Callot and M. Gross, Inorg. Chem.. 18 (1979) 201. 14 A. Gtraudeau, H.J. Callot. J. Jordan. I. Ezahr and M. Gross. J. Am. Chem. Sot., 101 (1979) 3857. 15 G. Rttzler, F. Peter and M. Gross, J. Electroanal Chem.. 146 (1983) 285. 16 C. Boudon, F. Peter and M. Gross, J. Electroanal. Chem.. 117 (1981) 65. 17 J.M. Lehn and P. Vierhng. Tetrahedron Lett.. 21 (1980) 1323. 18 F. Kotzyba-Hibert, These de Doctorat d’Etat, Umverstte Louis Pasteur. Strasbourg. 1983. 19 J.P. Kmtzmger, F. Kotzyba-Hibert, J.M. Lehn. A. Pagelot and K. Saigo, J. Chem. Sot. Chem. Commun., (1981) 833. 20 M. Gross and J. Jordan. J Electroanal. Chem., 75 (1977) 163. 21 (a) R.S. Ntcholson and I. Sham. Anal. Chem.. 36 (1964) 706; (b) E.R. Brown and R.F. Large m A. Weisberger and B.W Rosstter (Eds.), Techniques of Chemtstry, Vol. 1: Phystcal Methods of Chemistry, Part 2A. Electrochemtcal Methods. Wtley-Intersctence, New York, 1971, p. 425. 22 J.M. Lehn, Act. Chem Res., 11 (1978) 49. 23 F. Ammar and J.M. Saveant, J. Electroanal. Chem., 47 (1973) 215. 24 (a) M Kodama and E. Ktmura. J. Chem. Sot., Dalton Trans.. (1976) 2335: (b) ibtd. (1978) 1081. 25 G Anderegg. Helv. Chum. Acta. 58 (1975) 1218.
191 m The Netherlands
ELECTROCHEMICAL GENERATION MERCURIC CRYPTATES
C. BOUDON
l, J.P. GISSELBRECHT
and J.M. LEHN
l
OF MONO-
AND DINUCLEAR
l, M. GROSS l. F. KOTZYBA-HIBERT
l
*
*
DtQxwtement de Chrmre, LJmc~ers~t~LAWS Pasteur. 4, Rue Blarse Pa.ccul, 67000 Strushourg (Fruncr) (Received
26th September
1985. m revised form 23rd December
1985)
ABSTRACT
The electrochemical oxidation of mercury electrodes was studted m orgamc solvents m the presence of macrocychc ligands. It resulted m the generation of very stable mono- and dinuclear mercuric complexes. The redox couple based on the dimercuric complex with the hgand T(DPO), exhibits the unprecedented property of transferring almost stmultaneously four electrons at + 0.2 V (~5 SCE).
INTRODUCTION
Fifty years ago, Ravenda [l] reported that the electrochemical oxidation of mercury was facilitated in the presence of complexing anions. After the pioneering work of Kolthoff and Miller [2]. several authors studied extensively the electrochemtcal oxidation of mercury, thermodynamically assisted by complexing ligands, and, in some instances, this specific reaction was applied to indirect titrations of cations [3-91. More recently, it was observed that polyazamacrocycles, and among them the diazamacrobicycles, also facilitate the anodic oxidation of mercury [lO.ll]. allowing either direct titration of the macrocycles in the solution, indirect titration of metal ions [ll], or indirect titration of other mono- and dications [12]. Thus, considering the prominent interest in a better understanding of molecular receptor-substrate interactions and in identifying the critical parameters in these interactions, analysis of the characteristic parameters involved in the electrogeneration of mercuric complexes with the mono- and ditopic ligands shown in Fig. 1 was attempted. The corresponding results are reported and discussed in the present paper. These results reveal an actual control of the cooperativity between the two
* Laboratoire ** Laboratoue
d’Electrochimte et Chimie Physique du Corps Sohde. U.A. au C.N.R.S. de Chimie Orgamque Physique. U A au C.N.R.S No. 422.
0012-0728/86/$03.50
‘c’ 1986 Elsevter Sequoia
S.A
No. 405.
192
cN204
2 Me
3
[T(P-‘d]
Fig. 1. The ligands
[Tb-P’&]
used in the present
stud).
mercuric ions through the length ditopic ligands (Fig. 1).
and the flexibility
of the lateral
chains
in the
EXPERIMENTAL
The electrochemical measurements were carried out in dimethylformamide (DMF) or in a mixed solvent (l/3 1,2-dichloroethane + 2/3 propylene carbonate (PC)). The solvents were purified before use as described previously [13-151. The ligands studied (Fig. 1) are poorly soluble in PC and very soluble in 1,2-C>H,Cl,. The supporting electrolyte was either tetra-n-hexylammonium perchlorate (THAP. from Fluka) or tetra-n-butylammonium perchlorate (TBAP, from Eastman Kodak). and both were recrystallized before use [16]. Electrochemical measurements were carried out with a multipurpose PRG4 device (SOLEA, Tacussel), in a classical three-electrode cell. on a dropping mercury electrode as the working electrode (m = 0.178 mg/s. h = 40 cm Hg). The auxiliary electrode was a platinum wire. The reference electrode was a calomel electrode in an aqueous KCl-saturated solution, electrically connected to the electrolysis cell by a bridge filled with the same organic solution as that introduced into the cell (solvent + supporting electrolyte). It was repeatedly checked that, with this electrolytic junction, the junction potential remained constant during the series of experi-
193
ments. The followmg electrochemical methods were used in the present study: normal pulse polarography, dc polarography with controlled drop time, cyclic voltammetry and controlled potential coulometry. In the latter experiments (coulometry), the working electrode was a mercury pool electrode (area = 15 cm’) in a cell containing about 10 cm3 of solution (the exact volume was measured in each experiment). The syntheses of the ligands used (Fig. 1) have been described previously: [N,O, 2Me] and [N,O, 3Me] [17]. [T(o-Ph),] [18]. [T(p-Ph),] [19], [T(BPh)?] [18] and
[-UDPO)~l [181. RESULTS
Anodic oxidatron of mercury electrodes in the presence of macroqvclic Iigands When no complexing ligand was present in the solution, the electrochemical oxidation of the mercury electrode occurred from +0.80 V (vs. SCE) in the solvent (2/3 PC + l/3 C,H,C12) containing also 0.1 M THAP or TBAP. At this potential. +0.80 V (vs. SCE), i was 0.18 + 0.02 PA cm-‘. In the solvent DMF + 0.1 M THAP or TBAP, mercury oxidation was observed from + 0.45 V (vs. SCE) (I = 0.26 IO.03 I_IA cm-’ at this potential). In the presence of the complexing agents shown in Fig. 1. the anodic oxidation of mercury exhibited the characteristics gathered in Table 1. In normal pulse polarography, the limiting current of the observed wave was controlled by the diffusion of the ligand to the oxidizing electrode, as reported earlier with other macrocyclic ligands [lO,ll]. The method used to ascertain the diffusion control has been described previously [20] and, for a 2 s drop time and a pulse duration of 100 ms, /S In tP], for 10 < t,/ms ,< 90 (i.e. from 10 to the measured slopes [ 6 In I,,, cnrrccred 90% of the pulse) are given in Table 1. With the ligand [T( p-Ph),], however. two distinct oxidation steps were obtained. at variance with the single step observed with all the other ligands of Fig. 1. Controlled potential coulometry indicated that two or four electrons (Table 1) were involved in each of these oxidation steps. Cyclic voltammetric experiments performed between -0.300 and +0.700 V (vs. SCE) on the mercury electrode consistently indicated [21] ( Zpc,/r,‘~” constant. a positive shift of E,, which is non-linear, i.e. AE, more than proportional to log ~1. AE, = E,, - E,, increasing and non-linear with log I’, and Z,,/I,,, decreasing from unity with increasing cl) that the anodic oxidation of mercury, in the presence of all involved a quasi-reversible electrochemical the ligands of Fig. 1 but [T( p-Ph),]. oxidation of mercury. This quasi-reversible oxidation is followed by a reversible chemical reaction. as indicated by the decrease of I,,/Z,, on increasing I’ in cyclic voltammetry and by the negative shift of the oxidation potential of mercury when complex-forming ligands are present in the solution. Thus for each mercury atom the corresponding oxidation scheme is probably a two-electron oxidation followed by complexation of the resulting dication.
c
Mercury oxidation parameters
’
0.49
2.3 ?r 0.5
2 2.1
43+3
-
2.240.5
2 2.0
40&S
-
in the presence
lrgands
1.2kO.2
4 2.2
32+2
-0.55
1.3 * 0.4
2+2
3S&2 5S&5
- 0.54
+0.38 i (+o.ol) + 0.270 d (+0.015)
1.4+0.2
4 3.9
60+5
- 0.56
4
4.5&S
- 0.52
1.010.2 pfrom 80 to 90% of the pulse).
d
0.29 ’ (zkO.01)
FW'h),l
given in Fig. 1 a
not detectable
+O.26 b
(SOW
i-O.26 b
(+_O.Ol)
+0.27 E
I+W'W,I
of the maerocycltc
( + 0.01)
['Uo-Ph),]
electrodes
a Method used: normal pulse polarography (drop time = 2.000 s, delay = 1.950 s, pulse width = 0.040 s. current samphng h 2/3 PC+ l/3 C,H,C!a +O.l M THAI’. ’ 2/3 PC+ l/3 C,H,C12 CO.1 M TBAP. ’ DMF+O.l M THAP. fn this solvent, only the ftrst oxrdatton IS observable for the hgand [T(p-Ph),]. e in the determlnat~o~ of S In I NPPiorrecced/S In ~a the following expertmental parameters were used in normal pulse polarography (drop time = 2.000 s. delay = 1 890 s, pulse width = 0.100 5, current samplmg over 2% of the pulse (1.e. 2 ms) shtfted from the first 10 ms of the pulse to the end of the pulse duration). ’ In the coulometric determmations of n, the followrng exact results were obtained, respectively total charge (Q,,,,). background charge (Qh). amount of matertal ( VC). leading to the value n reported tn the table: Q,,, , Qh. VC m the presence: [N,O, ZMej (0.98 C: 1.4~10^~ C: 4.8 pmol), INsO, 3Mej (095 C; 1.4x lo-* C: 4.9 pmol). [T(o-Ph),) (1.26 C.1.4x 10e2 C; 5.8 firnoB. [T(DPO)J (1.73 C: 1.4~10~’ C: 4.6 pmol).
lo6 D/cm2 s-’
NPP Coufometry
n electrons per molecuie of complexed ligand:
fog[I/fI, - I)1 =f(E) /mV dec-”
Slope of
6 In I NPP corrected s In t
f 0.09 h (rtO.01) + 0.09 c (SO.01) + 0.085 d (+0.005)
+0.19 b
( z!z0.01) +0.16 = (i:O.Ol) i-o.19 d (t0.01)
4/2/V
3W
of mercury
WE)
oxtdatton
[N,O,
Ligand
the electrochemrcal
IN&‘., 2Mel
L201 c
characterizing
1
parameters
TABLE
195
In the presence of the ligand [T(p-Ph),] (Fig. l), the same overall oxidation-complexation scheme remains valid for the oxidation of the mercury electrode, but a specific feature here is the occurrence of two distinct two-electron steps (Table l), corresponding, respectively, to the oxidation-complexation of a first mercury at El,* = +0.25 V (vs. SCE) followed at E,,, = +0.39 V (vs. SCE) by the same reaction for a second mercury to be accommodated in the ligand [T( p-Ph)2], thus totalling four electrons exchanged in both oxidation steps (Table 1). If the size parameters of the ditopic ligand [T(p-Ph),] are taken into account, it is possible to calculate that the distance left between the two mercuric cations in this ligand will be only slightly more than 0.3 nm. Thus, it is easy to rationalize the above-reported behaviour, which indicated that the complexation of a second mercuric cation in the moiety [T( p-Ph),, Hg’+]‘+ was made slightly more difficult by the presence of the first complexed cation. This slightly negative cooperativity between the two mercuric cations can be ascribed to the limited size of the macrocyclic ligand [T( p-Ph),]. This explanation is also supported by the occurrence of a single four-electron process observed in the presence of [T(BPh),] and of [T(DPO),], in identical experimental conditions: in the ligand [T(DPO),], the distance between two complexed mercuric cations is almost twice the distance prevailing in the ligand [T( p-Ph)*]. Stoichiometry
of the resultmg mercuric complexes
In the present study, the oxidation potential of the mercury electrodes was found to be independent of the concentration of the ligand in solution. This means that each molecule of the electrogenerated mercuric complexes involves only one ligand. Thus, the stoichiometry of the electrogenerated mercuric complexes may not be higher than one ligand per mercuric cation [3]. It remained therefore to be determined which of the (Hg’+/ligand) stoichiometries, l/l or 2/l, prevailed. The stoichiometry of the mercuric complexes obtained was established in two different ways. The first was controlled potential coulometry applied to the global oxidation reaction: x Hg-2xe-+L+(Hg,L)“+
(x=1
or2)
For each of the ligands studied, controlled potential coulometry carried out at +0.450 V (vs. SCE) revealed that two electrons were consumed for each of the ligands [N,O, 2Me], [N,O, 3MeJ and [T(o-Ph),] (Fig. 1) whereas four electrons were consumed for each of the ligands [T( p-Ph),], [T(BPh),] and [T(DPO),] (Fig. 1) introduced into the solution. After consideration of the very great stability expected for the generated complexes from the large shift of the oxidation potential of mercury (i.e. generation of inclusion [22] complexes), considering also the available space inside the ligands (CPK space filling models), and in agreement with all previous experimental results which indicated that oxidation of mercury in the presence of complex-forming agents strongly stabilizes mercuric vs. mercurous complexes [3], the above results
196
indicate l/l stoichiometry for the complexes between the mercuric cations and the ligands [N204 2Me], [N,O, 3Me] and [T(o-Ph),]. They also reveal that two [Hg’+] are complexed in each of the ditopic ligands [T( p-Ph),], [T(BPh),] and [T(DPO),]. In a second set of experiments based on polarography, the stoichiometry 2 Hg’+/l [T(DP0)2] was confirmed independently as follows: when 0.15 ml of a lo-’ A4 Hg(NO,), solution in 2/3 PC + l/3 C,H,Cl, was added to a solution of [T(DPO),] (10 ml, c = 1.46 X 10d4 M), the original anodic wave changed to a composite anodic-cathodic wave. It was observed that, in these conditions corresponding to equal amounts of the ligand [T(DPO),] and of the cation [Hg”] introduced into the solution, the limiting current corresponding to the anodic oxidation of the mercury electrode (this current being proportional to the concentration of free ligand in the solution [11,12]) was half of the limiting current measured before any addition of [Hg”] cations to the solution of [T(DPO),]. Also, this anodic limiting current became zero as soon as the amount of [Hg”] cations introduced into the solution was twice that of the ligand. At this point, it might be worth noting that the conclusions drawn from this set of experiments are validated by the very high stability of the resulting mercuric complexes, which will be discussed later in the present paper. The results obtained for the anodic oxidation of mercury electrodes in the presence of the ligand [T(o-Ph),] deserve specific attention because exhaustive controlled potential coulometry indicated that two electrons were involved per mole of the resulting mercuric complex whereas experiments carried out on a shorter time-scale, by stationary and cyclic voltammetry, revealed (by intercomparisons with the limiting or peak currents in solutions of other ligands given in Fig. 1) that four electrons were exchanged, in the single oxidation observed, per mole of the resulting complex. However, the stable form of the mercuric complex of [T(o-PhJ has l/l stoichiometry, as revealed by increasing additions of Hg(CF,SO,), to a solution of [T(o-Ph),] and subsequent analysis of the decrease of the limiting anodic current according to the procedure described above in studying the stoichiometry of the other mercuric complexes. The following oxidation scheme provides a consistent interpretation of the body of the results observed with the ligand [T( o-Ph),]: 2 Hg + [T( o-Ph),]
{ Hg’+,
- 4 e- s { Hg2+, [T( o-Ph),] , Hg’+}4+ External complex 2/l
[T( o-Ph),] , Hg2+}4+ 2
Hg2+ + 2 e- + Hg”
(Hg*+ c [T( o-Ph),])2+ + Hg*+ Stable inclusion complex l/l
(1)
(2)
(3)
On the time-scale of the voltammetry, only the first step was observed, whereas the whole set of steps (l)-(3) occurred on the large time-scale of controlled potential coulometry at +0.4 or +0.5 V (vs. SCE), thus generating the stable mononuclear l/l complex.
797
Electrochemical
reduction of the dimercuric complex ((Hg’ ‘)?,
T(DPO), / ’ +
A solution of the dimercuric complex of [T(DPO),] was prepared by exhaustive controlled potential oxidation of mercury at +0.450 V (vs. SCE) in 2/3 PC + l/3 C,H,C12 containing 0.1 M TBAP and 4.6 X 10e4 M [T(DPO)I]. In the resulting solution, the electrochemical reduction of the complex [(Hg”)*. T(DP0),14+ was studied by cyclic voltammetry between +0.7 and - 0.5 V (vs. SCE). The results obtained (Fig. 2a) revealed that the reduction of this dimercuric complex was achieved in a single step on the time-scale of the method.
Fig. 2. Reduction of the dlmercuric complex [(Hg’+ J2, T(DP0)2]4+ (4.6X 10m4 M) m the solvent (2/3 PC + l/3 C, H,Cl z ) contammg 0.1 M TBAP, on a Hg electrode. (- - -_) Current-potential curve for the supportmg electrolyte. (a) Cyclic voltammetry. L’= 5 V/s; the asterisk mdlcates the start of the scan. (b) Tast polarography (drop time = 2 00 s: pulse width = 2 00 s: current sampling from 90 to 100% of the pulse).
198
Analysis of the obtained cyclic voltammograms led to the conclusion that this reduction occurred via a fast electron transfer, followed by a reversible chemical reaction (negative shift of Epc with the scan rate u was non-linear, E,, = +0.155 V (vs. SCE) at u = 2 V/s and +0.030 V (vs. SCE) at u = 100 V/s. Zpc/u”’ almost constant = 3.1 PA s ‘I2 V-‘12 from 2 to 100 V/s, and I,,/Z,,, departing from unity when u increases: from 1.0 at 2 V/s to 0.63 at 100 V/s) [21]. Polarographic analysis (Fig. 2b) of this reduction reaction (I?,,, = 0.205 V (vs. SCE)) confirmed the overall reversibility of the reduction on the polarographic time-scale. Thus the following reduction scheme consistently accounts for the above results: electrochemical chemical
step :
step: { (Hg’+),. { (Hd,.
[T(DPO),]}~++ [‘W-O),]
}3
4 e-a
{ (Hg),,
[T(DP0)2]
}
Hg + [TWO),]
with k, + k, 2 10” ( kf/k,,)’ This scheme proposed for the electrochemical reduction of the dimercuric complex ) deserves a short specific comment of ]T(DPO) 2I OLerslble followed by Creverblble with respect to the reaction scheme proposed for the oxidation of the mercury electrode in a solution containing free macrocyclic ligand (Equasl__, followed by C re,,ers,b,e)in order to make clear that both schemes do not correspond to the same reaction step (direct and reverse reaction). Therefore, their electrochemical (E) and chemical (C) steps do not correspond at all to identical elementary events. This is reflected in the different values obtained, for instance, from the analysis of the corresponding polarographic curves (for the oxidation of the mercury electrode in a solution of free [T(DPO),] ligand, E,,, = +0.26 V (vs. SCE) with a slope of 60 mV/dec; whereas for the reduction of the complex [(Hg2)T(DP0)2J4+, E,,z = +0.20 V (vs. SCE) with a slope of 36 mV/dec). Thus, the overall oxido-reduction process involving mercury and [T(DPO),] will be described as follows: (i) Reduction of [(Hg)ZT(DPO),]4’ to Hg and [T(DPO),]: reduction of the complexed mercury (E) followed by its decoordination (C) from [T(DPO),]. The chemical step then involves zerovalent mercury and the ligand (plus the solvent). (ii) Electrogeneration of [(Hg)2T(DPO),]4’ from the Hg electrode and free [T(DPO),]: oxidation of the mercury electrode (E) followed by complexation (C) of the mercuric cations into [T(DPO),]. The chemical step then involves. besides the solvent, mercuric cations and the ligand. The voltammetric and polarographic curves also exhibit additional features which provide interesting insight into the reduction mechanism of [(Hg)2T(DPO),]4’. First, in cyclic voltammetry, the difference between the peak and the half-peak reduction potentials ( Ep - Ep,z = -53 mV), independent of ~1, corresponds to a sequence of two distinct redox steps (distinct values of the corresponding E’s) if the electron transfers are Nernstian and in the absence of any other rate-determining step. On the other hand, the log plot analysis (E vs. log[ Z/( Id - Z)] of the single stationary wave gives 36 mV/dec for the slope of this linear plot (and not 15
199
mV/dec which would correspond to n = 4). From this slope. the difference AE” = EP - E,” between the two successive two-electron transfers may then be calculated [23] equal to 27 mV, if the two steps are Nernstian and rate-determining. As 17.4 mV corresponds [23] to two non-interacting dielectronic redox centres, this experimental value of AE” indicates that the two redox centres are almost independent. The reduction scheme is thus: { (Hg2+),[T(DPO),]
}4++ 2 em g { (Hg’+)(Hg’)[T(DPO),]
{ (Hg2+)(Hgo)[T(DP0)2]2’}
+2 em g { (Hg”),[T(DPO)2]
}‘+
)
It is clear, however, that the above scheme is correct only if neither nor any conformation change parallels the two redox steps. Stabilrty of the mercutw
a chemical
step
complexes
It has been a general observation that the stability of the mercuric complexes involving polyazamacrocyclic ligands is high in water as solvent [24]. With the diaza ligands, this stability remains almost unaffected by substituting groups on the two nitrogens, and even by the cyclization of the ligand [25]: for instance, the formation constant K is about 10” in H,O (25°C and ionic strength of 0.1) for the l/l mercuric complexes of the diaza ligands H,_,N(CH,CH,OCH,CH,OCH, CH,),NH,_, with n = 1, 2 or 3 [25]. This high stability was confirmed by the important shift in the oxidation potential of mercury, in the presence of diazamacrobicyclic ligands [ 111. In the presence of the mono- and ditopic ligands studied in the present paper (Fig. l), the electrochemical oxidation of the mercury electrode is also much facilitated (Tables 1 and 2). However, as the corresponding redox reactions are not reversible, the difference between the observed oxidation potential and that of mercury in the absence of any ligand in the solution has no straightforward thermodynamic significance because it results both from a thermodynamic contribution-proportional to the logarithm of the stability constant of the generated mercuric complex-and from a kinetic term. In this global effect, the net effect of the kinetic term is to decrease the shift of the oxidation potential of the mercury which would be observed with the thermodynamic contribution alone. On the other hand, the selected value 0.8 V (vs. SCE) for the oxidation potential of mercury in the absence of a complex-forming agent is an experimental one. which is actually E” + (RT/2F) In [Hg2’]. The value +0.8 V is therefore less positive than E”(Hg’+/Hg’). Thus, for a given ligand (Fig. 1) in solution, dividing the actual shift in the oxidation potential of mercury by (0.060/n) V gives only an approximate, minimum value for the logarithm of the corresponding stability constant (Table 2). However, as the oxidation processes depart only slightly from reversibility, the values calculated in Table 2 are likely to be reasonable approximations of
700 TABLE
2
Relative
stabihtles
Llgands
L
of the electrogenerated
%/mV (SCE)
W204 2Mel
’
mercwc
AE/mV (800-E,
complexes
in 2/3
10
h
PC + l/3
bOm”
Cz H,CI z + 0.1 M THAP
R stab
2)
+ 190
610
2 x 102” (n=2)
2 2x10’“‘
+90
710
4x10” (ll=2)
> 4x10’3’
+ 260 + 380
540 420
lolX (n = 2) 10’4(n=2)/
/T(o-Ph),l
+ 270
530
2x10’5r (n=4)
> 2x10’5d.r
[T(BPh)ll
+ 290
510
lox4 (n = 4)
> 10’“”
IT(DPO),I
+ 260
540
1036(n=4)
> 1036d
[N,O,
[T(
3Mel
p-Ph)zl
\
2 10’2 J
a Half-wave oxldatlon potential of the mercury electrode (normal pulse polarography as m Table 1). h 800 mV (vs. SCE) IS the oxidation potential of the mercury electrode m the absence of any hgand m solutmn [ HgL’+ ] ‘ K\,& =
’ Kstah= [Hg’+ ]‘[L] for the external
e Calculated
the stability constants chemical processes.
complex
attached
(2 mercury/l
ligand).
to the mercuric
complexes
generated
in this electro-
CONCLUSION
The electrochemical oxidation of mercury is much facilitated in the presence of both monotopic and ditopic macrocyclic ligands (Fig. 1). The resulting mono- and dinuclear complexes are extremely stable, with a clear-cut effect of the length of the lateral chains in the ditopic ligands on the stoichrometries and stabilities of the complexes. The dinuclear complex {(Hg’+),, [T(DPO),]}4’ exhtbits the remarkable characteristic of being able to undergo an almost simultaneous four-electron transfer at +0.2 V (vs. SCE), which may be reversed on the time-scale of voltammetric methods, thus implying the generation of the reduced form {(Hg”)?, [T(DPO), } at the electrode. REFERENCES 1 J. Ravenda, Collect. Czech. Chem. Commun . 6 (1934) 453. 2 I.M. Kolthoff and C.S. Mdler. J. Am. Chem. Sot . 63 (1941) 1405. 2732
201 3 J. Heyrovsky and J. Kuta, Prmctples of Polarography, Academic Press, New York, 1966. pp. 172-178 and refs. ctted therem. 4 J. Goffart. G Michel and G. Duyckaerts. Anal. Chtm. Acta, 9 (1953) 184 5 G. Mtchel. Anal Chtm. Acta. 10 (1954) 87. 6 C.N. Retlley. W.G Scribner and C. Temple, Anal. Chem.. 28 (1956) 450. 7 M. Morattlle and B. Tremillon. Bull. Sot. Chum. Fr., (1961) 506. 8 B. Fleet. Soe Wm and T.S. West. Analyst. 94 (1969) 269. 9 L.L Jackson. J. Osteryoung and R.A. Osteryoung, Anal. Chem.. 52 (1980) 66. 10 F. Peter. M. Gross, L. Posptsd and J. Kuta. J. Electroanal. Chem.. 90 (1978) 239. 11 C Boudon, F. Peter and M. Gross. J Electroanal. Chem.. 135 (1982) 93. 12 C. Boudon, M. Gross, F Kotzyba-Hibert. J.M. Lehn and K. Satgo. J. Electroanal. Chem.. 191 (1985) 201. 13 A. Gtraudeau. H.J. Callot and M. Gross, Inorg. Chem.. 18 (1979) 201. 14 A. Gtraudeau, H.J. Callot. J. Jordan. I. Ezahr and M. Gross. J. Am. Chem. Sot., 101 (1979) 3857. 15 G. Rttzler, F. Peter and M. Gross, J. Electroanal Chem.. 146 (1983) 285. 16 C. Boudon, F. Peter and M. Gross, J. Electroanal. Chem.. 117 (1981) 65. 17 J.M. Lehn and P. Vierhng. Tetrahedron Lett.. 21 (1980) 1323. 18 F. Kotzyba-Hibert, These de Doctorat d’Etat, Umverstte Louis Pasteur. Strasbourg. 1983. 19 J.P. Kmtzmger, F. Kotzyba-Hibert, J.M. Lehn. A. Pagelot and K. Saigo, J. Chem. Sot. Chem. Commun., (1981) 833. 20 M. Gross and J. Jordan. J Electroanal. Chem., 75 (1977) 163. 21 (a) R.S. Ntcholson and I. Sham. Anal. Chem.. 36 (1964) 706; (b) E.R. Brown and R.F. Large m A. Weisberger and B.W Rosstter (Eds.), Techniques of Chemtstry, Vol. 1: Phystcal Methods of Chemistry, Part 2A. Electrochemtcal Methods. Wtley-Intersctence, New York, 1971, p. 425. 22 J.M. Lehn, Act. Chem Res., 11 (1978) 49. 23 F. Ammar and J.M. Saveant, J. Electroanal. Chem., 47 (1973) 215. 24 (a) M Kodama and E. Ktmura. J. Chem. Sot., Dalton Trans.. (1976) 2335: (b) ibtd. (1978) 1081. 25 G Anderegg. Helv. Chum. Acta. 58 (1975) 1218.