Reactions of Silver(l), Mercury I), and Mercury(l1) with Halide Ions in Acetonitrile as Solvent J. F. Coetzee,' J. J. Campion, and D. R. Liberman Department of Chemistry, University of Pittsburgh, Pittsburgh, Pa. 15213 The following quantities are reported for acetonitrile as solvent: the standard potentials of the Ag+-Ag and Hg22+-Hg couples, the equilibrium constant for the disproportionation reaction H g 2 * + eHg HgZ+,solubility product constants for silver and mercury(1) halides, and formation constants for the following complexes: AgX?-, HgX2, HgX3-, and HgX42-. The existence of the polynuclear complexes Ag2Br3-, Ag2!3-, and Ag516-is demonstrated. Some comparisons with other solvents are presented. From these data the behavior of silversilver(1) and mercury-mercury(1) halide electrodes can be predicted quantitatively.
+
ACETONITRILE 1s A RELATIVELY INERT solvent in which the majority of inorganic and many organic ions are solvated more weakly and hence are more reactive than in water. This is particularly true of anions, except those having very low charge density (e.g., tetraphenylborate) and/or high polarizability (e.g., picrate). As a result, studies in acetonitrile and similar solvents have contributed new perspectives on many intrinsic properties of solutes that are masked in more reactive solvents. For a variety of reasons, it is interesting to compare the reactions of halide ions with silver@), mercury(I), and mercury(II), among them the following. Silver(1) is one of the few inorganic cations that are more strongly solvated in acetonitrile than in water; hence, we are considering here the interactions of a relatively stable cation with highly reactive anions. Second, there is still a dearth of quantitative information on differences in chemical properties of different oxidation states of the same element in nonaqueous solvents; hence, comparison of the properties of mercury(1) with those of mercury(II) should be instructive. Finally, silver-silver(1) and mercury-mercury(1) couples are of unusual interest in electrochemistry, one reason being their widespread application in reference electrodes. The solubility product constants for silver halides and the formation constants for AgX2- Complexes determined in this study (1) agree well with the results of a simultaneous, independent investigation by Luehrs, Iwamoto, and Kleinberg (2). Furthermore, under our experimental conditions additional silver halide complexes could be identified and some of their properties determined. Finally, certain analytical implications of the properties of silver(I), mercury(I), and mercury(I1) couples in acetonitrile will be pointed out. EXPERIMENTAL
Apparatus. Potentiometric experiments were carried out with a Leeds and Northrup Model 8687 precision potentiometer in an H-type cell constructed with IO-mm diameter fritted glass disks of fine porosity located 5 cm apart in the salt bridge section of the cell. The reference compartment of the Please address all correspondence to this author. (1) J. J. Campion, Ph.D. thesis, University of Pittsburgh, Pitts-
burgh. Pa., 1966. (2) D. C . Luehrs, R. T. Iwamoto, and J. Kleinberg, Inorg. Chem., 5,201 (1966).
cell contained an Ag/(O.OlOM AgN03 in acetonitrile) electrode, hereafter designated as AgRE, and the salt bridge solution was 0.1M Et4NC10, in acetonitrile. The working compartment of the cell contained either a silver or a platinum electrode, or a J-tube electrode ( 3 ) filled with mercury, and was fitted with a magnetic stirrer. Mercury was determined by atomic absorption, using a Perkin-Elmer Model 305 spectrophotometer. Reagents. Matheson, Coleman and Bell practical grade acetonitrile was purified by a procedure ( 4 ) in which the final steps are fractional distillation, first from phosphorus pentoxide and then from calcium hydride. The purified solvent was tested for significant impurities ( 4 ) ; upper limits of water, total amines (as ammonia), and acrylonitrile were 1 X 3 x lo+, and 1 X 10F3M,respectively. Tetraethylammonium chloride, bromide, iodide, and perchlorate were purified as described before (5). Fisher Certified silver nitrate and G. F. Smith Company mercury(1) and mercury(I1) perchlorates, as Hg2(C104)2.8H20 and Hg(C104)2.6H20, were dried in l;acuo at 50 OC for 72 hours. The mercury(I1) salt was converted into the dihydrate by this procedure. Baker Analyzed Reagent (anhydrous) mercury(I1) chloride, bromide, and iodide were dried at 120 O C . Mercury was a triple distilled product from Bethlehem Apparatus Company. RESULTS
Full details of results can be found elsewhere ( I ) and will not be repeated here. All potentials reported are reduction potentials on the molar scale. Standard Potential of the Ag+-Ag Couple. The standard potential was calculated from values of the potential of the silver electrode at various points beyond the main equivalence point (1 : 1 stoichiometry) in the titration of tetraethylammonium halides with silver nitrate. Typical potentiometric titration curves are presented in Figure 1. From conductance data the following ion pair formation constants are calculated (6): Et4N+N03-, 5; and Ag+NOs-, 70. From these values and the activity coefficient of silver ion, calculated from the standard potentials Debye-Huckel equation with a = 5 of 0.131, o.135, and 0.1& V us. AgRE were calculated from the titration curves for chloride, bromide, and iodide, respectively. The mean value of 0.133 V agrees well with values of 0.130 and 0.131 obtained under different conditions by Kolthoff and Thomas (7) and by Senne and Kratochvil (8), respectively. Standard Potentials of Mercury Couples. The standard potential of the Hg,2+-Hg couple was 0.490 V DS. AgRE. It was calculated from potentials of the mercury electrode in
A,
(3) I. M. Kolthoff and J. J. Lingane, J . Amer. Chem. Soc., 57,2377 (1935). (4) J. F. Coetzee, Pure Appl. Chem., 13, 427 (1966). ( 5 ) J. F. Coetzee and J. J. Campion, J . Amer. Chem. SOC.,89, 2513,2517 (1967). . , (6) B. 'Kratochvil and H. L. Yeager, Top. Currelit Chem., 27, l(1972). (7) I . M. Kolthoff and F. G. Thomas, J. Phys. Chem., 69, 3049 (1965). ( 8 ) J. K. Senne and B. Kratochvil, ANAL.CHEM.,44,585 (1972).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
343
Table I. Medium Effect Values on the Molar Scale for Single Ions in Acetonitrile at 25 "C" Ion, i Ag+ Hg,*+ Hg*+ C1BrIlog -2.0 8.4 9.0 5.3 3.1 1.0 a Based on an extrathermodynamic procedure described in ref. (5) and refined in ref. (13), according to which the standard reduction potential on the molar scale of the Rb+-Rb couple is 0.20 V more positive in acetonitrile than in water, and the potential of the AgRE(AN) = $0.55 V cs. NHE(W). Aquamolality corrections included before (5, 13) have been omitted here in order to facilitate computations of solubility product constants and comparisons with results of other workers. Values for halide ions have been obtained from solubilities of thallium(1) salts (5).
Ob
dl
012
d.5 d6 017 Of8 Volume of A g NO3 odded, ml
013
d.4
019
I(0
111
Figure 1. Potentiometric titration of halide ions (X-) with silver(1) in acetonitrile CI-: 10.0 ml 0.0238MEt4NClwith 0.251M AgNOa; Br-: 10.0 ml 0.0240M Et4NBr with 0.250M AgN03; !-: 10.0 ml 0.0250M Et,NI with 0.251M AgN03. First precipitation of Et4NAgnBr3 and EtrNAg*13indicated by u, that of Et4NAgj16 by a ' , and that of AgX by b
solutions containing from 3 X to 1.4 X 10-3M mercury(1) perchlorate as well as a constant concentration of 5 x 10-3M perchloric acid, added to prevent possible solvolysis of mercury(1) ion. Under these conditions, the formal potential was 0.48? V ; the activity coefficient of mercury(1) ion was calculated from the Debye-Huckel equation with a = 5 A. The concentration dependence of the potential was consistent with virtually complete dissociation of mercury(1) perchlorate. The standard potential of the Hg2+-Hg22+ couple was 0.644 V us. AgRE. It was calculated from potentials of the platinum electrode in solutions containing a constant concentration of 5.5 x ~ o - ~ mercury(1) M perchlorate and of 5 X 10-3M perchloric acid, as well as concentrations of mercury(I1) perchlorate varying from 6 x to 7 X 10-4M. It was assumed that the activity coefficients of mercury(1) and mercury(I1) ions, and hence the formal and standard potentials of the Hg2+-Hg22- couple, are equal. Finally, the equilibrium constant for the disproportionation reaction Hg22+ Hg(1) Hg2+ can be calculated from the formal potentials (which were measured under similar conditions) of the Hg2*+-Hg and Hg2+-HgZ2+ couples. Its value of 2 x is slightly lower than that in water, 7.7 x (9), which reflects the lower ability of acetonitrile, as compared t o water, to stabilize the higher charge density mercury(I1) ion relative t o mercury(1) ion. Medium Effect Values. It is instructive to compare the above standard potentials with the corresponding values in water-uiz., 0.799 and 0.789 V us. NHE(W) for the Ag+-Ag and Hg,*+-Hg couples, respectively. It should be stressed a t the outset that since such comparisons require extrathermodynamic assumptions it is difficult to assess their reliability (5,10-13). We report here the results of one such approach,
+
(9) W. Forsling, S. Hietanen, and L. G. Sillen, Acta Chem. Scund., 6,901 (1952). (10) D. H. Berne and 0. Popovych, ANAL.CHEM., 44,817 (1972). (11) B. G. Cox and A. J. Parker, J. Amer. Chem. SOC.,94, 3674 (1972). (12) I. M. Kolthoff and M. K. Chantooni, Jr.. J . Phys. Chem., 76, 2024 (1972). (13) J. F. Coetzee, J. M. Simon, and R. J. Bertozzi. ANAL.CHEM., 41,766 (1969). 344
one that is based on a modification of the Born equation (13). It is convenient t o introduce the medium eArecr ( I O ) or solvent activity coeficient ( I I ) , myi, which is related to quantities of interest here as follows. For the standard potential of the Mn+ - M couple, the relationship is Eo(AN) - Ea(W) = (0.05916/n)log,yi
(1)
and for the solubility product constant of MX, it is
Relevant results are listed in Table I. Parker's (11)and Kolth o p s (12) estimates, which are based on a reference electrolyte assumption, lead to values that are between n and 2n units less positive for Mn+ and by the equivalent amount of between 1 and 2 units more positive for X-. Whichever set of numbers is accepted, there is a qualitative consensus that silver(1) is the only ion listed that is more strongly solvated by acetonitrile than by water. We shall use the numbers in Table I primarily for estimates involving elecrrically neutral combinations of ions, for which the results are independent of the particular extrathermodynamic assumption applied in deriving the constituent single ion values. For example, for silver chloride our numbers predict that log &(W) - log K,,(AN) = -2.0 5.3 = 3.3, while Kolthoff and 7.4 = 3.6; Chantooni's corresponding numbers are -3.8 the (relatively small) difference between the two predictions is caused, not by the different extrathermodynamic assumptions, but by differences in the two sets of experimental data on which ,-fi values have been based. Formation Constants of Silver Halide Complexes. Formation constants were evaluated by analysis of potentiometric titration curves, such as those given in Figure 1. The curves obtained with chloride ion were consistent with the formation of one complex only-uiz., AgC12-. The same is true of the initial sections of the curves for bromide and iodide. From the value of the standard potential of the Ag+-Ag couple, assuming complete dissociation of Et4N+X- and Et4N+AgX2-, overall (activity) formation constants, p2, for the AgX2- complexes were calculated at various points on the titration curves. Mean values of log p2 (standard deviation +0.1) are listed in Table 11. The titration curves for bromide and iodide in Figure 1 are consistent with the formation of an additional complex, Ag2X3-, and in the case of iodide it appears that at least one further complex, Ag,16-, is also produced. Formation of these same complexes also has been observed in nitromethane (18) and propylene carbonate (19), and, as here, no evidence was obtained for the existence of the intermediate polynuclear complexes Ag&- and Ag415-. This puizling fact remains unexplained. Since the shape of a titration curve does not
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
+
+
Table 11. Comparison of Equilibrium Constants on the Molar Scale for Reactions of Silver(1) with Halide Ions in Nonaqueous Solvents and Water -log K O log P 2 log K82 Solvent C1BrIC1BrIc1BrIRef. A. Activity Constants at 25 "C Acetonitrile 13.0~ 13.3b 15.P 13.7 14.1 15.6 0.7 0.8 0.6 0 Sulfolane (30 "C) 18.5 18.9 ... 20.3 20.2 ... 1.8 1.3 ... (14) Propylene carbonate 20.2 ... ... 21.2 ... ... 1.0 ... ... (15) Dimethylformamide 15.2 ... ... 17.0 ... ... 1.8 ... ... (16) Water 9.7 12.2 16.0 5.0 7.2 10.6' -4.7 -5.0 -5.4' (17) B. Concentration Constants near 25 "C 14.6 0.2 12.6 13.4 0.2 13.2 14.2 Acetonitriled 12.4 0.4 (2) 19.7 22.0 0.3 0.0 20.5 19.5 19.7 Nitromethaned 19.2 1.5 (18) 23.5 1.1 0.7 22.2 22.5 21.8 22.6 Nitroethaned 21.1 0.9 (2) 22.8 0.9 0.7 20.9 21.2 20.5 21.8 Propylene carbonated 20.0 1.0 (19) 13.1 11.9 11.7 1.5 1.1 12.0 10.4 10.6 Dimethylsulfoxided 1.1 (2) 22.2 0.3 1 .o 16.7 19.7 18.7 20.9 Acetone" 16.4 1.3 (4 -4.3 14.8 -.5.0 8.0 10.9 15.2 18.2 Methanol/ 13.0 -3.4 (2) Direct experimental value from potentiometric titrations. -log KdAN) = -log b Indirect values calculated from medium effect numbers in Table I and solubility product constants in water: + log ,,&yx-.On this basis, the result for chloride ion is 13.0, in exact agreement with direct experimental value. &(W) + log e In 4 M NaC104. In 0.1M EtaNC1O4. e In 0.1 M LiC104. f In 1M LiC10,. This work. 0
Q
necessarily provide unambiguous evidence for the existence of a particular species, certain confirmatory experiments were carried out. The white precipitate which appeared when between 0.28 and 0.64 mole of silver was added per mole of EtlNBr was analyzed as follows. It was filtered off, washed with acetonitrile and then added to water, after which the supernatant liquid was titrated potentiometrically with aqueous silver nitrate solution. Finally, the total amount of silver bromide was filtered off and weighed. Exactly 3 moles of silver bromide was found per mole of silver ion used in the titration, in agreement with the following scheme : Et4NAg2Brs(s)% 2AgBr(s) Br-
+ Et4N++ Br-
+ Ag+ --+AgBr(s)
(3)
From the point on the titration curve at which precipitation of (white) EtrNAgzXafirst occurs, the maximum possible value of the concentration solubility product constant (i.e., if Ag2X3- were the only complex present, which in fact is not the case, and also if no supersaturation occurs) is found t o be 8x for Et4NAgzBr:3 and 6 X 10-j for Et4NAg213. However, the actual values are considerably lower. The solubility of Et4NAg213,which was precipitated when between 0.21 and 0.65 mole of silver was added per mole of Et4NI, was measured directly as follows. It was washed with acetonitrile and then equilibrated with fresh acetonitrile. Aliquots of the saturated solution were added to excess water and then titrated potentiometrically with aqueous silver nitrate solution. If Ag21a- is sufficiently stable in acetonitrile so that negligible (14) R. L. Benoit, A. L. Beauchamp, and M. Deneux, J. Phys. Chem., 73, 3268 (1969). (15) J. N. Butler, ANAL.CHEM., 39, 1799 (1967). (16) J. N. Butler, J. Phys. Clzem., 72, 3288 (1968). (17) L. G. Silltn and A. E. Martell, "Stability Constants of MetalIon Complexes," The Chemical Society, London, 1964. (18) J. Badoz-Lambling and J-C. Bardin, C . R. Acad. Sci., 266, 3288 (1968). (19) J. Courtot-Coupez and M. L"er, Bull. SOC.Chim. Fr., 1969, 675.
disproportionation into silver iodide and iodide ion occurs, then the reaction scheme will be analogous t o that given in Equation 3, from which a concentration solubility product constant of 5 x 10-6 is calculated. Essentially the same value was obtained from calculations based on a comparison of the point of first precipitation of Et4NAg213from 2.5 X 10-2 and 9.3 x 10-2M Et4NI, which occurred, respectively, at 0.21 and 0.08 mole of silver added per mole of Et4NI. Using this value for the solubility product constant in the analysis of the titration curve in Figure 1, a n approximate value of log p32 = 30.3 i 0.5 is estimated for AgZI3-. It follows that, in the titration of 0.025M Et4NI with 0.25M AgN03, precipitation of Et4NAg213first occurs when the concentrations of AgJ- and Ag12- are 2 X and 5 X lO-3M, respectively. The corresponding concentrations of the bromo complexes appear to be of the same order of magnitude. The remaining section of the titration curve for iodide could not be analyzed because, due to disproportionation of Ag,16-, the solubility of its tetraethylammonium salt could not be measured directly. N o precipitation of Et4NAg2C1, occurred under our conditions; undoubtedly the reason is partly that the solubility of this salt is higher, and partly that the stability of Ag2CI3-is lower, than is the case for the bromo and iodo complexes. The point of first precipitation of silver chloride (at 0.453 mole of silver added per mole of chloride) is in exact agreement with that calculated from the Ks2 value given in Table 11, where AgX(s)
K,z
=
+ X-
aa,xz-/ax-
AgXz=
Pz/Kso
(4)
Finally, we wish t o point out that our original analysis ( I ) of the titration curves for bromide and iodide was in part incorrect. It was erroneously concluded from the shapes of the curves that during the initial parts of the titrations the predominant species were AgzXs- and not AgX2-. It is now evident that the curves owe their shape more to the insolubility of Et4NAg2X3than to any unusual stability of AgzX3-.
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2 , FEBRUARY 1973
345
Table 111. Comparison of Equilibrium Constants for Reactions of Mercury(1) and Mercury(I1) with Halide Ions in Acetonitrile (AN) and Water (W) at 25 "C log K AN W Reaction K X -37.3 -17.7a c1 -37.4 -22.2a Br I
none
c1
Br I
c1 Br I
c1 Br
Ka
I
Pa
c1 Br I
c1 Br
(KsOKdpP$l
I
c1 Br I
-39.3 -2.7 35.1 35.4 38.2 6 7 8 4.2 3.8 2.8 45.5d 46.7d 48. 7d 4.9c 4.76 3.8a 1.3 2.8 3.9 5.5 6.6 6.7
-28.3. -2.lb 13.2~ 17.3c 23.8< 0. 8c 2.4c 3.8" 1. o c 1.3~ 2.2c 15. l c 21.0c 29.8" 6.6 7.0 6.6 -5.8 -4.6 -2.8 -4.7 -3.3 -0.6
Mean of several values (corrected to I = 0) listed in ref. (17). Ref. (9). In 0 . 5 k NaCIOa; ref. (20). Measured directlv: values of K , were calculated from B4 and are more uncertain; values of K 3 were calculated from f13 and the indirectly measured /32 and are still more uncertain (see text). e Measured directly. 5
Solubility Product Constants of Silver and Mercury(1) Halides. From the values of Eo(Ag+-Ag) and of 82, values for the solubility product constant of silver chloride were calculated at various points on the titration curve where a precipitate of silver chloride was present. The mean value of log Kso (standard deviation kO.1) is given in Table 11. Analogous calculations for silver bromide and iodide would require knowledge of the solubility product constants of Et4NAg2Br3and Et,NAg& and of the formation constants of the corresponding complex anions. The log K,o values listed in Table I1 for silver brorn.de and iodide were calculated from the medium effect numbers given in Table I, and have an estimated uncertainty of 3=0.2. The same procedure was followed to estimate the log K,o values listed in Table I11 for the mercury(1) halides, in this case with a rather large estimated uncertainty of *0.5 unit. Formation Constants of Mercury(I1) Halide Complexes. Values of the overall formation constants p4 and p 3 were calculated by analysis of the potentiometric titration curves of (typically) ca. 10-*M Et4NX with 10-lM mercury(I1) perchlorate in one set of experiments, and of ca. 10-*M HgX2 with 10-IM Et4NX in a second set. In the first set of titrations, large breaks corresponding to the formation of HgX2 were obtained, and in the second set well-defined but smaller breaks for the formation of HgX3- and HgXa2-. Values of ,B2could not be determined from the titration curves because in those regions of the curves where appreciable initial concentrations of HgX2 were present the following slow reaction occurred at the electrode: HgX2 346
+ Hg(1)
f-
HgrXz(s)
(5)
Some of the mercury(1) halide produced coated the electrode which became sluggish and acquired a time-dependent potential. Consequently, pZ values were determined by equilibrating 2 x and 5 X 10-aM solutions of HgX2 with mercury metal by shaking for 48 hours, and then determining the concentration of HgX2 left by atomic absorption. The equilibrium constant of reaction 5 is given by
The same final concentration of HgXz was found for the two initially different concentrations, indicating that equilibrium indeed had been established. The estimated uncertainty in the log pz values listed in Table 111 is determined mainly by that in the log Ksovalues, is.,*0.5 unit. For the case of Hg12, pz could also be estimated by an independent method. The solubility of (red) Hg12 in acetonitrile was found to be 7.0 X 10-aM, and in water it is 7.4 x lO+M (20); hence, logmyKgII= -2.0. Using this number and the corresponding values for Hg2f and I- in Table I, and also log /32 25 at in38 for acetofinite dilution in water (20), one obtains log p2 nitrile, in agreement with the number in Table 111.
--
DISCUSSION
General trends in the equilibrium constants for silver listed in Table I1 can be understood, at least qualitatively, in terms of the following hypotheses for which significant experimental support exists. 1) The Gibbs free energy of solvation of silver(1) ion becomes smaller (less negative) in the solvent (20) L. G. Sillin, Acta Chem. Scand., 3, 539 (1949).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
-
order dimethylsulfoxide > acetonitrile > dimethylformamide water > methanol > acetone sulfolane > propylene carbonate > nitromethane > nitroethane (11). 2) The free energy of solvation of chloride ion becomes smaller in the order water > methanol > nitromethane > dimethylsulfoxide acetonitrile dimethylformamide sulfolane propylene carbonate > acetone. For bromide and especially for iodide ion, while the order remains substantially the same, the spread in solvation energies is reduced. 3) As shown by Equation 4, the KSzvalues reflect the differences in the solvation energies of X- and AgXZ-. It is reasonble to conclude, as has been done before by others (2, 14, 15), that the dipolar aprotic solvents exhibit a generally moderate tendency to stabilize AgX2- (probably through a dispersion interaction with these relatively polarizable ions) relative to X-. In sharp contrast, in protic solvents X- is stabilized (undoubtedly through hydrogen bonding) relative to the lower charge density AgX2-. The fact that in dipolar aprotic solvents AgXz- is stabilized relative to X- has important analytical implications. Many of the most reliable thermodynamic data for water, the alcohols, and several other solvents have been obtained with internal silver-silver halide electrodes in cells without transference, whereby the uncertainties associated with liquid junction potentials were avoided. A major handicap in studying similar systems in dipolar aprotic solvents is that silversilver halide electrodes are unsuitable for internal use in such solvents, because of the large amount of silver going into solution as complex ions. For example, it follows from the Ks0value given in Table 11that the solubility of silver chloride in acetonitrile containing a total chloride ion concentration of 0.1M is O.O8M, which is far too high to permit internal use of the silver-silver chloride electrode. [It appears from our previous results (5) that thallium amalgam-thallium(1) chloride electrodes should be suitable as internal electrodes in acetonitrile.] Naturally, the use of this electrode as an external reference is not obviated by these considerations. The relative insolubility in acetonitrile, propylene carbonate, and nitromethane of the tetraalkylammonium salts of complex ions such as Ag2Br3-, Ag&-, and A g J - also has analytical implications. If (external) silver bromide or iodide reference electrodes are prepared from tetraalkylammonium halides, the solid phase may slowly change from AgX to R4NAg,X,+]. It is perhaps not generally realized that even in water, silver-silver halide electrodes are susceptible to similar specific interferences. For example, the emf of the folloning aqueous cell
>
-
-
-
-
1
Ag,AgI Et4NI(c)1 KCl(satd)1 1 KUc) 1 Ag1,Ag ~
was found to be much greater than could be attributed to differences in activity coefficients in the two solutions. This anomaly was shown (21) to be the result of conversion of (21) R . FernBndez-Prini and J. E. Prue, J . Phys. Chem., 69, 2793 (1965).
silver iodide in the left-hand side of the cell into a new solid phase, Et4NAgz13. It should be noted that the complex ion Ag&- is much less stable in water than in acetonitrile, but that in the presence of tetraethylammonium ion, complexation is promoted by the insolubility of the salt produced. Turning now to the results for mercury in Table 111, some of our data can be compared with those of Ellendt and Cruse (22) who calculated from conductimetric titrations of HgXz with X- log K3 values of 6.0 for all three halides and log Kd values of 2.2, 2.0, and 0.6 for chloride, bromide, and iodide, respectively. The lineage of our log K5values is such (Table 111, footnote d) that their estimated uncertainty is rather high (*1 unit), but agreement with the results of Ellendt and Cruse is good for chloride, less so for iodide ion. Our log K4 values are uniformly larger by ca. 2 units than those reported by Ellendt and Cruse, but they follow the same order of C1- > Br- > I-, which is the opposite of that observed in water, and also the opposite of the order of K3 and values in both acetonitrile and water. The last two equilibria given in Table I11 are of interest in connection with the use of mercury-mercury(1) halide reference electrodes. Over 30 years ago, Ulich and Spiegel (23) studied the behavior of a number of electrodes of the second kind in various solvents and observed that mercury(1) chloride and iodide were solubilized in acetonitrile solutions containing the corresponding anion. Cruse et al. (24) obserked that the potential of mercury-mercury(1) halide electrodes in acetonitrile drifted with time and attributed that to the slow formation of mercury(I1) halide complexes; in fact, a solid compound Et4NHgBr3 could be isolated. As already pointed out, we also found that mercury electrodes coated with mercury(1) halides acquired time-dependent potentials in the presence of halide ions. The equilibrium constants in Table I11 show that disproportionation of mercury(1) is promoted by complexation of mercury(I1) to a much greater extent in acetonitrile than in water. It is concluded that mercury electrodes of the second kind involving mercury(1) compounds quite generally are likely to be unsuitable for internal use in acetonitrile and other relatively inert dipolar aprotic solvents; furthermore, if such electrodes are used externally, sufficient time must be allowed after preparation to ensure equilibrium. Finally, silver(I), mercury(I), thallium(I), and all other electrodes of the second kind involving halide ions cannot be used in solutions of relatively strong acids (e.g., HClO,, HSbCls) in acetonitrile, sulfolane, propylene carbonate, nitromethane, or other solvents in which the hydrogen halides are weakly dissociated. RECEIVED for review August 28, 1972. Accepted October 25, 1972. We thank the National Science Foundation for financial aid under Grant Numbers GP-1479 and GP-16342. (22) G. Ellendt and K. Cruse, Z . Physik. Chem., 201, 130 (1952). (23) H. Ulich and G. Spiegel, ibid., A177,103 (1936). (24) K. Cruse, E. P. Goertz, and H. Petermoller, Z . Elektrochem., 55,405 (1951).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
347
