Electrochemistry of Mercury in Molten Lithiurn Chloride-Potassi um Chloride Eutectic K. W. Hanck and M. Lynn Deanhardt Department of Chemistry, North Carolina State University, Raleigh, N.C. 27607
THE ELECTROCHEMISTRY of mercury in molten LiCl-KCl was first studied by Laitinen, Liu, and Ferguson (1). Mercury(1) was found to disproportionate in the melt at 450 "C to mercury(I1) and mercury(0). A single reduction wave for mercury(I1) was observed but it was complicated by saw-toothlike fluctuations which were attributed to discharge of gaseous mercury at the electrode surface. Laitinen and Liu estimated the formal potential of the mercury(II)/mercury(O)couple to be -0.5 V os. 1 M Pt(II)/Pt from the voltammetric half wave potential of the mercury(I1) reduction (2). Hames and Plambeck observed two 1-electron steps for the reduction of HgCh in molten A1Cl3-NaC1-KC1 at 150 "C (3). Mamantov and coworkers observed three 2-electron reduction steps for Hg(I1) in acidic A1CI3-NaCl melts at 175 "C ( 4 , 5 ) . The first step produced Hg,z+, the second produced a new ion, Hg3z+,and the third yielded metallic mercury. The objective of this paper is to report the results of an extensive investigation of the reduction of Hg(I1) and Hg(1) in molten LiCl-KCl. The complicating effects observed by previous workers were minimized by making all measurements on a time scale of 1-2 seconds or less. Application of a recently developed model for EC processes in which the product is insoluble will also be described (6).
0.72 0.54 0.90
0.36
-
0.18
-
O.l8
1 0
425
lime
0.50
0.75
1.00
(Sec.1
Figure 1. Typical chronopotentiograms of Hg(I1) [Hg(II)J = 11.73 X lOW3M. T = 450 "C Curve A . Cathodic chronopotentiogram; I J / A = 9.81 mA/
cm2 Curve E . Chronopotentiogramwith current reversal; I f / A = I,/A = 9.81 mA/cm2
EXPERIMENTAL
A Lindberg Hevi-Duty Type 3012-S vertical three-zone split tube furnace was used. A solid state proportional temperature controller (Model MXC-5 with MXP-1 power pack; T & T Controls Company, Media, Pa.) was employed to regulate the furnace temperature. The chronopotentiometric constant current source was based on conventional operational amplifier circuitry and employed P65AU, P25AU, and P66B Philbrick/Nexus solid state operational amplifiers. A Tektronix 502-A oscilloscope equipped with a Model C-12 oscilloscope camera was employed for recording chronopotentiograms and linear sweep voltamrnograms. Pulse polarograms were obtained with a Princeton Applied Research Corporation Model 174 Polarographic Analyzer. The output from a triangle wave generator constructed from the circuit of Bull and Bull ( 7 ) was inputed into the external input of the Model 174 to obtain linear sweep voltammograms. The construction and area determination of the platinum flag electrodes have been previously described (8). The (1) H. A. Laitinen, C. H. Liu, and W. S . Ferguson, ANAL.CHEM., 30, 1266 (1958). (2) H. A. Laitinen and C . H. Liu, J. Amer. Chem. Soc., 80, 1015 (1958). (3) D. A. Harnes and James A. Plarnbeck. Can. J. Chem., 46, 1727 (1968). (4) G. Torsi and G. Mamantov, biorg. Nucl. Chrm. Lett., 6 , 843 (1970). (5) G. Torsi, K. W. Fung, G. M. Begun, and G. Marnantov, lnorg. Chem., 10,2285 (1971). (6) K. W. Hanck and M. Lynn Deanhardt, ANAL.CHEM., 45, 179 (1973). (7) R . H. Bull and G . C. Bull, ibid., 43, 1342 (1971). (8) K. W. Hanck and H. A. Laitinen, J . Electrochem. Soc., 118, 1123 (1971). 176
Pt(II)/Pt reference electrode was electrogenerated at constant current from a 2 cm2 platinum foil using a Sargent Model IV Coulometric Current Source. The molten salt cell assembly has been previously described ( 9 ) with the exception that the 71/60 standard taper joints were replaced by 102/75 O-ring joints. The procedures used for cleaning glassware and handling the molten eutectic have been described in detail (9). The LiCl-KC1 eutectic was obtained from Anderson Physics Laboratories, Inc., Urbana, Ill. Reagent grade HgC12 and Hg2C12(J. T. Baker Chemical Company) were vacuum dried over Mg(ClO& a few days before use. RESULTS AND DISCUSSION
Cathodic Chronopotentiometry. A typical cathodic chronopotentiogram of Hg(I1) is shown in Figure 1. The validity of the Sand equation was tested by making duplicate runs of several current densities at six different Hg(I1) concentrations using a fresh electrode for each run. The results are summarized in Table I and indicate that the reduction obeys the Sand equation with a value of 700 i 27 A sec1'2cm/mole for Ior1' 'IC. Plots of E m . log[(r1'2 - t1'2)/t1'2] were linear; the average slope of some 25 plots was 74.0 + 8.2 mV. The theoretical slope for a reversible 2-electron process at 450 "C is 71.7 mV. The quarter wave potential was found to be independent of applied current density and Hg(I1) concentration over the range studied; the average value being -0.622 i 0.006 V cs. the 1 M Pt(II)/Pt reference electrode. (9) K . W. Hanck, Ph.D. Thesis, University of Illinois. Urbana, Ill.. 1969.
ANALYTICAL CHEMISTRY, VOL. 45, NO. 1, JANUARY 1973
I
1
Table I. Variation of 1 , ~ with ~ ' Hg(I1) ~ Concentration tHdII)I (mmoles/l.)
I - A
1oT112(mA sec1'2/cm2) Range of Io(mA/cm2)
1.876 3.962 7.795 9.327 10.06 15.93
1.401 i 0.0670 2.818 i 0.112a 5.118 i 0.201a 6.275 =k 0 . 1 6 0 ~ 6.633 It 0.149a 11.33 i 0.35a
-
1.318-2.770 2,405-4.673 5.263-7.774 5.283-9.785 5.783-9.488 11.67 -19.77
+/(
All uncertainties quoted in this paper are 95% confidence intervals. Q
I
I
I
4.0
i E
2.0
I
I
I
I
I
I
02
03
04
05
06
07
-E,
VOLTS
Figure 3. Cyclic voltammograms of Hg(II)
- 0.4
-0.1
-
[Hg(II)] = 4.103 X 10-aM Curve A . Y = 0.968 V/sec; A = 0.0306 cm2 Curve B. Y = 25.3 V/sec; A = 0.0252 cm2
0.e
E , volts
Figure 2. Normal pulse polarogram of a 5.025 X lO+M solution of HgC12 Points are experimental; curve is theoretical and is based on a 2electron reversible process with D = 1.67 X cm*/sec, = -0.479 V cs. 0.010M Pt(Il)/Pt reference electrode and A = 0.513 cm2
The diffusion coefficient of Hg(I1) based on a 2-electron process is 1.67 k 0.13 x 10-5 cm2/sec. This compares favorably with those of Cd(I1) (1.68 X and Pb(I1) (2.18 X The chemical entity which diffuses to the electrode is a chloro complex of mercury, probably HgCI42-. The quarter wave potential and value of Z071/2/Cfor the reduction of mercurous chloride were identical, within experimental error, to the values found for the reduction of mercuric chloride. This observation is consistent with the hypothesis stated earlier by Laitinen, Liu, and Ferguson that Hg(1) rapidly disproportionates to Hg(I1) and Hg(0). Anodic chronopotentiograms of mercurous chloride were identical to those of pure eutectic; consequently we conclude that no Hg(1) species are present in detectable concentration and that Hg(1) rapidly disproportionates to Hg(I1) and Hg(0). Pulse Polarography. The normal pulse polarogram of a 5.025 X 10-3M solution of HgCl? is presented in Figure 2. The agreement between the theoretical curve and observed points is excellent, leading us to conclude that the reduction of HgCI? is reversible, involves the transfer of 2 electrons, and produces a soluble product. Linear Sweep Voltammetry. Voltammograms of HgClz undergo a change in shape as a function of sweep rate (Figure 3). At slow sweep rates, the anodic peak has the sharp, nearly symmetrical shape characteristic of the stripping of an insoluble deposit, while the cathodic peak has the shape expected for a 2-electron reversible process yielding a soluble
~
Table 11. Cyclic Voltammetry of HgCIz [HgCli] = 4.10 X 10-3M; all potentials cs. 1M Pt(II)/Pt
Scan rate (Visec)
( E d c(mv) -661 -655 -657 -661 -655 -660 -655
(Ep
- E p / 2 ) c a (Ep - E p / d a '
(mv)
(mv)
(Ep)e
- (Ep)cb
(mV)
15.2 75.9 -67.2 18.2 75.7 -68.5 -70.0 35.4 68.6 -71.9 35.0 68.5 50.6 68.9 -68.9 50.2 71.9 -68.6 -68.6 50.3 72.9 a Theoretical value for 2e- reversible process at 450 ' C (sol. product) = 68.5 mV. *Theoretical value for 2r- reversible process at 450 "C = 69.1 mV. 0.54 0.97 2.33 5.18 10.1 12.9 25.3
product (Table 11). As the sweep rate is increased, the anodic peak broadens and more closely approaches the shape expected for the reversible reoxidation of a soluble product; the peak potential and shape of the cathodic peak remain unchanged. These results indicate that the reduction product is initially soluble but depending on the time allowed before reoxidation, the reduction product may precipitate onto the electrode before it is reoxidized. In the limit of an extremely slow sweep ra+te, no evidence for solubility would be observed and the cathodic peak should have the narrow shape associated with the deposition of an insoluble product of unit activity (10). (10) Gleb Mamantov, D. L. Manning, and J. M. Dale, J . Electroaim/. Chern., 9,2531 (1965).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 1, JANUARY 1973
177
Table IV. X-Ray Diffraction Data for PtHg,
Electrolysis of HgCl@
4A) b
A
''
BBBBm8
2.173 1,262 1.382 1.543 1.095 1.781 1.649 1.318 3.057 2.507
e A
e
8
a
It I o 100 I00 95 90 90 80 80 75 60 50
Dipping pr0cedurea.c
4A.)
I1I o
2.18 1.26 1.38 1.54 1.09 1.78 1.65 1.32 3.07 2.51
vs vs S S S S
m m S
m
Only the tea strongest lines are listed. body centered cubic lattice. Data of Robbins and Enke (13).
a. = 6.204 A ; 1 .o
20
30
tf , rec.
Figure 4. Variation of +Jtf with tf [Hg(II)] = 0.218M
e e w e
2.00
0
=
A
= 1.50 = 1.00
Table 111. Analysis of Current Reversal Data Uncorrected for Solubility 0 kf(sec- 1) k j t , intercept 1 .o 1.5 2.0
0.124 0.124 0.123
0.133 0.103 0.078
Corrected for Solubility 1.0 1.5 2.0
0.126 0.124 0.122
-0.0279 -0.022 -0.017
No such peaks were observed because thermal convection in the melt prevented the collection of reliable data at sweep rates below 0.5 V/sec. The half wave potential calculated from the average value of (Ep)cwas -0.623 V, in excellent agreement with the pulse polarographic and chronopotentiometric results. Current Reversal Chronopotentiometry. In view of the voltammetric results, there is little doubt that the reduction of HgCh is a reversible process yielding a soluble product (at least initially) ; a number of chronopotentiometric experiments were run to further study the soluble-insoluble transformation tacitly suggested by the linear sweep results. A typical reverse current chronopotentiogram is shown in Figure 1. The difference between E 0 . 2 ~and 0 ~ Eo.215sr is greater than that expected for a reversible process (11). Most of this difference is due to iR drop effects similar to those observed by Delahay and Mattax (12). A series of experiments were performed in which the forward electrolysis time ( t i ) and the forward current ( i f )were chosen so that the forward charge was constant throughout a run but t , was varied from ca. 0.4 to 4 seconds. The reverse current (ir)was related to the forward current by the experimentally known ratio O(O = if/ir). The forward charge was also chosen ( 1 1 ) W. E. Palke, C. D. Russell, and F. C. Anson, ANAL.CHEM., 34, 1171 (1962). (12) P. Delahay and C. C. Mattax, J . Amer. Chem. SOC.,76, 874 (1 954).
178
to be as high as experimentally feasible (4 x l o 4pC) in order to minimize the solubility of the reduction product. The results of these experiments are presented in Figure 4. The consistent decrease in r r / t fas tf is increased suggests the presence of a chemical reaction following charge transfer. This reaction must involve an insoluble reduction product since r r / t fis greater than that expected for the reoxidation of a soluble product. We have recently described the mathematical relationships which describe the current reversal behavior of a system in which a reversible charge transfer reaction producing an insoluble product is followed by a firstorder chemical reaction involving the insoluble product (6). If such a mechanism is present, r r / t fwill be given by :
where k , is the first-order rate constant and 6 is if/&. The data points in Figure 4 were converted to kftf values using Equation 1. A plot of k,t, cs. t , was made for each data set; the plots should have a zero intercept with a slope of k,. The experimental plots were linear; their slopes and intercepts are tabulated in Table 111. The non-zero intercepts result from the partial solubility of the reduction product (6). Using the procedure outlined earlier (6),the fraction of the forward charge producing soluble product was estimated from the k t, intercept; an average value of 1 3 . 4 z soluble was obtained. Each data point was corrected for solubility using a procedure similar to that described previously (6); the apparent k , used in this correction was 0.0 sec-1 since the chemical reaction is believed to occur only in the deposit. Plots of kftf cs. tf using the solubility corrected data were linear with negligible intercepts (Table 111). It may be noted from Table 111 that the corrected rate constants appear to decrease as O is increased. This decrease is only an apparent one and results because of partial solubility; the results reported earlier (6) indicate the k , at '6 = 1 to be closest to the true value. Identification of Reaction Product. A 2-cm length of 26gauge Pt wire was electrolyzed in an HgC12 solution at a controlled potential of -0.7 V for 10 minutes. The wire was washed with deionized water, air dried, and mounted in a Debye-Scherrer X-ray diffraction camera. An 18-hour exposure to Ni filtered Cu K, radiation produced a powder diffraction pattern identical to that of PtHg, (Table IV). Robbins and Enke found that PtHg, is formed at room temperature when a clean Pt wire is allowed to stand in contact with mer -
ANALYTICAL CHEMISTRY, VOL. 45, NO. 1, JANUARY 1973
,
cury for several days (13). The role of PtHg and PtHg2in the reaction between platinum and mercury is probably small since they disproportionate at 180 and 240 “C, respectively (14).
Electrode Mechanism. The electroreduction of HgClz proceeds by a reversible, two-electron process to mercury(0). Based on our measurements of E,,*, the formal potential of the Hg(II)/Hg couple is -0.622 V cs. 1 M Pt(II)/Pt. At 450 “C, mercury should be in the vapor state. Most gases have a finite solubility in molten LiC1-KC1 and mercury vapor is not an exception (15). Attempts to measure the mercury(0) concentration of a 0.09M HgC12solution to which a droplet of mercury had been added were unsuccessful. We conclude, therefore, that the equilibrium solubility of mercury vapor in the melt is less than 10-3M which is the minimum concentration which could be reliably detected. After the electrolysis has produced enough mercury to saturate the diffusion layer, mercury deposits onto the electrode where it reacts with platinum producing PtHg, as the ultimate product. The production of some PtHg, probably begins when the electrolysis is initiated. The presence of traces of PtHg, on the electrode surface may be responsible for the deposition of mercury(0) onto the electrode, thus avoiding -
(13) G. D. Robbins and C. G. Enke, J . Electrouml. Chem., 23, 343 ( 1969). (14) Gerhard Jangg and Frany Steppan, 2. Metullk., 56, 172 (1965). (15) George J. Janz, “Molten Salts Handbook,” Academic Press, New York, N.Y., 1967.
the production of bubbles of mercury vapor with subsequent irratic distortion of E-t and i-E curves. The peculiar property of PtHg, to coherently bind mercury(0) to platinum has been cited by Robbins and Enke (13) and by Hartley, Hiebert, and Cox (16) in their studies of platinum-based thin mercury film electrodes. The reoxidation process is clearly that of mercury(0) rather than PtHg,. First the potential corresponds closely to that expected for the reoxidation of mercury(0); in aqueous solution it has been well established that the oxidation of PtHg, occurs several hundred millivolts positive of the oxidation of metallic mercury (13, 16). If the reoxidation process were that of PtHg,, r , / t f should increase rather than decrease as t f is made longer because of the longer time available to produce PtHg,. The rate constant obtained for the reaction between mercury and platinum should be regarded as a pseudo-first-order rate constant. The dependence of the rate constant on such parameters as HgC12 concentration and electrode area were not explored but may influence the rate. RECEIVED for review November 11, 1971. Accepted August 11, 1972. Presented in part at the 163rd National Meeting of the American Chemical Society, Boston, Mass., April 9-14, 1972.
(16) A. M. Hartley, A. G . Hiebert, and J. A. Cox, J . Electround. Chem., 17,81 (1968).
Current ReversaI Chronopotentiometry of EC Processes Involving an Insoluble Product K. W. Hanck and M. Lynn Deanhardt Department of Chemistry, North Carolina State University, Raleigh, N.C. 27607 CURRENTREVERSAL CHRONOPOTENTIOMETRY has been theoretically applied to many electrode mechanisms (1-12). Most of the theoretically derived relationships have been (1) 0. DraEka, Collect. Czech. Chem. Commw., 25, 338 (1960). (2) A. C. Testa and W. H. Reinmuth, ANAL.CHEM.,32, 1512 (1960). (3) C. Furlani and G. Morpurgo, J . Electrouuul. Cliem., 1, 351 (1960). (4) Harvey B. Herman and Henry N. Blount, ibid., 25, 165 (1970). (5) F. H. Beyerlein and R. S. Nicholson, ANAL.CHEM.,40, 286 (1968). (6) T. Berzins and P. Delahay, J. Amer. Chem. SOC.,75, 4205 (1 953). (7) H. B. Herman and A. J. Bard, J . Phys. Cliem., 70, 396 (1966). (8) Henry N. Blount and Harvey B. Herman, J . Electrochem. Soc., 117, 504 (1970). (9) J . H. Christie and George Lauer, ANAL. CHEM.,36, 2037 (1964). ( I O ) S. W. feldberg and Clemens Auerbach, ibid., p 505. (11) J. W. Bixler and Stanley Bruckenstein, ihid., 37, 791 (1965). (12) D. J. Macero and L. B. Anderson, J . Electrour7ul. Chem., 6 , 221 (1963).
experimentally verified. Although chronopotentiometry has certain limitations (13, 14), the current reversal technique is a convenient tool for determining whether the product of an electrode reaction is soluble or insoluble, or whether it is participating in a following chemical reaction. The existing models for mechanisms involving follow up chemical reactions require that the product of the forward electrolysis be soluble and that it undergo a homogeneous chemical reaction. Such an assumption frequently cannot be made when studying electrode reactions in molten salt media (15, 16). It is our purpose to present a model for an EC process in which the product of the charge transfer step is insoluble and is transformed by a first-order chemical reaction to an electroinactive product. The expected effects of partial solubility will also be presented. The application of the
(13) P. J. Lingane, CRC Crit. Rec. A d . Clzem., 1, 587 (1970). (14) R. S. Nicholson, ANAL.CHEM.,44, 478R (1972). (15) H . A. Laitinen and R. D . Bankert, ibid., 39, 1790 (1967). (16) K . W. Hanck and H. A. Laitinen, J . Electrochem. SOC.,118, 1123 (1971).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 1, JANUARY 1973
179
