Anodic Stripping Voltammetry of Mercury(l1) at the Graphite Electrode S. P. PERONE and W. J. KRETLOW Department of Chemistry, Purdue University, lofoyette, Ind.
b The extension of stripping analysis to the determination of trace quantities of mercuric ion has been investigated. The wax-impregnated graphite electrode was used in these studies. A thiocyanate medium was selected for best results. The concentration range studied extended from 1.0 X to 4.0 X 10% mercury(l1). Mercury was deposited at the graphite electrode with stirring at constant potential. The deposit was stripped by anodic voltammetry with linearly varying potential. Clearly defined stripping peaks were obtained. Multiple peaks were observed at intermediate concentrations. The total quantity of electricity involved in the stripping peak(s) was directly proportional to concentration and electrolysis time over the entire concentration range studied.
T
HIS WORK describes the application of stripping analysis to the trace determination of mercury(I1). The most sensitive previous voltammetric method reported for the determination of mercury(I1) involved the dropping mercury electrode (2). The maximum sensitivity reported was 5 X 10-6-If. The sensitivity of a voltammetric determination where an amalgam or deposit is formed can be increased greatly by employing stripping techniques (9). The utilization of the conventional approach, involving the hanging mercury drop electrode, for the stripping analysis of mercury(I1) is clearly an impossibility. The use of a platinum electrode did not appear promising, since the complicating effects of even slight amalgamation would probably interfere. Since the successful use of graphite electrodes in the study of anodic stripping voltammetry of noble metals had been demonstrated recently (3, 6), a graphite electrode was selected for use in the work reported here. The stripping voltammetry at the wax-impregnated graphite electrode of solutions of mercury(I1) as dilute as 4.0 X l O - g M was investigated in thiocyanate media. EXPERIMENTAL
Apparatus. Two voltammetric instruments were used in this work. One was the Sargent Model FS
968
ANALYTICAL CHEMISTRY
Polarograph (E. H. Sargent Co., Chicago, Ill.). The other instrument was a general-purpose all-electronic device utilizing operational amplifiers manufactured by G. A. Philbrick Researches, Inc., Boston, Mass. The pertinent characteristics of both instruments have been discussed previously (6-8). With the FS Poiarograph a two-electrode system was used; with the operational amplifier instrument, a three-electrode system was used. The recorder used in conjunction with the operational amplifier instrument was an Esterline Angus, Speed Servo, S601S, 1-mv., 1/8-second recorder. The planimeter used to determine peak areas was a Keuffel and Esser compensating polar planimeter, Yo. 62 0000. Cells and Electrodes. The working electrodes for all experiments were prepared from 12-inch spectroscopic graphite electrodes (Xational Carbon Co., N . Y., No. L4309). The preparation procedure involved impregnation with paraffin wax, insulation of the sides, and polishing of the plane circular exposed electrode tip. The details of the procedure have been described previously (6). The electrode surface area was approximately 0.32 sq. cm. The reference electrode was a large saturated calomel electrode isolated from the cell by a salt bridge containing 0.1M KCNS between two ultrafineporosity sintered-glass disks. The solution in the salt bridge was replenished periodically to minimize cross-contamination. The total cell resistance was of the order of 300 to 500 ohms, depending on electrolyte concentration. When a three-electrode system was used the counter electrode was a 3-inch graphite rod in 0.534 KNO3, separated from the sample solution by an ultrafine-porosity sintered-glass disk. The other details of the electrolysis cell used in this work have been discussed previously ( 8 ) . Reproducible stirring was provided at 600 r.p.m. by a Sargent synchronous rotator (E. H. Sargent Co., Chicago, Ill.) 'provided with a magnet attachment. The temperature was not controlled in these experiments. Materials. All chemicals were reagent grade and were used without further purification. All solutions were prepared in water purified by distillation and passage over a mixed cation-anion exchange resin bed (Mallinckrodt, Amberlite ME-3). Mercuric nitrate monohydrate (Baker Analyzed Reagent) was used to prepare stock solutions of 0.01OM mercury(I1)
in 0.1M KCNS, which were stored in polyethylene bottles. Stock solutions were standardized using an EDTA procedure (4). KO detectable change in stock solution concentration was observed over a period of two months, but fresh stock solutions were made u p arbitrarily a t two-month intervals, Solutions more dilute than lO-4M were prepared freshly as needed. The use and preparation of solutions more dilute than 10-6M required equilibration with the cell assembly and all other volumetric vessels with which the solutions would be in contact. A procedure analogous t o one previously reported was used (1). High purity nitrogen was used to remove oxygen from the sample solutions. It was bubbled through a gaswashing bottle containing the inert electrolyte solution and then dispersed in the cell through a coarse porosity sintered-glass disk. At least 15 minutes were required for initial de-aeration. I n addition, de-aeration was carried out for one minute between voltammetric runs which required more than two minutes. During a voltammetric run the nitrogen disperser was raised out of solution, while the nitrogen continued to flow over the surface. RESULTS A N D DISCUSSION
I n order t o select optimum conditions for the electrodeposition of mercury in a stripping analysis procedure, the cathodic voltammetric behavior of mercury(I1) at the graphite electrode was studied in various complexing s&-', media, including CNS-, "3, C1-. The most reproducible and bestdefined current-voltage behavior was obtained in a thiocyanate medium, and this medium was chosen for further studies. The selection of the proper electrolysis potential and the calculation of theoretical coulometric quantities are based on experimental convection-controlled current-voltage behavior. Figure l shows the current-voltage curve obtained for 1.0 X lO-'M mercury(I1) in O.1M KCNS under conditions identical to those used in the electrolysis step of the stripping analysis procedure. From this behavior an electrolysis potential of -0.70 volt us. S.C.E. was chosen. This potential is well enough along in the limiting current region to compensate for cathodic shifts in the Cathodic Voltammetry.
20
1 L -0.2
A -0.8
-0.6
-0.4
VOLTS
VI.
-1.0
S.C.E.
-
Figure 1 . Convection controlled polarographic behavior of mercury (11) at wax-impregnated graphite electrode 1 .O X 1 O-'M mercury(l1) in 0.1 M KCNS; scan rate, 5 mv./sec
half-wave potential with dilution, and yet not so cathodic as to cause significant interference from solvent reduction or codeposition of impurities. The maximum cathodic shift in E1/2 expected for this system would be 30 mv. per tenfold dilution (6). The shape of the convection polarogram indicates that the electrodeposition of mercury a t the graphite electrode is irreversible. This characteristic should have no adverse effect on a stripping analysis method, as long as the anodic dissolution does not occur a t too anodic a potential and is not too drawn-out. Moreover, previous studies of the stripping voltammetry of silver a t the graphite el-ctrode (an ostensibly irreversible process) indicated that irreversibility has no effect on the quantitative coulometric determinations (6). Anodic Stripping Voltammetry. PROCEDURE. Electrodeposition of mercury was carried out with stirring a t -0.70 volt us. S.C.E. At the end of the timed electrolysis period, the cell was switched out of the electrolysis circuit, and the stirring was stopped. (The authors recognized the possibility for error d u e to t h e spontaneous dissolution of the deposit when the cathode potential was VOLTS
vs.
SCE.
Figure 2. Comparison of stripping curves for dilute mercury solution with different electrolyte concentrations 1 .O X 10 -'M mercury(ll1, 20-minute electrolysis time A. 0.1M KCNS B. 0 . 0 2 M KCNS
removed. However, no difficulty of this sort was encountered, and the procedure of switching the cell out of the circuit was selected as the most convenient and accurate means of ending the timed electrolysis.) After 30 seconds, when the solution had settled, the cell was switched back in and an anodic potential sweep was initiated. The anodic potential sweep was carried out at 16.7 or 19.8 mv./second, depending on whether the FS Polarograph or the operational amplifier instrument was used. The initial potential for the anodic sweep was set a t the foot of the cathodic wave to minimize further deposition during the stripping step. Between runs the electrode was held a t a potential of f 0 . 8 volt us. S.C.E. for 1 minute. No other pretreatment was necessary. One serious problem encountered in this work was high residual currents in VOLTS vs. S.C.E.
0
0
z a
Figure 3. Current-voltage curve for stripping of mercury from graphite electrode 1 .O X lO-%4 mercury(ll), 2-minute electrolysis
the stripping curves obtained at low concentrations. However, the residual current appeared dependent on the thiocyanate concentration. Thus, by lowering the electrolyte concentration in dilute mercury(I1) solutions, the problem was minimized. For lop4 and lO-5-Tf solutions 0.111.1 electrolyte was used. At lower mercury(I1) concentrations, 0.02M electrolyte was used. Yo significant variation in the analytical data was observed as a result of using two different electrolyte concentrations. This was checked by using each of both electrolyte levels for intermediate concentrations with identical results. A comparison of stripping curves obtained a t lO-7M with 0.1 and 0.02M electrolyte is seen in Figure 2. (Dashed lines in this and other figures are experimental blanks.) For analytical purposes, then, the lower electrolyte level could have been used a t each mercury(I1) concentration studied. But, since some of these data were obtained partly to study the voltammetry of sub-monolayer deposits, conditions were optimized for making theoretical correlations. Thus, a higher electrolyte concentration was
VOLTS
vs S C E
Figure 4. Current-voltage curve for stripping of mercury from graphite electrode 1 .O X 10-OM rnercury(ll), 1 0-minute electrolysis
used a t higher mercury (11) concentrations, because higher currents necessitated a lower cell resistance to minimize I R drop. In addition, since each oxidized mercury(I1) ion could associate with four thiocyanate ions, a larger excess electrolyte than usual was necessary a t the high concentrations. Another problem encountered in this work was that the sizes of stripping peaks obtained in and 10-7M mercury(I1) solutions were smaller than expected for electrolysis times less than 10 minutes. The sizes increased rapidly with increasing electrolysis time, approaching the proper size for electrolysis times of 10 minutes or greater. No analogous problem was observed a t higher or lower concentrations, and no explanation for this behavior can be given a t this time. For analytical purposes, however, it can be stipulated simply that electrolysis times of 10 minutes or greater must be used below 10-5M. ANALYTICAL DATA. Typical anodic stripping curves are shown in Figures VOLTS vs.
-0.2
S.C.E. 0;2
0)0
0.4
Figure 5. Current-voltage curve for stripping of mercury from graphite electrode 4.0 X 1 O-gM mercury(ll), 30-minute electrolysis VOL. 37 NO.
a,
JULY 1965
0
969
~~
Table
I.
Anodic Stripping Analysis of Mercury(l1)" Electrolysis Av. quantity time, t, &/C*t of electricity, &,
Concn., C*, (moles/liter ) minutes pcoulombs x 10--6 0.33 1 . 0 x 10-4 345 10.4 1 . 0 x 10-5 2.0 206 10.3 1 . 0 x 10-6 10.0 108 10.8 1 . 0 x 10-7 20.0 20.7 10.4 1 . 0 x 10-8 20.0 2.35 11.7 4 . 0 x 10-9 30.0 1.35 11.2 Data refer to five replicate determinations at each concentration
3, 4, and 5 . Since the amount of material deposited should be proportional to the bulk concentration, C*, and the electrolysis time, t , a measurement of the number of coulombs, &, involved in the stripping peak(s) should reflect this dependency. Thus, assuming all deposited material is removed in the stripping step, &/ t should be proportional to C*. The mercury(I1) concentration was varied from 1.0 X 10-4L\f to 4.0 X l o - 9 ~ , and coulometric data are given in Table I. The quantity &/C*t remains constant over this large concentration range within 3~6.4%. This amount of deviation is not unusual for stripping analysis when the concentration is
Re]. std.
dev., Yo b3.4 2.1 4.1 5.8 4.9 3.4
varied over several orders of magnitude (9). The reproducibility a t a single concentration is somewhat better (Zk2.1 to 5.8%). Rased on limiting current data obtained a t 1.0 X 10-4Alf, the calculated constant, &/C*t, is 11.1 Zk 0.3 pcoulombs-l./mole-min. This quantity is in good agreement with the experimental data, indicating that all of the mercury deposited is subsequently stripped off, regardless of the deposit size. Although extremely sensitive, the technique has the primary disadvantage that, at intermediate concentrations (10-6 and lO-'M), multiple-peaked stripping curves are obtained (Figures 2 and 4). Thus, analysis of mixtures would be complicated. One possible
solution to the problem would be to use more dilute solutions when multiple peaks might interfere with a determination. The problem of multiple-peaked stripping curves is not unique with mercury. The stripping of micr.odeposits of silver from graphite (6) and of nickel from platinum ( 5 ) give similar behavior. Thus, it would be helpful to know more about the nature of the multiple peaks, and further studies are being carried out in this laboratory. LITERATURE CITED
(1) DeMars, R. D., Shain. I.. A N ~ L CHEM.29. 182.5 119.57), (2jIsrae1, Y.,Ibid.,31, 1473 (1959). (3) Jacobs, E. S., Zbzd., 35, 2112 (1963). (4) Korbl, J., Pribil, R., Chemist-Analyst, 45. 102 119561. (~, (5'1 5 ) r; Nicholson, M.>I., ANAL.CHEM.32, 1058 (1960). (6) Perone, S. P., Zbid., 35, 2091 (1963). (7) Perone, S. P., Mueller, T. R., Zbid., 37, 3 lQfi.5) 2 (1965). Perone, S. P., Oyster, T. J., Zbid., (8) F 36. 235 119641. (9) Shain, "Treatise on Analytical Chemistr;"' I. M. Kolthoff and P. J. Elving, eds., Part I, Section D-2, Chap. 50, Interscience, Yew York, 1963. RECEIVEDfor review March 8, 1965. Accepted April 16, 1965. \ - - -
An Electrochemical Study of Nitrite and Oxide in Sodium Nitrate-Potass um Nitrate Eutectic Melts H. S. SWOFFORD, Jr., and P. G. McCORMlCK Department of Chemistry, University of Minnesota, Minneapolis, Minn.
b The quantitative electrochemical determination of nitrite ion in fused alkali metal nitrate melts is discussed, and current-voltage curves are presented to demonstrate the presence of substantial residual nitrite in freshly fused melts. A method for removal of nitrite is also described. A wave has been observed in the eutectic melt at 250" C. which is attributed to the oxidation of oxide ion. Oxalate is proposed as a desirable species for the production of oxide ion in the melt in a concentration range suitable for electrochemical study. The literature regarding the subjects of nitrite and oxide in eutectic melts is also discussed in light of the present work.
T
determination of nitrite ion in fused alkali metal nitrates has been discussed by very few workers. Novik and Lyalikov ( I S ) reported that the oxidation of nitrite could be observed in a current-voltage curve following the addition of iodide to HE QUANTITATIVE
970
ANALYTICAL CHEMISTRY
NaN03-KN03 eutectic melts, nitrite not being seen before the halide addition. Oxide ion has received more consideration in the literature. Delarue (S,4has observed a wave in LiC1-KC1 eutectic a t 400" C. which he suggests is attributable to oxide; no other references have been found for the voltammetric determination of this species. Many methods, however, for adding oxide ion to nitrate melts have appeared in the literature (6, 7 , I S , 17-20); thus far no simple direct method has been suggested which does not leave some question as to the stoichiometry involved. This paper proposes oxalate as a species which is easily weighable and produces oxide in the NaN03-KN03 eutectic melt at 250' C. in the desired concentration range. Also described is a wave which has been observed in the eutectic as being attributable to oxide. Because of substantial residual nitrite, observed during the course of the
present work, an extensive electrochemical study was made of this species in the eutectic melt. A method for removal of residual nitrite is also described. EXPERIMENTAL
Equipment. The equipment used was similar t o that described by Swofford and Laitinen (21) with only minor modifications. The electrolytic cell had a capacity of 250 ml.; investigations were carried out in the bulk melt, rather than in compartments, unless otherwise noted. Temperature control was accomplished by manual adjustment of a variac in the heater circuit. A thermocouple immersed in the melt and connected to a dial-reading pyrometer was used to monitor the temperature; control to & 3 O C. was possible using this arrangement. All current-voltage curves were recorded on a Sargent Model XV Recording Polarograph using a potential scan rate of 0.20 volt per minute. Constant current was obtained from a
