Lanthanum Hexaboride as an Electrode Material for Electrochemical Studies D. J. Curran and K . S. Fletcher I11 Department of Cliemistry, Uniuersity of Massachusetts, Amherst, Mass. Lanthanum hexaboride, LaB6, has been studied as an electrode for electrochemical applications. A simple procedure has been developed for the analysis of LaB6 and the equivalent weight has been determined. An analytically useful potential range of approximately 1 V has been found. The cathodic limit is determined by the reduction of water at high pH or by the reduction of hydrogen ion at low pH and the anodic limit by the gross oxidation of LaB6. The reduction of hydrogen ion and the electrooxidation of LaB6 were studied in some detail.
SOLID ELECTRODE VOLTAMMETRY has received considerable attention in recent years, largely because of the desire to find electrodes suitable for application in anodic oxidations, where the utility of mercury is seriously hampered by oxidation and the use of noble metals is complicated by oxide formation. Further, while mercury shows a n extremely large overpotential for the reduction of hydrogen ion and water, few known solid electrode materials possess this property. From a n analytical point of view, a n ideal electrode material would be one which does not show oxide film formation and has a high cathodic and anodic overpotential in aqueous solutions or other solvent systems. Although the potential span of lanthanum hexaboride is too short to be considered ideal (approximately 1 V situated almost entirely cathodic of 0 V us. SCE), several interesting features of the material have been found. The known refractory binary compounds of boron with transition elements have attractive properties from a n electrochemical standpoint. As a class, these materials are extremely hard, chemically inert, and have high electrical conductivity. Reviews of preparation and chemical and physical properties of the refractory borides are available (1-3). Electrode fabrication of LaB6 was attempted using two techniques. The first employed a paste following the procedure developed by Mueller, Olson, and Adams for the carbon paste electrode ( 4 ) . The LaBs paste produced was found to have poor electrical conductivity (R > 10 Megohms) and was not employed further. La&, as its hot-pressed solid, provided an electrode whose physical and electrical properties were suitable for use as an electrochemical transducer. EXPERIMENTAL
La& samples were obtained from Cerac Chemicals, Butler, Wis., in the form of powder (5-8 microns) and hotpressed rods (1 inch long by 3/16-in~h diameter). The density of the rods was measured pycnometrically as 4.78 grams/cm3 which compared favorably to 4.75 grams/crn3 (x-ray) and (1) P. W. Gilles, “Borax to Boranes,” Advances in Chemistry
Series, No. 32, pp. 53-9, A.C.S., Washington, D. C., 1961. (2) N. N. Greenwood, R. V. Parish, and P. Thornton, Quart. Rev.,
4.70 grams/cm3 (pycnometric) reported by Samsonov (3). The calculated geometric area of one end of the rod was 0.179 cm2, based on a measured diameter of 0.188 inch. The electrodes were fabricated by mounting the sample material as its paste or hot-pressed solid into one end of a 6-inch length of 6-mm i.d. glass tubing. Electrical contact was provided at the back of the sample with a nickel wire, cemented into place in the case of the paste, and dipping into a mercury pool placed inside the glass tube in the case of the hot-pressed solid. The rod was sealed into the glass tube using 1/2-inch Teflon Ribbon Dope (Permacel, New Brunswick, N. J.) so only the end of the rod was exposed t o the solution. Current-time studies were performed with the same electrode assembly with the addition of an extra length of the Teflon tape projecting beyond the end of the electrode t o provide shielding. N o attempt was made t o polish the electrode surface. All current-potential studies were performed potentiostatically under quiet conditions. Instrumentation was accomplished with four K2-XA operational amplifiers and one K2-Bl booster, powered by a R-300 power supply (amplifiers and power supply manufactured by G. A. Philbrick Researches, Inc., Boston, Mass.). The input and feedback circuits were supplied by a Heath Model EUA-19-2 Polarography Module (Heath Co., Benton Harbor, Mich.) (5). All voltammetric curves were obtained using a Heath Model 20A servorecorder with a chart speed of 8 inches per minute. The cell used for linear sweep voltammetric experiments was a 150-ml glass beaker fitted with a Lucite cover through which counter, test, and reference electrodes were introduced. The counter electrode was a 2.6 cm2 platinum disk dipping into solution and separated from the main compartment of the cell by a medium porosity glass frit. The reference electrode was the saturated calomel electrode described by Lingane (6) with “J” tip filled with saturated KC1 to provide reproducible liquid junction potentials. The cell resistance was measured with a 1000-Hz conductivity bridge (Model RC 1682, Industrial Instruments, Cedar Grove, N. J.) as 80 ohms between the test electrode and counter electrode and 600 ohms between the test electrode and reference electrode in 0.1OON KCI. Constant currents were obtained with a Sargent Model IV Coulometric Current Source and were standardized by measuring the i R drop across a standard 10-ohm *0.05% resistor (General Radio, Concord, Mass., Type 500-B) with a Rubicon Portable Potentiometer, Model 2730 (Minneapolis-Honeywell Regulator Co., Rubicon Instruments, Philadelphia, Pa.). The cell used for the constant current studies was all glass and consisted of two compartments separated by a medium porosity glass frit. The auxiliary (cathode) compartment had a total volume of 25-ml and was fitted at the top with a 19/22 male ground glass joint. The auxiliary electrode was a 2.6 cm2 platinum disk mounted into a closed ground glass female joint t o fit the cathode compartment which contained 1.ON CuCI2 and 1.ON HC1. The test electrode compartment had a total volume of 25-ml and was fitted with two 10/18
20, 441 (1966). (3) G. V. Samsonov, “Refractory Transition Metal Compounds;
High Temperature Cermets,” Academic Press, New York, 1964. (4) T. R. Mueller, C. Olson, and R. N. Adams, 2nd Intern. Polarog. Congr., Cambridge, August 1959. 78
ANALYTICAL CHEMISTRY
(5) C. G. Enke and R. A. Baxter, J . Chem. Educ., 41, 202 (1964). (6) J. J. Lingane, “Electroanalytical Chemistry,” 2nd Ed., Interscience, 1958, p. 362.
female ground glass joints. The La& was sealed into the through end of a lOjl8 male joint with the Teflon tape and inserted through a n angled female joint. The second lOjl8 female ground glass joint was positioned directly over the end of the LaB6 electrode and permitted the introduction of a gas trap t o collect gaseous products which result o n electro-oxidation of LaB6. This trap consisted of a 10-inch length of 6-mm glass tubing fitted with 2 ground glass stopcocks 2 inches from each end, enabling isolation of approximately 2 ml of gas between them. One end of the tubing was sealed to a male 10118 ground glass joint t o fit the female joint on the electrolysis cell, and the other end was sealed t o a 10130 female ground glass joint t o allow the trap t o be attached to the gas sampling port of a Hitachi mass spectrometer for identification of collected gaseous products. A side arm from the test electrode compartment permitted connection of a 6-inch length of 6-mm glass tubing calibrated in milliliters, which was used t o measure the volume of solution displaced from the cell by the gas. Temperature control for current-potential and constant current studies was achieved with a P. M. Tamson (Zoetemer, Holland) Model T9 constant temperature bath a t 25 i 0.02” C. pH measurements were made using a Leeds & Northrup (Philadelphia, Pa.) Model 7401 p H meter. Reagent grade chemicals were used in all cases and supporting electrolytes were preelectrolyzed at a mercury pool held at - 1.25 V cs. SCE for several hours prior to use. Oxygen was found to produce a n ill-defined reduction wave a t La& and prepurified nitrogen was therefore used to purge all solutions prior to all runs. RESULTS AND DISCUSSION
Analysis of La&. Emission spectrographic examination of the powder showed the material to be quite pure. Quantitative estimates placed no impurities in the I-lOz range, only Fe in the O.l-l.O% range, only Ni in the 0.01-0.1 % range, and Si, AI, Ti, R.E., Mn, V, and Mg a t less than 0.01 %. When La& was oxidized in nitric acid solution, the products were La(III), identified by its precipitation with oxalate from acid solution (7), and H3B03,identified by the rose color produced by its complex with turmeric (7). For these qualitative studies a 1-gram sample of La& powder was’dissolved in 25 ml of 8 N ”0, and the resulting solution was examined with a Wallace Direct Vision Hand Spectroscope using a cell with a path length of 20 cm. N o absorption bands could be detected in the visible region, and it was concluded that other rare earth elements were not present in more than trace amounts-a result in agreement with the emission work on the powder. A small insoluble black residue remained after dissolution of LaB6 and it was concluded that this was carbon, as it was not affected by boiling aqua regia, which would eliminate the possibility of elemental boron (8). In the case of the powder, this residue was extremely small and could not be directly weighed. The hot-pressed solid had a higher percentage of this material and it was possible to separate and weigh the residue directly. A higher content in the hot-pressed solid is expected as the rod was no doubt fabricated from the powder in a carbon mold at high temperature (9). For quantitative analysis, samples of La& (0.1-0.25 gram) are dissolved in 50 ml of 8 N HNO, with heating as necessary. After solution, concentrated ammonia is added t o (7) A. A. Benedetti-Pichler, “Identification of Materials,” Academic Press, New York, 1964. (8) R. Kiessling, Acta Cliem. Scand., 2 , 707 (1948). (9) G. A. Meerson, R. M. Manelis, and T. M. Telyukova, Izc. Akad. Nauk., Neorg. Materialy, 2 , 291 (1966).
Analysis of Lanthanum Hexaboride Theoretical, Powder.Q Rod,b
Table I.
Lanthanum 68.17 67.0 3~ 0 . 1 66.8 Boron 31.83 32.82 f 0.13 32.59 Residue (carbon) ... 1.1 Total lod.00 99.82 f 0 . 2 100.5 a Reported percentages are average and standard deviation of four trials. Single determination.
adjust the p H to 2.0, as measured with a glass calomel electrode pair and standardized p H meter. The small residue of carbon is removed from the resulting solution by filtration through a weighed glass filter crucible with medium porosity glass frit. While the amount of carbon obtained from dissolution of the powdered samples could not be weighed, that collected from dissolution of the hot-pressed solid was sufficient to allow calculation of a per cent carbon in the rod and is reported in Table I. After filtration, the filtrate is heated to boiling and 25 ml of 0.2M H2C204is added with stirring. The solution with precipitate is allowed to sit for 8 to 10 hours with occasional stirring before the La2(C204)3is filtered using a n ashless filter paper (Whatman No. 40 or equivalent) and 4 t o 6 washings with 2% H2C2O4(10). The precipitate is ignited at 900” C. to constant weight as La203 to complete the determination of lanthanum. The filtrate and washings collected in the previous step are brought to boiling and sufficient NaOH is added to drive off all ammonia, as noted by a negative red litmus test. After cooling, the p H of the solution is adjusted to 6.0, 10 ml of mannitol solution (1 gram/ml) is added, and the boric acid is titrated as its mannitol complex to p H 10.0 with standard NaOH ( I ] ) . The sodium hydroxide solution is preferentially standardized against Borax, Na2B40710H20,prepared as primary standard following the procedure of Hurley (12) using the same p H range employed in the sample titration. Results of the analyses are shown in Table 1. Equivalent Weight of La&. Aqueous oxidative degradation of La& resulted in aquated boron and lanthanum species in their stable + 3 oxidation states [B(OH), and La(III)]. Using Ce(1V) and material of ideal stoichiometry, this process is expected to involve a loss of 21 electrons per mole of La& according to : +
La& $- 21Ce(IV)
+ 18Hn0
-+
La(II1) f 21Ce(III) f 6B(OH)3
+ 18H+
(1)
Based on the analysis of the powder shown in Table I, 204.14 grams would be expected to contain 1 mole of lanthanum hexaboride. The percentage composition of the powder and Equation 1 lead to the prediction that 21.54 eq/mole are required. To test this, weighed samples of the powder were heated moderately with a measured excess of standardized Ce(1V) in 2.ON H2S04 or 0.5N “0,. The Ce(IV) remaining after reaction was determined by back-titration (10) M. M. Woyski and R. E. Harris, “Treatise on Analytical
Chemistry,” Part 11, Vol. 8, I. M. Kolthoff, P. J. Elving, and E. B. Sandell, Eds., Interscience, 1963, p. 34. (11) I. M. Kolthoff and E. B. Sandell, “Textbook of Quantitative Inorganic Analysis,” 3rd Ed., pp. 534-5, Macmillan, New York, 1952. (12) F. H. Hurley, IND. ENG. CHEM., ANAL.ED., 8, 220 (1936); 9, 237 (1937). VOL 40, NO. 1, JANUARY 1968
79
Table 11. Current Efficiency for Electro-oxidation of La& in 0.1N HCI Calculated mmoles Current LaBB efficiency,r LaB6 found,a grams 0.0386 0.189 109 Lanthanum found,” grams 0.0263 0.189 109 Boron found,b mmoles 1.;51 0.1918 110.7 a Weight of La& oxidized was obtained by weighing the rod before and after electrolysis and is corrected for the carbon residue. * Lanthanum and boron were determined following the procedure for the analysis of La&. Current efficiencies were calculated using C.E. = (mmoles LaB, calculated/electrochemical mmoles) X 100 where the electrochemical mmoles was obtained using Faraday’s law with Qerp = 359.6 coulombs and / I = 21.50 eq/mole.
0 2
-02
0
-04
-06 VOLTS
V5
-08 SCE
Figure 1. Residual current-potential 1.00N KCI at various p H
-10
-12
-14
behavior of LaBs in
Numbers on each curve are pH values. Sweep rate was 14.4 mV/
second with standardized ferrous ammonium sulfate. In the sulfuric acid medium, two trials gave identical results of 20.85 moles of Ce(IV) consumed per mole of sample. The reaction was fairly slow and lanthanum sulfate was precipitated in this procedure. The possibility of some of the very fine grained powder being trapped in the precipitate and, therefore, not oxidized was suspected and could account for the low results. The experiment was repeated in the nitric acid medium and two trials gave 21.14 and 21.23 moles of Ce(1V) consumed per mole of sample. It was also possible that the nitric acid results were low as the analysis experiments showed that 8 N H N 0 3 , with moderate heating, would oxidize the powder. Two trials were repeated without the Ce(IV) and the La& powder was recovered from the 0.5N H N 0 3 solution. An average per cent weight loss of 0.4 was found, which would correspond to about 0.08 moles of Ce(1V) consumed per mole of sample. The results, while still somewhat low, have been taken as good evidence for Equation 1. Transition metal hexaborides are isomorphous and contain B6 groups lying in the centers of cubes of metal ions (13). There are two theoretical approaches to the electron bond formation in La&. The first, originally discussed by Kiessling (14) concludes that bond formation is accomplished by the transfer of electrons from the boron framework to the metal atom, The second and more popular view describes the bonding process as the transfer of electrons from the metal to the boron framework (3, 15). On the basis of the second approach, it is reasoned that each B6 octahedron requires 32 electrons to achieve its bonding electron requirements, 30 of which may be supplied by six boron atoms and two of which may be supplied by ionization of lanthanum. X-ray evidence suggests that lanthanum is present in its plus three oxidation state (16), and thus a 33rd electron is delocalized over the boron framework. In view of the second argument, a n assignment of a plus three oxidation number to lanthanum in the solid state would aptly illustrate the formalism of this concept. Analytical Potential Range. Current-potential curves were run a t the hot-pressed solid (in quiet solution) using a linear potential scan rate of 14.4 mV/second. The residual cur(13) R. W. Johnson and H. H. Daane, J. Chem. Phys., 38, 425 (1963). (14) R. Kiessling, Acta Clrem. Scaird., 4, 209 (1950). (15) H. C. Longuet-Higgins and M. deV. Roberts, Proc. Roy. Soc. (Loiidon), 230, 110 (1955). (16) E. E. Vainshtein, I. B. Staryi, S . M. Blokhin, and Yu. B. Paderno, 212. S f r u k t . Klrim., 3, 200 (1962). 80
ANALYTICAL CHEMISTRY
rents recorded using 1.00N KC1 as supporting electrolyte show a useful potential range dependent on pH, as shown in Figure 1. Cathodic polarization showed two distinct reduction processes both of which provided evolution of gas at the electrode. At low pH, reduction of hydrogen ion began a t approximately -0.90 V cs. SCE and resulted in the production of H2 gas at the electrode surface according to Equation 2. At high pH, the predominant cathodic reaction is assumed to be the reduction of water, yielding measurable current at approximately - 1.40 V cs. SCE according to Equation 3. 2H+ 2H20
+ 2e-
+ 2e-
.-t
-f
H2
H2
+ 20H-
(2)
(3)
Anodization of La& resulted in gross oxidation of the electrode and gave rise to a complex current-potential behavior. The Ce(1V) work indicated the following half-reaction for pure material of ideal stoichiometry: La&
+ 18H20
+
La(II1)
+ 6H3B03+ 18Hf + 21e-
(4)
From Figure 1, anodic polarization of La& in solutions of low p H shows a decomposition potential or anodic limit at approximately +O.lO V cs. SCE. Increasing pH shifts the decomposition potential cathodically, yielding a t p H 10 a n anodic limit of -0.40 V L’S. SCE. In the basic solution at least two current inflections prior to the massive anodic current excursion appeared. Further study of these peaks is necessary. A study of the electrochemical oxidation of La& was performed using a constant current of 48.24 mA. The current efficiency for electro-oxidation of La& was determined by comparing the weight loss of the electrode, the number of moles of HRBOIproduced, and the number of moles of La(II1) produced with the electrochemical value (Q/n F ) , based on 21.50 equivalents per mole (calculated using the analysis in Table I). The data in Table I1 show current efficiencies of 109, 109, and 110.7%, respectively. These values are believed to be equal within experimental precision. The oxidation was accompanied by the production of small amounts of gas in the vicinity of the electrode and a second series of runs was performed to obtain qualitative and quantitative analyses of the gas, The gas was collected in the gas trap connected to the electrolysis cell, isolated, and introduced into the mass spectrometer for identification. A substantial signal enhancement for m/e = 2 identified the gas as hydrogen. N o evidence could be found for the presence of 0: (above background)
Table 111. Current Efficiency for Electro-oxidation of La& and Stoichiometry of Hydrolysis of La&
Current, mA
Total coulombs passed
9.732
233.6
19.28
230.9
48.24
241.2
9,656
Supporting electrolyte 0 . ION NaOH 0 . 1 0 N NaCl 0.01M EDTA 0 . ION NaOH 0. ION NaCl 0.01M EDTA 0 . ION NaOH 0 . ION NaCl 0.01M EDTA 0 . 1 0 N HCI 0 , 9 0 N KCI
232.2
Weight lossa of La& 0,0252 gram 0.124 mmole
Currentb efficiency,
HPgas collectedc Volume Calculated collected, ml solubility, ml
Total mmoles
Mmoles H?/ Mmoles LaB,
-~
110
0.26
0.45
0.028
0.226
0.0251 gram 0.123 mmole
110
0.35
0.45
0.032
0.260
0.0259 gram 0.127 mmole
110
0.24
0.45
0.027
0.213
0.0253 gram 0.124 mmole
111
0.26
0.48
0.029
0.234
Weight of
oxidized was obtained by weighing the rod before and after electrolysis and is corrected for carbon residue. = (mrnoles LaB, observed/electrochemical mmoles) X 100 where electrochemical mmoles were calculated using Faraday’s law with Q,,, as noted and n = 21.50 eq/rnole. Calculated volume solubility is obtained from the Ostwald Solubility Expression. “1” values were estimated using data compiled by Seidell ( 1 7 ) as 0.018 and 0.019 for 0.1N NaOH and 0.1N HC1, respectively. The total mmoles of Ha were calculated for the temperature of the experiment and at the prevailing atmospheric pressure, corrected for the vapor pressure of water. a
* Current efficiencieswere calculated using CE
or of borohydrides in the gaseous products. The volume of hydrogen produced was determined by measuring the volume of supporting electrolyte displaced from the cell into the calibrated glass tube and adding this number to the volume solubility of Hz in the supporting electrolyte, which was obtained using the Ostwald Solubility Expression and data compiled by Seidell (I 7). The data for four runs are found in Table 111. The amount of hydrogen gas produced was small and apparently not affected by the acidity of the solution. The last column in Table 111 shows the ratio of moles of Hz produced t o moles of LaBGoxidized. The average for the four trials was 0.233. It is proposed that a competing hydrolysis reaction of boronboron fragments occurring as intermediates in the gross oxidative process produced the hydrogen gas. Inasmuch as the electrode material is stable in base and in nonoxidizing acids, the initial step or steps in the overall reaction must be electrochemical, but subsequent parallel electrochemical a n d chemical paths exist with the former path the predominent one. In both cases the oxidation products of the electrode material are the same, but the chemical route also produces some hydrogen gas. Equations 5 and 6 show a plausible route for the production of hydrogen gas from boron in the +2 oxidation state.
LaB,
+ 18H20
+
La(II1)
+
+
6H3B03 17.54H+ 0.23H2
+
+ 20.54e-
(7)
Similar equations could be written for boron in the +1 oxidation state and for other polymeric boron species in some stage of hydrolysis. If the production of hydrogen gas is the only side reaction, the sum of the electrochemical and chemical reactions would be equivalent to writing the equation:
The analysis of the rod shown in Table I indicates 21.50 eq/ mole are required for the electrode in the absence of any side reactions. A current efficiency would be expected, then, of (21.50/20.54) x 100 = 105% which is smaller than the 110% shown in Table 111 and may indicate the possibility of additional side reactions. However, no other evidence exists to support this idea and the difference in the percentages may rest in the accuracy of the analysis of the rod, the measurement of the volume of hydrogen evolved, and the estimation of the solubility of hydrogen gas in the electrolyte. Reduction of Hydrogen Ion. Pedk current-potential response, characteristic of a diffusion-limited reduction process, resulted on cathodic polarization of the rod in solutions of intermediate pH. This response was investigated in a series of 80 runs with hydrogen ion concentrations from 2.72 to 10.72 X 10-4Nin 0.100N KCI at a sweep rate of 14.4 mV/second. The various hydrogen ion concentrations were obtained by coulometric generation of base a t a platinum cathode immersed in a solution of 1.072 X 10-3N HC1 in 0.100N KCI. The anode compartment was separated from the rest of the cell by a medium porosity glass frit augmented by a n agar plug containing KC1. The peak currents were corrected for the residual current by subtracting the current obtained a t 14.4 mV/second in 0.1OON KCl at the potential of the current peak, and were a linear function of the hydrogen ion concentration over the range studied. The ratio of peak current to concentration for these runs was 42.7 amp-cm3/mole with a standard deviation of 2.5 A-cm3/mole. The averge peak and half-peak potentials were -1.053 + 0.010 and -0.937 + 0.012 V us. SCE, respectively. Assuming the reaction to be totally irreversible, the equations given by Meites (18) for the separation between peak and half-peak potentials, and for the peak current, were used to calculate the ana product and the diffusion coefficient of hydrogen ion. The results were: an, = 0.41, D = 10.8 X cm2/second. Using an equation from the same source, these results were
(17) A. Seidell, “Solubility of Compounds,” 2nd Ed., Van Nostrand, 1919.
(18) L. Meites, “Polarographic Techniques,” 2nd Ed., Interscience, 1956, pp. 415-16.
H O OH
I I + HzO I 1
B-B HO
-
OH
HB(OH)?
+ Hz0
B(OH),
-
+ HB(0H)z
B(0H)s
+ H?
(5)
(6)
VOL. 40, NO. 1 , JANUARY 1968
81
used to calculate the rate constant at the standard potential as 5.4 X 10-9 cm/second. The electrode configuration corresponded to unrestricted linear diffusion, however, so the equations for linear diffusion cannot be applied rigorously. Current-time curves were obtained with a shielded electrode in a solution 0.100N in K C L and 9.95 X 10-4N in hydrogen ion. The electrode was switched from -0.500 V to - 1.250 V us. SCE (approximately 200 mV cathodic of the peak current potential). The it1l2 product initially decreased but was then constant for a n interval and finally increased at longer times of electrolysis. The results of four runs, taken after the initial decrease, were: 89 pA-seconds1’2(from 10.50 to 66.0 seconds), 89 pA-secondsl’* (from 16.50 to 57.0 seconds), 88 pA-seconds1’2 (from 10.50 to 73.5 seconds), and 89 pA-secondsl’Z(from 18.00 to 66.0 seconds). The precision and accuracy of these results were + l % relative. Theoretical calculation of the itl’* product using n = 1 eq/mole, F = 9.65 x lo4coulombs/eq, A = 0.179 cm2, D = 8.68 X 10-5 cm2/second [obtained for 0.100N K C L supporting electrolyte from Reference ( I s ) ] and C = 9.95 X lo-’ moles/cm3 gave 90.4 pA-secondsl/*. The agreement between theoretical and experimental i f 1 I 2 products is very good. The initial decrease shown by the data is believed due to the surface roughness of the electrode, particularly as no attempt was made to polish the material. The increase observed at later times of electrolysis was no doubt due to convection. The results are also consistent with the linear potential sweep work. Although the geometric electrode area was the same in both cases, a n unshielded electrode was employed in the sweep studies while a shielded electrode was used in the current-time work. The peak current obtained in the former case should be larger than that obtained for restricted linear diffusion, and conseqeuntly the calculated diffusion coefficient should be larger than it is actually. The experimental results are in agreement with this conclusion. CONCLUSIONS
The analytical utility of LaB6 as a solid electrode for voltammetry has been investigated. While the electrode cannot be employed successfully for many oxidations, a span of 0 to - 1 V us. SCE compares very favorably with other solid electrodes for the study of reduction reactions. The reduction of hydrogen ion is well behaved and obeys the equation for linear diffusion to a plane electrode. While further studies are necessary, the results indicate that the electrode obeys fundamental electrochemical equations and can be used for both theoretical and chemical analysis purposes. Separation of the reduction of hydrogen ion from the reduc(19) M. von Stackelberg, M. Pilgram, and V. Toome, Z. Elekrrochem., 57, 342 (1953).
82
ANALYTICAL CHEMISTRY
tion of water is well defined and the electrode may be of some value in the study of hydrogen overpotential. It seems clear that the electrode can be used for the electrochemical generation of La(II1) and B(II1). Although it may be possible and desirable to find conditions of 100% current efficiency with respect to lanthanum and/or boron species, it is not necessary in a practical sense as the current efficiency for generation of either ion can be established by chemical analysis for a given electrochemical situation. Once the electrode has been calibrated in this fashion, the amount of La(II1) or B(II1) generated in subsequent experiments can be found from the number of coulombs of electricity passed through the cell. If only one of the ions is of interest, it is not necessary to know the composition of the material being used for the electrode because the current efficiency can be referred to pure material of ideal stoichiometry for purposes of calibration. It should perhaps be emphasized, however, that the current efficiency determined in this manner for one set of conditions, such as current level and solution composition, may not hold for another set of electrolysis conditions. Electrochemical generation of lanthanum ion could be useful for the precipitation titration of a number of anions such as oxalate (in acid solution) and fluoride. Little in the way of redox titrations seems possible with the lanthanum hexaboride electrode but the situation could be more interesting in this respect if similar voltammetric response extended to the cerium or europium hexaborides. There are a larger number of borides, carbides, nitrides, and silicides of transition and rare earth metals. The specific resistivities of these materials cover the range from a few microohm-cm to megohm-cm. Some of them are metallic conductors, others semiconductors, and others insulators. I t is interesting to speculate o n the possibility of a number of refractory materials being available for the electrochemical generation of cations and anions, and on their possibilities as inert electrodes for voltammetry and other techniques. We have examined the behavior of a few other materials as paste electrodes in our laboratory. While these pastes are not extremely well behaved, TiN, for example, showed a potential span of approximately + I to -1 V us. SCE in a number of supporting electrolytes. The anodic limit does not seem sufficiently positive for the oxidation of water, and we have evidence for the existence of titanium in solution and no evolution of oxygen at the electrode. ACKNOWLEDGMENTS
We would like to thank R. G. Russell for providing the emission spectrographic analysis of LaB6, I. S. Fagerson for mass spectrometric analysis, and the University of Massachusetts Research Council for partial support of this work.
RECEIVED for review May 9,1967. Accepted October 6,1967.
