purity if the separability is good enough. I n Table I1 are data which show the tremendous advantage gained in yield if the separation factor can be made large. With the separation factor as large as 3 the amalgam cathode method would require only a single-stage continuous electrolysis to obtain a reasonable yield of purified lanthanon from a binary mixture.
Onstott, B. I., J . A m . C’heui. ,bo( 77, 2129 (1955). ( 4 ) 16id.. 78. 2070 (19561. ( 5 ) Ibzd.: 81: 4451 (lY59j. ( 6 ) Ibzd., 82, 6297 (1960). ( 7 ) Onstott, E. I., Svverud, A. S , 1)ivision of ‘Inornanii CheAstrv. 133rd (:3)
Table II. Calculated Yield for Specified Purity Lnl-Ln2 Mixture
a
Meeting, ACS, San Francis&. Calif dpril 1958. ( 8 ) Warf, J. C., Donohoe, Jerry, Hardcastle, Kenneth, Office of Xaval Research, Rept. NP-6531 (October 1957) I 0 ) Warf. J. C.. Karst. W. L.. Ibzd.. NP5006 (Sovem’ber 1963). ’
LITERATURE CITED
~IECEIVEI) for review June 5, 1961 .4crepted August 14, 1961. Division of Analytical Chemistry, 139th hIeeting, ACS, St. Louis, 3\10.,,.Ifarch 1961. Xork done under auspices of the U. S r2tomic Energy Commission.
( 1 ) . Latimer,
W. M., “Oxidation Potentials,” p. 291, Prentice-Hall, New York, 1952. ( 2 ) Onstott, E. I., Division of Inorganic Chemistry, 138th Meeting, rl CS, New York, X . Y., September 1960
A Study of the Molybdenum-Catalyzed Reductio n of PerchIo rate G. A. RECHNITZ‘ and H. A. LAITINEN Department of Chemistry, University o f Illinois, Urbana, 111.
b The reduction of perchlorate ion to chloride has been studied with the objective of applying it to a quantitative determination of perchlorate. The kinetics of the reaction were studied and yielded rate expressions for the over-all reduction as well as the interaction between the catalyst and perchlorate. Controlled potential electrolysis showed that the catalytic cycle involves molybdenum in its +5 and +6 oxidation states. The reaction path seems to involve the reduction of Mo(VI) to Mo(V) dimer, which is then reaxidized b y Clod-. A possible mechanism is thus discussed in terms of a homogeneous rate-determining step preceded b y heterogeneous steps to produce the active catalytic species.
reaction orders appearing in the rathcr complex empirical rate law obtained Recently, Burns and biuraca ( I ) h a w developed a most effective means of reducing perchlorate to chloride involving reduction by excess titanous chloride in the presence of OsOl catalyst. The problem of the wet catalytic reduction of perchlorate was approached through a kinetic study of both the over-all reduction and some of the individual reaction steps involving the molybdenum catalyst. Efforts have been primarily directed toward the elucidation of the catalytic cycle with the hope of gaining an understanding of the over-all reaction mechanism. EXPERIMENTAL
T
catalytic reduction of perchlorate has been a topic of analytical interest for a number of years. Crowell, Yost, and Robert (3) used osmium salts to catalyze the reduction of perchlorate by hydrobromic acid. The most extensive work to date has been that of Haight (5-Q),who studied both the polarographic and chemical reduction of perchlorate in the presence of molybdenum(V1) and tungsten(V1) catalysts. Haight postulated species of molybdenum(1V) as active catalytic f o r m to explain certain fractional HE
Present address, Department of Chemistry, University of Pennsylvania, Philadelphia 4, Pa.
Apparatus. The Cary Model 14 recording spectrophotometer and, in a few cases, the Bausch & Lomb Spectronic 20 instrument were used for absorbance measurements. A multipurpose unit of operational amplifier type as described by Enke (4) was used for both controlled potential and controlled current electrolyses. The cell coiisisted of a 500-ml. round bottomed flask fitted with frittedglass entrance tubes to provide for anode and reference electrode compartments. The area of the mercury pool cathode was held at 100 sq. em. throughout. Stirring of the acid-mercury interface was provided by a Teflon paddle stirrer rotated at 33 r.p.m. The number of coulombs passed was measured with a nitrogen-hydrogen coulometer as described by Page and Lingane (f4).
A11 polarograms were taken with a Leeds & Northrup Type E polarograph while amperometric titrations a t the rotating platinum electrode were carried out with Sargent Model I11 Polarograph. Reagents. All reagents were of t h e best obtainable commercial quality and were used without further purification. Cadmium amalgams were prepared by dissolving pure cadmium metal in redistilled mercury. Samples of Mo(V) salts were kindly furnished by the Climax Molybdenum Co., Xew York. Procedures. The rate of t h e overall reduction of perchlorate was followed by determining t h e concentration of t h e final reduction product, i.e., chloride. A mercurimetric titration procedure (15) and an amperometric titration method (19) were used for this purpose. The concentration of Mo(V) was determined as a function of time by spectrophotometric measurements a t 430 mM. RESULTS A N D DISCUSSION
Preliminary experiments using 5 to 10% Cd(Hg) as the reducing agent in the presence of molybdenum(V1) as the catalyst showed t h a t the addition of phosphoric acid to a sulfuric acid medium was beneficial in raising the boiling point without increasing the sulfuric acid concentration to the point where sulfate reduction occurred. A 4M sulfuric acid-2M phosphoric acia was chosen on the basis of these preliminary experiments as optimum for quantitative perchlorate reductio~, VOL 35, NO 1
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520
WAVE LENGTH
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680
Studies of the effect of such experimental variables as temperature, concentration of reactants, products, and catalyst as well as the acidity of the medium at constant mole ratio of 2 H2S04:HaP04 led to the empirica: rate law dt
K [HClO,] 2'3 [ MotOt,~] l'*[H+] 2
where IC is equal to about 8 X 10-5, with a n apparent activation energy of 10.64 kcal./mole calculated over the temperahre interval 89' to 110' C. Experimental variations in the specific rate constant as well as the unusual fractional reaction orders indicated that this rate law does not truly represent the stoichiometry of a single ratedetermining step. The concentration of the reducing agent, however, does not appear in the rate law although the area of reducing agent exposed to the reaction medium does affect the magnitude of the specific rate const'ant. For purposes of elucidating the reaction mechanism, a more meaningful rate law can be obtained when the complicating effect of heterogeneity is eliminated. On the other hand, the use of a heterogeneous reducing agent has certain advantages from a n analytical viewpoint so that the over-all rate law is useful in the establishment of optimum conditions for the quantitative determination of perchlorate by this reaction. Molybdenum is known to exist in the MQ(VI), (V), (IV), (111)and Mo(I1) oxidation states. Stable compounds have been isolated for all of these forms, even Mo(1V). The solution chemistry of molybdenum compounds is further complicated, however, by the pronounced tendency of molybdenum to form complexes and polymeric products by intermtion with solvents and &her species present. While acid solutions of Mo(VIi show no characteristic abmrption features in the 350- to 800-mp range, l,To(?'I: undergoes a series of 1474 *
ANALYTICAL CHEMISTRY
I
I
I
I
-0.2
0
(mp)
E
Figure 1. Effect of acidity on spectra of Mo(V) in 2 : 1 mixtures of H2S04-HaPOh
d[C1 - = -I
I
0.2
I
L
1
.&0.6
( V O L T S vs. S.C.E.)
x
Figure 2. Polarograms of 2 H3P04media of varying acidity
pH-dependent equilibrium steps in acid solutions ( I S ) . Mo(V), on the other hand, yielded spectra in 4 M H&30a-2J1 H3P04solutions characterized by a stmng absorption peak at 430 mp. This peak has been identified with the dimeric form of Mo(V) (5-9, 16). Beer's lam was found to hold over a twenty-fold range a t 6-l4 total acidity with a molar absorbancy index of 450 a t 430 mp. X positive deviation was noted a t l 0 J 1 and a negative deviation at 2 X . These observations suggest a polymer-dimermonomer equilibrium shifted in favor of the monomer at high acidity and polymer a t low acidity. The dimer is greatly favored a t 6.11 total acidity. Figure 1 indicates that the 4.11 HzSO42M H3P04solvent strongly faT.ors the dimeric form of hIo(T'). The 430-mp maximum is partially obscured in the 2M acid medium so that measurements a t that wave length are no longer characteristic of the dimer concentration alone. The polarography of molybdenum is a complex subject which remains controversial and incompletely understood. It appears that both the ease of reduction and the reduction products are a function of the nature and acidity of the supporting electrolyte. The mechanism of the electrode reduction of molybdenum seems to be different in sulfuric and hydrochloric acid media (2,5,10, 11j. I n addition, catalytic waves are observed in perchloric or nitric acid media. Figure 2 shows a series of polarograms obtained for the reduction of Mo(V1) in solutions of varying concentrations of sulfuric-phosphoric acid mixtures. While the half-wave potential does not shift with changing acidity, the reduction wave at -0.172 volt us. S.C.E. divides into a n incompletely resolved doublet wave with decreasing acidity, accompanied by a n increase in total wave height, partly due to a reduction
I
-0.4
1 O-'M Mo(VI) in HzSOr
wave at anodic potentials. The effect of column height on polarographic current was investigated a t applied potentials of 0.00 and -0.40 volt 2's. S.C.E. Plots of log column height us. log current yielded slopes of 0.46 += 0.02 with zero intercept, clearly indicating diffusion control at both potentials. These findings and the coulometric evidence discussed below led t o the conclusion that the various p H dependent forms of Mo(V1) are in slow equilibrium with one another and that the changes in wave heights with changing acidity are related to the relative equilibrium concentrations of these species a t any given pH. The relative complexity of the polarograms obtained makes the identification of molybdenum oxidation states from polarograp5c evidence alone a questionable procedure. It was felt that controlled potential coulometry a t a mercury pool cathode would be less anibiguous in this connection. Complete controlled potential electrolysis of a millimolar solution of Mo(V1) in 451 H2SO4-2;11 &Po4 a t -0.40 volt vs. S.C.E. yielded an experimental value of n = 0.97. The electrode reduction a t this potential must thus be Mo(VI) e-+ Mo(Vj, a conclusion which finds further support in the fact that the final solution obtained had a spectrum identical with that of synthetic samples of Mo(V). Since the reduction of chlorate proceeds much more readily than that of perchlorate under the same conditions, i t is reasonable to postulate that the rate-controlling step involves the formation of chlorate which is reduced further to the chloride. Since both perchlorate and the catalyst are initially present in their highest oxidation states, i t would also seem reasonable that the reaction scheme might involve the reduction of the catalyst by the reducing agent, folloxed by some inter-
+
2o t
I
Y
I
0018 N ut104 0 002 u M g M a 0,
16
i
d
2
N O HC.104 0002 M Nd2MOO4
4-
, 0
I
I
2
3
‘
*
5 6 T I M E , HOURS
4
7
I
I
I
8
9
IO
o
!
20
40
60
m
I
100
120
140
,
160
I
180
zoo
TIME, MINUTES
Figure 3. Controlled potential reduction in 4M H2SO42 M H3P04 a t mercury pool cathode Temperature, 100’ C.
Figure 4. Constant current reduction of 0.018N HClOi and 0.004M Mo(VI) in 4M H2S04-2M HsP04 over mercury pool cathode Temperature, 100’ C.; stirring a t 33 r.p.m. (Note: some prereduction has occurred)
action between the reduced form of the catalyst and perchlorate to yield chlorate and reoxidized catalyst. Figure 3 shows the effect of HC1O4 on controlled potential reductions of No(V1). The applied potential is 0.00 volt us. S.C.E., a potential which corresponds to the early portion of the hto (VI) 2 Afo(V) reduction wave in the 4dP H2S04-2M H3P04 medium. It is immediately apparent that a steadystate condition is reached in the presence of excess HC104. Under these conditions, the total number of coulombs passed far exceeds the quantity necessary for the reduction of the initially present Mo(V1) alone. Since the only electrode process possible a t this potential in the absence of other reducible species is the reduction of llo(V1) to Rlo(V), i t can only be concluded that HC104interacts with Mo(V) to produce more Mo(V1) in solution. These conclusions are further substantiated by the fact that the solution obtained a t steady-state has a spectrum identical with that of Mo(V) in the same medium. No evidence for the presence of No(II1) could be found under these conditions, Since the steady-state current is reated to the concentrations of Mo(V) and HC104, i t should be possible to determine kinetic parameters from such measurements. HoLTever, the success of this approach is limited by difficulties of reproducing mass transfer conditions and the relative insensitivity of such measurements. These esperiments, homver, showed the nature of the shortcomings experienced in the study of the kinetics for the over-all reaction. Since hfo(V) is the active catalytic species, the rate of formation of C1- in the heterogeneous reaction is not only a function of the rate of reduction of Clod- in solution but also of the rate of production of Mo(V) from Mo (VI) a t the amalgam surface. Therefore, a rate law should be derived in terms of the active catalytic species’ concentration, not the total ooncentration of 1.10 in all oxidation dates.
A more practical way of approaching the kinetics of the Mo(V)-HC104 reaction is that of controlled current electrolysis where reproducible mass transfer conditions are not required since the total current passed is externally defined. Since the catalytic cycle involves the electrode reduction of Mo(V1) to AIo(V) and reoxidation of Mo(V) in solution by HC104, the rate of the homogeneous reaction can be controlled by the rate a t which ilIo(V) is made available a t the electrode. If the constant current electrolysis is carried out a t current densities far below those permitted by mass transfer considerations, the potential of the system will adjust itself in such a way as to permit only the Mo(V1) 5 Mo(V) reduction. A system containing initially onjy RIo(V1) and an excess of HClO, will, upon prolonged electrode reduction a t constant current, reach a steady-state condition in M hich the concentration of Mo(V) has adjusted itself to satisfy both the essentially constant perchlorate concentration and the reaction rate as externally imposed by the constant current. It should be possible to study the kinetic parameters of the homogeneous reaction without complications due to the heterogeneous production of Mo(V) by folloaing the Mo(V) concentration as a function of time. Figure 4 shows a typical plot of absorbance a t 430 mp us. time for two different applied currents a t identical total molybdenum and perchlorate concentrations. By varying the applied current and the initial perchlorate concentration the reaction orders with respect to ?\fo(V) and perchlorate, respectively, were both found to be first order by application of the Van’t Hoff differential method. Unexpected results were obtained n-hen the same techniques were applied to the determination of the effect of acidity on the homogeneous reaction rate. Figure 5 indicates that very
nearly the same steady-state concentration of Mo(V) is attained in 3 and 6M acid media, but that steady-state is reached more slowly a t the lower acidity. This is due to the fact that while the current passed and, therefore, the quantity of Mo(V) produced per unit time, is the same in both experiments; the fraction of Mo(V) in the dimeric form is much lower in 3114 than in 6 M acid. I n the less acidic medium the Mo(V) dimer concentration, which is related to the absorbance at 430 mp, will thus lag behind that in the more acidic medium until steady-state conditions are finally reached. This would seem to indicate that the rate-determining step of the homogeneous reaction is independent of acidity-a finding in direct contrast with those of the heterogeneous kinetic experiments. The differential rate law for the homogeneous reaction between Mo(V) and HC1O4may thus be expressed as Rate = K[HC1O4][Mo(V)]
where K = 2.8 X 10-4 liter mole-’ sec.-l for a typical rate of 4.4 X mole liter-’ see.-’ corresponding to an applied current of 5.35 ma. The determination of kinetic parameters by constant current techniques is made possible by the attainment of steady-state conditions through the opewtion of the catalytic cycle in this reaction. Such measurements have the important advantage that an accurate knowledge of instantaneous reaction rates is not critical, except insofar as the applied current must be constant and accurately measurable. On the other hand, these experiments provide only an indirect measure of the homogeneous interaction between perchlorate and Mo(V). This difficulty becomes especially critical when knowledge of the role of H + in the over-all reaction is desired. The problem was, therefore, ap; proached in an independent manner by measuring the disappearance of the VOL. 33, NO. 1 1 , OCTOBER 1 9 6 1
1475
i l l , , 0
40.
80
,I
,
,
,
,
,
,
,
120
I60
2w
240
280
320
360
Figure 6. Effect of acidity on homogeneous reaction between Mo(V) and HC104 at 100’ C.
400
T I M E , MINUTES
Figure 5. Constant current reduction of 0.01 8N HCIOd and 0.004M Mo(VI) over mercury pool cathode
Initial concentrations of Mo(VJmta~ and HClO, are constant
Temperature, 100’ C.; stirring a t 33 r.p.m.
hlo(V) dimer as a function of time in a reaction mi.xture containing only hIo(V) and HC104 initially. By preparing the Mo(V) externally and allowing i t to react with HC10, in the absence of any other reducing agent all variables of that part of the catalytic cycle involving the formation of Mo(V) are avoided and a direct measure of the honiogeneous interaction may be obtained. Figure 6 shows some typical plots of Mo(V) dimer concentration expressed as absorbance vs. time at three different acidities but identical perchlorate and total hlo(V) concentrations. It is immediately apparent that the effect of acidity on the rate of the over-all reaction is to determine the position of the monomerdimer equilibrium so that in 6111 acid essentially all of the hIo(V) exists as the dimer. Plots of log absorbance us. time show that the disappearance of RIo(V) dimer follows first-order kinetics a t constant perchlorate concentration in the 6-tl and 10M acid media. The apparent decrease in the rate of reaction for the 2 M acid medium is due to the fact that the absorbance at 430 mp is not related to hlo(s’) dimer concentration alone in this medium. A specific rate constant of 9.7 X 10- liter mole-‘ sec.-I is calculated from these data. The poor agreement of this value with that calculated from steady-state experiments may reflect the difference in experimental conditions insofar as species of
Mo(V) are in equilibrium for the honiogeneous experiments but not for the steady-state experiments. The role of H f , then, is one of controlling the concentration of the active Mo(V) species, namely the dimeric form, which in turn determines the rate of the homogeneous reduction of perchlorate. Thus an acidity term is to be expected in the empirical rate law which describes not only the interaction of i\lo(V) dimer with perchlorate but also the formation of the active Mo(V) species. On the other hand, the rate expression for the homogeneous reaction alone is indeed independent of acidity. at least over a moderate Concentration range. On the basis of these experimental findings a reaction mechanism may be proposed for reduction of perchlorate in the prmence of a molybdenum catalys t : H’
Mo(V1) polymer form A
Mo(V) monomer
Mo(1’I j polymer
1
\
H+
form B
H+
Mo(V) dimer
I
Mo(V) polymer
slow HCIO, ClOS-
I
+ Mo(V1) polymers
rapid Mo(V), Cd(Hg) C1Table
I.
Quantitative Reduction of Perchlorate
Relative
Standard
_ _ ~ Mmole ~
C104- in sample
C1- Found range of 8 detns.
Mean value 0 0919 0.0911-0.0922 0 0915 0 0198 0 0198-0.0211 0.0201
1476
ANALYTICAL
CHEMISTRY
Deviation
0.43 2 14
oxidation of Mo(V) becomes negligible, the hlo(V) present in the bulk of the solution can undergo further reduction at the surface of the reducing agent to produce Mo(II1). The above mechanism pertains to low concentrations of perchlorate in sulfuricphosphoric acid medium, and does not necessarily rule out other reaction paths, such as the catalytic cycle involving molybdenum(IV), under different experimental conditions. This reaction can be used for the determination of reasonably low concentrations of perchlorate by means of reduction to chloride. An amperometric titration procedure for the determination of C1- seems most useful for samples less than lO-3.U in Clod-. Higher C1- concentrations can be more rapidly determined by mercurimetric titration using a chemical indicator such as diphenylcarbazone. Table I summarizes some typical analytical data. These determinations were carried out by reducing the perchlorate in 100 ml. of 4M H2SOp2M H3P04 in the presence of 5 X 10-6M I\ianMo04 over 25 ml. of 4% Cd(Hg) reducing agent a t a temperature of 107’ C Under these conditions, reduction is completed in about 90 minutes.
+ Mo(V1) polymers
Polarographic evidence seems to indicate the presence of at least two forms of Mo(V1) in 61M acid. It seems likely that the Mo(V) dimer with its strong oxygen-bridging tendency would tend to react with Clod- via a n oxygen-atom transfer mechanism to produce C103and reoxidized catalyst. The resulting C103- is rapidly reduced by hfo(V) or excess reducing agent to finally yield chloride ion. After most of the perchlorate has been reduced, and the re-
LITERATURE CITED
(1) Burns, E. A., Muraca, R. F., ANAL. CHEM.32, 1316 (1960). (2) Carritt, D. E., dissertation, Harvard
University, 1947. (3) Crowell, W. R., Yost, D. M., Robert, J. D., J. Am. Chem. SOC. 62, 2176 (1940). (4) Enke, C. G., dissertation, University of Illinois, 1959. ( 5 ) Haight, G. P., ANAL.CHEW23, 1506 (1951). (6) Zbid., 25, 642 (1953). (7) Haight, G. P., Division of Analytical Chemistry, 134th Meeting, ACS, Chicago, Ill., 1958. (8) Haight, G. P., J . Am. Chem. SOC.76, 4718 (1954). (9) Haight, G. P., Sager, W. F., Zbid., 74, 6056 (1952)
( 1 0 ) Holtje, R., Geyer, R., 2. anorg. Chem. 246, 265 (1941). (11) Johnson, M. G., Robinson, R. J., -4NAL. CHEM. 24,366 (1952). (12) Laitinen. H. A,. Jennings. W. P., Parks. T. D.. IND. ENG.CH&.. A x . 4 ~ . ED. 18, 355 (1946). ( 1 3 ) Lindquist, Ingoar, Acta Chem. Scand. 5, 568 (1951). ~
(14) Page, J. A., Lingane, ,J. J., .anal. Chim. Acta 16, 175 (1957). (15) Roberts, I., IXD. Ex(; CIIEJI., ANAL.ED.8, 365 (1936). (16) Sacconi, L., Cini, R., J .Im. Chem. SOC.76, 4239 (1954).
alytical Chemistry, 139th Meeting, ACS, St. Louis, hlo., March 1961. Investigation carried out during the tenure of a Research Fellowship granted to G. A. Rechnitz by the National Heart Institute, United States Public Health Service. The financial support of the National RECEIVED for review May 10, 1961. .IC- Institutes of Health is gratefully accepted July 21, 1961. Division of Anknowledged.
EIectroa na lytical Techniques in Molten Lithium Sulfate-Potassium Sulfate Eutectic C. H. LIU' Brookhaven National Laboratory, Upton, Long Island, N. Y.
Three electroanalytical techniques for the determination of metal ions in molten lithium sulfate-potassium sulfate Couloat 625" C. were examined. metric titration with a potentiometric end point and chronopotentiometry with a solid electrode gave satisfactory results for the determination of copper (I), whereas determinations by direct potentiometry lacked precision. The chronopotentiometric technique involving the reduction of copper(1) to metallic copper gave much better results than the anodic process where copper(1) was oxidized to copper(l1).
E
LECTROANALYTICAL
TECHNIQUES
have been widely applied in molten salts in recent years including direct potentiometry (4, Q), polarography (3, I O , 13), chronopotentiometry (2, 8), and coulometric titration (7'). A list of literature references has been given by Van Norman in a recent article ( I S ) . I n most instances, alkali halide melts were employed as the solvents. Electrochemistry in the lithium sulfate-potassium sulfate eutectic melt was investigated in this laboratory and reported in a previous publication (11). The electrolytic decomposition of the melt was examined, and the osidationreduction potentialq of several electrode syhtemh were established. The objective of the present work is a critical evaluation of a few electroanalytical methods whirh are applicable in this melt. Direct potentiometry, chronopotentiometry, and coulometric titration with a potentiometric end point were compared in the determination of copper(1). APPARATUS A N D CHEMICALS
Furnace. Hevi-Duty crucible furnace, Type 51-506 (Hevi-Duty Electric Co., Milwaukee, Wis.) . Present address, Department of Chemistry, Polytechnic Institute of Brooklyn, Brooklyn 1, N. Y .
Temperature Controller. Wheelco indicating controller, Model 403 (Barber-Coleman Co., Rockford, Ill.). Constant Current Source. All-electronic current regulator constructed by the Instrumentation Division, Brookhaven National Laboratory. It is capable of delivering constant currents from 0.001 to 100 ma. with a precision of *0.5%. Potentiostat. Electronic controlled potential coulometric titrator built by the same Instrumentation Division according to the design of Kelley. Jones. and Fisher (6). Potential Recorder. High-speed Dynamaster recorder, Model 590 (The Bristol Co., Waterbury, Conn.). The pen speed is 0.4 second full scale, and the maximum chart speed is I inch per second. Solvent. The lithium sulfate-potassium sulfate eutectic (80'% lithium sulfate by mole; melting point 535" C.) was used at 625' C. The solvent was prepared in accordance with a procedure previously described (11). The eutectic mixture was first dehydrated and then filtered through a quartz frit while molten. Electrolytic Cell. The cell consisted of Vycor and quartz parts (11). Within the main cell, the melt was compartmented into separate portions by small quartz test tubes with fritted bottoms which served as salt bridges. Metal ions of interest were usually put into the melt by controlled anodization of the corresponding metals in foil or wire form. All experiments were performed under a dry argon atmosphere. The sulfate content of each compartment was determined at the end of a n experiment by converting sulfate into chloride through ion exchange and titration of the generated chloride by conventional argentometric methods. The concentrations of the solutes in the experiment could then be calculated. Reference Electrode. The reference electrode consisted of a silver coil in equilibrium with a solution of silver(I), generated by anodizing the silver coil in a fritted compartment a t a constant current for a measured period of time
(11). The silver(1) concentration was made approximately 0.05m, the exact value being calculated after analysis of the melt content of the compartment a t the end of the experiment. Indicator Electrodes. Copper foil indicator electrodes were used in direct potentiometry, and a palladium wire or foil served as both the generator and the indicator electrode in coulometric titration. Rectangular platinum indicator electrodes constructed from platinum foils welded to thin platinum contact wires were used in chronopotentiometry. Their areas, including the edges, ranged from 0.259 to 2.22 sq. em., measured micrometrically. The areas of the immersed portions of the contact wires were considered negligible. Chemicals. All chemicals were reagent. grade. EXPERIMENTAL RESULTS A N D DISCUSSIONS
Direct Potentiometry. Copper(1) solutions of various concentrations were prepared by a n electrolytic method which had been shown to be satisfactory (11). A copper foil electrode was anodized in a fritted compartment a t a constant current for a measured period of time. The current density was about 5 ma. per sq. cm. The potential at the electrode after generation was measured against a silver(1)-silver(0) reference electrode. The copper(1) concentration corresponding to each potential measurement was calculated after the determination of the salt content of the compartment. The copper(1)-copper(0) system obeyed the Nernst equation in the copper(1) concentration range from 4.8 X 10-6 to 4.7 X 10-2m. The measured potentials in each experiment vere corrected relative to the lm silver(1)-silver(0) electrode with the aid of the Nernst equation to facilitate the construction of a working curve. The logarithm of the concentration of copper(1) was plotted against VOL. 33, NO. 1 1 , OCTOBER 1961
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