STUDY OF ELECTROLYSIS OF METALLIC PERCHLORATES IN CELLOSOLVEL JOHN L. BEAL
AND
CHARLES A. MANN
Division of Chemical Engineering, University of Minnesota, Minneapolis, Minnesota Received July 85, 1057
The study of the electrolysis of solutions of metallic perchlorates in organic solvents conducted by Mann and Chaney (1) showed that in general these salts are very highly soluble, that their solutions in the monoethyl ether of ethylene glycol are of relatively high conductivity, and that from them many metallic elements can be electrodeposited. It is well known that from anhydrous solutions in organic solvents it is possible to electrodeposit some of the metals which are too reactive to be deposited from aqueous solutions. The alkalies and alkaline earths have been deposited from their halides and thiocyanates in alcohol, in acetone, or in pyridine solution. Patten and Mott (13, 14, 15, 16) and Muller and his associates (7 to 12) have investigated the electrolysis of many inorganic salts in common organic solvents by determining cathodic polarization curves, Le., curves showing the influence of current density upon cathode potential. From the forms of these curves and the values of the potentials they were able to observe the electrodeposition of metals even though they were so reactive that it was difficult to confirm the deposition by chemical analysis after the cathode had been removed from the cell and freed from its solution. During the course of the investigation by Mann and Chaney some evidence was obtained which indicated that aluminum might be electrodeposited from the solution of its perchlorate in the monoethyl ether of ethylene glycol. The scarcity of solutions, other than fused salts, from which aluminum can be deposited made this of interest. It was thought well to investigate deposition from this solvent further by using the polarization curve method of Patten and Mott and of Muller, This paper presents these results. In addition, the dielectric constant of the monoethyl ether of ethylene glycol and the conductivities of solutions of salts in it, even highly concentrated solutions, are reported, MATERIALS
The monoethyl ether of ethylene glycol was obtained from the Carbide and Carbon Chemicals Corporation under the trade name of Cellosolve. Submitted by John L. Beal to the Graduate School of the University of Minnesota in partial fulfillment of the requirements for the degree of Doctor of Philosophy. 283
284
JOHK L. BEAL AND CHARLES A . M A X 3
I t was dried with soluble anhydrite ( 5 ) , filtered, and vacuum-distilled from fresh turnings of metallic calcium. Fractionation with a SO-cm. packed column and discarding of the initial product gave a constant-boiling fraction whose boiling point a t 738 mm. was 1333°C. By use of the inanufacturer's vapor pressure-temperature chart this was corrected to 134.7OC. a t standard pressure. Solvent purified in this mariner had a conductivity as low as 9.3 X lop8 reciprocal ohms. Its physical propertied are given in the literature (17). Silver perchlorate was prepared from silver oxide and 70 per cent perchloric acid. The silver oxide was obtained by treating silver nitrate with sodium hydroxide. After two recrystallizations the silver perchlo-' rate was evaporated to dryness on a water bath and then further dried in a flask heated to 125-13OoC., through which dry purified air was drawn. The salt so obtained showed 52.02 per cent of silver as compared with the theoretical value of 52.03 per cent. Since it is very hygroscopic it was heated to abore 100°C. after each time that it had been exposed t o the atmosphere. Barium perchlorate was prepared by the method of Willard and Smith (18, 19). It was dried a t 250°C. in a current of dry air. By analysis it contained 40.83 per cent of barium as compared with the theoretical value of 40.85 per cent. Lead perchlorate was prepared from white lead and 70 per cent perchloric acid. After two recrystallizations as the trihydrate any excess acid 'i?'uq qteanied out by the method of Willard and Kassner (20). It was dried a t 160°C:. in a current of dry air, and was found to contain 51.01 percent of lead a:, comparrd n i t h the theoretical value of 51.02 per cent. Zinr perchlorate was obtained from the G. Frederick Smith Chemical i o . of Columbii;, Oliio, a i the. hexahydratr. The anhydrous salt is not known. ,iluniinum perchlorate it1 the anhydrous form is not knonn. All attempt> t o dry the hydratrd salt by heat and vacuum lrd t o decomposition, and drying agents, such as phosphorus pentoxide, do not remove thf-muisturr An electrolytic method of preparation of anhydrous solutionq in the monoethyl ether of ethylene glycol is reported by RIann and Chaney (1). This consisti in electrolyzing a solution of lead perchlorate in the solvc,nt using an alumiriuin anode and a lead cathode, the electric current causing the replacement of the diqsolred lead by aluminum from the anode. Preparation of such solutions was attempled by this method, but decomposition indicated by the presence of traces of chlorides in the resulting solutions. An alternative method for the preparation of solutions of a1urni1:uri~ perchlorate a as devised which depends upon the replacement of s i h w from an nnhydrous silver perchlorate solutim by contact with aluminum
ELECTROLYSIS OF METALLIC PERCHLORATES IN CELLOSOLVE
285
activated with mercury. This is a very slow process and several months are required for its completion. The solution remains colorless, chlorides are abwnt, and the solvent, after removal by evaporation in a vacuuni, is of unchangrcl boiling point. A solution made from 20 g. of silver perchlorate aiid 50 cc. of solverit in contact with an excess of granular amalgamated aluminum for thirteen months wa5 found t o contain no silver and the number of equivalents of aluminum was the same, within the limits of experimental error, as that of the perchlorate. This solution \\as used in the polarization measurements. DIELECTRIC CONSTAKT O F THE MOKOETHYL ETHER OF ETHYLENE GLYCOL
The dielectric constant of the monoethyl ether of ethylene glycol was determined by ineans of a capacity bridge supplied with alternating current of approximately 3500 frequency and tuned to the point of balance with two stage arpplified ear phones. One arm of the bridge, figure 1, contained two fixed condensers, C1 and Cz, each of 200 ppj’ capacity. The other arm contained a variable air condenser, C3, of approxiniately 400 ppf capacity with a sensitive vernier. The two variable air condensers, CZand Cg, in parallel were of 375 p p j and 35 p p j capacity. These latter two alone required calibration, and only the ratio of their capacities at one setting to that a t any other was needed. This calibration was accomplished by comparison with a small fixed condenser. The alternating current was supplied by the oscillator using a 171A tube in a circuit similar to the one described by Jones and Josephs (6). Its transformer consisted of three inductances,-L1 of 700 turns of KO. 28 enameled wire, LZof 800 turns tapped at 350, 500, and 650, and L3 of 100 turns, all wound on the same soft iron core but insulated from each other by several thicknesses of paper. A condenser of 0.1 p j was placed across L1. The power supplied to the bridge was controlled by varying the turns in LZby means of taps. The oscillator was shielded by a grounded copper box. Across C2 a variable resistance, R, was used to balance any current which leaked through the cell containing the solution. This resistance consisted of two binding posts beheen which pencil lines \ i n c x r i 4 .\rcording t o the interpretation of Patten and Nott (13 to TABLE 11 I ’ o l u i ~ i t a t i o n iiieasziremenis of 1.U
iv barium perchlorate
Platinum cathode
I
1.24 1.55 2.62 3.68
0.68 0.67 0.70
,
0.70
I
I1
-1.42 -1.52 -1.74 -1.93
F I 9~ Polarization curves of barium perchlorate with platinum cathode
16) a i d of Muller (7 to 12), the zinc reacts below the threshold value of current density, going into solution as rapidly as it is formed, and no reqidual amount is deposited. At higher current densities the rate of electrodeposition is greater than the rate of solution and zinc is actually deposited in a visible coat. The potential of the zinc is -1.40 xolts referred to the reference electrode used, and this can be converted to -0.62 yolt on the hydrogen scale.
ELECTROLTSIS OF METALLIC PERCHLORATES IN CELLOSOLVE
297
The curves for the electrolysis of barium perchlorate between platinum dertrodes are different from those of the previous metals. The cathode potential continues to increase with increasing current density, as shown in tahlc 11 and figure 9. The cathode decomposition potential, obtained by extrapolating to zero current density, is only -1.15 volts, and this is inurh less negative than would bc expected for the highly reactive metallic l~arium Thc potential of metallic barium cannot be measured directly in the wlution because its reactivity prevents the keeping of a clean metallic surfare. Muller, however, has cmployed the process of amalgamating TABLK 12 Polaizzalzon measurements of 1 0 N barium perchlorate ('admium cathode TOTAL VOLTAQE
CURREh-T DENSITY
ANODE POTENTIAL REFERRED TO Ag 0.1 AgClOd HALF-CELL
1
AV
CATHODE POTENTIAL REFERRED TO Ag 10.1 N AgClOl HALF-CELL
__.___-___--
volts
1.25 1.42 1.85 2.20 2.92 3.35 3.90 4.45 4.66
milliampere8 per spume centimeter
0.05 0.09 0.21 0.35 0.47 0.49 0.87 1.35 1.55 3.15 4.05 4.90 4.76 7.26
uolts
uok
0.25 0.35 0.57 0.68 0.68 0.68 0.68 0.69
-1.02 -1.13 -1.28 -1.56 -2.03 -2.24 -2.30 -2.28 -2.28 -2.29 -2.30 -2.40 -2.46 -2.48 -2.48 -2.49
0.70 0.71 0.72 0.73 0.73 0.77 0.77 0.78
reactive metals in ordrr t o obtain their potentials, and he identified a deposit as calcium by comparing its potential with that of amalgamated calcium in the same solution (9). In the 1.0 N barium perchlorate solution amalgamated barium was found to have a potential of -2.16 volts with respect to the reference half-cell electrode. This is so much more negative than the decomposition potential of the cathode that there is no possibility that barium was deposited during the electrolysis. The barium solution was electrolyzed using metal cathodes other than platinum. These were obtained by plating the desired metal upon the cell cathode. The cathode potential curves obtained with a copper cath-
298
JOHN L. BEAL AND CHARLES A. MANN
ode and with a zinc cathode were rather similar to those obtained with platinum, and the potentials were too positive to indicate the deposition of barium. The results with a cadmium cathode, given in table 12 and figure 10, do indicate the deposition of metallic barium. Above the threshold value of 0.5 milliampere per square centimeter barium is depositod a t a potential of -2.28 volts. The favorable influence of the cadmiurn cathode can be explained by the tendency of barium to form alloys with cadmium, as reported by Gautier (4). Chemical confirmation of the deposition was poor, for with removable cathodes any attempts to remove the cell liquor resulted in the removal of the deposit as well, and only traces of barium could be reported. But I
,
P I C
I
d d
”
FIG.10. Polarization curves of barium perchlorate with cadmium cathode
when the cathode was removed quickly and placed in distilled water, small amounts of hydrogen gas could be seen t o be evolved. For the deposition of aluminum one would expect copper to be a favorable cathode metal, for several compounds of copper and aluminum have been reported (3). Table 13 and figure 11 show the results of the electrolysis of a 0.5 N solution of aluminum perchlorate prepared as previously described. The irregularity of the curves is similar to that found by Muller for the electrolysis of aluminum iodide in pyridine (9). No cathode potential value is negative enough to indicate the deposition of metallic aluminum. For comparative purposes the potential of aluminum was measured in the same solution. When freshly cleaned and scraped, aluminum had a potential of -0.67 volt with respect to the 0.1 N AgClOrAg
ELECTROLYSIS OF METALLIC PERCHLORATES IN CELLOSOLVE
299
TABLE 13 Polarization measurements of 0.6 N aluminum perchlorate Copper cathode !ATEODE POTENTIALREFERREDTO A g 0.1 N AS104 HALF-CELL
TOTALVOLTAGE
CURRENT DENBITY
volts
nillianperes per square centimeter
volts
0.47 0.74 0.77 0.88 0.93 1.02 1.06 1.23 1.50 1.85 1.95 2.05 2.12 2.27 2.38 2.63 3.05 3.38 3.85 4.43 1.47
0.03 0.19 0.25 0.37 0.45 0.61 0.74 1.00 I .30 li 75 1.95 2.20 2.45 2.97 3.85 4.75 5.70 6.25 7.70 9.25 11.60
-0.26 -0.35 -0.35 -0.47 -0.48 -0.60 -0.65 -0.93 -0.85 -0.80 -1.08 -1.30 -1.20 -1.15 -0.95 -0.95 -1.10 -0.79 -0.80 -1.10 -0.90
1
FIG.11. Polarization curves of aluminum perchlorate
electrode. This value is obviously less negative than the true potential of the metal. Samples amalgamated by rubbing the cleaned surface with mercury had potentials ranging between -1.59 and - 1.61 volts. Samples
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JOHN L. BEAL AND CHARLES A. MANN
amalgamated by dipping in aqueous mercuric chloride and drying with alcohol and ether gave potentials between - 1.83 and - 1.94 voks. While these values are not in agreement, they are enough more negative than the cathode potentials to indicate the absence of aluminum deposition during the electrolysis. CONCLUSION
Both the dielectric ronstant' of the monoethyl ether of ethylene glycol and the conductance of its solutions are between those of methyl and ethyl alcohols. The monoethyl ether of ethylene glycol may be considered as an alcohol of high molecular weight, and as such the high molecular weight does not, in this case, show the usual effect of decreasing conductance. The polarization curves show that' metals even as reactive as barium may be elect'rodeposited from solutions in the monoethyl ether of ethylene glycol. Although anhydrous solutions of aluminum perchlorate may be prepared in this solvent, the polarization curves show no tendency toward the electrodeposition of aluminum from such solutions. REFERENCES (1) C H A N EAND Y NA~W: J . Phys. Chem. 36,2289 (1931). (2) D A V I E SConductivity : of Solutions, p. 51. John Wiley and Sons, Kew York (1929). (3) F R I E N DInorganic : Chemistry, Vol. IV, p. 61. Chas. Griffen and Co., London (1917). (4) GAUTIER: Compt. rend. 134, 1108 (1902). (5) HAMMOND AND WITHROW: Ind. Eng. Chem. 26, 653, 1112 (1933). (6) JONESAND JOSEPHS: 3. Am. Chem. Soc. 60,1049 (1928). (7) MULLER: hlonatsh. 43,67 (1922). (8) MULLERAXD D U S C H E KNonatsh. : 43, 75 (1922). AND (9) MULLER,HOLZL, KNAUS,P L A N Z I G , PRETT:Monatsh. 44, 219 (1923). (IO) MULLER,HOLZL, PONTOTI, . ~ N DWINTERSTEINER:Uonatsh. 43, 419 (1922). (11) l i C L L E R , HOXIG,A K D KONETGCHNIGG: Monatsh. 44, 236 (1923). (12) . \ I ~ % L E PIXTER, R, AND PRETT:Monatsh. 44, 525 (1924). (13) PATTEIV: J. Phys. Chem. 8,483 (1904). (14) PATTES:J . Phys. Chem, 8, 548 (1904). (15) PATTEX A K D MOTT:J. Phyo'. Chem. 8, 153 (1904). (16) PATTEN AND ~ ~ O T TJ.: PhyS. Chem. 12, 49 (1908). (17) REEDASD HOFFMAN: Ind. Eng. Chem. 20,497 (1928). (18) SMITHA N D WILLARD:Ind. Eng. Chern. 46,287 (1923). ASD WILLARD:Ind. Eng. Chern. 19, 411 (1927). (19) S M I T H (20) WILLARDAND K A S ~ N E J. R :Am. Chern. SOC.62, 2391 (1930).
