ELECTRODEPOSITION OF METALS FROM ASHYDROUS AhIMOSIA* BY HAROLD SIMMOSS BOOTH AND MEXAHEM MERLUB-SOBEL
That the aqueous world we chance to live in is but one of the many possible, was first indicated by the epoch-making studies of Franklin on the ”ammonia system;”’ with Franklin’s work the chemical equivalent of Ptolemaic geocentricism came t o an end. No less demolishing of previous trends of thought -and creative of new-were the further conclusions drawn by I(raus2 from his researches on certain electrochemical phenomena in liquid ammonia, for it is to Kraus we are indebted for the concept of the ionization of the elementary alkaline and alkaline earth metals in ammonia to yield normal metal ions and free (though generally solvated) electrons. It was this same research pyramid which yielded the concepts currently accepted regarding the nature of metallicity and the mechanism of electrical conduction in metals. The monumental researches on the electrochemical phases of ammonia solutions, by others as well as by Kraus, have, peculiarly enough, included extremely little concerning actual electrodeposit,ion of metals from this solvent. This fact is all the more surprising when viewed in contrast with electrodeposition studies made in the far less electrochemically understood pyridine3 or in contrast with Rohler’s4 work in the otherwise almost unstudied formamide. It was the purpose of this study to fill this gap. Theoretical Considerations As an electrolytic solvent, ammonia presents somewhat of an anomaly. Its relatively low dielectric constant ( 2 1-23) would normally be expected to allow for no great aptitude toward electrolytic conduction, yet its power in this respect is almost the equal of water. Early in their studies on this solvent Franklin and Kraus5 concluded that dissociation is far less in ammonia than * Abstracted from a thesis submitted to the faculty of the Graduate School of Western Reserve University hy Menahem hlerluh-Sohel in partial fulfillment of the requirements for the degree of Doctor of Philosophy, 1930. The experimental work reported upon here was done in the period of 1926-1928; previous publication was prevented, however, by the illness of the junior author. Recently, similar studies on ammonia electrode osition were announced by Taft and Barham in the May, I 930, issue of The Journal of Khysical Chemistry, representing partial duplication -hut partial only-of this work. Franklin: Am. Chem. J., 47, 285-317 (1912); J. Am. Chem. Soc., 27, 820-51 (1905); Proc. Eighth Int. Cong. Applied Chem., 6 , 119-30 (1912); Seealso Franklin: J. Am. Chem. SOC., 46, 2137-51 (1924). Kraus: J. Am. Chem. Soc., 30, 1323-44 (1908); 36, 864-77 (1914);43, 749-70 (192ri; 44, 1216-39 (1922); Trans. Am. Electrochem. Soc., 21, 119-20 (1912); Kraus and Lucasse: J. Am. Chem. SOC.,43, 2529-39 (1921). Muller et al.: Monatsheft., 44, 219-30 (1924). Rohler: Z. Elektrochemie, 16, 419-36 (1910). Franklin and Kraus: Am. Chem. J.,23, 297-8 (1900).
‘
3304
HAROLD SIMMONS BOOTH A S D MENAHEM MERLUB-SOBEL
in water, but that its great fluidity results in so much greater ionic migration velocities as to counterbalance completely the factor of a low dielectric constant. Comparing ammonia a t -33OC. with water a t I ~ O C . , the fluidity ratio is of quite significant a magnitude, 4.16: I , ~but even the interplay of dielectric constants and fluidities is admittedly insufficient to explain completely why some solvents prove excellent conducting media and others fail miserably in this regard. Consideration of the theoretical cause of high ammonia conductivities is of considerable consequence even in practical electrodeposition. Whether a metal will be deposited out of solution in any solvent where hydrogen ions are a product of the ionization of the solvent must depend upon the relative values of two potentials-that given by E M=RT/nF. In Pdp,,+nhf and that of E H= RT/F. In PHg/lJH+ fnH where Plf represents the electrolytic solution pressure of the metal. PH)represents the electrolytic solution pressure of hydrogen. pzf represents the osmotic pressure of the metal ions in solution. PH* represents the osmotic pressure of the hydrogen ions in the solution due to primary ionization of the solvent, and to any added ionogen dissociating to give €I+. nLl represents t,he overvoltage of nietal deposition. n H represents the overvoltage of hydrogen evolution. Since the osmotic pressure of the metal ions in solution will always be a function of their concentration, deposition will be facilitated, all other factors being equal, by increased ionization rather than by a mere high conductivity resulting from low viscosities. Conversely, the tendency toward metal deposition will rise and fall with decrease and increase of the actual hydrogen ion concentration. What the exact value of this last concentration is for pure ammonia is unknown, but it is undoubtedly very low.2 Some estimate may be made from the data of Carval10,~in which the conductivity of ammonia, after standing, is given as 5 x 10-10 at zo°C. and 3 . 7 x 10-l’ at -8o’C. By the method of Kohlrausch, and ut,ilizing the ion conductances, a t - 33’C., found by Kraus and Bray4 for the NHA+ and KH2- ions in ”3, we may compute the approximate elementary ionization of ammonia into Hf and XH2- as yielding concentrations of (5
x
I0-l’
x
1 0 3 ) / ( 1 3 1 f 133)
0.02
x
IO-’
as compared with the 0.8 x IO-? accepted for water a t 1 8 O c . The use of the S H 4 + ion conductance as an H+ value is amply justified by the fact that rc”,’.in liquid ammonia is merely a solvated hydrogen ion, and might better be written as H(NH3)+. Kraus: “The Properties of Electrically Conducting Systems,” p. 109. “The Properties of Electrically Conducting Systems,” p. 230. Carvallo: Compt. rend., 156, 175j-R (1913). Kraus and Bray: J. Am. Chem. Soc., 35, 1335 (1913).
* Kraus:
ELECTRODEPOSITIOX O F METALS FROM ANHYDROCS AMMONIA
3305
On this basis, the hydrogen ion concentration of ammonia is approximately 1/40 that of water, which would lead to the expectation that electrodeposition of metals would prove materially easier, even though the advantage of the lower hydrogen ion concentration in ammonia is decreased in part by the simultaneously lower ionic concentration of dissolved salts, as previously discussed. Current Status of Aqueous Electrodeposition Metals which have been deposited by electrolysis of aqueous solutions of their compounds are quite sharply delineated from those where failure results. Those which can be deposited include all metals in Group I-B of the periodic table (Cu, Ag, Au) and all Group 11-B if the uncertain beryllium and magnesium are assigned to the alkaline earths, leaving only zinc, cadmium and mercury. The same situation applies to Group 111-B, gallium, indium and thallium all being amenable to electrochemical treatment, as opposed to the failure registered by all other inherently trivalent metals, including the rare earths. In the fourth group of the periodic table, germanium is somewhat in doubt, but otherwise the rule continues t'o hold that only the minor, or B, division can be deposited (tin and lead, in this case) : the fifth group is normal to the extent that arsenic, antimony and bismuth deposit readily, while the other elements of the group are not obtainable in this manner.' In the whole of the sixth group, tellurium and chromium are the only metallic representatives which permit deposition, the failure of molybdenum and tungsten to respond to electrolytic operations being somewhat of an electrochemical mystery in view of the ease of deposition of the more active chromium. The seventh of the periodic table groups holds but one metal as yet subjected to any study-manganese; according to Allmand and Campbell? it deposits only because, fortuitously, an extremely high hydrogen overvoltage acts to prevent hydrogen being released (at least so completely as to exclude all metal deposition). Manganese, therefore, is about at the ve,ry border line for aqueous deposition. The transition, or eighth, group, with its three triads, represents elements all of which can be electrolyzed out of water solution. On the other side of the picture are the metals not amenable to deposition practice. They include the alkalies in Group I-A, the alkaline earths of 11--4, the boron-aluminum rare earth grouping of 111-A, the tetravalent metalloid congeners, the high melting point metals of the fifth group, and, finally, tungsten, molybdenum and uranium in the same group that harbors the relatively easy-to-deposit chromium. I t is immediately apparent that the high e.m.f. of deposition which explains readily enough our inability to obtain metallic deposition with the alkali or alkaline earth members, or aluminum and the rare earths, is not broad enough to cover all metals of the non-electrolyxable groups. For example, tungsten and molybdenum hold approximately the same position Personal unreported work by the junior author has negatived the claims of deposition made for tantalum. Allmand and Campbell: Trans. Faraday SOC., 19, j j 9 - 7 3 (1924).
3306
HAROLD SIMMOSDS BOOTH A S D MENAHEM MERLUB-SOBEL
in the normal aqueous series as does mercury, which is below hydrogen,' yet neither of these metals can be deposited from water solution; the work of Mann and Hahersen* seems to indicate that the presence of the slightest trace of the OH grouping, even if organic, is fatal to the deposition of t~ngsten.~ h completely satisfactory explanation of these phenomena is yet to be found. On the basis of unreported work, the authors, with data admittedly inadequate a t present to make their viewpoint much more than a working hypothesis a t best, are of the opinion that in cases where aqueous electrodeposition fails despite a relatively low position in the electromotive series, the situation is merely a reflection of the paucity or absence of elementary or simple metallic ions. Since even the alkali metals can be formed, in quasi-free state, by ammonia electrolysis, failure of other metals to deposit, if and when failure is evidenced in liquid ammonia, would be because of some such factor alsoor, a t least, because of factors similar in character to those operative in water, whatever may be their true nature. Previous Work in Anhydrous Ammonia of Electrodeposition Interest That ammonium nitrate absorbs ammonia a t room temperatures, forming an electrolytic solvent, has been known ever since the work of Divers4 indicated that this solvent (which we now recognize as merely a concentrated solution of nitric acid in anhydrous ammonia) electrolyzes to yield hydrogen a t the cathode and one third the volume of nitrogen a t the anode, if platinum and iron positive electrodes are utilized. When silver, mercury, lead, copper, zinc, or magnesium anodes are employed, however, electrolytic corrosion takes place. Kothing is indicated by Divers as to whether cathodic deposition of a metal would occur after sufficient metallic ion is electrolyzed into the solution. The work of Booth and Torreyj gives good reason to believe that such deposition would replace, in part a t least, hydrogen evolution. When first the systematic study of ammonia solutions began, Cady6 showed that the passage of current through a silver, lead, mercury, copper, or barium salt solution in anhydrous ammonia resulted in the deposition of the metal on the cathode. Nothing is stated by Cady regarding the nature of the deposit obtained. Since the alkali metals are so soluble in ammonia, it is but to be expected that electrolysis of salts of these metals would yield, cathodically, solutions of the metals, more or less concentrated as conditions permitted. Operating ~~
'Russell and Rowell: J . Chem. SOC.,130, 1881-92 (1926). * Mann and Halversen: Trans. Am. Electrochem. SOC.,45, 493-508 (1924). Since submission of this manuscript, Colin G. Fink and Frank L. Jones: Tram. Am. Electrochem. SOC.,April (1931), have succeeded in plating tungsten from aqueous alkaline solutions. Divers: Phil. Trans., 163, 359-75 (1873): R o c . Roy. Soc., 21, 109-11 (1873). Booth and Torrey: J. Am. Chem. Soc ,52,2581 (1930); J Phys. Chem. 35, 2465,2492, 3111 (1931). fi Cady: J. Phys. Chem., 1, 707-13 (1897).
ELECTRODEPOSITION O F METALS FROM ANHYDROUS A M M O M A
3307
a t - 7ooC. Ruff' electrolyzed a K I solution, and obtained drops of a coppercolored material known to be but a concentrated solution of potassium in ammonia; this, as soon as permitted out of the concentrated K I phase, redissolved in the ammonia with the blue color characteristic of the alkalies. Cottrell,*in his work on ammonia-acetylene reactions, has reported magnesium as slightly soluble in ammonia; on electrolysis, blue streaks developed from the cathode and an almost microscopic tree-like growth of bright metal grew out along the glass container. Quite evidently, magnesium is at the very threshold of alkalinity, if ammonia solubility is to be used as a criterion. Like the others of the top row of the periodic system, beryllium is sui generis. ?jot truly an alkaline earth, as evidenced by full insolubility in ammonia, it is yet too high in the aqueous e.m.f. series3 for electrodeposition. Booth and Torrey' have shown that beryllium can be deposited from anhydrous ammonia solutions of beryllium chloride and of dehydrated beryllium nitrate, their studies representing the first active effort at utilizing ammonia directly as an electrodeposition medium. Studies on electrode potentials in liquid ammonia, of interest to electrodeposition because electrode potential values determine the electromotive series in a given solvent, have been made by Cady,j who found a tendency of electrodes, copper and zinc particularly, to vary in their potentials when measured against their ammonia salt solutions. Johnson and Wilsmore,6 operating in the same field, found potential values in ammonia generally higher than in water; their tests included si'lver, mercury, copper, lead, nickel, cadmium, zinc, ammonium, magnesium, calcium, sodium and potassium. It is noteworthy that the potassium and sodium values are almost identical; these check quite closely with the values announced by Forbes and Sorton' for the oxidation potentials of sodium and potassium. In studying the true nature of oxidation, Cady and Tafts electrolyzed solutions of TI1 and CUI in ammonia. Electrolysis of thallous iodide resulted in oxidation to Tl+++ions, with simultaneous deposition of metallic thallium on the cathode. The deposited metal bridged over rapidly-a tendency of thallium no less in ammonia electrolyses than in aqueous, as we shall see later-although the use of I I O volts would obviously have made massive and coherent electrodeposition of any element impossible. The cuprous salt, on electrolysis, oxidized to yield the blue solution characteristic of cupric salts. 'Ruff:Ber., 34, 2604-7 (1901). Cottrell: J. Phys. Chem., 18, 85-100 (1914). Latimer: J. Phys. Chem., 31, 1267-9(rgq),is of interest in this connection, the data being based on thermodvnamic considerations rather than actual electrical measurements. See, however, Bodforss:-Z. physik. Chem., 124,66-82 (1926)and 130, 82-9 (1927),in which beryllium is placed between cadmium and zinc. Booth and Torrey: loc. cit. Cady: J. Phys. Chem., 9, 477-503 (1905). Johnson and Wilsmore: Trans. Faraday SOC.,3, 70-80 (1907). ' Forbes and Norton: J. Am. Chem. SOC.,48, 2278-85 (1926). Cady and Taft: J. Phys. Chem., 29, 1057-74 (1925).
3308
HAROLD SIMMONS BOOTH AND MENAHEM MERLUB-SOBEL
In strict analogy with the viewpoint that anodes are “oxidized” electrolytically in water solution, the very apposite term “nitridation” has been synthesized t o represent the analogous de-electronization process in liquid ammonia. Utilizing NH4K3 solutions (effectively, ammonated hydronitric acid solutions), Browne, Holmes and King’ found that copper, silver, cadmium, lead and antimony are electrolytically corroded, with current efficiencies slightly above I O O ~ which ~ , the authors explained as being due to either mechanical abrasion or purely chemical corrosion, or perhaps to a tendency toward the formation of compounds containing the metal in a lower state of valence. Closest in approach to actual electrodeposition studies in ammonia was the work on decomposition potentials and metal overvoltages, recently reported by Groening and Cady.* Nitrate and chloride decomposition voltages proved lower in ammonia; iodides and nitrites were higher than in water. In general, metal overvoltages were indicated as higher than in aqueous solution. I t was in conjunction with overvoltage measurements that metals were deposited on to cathodic surfaces; the metals studied were silver, nickel, cadmium, mercury, lead, zinc and iron. Scope of this Research What little was available in the literature regarding electrodeposition in ammonia, obtained as a by-product of other work, was far from encouraging. Groening and Cady3 declared, in explanation of some of their results, that “the metal seldom deposited in smooth form.” This, despite electrode rotation, which normally aids materially toward obtaining good plating. Kraus,* too, summarizing his vast experience in the field, pointed out that “metals in the electropositive condition, when precipitated from liquid ammonia solution, almost invariably appear in spongy form.” The problem here, therefore, integrated itself into the question of whether ammonia cathodic electrolyses inherently did not permit of bright, smooth, adherent deposits, such as are produceable in water, or whether, by proper manipulation of current densities, solutions, and perhaps temperatures, re. sults analogous to aqueous deposition, and on a par with it, could be obtained. The criterion of good deposition was, as always, a purely visual one: a deposit, to be considered satisfactory, had to possess, first and foremost, adherencethen compactness-and, finally, typical metallic lustre (the last two factors as opposed to the “burnt” spongy type of deposit encountered in ordinary electrodeposition when current densities are too high). The Apparatus and its Manipulation Since ammonia boils, a t normal pressure, at about -33’C., it was most logical to attempt all electrodepositions first at this temperature, and most of the studies were therefore carried out at about the boiling point of am-
’ Browne, Holmes and King: J. .4m. Chem. Soc., 41, 1769-76 (1919).
* Groening and Cady: J. Phys.
Chem., 30, I jgj-161j (1926). Groening and Cady: loc. cit. Kraus: Trans. Am. Electrochem. Soc., 45, 175-86 (1924).
ELECTRODEPOSITION O F METALS FROM ANHYDROUS AMMOXIA
3309
monia, to eliminate pressure problems. Only when complete failure a t low pressures and temperatures was definitely evidenced for certain metals was a change made to a room temperature, high pressure unit. The normal pressure arrangement is shown diagramatically in Fig. I . The anhydrous ammonia passed directly through the steel valve into the glass inlet tube of a barium oxide drying tube. Beyond the BaOl tube, stopcocked connections were made to a high vacuum pump and to a supply of dry nitrogen. Another stopcock permitted closing off the operating cell from this preliminary section. Pressures were measured by a manometer set between this point and the cell proper, both being further protected by
FIG.I Unit No. I Sormal Pressure, Low Temperature Cell
stopcocks. Ammonia passed into the cell through the first angle bend, and was liquefied by external cooling; it could be forced out a t the end of the run through the outlet tube (see cell detail diagram, Fig. 2 ) . The outlet was connected, through a stopcock anti an unwired rubber connection, to a trap holding about one inch of mercury, which acted to prevent influx of air during agitation. For the cell, a large-lipped glass tube was fitted with a four-hole rubber stopper and was set through a cushioning cork ring. Rubber stopper and cork were drawn together tightly by the use of iron plates and wing nut bolts. Through the stopper came the two electric leads, while a third tube brought the ammonia in, as gas, and otherwise connected the cell to the preliminary vacuum pumping and nitrogen system; the last of the tubes was the liquid l Booth and McIntyre, in Ind. Eng. Chem., Anal. Ed. 2, 1 3 (19301, have shown that barium oxide is a very effective desiccant for ammonia.
33'0
HAROLD SIMMONS BOOTH A N D M E S A H E M MERLUB-SOBEL
ammonia withdrawal tube, reaching to the bottom of the cell. The lower half of the electrolysis cell was encircled by a Dewar flask; for preliminary condensation] liquid air or COz-ether was used, while the electrolysis proper was conducted with the solution surrounded by an ammonia bath, thereby assuring operation in the vicinity of -33OC. For cell operation, a weighed amount of electrolyzing salt was added and the unit immediately assembled. After repeated evacuation and filling with
/RON
PLAT&
FIG.2 Sormal Low Pressure Cell
nitrogen, to remove adsorbed moisture fro? the system, pure anhydrous synthetic ammonia, dried by metallic sodium, was drawn from the cylinder into the cell. A COz-ether mixture is preferable for liquefying ammonia, since it does not cause solidification. After sufficient ammonia had been liquefied] the preliminary system was cut off from the cell, leaving only the manometer connected. Agitation of the liquid in the cell, so as to obtain complete solution of the salt or, a t least, saturation] was accomplished by a rapid sequence of opening and closing the stopcock controlling the liquid ammonia outlet tube in the cell. Under the few centimeters mercury pressure present, the liquid ammonia solution would rise in the tube, but after ascending about five or ten centimeters, it would reach a portion of $he tube considerably warmer than the liquid ammonia, and a powerful back pressure would therefore develop by the flashing of the ammonia into vapor, and this pressure
ELECTRODEPOSITION O F METALS FROM ASHYDROUS AMMOXIA
33 I I
would force the liquid back with some violence, making thorough agitation of the solution possible without external appliances. The mercury trap, of course, prevented any drawing back of air or moisture into the lines. After electrolysis had been cow>leted, the mercury trap was replaced by a loosely stoppered Dewar flask. By opening the stopcock, the pressure in the cell, increased by the removal of the external ammonia bath, forced the contents over into the Dewar, from which point the solution was discarded. The stopcock was closed and a fresh quantity of ammonia condensed in the cell in the manner originally used, in order to wash the electrodes. When this had been drawn off as before, the system was evacuated partially, and R
FIG.3 Unit No. 2 High Pressure, Room Temperature Cell
ethyl alcohol drawn into the cell for further washing, by placing the ammonia outlet tube tip in a beaker of alcohol, and opening the stopcock. Alcohol washings mere continued till the electrodes showed freedom from all salts, and then the cell was taken apart. After a large amount of experimentation had proven low temperature electrolysis ineffective for a number of metals, an entirely different type of cell had to be devised, to operate at room temperatures and therefore at pressures of several atmospheres. The preliminary part of this apparatus was identical with that used for low temperature work, but here the barium oxide tube was followed by a measuring tube into which the ammonia was first condensed before being revolatilized into the operating cell. The arrangement is shown in Fig. 3 .
3312
HAROLD SIMMONS BOOTH AND MENAHEM MERLUB-SOBEL
The complete high pressure unit consisted of a standard all-iron valve capable of withstanding 2 5 0 lbs. pressure, and specially packed so that it held vacuum also. A steel pipe cross connected the valve down to the cell and up to a standard ammonia pressure gauge. The remaining side of the cross was closed from the ammonia line, but was tapped to receive a supporting rod, enabling the whole unit to be clamped as desired. Relatively little difficulty was experienced in any part of the apparatus other than in the cell itself. It was obviously impossible to utilize the steel
FIG.4 High Pressure Cell
container as anything but a pressure-resistant jacket; electrolysis in a steel tube would have involved the risk of the tube acting as an intermediate cathode, and throwing iron into the deposit as well as into the solution. A number of attempts at enameling the inside of the tube proved utter failures. The arrangement shown in Fig. 4 was finally adopted, after considerable experimentation: About an inch of mercury was placed in the jacket, and a hard glass tube about one quarter of an inch smaller in diameter forced down into it, giving the tube a cushion of mercury. Into the tube, from above, passed the two electrodes, and pressure electrolysis took place therefore inside of a glass tube completely protected from the effects of such pressure by virtue of its being completely equalized.
ELECTRODEPOSITION OF METALS FROM ANHYDROUS AMMOSIA
33 13
One factor did continue to give trouble for a long period. The tube carrying the platinum seals had to be set through the cell cap, and no leakage was to be permitted. Repeated efforts a t using DeKhotinsky cement and its modifications, low melting alloys, and even adamant cements, all proved futile. The method finally adopted, simple enough admittedly, involved the use of tiny rubber stoppers through which the tubes passed. By setting these stoppers with taper upward, the cap being drilled accurately so as to make a perfect fit, increased pressure only served to tighten the seal. Complete sealing between the cap and the cell was accomplished by a rubber gasket, and proved quite adequate. Upon completion of a run in the high pressure unit, the cell was cooled to the normal boiling point of ammonia or lower, quickly taken apart, and the electrodes plunged immediately into anhydrous alcohol. Had individual runs been more fruitful of results, more elaborate washings, ammonia and other. wise, would have been called for, and the necessary devices could readily have been set into the cell to withdraw the ammonia and to wash the electrodes in situ. However, there was no occasion for this, as further reports on the individual metals will indicate. Experimental Results Khile many salts have been listed as ammonia soluble, most such determinations were made at room temperature. Solubilities at -33OC. proved significantly lower and thereby curtailed the choice of salts materially. In each case, those salts are reported upon which proved most satisfactory after preliminary tests. Copper. Preference was given to the cuprous salts, even though none were available which did not show some amount of cupric contamination (recognizable by color of resulting solutions). This latter factor was of little importance, since electrolysis would have oxidized the cuprous ion to cupric anyway, at least to some extent, as has been indicated by the work of Cady and Taft.' Best results were obtained in a run involving the solution of 0.3330 grams of CUI in 3 0 C . C . of ammonia, with a current of approximately I O milliamperes on a 6 square centimeter cathode face, equivalent to 167 milliamperes per square decimeter. The voltage required was 0 . 7 for an electrode separation of 1 2 mm. S o hydrogen or other gas was evolved. Based on cuprous copper, anode corrosion was 8 1 . 2 7 ~and ~ deposition only I 1 . 8 ~ ~ ? , , while from a cupric standpoint anode loss was 1 6 2 . 4 7 ~of theory and the cathode gain 2 3 . 5 % . High corrosions and low depositions were characteristic of all copper runs; even when current densities were too high for good plating this held true. Burnt deposits suffered further deterioration after the electrolysis was completed, irrespective of whether kept in a desiccator or in the open. Silver. As might have been expected, silver gave an excellent deposit. A weight of 0.3085 grams of pure silver iodide dissolved extremely easily Cady and Taft: J. Phgs. Chem., 29, 1057-74(1925).
3314
-
HAROLD SIMMONS BOOTH AND MENAHEY YERLUB-SOBEL
in 40 C . C . of ammonia, and with a silver rod anode and a gold cathode, a current of 2 1 j milliamperes per square decimeter was employed. Anode corrosion was somewhat under theory, 8 6 . 0 % , deposition slightly above, I O I . z % . Gold. Both AuI and the complex cyanide made in situ by the interaction of AuI with KCN gave good plating, the former operating perfectly a t 2 j o milliamperes per square decimeter, on a platinum cathode and a gold anode, while the latter required a lower density of current, 87 milliamperes to the square decimeter giving the best results. Voltage for the iodide was about 1.5, the cyanide requiring about 2.0 for the same electrode separation, 14 mm. For the best gold iodide run, 0.4963 grams were dissolved in z j C.C.of ammonia; the double cyanide received 0.2294 grams of AuI and 0.1142 grams of KCN. Both depositions were perfect, and throwing power complete; the cyanide bath gave results the equal of any aqueous cyanide electrolysis. Attack of the gold anode was practically nil in all iodide electrolyses; for the cyanide bath it was 3.6Yc. Cathode efficiencies were quite good, hovering around the 87,6Yc value found for the best iodide run, while cyanide gave 82.67,. Beryllium. The findings of Booth and Torrey, previously mentioned, were corroborated in this research; here beryllium iodide was used because of its greater solubility at low temperatures, though even this was far from adequately soluble. A weight of 1.5435 grams of salt, synthesized by reaction between iodine and beryllium metal flake, failed to dissolve to more than a small extent at first in 2 5 C . C . of ammonia. A cast beryllium plate, about one-eighth of an inch thick and I X 3 centimeters in area, was used as an anode and was set about 1 8 mm. from the gold cathode twice the area. As electrolysis continued, more of the salt went into solution, doubtlessly due to internal heating effects of the high currents used, varying from 12,300 to 14,200 milliamperes per square decimeter. The voltages required were from 8.2 to 2 7 . 0 , the rise in internal temperature making the higher voltage range necessary. Complete metal deposition, both front and back, resulted; the metal was dark and far from the usual steely color of polished beryllium, or the gray of the unpolished, a fact hardly remarkable considering the conditions of operation. Zinc. A dense deposit of fine, though somewhat matty, metal was obtained by the use of zinc cyanide with 0 . 2 0 7 0 grams of the salt dissolved in 40 C . C . of ammonia and with 20 mm. electrode separation. Current was equivalent to 74 milliamperes per square decimeter; voltage averaged somewhat under 2. Deposition was good, as was also t,hrowing power. Anode corrosion was I 1o.3YG, cathode deposition 9 0 . 4 7 ~ . Cadmium. Plating of cadmium was tried with two soluble salts, potassium cadmium cyanide and cadmium thiocyanate. The former gave bright metzl, although there was a tendency toward spottiness and burning, with 667 milliamperes per square decimeter, The thiocyanate bath, not nearly as soluble, gave warty edges, like copper from an aqueous sulphate bath, but otherwise the deposit was good, using the same current density. Anode corrosions were 1 o 1 - 1 o 3 ~ ~cathode , depositions from 87.5-9 j.5Tc. Voltages differed remark-
ELECTRODEPOSITION O F METALS FROM ANHYDROUS AMMONIA
33 I 5
ably, despite the fact that electrode separations were identical-zo mm. ; this was, of course, a reflection of the poorer solubility of the thiocyanate. With the latter, pressures of 3.0-3.5 volts were required-the double cyanide called for only 0.8 or so. Mercury. Our only normally liquid metal, mercury, was readily electrolyzed out of ammonia solution oi 0.8182 grams of HgIz in 30 C.C.of ammonia. h current of 433 milliamperes per square decimeter gave a deposit of beautiful matty metal. The metal alloyed on the gold surface, certain spots being very shiny due to excess of mercury (crystals which had melted completely before alloying). Some tendency to “tree,” as in the case of lead and thallium, was noticed. h platinum, and therefore insoluble, anode being employed, a relatively high voltage requirement of I . j was encountered, with an electrode separation of only 13 mm. Based on Hg++, the ion added, deposition was 89.3%. Thallium. In simple aqueous electrodepositions, both lead and thallium stand out because of their tendency to form spongy, tree-like growths of metal; ammonia solutions proved almost as bad in this respect. The individual crystals of the (‘tree” were somewhat smaller in the case of thallium than held true for lead, both iodide and nitrate giving almost identical results. Deposits were not weighed because the tiny size of the crystals made washing very difficult, even by the manipulation successful with lead. Anode corrosions were again well beyond theoretical requirements-106.67c and I 17.67~ for the nitrate and iodide, respectively. Tin. Since stannic chloride did not dissolve adequately in ammonia at -33OC., although giving a good deposit, with a stannous anode corrosion of 1 1 2 . 8 7 ~and a stannic deposition of 60.1%, it was decided to utilize stannic iodide as an electrolyzing salt. The dried salt weighed 0.6554 grams-30 C.C. of ammonia were used-but even here solution was far from complete; the red salt merely turned white, forming an ammoniate as a milky white liquid, the precipitate settling slowly but completely. Despite this relative insolubility, the solution yielded a fine metallic deposit, back no less than front, with a current of 333 milliamperes per square decimeter-about 1.1 voltsshowing a stannous anode corrosion of 9 5 . ~ 7and ~ cathode deposition of 40.17c (based on Sn++++, which was the ion a t first present, these were 1 9 0 . 9 7 ~and 80.17~ respectively). Lead. ,Just as water solutions of ordinary, simple salts fail to give firm adherent plating of lead, irrespective of current density, so ammonia solutions failed also, though the “treeing” tendency is not as sharp. Xeither the iodide nor the nitrate gave deposits of the desired characteristics, so sodium ammonoplumbite, the ammonia analogue of the ordinary, aqueous sodium plumbite,’ made in situ by the interaction of NaNH2 and PbI2, was tried. It gave deposits apparently a bit more compact, but the fern-like growth of metal was still a characteristic. Anode corrosions for lead varied from 97.j% to 1 1 2 7 and ~ cathodic deposition was 79.3YCfor the ammonoplumbite ~
‘ S e e Franklin: J. Phys. Chem., 15, 5 9 - 2 0 ( 1 9 1 1 ) .
33'6
HAROLD SIMYOSS BOOTH A S D M E N A H E N MERLUB-SOBEL
and 9;.5Ci: for the nitrate. I t may be of some point to mention the fact that the lead so obtained was of extreme purity; crystals were exposed to laboratory conditions for many months without losing their lustre or assuming any of the usual lead dullness. Arsenic. I n sharp distinction to the failure of its congeners, antimony and bismuth, to deposit a t all (as will be discussed later), arsenic gave a small amount of deposition, anhydrous AsBr3 being used as the electrolyte. The greater part of the salt added to 3 0 C . C . of ammonia failed to dissolve, though it was obvious that there had been some preliminary reaction between it and the ammonia, since the precipitate took on the appearance characteristic of aqueous Al(OH),. A density of 833 milliamperes per square decimeter, requiring 2 . 5 - 2 . 8 volts, with a lump of sublimed arsenic as anode, gave considerable gassing a t the cathode, but a thin, quit,e uniform deposit of dark metal did form, both front and back being covered. The deposit was definitely metallic in appearance, and dissolved very rapidly in concentrated HKO3 (aqueous). Because of anode disintegration, no anode loss could be computed; deposition was only j . 1 yGof coulombmeter theory. Chromiztm. Electrolysis of Cr(C98)3, using aqueously deposited chromium as anode, with a cathode current density of 16,700 milliamperes per square decimeter, required between 5 and 6 volts for operation, and gave a complete thick plate on the front of the cat'hode, with a somewhat scanty deposit on the back. Electrode separation was only 1 2 mm., yet the throwing power proved far better than for most aqueous chromium baths, though still far from good, particularly when compared with other metals. Hydrogen and metal were co-deposited continually during the electrolysis and the gas had to be blown off to prevent excess pressure in the cell (lower current densities also gave vigorous gassing at both electrodes). Based on Cr+-+ considerations, anode corrosion was 1 4 . 4 7 ~cathode ~ deposition only I .2yC.Electrode attack and deposition were of the same general order of magnitude with lower current densities, too, which did not give adequate plating. Chromium, therefore, acts much the same in ammonia as in water, with a sharp tendency toward anode passivity. ilfanganese. In direct contraposition to the difficulties encountered with aqueous deposition of manganese, electrolysis in ammonia gave excellent results, except for the grave tendency of the plate t o peel and flake off. Manganese metal of 9 7 . ~ purity 7~ was used as an anode; iron present being 1 . 4 8 7 ~and silicon 0 . 2 3 7 ~ . A current of 167 milliamperes per square decimeter was used, the low voltage of about 0.9 being explained to an extent by the small electrode separation-8 mm. As electrolyte, Mn(CKS)2 was used. The salt-o.z5;4 grams in weight-with 30 c.c of ammonia, did not dissolve completely, but formed n cloudy suspension. The deposit was not only complete on the front, but also covered most of the back of the cathode, indicating fairly good throwing power. Some of the cathode metal peeled off on the least jarring; it was washed and weighed separately. Anode corrosion 7 ~ cathode deposition 97.8yc. The metal obtained was of high was 9 ~ . 4 and purity-crystals exposed in the laboratory for some time retained their
ELECTRODEPOSITION O F METALS FROM ANHYDROUS AMMOXIA
33 I;
lustre, without the usual rusting, this freedom from corrosion being a characteristic of the pure metal.' Iron, if a t all present in the deposit, was so extremely low as to give no color with the delicate thiocyanate test. Iron. By the use of ferrous iodide a t a current rate of 116-133 milliamperes per square decimeter, perfect deposition of iron was obtained ; the voltage being approximately z .; 5 . Throwing power was excellent. Anode corrosion, rather surprisingly, ranged only from 5.4 to 6.0yG,while deposition dropped to 8.;Yc on the optimum current density, rising again sharply to 29.2% with higher values of current. Gassing a t both electrodes serves to explain the low current efficiencies. Nickel. As with aqueous solutions, ammonia electrolysis gives fine adherent deposits of nickel, platinum-like in brightness, with fine throwing power. Ni(CNS)2.4NH3was used, being synthesized by the method suggested by Bohart.* Ext,remely low current densities were required to prevent the deposit from peeling; only by dropping to j o milliamperes per square decimeter was perfect adhesion obtained. No gas was evolved in any electrolysis. In the best run, 1.1579 grams of the salt were dissolved in 40 C.C.of ammonia; electrode separation was 2 1 mm. and a voltage of approximately 1 . 2 5 was required. Anode corrosion was 104.57~;cathode deposition 92.3%. Cobalt. Cobalt was deposited in an analogous manner, 0.6534 grams of Co(CSS)2, not entirely nickel-free, dissolving completely in 35 C.C. of ammonia. As anode, aqueously electrodeposited cobalt was used. A square decimeter density of 100-1I 7 milliamperes required 0.7-0.8 amperes for 14 mm. separation. The deposit was considerably darker than in the case of nickel, with the slightest tendency toward blackness a t certain points. Adherence was perfect, as was also the throwing power, the back being plated equally with the front. Anode loss and cathode gain were identical and clove to theoretical, 98.07~. Pdadium. By the solution of 0.2578 grams of PdItin 30 C.C.of ammonia (solution was readily accomplished) and electrolysis with an insoluble platinum anode, a bright shiny deposit of palladium metal was obtained, completely covering both front and back of the cathode. The current density was about 116 milliamperes per square decimeter a t I.; volts. Electrode separation was 13 mm. Cathode deposition proved to be only 68.7y0; perhaps depletion of Pd++ will explain the low results, though no gassing as such was observed. Platinum. Rather surprisingly, platinum gave considerable difficulty, in this respect ammonia being radically different from water as a solvent. With the tetraiodide as the electrolyte I 6;-233 milliamperes per square decimeter failed to give any deposition a t all, and higher current densitiesI 700 milliamperes or so to the square decimeter-caused considerable burning, though even under these conditions deposition efficiencies were low. A Royce and Kahlenberg: Trans. Am. Electrochem. SOC.,50, 281-300(1926). S. Bohart: J. Phys. Chem., 19, 537-63 (1915).
* G.
3318
HAROLD S I M M O S S BOOTH AND YENAHEM MERLUB-SOBEL
thin greyish deposit was the best that could be obtained, with some burnt spots. In all electrolyses, hydrogen was evolved profusely. I t is interesting to note that in aqueous deposition of bright platinum, it is necessary to cut down the Pt++++concentration sharply by means of phosphates or similar ions, otherwise “platinized” types of plates will result. Here, hydrogen was inevitable with any deposition at all, yet the deposit was hardly a really good one and vas not very adherent. The best run indicated a 1 o . 0 7 ~cathode efficiency, 0.2339 grams of PtL being dissolved in z j C . C . of ammonia with anode and cathode 16 mm. apart. Current density was about 1,670 milliaamperes per square decimeter, and 8-12 rolts were required, rising toward the end, perhaps because of a tendency for depletion near the cathode. This brings to an end the metals where success was attained. Deposition failure occurred, among those tried, with two metals readily deposited in water, and mit’h four which are akin in their aqueous action to that of ammonia. The former included antimony and bismuth, while the latter were represented by aluminum, thorium, tungsten and molybdenum. Aluminum, Electrodeposition of beryllium raised great hopes for similar success wit,h aluminum, particularly in view of the generally accepted lower electromotive position of aluminum. Xothing of the sort materialized, however; proving, if nothing else, the danger of thinking strictly by analogy. Despite the high solubility of -411, a t room temperatures, such solubility proved very sharply lower at the normal boiling point of ammonia and only a small portion of the salt dissolved. K i t h the various current densities tried, there was vigorous gassing and a blue-black film seemed to form on the cathode, but this deposit, if such it was, disappeared whenever the current was turned off. Perhaps some subiodidc, subnitride, or similar compound was forming by reduction. Anode corrosions (the anodes were of Hoopes Process aluminum, 9 9 . 9 j i 5 pure) varied considerably, from 86.3
