T H E ELECTROLYSIS OF ACID SOLUTIONS OF COPPER SULPHATE B Y L. .V. REDMAN
The electrolysis of copper sulphat’esolutions has been the subject of much research; in 1800 Cruikshank’ shewed that copper could be separated from its solutions by means of the electric current, in 1834 Faraday2 determined its electrochemical equivalent to be 3 I . 6, in 1867 Regnault3 shewed t’hat in solut’ions of cuprous chloride the equivalent was double that value, and in 1897 Forster and Seide14established the existence of cuprous salt in solutions of cupric sulphate. As a result of these invest’igations, t’he quant’it,ativerelat’ionsbetween .the amount of copper deposited and tmhecurrent seemed to be fairly well established, when, in the Spring of 1901,Jos. Siegrists published the results of his researches “on the velocity of the electrolytic deposition of copper in the presence of sulphuric acid”. This author’s very numerous and careful experiments shew that although the amount’ of copper deposited per coulomb from concentrated solutions is about that calculated from Faraday’s law, yet, when the ratio of concentration t’o current-density falls below a certain critical value,6 the rate of deposition is almost independent of the current and is roughly proportional to the concentration of the copper sulphate in the solution. These results he explained by assuming that the reduction of copper sulphate by the cathode, like that by chemical reducing agents, is a process requiring time; and that the rate of t’hisreaction, like that of many others, is dependent on the concentration of t’hesolution; in support of this explanation he cites the work of Haber’ on the electrolytic reduction of nitrobenzene. Under given conditions of concentration, t’emperature, etc., then, the amount of copper that could be reduced in a given time would be limited; and if the current called for more, the deficit would have to be made up in some other way-by liberation of hydrogen, for instance, I n experiments carried out well below the critical concentration-density ratio, Siegrist found that the rate of deposition of copper is proportional to t,he concentration of copper in the solution; in the language of chemical kinetiw, the electrolytic reduct,ion of copper is a reaction “of the first order”. The researches of Noyes and Whitney8 and of Brunnez on the rate of solution of solids, and those of Salomon‘o,Vernst and Merriamll, and ot’hers Xicholson’s Journal, 4 , 187 (1800). “Experimental Researches”, VI1 Series, Sec. 846 (1834). 3Ann. Chim. Phys. (4) 11 137 (1867). Z. Elektrochem. 3 479 (1897). E Z. anorg. Chem. 26 273 (1901). 6 I use the term “limiting current” to indicate the greatest current that, under the oonditions of the experiment liberates copper only. Z. physik. Chem. 32 193 (1900). 8 Z.physik. Chem. 23 689 (1897). 9 Z. physik. Chem. 47 56 (1904). l o Z. physik. Chem. 24 54 (1897). Z. physik. Chem. 53 235 (190s). 2
P:LECTROLYSIS O F COPPER SULPHBTE SOI,U?‘IOiVS
I549
on the “residual current” suggest a wholly different explanation. These 811thors worked with cylinders of various solids rotating in a solution, or with flat plates and a mechanical stirrer; and their experiments justify the assumption that a thin film of liquid adheres to the surface of the solid, and that the constituents of the solution can reach the surface of the solid only by diffusion through the film. From this it follows that the concentration of copper sulphate at the surface of the cathode in Siegrist’s experiments must have been lower than it was in the rest of the solution, and that the amount of copper that could be deposited electrolytically in a given time would be limited by the amount that could reach the cathode by diffusion. This view explains a t once the great influence of stirring on the quantity of copper deposited; it also accounts for Siegrist’s observation that his “rate” is a linear function of the temperature and not an exponential function like the rates of most chemical reactions. It is, moreover, in zgreement with the generally accepted theory of decomposition-voltages, according t o which the concentration of the copper sulphate must fall almost to zero before the liberation of hydrogen is possible; and with the outcome of experiments by Sand1 undertaken to test this very point. I n the autumn of 1907, while Professors T . R. Rosebrugh and IT. Lash Miller were developing the mathematical theory of the changes of concentration at the electrode brought about by diffusion and by chemical reaction2, I began the work described below (as a student in the electrochemical laboratory of the University of Toronto) with solutions of copper sulphate in maximum-conducting sulphuric acid, to see in how far the conclusions of the theory were borne out by the experiment. After preliminary verification of Siegrist’s results-which shewed, incidentally, that the “limiting-current” could be determined much more easily by the use of a voltmeter than by Siegrist’s method of weighing-my attention was directed to measurements with interrupted current, because for this case the kinetic theory and the diffusion theory make contradictory predictions. The balance of probabilities was, no doubt, strongly in favour of the latter theory; but an esperimentum crucas seemed necessary, as the linear temperature coefficient, and the influence of stirring on the kinetic “constant!’, both of which are hard to explain on the kine tic view, were well known to Siegrist ; and that he did not regard the results of Sand to be fatal to his theory is proved by the following quotation from his paper: .‘TVie die T’erhaltnisse im Inneren des elektrolytischen Troges gestalten interessiert uns hier nicht ; besonders da im vorliegenden Falle das Kupfer nicht durch Ionenwanderung, sondern durch heftiges mechanisches Ruhren der Elektrode zugefuhrt wurde. (Wahrend der Korrektur dieser Arbeit erschient die Arbeit von H. J. S. Sand, melche sich speziell mit dieser Frage beschaftigt .)” If the liberation of hydrogen first occurs when the rate of deposition of copper called for by Faraday’s law exceeds the maximum rate at which copper
*
Phil. Mag. (6) 1 4 j (1901); Z. physik. Chem. 35 641 (1900). J. Phys. Chem 14 816-884 (1910); referred to as “R & M”.
can be chemically reduced from a given solution, then periodic interruptions of the current will be without effect and a current which when uninterrupted will liberate hydrogen, will also liberate it if periodically interrupted. If on the other hand the liberation of hydrogen is due to the exhaustion of copper salt in the solution at the cathode, then during the interruption the supply of salt will be replenished, and if the interruptions are frequent enough and of sufficient duration, nothing but copper may be deposited by a current which if uninterrupted would bring about liberation of hydrogen. As cathode I used a platinum cylinder 2 . 3 cm. in diameter and o 97 cm. high, rotating 1630 times per minute on a vertical axis between horizontal plates of ebonite in I z 5 litres of a solution containing I . o or 2 o g. of crystallized copper sulphate CuS01.;H20 per litre of maximum- conducting sulphuric acid (7.6 normal acid). The anode was a concentric copper cylinder, and the temperature 19°C. With this arrangement the minimum current that would liberate hydrogen was determined for different numbers of interruptions of the current per second, a sudden rise in the voltage over the cell being taken as evidence of liberation of hydrogen. The interruptor was constructed of one hundred plates of copper insulated by mica, like the commutator of a dynamo, and could interrupt the current from 0 . 5 to 3000 times per second; owing to the thickness of the insulation, the duration of each interruption was trifle longer than that of each beat of current. Current and voltage were determined by Weston instruments; the inertia of their moving parts was so great that, unless the number of jnterruptions fell below twelve per second, the needles were stationary enough to permit of accurate readings; calibration shewed that these readings represented 47’70 of the current and voltage respectively during the intervals when the currenh was flowing. When the interruptions were fewer than twelve per second, however, the oscillations of the voltmeter needle were too great and the voltmeter was replaced by two copper points dipping in maximum-conducting sulphuric acid in a small Erlenmeyer flask, the idea being that when the voltage rose high enough to liberate hydrogen in the cell, hydrogen would be liberated also at the point] connected to the cathode. The wall of the flask acted as a lens and enabled the gas to be easily detected, and the device worked well in practice; but on trying it out with known constant currents it was found that a considerable empirical correction had to be applied, thus the results with fewer than twelve interruptions per second are less reliabje than the others.
TABLE I 1.0 g. per 1. Int. per see. o 0.6 I.I 2 . 0 Amp.percm2.0.0044 j 2 56 70 2 . 0 g. per 1. Int. per sec. o Amp.percm2. 0 . 0 0 8 8
1.5 132
4.5 5.7
3.2
84 86
82
2.8
4.2
164 176
7.0 178
6.; 9.0 86 92 12.0
180
3000
126 212 2500
98
98
98 0.0098
1 2 6 2 1 2 2 5 0 0 3000
r78
178
178
0.0178
ELECTROLYSIS O F COPPER SULPHATE SOLUTIONS
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The results, given in Table I, shew that the limiting current rises continuously with the number of interruptions per second, and that when the interruptions are frequent enough the limiting current is about double that with uninterrupted current. This result not only directly contradicts the predictions of Siegrist’s theory, but it is in quantitative accord with the predictions of the diffusion theory; for equations 33 and 24 of the paper by Rosebrugh and Lash Miller lea,d to the conclusion that “the stationary state reached by electrolysis with a rapidly interrupted, alternated, or varied two-beat current is practically the same as would be reached by electrolysis with a constant current of the same number of coulombs, reckoned algebraically”.’ Siegrist’s theory being thus eliminated, I endeavoured to determine point by point the voltage-time curve during the first second or so of an electrolysis, with the help of a mechanical device which connected anode and cathode for exactly one-hundredth of a seccnd to the terminals of a previously calibrated ballistic galvanometer, and which could be arranged so that this connection was made at any desired interval after closing the current through the cell. Some of the curves constructed in this way agree fairly well with those obtained with the apparatus described further on; but the work was extremely tedious, for it was found necessary to clean and polish the cathode before each measurement, as the rough deposit formed a t the end of each electrolysis affected the voltage of the succeeding experiment. Recourse was therefore had to the oscillograph, and after some preliminary work with an instrument kindly loaned by the Department of Electrical Engineering, a two-element Siemens-Halske instrument was purchased for this work. Each element had an undamped frequency of 2900 per second, its resistance was 2 . 5 ohms, and a current of 4 milamperes through the element gave a deflection of 45 millimetres on the photographic paper. I n order that the 4 or 5 milamperes taken by the oscillograph might be negligible in comparison with the current passing through the cell, the (rotating) cathode was constructed of a solid copper cylinder 4.6 cm. in diameter and 6.4 cm. high, the ends covered with an insulating gum; contact was made through a mercury cup on the top of the shaft; the anode was a concentrir copper cylinder of 9.; em. diameter and the same height as the cathode, while two horizontal ebonite plates, the upper one perforated to admit the cathode shaft, ensured a uniform electrolytic field. One of the oscillograph elements, with a suitable resistance (about 280 ohms) in series, Ferved to measure the potential difference across the cell, the second element was connected in series with a resistance of 90 ohms and the secondary of a small induction coil (7.7 ohms) through the primary of which a constant current of 1.2 amperes was maintained by a storage battery. A tuning fork (middle C, 128 cycles per second) vibrating just above the core of’ the coil, induced 2 sinusoidal current in the secondary, which was recorded on the photographic paper of the oscillograph. This sinusoidal current loc. rit. page 840.
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L. TT. R E D M A S
was short-circuited once per revolution of the cathode shaft, so that the speed of revolution of the shaft was also recorded on the photographic paper. The electrolysis circuit is shewn in Fig. I . X is a heavy copper bar which in its normal position short-circuited the cell and the oscillograph circuits and thus enabled the current to be adjusted by means of the lamp rheostat R and the ammeter A ; when the bar was pulled back by an electromagnet the current flowed through the cell. The potential difference set up over the 0.01 ohm sent a trifling current through the second element (11)of the oscillograph, in parallel with the sinusoidal current from the coil; this displaced the time-wave on the photographic paper, and identified the moment when the electrolysis current was thrown on,
To get a good record it was necessary that the electrolysis should begin immediately after the photographic shutter of the oscillograph had opened; this result was secured by tapping the circuit that leads to the “openingmagnet’’ of the oscillograph, and sending part of the current through a relay which operated the magnet that drew back the short-circuiting bar. To secure a time-voltage curve with this apparatus the following operations were necessary :-The cathode was cleaned and polished,l the electrolyte analyzed and poured into the cell, the cathode set revolving a t the desired speed (estimated by means of an electrical speed-counter attached to the cathode shaft), the current adjusted, the tuning-fork set in vibration by a blow from a rubber hammer, the key of the oscillograph pressed and the photograph taken, the current read on ammeter B, and the short-circuiting bar restored to its normal position. All the data needed, except the composition of the electrolyte, its temperature, and the current used, were recorded on the photographic paper. Fig. 2 shews one of the records so obtained; the electrolyte contained 1.0g. of copper per litre, its temperature was r7.y°C, the current was 1.09 amperes, the cathode revolved once in 0.133 seconds ( 1 7 . 0 waves). The horizontal line a t A is the zero-voltage line, the line BCD gives the potential difference over the cell, while the horizontal at E (the “voltage calibration line”) corresponds to a voltage of 0.528 volts over the cell. If it be assumed that hyd1 This is very necessary; if three or four oscillograms be taken in close succession on the same paper, without changing the electrolyte or varying the current, the time required to liberate hydrogen in the last may be noticeably greater than in the first. This is due to the dark powdery deposit formed during the final moments of each electrolysis (while hydrogen is being evolved) which in erfect derreases the thickness of the diffusion film. If a proloi :ed electrolysis be carried out with a constant current somewhat greater than the limiting current, hydrogen is a t first evolved and a dark deposit formed, but later on the evolution of hydrogen ceases and the deposit turns red.
ELECTROLYSIS O F COPPEH. SULPHATE SOLUTIOSS
I553
rogen was first liberated when BCD ceased to become steeper (at the point of inflexion, marked with an arrow), the concentration of copper fell from I .o g. per litre to zero in 0 . 2 I I seconds. The limiting current -I‘ under the conditions of this experiment was 0.403 amperes, thus the ratio I’ : I = 0.37, which is less than 0.5, and the “parabolic approximation” can be used to determine the diffusion constant k , without knowledge of the thickness of the diffusion film. From the equation (R. & M. Eq. zzb) 96500 Ak(x-2,) = -1.129
.\/kt
setting A (area of cathode) = 92.5, x = 0 , x, = 1.0/31800,1 = 0 . 2 1 1 , = - 1.09 (the negative sign because the electrode is cathode) there follows
1 k
= 4.1x10-6.
A FIG. 2
I made some fifty or sixty determinations of this type, with solutions containing from one to six grams of copper per litre, currents from 0.9 to 6.2 amperes, and cathode speed from 0.03 to 0.15 sec. per revolution. The values of k varied from 3.5X IO-^ to 5.5X IO-^ a t 18OC,thehighervaluesbeingobtained with the most dilute solutions. While at first inclined to ascribe this regularity to change of k with the concentration of copper in the electrolyte, ultimately I found a source of error in my apparatus which made this conclusion unreliable, via: the solution of the copper electrodes in the strongly acid electrolyte which was unprotected from the air. The values of xo used in the calculations, being based on analyses made before the electrolyses, were, therefore, too low by a variable amount which depended on the interval between filling the cell and taking the photograph, and the error so introduced (which results in too high values of k ) must obviously be greater the smaller the origin31 concentration of the solution. By the time that blank experiments had established the magnitude of this error, it was too late to rebuild the apparatus; critical examination of the note book gave ro6k = 4.0 i 0.5 as an approximate value for the diffusion constant a t 18’C. A sharper check on the predictions of the diffusion theory was gained from time-voltage curves obtained with periodically interrupted currents. The interrupter used in my previous work could not handle the comparatively heavy currents needed by the large cathode without sparking, and was replaced by a sliding contact made and broken by a reciprocating piston. Fig. 3 shews one of the records; the current was 2.86 amperes, the electrolyPs contained 2.0 g. copper per litre, and its temperature was 18OC. Hydrogen was first liberated 0.359 seconds after the current first flowed through the cell (the point is marked with an arrow in the figure); assuming the value
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L . V. REDMAN
= 4.0x IO-^, equation (R & M. 29 d ) predicts 0.377 seconds. I made about a dozen such records, under different conditions of current and concentration, towards the end of my work, and in every case the hydrogen was first liberated in the “beat” of current predicted by the mathematical theory; owing to the uncertainty as to the diffusion constant closer agreement could not be expected.
IC
Overvoltage While the lapse of time before the first liberation of hydrogen could be predicted fairly well, the potential difference over the electrodes was very different from that calculated from the concentrations given by the diffusion theory if it be assumed that the cell is reversible, and that the “concentration E. M. F.” can be calculated from the logarithm of the concentration ratio by the ordinary formula. While I was making my first measurements with the oscillograph, LeBlanc’fi paper “On the E. M. F. of Polarization’’ appeared1, in
which he pointed out that the potential difference a t the moment of throwing on the current in acid solutions of copper sulphate is much greater than corresponds to the product of current into the resistance of the cell. LeBlanc calls the difference “overvoltage”; I prefer to call it “initial overvoltage”, because not only is the height A B (Fig. 2 ) too great, but the line BC rises much more than can be accounted for by the fall in copper concentration as calculated from the mathematical theory. I n addition, then, to the “initial overvoltage” noticed by LeBlanc, Reichinstein, and others, there is a further overvoltage which builds up during electrolysis before the liberation of hydrogen; this second sort of overvoltage was not noticed by LeBlanc, for it could not be recognized without knowledge of the concentrations a t the electrode furnished by the “mathematical theory”, and although the equations for calculating these concentrations were published fifteen years ago, I have seen no reference to it until Professor Lash Miller summarized the results of Mr. A. R. Gordon’s work at the Centenary of the Franklin Institute2. Summary
I think it may fairly be said that the experiments recorded in this paper dispose of Siegrist’s “kinetic theory”, and shew that the assumptions made deutsch. Bunsen-Ges., Kr 3 (1910). Franklin Inst. Centenary Publication.
1 Abh. 2
ELECTROLYSIS O F COPPER SULPHATE SOLUTIOKS
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by the “mathematical theory” are fairly applicable to the case of acid solutions of copper sulphate. At the time the work was done more could not be claimed; for it was uncertain whether the lack of sharper agreement between equations and experiments was due t o experimental errors, or to an error in the assumption made by the theory, that the diffusion constant is strictly independent of the concentration. This matter has since been thoroughly investigated by Mr. J. T. Burt-Gerrans and Mr. A. R. Gordon in the Toronto laboratory; before publishing their own results they wish to see in print an account of the preliminary studies on which their work was based, and it is at their request and that of Professor Lash Miller that this abridgment, of an old thesis and laboratory report is brought to light. University of Toronto July,1926
