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Volume 49, 1916
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Art XLVI.—A Study of the Electrical Deposition of Nickel in the Presence of Nitrate.

[Read before the Philosophical Institute of Canterbury, 6th December, 1916; received by Editors, 30th December, 1916; issued separately, 10th December, 1917.]

The modern method usually employed in the electrical separation of nickel originated in the use of the double sulphate of nickel and ammonium first suggested by Professor Boettger in 1843. The deposition, however, was very slow, and the conductivity of the electrolyte was extremely small. To remedy this latter defect a number of substances were suggested, ammonium sulphate being most frequently used for the purpose. Under suitable conditions a good coherent deposit can be obtained, and the process, though lengthy, admits of a quantitative estimation of the amount of nickel present in a solution.

Numerous experimenters have observed that nitrates exert a most disturbing effect on such estimations. In recent years Thiel* found that the presence of nitrites yields high results owing to deposition of nickel oxide on the cathode. The present investigation was undertaken to study the effect of the addition of known amounts of nitrate on the deposition of the metal, and to eliminate the disturbing products of the electrolysis.

Experimental Details.

The apparatus used consisted of a glass beaker containing the solution to be electrolyzed, in which were suspended the electrodes. These were of iridio-platinum, and were of the form devised by Perkin. The cathode was made of gauze, and possessed a total surface of 50 sq. cm.; the anode is opposed to both sides of the cathode during electrolysis, and consequently an even current-density is obtained on all parts of the cathode.

The conditions under which the investigations were conducted were those recommended by Marshall, viz., —

Amount of nickel to be deposited (approx.) 0.3 grm.
Volume of solution 150 c.c.
Ammonium sulphate 5 grm.
Ammonia 5 grm.

The experiments were made at the ordinary temperature, and about four hours was found necessary to complete the electrolysis.

The solution of nickel sulphate to be used in the subsequent course of the work was carefully standardized; a large number of determinations were made under precisely similar conditions, and the results showed close

[Footnote] * A. Thiel, Zeitsch. Elektrochem, vol. 14, 1908, pp. 201–8

[Footnote] † W. Perkin, A New Form of Electrode, Journ. Faraday Soc., vol. 1, 1903.

[Footnote] ‡ A. Marshall, Some Polarisation Phenomena, Proc. Roy. Soc. Edin., 1899.

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agreement, as the following figures indicate: Amount deposited after four hours, 0.3088 grm., 0.3087 grm., 0.3090 grm., 0.3084 grm., 0.3069 grm., 0.3085 grm.; yielding a mean value of 0.3084 grm. when 5 c c. of the standard solution with the ammonia and ammonium sulphate (to increase the conductivity) added was electrolyzed.

Series 1.

(Solution containing no nitrate.)

The first series of experiments was conducted with the solution as indicated above, in the absence of nitrate, with the object of obtaining a curve showing the rate of deposition of the metal. This was effected by interrupting the electrolysis after it had proceeded a definite time, and then washing, drying, and weighing the cathode in the usual manner. At the same time, during the course of the electrolysis, readings were taken continuously of (i) ammeter, (ii) P.D. across the terminals of the electrolytic cell, (iii) P.D. at the cathode.

Ammeter Readings. — These were noticed to exhibit slight variations, though they remained practically constant throughout each experiment. A small initial rise was invariably noticeable, due probably to increase in temperature as a result of the ohmic resistance encountered, and towards the conclusion of an experiment a diminution was always observable. The extent of these fluctuations never exceeded 0.1–0.15 amp., and to facilitate subsequent investigations it was decided to maintain the current at a constant value by inserting an adjustable resistance in the circuit. It did not seem possible to deduce anything of importance from a study of these small variations.

Voltmeter Readings.—The readings of the voltmeter were found to exhibit a much more marked fluctuation. A considerable number of experiments were conducted so as to confirm the unusual nature of the variations, and it was found that they could be divided into three stages: (a) Very slight rise during the first hour, followed by a return to the initial value; (b) period during which the voltmeter remained approximately constant; (c) towards the end of an experiment a distinct rise in P.D. could be detected, amounting to 0.2–0.4 volt. These stages were not always clearly defined, though the final rise was almost invariably observed. This would correspond to increased resistance between the electrodes, and, on account of the presence of the ammonium sulphate, it could not be explained by a diminution in the conductivity of the solution. Further, it is noticed after almost all the nickel has been deposited and only traces remain in solution. The removal, then, of these last traces of the metal must involve the expenditure of more electrical energy and indicate the presence of polarization phenomena. Marshall (loc. cit.) suggests that “a film of some other substance is deposited giving a higher polarization effect. At first this is replaced more or less by nickel, and occasionally breaks down, as shown by the lapses to the original potential, possibly on account of the richer solution being brought against it by convection currents, but it ultimately becomes permanent when the nickel deposition is completed.” In the case of this investigation, as in that of Marshall, the final rise appears to be an indication that the deposition is complete.

Potential Difference at the Cathode.—To study the changes in P.D at the cathode the arrangement depicted in fig. 1 was employed.

The composite cell consisted of calomel electrode, Ni-NiSO4 (ammon)

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From the known value of the E.M.F. of the Weston standard cell that of the composite cell could be calculated; and, taking 0.613 as the P.D. at the calomel electrode, it was possible to deduce the variation in the single P.D. Ni-NiSO4 at the cathode as deposition continued (figs. 2 and 3) in the absence and in the presence of ammonium nitrate.

Rate of Deposition.—It was noticed that when using the sulphate in the ordinary manner 80 per cent. of the nickel was deposited during the first hour of the electrolysis. The quantity of electricity and the time

Picture icon

Fig. 1.
B′—storage cells, giving a uniform fall of potential along the bridge wire XY; R′—adjustable resistance; K′—key.
B—battery connected to terminals of electrolytic cell M. By means of R, an adjustable rheostat, the current as indicated by the ammeter A is maintained at a constant value. K—key
C—calomel electrode (N/10 KC1) dipping into connecting vessel N, containing ammonium sulphate. A siphon, also containing ammonium sulphate, serves to connect N and the electrolytic cell M, one arm being placed in the immediate vicinity of the cathode.
V—voltmeter placed across the terminals of the electrolytic cell M
E—Lippmann capillary electrometer; T—short-circuiting key; S—switch.
W—Weston standard cadmium cell

necessary to remove the final traces of nickel from the solution are out of all proportion to the amount of metal actually present.

On the basis of Faraday's quantitative law, W = γηt (where W = weight in grammes of metal deposited, γ = current in amperes, η = electrochemical equivalent of metal, t = time in seconds), the theoretical amount of nickel was computed on the assumption of perfect efficiency, and on comparison with the experimental values the working efficiency was found to be 25–30 per cent.

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Results.

Each experiment of the series was repeated a number of times, and the agreement between the sets of values was found very satisfactory, and sufficient to confirm the general nature of the results, Typical series of values and curves are given below:—

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Time.
Hours
P D
Volts.
0 3.6
½ 3.65
¾ 3.6
1. Variation of P.D. between terminals 3.7
2 3.8
3.87
0 0.878
½ 0.874
2. Variation of P.D., cathode solution (fig. 2, a) 1 0.870
0.865
2 0.860
Time.
Hours
Amount in
Grammes.
1 0.2680
2 0.3050
3. Rate of deposition of nickel (fig. 4, b) 3 0.3077
4 0.3087

Series 2.

In this series of experiments varying amounts of ammonium nitrate were added to the solution, and the observations repeated as in the preceding series. The conditions were precisely similar, but in this case the current was kept at a constant value.

Results.

1. P.D. across terminals: The variations were not so general and definite as in the former series. During the larger part of an experiment the P.D. remained practically constant, and towards the end a fall, and not a rise, was observed. Whatever the explanation may be, the presence of the nitrate appears to have the effect of removing the polarization by destroying the film deposited on the cathode, which, as Marshall suggests, is the cause of the increased resistance observed in the previous case.

2. P.D. cathode solution: Here an interesting periodicity was found to be quite general. The curves indicating the variation of the P.D. with time are appended, and it will be seen that they consist of three distinct portions: (a.) At first a gradual diminution is observed, the extent of the diminution and the period during which it persists depending upon the actual amount of nitrate present in the solution. (b.) This is in every case followed by a sudden rise—very pronounced. (c.) After the maximum value of (b) is attained, the P.D. gradually diminishes as time proceeds.

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Time
Hours.
PD
Volts
0 0.660
1 0.635
0.646
2 0.820
(a.) One equivalent of ammonium nitrate (fig. 2, b) 0.851
3 0.837
0.815
4 0.793
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0 0.726
2 0.655
Values— 0.615
(b.) Two equivalents of ammonium nitrate (fig. 3) 3 0.585
4 0.579
0.616
5 0.859
6 0.825
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Fig. 2.
Fig. 3.

3. Rate of deposition: The rate of deposition of the nickel was found to be enormously affected by the addition of ammonium nitrate. The following typical values may be given from a number of closely agreeing experiments where one, two, and three equivalents of nitrate were added to the solution:—

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Fig. 4.
Fig. 5.

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Time
Hours.
Amount
deposited.
Grammes.
1 0.0065
(a) One equivalent of ammonium nitrate (fig. 4, c) 2 0.0382
3 0.2234
4 0.2861
5 0.3050
(b.) Two equivalents of ammonium nitrate (fig. 5, a) 2 0.0061
4 0.0153
5 0.0813
6 0.2748
(c.) Three equivalents of ammonium nitrate (fig. 5, b) 3 0.0055
4 0.0063
11 0.1469
12 0.2738
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Discussion.

From the results obtained it was evident that the electrical energy was utilized in effecting some reaction other than the separation of the metal from the solution. Accordingly a series of experiments was undertaken in order to ascertain the way in which the energy was used up.

Now, on comparing the sets of values showing the variation in the P.D. at the cathode and the rate of deposition of the metal, it will be noticed that a sudden rise in P.D. appears to occur when the metal is beginning to be separated in quantity. This was investigated, and from a number of experiments this rise in P.D. appeared to be an infallible criterion that deposition had commenced. Furthermore, during the course of the deposition subsequent to this sudden rise in the P.D. at the cathode the values for the latter exhibited a gradual diminution, as was previously observed in the absence of nitrate.

Series 3.

The preceding series of experiments was repeated, using a greater current-density, but similar results were obtained.

Series 4.

The most probable manner in which the energy would be utilized would be in effecting the reduction of the ammonium nitrate. This reduction may lead to a variety of products, according to the extent to which the reduction has taken place—e.g., nitrous acid, hydroxylamine, or ammonia may result.

In the case where one equivalent amount of ammonium nitrate (0.85 grm) was added, the amounts of hydrogen necessary to effect the reduction to the three stages were calculated from Faraday's law to be 0.0212 grm., 0.0637 grm, and 0.0849 grm. respectively.

The amounts of hydrogen liberated during the electrolysis, deducting the equivalent of the amount of nickel deposited, were calculated to be—At end of one hour, 0.037 grm.; at end of two hours, 0.075 grm.; at end of three hours, 0.1045 grm.; at end of four hours, 0.1398 grm

This series of experiments, then, was performed in the hope of isolating hydroxylamine, which one might reasonably expect to be one of the products of the reduction.

Estimation of Hydroxylamine.

At the end of a definite interval of time the electrolysis was interrupted and the solution evaporated to dryness. From the mixture of salts so obtained the hydroxylamine was extracted by trituration with absolute alcohol; on evaporation of the latter the salt crystallized out. The hydroxylamine so obtained was dissolved in carefully distilled water, and estimated in the usual way by the reduction of boiling Fehling's solution,* which proceeds according to the following equation: 2NH2OH + 4CuO = 2Cu2O + N2O + 3H2O. The cuprous oxide was filtered through a Gooch crucible, dried carefully, and weighed, and hence the corresponding amount of hydroxylamine could be readily computed.

[Footnote] * H. O. Jones and F. W. Carpenter, The Estimation of Hydroxlyamine, Journ. Chem. Soc., vol. 83, 1903, pp. 1394–1400.

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Results.

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Time
Hours.
Amount
Hydroxylamine.
Grammes.
1 0.059
(a.) One equivalent of ammonium nitrate 2 0.033
3 0.008
1 0.050
(b.) Two equivalents of ammonium nitrate 2 0.023
3 0.007
1 0.093
(c.) Experiment with ammonium nitrate alone 2 0.035

Series 5.

The preceding series of observations indicates that though hydroxylamine is undoubtedly a product of the reduction, yet it undergoes further change subsequent to its formation—probably being reduced to ammonia. The formation of hydroxylamine, however, can only be ascribed to reduction taking place at the cathode; and in order to study more closely this action it was decided to conduct a series of experiments in a divided cell where the anode and cathode compartments were kept separate, so that the solutions could be analysed separately.

Apparatus.

The apparatus consisted of a Grove porous pot and outer vessel; the inner compartment contained the same solution as in the previous cases—one-third quantities being used—and in it was placed the cathode. The outer vessel contained ammonia and ammonium sulphate solution.

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Results.
Time.
Hours.
Amount
deposited
Grammes.
(a.) No nitrate present ½ 0.0540
1 0.0686
1 0.056
(b.) One equivalent ammonium nitrate 2 0.068
1 0.054
(c.) Two equivalents ammonium nitrate 2 0.066
1 0.053
(d.) Three equivalents ammonium nitrate 2 0.063
Time
Hours.
Amount
formed.
Grammes
1 0.037
Amount of hydroxylamine 2 0.008

Discussion.

The results obtained with the divided cell were certainly more satisfactory in the presence of the nitrate, and seemed to indicate that increasing amounts of the nitrate did not produce such a decided effect on the rate of deposition as when the undivided cell was used. This would seem to suggest that perhaps some product formed at the anode had the effect of retarding the progress of the electrolysis.

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There appears to be no doubt that the hydroxylamine undergoes a further change. An experiment was performed as follows to test this:—

Experiment.

To the solution in the cathode compartment, containing no nitrate, were added successive quantities of hydroxylamine so as to produce conditions as nearly comparable as possible to those that would be expected to obtain from the basis of Faraday's law. Suppose, for instance, one equivalent of ammonium nitrate were present; then the amounts of hydroxylamine present in the solution at various times—e.g., ¼ hour, ½ hour, &c.—can be approximately computed. At intervals, then, such amounts of hydroxylamine were slowly added. It is at once apparent that the conditions in the two cases are by no means identical, for in one a certain amount of energy is expended in producing the hydroxylamine prior to its transformation, while in the other such an expenditure of energy is not involved. Hence, while one would expect to be able to isolate a greater amount of hydroxylamine in the latter case, yet a comparison of the results in the two cases was considered likely to prove instructive. This was repeated, and good agreement in each case was found. The results of one set of this experiment may be given as typical:—

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Nitrate present.
Grammes
Hydroxylamine
added
Grammes
Nickel deposited 0.056 0.056
Hydroxylamine 0.037 0.041

Series 6.

The results of the previous experiments indicated that energy was primarily utilized in the formation of hydroxylamine. It was, however, considered possible that nitrous acid might be formed as an intermediate substance according to equation 1, supra. Were this the case it was thought that some substance might be introduced which by chemical interaction would destroy the nitrous acid and convert it into harmless products, without otherwise interfering with the course of the electrolysis. This would tend to accelerate the removal of the nitrate, and would enable the electrical energy to be utilized more completely in effecting the deposition of the nickel.

The obvious desirability of introducing a substance that would not exert an effect upon the electrolyte and at the same time interact with the nitrous acid (if formed) suggested the use of urea.

Result.

The solution in the cathode compartment then was rendered faintly acid and the experiments conducted as before. The mean of a number of closely agreeing values gave at the end of one hour—

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No Urea

Grammes
Urea added—Acid
Solution
Grammes
Amount nickel 0.056 0.055
Amount hydroxylamine 0.037 0.038

This apparently indicated that nitrous acid is not formed as an intermediate product.

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Resumé.

The main results of this investigation may be summarized thus:—

1. In the electrical deposition of nickel from ammoniacal solution the efficiency is remarkably low. This may possibly be due to polarization effects. The nickel is perhaps present as Ni(NH3)x ions and a small proportion as Ni. Oxidation appears to take place at the anode with the formation of Ni(OH)3, for on removal of the anode it was often found to be practically coated with a black deposit, presumably of nickel hydroxide, On being allowed to stand in the solution it dissolved, forming Ni(NH3)x. Periodic phenomena are also to be observed at the anode.

2. In the presence of ammonium nitrate the rate of deposition is greatly diminished. The electrical energy appears to be utilized primarily in the reduction of the nitrate, and only when that has reached an advanced stage does deposition of the metal appear to take place in any quantity.

3. Periodic phenomena are observable at the cathode; and if the P.D. cathode solution is studied a sudden rise in the P.D. seems to be a safe criterion that electrical separation of the metal has begun.

4. The reduction of the nitrate appears to be first to hydroxylamine and finally to ammonia. The hydroxylamine undergoes some subsequent change, and nitrous acid does not seem to be formed.

5. Even in the presence of nitrate the use of the divided cell yielded distinctly better results. The deposit of metal was much more even than when the undivided cell was used. The writer hopes at a future date to investigate more fully the application of the divided cell to the separation of nickel; and the results thus far obtained seem to justify the possibility that conditions may be devised under which the deposition of the metal may be a commercial success.