Art. LXI.—On the Oxidation of Mercury in Air and Water, also of Iron, in Alkaline Solution.

Transactions and Proceedings of the Royal Society of New Zealand, Volume 29, 1896, Page 582

Art. LXI.—On the Oxidation of Mercury in Air and Water, also of Iron, in Alkaline Solution. By William Skey, Analyst to the Department of Mines. [Read before the Wellington Philosophical Society, 17th February, 1897.] The Oxidation of Mercury. About twenty years ago I stated before this Society † Trans. N.Z. Inst., vol. viii., p. 342. that, for certain reasons I at that time gave, the metal mercury should, like gold and platinum, oxidize in air and water conjointly, and I have now, as I believe, succeeded in proving that mercury does oxidize under these circumstances. The following is a short statement of the results upon which I base this conclusion:—

1. Clean mercury (as thrice distilled and then filtered twelve times) shaken up with a small quantity of spring water for a considerable time not only breaks up in semi-non-coalescing globules, as one would expect from our present knowledge, but, besides, imparts a slight though persistent turbidity to the water. 2. In distilled water the same effects follow, but they require a longer time to produce them. 3. When aqueous solution of the caustic or carbonated alkalies are substituted for the water these effects are rapidly produced 4. In weak sulphuric acid mercury breaks up considerably when shaken together after it has had contact for some time, and the liquid also becomes turbid. 5. If weak hydrochloric acid or iodide of potassium be substituted for sulphuric acid turbidity is not produced, and the mercury is not readily broken up. 6. Hydrochloric acid, potassic cyanide, and potassic iodide clear these turbid waters and saline solutions, and more or less agglomerate the mercury. 7. Mercury in strong solution of potash is weakly positive to this metal in both weak solutions of potash and solutions of sodic chloride. 8. Mercury in solutions of potash is strongly positive to itself in hydrochloric acid, and very strongly positive to itself in nitric acid. 9. Mercury in weak sulphuric acid is positive to mercury in nitric acid. 10. Mercury in potassic-cyanide solution is positive to mercury in potash solution. Now, the only explanation of these facts appears to me to be this: that in the experiments Nos. 1, 2, 3, and 4 the mercury has combined with certain of the elements present, forming in some cases oxides or carbonates, in other cases subchlorides and subsulphates, and these compounds became detached by the friction of the mercury on itself, thus producing the turbidity described; while the effect of the potassic cyanide and iodide, also of the hydrochloric acid in No. 5 experiment, is to dissolve these compounds, thus restoring the normal transparency of the supernatant liquids. The mercury, in fact, for the first series of experiments has (to use the miners' very expressive phrase) floured; had it not floured there would not have been any turbidity produced, and it would not have broken up permanently into small globules as it did. In the cases where only distilled water was used, or sulphuric acid, it appears certain that the mercury must have been directly oxidized by the air present therein.

The results. Nos. 7, 8, 9, and 10 were obtained last, and are strongly confirmatory of the deduction I have made from the results Nos. 1 to 4. It is not the least singular of these results that the oxidation of the mercury in the alkaline solution is assisted by the oxidation of what I must name the negative mercury by nitric acid. The reason of this is probably that the Hydrogen liberated at the negative pole combines with the nitrogen dissolved in the liquid to form ammonia, which, being easily soluble therein, leaves both poles clear of all impediment. Since writing the above I have become aware that Bellucci asserts that during the pulverisation of water ozone is formed, and that the quantity of it is greatly increased if there are solid substances in the water. Now, ozone is a substance that oxidizes mercury, so, to determine ‘as to whether or not the above results of mine were vitiated by its production, I made the following experiment. A quantity of finely-powdered glass was shaken up for a considerable time with a little boiled starch in water and potassic iodide in a phial, when it was seen that not the slightest coloration of the liquid had taken place. This result clearly showed that neither ozone nor nitrous acid in any sensible quantity had been produced. The Oxidation of iron in Alkaline Solutions. The facts above stated and referred to, showing that the noble metals generally are readily oxidized in alkaline solutions, naturally leads one to infer that iron, as being more easily oxidized in a general way than these metals, should also oxidize in such a solution; but, as is well known to chemists, iron—metallic iron—is not supposed to oxidize in a solution of this nature, nor yet in strong saline solutions. Thus Berzelius long ago affirmed this, and later, in 1871,* London Chemical News, vol. 23, page 90. Professor Grace Calvert reaffirmed it, and extended his experiments to show that alkaline carbonates act the same as the caustic alkalies in preventing oxidation. Again, and still later, “Wagner, carrying on further researches on this matter, also states, as the result of this, that “iron will not rust in alkaline waters,” evidently meaning that iron will not oxidize under these circumstances; for, if sesquioxide of iron is formed on the metal, the product is rust whether it exists in large or small quantity, whether it is visible or invisible to us. It has been this apparent anomaly that induced me to investigate the matter as rigorously as I could for myself, when results were obtained that appear to show, and very conclusively, that the popular opinions regarding the non-oxidizement of iron in these solutions are erroneous.

1. Iron in metallic connection with copper in an aqueous solution of ammonia prevents this metal from going into solution and so communicating a blue colour to the liquid, as it would do if placed therein, alone. This clearly indicates that the affinities of iron for the oxygen present in that solution are greater than those of copper for this substance, and it as clearly indicates that the iron is being oxidized. 2. Iron in a strong solution of potassic cyanide or caustic potash is positive to iron in weak solutions of these salts respectively. 3. Iron paired with gold or platinum in any alkaline solution is positive thereto, and slowly becomes tarnished, while, if the negative metal has a much smaller surface than the other has, hydrogen gas is evolved therefrom for about half an hour. These results make it very certain that when iron is electrically connected in a caustic solution with a metal negative to it, it becomes coated with a film of oxide of iron of a thickness sufficient to render it visible. It may be urged here, in defence of the old-established theory which I am combating, that these negative metals initiate the oxidation of the iron —that, in fact, there is as yet no proof given that iron when Alone in alkaline liquids will oxidize, which is the question under consideration. To this I would answer that, to my way of thinking, chemical action (e.g., oxidation) must in such cases as this always precede electrical action—that is, polarisation—as being its cause, and the only effect that should be ascribed to the negative metals (gold and platinum) is that of Accelerating the chemical action that has been already initiated by the iron and the oxygen of the water and of the air present. And it is effected, as I conceive, in this way: The hydrogen that is set free from the water by the abstraction of its oxygen by the iron is not allowed to remain upon the surface of this metal to clog it, and so retard oxidation, but is set free at a distance away—e.g., at the surface of the other metal—thus leaving the iron all bared to the exciting solution. But, whatever force there maybe in this argument, the following results of experiments especially designed to settle this point appear to be decisive, and, besides, show the great intensity of the currents that are produced by iron in these kinds of solutions. 4. Iron in a solution of caustic potash is positive to iron in sulphuric or hydrochloric, also in nitric, acid. In the case where the last acid is used the positive element is rapidly en-filmed with ferric oxide. 5. Iron in solution of potash paired with platina or iron in nitric acid is able to deposit copper from its sulphate.

The interpolar connection between the two vessels containing the acid and alkaline solutions was a third vessel charged with a strong solution of salt, from which solution two pieces of filter-paper rose on either side, and, one entering the acid and the other the alkali, the connection was thus completed. The object of interposing the saline solution was to escape the interference of those electric currents that Faraday discovered to be generated by the combination of acids with alkalies. These results (Nos. 4 and 5) demonstrate the fact that the affinity of iron for oxygen in alkaline solutions is even greater under these circumstances than its affinity for this substance in acid solutions, in which solutions, as we know, it is rapidly attacked and dissolved; and by so much they further demonstrate the fact that not only is a negative metal unnecessary to initiate the oxidizement of the iron, but that the tendency to oxidize on the part of this metal in alkaline solutions is so great that it can be very heavily handicapped without being overpowered. It should be stated here that, generally, the stronger the acid used in these experiments—the greater their action on the iron—the more electro-positive the iron in the potash is. 6. The best iron wire that I can obtain, when allowed to have contact with an alkaline solution for a considerable time, may be seen to have acquired a darkish colour when it is compared side by side with a piece of the same wire that has not been thus immersed, showing that the metal is oxidized as unassisted by coupling it with another metal. These results, taken as a whole, show very plainly that iron, like gold, platina, and silver, readily oxidizes in alkaline or strongly saline solutions in which it has hitherto been supposed to be unaffected; and they show besides that carbonic acid is not, as Professor Grace Calvert has stated, necessary for this action. In either of these solutions, in fact, it rapidly enfilms, but the film being insoluble therein, the process soon ceases, therefore the bulk of the iron is preserved, while the action may easily remain undetected. However, when the iron is polarised, either by itself or another metal, this oxidation goes on faster, and the evidences of it are quickly manifested. Thus all the metals are alike in this—that each oxidizes in air and water; the only difference being as to the degree of the intensity of their affinities for oxygen under these circumstances, some, as iron and zinc, being able to decompose water to satisfy their affinities, while others, such as platina, gold, and silver, are unable to do this, so take all their oxygen from the air present in the solution surrounding them. It will also be seen that, so far as these statements are accurate, they are favourable to the opinion I expressed over

twenty years ago,* Trans. N.Z. Inst., vol. viii., p. 342. that a great number of cades cited by Professor Becquerel and others of the polarisation of certain noble metals by the mechanical absorption of oxygen are really cases rather of polarisation by chemical absorption— —that is to say, by oxidation. To challenge investigation of the correctness of my assertions as to the general oxidation of all metals in solutions of this nature, and to give what I think may be a useful table, I append here one showing the relative affinity of twelve of the principal metals for oxygen in solutions of potash and sodic chloride respectively. Also, for comparison therewith, I append an extract of a table of mine from our Transactions showing the behaviour of these metals in potassic cyanide. The Electro-Motive Order Of Twelve Metals In Three Different Solutions, From Negative Downward To Positive. Potassic Hydrate. Sodic Chloride. Potassic Cyanide. Carbon (graphitic). Carbon (graphitic). Carbon (graphitic). Platinum. Platinum. Platinum. Gold. Gold. Iron. Mercury. Mercury. Arsenic. Silver. Silver. Antimony. Copper. Copper. Mercury. Iron. Arsenic. Lead. Lead. Antimony. Gold. Tin. Tin. Silver. Arsenic. Lead. Tin. Antimony. Iron. Copper. Zinc. Zinc Zinc. In potash, mercury is feebly positive to galena and strongly negative to chalcopyrites; while manganese binoxide is negative to the whole series in potash and potassic cyanide. In ammonia and sulphate of magnesia (for a conducting medium) platina paired with gold does not produce an electric current that I could detect.† I should state here that in this solution gold was at first momentarily negative to the platina. In concluding this section of my notes, I should not omit to remark upon the singularity of the results Nos. 4 and 5, where iron in alkaline solution is positive to itself in acid solutions. So singular indeed did this appear to me that I extended this investigation to include other metals in like solutions within its scope, to find that a number of metals, including platina, copper, lead, and zinc, gave like results to these. That metals should comport themselves in this manner is to me difficult to explain, and I would like to hear of the

matter being taken up by some one who can throw light upon it. In the meantime I will state here that, to my mind, the fact that in the direct oxidation of the iron (by free oxygen) in the one cell there is no loss of energy in undoing a chemical combination, as there is in the other cell, may be sufficient explanation. There is, too, the fact that the oxygen in this acidified cell will soon be much diluted with hydrogen, or even driven off entirely. While the fact that hydrogen may stand for the negative pole must not be overlooked. These experiments are well suited for a lecture-room, to demonstrate that it is the current of greatest intensity that dominates the galvanometer. Note.—In regard to this matter I have just ascertained that iron in a strong potash solution that has been boiled (to drive out air) is also electro-positive to itself, as in the mineral acids; also that gold in potash solution is positive to itself in these acids, and even in chlorine; while silver in the alkaline solution is also positive to itself in warm or strong nitric acid. These facts appear to prove that in all my experiments in this research water is decomposed in the potash cell, the action only ceasing when the metal therein is so coated with an oxide as to be impervious to the solution. (4th April, 1897.)