It is a common observation that the improvements introduced in practice since the first announcement of the cyanide process have been almost entirely mechanical. Although a good deal of study and research has been devoted to the chemical problems involved, the results obtained appear trifling in comparison. Few of the suggested modifications in the chemical treatment have come into practical use and those only in limited fields. Nevertheless, if we study the recent history of the cyanogen compounds from a somewhat broader standpoint, we shall find that considerable additions have been made, not only to our knowledge of their chemical properties, but also to their applications for industrial purposes.

The present review of these advances naturally falls into two divisions:

(1) processes involved in the treatment of ores; (2) manufacture of cyanides and other related compounds.

Advances in the Chemistry of Ore Treatment by Cyanide

In the attempt to extend the scope of the process, various ores have been encountered which contain gold and silver in such forms that they will not yield satisfactorily to ordinary methods of cyanide treatment. In some cases the difficulty is due to the excessive action on the solution of base metals present in the ore; in others, to the occurrence of minerals such as dyscrasite (the native alloy of antimony and silver), or a supposed oxidized manganese compound containing silver, which are either altogether insoluble in or very slightly attacked by cyanide. A particularly troublesome case is that of the so-called “graphitic” ores, containing carbon in some form which has the property of precipitating gold and silver already dissolved.

Other efforts have been directed to improving extraction by electrolysis of working solutions, and by addition of various chemicals with or without electrolysis; to the diminution of cyanide consumption by “regenerating” the spent solution by chemical or electrical means with the object of decomposing complex cyanides formed in the treatment, thus reproducing simple cyanides or other substances capable of dissolving further quantities of precious metal. ”

Treatment of Carboniferous Ores

Perhaps the most important recent papers bearing on the chemistry of ore treatment by cyanide are those of Morris Green and W. R. Feldtmann. (The references are to the bibliography at the end of this paper.) Green established the fact that the precipitation of gold by carbonaceous matter contained in certain ores and metallurgical products could be traced to the presence of occluded gases, of which carbon monoxide appears to be the most active. He also pointed out that graphite, anthracite and other dense forms of carbon, natural or artificial, have little or no precipitating effect. It was also found that charcoal gradually loses its precipitating power on exposure to air, and also loses it almost completely when heated to 500° C., if the occluded gases be extracted by a powerful Vacuum pump. Feldtmann has carried this investigation a step further, and thrown light on the reaction by which gold is deposited from solutions of its double cyanides by carbonaceous matter.

An investigation was made on a West African ore containing graphitic schist, in the treatment of which the phenomenon not very happily termed “re-precipitation”, had been a cause of poor extraction. Unsuccessful efforts were made to remedy the trouble by preliminary treatment of the ore with various metallic salts, oxidizing and reducing agents. Similar substances were added to the cyanide solution itself, but without effect. It was soon observed that an analogy existed between the effects of this graphitic schist, and those of charcoal on gold-bearing cyanide solutions, and that the action was entirely different from that of charcoal on gold chloride solutions: Gold deposited on charcoal from a chloride solution is visibly metallic, and readily soluble in cyanide, whereas gold precipitated by charcoal or graphitic schist from a cyanide solution is almost insoluble in fresh cyanide, and microscopic examination fails to reveal any metallic particles.

Finally it was discovered that a partial extraction of the precipitated gold could be made by treating the impregnated material with solutions of alkaline sulphides. Hydrosulphides of the type NaHS or KHS were the most effective; normal sulphides, as Na2S, acted nearly as well; polysulphides were, less effective and ammonium sulphide still less. It was shown that the gold so extracted appears in solution as an auro-cyanide and not as an aurosulphide. It can be readily recovered by precipitation on metallic copper, and also, but apparently with more difficulty, by precipitation with zinc or aluminum.

Feldtmann offers the surmise that the gold compound formed both in charcoal and in the graphitic schist consists of a carbonyl aurocyanide, possibly AuCN·CO(CN)2, insoluble in cyanide but soluble in alkaline sulphides with formation of thiocyanates and aurocyanides, presumably by such a reaction as the following:

Na2S + AuCN·CO(CN)2 = NaCNS + NaAu(CN)2 + CO

An observation which may have an important bearing on the subject is that cyanogen gas is readily occluded by charcoal.

It is to be regretted that the investigation did not also cover the reactions of silver double cyanides under similar conditions. Ores are known in which gold is readily soluble in cyanide whereas the associated silver is extracted very imperfectly. Experiments of R. K. Cowles indicate that the cyanide compounds of the two metals show some difference in their behavior in contact with carbonaceous matter. A selective precipitation was observed, first of the gold and later of the silver.

In the discussion on Feldtmann’s paper some alternative theories were put forward. S. J. Speak suggested that charcoal and the schist in question both probably contain oil, which may be the active precipitating agent. H. K. Picard thought the phenomenon might be due to oxidation, by a process similar to that occurring in the decolorization of organic matter by charcoal or platinum black, and stated that cyanates are always found in cyanide solutions after contact with carbonaceous matter of the kind in question. He suggested the reaction:

NaAu(CN)2 + O = AuCN + NaCNO

for the precipitation, and the reaction:

3AuCN + Na2S = NaCNS + NaAu(CN)2 + Au2

for the re-solution by alkaline sulphide, but admitted that it is difficult to see why, if the gold is precipitated as AuCN, it does not readily dissolve in cyanide. Moreover, if the precipitating effect is due to oxidation, it should be advantageous to pre-treat ores of this class with a strong reducing agent, but Feldtmann’s experiments do not bear out this view. Picard, however, pointed out that it would be necessary, after such pretreatment, to prevent the carbonaceous matter from coming in contact with air before the application of cyanide solution.

Several speakers threw doubt on the actual occurrence of graphite in the West African schist. Trewartha-James doubted the theory of adsorbed gases and explained the phenomenon by “surface energy which produces precipitation by purely physical means.” S. J. Truscott pointed out that in certain Westralian ores roasting did not completely eliminate the carbonaceous matter and that precipitation effects were still observed in the treatment of the roasted ore. He drew an analogy between this case and that of the manganiferous silver ores refractory to cyanide, in which pretreatment with a reducer has sometimes proved beneficial. F. P. Mennell thought that the association of this precipitation phenomenon with carbon was purely accidental and considered that other porous. material might be expected to act similarly, the chemical composition of the porous medium having nothing to do with the effect.

Electrolytic Regeneration of Cyanide

Turning now to another field of investigation, we find that the process introduced a few years ago by J. C. Clancy has stimulated research on the effects of electrolysis on cyanide solutions, though as yet the results obtained have not yielded any important practical benefits in ore treatment. The Clancy process mainly depends on the electrolysis of water and the secondary reaction of the nascent oxygen thus produced on the thiocyanates present in working solutions, which are converted into cyanides by the reaction:

KCNS + 2KOH + 3O = KCN + K2SO4 + H2O

Unfortunately, when the concentration of KCN in the solution is greater than that of KCNS, the former begins to be attacked, with the formation of cyanate, so that it is rarely possible, even under the most favorable circumstances, to “regenerate” to cyanide more than half the cyanogen which is present as a thiocyanate. Ferrocyanides are oxidized to some extent, by a similar reaction, but the conversion to cyanide is even less complete.

Other difficulties arise in connection with the material to be used as an anode in the process; all ordinary metals are rapidly corroded by the violent action of nascent oxygen at high current density. Peroxidized lead, “passive” iron, graphite, and fused magnetite were found to be the most resistant materials, though even these generally disintegrate in continued use.

In the course of some studies on the electrolysis of aqueous solutions of alkaline cyanides, E. F. Kern points out that a consumption of cyanide occurs when aqueous solutions are electrolyzed by direct current with insoluble anodes, the consumption being due to oxidation. Anodes of iron, nickel, and lead are dissolved and immediately precipitated, lead as hydroxide and the other metals as mixtures of hydroxides and cyanogen compounds. The consumption of cyanide increases with diminished current density; at a high current density oxygen is evolved and less metal dissolved. Anodes of peroxidized lead and “passive” iron are more permanent than those of pure metals, and are not corroded unless exposed to the air, passive iron being superior to peroxidized lead.

In the electrolysis of cyanide solutions used for leaching refractory gold and silver ores containing sulphides, no reduced cyanide consumption and no increased extraction were observed, as compared with the treatment of the same ores without electrolysis. I have had a similar experience. When a low current density at anode and cathode was used, some active solvent was produced, but not enough to compensate for loss in electrolysis. The lower the current density at the anode and the higher at the cathode, the greater was the relative cyanide consumption.

When solutions containing thiocyanates or ferrocyanides were electrolyzed, the loss of cyanide was greatly diminished. Increasing alkalinity also reduced loss of cyanide by increasing conductivity. Peroxidized lead and passive iron anodes were corroded in presence of thiocyanates at current densities above 32 amperes per square meter.

G. H. Clevenger and M. L. Hall also find that most of the decomposition of cyanide during electrolysis is due to oxygen liberated at the anode by decomposition of water. The final reaction results in the formation of carbonates, and when calcium salts are present, as is usually the case in solutions used for ore treatment, there is a precipitation of calcium carbonate. In some cases the latter might actually interfere with extraction by coating the ore particles.

On the whole, the results obtained by different investigators do not encourage the expectations entertained a few years ago that electrolysis will be of material benefit in the treatment of refractory ores by cyanide.

Effect of Mineral in Water on Cyanide Consumption

Studies on this subject have recently been made by Thomas B. Stevens and W. S. Bradley chiefly relating to the cyanicide effects of calcium and magnesium salts. They find that the common opinion with regard to the action of magnesium salts requires modification. It is known that these salts act as cyanicides by reactions such as

MgSO4 + KCN + 2H2O = Mg(OH)2 + HCN + KHSO4

and it is considered that Mg(OH)2 is useless as protective alkali owing to its insolubility. These writers find, however, that its presence prevents any lime in solution being converted into CaCO3 by atmospheric or dissolved CO2, by reason of the reaction

Mg(OH)2 + CO2 = MgCO3 + H2O

The Mg(OH)2 thus acts as a secondary protective alkali by allowing the lime to exert its full efficiency toward cyanicides present in the ore.

Another point noted by these writers, but not hitherto recognized, is the considerable solubility of calcium sulphate in cyanide solutions. Calcium sulphate and magnesium hydrate, though practically insoluble in water, are appreciably soluble in the presence of sodium chloride, a matter of importance in the case of the highly saline water sometimes used in Western Australia for making up solutions. It was found that the mixing of comparatively fresh water with solutions made from salt water, and consequently highly charged with calcium and magnesium salts dissolved in NaCl, results in the precipitation of CaSO4 andMg(OH)2. These deposits are injurious, as they coat the zinc and cause bad precipitation of the precious metals; but they are soluble in strong cyanide or in ammonium chloride.

Various salts contained in the water used for making up solutions were found to act as cyanicides, the order of destructiveness being as follows: Ca(HCO3)2, MgSO4, CaSO4, MgCl2, CaCl2. The calcium bi-carbonate acts much more rapidly than the other compounds. A test made by boiling the water in a reflux condenser, so as to expel CO2 without reducing volume, and then filtering off the precipitated CaCO3, showed that water thus treated had a greatly reduced cyanicide effect. When water containing calcium bicarbonate is used without previously being made alkaline or boiled, the reaction is as follows:

Ca(HCO3)2 + 2KCN = CaCO3 + K2CO3 + 2HCN

the cyanide acting as a softening agent in a similar manner to lime; compare:

Ca(HCO3)2 + Ca(OH)2 = 2CaCO3 + 2H2O

This reaction is the principal cause of the hardening of filter cloths, which occurs when a filter saturated with alkaline solution is washed with “fresh” water containing calcium salts.

The action of magnesium salts is much slower, and results in the gradual precipitation of Mg(OH)2. Dissolved magnesium salts cannot co-exist with protective alkali even in the presence of much sodium chloride.

Manufacturing of Cyanogen Compounds

Fixation of Atmospheric Nitrogen

Much activity has been shown of late in the study of methods for the fixation of atmospheric nitrogen, and a number of products are now made on a large scale both for direct use and as a basis for the production of other nitrogenous substances. The most remarkable development has been in the manufacture of calcium cyanamide, CaCN2. Impure commercial products consisting mainly of this substance, are now extensively used as fertilizers under the names of “lime nitrogen” and “nitrolim.” Alkali cyanides and other cyanogen compounds are obtained by further treatment of cyanamides in a variety of ways. Other processes are used to transform cyanamide into ammonia, ammonium salts, and other nitrogenous compounds of commercial value.

The earlier attempts at the fixation of atmospheric nitrogen were mostly based on the reactions taking place between barium oxide, carbon, and nitrogen at high temperatures. An important investigation of this process has been published by Thomas Ewan and Thomas Napier, who summarize the chief data obtained by previous investigators in this field. The formation of cyanides from a mixture of BaO, C, and N requires a minimum temperature of 1,200° C., but proceeds best at 1,400° C., when 40 per cent, of the Ba may be cyanized. Hempel (1890) showed that the percentage of Ba cyanized may be increased by pressure. Readman (1894) applied the electric furnace for this process, which has been carried out on a large scale by the Scottish Cyanide Co. Frank and Caro showed (1898) that a large part of the nitrogen is fixed in the form of barium cyanamide, BaCN2. Caro also found (1907) that the addition of alkali and alkaline earth fluorides accelerates the reaction, and enables it to proceed at temperatures of 900° to 1,100° C. Snodgrass (1907) showed that potassium carbonate has a similar action.

Ewan and Napier (1909) studied the action of catalysts to obtain the reaction at lower temperatures. In their first experiments, mixtures of barium carbonate and charcoal were placed in an iron boat in a porcelain tube, and heated in a current of dry nitrogen. The products finally obtained were barium cyanide Ba(CN)2, barium cyanamide BaCN2, carbon monoxide CO, and unaltered BaCO3, C, and N. The following reactions are supposed to occur:

They summarize their results as follows:

1. Absorption of nitrogen begins at 900° to 930° C.
2. The amount absorbed increases rapidly from 1 per cent. Ba combined at 930° to 40 per cent, at 1,000°, in tests made under the conditions that four molecules of N act on one molecule of BaCO3 for a period of 2 hr.
3. The greater part of the nitrogen is fixed as cyanide Ba(CN)2 and not as cyanamide. When calcium is used in place of barium, the nitrogen is fixed almost exclusively as cyanamide, CaCN2.
4. At 960° about 2½ per cent, of the nitrogen is fixed; at 1,000° about 10 per cent.
5. Addition of K2CO3 seems to improve the result, but the difference obtained falls within the limits of error in estimating temperatures.
6. No cyanide is produced until about 30 per cent, of the BaCO3 is converted into BaO, and until the percentage of CO has fallen to 30.

It is commonly supposed that barium carbide (BaC2) is formed as an intermediate product, but Ewan and Napier find no evidence of this. From theoretical considerations BaC2 could not be formed at all in presence of mixtures of CO and N from which N is freely absorbed by BaO and C. It is probable, however, that when calcium is used instead of barium the reaction:

CaC2 + CO ⇔ CaO + 3C

is a necessary intermediary.

Cyanid Gesellschaft reports that nitrogen has no action on a mixture of CaO and C at 1,100°, and that the reaction CaO + 2C + N2 = CaCN2 + CO takes place best at 2,000°. At this temperature nearly all the nitrogen is absorbed.

K. Kaiser observes that if barium oxide or carbonate, intimately mixed with carbon, be heated to 900° to 1,400° C. and allowed to cool in a dry atmosphere of nitrogen under pressure, as much as 90 to 95 per cent, of the theoretical amount of nitrogen combines with the barium.

A French patent of 1913 describes a method of carrying out the reaction between barium carbonate, carbon, and nitrogen in an electric furnace. Tar is used as binding material for the ingredients; these are molded into balls and part of the tar recovered by distillation. The residue is then heated to 1,600° C. in the electric furnace and acted on by “producer gas” or other gas containing nitrogen, to form barium cyanide. The electrodes consist of two rings of carbon placed at the ends of the middle reaction zone of an inclined cylindrical furnace; the fused mass serves to conduct the current. Barium oxide may be regenerated by decomposing the barium cyanide by means of steam:

Ba(CN)2 + 3H2O = BaO + 2NH3 + 2CO

When calcium is used instead of barium as the vehicle for the fixation of nitrogen, it was found by Barzano and Zanardo that the yield of calcium cyanamide is increased by mixing calcium carbide with sufficient calcium cyanamide from a previous operation to form a friable mass, before subjecting it to the action of nitrogen.

The manufacture of cyanamide by the process of Frank and Caro is carried out on a large scale at the works of the American Cyanamid Co., Niagara Falls, Ont. Nitrogen is obtained by passing air over heated copper. The copper oxide thus formed is again reduced by the action of natural gas, so that it may be used repeatedly. Operations began in January, 1910, with a 10,000-ton plant, which has since been largely extended.

Cyanides and cyanamides of the alkali metals may be produced without the preliminary formation of a calcium or barium compound. E. A. Ashcroft proposes to do this by using an alloy of the alkali metal with some heavy metal such as lead, and first fusing this with a portion of the product sought (e.g., NaCN or Na2CN2) at about 600° C., to separate the alkali from the heavy metal. The resulting melt, freed from the heavy metal, is now caused to react with cyanogen or a substance yielding cyanogen, with or without electrolysis of the melt.

Conversion of Cyanamides into Cyanides

We must now mention some of the recently proposed methods for converting calcium cyanamide into cyanides or other nitrogenous substances of commercial value. These may be roughly classified as fusion methods and wet methods.

Among the former, probably the best known is that of Erlwein and Frank by which the crude calcium cyanamide is converted into a product containing the equivalent of 25 to 30 per cent. KCN, by fusing with carbon and common salt. Careful regulation of the temperature appears to be necessary. I have found that the process is not easily imitated in laboratory experiments. The product has been placed on the market under the name of “surrogat.” For extraction of precious metals, it appears to be as effective, per unit of cyanogen contained, as ordinary sodium cyanide. It requires, however, to be digested with water and the insoluble carbonaceous residue filtered off before use.

J. C. Clancy proposes to heat calcium cyanamide with its own weight of a mixture of equal parts of sodium sulphide and chloride in presence of carbonaceous matter.

E. E. Naef treats crude calcium cyanamide at a temperature of’300° to 500° C. with sulphur or a substance yielding sulphur:

2CaCN2 + S = CaS + Ca(CN)2 + N2

The product is then stirred with water and filtered from the residue of CaS + C. The filtrate is boiled gently with zinc dust, and the zinc cyanide formed converted into alkali cyanide by known methods.

Calcium cyanamide may be used as a source of ammonia by acting on it with steam under pressure at a temperature of 170° C., using an agitator to prevent formation of lumps. C. Manuelli observes that under favorable conditions nearly theoretical yields of ammonia can be obtained at a cost of about 1.7c. per pound of nitrogen utilized. Another method (E. E. Naef) is to pass a current of dry hydrogen, either alone or mixed with CO2, CO, or N, at ordinary or higher pressure, over a heated mass of calcium cyanamide.

The wet methods are based on the fact that cyanamide compounds, such as CaCN2, are transformed by water into the sparingly soluble crystallizable salt dicyandiamide (CN·NH2)2. At a temperature of 13° to 14° C. about 90 per cent, of the nitrogen of calcium cyanamide may be dissolved in 25 times its weight of water, in 8 hr., the extraction proceeding rapidly at first, but very slowly later (C. Manuelli). Various salts of dicyandiamide may be prepared by digesting the substance in presence of acids with the oxides of the metals whose salts are required. The polymerization of cyanamide in aqueous solution is promoted by addition of fixed alkalies or ammonia. According to Grube and Kruger, the reaction appears to be due to the union of undissociated cyanamide (CN·NH2) with cyanamide ions (CN·NH)’, to form monobasic dicyandiamide ions H·(CNNH)’2, presumably in the following stages:

(a) Formation of cyanamide by hydrolysis of calcium cyanamide:

CaCN2+ 2H2O = Ca(OH)2 + CN·NH2

(b) Dissociation of a part of the cyanamide:

CN·NH2 = H’ + (CN·NH)’

(c) Combination of dissociated with undissociated cyanamide:

(CN·NH)’ + CN·NH2 = H(CN·NH)2′

(d) Formation of dicyandiamide:

H(CN·NH)2′ + H’ = (CN·NH2)2

It is found that at a given concentration of total cyanamide, polymerization proceeds most rapidly, when the concentrations of cyanamide and of dissociated cyanamide ions are equal.

For the preparation of dicyandiamide from commercial calcium cyanamide, the addition of a foreign base is unnecessary; the conversion may be brought about by precipitating the lime at intervals so as to maintain approximately equal concentrations of CN·NH2 and (CNNH)’.

Alkali salts of dicyandiamide are produced by heating the substance with caustic alkalies in presence of an absorbent of water:

4NaOH + (CNNH2)2 = (CN·NNa2)2 + 4H2O.

If the water is not absorbed, part of it reacts with the melt thus:

Na2CN2 + 3H2O = Na2CO3 + 2NH3

The absorbents used are alkali or alkaline earth metals or their oxides, amides, nitrides, carbides or alloys.

Dicyandiamide is; converted into cyanide by fusing with alkaline carbonates and carbon:

(CN·NH2)2 + Na2CO3+ 2C = 2NaCN + NH3 + N + H + 3CO

or it may be converted into a variety of nitrogenous products suitable for manures, explosives, etc.

The conversion. of cyanamide into dicyandiamide may also be brought about by heating its solution with a catalytic agent. H.

Immendorf and H. Kappen propose to carry out this process by means of ferric oxide or hydrate, manganous or manganic oxides or their hydrates, stannic acid, chromic oxide or hydrate, hydrated silicic acid or similar substances. The products obtained are ammonia, urea, and dicyandiamide. During the process the catalyst absorbs the bases liberated and loses its efficiency. This may be restored by treatment with acid or with water. Commercial calcium cyanamide contains calcium chloride, which hinders the separation of urea and dicyandiamide. This difficulty may be avoided by converting the calcium chloride into sulphate or carbonate, which may be separated by reason of their insolubility in alcohol.

H. Kappen obtains thiourea by treating dilute aqueous solutions of cyanamide with hydrogen sulphide at high temperatures, in closed vessels under pressure. The addition of acid and alkaline reagents and of soluble compounds of arsenic, antimony, or tin increases the efficiency of the process.

Production of Cyanogen Compounds

from Nitrogenous Gases by Catalysis

Another method for the manufacture of cyanides and related compounds, which has been largely developed of late years, consists in the combination, under suitable conditions, of volatile carbon and nitrogen compounds. Usually this takes place at a high temperature under the influence of a “catalyst,” i.e., of some substance which does not undergo any permanent chemical change in the process.

It has long been known that hydrocyanic acid or ammonium cyanide could be produced by passing mixtures of carboniferous and nitrogenous gases, over heated platinized pumice, as for example in Woltereck’s process:

CO + NH3 = HCN + H2O

C. Beindl proposes to form hydrocyanic acid and cyanogen compounds by catalytic combination of gaseous and volatile carbon compounds with gaseous and volatile nitrogen compounds, by the use of a contact material consisting of a metallic oxide or a mixture of metallic oxides. The gaseous mixture should contain not more than 7 volumes of the carbon compound to 1 volume of the nitrogen compound. It is claimed that oxides of metals of the iron group have the property of inducing the rapid formation of hydrocyanic acid at relatively low temperatures, from dry mixtures of these gases.

It has been found that the yield is increased by introducing into the arc of an electric furnace the vapors of metals and metallic compounds, such as those of copper or iron and their salts.

W. Moldenhauer and O. Wehrheim describe a method by which nitric oxide can be obtained by catalytic combustion of cyanogen compounds. Ammonia may also be used as the source of nitrogen, but the yield of nitric oxide is less than with cyanogen, and more attention must be paid to the condition of the catalyst and time of contact with it.

J. E. Bucher proposes to produce alkali cyanides by the action of atmospheric nitrogen on alkali carbonate and carbon dissolved in iron, the iron acting as a catalytic agent. The alkali cyanide formed is removed from the reaction zone by distillation under diminished pressure.

The Deutsche Gold und Silber Scheide Anstalt propose to pass gaseous mixtures containing nitrogen obtained by heating certain waste products such as “vinasses” through highly heated conduits lined with refractory material such as Dinas brick, fused quartz, or a fused mixture of zirconia and quartz which remains gas-tight and non-porous at the high temperature employed. It is said that practically the whole of the nitrogen is obtained in the forms of HCN and NH3, the relative proportions of these varying with the raw material used.

In the process of Barzano and Zanardo a mixture of nitrogen (say 60 per cent.), hydrogen (32 per cent.), and hydrocarbons (6 per cent.) is raised to a high temperature, for example in the electric arc, and hydrocyanic acid absorbed by circulating the issuing gases through cooling and absorption apparatus. Sufficient nitrogen is added to maintain, with the hydrogen liberated, the concentration of these gases most favorable to the reaction.

Catalysis may be employed for the production of ammonia, not only from gaseous mixtures but also from solid organic nitrogenous compounds. F. Schreiber proposes to act on the latter with contact masses containing hydrated iron oxide at temperatures below red heat. It is claimed that as much as 80 per cent, of the nitrogen in pyridine and in cyanogen salts may be thus converted into ammonia, while ordinary destructive distillation yields not more than 20 per cent.

Production of Cyanogen Compounds

from Gaseous Mixtures by Absorption with a Metallic Salt

The cyanogen contained in crude illuminating gas has for some years been utilized as a source of cyanides and cyanogen compounds on a commercial scale, as in the methods of Rowland and Bueb.

L. Bergfeld uses as purifying material dehydrated copper sulphate, or a mixture of salts, e.g., alkaline earth sulphates or chlorides, which combine with NH3 but not with H2S, with metallic oxides capable of combining with H2S. The mixture is treated with ammonia gas before use. The reaction is of the type:

H2S + NH3 + CaSO4·NH3 = CaS + (NH4)2SO4

At the same time some SO2 (which is fixed by the purifying material) and sulphur are also formed, thus:

SO3 + H2S = H2O + SO2 + S

When the purifying material is exhausted it is heated, with regulated access of air, and ammonia liberated as follows:

CaS + (NH4)2SO4 + 3O = CaSO4 + SO2 + H2O + 2NH3
SO3 + 2NH3 + 2H2O = (NH4)2SO3·H2O

By gentle warming in presence of air, the ammonium sulphite deposited is rapidly oxidized to sulphate.
Cyanogen in the crude gas is fixed by the purifying material as cuprous ammonium cyanide. On heating with air this yields NH3 and CO2, any thiocyanate giving NH3, CO2 and SO2. After passing the purifying material the coal gas may be further purified by leading it over alkali or alkaline earth sulphides, and if necessary over sodium amalgam.

C. A. Bergh uses a solution of a zinc salt, containing also ferrous chloride, as an absorbent. The free acid formed is neutralized, and zinc sulphide and cyanide are precipitated, by addition of calcium carbonate, or some other sparingly soluble substance. This is added in quantity corresponding to the amount of H2S and CN present. Iron remains in solution. By this method zinc is separated from iron and H2S from CN by a single operation.

W. H. Coleman passes coke-oven or other gases containing hydrocyanic acid through a tower packed with lumps of ferrous iron ore. A solution of sodium carbonate, or other suitable alkaline absorbent in solution or suspension, is circulated through the tower, and the resulting alkali or alkaline-earth ferrocyanide liquor is treated for recovery of the ferrocyanide.