PRUSSIC ACID, the familiar name for a dangerously poisonous, though chemically feeble, acid, known scientifically as “hydrocyanic acid,” or “cyanide of hydrogen,” is here taken as a convenient heading under which to treat of cyanides generally. This generic term (from κυανός, blue) is not meant to hint at any generic property; it is due simply to the fact that all cyanides, in an historical sense, are derivatives of a blue pigment which was discovered accidentally by Diesbach, a Berlin colourmaker, about the beginning of the 18th century.
The foundations of our present knowledge of cyanides were laid by Scheele (1783), whose discoveries were subsequently (from 1811) confirmed and supplemented, chiefly in the sense of quantitative determinations, by Gay-Lussac. Although we have no space for further historical notes, we must not omit to state that Gay-Lussac, as one result of his work, conceived and introduced into chemistry the notion of the “compound radical,” having shown that prussic acid and its salts are related to the group NC in precisely the same way as chlorides are to chlorine, or sulphides to sulphur. This idea, in his own eyes and in those of his contemporaries, was greatly fortified by his success in even isolating his “cyanogène” as a substance.
In preparing cyanogen or cyanides in the laboratory the operator now always starts from prussiate of potash, with which, accordingly, we begin.
Prussiate of Potash, (NC) β Fe.K 4 + 3H 2 O (syn. ferrocyanide of potassium; Germ. Blutlaτιgensalz).— This salt is being produced industrially from animal refuse (hide and horn clippings, old shoes, blood solids, &c. ), carbonate of potash, and iron filings or borings as raw materials. The carbonate of potash is fused at a red heat in an iron pear-shaped vessel suspended within a furnace, or on the cupel-shaped sole of a reverberatory furnace, and the animal matter, which should be as dry as possible, is then introduced in instalments along with the iron. The fusion is continued as long as inflammable gases are going off; then the still fluid mass is ladled out and allowed to cool, when it hardens into a black stone-like body known to the manufacturer as “metal.” When the broken-up metal is digested with water in an iron vessel prussiate of potash passes into solution, while a black residue of charcoal, metallic iron, sulphide of iron, &c., remains. The clarified solution, after sufficient concentration in the heat, deposits on cooling part of its prussiate in lemon-yellow quadratic crystals (generally truncated octahedra), which are purified by recrystallization. The last mother-liquors furnish an impure green salt, which is added to a fresh fuse and so utilized.
In former times it was believed that the prussiate was produced during the fusion process, and in the subsequent process of lixiviation simply passed into solution, until Liebig showed that this view was untenable. The fuse cannot contain ready-formed prussiate, because this salt at a red heat breaks up with formation of a residue of carbide of iron and cyanide of potassium. The metal in fact when treated with dilute alcohol gives up to it plain cyanide of potassium, and the fully exhausted residue yields no prussiate on treatment with water. The prussiate accordingly must be produced during the process of lixiviation by the action of the cyanide of potassium on some ferrous compound in the metal. Liebig thought that it was partly the metallic iron, partly the sulphide of iron present in the metal, which effected the conversion. According to more recent researches a double sulphide, K 2 S + Fe 2 S 3 , which is always produced during the fusion (from the reagents proper and the sulphur of the organic matter and that of the sulphate of potash present in the carbonate as an impurity), plays this important part. The double sulphide by the action of water breaks up into alkaline sulphide, sulphide of iron (FeS), and sulphur. This sulphide of iron is of a peculiar kind; it does what ordinary FeS does not effect, readily at least: it converts the cyanide into prussiate, thus, 6NC. K + FeS = K 2 S + (NC) β Fe.K 4 , while the eliminated sulphur of the original Fe 2 S 3 unites with another part of the cyanide of potassium into sulphocyanate, S + NCK = SNC.K, which latter salt is thus unavoidably produced as a (rather inconvenient) byeproduct. Pure prussiate of potash has the specific gravity 1*83; it is permanent in the air. It loses its water, part at 60° C., the rest at 100° C., but very slowly. The anhydrous salt is a white powder. The crystals dissolve in four parts of cold and in two parts of boiling water. It is insoluble in, and not dehydrated by, alcohol.
Prussiate of potash has the composition of a double salt, Fe(NC) 2 +4KNO, but the idea that it contains these two binary cyanides is entirely at variance with its reactions. Cyanide of potassium is readily decomposed by even the feeblest acids, and to some extent even by water, with elimination of hydrocyanic acid, and on this account perhaps is intensely poisonous. A solution of the prussiate remains absolutely unchanged on evaporation, and the action on it even of strong acids in the cold results in the formation of the hydrogen salt, (NC) β FeH 4 , which is decomposed, it is true, but only when the mixture is heated, with evolution of hydrocyanic acid. It is not poisonous. Its solution when mixed with nitrate of silver does not give a precipitate of cyanide of silver, NC. Ag, and a solution of the two nitrates, but yields a unitary precipitate of the composition (NC) 6 Fe.Ag 4 , which contains all the iron; only nitrate of potassium passes into solution. Other heavy metallic salts behave similarly. On the strength of these considerations chemists, following the lead of Liebig, view prussiate as a binary compound of potassium, K 4 , with a complex radical, N β C β Fe, ‘ ‘ ferrocyanogen.”
Hydrocyanic Acid, NC.H.—This acid is prepared most conveniently from prussiate of potash. Wöhler recommends the following method. Ten parts of powdered prussiate are placed in a retort, the neck of which is turned upwards, and a (cooled down) mixture of seven parts of oil of vitriol and fourteen parts of water is then added. If the aqueous acid is wanted, the exit-end of the retort is joined on direct to a Liebig’s condenser, which must be kept very cool by a current of cold water. If the anhydrous acid is desired, two wide-necked bottles (or two large U-tubes) charged with fused chloride of calcium and kept at 30“ C. by immersion in a water bath of this temperature, must be inserted between the retort and condenser. In this case more particularly it is indispensable to provide for a most efficient condensation of the vapours; the exitend of the condenser should be provided with an adapter going down to near the bottom of the receiver, which must be surrounded by a freezing mixture. The temperature of the latter, of course, must not be allowed to fall to the freezing-point of the distillate. The retort is heated by means of a sand bath and a brisk distillation maintained until the residue begins to dry up. The result of the reaction is in accordance with the assumption that the dilute vitriol, in the first instance, converts the prussiate, one-half into (NC) 6 Fe.H 4 , the other into (NC) β Fe.K 2 H 2 , and that through the effect of the heat these two bodies decompose each other into {(NC) 6 Fe} K 2 Fe, which remains in the residue as a precipitate, and (NC) β H 6 =6NCH, which distils over. Real NCH is a colourless liquid of 0 ’ 6967 specific gravity at 18° C., which freezes at -15° C. (Gay-Lussac) into a white fibrous solid. According to Schulz the acid, if really pure, remains liquid at - 37° C. It boils at 26° '5 C.; at 4° ’ 5 its vapour-tension already amounts to half an atmosphere. The vapour is inflammable and burns into carbonic acid, water, and nitrogen. The acid mixes with water in all proportions, with contraction and yet absorption of heat. The solution behaves on distillation like a mere mechanical mixture of its two components. Prussic acid has a very peculiar powerful smell; more characteristic still is a kind of choking action which even the highly attenuated vapour exerts on the larynx. Prussic acid is fearfully poisonous; a few drops of even the ordinary pharmaceutical preparation (of 2 per cent.) are sufficient to kill a large dog. It acts with characteristic promptitude, especially when inhaled as a vapour. Even a relatively large dose, if it has once found its way into the stomach without producing a fatal effect, is said to do relatively little harm there.
Prussic acid is characteristically prone to suffer “spontaneous decomposition.” Whether the pure anhydrous acid really is, in the strictest sense of the word, still requires to be found out; the ordinary preparation, when kept in a close bottle, soon turns brown and turbid from ‘ ‘ azulmic " acid, a substance of complex constitution. Other things are formed at the same time. The pure aqueous acid is liable to similar changes; in its case formiate of ammonia always forms the predominant product. This change is easily understood—
NC. H + 2H 2 O = NH 3 + H. COOH. Ammonia. Formic acid.
A strong aqueous prussic acid, when mixed with fuming hydrochloric acid, is soon converted into a magma of crystals of salammoniac, with formation of formic acid, which remains dissolved. And yet, most singularly, the addition to the preparation of a small proportion of hydrochloric or sulphuric acid is the best means for preventing, or at least greatly retarding, its spontaneous change in the very same direction. Aqueous prussic acid acts only very feebly (if at all) on blue litmus; it combines with aqueous caustic alkalis but does not decompose their carbonates; nor does it act upon the generality of insoluble basic metallic oxides or hydrates; mercuric oxide and oxide of silver form noteworthy exceptions to this rule.
Cyanogen, (NC) 2 ∙—When dry mercuric cyanide is heated it breaks up, below redness, into mercury and cyanogen gas; part of the latter, however, always suffers polymerization into a solid called “.paracyanogen,” and presumed to consist of molecules (NC) S . Cyanogen gas is colourless; it has the specific gravity demanded by its formula. It possesses a peculiar odour and has a characteristic irritating effect on the eyes and mucous membranes of the nose. It is poisonous. By strong pressure it is condensible into a liquid which freezes at — 34 o- 4 C., and has the following vapour-tensions P at the temperatures t stated— [table]
At ordinary temperatures water dissolves about 4 '5 times, alcohol about 23 times its volume of the gas. The solutions are liable to (very complex) spontaneous decomposition. The list of products includes oxalate of ammonia and urea. Cyanogen burns with a characteristically beautiful peach-blossom coloured flame into carbonic acid and nitrogen. This gas cyanogen, as already stated, is to cyanides what chlorine gas Cl 2 is to chlorides, but it is well to remember that the analogy, though perfect in regard to the corresponding formulae, does not, as a rule, extend to the conditions of formation of the bodies represented. Thus cyanogen does not unite with hydrogen into prιtssic acid, nor does it combine with ordinary metals in the chlorine fashion. When passed over heated potassium, it is true, it combines with it into cyanide; and caustic potash-ley absorbs it λvith formation of cyanide and cyanate (NCO.K), just as chlorine yields chloride and hypochlorite KC10; but this is about the sum-total of the analogies in action. Yet metallic cyanides of all kinds can be produced indirectly.
Cyanide of Potassium, NC.K.—An aqueous mixture of the quantities NCH and KHO no doubt contains this salt, but it smells of the acid, and on evaporation behaves more like a mixture of the two congeners than in any other way. An exhaustive union can be brought about by passing NCH vapour into an alcoholic solution of KHO; the salt NC.K then comes down as a crystalline precipitate, which must be washed with alcohol and dried, cold, over vitriol. A more convenient method is to dehydrate yellow prussiate and then decompose it by heating it to redness in an iron crucible. The Fe(NC) 2 part of the salt breaks up into cyanogen and nitrogen, which go off, and a carboniferous finely-divided iron, which remains, with cyanide of potassium, which at that temperature is a thin fluid. Yet the iron sometimes refuses to settle with sufficient promptitude to enable one to decant off the bulk even of the fused cyanide. According to private information received by the writer a French manufacturer uses a certain kind of very porous fireclay as an efficient filtering medium.
The ordinary “cyanide of potassium” of trade is not strictly that at all, but at best a mixture of the real salt with cyanate. It is produced by fusing a mixture of eight parts of anhydrous prussiate and three parts of anhydrous carbonate of potash, allowing the reaction
(NC) 6 Fe. K 4 + K 2 CO 3 = CO 2 + Fe + 5NCK + K. NCO
Cyanate to complete itself and the iron to settle, and decanting off the clear fuse. The product goes by the name of “Liebig’s cyanide,” but the process was really invented by Rodgers.
Fused cyanide of potassium assumes on cooling the form of a milky white stone-like solid. It fuses readily at a red heat, and at a white heat volatilizes without decomposition, provided that it is under the influence of heat alone; in the presence of air it gradually passes into cyanate; when heated in steam it is converted into carbonate of potash with evolution of ammonia, carbonic oxide, and hydrogen. When heated to redness with any of the more easily reducible metallic oxides it reduces them to the metallic state, while it passes itself into cyanate. It also reduces the corresponding sulphides with formation of sulphocyanate; for example, Pb(S or 0)+ NCK = Pb +NC(O or S)K. Hence its frequent application in blowpipe analysis. When heated with chlorates or nitrates it reduces them with violent explosion. The aqueous solution of the salt has a strongly alkaline reaction; it smells of hydrocyanic acid and is readily decomposed by even such feeble acids as acetic or carbonic. It readily dissolves precipitated chloride, bromide, and iodide of silver; this is the basis of its application in photography. Large quantities of the salt are used in electroplating.
Other Binary Cyanides.— Of these only a few can be noticed here. (1) Cyanide of sodium is very similar to the potassium salt. The same remark, in a more limited sense, holds for the cyanides of barium, strontium, and calcium. (2) Cyanide of ammonium (NC.NH 4 ) forms crystals volatile at 36° C. and smelling of ammonia and hydrocyanic acid. The solution in water decomposes spontaneously, pretty much like that of the free acid. But the anhydrous vapour by itself stands a high temperature, as is proved by the fact that it is produced largely when ammonia is passed over red-hot charcoal, C + 2NH 3 = H 2 + NCH. NH 3 . (3) Mercuric cyanide, Hg(NC) 2 , forms very readily when mercuric oxide is dissolved in aqueous prussic acid. The solution on evaporation and cooling deposits crystals soluble in eight parts of cold water. This salt is not at all decomposed, even when heated, by water, nor appreciably by dilute sulphuric or nitric acid; boiling hydrochloric acid eliminates the NC as hydrocyanic acid; sulphuretted hydrogen acts similarly in the cold. It gives no precipitate with nitrate of silver, nor is it changed visibly by caustic alkalis. It readily unites not only with other cyanides but also with a multitude of other salts into crystallizable double salts. Mercurous cyanide, Hg 2 (NC) 2 , seems to have no existence. When it is attempted to produce it by double decompositions, the mixture Hg + (NC) 2 Hg comes forth instead of the compound Hg 2 (NC) 2 ∙ (4) Heavy metallic cyanides are mostly insoluble in water, and the general method for their preparation is to decompose a solution of the respective sulphate, chloride, &c., with one of cyanide of potassium. The most important general property of these bodies is that they readily dissolve in solution of cyanide of potassium with formation of double cyanides, which in their capacity as double salts all exhibit, in a higher or lower degree, those anomalies which were fully explained above (see “prussiate of potash ”). These “metallo-cyanides,” as we will call them, being all, unlike plain cyanide of potassium, very stable in opposition to water and aqueous alkalis, are readily produced from almost any compound of the respective metallic radical—some from the metal itself— by treatment with solution of cyanide of potassium. In all we have said “potassium” may be taken as including sodium and in a limited sense ammonium, but the potassium compounds are best known, and we accordingly in the following section confine ourselves to these.
Metallo-cyanides. —(1) Silver.— Cyanide of silver, Ag.NC, is produced as a precipitate by addition of hydrocyanic acid or cyanide of potassium to solution of nitrate of silver. The precipitate is similar in appearance to chloride of silver and, like it, insoluble in cold dilute mineral acids, but soluble in ammonia. At a red heat it is decomposed with formation of a residue of carboniferous metallic silver. Precipitated cyanide of silver, though insoluble in hydrocyanic acid, dissolves readily in cyanide of potassium with formation of argentocyanide, AgK. (NC) 2 , which is easily obtained in crystals, permanent in the air and soluble in eight parts of cold water. Chloride of silver dissolves in cyanide of potassium solution as readily as the cyanide does and with formation of the same double salt— AgCl + 2 KN C = KC1 + AgK( NC) 2 . This salt is used very largely in electro-plating. (2) Lead.— From a solution of the acetate cyanide of lead is precipitated by addition of hydrocyanic acid or cyanide of potassium. The precipitate, Pb(NC) 2 , has the exceptional property of being insoluble in cyanide of potassium. (3) Zinc.— Cyanide of zinc, Zn(NC) 2 , is obtained by addition of hydrocyanic acid to a solution of the acetate, as a white precipitate readily soluble in cyanide of potassium with formation of a double salt, ZnK 2 (NC) 4 , which forms well-defined crystals. (4) Nickel.— The cyanide, Ni(NC) 2 , is an apple-green precipitate, which is obtained by methods similar to those given under ‘ ‘ zinc.” It readily dissolves in cyanide of potassium with formation of a crystallizable salt, NiK 2 (NC) 4 + H 2 O, the solution of which is stable in air and not convertible into one of a nickelic (Ni"') compound by chlorine (compare “cobalt” infra). The potassio-cyanides of silver, zinc, and nickel as solutions are not changed visibly by caustic alkalis, but their heavy metals can be precipitated by sulphuretted hydrogen or sulphide of ammonium, as from solutions of, for instance, the chlorides. Aqueous mineral acids (in the heat at least) decompose them exhaustively with elimination of all the NC as NCH. (5) Copper.— When cyanide of potassium solution is added to one of sulphate of copper, a yellow precipitate of cupι - ic cyanide, Cu(NC) 2 , comes down; but on boiling this precipitate loses cyanogen and is converted into a white precipitate of the cuprous salt Cu(NC). This white precipitate dissolves in cyanide of potassium with formation chiefly of two crystalline double salts, viz., CuNC + 6NCK, easily soluble in water, and CuNC + NCK. The latter is decomposed by water with elimination of Cu.NC. The solution of the 6NC.K salt is not precipitated by sulphuretted hydrogen. Solutions of potassio-cyanides of cuprosum are used in electro-plating. (6) Gold. —Metallic gold dissolves in cyanide of potassium solution in the presence of air, thus—
Au + 2KNC + ⅞O=⅜K 2 O+AuK.(NC) 2 .
This auro-cyanide of potassium is used largely for electro-gilding, for which purpose it is conveniently prepared as follows. Six parts of gold are dissolved in aqua regia and the solution is precipitated by ammonia. The precipitate (an explosive compound known as “fulminating gold ”) is dissolved in a solution of six parts of cyanide of potassium, when the double salt is formed with evolution of ammonia. The salt crystallizes in rhombic octahedra, soluble in seven parts of cold water.
In the following potassio - cyanides the heavy metals cannot be detected by means of their ordinary precipitants; these salts all behave like the potassium salts of complex radicals composed of the heavy metal and all the cyanogen. (7) Cobalt.— Cyanide of potassium when added to a solution of a cobaltous salt (CoCl 2 , &c. ) gives a precipitate soluble in excess of reagent. The solution presumably contains a cobalto-cyanide, Co(NC) 2 .≈KNC, but on exposure to air eagerly absorbs oxygen with formation of cobalti-cyanide, thus—
Co(NC), f-4KNC + ⅜O = ⅜K 2 O + Co"'(NC) 3 .3KNC.
Chlorine (Cl instead of ⅜0) acts more promptly with a similar effect If the alkaline solution is acidified and boiled, the same cobalti cyanide is produced with evolution of hydrogen—
Co(NC) 2 + 4KNC + HC1 = KC1 + ⅜H 2 + Co"'(NC) 3 .3KNC.
Cobalti-cyanide of potassium, (NC) β Co'". K 3 , forms yellow crystals isomorphous with those of red prussiate (see infra). It is a remarkably stable salt. In its behaviour to reagents it exhibits none of the characters of a cobalt salt or of a simple cyanide. Aqueous mineral acids convert it into the hydrogen salt (NC) 6 Co"Η 3 , which remains undecomposed on boiling. Heavy metallic salts produce precipitates of cobalti-cyanides; for example, (NC) β Co"'. Ag 3 . (8) Ferrosum.— See “prussiate of potash” above. (9) Fcrrieum.— Ferric hydrate and ferric compounds generally do not act upon cyanide of potassium in a manner analogous to that of ferrous compounds; but a ferri-cyanide analogous to the cobalti-salt referred to in (7) is readily produced by passing chlorine into a cold solution of ordinary prussiate, (NG) β Fe".K 4 + Cl∙=KCl + (NC) β fe"'.K 3 . In preparing the salt an excess of chlorine and elevation of temperature must be avoided, or else part of the salt is decomposed with formation of a green precipitate. The solution on evaporation and cooling yields splendid dark red crystals, soluble in 2’54 parts of water of 15 o- 6 C. (Wallace), forming a most intensely yellow solution. (Ordinary prussiate solution is only pale yellow even when saturated in the cold. ) This salt (discovered by L. Gmelin in 1822) is now being manufactured industrially and is known in commerce as “red prussiate.” In its reactions it is analogous to ordinary yellow prussiate. The same group, (NC) β Fe, which in the latter acts as a four-valent, in the red salt plays the part of a tri-valent radical, (NC) β fe. But the radical thus modified has a great tendency to assume the four-valent form; hence an alkaline solution of red prussiate is a powerful oxidizing agent, (NC) β fe.K 3 + KHO = (NC) β Fe. K 4 + HO. The HO goes to the reducing agent Like the yellow salt, red prussiate is not poisonous, at least when pure.
Ferro-and Ferri-cyanides of Iron.— The two prussiates are constantly being used in the laboratory as very delicate reagents for the detection of iron salt, and for the discrimination of ferrous and ferric compounds in solutions,—(1) ferro-cyanide and ferrous salt, white precipitate; (2) ferri-cyanide and ferric salt, intensely brown coloration; (3) ferro-cyanide and ferric salt, blue precipitate; (4) ferri-cyanide and ferrous salt, blue precipitate. These blue precipitates are being produced industrially and used as pigments, under the names of “prussian blue” and “Turnbull’s blue” for (3) and (4) respectively. The latter has been thus known for now half a century; yet the constitution of the precipitates and the true rationale of their formation have been fully cleared up only during the last few years. The main results of the researches referred to are included in the following paragraphs.
(1) Ferro-cyanide of Hydrogen, (NC) β Fe. H 4 , is obtained as a white crystalline precipitate when air-free concentrated solution of yellow precipitate is mixed with hydrochloric acid and ether. It is easily soluble in water and in alcohol. An aqueous solution of it is prepared for technical purposes by mixing a strong solution of yellow prussiate with enough tartaric acid to bring down the potassium as cream of tartar. When the solution of this ferro-hydrocyanic acid is boiled half the cyanogen goes off as N CH, while the other remains as part of a white, rather unstable, precipitate, (NC) β Fe .∩ θ . When the solution is exposed to the air, especially at higher temperatures, part of the cyanogen goes off as NCH, another part suffers oxidation into H 2 O + NC, and this latter combines with the Fe(NC) 2 of the original compound into blue bodies similar in their general properties to prussian blue. This latter change is utilized in calico-printing for producing patterns of, or dyeing with, prussian blue. The white precipitate (NC) β Fe.^ e may be looked upon as an acid of which p 2
(2) Everett’s Salt, (NC) β Fe. is the potash salt. This salt is produced in the ordinary process lor making prussic acid (see above). It is probably identical with the white precipitate produced when ferrous salt is decomposed by prussiate of potash. Everett’s salt when exposed to the air quickly absorbs oxygen and becomes blue; the reaction, as Williamson showed, assumes a simple form when the precipitate is boiled with nitric acid. One-half of the potassium is then oxidized away, and a blue double ferri-cyanide of potassium and ferrosum takes the place of the original precipitate :—
(NC) β Fe. K 2 Fe = ⅛(K 2 O as nitrate) + {(NC) β fe} "'Fe"K'. Williamson’s blue.
This blue when boiled with ferro-cyanide of potassium is reconverted into the original EveYett’s salt with formation of a solution of red prussiate—
(NC) β fe. KFe* + K 2 *K 2 . Fe(NC) β =(NC) β fe. K 3 + Fe(NC) β . FeK 2 , Red prussiate. Everett’s salt.
the asterisked radicals changing places.
(3) Soluble Prussian Blue is isomeric with Williamson’s blue. It is produced by mixing a solution of feme salt with excess of yellow prussiate, which, however, is an old process; what has been ascer tained lately is that the very same precipitate is produced by addition to a ferrous salt of an excess ς>f red prussiate.
I. (NC) β fe.K 3 + FeCl 2 = 2KCl + (NC) β fe. KFe = B'.
II. (NC) 6 Fe. K 4 + feCl 3 = 3KC1 + (NC) β Fe. Kfe = B".
B' and B" in the formulae look different, but the difference is only apparent; in either case the group (NC) β is combined with IFe and lie and IK; the bodies are identical (Skraup; Reindel). The precipitate B, though insoluble in salt solutions, is soluble in pure water, forming an intensely blue solution; hence the name.
Now the potassium in soluble prussian blue can be displaced by iron in two ways, namely, by digestion with solutions of ferrous or ferric salts. In the former case (NC) β feFeK becomes (NC) β feFe∣, or empirically (NC) 12 Fe 5 ; this is Gmelin’s (“Turnbull’s”) blue. In the lattercase (NC) β FefeKbec0mes(NC) β Fefe4, or empirically (NC) 18 Fe 7 ; this is prussian blue as discovered by Diesbach. Contrasting this latter formula with that of Gmelin’s blue (NC) 18 Fe 7 ∣, we see that the latter needs only lose ⅜Fe to become prussian blue; this surplus iron in fact can be withdrawn by means of nitric acid.
In the manufacture of prussian blue the general process is to first precipitate ferrous sulphate with yellow prussiate and then to fully oxidize the precipitate by means of nitric acid or chlorine as far as the oxygen of the air does not do it. The following receipt is recommended amongst others. Six parts each of green vitriol and yellow prussiate are dissolved separately, each in fifteen parts of water, and the solutions mixed. One part of concentrated sulphuric acid and twenty-four parts of fuming muriatic acid are then added, and after standing some hours also a solution of bleaching powder in instalments until the blue colour is fully developed. “Turnbull’s” blue is made by precipitating red prussiate of potash with excess of ferrous salt; but it is easily seen from what was said above that the use of this relatively expensive double cyanide might be dispensed with. The properties of the two pigments are {>retty much the same. They are sold in the form of solid cakes or umps, which, in addition to their blue colour, present a coppery lustre on fracture. They are stable against acids, but sensibly affected (bleached) on prolonged exposure to sunlight; and, although they stand neutral soap fairly well, they are decomposed promptly by solutions of even the carbonates of the alkalis with formation of hydrated oxides of iron. The cheaper commercial varieties are more or less largely diluted with clay, sulphate of baryta, &c. Pure prussian blue dissolves readily in a dilute solution of oxalic acid; the intensely blue solution used to serve as a blue ink, but has come to be superseded by the several more brilliant blues of the coal-tar series. These tar-blues have displaced prussian blue also in other applications, and as a commercial pigment it has besides to struggle against ultramarine. In short, it has gone very much out of use, and as a consequence the manufacture of yellow prussiate is no longer so remunerative as it used to be.
Analysis of Cyanides.— As hydrocyanic acid and cyanide of potassium are dangerously poisonous, and the latter at least is easily procured in commerce, the detection of cyanogen in this state of combination is one of the problems of forensic chemistry. To detect such cyanogen in, say, the contents of a stomach the first step is to distil the mass after acidification with tartaric acid, which decomposes cyanide of potassium but does not liberate prussic acid from prussian blue (or even prussiate of potash ?). If the distillate gives no precipitate with nitrate of silver hydrocyanic acid is absent, if it does the precipitate may have been produced by hydrochloric acid, which may then be eliminated by redistillation with borax or sulphate of soda, neither of which affects NCH. But even in the presence of chlorides the following two tests give perfect certainty. (1) A solution of hydrocyanic acid, when alkalinized with caustic potash and then mixed with, first ferroso-ferric salt and then excess of hydrochloric acid, gives a precipitate, or at least a green suspension, of prussian blue. (2) A solution of NCH, when mixed with ammonia and yellow sulphide of ammonium, is changed into one of sulphocyanate of ammonium, which, after removal of the excess of reagents by evaporation at a gentle heat, strikes an intense and very characteristic red colour with feιτic salts, which colour does not vanish (as that of ferric acetate does) on even strong acidification with mineral acid (Liebig’s test). The quantitative determination of cyanogen given as an aqueous solution of hydrocyanic acid or cyanide of potassium can (if haloids are absent) be effected by adding excess of nitrate of silver, then acidifying, if necessary, with nitric acid, filtering off, washing, drying, and weighing the cyanide of silver produced. AgNC=134 corresponds to NCH = 27 parts. Amoreex-peditious method has been invented by Liebig. A known quantity of the given prussic acid is alkalinized strongly with caustic potash and then diluted freely with water. The caustic alkali usually contains plenty of chloride as an impurity, else a little alkaline chloride must be added. A standard solution of nitrate of silver (conveniently adjusted so as to contain 6’30 grammes of fused nitrate per 1000 cubic centimètres, equivalent to 2 grammes of NCH) is now dropped in from a burette until the cloud of chloride of silver which appears locally from the first just fails to disappear on stirring, i.e., until the reaction 2KNC + AgNO 3 = KAg.(NC) 2 + KNO 3 has just been completed. One cub. cent, of silver solution used indicates 2 milligrammes of NCH. Liebig’s method lends itself particularly well for the assaying of the medicinal acid and of cyanide of potassium. The two tests for hydrocyanic acid given above apply as they stand to solutions of the cyanides of alkali and alkaline-earth metals, but not to mercuric cyanide. In regard to all other cyanides we have only space to say that from a certain set (which includes the cobalti-cyanides and the platinum cyanides) cyanogen cannot be extracted at all as NCH (or AgNC) by any known methods. Such bodies must be identified by their own specific reactions or by elementary analysis. All cyanides are decomposed by hot concentrated sulphuric acid; the carbon goes off as CO, the nitrogen remains as sulphate of ammonia and the metals as sulphates, which brings them within the range of the routine methods of analysis.
Cyanates.— These were discovered by Wöhler. The potassium salt NCO.K is produced by the oxidation of fused cyanide, for preparative purposes most conveniently by Wöhler’s method. An intimate mixture of two parts of absolutely anhydrous prussiate of potash and one part of equally dry binoxide of manganese is heated on an iron tray until the mass has become brownish black and just begun to fuse. It is now allowed to cool and exhausted by boiling 80 per cent, alcohol. The filtrate on cooling deposits crystals of the salt NCO.K. If only an aqueous solution of this salt is wanted for immediate use, the fuse may be extracted by cold water. From this solution the cyanate of silver, NCO.Ag, or lead, (NCO) 2 Pb, can be prepared by precipitation with solutions of the respective nitrates or acetates. Hot water decomposes cyanate of potash promptly with formation of carbonates of potash and ammonia, KNCO + 2H 2 O = NH 3 + KHO + CO 2 ∙ On addition of mineral acid to even the cold solution only a very little of the cyanic acid is liberated as such; the bulk breaks up at once with effervescence, thus, NCO. H + 2H 2 O = NH 3 + CO 2 + H 2 O. Very interesting is the action of the solution of cyanate of potash on sulphate of ammonia; its direct effect is the formation of cyanate of ammonia, NCO.NH 4 , but this salt almost immediately passes spontaneously into its isomer urea, which is not a cyanate at all but the amide of carbonic acid, i.e., CO(OH) 2 -2(OH) + 2NH 2 =CθiJS 2 . This reaction was discovered by Wöhler, who thus for the first time produced an organic substance from inorganic materials, or virtually from its elements. Singularly, it is this pseudo-cyanate urea which serves as a material for making cyanic acid. When hydrochlorate of urea, HC1.CON 2 H 4 , is heated to 145° C. the latter behaves as if it were cyanate of ammonia: the ammonia unites with the hydrochloric acid into sal-ammoniac and the cyanic acid is set free, but immediately suffers polymerization into cyanuric acid, a solid tri - basic acid of the composition N 3 C 3 O 3 H 3 , which, being difficultly soluble, can be freed from the sal-ammoniac by being washed with cold water. If perfectly anhydrous cyanuric acid be subjected to dry distillation it furnishes a distillate of (liquid) cyanic acid NCO. H, which must be condensed in a vessel surrounded by a freezing mixture.
Cyanic acid has a very appreciable vapour-tension even at ordinary temperatures, and the least trace of its vapour makes itself felt by a characteristically violent and dangerous action on the respiratory organs. With dry ammonia gas it unites into true cyanate of ammonia. We do not know much of its own properties, because as soon as it comes out of the freezing mixture it begins to suffer polymerization into “cyamelid” with great evolution of heat. This cyamelid is a porcelain-like mass, insoluble in all ordinary solvents and devoid of acid properties. Dry distillation reconverts it into cyanic acid.
Thiocyanates.— This term means bodies like cyanates, but containing sulphur instead of the oxygen of the latter. Thiocyanates are better known, however, as sulphocyanates or sulphocyanides. (1) The potassium salt NCS.K is formed when cyanide of potassium is fused with sulphur or certain metallic sulphides, e.g., PbS. The usual method of preparation is to fuse together forty-six parts of dehydrated yellow prussiate of potash, seventeen of dry carbonate of potash, and thirty-two of sulphur. The fuse is exhausted with boiling alcohol and the filtered solution allowed to cool, when crystals of the salt separate out. The salt is very soluble in water with characteristically large absorption of heat. (2) The ammonium salt NCS. NH 4 can be prepared by allowing a mixture of alcohol, strong aqueous ammonia, and bisulphide of carbon to stand for a time and then warming it. Thiocarbonate of ammonium, CS 2 .(NH 4 ) 2 S, is produced first, but subsequently it gives up 2H 2 S to the ammonia and becomes NCS. NH 4 , which is easily obtained in crystals. The tar water obtained in the manufacture of coal-gas sometimes contains sufficient quantities of this salt to make it worth while to recover it. Both the potassium and the ammonium salt are much used as reagents, and more especially as precipitants for copper and silver. Solutions of cupric salt when mixed with sulphocyanate assume the dark-brown colour of the cupric salt Cu(NCS) 2 , but on addition of sulphurous acid the colour disappears and a white precipitate of cuprous sulphocyanide,
NCS.Cu, comes down, which, if enough of reagent was used, contains all the copper. If sulphocyanate is added to nitrate of silver, all the silver is precipitated as Ag. NCS, similar in appearance to the chloride and, like it, insoluble in water and in nitric acid. Upon this and the fact that sulphocyanates strike a deep red colour with ferric salts Volhard has based an excellent titrim etric method for the determination of silver. (See Silver.) Syntheses of Cyanogen Compounds. — Synthetical organic chemistry dates from Wöhler’s discovery of the artificial formation of urea, and in the further development of this branch of the science cyanogen has played a prominent part. (For illustrations we may refer to certain passages in the present article and in those on Methyl and on Nitrogen.) Hence it is worth while to enumerate briefly the synthetical method for the making of cyanogen itself. (1) Hydrocyanic acid is produced when a current of electric sparks is made to cross a mixture of acetylene, C 2 H 2 , and nitrogen. (2) Cyanide of ammonium is formed when ammonia is passed over red-hot charcoal (see supra). (3) Metallic cyanides are produced when dry nitrogen gas is passed over a dry mixture of carbonate of potash or baryta and charcoal at a white heat. A similar reaction goes on spontaneously in the iron-smelting furnaces and gives rise to the formation of vapour of cyanide of potassium. (4) Sulphocyanide of ammonium is produced from bisulphide of carbon and ammonia, as explained above.
