# 1911 Encyclopædia Britannica/Prussic Acid

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PRUSSIC ACID, or Hydrocyanic Acid, HCN, an organic acid first prepared in 1782–1783 by C. Scheele and subsequently examined by J. Gay-Lussac. It is present in varying amounts in certain plants, being a product of the hydrolysis of the cyanogenetic glucosides, e.g. amygdalin (q.v.). It may be prepared by heating a mixture of cyanogen and hydrogen to 500°–550° C. (M. Berthelot, Ann. chim. phys., 1879 (5), 18, p. 380); by passing induction sparks through a mixture of acetylene and nitrogen; by the dry distillation of ammonium formate; by the decomposition of the simple cyanides with mineral acids; and by distilling potassium ferrocyanide with dilute sulphuric acid (F. Wöhler, Ann., 1850, 73, p. 219), the anhydrous acid being obtained by fractional distillation of the aqueous distillate, special precautions being necessary owing to the excessively poisonous nature of the free acid:
K$_4$Fe(NC)$_6$+3H$_2$SO$_4$＝2K$_2$SO$_4$+FeSO$_4$+6HCN.
The free acid is a colourless liquid with a smell resembling bitter almonds; it boils at 26·1° C., and may be solidified, in which condition it melts at −14° C. It burns with a blue flame, and is readily soluble in water, but the solution is unstable and decomposes on standing, giving amorphous insoluble substances, and ammonium formate, oxalic acid, &c. An aqueous solution of hydrogen peroxide converts it into oxamide, (CONH$_2$)$_2$, and reduction by zinc and hydrochloric acid gives methylamine. The anhydrous acid combines with hydrochloric, hydrobromic and hydriodic acids to form crystalline addition products, which are decomposed by water with the formation of the corresponding ammonium salt and formic acid. It combines with aldehydes and ketones to form the nitriles of oxy-acids, for example, CH$_3$CHO+HCN=CH 3 CH(OH)(CN). It is a very weak monobasic acid, and the aqueous solution has a very low electric conductivity.
Cyanides The salts of this acid, known as cyanides, may be prepared by the action of cyanogen or of gaseous hydrocyanic acid on a metal; by heating the carbonates or hydrooxides of the alkali metals in a current of hydrocyanic acid; by heating alkaline carbonates with carbon in the presence of free nitrogen: BaCO 3 + 4 C + N2 = Ba(NC) 2 + 3CO; by ignition of nitrogenous organic substances in the presence of alkaline carbonates or hydroxides; or by processes of double decomposition. The alkali and alkaline earth cyanides are soluble in water and in alcohol, and their aqueous solution, owing to hydrolytic dissociation, possesses an alkaline character. When heated in contact with air they undergo a certain amount of oxidation, being converted to some extent into the corresponding cyanate. The cyanides of other metals are decomposed by heat, frequently with liberation of cyanogen. The cyanides are usually reducing agents. Those of the heavy metals are mostly insoluble in water, but are soluble in a solution of potassium cyanide, forming more or less stable double salts, for example KAg(NC)2, KAu(NC) 2. Lead cyanide, Pb(NC) 2, however, does not form such a salt, and is insoluble in potassium cyanide solution.
Ammonium cyanide, NH 4 NC, a white solid found to some slight extent in illuminating gas, is easily soluble in water and alcohol, and is very poisonous. Its vapour is inflammable. It is obtained by passing ammonia gas over hot coal; by subliming a mixture of ammonium chloride and potassium cyanide; by passing a mixture of ammonia gas and chloroform vapour through a red hot tube; and by heating a mixture of ammonia and carbon monoxide: CO+2NH 3 = NH 4 NC+H 2 O. Barium cyanide, Ba(NC) 2, prepared by the action of potassium cyanide on baryta, or by passing air over a heated mixture of barium carbonate and coal, is a white solid, which when heated with water to 300° C. loses the whole of its nitrogen in the form of ammonia. Mercuric cyanide, Hg(NC)2, is a sparingly soluble salt formed by dissolving precipitated mercuric oxide in hydrocyanic acid, or by boiling potassium ferrocyanide with mercuric sulphate and water: 2K4Fe(NC)6+3HgSO4=3Hg(NC)2+ 3K 2 SO 4 -{-K 2 Fe[Fe(NC) 6 ]. Its aqueous solution is not an electrolyte, and consequently does not give the reactions of the mercury and cyanogen ions. When heated it yields mercury, cyanogen and paracyanogen. Silver cyanide, AgNC, is formed as a white precipitate by adding potassium cyanide to silver nitrate solution; or better, by adding silver nitrate to potassium silver cyanide, KAg(NC) 2, this double cyanide being obtained by the addition of one molecular proportion of potassium cyanide to one molecular proportion of silver nitrate, the white precipitate so formed being then dissolved by adding a second equivalent of potassium cyanide. On concentration the double salt separates as hexagonal tables. Dilute mineral acids decompose it with the formation of insoluble silver cyanide and hydrocyanic acid: KNC AgNC+HNO3=HCN+ KNO 3 +AgNC. A boiling solution of potassium chloride with the double cyanide gives silver chloride and potassium cyanide.
Potassium cyanide, KNC, and sodium cyanide, NaNC, are two of the most important of the salts of hydrocyanic acid, the former being manufactured in large quantities for consumption in the extraction of gold. Potassium cyanide may be obtained by fusing potassium ferrocyanide either alone - K4Fe(NC)6=4KNC+ FeC2+N2 - or with potassium carbonate [V. Alder, English patent 1 353 (1900)]; in the latter case the chief reaction probably is: K 4 Fe(NC) 6 + K 2 CO 3 =4KNC, ',+ 2KOCN + CO + Fe more potassium ferrocyanide is occasionally added in small quantities, in order to decompose the cyanate formed; 2KOCN+2K4Fe(NC)6= ioKNC + 2FeO + 4C + 2N2; 2FeOA+ 2C = 2CO + 2Fe. The reaction is accompanied by much frothing, and the whole is filtered when in a state of tranquil fusion. Rossler and Hasslacher prepare the double potassium sodium cyanide by fusing potassium ferrocyanide with sodium, the product of fusion being extracted with water and the solution evaporated: K 4 Fe(NC) 6 +2Na = Fe+ 4KNC 2NaNC. This process gives a product free from cyanate, which was always formed in the older fusion processes.
Many other processes have been devised. D. T. Playfair [Eng. pat. 7764 (1890)] decomposes sulphocyanides by fusing with zinc: the zinc is heated with a small quantity of carbon and when completely fused potassium sulphocyanide is added, the mass being well stirred and heated until it thickens and begins to turn red; finally it is allowed to cool out of contact with air, lixiviated with water, the solution decanted, and evaporated to a paste in vacuo. The potassium sulphocyanide is obtained from ammonium sulphocyanide, which is formed by washing crude coal gas with water containing suspended sulphur. Various processes involving the use of atmospheric nitrogen have been devised, but in most cases they do not yield good results. More successful results are obtained by the use of ammonia. The Stassfiirter Chem. Fabrik [Eng. pat. 935 0 - 2 (1900)] pass ammonia over a mixture of alkali or alkaline carbonate and charcoal, first at a dull red heat and then at a bright red heat: KHO + NH 3 + C = KNC + H 2 O. -{- H2. H. Y. Castner [Fr. pat. 242938 (1894)] passes anhydrous ammonia over heated sodium to form sodamide, which is then brought in a molten condition into contact with carbon: NaNH 2 +C = NaNC+H 2. The Deutsche Gold and Silber Scheide Anstalt [Eng. pat. 33 28, 3329 (1901)] prepare sodium cyanamide by melting sodium with carbons or some hydrocarbon, and passing ammonia over the melt at from 400 0 -600° C. The temperature is then raised to 700°-800° C., and the sodium cyanamide in contact with the residual carbon forms sodium cyanide. H. W. Crowther and E. C. Rossiter (Journ. Soc. Chem. Ind., 1893, 13, p. 887) digest carbon bisulphide with ammonia and lime in quantities slightly in excess of those demanded by the following equation: 2CS 2 + 2NH 3 + 2Ca(OH) 2 = Ca(SCN) 2 Ca(SH)2+4H2O; the product is then treated with a current of carbon dioxide, calcium carbonate being precipitated, sulphuretted hydrogen escaping, and calcium sulphocyanide remaining in solution. The sulphocyanide is converted into the potassium salt by adding potassium sulphate, and finally desulphurized by lead, zinc, or iron. Potassium cyanide is an excessively poisonous, colourless, deliquescent solid; it is readily soluble in water, but almost insoluble in absolute alcohol. It is stable ir, dry air, but is easily oxidized when fused, in which condition it is a powerful reducing agent. It dissolves gold (q.v.) in the presence of water and atmospheric oxygen. It is also largely used by the jeweler, electroplater and photographer.
Double Cyanides The double cyanides formed by the solution of the cyanide of a heavy metal in a solution of potassium cyanide are decomposed by mineral acids with liberation of hydrocyanic acid and formation of the cyanide of the heavy metal. Besides these, other double cyanides are known which do not suffer such decomposition, the heavy metal present being combined with the cyanogen radical in the form of a complexion. The most important members of these classes are the ferroand ferri-cyanides and the nitroprussides.
Potassium ferrocyanide, K 4 Fe(NC) 6, (yellow prussiate of potash), was first obtained by decomposing Prussian blue with caustic potash: Fe4[Fe(NC)6]3 + 12KHO = 3K 4 Fe(NC) 6 +4Fe(OH) 3; it may be also obtained by warming a solution of ferrous sulphate with an excess of potassium cyanide: FeSO4-I-6KNC = K4Fe(NC)6+ K2SO4. The older processes for the commercial preparation of this salt, which were based on the ignition of nitrogenous substances with an alkaline carbonate and carbon, have almost all been abandoned, since it is more profitable to prepare the salt from the byproducts obtained in the manufacture of illuminating gas. W. Fowlis [Eng. pat. 9474 (1892)] passes the gas (after freeing it from ammonia) through a solution of potassium carbonate containing ferric oxide or ferrous carbonate (actually ferrous sulphate and potassium carbonate) in suspension; the sulphuretted hydrogen in the gas probably converts the iron salts into ferrous sulphide which then, in the presence of the hydrocyanic acid in the gas, and the alkaline carbonate, forms the ferrocyanide, thus: FeS+6HCN+ 2K 2 CO 3 = K 4 Fe(NC) 6 + H 2 S + 2CO 2 + 2H 2 O. The salt is recovered by crystallization. The process is not very efficient, since the solutions are too dilute and large quantities of liquid have to be handled. A large quantity of the salt is now prepared from the "spent oxide" of the gas works, the cyanogen compounds formed in the manufacture of the gas combining with the ferric oxide in the purifiers to form insoluble iron ferrocyanides. The soluble salts are removed by lixiviation, and the residue is boiled with lime to form the soluble calcium ferrocyanide, which is finally converted into the potassium salt by potassium chloride or carbonate.
The salt crystallizes in large yellow plates, containing three molecules of water of crystallization. It is soluble in water, but insoluble in alcohol. It is not poisonous. When fused with potassium carbonate it yields potassium cyanide; warmed with dilute sulphuric acid it yields hydrocyanic acid, but with concentrated sulphuric acid it yields carbon monoxide: 6H 2 O + K 4 Fe(NC) 6 + 6H 2 SO 4 = 2K 2 SO 4 + FeSO 4 + 3(NH4)2SO4 + 6CO. Oxidizing agents (Cl, Br, H2O2, &c.) convert it into potassium ferricyanide (see below), a similar result being attained by the electrolysis of its aqueous solution: 2K 4 Fe(NC)s + 2H 2 O = 2KOH + H2 + 2K 3 Fe(NC) 6. Potassium ferrocyanide may be estimated quantitatively in acid solution by oxidation to ferricyanide by potassium permanganate (in absence of other reducing agents): 5K 4 Fe(NC)s + KMnO 4 + 4H2SO4= 5K 3 Fe(NC)s + 3K2SO4+MnSO4+4H2O.
Hydroferrocyanic acid, H 4 Fe(NC)s, is best obtained by decomposing the lead salt with sulphuretted hydrogen under water, or by passing hydrochloric acid gas into a concentrated ether solution of the potassium salt. In the latter case the precipitate is dissolved in water, reprecipitated by ether, and washed with ether-alcohol.
It is a tetrabasic acid, of markedly acid character, and readily decomposes carbonates and acetates. It dissolves unchanged in concentrated sulphuric acid, and oxidizes readily in moist air, forming Prussian blue.
Prussian blue, Fe 7 (NC) 18 or Fe4[Fe(NC)6]3, ferric ferrocyanide, was discovered in 1710 by a German manufacturer named Diesbach, who obtained it by the action of fused alkali and iron salts on nitrogenous organic matter (e.g. blood). It is now prepared from the calcium ferrocyanide formed in gas purifiers (see above) by decomposition with ferrous' sulphate. J. Bueb (Congress of German Gas Industries, March 1900) brings gas (free from tar) into intimate contact with a saturated solution of ferrous sulphate, when a "cyanogen mud" is obtained. This is heated to boiling, and the residue after filtration contains about 30% of Prussian blue. On the small scale it may be prepared by adding an acid solution of a ferrous salt to a solution of potassium ferrocyanide. The grey precipitate first formed is allowed to stand for some hours, well washed, and then oxidised by a warm solution of ferric chloride: 6K 2 Fe[Fe(NC) 6 ] + 3O = Fe7(NC)18 + 3K 4 Fe(NC) 6 + Fe2O3. It is a dark blue powder with a marked coppery lustre. It is insoluble in water and is not decomposed by acids.
Soluble Prussian blue, K2Fe2[Fe(NC)6]2, potassium ferric ferrocyanide, is formed when a solution of potassium ferrocyanide is added to an insufficiency of a solution of a ferric salt (t), or when potassium ferricyanide is added to a ferrous salt (2): (t) 2K 4 Fe(NC) 6 + 2FeC1 3 = 6KC1 + K2Fe2[Fe(NC)6]2 (2) 2K 3 Fe(NC)s + 2FeC1 2 = 4KC1 -{- K2Fe2[Fe(NC)s]z.
It is soluble in water, but is insoluble in salt solutions.
Potassium ferricyanide, K 3 Fe(NC)s, red prussiate of potash, is obtained by oxidizing potassium ferrocyanide with chlorine, bromine, &c., 2K 4 Fe(NC) 6 + C1 2 = 2K 3 Fe(NC) 6 + 2KC1. G. Kassner (Chem. Zeit., 1889, 13, p. 1701; 17, p. 1712) adds calcium plumbate to a solution of potassium ferrocyanide and passes carbon dioxide through the mixture: 2K 4 Fe(NC) -f-Ca 2 PbO 4 -} 4CO2= K3Fe(NC)6+ K2CO3+PbCO3+2CaCO3. The mixture of calcium and lead carbonates is filtered off and roasted at a low red heat in order to regenerate the calcium plumbate. It crystallizes in dark red monoclinic prisms which are readily soluble in water. The solution decomposes on standing, and in the presence of an alkali acts as an oxidizing agent: 2K 3 Fe(NC) 6 +2KHO = 2K4Fe(NC)6+H2O+O. With silver nitrate it gives an orange red precipitate of silver ferricyanide, Ag 3 Fe(NC)s. With a pure ferric salt it only gives a brown coloration. It can be estimated quantitatively by mixing a dilute solution with potassium iodide and hydrochloric acid in excess, adding excess of zinc sulphate, neutralizing the excess of free acid with sodium bicarbonate, and determining the amount of free iodine by a standard solution of sodium thiosulphate. The zinc sulphate is added in order to remove the ferrocyanide formed as an insoluble zinc salt: 2K 3 Fe(NC)6+2KI=2K 4 Fe(NC) 6 -O 2. As an alternative method it may be decomposed by hydrogen peroxide in alkaline solution and the amount of evolved oxygen measured: 2K 3 Fe(NC) 5 + 2KHO + H 2 O 2 = 2K 4 Fe(NC) 6 + 2H,O + O2.
Turnbull's blue, Fe5(NC)12 or Fe3[Fe(NC)6]2, ferrous ferricyanide, is best obtained by adding a hot solution of potassium ferricyanide to a ferrous salt, and allowing the mixture to stand some time in the presence of an iron salt: 2K 3 Fe(NC)s+3FeSO 4 = Fe3[Fe(NC)s]z+ 3K 2 SO 4. It is insoluble in dilute acids.
Hydroferricyanic acid, H 3 Fe(NC)s, obtained by adding concentrated hydrochloric acid to a cold saturated solution of potassium ferricyanide, crystallizes in brown needles, and is easily decomposed.
Nitroprussides The nitroprussides are salts of the type M2Fe(NC)5. NO. The free acid forms dark red deliquescent crystals and is obtained by decomposing the silver salt with hydrochloric acid, or the barium salt with dilute sulphuric acid.
Sodium nitroprusside, Na 2 Fe(NC) 5 NO2H 2 O, is the commonest salt. It is prepared by oxidizing potassium ferrocyanide with a diluted nitric acid. The solution is evaporated, separated from potassium nitrate, the free acid neutralized with soda, and the solution concentrated. It crystallizes in dark red prisms which are readily soluble in water; it is a valuable reagent for the detection of sulphur, this element when in the form of an alkaline sulphide giving a characteristic purple blue coloration with the nitroprusside. The potassium salt may be prepared by adding potassium cyanide to ferrous sulphate solution, the brown precipitate so formed being then heated with potassium nitrite: 5 KNC + 2 FeSO 4 = 2 K 2 SO 4 + KFe2(NC)5, 2 KFe 2 (NC) 5 + 2 KNO 2 = 2 FeO + 2 K2Fe(NC)5 NO.
Other complex cyanides are known which may be regarded as derived from the acids H2X(CN)4, X=Ni, Pd, Pt; H 4 X(CN) 6, X= Fe, Co, Ru; H 3 X(CN) 6, X=Fe, Co, Rh; and H 2 R(CN) 6 (see Abegg, Anorganischen Chemie). Organic Cyanides or Nitriles. - Hydrocyanic acid forms two series of derivatives by the exchange of its hydrogen atom for alkyl or aryl groups; namely the nitriles, of type R CN, and the isonitriles, of type R NC. The latter compounds may be considered as derivatives of the as yet unknown isohydrocyanic acid HNC.
Nitriles These substances were first isolated in 1 834 by J. Pblouze (Ann., 1834, to, p. 249). They may be prepared by heating the alkyl iodides with potassium cyanide; by heating sulphuric acid esters with potassium cyanide; by distilling the acid-amides with phosphorus pentoxide; and by distilling amines (containing more than five atoms of carbon) with bromine and potash (A. W. Hofmann), for example C7H15CH2NH2-->C7H15CH2NBr2--C7H35CN.
In addition to these methods, the nitriles of the aromatic series may be prepared by distilling the aromatic acids with potassium sulphocyanide: C$_6$H 5 CO 2 H -{- [Kcns = Hcns + C6h5CO2k, C 6 H 5 Co 2 H -}- Hcns = C 6 H 5 Cn]] -fH 2 S + CO2; from the primary aromatic amines by converting them into diazonium salts, which are then decomposed by boiling with potassium cyanide and copper sulphate; by fusing the potassium salts of the sulphonic acids with potassium cyanide; by leading cyanogen gas into a boiling hydrocarbon in the presence of aluminium chloride (A. Desgrex, Bull. soc. chim., 18 95, (3) 1 3, p. 735); and from the syn-aldoximes by the action of acetyl chloride or acetic anhydride.
They are mostly colourless liquids which boil without decomposition, or solids of low melting point. The lower members of the series are somewhat soluble in water. They behave in most respects as unsaturated compounds; they combine with hydrogen to form amines; with water to form acidamides; with sulphuretted hydrogen to form thio-amides; with alcohols, in the presence of acids, to form imido-ethers R C(:NH) OR'; with ammonia and primary amines to form amidines R C(:NH) NH 2 i and with hydroxylamine to form amidoximes, R C(:NOH) NH 2. When heated with sodium the y frequently polymerize. Heated with acids or alkalis they hydrolyse to acids: RCN + HC1 + 2H 2 O = R COOH NH4C1. This reaction shows that the alkyl or aryl group is attached to the carbon atom in the nitrile.
Acetonitrile boils at 81.6° C., and is readily miscible with water. Propionitrile boils at 97° C.; it is somewhat easily soluble in water, but is thrown out of solution by calcium chloride. It was obtained by E. Frankl and C. C. Graham (Journ. Chem. Soc., 1880, 37, p. 740) by the action of cyanogen gas on zinc ethyl. Allyl cyanide boils at 119° C. Benzonitrile boils at 190.6° C. When solidified it melts at t7° C. It is easily soluble in alcohol and ether.
The Isonitriles (isocyanides or carbylamines) were first prepared in 1866 by A. Gautier (Ann., 1869, 151, p. 239) by the action of alkyl iodides on silver cyanide, and the distillation of the resulting compound with potassium cyanide in concentrated aqueous solution: RIR Ag(NC) 2 -)R NC+KAg(NC)2. They may also be obtained by distilling a primary amine with alcoholic potash and chloroform: R NH 2 -{- CHC1 3 3KHO=3KC1 -}- 3H 2 O -}- R NC (A. W. Hofmann, Ann., 1868, 146, p. 107). They are colourless liquids, readily soluble in alcohol and in ether, but insoluble in water. They possess an exceedingly unpleasant smell and are poisonous. They boil at temperatures somewhat lower than those of the corresponding nitriles; and are stable towards alkalis, but in the presence of mineral acids they readily hydrolyse, forming primary amines and formic acid: RNC+2H 2 O = RNH2+H2CO2. This reaction shows that the alkyl or aryl group is linked to the nitrogen atom. The carbon atom in the isonitriles is assumed by J. U. Nef to be divalent, since these substances readily form addition compounds, such addition taking place on the carbon atom, as is shown by the products of hydrolysis; for example with ethyl carbylamine: C 2 H 5 NC -FCH 3 COC1--> C 2 H 5 NC(OOCH 3)CI -->HCI -{- C2H5NH3 -fCH3CO CO2H.
This view was confirmed by J. Wade (Journ. Chem. Soc., 1902, 81, p. 1 59 6) who showed that the products obtained by the action of alkyl iodides on the isonitriles in alcoholic solution at 100° C. yield amine hydroidides and formic acid when hydrolysed. Such a reaction can only take place if the addition of the alkyl group takes place on the nitrogen atom of the isonitrile, from which it follows that the nitrogen atom must be trivalent and consequently the carbon atom divalent. The reactions may probably be represented as follows: C 2 H5NC+C 2 H51+4C2H5OH=C2H5NH2 HI+HCO2C2H5+2 (C2H5)2O, C 2 H 5 NC C2H5N(C2H5 I)C(+3C2H5OH) -> (C 2 H 5) 2 NH HI-{- H CO2C2H5 + (C2H5)2O.
The isonitriles dissolve silver cyanide readily, forming a soluble silver salt (cf. KNC). At 200° C. the isonitriles are converted into nitriles.
Constitution of Metallic Cyanides. - Considerable discussion has taken place as to the structure of the metallic cyanides, since potassium cyanide and silver cyanide react with alkyl iodides to form nitriles and isonitriles respectively, thus apparently pointing to the fact that these two compounds possess the formulae KCN and AgNC. The metallic cyanides are analogous to the alkyl isocyanides, since they form soluble double silver salts, and the fact that ethyl ferrocyanide on distillation yields ethyl isocyanide also points to their isocyanide structure. J. Wade (loc. cit.) explains
the formation of nitriles from potassium cyanide, and of isonitriles from silver cyanide by the assumption that unstable addition products are formed, the nature of which depends on the relative state of unsaturation of the carbon and nitrogen atoms under the varying conditions: KNC--KN :C(:C 2 H 5 I) --SKI +C2H5CN, AgNC->AgN(:C2H51)C---AgI-f-C2H5NC; that is, when the metal is highly electro-positive the carbon atom is the more unsaturated, the addition takes place on the carbon atom, and nitriles are produced. The same type of reaction occurs when the metal is relatively electro-positive to the added radical, for example, with ethyl isocyanide and acetyl chloride (see above); compare also AgNC --AgN(:Cl OOOH 3)C -->AgCl+CH,OOON. On the other hand, when there is but little electro-chemical difference between the radical of the cyanide and that of the reacting compound then the nitrogen atom is the more unsaturated element and. isonitriles are produced. This explanation also accounts for the formation of nitriles by the diazo reaction, thus: C 6 H 5 N 2 C1+CuNC-)CuN :C Cl N 2 C 6 H 5 ->CuCl-{ N :C N2 C6H5C6H5CN-{- N2.
Detection
The metallic cyanides may be detected by adding ferrous sulphate, ferric chloride, and hydrochloric acid to their solution, when a precipitate of Prussian blue is produced; if the original solution contains free acid it must be neutralized by caustic potash before the reagents are added. As an alternative test the cyanide may be decomposed by dilute hydrochloric acid, and the liberated hydrocyanic acid absorbed in a little yellow ammonium sulphide. The excess of reagent is removed by evaporation and a small quantity of a ferric salt added, when a deep red colour is produced. Silver nitrate gives a white precipitate with cyanides, soluble in excess of potassium cyanide. The amount of hydrocyanic acid in a solution may be determined by adding excess of caustic potash and a small quantity of an alkaline chloride, and running into the dilute solution standard silver nitrate until a faint permanent turbidity (of silver chloride) is produced, that is, until the reaction, 2KNC+AgNO 3 = KAg(NC) 2 - -KNO 3, is completed.
See R. Robine and M. Lengler, The Cyanide Industry, 1906 (Eng. trans. by J. A. Le Clerc); W. Bertelsmann, Die Technologie der Cyanverbindungen, 1906.
Pharmacology, Therapeutics and Toxicology of Hydrocyanic Acid
The pharmacopoeial preparations of this acid are a 2% solution, which is given in doses of from two to six minims, the tinctura chloroformi et morphinae cornposita, which contains a half-minim of this solution in each ten minims, and the aqua laurocerasi, which owes its virtues to the presence of this acid, and is of inconstant strength, besides being superfluous. The acid is also the active ingredient of the preparations of Virginian Prune, to which the same strictures apply.
The simple cyanides share the properties of the acid, except those of platinum and iron. With these exceptions, the simple cyanides are readily decomposed even by carbonic acid, free prussic acid being liberated. The double cyanides are innocuous. Hydrocyanic acid is a protoplasmic poison, directly lethal to all living tissues, whether in a plant or an animal. It is by no means the most powerful poison known, for such an alkaloid as pseud-aconitine, which is lethal in dose of about 1/200 of a grain, is some hundreds of times more toxic, but prussic acid is by far the most rapid poison known, a single inhalation of it producing absolutely instantaneous death. The acid is capable of passing through the unbroken skin, whereupon it instantly paralyses the sensory nerves. It is very rapidly absorbed from raw surfaces and may thereby cause fatal consequences. It is naturally an antiseptic.
The therapeutic applications of the drug are based entirely upon its anaesthetic or anodyne power. A lotion containing ten minims of the dilute acid to an ounce of water and glycerin will relieve itching due to any cause; and is useful in some forms of neuralgia. It must never be employed when the skin is abraded. The diluted acid is used internally to relieve vomiting or gastric pain. It is also added to cough mixtures, when the cough is of the dry, painful kind, which serves no purpose, as nothing is expectorated. Such a cough is relieved by the sedative action on the central nervous system.
Toxicology
Instantaneous death results from taking the pure acid. The diluted form, in toxic quantities, will cause symptoms usually within a few seconds. The patient is quite unconscious, the eyes are motionless, the pupils dilated, the skin cold and moist, the limbs relaxed, the pulse is slow and barely perceptible, the respirations very slow and convulsive. Post mortem, the body is livid, and the blood very dark. There may be an odour of prussic acid, but this soon disappears.
Treatment is only rarely of use, owing to the rapidity of the toxic action. The patient who survives half-an-hour will probably recover, as the volatile acid is rapidly excreted by the lungs. The drug kills by paralysing the nervous arrangements of the heart and respiration. The appropriate drug is therefore atropine, which stimulates the respiration and prevents the paralysis of the heart.
One-fiftieth of a grain must be immediately injected subcutaneously. The stomach must be washed out and large doses of emetics given as soon as possible. Every second is of consequence. Ammonia should be given by inhalation, and artificial respiration must never be forgotten, as by it the paralysed breathing may be compensated for and the poison excreted. The use of chemical antidotes, such as iron salts, is futile, as the drug has escaped into the blood from the stomach long before they can be administered.