# 1911 Encyclopædia Britannica/Ammonia

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​AMMONIA (NH$_3$). Salts of ammonia have been known from very early times; thus the term Hammoniacus sal appears in the writings of Pliny (Nat. Hist. xxxi. 39), although it is not known whether the term is identical with the more modern sal-ammoniac (q.v.). In the form of sal-ammoniac, ammonia was known, however, to the alchemists as early as the 13th century, being mentioned by Albertus Magnus, whilst in the 15th century Basil Valentine showed that ammonia could be obtained by the action of alkalies on sal-ammoniac. At a later period when sal-ammoniac was obtained by distilling the hoofs and horns of oxen, and neutralizing the resulting carbonate with hydrochloric acid, the name spirits of hartshorn was applied to ammonia. Gaseous ammonia was first isolated by J. Priestley in 1774 and was termed by him “alkaline air.” In 1777 K. W. Scheele showed that it contained nitrogen, and C. L. Berthollet, in about 1785, ascertained its composition.
Ammonia is found in small quantities as the carbonate in the atmosphere, being produced from the putrefaction of nitrogenous animal and vegetable matter; ammonium salts are also found in small quantities in rain-water, whilst ammonium chloride (sal-ammoniac) and ammonium sulphate are found in volcanic districts; and crystals of ammonium bicarbonate have been found in Patagonian guano. Ammonium salts too are found distributed through all fertile soil, in sea-water, and in most plant and animal liquids, and also in urine.
Ammonia can be synthesized by submitting a mixture of nitrogen and hydrogen to the action of the silent electric discharge, the combination, however, being very imperfect. It is obtained by the dry distillation of nitrogenous vegetable and animal products; by the reduction of nitrous acid and nitrites with nascent hydrogen; and also by the decomposition of ammonium salts by alkaline hydroxides or by slaked lime, the salt most generally used being the chloride (sal-ammoniac, q.v.) thus 2NH$_4$Cl + Ca(OH)$_2$ = CaCl$_2$ + 2H$_2$O + 2NH$_3$. It also results on decomposing magnesium nitride (Mg$_3$N$_2$) with water, Mg$_3$N$_2$ + 6H$_2$O = 3Mg(OH)$_2$ + 2NH$_3$. Large quantities of ammonia and ammonium salts are now obtained from the ammoniacal liquor of gas-works.
Ammonia is a colourless gas possessing a characteristic pungent smell and a strongly alkaline reaction; it is lighter than air, its specific gravity being 0·589 (air = 1). It is easily liquefied and the liquid boils at −33·7° C., and solidifies at −75° C. to a mass of white crystals. It is extremely soluble in water, one volume of water at 0° C. and normal pressure absorbs 1148 volumes of ammonia (Roscoe and W. Dittmar). All the ammonia contained in an aqueous solution of the gas may be expelled by boiling. It does not support combustion; and it does not burn readily unless mixed with oxygen, when it burns with a pale yellowish-green flame. Ammonia gas has the power of combining with many substances, particularly with metallic halides; thus with calcium chloride it forms the compound CaCl$_2$·8NH$_3$, and consequently calcium chloride cannot be used for drying the gas. With silver chloride it forms two compounds (F. Isambert, Comptes rendus, 1868, lxvi. p. 1259)—one, AgCl·3NH$_3$ at temperatures below 15° C.; the other, 2AgCl·3NH$_3$ at temperatures above 20° C. On heating these substances, ammonia is liberated and the metallic chloride remains. It was by the use of silver chloride ammonia compounds that in 1823 M. Faraday was first able to liquefy ammonia. It can be shown by Isambert’s results that the compound AgCl·3NH$_3$ cannot be formed above 20° C., by the action of ammonia on silver chloride at atmospheric pressure; whilst 2AgCl·3NH$_3$; under similar conditions, cannot be formed above about 68° C. Liquid ammonia is used for the artificial preparation of ice. It readily dissolves sodium and potassium, giving in each case a dark blue solution. At a red heat ammonia is easily decomposed into its constituent elements, a similar decomposition being brought about by the passage of electric sparks through the gas. Chlorine takes fire when passed into ammonia, nitrogen and hydrochloric acid being formed, and unless the ammonia be present in excess, the highly explosive nitrogen chloride NCl$_3$ is also produced. With iodine it reacts to form nitrogen iodide. This compound was discovered in 1812 by Bernard Courtois, and was originally supposed to contain nitrogen and iodine only, but in 1840 R. F. Marchand showed that it contained hydrogen, whilst R. Bunsen showed that no oxygen was present. As regards its constitution, it has been given at different times the formulae NI$_3$, NHI$_2$, NH$_2$I, N$_2$H$_3$I$_3$, &c., these varying results being due to the impurities in the substance, owing to the different investigators working under unsuitable conditions, and also to the decomposing action of light. F. D. Chattaway determined its composition as N$_2$H$_3$I$_3$, by the addition of excess of standard sodium sulphite solution, in the dark, and subsequent titration of the excess of the sulphite with standard iodine. The constitution has been definitely determined by O. Silberrad (Jour. of Chem. Soc., 1905, lxxxvii. p. 55) by the interaction of nitrogen iodide with zinc ethyl, the products of the reaction being triethylamine and ammonia; the ammonia liberated was absorbed in hydrochloric acid, and 95% of the theoretical amount of the ammonium chloride was obtained. On these grounds O. Silberrad assigns the formula NH$_3$·NI$_3$ to the compound ​and explains the decomposition as taking place, 2NH$_3$·NI$_3$ + 6Zn(C$_2$H$_5$)$_2$ = 6ZnC$_2$H$_5$·I + 2NH$_3$ + 2N(C$_2$H$_5$)$_3$. The hydrogen in ammonia is capable of replacement by metals, thus magnesium burns in the gas with the formation of magnesium nitride Mg$_3$N$_2$, and when the gas is passed over heated sodium or potassium, sodamide, NaNH$_2$, and potassamide, KNH$_2$, are formed.
One of the most characteristic properties of ammonia is its power of combining directly with acids to form salts; thus with hydrochloric acid it forms ammonium chloride (sal-ammoniac); with nitric acid, ammonium nitrate, &c. It is to be noted that H. B. Baker (Journal of Chem. Soc., 1894, lxv. p. 612) has shown that perfectly dry ammonia will not combine with perfectly dry hydrochloric acid, moisture being necessary to bring about the reaction. The aqueous solution of ammonia is very basic in its reactions, and since it is a weak electrolyte, one must assume the solution to contain a certain amount of ammonium hydroxide NH$_4$OH, although it is probably chiefly composed of a solution of ammonia in water. (On the constitution of aqueous ammonia solutions see also Carl Frenzel, Zeit für angew. Chemie, xxxii. 3, p. 319.) Ammonia finds a wide application in organic chemistry as a synthetic reagent; it reacts with alkyl iodides to form amines (q.v.), with esters to form acid amides (q.v.), with halogen fatty acids to form amino-acids; while it also combines with isocyanic esters to form alkyl ureas and with the mustard oils to form alkyl thioureas. Aldehydes also combine directly with ammonia.
Liquid ammonia possesses strong ionizing powers, and solutions of salts in liquid ammonia have been much studied. For details see E. C. Franklin and C. A. Kraus, Amer. Chem. Jour., 1899, xxi. p. 8; 1900, xxiv. p. 83; 1902, xxviii. p. 277; also Carl Frenzel, Zeits für Elektrochemie, 1900, vi. p. 477.
The salts produced by the action of ammonia on acids are known as the ammonium salts and all contain the compound radical ammonium (NH$_4$). Numerous attempts have been made to isolate this radical, but so far none have been successful. By the addition of sodium amalgam to a concentrated solution of ammonium chloride, the so-called ammonium amalgam is obtained as a spongy mass which floats on the surface of the liquid; it decomposes readily at ordinary temperatures into ammonia and hydrogen; it does not reduce silver and gold salts, a behaviour which distinguishes it from the amalgams of the alkali metals, and for this reason it is regarded by some chemists as being merely mercury inflated by gaseous ammonia and hydrogen. M. le Blanc has shown, however, that the effect of ammonium amalgam on the magnitude of polarization of a battery is comparable with that of the amalgams of the alkali metals.
Many of the ammonium salts are made from the ammoniacal liquor of gas-works, by heating it with milk of lime and then absorbing the gas so liberated in a suitable acid. (See Gas: Manufacture.)
Ammonium bromide, NH$_4$Br, can be prepared by the direct action of bromine on ammonia. It crystallizes in colourless prisms, possessing a saline taste; it sublimes on heating and is easily soluble in water. On exposure to air it gradually assumes a yellow colour and becomes acid in its reaction.
Ammonium chloride, NH$_4$Cl. (See Sal Ammoniac.)
Ammonium fluoride, NH$_4$F, may be obtained by neutralizing ammonia with hydrofluoric acid. It crystallizes in small prisms, having a sharp saline taste, and is exceedingly soluble in water. It decomposes silicates on being heated with them.
Ammonium iodide, NH$_4$I, can be prepared by the action of hydriodic acid on ammonia. It is easily soluble in water, from which it crystallizes in cubes, and also in alcohol. It gradually turns yellow on standing in moist air, owing to decomposition with liberation of iodine.
Ammonium chlorate, NH$_4$ClO$_3$, is obtained by neutralizing chloric acid with either ammonia or ammonium carbonate, or by precipitating barium, strontium or calcium chlorates with ammonium carbonate. It crystallizes in small needles, which are readily soluble in water, and on heating, decompose at about 102° C., with liberation of nitrogen, chlorine and oxygen. It is soluble in dilute aqueous alcohol, but insoluble in strong alcohol.
Ammonium carbonates. The commercial salt is known as salvolatile or salt of hartshorn and was formerly obtained by the dry distillation of nitrogenous organic matter such as hair, horn, decomposed urine, &c., but is now obtained by heating a mixture of sal-ammoniac, or ammonium sulphate and chalk, to redness in iron retorts, the vapours being condensed in leaden receivers. The crude product is refined by sublimation, when it is obtained as a white fibrous mass, which consists of a mixture of hydrogen ammonium carbonate, NH$_4$·HCO$_3$, and ammonium carbamate, NH$_2$COONH$_4$, in molecular proportions; on account of its possessing this constitution it is sometimes called ammonium sesquicarbonate. It possesses a strong ammoniacal smell, and on digestion with alcohol the carbamate is dissolved and a residue of ammonium bicarbonate is left; a similar decomposition taking place when the sesquicarbonate is exposed to air. Ammonia gas passed into a strong aqueous solution of the sesquicarbonate converts it into normal ammonium carbonate, (NH$_4$)$_2$CO$_3$, which can be obtained in the crystalline condition from a solution prepared at about 30° C. This compound on exposure to air gives off ammonia and passes back to ammonium bicarbonate.
Ammonium bicarbonate, NH$_4$·HCO$_3$, is formed as shown above and also by passing carbon dioxide through a solution of the normal compound, when it is deposited as a white powder, which has no smell and is only slightly soluble in water. The aqueous solution of this salt liberates carbon dioxide on exposure to air or on heating, and becomes alkaline in reaction. The aqueous solutions of all the carbonates when boiled undergo decomposition with liberation of ammonia and of carbon dioxide.
Ammonium nitrate, NH$_4$NO$_3$, is prepared by neutralizing nitric acid with ammonia, or ammonium carbonate, or by double decomposition between potassium nitrate and ammonium sulphate. It can be obtained in three different crystalline forms, the transition points of which are 35° C., 83° C. and 125° C. It is easily soluble in water, a considerable lowering of temperature taking place during the operation; on this account it is sometimes used in the preparation of freezing mixtures. On gentle heating, it is decomposed into water and nitrous oxide. P. E. M. Berthelot in 1883 showed that if ammonium nitrate be rapidly heated the following reaction takes place with explosive violence:—2NH$_4$NO$_3$ = 4H$_2$O + 2N$_2$ + O$_2$.
Ammonium nitrite, NH$_4$NO$_2$, is formed by oxidizing ammonia with ozone or hydrogen peroxide; by precipitating barium or lead nitrites with ammonium sulphate, or silver nitrite with ammonium chloride. The precipitate is filtered off and the solution concentrated. It forms colourless crystals which are soluble in water and decompose on heating, with the formation of nitrogen.
Ammonium phosphates. The normal phosphate, (NH$_4$)$_3$PO$_4$, is obtained as a crystalline powder, on mixing concentrated solutions of ammonia and phosphoric acid, or on the addition of excess of ammonia to the acid phosphate (NH$_4$)$_2$HPO$_4$. It is soluble in water, and the aqueous solution on boiling loses ammonia and the acid phosphate NH$_4$H$_2$PO$_4$ is formed. Diammonium hydrogen phosphate, (NH$_4$)$_2$HPO$_4$, is formed by evaporating a solution of phosphoric acid with excess of ammonia. It crystallizes in large transparent prisms, which melt on heating and decompose, leaving a residue of metaphosphoric acid, (HPO$_3$). Ammonium dihydrogen phosphate, NH$_4$·H$_2$PO$_4$, is formed when a solution of phosphoric acid is added to ammonia until the solution is distinctly acid. It crystallizes in quadratic prisms.
Ammonium sodium hydrogen phosphate, NH$_4$·NaHPO$_4$·4H$_2$O. (See Microcosmic Salt.)
Ammonium sulphate (NH$_4$)$_2$SO$_4$ is prepared commercially from the ammoniacal liquor of gas-works (see Gas: Manufacture) and is purified by recrystallization. It forms large rhombic prisms, has a somewhat saline taste and is easily soluble in water. The aqueous solution on boiling loses some ammonia and forms an acid sulphate. It is used largely as an artificial manure, and also for the preparation of other ammonium salts.
Ammonium persulphate (NH$_4$)$_2$S$_2$O$_8$ has been prepared by H. Marshall (Jour. of Chem. Soc., 1891, lix. p. 777) by the method used for the preparation of the corresponding potassium salt ​(see Sulphur). Pure specimens are difficult to obtain. It is very soluble in cold water, a large fall of temperature accompanying solution.
Ammonium sulphide, (NH$_4$)$_2$S, is obtained, in the form of micaceous crystals, by passing sulphuretted hydrogen mixed with a slight excess of ammonia through a well-cooled vessel; the hydrosulphide NH$_4$·HS is formed at the same time. It dissolves readily in water, but is probably partially dissociated in solution. The hydrosulphide NH$_4$·HS can be obtained as a white solid, by mixing well-cooled ammonia with a slight excess of sulphuretted hydrogen. According to W. P. Bloxam (Jour. of Chem. Soc., 1895, lxvii. p. 283), if sulphuretted hydrogen is passed into strong aqueous ammonia at ordinary temperature, the compound (NH$_4$)$_2$S·2NH$_4$HS is obtained, which, on cooling to 0° C. and passing more sulphuretted hydrogen, forms the compound (NH$_4$)$_2$S·12NH$_4$HS. An ice-cold solution of this substance kept at 0° C. and having sulphuretted hydrogen continually passed through it gives the hydrosulphide. Several complex polysulphides of ammonium have been isolated, for details of which see Bloxam’s paper quoted above. Compounds are known which may be looked upon as derived from ammonia by the replacement of its hydrogen by the sulpho-group (HSO$_3$); thus potassium ammon-trisulphonate, N(SO$_3$K)$_3$·2H$_2$O, is obtained as a crystalline precipitate on the addition of excess of potassium sulphite to a solution of potassium nitrite, KNO$_2$ + 3K$_2$SO$_3$ + 2H$_2$O = N(SO$_3$K)$_3$ + 4KHO. It can be recrystallized by solution in alkalies. On boiling with water, it is converted, first into the disulphonate NH(SO$_3$K)$_2$ thus, N(SO$_3$K)$_3$ + H$_2$O = NH(SO$_3$K)$_2$ + KHSO$_4$, and ultimately into the monosulphonate NH$_2$·SO$_3$K. The disulphonate is more readily obtained by moistening the nitrilosulphonate with dilute sulphuric acid and letting it stand for twenty-four hours, after which it is recrystallized from dilute ammonia. It forms monosymmetric crystals which by boiling with water yield amidosulphonic acid. (See also E. Divers, Jour. of Chem. Soc., 1892, lxi. p. 943.) Amidosulphoiric acid crystallizes in prisms, slightly soluble in water, and is a stable compound.
Ammonia and ammonium salts can be readily detected, in very minute traces, by the addition of Nessler’s solution, which gives a distinct yellow coloration in the presence of the least trace of ammonia or ammonium salts. Larger quantities can be detected by warming the salts with a caustic alkali or with quicklime, when the characteristic smell of ammonia will be at once apparent. The amount of ammonia in ammonium salts can be estimated quantitatively by distillation of the salts with sodium or potassium hydroxide, the ammonia evolved being absorbed in a known volume of standard sulphuric acid and the excess of acid then determined volumetrically; or the ammonia may be absorbed in hydrochloric acid and the ammonium chloride so formed precipitated as ammonium chlorplatinate, (NH$_4$)$_2$PtCl$_6$.