AMMONIA and its salts have been known from very early times, sal-ammoniac being mentioned by Pliny. By the action of alkalis on this, Basil Valentine obtained free ammonia in the 15th century; later, ammonia was obtained by distilling the hoofs and horns of oxen, and was therefore called "spirits of hartshorn." J. Priestley was the first to isolate gaseous ammonia (1774), which he called "alkaline air." In 1777 K. W. Scheele showed it to contain nitrogen, and shortly afterwards C. L. Berthollet ascertained its composition to be 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 rainwater, whilst the chloride and sulphate are found in volcanic districts. Fertile soils, seawater, and plant and animal liquids (such as urine) also contain ammonium salts.
Ammonia is obtained by the dry distillation of animal and vegetable products, and also by the decomposition of its salts (usually the chloride or sulphate) by alkaline hydroxides or slaked lime: It may be obtained by the action of water on, e.g., magnesium nitride. Large quantities of ammonia and its salts (usually the sulphate) are obtained from the ammoniacal liquor of the gas works. Such ammonia is, however, difficult to purify from the pyridine which it contains, and the purest ammonia is now obtained synthetically.
The reversible reaction evolves heat when it produces ammonia, and therefore it becomes less favourable the higher the temperature ; but as it also takes place with a decrease in volume, it is favoured by an increase of pressure. In practice a compromise is effected, for, although at 45o° C. it would only be possible to effect combination of o.24% of the gases (if present in the theoretical proportion) under atmospheric pressure, and at C. only o.o8%, yet at the lower temperature reaction is so slow as to be unprofitable in working. The higher temperature is therefore chosen, and under 200 atmospheres pressure (as used by Haber) the equilibrium concentration is 12% and under i,000 atmospheres (as used by Claude) it is 4o%. The nitrogen (often from liquid air) and hydrogen (often from purified "water gas") are dried, compressed, and circulated at 55o° C. over a catalyst, the best being a very pure iron, obtained by reducing the oxide in a current of hydrogen at this temperature. The ammonia pro duced is either passed through a refrigerator and collected as a liquid, or absorbed in water.
Another industrial process of some importance in the produc tion of ammonia is that whereby calcium cyanamide, is heated under 3 atmospheres pressure with water in order to liberate traces of acetylene, and then under II atmos pheres at 18o° C. with a current of superheated steam, the result ing ammonia being blown over and collected. For fuller details of these processes see NITROGEN, FIXATION OF.
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= I). It is easily liquefied and the liquid boils at —33.7° C., and solidifies at C. to a mass of white crystals. It is extremely soluble in water, one volume at o° C. and normal pressure absorbing 1,300 volumes of ammonia, and a saturated solution at 15° C. contains 36%. It f cvms two hydrates with water, and both of which melt at —79° C., i.e., only one degree below the melting point of pure ammonia. All the ammonia contained in an aqueous solution of the gas may be expelled by boiling. It does not support com bustion and it does not burn readily unless mixed with oxygen, in which case it burns with a pale greenish flame. Ammonia gas has the power of combining with many substances, particularly with metallic halides ; thus, with calcium chloride it forms the com pound and consequently calcium chloride cannot be used for drying the gas. With silver chloride it forms two at temperatures below 15° C., and above 20° C. When heated, these substances liberate ammonia and the silver chloride remains; by this method M. Faraday was able to liquefy ammonia for the first time (1823). Ammonia is decomposed into its elements at a red heat or by the passage of electric sparks. Chlorine takes fire when passed into strong ammonia solutions, nitrogen and hydrochloric acid (or ammonium chloride) being formed, and unless the ammonia is present in excess the highly explosive nitrogen chloride is also produced. With iodine it reacts to form nitrogen iodide, which was originally supposed to contain nitrogen and iodine only; later, however, it was found to contain hydrogen also. F. D. Chattaway showed it to be and 0. Silberrad further showed its constitution to be The hydrogen in ammonia is capable of replacement by metals, thus magnesium burns in the gas with formation of the nitride, and when the gas is passed over heated sodium or potassium, sodamide, or potassamide, is formed.
Liquid ammonia is used for the artificial preparation of ice. It shows a slight but definite electrical conductivity, though it is not certain what ions are responsible for this, and salts dissolved in it are fairly strongly dissociated into ions, but not nearly so extensively as in aqueous solutions. Moreover, certain metals dissolve in it to give very deep blue solutions which may possi bly contain, e.g., etc. The works of E. C. Franklin (Amer. Cliem. J., 1899 et seq.), C. A. Kraus (e.g., J. Amer. Chem. Soc. 1922, 44, P. 5925, P. 749) and C. Frenzel should be consulted for fuller details as to solutions in Liquid ammonia.
One of the most characteristic properties of ammonia is its power of combining directly with acids to form salts ; thus with hydrochloric acid it gives ammonium chloride (sal-ammoniac), and with nitric acid ammonium nitrate, etc. It is of great interest, however, that ammonia and hydrochloric acid will • not react if they are both perfectly dry (H. B. Baker; see DRYNESS, CHEMI CAL) ; moreover, perfectly dry ammonium chloride does not exhibit the reverse phenomenon of dissociation into its com ponents, which is so characteristic of ammonium salts.
The aqueous solution of ammonia is very basic in its reactions, but it is not very definite as to whether this is due to an ammo nium hydroxide, NH2OH, to free ammonia as such, or to ammonia allied in some way with water, for there is doubtless an equilib rium involved: According to A. Hantzsch and T. S. Moore, the apparent "strength" of ammonia as a base is very low, for only about 1% of the total ammonia is present in the form of ions in decinormal (i.e., o.17%) aqueous solution, but as these may be derived from a relatively small proportion of it is possible that the latter molecule is fairly strongly ionized. When boiled, solutions of ammonium salts tend to lose ammonia, and if the acid is "weak" both it and the ammonia may be boiled away completely, but as most of the ammonia may be present in a form other than this again is no evidence that the latter is a weak base, especially as the above equilibria would tend to replace the volatile at the expense of the From a consideration of the ease of dissociation of the double sulphates CuSO,,M_. S0,6H,0 (where M = NH,,K,Rb or Cs), R. M. Caven places ammonium hydroxide as a base between rubidium and caesium hydroxides in strength.
Numerous attempts have been made to isolate the radical NH,, but the most that has been achieved is to obtain evidence as to its existence in an amalgam or in solution. The addition of sodium amalgam to a cold concentrated solution of ammonium chloride gives a spongy mass which resembles the amalgams of the alkali metals in some respects but not in others ; it readily decomposes into ammonia and hydrogen. H. H. Schlubach and F. Ballouf added very cold ammonium chloride to a solution of potassium in liquid ammonia at — 7o° C, and, as only one-third of the theoreti cal amount of hydrogen was evolved and solubility would not account for more than a trace of the rest, it is assumed that the free ammonium, remains in solution; the remainder of the hydrogen is given off at —4o° C.
The precise constitution of the ammonium radical or ion has been a matter of much speculation, but W. H. Mills and E. H. Warren have now shown that the nitrogen atom is situated at the centre of a regular tetrahedron with the hydrogen atoms at the four corners, and this has been confirmed by Mills, J. D. Parkin and W. J. D. Ward (see STEREOCHEMISTRY for further details) . The ammonium ion and the methane molecule therefore have similar structures, as might be anticipated on the electronic theory, for, if they are written as (I) and (2) respectively (where, in each case, the crosses represent electrons originally belonging to the N and C atoms, and the dots those belonging to the H atoms), it is seen that the nitrogen atom shares 8 electrons equally with hydrogen atoms and has lost one of its five electrons, thereby becoming positively charged, whereas the carbon atom has a half share in its own four electrons and in four others, thereby remain ing neutral.
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 amides (q.v.), and with halogen fatty acids to form amino-acids ; it combines with isocyanic esters to give alkyl-ureas, and with the mustard oils to give alkyl-thioureas. Aldehydes also combine directly with ammonia.
The alkyl-amines, which may be regarded as derived from ammonia by successive replacement of hydrogen atoms by alkyl groups, are not very strong bases, but the trialkyl-amines react with alkyl iodides to give tetra-alkyl-ammonium iodides, NR3+ RI = NR,I, which resemble the alkali-metal iodides, and this re semblance between alkali metals and the tetra-alkyl-ammonium group is still more pronounced in the corresponding hydroxides (obtained by the action of moist silver oxide on the iodide of the base), for these are strongly caustic bases which absorb carbon dioxide and generally resemble potassium hydroxide. The relative stability of the quaternary ammonium grouping is further illus trated by the fact that the corresponding amalgams are much more stable than that of ammonium. H. N. McCoy and W. C. Moore obtained a tetramethylammonium amalgam which did not decompose below Io° C. Moreover, Schlubach and Ballouf, using a method analogous to that described in the case of the ammo nium radical, found evidence for the existence of a tetraethylam monium which was stable in liquid ammonia at the ordinary tem perature if kept in an atmosphere of nitrogen. When all four alkyl groups of a quaternary ammonium salt are different, the compound can be resolved into optical isomerides, and W. J. Pope and S. J. Peachey resolved benzylphenylallylmethylammo nium salts into dextro- and laevo-rotatory forms (see STEREO