solution of the sesquicarbonate converts it into normal ammonium carbonate,
(NH4)2CO3, 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, NH4HCO3, 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:
NH4HCO3 --> NH3 + H2O + CO2
It is therefore occassionally used as baking powder, e.g. for gingerbread.
Ammonium nitrate, NH4NO3, is prepared by neutralizing nitric acid with
ammonia, or ammonium carbonate, or by double decomposition between potassium
nitrate and ammonium sulfate. 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:--2NH4NO3 = 4H2O + 2N2 + O2. In combination with gasoline it is a
widely used explosive.
Ammonium nitrite, NH4NO2, is formed by oxidizing ammonia with ozone or
hydrogen peroxide; by precipitating barium or lead nitrites with ammonium
sulfate, 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, (NH4)3PO4,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 (NH4)2HPO4. It is soluble in water, and the aqueous solution on
boiling loses ammonia and the acid phosphate NH4H2PO4 is formed. Diammonium
hydrogen phosphate, (NH4)2HPO4, 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, (HPO3). Ammonium dihydrogen phosphate, NH4.H2PO4, 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, NH4.NaHPO4.4H2O. (See microcosmic salt.)
Ammonium sulfate (NH4)2SO4 is prepared commercially from the ammoniacal
liquor of gas-works 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
sulfate. It is used largely as an artificial manure, and also for the
preparation of other ammonium salts.
Ammonium persulfate (NH4)2S2O8 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 sulfur). It is very soluble in cold water,
a large fall of temperature accompanying solution. It is a mighty strong
oxidizing agent.
Ammonium sulfide, (NH4)2S, is obtained, in the form of micaceous crystals,
by passing sulfuretted hydrogen mixed with a slight excess of ammonia
through a well-cooled vessel; the hydrosulfide NH4.HS is formed at the same
time. It dissolves readily in water, but is probably partially dissociated
in solution. The hydrosulfide NH4.HS can be obtained as a white solid, by
mixing well-cooled ammonia with a slight excess of sulfuretted hydrogen.
According to W. P. Bloxam (Jour. of Chem. Soc., 1895, lxvii. p. 283), if
sulfuretted hydrogen is passed into strong aqueous ammonia at ordinary
temperature, the compound (NH4)2S.2NH4HS is obtained, which, on cooling to
0¡C and passing more sulfuretted hydrogen, forms the compound
(NH4)2S.12NH4HS. An ice-cold solution of this substance kept at 0¡C and
having sulfuretted hydrogen continually passed through it gives the
hydrosulfide. Several complex polysulfides 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 sulfo-group (HSO3); thus potassium ammon-trisulfonate,
N(SO3K)3.2H2O, is obtained as a crystalline precipitate on the addition of
excess of potassium sulfite to a solution of potassium nitrite, KNO2 +
3K2SO3 + 2H2O = N(SO3K)3 + 4KHO. It can be recrystallized by solution in
alkalies. On boiling with water, it is converted, first into the disulfonate
NH(SO3K)2 thus, N(SO3K)3 + H2O = NH(SO3K)2 + KHSO4, and ultimately into the
monosulfonate NH2.SO3K. The disulfonate is more readily obtained by
moistening the nitrilosulfonate with dilute sulfuric 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
amidosulfonic acid. (See also E. Divers, Jour. of Chem. Soc., 1892, lxi. p.
943.) Amidosulfonic acid crystallizes in prisms, slightly soluble in water,
and is a stable compound.
Other compounds
Ammonia finds a wide application in organic chemistry as a synthetic
reagent; it reacts with alkyl iodides to form amines, with esters to form
acid amides, 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.
Ammonia gas has the power of combining with many substances, particularly
with metallic halides; thus with calcium chloride it forms the compound
CaCl2.8NH3, and consequently calcium chloride cannot be used for drying the
gas. With silver chloride it forms two compounds -- one, AgCl.3NH3 at
temperatures below 15¡C; the other, 2AgCl.3NH3 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
Michael Faraday was first able to liquefy ammonia. It can be shown by
Isambert's results that the compound AgCl.3NH3 cannot be formed above 20¡C,
by the action of ammonia on silver chloride at atmospheric pressure; whilst
2AgCl.3NH3, 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 trichloride NCl3 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 NI3, NHI2,
NH2I, N2H3I3, &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 N2H3I3, by the addition of excess of standard
sodium sulfite solution, in the dark, and subsequent titration of the excess
of the sulfite 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
NH3.NI3 to the compound, and explains the decomposition as taking place,
2NH3.NI3 + 6Zn(C2H5)2 = 6ZnC2H5.I + 2NH3 + 2N(C2H5)3.
The hydrogen in ammonia is capable of replacement by metals, thus magnesium
burns in the gas with the formation of magnesium nitride Mg3N2, and when the
gas is passed over heated sodium or potassium, sodamide, NaNH2, and
potassamide, KNH2, are formed.
Detection
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
sulfuric 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, (NH4)2PtCl6.
History
Salts of ammonia have been known from very early times; thus the term
Hammoniacus sal appears in the writings of Pliny, although it is not known
whether the term is identical with the more modern sal-ammoniac.
In the form of sal-ammoniac, ammonia was known, however, to the alchemists
as early as the 13th century, being mentioned by Albertus Magnus, while 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 Karl Wilhelm Scheele showed that it contained
nitrogen, and C. L. Berthollet, in about 1785, ascertained its composition.
The Haber process to produce ammonia from the nitrogen contained in the air
was developed by Fritz Haber and Carl Bosch in 1909 and patented in 1910. It
was first used on an industrial scale by the Germans during WWI. The ammonia
was used to produce explosives to sustain their war effort.
Etymology of "Ammonia"
In classical times, sal ammoniac was discovered by accident through burning
the dung of camels in the temple of Jupiter Ammon at Siwa oasis in Libya.
"Ammonia" is a genus name in the Foraminifera (marine planktonic protozoa
with a calcium carbonate shell, whose remains have contributed to limestone
and chalk deposits), and "ammonites" are an extinct group of cephalopod
whose fossil shells are abundant from the Paleozoic. In both cases, the
shell is formed of a series of chambers, arranged in a spiral, and the name
is given for the "Horn of Ammon", the ram's horns that the god by whose
temple the ammoniacal camel dung was to be found (see above) was supposed to
have had