Ammonia
{| width="300" border="1" cellpadding="2" cellspacing="0" align="right" style="margin-left: 1em; margin-right: 1em"\n|+
Properties\n|-\n! colspan="2" align=center bgcolor="#FFDEAD" |
General\n|-\n|Name || Ammonia\n|-\n|
Chemical formula ||
NH3\n|-\n|
Appearance || Colourless gas\n|-\n! colspan="2" align="center" bgcolor="#FFDEAD" |
Physical\n|-\n|
Formula weight || 17.0
amu\n|-\n|
Melting point || 195
K (-78
°C)\n|-\n|
Boiling point || 240
K (-33
°C)\n|-\n|
Density || 8.0 ×10
3 kg/
m³ (liquid)\n|-\n|
Solubility || 46
g in 100 g water\n|-\n! colspan="2" align="center" bgcolor="#FFDEAD" |
Thermochemistry\n|-\n|
ΔfH0gas || -45.9
kJ/
mol\n|-\n|Δ
fH
0liquid || -40.2 kJ/mol\n|-\n|Δ
fH
0solid || ? kJ/mol\n|-\n|
S0gas, 1 bar || 192.77 J/mol·K\n|-\n|S
0liquid, 1 bar || ? J/mol·K\n|-\n|S
0solid || ? J/mol·K\n|-\n! colspan="2" align="center" bgcolor="#FFDEAD" |
Safety\n|-\n|Ingestion || Dangerous. Symptoms include nausea & vomiting; damage to lips, mouth and esophagus.\n|-\n|Inhalation || Vapours are extremely irritating and corrosive.\n|-\n|Skin || Concentrated solutions may produce severe burns and necrosis.\n|-\n|Eyes || May cause permanent damage, even in small quantities.\n|-\n|More info ||
Hazardous Chemical Database\n|-\n! colspan="2" align="center" bgcolor="#FFDEAD" |
SI units were used where possible. Unless otherwise stated, standard conditions were used.Disclaimer and references
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Ammonia is a
chemical compound whose
molecule consists of one
atom of
nitrogen (N) and three atoms of
hydrogen (H) with the formula NH
3 and the structure:
The molecule is not flat, but has the shape of a flattened
tetrahedron known as
trigonal pyramidal. In solution it forms the positively charged
ammonium ion NH
4+ with the shape of a regular tetrahedron.
At
standard temperature and pressure, ammonia is a gas with a characteristic pungent smell. Its main uses are in the production of
fertilizers,
explosives and
polymers.
Ammonia is very well suited as a
refrigerant, since it liquefies readily under pressure, and was used in virtually all
refrigeration units prior to the advent of freons. Since the implication of freons as major
greenhouse gases during the
1990s, ammonia is again seeing increasing use as a refrigerant.
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 sulfate are found in volcanic districts; and crystals of ammonium
bicarbonate have been found in Patagonian guano. Ammonium salts also are found distributed through all fertile soil, in sea-water, and in most plant and animal liquids, and also in
urine.
Production
\nBecause of its many uses, ammonia is one of the most highly-produced inorganic chemicals. Before the start of WWI most ammonia was 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 unslaked lime (quicklime), the salt most generally used being the chloride (sal-ammoniac) thus
2NH4Cl + 2CaO 
CaCl2 + Ca(OH)2 + 2NH3.
It was also obtained by decomposing magnesium nitride (Mg3N2) with water,
Mg3N2 + 6H2O 3Mg(OH)2 + 2NH3.
Today the Haber process is the most important method for production of ammonia. The main advantage of the Haber process is that relatively cheap nitrogen and hydrogen gas are the primary feedstocks. They are reacted over an iron catalyst at high pressure (3000 psi or 20 MPa) and temperature (500 °C) to produce the ammonia.
Properties
\nAmmonia is a colourless gas possessing a characteristic pungent smell and a strongly alkaline reaction; it is lighter than air, its density being 0.589 times that of air. It is easily liquefied and the liquid boils at -33.7 °C, and solidifies at -75°C to a mass of white crystals. \nLiquid ammonia possesses strong ionizing powers, and solutions of salts in liquid ammonia have been much studied.
It is extremely soluble in water, one volume of water at 0°C and normal pressure absorbs 1148 volumes of ammonia. All the ammonia contained in an aqueous solution of the gas may be expelled by boiling. The aqueous solution of ammonia is very basic, and since it is a weak electrolyte, the solution will contain a small amount of ammonium hydroxide (NH4OH).
It does not support combustion, and it does not burn readily unless mixed with oxygen, when it burns with a pale yellowish-green flame. However it can form an explosive mixture with air.
Salts
\nOne of the most characteristic properties of ammonia is its power\nof combining directly with acids to form salts; thus with\nhydrochloric acid it forms ammonium chloride (sal-ammoniac); with\nnitric acid, ammonium nitrate, etc.\nIt 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 salts produced by the action of ammonia on acids are known as\nthe ammonium salts and all contain the compound radical ammonium\n(NH4+). Numerous attempts have been made to isolate\nthis radical, but so far none have been successful. By the\naddition of sodium amalgam to a concentrated solution of\nammonium chloride, the so-called ammonium amalgam is obtained as\na spongy mass which floats on the surface of the liquid; it\ndecomposes readily at ordinary temperatures into ammonia and\nhydrogen; it does not reduce silver and gold salts, a behaviour\nwhich distinguishes it from the amalgams of the alkali metals,\nand for this reason it is regarded by some chemists as being\nmerely mercury inflated by gaseous ammonia and hydrogen. M. le Blanc has shown, however, that the effect of ammonium amalgam on\nthe magnitude of polarization of a battery is comparable with\nthat of the amalgams of the alkali metals.
Ammonium bromide, NH4Br, can be prepared by the direct\naction of bromine on ammonia. It crystallizes in colourless\nprisms, possessing a saline taste; it sublimes on heating and is\neasily soluble in water. On exposure to air it gradually assumes\na yellow colour and becomes acid in its reaction.
Ammonium chloride, NH4Cl. (See sal-ammoniac.)
Ammonium fluoride, NH4F, may be obtained by\nneutralizing ammonia with hydrofluoric acid. It crystallizes in\nsmall prisms, having a sharp saline taste, and is exceedingly\nsoluble in water. It decomposes silicates on being heated with\nthem.
'\Ammonium iodide', NH4I, can be prepared by the action\nof hydriodic acid on ammonia. It is easily soluble in water,\nfrom which it crystallizes in cubes, and also in alcohol. It\ngradually turns yellow on standing in moist air, owing to\ndecomposition with liberation of iodine.
Ammonium chlorate, NH4ClO3, is obtained by\nneutralizing chloric acid with either ammonia or ammonium\ncarbonate, or by precipitating barium, strontium or\ncalcium chlorates with ammonium carbonate. It crystallizes in\nsmall needles, which are readily soluble in water, and on\nheating, decompose at about 102°C, with liberation of\nnitrogen, chlorine and oxygen. It is soluble in dilute aqueous\nalcohol, but insoluble in strong alcohol.
Ammonium carbonates. The commercial salt was formerly known as\nsal-volatile or salt of hartshorn and was formerly obtained by\nthe dry distillation of nitrogenous organic matter such as hair,\nhorn, decomposed urine, etc., but is now obtained by heating a\nmixture of sal-ammoniac, or ammonium sulfate and chalk, to\nredness in iron retorts, the vapours being condensed in leaden\nreceivers. The crude product is refined by sublimation, when it\nis obtained as a white fibrous mass, which consists of a mixture\nof hydrogen ammonium carbonate, NH4.HCO3,\nand ammonium carbamate, NH2COONH4, in\nmolecular proportions; on account of its possessing this\nconstitution it is sometimes called ammonium sesquicarbonate. It\npossesses a strong ammoniacal smell, and on digestion with\nalcohol the carbamate is dissolved and a residue of ammonium\nbicarbonate is left; a similar decomposition taking place when\nthe sesquicarbonate is exposed to air. Ammonia gas passed into a\nstrong aqueous solution of the sesquicarbonate converts it into\nnormal ammonium carbonate,\n(NH4)2CO3, which can be obtained\nin 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\nas shown above and also by passing carbon dioxide through a\nsolution of the normal compound, when it is deposited as a white\npowder, which has no smell and is only slightly soluble in water.\nThe aqueous solution of this salt liberates carbon dioxide on\nexposure to air or on heating, and becomes alkaline in reaction.\nThe aqueous solutions of all the carbonates when boiled undergo\ndecomposition with liberation of ammonia and of carbon dioxide:
NH4HCO3 NH3 + H2O + CO2
It is therefore occasionally used as baking powder, e.g. for gingerbread.
\nAmmonium nitrate, NH4NO3, is prepared by\nneutralizing nitric acid with ammonia, or ammonium carbonate, or\nby double decomposition between potassium nitrate and ammonium\nsulfate. It can be obtained in three different crystalline\nforms, the transition points of which are 35°C, 83°C and 125°C. It is easily soluble in water, a considerable\nlowering of temperature taking place during the operation; on\nthis account it is sometimes used in the preparation of freezing\nmixtures. On gentle heating, it is decomposed into water and\nnitrous oxide. P. E. M. Berthelot in 1883 showed that if\nammonium nitrate be rapidly heated the following reaction takes\nplace with explosive violence:--2NH4NO3 =\n4H2O + 2N2 + O2. In combination with gasoline it is a widely used explosive.
Ammonium nitrite, NH4NO2, is formed by\noxidizing ammonia with ozone or hydrogen peroxide; by\nprecipitating barium or lead nitrites with ammonium sulfate, or\nsilver nitrite with ammonium chloride. The precipitate is\nfiltered off and the solution concentrated. It forms colourless\ncrystals which are soluble in water and decompose on heating,\nwith the formation of nitrogen.
Ammonium phosphates. The normal phosphate,\n(NH4)3PO4,is obtained as a crystalline\npowder, on mixing concentrated solutions of ammonia and\nphosphoric acid, or on the addition of excess of ammonia to the\nacid phosphate (NH4)2HPO4. It is soluble\nin water, and the aqueous solution on boiling loses ammonia and\nthe acid phosphate NH4H2PO4 is\nformed. Diammonium hydrogen phosphate,\n(NH4)2HPO4, is formed by evaporating a\nsolution of phosphoric acid with excess of ammonia. It\ncrystallizes in large transparent prisms, which melt on heating\nand decompose, leaving a residue of metaphosphoric acid,\n(HPO3). Ammonium dihydrogen phosphate,\nNH4.H2PO4, is formed when a\nsolution of phosphoric acid is added to ammonia until the\nsolution is distinctly acid. It crystallizes in quadratic\nprisms.
Ammonium sodium hydrogen phosphate,\nNH4.NaHPO4.4H2O. (See microcosmic salt.)
Ammonium sulfate (NH4)2SO4 is\nprepared commercially from the ammoniacal liquor of gas-works and is purified by recrystallization. It\nforms large rhombic prisms, has a somewhat saline taste and is\neasily soluble in water. The aqueous solution on boiling loses\nsome ammonia and forms an acid sulfate. It is used largely as an\nartificial manure, and also for the preparation of other ammonium\nsalts.
Ammonium persulfate (NH4)2S2O8\nhas been prepared by H. Marshall (Jour. of Chem. Soc., 1891,\nlix. p. 777) by the method used for the preparation of the\ncorresponding potassium salt (see sulfur). It is very soluble in cold water, a large fall of temperature accompanying solution. It is a very strong oxidizing agent.
Ammonium sulfide, (NH4)2S, is obtained, in the form of micaceous crystals, by passing sulfuretted hydrogen\nmixed with a slight excess of ammonia through a well-cooled\nvessel; the hydrosulfide NH4.HS is formed at the same\ntime. It dissolves readily in water, but is probably partially\ndissociated in solution. The hydrosulfide NH4.HS can\nbe obtained as a white solid, by mixing well-cooled ammonia with\na slight excess of sulfuretted hydrogen. According to W. P.\nBloxam (Jour. of Chem. Soc., 1895, lxvii. p. 283), if\nsulfuretted hydrogen is passed into strong aqueous ammonia at\nordinary temperature, the compound\n(NH4)2S.2NH4HS is obtained,\nwhich, on cooling to 0°C and passing more sulfuretted\nhydrogen, forms the compound\n(NH4)2S.12NH4HS. An ice-cold\nsolution of this substance kept at 0°C and having\nsulfuretted hydrogen continually passed through it gives the\nhydrosulfide. Several complex polysulfides of ammonium have been\nisolated, for details of which see Bloxam's paper quoted above.\nCompounds are known which may be looked upon as derived from\nammonia by the replacement of its hydrogen by the sulfo-group\n(HSO3); thus potassium ammon-trisulfonate, \nN(SO3K)3.2H2O, is obtained as a\ncrystalline precipitate on the addition of excess of potassium\nsulfite to a solution of potassium nitrite, KNO2 +\n3K2SO3 + 2H2O =\nN(SO3K)3 + 4KHO. It can be recrystallized\nby solution in alkalies. On boiling with water, it is converted,\nfirst into the disulfonate NH(SO3K)2 thus,\nN(SO3K)3 + H2O =\nNH(SO3K)2 + KHSO4, and\nultimately into the monosulfonate NH2.SO3K.\nThe disulfonate is more readily obtained by moistening the\nnitrilosulfonate with dilute sulfuric acid and letting it stand\nfor twenty-four hours, after which it is recrystallized from\ndilute ammonia. It forms monosymmetric crystals which by boiling\nwith water yield amidosulfonic acid. (See also E. Divers, Jour.\nof Chem. Soc., 1892, lxi. p. 943.) Amidosulfonic acid\ncrystallizes in prisms, slightly soluble in water, and is a\nstable compound.
Other compounds
\nAmmonia\nfinds a wide application in organic chemistry as a synthetic\nreagent; it reacts with alkyl iodides to form amines, with\nesters to form acid amides, with halogen fatty acids to\nform amino acids; while it also combines with isocyanic esters to\nform alkyl ureas and with the mustard oils to form alkyl\nthioureas. 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, 2AgClCl.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\ntakes 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 +\n6Zn(C2H5)2 \n6ZnC2H5.I + 2NH3 +\n2N(C2H5)3.
The hydrogen in ammonia is capable of replacement by metals, thus\nmagnesium burns in the gas with the formation of magnesium nitride Mg3N2, and when the gas is passed\nover heated sodium or potassium, sodamide, NaNH2, and\npotassamide, KNH2, are formed.
Liquid ammonia as a solvent
Liquid ammonia is the best-known and most widely studied non-aqueous ionizing solvent. Its most conspicuous property is its ability to dissolve alkali metals to form highly coloured, electrically conducting solutions containing solvated electrons. Apart from these remarkable solutions, much of the chemistry in liquid ammonia can be classified by analogy with related reactions in aqueous solutions. Comparison of the physical proprties of NH3 with those of water shows that NH3 has the lower mp, bp, density, viscosity, dielectric constant and electrical conductivity; this is due at least in part to the weaker H bonding in NH3 and the fact that such bonding cannot form cross-linked networks since each NH3 molecule has only 1 lone-pair of electrons compared with 2 for each H2O molecule. The ionic self-dissociation constant of liquid NH3 at -50 degrees C is approx. 10-33 mol2l-2.
\nDetection
\nAmmonia 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\nby warming the salts with a caustic alkali or with quicklime, when the characteristic smell of ammonia will be at once\napparent. The amount of ammonia in ammonium salts can be estimated quantitatively by distillation of the salts with sodium\nor potassium hydroxide, the ammonia evolved being absorbed in a known volume of standard sulfuric acid and the excess of acid\nthen 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.
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Initial text from 1911 encyclopedia
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Category:Inorganic compounds\nCategory: Bases
