| General | |
|---|---|
| Name | Ammonia |
| Chemical formula | NH3 |
| Appearance | Colourless gas |
| Physical | |
| Formula weight | 17.0 amu |
| Melting point | 195 K (-78 °C) |
| Boiling point | 240 K (-33 °C) |
| Density | 8.0 ×103 kg/m3 (liquid) |
| Solubility | 46 g in 100g water |
| Thermochemistry | |
| ΔfH0gas | -45.9 kJ/mol |
| ΔfH0liquid | -40.2 kJ/mol |
| ΔfH0solid | ? kJ/mol |
| S0gas, 1 bar | 192.77 J/mol·K |
| S0liquid, 1 bar | ? J/mol·K |
| S0solid | ? J/mol·K |
| Safety | |
| Ingestion | Dangerous. Symptoms include nausea & vomiting; damage to lips, mouth and esophagus. |
| Inhalation | Vapours are extremely irritating and corrosive. |
| Skin | Concentrated solutions may produce severe burns and necrosis. |
| Eyes | May cause permanent damage, even in small quantities. |
| More info | Hazardous Chemical Database (http://ull.chemistry.uakron.edu/erd/chemicals/7/6014.html) |
| SI units were used where possible. Unless otherwise stated, standard conditions were used. | |
The molecule is not flat, instead it has the shape of a tetrahedron with one empty corner. In solution it forms the positively charged ammonium ion NH4+ with a hydrogen atom on all four corners of the 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 is readily liquified 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.
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| Production |
Because of its many uses, ammonia is one of the most highly-produced inorganic chemicals. Prior to 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 |
Ammonia 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. Liquid 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 in its reactions, and since it is a weak electrolyte, one must assume the solution to contain a certain amount of ammonium hydroxide NH4OH, although it is probably chiefly composed of a solution of ammonia in water.
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 |
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, etc. 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 salts produced by the action of ammonia on acids are known as the ammonium salts[?] and all contain the compound radical ammonium (NH4). 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.
Ammonium bromide, NH4Br, 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, NH4Cl. (See sal-ammoniac[?].)
Ammonium fluoride, NH4F, 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, NH4I, 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, NH4ClO3, 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 was formerly known as sal-volatile[?] or salt of hartshorn and was formerly obtained by the dry distillation of nitrogenous organic matter such as hair, horn, decomposed urine, etc., but is now obtained by heating a mixture of sal-ammoniac, or ammonium sulfate 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, NH4.HCO3, and ammonium carbamate, NH2COONH4, 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, (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, NH4.HCO3, 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, 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 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" |
Apparently, in classical times, Sal Ammoniac was prepared from the dung of camels found near 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.
In the Bible it is related that, Jephthah smote the Ammonites, with the help of God, who accepted Jepthah's daughter, as a burned offering, as compensation (Book of Judges). Presumably, these Ammonites (descendants of the offspring of Lot[?] and his daughter) were followers of the same horny deity.
Initial text from 1911 encyclopedia