chemical of the month Ammonia Gregory C. Miller The Culver Academies Culver. IN 46511
Ammonia. NH3; formerly called alkaline air and spirit of hartshorn. Methods of Preparation Methods in Nature Any nitrogen containing organic substance (dead plants, animals, animal excretions) will decompose in soil through the help of anaerohic-acting hacterla to produce ammonia. Bacteria also have the ability to convert (fix) atmospheric nitrogen to ammonia. This is an integral step in the nitrogen cycle where bacteria living in nodules on the roots of leguminous plants (peas, beans, clover, alfalfa) produce NH3 by enzyme catalyzation. Methods of Man One very old method of producing ammonia in mass was from the distillation of hoofs, horns, and hide scraps of slaughtered animals. Hence the name spirit of hartshorn. More recent methods include the destructive distillation of coal (which is 1% combined nitrogen from the proteins of its original plants). When undergoing distillation, the coal will yield ahuut 20% of its nitrogen as ammonia. This ammonia $as mixes with the other coal products as it forms and can he separated by dissolving the mix in water and treating with lime (CaO) to produce isolated ammonia. This was the predominant niethod of ammonia production prior to World War 1.
The Haber Process is named in honor of the German chemist, Fritz Haher (1868-1934), who solved the prohlem of producing NH3 directly from its elements in 1912. He received the Nobel Prize in chemistry in 1918 for this process. Nz(g)
+ 3H&)
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2NHdg) AHo,,
= -22.04 kcal
By applying Le Chatelier's Principle to the problem, Haber determined that the best conditions under which to produce an appreciable yield of NH3 were high pressure (approximately 200 atmos) and a temperature of up to 550'C. These were technologically extreme conditiods for the day and Karl Bosch, a research engineer for the Bodische Anilin urid Soda Fahrik pioneered the equipment for industrial production of ammonia by this process; consequently, the method is also often called the Haher-Bosch Process. In the laboratory, ammonia gas may he prepared by heating an ammonium salt in the presence of slaked lime (Ca(0H)d or any other strong base.
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~ 2H20 + CaClz 2NHhCI + Ca(0H)z Z N H + Adding water to ionic nitrides will also produce ammonia.
History References to aaueous solutions of ammonia are found in Raymond Lully's (1235-1315, a Catalan scholar and missionary) work. Robert Boyle's "Skeptical Chymist" (1661) also 424
Journal of Chemical Education
DARRELL H. BEACH The Culver Academies Culver. Indiana 4651 1
mentions the aqueous solution and Johann Kunkel van Lowenstern (1630-1702) briefly mentions the gas itself from the addition of lime to sal ammoniac (NH4C1). Credit for the discovery of ammonia, however, goes to Joseph Priestly (1733-1804), who first isolated and characterized the compound in 1773 hy heating "volatile spirit of sal arnmoniac" (aqueous NH3); he referred to the gas as alkaline air. In 1782, Swedish chemist Torhurn Olaf Rergman (1735-1784) suggested the name ammonia for the vapor, and one year after his death Claude Louis Berthollet, a French chemist, determined the compound to he composed of hydrogen and nitrogen (or azote as the French called nitrogen). Physical Properties Ammonia is a colorless gas having a pungent, choking odor, stimulating both respiration and circulation. Its molar mass, 17.03 g mol-I is very close to that of water. It is easily liquefied (although Priestly could not do it, it was accomplished in 1799 by cooling and applying pressure), having a hoiling point of -33.35"C and a freezing point of -77.7"C. At S T P ammonia has a density of 0.7710 g/l-'. The soluhility of ammonia is 91.24 g/100 ml in water at O0C, which may not he readily appreciated until one realizes this means that 1185 1 of ammonia will dissolve in 11 of water a t O°C. This great solubility is often demonstrated by the ammonia fountain (for details of this demonstration see reference ( I ) and for an interesting modification see Cates and Moore, p. 498 of the upcoming June issue. In hot water, as we expect, the soluhility is far less, 7.4 gIl00 ml. Ammonia is soluble in most organic solvents such as alcohol, ether, etc. and is itself recognized as an excellent solvent when in the liquid state, second only to water in dissolving salts. The resulting solutions are very good conductors of electricity. The heat of vaporization of NH3(1), 327 cal g-I is very high, shadowed only by that of water and hydrogen fluoride; this unique and unexpected property arises because liquid ammonia is a highly associated liquid (having extensive hydrogen bonding). These bonds have a bond strength of 1.3 kcal mol-' and a hond length of 3.38 X 10-lo m; within the molecule are bond lengths of 1.008 X 10-1°m and hond strength of 110 kcal mol-'. The specific heat of ammonia, 1.07 cal g-I a t its hoiling point is slightly larger than that of water. At 700°C, ammonia decomposes to its atomic constituents. Occurrence Although ammonia is not a primary constituent of today's atmosphere, where it occurs in the troposphere a t a concentration of anywhere from 0 to 0.02 ppm (one reason for the large variance lies in the fact that rainfall readily washes the gas out of the atmosphere), it is believed to have been present in earth's primitive atmosphere along with hydrogen and methane. In photochemical smog, of which ammonia is a trace constituent, it occurs at concentrations of 2 parts per hundred million (pphm). Two sources of atmospheric NH3 are volcanic gases (as determined by the constituents of huhhles in Hawaiian glass) and the nitrogen cycle. Some waters contain ammonia hut not hy nature's own doing. Much NH3 enters water through industrial waste, cattle feedlots (areas for fattening cows prior to slaughter) and fertilizers. In water, ammonia is toxic to fish and phytoplankton a t rather low levels and is also believed to interfere with photosynthesis.
Travelling away from this planet, ammonia is found in the interstellar medium. It was discovered there in 1968 by researchers from Berkeley, California. Because of intense ultraviolet radiation from the sun, however, each ammonia molecule is estimated to have a lifetime in space of less than 100 years. As for the planets, ammonia has been found in the atmosphere of Jupiter through spectroscopic study, and it has been postulated that the other Jovian planets, Saturn, Uranus, and Neptune also have ammonia present in their environments: but because of the extremelycold conditions on them, NH3 would be frozen out of their atmospheres. It has been postulated that sodium in ammonia droplets are one cause for the bands of coloration around Jupiter (more on this shortly). Comets also contain frozen ammonia which is vaporized when the comets approach within two astronomical units (A.U.) from the source of solar radiation (1 AU = 149,597,892 km). Structure
Ammonia is a pyramid01 molecule, but its structure is often regarded as a tetrahedron due to the lone pair of electrons possessed by nitrogen.
It is this unbonded pair which gives ammonia many of its unique and unexpected properties, some of which have been mentioned already. The molecule's tetrahedron is not perfect because the pair of electrons take up more room than a bonding electron would, consequently, the bond angle of 106"45' is sliehtlv less than a tetrahedron's 109"28'. (Soec-
.
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energy for snch transformation is provided by molecular collisions.) Ammonia Chemistry
Because of its lone pair of electrons, ammonia is a good reducing agent and may act as Bronsted and Lewis bases. Reactions illustrating these tendencies are shown below
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reducing agent 4NH3 + 302 2Nz + 6Hz0 NH3 + HCI NH&l Bronsted base HBN:+ BFs H3NBF8 Lewis base Ammonia also acts as a base in the Arrhenins sense, providing hydroxide ions when dissolved in water
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NH, + H 2 0 NH,+ + OHHowever. because formation of the ammonium ion is verv at 25'C), it is a very weak ~ r r h e n i u i slight (K:, = 1.8 X base (as weak a base as acetic acid is an acid). Altbonah molecular ammonium hydroxide does not occur to anyappreciahle extent, solutions of aqueous ammonia have been traditionally labelled ammonium hydroxide. This form of ammonia, however, is far from useless as it is the most convenient form of reacting ammonia to produce other nitrogen bearing compounds. Linuid ammonia exhibits anite a bit of chemistrv. Its similarities to water have been noted already. Many reactions of ammonia (1) are analogous to water reactions and, therefore, life based on ammonia (I) is conceivable, but certainly not on this planet due to its physical properties. Like water, liquid ammonia undergoes self dissociation producing the ammonium and amide ions, respectively; consequently, liquid ammonia is amphoteric. Liquid ammonia is an excellent solvent, dissolving most of the salts that water does and some that water does not snch
as AgCl (which dissolves because it reacts with the ammonia).
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AgCl + 2NH3 Ag(NH&+ + ClThe alkali metals dissolve in ammonia to produce bronzecolored solutions when concentrated (such solutions conduct like metals) and blue-colored solutions when dilute (conducting like aqueous salt solutions). The blue color results from ammonation of the electron lost by the metal upon dissolution (recall Jupiter's color).
The metal ions are also ammonated, yet upon evaporation of the NH3, the pure metal is left behind. Beautiful crystals (dodecahedron for sodium) may be obtained by slow evaporation of the solution. The alkali earth metals react with ammonia to produce ammonates of the general formula M(NH3)e. A complete listing of substances soluble and insoluble in ammonia is too extensive to list here but may be found in reference (11) (see pp. 20-33). Uses of NH3
The most important industrial nitrogen fixation process today is the production of synthetic ammonia; 26.6% of all ammonia used today goes toward fertilizer production where it is used either directly or as ammonium salts such as ammonium sulfate and ammonium phosphate. Ammonia plays an important role in industrial organic synthesis and is an ingredient in some cleaning and bleaching compounds. Because of its physical properties, NH3(g) is used extensively as a refrigerant for ice houses, most recently in the CanadianAlaskan pipeline to cool the supporting stands of the oil pipes thus preventing damage to the permafrost through heat conduction from the friction of the flowing oil. I t is an agent for saponification of fats and oils and an etching compound on metals, particularly aluminum. Ammonia is the primary source for nitrogen in all explosives (except the nuclear bombs). And finally, ammonia finds its way into the household in dilute solution as a disinfectant and deodorant. Liquid ammonia in comhination with an oxidant is sometimes used as a rocket fuel. General Comments
Short exposure to ammonia can cause serious temporary, or residual, injury even if promptly treated. It is classified as a primary skin irritant causing severe eruptions and burns. The safety respiratory concentration limits are 50 ppm; however, one should inhale ammonia if one has accidently inhaled toxic halogen vapors. Liquid ammonia should not come in contact with Hg, halogens, calcium hypochlorite, or hvdroeen fluoride. - I n 1970, the United States made enough ammonia to rank it third in production amounts: in 1973 ammonia ranked number 2, a i d in 1979 i t ranked number 3. General References (1) Alyea. Dutton (Editors), 'T&edDemonstrstions in Chemistri? JournalofChemieal Education, Earton. Pa. 1965, p. 13. (2) Aaimov, laaac, "Asimovon Chsmistry: Anchor PressiDoubieday, GardenCity,N.Y.. 1974, p. 137,204,205. (3) Admov, Isaac. "AShort Hiitoryai Chemistry: Anchor Bwb Doubleday and Co.,Inc, Garden City. N.Y., 1965.p. 165. ( 4 ) Baker, Dsuid."TheLarousseGuideto kstronomy." LarouaeandCo.,Ine., N.Y., 1978. 117 ---..
( 5 ) Biddle, and Vaughn, "ChemisVyin H~a1thandDisease"F.A. Davis Co.,Philadslphia,
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D i v i ~ i md1natrmtion oakland Schools, Pontisc, MI., 1977, p. 85.
(9)Ehret, Willism. "Smith's College Chemistry? D. Appleton-Century Co., Inc.. N.Y., 1947. nn. 344-96" (10) Findlay, Alexander, "A Hundred Yeam of Chemistry: Gerald Duckworth and Co., LTD, London, 1965, p. 308. ~~
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Volume 58
Number 5
May 1981
425
426
Journal of Chemical Education
