Formula; NH3; MW 17.03; tetrahedral planar geometry, H—N—H bond angle 107.3°; N—H bond distance 1.016Å; dipole moment of the gas 1.46 x 10–18 esu; a Lewis base.
Occurrence and Uses
Ammonia occurs in nature, being constantly formed by putrefaction of the protein of dead animals and plants. While some of it is washed away by the rain into rivers and oceans where it is recycled and converted into proteins by microorganisms, much of it is rapidly absorbed from the earth by living plants making new proteins. Ammonia occurs in urine from which it was produced earlier by chemists and alchemists for use as a soluble base. It occurs in gas liquor obtained from coal gas and producer gas plants and coke ovens. Gas liquor was a major source for producing ammonia before Haber-Bosch process was developed. Combustion of coal, fuel oil, wood and natural gas, as well as forest fires produce ammonia in small amounts in the range 1 to 10 lb per ton. It occurs in many industrial effluents, wastewaters, and groundwaters at trace concentrations. It is also found at trace levels in varying concentrations in the air in most metropolitan cities.
The single largest use of ammonia is its direct application as fertilizer, and in the manufacture of ammonium fertilizers that have increased world food production dramatically. Such ammonia-based fertilizers are now the primary source of nitrogen in farm soils. Ammonia also is used in the manufacture of nitric acid, synthetic fibers, plastics, explosives and miscellaneous ammonium salts. Liquid ammonia is used as a solvent for many inorganic reactions in non-aqueous phase. Other applications include synthesis of amines and imines; as a fluid for supercritical fluid extraction and chromatography; and as a reference standard in 15N–NMR.
Physical Properties
Colorless gas; pungent suffocating odor; human odor perception 0.5 mg/m3; liquefies by compression at 9.8 atm at 25°C, or without compression at –33.35°C (at 1 atm); solidifies at –77.7°C; critical temperature and pressure, 133°C and 112.5 atm, respectively; vapor density 0.59 (air=1); density of liquid ammonia 0.677 g/mL at –34°C; dielectric constant at –34°C is about 22; extremely soluble in water; solution alkaline; pKa 9.25 in dilute aqueous solution at 25°C; the gas does not support ordinary combustion, but burns with a
yellow flame when mixed in air at 16—27% composition.
Thermochemical Properties
ΔΗ°ƒ (g) –11.02 kcal/mol
ΔΗ°ƒ (aq) –19.19 kcal/mol
ΔΗ°ƒ [NH4+(aq)] –31.67 kcal/mol
ΔG°ƒ (g) –3.94 kcal/mol
G°ƒ (aq) –6.35kcal/mol
ΔG°ƒ [NH4+(aq)] –18.97 kcal/mol
S°(g) 45.97 cal/degree mol
S°(aq) 26.6 cal/degree mol
S° [NH4+(aq)] 27.1 cal/degree mol
Cρ° (g) 8.38 cal/degree mol
Cρ° [NH4+(aq)] 19.1 cal/degree mol
ΔHvap 5.57 kcal/mol
Synthesis
Ammonia is produced from nitrogen and hydrogen at elevated temperature (500 to 550°C) and pressure (200–350 atm) (Haber–Bosch process), using a promoted iron catalyst
N2 + 3H2 → 2NH3 + heat
In the above process, finely divided iron oxide combined with sodium oxide and silica or alumina is used as the catalyst. The reaction is favored (as per Le Chatelier’s principle) by high pressure and low temperature. However, a temperature of 500 to 550°C is employed to enhance the reaction rate and prevent catalyst deactivation. Although at 200°C and 250 atm the equilibrium may yield up to 90% ammonia, the product yield is too slow. The sources of hydrogen in commercial processes include natural gas, refinery gas, water gas, coal gas, water (electrolysis) and fuel oil, and the nitrogen source is liquefied air.
Most other synthetic processes are modifications of the Haber–Bosch process, using different pressures, temperatures, gas velocities, and catalysts.
Ammonia may be obtained by decomposition of ammonium carbonate or bicarbonate. Such reactions, however, are not applied in commercial production.
(NH4)2CO3→2NH3 + CO2 + H2O
NH4HCO3 → NH3 + CO2 + H2O
Ammonia also may be produced as a by-product from gas liquor obtained from coal, gas, and coke ovens. Organic nitrogen in the coal converts to ammonium compounds which are separated from tar and distilled with an aqueous suspension of Ca(OH)2 to produce ammonia.
(NH4)2CO3 + Ca(OH)2 → CaCO3 + 2H2O + 2 NH3
Reactions
Ammonia is stable at ordinary temperatures but begins to decompose to H2 and N2 at 450°C. Decomposition is catalyzed by porcelain, pumice and metal surfaces (but not glass) in presence of which the dissociation starts at 300°C and completes around 500 to 600°C. Ammonia reacts with water producing NH4OH. The reaction is reversible; NH4OH dissociates into NH4+ and OH– ions in solution;
NH3 + H2O → [NH4OH] → NH4+ + OH–
NH4OH is probably unstable in the molecular form, dissociating into ions. There is evidence of existence of NH3•H2O and 2NH3•H2O species in aqueous solution ( J.R. LeBlanc, (Jr), Madhavan, S. and R.E. Porter. 1978. Ammonia. In Kirk-Othmer Encyclopedia of Chemical Technology, 3rd ed., Vol. 2 p. 474, New York: Wiley Interscience). Formation of such adducts may be attributed to hydrogen bonding. Gaseous NH3 and its aqueous solution is weakly basic, undergoing neutralization reactions with acids. It reacts with HCl, H2SO4, HNO3 to form corresponding ammonium salts (after the loss of water from evaporation):
NH3•H2O + HCl → NH4Cl + H2O
2NH3•H2O + H2SO4 → (NH4)2SO4 + H2O
Similar neutralization reactions occur with phosphoric, acetic and other acids. Liquid ammonia reacts with alkali metals forming amides and liberating H2. The reaction occurs in presence of a catalyst (e.g., Pt black). Alternatively, heating alkali metals in a stream of ammonia yields their amides.
2Na + 2NH3 → 2NaNH2 + H2
Reacts with Mg to form magnesium nitride, Mg3N2 liberating H2:
3Mg + 2 NH3 → Mg3N2 + 3H2
Aqueous ammonia reacts with solutions of many metal salts forming precipitates of metal hydroxides:
2NH3•H2O + ZnSO4 → Zn(OH)2 + (NH4)2SO4
Forms cupric hydroxide, Cu(OH)2 with CuSO4; the precipitate, however, dissolves in excess ammonia, forming a tetrammine copper (II) complex ion.
Cu2+ + 4NH3 → [Cu(NH3)4]2+
Reacts with chlorine forming chloramines: monochloramine, dichloramine and nitrogen trichloride:
NH3 + Cl2 → NH2Cl + HCl
NH2Cl + Cl2 → NHCl2 + HCl
NHCl2 + Cl2 → NCl3 + HCl
Such chloramines may occur in trace quantities in many chlorine-treated wastewaters that also contain trace ammonia. NCl3 combines with ammonia to form an unstable adduct, NCl3•NH3 which reacts with excess NH3 producing NH4Cl and liberating N2.
NCl3•NH3 + 3NH3 → 3NH4Cl + N2
Chloramine is also formed when chlorine is passed into liquid ammonia; furtherc reaction with ammonia produces hydrazine:
NH2Cl + NH3 → N2H4 + HCl
However, with excess ammonia, chlorine and bromine form ammonium chloride and bromide, respectively, liberating N2:
8NH3 + 3Cl2 → N2 + 6NH4Cl
Reaction with hypochlorite solution also produces chloramine. Ammonia reacts with iodine to form nitrogen triiodide, which further combines with a molecule of NH3 to form an adduct NI3•NH3, an insoluble brown-black solid which decomposes upon exposure to light in the presence of NH3:
NH3 + 3I2 → NI3 + 3HI
NI3 + NH3 → NI3 • NH3
Reacts with carbon at red heat to give ammonium cyanide, NH4CN; forms phosphine and nitrogen upon reaction with phosphorus vapor at red heat:
2NH3 + 2P→2PH3 + N2
Liquid ammonia reacts with sulfur forming nitrogen sulfide and H2S:
10S + 4NH3 —→ N4S4 + 6H2S
whereas gaseous ammonia and sulfur vapor react to form ammonium sulfide and N2:
8NH3 + 3S → 3(NH4)2S + N2
Heating with oxygen or air produces nitrogen and water:
4NH3 + 3O2 → 2N2 + 6H2O
However, reaction at 750°C to 900°C in presence of platinum or platinumrhodium catalyst produces nitric oxide and water:
4NH3 + 5O → 4NO + 6H2O
Reacts with oxides of copper, zinc, silver and many metals other than those of Group 1A and Mg at high temperatures, decomposing to N2 and water. At ambient temperatures strong oxidants oxidize ammonia:
2 NH3 + 2 KMnO4 → 2 KOH + 2 MnO2 + 2H2O + N2
K2S2O8 + 2NH3 → 2KOH + 2SO2 + 2K2O + N2
Reactions with H2S at different stoichiometric ratios may produce ammonium sulfide, hydrosulfide, NH4HS and polysulfide (NH4)2S3 having varying S contents, depending on temperature and stoichiometric ratios.
Forms ammonium carbamate, NH2•COO•NH4 with CO2 and ammonium dithiocarbamate, NH2•CSS•NH4 with CS2:
2NH3 + CO2 → NH2•COO•NH4
2NH3 + CS2 → NH2•CSS•NH4
The carbamate decomposes to urea and water when heated. Reaction with chromic acid forms ammonium dichromate, (NH4)2Cr2O7:
2NH3 + 2CrO3 + H2O → (NH4)2Cr2O7
Reactions with organic acids such as formic, acetic, benzoic, oxalic, and salicylic acids produce their corresponding ammonium salts; concentrated ammonia solution in excess forms ammonium stearate,
CH3•(CH2)16•COONH4 with stearic acid. Forms a red-colored double salt, ammonium ferric chromate, NH4Fe(CrO4)2 when added to an aqueous solution of Fe(NO3)3•6H2O and CrO3. Forms a number of coordination compounds (ammonia complex) with several metals; adds to AgCl forming soluble complex [Ag(NH3)2]Cl; forms tetraamine complex [Cu(NH3)4]SO4 with CuSO4; and forms many hexaamine complexes with cobalt, chromium, palladium, platinum and other metals. Ammonia undergoes “ammonolysis” reactions with many classes of organics including alcohols, ketones, aldehydes, phenols, and halogenated hydrocarbons. Addition and substitution reactions of ammonia are utilized in many organic syntheses. Reactions of liquid ammonia with ethanol, or gaseous ammonia with ethyl iodide, produce diethylamine, monoethylamine, and triethylamine in lesser amounts. Many organic amines and imines are synthesized
using ammonia. For example, reaction with ethylene dichloride gives ethylenediamine.
Analysis
Ammonia may be readily identified from its odor. It may be measured by titrimetry. It is absorbed in an excess amount of a standard solution of dilute sulfuric acid and the excess unreacted acid is back titrated against a standard solution of caustic soda using methyl orange indicator. Alternatively, potentiometric titration may be used to find the end point. Concentrations at trace levels in wastewaters, groundwaters, drinking waters, and air may be measured by various colorimetric techniques or by the ammonia–selective electrode method (APHA, AWWA and WEF, 1999. Standard Methods for the Examination of Water and Wastewater, 20th ed. Washington, DC, American Public Health Association). Ammonia reacts with Nessler reagent under alkaline conditions, forming a yellow color. The intensity of color is measured by spectrophotometer, absorbance being proportional to concentration of ammonia in the solution. Alternatively, it may be analyzed by the indophenol bluemethod. Ammonia reacts with hypochlorite to form monochloramine which reacts with phenol in the presence of manganous sulfate catalyst to produce
blue indophenol (Patnaik, P. 1997. Handbook of Environmental Analysis. Boca Raton, FL, Lewis Publishers). Solutions at high concentrations may be appropriately diluted to measure ammonia within the calibration range in colorimetric and electrode methods.
Hazard
Ammonia causes intense irritation of eyes, nose and respiratory tract which can lead to tears, respiratory distress, chest pain, and pulmonary edema. A few minutes exposure to 3,000 ppm can cause severe blistering of skin, lung edema, and asphyxia which can lead to death (Patnaik, P. 1992. A Comprehensive Guide to the Hazardous Properties of Chemical Substances, p. 304. New York, Van Nostrand Reinhold). Contact with liquid ammonia can cause serious blistering and destruction of skin tissues. LC50 inhalation
(mouse): 4,200 ppm/hr. Fire or explosion hazard may arise from the following ammonia reactions: Reaction with halogens produces nitrogen trihalides which explode on heating; its mixture with fluorine bursts into flame; reacts with gold, silver, or mercury to form unstable fulminate-type shock-sensitive compounds; similarly, shock-sensitive nitrides are formed when ammonia reacts with sulfur or certain metal chlorides, such as mercuric, or silver chloride; liquid ammonia reacts violently with alkali metal chlorates and ferricyanides.
