Anhydrous (water-free) ammonia gas is easily liquefied under pressure (at 20°C; liquid ammonia has a vapor pressure of about 120 lb per sq in.) It is extremely soluble in water; one volume of water dissolves about 1,200 volumes of the gas at 0°C; (90 grams of ammonia in 100 cc of water), but only about 700 volumes at room temperature and still less at higher temperatures. The solution is alkaline because much of the dissolved ammonia reacts with water, H2O, to form ammonium hydroxide, NH4OH, a weak base. Liquid ammonia is used in the chemical laboratory as a solvent. It is a better solvent for ionic and polar compounds than ethanol, but not as good as water; it is a better solvent for nonpolar covalent compounds than water, but not as good as ethanol. It dissolves alkali metals and barium, calcium, and strontium by forming an unstable blue solution containing the metal ion and free electrons that slowly decomposes, releasing hydrogen and forming the metal amide. Compared to water, liquid ammonia is less likely to release protons (H+ ions), is more likely to take up protons (to form NH4+ ions), and is a stronger reducing agent. Because strong acids react with it, it does not allow strongly acidic solutions, but it dissolves many alkalies to form strongly basic solutions.
Ammonia takes part in many chemical reactions. Ammonia reacts with strong acids to form stable ammonium salts: with hydrogen chloride it forms ammonium chloride; with nitric acid, ammonium nitrate; and with sulfuric acid, ammonium sulfate. Ammonium salts of weak acids are readily decomposed into the acid and ammonia. Ammonium carbonate, (NH3)2CO3·H2O, is a colorless-to-white crystalline solid commonly known as smelling salts; in water solution it is sometimes called aromatic spirits of ammonia. Ammonia reacts with certain metal ions to form complex ions called ammines. Ammonia also reacts with Lewis acids (electron acceptors), e.g., sulfur dioxide or trioxide or boron trifluoride.
Another kind of reaction, commonly called ammonolysis, occurs when one or more of the hydrogen atoms in the ammonia molecule is replaced by some other atom or radical. Chlorine gas, Cl2, reacts directly with ammonia to form monochloramine, NH2Cl, and hydrogen chloride, HCl. Products of such ammonolyses include amides, amines, imides, imines, and nitrides. Ammonia also takes part in oxidation and reduction reactions. It burns in oxygen to form nitrogen gas, N2, and water. In the presence of a catalyst (e.g., platinum) it is oxidized in air to form water and nitric oxide, NO. It reduces hot-metal oxides to the metal (e.g., cupric oxide to copper).
Ammonia is prepared commercially in vast quantities. The major method of production is the Haber process, in which nitrogen is combined directly with hydrogen at high temperatures and pressures in the presence of a catalyst. It is obtained as a byproduct of the destructive distillation of coal. Ammonia is also prepared synthetically by the cyanamide process: nitrogen gas combines with calcium carbide, CaC2, at high temperatures to form calcium cyanamide, CaCN2, and carbon; the calcium cyanamide reacts with steam to form calcium carbonate, CaCO3, and ammonia. For use in the laboratory, ammonia is prepared by heating an ammonium salt with a strong base. It can also be prepared by reacting a metal nitride with water.
Ammonia solutions are used to clean, bleach, and deodorize; to etch aluminum; to saponify (hydolyze) oils and fats; and in chemical manufacture. The ammonia sold for household use is a dilute water solution of ammonia in which ammonium hydroxide is the active cleansing agent. It should be used with caution since it can attack the skin and eyes. The vapors are especially irritating—prolonged exposure and inhalation cause serious injury and may be fatal. Water solutions of ammonia are also called ammonium hydrate, aqua ammonia, or ammonia water; the solution may contain up to 30% ammonium hydroxide by weight at room temperature and pressure.
The major use of ammonia and its compounds is as fertilizers. Ammonia is also used in large amounts in the Ostwald process (see Ostwald, Wilhelm) for the synthesis of nitric acid; in the Solvay process for the synthesis of sodium carbonate; in the synthesis of numerous organic compounds used as dyes, drugs, and in plastics; and in various metallurgical processes.
Modern method of manufacturing sodium carbonate (soda ash), devised and commercialized in Belgium by Ernest Solvay (1838–1922). Common salt (sodium chloride) is treated with ammonia and then carbon dioxide, under carefully controlled conditions, to form sodium bicarbonate and ammonium chloride. When heated, the bicarbonate yields sodium carbonate, the desired product; the ammonium chloride is treated with lime to produce ammonia (for reuse) and calcium chloride. The process proved of great commercial value, since large quantities of soda ash are used in making glass, detergents, and cleansers. Seealso caustic soda.
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Colourless, pungent gas composed of nitrogen and hydrogen, chemical formula NH3. Easily liquefied by compression or cooling for use in refrigerating and air-conditioning equipment, it is manufactured in huge quantities. Ammonia is made by the Haber-Bosch process (see Fritz Haber). Its major use is as a fertilizer, applied directly to soil from tanks of the liquefied gas. Also employed as fertilizers are salts of ammonia, including ammonium phosphate and ammonium nitrate (the latter used in high explosives as well). Ammonia has many other industrial uses as a raw material, catalyst, and alkali. It dissolves readily in water to form ammonium hydroxide, an alkaline solution (see base) familiar as a household cleaner.
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| Section8 = }} Ammonia is a compound with the formula NH3. It is normally encountered as a gas with a characteristic pungent odor. Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to foodstuffs and fertilizers. Ammonia, either directly or indirectly, is also a building block for the synthesis of many pharmaceuticals. Although in wide use, ammonia is both caustic and hazardous. In 2006, worldwide production was estimated at 146.5 M tonnes.
Ammonia, as used commercially, is often called anhydrous ammonia. This term emphasizes the absence of water in the material. Because NH3 boils at -33 °C, the liquid must be stored under high pressure or at low temperature. Its heat of vaporization is, however, sufficiently great that NH3 can be readily handled in ordinary beakers in a fume hood. "Household ammonia" or "ammonium hydroxide" is a solution of NH3 in water. The strength of such solutions is measured in units of baume (density), with 26 degrees baume (about 30 weight percent ammonia at 15.5 °C) being the typical high concentration commercial product. Household ammonia ranges in concentration from 5 to 10 weight percent ammonia.
The main use of ammonia is for fertilizer (83% in 2003). Another major application is its conversion to explosives, because nitric acid is made via oxidation of ammonia. The entire nitrogen content of all manufactured organic compounds is derived from ammonia.
In the form of sal-ammoniac, ammonia was known to the Arabic alchemists as early as the 8th century, first mentioned by Geber (Jabir ibn Hayyan), and to the European alchemists since the 13th century, being mentioned by Albertus Magnus. It was also used by dyers in the Middle Ages in the form of fermented urine to alter the colour of vegetable dyes. In the 15th century, Basilius Valentinus showed that ammonia could be obtained by the action of alkalis 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 "spirit of hartshorn" was applied to ammonia.
Gaseous ammonia was first isolated by Joseph Priestley in 1774 and was termed by him alkaline air; however it was acquired by the alchemist Basil Valentine. Eleven years later in 1785, Claude Louis Berthollet ascertained its composition.
The Haber process to produce ammonia from the nitrogen 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 World War I, following the allied blockade that cut off the supply of nitrates from Chile. The ammonia was used to produce explosives to sustain their war effort.
Prior to the advent of cheap natural gas, hydrogen as a precursor to ammonia production was produced via the electrolysis of water. The Vemork 60 MW hydroelectric plant in Norway constructed in 1911 was used purely for this purpose and up until the second world war provided the majority of Europe's ammonia.
Because of its many uses, ammonia is one of the most highly produced inorganic chemicals. Dozens of chemical plants worldwide produce ammonia. The worldwide ammonia production in 2004 was 109 million metric tonnes. The People's Republic of China produced 28.4% of the worldwide production followed by India with 8.6%, Russia with 8.4%, and the United States with 8.2%. About 80% or more of the ammonia produced is used for fertilizing agricultural crops.
Before the start of World War I, most ammonia was obtained by the dry distillation of nitrogenous vegetable and animal waste products, including camel dung, where it was distilled by the reduction of nitrous acid and nitrites with hydrogen; in addition, it was produced by the distillation of coal, and also by the decomposition of ammonium salts by alkaline hydroxides such as quicklime, the salt most generally used being the chloride (sal-ammoniac) thus:
(Two molecules of ammonium chloride plus two calcium oxide yields calcium chloride and calcium hydroxide and two molecules of ammonia)
Today, the typical modern ammonia-producing plant first converts natural gas (i.e., methane) or liquified petroleum gas (such gases are propane and butane) or petroleum naphtha into gaseous hydrogen. The processes used in producing the hydrogen begins with removal of sulfur compounds from the natural gas (because sulfur deactivates the catalysts used in subsequent steps). Catalytic hydrogenation converts organosulfur compounds into gaseous hydrogen sulfide:
Hydrogen required for ammonia synthesis could in principle be obtained from other sources, but these alternatives - apart from the electrolysis of water into oxygen + hydrogen - are presently impractical. At one time, most of Europe's ammonia was produced from the Hydro plant at Vemork, via the electrolysis route. Various renewable energy electricity sources are also potentially applicable.
Ammonia is also a metabolic product of amino acid deamination. Ammonia excretion is common in aquatic animals. In humans, it is quickly converted to urea, which is much less toxic. This urea is a major component of the dry weight of urine. Most reptiles, including birds, as well as insects and snails solely excrete uric acid as nitrogenous waste.
It is miscible with water. Ammonia in an aqueous solution can be expelled by boiling. The aqueous solution of ammonia is basic. The maximum concentration of ammonia in water (a saturated solution) has a density of 0.880 g /cm³ and is often known as '.880 Ammonia'. Ammonia does not burn readily or sustain combustion, except under narrow fuel-to-air mixtures of 15-25% air. When mixed with oxygen, it burns with a pale yellowish-green flame. At high temperature and in the presence of a suitable catalyst, ammonia is decomposed into its constituent elements. Ignition occurs when chlorine is passed into ammonia, forming nitrogen and hydrogen chloride; if ammonia is present in excess, then the highly explosive nitrogen trichloride (NCl3) is also formed.
The ammonia molecule readily undergoes nitrogen inversion at room temperature; a useful analogy is an umbrella turning itself inside out in a strong wind. The energy barrier to this inversion is 24.7 kJ/mol, and the resonance frequency is 23.79 GHz, corresponding to microwave radiation of a wavelength of 1.260 cm. The absorption at this frequency was the first microwave spectrum to be observed.
The salts produced by the action of ammonia on acids are known as the Ammonium compounds and all contain the ammonium ion (NH4+). Anhydrous ammonia is often used for the production of methamphetamine.
In this Brønsted-Lowry acid-base reaction, ammonia serves as an acid.
The combustion of ammonia in air is very difficult in the absence of a catalyst (such as platinum gauze), as the temperature of the flame is usually lower than the ignition temperature of the ammonia-air mixture. The flammable range of ammonia in air is 16–25%.
Amides can be prepared by the reaction of ammonia with a number of carboxylic acid derivatives. Acyl chlorides are the most reactive, but the ammonia must be present in at least a twofold excess to neutralise the hydrogen chloride formed. Esters and anhydrides also react with ammonia to form amides. Ammonium salts of carboxylic acids can be dehydrated to amides so long as there are no thermally sensitive groups present: temperatures of 150–200 °C are required.
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. Where necessary in substitutive nomenclature, IUPAC recommendations prefer the name azane to ammonia: hence chloramine would be named chloroazane in substitutive nomenclature, not chloroammonia.
Ammonia can act as a ligand in transition metal complexes. It is a pure σ-donor, in the middle of the spectrochemical series, and shows intermediate hard-soft behaviour. For historical reasons, ammonia is named ammine in the nomenclature of coordination compounds. Some notable ammine complexes include:
Ammine complexes of chromium(III) were known in the late 19th century, and formed the basis of Alfred Werner's theory of coordination compounds. Werner noted that only two isomers (fac- and mer-) of the complex [CrCl3(NH3)3] could be formed, and concluded that the ligands must be arranged around the metal ion at the vertex of an octahedron. This has since been confirmed by X-ray crystallography.
An ammine ligand bound to a metal ion is markedly more acidic than a free ammonia molecule, although deprotonation in aqueous solution is still rare. One example is the Calomel reaction, where the resulting amidomercury(II) compound is highly insoluble.
The calorific value of ammonia is 22.5 MJ/kg (9690 BTU/lb) which is about half that of diesel. In a normal engine, in which the water vapour is not condensed, the calorific value of ammonia will be about 21% less than this figure.
Ammonia also plays a role in both normal and abnormal animal physiology. Ammonia is created through normal amino acid metabolism and is toxic in high concentrations. The liver converts ammonia to urea through a series of reactions known as the urea cycle. Liver dysfunction, such as that seen in cirrhosis, may lead to elevated amounts of ammonia in the blood (hyperammonemia). Likewise, defects in the enzymes responsible for the urea cycle, such as ornithine transcarbamylase, lead to hyperammonemia. Hyperammonemia contributes to the confusion and coma of hepatic encephalopathy as well as the neurologic disease common in people with urea cycle defects and organic acidurias.
Ammonia is important for normal animal acid/base balance. After formation of ammonium from glutamine, α-ketoglutarate may be degraded to produce two molecules of bicarbonate which are then available as buffers for dietary acids. Ammonium is excreted in the urine resulting in net acid loss. Ammonia may itself diffuse across the renal tubules, combine with a hydrogen ion, and thus allow for further acid excretion.
|Solubility (g of salt per 100 g liquid NH3)|
Liquid ammonia is an ionizing solvent, although less so than water, and dissolves a range of ionic compounds including many nitrates, nitrites, cyanides and thiocyanates. Most ammonium salts are soluble, and these salts act as acids in liquid ammonia solutions. The solubility of halide salts increases from fluoride to iodide. A saturated solution of ammonium nitrate contains 0.83 mol solute per mole of ammonia, and has a vapour pressure of less than 1 bar even at 25 °C.
These solutions are very useful as strong reducing agents. At higher concentrations, the solutions are metallic in appearance and in electrical conductivity. At low temperatures, the two types of solution can coexist as immiscible phases.
|E° (V, ammonia)||E° (V, water)|
|Li+ + e− ⇌ Li||−2.24||−3.04|
|K+ + e− ⇌ K||−1.98||−2.93|
|Na+ + e− ⇌ Na||−1.85||−2.71|
|Zn2+ + 2e− ⇌ Zn||−0.53||−0.76|
|NH4+ + e− ⇌ ½ H2 + NH3||0.00||–|
|Cu2+ + 2e− ⇌ Cu||+0.43||+0.34|
|Ag+ + e− ⇌ Ag||+0.83||+0.80|
The range of thermodynamic stability of liquid ammonia solutions is very narrow, as the potential for oxidation to dinitrogen, E° (N2 + 6NH4+ + 6e− ⇌ 8NH3), is only +0.04 V. In practice, both oxidation to dinitrogen and reduction to dihydrogen are slow. This is particularly true of reducing solutions: the solutions of the alkali metals mentioned above are stable for several days, slowly decomposing to the metal amide and dihydrogen. Most studies involving liquid ammonia solutions are done in reducing conditions: although oxidation of liquid ammonia is usually slow, there is still a risk of explosion, particularly if transition metal ions are present as possible catalysts.
The following isotopic species of ammonia have been detected:deuterated ammonia was considered a surprise as deuterium is relatively scarce. It is thought that the low-temperature conditions allow this molecule to survive and accumulate. The ammonia molecule has also been detected in the atmospheres of the gas giant planets, including Jupiter, along with other gases like methane, hydrogen, and helium. The interior of Saturn may include frozen crystals of ammonia.
The toxicity of ammonia solutions does not usually cause problems for humans and other mammals, as a specific mechanism exists to prevent its build-up in the bloodstream. Ammonia is converted to carbamoyl phosphate by the enzyme carbamoyl phosphate synthetase, and then enters the urea cycle to be either incorporated into amino acids or excreted in the urine. However fish and amphibians lack this mechanism, as they can usually eliminate ammonia from their bodies by direct excretion. Ammonia even at dilute concentrations is highly toxic to aquatic animals, and for this reason it is classified as dangerous for the environment. Ammonium compounds should never be allowed to come in contact with bases (unless in an intended and contained reaction), as dangerous quantities of ammonia gas could be released.
by weight (w/w)
|5–10%||2.87–5.62 mol/L||48.9–95.7 g/L||Irritant (Xi)|
|10–25%||5.62–13.29 mol/L||95.7–226.3 g/L||Corrosive (C)|
|>25%||>13.29 mol/L||>226.3 g/L|| Corrosive (C)|
the environment (N)
The ammonia vapour from concentrated ammonia solutions is severely irritating to the eyes and the respiratory tract, and these solutions should only be handled in a fume hood. Saturated ("0.880") solutions can develop a significant pressure inside a closed bottle in warm weather, and the bottle should be opened with care: this is not usually a problem for 25% ("0.900") solutions.
Ammonia solutions should not be mixed with halogens, as toxic and/or explosive products are formed. Prolonged contact of ammonia solutions with silver, mercury or iodide salts can also lead to explosive products: such mixtures are often formed in qualitative chemical analysis, and should not be acidified and concentrated before disposal once the test is completed.
Ammonia reacts violently with the halogens. Nitrogen triiodide is formed when ammonia comes in contact with iodine. It causes the explosive polymerization of ethylene oxide. It also forms explosive fulminating compounds with compounds of gold, silver, mercury, germanium or tellurium, and with stibine. Violent reactions have also been reported with acetaldehyde, hypochlorite solutions, potassium ferricyanide and peroxides.