Ammonia

For other uses, see Ammonia (disambiguation).

Template:Chembox new 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, also is 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. (See Baumé scale)

Structure and basic chemical properties

The ammonia molecule has a trigonal pyramidal shape, as predicted by VSEPR theory. The nitrogen atom in the molecule has a lone electron pair, and ammonia acts as a base, a proton acceptor. This shape gives the molecule a dipole moment and makes it polar so that ammonia readily dissolves in water. The degree to which ammonia forms the ammonium ion increases upon lowering the pH of the solution— at "physiological" pH (~7), about 99% of the ammonia molecules are protonated. Temperature and salinity also affect the proportion of NH₄⁺. NH₄⁺ has the shape of a regular tetrahedron.

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.

Natural occurrence

Ammonia is found in small quantities in the atmosphere, being produced from the putrefaction of nitrogenous animal and vegetable matter. Ammonia and ammonium salts are also found in small quantities in rainwater, whereas ammonium chloride (sal-ammoniac), and ammonium sulfate are found in volcanic districts; crystals of ammonium bicarbonate have been found in Patagonian guano. The kidneys secrete NH₃ to neutralize excess acid. Ammonium salts also are found distributed through all fertile soil and in seawater. Substances containing ammonia, or those that are similar to it, are called ammoniacal.

History

The Romans called the ammonium chloride deposits they collected from near the Temple of Jupiter Amun (Greek Ἄμμων Ammon) in ancient Libya 'sal ammoniacus' (salt of Amun) because of proximity to the nearby temple. 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 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.

Synthesis and production

see Haber Process

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:

2 NH₄Cl + 2 CaO → CaCl₂ + Ca(OH)₂ + 2 NH₃

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:

H₂ + RSH → RH + H₂S(g)

• The hydrogen sulfide is then removed by passing the gas through beds of zinc oxide where it is absorbed and converted to solid zinc sulfide:

H₂S + ZnO → ZnS + H₂O

• Catalytic steam reforming of the sulfur-free feedstock is then used to form hydrogen plus carbon monoxide:

CH₄ + H₂O → CO + 3 H₂

• In the next step, the water gas shift reaction is used to convert the carbon monoxide into carbon dioxide and more hydrogen:

CO + H₂O → CO₂ + H₂

• The carbon dioxide is then removed either by absorption in aqueous ethanolamine solutions or by adsorption in pressure swing adsorbers (PSA) using proprietary solid adsorption media.

• The final step in producing the hydrogen is to use catalytic methanation to remove any small residual amounts of carbon monoxide or carbon dioxide from the hydrogen:

CO + 3 H₂ → CH₄ + H₂O

CO₂ + 4 H₂ → CH₄ + 2 H₂O

• To produce the desired end-product ammonia, the hydrogen is then catalytically reacted with nitrogen (derived from process air) to form anhydrous liquid ammonia. This step is known as the ammonia synthesis loop (also referred to as the Haber-Bosch process):

3 H₂ + N₂ → 2 NH₃

Hydrogen required for ammonia synthesis could in principle be obtained from other sources, but these alternatives, apart from the electrolysis of water to hydrogen route 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.

Biosynthesis

In certain organisms, ammonia is produced from atmospheric N₂ by enzymes called nitrogenases. The overall process is called nitrogen fixation. Although it is unlikely that biomimetic methods will be developed that are competitive with the Haber process, intense effort has been directed toward understanding the mechanism of biological nitrogen fixation. The scientific interest in this problem is motivated by the unusual structure of the active site of the enzyme, which consists of an Fe₇MoS₉ ensemble.

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.

Properties

Ammonia is a colorless gas with a characteristic pungent smell. It is lighter than air, its density being 0.589 times that of air. It is easily liquefied due to the strong hydrogen bonding between molecules; the liquid boils at -33.3 °C, and solidifies at -77.7 °C to a white crystals. Liquid ammonia possesses strong ionizing powers reflecting its high ε of 22. Liquid ammonia has a very high standard enthalpy change of vaporization (23.35 kJ/mol, cf. water 40.65 kJ/mol, methane 8.19 kJ/mol, phosphine 14.6 kJ/mol) and can therefore be used in laboratories in non-insulated vessels without additional refrigeration.

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.

Basicity

One of the most characteristic properties of ammonia is its basicity. It combines with acids to form salts; thus with hydrochloric acid it forms ammonium chloride (sal-ammoniac); with nitric acid, ammonium nitrate, etc. However perfectly dry ammonia will not combine with perfectly dry hydrogen chloride: moisture is necessary to bring about the reaction.

NH₃ + HCl → NH₄Cl

The salts produced by the action of ammonia on acids are known as the ammonium salts and all contain the ammonium ion (NH₄⁺). Anhydrous ammonia is often used for the production of methamphetamine.

Acidity

Although ammonia is well-known as a base, it can also act as an extremely weak acid. It is a protic substance, and is capable of formation of "amides" (NH₂⁻) ion, for example when solid lithium nitride is added to liquid ammonia, forming a lithium amide solution:

Li₃N_(s)+ 2 NH_{3 (l)} → 3 Li⁺_(am) + 3 NH₂⁻_(am)

In this Brønsted-Lowry acid-base reaction, ammonia serves as an acid.

Formation of other compounds

In organic chemistry, ammonia can act as a nucleophile in substitution reactions. Amines can be formed by the reaction of ammonia with alkyl halides, although the resulting –NH₂ group is also nucleophilic and secondary and tertiary amines are often formed as by-products. An excess of ammonia helps minimise multiple substitution, and neutralises the hydrogen halide formed. Methylamine is prepared commercially by the reaction of ammonia with chloromethane, and the reaction of ammonia with 2-bromopropanoic acid has been used to prepare racemic alanine in 70% yield. Ethanolamine is prepared by a ring-opening reaction with ethylene oxide: the reaction is sometimes allowed to go further to produce diethanolamine and triethanolamine.

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 as a ligand

Ball-and-stick model of the tetraamminediaquacopper(II) cation, [Cu(NH₃)₄(H₂O)₂]²⁺

Ball-and-stick model of the diamminesilver(I) cation, [Ag(NH₃)₂]⁺

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:

• Tetraamminediaquacopper(II), [Cu(NH₃)₄(H₂O)₂]²⁺, a characteristic dark blue complex formed by adding ammonia to solution of copper(II) salts.
• Diamminesilver(I), [Ag(NH₃)₂]⁺, the active species in Tollens' reagent. Formation of this complex can also help to distinguish between precipitates of the different silver halides: AgCl is soluble in dilute (2M) ammonia solution, AgBr is only soluble in concentrated ammonia solution while AgI is insoluble in aqueous solution of ammonia.

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 vertices 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.

Hg₂Cl₂ + 2 NH₃ → Hg + HgCl(NH₂) + NH₄⁺ + Cl⁻

Uses

Fertilizer

Approximately 83% (as of 2003) of ammonia is used as fertilizers either as its salts or as solutions. Consuming more than 1% of the man-made power, the production of ammonia is significant component of the world energy budget.

Precursor to nitrogenous compounds

Ammonia is directly or indirectly the precursor to most nitrogen-containing compounds. Practically all synthetic and natural organonitrogen and all inorganic nitrogen compounds are prepared from ammonia. An important derivative is nitric acid. This key material is generated via the Ostwald process by oxidisation of ammonia with air over a platinum catalyst at 700 - 850 °C, ~9 atm. Nitric oxide is an intermediate:

4 NH₃ + 8 O₂ → 4 HNO₃ + 4 H₂O

Nitric acid is used for the production of fertilizers, explosives, and other chemical compounds.

Minor and emerging uses

Refrigeration - R717

Ammonia's thermodynamic properties made it one of the refrigerants commonly used prior to the discovery of dichlorodifluoromethane. Ammonia's toxicity complicates this application. The energy efficiency of such refrigerators is relatively low.

Ammonia continues to be used as a refrigerant in industrial processes such as bulk icemaking and industrial food processing. Ammonia is increasingly popular in commercial applications, such as in grocery store freezer cases and refrigerated displays.

For remediation of gaseous emissions

Ammonia used to scrub SO₂ from the burning of fossil fuels, the resulting product is converted to ammonium sulfate for use as fertilizer. Ammonia neutralizes the nitrogen oxides (NOx) pollutants emitted by diesel engines. This technology, called SCR (selective catalytic reduction), relies on a vanadia-based catalyst.

As a fuel

Ammonia was used during World War II fuel shortages to power buses in Belgium and used in engine and solar energy applications prior to 1900. Liquid ammonia was used as the fuel of the rocket airplane, the X-15. Although not as powerful as other fuels, it left no soot in the reusable rocket engine and its density approximately matches that for the oxidizer, liquid oxygen, which simplified the aircraft's design. Ammonia is proposed as a practical, clean (CO₂-free), alternative to fossil fuel for internal combustion engines. In 1981 a Canadian company converted a 1981 Chevrolet Impala to operate using ammonia as fuel. The use of ammonia as fuel continues to be discussed.

The calorific value of ammonia is 22.5 MJ/kg (9690 BTU/lb) which is about half that of diesel. In addition because of the high hydrogen content of ammonia, in a normal engine, in which the moisture is not condensed, the calorific value of ammonia will be about 20% less than this figure.

Use as a diesel / gasoline / petrol replacement

If the ammonia is produced from renewable electricity (as most of Europe's was until 1939 by the Vemork hydro station) the potential attractions are large:

Upside:

It's an established technology with the compound available in large quantities.

It can be burnt in modern high speed car engines with little modification, other than the addition of 5% pilot fuel - which could be hydrocarbon or hydrogen, of which the latter could be stored on board. [25]

It is compressible/liquefiable/dense enough to fit into existing cars.

Retention of and familiarity with internal combustion engine and maintenance and production infrastructure

No greenhouse emissions emitted during combustion. (Though CO2 emissions of contemporary natural gas based production processes are on par with petrol combustion).

Huge manufacturing infrastructure exists based on natural gas, but only the front end of the process needs to change.

Ammonia distribution infrastructure already existing.

Could be produced using electricity from renewable and fluctuating wind power.

Could be produced from remote renewable energy sources ( eg in the the Gobi desert / Sahara ) and shipped to markets.

It can readily be used in modern engines (with the 5% pilot hydrogen or hydrocarbon) since unlike methanol it will not damage seals and pumps.

Battery powered cars will still require some form of stored fuel to heat the occupants in winter whereas as ammonia means drivers could carry on using the existing engine waste heat method.

Downside:

The efficiency will be less than a battery powered car charge / discharge efficiency at 70%. It seems that ammonia can be burnt at around 40% [26]

Hydrogen can be electrolysed at say 70% efficiency and burnt in a car at around 40% [27] given 28% overall, which is half that of a battery round trip efficiency.

One process claims and energy to ammonia conversion efficiency of 83% - [28]

However, if the waste heat from the ammonia production process can be used in say district heating this increases the effective efficiency.

Clearly the energy balance for producing ammonia from electricity, and the offset oxygen production cost, use of waste heat offset energy cost, and at what temperature the waste heat is available at are crucial.

Potential production from renewable energy sources:

The 60 MW hydro station at Vemork, Norway provided most of Europe's ammonia fertilizer, from 1911 onwards to the '30s, by electrolysis of water, to hydrogen, and then synthesis from hydrogen plus air to ammonia. So "green" ammonia could be produced in a similar fashion from renewable energy generation schemes.

If the Hydrogen Engine Center ammonia fueled commercial internal combustion engines are as high in efficiency (50%) as Ted Hollinger indicates, it will be difficult for fuel cells to compete," commented Norm Olsen, P.E., Manager of Iowa State University's BECON (Biomass Energy Conversion) Facility in Nevada, IA. [31]

Ammonia fuel: the key to hydrogen-based transportation MacKenzie, J.J. Avery, W.H. World Resources Inst., Washington, DC;

Ammonia (NH3) is a high octane fuel (110) that can replace CO2 producing fuels in automobile transportation. It shares with hydrogen the virtue of yielding only water and nitrogen as combustion products when burned in internal combustion engines but avoids the packaging, safety and logistic problems of using hydrogen fuels in motor vehicles. Ammonia can be stored under moderate pressure at ambient temperatures. (Its physical properties are closely similar to those of liquid propane.) It can be packaged in a volume compatible with present automobiles. It is used as a fertilizer in quantities of over 100 million tons per year so that facilities for its storage, safe handling, transportation and distribution are available worldwide. It could be an economical replacement for gasoline if the foreseen costs of air pollution and global warming caused by fossil fuels are included in the economic evaluation [32] [33]

Poison treatment

Solutions of ammonia in water can be applied on the skin to lessen the effects of acidic animal poisons, expecially insect poison and jellyfish poison.

Ammonia's role in biologic systems and human disease

Ammonia is an important source of nitrogen for living systems. Although atmospheric nitrogen abounds, few living creatures are capable of utilizing this nitrogen. Nitrogen is required for the synthesis of amino acids, which are the building blocks of protein. Some plants rely on ammonia and other nitrogenous wastes incorporated into the soil by decaying matter. Others, such as nitrogen-fixing legumes, benefit from symbiotic relationships with rhizobia which create ammonia from atmospheric nitrogen.

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.

Excretion

Main article: Excretion

Ammonium ions are a toxic waste product of the metabolism in animals. In fishes and aquatic invertebrates, it is excreted directly into the water. In mammals, sharks, and amphibians, it is converted in the urea cycle to urea, because it is less toxic and can be stored more efficiently. In birds, reptiles, and terrestrial snails, metabolic ammonium is converted into uric acid, which is solid, and can therefore be excreted with minimal water loss.

Theoretical role in alternative biochemistry

Ammonia has been proposed as a possible replacement for water as a bodily solvent in the theoretical alternative biochemistries of lifeforms that do not use carbon for cellular structure and water as a solvent to dissolve bodily solutes and allow essential parts of metabolic processes to occur. It has been suggested that ammonia would be most favorable for lifeforms that live in temperatures below the freezing point of water.

Liquid ammonia as a solvent

See also: Inorganic nonaqueous 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 properties of NH₃ with those of water shows that NH₃ has the lower melting point, boiling point, 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 °C is approx. 10^-33 mol²·l^-2.

Solubility of salts

	Solubility (g of salt per 100 g liquid NH3)
Ammonium acetate	253.2
Ammonium nitrate	389.6
Lithium nitrate	243.7
Sodium nitrate	97.6
Potassium nitrate	10.4
Sodium fluoride	0.35
Sodium chloride	3.0
Sodium bromide	138.0
Sodium iodide	161.9
Sodium thiocyanate	205.5

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.

Solutions of metals

See also: Solvated electron, metallic solution

Liquid ammonia will dissolve the alkali metals and other electropositive metals such as calcium, strontium, barium, europium and ytterbium. At low concentrations (<0.06 mol/L), deep blue solutions are formed: these contain metal cations and solvated electrons, free electrons which are surrounded by a cage of ammonia molecules.

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.

Redox properties of liquid ammonia

See also: Redox.

	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.

Detection and determination

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. Sulfur sticks are burnt to detect small leaks in industrial ammonia refrigeration systems. 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 hexachloroplatinate, (NH₄)₂PtCl₆.

Interstellar space

Ammonia was first detected in interstellar space in 1968, based on microwave emissions from the direction of the galactic core. This was the first polyatomic molecule to be so detected. The sensitivity of the molecule to a broad range of excitations and the ease with which it can be observed in a number of regions has made ammonia one of the most important molecules for studies of molecular clouds. The relative intensity of the ammonia lines can be used to measure the temperature of the emitting medium.

The following isotopic species of ammonia have been detected:

NH₃, ¹⁵NH₃, NH₂D, NHD₂, and ND₃

The detection of triply-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.

Safety precautions

Toxicity and storage information

Hydrochloric acid sample releasing HCl fumes which are reacting with ammonia fumes to produce a white smoke of ammonium chloride.

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.

Household use

Solutions of ammonia (5–10% by weight) are used as household cleaners, particularly for glass. These solutions are irritating to the eyes and mucous membranes (respiratory and digestive tracts), and to a lesser extent the skin. Mixing with chlorine-containing products or strong oxidants, for example household bleach can lead to hazardous compounds such as chloramines.

Laboratory use of ammonia solutions

The hazards of ammonia solutions depend on the concentration: "dilute" ammonia solutions are usually 5–10% by weight (<5.62 mol/L); "concentrated" solutions are usually prepared at >25% by weight. A 25% (by weight) solution has a density of 0.907 g/cm³, and a solution which has a lower density will be more concentrated. The European Union classification of ammonia solutions is given in the table.

Concentration by weight	Molarity	Density Mass/Volume	Classification	R-Phrases
5–10%	2.87–5.62 mol/L	48.9–95.7 g/L	Irritant (Xi)	Template:R36/37/38
10–25%	5.62–13.29 mol/L	95.7–226.3 g/L	Corrosive (C)	Template:R34
>25%	>13.29 mol/L	>226.3 g/L	Corrosive (C) Dangerous for the environment (N)	Template:R34, Template:R50

S-Phrases: Template:S1/2, Template:S16, Template:S36/37/39, Template:S45, Template:S61.

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.

Laboratory use of anhydrous ammonia (gas or liquid)

Anhydrous ammonia is classified as toxic (T) and dangerous for the environment (N). The gas is flammable (autoignition temperature: 651 °C) and can form explosive mixtures with air (16–25%). The permissible exposure limit (PEL) in the United States is 50 ppm (35 mg/m³), while the IDLH concentration is estimated at 300 ppm. Repeated exposure to ammonia lowers the sensitivity to the smell of the gas: normally the odour is detectable at concentrations of less than 0.5 ppm, but desensitized individuals may not detect it even at concentrations of 100 ppm. Anhydrous ammonia corrodes copper- and zinc-containing alloys, and so brass fittings should not be used for handling the gas. Liquid ammonia can also attack rubber and certain plastics.

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.

Safety

The U. S. Occupational Safety and Health Administration (OSHA) has set a 15-minute exposure limit for gaseous ammonia of 35 ppm by volume in the environmental air and an 8-hour exposure limit of 25 ppm by volume. Exposure to very high concentrations of gaseous ammonia can result in lung damage and death. Although ammonia is regulated in the United States as a non-flammable gas, it still meets the definition of a material that is toxic by inhalation and requires a hazardous safety permit when transported in quantities greater than 13,248 L (3,500 gallons).

See also

• Ammonia (data page)
• Ammonia production
• Chlorination
• Water purification

Bibliography

• This article incorporates text from a publication now in the public domain: Chisholm, Hugh, ed. (1911). Encyclopædia Britannica (11th ed.). Cambridge University Press. {{cite encyclopedia}}: Missing or empty |title= (help)
• Greenwood, N. N. (1997). Chemistry of the Elements (2nd Edn. ed.). Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4. {{cite book}}: |edition= has extra text (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)
• Housecroft, C. E. (2001). Inorganic Chemistry. Harlow (UK): Prentice Education. ISBN 0-582-31080-6. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
• Bretherick, L., ed. (1986). Hazards in the Chemical Laboratory (4th Edn. ed.). London: Royal Society of Chemistry. ISBN 0-85186-489-9. {{cite book}}: |edition= has extra text (help)
• Weast, R. C., ed. (1972). Handbook of Chemistry and Physics (53rd ed.). Cleveland, OH: Chemical Rubber Co.

External links

• International Chemical Safety Card 0414 (anhydrous ammonia)
• International Chemical Safety Card 0215 (aqueous solutions)
• National Pollutant Inventory - Ammonia
• CID 222 from PubChem
• Template:InFrench
• Emergency Response to Ammonia Fertilizer Releases (Spills) for the Minnesota Department of Agriculture
• National Institute for Occupational Safety and Health - Ammonia Page