|Molar mass||17.031 g/mol|
|Odor||strong pungent odour|
|Density||0.86 kg/m3 (1.013 bar at boiling point)
|Melting point||−77.73 °C (−107.91 °F; 195.42 K)|
|Boiling point||−33.34 °C (−28.01 °F; 239.81 K)|
|47% w/w (0 °C)
31% w/w (25 °C)
18% w/w (50 °C)
|Solubility||soluble in chloroform, ether, ethanol, methanol|
|Vapor pressure||8573 h Pa|
|Acidity (pKa)||32.5 (−33 °C), 10.5 (DMSO)|
Refractive index (nD)
Std enthalpy of
|Safety data sheet||See: data page
ICSC 0414 (anhydrous)
|H221, H280, H314, H331, H400|
|P210, P261, P273, P280, P305+351+338, P310|
|EU classification||T C N|
|R-phrases||R10, R23, R34, R50|
|S-phrases||(S1/2), S9, S16, S26, S36/37/39, S45, S61|
|Flash point||flammable gas|
|651 °C (1,204 °F; 924 K)|
|Lethal dose or concentration (LD, LC):|
LD50 (Median dose)
|0.015 mL/kg (human, oral)|
LC50 (Median concentration)
|40,300 ppm (rat, 10 min)
28595 ppm (rat, 20 min)
20300 ppm (rat, 40 min)
11590 ppm (rat, 1 hr)
7338 ppm (rat, 1 hr)
4837 ppm (mouse, 1 hr)
9859 ppm (rabbit, 1 hr)
9859 ppm (cat, 1 hr)
2000 ppm (rat, 4 hr)
4230 ppm (mouse, 1 hr)
LCLo (Lowest published)
|5000 ppm (mammal, 5 min)
5000 ppm (human, 5 min)
|US health exposure limits (NIOSH): |
|50 ppm (25 ppm ACGIH- TLV; 35 ppm STEL)|
|TWA 25 ppm (18 mg/m3) ST 35 ppm (27 mg/m3)|
IDLH (Immediate danger
Related nitrogen hydrides
|Supplementary data page|
|Refractive index (n),
Dielectric constant (εr), etc.
|UV, IR, NMR, MS|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is: / ?)(|
Ammonia or azane is a compound of nitrogen and hydrogen with the formula NH3. It is a colourless gas with a characteristic pungent smell. Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilizers. Ammonia, either directly or indirectly, is also a building block for the synthesis of many pharmaceuticals and is used in many commercial cleaning products. Although in wide use, ammonia is both caustic and hazardous.
The global industrial production of ammonia for 2012 was anticipated to be 198,000,000 tonnes (195,000,000 long tons; 218,000,000 short tons), a 35% increase over the estimated 2006 global output of 146,500,000 tonnes (144,200,000 long tons; 161,500,000 short tons).
NH3 boils at −33.34 °C (−28.012 °F) at a pressure of one atmosphere, so the liquid must be stored under pressure or at low temperature. Household ammonia or ammonium hydroxide is a solution of NH3 in water. The concentration of such solutions is measured in units of the Baumé scale (density), with 26 degrees baumé (about 30% (by weight) ammonia at 15.5 °C or 59.9 °F) being the typical high-concentration commercial product.
- 1 Natural occurrence
- 2 Properties
- 3 Detection and determination
- 4 History
- 5 Uses
- 5.1 Fertilizer
- 5.2 Precursor to nitrogenous compounds
- 5.3 Cleaner
- 5.4 Fermentation
- 5.5 Antimicrobial agent for food products
- 5.6 Minor and emerging uses
- 6 Safety precautions
- 7 Synthesis and production
- 8 Liquid ammonia as a solvent
- 9 Ammonia's role in biological systems and human disease
- 10 In astronomy
- 11 See also
- 12 Notes
- 13 References
- 14 Further reading
- 15 External links
Ammonia is found in trace quantities in the atmosphere, being produced from the putrefaction (decay process) 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 ammonia to neutralize excess acid. Ammonium salts are found distributed through fertile soil and in seawater.
Ammonia is a colourless 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 (−27.94 °F), and freezes at −77.7 °C (−107.86 °F) to white crystals.
Ammonia may be conveniently deodorized by reacting it with either sodium bicarbonate or acetic acid. Both of these reactions form an odourless ammonium salt.
- The crystal symmetry is cubic, Pearson symbol cP16, space group P213 No.198, lattice constant 0.5125 nm.
- Liquid ammonia possesses strong ionising 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 uninsulated vessels without additional refrigeration. See liquid ammonia as a solvent.
- Solvent properties
- Ammonia is miscible with water. In an aqueous solution, it 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 chlorine is present in excess, then the highly explosive nitrogen trichloride (NCl3) is also formed.
The ammonia molecule has a trigonal pyramidal shape as predicted by the valence shell electron pair repulsion theory (VSEPR theory) with an experimentally determined bond angle of 106.7°. The central nitrogen atom has five outer electrons with an additional electron from each hydrogen atom. This gives a total of eight electrons, or four electron pairs that are arranged tetrahedrally. Three of these electron pairs are used as bond pairs, which leaves one lone pair of electrons. The lone pair of electrons repel more strongly than bond pairs, therefore the bond angle is not 109.5°, as expected for a regular tetrahedral arrangement, but 106.7°. The nitrogen atom in the molecule has a lone electron pair, which makes ammonia a base, a proton acceptor. This shape gives the molecule a dipole moment and makes it polar. The molecule's polarity and, especially, its ability to form hydrogen bonds, makes ammonia highly miscible with water. Ammonia is moderately basic, a 1.0 M aqueous solution has a pH of 11.6 and if a strong acid is added to such a solution until the solution is neutral (pH = 7), 99.4% of the ammonia molecules are protonated. Temperature and salinity also affect the proportion of NH4+. The latter has the shape of a regular tetrahedron and is isoelectronic with methane.
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.
One of the most characteristic properties of ammonia is its basicity. Ammonia is considered to be a weak base. It combines with acids to form salts; thus with hydrochloric acid it forms ammonium chloride (sal ammoniac); with nitric acid, ammonium nitrate, etc. Perfectly dry ammonia will not combine with perfectly dry hydrogen chloride; moisture is necessary to bring about the reaction. As a demonstration experiment, opened bottles of concentrated ammonia and hydrochloric acid produce clouds of ammonium chloride, which seem to appear "out of nothing" as the salt forms where the two diffusing clouds of molecules meet, somewhere between the two bottles.
- NH3 + HCl → NH4Cl
Although ammonia is well known as a weak base, it can also act as an extremely weak acid. It is a protic substance and is capable of formation of amides (which contain the NH2− ion). For example, lithium dissolves in liquid ammonia to give a solution of lithium amide:
- Li + NH3 → LiNH2 + ½ H2
Like water, ammonia undergoes molecular autoionisation to form its acid and base conjugates:
- 2 NH
3 (aq) NH+
4 (aq) + NH−
At standard pressure and temperature, K=[NH+
2] = 10−30
- 4 NH3 + 3 O2 → 2 N2 + 6 H2O (g) (ΔH°r = −1267.20 kJ/mol)
The standard enthalpy change of combustion, ΔH°c, expressed per mole of ammonia and with condensation of the water formed, is −382.81 kJ/mol. Dinitrogen is the thermodynamic product of combustion: all nitrogen oxides are unstable with respect to N2 and O2, which is the principle behind the catalytic converter. Nitrogen oxides can be formed as kinetic products in the presence of appropriate catalysts, a reaction of great industrial importance in the production of nitric acid:
- 4 NH3 + 5 O2 → 4 NO + 6 H2O
A subsequent reaction leads to NO2
- 2 NO + O2 → 2 NO2
The combustion of ammonia in air is very difficult in the absence of a catalyst (such as platinum gauze), because 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%.
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 byproducts. 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 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".
Pentavalent ammonia is known as λ5-amine, or more commonly, ammonium hydride. This crystalline solid is only stable under high pressure, and decomposes back into trivalent ammonia and hydrogen gas at normal conditions. This substance was once investigated as a possible solid rocket fuel in 1966.
Ammonia as a ligand
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 dark blue complex formed by adding ammonia to a solution of copper(II) salts. Tetraamminediaquacopper(II) hydroxide is known as Schweizer's reagent, and has the remarkable ability to dissolve cellulose. Diamminesilver(I) ([Ag(NH3)2]+) is the active species in Tollens' reagent. Formation of this complex can also help to distinguish between precipitates of the different silver halides: silver chloride (AgCl) is soluble in dilute (2M) ammonia solution, silver bromide (AgBr) is only soluble in concentrated ammonia solution, whereas silver iodide (AgI) is insoluble in aqueous ammonia.
Ammine complexes of chromium(III) were known in the late 19th century, and formed the basis of Alfred Werner's revolutionary theory on the structure of coordination compounds. Werner noted only two isomers (fac- and mer-) of the complex [CrCl3(NH3)3] could be formed, and concluded the ligands must be arranged around the metal ion at the vertices of an octahedron. This proposal 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−
Detection and determination
This section is about detection in the laboratory. For detection in astronomy, see chapter In astronomy.
Ammonia in solution
Ammonia and ammonium salts can be readily detected, in very minute traces, by the addition of Nessler's solution, which gives a distinct yellow colouration in the presence of the least trace of ammonia or ammonium salts. 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.
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. Ammonia is an irritant and irritation increases with concentration; the Permissible Exposure Limit is 25 ppm, and lethal above 500 ppm. Higher concentrations are hardly detected by conventional detectors, the type of detector is chosen according to the sensitivity required (e.g. semiconductor, catalytic, electrochemical). Holographic sensors have been proposed for detecting concentrations up to 12.5% in volume.
Ammoniacal nitrogen (NH3-N)
Ammoniacal nitrogen (NH3-N) is a measure commonly used for testing the quantity of ammonium ions, derived naturally from ammonia, and returned to ammonia via organic processes, in water or waste liquids. It is a measure used mainly for quantifying values in waste treatment and water purification systems, as well as a measure of the health of natural and man made water reserves. It is measured in units of mg/L (milligram per litre).
The Romans gave the name sal ammoniacus (salt of Amun) to the ammonium chloride deposits that they collected near the Temple of Amun (Greek Ἄμμων Ammon) in ancient Libya. 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 (ammonium chloride).
In the form of sal ammoniac (نشادر, nushadir) ammonia was important to the Muslim alchemists as early as the 8th century, first mentioned by the Persian chemist Jābir ibn Hayyān, 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 hooves and horns of oxen and neutralizing the resulting carbonate with hydrochloric acid, the name "spirit of hartshorn" was applied to ammonia.
The Haber–Bosch 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 in Germany 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 war efforts.
Approximately 83% (as of 2004) of ammonia is used as fertilizers either as its salts, solutions or anhydrously. When applied to soil, it helps provide increased yields of crops such as maize and wheat. 30% of agricultural nitrogen applied in the USA is in the form of anhydrous ammonia and worldwide 110 million tonnes are applied each year.
Precursor to nitrogenous compounds
Ammonia is directly or indirectly the precursor to most nitrogen-containing compounds. Virtually all synthetic nitrogen compounds are derived from ammonia. An important derivative is nitric acid. This key material is generated via the Ostwald process by oxidation of ammonia with air over a platinum catalyst at 700–850 °C (1,292–1,562 °F), ~9 atm. Nitric oxide is an intermediate in this conversion:
- NH3 + 2 O2 → HNO3 + H2O
Ammonia is also used to make the following compounds:
- Hydrazine, in the Olin Raschig process and the peroxide process
- Hydrogen cyanide, in the BMA process and the Andrussow process
- Hydroxylamine and ammonium carbonate, in the Raschig process
- Phenol, in the Raschig–Hooker process
- Urea, in the Bosch–Meiser urea process and in Wöhler synthesis
- Amino acids, using Strecker amino-acid synthesis
- Acrylonitrile, in the Sohio process
Ammonia can also be used to make compounds in reactions which are not specifically named. Examples of such compounds include: ammonium perchlorate, ammonium nitrate, formamide, dinitrogen tetroxide, alprazolam, ethanolamine, ethyl carbamate, hexamethylenetetramine, and ammonium bicarbonate.
Household ammonia is a solution of NH3 in water (i.e., ammonium hydroxide) used as a general purpose cleaner for many surfaces. Because ammonia results in a relatively streak-free shine, one of its most common uses is to clean glass, porcelain and stainless steel. It is also frequently used for cleaning ovens and soaking items to loosen baked-on grime. Household ammonia ranges in concentration by weight from 5 to 10% ammonia.
Solutions of ammonia ranging from 16% to 25% are used in the fermentation industry as a source of nitrogen for microorganisms and to adjust pH during fermentation.
Antimicrobial agent for food products
As early as in 1895, it was known that ammonia was "strongly antiseptic ... it requires 1.4 grams per litre to preserve beef tea." In one study, anhydrous ammonia destroyed 99.999% of zoonotic bacteria in 3 types of animal feed, but not silage.[non-primary source needed] Anhydrous ammonia is currently used commercially to reduce or eliminate microbial contamination of beef. Pink slime (or lean finely textured beef in the beef industry) is made from fatty beef trimmings (about 50–70% fat) by removing the fat using heat and centrifugation, then treating it with ammonia to kill E. coli. The process was deemed effective and safe by the US Department of Agriculture based on a study (funded by a producer of pink slime) that found that the treatment reduces E. coli to undetectable levels. There have been safety concerns about the process as well as consumer complaints about the taste and smell of beef treated at optimal levels of ammonia. The level of ammonia in any final product has not come close to toxic levels to humans.
Minor and emerging uses
Refrigeration – R717
Because of ammonia's vaporization properties, it is a useful refrigerant. It was commonly used prior to the popularisation of chlorofluorocarbons (Freons). Anhydrous ammonia is widely used in industrial refrigeration applications and hockey rinks because of its high energy efficiency and low cost. It suffers from the disadvantage of toxicity, which restricts its domestic and small-scale use. Along with its use in modern vapor-compression refrigeration it was used in a mixture along with hydrogen and water in absorption refrigerators. The Kalina cycle, which is of growing importance to geothermal power plants, depends on the wide boiling range of the ammonia–water mixture.
For remediation of gaseous emissions
Ammonia is used to scrub SO2 from the burning of fossil fuels, and 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.
Ammonia may be used to mitigate gaseous spills of phosgene.
As a fuel
Ammonia was used during World War II to power buses in Belgium, and in engine and solar energy applications prior to 1900. Liquid ammonia also fuelled the Reaction Motors XLR99 rocket engine that powered the X-15 hypersonic research aircraft. Although not as powerful as other fuels, it left no soot in the reusable rocket engine and its density approximately matches the density of the oxidizer, liquid oxygen, which simplified the aircraft's design.
Ammonia has been proposed as a practical alternative to fossil fuel for internal combustion engines. 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 cannot be easily or efficiently used in existing Otto cycle engines because of its very low octane rating, although with only minor modifications to carburettors/injectors and a drastic reduction in compression ratio, which would require new pistons, a gasoline engine could be made to work exclusively with ammonia, at a low fraction of its power output before conversion and much higher fuel consumption.
An automobile fuel tank could store ammonia as a liquid as long as the tank was pressurized appropriately, depending on the temperature. Ammonia's thermodynamic properties are such that at -30 C, the tank pressure would only have to be 27.5 psi, about the same as a car tire. At 30 °C (86 °F) the pressure in the tank would need to be 170 psi to keep the ammonia liquid. If tank pressure was released, the liquid ammonia would turn gaseous and raise the pressure again to that level. Common pneumatic tool air compressors operate at this pressure, so fuel tank pressure is not a barrier to adoption of automobile fuel usage.
There are other barriers to widespread automobile usage. In terms of raw ammonia supplies, plants would have to be built to increase production levels, requiring significant capital and energy sources. Although it is the second most produced chemical, the scale of ammonia production is a small fraction of world petroleum usage. It could be manufactured from renewable energy sources, as well as coal or nuclear power. The 60 MW Rjukan dam in Telemark, Norway produced ammonia via electrolysis of water for many years from 1913 producing fertilizer for much of Europe. If produced from coal, the CO2 could be sequestered, but carbon capture and storage from coal power plants is not yet beyond prototype stages.
In 2007, a University of Michigan pickup powered by ammonia drove from Detroit to San Francisco as part of a demonstration, requiring only one fill-up in Wyoming.
Ammonia engines or ammonia motors, using ammonia as a working fluid, have been proposed and occasionally used. The principle is similar to that used in a fireless locomotive, but with ammonia as the working fluid, instead of steam or compressed air. Ammonia engines were used experimentally in the 19th century by Goldsworthy Gurney in the UK and in streetcars in New Orleans.
As a stimulant
Ammonia, as the vapor released by smelling salts, has found significant use as a respiratory stimulant. Ammonia is commonly used in the illegal manufacture of methamphetamine through a Birch reduction. The Birch method of making methamphetamine is dangerous because the alkali metal and liquid ammonia are both extremely reactive, and the temperature of liquid ammonia makes it susceptible to explosive boiling when reactants are added.
At standard temperature and pressure, ammonia is less dense than atmosphere, and has approximately 60% of the lifting power of hydrogen or helium. Ammonia has sometimes been used to fill weather balloons as a lifting gas. Because of its relatively high boiling point (compared to helium and hydrogen), ammonia could potentially be refrigerated and liquefied aboard an airship to reduce lift and add ballast (and returned to a gas to add lift and reduce ballast).
Ammonia has been used to darken quartersawn white oak in Arts & Crafts and Mission-style furniture. Ammonia fumes react with the natural tannins in the wood and cause it to change colours.
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. NIOSH recently reduced the IDLH from 500 to 300 based on recent more conservative interpretations of original research in 1943. IDLH (Immediately Dangerous to Life and Health) is the level to which a healthy worker can be exposed for 30 minutes without suffering irreversible health effects. Other organizations have varying exposure levels. U.S. Navy Standards [U.S. Bureau of Ships 1962] maximum allowable concentrations (MACs):continuous exposure (60 days): 25 ppm / 1 hour: 400 ppm Ammonia vapour has a sharp, irritating, pungent odour that acts as a warning of potentially dangerous exposure. The average odour threshold is 5 ppm, well below any danger or damage. 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).
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. 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.
Ammonia toxicity is believed to be a cause of otherwise unexplained losses in fish hatcheries. Excess ammonia may accumulate and cause alteration of metabolism or increases in the body pH of the exposed organism. Tolerance varies among fish species. At lower concentrations, around 0.05 mg/L, un-ionised ammonia is harmful to fish species and can result in poor growth and feed conversion rates, reduced fecundity and fertility and increase stress and susceptibility to bacterial infections and diseases. Exposed to excess ammonia, fish may suffer loss of equilibrium, hyper-excitability, increased respiratory activity and oxygen uptake and increased heart rate. At concentrations exceeding 2.0 mg/L, ammonia causes gill and tissue damage, extreme lethargy, convulsions, coma, and death. Experiments have shown that the lethal concentration for a variety of fish species ranges from 0.2 to 2.0 mg/l.
During winter, when reduced feeds are administered to aquaculture stock, ammonia levels can be higher. Lower ambient temperatures reduce the rate of algal photosynthesis so less ammonia is removed by any algae present. Within an aquaculture environment, especially at large scale, there is no fast-acting remedy to elevated ammonia levels. Prevention rather than correction is recommended to reduce harm to farmed fish and in open water systems, the surrounding environment.
Similar to propane, anhydrous ammonia boils below room temperature when at atmospheric pressure. A storage vessel capable of 250 psi (1.7 MPa) is suitable to contain the liquid. 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.
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. Caution should be used that the chemical is never mixed into any liquid containing bleach, as a poisonous gas may result. Mixing with chlorine-containing products or strong oxidants, such as 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 that has a lower density will be more concentrated. The European Union classification of ammonia solutions is given in the table.
by weight (w/w)
|5–10%||2.87–5.62 mol/L||48.9–95.7 g/L||Irritant (Xi)||R36/37/38|
|10–25%||5.62–13.29 mol/L||95.7–226.3 g/L||Corrosive (C)||R34|
|>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 inorganic analysis, and should be lightly acidified but not concentrated (<6% w/v) 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 50 ppm, but desensitised 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, a primary high explosive, is formed when ammonia comes in contact with iodine. Ammonia causes the explosive polymerisation 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.
Synthesis and production
This section is about industrial synthesis. For synthesis in certain organisms, see section § Biosynthesis below.
Because of its many uses, ammonia is one of the most highly produced inorganic chemicals. Dozens of chemical plants worldwide produce ammonia. Consuming more than 1% of all man-made power, ammonia production is a significant component of the world energy budget. Market research reports total ammonia production in 2012 is 198 million tonnes and is predicted to increase by about 35 million tonnes by 2019. China produced 28.4% of the worldwide production (increasingly from coal as part of urea synthesis) 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
Hydrogen for ammonia synthesis could also be produced economically by using the water gas reaction followed by the water gas shift reaction, produced by passing steam through red-hot coke, to give a mixture of hydrogen and carbon dioxide gases, followed by removal of the carbon dioxide "washing" the gas mixture with water under pressure (25 standard atmospheres (2,500 kPa)); or by using other sources like coal or coke gasification.
Modern ammonia-producing plants depend on industrial hydrogen production to react with atmospheric nitrogen using a magnetite catalyst or over a promoted Fe catalyst under high pressure (100 standard atmospheres (10,000 kPa)) and temperature (450 °C) 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 also be produced economically using other sources like coal or coke gasification or less economically from the electrolysis of water into oxygen + hydrogen and other alternatives that are presently impractical for large scale. 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.
As a sustainable alternative to the relatively inefficient electrolysis, hydrogen can be generated from organic wastes (such as biomass or food-industry waste), using catalytic reforming. This releases hydrogen from carbonaceous substances at only 10-20% of energy used by electrolysis and may lead to hydrogen being produced from municipal wastes at below zero cost (allowing for the tipping fees and efficient catalytic reforming, such as cold-plasma). Catalytic (thermal) reforming is possible in small, distributed (even mobile) plants, to take advantage of low-value, stranded biomass/biowaste or natural gas deposits. Conversion of such wastes into ammonia solves the problem of hydrogen storage, as hydrogen can be released economically from ammonia on-demand, without the need for high-pressure or cryogenic storage.
It is also easier to store ammonia on board vehicles than to store hydrogen, as ammonia is less flammable than petrol or LPG.
Liquid ammonia as a solvent
Liquid ammonia is the best-known and most widely studied nonaqueous ionising solvent. Its most conspicuous property is its ability to dissolve alkali metals to form highly coloured, electrically conductive 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 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 hydrogen bonding in NH3 and because such bonding cannot form cross-linked networks, since each NH3 molecule has only one lone pair of electrons compared with two for each H2O molecule. The ionic self-dissociation constant of liquid NH3 at −50 °C is about 10−33 mol2·l−2.
Solubility of salts
|Solubility (g of salt per 100 g liquid NH3)|
Liquid ammonia is an ionising 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 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 (77 °F).
Solutions of metals
Liquid ammonia will dissolve the alkali metals and other electropositive metals such as magnesium, 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 that 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
|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.
Ammonia's role in biological systems and human disease
Ammonia is an important source of nitrogen for living systems. Although atmospheric nitrogen abounds (more than 75%), few living creatures are capable of using this atmospheric nitrogen in its diatomic form, N2 gas. Therefore, nitrogen fixation 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 that create ammonia from atmospheric nitrogen.
In certain organisms, ammonia is produced from atmospheric nitrogen by enzymes called nitrogenases. The overall process is called nitrogen fixation. Although it is unlikely that biomimetic methods that are competitive with the Haber process will be developed, 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 catalyzed by enzymes such as glutamate dehydrogenase 1. Ammonia excretion is common in aquatic animals. In humans, it is quickly converted to urea, which is much less toxic, particularly less basic. This urea is a major component of the dry weight of urine. Most reptiles, birds, insects, and snails excrete uric acid solely as nitrogenous waste.
Ammonia also plays a role in both normal and abnormal animal physiology. It is biosynthesised 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.
Ammonium ions are a toxic waste product of metabolism in animals. In fish 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.
Ammonia has 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. It is naturally found on Deimos and Phobos – the two moons of Mars.
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.
Since its interstellar discovery, NH3 has proved to be an invaluable spectroscopic tool in the study of the interstellar medium. With a large number of transitions sensitive to a wide range of excitation conditions, NH3 has been widely astronomically detected – its detection has been reported in hundreds of journal articles. Listed below is a sample of journal articles that highlights the range of detectors that have been used to identify ammonia.
The study of interstellar ammonia has been important to a number of areas of research in the last few decades. Some of these are delineated below and primarily involve using ammonia as an interstellar thermometer.
Interstellar formation mechanisms
The interstellar abundance for ammonia has been measured for a variety of environments. The [NH3]/[H2] ratio has been estimated to range from 10−7 in small dark clouds up to 10−5 in the dense core of the Orion Molecular Cloud Complex. Although a total of 18 total production routes have been proposed, the principal formation mechanism for interstellar NH3 is the reaction:
- NH4+ + e− → NH3 + H·
The rate constant, k, of this reaction depends on the temperature of the environment, with a value of 5.2×10−6 at 10 K. The rate constant was calculated from the formula k = a(T/300)B. For the primary formation reaction, a = 1.05×10−6 and B = −0.47. Assuming an NH4+ abundance of 3×10−7 and an electron abundance of 10−7 typical of molecular clouds, the formation will proceed at a rate of 1.6×10−9 cm−3s−1 in a molecular cloud of total density 105 cm−3.
All other proposed formation reactions have rate constants of between 2 and 13 orders of magnitude smaller, making their contribution to the abundance of ammonia relatively insignificant. As an example of the minor contribution other formation reactions play, the reaction:
- H2 + NH2 → NH3 + H
has a rate constant of 2.2×10−15. Assuming H2 densities of 105 and NH2/H2 ratio of 10−7, this reaction proceeds at a rate of 2.2×10−12, more than 3 orders of magnitude slower than the primary reaction above.
Some of the other possible formation reactions are:
- H− + NH4+ → NH3 + H2
- PNH3+ + e− → P + NH3
Interstellar destruction mechanisms
There are 113 total proposed reactions leading to the destruction of NH3. Of these, 39 were tabulated in extensive tables of the chemistry among C, N, and O compounds. A review of interstellar ammonia cites the following reactions as the principal dissociation mechanisms:
- (1) NH3 + H3+ → NH4+ + H2
- (2) NH3 + HCO+ → NH4+ + CO
with rate constants of 4.39×10−9 and 2.2×10−9, respectively. The above equations (1,2) run at a rate of 8.8×10−9 and 4.4×10−13, respectively. These calculations assumed the given rate constants and abundances of [NH3]/[H2] = 10−5, [H3+]/[H2] = 2×10−5, [HCO+]/[H2] = 2×10−9, and total densities of n = 105, typical of cold, dense, molecular clouds. Clearly, between these two primary reactions, equation (1) is the dominant destruction reaction, with a rate ~10,000 times faster than equation (2). This is due to the relatively high abundance of H3+.
Single antenna detections
Radio observations of NH3 from the Effelsberg 100-m Radio Telescope reveal that the ammonia line is separated into two components – a background ridge and an unresolved core. The background corresponds well with the locations previously detected CO. The 25 m Chilbolton telescope in England detected radio signatures of ammonia in H II regions, HNH2O masers, H-H objects, and other objects associated with star formation. A comparison of emission line widths indicates that turbulent or systematic velocities do not increase in the central cores of molecular clouds.
Microwave radiation from ammonia was observed in several galactic objects including W3(OH), Orion A, W43, W51, and five sources in the galactic centre. The high detection rate indicates that this is a common molecule in the interstellar medium and that high-density regions are common in the galaxy.
VLA observations of NH3 in seven regions with high-velocity gaseous outflows revealed condensations of less than 0.1 pc in L1551, S140, and Cepheus A. Three individual condensations were detected in Cepheus A, one of them with a highly elongated shape. They may play an important role in creating the bipolar outflow in the region.
Extragalactic ammonia was imaged using the VLA in IC 342. The hot gas has temperatures above 70 K, which was inferred from ammonia line ratios and appears to be closely associated with the innermost portions of the nuclear bar seen in CO. NH3 was also monitored by VLA toward a sample of four galactic ultracompact HII regions: G9.62+0.19, G10.47+0.03, G29.96−0.02, and G31.41+0.31. Based upon temperature and density diagnostics, it is concluded that in general such clumps are likely to be the sites of massive star formation in an early evolutionary phase prior to the development of an ultracompact HII region.
Absorption at 2.97 micrometres due to solid ammonia was recorded from interstellar grains in the Becklin-Neugebauer Object and probably in NGC 2264-IR as well. This detection helped explain the physical shape of previously poorly understood and related ice absorption lines.
A spectrum of the disk of Jupiter was obtained from the Kuiper Airborne Observatory, covering the 100 to 300 cm−1 spectral range. Analysis of the spectrum provides information on global mean properties of ammonia gas and an ammonia ice haze.
A total of 149 dark cloud positions were surveyed for evidence of 'dense cores' by using the (J,K) = (1,1) rotating inversion line of NH3. In general, the cores are not spherically shaped, with aspect ratios ranging from 1.1 to 4.4. It is also found that cores with stars have broader lines than cores without stars.
Ammonia has been detected in the Draco Nebula and in one or possibly two molecular clouds, which are associated with the high-latitude galactic infrared cirrus. The finding is significant because they may represent the birthplaces for the Population I metallicity B-type stars in the galactic halo that could have been borne in the galactic disk.
Observations of nearby dark clouds
By balancing and stimulated emission with spontaneous emission, it is possible to construct a relation between excitation temperature and density. Moreover, since the transitional levels of ammonia can be approximated by a 2-level system at low temperatures, this calculation is fairly simple. This premise can be applied to dark clouds, regions suspected of having extremely low temperatures and possible sites for future star formation. Detections of ammonia in dark clouds show very narrow lines—indicative not only of low temperatures, but also of a low level of inner-cloud turbulence. Line ratio calculations provide a measurement of cloud temperature that is independent of previous CO observations. The ammonia observations were consistent with CO measurements of rotation temperatures of ~10 K. With this, densities can be determined, and have been calculated to range between 104 and 105 cm−3 in dark clouds. Mapping of NH3 gives typical clouds sizes of 0.1 pc and masses near 1 solar mass. These cold, dense cores are the sites of future star formation.
UC HII regions
Ultra-compact HII regions are among the best tracers of high-mass star formation. The dense material surrounding UCHII regions is likely primarily molecular. Since a complete study of massive star formation necessarily involves the cloud from which the star formed, ammonia is an invaluable tool in understanding this surrounding molecular material. Since this molecular material can be spatially resolved, it is possible to constrain the heating/ionising sources, temperatures, masses, and sizes of the regions. Doppler-shifted velocity components allow for the separation of distinct regions of molecular gas that can trace outflows and hot cores originating from forming stars.
Ammonia has been detected in external galaxies, and by simultaneously measuring several lines, it is possible to directly measure the gas temperature in these galaxies. Line ratios imply that gas temperatures are warm (~50 K), originating from dense clouds with sizes of tens of pc. This picture is consistent with the picture within our Milky Way galaxy—hot dense molecular cores form around newly forming stars embedded in larger clouds of molecular material on the scale of several hundred pc (giant molecular clouds; GMCs).
- Ammonia (data page)
- Ammonia fountain
- Ammonia production
- Forming gas
- Haber process
- Relative cost of electricity generated by different sources
- Water purification
- Yost, Don M. (2007). "Ammonia and Liquid Ammonia Solutions". Systematic Inorganic Chemistry. READ BOOKS. p. 132. ISBN 1-4067-7302-6.
- Blum, Alexander (1975). "On crystalline character of transparent solid ammonia". Radiation Effects and Defects in Solids 24 (4): 277. doi:10.1080/00337577508240819.
- Budavari, Susan, ed. (1996). The Merck Index: An Encyclopedia of Chemicals, Drugs, and Biologicals (12th ed.). Merck. ISBN 0-911910-12-3.
- Perrin, D. D., Ionisation Constants of Inorganic Acids and Bases in Aqueous Solution; 2nd Ed., Pergamon Press: Oxford, 1982.
- Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
- Sigma-Aldrich Co., Ammonia. Retrieved on 20 July 2013.
- "Ammonia". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH).
- "NIOSH Pocket Guide to Chemical Hazards #0028". National Institute for Occupational Safety and Health (NIOSH).
- Ceresana. "Market Study Ammonia". Ceresana. Retrieved 7 November 2012.
- Max Appl (2006). Ammonia, in Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a02_143.pub2.
- "Ammonium hydroxide physical properties" (PDF). Archived from the original (PDF) on 27 November 2007.
- Kirschbaum, B; Sica, D; Anderson, F. P. (1999). "Urine electrolytes and the urine anion and osmolar gaps.". The Journal of laboratory and clinical medicine 133 (6): 597–604. doi:10.1016/S0022-2143(99)90190-7. ISSN 0022-2143. PMID 10360635.
- Chisholm 1911.
- Hewat, A. W.; Riekel, C. (1979). "The crystal structure of deuteroammonia between 2 and 180 K by neutron powder profile refinement". Acta Crystallographica Section A 35 (4): 569. Bibcode:1979AcCrA..35..569H. doi:10.1107/S0567739479001340.
- CRC Handbook of Chemistry and Physics, 94th ed. http://www.hbcpnetbase.com. Page 9-26. Retrieved 18 June 2013.
- Cleeton, C. E.; Williams, N. H. (1934). "Electromagnetic Waves of 1.1 cm (0 in). Wave-Length and the Absorption Spectrum of Ammonia". Physical Review 45 (4): 234. Bibcode:1934PhRv...45..234C. doi:10.1103/PhysRev.45.234.
- Baker, H. B. (1894). "Influence of moisture on chemical change". J. Chem. Soc. 65: 611–624. doi:10.1039/CT8946500611.
- Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 485. ISBN 0-08-022057-6.
- Sterrett, K. F. ; Caron, A. P. (1966). "High pressure chemistry of hydrogenous fuels". Northrop Space Labs.
- (OSHA) Source: Dangerous Properties of Industrial Materials (Sixth Edition) by N. Irving Sax
- J. L. Martinez Hurtado and C. R. Lowe (2014), Ammonia-Sensitive Photonic Structures Fabricated in Nafion Membranes by Laser Ablation, ACS Applied Materials & Interfaces 6 (11), 8903-8908. http://pubs.acs.org/doi/abs/10.1021/am5016588
- "Ammonia". h2g2 Eponyms. BBC.CO.UK. 11 January 2003. Archived from the original on 2 November 2007. Retrieved 8 November 2007.
- Chisholm 1911 cites Pliny Nat. Hist. xxxi. 39
- "Sal-ammoniac". Webmineral. Retrieved 7 July 2009.
- Haq, Syed Nomanul (28 February 1995). Names, Natures and Things: The Alchemist Jabir Ibn Hayyan and His Kitab Al-Ahjar (Book of Stones). Springer. ISBN 978-0-7923-3254-1. Retrieved 22 June 2010.
- Maurice P. Crosland (2004). Historical Studies in the Language of Chemistry. Courier Dover Publications. p. 72. ISBN 0-486-43802-3.
- Abraham, Lyndy (1990). Marvell and alchemy. Aldershot Scolar. ISBN 0-85967-774-5.
- Smith, Roland (2001). Conquering Chemistry. Sydney: McGraw-Hill. ISBN 0-07-470146-0.
- David Brown (18 April 2013). "Anhydrous ammonia fertilizer: abundant, important, hazardous". Washington Post. Retrieved 23 April 2013.
- Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
- Samuel Rideal (1895). Disinfection and Disinfectants: An Introduction to the Study of. London: Charles Griffin and Company. p. 109.
- Tajkarimi, Mehrdad; Riemann, H. P.; Hajmeer, M. N.; Gomez, E. L.; Razavilar, V.; Cliver, D. O. et al. (2008). "Ammonia disinfection of animal feeds — Laboratory study". International Journal of Food Microbiology 122 (1–2): 23–28. doi:10.1016/j.ijfoodmicro.2007.11.040. PMID 18155794.
- "Evaluation of Treatment Methods for Reducing Bacteria in Textured Beef", Jensen, Jean L et al., American Society of Agricultural and Biological Engineers Annual Meeting 2009
- Reference Document: Antimicrobial Interventions for Beef, Dawna Winkler and Kerri B. Harris, Center for Food Safety, Department of Animal Science, Texas A&M University, May 2009, page 12
- Moss, Michael (3 October 2009). "The Burger That Shattered Her Life". The New York Times.
- Moss, Michael (31 December 2009). "Safety of Beef Processing Method Is Questioned". The New York Times.
- "Diesel: Greener Than You Think". Retrieved 7 July 2009.
- "Phosgene: Health and Safety Guide". International Programme on Chemical Safety. 1998.
- "Ammonia as a Transportation Fuel IV" (PDF). Norm Olson – Iowa Energy Center. 15–16 October 2007.
- Ammonia Powered Car on YouTube
- "Watch 'Ammonia Fuel'". Greg Vezina. Retrieved 7 July 2009.
- "Ammonia Motors". aqpl43.dsl.pipex.com. 1 October 2007. Retrieved 28 November 2010.
- The Gazette Staff (6 October 2009). "Anhydrous ammonia tank locks have flaws". Cedar Rapids Gazette.
- "Illinois Attorney General | Basic Understanding Of Meth". Illinoisattorneygeneral.gov. Retrieved 21 May 2011.
- Włochowicz, A.; Stelmasiak, E. (1983). "Change in thermal properties of wool after treatment with liquid ammonia". Journal of Thermal Analysis and Calorimetry 26 (1): 17. doi:10.1007/BF01914084.
- "Fuming white oak".
- "Toxic FAQ Sheet for Ammonia" (PDF). Agency for Toxic Substances and Disease Registry (ATSDR). September 2004.
- Ammonia, IDLH Documentation
- Is Anhydrous Ammonia covered under the Hazardous Materials Safety Permit Program? from the website of the United States Department of Transportation (DOT)
- Oram, Brian. "Ammonia in Groundwater, Runoff, and Streams". The Water Centre. Retrieved 3 December 2014.
- Hargreaves, J.A.; Tucker, C.S. (2004). Managing ammonia in fish ponds. Southern Regional Aquaculture Center.
- Sergeant, Chris (5 February 2014). "The Management of Ammonia Levels in an Aquaculture Environment". Water/Wastewater. Retrieved 3 December 2014.
- Electronic Code of Federal Regulations:. Ecfr.gpoaccess.gov. Retrieved on 22 December 2011.
- Rizk-Ouaini, Rosette; Ferriol, Michel; Gazet, Josette; Saugier-Cohen Adad, Marie Therese (2006). "Oxidation reaction of ammonia with sodium hypochlorite. Production and degradation reactions of chloramines". Bulletin de la Societe Chimique de France 4: 512. doi:10.1002/14356007.a02_143.pub2. ISBN 3527306730.
- "New coal-based ammonia and urea capacity.(China)". January 2004. Retrieved 7 July 2009.
- "United States Geological Survey publication" (PDF). Retrieved 7 July 2009.
- "Nobel Prize in Chemistry (1918) – Haber-Bosch process". Retrieved 7 July 2009.
- "Chemistry of the Group 2 Elements – Be, Mg, Ca, Sr, Ba, Ra". BBC.co.uk. Retrieved 7 July 2009.
- Parkes, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry (New Edition, Revised and Rewritten ed.). London: Longmans, Green and Co. pp. 255–256 & 362.
- Shriver and Atkins. Inorganic Chemistry (Fifth Edition). W. H. Freeman and Company, New York, 2010, pp 383.
- Karl S. Roth, MD. "eMedicine Specialties > Metabolic Diseases > Hyperammonemia". Retrieved 7 July 2009.
- Adjei, M. B.; Quesenberry, K. H. and Chamblis, C. G. (June 2002). "Nitrogen Fixation and Inoculation of Forage Legumes". University of Florida IFAS Extension. Archived from the original on 20 May 2007.
- "PubChem Substance Summary". Retrieved 7 July 2009.
- Zschocke, Johannes, and Georg Hoffman (2004). Vademecum Metabolism. Friedrichsdorf, Germany: Milupa GmbH.
- Rose, Burton, and Helmut Rennke (1994). Renal Pathophysiology. Baltimore: Williams & Wilkins. ISBN 0-683-07354-0.
- Campbell, Neil A.; Jane B. Reece (2002). "44". Biology (6th ed.). San Francisco: Pearson Education, Inc. pp. 937–938. ISBN 0-8053-6624-5.
- Edited by Kirk Munsell. Image page credit Lunar and Planetary Institute. NASA. "NASA's Solar Exploration: Multimedia: Gallery: Gas Giant Interiors". Retrieved 26 April 2006.
- Cheung, A. C.; Rank, D. M.; Townes, C. H.; Thornton, D. D. and Welch, W. J. (1968). "Detection of NH3 molecules in the interstellar medium by their microwave emission". Phys. Rev. Lett. 21 (25): 1701. Bibcode:1968PhRvL..21.1701C. doi:10.1103/PhysRevLett.21.1701.
- Ho, P. T. P. and Townes, C. H. (1983). "Interstellar ammonia". Ann. Rev. Astron. Astrophys. 21 (1): 239–70. Bibcode:1983ARA&A..21..239H. doi:10.1146/annurev.aa.21.090183.001323.
- Millar, T. J. (2003). "Deuterium Fractionation in Interstellar Clouds". Space Science Reviews 106 (1): 73–86. Bibcode:2003SSRv..106...73M. doi:10.1023/A:1024677318645.
- Ungerechts, H., Walmsley, C. M., Winnewisser, G.; Walmsley; Winnewisser (1980). "Ammonia and cyanoacetylene observations of the high-density core of L-183 (L-134-N)". Astron. Astrophys. 88: 259. Bibcode:1980A&A....88..259U.
- Genzel, R.; Downes, D.; Ho, P. T. P. (1982). "NH3 in Orion-KL – A new interpretation". Astrophysics Journal 259: L103. Bibcode:1982ApJ...259L.103G. doi:10.1086/183856.
- "The UMIST data for Astrochemistry". Retrieved 7 July 2009.
- Vikor, L.; Al-Khalili, A.; Danared, H.; Djuric, N.; Dunn, G. H.; Larsson, M.; Le Padellec, A.; Rosaen, S.; Af Ugglas, M. (1999). "Branching fractions of dissociative recombination of NH4+ and NH2+ molecular ions". Astronomy and Astrophysics 344: 1027. Bibcode:1999A&A...344.1027V.
- van Dishoeck, E. F.; and Black, J. H. (1986). "Comprehensive models of diffuse interstellar clouds – Physical conditions and molecular abundances". Astrophys. J. Suppl. Ser. 62: 109–145. Bibcode:1986ApJS...62..109V. doi:10.1086/191135.
- "astrochemistry.net". astrochemistry.net. Retrieved 21 May 2011.
- Prasad, S. S. and Huntress, W. T. (1980). "A model for gas phase chemistry in interstellar clouds". The Astrophysical Journal Supplement Series 43: 1. Bibcode:1980ApJS...43....1P. doi:10.1086/190665.
- Lininger, W.; Albritton, D. L.; Fehsenfeld, F. C.; Schmeltekopf, A. L.; Ferguson, E. E. (1975). "Flow–drift tube measurements of kinetic energy dependences of some exothermic proton transfer rate constants". J. Chem. Phys. 62 (9): 3549. Bibcode:1975JChPh..62.3549L. doi:10.1063/1.430946.
- Smith, D.; Adams, N. G. (1977). "Reactions of CH+n IONS with ammonia at 300 K". Chemical Physics Letters 47 (1): 145. Bibcode:1977CPL....47..145S. doi:10.1016/0009-2614(77)85326-8.
- Wooten, A.; Bozyan, E. P.; Garrett, D. B. (1980). "Detection of C2H in cold dark clouds". Astrophysical Journal 239: 844. Bibcode:1980ApJ...239..844W. doi:10.1086/158168.
- Wilson, T. L., Downes, D., Bieging, J.; Downes; Bieging (1979). "Ammonia in Orion". AAp 71: 275. Bibcode:1979A&A....71..275W.
- MacDonald, G. H., Little, L. T., Brown, A. T., Riley, P. W., Matheson, D. N., Felli, M.; Little; Brown; Riley; Matheson; Felli (1981). "Detection of new ammonia sources". MNRAS 195: 387. Bibcode:1981MNRAS.195..387M.
- Morris, M., Zuckerman, B., Palmer, P., Turner, B. E. (1973). "Interstellar ammonia". APJ 186: 501. Bibcode:1973ApJ...186..501M. doi:10.1086/152515.
- Torrelles, J. M., Ho, P. T. P., Rodriguez, L. F., Canto, J. (1985). "VLA observations of ammonia and continuum in regions with high-velocity gaseous outflows". APJ 288: 595. Bibcode:1985ApJ...288..595T. doi:10.1086/162825.
- Ho, P. T. P., Martin, R. N.,Turner, J. L., Jackson, J. M. (1990). "VLA imaging of extragalactic ammonia – Hot gas in the nucleus of IC 342". APJL 355: L19. Bibcode:1990ApJ...355L..19H. doi:10.1086/185728.
- Cesaroni, R., Churchwell, E., Hofner, P., Walmsley, C. M., Kurtz, S.; Churchwell; Hofner; Walmsley; Kurtz (1994). "Hot ammonia toward compact HII regions". AAP 288: 903. Bibcode:1994A&A...288..903C.
- Knacke, R. F., McCorkle, S., Puetter, R. C., Erickson, E. F., Kraetschmer, W. (1982). "Observation of interstellar ammonia ice". APJ 260: 141. Bibcode:1982ApJ...260..141K. doi:10.1086/160241.
- Orton, G. S., Aumann, H. H., Martonchik, J. V., Appleby, J. F. (1982). "Airborne spectroscopy of Jupiter in the 100- to 300-cm−1 region: Global properties of ammonia gas and ice haze". Icarus 52 (1): 81. Bibcode:1982Icar...52...81O. doi:10.1016/0019-1035(82)90170-1.
- Benson, P. J., Myers, P. (1989). "A survey for dense cores in dark clouds". APJS 71: 89. Bibcode:1989ApJS...71...89B. doi:10.1086/191365.
- Mebold, U., Heithausen, A., Reif, K.; Heithausen; Reif (1987). "Ammonia in the galactic halo and the infrared cirrus". AAP 180: 213. Bibcode:1987A&A...180..213M.
- "Aqua Ammonia". airgasspecialtyproducts.com. Archived from the original on 19 November 2010. Retrieved 28 November 2010.
- This article incorporates text from a publication now in the public domain: Chisholm, Hugh, ed. (1911). "Ammonia". Encyclopædia Britannica (11th ed.). Cambridge University Press.
- Bretherick, L., ed. (1986). Hazards in the Chemical Laboratory (4th ed.). London: Royal Society of Chemistry. ISBN 0-85186-489-9. OCLC 16985764.
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419.
- Housecroft, C. E.; Sharpe, A. G. (2000). Inorganic Chemistry (1st ed.). New York: Prentice Hall. ISBN 978-0582310803.
- Weast, R. C., ed. (1972). Handbook of Chemistry and Physics (53rd ed.). Cleveland, OH: Chemical Rubber Co.
|Wikimedia Commons has media related to Ammonia.|
- International Chemical Safety Card 0414 (anhydrous ammonia), ilo.org.
- International Chemical Safety Card 0215 (aqueous solutions), ilo.org.
- CID 222 from PubChem
- "Ammoniac et solutions aqueuses" (in French). Institut National de Recherche et de Sécurité.
- Emergency Response to Ammonia Fertilizer Releases (Spills) for the Minnesota Department of Agriculture.ammoniaspills.org
- National Institute for Occupational Safety and Health – Ammonia Page, cdc.gov
- NIOSH Pocket Guide to Chemical Hazards - Ammonia, cdc.gov
- Ammonia, video