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Ammonia
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For NH4+, see Ammonium. For other uses, see Ammonia (disambiguation).
Ammonia
Names
IUPAC name
Azane
Other names
Hydrogen nitride
Trihydrogen nitride
Nitrogen trihydride
Identifiers
CAS Number
7664-41-7 Y
3D model (JSmol)
Interactive image
3DMet B00004
Beilstein Reference
3587154
ChEBI
CHEBI:16134 Y
ChemSpider
217 Y
ECHA InfoCard 100.028.760
EC Number 231-635-3
Gmelin Reference
79
KEGG
D02916 Y
MeSH Ammonia
PubChem CID
222
RTECS number BO0875000
UNII
5138Q19F1X Y
UN number 1005
InChI[show]
InChI=1S/H3N/h1H3 Y
Key: QGZKDVFQNNGYKY-UHFFFAOYSA-N Y
InChI=1/H3N/h1H3
Key: QGZKDVFQNNGYKY-UHFFFAOYAF
SMILES[show]
N
Properties
Chemical formula
NH3
Molar mass 17.031 g/mol
Appearance Colourless gas
Odor strong pungent odour
Density 0.86 kg/m3 (1.013 bar at boiling point)
0.769 kg/m3 (STP)[1]
0.73 kg/m3 (1.013 bar at 15 °C)
681.9 kg/m3 at −33.3 °C (liquid)[2] See also Ammonia (data page)
817 kg/m3 at −80 °C (transparent solid)[3]
Melting point −77.73 °C (−107.91 °F; 195.42 K) (Triple point at 6.060 kPa, 195.4 K)
Boiling point −33.34 °C (−28.01 °F; 239.81 K)
Solubility in water
47% w/w (0 °C)
31% w/w (25 °C)
18% w/w (50 °C)[4]
Solubility soluble in chloroform, ether, ethanol, methanol
Vapor pressure 857.3 kPa
Acidity (pKa) 32.5 (−33 °C),[5] 10.5 (DMSO)
Basicity (pKb) 4.75
Magnetic susceptibility (χ)
-18.0·10−6 cm3/mol
Refractive index (nD)
1.3327
Viscosity 0.276 cP (−40 °C)
Structure
Point group
C3v
Molecular shape
Trigonal pyramid
Dipole moment
1.42 D
Thermochemistry
Std molar
entropy (So298)
193 J·mol−1·K−1[6]
Std enthalpy of
formation (ΔfHo298)
−46 kJ·mol−1[6]
Hazards
Safety data sheet See: data page
ICSC 0414 (anhydrous)
GHS pictograms [7]
GHS hazard statements
H221, H280, H314, H331, H400[7]
GHS precautionary statements
P210, P261, P273, P280, P305+351+338, P310[7]
NFPA 704
1
3
0
Flash point flammable gas
Autoignition
temperature
651 °C (1,204 °F; 924 K)
Explosive limits 15–28%
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)[8]
LCLo (lowest published)
5000 ppm (mammal, 5 min)
5000 ppm (human, 5 min)[8]
US health exposure limits (NIOSH):[9]
PEL (Permissible)
50 ppm (25 ppm ACGIH- TLV; 35 ppm STEL)
REL (Recommended)
TWA 25 ppm (18 mg/m3) ST 35 ppm (27 mg/m3)
IDLH (Immediate danger)
300 ppm
Related compounds
Other cations
Phosphine
Arsine
Stibine
Related nitrogen hydrides
Hydrazine
Hydrazoic acid
Related compounds
Ammonium hydroxide
Supplementary data page
Structure and
properties
Refractive index (n),
Dielectric constant (εr), etc.
Thermodynamic
data
Phase behaviour
solid–liquid–gas
Spectral data
UV, IR, NMR, MS
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Y verify (what is Y N ?)
Infobox references
Ammonia or azane is a compound of nitrogen and hydrogen with the formula NH3. The simplest pnictogen hydride, ammonia is a colourless gas with a characteristic pungent smell. It is a common nitrogenous waste, particularly among aquatic organisms, and it 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 pharmaceutical products and is used in many commercial cleaning products.
Although common in nature and in wide use, ammonia is both caustic and hazardous in its concentrated form. It is classified as an extremely hazardous substance in the United States as defined in Section 302 of the U.S. Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities.[10]
The global industrial production of ammonia in 2014 was 176 million tonnes (173,000,000 long tons; 194,000,000 short tons),[11] a 16% increase over the 2006 global industrial production of 152 million tonnes (150,000,000 long tons; 168,000,000 short tons).[12] Industrial ammonia is sold either as ammonia liquor (usually 28% ammonia in water) or as pressurized or refrigerated anhydrous liquid ammonia transported in tank cars or cylinders.[13]
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.[14]
Contents
[hide]
1 Natural occurrence
2 Properties
2.1 Structure
2.2 Amphotericity
2.3 Self-dissociation
2.4 Combustion
2.5 Formation of other compounds
2.6 Ammonia as a ligand
3 Detection and determination
3.1 Ammonia in solution
3.2 Gaseous ammonia
3.3 Ammoniacal nitrogen (NH3-N)
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
5.6.1 Refrigeration – R717
5.6.2 For remediation of gaseous emissions
5.6.3 As a fuel
5.6.4 As a stimulant
5.6.5 Textile
5.6.6 Lifting gas
5.6.7 Woodworking
6 Safety precautions
6.1 Toxicity
6.1.1 Coking wastewater
6.1.2 Aquaculture
6.2 Storage information
6.3 Household use
6.4 Laboratory use of ammonia solutions
6.5 Laboratory use of anhydrous ammonia (gas or liquid)
7 Synthesis and production
8 Liquid ammonia as a solvent
8.1 Solubility of salts
8.2 Solutions of metals
8.3 Redox properties of liquid ammonia
9 Ammonia's role in biological systems and human disease
9.1 Biosynthesis
9.2 In physiology
9.3 Excretion
10 In astronomy
10.1 Interstellar space
10.1.1 Interstellar formation mechanisms
10.1.2 Interstellar destruction mechanisms
10.1.3 Single antenna detections
10.1.4 Interferometric studies
10.1.5 Infrared detections
10.1.6 Observations of nearby dark clouds
10.1.7 UC HII regions
10.1.8 Extragalactic detection
11 See also
12 Notes
13 References
14 Further reading
15 External links
Natural occurrence[edit]
Ammonia is found in trace quantities in nature, being produced from the 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.[15] The kidneys secrete ammonia to neutralize excess acid.[16] Ammonium salts are found distributed through fertile soil and in seawater.
Ammonia is also found throughout the Solar System on Mars, Jupiter, Saturn, Uranus, Neptune, and Pluto, among other places: on smaller, icy planets such as Pluto, ammonia can act as a geologically important antifreeze, as a mixture of water and ammonia can potentially have a melting point of as low as 173 kelvins if the ammonia concentration is high enough and thus allow such planets to retain internal oceans and active geology at a far lower temperature than would be possible with water alone.[17][18] Substances containing ammonia, or those that are similar to it, are called ammoniacal.
Properties[edit]
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.[15]
Ammonia may be conveniently deodorized by reacting it with either sodium bicarbonate or acetic acid. Both of these reactions form an odourless ammonium salt.
Solid
The crystal symmetry is cubic, Pearson symbol cP16, space group P213 No.198, lattice constant 0.5125 nm.[19]
Liquid
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.
Combustion
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.
Structure[edit]
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°.[20] 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°.[20] 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.[21]
Amphotericity[edit]
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.[22][23] 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
The salts produced by the action of ammonia on acids are known as the ammonium salts and all contain the ammonium ion (NH4+).[22]
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:
2Li + 2NH3 → 2LiNH2 + H2
Self-dissociation[edit]
Like water, ammonia undergoes molecular autoionisation to form its acid and base conjugates:
2 NH
3 (aq) ⇌ NH+
4 (aq) + NH−
2 (aq)
Ammonia often functions as a weak base, so it has some buffering ability. Shifts in pH will cause more or fewer ammonium cations (NH+
4) and amide anions (NH−
2) to be present in solution. At standard pressure and temperature, K=[NH+
4][NH−
2] = 10−30
Combustion[edit]
The combustion of ammonia to nitrogen and water is exothermic:
4 NH3 + 3 O2 → 2 N2 + 6 H2O (g) ΔH°r = −1267.20 kJ/mol (or −316.8 kJ/mol if expressed per mol of NH3)
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 or warm chromium(III) oxide), 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%.[24]
Formation of other compounds[edit]
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.[22] 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.[25]
Ammonia as a ligand[edit]
Main article: Metal ammine complex
Ball-and-stick model of the tetraamminediaquacopper(II) cation, [Cu(NH3)4(H2O)2]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 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[edit]
This section is about detection in the laboratory. For detection in astronomy, see § In astronomy.
Ammonia in solution[edit]
Main article: Ammonia 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.[26]
Gaseous ammonia[edit]
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.[26] Ammonia is an irritant and irritation increases with concentration; the permissible exposure limit is 25 ppm, and lethal above 500 ppm.[27] 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.[28]
Ammoniacal nitrogen (NH3-N)[edit]
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).
History[edit]
This high-pressure reactor was built in 1921 by BASF in Ludwigshafen and was re-erected on the premises of the University of Karlsruhe in Germany.
The ancient Greek historian Herodotus mentioned that there were outcrops of salt in an area of Libya that was inhabited by a people called the "Ammonians" (now: the Siwa oasis in northwestern Egypt, where salt lakes still exist).[29][30] The Greek geographer Strabo also mentioned the salt from this region. However, the ancient authors Dioscorides, Apicius, Arrian, Synesius, and Aëtius of Amida described this salt as forming clear crystals that could be used for cooking and that were essentially rock salt.[31] Hammoniacus sal appears in the writings of Pliny,[32] although it is not known whether the term is identical with the more modern sal ammoniac (ammonium chloride).[15][33][34]
The fermentation of urine by bacteria produces a solution of ammonia; hence fermented urine was used in Classical Antiquity to wash cloth and clothing, to remove hair from hides in preparation for tanning, to serve as a mordant in dying cloth, and to remove rust from iron.[35]
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,[36] and to the European alchemists since the 13th century, being mentioned by Albertus Magnus.[15] 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.[37] 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.[15][38]
Gaseous ammonia was first isolated by Joseph Black in 1756 by reacting sal ammoniac (Ammonium Chloride) with calcined magnesia (Magnesium Oxide).[39][40] It was isolated again by Peter Woulfe in 1767,[41][42] by Carl Wilhelm Scheele in 1770[43] and by Joseph Priestley in 1773 and was termed by him "alkaline air".[15][44] Eleven years later in 1785, Claude Louis Berthollet ascertained its composition.[45][15]
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,[46] following the allied blockade that cut off the supply of nitrates from Chile. The ammonia was used to produce explosives to sustain war efforts.[47]
Prior to the availability of natural gas, hydrogen as a precursor to ammonia production was produced via the electrolysis of water or using the chloralkali process.
With the advent of the steel industry in the 20th century, ammonia became a byproduct of the production of coking coal.
Uses[edit]
Fertilizer[edit]
Globally, approximately 88% (as of 2014) of ammonia is used as fertilizers either as its salts, solutions or anhydrously.[11] When applied to soil, it helps provide increased yields of crops such as maize and wheat.[citation needed] 30% of agricultural nitrogen applied in the USA is in the form of anhydrous ammonia and worldwide 110 million tonnes are applied each year.[48]
Precursor to nitrogenous compounds[edit]
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:[49]
NH3 + 2 O2 → HNO3 + H2O
Nitric acid is used for the production of fertilizers, explosives, and many organonitrogen compounds.
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.
Cleaner[edit]
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.
Fermentation[edit]
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[edit]
As early as in 1895, it was known that ammonia was "strongly antiseptic ... it requires 1.4 grams per litre to preserve beef tea."[50] In one study, anhydrous ammonia destroyed 99.999% of zoonotic bacteria in 3 types of animal feed, but not silage.[51][non-primary source needed] Anhydrous ammonia is currently used commercially to reduce or eliminate microbial contamination of beef.[52][53] Lean finely textured beef in the beef industry is made from fatty beef trimmings (c. 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 that found that the treatment reduces E. coli to undetectable levels.[54] 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.[55] The level of ammonia in any final product has not come close to toxic levels to humans.
Minor and emerging uses[edit]
Refrigeration – R717[edit]
Because of ammonia's vaporization properties, it is a useful refrigerant.[46] 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 is 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. Ammonia coolant is also used in the S1 radiator aboard the International Space Station in two loops which are used to regulate the internal temperature and enable temperature dependent experiments.[56][57]
For remediation of gaseous emissions[edit]
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.[58]
Ammonia may be used to mitigate gaseous spills of phosgene.[59]
As a fuel[edit]
Ammoniacal Gas Engine Streetcar in New Orleans drawn by Alfred Waud in 1871.
The X-15 aircraft used ammonia as one component fuel of its rocket engine
The raw energy density of liquid ammonia is 11.5 MJ/L,[60] which is about a third that of diesel. Although it can be used as a fuel, for a number of reasons this has never been common or widespread. In addition to direct utilization of ammonia as a fuel in combustion engines there is also the opportunity to convert ammonia back to hydrogen where it can be used to power hydrogen fuel cells or it can be directly used within high temperature fuel cells[61].
Ammonia engines or ammonia motors, using ammonia as a working fluid, have been proposed and occasionally used.[62] 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 the St. Charles Avenue Streetcar line in New Orleans in the 1870s and 1880s,[63] and during World War II ammonia was used to power buses in Belgium.[64]
Ammonia is sometimes proposed as a practical alternative to fossil fuel for internal combustion engines.[64] Its high octane rating of 120[65] and low flame temperature allows the use of high compression ratios without a penalty of high NOx production. Since ammonia contains no carbon, its combustion cannot produce carbon monoxide, hydrocarbons or soot.
However ammonia cannot be easily used in existing Otto cycle engines because of its very narrow flammability range and there are also 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 for many years from 1913 producing fertilizer for much of Europe.
Despite this, several tests have been done. In 1981, a Canadian company converted a 1981 Chevrolet Impala to operate using ammonia as fuel.[66][67] 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.[68]
Compared to hydrogen as a fuel, ammonia is much more energy efficient, and it would be a much lower cost to produce, store, and deliver hydrogen as ammonia than as compressed and/or cryogenic hydrogen.[60] The conversion of ammonia to hydrogen via the sodium-amide process,[69] either as a catalyst for combustion or as fuel for a proton exchange membrane fuel cell,[60] is another possibility. Conversion to hydrogen would allow the storage of hydrogen at nearly 18 wt% compared to ~5% for gaseous hydrogen under pressure.
Rocket engines have also been fueled by ammonia. The Reaction Motors XLR99 rocket engine that powered the X-15 hypersonic research aircraft used liquid ammonia. 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.
As a stimulant[edit]
Anti-meth sign on tank of anhydrous ammonia, Otley, Iowa. Anhydrous ammonia is a common farm fertilizer that is also a critical ingredient in making methamphetamine. In 2005, Iowa used grant money to give out thousands of locks to prevent criminals from getting into the tanks.[70]
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.[71] 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.[citation needed]
Textile[edit]
Liquid ammonia is used for treatment of cotton materials, giving properties like mercerisation, using alkalis. In particular, it is used for prewashing of wool.[72]
Lifting gas[edit]
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).
Woodworking[edit]
See also: Ammonia fuming
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.[73]
Safety precautions[edit]
The world's longest ammonia pipeline, running from the TogliattiAzot plant in Russia to Odessa in Ukraine.
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.[74] 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[75] 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.[74] 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).[76] Household products containing ammonia (i.e., Windex) should never be used in conjunction with products containing bleach, as the resulting chemical reaction produces highly toxic fumes.[77]
Liquid ammonia is dangerous because it is hygroscopic and because it can freeze flesh. See Gas carrier#Health effects of specific cargoes carried on gas carriers for more information.
Toxicity[edit]
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[citation needed]. 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 is a constituent of tobacco smoke.[78]
Coking wastewater[edit]
Ammonia is present in coking wastewater streams, as a liquid by-product of the production of coke from coal.[79] In some cases, the ammonia is discharged to the marine environment where it acts as a pollutant. The Whyalla steelworks in South Australia is one example of a coke-producing facility which discharges ammonia into marine waters.[80]
Aquaculture[edit]
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.[81] 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.[82] Exposed to excess ammonia, fish may suffer loss of equilibrium, hyper-excitability, increased respiratory activity and oxygen uptake and increased heart rate.[81] At concentrations exceeding 2.0 mg/L, ammonia causes gill and tissue damage, extreme lethargy, convulsions, coma, and death.[81][83] Experiments have shown that the lethal concentration for a variety of fish species ranges from 0.2 to 2.0 mg/l.[83]
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[83] and in open water systems, the surrounding environment.
Storage information[edit]
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.[84] 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[edit]
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.[85]
Laboratory use of ammonia solutions[edit]
Hydrochloric acid sample releasing HCl fumes, which are reacting with ammonia fumes to produce a white smoke of ammonium chloride.
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.
Concentration
by weight (w/w)
Molarity
Concentration
mass/volume (w/v)
Classification
R-Phrases
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)
Dangerous for
the environment (N) R34, R50
S-Phrases: (S1/2), S16, S36/37/39, S45, 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" — see #Properties) 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)[edit]
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[edit]
This section is about industrial synthesis. For synthesis in certain organisms, see § Biosynthesis.
See also: Haber process
Production trend of ammonia between 1947 and 2007
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.[46] The USGS reports global ammonia production in 2014 was 176 million tonnes.[11] China accounted for 32.6% of that (increasingly from coal as part of urea synthesis), followed by Russia at 8.1%, India at 7.6%, and the United States at 6.4%.[11] About 88% of the ammonia produced was used for fertilizing agricultural crops.[11]
Before the start of World War I, most ammonia was obtained by the dry distillation[86] 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[87] such as quicklime, the salt most generally used being the chloride (sal ammoniac) thus:[15]
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));[88] 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):[89]
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.
For small scale laboratory synthesis, one can heat urea and Ca(OH)2
(NH2)2CO + Ca(OH)2 → CaCO3 + 2 NH3
Liquid ammonia as a solvent[edit]
See also: Inorganic nonaqueous 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[edit]
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 157.0
Sodium bromide 138.0
Sodium iodide 161.9
Sodium thiocyanate 205.5
Liquid ammonia is an ionising solvent, although less so than water, and dissolves a range of ionic compounds, including many nitrates, nitrites, cyanides, thiocyanates, metal cyclopentadienyl complexes and metal bis(trimethylsilyl)amides.[90] 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[edit]
See also: Solvated electron
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[edit]
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.
Ammonia's role in biological systems and human disease[edit]
Main symptoms of hyperammonemia (ammonia reaching toxic concentrations).[91]
Ammonia is both a metabolic waste and a metabolic input throughout the biosphere. It 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.[92]
Biosynthesis[edit]
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,[citation needed] 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.
In physiology[edit]
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.[93]
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.[94]
Excretion[edit]
Main article: 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.[95]
Reference ranges for blood tests, comparing blood content of ammonia (shown in yellow near middle) with other constituents
In astronomy[edit]
Ammonia occurs in the atmospheres of the outer gas planets such as Jupiter (0.026% ammonia) and Saturn (0.012% ammonia).
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.[96] It is naturally found on Deimos and Phobos – the two moons of Mars.
Interstellar space[edit]
Ammonia was first detected in interstellar space in 1968, based on microwave emissions from the direction of the galactic core.[97] 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.[98] 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.[99]
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[edit]
Ball-and-stick model of the diamminesilver(I) cation, [Ag(NH3)2]+
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[100] up to 10−5 in the dense core of the Orion Molecular Cloud Complex.[101] Although a total of 18 total production routes have been proposed,[102] 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.[103] The rate constant was calculated from the formula k = a ( T / 300 ) B {\displaystyle k=a(T/300)^{B}} . For the primary formation reaction, a = 6994104999999999999♠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 6991160000000000000♠1.6×10−9 cm−3s−1 in a molecular cloud of total density 7005100000000000000♠105 cm−3.[104]
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.[105] 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[edit]
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.[106] A review of interstellar ammonia cites the following reactions as the principal dissociation mechanisms:[98]
NH3 + H3+ → NH4+ + H2
(1)
NH3 + HCO+ → NH4+ + CO
(2)
with rate constants of 4.39×10−9[107] and 2.2×10−9,[108] 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.[109] 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[edit]
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.[110] 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.[111]
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.[112]
Interferometric studies[edit]
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.[113]
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.[114] 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.[115]
Infrared detections[edit]
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.[116]
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.[117]
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.[118]
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.[119]
Observations of nearby dark clouds[edit]
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[edit]
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.
Extragalactic detection[edit]
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).
See also[edit]
Ammonia (data page)
Ammonia fountain
Ammonia production
Cost of electricity by source
Forming gas
Haber process
Hydrazine
Water purification
Notes[edit]
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Jump up ^ "NIOSH Pocket Guide to Chemical Hazards #0028". National Institute for Occupational Safety and Health (NIOSH).
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^ Jump up to: a b c Chisholm 1911, p. 862.
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^ Jump up to: a b Chisholm 1911, p. 863.
Jump up ^ (OSHA) Source: Sax, N. Irving (1984) Dangerous Properties of Industrial Materials. 6th Ed. Van Nostrand Reinhold. ISBN 0-442-28304-0.
Jump up ^ Hurtado, J. L. Martinez; Lowe, C. R. (2014). "Ammonia-Sensitive Photonic Structures Fabricated in Nafion Membranes by Laser Ablation". ACS Applied Materials & Interfaces. 6 (11): 8903–8908. doi:10.1021/am5016588. ISSN 1944-8244.
Jump up ^ Herodotus with George Rawlinson, trans., The History of Herodotus (New York, New York: Tandy-Thomas Co., 1909), vol.2, Book 4, § 181, pp. 304–305.
Jump up ^ The land of the Ammonians is mentioned elsewhere in Herodotus' History and in Pausanias' Description of Greece:
Herodotus with George Rawlinson, trans., The History of Herodotus (New York, New York: Tandy-Thomas Co., 1909), vol. 1, Book 2, § 42, p. 245, vol. 2, Book 3, § 25, p. 73, and vol. 2, Book 3, § 26, p. 74.
Pausanias with W.H.S. Jones, trans., Description of Greece (London, England: William Heinemann Ltd., 1979), vol. 2, Book 3, Ch. 18, § 3, pp. 109 and 111 and vol. 4, Book 9, Ch. 16, § 1, p. 239.
Jump up ^ Kopp, Hermann, Geschichte der Chemie [History of Chemistry] (Braunschweig, (Germany): Friedrich Vieweg und Sohn, 1845), Part 3, p. 237. [in German]
Jump up ^ Chisholm 1911 cites Pliny Nat. Hist. xxxi. 39. See: Pliny the Elder with John Bostock and H. T. Riley, ed.s, The Natural History (London, England: H. G. Bohn, 1857), vol. 5, Book 31, § 39, p. 502.
Jump up ^ "Sal-ammoniac". Webmineral. Retrieved 7 July 2009.
Jump up ^ Pliny also mentioned that when some samples of what was purported to be natron (Latin: nitrum, impure sodium carbonate) were treated with lime (calcium carbonate) and water, the natron would emit a pungent smell, which some authors have interpreted as signifying that the natron either was ammonium chloride or was contaminated with it. See:
Pliny with W.H.S. Jones, trans., Natural History (London, England: William Heinemann Ltd., 1963), vol. 8, Book 31, § 46, pp. 448–449. From pp. 448–449: "Adulteratur in Aegypto calce, deprehenditur gusto. Sincerum enim statim resolvitur, adulteratum calce pungit et asperum [or aspersum] reddit odorem vehementer." (In Egypt it [i.e., natron] is adulterated with lime, which is detected by taste ; for pure natron melts at once, but adulterated natron stings because of the lime, and emits a strong, bitter odour [or: when sprinkled [(aspersum) with water] emits a vehement odour])
Kidd, John, Outlines of Mineralogy (Oxford, England: N. Bliss, 1809), vol. 2, p. 6.
Moore, Nathaniel Fish, Ancient Mineralogy: Or, An Inquiry Respecting Mineral Substances Mentioned by the Ancients: … (New York, New York: G. & C. Carvill & Co., 1834), pp. 96–97.
Jump up ^ See:
Forbes, R.J., Studies in Ancient Technology, vol. 5, 2nd ed. (Leiden, Netherlands: E.J. Brill, 1966), pp. 19, 48, and 65.
Moeller, Walter O., The Wool Trade of Ancient Pompeii (Leiden, Netherlands: E.J. Brill, 1976), p. 20.
Faber, G.A. (pseudonym of: Goldschmidt, Günther) (May 1938) "Dyeing and tanning in classical antiquity," Ciba Review, 9 : 277–312. Available at: Elizabethan Costume
Smith, William, A Dictionary of Greek and Roman Antiquities (London, England: John Murray, 1875), article: "Fullo" (i.e., fullers or launderers), pp. 551–553.
Rousset, Henri (March 31, 1917) "The laundries of the Ancients," Scientific American Supplement, 83 (2152) : 197.
Bond, Sarah E., Trade and Taboo: Disreputable Professions in the Roman Mediterranean (Ann Arbor, Michigan: University of Michigan Press, 2016), p. 112.
Binz, Arthur (1936) "Altes und Neues über die technische Verwendung des Harnes" (Ancient and modern [information] about the technological use of urine), Zeitschrift für Angewandte Chemie, 49 (23) : 355–360. [in German]
Witty, Michael (December 2016) "Ancient Roman urine chemistry," Acta Archaeologica, 87 (1) : 179–191. Witty speculates that the Romans obtained ammonia in concentrated form by adding wood ash (impure potassium carbonate) to urine that had been fermented for several hours. Struvite (magnesium ammonium phosphate) is thereby precipitated, and the yield of struvite can be increased by then treating the solution with bittern, a magnesium-rich solution that is a byproduct of making salt from sea water. Roasting struvite releases ammonia vapors.
Jump up ^ 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.
Jump up ^ Spiritus salis urinæ (spirit of the salt of urine, i.e., ammonium carbonate) had apparently been produced prior to Valentinus, although he presented a new, simpler method for preparing it in his book: Valentinus, Basilius, Vier Tractätlein Fr. Basilii Valentini … [Four essays of Brother Basil Valentine … ] (Frankfurt am Main, (Germany): Luca Jennis, 1625), "Supplementum oder Zugabe" (Supplement or appendix), pp. 80–81: "Der Weg zum Universal, damit die drei Stein zusammen kommen." (The path to the Universal, so that the three stones come together.). From p. 81: "Der Spiritus salis Urinæ nimbt langes wesen zubereiten / dieser proceß aber ist waß leichter unnd näher auß dem Salz von Armenia, … Nun nimb sauberen schönen Armenischen Salz armoniac ohn alles sublimiren / thue ihn in ein Kolben / giesse ein Oleum Tartari drauff / daß es wie ein Muß oder Brey werde / vermachs baldt / dafür thu auch ein grosen vorlag / so lege sich als baldt der Spiritus Salis Urinæ im Helm an Crystallisch … " (Spirit of the salt of urine [i.e., ammonium carbonate] requires a long method [i.e., procedure] to prepare; this [i.e., Valentine’s] process [starting] from the salt from Armenia [i.e., ammonium chloride], however, is somewhat easier and shorter … Now take clean nice Armenian salt, without sublimating all [of it]; put it in a [distillation] flask; pour oil of tartar [i.e., potassium carbonate that has dissolved only in the water that it has absorbed from the air] on it, [so] that it [i.e., the mixture] becomes like a mush or paste; assemble it [i.e., the distilling apparatus (alembic)] quickly; for that [purpose] connect a large receiving flask; then soon spirit of the salt of urine deposits as crystals in the "helmet" [i.e., the outlet for the vapors, which is atop the distillation flask] … )
See also: Kopp, Hermann, Geschichte der Chemie [History of Chemistry] (Braunschweig, (Germany): Friedrich Vieweg und Sohn, 1845), Part 3, p. 243. [in German]
Jump up ^ Maurice P. Crosland (2004). Historical Studies in the Language of Chemistry. Courier Dover Publications. p. 72. ISBN 0-486-43802-3.
Jump up ^ Black, Joseph (1893) [1755]. Experiments upon magnesia alba, quick-lime, and other alcaline substances. Edinburgh: W.F. Clay.
Jump up ^ Jacobson, Mark Z. (2012-04-23). Air Pollution and Global Warming: History, Science, and Solutions. Cambridge University Press. ISBN 9781107691155.
Jump up ^ "Woulfe's bottle". Chemistry World. Retrieved 2017-07-01.
Jump up ^ Woulfe, Peter (1767-01-01). "Experiments on the Distillation of Acids, Volatile Alkalies, &c. Shewing How They May be Condensed without Loss, and How Thereby We May Avoid Disagreeable and Noxious Fumes: In a Letter from Mr. Peter Woulfe, F. R. S. to John Ellis, Esq; F. R. S". Philosophical Transactions. 57: 517–536. doi:10.1098/rstl.1767.0052. ISSN 0261-0523.
Jump up ^ https://kb.osu.edu/dspace/bitstream/handle/1811/28946/Pictorial%20Life%20History_Scheele.pdf?sequence=1
Jump up ^ See:
Priestley, Joseph (1773) "Extrait d'une lettre de M. Priestley, en date du 14 Octobre 1773" (Extract of a letter from Mr. Priestley, dated 14 October 1773), Observations sur la Physique …, 2 : 389.
Priestley, Joseph, Experiments and Observations on Different Kinds of Air, vol. 1, 2nd ed. (London, England: 1775), Part 2, § 1: Observations on Alkaline Air, pp. 163–177.
Schofield, Robert E., The Enlightened Joseph Priestley: A Study of His Life and Work from 1773 to 1804 (University Park, Pennsylvania: Pennsylvania State University Press, 2004), pp. 93–94.
By 1775, Priestley had observed that electricity could decompose ammonia ("alkaline air"), yielding a flammable gas (hydrogen). See: Priestley, Joseph, Experiments and Observations on Different Kinds of Air, vol. 2 (London, England: J. Johnson, 1775), pp. 239–240.
Jump up ^ Berthollet (1785) "Analyse de l'alkali volatil" (Analysis of volatile alkali), Mémoires de l'Académie Royale des Sciences, 316–326.
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Jump up ^ Smith, Roland (2001). Conquering Chemistry. Sydney: McGraw-Hill. ISBN 0-07-470146-0.
Jump up ^ David Brown (18 April 2013). "Anhydrous ammonia fertilizer: abundant, important, hazardous". Washington Post. Retrieved 23 April 2013.
Jump up ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
Jump up ^ Samuel Rideal (1895). Disinfection and Disinfectants: An Introduction to the Study of. London: Charles Griffin and Company. p. 109.
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Jump up ^ "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
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Jump up ^ Ammonia Powered Car on YouTube
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Jump up ^ Fuming white oak. woodweb.com
^ Jump up to: a b "Toxic FAQ Sheet for Ammonia" (PDF). Agency for Toxic Substances and Disease Registry (ATSDR). September 2004.
Jump up ^ Ammonia, IDLH Documentation
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Jump up ^ "Why You Shouldn't Mix Bleach and Ammonia". ThoughtCo. Retrieved 2017-07-01.
Jump up ^ Talhout, Reinskje; Schulz, Thomas; Florek, Ewa; Van Benthem, Jan; Wester, Piet; Opperhuizen, Antoon (2011). "Hazardous Compounds in Tobacco Smoke". International Journal of Environmental Research and Public Health. 8 (12): 613–628. doi:10.3390/ijerph8020613. ISSN 1660-4601. PMC 3084482  . PMID 21556207.
Jump up ^ "Cutting-Edge Solutions For Coking Wastewater Reuse To Meet The Standard Of Circulation Cooling Systems". www.wateronline.com. Retrieved 2016-01-16.
Jump up ^ Vasudevan Rajaram; Subijoy Dutta; Krishna Parameswaran (30 June 2005). Sustainable Mining Practices: A Global Perspective. CRC Press. pp. 113–. ISBN 978-1-4398-3423-7.
^ Jump up to: a b c Oram, Brian. "Ammonia in Groundwater, Runoff, and Streams". The Water Centre. Retrieved 3 December 2014.
Jump up ^ Hargreaves, J.A.; Tucker, C.S. (2004). Managing ammonia in fish ponds. Southern Regional Aquaculture Center.
^ Jump up to: a b c Sergeant, Chris (5 February 2014). "The Management of Ammonia Levels in an Aquaculture Environment". Water/Wastewater. Retrieved 3 December 2014.
Jump up ^ Electronic Code of Federal Regulations: Archived 4 November 2011 at the Wayback Machine.. Ecfr.gpoaccess.gov. Retrieved on 22 December 2011.
Jump up ^ 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.
Jump up ^ "Nobel Prize in Chemistry (1918) – Haber-Bosch process". Retrieved 7 July 2009.
Jump up ^ "Chemistry of the Group 2 Elements – Be, Mg, Ca, Sr, Ba, Ra". BBC.co.uk. Retrieved 7 July 2009.
Jump up ^ 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.
Jump up ^ Atkins, P.W.; Overton, T.L.; Rourke, J.P.; Weller, M.T. and Armstrong, F.A. (2010) Shriver and Atkins Inorganic Chemistry. 5th Edi. W. H. Freeman and Company, New York. p. 383. ISBN 978-1-42-921820-7
Jump up ^ Neufeld, R.; Michel, R.; Herbst-Irmer, R.; Schöne, R.; Stalke, D. (2016). "Introducing a Hydrogen-Bond Donor into a Weakly Nucleophilic Brønsted Base: Alkali Metal Hexamethyldisilazides (MHMDS, M = Li, Na, K, Rb and Cs) with Ammonia". Chem. Eur. J. 22 (35): 12340–12346. doi:10.1002/chem.201600833.
Jump up ^ Roth, Karl S. "eMedicine Specialties > Metabolic Diseases > Hyperammonemia". Retrieved 7 July 2009.
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Jump up ^ Rose, Burton; Helmut Rennke (1994). Renal Pathophysiology. Baltimore: Williams & Wilkins. ISBN 0-683-07354-0.
Jump up ^ Campbell, Neil A.; Jane B. Reece (2002). "44". Biology (6th ed.). San Francisco: Pearson Education, Inc. pp. 937–938. ISBN 0-8053-6624-5.
Jump up ^ 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.
Jump up ^ Cheung, A. C.; Rank, D. M.; Townes, C. H.; Thornton, D. D.; 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.
^ Jump up to: a b Ho, P. T. P.; Townes, C. H. (1983). "Interstellar ammonia". Annu. Rev. Astron. Astrophys. 21 (1): 239–70. Bibcode:1983ARA&A..21..239H. doi:10.1146/annurev.aa.21.090183.001323.
Jump up ^ 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.
Jump up ^ Ungerechts, H.; Walmsley, C. M.; Winnewisser, G. (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.
Jump up ^ Genzel, R.; Downes, D.; Ho, P. T. P. (1982). "NH3 in Orion-KL – A new interpretation". Astrophysical Journal. 259: L103. Bibcode:1982ApJ...259L.103G. doi:10.1086/183856.
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Jump up ^ van Dishoeck, E. F.; 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.
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Jump up ^ 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.
Jump up ^ 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.
Jump up ^ 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.
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References[edit]
"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. 1 (11th ed.). Cambridge University Press. pp. 861–863.
Further reading[edit]
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 0-08-037941-9.
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.
External links[edit]
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 {{{1}}} from PubChem
"Ammoniac et solutions aqueuses" (in French). Institut National de Recherche et de Sécurité. Archived from the original on 11 December 2010.
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
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v
t
e
Molecules detected in outer space
Molecules
Diatomic
Aluminium monochloride
Aluminium monofluoride
Aluminium monoxide
Argonium
Carbon monophosphide
Carbon monosulfide
Carbon monoxide
Carborundum
Cyanogen radical
Diatomic carbon
Fluoromethylidynium
Hydrogen chloride
Hydrogen fluoride
Hydrogen (molecular)
Hydroxyl radical
Iron(II) oxide
Magnesium monohydride cation
Methylidyne radical
Nitric oxide
Nitrogen (molecular)
Nitrogen monohydride
Nitrogen sulfide
Oxygen (molecular)
Phosphorus monoxide
Phosphorus mononitride
Potassium chloride
Silicon carbide
Silicon mononitride
Silicon monoxide
Silicon monosulfide
Sodium chloride
Sodium iodide
Sulfur monohydride
Sulfur monoxide
Titanium oxide

Triatomic
Aluminium hydroxide
Aluminium isocyanide
Amino radical
Carbon dioxide
Carbonyl sulfide
CCP radical
Chloronium
Diazenylium
Dicarbon monoxide
Disilicon carbide
Ethynyl radical
Formyl radical
Hydrogen cyanide (HCN)
Hydrogen isocyanide (HNC)
Hydrogen sulfide
Hydroperoxyl
Iron cyanide
Isoformyl
Magnesium cyanide
Magnesium isocyanide
Methylene radical
N2H+
Nitrous oxide
Nitroxyl
Ozone
Phosphaethyne
Potassium cyanide
Protonated molecular hydrogen
Sodium cyanide
Sodium hydroxide
Silicon carbonitride
c-Silicon dicarbide
Silicon naphthalocyanine
Sulfur dioxide
Thioformyl
Thioxoethenylidene
Titanium dioxide
Tricarbon
Water
Four
atoms
Acetylene
Ammonia
Cyanic acid
Cyanoethynyl
Cyclopropynylidyne
Formaldehyde
Fulminic acid
HCCN
Hydrogen peroxide
Hydromagnesium isocyanide
Isocyanic acid
Isothiocyanic acid
Ketenyl
Methylene amidogen
Methyl radical
Propynylidyne
Protonated carbon dioxide
Protonated hydrogen cyanide
Silicon tricarbide
Thioformaldehyde
Tricarbon monoxide
Tricarbon sulfide
Thiocyanic acid
Five
atoms
Ammonium ion
Butadiynyl
Carbodiimide
Cyanamide
Cyanoacetylene
Cyanoformaldehyde
Cyanomethyl
Cyclopropenylidene
Formic acid
Isocyanoacetylene
Ketene
Methane
Methoxy radical
Methylenimine
Propadienylidene
Protonated formaldehyde
Protonated formaldehyde
Silane
Silicon-carbide cluster
Six
atoms
Acetonitrile
Cyanobutadiynyl radical
E-Cyanomethanimine
Cyclopropenone
Diacetylene
Ethylene
Formamide
HC4N
Ketenimine
Methanethiol
Methanol
Methyl isocyanide
Pentynylidyne
Propynal
Protonated cyanoacetylene
Seven
atoms
Acetaldehyde
Acrylonitrile
Vinyl cyanide
Cyanodiacetylene
Ethylene oxide
Hexatriynyl radical
Methylacetylene
Methylamine
Methyl isocyanate
Vinyl alcohol
Eight
atoms
Acetic acid
Aminoacetonitrile
Cyanoallene
Ethanimine
Glycolaldehyde
Heptatrienyl radical
Hexapentaenylidene
Methylcyanoacetylene
Methyl formate
Propenal
Nine
atoms
Acetamide
Cyanohexatriyne
Cyanotriacetylene
Dimethyl ether
Ethanol
Methyldiacetylene
Octatetraynyl radical
Propene
Propionitrile
Ten
atoms
or more
Acetone
Benzene
Buckminsterfullerene (C60 fullerene, buckyball)
C70 fullerene
Cyanodecapentayne
Cyanopentaacetylene
Cyanotetra-acetylene
Ethylene glycol
Ethyl formate
Methyl acetate
Methyl-cyano-diacetylene
Methyltriacetylene
Propanal
n-Propyl cyanide
Pyrimidine
Deuterated
molecules
Ammonia
Ammonium ion
Formaldehyde
Formyl radical
Heavy water
Hydrogen cyanide
Hydrogen deuteride
Hydrogen isocyanide
Methylacetylene
N2D+
Trihydrogen cation
Unconfirmed
Anthracene
Dihydroxyacetone
Ethyl methyl ether
Glycine
Graphene
H2NCO+
Linear C5
Naphthalene cation
Phosphine
Pyrene
Silylidine
Related
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Chemical formula
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Forbidden mechanism
Helium hydride ion
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Iron–sulfur world theory
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[show]
v
t
e
Nitrogen species
Hydrides:
NH3
NH4+
NH2−
N3−
N2H4
N3–
Organic:
-NH2
HCN
CN−
C2N2
H2NCN
CH2N2
Oxides:
N2O
NO
NO2
N2O3
N2O4
N2O5
HNO2
HNO3
NO2−
NO3−
Halides:
NF3
NCl3
ClN3
Chemical formulas
[show]
v
t
e
Binary compounds of hydrogen
Alkali metal hydrides
LiH
NaH
KH
RbH
CsH

Lithium hydride, LiH
ionic metal hydride

Beryllium hydride
Left (gas phase): BeH2
covalent metal hydride
Right: (BeH2)n (solid phase)
polymeric metal hydride

Borane and diborane
Left: BH3 (special conditions), covalent metalloid hydride
Right: B2H6 (standard conditions), dimeric metalloid hydride

Methane, CH4
covalent nonmetal hydride

Ammonia, NH3
covalent nonmetal hydride

Water, H2O
covalent nonmetal hydride

Hydrogen fluoride, HF
covalent nonmetal hydride
Alkaline earth hydrides
Monohydrides
BeH
MgH
CaH
SrH
BaH
BeH2
MgH2
CaH2
SrH2
BaH2
Group 13 hydrides
Boranes
BH3
B2H6
B2H2
B2H4
Alanes
AlH3
Al2H6
Gallanes
GaH3
Ga2H6
Indiganes
InH3
In2H6
Thallanes
TlH3
Tl2H6
B2H2
B2H4
B4H10
B5H9
B5H11
B6H10
B6H12
B10H14
B18H22
Group 14 hydrides
Linear alkanes
CH4
C2H6
C3H8
C4H10
C5H12
C6H14
C7H16
C8H18
C9H20
C10H22
more...
Linear alkenes
C2H4
C3H6
C4H8
C5H10
C6H12
C7H14
C8H16
C9H18
C10H20
more...
Linear alkynes
C2H2
C3H4
C4H6
C5H8
C6H10
C7H12
C8H14
C9H16
C10H18
more...
Silanes
SiH4
Si2H6
Si3H8
Si4H10
Si5H12
Si6H14
Si7H16
Si8H18
Si9H20
Si10H22
more...
Silenes
Si2H4
Silynes
Si2H2
Germanes
GeH4
Ge2H6
Ge3H8
Ge4H10
Ge5H12
Stannanes
SnH4
Sn2H6
Plumbanes
PbH4
CH
CH2
CH3
C2H
Cycloalkanes
Cycloalkenes
Annulenes
Many more
Pnictogen hydrides
Azanes
NH3
N2H4
N3H5
N4H6
N5H7
N6H8
N7H9
N8H10
N9H11
N10H12
more...
Azenes
N2H2
N3H3
N4H4
Phosphanes
PH3
P2H4
P3H5
P4H6
P5H7
P6H8
P7H9
P8H10
P9H11
P10H12
more...
Phosphenes
P2H2
P3H3
P4H4
Arsanes
AsH3
As2H4
Stibanes
SbH3
Bismuthanes
BiH3
HN3
NH
radical
Hydrogen chalcogenides
Polyoxidanes
H2O
H2O2
H2O3
H2O4
H2O5
H2O6
H2O7
H2O8
H2O9
H2O10
more...
Polysulfanes
H2S
H2S2
H2S3
H2S4
H2S5
H2S6
H2S7
H2S8
H2S9
H2S10
more...
Selanes
H2Se
H2Se2
Tellanes
H2Te
H2Te2
Polanes
PoH2
HO
HO2
HO3
H2O+–O–
H2S=S
(HS)2S+–S–
HS
HDO
D2O
T2O
Hydrogen halides
HF
HCl
HBr
HI
HAt
Transition metal hydrides
ScH2
YH2
YH3
TiH2
ZrH2
HfH2
VH
VH2
NbH
NbH2
TaH
CrH
CrH2
CrHx
NiH
PdHx (x < 1)
FeH
FeH2
FeH5
CuH
ZnH2
CdH2
HgH2
Lanthanide hydrides
LaH2
LaH3
CeH2
CeH3
PrH2
PrH3
NdH2
NdH3
SmH2
SmH3
EuH2
GdH2
GdH3
TbH2
TbH3
DyH2
DyH3
HoH2
HoH3
ErH2
ErH3
TmH2
TmH3
YbH2
YbH2.5
LuH2
LuH3
Actinide hydrides
AcH2
ThH2
Th4H15
PaH3
UH3
NpH2
NpH3
PuH2
PuH3
AmH2
AmH3
CmH2
Authority control
LCCN: sh85004538
GND: 4001735-7
NDL: 00560264
Retrieved from "https://en.wikipedia.org/w/index.php?title=Ammonia&oldid=810516359"
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Ammonia (disambiguation)

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Trigonal pyramid (chemistry)

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GHS hazard statement

values: Contains gas under pressure: may explode if heated

values:
abbr
Causes severe skin burns and eye damage

values: Keep away from heat/sparks/open flames/hot surfaces – No smoking

values: Wear protective gloves/protective clothing/eye protection/face protection

values: IF IN EYES: Rinse continuously with water for several minutes. Remove contact lenses if present and easy to do – continue rinsing

values: Immediately call a POISON CENTER or doctor/physician

values: Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil

values: Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas

values: Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen

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Explosive limit

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Ammonium hydroxide

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Dielectric constant

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UV/VIS spectroscopy

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NMR spectroscopy

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Caustic (substance)

values: Emergency Planning and Community Right-to-Know Act

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/wiki/Ammonia_liquor
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Ammonia liquor

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Earth's atmosphere

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Polar molecule

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Resonance frequency

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Microwave spectroscopy

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Molecular autoionisation

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Alkyl halide

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IUPAC nomenclature

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Complex (chemistry)

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HSAB concept

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Coordination compound

values: Edit section: Detection and determination

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/wiki/Nessler%27s_solution
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Nessler's solution

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Volumetric analysis

values: Edit section: Ammoniacal nitrogen (NH3-N)

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Classical Antiquity

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/wiki/J%C4%81bir_ibn_Hayy%C4%81n
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Jābir ibn Hayyān

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/wiki/Basilius_Valentinus
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Basilius Valentinus

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Ammonium Chloride

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Magnesium Oxide

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Haber–Bosch process

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Bosch–Meiser urea process

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Strecker amino-acid synthesis

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Sohio process

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/wiki/Zoonotic_bacteria
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Zoonotic bacteria

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Compound feed

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Beef trimmings

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US Department of Agriculture

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St. Charles Avenue Streetcar

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Hydrogen as a fuel

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Directive 67/548/EEC

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Marine environment

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Whyalla steelworks

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values: R36/37/38: Irritating to eyes, respiratory system and skin

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R34: Causes burns

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R50: Very toxic to aquatic organisms

values: (S1/2): Keep locked up and out of the reach of children

values: S16: Keep away from sources of ignition - No smoking

values: S36/37/39: Wear suitable protective clothing, gloves and eye/face protection

values: S45: In case of accident or if you feel unwell seek medical advice immediately (show the label where possible)

values: S61: Avoid release to the environment. Refer to special instructions/safety data sheet

values: Edit section: Laboratory use of anhydrous ammonia (gas or liquid)

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Parts per million

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Primary explosive

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High explosive

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Water gas shift

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Electrical conductivity

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Vapour pressure

values: Edit section: Redox properties of liquid ammonia

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values: This claim needs references to reliable sources. (May 2015)

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Organic aciduria

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Celestial body atmosphere

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Very Large Array

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Becklin-Neugebauer Object

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PubMed Identifier

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Interstellar nitrogen monohydride

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Mononitrogen monosulfide

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Tricarbon sulfide

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Vinyl cyanide

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Iron–sulfur world theory

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Alkaline earth hydride

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Group 15 hydride

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Group 16 hydride

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Hydrogen tetroxide

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Hydrogen pentoxide

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Hydrogen hexoxide

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Hydrogen heptoxide

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Hydrogen octoxide

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Hydrogen nonoxide

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Hydrogen decoxide

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values: Hydrogen hexasulfide (page does not exist)

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values: Hydrogen octasulfide (page does not exist)

values: Hydrogen nonasulfide (page does not exist)

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Group 17 hydride

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Iron dihydride

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https://af.wikipedia.org/wiki/Ammoniak
Ammoniak – Afrikaans
af
af
interlanguage-link-target

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Ammonyak – Azerbaijani
az
az
interlanguage-link-target

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Amonijak – Bosnian
bs
bs
interlanguage-link-target

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https://cs.wikipedia.org/wiki/Amoniak
Amoniak – Czech
cs
cs
interlanguage-link-target

values:
https://sn.wikipedia.org/wiki/Ammonia
Ammonia – Shona
sn
sn
interlanguage-link-target

values:
https://cy.wikipedia.org/wiki/Amonia
Amonia – Welsh
cy
cy
interlanguage-link-target

values:
https://da.wikipedia.org/wiki/Ammoniak
Ammoniak – Danish
da
da
interlanguage-link-target

values:
https://de.wikipedia.org/wiki/Ammoniak
Ammoniak – German
de
de
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https://et.wikipedia.org/wiki/Ammoniaak
Ammoniaak – Estonian
et
et
interlanguage-link-target

values: Αμμωνία – Greek

values:
https://eo.wikipedia.org/wiki/Amoniako
Amoniako – Esperanto
eo
eo
interlanguage-link-target

values:
https://eu.wikipedia.org/wiki/Amoniako
Amoniako – Basque
eu
eu
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values: آمونیاک – Persian

values:
https://fr.wikipedia.org/wiki/Ammoniac
Ammoniac – French
fr
fr
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values:
https://gom.wikipedia.org/wiki/Ammonia
Ammonia – Goan Konkani
gom
gom
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Amonijak – Croatian
hr
hr
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https://io.wikipedia.org/wiki/Amono
Amono – Ido
io
io
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values:
https://id.wikipedia.org/wiki/Amonia
Amonia – Indonesian
id
id
interlanguage-link-target

values:
https://it.wikipedia.org/wiki/Ammoniaca
Ammoniaca – Italian
it
it
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values: אמוניה – Hebrew

values:
https://jv.wikipedia.org/wiki/Amonia
Amonia – Javanese
jv
jv
interlanguage-link-target

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values:
https://sw.wikipedia.org/wiki/Amonia
Amonia – Swahili
sw
sw
interlanguage-link-target

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https://ku.wikipedia.org/wiki/Amonyak
Amonyak – Kurdish
ku
ku
interlanguage-link-target

values: Аммиак – Kyrgyz

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https://la.wikipedia.org/wiki/Ammonia
Ammonia – Latin
la
la
interlanguage-link-target

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https://lv.wikipedia.org/wiki/Amonjaks
Amonjaks – Latvian
lv
lv
interlanguage-link-target

values:
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Amoniakas – Lithuanian
lt
lt
interlanguage-link-target

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Ammoniak – Limburgish
li
li
interlanguage-link-target

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Ammonia – Malay
ms
ms
interlanguage-link-target

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Ammoniak – Dutch
nl
nl
interlanguage-link-target

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https://no.wikipedia.org/wiki/Ammoniakk
Ammoniakk – Norwegian
no
no
interlanguage-link-target

values:
https://nn.wikipedia.org/wiki/Ammoniakk
Ammoniakk – Norwegian Nynorsk
nn
nn
interlanguage-link-target

values:
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Amoniac – Occitan
oc
oc
interlanguage-link-target

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Ammiak – Uzbek
uz
uz
interlanguage-link-target

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Ammoniak – Low German
nds
nds
interlanguage-link-target

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https://pl.wikipedia.org/wiki/Amoniak
Amoniak – Polish
pl
pl
interlanguage-link-target

values:
https://ro.wikipedia.org/wiki/Amoniac
Amoniac – Romanian
ro
ro
interlanguage-link-target

values: Аммиак – Russian

values:
https://sco.wikipedia.org/wiki/Ammonia
Ammonia – Scots
sco
sco
interlanguage-link-target

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Amoniaku – Albanian
sq
sq
interlanguage-link-target

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values:
https://sk.wikipedia.org/wiki/Amoniak
Amoniak – Slovak
sk
sk
interlanguage-link-target

values:
https://sl.wikipedia.org/wiki/Amonijak
Amonijak – Slovenian
sl
sl
interlanguage-link-target

values: Амонијак – Serbian

values:
https://sh.wikipedia.org/wiki/Amonijak
Amonijak – Serbo-Croatian
sh
sh
interlanguage-link-target

values:
https://su.wikipedia.org/wiki/Amonia
Amonia – Sundanese
su
su
interlanguage-link-target

values:
https://sv.wikipedia.org/wiki/Ammoniak
Ammoniak – Swedish
sv
sv
interlanguage-link-target

values:
https://tl.wikipedia.org/wiki/Amonya
Amonya – Tagalog
tl
tl
interlanguage-link-target

values: அமோனியா – Tamil

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values:
https://tr.wikipedia.org/wiki/Amonyak
Amonyak – Turkish
tr
tr
interlanguage-link-target

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values:
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Amoniac – Vietnamese
vi
vi
interlanguage-link-target

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Amonya – Waray
war
war
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Ammonia – Yoruba
yo
yo
interlanguage-link-target

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氨 – Chinese
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zh
interlanguage-link-target

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