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Ammonia

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For other uses, see Ammonia (disambiguation).
Ammonia
Names
IUPAC names
Ammonia
Azane (see text)
Other names
Ammonia
Hydrogen nitride
Spirit of Hartshorn
Nitro-Sil
Vaporole
Identifiers
ChemSpider
ECHA InfoCard 100.028.760 Edit this at Wikidata
EC Number
  • 231-635-3
RTECS number
  • BO0875000 (anhydrous)
    BQ9625000 (solutions)
UN number 1005 (anhydrous)
2672, 2073, 3318 (solutions)
  • InChI=1/H3N/h1H3
Properties
NH3
Molar mass 17.031 g/mol
Appearance Colorless gas with strong pungent odor
Density 0.86 kg/m3 (1.013 bar at boiling point)
0.73 kg/m3 (1.013 bar at 15 °C)
681.9 kg/m3 at −33.3 °C (liquid)[1]
820 kg/m3 at -80 °C (crystal solid)
817 kg/m3 at -80 °C (transparent solid)[2]
Melting point −77.73 °C (195.42 K)
Boiling point −33.34 °C (239.81 K)
1176 g/100 mL (0 °C)
702 g/100 mL (20 °C)
88 g/100 mL (100 °C)
Acidity (pKa) 9.75
Basicity (pKb) 4.75
Structure
Trigonal pyramid
1.42 D
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
1
0
Flash point flammable gas (see text)
Explosive limits 15–28%
NIOSH (US health exposure limits):
PEL (Permissible)
 50 ppm
Related compounds
Other cations
 Phosphine
Arsine
Stibine
Supplementary data page
Ammonia (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Ammonia is a compound of nitrogen and hydrogen with the formula NH3. It is normally encountered as a gas with a characteristic pungent odor. Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to foodstuffs and fertilizers. Ammonia, either directly or indirectly, is also a building block for the synthesis of many pharmaceuticals. Although in wide use, ammonia is both caustic and hazardous. In 2006, worldwide production was estimated at 146.5 million tons.[3] It is used in commercial cleaning products.

Ammonia, as used commercially, is often called anhydrous ammonia. This term emphasizes the absence of water in the material. Because NH3 boils at −33.34 °C, the liquid must be stored under high pressure or at low temperature. Its heat of vaporization is, however, sufficiently high so that NH3 can be readily handled in ordinary beakers, in a fume hood. "Household ammonia" or "ammonium hydroxide" is a solution of NH3 in water. The strength of such solutions is measured in units of baume (density), with 26 degrees baume (about 30 weight percent ammonia at 15.5 °C) being the typical high concentration commercial product.[4] Household ammonia ranges in concentration from 5 to 10 weight percent ammonia.

Structure and basic chemical properties

The ammonia molecule has a trigonal pyramidal shape with a bond angle of 107.8° as shown above, as predicted by Valence shell electron pair repulsion theory (VSEPR). The central nitrogen atom has five outer electrons with an additional electron from each hydrogen. This gives a total of eight electrons, or four electron pairs which 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 is measured at 107.8°. 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 so that ammonia readily dissolves in water. The degree to which ammonia forms the ammonium ion increases upon lowering the pH of the solution— at "physiological" pH (~7), about 67% of the ammonia molecules are protonated. Temperature and salinity also affect the proportion of NH4+. NH4+ has the shape of a regular tetrahedron and is isoelectronic with methane.

The main use of ammonia is for production of fertilizer (83% in 2003). Another major application is its conversion to explosives, because nitric acid is made via oxidation of ammonia. The entire nitrogen content of all manufactured organic compounds is derived from ammonia.[3]

Natural occurrence

Ammonia is found in small quantities in the atmosphere, being produced from the putrefaction of nitrogenous animal and vegetable matter. Ammonia and ammonium salts are also found in small quantities in rainwater, whereas ammonium chloride (sal-ammoniac), and ammonium sulfate are found in volcanic districts; crystals of ammonium bicarbonate have been found in Patagonian guano. The kidneys secrete NH3 to neutralize excess acid.[5] Ammonium salts also are found distributed through all fertile soil and in seawater. Substances containing ammonia, or those that are similar to it, are called ammoniacal.

History

The Romans called the ammonium chloride deposits they collected from near the Temple of Jupiter Amun (Greek Ἄμμων Ammon) in ancient Libya 'sal ammoniacus' (salt of Amun) because of proximity to the nearby temple.[6] Salts of ammonia have been known from very early times; thus the term Hammoniacus sal appears in the writings of Pliny, although it is not known whether the term is identical with the more modern sal-ammoniac.[7]

In the form of sal-ammoniac, ammonia was known to the Arabic alchemists as early as the 8th century, first mentioned by Geber (Jabir ibn Hayyan),[8] and to the European alchemists since the 13th century, being mentioned by Albertus Magnus. It was also used by dyers in the Middle Ages in the form of fermented urine to alter the colour of vegetable dyes. In the 15th century, Basilius Valentinus showed that ammonia could be obtained by the action of alkalis on sal-ammoniac. At a later period, when sal-ammoniac was obtained by distilling the hooves and horns of oxen and neutralizing the resulting carbonate with hydrochloric acid, the name "spirit of hartshorn" was applied to ammonia.[9]

Gaseous ammonia was first isolated by Joseph Priestley in 1774 and was termed by him alkaline air; however it was acquired by the alchemist Basil Valentine.[10] Eleven years later in 1785, Claude Louis Berthollet ascertained its composition.

The Haber process to produce ammonia from the nitrogen in the air was developed by Fritz Haber and Carl Bosch in 1909 and patented in 1910. It was first used on an industrial scale by the Germans during World War I,[3] following the allied blockade that cut off the supply of nitrates from Chile. The ammonia was used to produce explosives to sustain their war effort.[11]

Prior to the advent of cheap natural gas, hydrogen as a precursor to ammonia production was produced via the electrolysis of water. The Vemork 60 MW hydroelectric plant in Norway constructed in 1911 was used purely for this purpose and up until the second world war provided the majority of Europe's ammonia.

Synthesis and production

See also: Haber Process

Because of its many uses, ammonia is one of the most highly produced inorganic chemicals. Dozens of chemical plants worldwide produce ammonia. The worldwide ammonia production in 2004 was 109 million metric tonnes.[12] The People's Republic of China produced 28.4% of the worldwide production (increasingly from coal as part of urea synthesis[13]) followed by India with 8.6%, Russia with 8.4%, and the United States with 8.2%.[12] About 80% or more of the ammonia produced is used for fertilizing agricultural crops.[12]

Before the start of World War I, most ammonia was obtained by the dry distillation[14] 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[15] such as quicklime, the salt most generally used being the chloride (sal-ammoniac) thus:

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

(Two molecules of ammonium chloride plus two calcium oxide yields calcium chloride and calcium hydroxide and two molecules of ammonia)

Today, the typical modern ammonia-producing plant first converts natural gas (i.e., methane) or liquified petroleum gas (such gases are propane and butane) or petroleum naphtha into gaseous hydrogen. The process used in producing the hydrogen begins with removal of sulfur compounds from the natural gas (because sulfur deactivates the catalysts used in subsequent steps). Catalytic hydrogenation converts organosulfur compounds into gaseous hydrogen sulfide:

H2 + RSH → RH + H2S(g)
  • The hydrogen sulfide is then removed by passing the gas through beds of zinc oxide where it is adsorbed and converted to solid zinc sulfide:
H2S + ZnO → ZnS + H2O
CH4 + H2O → CO + 3 H2
CO + H2O → CO2 + H2
  • The final step in producing the hydrogen is to use catalytic methanation to remove any small residual amounts of carbon monoxide or carbon dioxide from the hydrogen:
CO + 3 H2 → CH4 + H2O
CO2 + 4 H2 → CH4 + 2 H2O
  • To produce the desired end-product ammonia, the hydrogen is then catalytically reacted with nitrogen (derived from process air) to form anhydrous liquid ammonia. This step is known as the ammonia synthesis loop (also referred to as the Haber-Bosch process):
3 H2 + N2 → 2 NH3

Hydrogen required for ammonia synthesis could also be produced economically using other sources like coal or coke gasification, less economically from the electrolysis of water into oxygen + hydrogen and other alternatives which 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.

Biosynthesis

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

Ammonia is also a metabolic product of amino acid deamination. Ammonia excretion is common in aquatic animals. In humans, it is quickly converted to urea, which is much less toxic. This urea is a major component of the dry weight of urine. Most reptiles, birds, as well as insects and snails solely excrete uric acid as nitrogenous waste.

Properties

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

It is miscible with water. Ammonia in an aqueous solution can be expelled by boiling. The aqueous solution of ammonia is basic. The maximum concentration of ammonia in water (a saturated solution) has a density of 0.880 g /cm³ and is often known as '.880 Ammonia'. Ammonia does not burn readily or sustain combustion, except under narrow fuel-to-air mixtures of 15-25% air. When mixed with oxygen, it burns with a pale yellowish-green flame. At high temperature and in the presence of a suitable catalyst, ammonia is decomposed into its constituent elements. Ignition occurs when chlorine is passed into ammonia, forming nitrogen and hydrogen chloride; if ammonia is present in excess, then the highly explosive nitrogen trichloride (NCl3) is also formed.

The ammonia molecule readily undergoes nitrogen inversion at room temperature; a useful analogy is an umbrella turning itself inside out in a strong wind. The energy barrier to this inversion is 24.7 kJ/mol, and the resonance frequency is 23.79 GHz, corresponding to microwave radiation of a wavelength of 1.260 cm. The absorption at this frequency was the first microwave spectrum to be observed.[16]

Basicity

One of the most characteristic properties of ammonia is its basicity. It combines with acids to form salts; thus with hydrochloric acid it forms ammonium chloride (sal-ammoniac); with nitric acid, ammonium nitrate, etc. However, perfectly dry ammonia will not combine with perfectly dry hydrogen chloride: moisture is necessary to bring about the reaction.[17]

NH3 + HCl → NH4Cl

The salts produced by the action of ammonia on acids are known as the ammonium salts and all contain the ammonium ion (NH4+). Anhydrous ammonia is often used for the production of methamphetamine. Aqueous ammonia can be applied on the skin to lessen the effects of acidic animal poisons, such as from insect and jellyfish.

Acidity

Although ammonia is well-known as a base, it can also act as an extremely weak acid. It is a protic substance and is capable of formation of amides (which contain the NH2− ion), for example lithium and ammonia react to give a solution of lithium amide:

2 Li + 2 NH3 → 2 LiNH2 + H2

Combustion

The combustion of ammonia to nitrogen and water is exothermic:

4 NH3 + 3 O2 → 2 N2 + 6 H2O (g) (ΔHºr = –1267.20 kJ/mol)

The standard enthalpy change of combustion, ΔHºc, 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 nitrogen and oxygen, which is the principle behind the catalytic converter. However, 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

The combustion of ammonia in air is very difficult in the absence of a catalyst (such as platinum gauze), as the temperature of the flame is usually lower than the ignition temperature of the ammonia-air mixture. The flammable range of ammonia in air is 16–25%.[18]

Formation of other compounds

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

Amides can be prepared by the reaction of ammonia with a number of carboxylic acid derivatives. Acyl chlorides are the most reactive, but the ammonia must be present in at least a twofold excess to neutralise the hydrogen chloride formed. Esters and anhydrides also react with ammonia to form amides. Ammonium salts of carboxylic acids can be dehydrated to amides so long as there are no thermally sensitive groups present: temperatures of 150–200 °C are required.

The hydrogen in ammonia is capable of replacement by metals, thus magnesium burns in the gas with the formation of magnesium nitride Mg3N2, and when the gas is passed over heated sodium or potassium, sodamide, NaNH2, and potassamide, KNH2, are formed. Where necessary in substitutive nomenclature, IUPAC recommendations prefer the name azane to ammonia: hence chloramine would be named chloroazane in substitutive nomenclature, not chloroammonia.

Ammonia as a ligand

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

Ammonia can act as a ligand in transition metal complexes. It is a pure σ-donor, in the middle of the spectrochemical series, and shows intermediate hard-soft behaviour. For historical reasons, ammonia is named ammine in the nomenclature of coordination compounds. Some notable ammine complexes include:

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

Ammine complexes of chromium(III) were known in the late 19th century, and formed the basis of Alfred Werner's theory of coordination compounds. Werner noted that only two isomers (fac- and mer-) of the complex [CrCl3(NH3)3] could be formed, and concluded that the ligands must be arranged around the metal ion at the vertices of an octahedron. This 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−

Rotational spectrum

The energy level diagram depicting levels below 350 cm−1 for ammonia is in the figure to the right. Energies for J=0 to J=5 are included. Note the decreasing value of the energy with increasing value of K along a single value of J. This is due to the fact that ammonia is an oblate symmetric top, meaning A=B<C where A, B, and C are inversely related to the angular momenta along each molecular axis. The quantum numbers of each level are provided to the right of the level and the percentage of the molecule in each state assuming LTE at 30 K is given above each level when statistically significant.

A spectral simulation is provided below for ammonia at 30 K. Transitions of the lines are indicated as follows: (J1,K1) - (J2,K2). The J1=2,3,4 lines are too close together to distinguish individually. For clarity, the J1=4 lines have been expanded. A spectral line list of this simulation is below for J1<6. Frequencies are given in units of cm−1.

Frequency (cm−1) J1 K1 J2 K2
19.89634259 1 0 0 0
39.81313621 2 0 1 0
39.8067724 2 1 1 1
59.73269224 3 2 2 2
59.7614132 3 1 2 1
59.7710093 3 0 2 0
79.67548414 4 3 3 3
79.67548414 4 3 3 3
79.7394795 4 2 3 2
79.77805626 4 1 3 1
79.7909451 4 0 3 0
99.63650816 5 4 4 4
99.74877258 5 3 4 3
99.74877258 5 3 4 3
99.82952252 5 2 4 2
99.8781969 5 1 4 1
99.8944591 5 0 4 0

A spectrum is also provided of ammonia a T=300 K for comparison. The frequencies of the simulation have been compared to and match Poynter & Kakar (1975).[19]

These spectra and energy level diagrams were produced using the molecular spectrum simulator package PGopher. The constants used are taken from Table 4 of Poynter et al. (1975).[19]

Interstellar formation and destruction

Formation mechanisms

The interstellar abundance for ammonia has been measured for a variety of environments. The [NH3]/[H2] ratio has been estimated to range from 10−7 in small dark clouds (cf. Ungerechts et al. 1980) up to 10−5 in the dense core of the Orion Molecular Cloud (Genzel et al. 1982). Although a total of 18 total production routes have been proposed,[20] 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 (see Vikor, Al-Khalili, Danared et al., 1999, A&A, 344, 1027). The rate constant was calculated from the formula . For the primary formation reaction, a = 1.05×10−6 and B = −0.47. Assuming an NH4+ abundance of 3×10−7 (van Dishoeck & Black 1986) and an electron abundance of 10−7 typical of molecular clouds, the formation will proceed at a rate of 1.6×10−9 cm−3s−1 in a molecular cloud of total density 105 cm−3.

All other proposed formation reactions have rate constants of between 2 and 13 orders of magnitude smaller, making their contribution to the abundance of ammonia relatively insignificant.[21] 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 that the primary reaction above.

Some of the other possible formation reactions are:

H− + NH4+ → NH3 + H2
PNH3+ + e− → P + NH3

Destruction mechanisms

According to the online database UDFA.net, there are 113 total proposed reactions leading to the destruction of NH3. Of these, 39 were tabulated by Prasad and Huntress [22] who compiled extensive tables of the chemistry among C, N, and O compounds. A review of interstellar ammonia by Ho & Townes (1983) cites the following reactions as the principal dissociation mechanisms:[23]

(5) NH3 + H3+ → NH4+ + H2

(6) NH3 + HCO+ → NH4+ + CO

with rate constants a of 4.39×10−9 (Lininger et al. 1975) and 2.2×10−9 (Smith & Adams 1977), respectively. For both reactions, B and gamma are 0, therefore k = a. Equation (5) and (6) 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 (Lepp et al. 1986), [HCO+]/[H2] = 2×10−9 (Wooten et al. 1980), and total densities of n = 105, typical of cold, dense, molecular clouds. Clearly, between these two primary reactions, equation (5) is the dominant destruction reaction, with a rate ~10,000 times faster than equation (6). This is due to the relatively high abundance of H3+.

Uses

Fertilizer

Approximately 83% (as of 2003) of ammonia is used as fertilizers either as its salts or as solutions. Consuming more than 1% of all man-made power, the production of ammonia is a significant component of the world energy budget.[3]

Cleaner

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 or caked-on grime. Household ammonia ranges in concentration from 5 to 10 weight percent ammonia.

Precursor to nitrogenous compounds

Ammonia is directly or indirectly the precursor to most nitrogen-containing compounds. Practically all synthetic and all inorganic nitrogen compounds are prepared from ammonia. An important derivative is nitric acid. This key material is generated via the Ostwald process by oxidisation of ammonia with air over a platinum catalyst at 700 - 850 °C, ~9 atm. Nitric oxide is an intermediate:[24]

NH3 + 2 O2 → HNO3 + H2O

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

Minor and emerging uses

Refrigeration - R717

Ammonia's thermodynamic properties made it one of the refrigerants commonly used prior to the discovery of dichlorodifluoromethane, normally known as Freon.[25] Ammonia's toxicity complicates this application. Anhydrous ammonia is widely used in industrial refrigeration applications because of its high energy efficiency and low cost. Ammonia is used less frequently in commercial applications, such as in grocery store freezer cases and refrigerated displays due to its earlier mentioned toxicity.

For remediation of gaseous emissions

Ammonia used to scrub SO2 from the burning of fossil fuels, the resulting product is converted to ammonium sulfate for use as fertilizer. Ammonia neutralizes the nitrogen oxides (NOx) pollutants emitted by diesel engines. This technology, called SCR (selective catalytic reduction), relies on a vanadia-based catalyst.[26]

As a fuel

Ammonia was used during World War II to power buses in Belgium, and in engine and solar energy applications prior to 1900. Liquid ammonia was used as the fuel of the rocket airplane, the X-15. Although not as powerful as other fuels, it left no soot in the reusable rocket engine and its density approximately matches that for the oxidizer, liquid oxygen, which simplified the aircraft's design.

As a vehicle fuel

Ammonia has been proposed as a practical alternative to fossil fuel for internal combustion engines.[27] The calorific value of ammonia is 22.5 MJ/kg (9690 BTU/lb) which is about half that of diesel. In a normal engine, in which the water vapor is not condensed, the calorific value of ammonia will be about 21% less than this figure. It can be used in existing engines with only minor modifications to carburetors/injectors.

To meet these demands, significant capital would be required to increase present production levels. Although 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. It is however significantly less efficient than batteries. The 60 MW Rjukan dam in Telemark Norway, produced ammonia via electrolysis of water for many years from 1913 producing fertilizer for much of Europe. If produced from coal, the CO2 can be readily sequestrated.[27][28] (the combustion products are nitrogen and water). In 1981 a Canadian company converted a 1981 Chevrolet Impala to operate using ammonia as fuel.[29][30]

Textile

Liquid ammonia is used for treatment of cotton materials, give a properties like mercerisation using alkalies. And also used for pre-washing of wool.[citation needed]

Ammonia's role in biological systems and human disease

Ammonia is an important source of nitrogen for living systems. Although atmospheric nitrogen abounds, few living creatures are capable of utilizing this nitrogen. Nitrogen is required for the synthesis of amino acids, which are the building blocks of protein. Some plants rely on ammonia and other nitrogenous wastes incorporated into the soil by decaying matter. Others, such as nitrogen-fixing legumes, benefit from symbiotic relationships with rhizobia which create ammonia from atmospheric nitrogen.[31]

Main symptoms of hyperammonemia (ammonia reaching toxic concentrations).[32]

Ammonia also plays a role in both normal and abnormal animal physiology. Ammonia is biosynthesized through normal amino acid metabolism and is toxic in high concentrations.[33] 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.[34]

Reference ranges for blood tests, comparing blood content of ammonia (shown in yellow near middle) with other constituents.

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.[35]

Excretion

Main article: Excretion

Ammonium ions are a toxic waste product of the metabolism in animals. In fishes and aquatic invertebrates, it is excreted directly into the water. In mammals, sharks, and amphibians, it is converted in the urea cycle to urea, because it is less toxic and can be stored more efficiently. In birds, reptiles, and terrestrial snails, metabolic ammonium is converted into uric acid, which is solid, and can therefore be excreted with minimal water loss.[36]

Theoretical role in alternative biochemistry

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

Liquid ammonia as a solvent

See also: Inorganic nonaqueous solvent

Liquid ammonia is the best-known and most widely studied non-aqueous ionizing solvent. Its most conspicuous property is its ability to dissolve alkali metals to form highly coloured, electrically conducting solutions containing solvated electrons. Apart from these remarkable solutions, much of the chemistry in liquid ammonia can be classified by analogy with related reactions in aqueous solutions. Comparison of the physical properties of NH3 with those of water shows that NH3 has the lower melting point, boiling point, density, viscosity, dielectric constant and electrical conductivity; this is due at least in part to the weaker H bonding in NH3 and the fact that such bonding cannot form cross-linked networks since each NH3 molecule has only 1 lone-pair of electrons compared with 2 for each H2O molecule. The ionic self-dissociation constant of liquid NH3 at −50 °C is approx. 10−33 mol2·l−2.

Solubility of salts

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

Liquid ammonia is an ionizing solvent, although less so than water, and dissolves a range of ionic compounds including many nitrates, nitrites, cyanides and thiocyanates. Most ammonium salts are soluble, and these salts act as acids in liquid ammonia solutions. The solubility of halide salts increases from fluoride to iodide. A saturated solution of ammonium nitrate contains 0.83 mol solute per mole of ammonia, and has a vapour pressure of less than 1 bar even at 25 °C.

Solutions of metals

See also: Solvated electron, metallic solution

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

These solutions are very useful as strong reducing agents. At higher concentrations, the solutions are metallic in appearance and in electrical conductivity. At low temperatures, the two types of solution can coexist as immiscible phases.

Redox properties of liquid ammonia

See also: Redox.
  E° (V, ammonia) E° (V, water)
Li+ + e− ⇌ Li −2.24 −3.04
K+ + e− ⇌ K −1.98 −2.93
Na+ + e− ⇌ Na −1.85 −2.71
Zn2+ + 2e− ⇌ Zn −0.53 −0.76
NH4+ + e− ⇌ ½ H2 + NH3 0.00 –
Cu2+ + 2e− ⇌ Cu +0.43 +0.34
Ag+ + e− ⇌ Ag +0.83 +0.80

The range of thermodynamic stability of liquid ammonia solutions is very narrow, as the potential for oxidation to dinitrogen, E° (N2 + 6NH4+ + 6e− ⇌ 8NH3), is only +0.04 V. In practice, both oxidation to dinitrogen and reduction to dihydrogen are slow. This is particularly true of reducing solutions: the solutions of the alkali metals mentioned above are stable for several days, slowly decomposing to the metal amide and dihydrogen. Most studies involving liquid ammonia solutions are done in reducing conditions: although oxidation of liquid ammonia is usually slow, there is still a risk of explosion, particularly if transition metal ions are present as possible catalysts.

Detection and determination

Ammonia and ammonium salts can be readily detected, in very minute traces, by the addition of Nessler's solution, which gives a distinct yellow coloration in the presence of the least trace of ammonia or ammonium salts. Sulfur sticks are burnt to detect small leaks in industrial ammonia refrigeration systems. Larger quantities can be detected by warming the salts with a caustic alkali or with quicklime, when the characteristic smell of ammonia will be at once apparent. The amount of ammonia in ammonium salts can be estimated quantitatively by distillation of the salts with sodium or potassium hydroxide, the ammonia evolved being absorbed in a known volume of standard sulfuric acid and the excess of acid then determined volumetrically; or the ammonia may be absorbed in hydrochloric acid and the ammonium chloride so formed precipitated as ammonium hexachloroplatinate, (NH4)2PtCl6.

Interstellar space

Ammonia was first detected in interstellar space in 1968, based on microwave emissions from the direction of the galactic core.[37] 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.[23] 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.[38] The ammonia molecule has also been detected in the atmospheres of the gas giant planets, including Jupiter, along with other gases like methane, hydrogen, and helium. The interior of Saturn may include frozen crystals of ammonia.[39]

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.

Single Antenna Detections

  • Wilson et al. 1979[40] Radio observations of NH3 from the 100 m Effelsberg Telescope are reported. 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.
  • MacDonald et al. 1981[41] Radio observations of NH3 from the 25 m Chilbolton telescope in England are presented. Among the observations are 35 new detections of ammonia in HII 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.
  • Morris et al. 1973[42] Microwave radiation from ammonia was observed in several galactic objects including W3(OH), Orion A, W43, W51, and five sources in the galactic center. 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.

Interferometric studies

  • Torrelles et al. 1985[43] VLA observations of NH3 in seven regions with high-velocity gaseous outflows is presented. Condensations of less than 0.1 pc were detected in L1551, S140, and Cepheus A. Three individual condensations were detected in Cepheus A, one of them with a highly elongated shape. These condensations may play an important role in creating the bipolar outflow in the region.
  • Ho et al. 1990[44] Extragalactic ammonia is imaged using the VLA in IC 342. The hot gas has temperatures above 70 K inferred from ammonia line ratios and appears to be closely associated with the innermost portions of the nuclear bar seen in CO.
  • Cesaroni et al. 1994[45] The authors present VLA measurements of NH3 towards 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.

Infrared detections

  • Knacke et al. 1982[46] The authors report the detection of absorption at 2.97 micrometres due to solid ammonia on interstellar grains in the Becklin-Neugebauer object and probably in NGC 2264-IR as well. This detection helps explain the physical shape of previously poorly-understood related ice absorption lines.
  • Orton et al. 1982[47] 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.
  • Benson and Meyers 1989[48] 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. The cores are not generally 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.
  • Mebold et al. 1987[49] NH3 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 birth places for the Population I metallicity B-type stars in the galactic halo which could have been borne in the galactic disk.

Astronomical observations and research applications

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.

Observations of nearby dark clouds

By balancing and stimulated emission with spontaneous emission, it is possible to construct a relation between excitation temperature and density. Moreover, since the transitional levels of ammonia can be approximated by a 2-level system at low temperatures, this calculation is fairly simple. This premise can be applied to dark clouds, regions suspected of having extremely low temperatures and possible sites for future star formation. Detections of ammonia in dark clouds show very narrow lines – indicative not only of low temperatures, but also of a low level of inner-cloud turbulence. Line ratio calculations provide a measurement of cloud temperature that is independent of previously-done CO observations. The ammonia observations were consistent with CO measurements of rotation temperatures of ~10 K. With this, densities can be determined, and have been calculated to range between 104 and 105 cm−3 in dark clouds. Mapping of NH3 gives typical clouds sizes of 0.1 pc and masses near 1 solar mass. These cold, dense cores are the sites of future star formation.

UC HII regions

Ultra-compact HII regions are among the best tracers of high-mass star formation. The dense material surrounding UCHII regions is likely primarily molecular. Since a complete study of massive star formation necessarily involves the cloud from which the star formed, ammonia is an invaluable tool in understanding this surrounding molecular material. Since this molecular material can be spatially resolved, it is possible to constrain the heating/ionizing sources, temperatures, masses, and sizes of the regions. Doppler-shifted velocity components allow for the separation of distinct regions of molecular gas which can trace outflows and hot cores originating from forming stars.

Extragalactic detection

NH3 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).

Safety precautions

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.[50] Exposure to very high concentrations of gaseous ammonia can result in lung damage and death.[50] 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).[51]

Toxicity and storage information

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

The toxicity of ammonia solutions does not usually cause problems for humans and other mammals, as a specific mechanism exists to prevent its build-up in the bloodstream. Ammonia is converted to carbamoyl phosphate by the enzyme carbamoyl phosphate synthetase, and then enters the urea cycle to be either incorporated into amino acids or excreted in the urine. However fish and amphibians lack this mechanism, as they can usually eliminate ammonia from their bodies by direct excretion. Ammonia even at dilute concentrations is highly toxic to aquatic animals, and for this reason it is classified as dangerous for the environment. Ammonium compounds should never be allowed to come in contact with bases (unless in an intended and contained reaction), as dangerous quantities of ammonia gas could be released.

Household use

Solutions of ammonia (5–10% by weight) are used as household cleaners, particularly for glass. These solutions are irritating to the eyes and mucous membranes (respiratory and digestive tracts), and to a lesser extent the skin. Caution should be used that the chemical is never mixed into any liquid containing bleach, or a poisonous gas may result. Mixing with chlorine-containing products or strong oxidants, for example household bleach can lead to hazardous compounds such as chloramines.[52]

Laboratory use of ammonia solutions

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

Concentration
by weight (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) Template:R36/37/38
10–25% 5.62–13.29 mol/L 95.7–226.3 g/L Corrosive (C) Template:R34
>25% >13.29 mol/L >226.3 g/L Corrosive (C)
Dangerous for
the environment (N) 		Template:R34, Template:R50
S-Phrases: Template:S1/2, Template:S16, Template:S36/37/39, Template:S45, Template:S61.

The ammonia vapour from concentrated ammonia solutions is severely irritating to the eyes and the respiratory tract, and these solutions should only be handled in a fume hood. Saturated ("0.880") solutions can develop a significant pressure inside a closed bottle in warm weather, and the bottle should be opened with care: this is not usually a problem for 25% ("0.900") solutions.

Ammonia solutions should not be mixed with halogens, as toxic and/or explosive products are formed. Prolonged contact of ammonia solutions with silver, mercury or iodide salts can also lead to explosive products: such mixtures are often formed in qualitative chemical analysis, and should be lightly acidified but not concentrated (<6%w/v) before disposal once the test is completed.

Laboratory use of anhydrous ammonia (gas or liquid)

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

Ammonia reacts violently with the halogens. Nitrogen triiodide, a primary high explosive, is formed when ammonia comes in contact with iodine. It causes the explosive polymerization of ethylene oxide. It also forms explosive fulminating compounds with compounds of gold, silver, mercury, germanium or tellurium, and with stibine. Violent reactions have also been reported with acetaldehyde, hypochlorite solutions, potassium ferricyanide and peroxides.

See also

References

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  43. ^ Torrelles, J.M., Ho, P.T.P., Rodriguez, L.F., Canto, J. (1985). APJ. 288: 595 http://adsabs.harvard.edu/abs/1985ApJ...288..595T. {{cite journal}}: Missing or empty |title= (help)CS1 maint: multiple names: authors list (link)
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  45. ^ Cesaroni, R., Churchwell, E., Hofner, P., Walmsley, C. M., Kurtz, S. (1994). AAP. 288: 903 http://adsabs.harvard.edu/abs/1994A%26A...288..903C. {{cite journal}}: Missing or empty |title= (help)CS1 maint: multiple names: authors list (link)
  46. ^ Knacke, R.F., Mc Corkle, S., Puetter, R.C., Erickson, E.F., Kraetschmer, W. (1982). APJ. 260: 141 http://adsabs.harvard.edu/abs/1982ApJ...260..141K. {{cite journal}}: Missing or empty |title= (help)CS1 maint: multiple names: authors list (link)
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  50. ^ a b "Toxic FAQ Sheet for Ammonia" (PDF). Agency for Toxic Substances and Disease Registry (ATSDR). September 2004.
  51. ^ Hazardous Materials (HM) Safety Permits from the website of the United States Department of Transportation (DOT)
  52. ^ 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)

Bibliography

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