Ammonia

Ammonia
Ammonia

General
Systematic name	Ammonia Azane (see text)
Trivial names	Spirit of hartshorn Nitrosil Vaporole
Molecular formula	NH₃
Molar mass	17.03 g/mol
Appearance	Colourless gas with strong pungent odor
CAS number	[7664-41-7]
Properties
Density and phase	0.6813 g/L, gas
Solubility in water	54 g/100 ml
Melting point	-78.27 °C (195.42 K)
Boiling point	-33.49 °C (240.74 K)
Basicity (pK_b)	4.75
Acidity (pK_a)	approx. 34
Thermodynamic data
Std enthalpy of formation Δ_fH⁰_solid	-45.92 kJ/mol
Standard molar entropy S⁰_solid	192.77 J·K⁻¹·mol⁻¹
Hazards
EU classification	Conc. dependent. See text
R-phrases	Conc. dependent See text
S-phrases	1/2, 16, 36/37/39, 45, 61
NFPA 704
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Regulatory data	Flash point, RTECS number, etc.
Related compounds
Related Amines	See Amine
Related Hydrides	Phosphine Arsine
Related compounds	Hydrazine Hydrazoic acid Hydroxylamine Chloramine
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Ammonia is a compound of nitrogen and hydrogen with the formula NH₃. At standard temperature and pressure ammonia is a gas. It is toxic and corrosive to some materials, and has a characteristic pungent odor.

An ammonia molecule is not flat, but has the shape of a compressed tetrahedron known as a trigonal pyramid, as would be expected from VSEPR theory. This shape gives the molecule an overall dipole moment and makes it polar so that ammonia very readily dissolves in water. The nitrogen atom in the molecule has a lone electron pair, and ammonia acts as a base. In acidic or even neutral aqueous solutions, it can bond to a hydronium ion (H₃O⁺), releasing a water molecule (H₂O) to form the positively charged ammonium ion (NH₄⁺), which has the shape of a regular tetrahedron. The degree to which ammonia forms the ammonium ion depends on the pH of the solution.

The main uses of ammonia are in the production of fertilizers, explosives and polymers. It is also an ingredient in certain household glass cleaners. 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, while ammonium chloride (sal-ammoniac) and ammonium sulfate are found in volcanic districts; crystals of ammonium bicarbonate have been found in Patagonian guano. Ammonium salts also are found distributed through all fertile soil and in seawater.

History

Salts of ammonia have been known from very early times; thus the term Hammoniacus sal appears in the writings of Pliny, although it is not known whether the term is identical with the more modern sal-ammoniac.

In the form of sal-ammoniac, ammonia was known to the alchemists as early as 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, Basil Valentine showed that ammonia could be obtained by the action of alkalis on sal-ammoniac. At a later period, when sal-ammoniac was obtained by distilling the hoofs and horns of oxen and neutralizing the resulting carbonate with hydrochloric acid, the name Spirit of hartshorn was applied to ammonia.

Gaseous ammonia was first isolated by Joseph Priestley in 1774 and was termed by him alkaline air. In 1777 Karl Wilhelm Scheele showed that it contained nitrogen, and Claude Louis Berthollet, in about 1785, ascertained its composition.

The Haber process to produce ammonia from the nitrogen contained 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. The ammonia was used to produce explosives to sustain their war effort.

Production

Because of its many uses, ammonia is one of the most highly-produced inorganic chemicals. Before the start of WWI most ammonia was obtained by the dry distillation of nitrogenous vegetable and animal products; by the reduction of nitrous acid and nitrites with nascent hydrogen; and also by the decomposition of ammonium salts by alkaline hydroxides or by quicklime, the salt most generally used being the chloride (sal-ammoniac) thus

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

It has also been obtained by decomposing magnesium nitride (Mg₃N₂) with water,

Mg₃N₂ + 6H₂O → 3Mg(OH)₂ + 2NH₃

Today the Haber process is the most important method for production of ammonia. In this process, nitrogen and hydrogen gases combine directly on an iron catalyst at a pressure of 200 bar (20 MPa, 3000 lbf/in²) and a temperature of 500 °C to produce ammonia.

N₂ + 3H₂ → 2 NH₃

Compared to older methods, the feedstocks of the Haber process are relatively inexpensive—nitrogen makes up 78% of the atmosphere, while hydrogen can be readily produced from natural gas.

Properties

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 and the liquid boils at -33.7 °C, and solidifies at -75 °C to a mass of white crystals. Liquid ammonia possesses strong ionizing powers (ε = 22), and solutions of salts in liquid ammonia have been much studied. Liquid ammonia has a very high enthalpy change of vaporization (23.35 kJ/mol, c.f. 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 at room temperature, even though it is well above its boiling point.

It is miscible with water. All the ammonia contained in an aqueous solution of the gas may 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 cm^-3 and is often known as '.880 Ammonia'.

It does not sustain combustion, and it does not burn readily unless mixed with oxygen, when 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. Chlorine catches fire when passed into ammonia, forming nitrogen and hydrochloric acid; unless the ammonia is present in excess, the highly explosive nitrogen trichloride (NCl₃) is also formed.

The ammonia molecule readily undergoes nitrogen inversion at normal pressures, that is to say that the nitrogen atom passes through the plane of the three hydrogen atoms as if an umbrella turns inside out in strong winds. The energy barrier to this inversion is 24.7 kJ/mol in ammonia, 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 (C. E. Cleeton & N. H. Williams, 1934).

Formation of salts

One of the most characteristic properties of ammonia is its power of combining directly 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 being necessary to bring about the reaction.^[1]

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⁺).

Ammonia as a (very weak) acid

Although ammonia is well-known as a base, ammonia can also act as an extremely weak Bronsted-Lowry acid. It is a protic substance, and is capable of dissociation into the amide (NH₂^-) ion. However, this behavior can never be observed in aqueous solution or in the presence of water, because water is a much stronger Bronsted-Lowry acid than it. An example of ammonia behaving as an acid would be when solid lithium nitride is added to liquid ammonia, forming a lithium amide solution:

Li₃N(s)+ 2 NH₃(l) ---> 3 Li⁺(am) + 3 NH₂^-(am).

This is a Bronsted-Lowry acid-base reaction where ammonia is acting as an acid.

Formation of other compounds

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. Using 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 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 two-fold 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

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:

Hexamminecopper(II), [Cu(NH₃)₆]²⁺, a characteristic dark blue complex formed by adding ammonia to solution of copper(II) salts.
Diamminesilver(I), [Ag(NH₃)₂]⁺, the active species in Tollens' reagent. Formation of this complex can also help to distinguish between precipitates of the different silver halides: AgCl is soluble in dilute (2 M) ammonia solution, AgBr is only soluble in concentrated ammonia solution while AgI is insoluble in aqueous solution of ammonia.

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

An ammine ligand bound to a metal ion is markedly more acidic than a free ammonia molecule, although deprotonation in aqueous solution is still rare. One example is the Calomel reaction, where the resulting amidomercury(II) compound is highly insoluble.

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

Uses

The most important single use of ammonia is in the production of nitric acid. A mixture of one part ammonia to nine parts air is passed over a platinum gauze catalyst at 850 °C, whereupon the ammonia is oxidized to nitric oxide.

4NH₃ + 5O₂ → 4NO + 6H₂O

The catalyst is essential, as the normal oxidation (or combustion) of ammonia gives dinitrogen and water: the production of nitric oxide is an example of kinetic control. As the gas mixture cools to 200–250 °C, the nitric oxide is in turn oxidized by the excess of oxygen present in the mixture, to give nitrogen dioxide. This is reacted with water to give nitric acid for use in the production of fertilizers and explosives.

In addition to serving as a fertilizer ingredient, ammonia can also be used directly as a fertilizer by forming a solution with irrigation water, without additional chemical processing. This later use allows the continuous growing of nitrogen dependent crops such as maize (corn) without crop rotation but this type of use leads to poor soil health.

Ammonia has thermodynamic properties that make it very well suited as a refrigerant, since it liquefies readily under pressure, and was used in virtually all refrigeration units prior to the advent of haloalkanes such as Freon. However, ammonia is a toxic irritant and its corrosiveness to any copper alloys increases the risk that an undesirable leak may develop and cause a noxious hazard. Its use in small refrigeration units has been largely replaced by haloalkanes, which are not toxic irritants and are practically not flammable. (Note: Butane and isobutane, which have very suitable thermodynamic properties for refrigerants, are extremely flammable.) Ammonia continues to be used as a refrigerant in large industrial processes such as bulk icemaking and industrial food processing. Ammonia is also useful as a component in absorption-type refrigerators, which do not use compression and expansion cycles but can exploit heat differences. Since the implication of haloalkane being major contributors to ozone depletion, ammonia is again seeing increasing use as a refrigerant.

Ammonia is a primary ingredient in old-style household cleaners.

It is also sometimes added to drinking water along with chlorine to form chloramine, a disinfectant. Unlike chlorine on its own, chloramine does not combine with organic (carbon containing) materials to form carcinogenic halomethanes such as chloroform.

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 per 100 g)
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 Ca, Sr, Ba Eu and Yb. 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

	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⁻ ⇌ ½H₂ + 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.

Safety precautions

Toxicity

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 synthase, 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.

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. They should never be mixed with chlorine-containing products, for example household bleach, as a variety of toxic and carcinogenic compounds are formed (e.g., chloramine, hydrazine).

Laboratory use of ammonia solutions

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

Concentration by weight	Molarity	Classification	R-Phrases
5–10%	2.87–5.62 mol/L	Irritant (I)	R36/37/38
10–25%	5.62–13.29 mol/L	Corrosive (C)	R34
>25%	>13.29 mol/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") 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 analysis, and should be acidified and diluted before disposal once the test is completed.

Laboratory use of anhydrous ammonia (gas or liquid)

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

Ammonia reacts violently with the halogens, and causes the explosive polymerization of ethylene oxide. It also forms explosive 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.

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.

Note

* The notion of ammonium salts in liquid ammonia is shaky, as they are analogous to strong acid solutions in water. These solutions are not normally referred to as hydronium salt solutions. (See talk page for discussion)

Reference

^ Baker, H. B. (1894). J. Chem. Soc. 65: 612.

Bibliography

This article incorporates text from a publication now in the public domain: Chisholm, Hugh, ed. (1911). Encyclopædia Britannica (11th ed.). Cambridge University Press. {{cite encyclopedia}}: Missing or empty |title= (help)
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External links

NIST Standard Reference Database
European Chemicals Bureau
International Chemical Safety Card 0414 (anhydrous ammonia)
International Chemical Safety Card 0215 (aqueous solutions)
NIOSH Pocket Guide to Chemical Hazards
Institut national de recherche et de securite (in French)
Computational Chemistry Wiki