The Online Encyclopedia and Dictionary







Chemical formula
Molecular weight 17.03 u
Appearance colorless gas
CAS number 7664-41-7
Other names
  • anhydrous ammonia
  • aromatic ammonia
  • nitro-sil
  • spirit of hartshorn
  • Vaporole
Physical properties
Density and phase at STP 0.7714 kg/m3 (gas)
Solubility miscible with water
Specific gravity (gas vs. air) 0.63
Crystal structure  ?
pH (10% solution in water)
Thermal decomposition  ? K (? C)
Phase behavior
Melting point 195 K (−78 C)
Boiling point 240 K (−33 C)
Triple point 195.40 K (−77.75 C)
6.07 kPa
Critical point 405.5 K (132.4 C) (270.3 F)
11.42 MPa
Heat of fusion
5.655 kJ/mol
Entropy of fusion
 ? J/(mol·K)
Heat of vaporization
23.350 kJ/mol
Ingestion Dangerous. Symptoms include nausea and vomiting; damage to lips, mouth, and esophagus.
Inhalation Vapors are extremely irritating and corrosive.
Skin Concentrated solutions may produce severe burns and necrosis.
Eyes May cause permanent damage, even in small quantities.
Flash point 11 C
Autoignition temperature 651.1 C
Explosive limits 16 to 25%
OSHA Permissible Exposure Limit
50 ppmv
NIOSH Immediate Danger to Life and Health
500 ppmv
Solid properties
Standard enthalpy change of formation
−46.1 kJ/mol
Heat capacity
 ? J/(mol·K)
Density  ? g/cm3
Liquid properties
ΔfH0liquid −40.2 kJ/mol
Cp  ? J/(mol·K)
Density 8.0 g/cm3
Gas properties
ΔfH0gas −45.9 kJ/mol
Standard molar entropy
192.77 J/(molK)
Cp  ? J/(mol·K)

SI units and standard conditions used unless otherwise stated.
Disclaimer and references

Ammonia is a chemical compound with the formula NH3. The molecule is not flat, but has the shape of a flattened tetrahedron known as a trigonal pyramid. In solution it forms the positively charged ammonium ion NH4, which has the shape of a regular tetrahedron.

At standard temperature and pressure, ammonia is a gas with a characteristic pungent odor. Its main uses are in the production of fertilizers, explosives and polymers. 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 is 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 freons. This use was largely replaced in small refrigeration units by freons as these are not toxic irritants as is ammonia, while continuing in use 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 freons being major contributors to ozone depletion, ammonia is again seeing increasing use as a refrigerant.

Ammonia is found in small quantities as the carbonate in the atmosphere, being produced from the putrefaction of nitrogenous animal and vegetable matter. Ammonium salts are also found in small quantities in rain-water, while ammonium chloride (sal-ammoniac) and ammonium sulfate are found in volcanic districts; and crystals of ammonium bicarbonate have been found in Patagonian guano. Ammonium salts also are found distributed through all fertile soil, in sea-water, and in most plant and animal liquids, and also in urine.



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 unslaked lime (quicklime), the salt most generally used being the chloride (sal-ammoniac) thus

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

A similar reaction yields

2NH4Cl + CaO → CaCl2 + H2O +2NH3

It was also 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 high pressure of 3000 lbf/in² (20 MPa) and temperature (500 C) to produce ammonia.

N2 + 3 H2 → 2 NH3

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


Ammonia is a colourless gas possessing a characteristic pungent smell and a strongly alkaline reaction; 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, and solutions of salts in liquid ammonia have been much studied. Liquid ammonia has a very high specific heat capacity 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 very basic, and since it is a weak electrolyte, the solution will contain a small amount of ammonium hydroxide (NH4OH). The maximum concentration of ammonia in water (a saturated solution) has a density of 880 kg m-3 and is often known as '880 Ammonia'.

It does not support combustion, and it does not burn readily unless mixed with oxygen, when it burns with a pale yellowish-green flame. However it can form an explosive mixture with air.

Household use danger

As both ammonia in water solution and chlorine bleach (sodium hypochlorite in water solution) are common household cleaners there is considerable danger that these may be used in combination in order to obtain a more active cleaning agent. This is extremely dangerous as the combination of the two solutions will release the toxic gas chloramine. The intentional combination and reaction of chlorine with ammonia is used in the industrial production of chloramine, a product useful as a disinfectant of water supplies to be distributed for household use. Unlike the use of gaseous chlorine for this purpose it does not combine with organic (carbon containing) materials to form halomethanes such as carbon tetrachloride, which is a long–term hazard to human health, even in minuscule concentrations, as it is considered to be carcinogenic.


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. It is to be noted that H. B. Baker (Journal of Chem. Soc., 1894, lxv. p. 612) has shown that perfectly dry ammonia will not combine with perfectly dry hydrochloric acid, moisture being necessary to bring about the reaction.

The salts produced by the action of ammonia on acids are known as the ammonium salts and all contain the compound radical ammonium (NH4+). Numerous attempts have been made to isolate this radical, but so far none have been successful. By the addition of sodium amalgam to a concentrated solution of ammonium chloride, the so-called ammonium amalgam is obtained as a spongy mass which floats on the surface of the liquid; it decomposes readily at ordinary temperatures into ammonia and hydrogen; it does not reduce silver and gold salts, a behaviour which distinguishes it from the amalgams of the alkali metals, and for this reason it is regarded by some chemists as being merely mercury inflated by gaseous ammonia and hydrogen. M. le Blanc has shown, however, that the effect of ammonium amalgam on the magnitude of polarization of a battery is comparable with that of the amalgams of the alkali metals.

Ammonium bromide, NH4Br, can be prepared by the direct action of bromine on ammonia. It crystallizes in colourless prisms, possessing a saline taste; it sublimes on heating and is easily soluble in water. On exposure to air it gradually assumes a yellow colour and becomes acid in its reaction.

Ammonium chloride, NH4Cl. (See sal-ammoniac.)

Ammonium fluoride, NH4F, may be obtained by neutralizing ammonia with hydrofluoric acid. It crystallizes in small prisms, having a sharp saline taste, and is exceedingly soluble in water. It decomposes silicates on being heated with them.

Ammonium iodide, NH4I, can be prepared by the action of hydriodic acid on ammonia. It is easily soluble in water, from which it crystallizes in cubes, and also in alcohol. It gradually turns yellow on standing in moist air, owing to decomposition with liberation of iodine.

Ammonium chlorate, NH4ClO3, is obtained by neutralizing chloric acid with either ammonia or ammonium carbonate, or by precipitating barium, strontium or calcium chlorates with ammonium carbonate. It crystallizes in small needles, which are readily soluble in water, and on heating, decompose at about 102 C, with liberation of nitrogen, chlorine and oxygen. It is soluble in dilute aqueous alcohol, but insoluble in strong alcohol.

Ammonium carbonates. The commercial salt was formerly known as sal-volatile or salt of hartshorn and was formerly obtained by the dry distillation of nitrogenous organic matter such as hair, horn, decomposed urine, etc., but is now obtained by heating a mixture of sal-ammoniac, or ammonium sulfate and chalk, to redness in iron retorts, the vapours being condensed in leaden receivers. The crude product is refined by sublimation, when it is obtained as a white fibrous mass, which consists of a mixture of hydrogen ammonium carbonate, H4"HCO3, and ammonium carbamate, NH2COONH4, in molecular proportions; on account of its possessing this constitution it is sometimes called ammonium sesquicarbonate. It possesses a strong ammoniacal smell, and on digestion with alcohol the carbamate is dissolved and a residue of ammonium bicarbonate is left; a similar decomposition taking place when the sesquicarbonate is exposed to air. Ammonia gas passed into a strong aqueous solution of the sesquicarbonate converts it into normal ammonium carbonate, (NH4)2CO3, which can be obtained in the crystalline condition from a solution prepared at about 30 C. This compound on exposure to air gives off ammonia and passes back to ammonium bicarbonate.

Ammonium bicarbonate, NH4HCO3, is formed as shown above and also by passing carbon dioxide through a solution of the normal compound, when it is deposited as a white powder, which has no smell and is only slightly soluble in water. The aqueous solution of this salt liberates carbon dioxide on exposure to air or on heating, and becomes alkaline in reaction. The aqueous solutions of all the carbonates when boiled undergo decomposition with liberation of ammonia and of carbon dioxide:

NH4HCO3 → NH3 + H2O + CO2

It is therefore occasionally used as baking powder, e.g. for gingerbread.

Ammonium nitrate, NH4NO3, is prepared by neutralizing nitric acid with ammonia, or ammonium carbonate, or by double decomposition between potassium nitrate and ammonium sulfate. It can be obtained in three different crystalline forms, the transition points of which are 35 C, 83 C and 125 C. It is easily soluble in water, a considerable lowering of temperature taking place during the operation; on this account it is sometimes used in the preparation of freezing mixtures. On gentle heating, it is decomposed into water and nitrous oxide. Berthelot showed in 1883 that if ammonium nitrate is rapidly heated the following reaction takes place with explosive violence: 2NH4NO3 = 4H2O + 2N2 + O2. In combination with gasoline or other liquid hydrocarbons it is a widely used industrial explosive, being particularly useful in open pit mining and is known as ANFO. The detonation rate is about 3000 feet per second (900 m/s); relatively slow compared to high explosives, which detonate at over 25,000 ft/s (7,600 m/s). This explosive combines the advantages of low cost and stability, requiring a high velocity explosive primer to begin detonation. It is sometimes used in small packets to break up snow cornices in avalanche control. Ammonium nitrate confined in large quantities (such as might be found in a ship's cargo hold) can detonate explosively if combined with hydrocarbons and heated sufficiently by a fire. A fire in a ship carrying ammonium nitrate waterproofed with wax was the cause of a devastating explosion resulting in the Texas City, Texas disaster.

Ammonium nitrite, NH4NO2, is formed by oxidizing ammonia with ozone or hydrogen peroxide; by precipitating barium or lead nitrites with ammonium sulfate, or silver nitrite with ammonium chloride. The precipitate is filtered off and the solution concentrated. It forms colourless crystals which are soluble in water and decompose on heating, with the formation of nitrogen.

Ammonium phosphates. The normal phosphate, (NH4)3PO4,is obtained as a crystalline powder, on mixing concentrated solutions of ammonia and phosphoric acid, or on the addition of excess of ammonia to the acid phosphate (NH4)2HPO4. It is soluble in water, and the aqueous solution on boiling loses ammonia and the acid phosphate NH4H2PO4 is formed. Diammonium hydrogen phosphate, (NH4)2HPO4, is formed by evaporating a solution of phosphoric acid with excess of ammonia. It crystallizes in large transparent prisms, which melt on heating and decompose, leaving a residue of metaphosphoric acid, (HPO3). Ammonium dihydrogen phosphate, NH4•H2PO4, is formed when a solution of phosphoric acid is added to ammonia until the solution is distinctly acid. It crystallizes in quadratic prisms.

Ammonium sodium hydrogen phosphate, NH4"NaHPO4•4H2O. (See microcosmic salt.)

Ammonium sulfate (NH4)2SO4 is prepared commercially from the ammoniacal liquor of gas-works and is purified by recrystallization. It forms large rhombic prisms, has a somewhat saline taste and is easily soluble in water. The aqueous solution on boiling loses some ammonia and forms an acid sulfate. It is used largely as an artificial manure, and also for the preparation of other ammonium salts.

Ammonium persulfate (NH4)2S2O8 has been prepared by H. Marshall (Jour. of Chem. Soc., 1891, lix. p. 777) by the method used for the preparation of the corresponding potassium salt (see sulfur). It is very soluble in cold water, a large fall of temperature accompanying solution. It is a very strong oxidizing agent.

Ammonium sulfide, (NH4)2S, is obtained, in the form of micaceous crystals, by passing hydrogen sulfide mixed with a slight excess of ammonia through a well-cooled vessel; the hydrosulfide NH4•HS is formed at the same time. It dissolves readily in water, but is probably partially dissociated in solution. The hydrosulfide NH4•HS can be obtained as a white solid, by mixing well-cooled ammonia with a slight excess of hydrogen sulfide. According to W. P. Bloxam (Jour. of Chem. Soc., 1895, lxvii. p. 283), if hydrogen sulfide is passed into strong aqueous ammonia at ordinary temperature, the compound (NH4)2S.2NH4HS is obtained, which, on cooling to 0 C and passing more hydrogen sulfide, forms the compound (NH4)2S.12NH4HS. An ice-cold solution of this substance kept at 0 C and having hydrogen sulfide continually passed through it gives the hydrosulfide. Several complex polysulfides of ammonium have been isolated, for details of which see Bloxam's paper quoted above. Compounds are known which may be looked upon as derived from ammonia by the replacement of its hydrogen by the sulfo-group (HSO3); thus potassium ammon-trisulfonate, N(SO3K)3•2H2O, is obtained as a crystalline precipitate on the addition of excess of potassium sulfite to a solution of potassium nitrite, KNO2 + 3K2SO3 + 2H2O = N(SO3K)3 + 4KHO. It can be recrystallized by solution in alkalies. On boiling with water, it is converted, first into the disulfonate NH(SO3K)2 thus, N(SO3K)3 + H2O = NH(SO3K)2 + KHSO4, and ultimately into the monosulfonate NH2•SO3K. The disulfonate is more readily obtained by moistening the nitrilosulfonate with dilute sulfuric acid and letting it stand for twenty-four hours, after which it is recrystallized from dilute ammonia. It forms monosymmetric crystals which by boiling with water yield amidosulfonic acid. (See also E. Divers, Jour. of Chem. Soc., 1892, lxi. p. 943.) Amidosulfonic acid crystallizes in prisms, slightly soluble in water, and is a stable compound.

Other compounds

Ammonia finds a wide application in organic chemistry as a synthetic reagent; it reacts with alkyl iodides to form amines, with esters to form acid amides , with halogen fatty acids to form amino acids; while it also combines with isocyanic esters to form alkyl ureas and with the mustard oils to form alkyl thioureas. Aldehydes also combine directly with ammonia.

Ammonia gas has the power of combining with many substances, particularly with metallic halides; thus with calcium chloride it forms the compound CaCl2•8NH3, and consequently calcium chloride cannot be used for drying the gas. With silver chloride it forms two compounds -- one, AgCl•3NH3 at temperatures below 15 C; the other, 2AgCl•3NH3 at temperatures above 20 C. On heating these substances, ammonia is liberated and the metallic chloride remains. It was by the use of silver chloride ammonia compounds that in 1823 Michael Faraday was first able to liquefy ammonia. It can be shown by Isambert's results that the compound AgCl•3NH3 cannot be formed above 20 C, by the action of ammonia on silver chloride at atmospheric pressure; while 2AgCl•3NH3, under similar conditions, cannot be formed above about 68 C.

Liquid ammonia is used for the artificial preparation of ice. It readily dissolves sodium and potassium, giving in each case a dark blue solution. At a red heat ammonia is easily decomposed into its constituent elements, a similar decomposition being brought about by the passage of electric sparks through the gas. Chlorine takes fire when passed into ammonia, nitrogen, and hydrochloric acid being formed, and unless the ammonia be present in excess, the highly explosive nitrogen trichloride NCl3 is also produced.

With iodine it reacts to form nitrogen iodide. This compound was discovered in 1812 by Bernard Courtois, and was originally supposed to contain nitrogen and iodine only, but in 1840 R.F. Marchand showed that it contained hydrogen, while R. Bunsen showed that no oxygen was present. As regards its constitution, it has been given at different times the formulae NI3, NHI2, NH2I, N2H3I3, etc., these varying results being due to the impurities in the substance, owing to the different investigators working under unsuitable conditions, and also to the decomposing action of light. F. D. Chattaway determined its composition as N2H3I3, by the addition of excess of standard sodium sulfite solution, in the dark, and subsequent titration of the excess of the sulfite with standard iodine. The constitution has been definitely determined by O. Silberrad (Jour. of Chem. Soc., 1905, lxxxvii. p. 55) by the interaction of nitrogen iodide with zinc ethyl, the products of the reaction being triethylamine and ammonia; the ammonia liberated was absorbed in hydrochloric acid, and 95% of the theoretical amount of the ammonium chloride was obtained. On these grounds O. Silberrad assigns the formula NH3•NI3 to the compound, and explains the decomposition as taking place,

2NH3•NI3 +

6Zn(C2H5)2 → 6ZnC2H5•I + 2NH3 + 2N(C2H5)3.

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.

Liquid ammonia as a 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 mp, bp, 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.


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. 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 chlorplatinate, (NH4)2PtCl6.


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, however, to the alchemists as early as the 13th century, being mentioned by Albertus Magnus, while in the 15th century Basil Valentine showed that ammonia could be obtained by the action of alkalies 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 spirits of hartshorn was applied to ammonia.

Gaseous ammonia was first isolated by J. 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 WWI. The ammonia was used to produce explosives to sustain their war effort.

Data on the heat of fusion and heat of vaporization are from The Planetary Scientist's Companion, by Katharina Lodders and Bruce Fegley, Jr. (New York: Oxford UP Inc., 1998).

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