Ammonia

Ammonia is a compound with the formula NH₃. It is normally encountered as a gas with a characteristic pungent odor. Ammonia contributes significantly to the nutritional needs of the planet as a precursor to foodstuffs and fertilizers. Ammonia, either directly or indirectly, is a building block for the synthesis of many pharmaceuticals. Although in wide use, ammonia is caustic and hazardous.

Ammonia, used commercially is usually named anhydrous ammonia. This term emphasizes the absence of water. Because NH₃ boils at -33 °C, the liquid must be stored under pressure or at low temperature. Its heat of vaporization is, however, sufficiently high 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. See Baumé scale.

Additional recommended knowledge

Correct Test Weight Handling Guide: 12 Practical Tips

Guide to balance cleaning: 8 simple steps

Daily Visual Balance Check

Structure and basic chemical properties

The ammonia molecule has a trigonal pyramid shape, as predicted by VSEPR theory. The nitrogen atom in the molecule has a lone electron pair, and ammonia acts as a base, a proton acceptor. This shape gives the molecule an overall dipole moment and makes it polar so that ammonia readily dissolves in water. In water, a very small percentage of NH₃ is converted into the ammonium cation (NH₄⁺). Thus, the term ammonium hydroxide is a misnomer. The degree to which ammonia forms the ammonium ion increases upon lowering the pH of the solution— at "physiological" pH (~7), about 99% of the ammonia molecules are protonated. Temperature and salinity also affect the proportion of NH₄⁺. NH₄⁺ has the shape of a regular tetrahedron.

The main use of ammonia is for 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.^[5]

Natural occurence

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.^[6] 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^[7]. Salts of ammonia have been known from very early times; thus the term Hammoniacus sal^[8] appears in the writings of Pliny, although it is not known whether the term is identical with the more modern sal-ammoniac.^[8]

In the form of sal-ammoniac, ammonia was known to the alchemists as early as the 13th century, being mentioned by Albertus Magnus.^[9] It was also used by dyers in the Middle Ages in the form of fermented urine^[9] 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 hoofs 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,^[5]

Synthesis and production

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.^[11] The People's Republic of China produced 28.4% of the worldwide production followed by India with 8.6%, Russia with 8.4%, and the United States with 8.2%.^[11] About 80% or more of the ammonia produced is used for fertilizing agricultural crops.^[11]

Before the start of World War I, most ammonia was obtained by the dry distillation^[12] of nitrogenous vegetable and animal waste products, including camel dung, where it was distilled^[13] by the reduction of nitrous acid and nitrites with hydrogen; in addition, it was produced by the distillation of coal,^[13] and also by the decomposition of ammonium salts by alkaline hydroxides^[14] such as quicklime, the salt most generally used being the chloride (sal-ammoniac) thus:

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

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. Starting with a natural gas feedstock, the processes used in producing the hydrogen are:

• The first step in the process entails removal of sulfur compounds from the feedstock, 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 absorbed and converted to solid zinc sulfide:

H₂S + ZnO → ZnS + H₂O

• Catalytic steam reforming of the sulfur-free feedstock is then used to form hydrogen plus carbon monoxide:

CH₄ + H₂O → CO + 3 H₂

• In the next step, the water gas shift reaction is used to convert the carbon monoxide into carbon dioxide and more hydrogen:

CO + H₂O → CO₂ + H₂

• The carbon dioxide is then removed either by absorption in aqueous ethanolamine solutions or by adsorption in pressure swing adsorbers (PSA) using proprietary solid adsorption media.

• 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

The steam reforming, shift conversion, carbon dioxide removal and methanation steps each operate at absolute pressures of about 25 to 35 bar, and the ammonia synthesis loop operates at absolute pressures ranging from 60 to 180 bar, depending upon which proprietary design is used. There are many engineering and construction companies that offer proprietary designs for ammonia synthesis plants. Haldor Topsoe of Denmark, Lurgi AG of Germany, Uhde of Germany, and Kellogg, Brown and Root of the United States are among the most experienced companies in that field.^[15]

As the availability and usage of fossil fuel become problematic (see peak oil and climate change), the hydrogen required for ammonia synthesis could in principle be obtained from electrolysis (currently 4% of hydrogen production is from electrolysis) or thermal chemical cracking of water, but these alternatives are presently impractical. The heat needed for thermal cracking can be obtained from nuclear reaction, while the electricity needed for electrolysis can be obtained from various renewable energy sources such as wind, solar, hydroelectricity, and various forms of ocean energy especially that of OTEC. On really windy days on wind farms the power lines can't handle all the electricity. A possible use for the excess electricity would be to use electrolysis on water to acquire the needed hydrogen. Again alternatives to the production of ammonia from natural gas and air are uneconomic and the environmental benefits have not been established.

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. 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.

Properties

Ammonia is a colorless gas with a characteristic pungent smell similar to human urine, as urine decomposes to release ammonia. 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 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 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 at room temperature, even though it is well above its boiling point.

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 from 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. 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 room temperature - that is, the nitrogen atom passes through the plane of symmetry of the three hydrogen atoms; 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 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.^[16]

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, a gas, moisture being 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⁺).

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 dissociation into the amide (NH₂⁻) ion, for example when solid lithium nitride is added to liquid ammonia, forming a lithium amide solution:

Li₃N_(s)+ 2 NH_{3 (l)} → 3 Li⁺_(am) + 3 NH₂⁻_(am)

This is a Brønsted-Lowry acid-base reaction in which ammonia is acting as an acid.

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

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.
• 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 (2M) 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 + 2 NH₃ → Hg + HgCl(NH₂) + NH₄⁺ + Cl⁻

Uses

Nitric Acid production

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.

4 NH₃ + 5 O₂ → 4 NO + 6 H₂O

2 NO + O₂ → 2 NO₂

2 NO₂ + 2 H₂O → 2 HNO₃ + H₂

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.

Fertilizer

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.

Refrigeration

Ammonia's thermodynamic properties made it one of the refrigerants commonly used in refrigeration units prior to the discovery of dichlorodifluoromethane^[18] in 1928, also known as Freon or R12.

But ammonia is toxic, gaseous, irritant, and corrosive to copper alloys, and over a kilo is needed for even a miniature fridge. With an ammonia refrigerant, the ever present risk of an escape brings with it a risk to life. However data on ammonia escapes has shown this to be an extremely small risk in practice, and there is consequently no control on the use of ammonia refrigeration in densely populated areas and buildings in almost all jurisdictions in the world.

Its use in domestic refrigeration has been mostly replaced by CFCs and HFCs in the first world, which are more or less non-toxic and non-flammable, and butane and propane in the 3rd world, which despite their high flammability do not seem to have produced any significant level of accidents. Ammonia has continued to be used for miniature and multifuel fridges, such as minibars and caravan refridgerators.

These ammonia absorption cycle domestic refrigerators do not use compression and expansion cycles, but are driven by temperature differences. However the energy efficiency of such refrigerators is relatively low. Today the smallest refrigerators mostly use solid state peltier thermopile heat pumps rather than the ammonia absorption cycle.

Ammonia continues to be used as a refrigerant in large industrial processes such as bulk icemaking and industrial food processing. Since the implication of haloalkanes being major contributors to ozone depletion, ammonia is again seeing increasing use as a refrigerant. Ammonia is increasingly popular in commercial applications, such as in grocery store freezer cases and refrigerated displays.

Disinfectant

It is also sometimes added to drinking water along with chlorine to form chloramine, a disinfectant. Unlike chlorine alone, chloramine does not combine with organic (carbon containing) materials to form carcinogenic halomethanes such as chloroform. However, chlorine and ammonia should never be mixed in an uncontrolled environment because they cause a chemical reaction that releases toxic gas. See Safety precautions for more information.

Fuel

Ammonia was used during World War II fuel shortages to power buses in Belgium and used 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. Ammonia is proposed as a practical, clean (CO₂-free), alternative to fossil fuel for internal combustion engines.^[19] In 1981 a Canadian company converted a 1981 Chevrolet Impala to operate using ammonia as fuel.^[20][21] Ammonia is marketed as a low-emission fuel.^[22]

Cigarettes

During the 1960s, tobacco companies such as Brown & Williamson and Philip Morris began using ammonia in cigarettes.^{[citation needed]} Many hypotheses have appeared in the media and technical literature that ammonia enhances the amount of nicotine available to the smoker, nicotine's bioavailability, and the reinforcing or addictive ability^[23]. In contrast, a number of recent publications in the scientific literature have demonstrated that ammonia-forming compounds in tobacco and ammonia in mainstream tobacco smoke do not increase either the total amount or total rate of nicotine to the bloodstream or brains of smokers.^[24].

Illicit Drug Manufacture

• Largely before the popularization of crack cocaine, Ammonium hydroxide was commonly used in the production of " freebase cocaine"
• Conversely (when noting the presence or absence of water in the solution), Anhydrous ammonia is often used in one of the most dangerous methods for the production of methamphetamine

Ammonia's role in biologic 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.^[25]

Ammonia also plays a role in both normal and abnormal animal physiology. Ammonia is created through normal amino acid metabolism and is toxic in high concentrations.^[26] 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.^[27]

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.^[28]

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 lifeforms 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 lifeforms 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

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.

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.^[29] 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.^[30] 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.^[31] 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.^[32]

Safety precautions

Toxicity and storage information

  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. Ammonium compounds should never be allowed to come in contact with bases (unless 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. They should never be mixed with chlorine-containing products or strong oxidants, for example household bleach, as a variety of toxic and carcinogenic compounds are formed (e.g., chloramine, hydrazine, and chlorine gas). Ammonia and sodium hypochlorite react to form a number of products, depending on the temperature, concentration, and how they are mixed. ^[33]. The main reaction is chlorination of ammonia, first giving chloramine (NH₂Cl), then NHCl₂ and finally nitrogen trichloride (NCl₃). These materials are very irritating to eyes and lungs and are toxic above certain concentrations.

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.

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 chemical 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. 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, 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.

Safety

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.^[34] Exposure to very high concentrations of gaseous ammonia can result in lung damage and death.^[34] 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).^[35]

See also

• Ammonia (data page)
• Ammonia production
• Chlorination
• Water purification

References

Bibliography

• This article incorporates text from the Encyclopædia Britannica Eleventh Edition, a publication now in the public domain.
• Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements, 2nd Edn., Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4. 
• Housecroft, C. E.; Sharpe, A. (2001). Inorganic Chemistry. Harlow (UK): Prentice Education. ISBN 0-582-31080-6. 
• (1986) in Bretherick, L.: Hazards in the Chemical Laboratory, 4th Edn., London: Royal Society of Chemistry. ISBN 0-85186-489-9. 
• Weast, R. C. (Ed.) (1972). Handbook of Chemistry and Physics (53rd Edn.). Cleveland: Chemical Rubber Co.

• National Institute for Occupational Safety and Health - Ammonia Page