Ammonia

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Ammonia
AmmoniaAmmonia 3D representation
General
Systematic name Ammonia
Azane (See Text)
Other names Hydrogen nitride
Spirit of hartshorn
Nitrosil
Vaporole
Molecular formula NH3
Molar mass 17.0304 g/mol
Appearance Colorless gas with
strong pungent odor
CAS number [7664-41-7]
Properties
Density and phase 0.6813 g/L, gas.
Solubility in water 89.9 g/100 ml at 0 °C.
Melting point -77.73 °C (195.42 K)
Autoignition temperature 651 °C
Boiling point -33.34 °C (239.81 K)
Acidity (pKa) 9.25
Basicity (pKb) 4.75
Structure
Molecular shape Terminus
Dipole moment 1.42 D
Bond angle 107.5°
Hazards
MSDS External MSDS
Main hazards Toxic and corrosive.
NFPA 704

1
3
0
 
Flash point None[1]
R/S statement R: R10, R23, R34, R50
S: S1/2, S16, S36/37/39,
S45, S61
RTECS number BO0875000
Supplementary data page
Structure and
properties		 n, εr, etc.
Thermodynamic
data		 Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Related compounds
Other ions Ammonium (NH4+)
hydroxide (NH4OH)
chloride (NH4Cl)
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 with the formula NH3. It is normally encountered as a gas with a characteristic pungent odor. Although ammonia contributes significantly to the nutritional needs of the planet, the gas itself is toxic.

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 and/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.[2] Household ammonia ranges in concentration from 5 to 10 weight percent ammonia. See Baumé scale.

Contents

[edit] 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 uses of ammonia are in the production of fertilizers, explosives, and synthesis of organonitrogen compounds. It is also the active ingredient in 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. The kidneys secrete NH4+ to neutralize excess acid.[3] Ammonium salts also are found distributed through all fertile soil and in seawater. Substances containing ammonia, or that are similar to it, are called ammoniacal.

[edit] History

Salts of ammonia have been known from very early times; thus the term Hammoniacus sal[4] appears in the writings of Pliny, although it is not known whether the term is identical with the more modern sal-ammoniac.[4]

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

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.[6] 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,[7] 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.[8]

[edit] 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.[9] 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%.[9] About 80% or more of the ammonia produced is used for fertilizing agricultural crops.[9]

Before the start of World War I most ammonia was obtained by the dry distillation[10] of nitrogenous vegetable and animal waste products, including camel dung where it was distilled[8] by the reduction of nitrous acid and nitrites with hydrogen; additionally, it was produced by the distillation of coal;[8] and also by the decomposition of ammonium salts by alkaline hydroxides[11] 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
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

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.[12]

[edit] 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.

[edit] 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 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, 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 /cm³ and is often known as '.880 Ammonia'. Ammonia 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 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.[13]

[edit] 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.[14]

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+).

[edit] 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 NH3 (l) → 3 Li+(am) + 3 NH2−(am)

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

[edit] 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 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.

[edit] Ammonia as a ligand

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

  • Tetraamminecopper(II), [Cu(NH3)4]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−

[edit] Uses

[edit] 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

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.

[edit] Universal Indicator

Ammonia solution is also used as universal indicator that could be used to test for different gases that require a universal indicator solution to show the gases were present.

[edit] 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.

[edit] Refrigeration

Ammonia's thermodynamic properties made it one of the refrigerants commonly used in refrigeration units prior to the discovery of dichlorodifluoromethane[15] 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 fridges.

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.

[edit] Disinfectant

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.

[edit] Fuel

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 has about the same density as the oxidizer, liquid oxygen, which simplified the aircraft's keeping the same center of gravity in flight.

[edit] Cigarettes

During the 1960s, tobacco companies such as Brown & Williamson and Philip Morris began using ammonia in cigarettes. The addition of ammonia serves to enhance the delivery of nicotine into the blood stream. As a result, the reinforcement effect of the nicotine was enhanced, increasing its addictive ability without actually increasing the portion of nicotine.[16]

[edit] 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.[17]

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. 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.[18]

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.[19]

[edit] 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.

[edit] 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.

[edit] 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.

[edit] 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.

[edit] 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.

[edit] Interstellar space

Ammonia was first detected in interstellar space in 1968, based on microwave emissions from the direction of the galactic core.[20] 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.[21] 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.[22] 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.[23]

[edit] Safety precautions

[edit] Toxicity and storage information

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

[edit] 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).

[edit] 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 (Xi) 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 chemical analysis, and should be acidified and diluted before disposal once the test is completed.

[edit] 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.

[edit] See also

[edit] References

  1. ^ MSDS Sheet from W.D. Service Co.
  2. ^ Ammonium hydroxide physical properties
  3. ^ electrolytes and the urine anion and osmolar gaps.
  4. ^ a b Webmineral website URL last accessed August 27 2006
  5. ^ a b c Absolouteastronomy.com URL last accessed April 24 2006
  6. ^ Abraham, Lyndy. Marvell and alchemy. Aldershot Scolar 1990.
  7. ^ BBC.co.uk URL last accessed April 24 2006
  8. ^ a b c Smith, Roland. Conquering Chemistry 2001
  9. ^ a b c United States Geological Survey publication
  10. ^ Nobel Prize in Chemistry (1918) - Haber process. URL last accessed April 24 2006
  11. ^ BBC.co.uk URL last accessed April 24 2006
  12. ^ Kellogg Brown's Ammonia Process URL last accessed April 24 2006
  13. ^ C. E. Cleeton & N. H. Williams, 1934 - Online version; archive. URL last accessed May 8, 2006
  14. ^ Baker, H. B. (1894). J. Chem. Soc. 65: 612.
  15. ^ Dichlorodifluoromethane by Aaron Vorderstrasse, Western Oregon University.
  16. ^ Alix M. Freedman, "'Impact Booster': Tobacco Firm Shows How Ammonia Spurs Delivery of Nicotine", The Wall Street Journal, Dec. 28, 1995.
  17. ^ M.B. Adjei, K.H. Quesenberry and C.G. Chamblis. Nitrogen Fixation and Inoculation of Forage Legumes University of Florida IFAS Extension June 2002.
  18. ^ Zschocke, Johannes, and Georg Hoffman. Vademecum Metabolism. Friedrichsdorf, Germany: Milupa GmbH, 2004.
  19. ^ Rose, Burton, and Helmut Rennke. Renal Pathophysiology. Baltimore, Maryland: Williams & Wilkins, 1994.
  20. ^ A.C. Cheung, D.M. Rank, C.H. Townes, D.D. Thornton, and W.J. Welch, 1968, "Detection of NH3 molecules in the interstellar medium by their microwave emission," Phys. Rev. Lett. 21, 1701.
  21. ^ P. T. P. Ho and C.H. Townes, 1983, "Interstellar ammonia, Ann. Rev. Astron. Astrophys., vol. 21, pp. 239-70.
  22. ^ T. J. Millar, "Deuterium Fractionation in Interstellar Clouds", Space Science Reviews, Vol. 106, Issue 1, pp 73-86.
  23. ^ Edited by Kirk Munsell. Image page credit Lunar and Planetary Institute. NASA. "NASA's Solar Exploration: Multimedia: Gallery: Gas Giant Interiors". URL accessed April 26, 2006.

[edit] 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.

[edit] External links

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