Information AboutAmmonia |
| CATEGORIES ABOUT AMMONIA | |
| hydrogen compounds | |
| hydrides | |
| nitrogen compounds | |
| bases | |
| nitrogen metabolism | |
| household chemicals | |
| refrigerants | |
| SHOPPER'S DELIGHT | |
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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 . Ammonia used commercially is called ''anhydrous ammonia'' to distinguish it from Ammonium Hydroxide solution, which is household ammonia. An ammonia molecule has a Trigonal Pyramid shape, 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 . That means that, when in aqueous solution, it can take a Proton from water; this produces a Hydroxide Anion and an Ammonium Cation (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 —at "physiological" pH (~7), about 99% of the ammonia molecules are protonated. The main uses of ammonia are in the production of Fertilizer s, Explosive s and Polymer s. 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 Patagonia n Guano . 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. HISTORY Salts of ammonia have been known from very early times; thus the term ''Hammoniacus sal'' Ammonia at Encarta . URL last accessed April 27 2006 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 s in the Middle Ages in the form of fermented Urine to alter the colour of vegetable dyes. In the 15th century, Basilius Valentinus showed that ammonia could be obtained by the action of alkalis on sal-ammoniac. At a later period, when sal-ammoniac was obtained by distilling the 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''; however it was acquired by the alchemist Basil Valentine . Eleven years later in 1785, Claude Louis Berthollet ascertained its composition. The . The ammonia was used to produce explosives to sustain their war effort. SYNTHESIS AND PRODUCTION Because of its many uses, ammonia is one of the most highly-produced inorganic chemicals. There are literally dozens of chemical plants worldwide that produce ammonia. The worldwide Ammonia Production in 2004 was 109 million metric tons. United States Geological Survey publication 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%. About 80% or more of the ammonia produced is used for fertilizing agricultural crops. Before the start of . URL last accessed April 24 2006 of nitrogenous vegetable and animal waste products, including Camel Dung where it was Distilled by the reduction of Nitrous Acid and Nitrite s with Hydrogen ; additionally, it was produced by the distillation of Coal ; and also by the decomposition of ammonium salts by alkaline hydroxides BBC.co.uk URL last accessed April 24 2006 or by 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:
::H2 + RSH → RH + H2S(''g'')
::H2S + ZnO → ZnS + H2O
::CH4 + H2O → CO + 3 H2
::CO + H2O → CO2 + H2
::CO + 3 H2 → CH4 + H2O ::CO2 + 4 H2 → CH4 + 2 H2O
::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, and Kellogg, Brown And Root of the United States are among the most experienced companies in that field. Kellogg Brown's Ammonia Process URL last accessed April 24 2006 BIOSYNTHESIS In certain organisms, ammonia is produced from atmospheric N2 by enzymes called Nitrogenase s. 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 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 Ion izing powers ( ε = 22), and Solution s of Salt s 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 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 it were an umbrella turning 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 (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 Acid s to form Salt s; 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.Baker, H. B. (1894). ''J. Chem. Soc.'' 65: 612. ::NH3 + HCl → NH4Cl The salts produced by the action of ammonia on acids are known as the 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 NH3(l) → 3 Li+(am) + 3 NH2−(am). This is a Bronsted-Lowry acid-base reaction in which ammonia is acting as an acid. Formation of other compounds Ammonia can act as a Nucleophile in Substitution reactions. Amine s can be formed by the reaction of ammonia with Alkyl Halide s, 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. Amide s can be prepared by the reaction of ammonia with a number of Carboxylic Acid derivatives. Acyl Chloride s are the most reactive, but the ammonia must be present in at least a two-fold excess to neutralise the Hydrogen Chloride formed. Ester s and Anhydride s 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 Metal s, 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 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 Compound s. Some notable ammine complexes include:
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 {Link without Title} 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 Acid ic 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 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 . 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 Fertilizer s and Explosive s. 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 Haloalkane s such as Freon . However, ammonia is a toxic irritant and its corrosiveness to any Copper Alloy s 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 . 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 Carcinogen ic Halomethane s such as Chloroform . During the 1960s, Tobacco companies such as '' Brown & Williamson '' and '' Philip Morris '' began using ammonia in Cigarette s. The addition of ammonia serves to enhance the delivery of Nicotene 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 nicotene. Alix M. Freedman, " 'Impact Booster': Tobacco Firm Shows How Ammonia Spurs Delivery of Nicotene ", '' The Wall Street Journal '', Dec. 28, 1995. 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 Electron s. 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 Nitrate s, Nitrite s, Cyanide s and Thiocyanate s. Most Ammonium salts are soluble, and these salts act as Acid s 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 phases. Redox properties of liquid ammonia The range of thermodynamic stability of liquid ammonia solutions is very narrow, as the potential for oxidation to 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 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.P. T. P. Ho and C.H. Townes, 1983, " Interstellar ammonia , ''Ann. Rev. Astron. Astrophys.'', vol. 21, pp. 239-70. 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, 15N H3, NH2 D , NHD2, 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.T. J. Millar, "Deuterium Fractionation in Interstellar Clouds", ''Space Science Reviews'', Vol. 106, Issue 1, pp 73-86. 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 . It is often found as a Liquid on the surface, as temperatures are likely to be very low.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. 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 Acid s or excreted in the urine. However Fish and Amphibian s 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 Membrane s (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 Carcinogen ic compounds are formed (''e.g.'', Chloramine , Hydrazine , and chlorine gas). 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 : , , , , . 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 , 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 ( (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 Peroxide s. Anhydrous ammonia corrodes Copper - and Zinc -containing Alloy s, and so Brass fittings should not be used for handling the gas. Liquid ammonia can also attack rubber and certain plastics. Handling and storage of ammonium compounds 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. SEE ALSO REFERENCES BIBLIOGRAPHY EXTERNAL LINKS
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