NH3, boiling point ?33°C. A colourless gas with a pungent odour; its molecule is pyramidal with bond angles c.107°C. It is a weak base; aqueous solutions partially neutralized with strong acid have a pH c.9·5. It reacts with acids to form ammonium ions. An important industrial chemical, it is mainly prepared by the Haber process.
Ammonia used commercially can be anhydrous ammonia (not dissolved in water) or an aqueous solution of ammonia and water referred to as ammonium hydroxide. Ammonium hydroxide strength 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. Household ammonia ranges in concentration from 5 to 10 weight percent ammonia.An ammonia molecule has a trigonal pyramid shape, as predicted by VSEPR theory. 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 fertilizers, explosives and polymers. 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;
History
Salts of ammonia have been known from very early times;
In the form of sal-ammoniac, ammonia was known to the alchemists as early as the 13th century, being mentioned by Albertus Magnus.
Gaseous ammonia was first isolated by Joseph Priestley in 1774 and was termed by him alkaline air;
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.
Synthesis and production
Because of its many uses, ammonia is one of the most highly produced inorganic chemicals.
Before the start of World War I most ammonia was obtained by the dry distillation of nitrogenous vegetable and animal waste products, including camel dung where it was distilled by the reduction of nitrous acid and nitrites with hydrogen; Sulfur removal requires catalytic hydrogenation to convert sulfur compounds in the feedstocks to gaseous hydrogen sulfide: H2 + RSH → RH + H2S(g) The gaseous hydrogen sulfide is then absorbed and removed by passing it through beds of zinc oxide where it is 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 The next step then uses catalytic shift conversion to convert the carbon monoxide to 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.
Biosynthesis
In certain organisms, ammonia is produced from atmospheric N2 by enzymes called nitrogenases.
Ammonia is also a metabolic product of amino acid deamination.
Properties
Ammonia is a colourless gas with a characteristic pungent smell; 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. 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;
Formation of salts
One of the most characteristic properties of ammonia is its power of combining directly with acids to form salts;
NH3 + HCl → NH4ClThe 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 NH3 (l) → 3 Li(am)This is a Brønsted-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. 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.
Amides can be prepared by the reaction of ammonia with a number of carboxylic acid derivatives. Acyl chlorides are the most reactive, but the ammonia must be present in at least a two-fold excess to neutralise the hydrogen chloride formed.
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.
Ammonia as a ligand
Ammonia can act as a ligand in transition metal complexes. Some notable ammine complexes include:
Tetraaminecopper(II), [Cu(NH3)4]2+, a characteristic dark blue complex formed by adding ammonia to solution of copper(II) salts. 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.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.
Hg2Cl2 + 2 NH3 → Hg + HgCl(NH2) + NH4Uses
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 H2OThe 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.
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.
Ammonia has thermodynamic properties that make it very well suited as a refrigerant, since it liquefies readily under pressure, and was used in virtually all refrigeration units prior to the advent of haloalkanes such as Freon.
Liquid ammonia was used as the fuel of the rocket airplane, the X-15.
Ammonia's role in biologic systems and human disease
Ammonia is an important source of nitrogen for living systems.
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.
Ammonia is important for normal animal acid/base balance. Ammonia may itself diffuse across the renal tubules, combine with a hydrogen ion, and thus allow for further acid excretion.
Liquid ammonia as a solvent
Liquid ammonia is the best-known and most widely studied non-aqueous ionizing solvent. Apart from these remarkable solutions, much of the chemistry in liquid ammonia can be classified by analogy with related reactions in aqueous solutions.
| 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. 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
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.
Redox properties of liquid ammonia
| E° (V, ammonia) | E° (V, water) | |
|---|---|---|
| Li ⇌ Li | −2.24 | −3.04 |
| K ⇌ K | −1.98 | −2.93 |
| Na ⇌ Na | −1.85 | −2.71 |
| Zn ⇌ Zn | −0.53 | −0.76 |
| NH4 ⇌ ½ H2 + NH3 | 0.00 | – |
| Cu ⇌ Cu | +0.43 | +0.34 |
| Ag ⇌ 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 ⇌ 8NH3), is only +0.04 V. 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. 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.
The following isotopic species of ammonia have been detected:
NH3, 15NH3, NH2D, NHD2, and ND3The detection of triply-deuterated ammonia was considered a surprise as deuterium is relatively scarce. 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.
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. 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.
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);
|
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 |
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.
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). 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.
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