What are the inorganic compounds of nitrogen? Information about the properties and uses of inorganic nitrogen compounds.
Inorganic Nitrogen Compounds; The two principal inorganic nitrogen compounds are ammonia and nitric acid. A number of substances are either made from them or may be regarded as closely related to them.
Ammonia.—Ammonia has the formula NHS and is one of the earliest nitrogen compounds known to man, having been produced originally by destructive distillation of such organic materials as the horns of animals. With the development of the process for the carbonization of coal, ammonia was scrubbed from coal gas and became an important commercial chemical. In industry, ammonia is made almost entirely by the direct combination of nitrogen and hydrogen at pressures from 200 to 1000 atmospheres and temperatures between 450°C. and 700°C. in the presence of an iron catalyst containing other elements as promoters.
Ammonia is a gas at ordinary temperatures, but liquefies at —33.4°C., and is commonly handled in cylinders and tank cars, the liquid being under approximately 10 atmospheres pressure at 25 °C. It is extremely soluble in water, the commercial product being known as “aqua ammonia,” a 28 per cent solution.
Gaseous ammonia may be introduced directly into the soil, but is more commonly an ingredient of commercial fertilizer in the form of an ammonium salt. Liquid ammonia is an industrial refrigerant, being employed mainly in large installations. Its chief use is that of a starting substance in the production of a number of industrial chemicals. These include ammonium sulphate, ammonium nitrate, ammonium chloride, ammonium phosphates, urea, and sulfamic acid.
Hydrazine.—Isolated in 1887 by the German chemist, Theodor Curtius, hydrazine has the formula NH2NH2, and is regarded as having the same relation to ammonia that hydrogen peroxide, H2O2, bears to water. Until World War II, hydrazine was primarily a laboratory chemical. Within a few years, hydrazine became a large-scale product of chemical industry. In addition to use directly as a foaming agent in the rubber industry, as a reducing agent in depositing metal films on glass and plastics, and as a rocket fuel, it is a starting substance for soldering fluxes, plant growth retardants, and antituberculosis drugs and other pharmaceuticals. Hydrazine synthesis starts with chlorine and caustic soda, followed by treatment of the resulting hypochlorite with ammonia. Anhydrous hydrazine is obtained from the dilute solution. (See also Hydrazine.)
Hydroxylamine.—This substance is a weak base with the formula NH2OH. It is produced as a salt by electrolytic reduction of nitrates in SO per cent sulphuric acid solution. It is a commercial chemical in the form of its hydrochloride and sulphate, and is used as a reagent in synthesis, and as an antiseptic, a photographic developer, and an oxidation inhibitor in fatty acids.
Hydrazoic Acid.—A poisonous and highly explosive acid, it has the formula HN3, and is made by the action of hydrazine on esters of organic compounds, followed by treatment with nitrous acid. It forms salts with metals such as sodium azide, NaNs, silver azide, AgNa, and lead azide, Pb(N3)2, the last named being an important detonator in the explosives industry. See Lead Azide.
Nitrides.—Nitrogen reacts with a number of free elements to form nitrides in which the valence of nitrogen is three. For example, when magnesium is burned in air, in addition to magnesium oxide, a certain amount of magnesium nitride, Mg3N2, is formed. Other typical nitrides include those of lithium, calcium, boron, aluminum, silicon, titanium, and chromium. Of these, boron nitride is the most important industrially, being an excellent refractory. When employed in the researches leading to the discovery of the rare gases of the atmosphere, magnesium removed nitrogen from a mixture of these gases by formation of the nitride. However, the development of processes of distilling liquid air and adsorption of the rare gases by activated carbon replaced this technique.
Metal Amides.—Active metals, and a few others which are less active, form amides with ammonia, the most important being sodamide, NaNH2, which is extensively used in chemical synthesis.
Nitrogen Halides.—Chlorine reacts with ammonium chloride to form nitrogen trichloride, NCls, an explosive, oily liquid. As a gas, and when greatly diluted with air, this substance is a valuable bleaching agent. The lower fluorides, NF2 and N2F2, have also been made. What is known as nitrogen iodide is a complex compound having the formula NI« ■ 2 NH3. Stable only when wet, this substance after drying explodes at the slightest contact with a moving solid object, and can be set off by the touch of a feather.
Nitrogen Oxides.—All five oxides of nitrogen are derived either directly or indirectly from nitric oxide, NO ; the chemistry of these oxides is best understood by taking them up in increasing order of oxygen content.
Nitrous Oxide, N20, is a colorless and almost odorless gas which liquefies at — 88.S°C. and is sold as a liquid in strong steel cylinders, its vapor
pressure at 20°C. being 50 atmospheres. It is quite soluble in water, and its solution has a slightly sweet taste. Nitrous oxide is made by gently heating ammonium nitrate: NH4NOa = N2O + 2 H2O. The amount of heat liberated in this reaction is not large, but the decomposition of nitrous oxide into nitrogen and oxygen is accompanied by much greater heat, and rapid and intense heating of the ammonium salt may result in a serious explosion. Because nitrous oxide breaks up readily with the liberation of oxygen, it responds to the same test as that used for the presence of free oxygen, namely the bursting into flame of a glowing splinter when it is introduced into the gas.
The principal use of nitrous oxide is as an anesthetic. Before the development of local anesthesia, it had more extensive use in dentistry than at present. It is still among the anesthetics in common use in surgery. Since one of its effects before consciousness is lost is that of exhilaration or mild hysteria, this substance long ago acquired the name of “laughing gas.”
Sir Humphry Davy in 1801 observed that nitrous oxide produced insensibility on prolonged inhalation and suggested its use in surgery. It was more than forty years later that Dr. Horace Wells, a dentist of Hartford, Conn., employed nitrous oxide as an anesthetic in dental surgery.
Another interesting but minor use of this substance is to dissolve it under slight pressure in cream in a container. When the valve is opened, a foam is produced of the same consistency as that obtained by mechanical inclusion of air by rapid agitation or “whipping.”
Nitric Oxide, NO, is a colorless gas which liquefies at —1S2°C. and solidifies at —164°C. and is only slightly soluble in water. Its most striking property is the readiness with which it combines with oxygen in the cold; its mixture with air at any temperature below 150°C. results in immediate combination : 2NO + O2 = 2N02. At higher temperatures the resulting nitrogen dioxide dissociates, and at 620° C. the reaction is completely reversed. Nitric oxide forms unstable compounds with some salts; the addition compound with ferrous sulphate, FeS04 ■ NO, which has a brown color, is a test for the presence of nitrates, the ferrous sulphate first reducing pentavalent to divalent nitrogen.
Nitric oxide is made in quite pure form by the action of ferrous sulphate on nitric acid in sulphuric acid solution. It is the principal oxide of nitrogen formed when 30 to 35 per cent nitric acid reacts with copper. Industrially it is produced in quantity by passing a mixture of ammonia and air over heated platinum gauze: 4NH3 + 50_2 = 4NO + 6H2O. The nitric oxide so formed is commonly oxidized at lower temperatures to nitrogen dioxide, N02. Direct combination of oxygen and nitrogen in an electric arc at temperatures of 2,500° to 3,200°C., followed by rapid cooling of the gas, results in the formation of nitric oxide to a sufficient extent that the gas at 700° C. contains a little over I per cent of this oxide. The arc is spread either by a magnetic field or by introducing the gas through special nozzles. This process was operated successfully in Norway, where water power produced cheap electric current, but is now of minor importance. It has been proved feasible to make nitric oxide by passing air over refractory material previously subjected to intense heat by burning hydrocarbon gases. The further development of this process is largely dependent on economic considerations. Nitric oxide is primarily employed in the manufacture of nitric acid.
Nitrogen Trioxide, N2Os. At low temperature, nitrogen trioxide is a blue liquid which freezes at —103°C. and boils with decomposition at 3.5°C. At ordinary temperatures nitrogen dioxide is incapable of independent existence, and may be regarded as an equimolecular mixture of nitric oxide and nitrogen dioxide: N2O3 = NO + NOa. This mixture is the carrier of oxygen in the nitration or lead-chamber process for the manufacture of sulphuric acid (see Sulphur, Applications of—Sulphuric Acid). It is the anhydride of nitrous acid, HNO2, and forms this acid when dissolved in water.
Nitrogen Dioxide, N02(and Nitrogen Tetrox-ide, N2O1). This substance is a toxic gas with a pungent and irritating odor and a deep reddish-brown color when warm, the color becoming almost black at still higher temperatures. Dissociation into oxygen and nitric oxide begins at around 150°C. and continues until it is complete at 620° C. As nitrogen dioxide gas is cooled, it begins to polymerize to nitrogen tetroxide (sometimes called dinitrogen tetroxide) ,N204. At the liquefaction point, 21.1 °C., the yellow liquid is made up primarily of this polymer. Nitrogen dioxide is formed quantitatively in the cold by the action of oxygen of the air on nitric oxide, and is also the principal product of the reaction between concentrated nitric acid and such metals as copper, mercury, and silver. Nitrogen dioxide reacts with cold water to form nitrous acid, HNO2, and nitric acid, HNO3. The former acid is unstable, and if water is not kept cold during the absorption, the net result is represented by the equation 3 N02 + H20 = 2 _ HNOs + NO. Nitrogen dioxide is primarily an intermediate in the manufacture of nitric acid, but is employed as an oxidation and nitration agent in the organic chemical industry.
Nitrogen Pcntoxide,N2Os, is a white crystalline solid, melting at 30°C. and decomposing spontaneously at 50°C. into nitrogen dioxide and oxygen. This oxide is produced by the action of phosphorus pentoxide on anhydrous nitric acid. Concentrated nitric acid fumes in moist air because of the escape of the pentoxide and its reaction with the water vapor. The pentoxide has no commercial use.
Nitrogen-Oxygen Acids.—There are six acids containing nitrogen, oxygen, and hydrogen, but of these only three are of importance.
Hyponitrous Acid, H202N2, is regarded as having the structure HO—N=N—OH. It is an unstable white solid. Sodium nitriate, NaN02, is reduced to sodium hyponitrite by sodium amalgam, from which the free acid can be obtained. Neither the acid nor its salts have industrial value.
Nitrous Acid, HN02, has never been isolated in the free state, but can be formed in solution by the action of an acid upon dilute solutions of sodium nitrite, a blue color being imparted by the acid. Warming causes the acid to break up into nitric acid, nitric oxide, and water. Solutions of nitrous acid are active oxidizing agents, releasing iodine from iodides and reacting readily with organic compounds capable of easy oxidation. With strong oxidizing agents, such as permanganates, nitrous acid functions as a reducing agent.
The chief value of solutions of nitrous acid is in the synthetic organic ‘ chemical industry, particularly in the manufacture of dyes. A considerable percentage of dyes belong to classes which are characterized by the azo group, —N=N—, while a small number are nitroso compounds containing the —NO group. Other nitroso compounds are dye intermediates. These substances depend on nitrous acid from sodium nitrite in acid solution.
Nitric Acid, HNOs, the most important of the nitrogen acids, is one of the earliest inorganic chemicals known (since the 9th century). Its early name, given by the alchemists, was “aqua fortis” or “strong water.” Pure hydrogen nitrate is a colorless liquid (specific gravity 1.S4) which solidifies at —41.6°F. It distills with some decomposition at 83°C. at atmospheric pressure. Commercial nitric acid is commonly colored by the presence of dissolved nitrogen dioxide, which is soluble up to the extent of 12 per cent at ordinary temperatures. Nitric acid and water form a constant-boiling mixture made up of 68 per cent nitric acid and 32 per cent water, the boiling point being 110°C. at 760 mm. (millimeters) pressure. Concentrations of nitric acid up to 100 per cent can be made by heating a mixture of sodium nitrate and sulphuric acid, but the common commercial process in which nitrogen dioxide is absorbed in water yields an acid of about 60 per cent concentration. Addition of concentrated sulphuric acid and heating produces a 90 to 95 per cent acid.
Nitric acid is a powerful oxidizing agent; its oxidation potential depends on the concentration of its solutions. The course and final products of the reactions also depend on the reducing potentials of the substances with which nitric acid reacts. Reduction products of nitric acid include nitrogen dioxide, nitric oxide, nitrous oxide, nitrogen, and ammonia. It is generally considered that there are only two principal products, nitric oxide and the ammonium ion, NH4+, the other substances being formed by secondary reactions. With dilute nitric acid and active metals such as magnesium, hydrogen and nitric oxide are the products. Somewhat less active metals like zinc and iron produce the ammonium ion. Only moderately active metals such as copper, silver, and mercury yield nitric oxide, the acid being from 25 to 45 per cent concentration. An acid of 60 to 68 per cent strength produces nitrogen dioxide. This difference is accounted for by the reaction, 3 NO2 + H20= 2 HNOs + NO. If the concentration of water is high, as in the more dilute acid, the product is nitric oxide, while the higher concentration of nitric acid oxidizes the nitric oxide to nitrogen dioxide. In all of these cases a metallic nitrate is formed. Metals whose oxides are insoluble, like tin, antimony, tungsten, and molybdenum, are reacted upon by nitric acid with the formation of these oxides rather than their nitrates. Some metals of very low reducing potential are not affected by nitric acid, some of these being gold, indium, rhodium, and platinum. Some metals, such as aluminum and chromium, are protected by an oxide coating, and thus resist the action of nitric acid. Nonmetals, of which sulphur, phosphorus, and iodine are representative, react with nitric acid to form their oxygen acids, along with either nitric oxide or nitrogen dioxide, depending upon the concentration of the nitric acid. Compounds such as