Nitric acid (4) - Abstract. Special purity nitric acid

Nitric acid page 3

Oxidative properties of nitric acid page 3

Nitrates page 6

Industrial production of nitric acid page 7

Nitrogen cycle in nature page 8

6. Bibliography page 10

1. Nitric acid. Pure nitric acid HNO is a colorless liquid with a density of 1.51 g/cm at - 42 °C, solidifying into a transparent crystalline mass. In the air, it, like concentrated hydrochloric acid, “smoke”, since its vapors form small droplets of fog with the moisture in the air,

Nitric acid is not durable. Even under the influence of light, it gradually decomposes:

The higher the temperature and the more concentrated the acid, the faster the decomposition occurs. The released nitrogen dioxide dissolves in the acid and gives it a brown color.

Nitric acid is one of the most powerful acids; in dilute solutions it completely disintegrates into H and NO ions.

2. Oxidizing properties of nitric acid. A characteristic property of nitric acid is its pronounced oxidizing ability. Nitric acid is one of the most energetic oxidizing agents. Many non-metals are easily oxidized by it, turning into the corresponding acids. Thus, sulfur, when boiled with nitric acid, is gradually oxidized into sulfuric acid, phosphorus into phosphoric acid. A smoldering ember immersed in concentrated HNO flares up brightly.

Nitric acid acts on almost all metals (with the exception of gold, platinum, tantalum, rhodium, iridium), turning them into nitrates, and some metals into oxides.

Concentrated HNO passivates some metals. Lomonosov also discovered that iron, which easily dissolves in dilute nitric acid, does not dissolve in cold concentrated HNO. Later it was found that nitric acid has a similar effect on chromium and aluminum. These metals pass under the influence of concentrated nitric acid into a passive state.

The oxidation degree of nitrogen in nitric acid is 4-5. Acting as an oxidizing agent, HNO can be reduced to various products:

Which of these substances is formed, i.e., how deeply nitric acid is reduced in a given case, depends on the nature of the reducing agent and on the reaction conditions, primarily on the concentration of the acid. The higher the concentration of HNO, the less deeply it is reduced. When reacting with concentrated acid, it is most often released. When dilute nitric acid reacts with low-active metals, for example, copper, it releases NO. In the case of more active metals - iron, zinc - is formed. Strongly diluted nitric acid reacts with active metals - zinc, magnesium, aluminum - to form ammonium ion, which gives ammonium nitrate with the acid. Usually several products are formed simultaneously.

For illustration, we present diagrams of the oxidation reactions of some metals with nitric acid;

When nitric acid acts on metals, hydrogen, as a rule, is not released.

When non-metals are oxidized, concentrated nitric acid, as in the case of metals, is reduced to, for example

A more dilute acid is usually reduced to NO, for example:

The given diagrams illustrate the most typical cases of interaction of nitric acid with metals and non-metals. In general, redox reactions involving oxidation are complex.

A mixture consisting of 1 volume of nitric and 3-4 volumes of concentrated hydrochloric acid is called royal vodka. Aqua regia dissolves some metals that do not react with nitric acid, including the “king of metals” - gold. Its action is explained by the fact that nitric acid oxidizes hydrochloric acid, releasing free chlorine and forming nitrogen chloroxide a(III), or nitrosyl chloride, :

Nitrosyl chloride is a reaction intermediate and decomposes:

Chlorine at the moment of release consists of atoms, which determines the high oxidizing ability of aqua regia. The oxidation reactions of gold and platinum proceed mainly according to the following equations.

With excess hydrochloric acid, gold(III) chloride and platinum(IV) chloride form complex compounds

For many organic matter nitric acid acts by replacing one or more hydrogen atoms in an organic compound molecule with nitro groups. This process is called nitration and has great value in organic chemistry.

Nitric acid is one of the most important nitrogen compounds: it is used in large quantities in the production of nitrogen fertilizers, explosives and organic dyes, serves as an oxidizing agent in many chemical processes, is used in the production of sulfuric acid using the nitrous method, and is used for the manufacture of cellulose varnishes and film.

3. Nitrates. Salts of nitric acid are called nitrates. All of them dissolve well in water, and when heated, they decompose, releasing oxygen. In this case, the nitrates of the most active metals turn into nitrites:

Nitrates of most other metals decompose when heated into metal oxide, oxygen and nitrogen dioxide. For example:

Finally, nitrates of the least active metals (for example, silver, gold) decompose when heated to the free metal:

Easily stripping off oxygen, nitrates are energetic oxidizing agents at high temperatures. Their aqueous solutions, on the contrary, exhibit almost no oxidizing properties.

Most important have sodium, potassium, ammonium and calcium nitrates, which in practice are called nitrate.

Sodium nitrate or sodium nitrate, sometimes also called Chilean saltpeter, found in in large quantities in nature only in Chile.

Potassium nitrate, or potassium nitrate, It also occurs in nature in small quantities, but is mainly produced artificially by reacting sodium nitrate with potassium chloride.

Both of these salts are used as fertilizers, and potassium nitrate contains two elements necessary for plants: nitrogen and potassium. Sodium and potassium nitrates are also used in glass melting and food industry for canning food.

Calcium nitrate or calcium nitrate, obtained in large quantities by neutralizing nitric acid with lime; used as fertilizer.

4. Industrial production of nitric acid. Modern industrial methods for producing nitric acid are based on the catalytic oxidation of ammonia with atmospheric oxygen. When describing the properties of ammonia, it was indicated that it burns in oxygen, and the reaction products are water and free nitrogen. But in the presence of catalysts, the oxidation of ammonia with oxygen can proceed differently. If a mixture of ammonia and air is passed over a catalyst, then at 750 °C and a certain composition of the mixture, almost complete conversion occurs

The resulting mixture easily passes into, which, with water in the presence of atmospheric oxygen, gives nitric acid.

Platinum-based alloys are used as catalysts for the oxidation of ammonia.

The nitric acid obtained by the oxidation of ammonia has a concentration not exceeding 60%. If necessary, it is concentrated,

The industry produces diluted nitric acid with a concentration of 55, 47 and 45%, and concentrated nitric acid - 98 and 97%. Concentrated acid is transported in aluminum tanks, diluted - in acid-resistant steel tanks.

5. Nitrogen cycle in nature. When organic matter rots, a significant part of the nitrogen contained in them is converted into ammonia, which, under the influence of nitrifying bacteria living in the soil, is then oxidized into nitric acid. The latter, reacting with carbonates in the soil, for example calcium carbonate, forms nitrates:

Some part of the nitrogen is always released during decay in free form into the atmosphere. Free nitrogen is also released during the combustion of organic substances, during the combustion of wood, coal, and peat. In addition, there are bacteria that, if there is insufficient air access, can take oxygen away from nitrates, destroying them and releasing free nitrogen. The activity of these denitrifying bacteria leads to the fact that part of the nitrogen from the form available to green plants (nitrates) becomes inaccessible (free nitrogen). Thus, not all the nitrogen that was part of the dead plants returns back to the soil; part of it is gradually released in free form.

The continuous loss of mineral nitrogen compounds should have long ago led to the complete cessation of life on Earth if processes did not exist in nature to compensate for the loss of nitrogen. Such processes include, first of all, electrical discharges occurring in the atmosphere, during which a certain amount of nitrogen oxides is always formed; the latter produce nitric acid with water, which is converted into nitrates in the soil. “Another source of replenishment of nitrogen compounds in the soil is the vital activity of the so-called azotobacteria, which are capable of assimilating atmospheric nitrogen. Some of these bacteria settle on the roots of plants from the legume family, causing the formation of characteristic swellings - “nodules”, which is why they are called nodule bacteria. Assimilating atmospheric nitrogen , nodule bacteria process it into nitrogen compounds, and plants, in turn, convert the latter into proteins and other complex substances.

Thus, a continuous nitrogen cycle occurs in nature. However, every year, the most protein-rich parts of plants, such as grain, are removed from the fields with the harvest. Therefore, it is necessary to add fertilizers to the soil to compensate for the loss of essential plant nutritional elements.

The study of plant nutrition and increasing the productivity of the latter through the use of fertilizers is the subject of a special branch of chemistry, called agrochemistry.

Nitric acid (HNO3) is a strong monobasic acid, one of the most important mineral acids. The substance has a sharp suffocating odor, is sensitive to light and decomposes in bright light. Nitric acid mixes with water in any ratio. In aqueous solutions, it almost completely dissociates into ions.

It is not found in nature in a free state and occurs only in the form of nitrate salts (in the form of ammonium nitrate in the air and rainwater, especially after thunderstorms, then in the form of sodium nitrate in Chilean or Peruvian nitrate and potassium and calcium nitrate in the upper layers of arable land, on walls of stables, in the lowlands of the Ganges and other rivers of India).

Nitric acid is obtained by decomposing sodium nitrate with concentrated sulfuric acid when heated.

Easily gives oxygen to other bodies and is a strong oxidizing agent and solvent for many substances. Moreover, nitric acid in any concentration exhibits the properties of an oxidizing acid. However, nitric acid, even concentrated, does not interact with gold and platinum, but is part of the so-called “regia vodka”, which dissolves these metals.

Nitric acid is widely used in various industries:

in production
in the military industry (fuming - in the production of explosives, as an oxidizer for rocket fuel, diluted - in synthesis various substances, including poisonous ones)
in easel graphics - for etching printing forms (etching boards, zincographic printing forms and magnesium clichés)
in the production of dyes
in pharmacology (nitroglycerin)
V jewelry- the main method for determining gold in a gold alloy
mostly organic synthesis(nitroalkanes, aniline, nitrocellulose, TNT).

Nitric acid is very harmful to humans. Her fumes are annoying respiratory tract, and the acid itself leaves ulcers on the skin that do not heal well. When heated or exposed to light, the acid decomposes to form highly toxic nitrogen dioxide.

Transportation can be carried out by all modes of transport, except air, in accordance with the rules for the transportation of dangerous goods. Since the acid is sensitive to light, acid can only be stored and transported in opaque containers. Concentrated nitric acid, special purity grade, is poured into glass bottles, barrels, vessels and tanks made of stainless steel. Storage is carried out at a temperature of no more than +40°C.

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NITRIC ACID, HNO 3, is obtained by dissolving nitrogen oxides in water:

3NO 2 + H 2 O = 2HN 3 + NO
N 2 O 3 + H 2 O = HNO 3 + NO
N2O5 + H2O = 2HNO3

Physical properties of nitric acid. Molar weight- 63.016; colorless liquid with a characteristic odor; boiling point 86°, melting point -47°; specific gravity 1.52 at 15°; during distillation, due to the decomposition of 2HNO 3 = N 2 O 3 + 2O + H 2 O, nitric acid immediately releases oxygen, N 2 O 3 and water; absorption of the latter causes an increase in boiling point. In aqueous solution, strong nitric acid usually contains nitrogen oxides, and the preparation of completely anhydrous nitric acid presents significant difficulties. It is impossible to obtain anhydrous nitric acid by distillation, since they have a minimum elasticity aqueous solutions nitric acid, i.e. adding water to the acid and vice versa lowers the vapor pressure (and increases the boiling point). Therefore, as a result of distillation of a weak acid (D< 1,4) получается постоянно кипящий остаток D = 1,415, с содержанием 68% HNО 3 и с температурой кипения 120°,5 (735 мм). Перегонка при пониженном давлении дает остаток с меньшим содержанием HNО 3 , при high blood pressure- with a high content of HNO 3. Acid D = 1.503 (85%), purified by blowing air from N 2 O 4, gives a residue with 77.1% HNO 3 during distillation. During distillation, acid D = 1.55 (99.8%) first gives a solution D = 1.62, strongly colored by nitrogen oxides, and the remainder acid D = 1.49. That. The residue from the distillation of nitric acid always contains the acid corresponding to the minimum elasticity (maximum boiling point). Anhydrous acid can be obtained only by mixing strong (99.1%) nitric acid with nitric anhydride.

By freezing, apparently, it is impossible to obtain acid above 99.5%. With the new methods (Valentiner) of extracting nitric acid from saltpeter, the acid is quite pure, but with the old ones it was necessary to purify it mainly from chloride compounds and from N 2 O 4 vapors. The strongest acid has D0 = 1.559, D15 = 1.53, and 100% HNO3 - D4 = 1.5421 (Veley and Manley); 100% acid fumes in air and attracts water vapor as strongly as sulfuric acid. An acid with D = 1.526 heats up when mixed with snow.

Heat of formation (from 1/2 H 2 + 1/2 N 2 + 3/2 O 2):

HNO 3 – steam + 34400 cal
HNO 3 – liquid + 41600 cal
HNO 3 – crystals + 42200 cal
HNO 3 – solution + 48800 cal

Heat of dilution: when adding one particle of H 2 O to HNO 3 - 3.30 Cal, two particles - 4.9 Cal, five particles - 6.7 Cal, ten - 7.3 Cal. Further addition gives an insignificant increase in the thermal effect. In the form of crystals you get:
1) HNO 3 ·H 2 O = H 3 NO 4 - rhombic tablets reminiscent of AgNO 3, melting point = -34° (-38°);
2) HNO 3 (H 2 O) 2 = H 5 NO 5 - needles, melting point -18°.2, stable only below -15°. The crystallization temperature curve of aqueous acid has three eutectics (at -66°.3, at -44°.2, at -43°) and two maxima (HNO 3 H 2 O -38°, HNO 3 3H 2 O -18 °,2). The same special points are observed for the heats of solution and for the turns of the electrical conductivity curve, but on the latter 2HNO 3 ·H 2 O and HNO 3 ·10H 2 O are also noticed. From what has just been said and by analogy with phosphoric acids, it follows that in solutions of nitric acid there is its hydrate HNO 3, but it decomposes very easily, which determines the high reactivity HNO3. Nitric acid containing NO 2 in solution is called smoking(red).

Chemical properties. Pure HNO 3 easily decomposes and turns yellowish due to the reaction 2HNO 3 = 2NO 2 + O 2 + H 2 O and the absorption of the resulting nitrous anhydride. Pure nitric acid and strong nitric acid in general are stable only at low temperatures. The main feature of nitric acid is its extremely strong oxidizing ability due to the release of oxygen. Thus, when acting on metals (except Pt, Rh, Ir, Au, on which HNO 3 has no effect in the absence of chlorine), nitric acid oxidizes the metal, releasing nitrogen oxides, the lower the degree of oxidation, the more energetic the oxidized metal was as a reducing agent. For example, lead (Pb) and tin (Sn) give N 2 O 4; silver - mainly N 2 O 3. Sulfur, especially freshly precipitated, oxidizes easily; phosphorus, when slightly heated, turns into phosphorous acid. Red-hot coal ignites in the vapor of nitric acid and in the nitric acid itself. The oxidizing effect of fuming red acid is greater than that of colorless acid. Iron immersed in it becomes passive and is no longer susceptible to the action of acid. Anhydrous nitric acid or mixed with sulfuric acid has a very strong effect on cyclic organic compounds (benzene, naphthalene, etc.), giving nitro compounds C 6 H 5 H + HNO 3 = C 6 H 5 NO 2 + HOH. Nitration of paraffins occurs slowly, and only under the action of a weak acid ( high degree ionization). As a result of the interaction of substances containing hydroxyl (glycerin, fiber) with nitric acid, nitrate esters are obtained, incorrectly called nitroglycerin, nitrocellulose, etc. All experiments and all work with nitric acid must be carried out in a well-ventilated room, but preferably under a special draft .

Analysis . To detect traces of nitric acid, use: 1) diphenylenedanyl dihydrotriazole (commercially known as “nitron”); 5 or 6 drops of a 10% solution of nitron in 5% acetic acid pour to 5-6 cm 3 of the test solution, adding one drop of H 2 SO 4 to it in advance: in the presence of noticeable amounts of NO 3 ions, a large precipitate is released, with very weak solutions, needle-shaped crystals are released; at 0° even 1/80000 HNO 3 can be opened with nitron; 2) brucine in solution; mix with the test solution and carefully pour it along the wall of the test tube to strong sulfuric acid; at the point of contact of both layers in the test tube, a pinkish-red color is formed, turning from below to greenish.

To determine the amount of HNO 3 in a solution of fuming nitric acid, you need to titrate N 2 O 4 with a solution of KMnO 4, determine the density of the liquid with a hydrometer and subtract the correction for the N 2 O 4 content indicated in a special table.

Industrial methods nitric acid production. Nitric acid is extracted. arr. from saltpeter. Previously, saltpeter mining was carried out in the so-called. “salpetriere”, or “burts”, where, as a result of mixing manure, urine, etc. with old plaster, gradually, partly due to the action of bacteria, oxidation of urea and other organic compounds nitrogen (amines, amides, etc.) in nitric acid, forming calcium nitrate with limestone. On hot days, especially in the south (for example, in India and in Central Asia), the process goes very quickly.

In France in 1813, up to 2,000,000 kg of saltpeter were extracted from saltpeter. 25 large animals produce about 500 kg of saltpeter per year. In some areas, with basic soil rich in animal remains (for example, the Kuban region), there may be a noticeable amount of nitrate in the soil, but not sufficient for extraction. Noticeable quantities were mined in the Ganges valley and are found in our Central Asian fortresses, where reserves of soil containing saltpeter reach up to 17 tons in each place, but the content of saltpeter in it is no more than 3%. Deposits of sodium nitrate - Chilean - were discovered in 1809; they are found mainly in the province of Tarapaca, between 68° 15" and 70° 18" east longitude and 19° 17" and 21° 18" south latitude, but are found both to the south and north (in Peru and Bolivia); their deposit is located at an altitude of 1100 m above sea level. The deposits are about 200 km long, 3-5 km wide, and have an average NaNO 3 content of 30-40%. Reserves, assuming an annual increase in consumption of 50,000 tons, may last for 300 years. In 1913, 2,738,000 tons were exported, but exports to Europe decreased somewhat, although, after a very noticeable drop in exports during the war, they increased slightly again from 1920. Usually on top lies a “fire” (50 cm - 2 m thick), consisting of quartz and feldspathic sand, and under it “kalihe” (25 cm - 1.5 m), containing saltpeter (the deposits are located in the desert next to deposits of salt and boron calcium salt). The composition of "kalihe" is very diverse; it contains NaNO 3 - from 30% to 70%, iodide and iodine salts - up to 2%, sodium chloride - 16-30%, sulfate salts - up to 10%, magnesium salts - up to 6%. The best varieties contain on average: NaNO 3 - 50%, NaCl - 26%, Na 2 SO 4 - 6%, MgSO 4 - 3%. NaNO 3 is dissolved at high temperatures so that much more NaNO 3 goes into the solution than NaCl, the solubility of which increases slightly with temperature. From 3 tons of “kalihe” you get 1 ton of raw saltpeter with an average content of 95-96% saltpeter. From 1 liter of mother brine, 2.5-5 g of iodine is usually obtained. Typically, raw saltpeter is brown in color, due to the admixture of iron oxide. Saltpeter containing up to 1-2% chloride compounds is used for fertilizer. Pure sodium nitrate is colorless, transparent, and non-hygroscopic if it does not contain chloride compounds; crystallizes in cubes. To obtain nitric acid, saltpeter is heated with sulfuric acid; the interaction follows the equation:

NaNO 3 + H 2 SO 4 = HNO 3 + NaSO 4

i.e. acid sulfate is obtained. The latter can be used to produce hydrogen chloride by calcining a mixture of NaHSO 4 and NaCl in muffles. For interaction according to the equation

theoretically, it is necessary to take 57.6 kg of H 2 SO 4 or 60 kg of acid 66° Bẻ per 100 kg of NaNO 3. In fact, to avoid decomposition, 20-30% more sulfuric acid is taken. The interaction is carried out in horizontal cylindrical iron retorts 1.5 m long, 60 cm in diameter, with walls 4 cm thick. Each cylinder contains 75 kg of saltpeter and 75 kg of H 2 SO 4. The vapors are first passed through a ceramic refrigerator, cooled by water, or through an inclined ceramic pipe, then through absorbers: “cylinders” or “bonbons,” i.e., large ceramic “Wulf flasks.” If sulfuric acid 60° Вẻ (71%) is taken and 4 kg of water per 100 kg of saltpeter is placed in the first absorber, then an acid of 40-42° Вẻ (38-41%) is obtained; using acid at 66° Вẻ (99.6%) and dry saltpeter, we get 50° Вẻ (53%); to obtain acid at 36° Вẻ, 8 liters of water are placed in the first absorber, 4 liters in the second, and 2.6 liters in the next ones. Fuming nitric acid is obtained by reacting saltpeter with half the amount of sulfuric acid required by calculation. Therefore, the method produces acid contaminated with nitrosyl chloride and other substances leaving at the beginning of the process, and with nitrogen oxides at the end of distillation. Nitrogen oxides are relatively easy to drive off by blowing a current of air through the acid. It is much more profitable to work in retorts, surrounded by fire on all sides and having a pipe at the bottom for releasing bisulfate containing a noticeable amount of acid. The fact is that cast iron is not corroded by acid if it is sufficiently heated and if contact with fire on all sides ensures that no drops of acid are deposited. In such retorts (1.20 wide and 1.50 m in diameter, with a wall thickness of 4-5 cm), saltpeter is treated with sulfuric acid at the rate of 450 kg and even 610 kg of saltpeter per 660 kg of H 2 SO 4 (66 ° Bẻ). Instead of cylinders, vertical pipes are now often used or these pipes are connected to cylinders.

According to the Guttmann method, decomposition is carried out in cast iron retorts composed of several parts (Fig. 1 and 1a); The parts are connected with putty, usually consisting of 100 parts. iron filings, 5 parts of sulfur, 5 parts of ammonium chloride with as little water as possible; retorts and, if possible, the loading hatch are enclosed in brickwork and heated by furnace gases.

800 kg of saltpeter and 800 kg of 95% sulfuric acid are loaded into the retort and distillation is carried out for 12 hours; this consumes about 100 kg of coal. Cylindrical retorts are also used. The released vapors first enter cylinder 8; then pass a series of ceramic pipes, 12 and 13, placed in a wooden box with water; here the vapors are condensed into nitric acid, which flows through pipes 22 of the Gutman installation, and 23 into collection 28, and condensate from cylinder 8 also enters here; nitric acid that has not condensed in pipes 12 enters through 15a into a tower filled with balls and washed with water; the last traces of acid not absorbed in the tower are captured in cylinder 43a; the gases are carried away through pipe 46a into the chimney. To oxidize the nitrogen oxides formed during distillation, air is mixed into the gases directly at the exit from the retort. If strong sulfuric acid and dried saltpeter are used in production, then colorless 96-97% nitric acid is obtained. Almost all the acid condenses in the pipes, only a small part (5%) is absorbed in the tower, giving 70% nitric acid, which is added to the next load of nitrate. That. the result is colorless nitric acid, devoid of chlorine, with a yield of 98-99% of theory. Gutman's method received widespread due to the simplicity and low cost of installation.

96-100% acid is extracted from saltpeter according to the Valentiner method, by distillation under reduced pressure (30 mm) in cast iron retorts of a mixture of 1000 kg NaNO 3, 1000 kg H2SO 4 (66 ° Вẻ) and such an amount of weak acid HNO 3 that add 100 kg of water with it. The distillation lasts 10 hours, with air being introduced into the alloy all the time. The interaction occurs at 120°, but at the end of the process a “crisis” occurs (1 hour) and strong shocks are possible (at 120-130°). After this, the heating is brought to 175-210°. Proper thickening and acid capture is very important. Vapors from the retort enter the cylinder, from it into 2 highly cooled coils, from them into a collection (such as a Wulf flask), followed by a coil again and then 15 cylinders, behind which a pump is placed. With a 1000 kg load of NaNO 3 in 6-8 hours, 600 kg of HNO 3 (48° Вẻ) is obtained, i.e. 80% of the norm.

To obtain nitric acid from Norwegian nitrate (calcium), the latter is dissolved, strong nitric acid is added and mixed sulfuric acid, after which the nitric acid is filtered from the gypsum.

Storage and packaging. To store nitric acid, you can use glass, fireclay and pure aluminum (no more than 5% impurities) dishes, as well as dishes made of special silicon acid-resistant Krupp steel (V2A). Because when strong nitric acid acts on wood, sawdust, rags, wetted vegetable oil, etc. outbreaks and fires are possible (for example, if a bottle bursts during transportation), then nitric acid can only be transported in special trains. Turpentine ignites especially easily when heated when it comes into contact with strong nitric acid.

Application: 1) in the form of salts for fertilizer, 2) for the production of explosives, 3) for the production of semi-finished products for dyes, and partly the dyes themselves. Ch. arr. salts of nitric acid or nitrate (sodium, ammonium, calcium and potassium) are used for fertilizers. In 1914, world consumption of nitrogen in the form of Chilean nitrate reached 368,000 tons and in the form of nitric acid from the air - 10,000 tons. In 1925, consumption should have reached 360,000 tons of nitric acid from the air. The consumption of nitric acid increases greatly during war due to the expenditure on explosives, the main of which are nitroglycerin and nitrocellulose. different types, nitro compounds (nitrotoluene, TNT, melinite, etc.) and substances for fuses (mercury fulminate). IN peacetime nitric acid is spent on the production of nitro compounds, for example, nitrobenzene, for the transition to dyes through aniline, obtained from nitrobenzene by reduction. Significant amount nitric acid is used for etching metals; salts of nitric acid (saltpeter) are used for explosives (ammonium nitrate - in smokeless, potassium nitrate - in black powder) and for fireworks (barium nitrate - for green).

Nitric acid standard. The nitric acid standard exists so far only in the USSR and was approved by the Standardization Committee at the STO as an all-Union mandatory standard (OST-47) for acid at 40° Bẻ. The standard sets the HNO 3 content in nitric acid to 61.20% and limits the content of impurities: sulfuric acid no more than 0.5%, chlorine no more than 0.8%, iron no more than 0.01%, solid residue no more than 0.9 %; standard nitric acid should not contain sediment. The standard regulates the relationship between the seller and the buyer, strictly regulating the sampling and analysis methods. The content of nitric acid is determined by adding NaOH to the acid and back titrating with the acid. The content of sulfuric acid is determined in the form of BaSO 4 by precipitation of BaCl 2. The chlorine content is determined by titration in an alkaline medium with silver nitrate. The iron content is determined by precipitation of sesquioxides with ammonia, reduction of oxide iron to ferrous iron and subsequent titration of KMnO 4. The packaging of nitric acid is not yet standard. Without touching on the size, weight and quality of the container, the standard stipulates the packaging of nitric acid in glass containers and gives instructions on how to pack and seal it.

Preparation of nitric acid.

I. From the air. The synthesis of nitric acid from air under the action of a voltaic arc is repeated until to a certain extent a process occurring in nature under the influence of discharges atmospheric electricity. Cavendish was the first to observe (in 1781) the formation of nitrogen oxides during the combustion of H 2 in air, and then (in 1784) during electric spark through the air. Mutman and Gopher in 1903 were the first to try to study the equilibrium: N 2 + O 2 2NO. Passing a voltaic arc through the air AC at 2000-4000 V, they practically achieved a NO concentration of 3.6 to 6.7 vol%. Their energy consumption per 1 kg of HNO 3 reached 7.71 kWh. Nernst then studied this equilibrium by passing air through an iridium tube. Further, Nernst, Jellinek and other researchers worked in the same direction. By extrapolation experimental results By studying the equilibrium between air and nitrogen oxide, Nernst was able to calculate that on the right side of the equation, at a temperature of 3750° (i.e., approximately at the temperature of a voltaic arc), a content of 7 volume % NO is established.

The priority of the idea of technical use of a voltaic arc for fixing atmospheric nitrogen belongs to the French researcher Lefebre, who back in 1859 patented her method of producing nitric acid from air in England. But at that time the cost electrical energy was too high for Lefebre's method to be of practical use. It is also worth mentioning the patents of McDougal (An. P. 4633, 1899) and the Bradley and Lovejoy method, implemented on a technical scale, operated in 1902 by the American company Atmospheric Products С° (with 1 million dollars of capital) with using the energy of Niagara Falls. The attempts to use a voltage of 50,000 V to fix atmospheric nitrogen, made by Kowalski and his collaborator I. Moscytski, should also be dated to this time. But the first significant success in the fabrication of nitric acid from air was brought by the historical idea of the Norwegian engineer Birkeland, which was to use the ability of the latter to stretch in a strong electromagnetic field to increase the yield of nitrogen oxides when passing a voltaic arc through the air. Birkeland combined this idea with another Norwegian engineer, Eide, and translated it into a technical installation that immediately provided a cost-effective opportunity to obtain nitric acid from air. Due to the constant change in the direction of the current and the action of the electromagnet, the resulting voltaic arc flame has a constant tendency to swell in different directions, which leads to the formation of a voltaic arc that moves rapidly all the time at a speed of up to 100 m/sec, creating the impression of a calmly burning wide electric sun with a diameter of 2 m or more. A strong stream of air is continuously blown through this sun, and the sun itself is enclosed in a special furnace made of refractory clay, bound in copper (Fig. 1, 2 and 3).

The hollow electrodes of the voltaic arc are cooled from the inside by water. Air through channels A in the fireclay lining of the furnace it enters the arc chamber b; through the oxidized gas leaves the furnace and is cooled using its heat to heat the boilers of the evaporators. After this, NO enters the oxidation towers, where it is oxidized by atmospheric oxygen to NO 2. Last process is an exothermic process (2NO + O 2 = 2NO 2 + 27Cal), and therefore conditions that increase heat absorption significantly promote the reaction in this direction. Next, nitrogen dioxide is absorbed by water according to the following equations:

3NO 2 + H 2 O = 2HNO 3 + NO
2NO 2 + H 2 O = HNO 3 + HNO 2

In another method, the reacting mixture of gases is cooled below 150° before absorption; at this temperature inverse expansion– NO 2 = NO + O has almost no place. Bearing in mind that under certain conditions the equilibrium NO + NO 2 N 2 O 3 is established with a maximum content of N 2 O 3, it can be obtained by pouring hot nitrite gases even before their complete oxidation, at a temperature of 200 to 300 °, with a solution of soda or caustic soda, instead of nitrate salts - pure nitrites (Norsk Hydro method). When leaving the furnace, the blown air contains from 1 to 2% nitrogen oxides, which are immediately captured by counter jets of water and then neutralized with lime to form calcium, the so-called. "Norwegian" saltpeter. Carrying out the process itself N 2 + O 2 2NO - 43.2 Cal requires the expenditure of a relatively small amount of electrical energy, namely: to obtain 1 ton of bound nitrogen in the form of NO only 0.205 kW-year; Meanwhile, in the best modern installations it is necessary to spend 36 times more, i.e. about 7.3 and up to 8 kW-years per 1 ton. In other words, over 97% of the energy expended does not go to the formation of NO, but to create for this process favorable conditions. To shift the equilibrium towards the highest possible NO content, it is necessary to use a temperature from 2300 to 3300° (NO content at 2300° is 2 vol% and for 3300° - 6 vol%), but at such temperatures 2NO quickly decomposes back into N 2 + O 2. Therefore, in a small fraction of a second it is necessary to remove gas from hot regions to colder ones and cool it to at least 1500°, when the decomposition of NO proceeds more slowly. Equilibrium N 2 + O 2 2NO is established at 1500° in 30 hours, at 2100° in 5 seconds, at 2500° in 0.01 seconds. and at 2900° - in 0.000035 sec.

The method of Schonherr, an employee of BASF, differs from the method of Birkeland and Eide in significant improvements. In this method, instead of a pulsating and intermittently acting voltaic arc flame variable current, apply a calm flame of high permanent current This prevents frequent blowing out of the flame, which is very harmful to the process. The same result, however, can be achieved with an alternating current voltaic arc, but by blowing air through the burning flame not in a straight line, but in the form of a vortex wind along the voltaic arc flame. Therefore, the oven could designed in the form of a rather narrow metal tube, moreover, so that the arc flame does not touch its walls. The design diagram of the Schongherr furnace is shown in Fig. 4.

A further improvement in the arc method is made by the Pauling method (Fig. 5). The electrodes in the combustion furnace look like horn dischargers. The voltaic arc 1 m long formed between them is blown upward by a strong stream of air. In the most bottleneck If the flame breaks, the arc is re-ignited using additional electrodes.

A slightly different design of a furnace for the oxidation of air nitrogen was patented by I. Mościtsky. One of both electrodes (Fig. 6) has the shape of a flat disk and is located very far from the other electrode close range. The upper electrode is tubular, and neutral gases flow through it in a fast stream, then spreading in a cone.

The flame of the voltaic arc is given in Roundabout Circulation under the influence electromagnetic field, and a fast cone-shaped gas stream prevents short circuits. Detailed description the entire installation is given in W. Waeser, Luftstickstoff-Industrie, p. 475, 1922. One plant in Switzerland (Chippis, Wallis) operates according to the method of I. Moscicki, producing 40% HNO 3. Another plant in Poland (Bory-Jaworzno) is designed for 7000 kW and should produce concentrated HNO 3 and (NH 4) 2 SO 4. To improve the yield of nitrogen oxides and to increase the flame of the voltaic arc, in lately used as original product not air, but a more oxygen-rich mixture of nitrogen and oxygen, with a ratio of 1: 1. The French plant in Laroche-de-Rham works with this mixture with very good results.

It is advisable to condense the resulting nitrogen tetroxide N 2 O 4 into a liquid by cooling to -90°. Such liquid nitrogen tetroxide, obtained from pre-dried gases - oxygen and air, does not react with metals and therefore can be transported in steel bombs and used for the production of HNO 3 in strong concentrations. Toluene was used as a coolant in this case at one time, but due to the inevitable seepage of nitrogen oxides and their effect on toluene, terrible explosions occurred at the Tschernewitz (in Germany) and Bodio (in Switzerland) plants, destroying both enterprises. Extraction of N 2 O 4 from a gas mixture. also achieved through the absorption of N 2 O 4 by silica gel, which releases the absorbed N 2 O 4 back when heated.

II. Contact oxidation of ammonia. All the described methods for producing synthetic nitric acid directly from the air, as already indicated, are profitable only if cheap hydroelectric energy is available. The problem of bound nitrogen (see Nitrogen) could not be considered finally resolved if a method for producing relatively cheap synthetic nitric acid had not been found. The absorption of bound nitrogen from fertilizers by plants is especially facilitated if these fertilizers are salts of nitric acid. Ammonium compounds introduced into the soil must first undergo nitrification in the soil itself (see Nitrogen fertilizers). In addition, nitric acid, along with sulfuric acid, is the basis of numerous industries chemical industry and military affairs. Production of explosives and smokeless powder(TNT, nitroglycerin, dynamite, picric acid and many others), aniline dyes, celluloid and rayon, many medicines, etc. are impossible without nitric acid. That is why in Germany, which was cut off from the source of Chilean saltpeter by a blockade during the World War and at the same time did not have cheap hydroelectric energy, the production of synthetic nitric acid developed to a large extent using the contact method, starting from coal coal or synthetic ammonia by oxidizing it with atmospheric oxygen with the participation of catalysts. During the war (1918), Germany produced up to 1000 tons of nitric acid and ammonium nitrate per day.

Back in 1788, Milner in Cambridge established the possibility of the oxidation of NH 3 into nitrogen oxides under the action of manganese peroxide when heated. In 1839, Kuhlman established the contact action of platinum during the oxidation of ammonia with air. Technically, the method of oxidizing ammonia to nitric acid was developed by Ostwald and Brouwer and patented by them in 1902 (Interestingly, in Germany, Ostwald’s application was rejected due to recognition of priority for French chemist Kulman.) Under the action of finely crushed platinum and the slow flow of the gas mixture, oxidation proceeds according to the reaction 4NH 3 + ZO 2 = 2N 2 + 6H 2 O. Therefore, the process should be strictly regulated both in the sense of the significant speed of movement of the gas jet blown through the contact “converter”, and in the sense of the composition of the gas mixture. The mixture of gases entering the “converters” should. previously thoroughly cleaned of dust and impurities that could “poison” the platinum catalyst.

It can be assumed that the presence of platinum causes the decomposition of the NH 3 molecule and the formation of an unstable intermediate compound of platinum with hydrogen. In this case, nitrogen in statu nascendi is subject to oxidation by atmospheric oxygen. The oxidation of NH 3 to HNO 3 proceeds through the following reactions:

4NH 3 + 5O 2 = 4NO + 6H 2 0;

cooled colorless NO gas, being mixed with a new portion of air, spontaneously oxidizes further to form NO 2 or N 2 O 4:

2NO + O 2 = 2NO 2, or N 2 O 4;

the dissolution of the resulting gases in water in the presence of excess air or oxygen is associated with further oxidation according to the reaction:

2NO 2 + O + H 2 O = 2HNO 3,

after which HNO 3 is obtained, with a strength of approximately 40 to 50%. By distilling the resulting HNO 3 with strong sulfuric acid, concentrated synthetic nitric acid can finally be obtained. According to Ostwald, the catalyst must consist of metallic platinum coated with part or completely spongy platinum or platinum black.

The reaction should take place when the red heat has barely begun and at a significant flow rate of the gas mixture, consisting of 10 or more parts of air per 1 hour NH 3. The slow flow of the gas mixture promotes complete collapse NH 3 to elements. With a platinum contact grid of 2 cm, the gas flow velocity should be 1-5 m/sec, i.e. the time of contact of gas with platinum should not exceed 1/100 sec. Optimum temperatures are around 300°. The gas mixture is preheated. The higher the flow rate of the gas mixture, the greater the NO output. Working with a very thick platinum mesh (catalyst) with a mixture of ammonia and air containing about 6.3% NH 3, Neumann and Rose obtained the following results at a temperature of 450 ° (with a contact surface of platinum of 3.35 cm 2):

A higher or lower NH 3 content is also of great importance for the direction of the chemical process, which can proceed either according to the equation: 4NH 3 + 5O 2 = 4NO + 6H 2 O (with a content of 14.38% NH 3), or according to the equation: 4NH 3 + 7O 2 = 4NO 2 + 6H 2 O (with a mixture content of 10.74% NH 3). With less success than platinum, maybe. Other catalysts were also used (iron oxide, bismuth, cerium, thorium, chromium, vanadium, copper). Of these, only the use of iron oxide at a temperature of 700-800°, with a yield of 80 to 85% NH 3, deserves attention.

Significant role in oxidative process The transition of NH 3 to HNO 3 is influenced by temperature. The ammonia oxidation reaction itself is exothermic: 4NH 3 + 5O 2 = 4NO + 6H 2 O + 215.6 Cal. Only initially it is necessary to warm up the contact apparatus, further reaction is underway due to its own heat. Technical design of “converters” for ammonia oxidation different systems is clear from the given figures (Fig. 7-8).

The scheme for the production of HNO 3 according to the currently accepted Franck-Caro method is shown in Fig. 9.

In fig. 10 shows a diagram of the oxidation of NH 3 at the factory of Meister Lucius and Brünning in Hechst.

In modern installations, the oxidation of NH 3 to NO is carried out with a yield of up to 90%, and the subsequent oxidation and absorption of the resulting nitrogen oxides by water - with a yield of up to 95%. Thus, the whole process gives a yield of bound nitrogen of 85-90%. Obtaining HNO 3 from nitrate currently costs (in terms of 100% HNO 3) $103 per 1 ton, using the arc process, $97.30 per 1 ton, while 1 ton of HNO 3 obtained by oxidation of NH -3 costs only $85.80. It goes without saying that these numbers could be are only approximate and largely depend on the size of the enterprise, the cost of electrical energy and raw materials, but still they show that the contact method for producing HNO 3 is destined to occupy a dominant position in the near future compared to other methods.

34. Phosphorus

Phosphorus(P) is in the 3rd period, in group V, of the main subgroup of the periodic system of D.I. Mendeleev. Serial number 15, nuclear charge +15, Ar = 30.9738 a.u. m... has 3 energy levels, there are 15 electrons on the energy shell, of which 5 are valence. Phosphorus has a d-sublevel. Electronic configuration P: 1 s2 2s2 2p63 s2 3p33d0. Characteristic is sp3 hybridization, less commonly sp3d1. The valence of phosphorus is III, V. The most characteristic oxidation state is +5 and -3, less characteristic: +4, +1, -2, -3. Phosphorus can exhibit both oxidizing and reducing properties: accepting and donating electrons.

Molecule structure: the ability to form β-bonds is less pronounced than that of nitrogen - at ordinary temperatures in the gas phase, phosphorus is presented in the form of P4 molecules, having the shape of equilateral pyramids with angles of 60°. The bonds between atoms are covalent, non-polar. Each P atom in the molecule is connected by other atoms by ?-bonds.

Physical properties: Phosphorus forms three allotropes: white, red and black. Each modification has its own melting and freezing point.

Chemical properties:

1) when heated, P4 reversibly dissociates:

2) above 2000 °C P2 disintegrates into atoms:

3) phosphorus forms compounds with non-metals:

Directly combines with all halogens: 2P + 5Cl2 = 2PCl5.

When interacting with metals, phosphorus forms phosphides:

Combining with hydrogen, it forms phosphine gas: Р4 + 6Н2 = 4РН3?.

When interacting with oxygen, it forms the anhydride P2O5: P4 + 5O2 = 2P2O5.

Receipt: phosphorus is obtained by calcining the mixture Ca3(P O4 )2 with sand and coke in an electric furnace at a temperature of 1500 °C without air access: 2Ca3(PO4)2 + 1 °C + 6SiO2 = 6CaSiO3 + 1 °CO + P4?.

Phosphorus does not occur in nature in its pure form, but is formed as a result of chemical activity. The main natural phosphorus compounds are the following minerals: Ca3(PO4)2 – phosphorite; Ca3(PO4)2?CaF2 (or CaCl) or Ca3(PO4)2?Ca(OH)2 – apatite. The biological significance of phosphorus is great. Phosphorus is part of some plant and animal proteins: protein in milk, blood, brain and nervous tissue. A large amount of it is contained in the bones of vertebrates in the form of compounds: 3Ca3(PO4)2?Ca(OH)2 and 3Ca3(PO4)2?CaCO3?H2O. Phosphorus is an essential component of nucleic acids, playing a role in the transmission of hereditary information. Phosphorus is found in tooth enamel and in tissues in the form of lecithin - a compound of fats with phosphoroglycerol esters.