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 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 aqueous solutions of nitric acid have a minimum elasticity, i.e., adding water to the acid and vice versa lowers the vapor elasticity (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 повышенном давлении - с большим содержанием HNO 3 . Кислота D = 1,503 (85%), очищенная продуванием воздуха от N 2 О 4 , дает при перегонке остаток с 77,1% HNО 3 . Кислота D = 1,55 (99,8%) дает при перегонке сначала сильно окрашенный окислами азота раствор D = 1,62, а в остатке кислоту D = 1,49. Т. о. в остатке при перегонке азотной кислоты всегда оказывается кислота, соответствующая минимуму упругости (максимуму температуры кипения). Безводную кислоту можно получить лишь при смешивании крепкой (99,1%) азотной кислоты с азотным ангидридом.
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 of HNO 3. 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 are poured into 5-6 cm 3 of the test solution, adding to it in advance one drop of H 2 SO 4: in the presence of noticeable amounts of NO 3 ions, a copious precipitate is released, in 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 for producing nitric acid. 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 nitrogen compounds (amines, amides, etc.) occurs 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 also found further 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 Guttman 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 of iron filings, 5 parts of sulfur, 5 parts of ammonium chloride with as little water as possible; The 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 has become widespread due to its 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, and air is 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 amounts of nitric acid are 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 containers, the standard stipulates the packaging of nitric acid in glassware 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 repeats to a certain extent the process that occurs in nature under the influence of discharges of 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) when an electric spark passes through the air. Mutman and Gopher in 1903 were the first to try to study the equilibrium: N 2 + O 2 2NO. By passing a voltaic arc of alternating current at 2000-4000 V through the air, they practically achieved an 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 extrapolating the experimental results of studying the equilibrium between air and nitrogen oxide, Nernst was able to calculate that on the right side of the equation a content of 7 volume % NO is established at a temperature of 3750 ° (i.e., approximately at the temperature of the voltaic arc).
The priority of the idea of technically using 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 of electrical energy was too high for Lefebre's method to achieve practical significance. It is also worth mentioning the patents of McDougal (An. P. 4633, 1899) and the Bradley and Lovejoy method implemented on a technical scale, exploited in 1902 by the American company Atmospheric Products С° (with 1 million dollars of capital) with energy use 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 attributed 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 with 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, the reverse decomposition – NO 2 = NO + O – almost does not take 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 towards the formation of NO, but towards creating favorable conditions for this process. 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, a BASF employee, is a significant improvement over the Birkeland and Eide method. 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 narrowest place of the broken flame, 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 at a very close distance from the other electrode. The upper electrode is tubular, and neutral gases flow through it in a fast stream, then spreading in a cone.
The flame of a voltaic arc is set in a circular motion under the influence of an 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 The starting product is 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 such a 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 branches of the chemical industry and military affairs. The production of explosives and smokeless gunpowder (TNT, nitroglycerin, dynamite, picric acid, and many others), aniline dyes, celluloid and rayon, many medicines, etc. is 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 the French chemist Kuhlmann.) Under the action of finely divided 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 be 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 occurs according to 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 the complete decomposition of 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):
The greater or lesser content of NH 3 is also of great importance for the direction chemical process, which can go 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 content of mixture 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.
Temperature plays a significant role in the oxidative process of the transition of NH 3 to HNO 3. The ammonia oxidation reaction itself is exothermic: 4NH 3 + 5O 2 = 4NO + 6H 2 O + 215.6 Cal. Only initially it is necessary to heat up the contact apparatus; then the reaction occurs due to its own heat. The technical design of “converters” for the oxidation of ammonia of different systems is clear from the figures given (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.
See also
Nitric acid is one of the main nitrogen compounds. Chemical formula - HNO 3. So what physical and chemical properties does this substance have?
Physical properties
Pure nitric acid is colorless, has a pungent odor, and has the property of “smoking” when exposed to air. The molar mass is 63 g/mol. At a temperature of -42 degrees it becomes solid physical state and turns into a snow-white mass. Anhydrous nitric acid boils at 86 degrees. When mixed with water, it forms solutions that differ from each other in concentration.
This substance is monobasic, that is, it always has one carboxyl group. Among the acids that are powerful oxidizing agents, nitric acid is one of the strongest. It reacts with many metals and non-metals, organic compounds due to nitrogen reduction
Nitrates are salts of nitric acid. They are most often used as fertilizers in agriculture.
Chemical properties
The electronic and structural formula of nitric acid is depicted as follows:
Rice. 1. Electronic formula of nitric acid.
Concentrated nitric acid is exposed to light and, under its influence, is capable of decomposing into nitrogen oxides. The oxides, in turn, interact with the acid, dissolve in it and give the liquid a yellowish tint:
4HNO 3 =4NO 2 +O 2 +2H 2 O
The substance should be stored in a cool and dark place. As its temperature and concentration increase, the decomposition process occurs much faster. Nitrogen in a nitric acid molecule always has a valence of IV, an oxidation state of +5, and a coordination number of 3.
Since nitric acid is a very strong acid, in solutions it completely decomposes into ions. It reacts with basic oxides, with bases, and with salts of weaker and more volatile acids.
Rice. 2. Nitric acid.
This monobasic acid is a strong oxidizing agent. Nitric acid attacks many metals. Depending on the concentration, activity of the metal and reaction conditions, it can be reduced with the simultaneous formation of a nitric acid salt (nitrate) to compounds.
When nitric acid reacts with low-active metals, NO 2 is formed:
Cu+4HNO 3 (conc.)=Cu(NO 3) 2 +2NO 2 +2H 2 O
Dilute nitric acid in this situation is reduced to NO:
3Cu+8HNO 3 (diluted)=3Сu(NO 3) 2 +2NO+4H 2 O
If more active metals react with dilute nitric acid, NO 2 is released:
4Mg+10HNO 3 (diluted)=4Mg(NO 3) 2 +N 2 O+5H 2 O
Very dilute nitric acid, when interacting with active metals, is reduced to ammonium salts:
4Zn+10HNO 3 (very dilute)=4Zn(NO 3) 2 +NH 4 NO 3 +3H 2 O
Au, Pt, Rh, Ir, Ta, Ti are stable in concentrated nitric acid. It “passivates” the metals Al, Fe, Cr as a result of the formation of oxide films on the surface of the metals.
A mixture formed from one volume of concentrated nitric and three volumes of concentrated hydrochloric (hydrochloric) acid is called “aqua regia”.
Rice. 3. Royal vodka.
Non-metals are oxidized with nitric acid to the corresponding acids, and nitric acid, depending on the concentration, is reduced to NO or NO 2:
C + 4HNO 3 (conc.) = CO 2 +4NO 2 +2H 2 O
S+6HNO 3 (conc.)=H 2 SO 4 +6NO 2 +2H 2 O
Nitric acid is capable of oxidizing some cations and anions, as well as inorganic covalent compounds, such as hydrogen sulfide.
3H 2 S+8HNO 3 (diluted)= 3H 2 SO 4 +8NO+4H 2 O
Nitric acid interacts with many organic substances, and one or more hydrogen atoms in the molecule of the organic substance are replaced by nitro groups - NO 2. This process is called nitration.
Structural formula
True, empirical, or gross formula: HNO3
Chemical composition of Nitric acid
Molecular weight: 63.012
Nitric acid ( HNO3) is a strong monobasic acid. Solid nitric acid forms two crystal modifications with monoclinic and orthorhombic lattices.
Nitric acid mixes with water in any ratio. In aqueous solutions, it almost completely dissociates into ions. Forms an azeotropic mixture with water with a concentration of 68.4% and boiling point 120 °C at normal atmospheric pressure. Two solid hydrates are known: monohydrate (HNO 3 ·H 2 O) and trihydrate (HNO 3 ·3H 2 O).
Nitrogen in nitric acid is tetravalent, oxidation state +5. Nitric acid is a colorless liquid that fumes in air, melting point −41.59 °C, boiling point +82.6 °C (at normal atmospheric pressure) with partial decomposition. Nitric acid mixes with water in all proportions. Aqueous solutions of HNO 3 with a mass fraction of 0.95-0.98 are called “fuming nitric acid”, with a mass fraction of 0.6-0.7 - concentrated nitric acid. Forms an azeotropic mixture with water (mass fraction 68.4%, d20 = 1.41 g/cm, T bp = 120.7 °C)
Highly concentrated HNO 3 is usually brown in color due to the decomposition process that occurs in the light. When heated, nitric acid decomposes according to the same reaction. Nitric acid can be distilled without decomposition only under reduced pressure (the indicated boiling point at atmospheric pressure was found by extrapolation).
Gold, some platinum group metals and tantalum are inert to nitric acid over the entire concentration range, other metals react with it, the course of the reaction being determined by its concentration.
Nitric acid at any concentration exhibits the properties of an oxidizing acid, with nitrogen being reduced to an oxidation state from +5 to −3. The depth of reduction depends primarily on the nature of the reducing agent and the concentration of nitric acid.
A mixture of nitric and sulfuric acids is called “melange”.
Nitric acid is widely used to obtain nitro compounds.
A mixture of three volumes of hydrochloric acid and one volume of nitric acid is called “aqua regia.” Aqua regia dissolves most metals, including gold and platinum. Its strong oxidizing abilities are due to the resulting atomic chlorine and nitrosyl chloride.
Nitric acid is a strong acid. Its salts - nitrates - are obtained by the action of HNO 3 on metals, oxides, hydroxides or carbonates. All nitrates are highly soluble in water. Nitrate ion does not hydrolyze in water. Nitrates are widely used as fertilizers. Moreover, almost all nitrates are highly soluble in water, so there are extremely few of them in nature in the form of minerals; the exceptions are Chilean (sodium) nitrate and Indian nitrate (potassium nitrate). Most nitrates are obtained artificially.
In terms of the degree of impact on the body, nitric acid belongs to substances of the 3rd hazard class. Its fumes are very harmful: the fumes cause irritation of the respiratory tract, and the acid itself leaves long-healing ulcers on the skin. When exposed to skin, a characteristic yellow coloration of the skin occurs due to the xanthoprotein reaction. When heated or exposed to light, the acid decomposes to form highly toxic nitrogen dioxide NO 2 (a brown gas). MPC for nitric acid in the air of the working area for NO 2 2 mg/m 3.
A monobasic strong acid, which is a colorless liquid under standard conditions, which turns yellow during storage, can be in a solid state, characterized by two crystalline modifications (monoclinic or rhombic lattice), at temperatures below minus 41.6 °C. This substance with chemical formula— HNO3 — is called nitric acid. It has a molar mass of 63.0 g/mol, and its density corresponds to 1.51 g/cm³. The boiling point of the acid is 82.6 °C, the process is accompanied by decomposition (partial): 4HNO3 → 2H2O + 4NO2 + O2. An acid solution with a mass fraction of the main substance equal to 68% boils at a temperature of 121 °C. pure substance corresponds to 1.397. The acid can be mixed with water in any ratio and, being a strong electrolyte, almost completely decomposes into H+ and NO3- ions. Solid forms - trihydrate and monohydrate have the formula: HNO3. 3H2O and HNO3. H2O respectively.
Nitric acid is corrosive, toxic substance and a strong oxidizing agent. Since the Middle Ages, the name “strong water” (Aqua fortis) has been known. Alchemists who discovered the acid in the 13th century gave it this name, convinced of its extraordinary properties (it corroded all metals except gold), which were a million times greater than the strength of acetic acid, which in those days was considered the most active. But three centuries later it was found that even gold can be corroded by a mixture of acids such as nitric and hydrochloric in a volume ratio of 1:3, which for this reason was called “aqua regia.” The appearance of a yellow tint during storage is explained by the accumulation of nitrogen oxides in it. On sale, acid is often found with a concentration of 68%, and when the content of the main substance is more than 89%, it is called “fuming”.
The chemical properties of nitric acid distinguish it from dilute sulfuric or hydrochloric acids in that HNO3 is a stronger oxidizing agent, so hydrogen is never released in reactions with metals. Due to its oxidizing properties, it also reacts with many non-metals. In both cases, nitrogen dioxide NO2 is always formed. In redox reactions, nitrogen reduction occurs to varying degrees: HNO3, NO2, N2O3, NO, N2O, N2, NH3, which is determined by the acid concentration and the activity of the metal. The molecules of the resulting compounds contain nitrogen with the oxidation state: +5, +4, +3, +2, +1, 0, +3, respectively. For example, copper is oxidized with concentrated acid to copper (II) nitrate: Cu + 4HNO3 → 2NO2 + Cu(NO3)2 + 2H2O, and phosphorus to metaphosphoric acid: P + 5HNO3 → 5NO2 + HPO3 + 2H2O.
Otherwise, dilute nitric acid interacts with non-metals. Using the example of the reaction with phosphorus: 3P + 5HNO3 + 2H2O → 3H3PO4 + 5NO, it can be seen that nitrogen is reduced to the divalent state. As a result, nitrogen monoxide is formed, and phosphorus is oxidized to Concentrated nitric acid mixed with hydrochloric acid dissolves gold: Au + 4HCl + HNO3 → NO + H + 2H2O and platinum: 3Pt + 18HCl + 4HNO3 → 4NO +3H2 + 8H2O. In these reactions to initial stage hydrochloric acid is oxidized by nitric acid with the release of chlorine, and then the metals form complex chlorides.
Nitric acid is produced on an industrial scale in three main ways:
- The first is the interaction of salts with sulfuric acid: H2SO4 + NaNO3 → HNO3 + NaHSO4. Previously, this was the only method, but with the advent of other technologies, it is now used in laboratory conditions to obtain fuming acid.
- The second is the arc method. When air is blown through at a temperature of 3000 to 3500 °C, part of the nitrogen in the air reacts with oxygen, resulting in the formation of nitrogen monoxide: N2 + O2 → 2NO, which, after cooling, is oxidized to nitrogen dioxide (at high temperatures, the monoxide does not interact with oxygen): O2 + 2NO → 2NO2. Then, practically all nitrogen dioxide, with an excess of oxygen, dissolves in water: 2H2O + 4NO2 + O2 → 4HNO3.
- The third is the ammonia method. Ammonia is oxidized on a platinum catalyst to nitrogen monoxide: 4NH3 + 5O2 → 4NO + 6H2O. The resulting nitrous gases cool and form nitrogen dioxide, which is absorbed by water. This method produces acid with a concentration of 60 to 62%.
Nitric acid is widely used in industry to produce drugs, dyes, nitrogen fertilizers and nitric acid salts. In addition, it is used to dissolve metals (eg copper, lead, silver) that do not react with other acids. In jewelry it is used to determine gold in an alloy (this is the main method).
Nitric acid- a colorless, “smoking” liquid in air with a pungent odor. Chemical formula HNO3.
Physical properties. At a temperature of 42 °C it hardens in the form of white crystals. Anhydrous nitric acid boils at atmospheric pressure and 86 °C. Mixes with water in arbitrary proportions.
When exposed to light, concentrated HNO3 decomposes into nitrogen oxides:
HNO3 is stored in a cool and dark place. The valency of nitrogen in it is 4, the oxidation state is +5, the coordination number is 3.
HNO3 is a strong acid. In solutions it completely disintegrates into ions. Interacts with basic oxides and bases, and with salts of weaker acids. HNO3 has strong oxidizing ability. Capable of being reduced with the simultaneous formation of nitrate to compounds, depending on the concentration, activity of the interacting metal and conditions:
1) concentrated HN03, interacting with low-active metals, is reduced to nitrogen oxide (IV) NO2:
2) if the acid is dilute, then it is reduced to nitric oxide (II) NO:
3) more active metals reduce dilute acid to nitrogen oxide (I) N2O:
A very dilute acid is reduced to ammonium salts:
Au, Pt, Rh, Ir, Ta, Ti do not react with concentrated HNO3, and Al, Fe, Co and Cr are “passivated”.
4) HNO3 reacts with non-metals, reducing them to the corresponding acids, and itself is reduced to oxides:
5) HNO3 oxidizes some cations and anions and inorganic covalent compounds.
6) interacts with many organic compounds - nitration reaction.
Industrial production of nitric acid: 4NH3 + 5O2 = 4NO + 6H2O.
Ammonia– NO transforms into NO2, which, with water in the presence of atmospheric oxygen, produces nitric acid.
Catalyst – platinum alloys. The resulting HNO3 is no more than 60%. If necessary, it is concentrated. The industry produces diluted HNO3 (47–45%) and concentrated HNO3 (98–97%). Concentrated acid is transported in aluminum tanks, diluted acid is transported in tanks made of acid-resistant steel.
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?.
In nature, phosphorus is pure form does not occur, 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.