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 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 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 a negligible increase 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 is easily decomposed and colored yellowish color 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. To cyclic organic compounds(benzene, naphthalene, etc.) Anhydrous nitric acid or mixed with sulfuric acid acts very strongly, giving nitro compounds C 6 H 5 H + HNO 3 = C 6 H 5 NO 2 + NOH. 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 into 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; 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), it is possible that 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 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 temperature so that much more NaNO 3 passes 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, 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 prevents acid drops from settling. 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 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 using 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 nitrocelluloses of various types, nitro compounds (nitrotoluene, TNT, melinite, etc.) and substances for fuses (mercury fulminate). IN Peaceful time 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 standard mandatory standard(OST-47) for acid at 40° Вẻ. 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 alkaline environment 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 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 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.
Idea priority technical use voltaic arc for fixation atmospheric nitrogen belongs to the French researcher Lefebre, who back in 1859 patented in England her method of producing nitric acid from the air. But at that time the cost electrical energy was too high for the Lefebre method to obtain 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 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 attributed to this time. But the first significant success in the fabrication of nitric acid from the air came from historical idea 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 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 flame of the voltaic arc has a constant tendency to, as it were, inflate into different sides, 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 evaporation apparatus. After this, NO enters the oxidation towers, where it is oxidized by atmospheric oxygen to NO 2. The latter process is an exothermic process (2NO + O 2 = 2NO 2 + 27Cal), and therefore conditions that increase heat absorption significantly favor 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
2NO2 + H2O = HNO3 + HNO2
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 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 areas 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 still intermittent flame of a voltaic arc 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. A 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 nitrogen in the air was patented by I. Moscicki. 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 influenced electromagnetic field, and the fast cone-shaped gas jet 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 I. Mościcki’s method, 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, not air, but a more oxygen-rich mixture of nitrogen and oxygen, with a ratio of 1: 1, has recently been used as the starting product. 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 leakage 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 gas mixture m.b. 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 cost-effective 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 nitrate during the World War by a blockade 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 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 the 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 gas NO, 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):
More or less NH 3 content also has great importance for directions 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%. The production of HNO 3 from nitrate currently costs (in terms of 100% HNO 3) $103 per 1 ton, according to 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
Nitric acid(HNO 3) is a strong monobasic acid. Solid nitric acid forms two crystalline modifications: monoclinic and orthorhombic lattices.
Nitric acid mixes with water in any ratio. In aqueous solutions, it almost completely dissociates into ions. Forms with water azeotropic mixture with a concentration of 68.4% and boiling point 120 °C at atmospheric pressure. Two solid hydrates are known: monohydrate (HNO 3 ·H 2 O) and trihydrate (HNO 3 ·3H 2 O).
Chemical properties
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 is 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.
HNO 3 as a strong monobasic acid interacts:
a) with basic and amphoteric oxides:
b) with reasons:
c) displaces weak acids from their salts:
When boiling or exposed to light, nitric acid partially decomposes:
Nitric acid at any concentration exhibits the properties of an oxidizing acid, with nitrogen being reduced to an oxidation state from +4 to −3. The depth of reduction depends primarily on the nature of the reducing agent and the concentration of nitric acid. As an oxidizing acid, HNO 3 interacts:
a) with metals standing in the voltage series to the right of hydrogen:
Concentrated HNO3
Dilute HNO 3
b) with metals standing in the voltage series to the left of hydrogen:
All the above equations reflect only the dominant course of the reaction. This means that under given conditions there are more products of this reaction than products of other reactions, for example, when zinc reacts with nitric acid ( mass fraction nitric acid in a solution of 0.3), the products will contain the most NO, but will also contain (only in smaller quantities) NO 2, N 2 O, N 2 and NH 4 NO 3.
The only general pattern in the interaction of nitric acid with metals: the more dilute the acid and the more metal is more active, the deeper the nitrogen is reduced:
Increasing acid concentration increasing metal activity
Products of interaction of iron with HNO 3 of different concentrations
Nitric acid, even concentrated, does not interact with gold and platinum. Iron, aluminum, chromium are passivated with cold concentrated nitric acid. Iron reacts with dilute nitric acid, and depending on the concentration of the acid, not only various products nitrogen reduction, but also various products of iron oxidation:
Nitric acid oxidizes nonmetals, and nitrogen is usually reduced to NO or NO 2:
And complex substances, For example:
Some organic compounds (for example, amines and hydrazine, turpentine) spontaneously ignite when in contact with concentrated nitric acid.
Nitric acid
Some metals (iron, chromium, aluminum, cobalt, nickel, manganese, beryllium), which react with dilute nitric acid, are passivated by concentrated nitric acid and are resistant to its effects.
A mixture of nitric and sulfuric acids is called “melange”. Thanks to the presence of amyl, a concentration of 104% is achieved [ source not specified 150 days] (that is, when adding 4 parts of distillate to 100 parts of melange, the concentration remains at 100% due to the absorption of water by amyl [ source not specified 150 days]).
Nitric acid is widely used to produce nitro compounds.
Mixture of three volumes of hydrochloric acid and one volume of nitrogen is called “royal vodka”. Aqua regia dissolves most metals, including gold and platinum. Its strong oxidizing abilities are due to the atomic chlorine and nitrosyl chloride formed:
Nitrates
HNO 3 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.
Salts of nitric acid - nitrates - decompose irreversibly when heated, the decomposition products are determined by the cation:
a) nitrates of metals located in the voltage series to the left of magnesium:
2NaNO3 = 2NaNO2 + O2
b) nitrates of metals located in the voltage range between magnesium and copper:
4Al(NO 3) 3 = 2Al 2 O 3 + 12NO 2 + 3O 2
c) nitrates of metals located in the voltage series to the right of mercury:
2AgNO3 = 2Ag + 2NO2 + O2
d) ammonium nitrate:
NH 4 NO 3 = N 2 O + 2H 2 O
Nitrates in aqueous solutions practically do not show oxidative properties, but at high temperatures in the solid state, nitrates are strong oxidizing agents, for example:
Fe + 3KNO 3 + 2KOH = K 2 FeO 4 + 3KNO 2 + H 2 O - during fusion solids.
Zinc and aluminum in alkaline solution reduce nitrates to NH 3:
Salts of nitric acid - 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.
Glass and fluoroplastic-4 do not react with nitric acid.
Historical information
Method for obtaining dilute nitric acid by dry distillation of saltpeter with alum and copper sulfate was apparently first described in the treatises of Jabir (Geber in Latinized translations) in the 8th century. This method, with various modifications, the most significant of which was the replacement of copper sulfate with iron sulfate, was used in European and Arab alchemy until the 17th century.
In the 17th century, Glauber proposed a method for producing volatile acids by reacting their salts with concentrated sulfuric acid, including nitric acid from potassium nitrate, which made it possible to introduce concentrated nitric acid into chemical practice and study its properties. The Glauber method was used until the beginning of the 20th century, and its only significant modification was the replacement of potassium nitrate with cheaper sodium (Chilean) nitrate.
In the time of M.V. Lomonosov, nitric acid was called strong vodka.
Industrial production, use and effect on the body
Nitric acid is one of the largest volume products of the chemical industry.
Nitric acid production
The modern method of its production is based on the catalytic oxidation of synthetic ammonia on platinum-rhodium catalysts (Ostwald process) to a mixture of nitrogen oxides (nitrous gases), with their further absorption by water
4NH 3 + 5O 2 (Pt) → 4NO + 6H 2 O 2NO + O 2 → 2NO 2 4NO 2 + O 2 + 2H 2 O → 4HNO 3.
The concentration of nitric acid obtained by this method varies depending on the technological design of the process from 45 to 58%. Alchemists were the first to obtain nitric acid by heating a mixture of saltpeter and iron sulfate:
4KNO 3 + 2(FeSO 4 7H 2 O) (t°) → Fe 2 O 3 + 2K 2 SO 4 + 2HNO 3 + NO 2 + 13H 2 O
Pure nitric acid was first obtained by Johann Rudolf Glauber by treating nitrate with concentrated sulfuric acid:
KNO 3 + H 2 SO 4 (conc.) (t°) → KHSO 4 + HNO 3
By further distillation the so-called “fuming nitric acid”, containing virtually no water.
· Industrial production, application and effect on the body · Related articles · Notes · Literature · Official website ·
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 is 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 is also determined by its concentration.
HNO 3 as a strong monobasic acid interacts:
a) with basic and amphoteric oxides:
c) displaces weak acids from their salts:
When boiling or exposed to light, nitric acid partially decomposes:
Nitric acid at any concentration exhibits the properties of an oxidizing acid; in addition, nitrogen is reduced to an oxidation state from +4 to 3. The depth of reduction depends primarily on the nature of the reducing agent and the concentration of nitric acid. As an oxidizing acid, HNO 3 interacts:
a) with metals standing in the voltage series to the right of hydrogen:
Concentrated HNO3
Dilute HNO 3
b) with metals standing in the voltage series to the left of hydrogen:
All the above equations reflect only the dominant course of the reaction. This means that under given conditions there are more products of this reaction than products of other reactions, for example, when zinc reacts with nitric acid (mass fraction of nitric acid in solution 0.3), the products will contain the most NO, but will also contain ( only in smaller quantities) and NO 2, N 2 O, N 2 and NH 4 NO 3.
The only general pattern in the interaction of nitric acid with metals is: the more dilute the acid and the more active the metal, the deeper the nitrogen is reduced:
Increasing acid concentration increasing metal activity
Nitric acid, even concentrated, does not interact with gold and platinum. Iron, aluminum, chromium are passivated with cold concentrated nitric acid. Iron reacts with dilute nitric acid, and based on the concentration of the acid, not only various nitrogen reduction products are formed, but also various iron oxidation products:
Nitric acid oxidizes nonmetals, and nitrogen is usually reduced to NO or NO 2:
and complex substances, for example:
Some organic compounds (for example, amines, turpentine) spontaneously ignite when in contact with concentrated nitric acid.
Some metals (iron, chromium, aluminum, cobalt, nickel, manganese, beryllium), which react with dilute nitric acid, are passivated by concentrated nitric acid and are resistant to its effects.
A mixture of nitric and sulfuric acids is called “melange”.
Nitric acid is widely used to produce 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:
Nitrates
Nitric acid is 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.
Salts of nitric acid decompose irreversibly when heated, and the composition of the decomposition products is determined by the cation:
a) nitrates of metals located in the voltage series to the left of magnesium:
b) nitrates of metals located in the voltage range between magnesium and copper:
c) nitrates of metals located in the voltage series to the right of mercury:
d) ammonium nitrate:
Nitrates in aqueous solutions exhibit practically no oxidizing properties, but at high temperatures in the solid state they are strong oxidizing agents, for example, when fusing solids:
Zinc and aluminum in an alkaline solution reduce nitrates to NH 3:
Salts of nitric acid - nitrates - are widely used as fertilizers. In addition, 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 of nitrates are obtained artificially.
Glass and fluoroplastic-4 do not react with nitric acid.
Nitric acid- colorless liquid with a pungent odor, density 1.52 g/cm3, boiling point 84°C, at a temperature of -41°C hardens into colorless crystalline substance. Typically used in practice, concentrated nitric acid contains 65 - 70% HNO3 (maximum density 1.4 g/cm3); Acid mixes with water in any ratio. There is also fuming nitric acid with a concentration of 97 - 99%.
Nitric acid high concentrations release gases in air, which in a closed bottle are detected in the form of brown vapors (nitrogen oxides). These gases are very poisonous, so you need to be careful not to inhale them. Nitric acid oxidizes many organic matter. Paper and fabrics are destroyed due to the oxidation of the substances that form these materials. Concentrated nitric acid causes severe burns with prolonged contact and yellowing of the skin for several days with short contact. Yellowing of the skin indicates the destruction of protein and the release of sulfur (a qualitative reaction to concentrated nitric acid - yellow coloring due to the release of elemental sulfur when the acid acts on protein - xanthoprotein reaction). That is, it is a skin burn.
To prevent burns, you should work with concentrated nitric acid while wearing rubber gloves. At the same time, handling nitric acid is less dangerous than, for example, sulfuric acid; it evaporates quickly and does not remain in unexpected places. Splashes of nitric acid should be washed off with plenty of water, or even better, moistened with a soda solution.
Fuming nitric acid, when stored under the influence of heat and light, partially decomposes:
4HNO3 = 2H2O + 4NO2 + O2.
The higher the temperature and the more concentrated the acid, the faster the decomposition occurs. Therefore, store it in a cool and dark place. The released nitrogen dioxide dissolves in the acid and gives it a brown color.
Dilute acid can be easily prepared by pouring concentrated acid in water.
Dilute nitric acid is stored and transported in chromium steel containers, concentrated - in aluminum containers, because concentrated acid passivates aluminum, iron and chromium due to the formation of insoluble oxide films:
2Al + 6HNO3 = Al2O3 + 6NO2 + 3H2O.
Small quantities are stored in glass bottles. Nitric acid strongly corrodes rubber. Therefore, bottles must have ground or polyethylene stoppers.
Nitric acid is used mainly in the form of aqueous solutions and is one of the components aqua regia, found in assay acids. In industry, it is used to produce combined nitrogen fertilizers, for dissolving ores and concentrates, in the production of sulfuric acid, various organic nitroproducts, in rocket technology as a fuel oxidizer, etc.
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 NO easily transforms into NO2, which, with water in the presence of atmospheric oxygen, produces 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%,
Application of nitric acid
Nitric acid is used in the production of nitrogen and combined fertilizers (sodium, ammonium, calcium and potassium nitrate, nitrophos, nitrophoska), various sulfuric acid salts, explosives (trinitrotoluene, etc.), organic dyes.
IN organic synthesis A mixture of concentrated nitric acid and sulfuric acid - a “nitrating mixture” - is widely used.
In metallurgy, nitric acid is used to dissolve and pickle metals, as well as to separate gold and silver. Nitric acid is also used in the chemical industry, in the production of explosives, and in the production of intermediates for the production of synthetic dyes and other chemicals.
Technical nitric acid is used for nickel plating, galvanizing and chrome plating of parts, as well as in the printing industry. Nitric acid is widely used in the dairy and electrical industries.
Density of solutions of different concentrations of nitric acid
|
Density, g/cm 3 |
Concentration |
Density, |
Concentration |
||
|
g/l. |
g/l. |
||||
|
1, 000 |
0, 3296 |
3, 295 |
1, 285 |
46, 06 |
591, 9 |
|
1, 005 |
1, 255 |
12, 61 |
1, 290 |
46, 85 |
604, 3 |
|
1, 010 |
2, 164 |
21, 85 |
1, 295 |
47, 63 |
616, 8 |
|
1, 015 |
3, 073 |
31, 19 |
1, 300 |
48, 42 |
629, 5 |
|
1, 020 |
3, 982 |
40, 61 |
1, 305 |
49, 21 |
642, 1 |
|
1, 025 |
4, 883 |
50, 05 |
1, 310 |
50, 00 |
644, 7 |
|
1, 030 |
5, 784 |
59, 57 |
1, 315 |
50, 85 |
668, 5 |
|
1, 035 |
6, 661 |
68, 93 |
1, 320 |
51, 71 |
682, 4 |
|
1, 040 |
7, 530 |
78, 32 |
1, 325 |
52, 56 |
696, 3 |
|
1, 045 |
8, 398 |
87, 77 |
1, 330 |
53, 41 |
710, 1 |
|
1, 050 |
9, 259 |
97, 22 |
1, 335 |
54, 27 |
724, 0 |
|
1, 055 |
10, 12 |
106, 7 |
1, 340 |
55, 13 |
738, 5 |
|
1, 060 |
10, 97 |
116, 3 |
1, 345 |
56, 04 |
753, 6 |
|
1, 065 |
11, 81 |
125, 8 |
1, 350 |
56, 95 |
768, 7 |
|
1, 070 |
12, 65 |
135, 3 |
1, 355 |
57, 87 |
783, 8 |
|
1, 075 |
13, 48 |
145, 0 |
1, 360 |
58, 78 |
799, 0 |
|
1, 080 |
14, 31 |
154, 6 |
1, 365 |
59, 69 |
814, 7 |
|
1, 085 |
15, 13 |
164, 1 |
1, 370 |
60, 67 |
831, 1 |
|
1, 090 |
15, 95 |
173, 8 |
1, 375 |
61, 69 |
848, 1 |
|
1, 095 |
16, 76 |
183, 5 |
1, 380 |
62, 70 |
865, 1 |
|
1, 100 |
17, 58 |
193, 3 |
1, 385 |
63, 72 |
882, 8 |
|
1, 105 |
18, 39 |
203, 1 |
1, 390 |
64, 74 |
900, 4 |
|
1, 110 |
19, 19 |
213, 0 |
1, 395 |
65, 84 |
918, 1 |
|
1, 115 |
20, 00 |
223, 0 |
1, 400 |
66, 97 |
937, 6 |
|
1, 120 |
20, 79 |
232, 9 |
1, 405 |
68, 10 |
956, 6 |
|
1, 125 |
21, 59 |
242, 8 |
1, 410 |
69, 23 |
976, 0 |
|
1, 130 |
22, 38 |
252, 8 |
1, 415 |
70, 34 |
996, 2 |
|
1, 135 |
23, 16 |
262, 8 |
1, 420 |
71, 63 |
1017 |
|
1, 140 |
23, 94 |
272, 8 |
1, 425 |
72, 86 |
1038 |
|
1, 145 |
24, 71 |
282, 9 |
1, 430 |
74, 09 |
1059 |
|
1, 150 |
25, 48 |
292, 9 |
1, 435 |
74, 35 |
1081 |
|
1, 155 |
26, 24 |
303, 1 |
1, 440 |
76, 71 |
1105 |
|
1, 160 |
27, 00 |
313, 2 |
1, 445 |
78, 07 |
1128 |
|
1, 165 |
27, 26 |
323, 4 |
1, 450 |
79, 43 |
1152 |
|
1, 170 |
28, 51 |
333, 5 |
1, 455 |
80, 88 |
1177 |
|
1, 175 |
29, 25 |
343, 7 |
1, 460 |
82, 39 |
1203 |
|
1, 180 |
30, 00 |
354, 0 |
1, 465 |
83, 91 |
1229 |
|
1, 185 |
30, 74 |
364, 2 |
1, 470 |
8550 |
1257 |
|
1, 190 |
31, 47 |
374, 5 |
1, 475 |
87, 29 |
1287 |
|
1, 195 |
32, 21 |
385, 0 |
1, 480 |
89, 07 |
1318 |
|
1, 200 |
32, 94 |
395, 3 |
1, 485 |
91, 13 |
1353 |
|
1, 205 |
33, 68 |
405, 8 |
1, 490 |
93, 19 |
1393 |
|
1, 210 |
34, 41 |
416, 3 |
1, 495 |
95, 46 |
1427 |
|
1, 215 |
35, 16 |
427, 1 |
1, 500 |
96, 73 |
1450 |
|
1, 220 |
35, 93 |
438, 3 |
1, 501 |
96, 98 |
1456 |
|
1, 225 |
36, 70 |
449, 6 |
1, 502 |
97, 23 |
1461 |
|
1, 230 |
37, 48 |
460, 9 |
1, 503 |
97, 49 |
1465 |
|
1, 235 |
38, 25 |
472, 4 |
1, 504 |
97, 74 |
1470 |
|
1, 240 |
39, 02 |
483, 8 |
1, 505 |
97, 99 |
1474 |
|
1, 245 |
39, 80 |
495, 5 |
1, 506 |
98, 25 |
1479 |
|
1, 250 |
40, 58 |
505, 2 |
1, 507 |
98, 50 |
1485 |
|
1, 255 |
41, 36 |
519, 0 |
1, 508 |
98, 76 |
1490 |
|
1, 260 |
42, 14 |
530, 9 |
1, 509 |
99, 01 |
1494 |
|
1, 265 |
42, 92 |
542, 9 |
1, 510 |
99, 26 |
1499 |
|
1, 270 |
43, 70 |
555, 0 |
1, 511 |
99, 52 |
1503 |
|
1, 275 |
44, 48 |
567, 2 |
1, 512 |
99, 74 |
1508 |
|
1, 280 |
45, 27 |
579, 4 |
1, 513 |
100, 00 |
1513 |
And water.
In twilight, acid and water are easily mixed in any proportions. The substance also has a crystalline state.
It can be monoclinic or rhombic. This indicates the shape of the crystal lattice cells.
The monoclinic one is made up of inclined parallelepipeds, and the rhombic one, respectively, from rhombuses.
Do the properties of solutions differ from theirs, how is the substance obtained and where is it used? The questions have been asked, all that remains is to give answers below.
Properties of nitric acid
IN normal conditions crystalline acid can only be seen in hot countries.
It becomes a colorless liquid only at 42 degrees Celsius. Up to this point, the substance remains liquid and floats.
At the same time, the reagent emits a sharp, suffocating odor. Actually connected with him history of the discovery of nitric acid. It was discovered by Daniel Rutherford.
The Scotsman studied combustion products. During the work, gas was released, which the chemist called suffocating air.
The scientist noted that the substance does not support combustion and is unsuitable for breathing.
Later, it turned out nitric acid formula: - HNO 3 . It turns out that the substance is monobasic.
So called, which contain only one hydrogen atom. The substance is mixed with water in any proportions.
Therefore, there is concentrated nitric acid and unconcentrated.
The first one actively smokes, that is, it is volatile. The chemical properties of the concentrate differ from the diluted version.
If the acid in the solution is about 60%, it will react with all metals except , , , , , and .
Hence the conclusion in what container the substance should be stored. and flasks, of course, are not profitable.
But containers made of iron and aluminum are both inexpensive and reliable, since they block the acid from light. The main thing is not to choose a container from copper Nitric acid will dissolve it.
Reacting with metals, concentrated nitric acid solution releases brown gas. Its formula: - NO 2.
At the same time, acids are formed. Depending on the dissolved metal, the reactions vary.
When interacting with a number of up, dioxides are formed and oxygen is released.
Reaction with metal salts located after magnesium before produces brown gas, nitric oxide and oxygen.
If a salt of any metal after the copper is added to the acid, the metal will separate. Along with it, brown gas and oxygen are released.
Dilute nitric acid reacts with most of the same metals, but is oxidized to ammonia.
This outcome is caused by interaction, for example, with elements of the alkaline earth group. Iron also reacts.
So, it is better not to store diluted acid in containers containing ferrum.
The result interaction with nitric acid Not only ammonia, but also ammonium nitrate can become a diluted type.
The rarest option is nitrous oxide. It will be given, for example, by a reaction with magnesium. With other metals, nitric acid forms nitric oxide.
It can be obtained, in particular, by interacting with . Argentum oxide will precipitate, forming water and nitric oxide.
Reactions of acids with non-metals follow the same scheme, only sulfuric acid is formed instead.
Of the reactions with other acids, mixing with hydrochloric acid is noteworthy. The last one takes 3 parts, and the first one - one. It turns out .
It was named so because the substance dissolves even the metal of rulers, powerful of the world this.
None of the pure acids are capable of this. Noble metals They are rarely succumbed to, and never at all.
Nitric acid production
In small quantities, the substance can be extracted even from the air, and in the literal sense. It is no secret that nitrogen is one of the components of the atmosphere.
The 15th gas in it accounts for 78%. Nitrogen reacts with oxygen to form an oxide. Further oxidation produces nitrogen dioxide. This is the same brown gas.
It is this that reacts with water, a suspension of which, as is known, is present in the air. Coming into contact with clouds and fog, the brown gas turns into nitric acid.
Mass fraction of nitric acid in the atmosphere is so small that the substance does not harm humans or other living organisms.
Acid from the air is also not suitable for industrial production. Factories use different schemes.
First: - nitric acid production from ammonia. First, its conversion is carried out, that is, the composition of the initial gas mixture is crushed.
The reaction takes place on platinum-rhodium grids at a temperature of about 1000 degrees Celsius. This is how nitric oxide is obtained. It is oxidized to dioxide.
This is the second stage of the process. Afterwards, nitrogen oxides are absorbed by water. The result is nitric acid and pure water.
The described method results in the formation of dilute acid. Subsequent concentration is possible.
Therefore, the method is the most popular, because consumers need both saturated and unsaturated acids.
When working with ammonia, industrialists “kill two birds with one stone.”
The second method of producing the reagent leads directly to the production of a concentrate. We are talking about direct synthesis from nitrogen oxides. Take liquid ones.
They interact with water and oxygen. Such reactions with nitric acid pass under a pressure of 5 megapascals.
This produces nitrogen dioxide. Under normal conditions it goes into liquid state. Oxidation of ammonia produces double nitric oxide.
It is about 11% in the gas mixture. The dioxide is liquefied under pressure. Under standard conditions the transition is not possible.
Application of nitric acid
As a constituent of aqua regia, nitric acid is part of the acids. With their help, quality is studied.
Without appropriate research, they will not go to the market, but to the shelves.
Before a precious metal can be tested and sold, it must be mined. Nitric acid and aqua regia also help with this.
They process ores, bringing the necessary elements into solution. All that remains is to precipitate the metals and dry them, clean them of impurities. This is how not only noble but also ignoble elements are mined.
As you know, they make metals, and from them, for example, equipment. If we consider air and space, they contain pure acid.
It is mixed with fuel, obtaining oxide. Nitric acid acts as an oxidizing agent. .
All these are salts, united by the name “saltpeter”. Nitrogen allows plants to develop quickly and increases productivity.
The fact is that the 15th element is part of chlorophyll. This is a green plant pigment responsible for energy absorption.
The more energy harnessed, the better development herbs, shrubs, trees.
The word “saltpeter” is also well known among pyrotechnicians. Nitric acid is the basis of explosives.
Ammonium nitrate in most of them is about 60%. The remainder is diesel fuel or other fuel. You can get both harmless fireworks and a military bomb.
Nitric acid price
Nitric acid, like most popular acids, can be pure or technical, laden with impurities. The latter is cheaper.
Pure reagent is more expensive. For reference, GOST 4461-77 is the standard for purified acid.
The Russian-made reagent costs around 30-55 rubles per kilogram. The price depends on the concentration of the solution.
For technical acid, the upper price limit is usually 40 per kilo. Large packaging is also available.
There are, for example, 25-liter canisters into which Nitric acid.
Buy reagent with maximum benefit allows bulk orders. These go to enterprises where they know the rules for handling the reagent.
It corrodes not only metals, but also mucous membranes. Vapors of the substance can make breathing difficult and damage the trachea lining the tissues of the nose.
Therefore, people only work with acid wearing masks. If the rules are violated, in addition to difficulty breathing, poisoning occurs.
Intoxication is expressed in vomiting, scabies, visual impairment, and smell. Only weak solutions of the reagent are more or less harmless.
These are the ones that, for example, are used in school laboratories. It is worth learning how to handle chemicals from an early age.