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Making nitrous oxide at home
There are several ways to obtain laughing gas. The most accessible at home is G. Davy’s method - the thermal decomposition of ammonium nitrate (ammonium nitrate) according to the reaction:
NH 4 NO 3 = N 2 O + 2H 2 O.
In laboratory conditions, it is more convenient to heat sulfamic acid with nitric acid:
NH 2 SO 2 OH + HNO 3 (73%) = N 2 O + H 2 SO 4 + H 2 O.
However, sulfamic and nitric acids are more difficult to obtain, so we will focus on the decomposition of ammonium nitrate. By the way, the decomposition of ammonium nitrate is used to synthesize nitrous oxide on an industrial scale.
When ammonium nitrate is heated, several reactions take place. Here are excerpts from the book L.I. Bagal Chemistry and Technology of Initiating Explosives (1975)
Ammonium nitrate, when heated slightly above its melting point (dry ammonium nitrate melts at 169.6°C), decomposes according to the reaction
NH 4 NO 3 NH 3 + HNO 3 (1)
<...>
The reaction of decomposition to nitrous oxide and water was studied by Berthelot, Thomsen and Velay. The first two researchers found that the reaction was exothermic
NH 4 NO 3 => N 2 O + 2H 2 O + 8.8 kcal (2)
The main reactions of ammonium nitrate decomposition at temperatures up to 270°C are (1) and (2). Molten ammonium nitrate, when heated above 250-260°C, can release nitrogen oxides, nitrogen and water:
NH 4 NO 3 => 0.5N 2 + NO + 2H 2 O
4NH 4 NO 3 => 3N 2 + N 2 O 4 + 8H 2 O
3NH 4 NO 3 => 2N 2 + N 2 O 3 + 6H 2 O
Saunders (1922), based on the results of gas analysis, came to the conclusion that the main decomposition reactions at temperatures up to 260°C are (1) and (2), as well as the reaction
5NH 3 + 3HNO 3 => 4N 2 + 9H 2 O
In his opinion, decomposition during an explosion proceeds according to the reaction
8NH 4 NO 3 => 16H 2 O + 2NO 2 + 4NO + 5N 2
<...>
For the normal process of formation of nitrous oxide by decomposition of ammonium nitrate, its temperature regime and degree of purity are of exceptional importance.
As can be seen from the above data, ammonium nitrate, when heated to 240-250°C, decomposes to form nitrous oxide and water, however, even at this temperature the resulting “raw” gas contains nitric acid vapor, nitrogen oxides NO and NO 2, ammonia, chlorine (due to chloride impurities), nitrogen and “fog” of sublimated ammonium nitrate. It is clear that such a mixture cannot be inhaled (if the idea arises of repeating Davy’s experiments), since it deadly! Moreover, if the flask is closed with a rubber stopper, then even after short-term use it gradually collapses (with the formation of completely harmless products).
Therefore, the method of producing laughing gas by heating ammonium nitrate in a frying pan (which is often recommended by “gurus” to laugh at “laymen”) looks like black humor at best.
Let's move on to the installation. Ammonium nitrate is decomposed in a Wurtz flask under gentle heating. It is better to use a thermometer, but you can do without it if necessary. As experience has shown, it is better to use heating to approximately 220°C, in which case a slight “boiling” of the melt is observed. The resulting “raw gas” for purification is first passed through an ice-cooled trap to collect distilled water mixed with nitric acid. Next, the gas passes through a Drexel flask with a solution of iron sulfate; it also serves as a kind of indicator of the rate of gas release. Then the gas is washed in an improvised washing machine (with a porous spray) with a solution of 5-7% alkali (sodium or potassium hydroxide), where it is cleared of NO 2, nitric acid, and chlorine. And finally, in the third wash with a porous spray, into which a solution of iron (II) sulfate is poured, nitrous oxide is cleared of NO and traces of remaining impurities. After this, the gas contains nitrous oxide with some water and nitrogen, as well as traces of NO 2 and NO.
It should be remembered that the purification of nitrous oxide, if it is used to repeat Davy's experiments, should be given Special attention, otherwise the gas will be toxic.
Ammonium nitrate fertilizer (ammonium nitrate) was used as a reaction load.
Nitrate is the name given to nitrate salts (nitrates) of ammonium, sodium, calcium and potassium. They are mainly used in agriculture, as mineral fertilizers, and in the industrial production of pyrotechnic products and explosives.
Potassium nitrate is considered a very valuable fertilizer, as it simultaneously contains two substances important for plant life - nitrogen and potassium. But, at the same time, potassium nitrate is the basis of black gunpowder and is simply irreplaceable in the manufacture of various pyrotechnics. However, experiments by home-grown craftsmen to create rockets, smoke bombs and other “explosives” often end very disastrously. Therefore, the sale of potassium nitrate has recently been limited, and with the onset of spring, summer residents are increasingly forced to think about how to make nitrate themselves. Our advice is intended for amateur gardeners who use potassium nitrate exclusively for peaceful purposes.
How to make potassium nitrate
- Buy potassium carbonate, also known as potash, and ammonium nitrate at the hardware store.
- Dissolve them separately in warm water, preferably distilled. Use equal parts by weight of reagents. Mix both solutions in an unnecessary container, pouring the potash solution into the ammonium nitrate solution.
- Place the pan over low heat. The pan must be large enough, since during the reaction the mixture foams and increases in volume. Stir the mixture regularly. Soon ammonia gas with a sharp characteristic odor will begin to be released from it - this means that the reaction has begun. Due to the pungent odor of gas, it is better to carry out the process outdoors or indoors with good ventilation.
- After the gas evolution stops, remove the pan from the heat and leave in a cool place for a day. After this, large needle-shaped crystals of potassium nitrate form at the bottom, which can only be removed by draining the liquid and dried.
Ammonium nitrate is one of the most common fertilizers; it is applied when sowing almost all agricultural crops, both grains and vegetables, and is also used as a top dressing for adult plants. In mining, ammonium nitrate is widely used as the main component of high explosives - ammonal, ammonite or ammotol. Ammonium nitrate is sold in all hardware stores in the “Fertilizer” departments, where it can be easily purchased. Making ammonium nitrate in artisanal conditions is extremely dangerous and completely unprofitable! You can try to synthesize it yourself only in small doses, observing all safety rules, for educational purposes.
How to make ammonium nitrate
- At the hardware store you need to buy: ammonia, copper sulfate, calcium nitrate.
- Mix ammonia with copper sulfate until a blue solution is obtained. As a result of the substitution reaction, we get copper hydroxide precipitated and ammonium sulfate remaining in solution.
- Drain the ammonium sulfate solution from the sediment and mix it with calcium nitrate. As a result, we obtain calcium sulfate in the form of a precipitate and a solution containing our ammonium nitrate.
We have described the main methods for obtaining saltpeter, and it’s up to you to decide what can be made from saltpeter produced at home.
Nitric acid is a strong acid. Its salts - nitrates- 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:
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.
Liquid nitrogen is used as a refrigerant and for cryotherapy. In petrochemistry, nitrogen is used to purge tanks and pipelines, check the operation of pipelines under pressure, and increase the production of fields. In mining, nitrogen can be used to create an explosion-proof environment in mines and to expand rock layers.
An important area of application of nitrogen is its use for the further synthesis of a wide variety of compounds containing nitrogen, such as ammonia, nitrogen fertilizers, explosives, dyes, etc. Large quantities of nitrogen are used in coke production (“dry quenching of coke”) during unloading coke from coke oven batteries, as well as for “pressing” fuel in rockets from tanks to pumps or engines.
In the food industry, nitrogen is registered as a food additive E941, as a gaseous medium for packaging and storage, a refrigerant, and liquid nitrogen is used when bottling oils and non-carbonated drinks to create excess pressure and an inert environment in soft containers.
The tire chambers of aircraft landing gear are filled with nitrogen gas.
31. Phosphorus – production, properties, application. Allotropy. Phosphine, phosphonium salts – preparation and properties. Metal phosphides, preparation and properties.
Phosphorus- chemical element of the 15th group of the third period of the periodic system of D. I. Mendeleev; has atomic number 15. The element is part of the pnictogen group.
Phosphorus is obtained from apatites or phosphorites as a result of interaction with coke and silica at a temperature of about 1600 ° C:
The resulting phosphorus vapors condense in the receiver under a layer of water into an allotropic modification in the form of white phosphorus. Instead of phosphorites, to obtain elemental phosphorus, other inorganic phosphorus compounds can be reduced with coal, for example, metaphosphoric acid:
The chemical properties of phosphorus are largely determined by its allotropic modification. White phosphorus is very active; in the process of transition to red and black phosphorus, the chemical activity decreases. White phosphorus in the air, when oxidized by air oxygen at room temperature, emits visible light; the glow is due to the photoemission reaction of phosphorus oxidation.
Phosphorus is easily oxidized by oxygen:
(with excess oxygen)
(with slow oxidation or lack of oxygen)
It interacts with many simple substances - halogens, sulfur, some metals, exhibiting oxidizing and reducing properties: with metals - an oxidizing agent, forms phosphides; with non-metals - a reducing agent.
Phosphorus practically does not combine with hydrogen.
In cold concentrated solutions of alkalis, the disproportionation reaction also occurs slowly:
Strong oxidizing agents convert phosphorus into phosphoric acid:
The oxidation reaction of phosphorus occurs when matches are lit; Berthollet salt acts as an oxidizing agent:
The most chemically active, toxic and flammable is white (“yellow”) phosphorus, which is why it is very often used (in incendiary bombs, etc.).
Red phosphorus is the main modification produced and consumed by industry. It is used in the production of matches, explosives, incendiary compositions, various types of fuels, as well as extreme pressure lubricants, as a getter in the production of incandescent lamps.
Under normal conditions, elemental phosphorus exists in the form of several stable allotropic modifications. All possible allotropic modifications of phosphorus have not yet been fully studied (2016). Traditionally, four modifications are distinguished: white, red, black and metallic phosphorus. Sometimes they are also called main allotropic modifications, implying that all other described modifications are a mixture of these four. Under standard conditions, only three allotropic modifications of phosphorus are stable (for example, white phosphorus is thermodynamically unstable (quasi-stationary state) and transforms over time under normal conditions into red phosphorus). Under conditions of ultra-high pressures, the metallic form of the element is thermodynamically stable. All modifications differ in color, density and other physical and chemical characteristics, especially chemical activity. When the state of a substance transitions to a more thermodynamically stable modification, chemical activity decreases, for example, during the sequential transformation of white phosphorus into red, then red into black (metallic).
Phosphine (hydrogen phosphide, hydrogen phosphide, phosphorus hydride, phosphane PH 3) is a colorless, poisonous gas (under normal conditions) with a specific smell of rotten fish.
Phosphine is obtained by reacting white phosphorus with hot alkali, for example:
It can also be obtained by treating phosphides with water or acids:
When heated, hydrogen chloride reacts with white phosphorus:
Decomposition of phosphonium iodide:
Decomposition of phosphonic acid:
or restoring it:
Chemical properties.
Phosphine is very different from its counterpart, ammonia. Its chemical activity is higher than that of ammonia; it is poorly soluble in water, as a base is much weaker than ammonia. The latter is explained by the fact that the H–P bonds are weakly polarized and the activity of the lone pair of electrons in phosphorus (3s 2) is lower than that of nitrogen (2s 2) in ammonia.
In the absence of oxygen, when heated, it decomposes into elements:
spontaneously ignites in air (in the presence of diphosphine vapor or at temperatures above 100 °C):
Shows strong restorative properties:
When interacting with strong proton donors, phosphine can produce phosphonium salts containing the PH 4 + ion (similar to ammonium). Phosphonium salts, colorless crystalline substances, are extremely unstable and easily hydrolyze.
Phosphonium salts, like phosphine itself, are strong reducing agents.
Phosphides- binary compounds of phosphorus with other less electronegative chemical elements in which phosphorus exhibits a negative oxidation state.
Most phosphides are compounds of phosphorus with typical metals, which are obtained by direct interaction of simple substances:
Na + P (red) → Na 3 P + Na 2 P 5 (200 °C)
Boron phosphide can be obtained either by direct interaction of substances at a temperature of about 1000 °C, or by the reaction of boron trichloride with aluminum phosphide:
BCl 3 + AlP → BP + AlCl 3 (950 °C)
Metal phosphides are unstable compounds that decompose with water and dilute acids. This produces phosphine and, in the case of hydrolysis, metal hydroxide; in the case of interaction with acids, salts.
Ca 3 P 2 + 6H 2 O → 3Ca(OH) 2 + 2PH 3
Ca 3 P 2 + 6HCl → 3CaCl 2 + 2PH 3
When heated moderately, most phosphides decompose. Melts under excess pressure of phosphorus vapor.
Boron phosphide BP, on the contrary, is refractory (melting point 2000 °C, with decomposition), a very inert substance. It decomposes only with concentrated oxidizing acids, reacts when heated with oxygen, sulfur, and alkalis during sintering.
32. Phosphorus oxides - structure of molecules, preparation, properties, application.
Phosphorus forms several oxides. The most important of them are phosphorus oxide (V) P 4 O 10 and phosphorus oxide (III) P 4 O 6. Often their formulas are written in a simplified form - P 2 O 5 and P 2 O 3. The structure of these oxides retains the tetrahedral arrangement of phosphorus atoms.
Phosphorus (III) oxide P 4 O 6- a waxy crystalline mass that melts at 22.5°C and turns into a colorless liquid. Poisonous.
When dissolved in cold water it forms phosphorous acid:
P 4 O 6 + 6H 2 O = 4H 3 PO 3,
and when reacting with alkalis - the corresponding salts (phosphites).
Strong reducing agent. When interacting with oxygen, it is oxidized to P 4 O 10.
Phosphorus (III) oxide is obtained by the oxidation of white phosphorus in the absence of oxygen.
Phosphorus (V) oxide P 4 O 10- white crystalline powder. Sublimation temperature 36°C. It has several modifications, one of which (the so-called volatile) has the composition P 4 O 10. The crystal lattice of this modification is composed of P 4 O 10 molecules connected to each other by weak intermolecular forces, which are easily broken when heated. Hence the volatility of this variety. Other modifications are polymeric. They are formed by endless layers of PO 4 tetrahedra.
When P 4 O 10 interacts with water, phosphoric acid is formed:
P 4 O 10 + 6H 2 O = 4H 3 PO 4.
Being an acidic oxide, P 4 O 10 reacts with basic oxides and hydroxides.
It is formed during high-temperature oxidation of phosphorus in excess oxygen (dry air).
Due to its exceptional hygroscopicity, phosphorus (V) oxide is used in laboratory and industrial technology as a drying and dehydrating agent. In its drying effect it surpasses all other substances. Chemically bound water is removed from anhydrous perchloric acid to form its anhydride:
4HClO4 + P4O10 = (HPO3)4 + 2Cl2O7.
P 4 O 10 is used as a desiccant for gases and liquids.
Widely used in organic synthesis in dehydration and condensation reactions.
The invention relates to the production of nitric acid salts. The essence of the method is that nitrite-nitrate solutions obtained by absorption of nitrogen oxides with soda or caustic soda are evaporated to a total salt concentration of 750-900 g/l without isolating the solid phase and at a temperature of 70-90 o C are sent for inversion, inversion the gases are diluted with air and returned to the absorption stage, and the contact separation of ammonia to produce nitrous gases is turned on periodically as sodium nitrite is processed, and the product solutions of sodium nitrate are processed into a salt product in a known manner, including crystallization and drying of the product. The technical result is that the method makes it possible to obtain sodium nitrate without producing sodium nitrite, and also to use nitric acid as a donor of nitrogen oxides at the inversion stage instead of the catalytic oxidation of ammonia. 1 salary f-ly, 1 ill.
The invention relates to the chemical industry and can be used in enterprises producing nitric acid salts. There is a known method for producing sodium nitrate by neutralizing a solution of soda and (or) sodium hydroxide with nitric acid (V.A. Klevke, N.N. Polyakov, L.Z. Arsenyeva. Technology of nitrogen fertilizers. - M.: Goskhimizdat, 1956, p. 94; RF patent 2159738 dated December 3, 1999. Method for producing sodium nitrate). The disadvantage of the known methods is the low concentration of sodium nitrate in the product solution (320-360 g/l) and the associated high consumption of steam for its concentration before crystallization of the finished product. The closest in technical essence is the method of obtaining sodium nitrate from nitrite-nitrate solutions by inverting the latter with nitric acid (M.A. Miniovich, V.M. Miniovich. Salts of nitrous acid. - M.: Chemistry, 1997, pp. 100-101 ). The disadvantage of this method is the need for simultaneous production of sodium nitrite and the use of expensive platinum catalysts for the conversion of ammonia to nitrogen oxides. The objective of this proposed invention is to develop a method for producing sodium nitrate from nitrite-nitrate solutions without the production of sodium nitrite, the demand for which is highly seasonal, as well as the use of nitric acid to produce nitrogen oxides at the inversion stage instead of the catalytic oxidation of ammonia with atmospheric oxygen. This goal is achieved by the fact that after the stage of ammonia oxidation with atmospheric oxygen, cooling of nitrous gases, their absorption with a solution of soda or caustic soda, the nitrite-nitrate solution is evaporated without isolating the solid phase to a total salt content of 750-900 g/l and sent for inversion. When a solution is mixed with non-concentrated nitric acid, a well-known reaction occurs intensively: 3NaNO 2 + 2HNO 3 = 3NaNO 3 + 2NO + H 2 O. The ratio of the flows of nitric acid and a solution of nitrite-nitrate salts, in which the sum of salts is 750-900 g/l, is maintained in such a way that the acidity of the intermediate production solution is in the range of 30-80 g/l HNO 3 . The resulting nitrogen oxides are blown off with air in an inversion column. Since the inversion is carried out in the presence of an increased concentration of sodium nitrite, the inversion gases are diluted with additional air before entering the absorption stage. During absorption, they are absorbed by a circulating solution containing excess alkalinity in the form of soda or sodium hydroxide. In this case, sodium nitrite and nitrate are formed again. The drawing shows a diagram of the implementation of the method for producing sodium nitrate. The start-up of the technological scheme is carried out in the traditional way: the contact apparatus is ignited at the ammonia conversion stage, for which ammonia gas and air are used. Nitrous gases pass through a waste heat boiler, where they are cooled to 200-220 o C and enter an absorber irrigated with a circulating nitrite-nitrate solution containing an excess amount of soda or caustic soda. This solution (the sum of salts is 320-400 g/l) is periodically taken for evaporation, where by evaporation the sum of salts (NaNO 2 + NaNO 3 + Na 2 CO 3) increases to 750-900 g/l. The temperature of the evaporated solution is maintained within 70-90 o C to prevent precipitation of the solid phase. This solution is sent to a continuous inversion reactor column, into which non-concentrated nitric acid containing 56-58 wt.% HNO 3 is simultaneously dosed. The ratio of the flows of nitric acid and a solution of nitrite-nitrate salts is selected in such a way that an acidic environment is maintained in the column and the intermediate product solution has an acidity of 30-80 g/l HNO 3 . Air is continuously supplied to the reactor, which, in addition to improving the mixing of the reagents, ensures the removal of nitrous gases from the reaction zone. The intermediate acidic solution of sodium nitrate is sent to a neutralizer, where it is neutralized to a pH of 8-10 by mixing with a solution of soda or caustic soda. The inverted nitrous gases are diluted with additional air and sent to an absorption column. After the accumulation of a certain amount of nitrite-nitrate solutions, the contact department is stopped, and the production of sodium nitrate continues due to inversion nitrous gases. In this case, as can be seen from the above reaction, the donor of nitric oxide is nitric acid, and the process continues as long as the reducing agent sodium nitrite is present in the system. As sodium nitrite is used up and the concentration of nitrous gases decreases, it becomes necessary to connect contact oxidation of ammonia. The resulting product solution of sodium nitrate is processed into a salt product according to a known method by evaporation, crystallization and separation of sodium nitrate, followed by drying. Recycled mother liquors, after the accumulation of impurities (Cl-ions) in them, are used in the process of conversion of potassium nitrate from potassium chloride.
Claim
1. A method for producing sodium nitrate, including the oxidation of ammonia with atmospheric oxygen, cooling of nitrous gases, their absorption with a solution of soda or caustic soda, evaporation of the resulting nitrite-nitrate solutions, inversion of sodium nitrite with nitric acid at a temperature of 70-90 o C with the return of inversion gases to the stage absorption, neutralization of the intermediate solution of sodium nitrate, evaporation, crystallization and drying of the finished product, characterized in that nitrite-nitrate solutions evaporated to a total salt concentration of 750-900 g/l without isolating the solid phase are sent for inversion and the acidity of the intermediate solution is maintained at 30-80 g/l HNO 3, and the inversion gases are diluted with air before returning to the absorption stage. 2. The method according to claim 1, characterized in that the ammonia oxidation stage is switched on periodically as sodium nitrite is processed in inversion and the concentration of inversion nitrous gases decreases.
Similar patents:
The invention relates to the field of recycling components of liquid rocket fuels, in particular the processing of special nitro mixtures, which are an oxidizer of rocket fuel, into fertilizers and salts
The invention relates to a method for producing alkali metal nitrate and alkali metal phosphate in the same technological process from phosphate raw materials and nitrate raw materials, including the following steps: a) interaction of phosphate raw materials with nitrate raw materials to form an aqueous nitrophosphate reaction mixture, followed by optional separation of solid material, b) introducing an aqueous nitrophosphate reaction mixture into the first ion exchange stage, carried out in the presence of a cation exchange resin saturated with alkali metal ions, to exchange cations present in the reaction mixture for alkali metal ions present in this resin, obtaining a stream enriched with ions alkali metal, c) carrying out the first crystallization of the stream obtained in step (b), under conditions ensuring crystallization of the alkali metal nitrate, and separating the crystallized alkali metal nitrate from the mother liquor, d) introducing the mother liquor formed in step (c) into a second ion exchange step carried out in the presence of an alkali metal ion-rich cation exchange resin to exchange cations present in the mother liquor for alkali metal ions present in the resin to produce an alkali metal ion-rich phosphate containing stream, and e) performing second crystallization of the stream obtained in step (d), under conditions ensuring crystallization of the alkali metal phosphate, and separation of the crystallized alkali metal phosphate from the mother liquor
The invention relates to the production of nitric acid salts
Introduction
It is unlikely that any of you have heard anything about sodium nitrate. Its name is often mentioned in school, let alone in industry. But only the name! What else is known about sodium nitrate? That's what we'll talk about in today's article.
DefinitionSodium nitrate (formula NaNO 3) is the sodium salt of nitric acid. It may be called "sodium nitrate" or "sodium/sodium/Chile nitrate".
Properties
Sodium nitrate is represented by colorless long crystals having a rhombohedral or trigonal crystal lattice. They taste very salty. They dissolve differently in different substances, but sodium nitrate “melts” best in water. At a temperature of 380 o C, this compound decomposes into sodium nitrate and oxygen. The reaction looks like this: 2NaNO 3 => 2NaNO 2 + O 2. Also, sodium nitrate can enter into exchange reactions, the second reagent of which is alkali metal salts. One of the products will always be nitrate with a solubility value that will be much lower than that of the substance now being discussed. For example, when sodium nitrate reacts with potassium chloride, potassium nitrate and table salt (sodium chloride) are formed. In a melt, the nitrate under discussion exhibits strong oxidizing properties, and in solution, reducing properties. When it decomposes, oxygen is released, and due to this, this compound can react with non-metals.
This nitrate can be obtained in several ways:
Reaction of nitric acid with metal or sodium oxide
When pure sodium is added to nitric acid, a neutralization reaction occurs. Its products will be the desired substance, water, as well as gaseous nitrogen and its oxides (I, II). If sodium oxide is added to the same acid, the result is the compound now discussed and water.
Reaction of nitric acid with acid salts or sodium hydroxide
If an acid sodium salt (for example, its bicarbonate) is added to HNO 3, the desired substance, water and carbon dioxide, are formed, which quickly evaporates. If the second reagent is sodium hydroxide, then, as is the case with its oxide and nitric acid, only sodium nitrate and H 2 O are obtained.
Reaction of ammonium nitrate with acid salts or sodium hydroxide
Ammonium nitrate can successfully replace nitric acid. During its interaction with sodium hydroxide, the desired substance, water and ammonia gas are formed, and when reacting with sodium bicarbonate, the products will be these compounds and carbon dioxide.
Reaction of table salt with silver nitrate
In this case, an exchange reaction occurs, the products of which are sodium nitrate and silver chloride.
Application
This substance is used as a valuable nitrogen fertilizer. The pyrotechnic, food, glass and metalworking industries cannot do without sodium nitrate. Sodium nitrate is extracted from natural deposits in several ways:
Leaching of this substance using hot water and crystallization;
Absorption of nitrogen oxides using a baking soda solution;
An exchange reaction involving sodium sulfate/chloride/carbonate and calcium/ammonium nitrate.
Conclusion
This is the important role sodium nitrate plays. There are also other substances that a person cannot do without, but we’ll talk about them another time.
