Nitrogen and phosphorus compounds of nitrogen and phosphorus. Ammonium nitrates decompose
Nitric acid HNO3 in its pure form is a colorless liquid with a sharp, suffocating odor. In small quantities, it is formed during lightning discharges and is present in rainwater. Under the action of light, nitric acid partially decomposes with the release of NO2 and due to this it acquires a light brown color: 4HNO3 \u003d 4NO2 + 2H2O + O2. Nitric acid is one of the...
When solid nitrates are heated, they all decompose with the release of oxygen (the exception is ammonium nitrate), while they can be divided into four groups. The first group consists of alkali metal nitrates, which, when heated, decompose into nitrites and oxygen: 2KNO3 = 2KNO2 + O2. The second group consists of the majority of nitrates (from alkaline earth metals to copper inclusive), decomposing into metal oxide, NO2 and oxygen: 2Cu (NO3) 2 \u003d 2CuO + 4NO2 + O2, The third group consists of nitrates of the most heavy metals (AgNO3 and Hg (NO3) 2 ), decomposing to free metal, NO2 and oxygen: Hg (NO3) 2 \u003d Hg + 2NO2 + O2. The fourth "group" is ammonium nitrate: NH4NO3 \u003d N2O + 2H2O.
Nitrous acid HNO2 belongs to weak acids (K = 6.10-4 at 25 °C), is unstable and is known only in dilute solutions in which the equilibrium 2HNO2 NO + NO2 + H2O occurs. Nitrites, unlike the acid itself, are stable even when heated. An exception is crystalline ammonium nitrite, which decomposes into free nitrogen and water when heated.
Of the three phosphoric acids, orthophosphoric acid H3PO4 (often referred to simply as phosphoric acid) is of greatest practical importance, a white solid that is highly soluble in water. In aqueous solution, it dissociates in steps. As tribasic, phosphoric acid forms three types of salts: dihydrophosphates (NaH2PO4); hydrophosphates (Na2HPO4); phosphates (Na3PO4). All dihydrogen phosphates are soluble in water. Of the hydrophosphates and phosphates, only alkali metal and ammonium salts are soluble in water. Salts of phosphoric acid are valuable mineral fertilizers. The most common among them are superphosphate, precipitate and phosphate rock. Simple superphosphate is a mixture of calcium dihydrogen phosphate Ca (H2PO4) 2 and "ballast" CaSO4. It is obtained by treating phosphorites and apatites with sulfuric acid. When mineral phosphates are treated with phosphoric acid, double superphosphate Ca(H2PO4)2 is obtained. When phosphoric acid is quenched with lime, the CaHPO4.2H2O precipitate is obtained. Complex fertilizers (i.e. containing both nitrogen and phosphorus; or nitrogen, phosphorus and potassium) are important. Of these, ammophos is the best known - a mixture of NH4H2PO4 and (NH4)2HPO4.
Under normal conditions - a colorless gas, with a pungent odor (the smell of "ammonia"); liquefies at -33.4°C and solidifies at -77.7°C. The ammonia molecule has the shape of a pyramid; in liquid ammonia, the NH3 molecules are linked by hydrogen bonds, thereby causing an abnormally high boiling point. Polar NH3 molecules are very soluble in water (700 volumes of NH3 in one volume of H2O)…
Phosphorus forms two chlorides: phosphorus trichloride PCl3 and phosphorus pentachloride PCl5. Phosphorus trichloride is produced by passing chlorine over the surface of white phosphorus. In this case, phosphorus burns with a pale green flame, and the resulting phosphorus chloride condenses as a colorless liquid. Phosphorus trichloride is hydrolyzed by water to form phosphorous acid and hydrogen chloride: PCl3 + ZH2O = H3PO3 + ZHCl. Phosphorus pentachloride can be obtained in the laboratory ...
In oxides, the oxidation state of nitrogen varies from +1 to +5. Oxides N2O and NO are colorless gases, nitric oxide (IV) NO2 is a brown gas, which has received the name “fox tail” in the industry. Nitric oxide (III) N2O3 is a blue liquid, nitric oxide (V) N2O5 under normal conditions is transparent colorless crystals. The trivial name of nitric oxide (I) is often used ...
Phosphoric anhydride P2O5 (the "simplest" formula) is the most stable phosphorus oxide under normal conditions. It is a solid white substance of the composition P4O10. Phosphoric anhydride is described by the simplest formula P2O3 and the true formula P4O6. It has been shown that phosphorus in P4O6 is coordinatively unsaturated and therefore unstable. The interaction of P4O6 with hot water leads to disproportionation of P4O6 + 6H2O = PH3 + 3H3PO4; Gaseous HCl decomposes P4O6: P4O6 + 6HCl = 2H3PO3 + 2PCl3. P4O10 actively interacts with water, and also takes it away from other compounds, forming, depending on the conditions, either metaphosphoric HPO3, or orthophosphoric H3PO4, or pyrophosphoric H4P2O7 acid. That is why P4O10 is widely used as a desiccant of various substances from water vapor.
Nitric acid is a strong acid. Her salts nitrates- obtained by the action of HNO 3 on metals, oxides, hydroxides or carbonates. All nitrates are highly soluble in water. The 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 standing in the series of voltages to the left of magnesium:
b) nitrates of metals located in a series of voltages between magnesium and copper:
c) nitrates of metals located in a series of voltages to the right of mercury:
d) ammonium nitrate:
Nitrates in aqueous solutions practically do not show oxidizing properties, but at high temperatures in the solid state they are strong oxidizing agents, for example, when solids are fused:
Zinc and aluminum in an alkaline solution reduce nitrates to NH 3:
Nitrates are widely used as fertilizers. At the same time, almost all nitrates are highly soluble in water, therefore, in the form of minerals, they are extremely small in nature; 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 the petrochemical industry, nitrogen is used to purge tanks and pipelines, test the operation of pipelines under pressure, and increase the production of deposits. In mining, nitrogen can be used to create an explosion-proof environment in mines, to burst rock layers.
An important field 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 amounts of nitrogen are used in coke production (“dry coke quenching”) during unloading coke from coke oven batteries, as well as for "squeezing" fuel in rockets from tanks to pumps or engines.
In the food industry, nitrogen is registered as a food additive. E941, as a gas medium for packaging and storage, a refrigerant, and liquid nitrogen is used when bottling oils and non-carbonated drinks to create overpressure and an inert atmosphere in soft containers.
Nitrogen gas fills the tire chambers of the landing gear of aircraft.
31. Phosphorus - obtaining, properties, application. Allotropy. Phosphine, phosphonium salts - preparation and properties. Metal phosphides, preparation and properties.
Phosphorus- a chemical element of the 15th group of the third period of the periodic system of D. I. Mendeleev; has atomic number 15. The element belongs to the group of pnictogens.
Phosphorus is obtained from apatite or phosphorite as a result of interaction with coke and silica at a temperature of about 1600 ° C:
The resulting phosphorus vapor condenses 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, including 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, chemical activity decreases. White phosphorus in the air, when oxidized by atmospheric 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)
Interacts with many simple substances - halogens, sulfur, some metals, showing oxidizing and reducing properties: with metals - an oxidizing agent, forms phosphides; with non-metals - reducing agent.
Phosphorus practically does not combine with hydrogen.
In cold concentrated alkali solutions, the disproportionation reaction also proceeds slowly:
Strong oxidizing agents convert phosphorus to phosphoric acid:
The oxidation reaction of phosphorus occurs when matches are ignited, Berthollet salt acts as an oxidizing agent:
White ("yellow") phosphorus is the most chemically active, toxic and flammable, and therefore 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 manufacture of matches, explosives, incendiary compositions, various types of fuels, as well as extreme pressure lubricants, as a getter in the manufacture of incandescent lamps.
Elemental phosphorus under normal conditions exists in the form of several stable allotropic modifications. All possible allotropic modifications of phosphorus have not yet been fully studied (2016). Traditionally, four of its 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 ultrahigh pressures, the metallic form of the element is thermodynamically stable. All modifications differ in color, density and other physical and chemical characteristics, especially in chemical activity. When the state of a substance passes into a more thermodynamically stable modification, the 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 the action of water or acids on phosphides:
Hydrogen chloride, when heated, interacts with white phosphorus:
Decomposition of phosphonium iodide:
Decomposition of phosphonic acid:
or restore 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 the base is much weaker than ammonia. The latter is explained by the fact that the H–P bonds are weakly polarized and the lone pair activity of 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 give phosphonium salts containing the PH 4 + ion (similar to ammonium). Phosphonium salts, colorless crystalline substances, are extremely unstable, easily hydrolyzed.
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 both by direct interaction of substances at a temperature of about 1000 ° C, and 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. In this case, a phosphine is obtained and, in the case of hydrolysis, a 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
With moderate heating, most phosphides decompose. Melted under excess pressure of phosphorus vapor.
Boron phosphide BP, on the contrary, is refractory (t pl. 2000 ° C, with decomposition), a very inert substance. It decomposes only with concentrated oxidizing acids, reacts when heated with oxygen, sulfur, alkalis during sintering.
32. Phosphorus oxides - molecular structure, production, properties, application.
Phosphorus forms several oxides. The most important of these 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- waxy crystalline mass, melting at 22.5 ° C and turning into a colorless liquid. Poisonous.
When dissolved in cold water, it forms phosphorous acid:
P 4 O 6 + 6H 2 O \u003d 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. The sublimation temperature is 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 interconnected by weak intermolecular forces, which are easily broken when heated. Hence the volatility of this variety. Other modifications are polymeric. They are formed by infinite layers of PO 4 tetrahedra.
When P 4 O 10 interacts with water, phosphoric acid is formed:
P 4 O 10 + 6H 2 O \u003d 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 with the formation of its anhydride:
4HClO 4 + P 4 O 10 \u003d (HPO 3) 4 + 2Cl 2 O 7.
P 4 O 10 is used as a dryer for gases and liquids.
It is widely used in organic synthesis in dehydration and condensation reactions.
Nitrogen and Phosphorus
The elements Nitrogen and Phosphorus are located in Group V of the Periodic Table, Nitrogen in the 2nd period, Phosphorus - in the 3rd. Electronic configuration of the nitrogen atom:
Nitrogen valency: III and IV, oxidation state in compounds: from -3 to +5.
The structure of the nitrogen molecule:,.
Electronic configuration of the Phosphorus atom: 
Electronic configuration of a Phosphorus atom in an excited state: 
Phosphorus valency: III and V, oxidation state in compounds: -3, 0, +3, +5.
Physical properties of nitrogen. Colorless gas, tasteless and odorless, slightly lighter than air g/mol, g/mol), poorly soluble in water. Melting point -210 °C, boiling point -196 °C.
Allotropic modifications of Phosphorus. Among the simple substances that form the element Phosphorus, the most common are white, red and black phosphorus.
Distribution of nitrogen in nature. Nitrogen occurs in nature mainly in the form of molecular nitrogen. In the air, the volume fraction of nitrogen is 78.1%, mass - 75.6%. Nitrogen compounds are found in soil in small amounts. As part of organic compounds (proteins, nucleic acids, ATP), nitrogen is found in living organisms.
Distribution of phosphorus in nature. Phosphorus is found in a chemically bound state in the composition of minerals: phosphorites, apatites, the main component of which is . Phosphorus is a vital element that is part of lipids, nucleic acids, ATP, calcium orthophosphate (in bones and teeth).
Obtaining nitrogen and phosphorus.
Nitrogen obtained in industry from liquid air: since nitrogen has the lowest boiling point of all atmospheric gases, it evaporates first from liquid air. In the laboratory, nitrogen is obtained by thermal decomposition of ammonium nitrite: . Phosphorus obtained from apatites or phosphorites by calcining them with coke and sand at a temperature of:
Chemical properties of nitrogen.
1) Interaction with metals. The substances formed as a result of these reactions are called nitrides and.
At room temperature, nitrogen reacts only with lithium:
Nitrogen reacts with other metals at high temperatures:
- aluminum nitride
Nitrogen interacts with hydrogen in the presence of a catalyst at high pressure and temperature:
- ammonia
At very high temperatures (about ) nitrogen reacts with oxygen:
- nitrogen(II) oxide
Chemical properties of phosphorus.
1) Interaction with metals.
When heated, phosphorus reacts with metals:
- calcium phosphide
2) Interaction with non-metals.
White phosphorus ignites spontaneously, while red phosphorus burns when ignited:
- phosphorus(V) oxide
With a lack of oxygen, phosphorus (III) oxide is formed (a very toxic substance):
Interaction with halogens:
Interaction with sulfur:
Ammonia
Molecular formula of ammonia: . Electronic formula:
Structural formula:
Physical properties of ammonia. A colorless gas with a characteristic pungent odor, almost twice as light as air, poisonous. When pressure is increased or cooled, it easily condenses into a colorless liquid, boiling point, melting point. Ammonia dissolves very well in water: at 1 volume of water, up to 700 volumes of ammonia dissolve, at - 1200 volumes.
Getting ammonia.
1) Ammonia is obtained in the laboratory by heating a dry mixture of calcium hydroxide (slaked lime) and ammonium chloride (ammonia):
2) Ammonia in industry is obtained from simple substances - nitrogen and hydrogen:
Chemical properties of ammonia. Nitrogen in ammonia has the lowest oxidation state and therefore exhibits only reducing properties.
1) Combustion in an atmosphere of pure oxygen or in heated air:
2) Oxidation to nitrogen (II) oxide in the presence of a catalyst (hot platinum):
3) Reverse interaction with water:
The presence of ions determines the alkaline environment of the ammonia solution. The resulting solution is called ammonia or ammonia water. Ammonium ions exist only in solution. It is impossible to isolate ammonium hydroxide as an independent compound.
4) Recovery of metals from their oxides:
5) Interaction with acids to form ammonium salts (compound reaction):
- ammonium nitrate.
Application of ammonia. A large amount of ammonia is spent on the production of nitric acid, nitrogen salts, urea, soda by the ammonia method. Its use in refrigeration plants is based on light accumulation and subsequent evaporation with heat absorption. Aqueous solutions of ammonia are used as nitrate fertilizers.
ammonium salts
ammonium salts- salts containing a cation group. For example, - ammonium chloride, - ammonium nitrate, - ammonium sulfate. Physical properties of ammonium salts. White crystalline substances, highly soluble in water.
Getting ammonium salts. Ammonium salts are formed by the interaction of gaseous ammonia or its solutions with acids:
Chemical properties of ammonium salts.
1) Dissociation:
2) Interaction with other salts:
3) Interaction with acids:
4) Interaction with alkalis:
This reaction is qualitative for ammonium salts. Ammonia released is determined by the smell or by the blue color of wet indicator paper.
5) Decomposition when heated:
The use of ammonium salts. Ammonium salts are used in the chemical industry and as mineral fertilizers in agriculture.
Nitrogen oxides and phosphorus oxides
Nitrogen forms oxides in which it exhibits an oxidation state of +1 to +5: ; NO; ; ; ; . All nitrogen oxides are poisonous. The oxide has narcotic properties, which at the initial stage are indicated by euphoria, hence the name - "laughing gas". The oxide irritates the respiratory tract and mucous membranes of the eyes. A harmful consequence of chemical production, it enters the atmosphere in the form of a "fox tail" - a red-brown color.
Phosphorus oxides: and. Phosphorus(V) oxide is the most stable oxide under normal conditions.
Obtaining nitrogen oxides and phosphorus oxides.
With a direct combination of molecular nitrogen and oxygen, only nitrogen (II) oxide is formed:
Other oxides are obtained indirectly.
Phosphorus (V) oxide is obtained by burning phosphorus in excess oxygen or air:
Chemical properties of nitrogen oxides.
1) - oxidizing agent, can support combustion:
2) NO - easily oxidized:
Does not react with water and alkalis.
3) acid oxide:
4) - strong oxidizing agent, acidic oxide:
In the presence of excess oxygen:
Dimerizes, forming an oxide - a colorless liquid: . The reaction is reversible. At -11 °С, the equilibrium is practically shifted towards formation, and at 140 °С - towards formation.
5) - acid oxide:
Chemical properties of phosphorus(V) oxide. Phosphorus-containing acids.
is typically an acidic oxide. Three acids correspond to it: meta-,ortho- and diphosphate a. When dissolved in water, metaphosphate acid is first formed:
With prolonged boiling with water - orthophosphate acid:
When orthophosphate acid is carefully calcined, diphosphate acid is formed:
The use of nitrogen oxides and phosphorus oxides.
Nitrogen (IV) oxide is used in the production of nitric acid, nitrogen (I) oxide - in medicine.
Phosphorus(V) oxide is used for drying gases and liquids, and in some cases for splitting chemically bound water from substances.
Nitric and phosphate acids
Physical properties of orthophosphate (phosphoric) acid. Under normal conditions - a solid, colorless, crystalline substance. Melting point +42.3. In solid and liquid acids, molecules are combined by hydrogen bonds. This is due to the increased viscosity of concentrated solutions of phosphoric acid. It is highly soluble in water, its solution is an electrolyte of medium strength. Physical properties of nitric acid. Anhydrous (100%) acid is a colorless liquid, smells strongly, boiling point. In the case of storage in the light, it gradually turns brown due to decomposition and the formation of higher nitrogen oxides, including brown gas. It mixes well with water in any ratio.
Getting phosphate acid.
1) From its salts contained in phosphate minerals (apatites and phosphorites), under the action of sulfuric acid:
2) Hydration of phosphorus(V) oxide:
Obtaining nitrate acid.
1) From dry salts of nitric acid when exposed to concentrated sulfuric acid:
2) With nitrogen oxides:
3) Industrial synthesis of nitric acid:
- catalytic oxidation of ammonia, catalyst - platinum.
- Oxidation by atmospheric oxygen.
- absorption by water in the presence of oxygen.
Chemical properties of phosphoric acid. Shows all the typical properties of acids. Phosphate acid - tribasic, forms two series of acid salts - dihydrophosphate and hydrophosphate s.
1) Dissociation:
4) Interaction with salts. The reaction with argentum nitrate is qualitative for the ion - a yellowish precipitate of argentum phosphate precipitates:
5) Interaction with metals standing in the electrochemical series of voltages up to Hydrogen:
Chemical properties of nitric acid. Nitric acid is a strong oxidizing agent.
1) Dissociation:
2) Interaction with metal oxides:
3) Interaction with bases:
4) Interaction with salts:
5) Interaction with metals. When concentrated and dilute nitric acid reacts with metals, salt (nitrate), nitrogen oxides, nitrogen or ammonia and water are formed.
The use of orthophosphate and nitric acids.
Orthophosphate acid widely used in the production of mineral fertilizers. It is not poisonous and is used in the food industry for the manufacture of syrups, drinks (Coca-Cola, Pepsi-Cola).
Nitric acid spent on the production of nitrogen fertilizers, explosives, medicines, dyes, plastics, artificial fibers and other materials. Concentrated nitric acid is used in rocket technology as a propellant oxidizer.
Nitrates
Salts of nitric acid - nitrate s. These are solid crystalline NH 4 NO 3 → N 2 O + 2H 2 O
NH 4 NO 3 → N 2 + NO + H 2 O
Nitrites do not decompose, except for NH 4 NO 2
NH 4 NO 2 → N 2 + 2H 2 O
Obtaining nitric acid
In laboratory conditions - KNO 3tv + H 2 SO 4 k \u003d KHSO 4 + HNO 3
In industry: ammonia or contact method.
Catalytic oxidation in a contact apparatus (catalyst - platinum-rhodium grids)
1) 4NH 3 + 5O 2 → 4NO + 6H 2 O
2) NO + O 2 → NO 2 at normal t and increased P ≈ 600 - 1100 kPa
3)4NO 2 + O 2 + H 2 O → 4HNO 3 ω (50 - 60%)
Salts of nitric acid. nitrogen fertilizers
Nitrates - almost all are highly soluble in H 2 O, so natural deposits are rare. The main amount is obtained artificially in chemical plants, from HNO 3 and hydroxides.
Get:
1) Interaction with metals, bases, amphoteric bases, alkalis, insoluble bases, ammonia or its aqueous solution, with some salts.
2) NO 2 with alkali solutions
2Ca(OH) 2 + NO 2 = Ca(NO 3) 2 + Ca(NO 2) 2 + 2H 2 O
AT sour environment nitrates exhibit oxidizing properties like dilute HNO 3
3FeCl 2 + KNO 3 + 4HCl = 3FeCl 3 + KCl + NO + H 2 O
AT alkaline oxidize active metals (Mg, Al, Zn)
4Zn + NaNO 3 + 7NaOH + 6H 2 O = 4Na 2 + NH 3
Nitrates show the strongest oxidizing properties when fused
Cr 2 O 3 + 3NaNO 3 + 4KOH = 3K 2 CrO 4 + 3NaNO 2 + 2H 2 O
The most important nitrogen fertilizers are:
Sodium, potassium, ammonium and calcium nitrates are mainly used as mineral nitrogen fertilizers and are called saltpeter.
NH 4 NO 3 (NH 4) 2 SO 4 ammonium sulfate
KNO 3 saltpeter NH 3 H 2 O ammonia water
NaNO 3 NH 4 H 2 PO 4 ammophos
Ca (NO 3) 2 (NH 4) 2 HPO 4 diammophos
CO (NH 2) 2 urea, carbamide
The nutritional value of the fertilizer is dissolved by ω(N) in it.
In urea ω(N) = (2 14)/ (12 + 16 + 28 + 4) = 28/60 = 0.47 (47%).
In NH 4 NO 3 - nitrogen in nitrate and ammonia form (35%), (NH 4) 2 SO 4 - the most valuable fertilizer, since nitrogen is most of all in a well-assimilated form.
Nitrogen fertilizers, as sources of nitrogen nutrition for plants to increase productivity, also include organic fertilizers (manure, compost, etc.), as well as green fertilizers (lupine).
Chemistry of phosphorus
Phosphorus(lat. Phosphorus) - one of the most common elements in the earth's crust. It is not found in the free state in nature due to its high chemical activity. In a bound form, it is part of about 200 minerals, mainly apatites 3Ca 3 (PO 4) 2 * CaX 2 (X \u003d Cl, F, OH) 2 and phosphorites Ca 3 (PO 4) 2.
There are 11 known allotropic modifications of phosphorus, the most studied are white, red and black phosphorus. White phosphorus has the molecular formula P 4 and is a regular tetrahedron with a bond angle of 60 O. 
White phosphorus is highly toxic. The lethal dose for humans is 0.15 g. Already at room temperature, white phosphorus easily evaporates and its vapors are oxidized. The energy of these reactions is partially converted into light, which is the reason for the glow of white phosphorus in the dark.
It ignites easily (self-ignition is possible). It must be handled with extreme caution. Must be stored under water.
red phosphorus obtained by prolonged heating of white phosphorus at a temperature of 280-340 ° C under pressure and without air access. It is a dark red, finely crystalline, low-volatile substance.
280 - 340°C 200°C
P white → P red P white → R black
Red phosphorus is almost non-toxic and less flammable than white phosphorus. Spontaneous combustion does not occur, but it is easy to ignite and combustion proceeds very rapidly.
At the core, polymers are obtained by opening the tetrahedron P 4 .
The most stable form of phosphorus is black phosphorus. In appearance and properties, it resembles graphite, greasy to the touch, is divided into scales, and conducts electric current. It is not poisonous, chemically the least inactive, ignites only at a temperature of 490 ° C.
Although phosphorus is an electronic analogue of nitrogen, the presence of free d orbitals in the valence electron layer of the atom makes phosphorus compounds unlike nitrogen compounds.
The difference between nitrogen and phosphorus compounds is associated with the formation of donor-acceptor π-bonds between phosphorus atoms and electron pair donors, especially oxygen. Therefore, when moving from N to P, the strength of the E-H bonds decreases due to an increase in the size of the atom, but the E-O bonds become significantly stronger.
The formation of donor-acceptor bonds explains the intense interaction of phosphorus with oxygen, the stability and diversity of oxygen compounds of phosphorus.
The most stable oxidation state is +5. In this oxidation state, phosphorus compounds do not exhibit oxidizing properties due to its stability, unlike nitrogen. Because there are free 3d - orbitals, then compared to nitrogen, there are more valence possibilities and the maximum valence of phosphorus can be 5, rarely 6.
Receipt:
1. From phosphate rock by fusion with carbon and silicon oxide
Ca 3 (PO 4) 2 + C + SiO 2 → P 4 + CaSiO 3 + CO
2. From Ca phosphate, at temperatures above 1500 ° C
Ca 3 (PO 4) 2 + C → CaO + P 4 + CO
Chemical properties:
P + O 2 \u003d P 2 O 3
P + O 2 \u003d P 2 O 5
P + S = P 2 S 3
P + Cl 2 = PCl 3
