Ca 3 (PO 4) 2 + 3SiO 2 + 5C = 3CaSiO 3 + 5CO + P 2

Phosphorus vapor at this temperature consists almost entirely of P 2 molecules, which condense into P 4 molecules upon cooling.

When the vapor condenses, it forms white (yellow) phosphorus, which consists of P 4 molecules having the shape of a tetrahedron. It is a highly reactive, soft, waxy, pale yellow substance, soluble in carbon disulfide and benzene. In air, phosphorus ignites at 34 ° C. It has the unique ability to glow in the dark due to slow oxidation to lower ones. It was white phosphorus that was once isolated by Brand.

If white phosphorus is heated without access to air, it turns into red (it was first obtained only in 1847). Name red phosphorus refers to several modifications at once, differing in density and color: it ranges from orange to dark red and even purple. All varieties of red phosphorus are insoluble in organic solvents, compared to white phosphorus, they are less reactive (ignite in air at t> 200 ° C) and have a polymeric structure: these are P 4 tetrahedra connected to each other in endless chains. The "violet phosphorus" is somewhat different from them, which consists of P 8 and P 9 groupings laid in long tubular structures with a pentagonal cross section.

At high blood pressure white phosphorus is converted to black phosphorus, built from volumetric hexagons with phosphorus atoms at the vertices, connected to each other in layers. For the first time this transformation was carried out in 1934 by the American physicist Percy Williams Bridgman. The structure of black phosphorus resembles graphite, with the only difference that the layers formed by phosphorus atoms are not flat, but “corrugated”. Black phosphorus is the least active modification of phosphorus. When heated without access to air, it, like red, passes into vapor, from which white phosphorus condenses.

White phosphorus is very toxic: the lethal dose is about 0.1 g. Due to the danger of self-ignition in air, it is stored under a layer of water. Red and black phosphorus are less toxic, since they are non-volatile and practically insoluble in water.

Chemical properties

The most chemically active is white phosphorus (for simplicity, in the equations of reactions involving white phosphorus, they are written as P, and not P 4, especially since similar reactions are possible with the participation of red phosphorus, the molecular composition of which is undetermined). Phosphorus combines directly with many simple and complex substances. In chemical reactions, phosphorus, like , can be both an oxidizing agent and a reducing agent.

How oxidizer phosphorus interacts with many to form phosphides, for example:

2P + 3Ca = Ca 3 P 2

P + 3Na = Na 3 P

Please note that phosphorus practically does not combine directly with phosphorus.

How reducing agent phosphorus interacts with, halogens, sulfur (i.e. with more electronegative non-metals). In this case, depending on the reaction conditions, both phosphorus (III) compounds and phosphorus (V) compounds can be formed.

a) with slow oxidation or with a lack of oxygen, phosphorus is oxidized to phosphorus (III) oxide, or phosphorous anhydride P 2 O 3:

4P + 3O 2 \u003d 2P 2 O 3

When phosphorus is burned in excess (or air), phosphorus oxide (V) is formed, or phosphoric anhydride P 2 O 5:

4P + 5O 2 \u003d 2P 2 O 5

b) depending on the ratio of reagents, the interaction of phosphorus with halogens and sulfur forms, respectively, halides and sulfides of tri- and pentavalent phosphorus; for example:

2P + 5Cl 2 (ex.) \u003d 2PCl 5

2P + 3Cl 2 (insufficient) = 2PCl 3

2P + 5S(e) = P 2 S 5

2P + 3S (insufficient) = P 2 S 3

It should be noted that phosphorus forms only the PI3 compound with iodine.

Phosphorus plays the role of a reducing agent in reactions with oxidizing acids:

3P + 5HNO 3 + 2H 2 O = 3H 3 PO 4 + 5NO

- with concentrated nitric acid:

P + 5HNO 3 \u003d H 3 PO 4 + 5NO 2 + H 2 O

- with concentrated sulfuric acid:

2P + 5H 2 SO 4 \u003d 2H 3 PO 4 + 5SO 2 + 2H 2 O

Phosphorus does not interact with other acids.

When heated with aqueous solutions, phosphorus undergoes disproportionation, for example:

4P + 3KOH + 3H 2 O \u003d PH 3 + 3KH 2 PO 2

8P + 3Ba(OH) 2 + 6H 2 O = 2PH 3 + 3Ba(H 2 PO 2) 2

In addition to phosphine PH 3, as a result of these reactions, salts of hypophosphorous acid H 3 PO 2 are formed - hypophosphites, in which phosphorus has a characteristic oxidation state of +1.

The use of phosphorus

The main part of the phosphorus produced in the world is spent on the production of phosphoric acid, from which fertilizers and other products are obtained. Red phosphorus is used in the manufacture of matches, it is contained in the mass, which is applied to the matchbox.

Phosphine

The best known hydrogen compound of phosphorus is phosphine PH 3 . Phosphine is a colorless gas with a garlic odor and is highly toxic. Highly soluble in organic solvents. Unlike ammonia, it is sparingly soluble in water. Phosphine has no practical value.

Receipt

Above was considered a method for obtaining phosphine by the interaction of phosphorus with aqueous solutions. Another way is the action of hydrochloric acid on metal phosphides, for example:

Zn 3 P 2 + 6HCl \u003d 2PH 3 + 3ZnCl 2

Chemical properties

1. Acid - basic properties

Being sparingly soluble in water, phosphine forms an unstable hydrate with it, which exhibits very weak basic properties:

PH 3 + H 2 O ⇄ PH 3 ∙H 2 O ⇄ PH 4 + + OH -

Phosphonium salts are formed only with:

PH 3 + HCl = PH 4 Cl

PH 3 + HClO 4 = PH 4 ClO 4

1. Redox properties

The full list of abstracts can be viewed

*in the recording image is a photograph of white phosphorus

DEFINITION

Phosphine(phosphorus hydride, monophosphane) under normal conditions is a colorless gas, poorly soluble in water and does not react with it.

The gross formula is PH 3 (the structure of the molecule is shown in Fig. 1). The molar mass of phosphine is 34.00 g/mol.

Rice. 1. The structure of the phosphine molecule, indicating the bond angle and the length of the chemical bond.

At low temperatures, it forms a solid clarate 8PH 3 ×46H 2 O. Density - 1.5294 g / l. Boiling point - (-87.42 o C), melting point - (-133.8 o C).

In OVR, it is a strong reducing agent; it is oxidized by concentrated sulfuric and nitric acids, iodine, oxygen, hydrogen peroxide, and sodium hypochlorite. Donor properties are much less pronounced than those of ammonia.

PH3, the oxidation states of the elements in it

To determine the oxidation states of the elements that make up the phosphine, you first need to figure out for which elements this value is exactly known.

Phosphine is the trivial name for phosphorus hydride, and, as you know, the oxidation state of hydrogen in hydrides is (+1). To find the oxidation state of phosphorus, let's take its value as "x" and determine it using the electroneutrality equation:

x + 3×(+1) = 0;

So the oxidation state of phosphorus in phosphine is (-3):

Examples of problem solving

EXAMPLE 1

Exercise	Determine the oxidation states of acid-forming elements in the following compounds: HNO 2 , H 2 CO 3 , H 4 SiO 4 , HPO 3 .
Solution	In these compounds, the acid-forming elements are nitrogen, carbon, silicon and phosphorus. The oxidation state of oxygen is (-2), and hydrogen - (+1). Let us take the oxidation state of the acid-forming element as “x” and use the electroneutrality equation to find its value: 1 + x + 2×(-2) = 0; The oxidation state of nitrogen is (+3). 2×(+1) + x + 3×(-2) = 0; The oxidation state of carbon is (+4). 4×(+1) + x + 4×(-2) = 0; The oxidation state of silicon is (+4). 1 + x + 3×(-2) = 0; The oxidation state of phosphorus is (+5).
Answer	HN +3 O 2, H 2 C +4 O 3, H 4 Si +4 O 4, HP +5 O 3

EXAMPLE 2

Exercise	Iron exhibits the highest oxidation state in the compound: K4; K3; Fe(OH)2.
Solution	In order to give the correct answer to the question posed, we will alternately determine the degree of oxidation of iron in each of the proposed compounds using the electrical neutrality equation. a) The oxidation state of potassium is always (+1). The oxidation state of carbon in the cyanide ion is (+2), and nitrogen - (-3). Let's take for "x" the value of the oxidation state of iron: 4x1 + x + 6x2 + 6x (-3) = 0; b) The oxidation state of potassium is always (+1). The oxidation state of carbon in the cyanide ion is (+2), and nitrogen - (-3). Let's take for "x" the value of the oxidation state of iron: 3x1 + x + 6x2 + 6x (-3) = 0; c) The degree of oxidation of oxygen in oxides (-2). Let's take for "x" the value of the oxidation state of iron: d) The oxidation states of oxygen and hydrogen are (-2) and (+1), respectively. Let's take for "x" the value of the oxidation state of iron: x + 2×(-2) + 2× 1 = 0; The highest oxidation state of iron is (+3) and it manifests it in a compound of composition K 3 .
Answer	Option 2

Nearest source of stone containing phosphine, was indicated on the maps, and David sent a working group of blue and green horsemen there, who were supposed to start harvesting the fire stone.

Now they were aware of all the tricks of the enemy, they learned to evaluate the features of attacks, learned how to save the strength of horsemen and animals, how to protect themselves from vapors. phosphine and strokes of Threads.

fire jets phosphine, spewing out dragons, formed a constantly changing pattern of light in the air.

The riders discovered deposits phosphine on a plateau somewhere between the Malay River and Sadrid.

As the dragon settled its bulky body on such an unsuitable landing site, its broad wings were driven along the yard smelling of phosphine air.

Then he washed off the stinking phosphine pants and shirt and dried them in the sun, hanging them in the bushes.

When Jaxom entered his room, on his way to change the reeking phosphine flight suit, he caught sight of a sketch of the bay, still laid out on his work table.

Jacksom shoved Ruta's portion into his mouth and, as always experiencing inner trepidation, began to listen to the dragon's powerful teeth crushing the saturated phosphine stone.

Oxidation state in PH3

General information about phosphine and oxidation states in PH3

The gross formula is PH3 (the structure of the molecule is shown in Fig. 1). The molar mass of phosphine is 34.00 g/mol.

The meaning of the word phosphine

1. The structure of the phosphine molecule, indicating the bond angle and the length of the chemical bond.

At low temperatures it forms a solid clarate 8PH3×46H2O. Density - 1.5294 g / l. Boiling point - (-87.42oC), melting point - (-133.8oC).

In OVR, it is a strong reducing agent; it is oxidized by concentrated sulfuric and nitric acids, iodine, oxygen, hydrogen peroxide, and sodium hypochlorite. Donor properties are much less pronounced than those of ammonia.

PH3, the oxidation states of the elements in it

To determine the oxidation states of the elements that make up the phosphine, you first need to figure out for which elements this value is exactly known.

Phosphine is the trivial name for phosphorus hydride, and, as you know, the oxidation state of hydrogen in hydrides is (+1). To find the oxidation state of phosphorus, let's take its value as "x" and determine it using the electroneutrality equation:

x + 3×(+1) = 0;

So the oxidation state of phosphorus in phosphine is (-3):

Examples of problem solving

3. Molecules. Chemical bond. The structure of substances

Chemical particles formed from two or more atoms are called molecules(real or conditional formula units polyatomic substances). Atoms in molecules are chemically bonded.

A chemical bond is an electrical force of attraction that holds particles together. Each chemical bond in structural formulas seems valence line, for example:

H - H (bond between two hydrogen atoms);

H3N - H + (bond between the nitrogen atom of the ammonia molecule and the hydrogen cation);

(K+) - (I-) (bond between potassium cation and iodide ion).

A chemical bond is formed by a pair of electrons ( ), which in the electronic formulas of complex particles (molecules, complex ions) is usually replaced by a valence line, in contrast to their own, unshared electron pairs of atoms, for example:

The chemical bond is called covalent, if it is formed by the socialization of a pair of electrons by both atoms.

In the F2 molecule, both fluorine atoms have the same electronegativity, therefore, the possession of an electron pair is the same for them. Such a chemical bond is called non-polar, since each fluorine atom has electron density the same in electronic formula molecules can be conditionally divided between them equally:

In the HCl molecule, the chemical bond is already polar, since the electron density on the chlorine atom (an element with greater electronegativity) is much higher than on the hydrogen atom:

A covalent bond, for example H - H, can be formed by sharing the electrons of two neutral atoms:

H + H > H – H

H H

This bonding mechanism is called exchange or equivalent.

According to another mechanism, the same covalent bond H – H arises when the electron pair of the hydride ion H is shared by the hydrogen cation H+:

H+ + (:H)-> H – H

H H

The H+ cation in this case is called acceptor and the anion H - donor electron pair. The mechanism of formation of a covalent bond in this case will be donor-acceptor, or coordinating.

Single bonds (H - H, F - F, H - CI, H - N) are called a-links, they determine the geometric shape of the molecules.

Double and triple bonds () contain one?-component and one or two?-components; ?-component, which is the main and conditionally formed first, is always stronger than?-components.

The physical (actually measurable) characteristics of a chemical bond are its energy, length, and polarity.

Chemical bond energy (E cv) is the heat that is released during the formation of this bond and is spent on breaking it. For the same atoms, a single bond is always weaker than a multiple (double, triple).

Chemical bond length (l s) - internuclear distance. For the same atoms, a single bond is always longer than a multiple.

Polarity communication is measured electric dipole moment p- the product of a real electric charge (on the atoms of a given bond) by the length of the dipole (i.e.

Phosphorus. Phosphine

bond length). The larger the dipole moment, the higher the polarity of the bond. The real electric charges on atoms in a covalent bond are always smaller in value than the oxidation states of the elements, but they coincide in sign; for example, for the H + I-Cl-I bond, the real charges are H + 0'17-Cl-0'17 (bipolar particle, or dipole).

Polarity of molecules determined by their composition and geometric shape.

Non-polar (p = O) will be:

a) molecules simple substances, since they contain only non-polar covalent bonds;

b) polyatomic molecules complex substances, if their geometric shape symmetrical.

For example, CO2, BF3 and CH4 molecules have the following directions of equal (along length) bond vectors:

When bond vectors are added, their sum always vanishes, and the molecules as a whole are non-polar, although they contain polar bonds.

Polar (p> O) will be:

a) diatomic molecules complex substances, since they contain only polar bonds;

b) polyatomic molecules complex substances, if their structure asymmetrically, i.e., their geometric shape is either incomplete or distorted, which leads to the appearance of a total electric dipole, for example, in the molecules of NH3, H2O, HNO3 and HCN.

Complex ions, such as NH4+, SO42- and NO3-, cannot be dipoles in principle, they carry only one (positive or negative) charge.

Ionic bond arises during the electrostatic attraction of cations and anions with almost no socialization of a pair of electrons, for example, between K+ and I-. The potassium atom has a lack of electron density, the iodine atom has an excess. This connection is considered limiting case of a covalent bond, since a pair of electrons is practically in the possession of the anion. Such a relationship is most typical for compounds of typical metals and non-metals (CsF, NaBr, CaO, K2S, Li3N) and substances of the salt class (NaNO3, K2SO4, CaCO3). All these compounds under room conditions are crystalline substances, which are united by the common name ionic crystals(crystals built from cations and anions).

There is another type of connection called metallic bond, in which valence electrons are so loosely held by metal atoms that they do not actually belong to specific atoms.

Atoms of metals, left without external electrons clearly belonging to them, become, as it were, positive ions. They form metal crystal lattice. The set of socialized valence electrons ( electron gas) holds positive metal ions together and at specific lattice sites.

In addition to ionic and metallic crystals, there are also atomic and molecular crystalline substances, in the lattice sites of which there are atoms or molecules, respectively. Examples: diamond and graphite - crystals with an atomic lattice, iodine I2 and carbon dioxide CO2 (dry ice) - crystals with a molecular lattice.

Chemical bonds exist not only inside the molecules of substances, but can also be formed between molecules, for example, for liquid HF, water H2O and a mixture of H2O + NH3:

hydrogen bond is formed due to the forces of electrostatic attraction of polar molecules containing atoms of the most electronegative elements - F, O, N. For example, hydrogen bonds are present in HF, H2O and NH3, but they are not in HCl, H2S and PH3.

Hydrogen bonds are unstable and break quite easily, for example, when ice melts and water boils. However, some additional energy is expended on breaking these bonds, and therefore the melting points (Table 5) and boiling points of substances with hydrogen bonds

(for example, HF and H2O) are significantly higher than for similar substances, but without hydrogen bonds (for example, HCl and H2S, respectively).

Many organic compounds also form hydrogen bonds; The hydrogen bond plays an important role in biological processes.

Examples of Part A assignments

1. Substances with only covalent bonds are

1) SiH4, Cl2O, CaBr2

2) NF3, NH4Cl, P2O5

3) CH4, HNO3, Na(CH3O)

4) CCl2O, I2, N2O

2–4. covalent bond

2. single

3. double

4. triple

present in matter

5. Multiple bonds are present in molecules

6. The particles called radicals are

7. One of the bonds is formed by the donor-acceptor mechanism in the set of ions

8. The most durable and short bond - in a molecule

9. Substances with only ionic bonds - in the set

10–13. The crystal lattice of matter

1) metal

3) nuclear

4) molecular

Phosphorus compounds.

R-3. Metal phosphides are ionic-covalent compounds. Phosphides of s-metals (except Be) and lanthanides are ionic salt-like compounds, they are easily hydrolyzed by water and acids: Mg3P2 + 6H2O = 3Mg(OH)2↓ + 2PH3 Na3P + 3HCl = 3NaCl + PH3. Phosphides of d-elements are metal-like chemically inert compounds. The exception is the phosphides of metals of groups I and II, secondary subgroups, which are also salt-like, but with a large admixture of covalence. Phosphorus does not form stable compounds with antimony, bismuth, lead and mercury.

The combination of phosphorus with hydrogen is called hydrogen phosphide, although the electronegativity of these elements is almost equal. The compound has the formula PH3, called phosphine. It is an extremely poisonous gas with an unpleasant garlic odor, bp=-88°C. There are no hydrogen bonds between phosphine molecules in a liquid and between water and phosphine molecules during dissolution, therefore the boiling point is low and phosphine practically does not dissolve in water. The molecule is a pyramid with a phosphorus atom at the top and an angle of 93.5° between the P-H bonds, which indicates the absence of hybridization of phosphorus atomic orbitals during the formation of this compound. The bonds are formed by almost pure p-orbitals. The lone electron pair of phosphorus remains in the 3s orbital, so phosphine is a weak base and a weak complexing agent in general. The phosphonium cation is formed only with the strongest acids in an anhydrous medium (HJ, HClO4, HBF4), for example PH3 + HJ = PH4J. Water easily decomposes phosphonium salts. Phosphine exhibits strong reducing properties: PH3 + 2O2 = H3PO4 (at 150°C this reaction occurs with an explosion), PH3 + 6AgNO3 + 3H2O = 6Ag↓ + H2(PHO3) + 6AgNO3 PH3 + 3J2 + 3H2O = H2(PHO3) + 6HJ . The synthesis of phosphine from simple substances cannot be carried out, since the P-H bond is not strong enough due to its length and due to the insignificant contribution of the electrostatic component. Therefore, phosphine is obtained by hydrolysis of metal phosphides or by dissolving phosphorus in alkali (the reactions are given above).

The main compounds of phosphorus in its positive oxidation states are oxides, oxygen-containing acids and halides. It is advisable to consider them separately.

Phosphorus oxides– P4O6 and P4O10 are acidic oxides, have a molecular structure, are solids (tmelt (P4O6) = 23.8 ° C, the molecular modification of P4O10 sublimes at 3590 ° C, and the polymer modification melts at 580 ° C), both dissolve in water, giving hydroxides, which are acids, phosphorous and orthophosphoric, respectively. Phosphorus (V) oxide is very hygroscopic, it absorbs moisture from the air, therefore it is used as a desiccant and also as a water-removing agent: P2O5 + HNO3 = HPO3 + N2O5, this forms metaphosphoric acid or polyphosphoric acids - (HPO3) 3-4. Phosphorus (III) oxide, in which phosphorus is in an intermediate oxidation state, is capable of further oxidation reactions and disproportionation reactions, for example: the reaction 5P4O6 = 2P4 + 3P4O10 occurs. Phosphorus (V) oxide does not have oxidizing properties, and can itself be obtained by oxidizing phosphorus under anhydrous conditions, for example, by thermal decomposition of some salts: 6P + 5KClO3 = 3P2O5 + 5KCl

Oxygen acids of phosphorus. The variety of oxygen acids of phosphorus is caused by the following reasons: 1. Valency of phosphorus can be III or V. 2. In the case of valence V, the formation of ortho and meta acids, which differ in the number of attached water molecules, is possible. 3. In all hydroxides, phosphorus exhibits a coordination number of 4, such hydroxides are more stable for it, if there are not enough oxygen atoms, then a P-H bond is formed ((HO) 2PHO, and not P (OH) 3, etc.). 4. Phosphoric acids tend to form linear or cyclic polymers. 5. Under certain conditions, the formation of a P-P bond is possible. 6. As for all hydroxides, peroxo acids are formed during further oxidation. Let us give the structure and properties of the most famous phosphorus acids.

H3PO4 is orthophosphoric acid. This is a tribasic acid, medium in dissociation in the first stage (Ka = 7.52.10-3) and weak in the other two stages. In the anhydrous state, it forms transparent hygroscopic crystals with mp=42°C. It dissolves in water in any concentration. Orthophosphoric acid is obtained by dissolving phosphorus (V) oxide in water, by burning phosphine, by oxidizing any form of phosphorus in an acidic environment, by hydrolyzing binary phosphorus (V) compounds: P4S10 + 16H2O = 4H3PO4 + 10H2S. The industry uses the method of burning phosphorus with subsequent dissolution of the oxide, as well as the displacement of orthophosphoric acid from calcium phosphate with concentrated sulfuric acid when heated: Ca3(PO4)2 + 3H2SO4 = 3CaSO4↓ + 2H3PO4. This acid corresponds to three series of salts - medium (phosphates or orthophosphates) and acidic (hydrophosphates and dihydrophosphates). Phosphates and hydrophosphates of all metals except sodium, potassium, rubidium and cesium are insoluble in water. Dihydrogen phosphates are soluble. Soluble phosphates undergo strong anion hydrolysis, the phosphate anion has the highest hydrolysis constant, and the dihydrophosphate has the lowest. Hydrolysis of the anion leads to an alkaline environment of salt solutions. Acid anions, simultaneously with hydrolysis, participate in the dissociation equilibrium, which leads to an acidic solution environment, for dihydrophosphate to a greater extent, for hydrophosphate to a lesser extent. As a result of these processes, the sodium dihydrogen phosphate solution has a slightly acidic environment, the hydrophosphate solution has a slightly alkaline environment, and the phosphate solution has a strongly alkaline environment. Ammonium phosphate as a salt formed weak acid and base, completely decomposed by water. Orthophosphates melt without decomposition at very high temperatures. Hydrophosphates give diphosphates when heated: 2K2HPO4 = K4P2O7 + H2O. When heated, dihydrophosphates turn into polymetaphosphates: xKH2PO4 = (KPO3)x + H2O. Phosphates do not have strong oxidizing properties, but can be reduced by carbon when heated. In the presence of silicon dioxide, this reaction leads to the production of phosphorus (the reaction equation was given), in the absence of SiO2, the process proceeds as follows: Ca3(PO4)2 + 8C = Ca3P2 + 8CO. Heating ammonium phosphate leads to a gradual loss of ammonia molecules with the formation eventually at temperatures above 300°C polymetaphosphoric acid.

Dehydration of phosphoric acid produces condensed phosphoric acids, which contain one or more bridging oxygen atoms. In this case, chain, cyclic and mixed structures are formed. Let's consider the simplest of them.

Diphosphoric (pyrophosphoric) acid - H4P2O7. It is obtained by heating phosphoric acid to 2000C. In the anhydrous state, it is colorless crystals with mp=61°C, which are highly soluble in water with the formation of a much stronger acid than phosphoric acid. This acid is especially strong in the first two stages. Any condensed acid is stronger than a single acid, since its dissociation produces a more stable anion. Solutions of pyrophosphoric acid are unstable, since a water molecule is gradually added to form two molecules of orthophosphoric acid. More stable are salts - pyrophosphates, which, as already mentioned, can be obtained by heating hydrophosphates.

Metaphosphoric acids - (HPO3) x, where x \u003d 3.4.6. Cyclic condensed acids containing a cycle of alternating phosphorus and oxygen atoms. Obtained by dissolving phosphorus (V) oxide in orthophosphoric acid, as well as by heating pyrophosphoric acid to 300 ° C: 3H4P2O7 \u003d 2 (HPO3) 3 + H2O. All metaphosphoric acids are very strong, for trimetaphosphoric acid Ka2 = 0.02. All these acids are also gradually converted in aqueous solution into phosphoric acid. Their salts are called, respectively, tri-, tetra- and hexametaphosphates.

Oxidation of phosphorus (V) oxide can be obtained peroxophosphoric acid: P4O10 + 4H2O2 + 2H2O = 4H3RO5.

Phosphoric (hypophosphoric) acid H4P2O6 has a P-R connection. The structural formula can be represented as (OH)2OP-RO(OH)2.

Phosphine properties

It can be seen from the formula that the valency of phosphorus is 5, and the oxidation state +4 is a formal value associated with the presence of a bond between identical atoms. This is a tetrabasic acid, the strength of which corresponds to orthophosphoric. It is obtained by the reaction: PbP2O6 + 2H2S = 2PbS↓ + H4P2O6 and is isolated from the solution in the form of a dihydrate with mp=62°C. In an acidic solution, it disproportionates into orthophosphoric and phosphorous acids.

Phosphorous acid H3PO3 or H2. It is a dibasic acid of medium strength, in an anhydrous state - solid with tmelt=74°C. It is obtained by the hydrolysis of phosphorus (III) halides, as well as by the oxidation of white phosphorus with chlorine under water: P4 + 6Cl2 + 12H2O = 4H2 + 12HCl. As mentioned above, the compound of the composition P(OH)3 is less stable, therefore, isomerization occurs with the formation of a P-H bond, which no longer dissociates in an aqueous solution. Salts of phosphorous acid are called phosphites, acidic salts are called hydrophosphites. Most phosphites (except alkali metal salts) are insoluble in water. Like all phosphorus (III) compounds, phosphorous acid is a strong reducing agent, it is oxidized to phosphoric acid by halogens, nitrogen dioxide and other oxidizing agents, and also restores low-active metals from a solution of their salts, for example: HgCl2 + H2 + H2O = H3PO4 + 2HCl + Hg↓. When heated, it disproportionates: 4H2 = 3H3PO4 + PH3.

Phosphorous (phosphinic) acid H3PO2 or H. This is a solid substance with mp=26.5°C, the aqueous solution of which is a fairly strong (Ka=7.9.10-2) monobasic acid. Phosphorus in this compound also has five bonds, two of which are with hydrogen atoms. Only undergoes dissociation N-O bond. The formal oxidation state of phosphorus in this compound is +1. Phosphorous acid and its salts, hypophosphites, are strong reducing agents. Metal cations, even those standing in the voltage series before hydrogen, can be reduced to metal: NiCl2 + Na + 2H2O = H3PO4 + HCl + NaCl + H2 + Ni↓. When heated, phosphorous acid disproportionates: 3H = PH3 + 2H2. With increasing temperature, phosphorous acid has also been shown to decompose into phosphoric acid and phosphine. Hypophosphites of alkali and alkaline earth metals are obtained by the interaction of phosphorus and alkali (see above). Oxidation of phosphine with a mild oxidizer: PH3 + SO2 = H + S↓ (catalysts are mercury and traces of water).

Phosphorus halides PX3 and PX5. All phosphorus halides are known except PJ5. In the case of phosphorus (III), these are pyramidal molecules with a phosphorus atom at the top and with angles between P-X bonds equal to 100°. Phosphorus(V) halides are trigonal bipyramids with sp3d hybridization of phosphorus atomic orbitals. Both phosphorus fluorides under normal conditions are gases, PCl3 and PBr3 are liquids, and triiodide, pentachloride and pentabromide are solids. The last two compounds are salts with complex ions PCl5: +-, PBr5: +Br-. When heated, both compounds cleave off a halogen molecule and turn into a trihalogenide. Phosphorus halides are obtained by direct synthesis. Only PF3 - indirectly: PCl3 + AsF3 = PF3 + AsCl3. All phosphorus halides are subject to hydrolysis, and trihalides are also capable of oxidation: 2PCl3 + O2 = 2POCl3 - phosphorus oxychloride, can also be obtained by other reactions: PCl3 + 2CrO3 = POCl3 + Cr2O3↓ + O2, 6PCl5 + P4O10 = 10POCl3. Trihalides also add sulfur : PCl3 + S = PSCl3. In non-aqueous solutions, reactions are possible: KF + PF5 = K HF (liquid) + PF5 = H - hexafluorophosphoric acid, stable only in aqueous solution, comparable in strength to perchloric acid.

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Phosphine. Phosphorus oxides and phosphoric acids: properties, preparation.

Phosphine word

Medico-biological significance of phosphorus.

Phosphine (hydrogen phosphorous, phosphorus hydride, according to the IUPAC nomenclature - phosphane PH3) is a colorless, very toxic, rather unstable gas (under normal conditions) with a specific smell of rotten fish.

Physical properties

colorless gas. Poorly soluble in water, does not react with it. At low temperatures it forms a solid clathrate 8РН3·46Н2О. Soluble in benzene, diethyl ether, carbon disulfide. At −133.8 °C, it forms crystals with a face-centered cubic lattice.

The phosphine molecule has the shape of a trigonal pyramid with C3v molecular symmetry (dPH = 0.142 nm, HPH = 93.5o). The dipole moment is 0.58 D, significantly lower than that of ammonia. Hydrogen bonding between PH3 molecules is practically not manifested and therefore phosphine has lower melting and boiling points.

]Receive

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:

Phosphonic acid decomposition:

or restore it:

Chemical properties

Phosphine is very different from its ammonia counterpart. 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 (3s2) is lower than that of nitrogen (2s2) 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 PH4+ ion (similar to ammonium). Phosphonium salts, colorless crystalline substances, are extremely unstable, easily hydrolyzed.

Phosphine salts, like phosphine itself, are strong reducing agents.

Toxicity

Phosphine is highly toxic, acts on the nervous system, disrupts metabolism. MAC = 0.1 mg/m³. The smell is felt at a concentration of 2-4 mg / m³, prolonged inhalation at a concentration of 10 mg / m³ is fatal. In human blood, the content of phosphine is not more than 0.001 mg/m³.

The following phosphorus oxides are known:

Phosphorus(III) oxide - a binary inorganic compound, phosphorus oxide with the formula P4O6, white flakes or crystals with an unpleasant odor, react with water.

Receipt

• Careful oxidation of white phosphorus with nitrous oxide or carbon dioxide:

• Reverse disproportionation of phosphorus(V) oxide and white phosphorus:

[edit] Physical properties

Phosphorus(III) oxide forms white flakes or crystals with an unpleasant odor.

It dissolves well in organic solvents (benzene, carbon disulfide).

Unstable in the light, first turns yellow, and then reddens.

Properties

P4O10 interacts very actively with water (the H-form absorbs water even with an explosion), forming mixtures of phosphoric acids, the composition of which depends on the amount of water and other conditions:

It is also capable of extracting water from other compounds, making it a powerful dehydrator:

Phosphorus(V) oxide is widely used in organic synthesis. It reacts with amides, converting them to nitriles:

Carboxylic acids are converted to the corresponding anhydrides:

Phosphorus(V) oxide also interacts with alcohols, ethers, phenols and other organic compounds. In this case, the P-O-P bonds are broken and organophosphorus compounds are formed. Reacts with NH3 and hydrogen halides to form ammonium phosphates and phosphorus oxyhalides:

When P4O10 is fused with basic oxides, it forms various solid phosphates, the nature of which depends on the reaction conditions.

Receipt

Phosphorus(V) oxide is obtained by burning phosphorus. The technological process takes place in the combustion chamber and includes the oxidation of elemental P with pre-dried air, the precipitation of P4O10 and the purification of exhaust gases. The resulting pentoxide is purified by sublimation.

The technical product has the appearance of a white snow-like mass, consisting of a mixture of different forms of P4O10.

Application

P4O10 is used as a dryer for gases and liquids. It is also an intermediate product in the thermal production of phosphoric acid H3PO4.

It is widely used in organic synthesis in dehydration and condensation reactions.

The value of phosphorus

• phosphorus is included nucleic acids, which take part in the processes of growth, cell division, storage and use of genetic information
• phosphorus is found in the bones of the skeleton (about 85% of the total amount of phosphorus in the body)
• phosphorus is essential for the normal structure of teeth and gums
• ensures the proper functioning of the heart and kidneys
• phosphorus is involved in the processes of accumulation and release of energy in cells
• involved in the transmission of nerve impulses
• helps the metabolism of fats and starches.

The inorganic element phosphorus, P, is found in the human body in the form of phosphorus compounds - inorganic phosphates and lipids or nucleotides.

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Physical properties

Phosphorus P has several allotropic modifications: white, red, black.

Obtaining phosphorus P

Free phosphorus P obtained from natural calcium phosphate by heating it with sand ( SiO2) and coal in an electric furnace at high temperature:

Chemical properties of phosphorus - P

White phosphorus more reactive than red.

Beware of Phosphine!

It oxidizes easily and ignites spontaneously in air.

When oxidized, white phosphorus glows in the dark, chemical energy is converted into light energy.

Phosphorus compounds P with metals are called phosphides. They are easily decomposed by water to form gas. phosphine (PH3).

Phosphine - PH3

4. With a large excess of chlorine, phosphorus pentachloride is formed:

Oxides and acids of phosphorus

Phosphorus forms with oxygen three oxides :

P2O3 - phosphorous anhydride - phosphorus oxide (SH);

P2O5 - phosphoric anhydride - phosphorus (V) oxide;

(P2O4 is phosphorus tetroxide).

P2O3 obtained by slow oxidation of phosphorus (with a lack of oxygen):

When exposed to cold water, it forms phosphorous acid – H3PO3.

P2O5 is formed during the combustion of phosphorus in air (with an excess of oxygen):

acids

Phosphoric anhydride P2O5, depending on the temperature, can attach a different amount of water, forming acids of various composition:

Of greatest importance is ortho phosphoric acid -H3PO4.

It can be obtained in the following way:

1. Boiling metaphosphoric acid:

2. Oxidation of red phosphorus:

3. The action of sulfuric acid on calcium phosphate:

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Chemistry Tutor

Continuation. See in No. 22/2005; 1, 2, 3, 5, 6, 8, 9, 11, 13, 15, 16, 18, 22/2006;
3, 4, 7, 10, 11, 21/2007;
2, 7, 11, 18, 19, 21/2008;
1, 3, 10, 11/2009

ACTIVITY 30

10th grade(first year of study)

Phosphorus and its compounds

1. Position in the table of D.I. Mendeleev, the structure of the atom.

2. Short story discoveries and the origin of the name.

3. Physical properties.

4. Chemical properties.

5. Being in nature.

6. Main methods of obtaining

7. The most important compounds of phosphorus.

Phosphorus is in the main subgroup of group V periodic system D.I. Mendeleev. Its electronic formula is 1 s 2 2s 2 p 6 3s 2 p 3 is R-element. Characteristic oxidation states of phosphorus in compounds –3, +3, +5; the most stable is the oxidation state +5. In compounds, phosphorus can be included both in the composition of cations and in the composition of anions, for example:

Phosphorus got its name from the property of white phosphorus to glow in the dark. The Greek word translates as "bringing light." Phosphorus owes this name to its discoverer - the alchemist Brand, who, fascinated by the glow of white phosphorus, came to the conclusion that he had received the philosopher's stone.

Phosphorus can exist in the form of several allotropic modifications, the most stable of which are white, red and black phosphorus.

Molecule white phosphorus (the most active allotrope) has a molecular crystal lattice, in the nodes of which there are four-atomic P 4 molecules of a tetrahedral structure.

White phosphorus is soft, like wax, melts and boils without decomposition, has a garlic smell. In air, white phosphorus is rapidly oxidized (glows greenish), self-ignition of finely dispersed white phosphorus is possible. It is insoluble in water (stored under a layer of water), but readily soluble in organic solvents. Poisonous (even in small doses, MPC = 0.03 mg / m 3). It has a very high chemical activity. When heated without air access to 250-300 ° C, it turns into red phosphorus.

red phosphorus is an inorganic polymer; macromolecules P n can have both cyclic and acyclic structures. It differs sharply from white phosphorus in its properties: it is not poisonous, does not glow in the dark, does not dissolve in carbon disulfide and other organic solvents, and does not have high chemical activity. At room temperature, it slowly turns into white phosphorus; when heated to 200 ° C under pressure, it turns into black phosphorus.

black phosphorus looks like graphite. By structure, it is an inorganic polymer, the molecules of which have a layered structure. Semiconductor. Not poisonous. The chemical activity is much lower than that of white phosphorus. Air resistant. When heated, it turns into red phosphorus.

Chemical properties

The most active chemically is white phosphorus (but in practice they prefer to work with red phosphorus). It can exhibit the properties of both an oxidizing agent and a reducing agent in reactions, for example:

4P + 3O 2 2P 2 O 3,

4P + 5O 2 2P 2 O 5.

Metals (+/-)*:

3Ca + 2P Ca 3 P 2 ,

3Na + P Na 3 P,

Cu + P does not react.

Nonmetals (+):

2P + 3I 2PI 3,

6P + 5N 2 2P 2 N 5 .

Basic oxides (-).

Acid oxides (-).

Alkalis (+):

Acids (not oxidizing agents) (-).

Oxidizing acids (+):

3P (cr.) + 5HNO 3 (razb.) + 2H 2 O \u003d 3H 3 PO 4 + 5NO,

P (cr.) + 5HNO 3 (conc.) H 3 PO 4 + 5NO 2 + H 2 O,

2P (cr.) + H 2 SO 4 (conc.) 2H 3 PO 4 + 5SO 2 + 2H 2 O.

Salts (-)**.

In nature, phosphorus occurs in the form of compounds (salts), the most important of which are phosphorite (Ca 3 (PO 4) 2), chlorapatite (Ca 3 (PO 4) 2 CaCl 2) and fluorapatite (Ca 3 ( PO 4) 2 CaF 2). Calcium phosphate is found in the bones of all vertebrates, causing their strength.

Phosphorus is obtained in electric furnaces by fusing calcium phosphate, sand and coal without air access:

Ca 3 (PO 4) 2 + 3SiO 2 + 5C 2P + 5CO + 3CaSiO 3.

The most important phosphorus compounds are: phosphine, phosphorus(III) oxide, phosphorus(V) oxide, phosphoric acids.

F o s f i n

This hydrogen compound of phosphorus, a colorless gas with a garlic-fish odor, is highly toxic. Let's badly dissolve in water, but we will well dissolve in organic solvents. Much less stable than ammonia, but a stronger reducing agent. Has no practical value.

To obtain phosphine, a direct synthesis reaction from simple substances is usually not used; The most common way to obtain phosphine is the hydrolysis of phosphides:

Ca 3 P 2 + 6HOH \u003d 3Ca (OH) 2 + 2PH 3.

In addition, phosphine can be obtained by a disproportionation reaction between phosphorus and alkali solutions:

4P + 3KOH + 3H 2 O PH 3 + KPO 2 H 2,

or from phosphonium salts:

PH 4 I PH 3 + HI,

PH 4 I + NaOH PH 3 + NaI + H 2 O.

It is advisable to consider the chemical properties of phosphine from two sides.

Acid-base properties. Phosphine forms an unstable hydrate with water, which exhibits very weak basic properties:

PH 3 + H 2 O PH 3 H 2 O (PH 4 OH),

PH 3 + HCl PH 4 Cl,

2PH 3 + H 2 SO 4 (PH 4) 2 SO 4.

redox properties. Phosphine is a strong reducing agent:

2PH 3 + 4O 2 P 2 O 5 + 3H 2 O,

PH 3 + 8AgNO 3 + 4H 2 O \u003d H 3 PO 4 + 8Ag + 8HNO 3.

O x i d f o s f o r a (III)

Oxide P 2 O 3 (true formula - P 4 O 6) is a white crystalline substance, a typical acid oxide. When interacting with water in the cold, it forms phosphorous acid (medium strength):

P 2 O 3 + 3H 2 O \u003d 2H 3 PO 3

Since phosphorous acid is dibasic, the interaction of phosphorus trioxide with alkalis forms two types of salts - hydrophosphites and dihydrophosphites.

For example:

P 2 O 3 + 4NaOH \u003d 2Na 2 HPO 3 + H 2 O,

P 2 O 3 + 2NaOH + H 2 O \u003d 2NaH 2 PO 3.

Phosphorus dioxide P 2 O 3 is oxidized by atmospheric oxygen to pentoxide:

P 2 O 3 + O 2 P 2 O 5 .

Phosphorus trioxide and phosphorous acid are fairly strong reducing agents. Phosphorus(III) oxide is obtained by slow oxidation of phosphorus in the absence of oxygen:

4P + 3O 2 2P 2 O 3 .

Phos phora(V) oxide and phos phoric acids

Phosphorus pentoxide P 2 O 5 (true formula - P 4 O 10) is a white hygroscopic crystalline substance. In the solid and gaseous states, the molecule exists in the form of a dimer, and at high temperatures it monomerizes. A typical acidic oxide. It is very soluble in water, forming a number of phosphoric acids:

metaphosphoric:

P 2 O 5 + H 2 O \u003d 2HPO 3

pyrophosphoric (diphosphoric):

P 2 O 5 + 2H 2 O \u003d H 4 P 2 O 7

orthophosphoric (phosphoric):

P 2 O 5 + 3H 2 O \u003d 2H 3 PO 4

Phosphorus pentoxide exhibits all the properties characteristic of acidic oxides, for example:

P 2 O 5 + 3H 2 O \u003d 2H 3 PO 4,

P 2 O 5 + 3CaO 2Ca 3 (PO 4) 2;

can form three types of salts:

Oxidizing properties are not typical for it, because. +5 oxidation state is very stable for phosphorus. Phosphorus pentoxide is obtained by burning phosphorus in a sufficient amount of oxygen:

4P + 5O 2 2P 2 O 5 .

Orthophosphoric acid H 3 RO 4 is a colorless crystalline substance, very soluble in water, hygroscopic. It is a tribasic acid of medium strength; does not have pronounced oxidizing properties. Shows all the chemical properties characteristic of acids, forms three types of salts (phosphates, hydrophosphates and dihydrophosphates):

2H 3 PO 4 + 3Ca = Ca 3 (PO 4) 2 + 3H 2,

H 3 PO 4 + Cu,

2H 3 PO 4 + 3CaO = Ca 3 (PO 4) 2 + 3H 2 O,

2H 3 PO 4 + K 2 CO 3 \u003d 2KH 2 PO 4 + CO 2 + H 2 O.

In industry, phosphoric acid is obtained by extraction:

Ca 3 (PO 4) 2 + 3H 2 SO 4 \u003d 2H 3 PO 4 + 3CaSO 4,

as well as thermal method:

Ca 3 (PO 4) 2 + 3SiO 2 + 5C 3СaSiO 3 + 2P + 5CO,

4P + 5O 2 2P 2 O 5,

P 2 O 5 + 3H 2 O \u003d 2H 3 PO 4.

Laboratory methods for obtaining phosphoric acid include the action of dilute nitric acid on phosphorus:

3P (cr.) + 5HNO 3 (razb.) + 2H 2 O \u003d 3H 3 PO 4 + 5NO,

interaction of metaphosphoric acid with water when heated:

HPO 3 + H 2 O H 3 PO 4 .

In the human body, orthophosphoric acid is formed by the hydrolysis of adenosine triphosphate (ATP):

ATP ADP + H 3 PO 4.

Qualitative reaction to phosphate ion is the reaction with the silver cation; a yellow precipitate is formed, insoluble in slightly acidic media:

3Ag + + \u003d Ag 3 PO 4,

3AgNO 3 + K 3 PO 4 = Ag 3 PO 4 + 3KNO 3.

In addition to the above phosphoric acids (containing phosphorus in the +5 oxidation state), many other oxygen-containing acids are known for phosphorus. Here are some of the most important representatives.

Phosphorous(HPO 2 H 2) is a monobasic acid of medium strength. Its second name is phosphine:

Salts of this acid are called hypophosphites, or phosphites, for example KPO 2 H 2 .

Phosphorous(H 3 RO 3) - dibasic acid of medium strength, slightly weaker than hypophosphorous. It also has a second name - phosphonic:

Its salts are called phosphites, or phosphonates, for example K 2 PO 3 H.

Diphosphoric (pyrophosphoric)(H 4 P 2 O 7) - a tetrabasic acid of medium strength, slightly stronger than orthophosphoric:

Salts are diphosphates, for example K 4 P 2 O 7.

Test on the topic "Phosphorus and its compounds"

1. Eliminate the "extra" element from those listed according to the principle of the possibility of forming allotropic modifications:

a) oxygen; b) nitrogen;

c) phosphorus; d) sulfur.

2. When interacting 42.6 g of phosphoric anhydride and 400 g of a 15% sodium hydroxide solution, the following is formed:

a) sodium phosphate;

b) sodium hydrogen phosphate;

c) a mixture of phosphate and sodium hydrogen phosphate;

d) a mixture of sodium hydro- and dihydrogen phosphate.

3. The sum of the coefficients in the equation electrolytic dissociation potassium phosphate is:

a) 5; b) 3; at 4; d) 8.

4. The number of electrons in the outer level of the phosphorus atom:

a) 2; b) 3; at 5; d) 15.

5. Phosphorus, obtained from 33 g of technical calcium phosphate, was burned in oxygen. The formed phosphorus(V) oxide reacted with 200 ml of 10% sodium hydroxide solution (density 1.2 g/ml) to form a medium salt. The mass of impurities in the technical sample of calcium phosphate (in g) is:

a) 3.5; b) 1.5; in 2; d) 4.8.

6. The number of -bonds in a molecule of pyrophosphoric acid:

a) 2; b) 12; c) 14; d) 10.

7. The number of hydrogen atoms contained in 4.48 L (N.O.) of phosphine is:

a) 1.2 10 23; b) 0.6 10 23;

c) 6.02 10 23; d) 3.6 10 23 .

8. At a temperature of 30 ° C, a certain reaction proceeds in 15 s, and at 0 ° C - in 2 minutes. Van't Hoff coefficient for this reaction:

a) 2.4; b) 2; c) 1.8; d) 3.

9. Orthophosphoric acid can react with the following substances:

a) copper(II) oxide; b) potassium hydroxide;

c) nitric acid; d) zinc.

10. The sum of the coefficients in the reaction between phosphorus and Bertolet's salt is:

a) 9; b) 6; c) 19; d) such a reaction is impossible.

Key to the test

1	2	3	4	5	6	7	8	9	10
b	in	a	in	in	b	G	b	a, b, d	in

Tasks and exercises for phosphorus and its compounds

Chains of rotations:

1. Phosphorus -> phosphorus pentoxide -> phosphoric acid -> calcium phosphate ® phosphoric acid.

2. Calcium phosphate -> phosphorus -> calcium phosphide -> phosphine -> phosphorus pentoxide -> phosphoric acid -> calcium dihydrogen phosphate.

3. Calcium phosphate -> A -> B -> C -> D -> E -> calcium phosphate. All substances contain phosphorus, in the scheme there are three OVRs in a row.

4. Phosphorus -> phosphorus pentoxide -> calcium phosphate -> phosphorus -> phosphine -> phosphoric acid -> calcium dihydrogen phosphate.

5. Calcium phosphide (+ hydrochloric acid solution) -> A (+ oxygen) -> B (+ sodium hydroxide, deficiency) -> C (+ sodium hydroxide, excess) -> D (+ calcium hydroxide) -> E.

A level

1. With complete combustion of 6.8 g of the substance, 14.2 g of phosphorus pentoxide and 5.4 g of water were obtained. 37 ml of 32% sodium hydroxide solution (density 1.35 g/ml) was added to the reaction products obtained. Set the formula of the starting substance and determine the concentration of the resulting solution.

Solution

Reaction equation:

(P 2 O 5) = 0.1 mol, (H 2 O) = 0.3 mol.

(P) = 0.2 mol, (H) = 0.6 mol.

m(P) = 6.2 g, m(H) = 0.6 g.

m= 6.8 g.

(P): (H) = 0.2: 0.6 = 1: 3.

Therefore, the formula of the starting substance is PH 3, and the reaction equation:

then phosphoric acid is formed:

(H 3 PO 4) \u003d 2 (P 2 O 5) \u003d 0.2 mol.

With alkali, phosphoric acid can react as follows:

Let us determine the amount of substance NaOH according to the condition of the problem:

(H 3 PO 4): (NaOH) \u003d 0.2: 0.4 \u003d 1: 2,

so reaction 2 takes place.

(Na 2 HPO 4) \u003d (H 3 PO 4) \u003d 0.2 mol;

m(Na2HPO4) = M(Na 2 HPO 4) (Na 2 HPO 4) = 142 0.2 = 28.4 g;

m(r-ra) = m(P 2 O 5) + m(H 2 O) + m(p-ra NaOH) \u003d 14.2 + 5.4 + 37 1.35 \u003d 69.55 g.

(Na2HPO4) = m(Na2HPO4)/ m(solution) = 28.4 / 69.55 = 0.4083, or 40.83%.

Answer. PH 3 ; (Na 2 HPO 4) = 40.83%.

2. With complete electrolysis of 1 kg of iron(II) sulfate solution, 56 g of metal was released on the cathode. What mass of phosphorus can react with the substance released at the anode, and what will be the composition of the salt if the resulting reaction product is dissolved in 87.24 ml of a 28% sodium hydroxide solution (solution density 1.31 g / ml)?

Answer. 12.4 g phosphorus; sodium hydrogen phosphate.

3. 20 g of a mixture of barium sulfate, calcium phosphate, calcium carbonate and sodium phosphate was dissolved in water. The mass of the insoluble part was 18 g. Under the action of hydrochloric acid on it, 2.24 l of gas (n.o.) was released and the mass of the insoluble residue was 3 g. Determine the composition of the initial mixture of salts by mass.

Answer. Na 3 PO 4 - 2 g; BaCO 3 - 3 g;
CaCO 3 - 10 g; Ca 3 (PO 4) 3 - 5 g.

4. How many kg of phosphorus can be obtained from 1 ton of phosphorite containing 40% impurities? What is the volume at n.o. take phosphine derived from this phosphorus?

Answer. 120 kg P; 86.7 m 3 PH 3 .

5. 40 g of a mineral containing 77.5% calcium phosphate was mixed with an excess of sand and coal and heated without air in an electric furnace. Received simple matter dissolved in 140 g of 90% nitric acid. Determine the mass of sodium hydroxide required to completely neutralize the oxidation product of a simple substance.

Answer. 24 g NaOH.

Level B

1. To completely neutralize the solution obtained by hydrolysis of 1.23 g of some phosphorus halide, 35 ml of a 2M potassium hydroxide solution were required. Determine the formula for the halide.

Answer. Phosphorus trifluoride.

2. A sample of anhydrous ethanol containing 0.5% phosphorus(V) oxide as an impurity was burned in sufficient oxygen. The resulting gases were separated, and the resulting solution was heated until the evolution of gas ceased, after which a 0.5% potassium hydroxide solution equal in mass was added to it. Determine the mass fractions of substances in the resulting solution.

Answer. K 2 HPO 4 - 0.261%;
KH 2 PO 4 - 0.204%.

3. To 2 g of a mixture of hydrophosphate and potassium dihydrogen phosphate, in which the mass fraction of phosphorus is 20%, was added 20 g of a 2% solution of phosphoric acid. Calculate the mass fractions of substances in the resulting solution.

Answer. KH 2 PO 4 - 9.03%;
K 2 HPO 4 (remaining) - 1.87%.

4. When a mixture of hydride and phosphide of an alkali metal with equal mass fractions was treated with water, a gas mixture with a nitrogen density of 0.2926. Determine which metal was included in the compounds.

Answer. Sodium.

5. 50 g of a mixture of calcium phosphate and calcium and ammonium carbonates was calcined, resulting in 25.2 g of a solid residue, to which water was added, and then an excess of carbon dioxide was passed through. The mass of the undissolved residue was 14 g. Determine the mass of ammonium carbonate in the initial mixture.

Solution

When the mixture is calcined, the following processes take place:

1) Ca 3 (PO 4) 2;

2)

3) (NH 4) 2 CO 3 2NH 3 + CO 2 + H 2 O.

In the solid residue - Ca 3 (PO 4) 2 and CaO.

After adding water:

4) Ca 3 (PO 4) 2 + H 2 O;

5) CaO + H 2 O \u003d Ca (OH) 2.

After passing carbon dioxide:

6) Ca (OH) 2 + H 2 O + CO 2 \u003d Ca (HCO 3) 2.

The undissolved residue is Ca 3 (PO 4) 2, therefore, m(Ca 3 (PO 4) 2) = 14 g.

Find the mass of CaO:

m(CaO) \u003d 25.2 - 14 \u003d 11.2 g.

(CaO) \u003d 11.2 / 56 \u003d 0.2 mol,

(CaCO 3) \u003d (CaO) \u003d 0.2 mol,

m(CaCO 3) \u003d 0.2 100 \u003d 20 g.

m(NH 4) 2 CO 3 = m(mixes) - m(Ca 3 (PO 4) 2) - m(CaCO 3) \u003d 50 - 14 - 20 \u003d 16 g.

Answer. m(NH 4) 2 CO 3 \u003d 16 g.

Qualitative tasks

1. Solid, white, highly water-soluble compound A is an acid. When oxide B is added to an aqueous solution A, a white, water-insoluble compound C is formed. As a result of calcining substance C at a high temperature in the presence of sand and coal, a simple substance is formed that is part of A. Identify the substances, write reaction equations.

Answer. Substances: A - H 2 PO 4, B - CaO,
C - Ca 3 (PO 4) 2 .

2. A mixture of two red solids (A) and white color(B) ignites with little friction. The reaction produces two white solids, one of which (C) dissolves in water to form an acidic solution. If calcium oxide is added to substance C, a white, water-insoluble compound is formed. Identify substances, write reaction equations.

Answer. Substances: A - P (cr.), B - KClO 3,
C - P 2 O 5.

3. The water-insoluble compound A of white color, as a result of calcination at high temperature with coal and sand in the absence of oxygen, forms a simple substance B, which exists in several allotropic modifications. When substance B is burned, compound C is formed, which dissolves in water to form acid E, which is capable of forming three types of salts. Identify substances, write reaction equations.

Answer. Substances: A - Ca 3 (PO 4) 2, B - P,
C - P 2 O 5, E - H 3 PO 4.

* The +/– sign means that this reaction does not proceed with all reagents or under specific conditions.

** Of interest is the redox reaction (ORD) that occurs when the matches are ignited:

To be continued

The story about gaseous compounds of phosphorus, and first of all about phosphine, should probably be started with the words: “the flickering light that appears in the swamps (the famous “wandering lights”) is the result of spontaneous ignition of phosphine.” Well, the following definition is already of an encyclopedic sense: “phosphine, or hydrogen phosphide (PH 3) is a colorless gas with an unpleasant odor (rotting fish, garlic or industrial carbide), poisonous, formed during the biochemical reduction of phosphoric acid esters, mainly under anaerobic conditions , i.e. without access to oxygen.

Phosphorus compounds in nature

There are many other gaseous organophosphorus compounds in nature, in the molecules of which the phosphorus atom P is connected to the carbon atom C. There are thousands of them. Many of them are part of ecosystems, including living cells of plants and microorganisms. The largest group of compounds with C-P bonds was discovered about fifty years ago in living objects.

There are also phosphonates in soils - derivatives of organophosphorus compounds with preserved C-P bonds. True, they are few, no more than 1-2% of the phosphorus contained in organic matter, therefore they can not always be detected on arable land, but in swampy soils and meadows their content rises to 3-4%.

Under normal (aerobic) conditions, natural compounds of organic and mineral phosphorus are phosphates (orthophosphates). There are a great many of them. For organic phosphates, C-O-R connection, in other words, carbon and phosphorus are connected through an oxygen atom.

One of the amazing mysteries of nature is that organic phosphates in living systems (for example, in algae and microorganisms) are synthesized and decomposed not arbitrarily, but according to the “golden section” rule, obeying a certain law described by the famous series of Fibonacci numbers (1, 1 , 2, 3, 5, 8...), in which each next term is equal to the sum of the two previous ones. The harmony of nature is incomprehensibly manifested here in the accumulation and consumption of energy and matter (in particular, phosphorus) in ecosystems, described by a ratio that is approximately given by the classical “golden section” coefficient of 1.618 (5/3, 8/5, 13/8, etc.). etc.), i.e., 62% of the mentioned compounds should be bound and accumulated, and only 38% should be destroyed or volatilized. These patterns subsequently affect the accumulation of humus, and the cycle of phosphorus and nitrogen, and gaseous flows determined by the emissions and "sinks" of carbon dioxide CO 2, and the "respiration" of the soil (the release of CO 2 and the assimilation of oxygen O 2). In fact, in nature there are fluctuations in the numerical values ​​of this ratio within 1.3-1.7. But, as noted more than once in the writings of the author and other scientists, it turns out to be much more terrible that main reason deviations and even violations of this pattern has become anthropogenic activity.

Some experts have already drawn attention to the fact that new dangers may lie in wait for us if this ratio tends to unity, i.e., accumulation and decomposition proceed with the same intensity, as happens, for example, in the carbon cycle, where due to "intervention" of the global economy, the ocean and biosphere now absorb only half of carbon emissions (62% should be).

But let us return to phosphine and its derivatives, in other words, to those organophosphorus compounds in which various elements (nitrogen, sulfur, silicon, molybdenum, etc.) and their complexes are found together with phosphorus and carbon. Under favorable conditions for the growth of microorganisms (in particular, in the conditions of swamps and tundra during the observed warming), organophosphorus compounds decompose with the help of the enzyme (catalyst) C-P-lyase. Now it is found in 9 groups of bacteria that feed on phosphorus, extracting it from the breakdown of organophosphorus compounds. But fungi and yeast, which account for 50-70% of the total microflora in ecosystems, do not break down these compounds. On the contrary, protozoa, mollusks and fungi synthesize them. Mushrooms can grow even at fairly high concentrations of phosphine, only their mycelium turns yellow.

Application, properties, dangers

Phosphine is poisonous (a dangerous concentration that can lead to death is 0.05 mg / l), and at a concentration of 2000 ml / m 3 (2 l / m 3, or 2 10 -3) it causes instant death. It is encountered primarily in agriculture during the disinfection of granaries and protection from ticks and other pests during the transportation of crops, especially grain crops. Previously, it was actively used against rats and mice in barns. In Australia, they resort to his help even in the fight against overly rapidly breeding rabbits. In addition, a number of herbicides and insecticides contain organophosphorus compounds based on phosphine and its derivatives. And, finally, in recent times it has been increasingly necessary to deal with it in connection with the large-scale destruction chemical weapons, providing for the neutralization of poisonous organophosphorus compounds of sarin and soman - phosphine derivatives.

Pure phosphine (without impurities) ignites at a temperature of 150 ° C, burns out with the formation of toxic phosphoric acid, but in the presence of impurities of diphosphine P 2 H 4 or gaseous phosphorus P 4 it can spontaneously ignite in air. The reaction of phosphine with oxygen (as well as the oxidation of similar methane - CH 4 and silane - SiH 4) refers to branched chain reactions. chemical reactions, i.e., it flows faster and can lead to an explosion. Phosphine oxidation occurs at room temperature, but the gas can be stable at low temperature. The oxidation of phosphine can be accelerated by irradiating it with ultraviolet light. Its self-ignition in air is possible at concentrations of 1.7-1.9% (17-19 l / m 3), or 26-27 g / m 3. So, in marsh ecosystems, one often has to deal not only with the mentioned “stray fires”, but also with spontaneous combustion (by the way, widespread peat fires are of the same nature).

For fumigation (to rid storages of grain and agricultural products from mites and other pests), phosphides are usually used, in particular, phosphorus compounds with metals. Reacting with air moisture, phosphides release phosphine. Tablets and tapes containing phosphides are laid out in storage facilities at the rate of 9 g/t of grain or other products subject to long-term storage, they are even added to apples. Phosphine is believed to volatilize when aerated, although according to the data available in the scientific literature, up to 13% of the poisonous gas is absorbed in feed grain. Shouldn't this circumstance alone make one treat such "disinfection" with extreme caution?!

Now, for the fumigation of grain during transportation and storage, two compounds are allowed for use - methylbromine and methylphosphine, and the first is an order of magnitude less toxic (and effective) than the second. Using the latter, it is tacitly assumed that the poisonous phosphine, after being absorbed by the contents of the vault, is miraculously extracted and volatilized, poisoning only ticks and other pests. It seems that earlier it was not customary to think about how this picture corresponds to reality. Meanwhile, almost half a century ago, it was found that methylphosphine (a mixture of two gases - methane CH 4 and phosphine PH 3) is extremely toxic, almost like phosphine itself.

Methane and phosphine in the biosphere

It is no secret that methane emitted from swamps is considered one of the main greenhouse gases and remains the subject of active discussion and research in connection with the problems of global climate change. Alas, in Russia its concentration in the atmosphere is determined only at one weather station (Teriberka on the Kola Peninsula). But it would not hurt to measure it over the Siberian swamps!

As is known, huge reserves of methane (7·10 11 -3·10 13 tons) have been conserved in the depths of the earth, and 4·10 11 tons of them are in the Arctic permafrost zone. On land, methane is found in organic compounds of swamps, sediments and detritus, and in the World Ocean - in gas hydrates occurring under the bottom, under conditions of low temperatures. In the UN Climate Change Report, experts report that in Siberia, the release of methane from swamps and permafrost in last years is growing rapidly. The maximum emission of methane from tundra soils is reached at 8-10°C, and at 5°C its oxidation to CO 2 and water prevails. It is formed in all soil horizons. As a result of recent studies, it turned out that, for example, our southern shrub tundra (near Vorkuta) served as a carbon sink only two of the last five years.

This is a rather dangerous trend, especially if we take into account that our country accounts for 2/3 of all swamps on Earth. Our areas of wetlands exceed the area of ​​all agricultural land: according to 2003 data, 343 million hectares of swamps (of which 130 million hectares are not overgrown with forests) and 221 million hectares of agricultural land (of which 123 million hectares are arable land).

And here is how the employees of Moscow State University assessed the release of methane in 2007 based on the results of measurements in swamps in the Tomsk region. According to their estimates, the average value of the methane flux was about 10 mg/m 2 per hour. In summer, 2.4 kg/ha can be released per day, and 432 kg/ha per season (6 months). And from 130 million hectares of swamps - almost 60 million tons. The oxidation of such an amount of methane will require twice as much oxygen - 120 million tons.

The main “side” effect of methane emission should be recognized as the fact that in tundra and marsh ecosystems at low temperatures, methane not only represents a fair amount of carbon that can significantly change its content in the atmosphere, but is also closely associated with organophosphorus compounds, which are invariably present. in plants, microflora of swamps and sediments (mainly due to the mentioned C-P connection). And its isolation from those places where it was previously synthesized, due to the intensification of biochemical fermentation processes with increasing temperature, occurs not least due to the decomposition of phosphine-based compounds. In other words, CH 4 and PH 3 gases are emitted in parallel. Meanwhile, while environmentalists and climatologists are only monitoring changes in the content of CO 2 and CH 4 in the atmosphere, and the content of PH 3 is not taken into account by anyone. But in vain!

This omission is due, in part, to the fact that only a few experts are aware of methods for measuring the content of phosphorus in the atmosphere in the gaseous state. After all, even in scientific world there is still an opinion that phosphorus in nature exists mainly in the form of phosphates and after hydrolysis P-O-R connections, P-O-C and even P-C turns into a solid. Fluxes of phosphorus into the atmosphere in the form of volatile compounds of the PH 3 type are considered negligible and are neglected. Determining the content of phosphorus released into the atmosphere with phosphine using only the usual methods used to detect phosphorus in solid compounds significantly distorts the real picture of the phosphorus cycle in ecosystems. At the same time, the appearance of poisonous and spontaneously combustible phosphine in the atmosphere is ignored.

Phosphine Threat: Simple Estimates

Meanwhile, the simplest quantitative assessment of the release of phosphine in ecosystems can be obtained by studying areas flooded with water, simulating water meadows or rice paddies. As was established in the Moscow Agricultural Academy, held back in 1926. K. A. Timiryazev, a series of six experiments carried out under strictly controlled conditions, 9.7 mg of phosphorus from 1 kg of soil per hour passes into the gaseous form (phosphine). A not too complicated calculation gives 2.13 kg/ha per day. But this is almost as much as methane is released from swamps! Therefore, for the season we get 383 kg/ha, and from the entire area of ​​treeless swamps (130 million hectares) - about 50 million tons of PH 3 . On its oxidation to phosphoric acid according to the formula

PH 3 + 2O 2 → H 3 PO 4

it is easy to see that twice as much oxygen will be required - almost 100 million tons (for methane, these values ​​were 60 and 120 million tons, respectively).

An indirect confirmation of the release of phosphine from soils is the study of phosphorus fluxes in rice paddies - from planting to harvesting, the loss of phosphorus in flooded soils is 3-8 times higher than its content in grain and straw. The maximum removal of Р 2 O 5 reaches 100 kg/ha. Organic phosphorus compounds are excreted from soils 4 times more than stored in plants. The total loss of phosphorus from the upper (20 cm) soil layer, according to various estimates, is 960-2940 kg/ha. There is evidence that when rice is grown on flooded checks for 32 years, more than half of the humus is lost from the soil, and with it, of course, nitrogen and phosphorus are carried out.

This can also occur due to the release of their gaseous forms - ammonia (NH 3) and phosphine (PH 3). It has long been known that, in terms of chemical properties, they are chemical structural analogues. I repeat, the determination of phosphorus and nitrogen only in the mineral form, ignoring the gas components does not reflect the true processes in ecosystems, especially under anaerobic conditions. In particular, direct confirmation that phosphorus is released along with methane in swamp ecosystems has been obtained in recent studies.

Returning to discussions about the possible underestimation of the phosphine content in the atmosphere, it should be noted that not only the swamps of the North or the tropics, but also extensive rice plantations (primarily in India, China, Japan and the countries of Southeast Asia) can make a quite tangible contribution.

In the scientific literature there is evidence that up to 3.5 kg/ha of phosphorus falls on the ground with precipitation. In other words, this is only about 1% of the phosphorus that is estimated to be removed from swamp systems or flooded soils by phosphine to the atmosphere (383 kg/ha), the remaining 99% seems to be rapidly oxidized, precipitated or decomposed (for example, in as a result of hydrolysis) in the surface layers of air, the lithosphere and the biosphere, ensuring the redistribution of phosphorus on the surface of the earth.

Of course, phosphine, like methane, is in the atmosphere, but it must be admitted that the phosphorus cycle has been studied much worse than the nitrogen or carbon cycle. Highly active phosphorus compounds in the presence of oxygen quickly turn into neutral complexes, "harmless" phosphates. In addition, phosphorus is usually scarce in ecosystems, i.e. it is present in low concentrations. Therefore, I repeat, attempts to take into account phosphorus only in the form of phosphates can lead to a noticeable distortion of its true role in ecosystems. And what an underestimation of this role can lead to can be clearly seen, for example, from previously thoughtlessly drained swamps, which ignite easily in dry years due to methane (CH 4), silane (SiH 4) and phosphine (PH 3).

According to the results of measurements at the above-mentioned Teriberka meteorological station, it was found that in 1990 48.8 million tons of methane were emitted into the atmosphere from the territory of Russia (recall, our estimates for the entire area of ​​treeless swamps amounted to about 60 million tons). For 1996-2003 the highest concentration was recorded in 2003. This year was the warmest for all of Russia, especially in summer and autumn in the swamp and tundra zones (Yakutia, Western Siberia) - on average, the temperature here turned out to be almost 6 ° C higher than the long-term one. Under these conditions, a summer decrease in the content of upstream ozone O 3 over the North of Russia by 5–10% was simultaneously observed. But in the summer, the processes of photosynthesis and the formation of oxygen are also accelerated here. Therefore, it is obvious that ozone was intensively consumed here to oxidize the increased amount of methane and phosphine under the conditions of warm 2003.

From Phosphine to Oxygen: Some Statistics and Philosophy

It is no secret that because of the richest biological resources, Russia has already become accustomed to be considered the world's oxygen donor. According to experts, 8130 million tons of O 2 are formed annually over its territory. It seems that we will not sin too much against the truth, assuming that the process of photosynthesis, which is responsible for the formation of this mass of oxygen, obeys the aforementioned "law of universal harmony" - the rule of the "golden section". After all, 1.47 tons of carbon dioxide, 0.6 tons of water and 3.84 Gcal of solar energy are spent on the formation of 1 ton of organic matter during photosynthesis, and 1.07 tons of oxygen are released. The ratio between the amount of absorbed CO 2 and the released O 2 (1.47: 1.07) is not so different from the “golden one”.

According to some published estimates, oxygen consumption in Russia (breathing, fuel combustion and other industrial needs) is 2784 million tons. Then its “production” by Russia exceeds its consumption by 5346 million tons. But in other calculations, which take into account the consumption of oxygen by microflora (formerly total soil) for "breathing", the Russian excess of oxygen production over its consumption is already an order of magnitude lower - 560 million tons. gas and consumed oxygen. On virgin lands, the value of this value is close to 1.58, and on arable land it fluctuates between 1.3-1.75 - in other words, oxygen is spent "economically" (42-37%) in the process of "breathing" of the soil (42-37%), and carbon dioxide is released more (58-63%). If we proceed from the average value of the "golden section" of 1.52 for the CO 2: O 2 ratio, then with the emission of CO 2 from the soils of Russia 10409 million tons of oxygen, another 6848 million tons of oxygen are consumed for the "breathing" of Russian soils (2004 estimates based on data employees of the Institute of Fundamental Problems of Biology of the Russian Academy of Sciences, in particular, V. N. Kudeyarov).

A kind of "golden proportion" is also observed between the sink of CO 2 and its emission on the scale of Russia. The ratio between the sink, which is 4450 million tons per year (in terms of carbon), and the emission (2800 million tons - in the same units) turns out to be equal to 1.59, i.e. surprisingly close to "golden". Well, as long as there is no excess of CO 2 over Russia as a whole, our ecosystems absorb more than we emit, our forests save us and cover our “sins”. But in recent years (primarily in the North), it has been increasingly noted that ecosystems cannot cope with the "plan" for absorption and the noted ratio is violated.

However, it is much more important that, as follows from a number of estimates, in Russia the total oxygen consumption per year for our needs (2784 million tons), soil respiration (6848 million tons) and the oxidation of methane and phosphine (220 million tons) is approaching 10 billion tons, which is almost 2 billion tons more than all our forests produce. And this sad balance seems to me a much more serious problem than the expected trading in quotas. For the sake of conservation environment and the biosphere of the planet, whose resources we today spend 25% more than they have time to recover, we must finally realize that without limiting consumption, we and our descendants simply cannot survive. And last but not least, it concerns oxygen. There seems to be a lot of it in the atmosphere (21%), but it should not be allowed that more of it is consumed on Earth than is produced.

Summing up

It is no secret that over the past 100 years, as a result of thoughtless human activity and ignoring the laws of nature, carbon dioxide emissions into the atmosphere (and its content there), according to various estimates, have increased by 25-35%. One of the poorly calculated consequences of global warming may be a sharp intensification of biochemical processes in natural areas of swamps and permafrost. At the same time, the emission of not only methane (this is already almost obvious) can sharply increase, but also gases that are little studied in terms of their effect on the biosphere: ammonia, silane and phosphine, which will require a lot of oxygen for oxidation and neutralization. But there are also not fully analyzed feedback effects (for example, a more intense release of methane will accelerate a further increase in the concentration of CO 2 in the atmosphere, which, in turn, can lead to a sharp slowdown in photosynthesis). As follows from recent studies, the compensatory role of photosynthesis in boreal forests noticeably weakened in the 1990s. But before it was firmly established that trees at all latitudes reliably contributed to photosynthesis and CO 2 assimilation. Dangerous trend! And examples of such "metamorphoses" of forests multiply year by year.

At present, we know almost nothing about the isolation and oxidation of the silane (SiH 4) mentioned more than once in this article. Meanwhile, all marsh plants, cereals and microorganisms are rich in organic silicon. In the peat of raised bogs - 43% SiO 2, transitional - 28%, lowland - 21%. So far, there is only fragmentary evidence that silane in combination with phosphine forms insufficiently studied complexes - silylphosphines. The processes of silane isolation, its oxidation and combination with other elements require serious study.

And in conclusion - a fantastic-looking plot that should make everyone think who has not yet lost this ability. In the surface layer of the atmosphere, due to the rapid increase in the content of carbon dioxide and some other "dead" gases, in the foreseeable future, there may be a shortage of oxygen not only due to a slowdown in photosynthesis, an increase in consumption for oxidation, combustion and respiration, but also because of the "screen" poisonous gases that interfere with the influx of O 2 from higher layers of the atmosphere.

For billions of years, the basis of all life on Earth was photosynthesis, which regularly supplied the planet with oxygen. Alas, as some researchers rightly point out, for the first time in history, modern civilization seems to have managed to slow down the replenishment of the atmosphere with oxygen, and brought nature to the point of bifurcation. Will she survive?

See, for example: Yeldyshev Yu.N. Is methane the culprit of global warming? // Ecology and Life, 2007, No. 11, p. 45; Climate change: facts and factors // Ecology and Life, 2008, No. 3, p. 44.
See, for example, the article Kravchenko I.K. in the journal "Microbiology", No. 6, 2007.