Phosphorus oxides. Phosphorus forms several oxides. The most important of them are P4O6 and P4O10. Often their formulas are written in simplified form as P2O3 and P2O5 (the previous indices are divided by 2).
Phosphorus (III) oxide P4O6 is a waxy crystalline mass that melts at 22.5° C. It is obtained by burning phosphorus with a lack of oxygen. Strong reducing agent. Very poisonous.
Phosphorus (V) oxide P4O10 is a white hygroscopic powder. It is obtained by burning phosphorus in excess air or oxygen. It combines very vigorously with water and also removes water from other compounds. Used as a dehumidifier for gases and liquids.
Oxides and all oxygen compounds Phosphorus is much stronger than similar nitrogen compounds, which should be explained by the weakening of the non-metallic properties of phosphorus compared to nitrogen.
Phosphorus (V) oxide. P2O5 interacts vigorously with water and also removes water from other compounds. That is why P2O5 is widely used as a desiccant for various substances from water vapor.
Phosphoric anhydride, interacting with water, primarily forms metaphosphoric acid HPO3:
When a solution of metaphosphoric acid is boiled, orthophosphoric acid H3PO4 is formed:
When H3PO4 is heated, pyrophosphoric acid H4P2O7 can be obtained:
P2O5 is a white snow-like substance that greedily absorbs
It contains water and is used for drying gases and liquids, and in some cases
yah for the removal of chemically bound water from substances:
2 HNO3 + P2O5 = N2O5 + 2 HPO3
4HClO4 + P4O10 → (HPO3)4 + 2Cl2O7.
Phosphorus(V) oxide is widely used in organic synthesis. It reacts with amides, turning them into nitriles:
P4O10 + RC(O)NH2 → P4O9(OH)2 + RCN
Carboxylic acids converts into the corresponding anhydrides:
P4O10 + 12RCOOH → 4H3P04 + 6(RCO)2O
P2O5 + 6RCOOH → 2H3P04 + 3(RCO)2O
Also interacts with alcohols, ethers, phenols and other organic compounds. In this case, a gap occurs P-O-P connections and organophosphorus compounds are formed. Reacts with NH3 and hydrogen halides, forming ammonium phosphates and phosphorus oxyhalides:
P4O10 + 8PCl3 + O2 → 12Cl3PO
When P4O10 fuses with basic oxides, it forms various solid phosphates, the nature of which depends on the reaction conditions.
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Phosphorus oxide is a colorless amorphous or glassy substance that exists in three crystalline, two amorphous and two liquid forms.
Toxic substance. Causes skin burns and irritation of the mucous membrane.
Phosphorus pentoxide is very hygroscopic. Reacts with alcohols, ethers, phenols, acids and other substances. During the reaction with organic substances The bonds between phosphorus and oxygen are broken, and organophosphorus compounds are formed. It enters into chemical reactions with ammonia (NH 3) and hydrogen halides to form ammonium phosphates and phosphorus oxyhalides. Forms phosphates with basic oxides.
Three-dimensional model of a molecule
Phosphorus pentoxide content in soil and fertilizers
In fact, in the soil there are only salts of orthophosphoric acid H 3 PO 4, but complex fertilizers can also contain salts of meta-, pyro- and polyphosphoric acids.
The basis for the formation of orthophosphoric acid is phosphorus pentoxide. That is why, and also due to the fact that plants do not absorb elemental phosphorus, it is agreed to indicate the concentration of phosphorus through the content of phosphorus pentoxide.
P2O5+3H 2O→ 2 H 3PO 4
All salts of orthophosphoric acid and monovalent cations (NH 4 +, Na +, K +) and monosubstituted salts of divalent cations (Ca(H 2 PO 4) 2 and Mg(H 2 PO 4) 2) found in the soil are soluble in water.
Disubstituted salts of divalent cations are insoluble in water, but easily dissolve in weak acids of root secretions and organic acids of microorganisms. In this regard, they are also a good source of P 2 O 5 for plants.
Phosphorus pentoxide uptake by plants
As mentioned above, in nature the main source of phosphorus is salts of orthophosphoric acid H 3 PO 4. However, after hydrolysis, pyro-, poly- and metaphosphates are also used by almost all crops.
Hydrolysis of sodium pyrophosphate:
Na 4 P 2 O 7 + H 2 O + 2H + → 2NaH 2 PO 4 +2Na +
Hydrolysis of sodium tripolyphosphate:
Na 5 P 3 O 10 + 2H 2 O + 2H + → 3NaH 2 PO 4 +2Na +
Hydrolysis of metaphosphate ion (in acidic environment):
(PO 3) 6 6- + 3H 2 O → H 2 P 3 O 10 3- + H 2 P 2 O 7 2- + H 2 PO 4 -
Orthophosphoric acid, being tribasic, dissociates three H 2 PO anions - 4, HPO 4 2-, PO 4 3-. Under slightly acidic environmental conditions, where plants are grown, the first ion is most common and accessible, the second to a lesser extent, and the third practically inaccessible. However, lupine, buckwheat, mustard, peas, sweet clover, hemp and other plants are able to absorb phosphorus from tribasic phosphates.
Some plants have adapted to absorb phosphate ion from phosphorus organic compounds(phytin, glycephosphates, etc.). The roots of these plants secrete a special enzyme (photophthase), which splits off the phosphoric acid anion from organic compounds, and then the plants absorb this anion. Plants of this kind include peas, beans, and corn. Moreover, phosphatase activity increases under conditions of phosphorus starvation.
Many plants can feed phosphorus from very dilute solutions, up to 0.01 mg/l P 2 O 5. Naturally, plants can satisfy the need for phosphorus only if the concentration in it is constantly restored to at least the same low level.
It has been experimentally established that the phosphorus absorbed by the roots is primarily used for the synthesis of nucleotides, and for further advancement to the above-ground part, the phosphates again enter the conducting vessels of the root in the form of mineral compounds.
Recalculation of phosphorus content in fertilizers
y = x.% × 30.974 ( molar mass) × 2 / 30.974 (molar mass) × 2 + 15.999 (molar mass of O) × 5
X- P 2 O 5 content in fertilizer, %;
y- P content in fertilizer, %
y = x, % × 0.43643
For example:
the fertilizer contains 40% phosphorus oxide
to recalculate the percentage of the element phosphorus in the fertilizer, you need to multiply mass fraction oxide in the fertilizer per mass fraction of the element in the oxide (for P 2 O 5 - 0.43643): 40 * 0.43643 = 17.4572%
Phosphorus was discovered and isolated in 1669 by the German chemist H. Brand. In nature, this element is found only in the form of compounds. The main minerals are phosphorite Ca3(PO4)2 and apatite 3Ca3(PO4)2. CaF2 or Ca5F(PO4)3. In addition, the element is part of proteins and is also found in teeth and bones. Phosphorus reacts most easily with oxygen and chlorine. With an excess of these substances, compounds with (for P) +5 are formed, and with a deficiency - with an oxidation state of +3. Phosphorus oxide can be represented by several formulas representing different chemical substances. Among them, the most common are P2O5 and P2O3. Other rare and little-studied oxides include: P4O7, P4O8, P4O9, PO and P2O6.
The oxidation reaction of elemental phosphorus with oxygen proceeds slowly. Its various sides are interesting. Firstly, in the dark the glow that accompanies it is clearly visible. Secondly, the oxidation process always occurs with the formation of ozone. This is due to the production of an intermediate compound - phosphoryl PO - according to the scheme: P + O2 → PO + O, and then: O + O2 → O3. Thirdly, oxidation is associated with a sharp change in the electrical conductivity of the surrounding air due to its ionization. Emission of light without noticeable heating when flowing chemical reactions, is called chemiluminescence. In humid environments, green chemiluminescence is due to the formation of the intermediate PO.
Phosphorus oxidation occurs only at a certain oxygen concentration. It should not be below the minimum and above the maximum O2 thresholds. The interval itself depends on temperatures and a number of other factors. For example, under standard conditions, the oxidation of phosphorus with pure oxygen increases to reach 300 mm Hg. Art. Then it decreases and drops almost to zero when the partial pressure of oxygen reaches 700 mmHg. Art. and higher. Thus, an oxide is not formed under normal conditions, since phosphorus is practically not oxidized.
Phosphorus pentoxide
The most characteristic oxide is phosphoric anhydride, or phosphorus, P2O5. It is a white powder with a pungent odor. When determining its molecular weight in vapor, it was found that the more correct notation of its formula is P4O10. This is a non-flammable substance, it melts at a temperature of 565.6 C. Anhydride P2O5 is an acidic oxide with all characteristic properties, but it greedily absorbs moisture, so it is used as a desiccant for liquids or gases. Phosphorus oxide can take away water, which is included in the composition chemical substances. Anhydride is formed as a result of the combustion of phosphorus in an atmosphere of oxygen or air, with a sufficient amount of O2 according to the scheme: 4P + 5O2 → 2P2O5. It is used in the production of H3PO4 acid. When interacting with water, it can form three acids:
- metaphosphoric: P2O5 + H2O → 2HPO3;
- pyrophosphoric: P2O5 + 2H2O → H4P2O7;
- orthophosphoric: P2O5 + 3H2O → 2H3PO4.
Phosphorus pentoxide reacts violently with water and substances containing water, such as wood or cotton. This creates a large number of heat, which may even cause a fire. It causes metal corrosion and is very irritating (serious burns to eyes and skin) to the respiratory tract and mucous membranes, even at concentrations as low as 1 mg/m³.
Phosphorus trioxide
Phosphorous anhydride, or phosphorus trioxide, P2O3 (P4O6) is a white crystalline substance(outwardly similar to wax), which melts at a temperature of 23.8 C and boils at a temperature of 173.7 C. Like P2O3, it is a very toxic substance. This is an acidic oxide, with all the inherent properties. Phosphorus oxide 3 is formed due to the slow oxidation or combustion of free substance (P) in an environment where there is a lack of oxygen. Phosphorus trioxide reacts slowly with cold water to form an acid: P2O3 + 3H2O → 2H3PO3. This phosphorus oxide reacts vigorously with hot water, and the reactions proceed in different ways, which can result in the formation of red phosphorus (an allotropically modified product), phosphorus hydride, and acids: H3PO3 and H3PO4. The thermal decomposition of the anhydride P4O6 is accompanied by the elimination of phosphorus atoms, resulting in the formation of mixtures of oxides P4O7, P4O8, P4O9. In structure they resemble P4O10. The most studied of them is P4O8.
Phosphorus(V) oxide
Phosphorus forms several oxides. The most important of them is phosphorus oxide (V) P 4 O 10 (Fig. 4). Often its formula is written in a simplified form - P 2 O 5. The structure of this oxide retains the tetrahedral arrangement of phosphorus atoms.
P2+5O5 Phosphoric anhydride (phosphorus (V) oxide)
White crystals, t 0 pl. = 570 0 C, t 0 boil. = 600 0 C, = 2.7 g/cm 3. Has several modifications. In vapor it consists of P4H10 molecules and is very hygroscopic (used as a desiccant for gases and liquids).
Receipt
4P + 5O 2 2P 2 O 5
Chemical properties
All Chemical properties acid oxides: reacts with water, basic oxides and alkalis
1) P 2 O 5 +H 2 O2HPO 3 (metaphosphoric acid)
P 2 O 5 + 2H 2 O H 4 P 2 O 7 (pyrophosphoric acid)
P 2 O 5 + 3H 2 O2H 3 P.O. 4 (orthophosphoric acid)
2) P 2 O 5 + 3BaO Ba 3 (P.O. 4 ) 2
Depending on the excess of alkali, it forms medium and acidic salts:
sodium hydrogen phosphate
sodium dihydrogen phosphate
Due to its exceptional hygroscopicity, phosphorus (V) oxide is used in laboratory and industrial technology as a drying and dehydrating agent. In its drying effect it surpasses all other substances. Chemically bound water is removed from anhydrous perchloric acid to form its anhydride:
Orthophosphoric acid. Several acids containing phosphorus are known. The most important of them is orthophosphoric acid H 3 PO 4 (Fig. 5).
Anhydrous orthophosphoric acid is light transparent crystals, with room temperature spreading in the air. Melting point 42.35 0 C. Phosphoric acid forms solutions of any concentration with water.
Orthophosphoric acid corresponds to the following structural formula:
In the laboratory phosphoric acid get oxidation of phosphorus with 30% nitric acid:
In industry, orthophosphoric acid is produced in two ways: extraction and thermal.
1. The extraction method is based on the treatment of crushed natural phosphates with sulfuric acid:
The phosphoric acid is then filtered and concentrated by evaporation.
2. The thermal method consists of reducing natural phosphates to free phosphorus, followed by burning it to P4O10 and dissolving the latter in water. Produced according to this method orthophosphoric acid is characterized by higher purity and increased concentration (up to 80% by weight).
Physical properties. Orthophosphoric acid is a solid, colorless, crystalline substance, highly soluble in water.
Chemical properties orthophosphoric acid are presented in Table 2:
table 2
Chemical properties of orthophosphoric acid
|
Common with other acids |
Specific |
|
1. An aqueous solution of acid changes the color of the indicators. Dissociation occurs in stages: The easiest way to dissociate is the first one. steps and the most difficult is the third
sodium hydrogen phosphate sodium dihydrogen phosphate 5. Reacts with salts of weak acids: |
 1. When heated, it gradually turns into metaphosphoric acid: biphosphorus acid 2. When exposed to a solution of silver (I) nitrate, a yellow precipitate appears: yellow sediment  3. Orthophosphoric acid plays an important role in the life of animals and plants. Its residues are part of adenosine triphosphoric acid ATP. When ATP breaks down, a large amount of energy is released. |
Orthophosphates. Phosphoric acid forms three series of salts. If we denote metal atoms by the letters Me, we can depict them in general view composition of its salts (Table 3).
Table 3
Chemical formulas of orthophosphates containing metals
Instead of a monovalent metal, the composition of orthophosphate molecules may include an ammonium group: (NH 4) 3 PO 4 - ammonium orthophosphate;
(NH 4) 2 HPO 4 --ammonium hydrogen orthophosphate; NH 4 H 2 PO 4 - ammonium dihydrogen orthophosphate.
Calcium and ammonium orthophosphates and hydroorthophosphates are widely used as fertilizers, and sodium orthophosphate and sodium hydrogenorthophosphate are used for the precipitation of calcium salts from water.
Phosphorus- element of the 3rd period and VA group Periodic table, serial number 15. Electronic formula of the atom [ 10 Ne]3s 2 3p 3, stable oxidation state in compounds +V.
Phosphorus oxidation state scale:
The electronegativity of phosphorus (2.32) is significantly lower than that of typical nonmetals and slightly higher than that of hydrogen. Forms various oxygenated acids, salts and binary compounds, exhibits non-metallic (acidic) properties. Most phosphates are insoluble in water.
In nature - thirteenth element by chemical abundance (sixth among non-metals), found only in a chemically bound form. Vital element.
The lack of phosphorus in the soil is compensated by the introduction of phosphorus fertilizers - mainly superphosphates.
Allotropic modifications of phosphorus
Red and white phosphorus P. Several allotropic forms of phosphorus are known in free form, the main ones being white phosphorus R 4 and red phosphorus Pn. In reaction equations, allotropic forms are represented as P (red) and P (white).
Red phosphorus consists of Pn polymer molecules of different lengths. Amorphous, at room temperature it slowly turns into white phosphorus. When heated to 416 °C, it sublimes (when the steam cools, white phosphorus condenses). Insoluble in organic solvents. Chemical activity is lower than that of white phosphorus. In air it ignites only when heated.
It is used as a reagent (safer than white phosphorus) in inorganic synthesis, a filler for incandescent lamps, and a component of box lubricant in the manufacture of matches. Not poisonous.
White phosphorus consists of P4 molecules. Soft like wax (cut with a knife). Melts and boils without decomposition (melt 44.14 °C, boil 287.3 °C, p 1.82 g/cm3). Oxidizes in air (green glow in the dark); with a large mass, self-ignition is possible. IN special conditions converted to red phosphorus. Well soluble in benzene, ethers, carbon disulfide. Does not react with water, stored under a layer of water. Extremely chemically active. Exhibits redox properties. Restores noble metals from solutions of their salts.
It is used in the production of H 3 P0 4 and red phosphorus, as a reagent in organic syntheses, a deoxidizing agent for alloys, and an incendiary agent. Burning phosphorus should be extinguished with sand (but not water!). Extremely poisonous.
Equations of the most important reactions of phosphorus:
Production of phosphorus in industry
- reduction of phosphorite with hot coke (sand is added to bind calcium):
Ca 3 (PO4)2 + 5C + 3SiO2 = 3CaSiO3 + 2 R+ 5СО (1000 °С)
The phosphorus vapor is cooled and solid white phosphorus is obtained.
Red phosphorus is prepared from white phosphorus (see above); depending on the conditions, the degree of polymerization n (P n) can be different.
Phosphorus compounds
Phosphine PH 3. Binary compound, the oxidation state of phosphorus is III. Colorless gas with an unpleasant odor. The molecule has the structure of an incomplete tetrahedron [: P(H) 3 ] (sp 3 hybridization). Slightly soluble in water, does not react with it (unlike NH 3). A strong reducing agent, burns in air, oxidizes to HNO 3 (conc.). Attaches HI. Used for synthesis organophosphorus compounds. Highly poisonous.
Equations of the most important reactions of phosphine:
Obtaining phosphine in laboratories:
Casp2 + 6HCl (dil.) = 3CaCl + 2 RNZ
Phosphorus (V) oxide P 2 O 5. Acidic oxide. White, thermally stable. In the solid and gaseous states, the P 4 O 10 dimer has a structure of four tetrahedra connected along three vertices (P - O-P). At very high temperatures it monomerizes to P 2 O 5 . There is also a glassy polymer (P 2 0 5) n. It is extremely hygroscopic, reacts vigorously with water and alkalis. Restored with white phosphorus. Removes water from oxygen-containing acids.
Used as a very effective dehydrating agent for drying solids, liquids and gas mixtures, reagent in the production of phosphate glasses, catalyst for the polymerization of alkenes. Poisonous.
Equations for the most important reactions of phosphorus oxide +5:
Receipt: burning phosphorus in excess dry air.
Orthophosphoric acid H 3 P0 4. Oxoacid. White matter, hygroscopic, final product interaction of P 2 O 5 with water. The molecule has the structure of a distorted tetrahedron [P(O)(OH) 3 ] (sp 3 -hybridisadium), contains covalent σ-bonds P - OH and σ, π-bond P=O. Melts without decomposition, and decomposes upon further heating. It is highly soluble in water (548 g/100 g H 2 0). A weak acid in solution, it is neutralized by alkalis, and not completely by ammonia hydrate. Reacts with typical metals. Enters into ion exchange reactions.
A qualitative reaction is the precipitation of a yellow precipitate of silver (I) orthophosphate. It is used in the production of mineral fertilizers, for the clarification of sucrose, as a catalyst in organic synthesis, and as a component of anti-corrosion coatings on cast iron and steel.
Equations of the most important reactions of orthophosphoric acid:
Production of phosphoric acid in industry:
boiling phosphate rock in sulfuric acid:
Ca3(PO4)2 + 3H2SO4 (conc.) = 2 H3PO4+ 3CaSO4
Sodium orthophosphate Na 3 PO 4. Oxosol. White, hygroscopic. Melts without decomposition, thermally stable. It is highly soluble in water, hydrolyzes at the anion, and creates a highly alkaline environment in solution. Reacts in solution with zinc and aluminum.
Enters into ion exchange reactions.
Qualitative reaction to the PO 4 3- ion
— formation of a yellow precipitate of silver(I) orthophosphate.
Used to eliminate “permanent” hardness fresh water, as a component detergents and photographic developers, a reagent in the synthesis of rubber. Equations of the most important reactions:
Receipt: complete neutralization of H 3 P0 4 with sodium hydroxide or according to the reaction:
Sodium hydrogen phosphate Na 2 HPO 4. Acid oxo salt. White, decomposes without melting when heated moderately. It is highly soluble in water and hydrolyzes at the anion. Reacts with H 3 P0 4 (conc.), neutralized by alkalis. Enters into ion exchange reactions.
Qualitative reaction to the HPO 4 2- ion— formation of a yellow precipitate of silver (I) orthophosphate.
It is used as an emulsifier for condensing cow's milk, a component of food pasteurizers and photo-bleaches.
Equations of the most important reactions:
Receipt: incomplete neutralization of H 3 P0 4 with sodium hydroxide in a dilute solution:
2NaOH + H3PO4 = Na2HPO4 + 2H2O
Sodium dihydrogen orthophosphate NaH 2 PO 4. Acid oxo salt. White, hygroscopic. When heated moderately, it decomposes without melting. Highly soluble in water, the H 2 P0 4 anion undergoes reversible dissociation. Neutralized by alkalis. Enters into ion exchange reactions.
Qualitative reaction to the H 2 P0 4 ion - formation of a yellow precipitate of silver orthophosphate (1).
Used in glass production, to protect steel and cast iron from corrosion, and as a water softener.
Equations of the most important reactions:
Receipt: incomplete neutralization of H 3 PO 4 with sodium hydroxide:
H3PO4 (conc.) + NaOH (dil.) = NaH2PO4+ H2O
Calcium orthophosphate Ca 3(PO 4)2— Oxosol. White, refractory, thermally stable. Insoluble in water. Decomposes concentrated acids. Restored by coke during fusion. The main component of phosphorite ores (apatite, etc.).
It is used to obtain phosphorus in the production of phosphorus fertilizers (superphosphates), ceramics and glass; the precipitated powder is used as a component of toothpastes and a polymer stabilizer.
Equations of the most important reactions:
Phosphorus fertilizers
The mixture of Ca(H 2 P0 4) 2 and CaS0 4 is called simple superphosphate, Ca(H 2 P0 4) 2 with an admixture of CaНР0 4 - double superphosphate, they are easily absorbed by plants when feeding.
The most valuable fertilizers are ammophos(contain nitrogen and phosphorus), are a mixture of ammonium acid salts NH 4 H 2 PO 4 and (NH 4) 2 HPO 4.
Phosphorus (V) chloride PCI5. Binary connection. White, volatile, thermally unstable. The molecule has the structure of a trigonal bipyramid (sp 3 d-hybridization). In the solid state, the dimer P 2 Cl 10 with the ionic structure PCl 4 + [PCl 6 ] - . “Smoke” in humid air. Very reactive, completely hydrolyzed by water, reacts with alkalis. Restored with white phosphorus. It is used as a chlorine agent in organic synthesis. Poisonous.
Equations of the most important reactions:
Receipt: chlorination of phosphorus.
