Nitrogen enters the earth's atmosphere in an unbound form in the form of diatomic molecules. Approximately 78% of the total volume of the atmosphere is nitrogen. In addition, nitrogen is included in plants and animal organisms in the form of proteins. Plants synthesize proteins using nitrates from the soil. Nitrates are formed there from atmospheric nitrogen and ammonium compounds present in the soil. The process of converting atmospheric nitrogen into a form that can be used by plants and animals is called nitrogen fixation.
Nitrogen fixation can occur in two ways:
1) During a lightning strike, some nitrogen and oxygen in the atmosphere combine to form nitrogen oxides. They dissolve in water, forming dilute nitric acid, which in turn forms nitrates in the soil.
2) Atmospheric nitrogen is converted to ammonia, which is then converted by bacteria to nitrates in a process called nitrification. Some
of these bacteria are present in the soil, while others exist in the nodules of the root system of nodule plants, such as clover.
Nitrosamine. Recently, there has been an increase in the content of nitrates in drinking water, mainly due to the increased use of artificial ones. nitrogen fertilizers in agriculture. Although nitrates themselves are not that dangerous for adults, they can be converted into nitrites in the human body. In addition, nitrates and nitrites are used to process and preserve many foods, including ham, bacon, corned beef, and some cheeses and fish. Some scientists believe that in the human body nitrates can be converted into nitrosamines:
It is known that nitrosamines can cause cancer in animals. Most of us are already exposed to nitrosamines, which are found in small amounts in air pollution, cigarette smoke and some pesticides. It is believed that nitrosamines may be the cause of 70-90% of cases of cancer, the occurrence of which is attributed to the action of environmental factors.
(see scan)
Rice. 15.15. Nitrogen cycle in nature.
Nitrates are also added to the soil in the form of fertilizers. In ch. 13 nitrogen-containing fertilizers such as calcium nitrate, ammonium nitrate, sodium nitrate and potassium nitrate have already been described.
Plants absorb nitrates from the soil through their root system.
After plants and animals die, their proteins decompose to form ammonium compounds. These compounds are eventually converted by putrefactive bacteria into nitrates, which remain in the soil, and nitrogen, which is returned to the atmosphere.
All these processes are components of the nitrogen cycle in nature (see Fig. 15.15).
Every year, more than 50 million tons of nitrogen are produced worldwide. Pure nitrogen, along with oxygen and other gases, including argon, is produced industrially using fractional distillation of liquefied air. This process includes three stages. At the first stage, dust particles, water vapor and carbon dioxide are removed from the air. The air is then liquefied by cooling it and compressing it to
high pressures. At the third stage, nitrogen, oxygen and argon are separated by fractional distillation of liquid air.
Approximately three quarters of all nitrogen produced annually in the UK is converted to ammonia (see section 7.2), a third of which is then converted to nitric acid (see below).
Nitric acid has a number of important uses:
1) approximately 80% of the synthesized nitric acid - to obtain ammonium nitrate fertilizer;
2) in the production of synthetic yarn, such as nylon;
3) for the manufacture of explosives, for example trinitrotoluene (tol) or trinitroglycerin (dynamite);
4) for the nitration of aromatic amines in the production of dyes.
Nitrates are used to produce fertilizers and explosives. For example, gunpowder is a mixture of sulfur, charcoal and sodium nitrate. Strontium nitrate and barium nitrate are used in pyrotechnics to produce red and pale green lights, respectively.
Tol and dynamite. Tol is a shortened name for trinitrotoluene. Dynamite contains trinitroglycerin, which is impregnated with kieselguhr. Nitric acid is used to produce this and other explosives.
Silver nitrate is used to produce silver halides used in photography.
Nitrogen is used to create an inert atmosphere in the production of plate glass, semiconductors, vitamin A, nylon and sodium lead alloy, which is used to make. Liquid nitrogen is used for refrigerated storage of blood, bovine semen (for breeding purposes) and some food products.
Phosphorus, like nitrogen, is also one of the essential elements for life and is part of all living organisms. It is found in bone tissue and is necessary for animals in metabolic processes to accumulate energy.
Phosphorus is found naturally in minerals such as apatite, which contains calcium phosphate. Approximately 125 million tons of phosphate ore are mined worldwide each year. Most of it is spent on the production of phosphate fertilizers (see Chapter 13).
White phosphorus is obtained from phosphate ore by calcining it in a mixture with coke and silica in an electric furnace at a temperature of about 1500°C. This produces an oxide which is then reduced to white phosphorus by heating in a mixture with coke. Red phosphorus is obtained by heating white phosphorus without access to air at a temperature of about 270 ° C for several days.
Red phosphorus is used to make matches. They cover the sides of a matchbox. Match heads are made from potassium, manganese (IV) oxide and sulfur. When a match rubs against the box, phosphorus oxidizes. Most of the white phosphorus produced today is consumed in the production of phosphoric acid. Phosphoric acid is used in production
stainless steel and for chemical polishing of aluminum and copper alloys. Dilute phosphoric acid is also used in the food industry to regulate the acidity of jelly products and soft drinks.
Pure calcium phosphate is also used in the food industry, for example in baking powder. One of the most important phosphate compounds is sodium tripolyphosphate. It is used to make synthetic detergents and other types of water softeners. Polyphosphates are also used to increase the water content of some foods.
Nitrogen and Phosphorus
The elements Nitrogen and Phosphorus are located in Group V of the Periodic Table, Nitrogen in the 2nd period, Phosphorus in the 3rd. Electronic configuration of the Nitrogen atom:
Nitrogen valency: III and IV, oxidation state in compounds: from -3 to +5.
Structure of a nitrogen molecule: , .
Electron configuration of Phosphorus atom: 
Electronic configuration of the Phosphorus atom in the excited state: 
Phosphorus valence: III and V, oxidation state in compounds: -3, 0, +3, +5.
Physical properties of nitrogen. A colorless gas, tasteless and odorless, slightly lighter than air (g/mol, g/mol), poorly soluble in water. Melting point -210 °C, boiling point -196 °C.
Allotropic modifications of Phosphorus. Among the simple substances that form the element Phosphorus, the most common are white, red and black phosphorus.
Distribution of Nitrogen in nature. Nitrogen occurs in nature primarily as molecular nitrogen. In air, the volume fraction of nitrogen is 78.1%, mass - 75.6%. Nitrogen compounds are found in small quantities in the soil. Nitrogen is found in living organisms as part of organic compounds (proteins, nucleic acids, ATP).
Distribution of Phosphorus in nature. Phosphorus is found in a chemically bound state in the composition of minerals: phosphorites, apatites, the main component of which is . Phosphorus is a vital element; it is part of lipids, nucleic acids, ATP, calcium orthophosphate (in bones and teeth).
Obtaining nitrogen and phosphorus.
Nitrogen obtained industrially from liquid air: since nitrogen has the lowest boiling point of all atmospheric gases, it evaporates first from liquid air. In the laboratory, nitrogen is obtained from the thermal decomposition of ammonium nitrite: . Phosphorus obtained from apatites or phosphorites by calcining them with coke and sand at a temperature:
Chemical properties of nitrogen.
1) Interaction with metals. The substances formed as a result of these reactions are called nitrides And.
At room temperature, nitrogen reacts only with lithium:
Nitrogen reacts with other metals at high temperatures:
- aluminum nitride
Nitrogen reacts with hydrogen in the presence of a catalyst at high pressure and temperature:
- ammonia
At very high temperatures (about ) nitrogen reacts with oxygen:
- nitrogen(II) oxide
Chemical properties of phosphorus.
1) Interaction with metals.
When heated, phosphorus reacts with metals:
- calcium phosphide
2) Interaction with non-metals.
White phosphorus ignites spontaneously, while red phosphorus burns when ignited:
- phosphorus(V) oxide
When there is a lack of oxygen, phosphorus(III) oxide is formed (a very toxic substance):
Interaction with halogens:
Interaction with sulfur:
Ammonia
Molecular formula of ammonia: . Electronic formula:
Structural formula:
Physical properties of ammonia. A colorless gas with a characteristic pungent odor, almost twice as light as air, poisonous. When pressure increases or cools, it easily scrapes into a colorless liquid, boiling point, melting point. Ammonia dissolves very well in water: with 1 volume of water, up to 700 volumes of ammonia dissolves, with 1200 volumes.
Production of ammonia.
1) Ammonia is obtained in the laboratory by heating a dry mixture of calcium hydroxide (slaked lime) and ammonium chloride (ammonia):
2) Ammonia in industry is obtained from simple substances - nitrogen and hydrogen:
Chemical properties of ammonia. Nitrogen in ammonia has the lowest oxidation state and therefore exhibits only reducing properties.
1) Combustion in an atmosphere of pure oxygen or in heated air:
2) Oxidation to nitrogen(II) oxide in the presence of a catalyst (hot platinum):
3) Reverse interaction with water:
The presence of ions determines the alkaline environment of the ammonia solution. The resulting solution is called ammonia or amoniacal water. Ammonium ions exist only in solution. It is impossible to isolate ammonium hydroxide as an independent compound.
4) Recovery of metals from their oxides:
5) Interaction with acids to form ammonium salts (compound reaction):
- ammonium nitrate.
Application of ammonia. A large amount of ammonia is consumed to produce nitric acid, nitrogenous salts, urea, and soda using the ammonium method. Its use in refrigeration units is based on light scraping and subsequent evaporation with heat absorption. Aqueous solutions of ammonia are used as nitrate fertilizers.
Ammonium salts
Ammonium salts- salts containing a cation group. For example, - ammonium chloride, - ammonium nitrate, - ammonium sulfate. Physical properties of ammonium salts. White crystalline substances, highly soluble in water.
Preparation of ammonium salts. Ammonium salts are formed when gaseous ammonia or its solutions react with acids:
Chemical properties of ammonium salts.
1) Dissociation:
2) Interaction with other salts:
3) Interaction with acids:
4) Interaction with alkalis:
This reaction is qualitative for ammonium salts. Ammonia released is determined by its odor or the blueness of wet indicator paper.
5) Heat decomposition:
Application of ammonium salts. Ammonium salts are used in the chemical industry and as mineral fertilizers in agriculture.
Nitrogen oxides and phosphorus oxides
Nitrogen forms oxides in which it exhibits an oxidation state from +1 to +5: ; NO; ; ; ; . All nitrogen oxides are poisonous. The oxide has narcotic properties, which at the initial stage are indicated by euphoria, hence the name “laughing gas”. The oxide irritates the respiratory tract and mucous membranes of the eyes. A harmful consequence of chemical production, it enters the atmosphere in the form of a “fox tail” - red-brown in color.
Phosphorus oxides: and. Phosphorus(V) oxide is the most stable oxide under normal conditions.
Obtaining nitrogen oxides and phosphorus oxides.
With the direct combination of molecular nitrogen and oxygen, only nitrogen(II) oxide is formed:
Other oxides are obtained indirectly.
Phosphorus(V) oxide is obtained by burning phosphorus in excess oxygen or air:
Chemical properties of nitrogen oxides.
1) - oxidizer, can support combustion:
2) NO - easily oxidized:
Does not react with water and alkalis.
3) acid oxide:
4) - strong oxidizing agent, acid oxide:
In the presence of excess oxygen:
Dimerizes, forming an oxide - a colorless liquid: . The reaction is reversible. At -11 °C the equilibrium is practically shifted towards the formation of , and at 140 °C - towards the formation of .
5) - acid oxide:
Chemical properties of phosphorus(V) oxide. Phosphorus-containing acids.
- typically acidic oxide. Three acids correspond to it: meta-,ortho- And diphosphate A. When dissolved in water, metaphosphate acid is first formed:
During prolonged boiling with water - orthophosphate acid:
When orthophosphate acid is carefully calcined, diphosphate acid is formed:
Application of nitrogen oxides and phosphorus oxides.
Nitrogen(IV) oxide is used in the production of nitric acid, nitrogen(IV) oxide is used in medicine.
Phosphorus(V) oxide is used for drying gases and liquids, and in some cases for removing chemically bound water from substances.
Nitric and phosphate acids
Physical properties of orthophosphate (phosphoric) acid. Under normal conditions, it is a solid, colorless, crystalline substance. Melting point +42.3. In solid and liquid acids, molecules are joined together by hydrogen bonds. This is due to the increased viscosity of concentrated solutions of phosphoric acid. It is highly soluble in water, its solution is an electrolyte of medium strength. Physical properties of nitric acid. Anhydrous (100%) acid is a colorless liquid with a strong odor, boiling point . If stored in the light, it gradually turns brown due to the decomposition and formation of higher nitrogen oxides, including brown gas. Mixes well with water in any ratio.
Preparation of phosphate acid.
1) From its salts contained in phosphate minerals (apatites and phosphorites), under the action of sulfuric acid:
2) Hydration of phosphorus(V) oxide:
Preparation of nitrate acid.
1) From dry salts of nitric acid under the action of concentrated sulfuric acid:
2) With nitrogen oxides:
3) Industrial synthesis of nitric acid:
- catalytic oxidation of ammonia, catalyst - platinum.
- oxidation by atmospheric oxygen.
- absorption by water in the presence of oxygen.
Chemical properties of phosphoric acid. Exhibits all typical properties of acids. Phosphate acid is tribasic and forms two series of acid salts - dihydrophosphate And hydrogen phosphate s.
1) Dissociation:
4) Interaction with salts. The reaction with argentum nitrate is qualitative for the ion - a yellowish precipitate of argentum phosphate forms:
5) Interaction with metals in the electrochemical voltage range up to Hydrogen:
Chemical properties of nitric acid. Nitric acid is a strong oxidizing agent.
1) Dissociation:
2) Interaction with metal oxides:
3) Interaction with bases:
4) Interaction with salts:
5) Interaction with metals. When concentrated and dilute nitric acid reacts with metals, salt (nitrate), nitrogen oxides, nitrogen or ammonia and water are formed.
Application of orthophosphate and nitric acids.
Orthophosphate acid widely used in the production of mineral fertilizers. It is non-toxic and is used in the food industry to make syrups and drinks (Coca-Cola, Pepsi-Cola).
Nitric acid is spent on the production of nitrogen fertilizers, explosives, medicines, dyes, plastics, artificial fibers and other materials. Concentrated nitric acid is used in rocket technology as a rocket fuel oxidizer.
Nitrates
Nitric acid salts - nitrate s. These are crystalline solids Lecture outline
1. Nitrogen. Position in the PS. Oxidation states. Being in nature. Physical and chemical properties.
2. Hydrogen compounds of nitrogen (ammonia, hydrazine, hydroxylamine, hydronitrous acid).
3. Oxygen compounds of nitrogen (nitrogen oxides, nitrous, nitrous and nitric acids).
4. Phosphorus. Physical and chemical properties. Hydrogen and oxygen compounds.
5. Nitrogen and phosphorus fertilizers.
14.1 Nitrogen. Position in the PS. Oxidation states. Being in nature. Physical and chemical properties
Nitrogen is a p-element of group 5 PS. It has 5 electrons in its valence layer (2s 2 2p 3). Oxidation states -3, -2, -1, 0, +1, +2, +3, +4, +5. This is a typical non-metal.
The total nitrogen content of the earth's crust is about 0.03%. The largest part of it is concentrated in the atmosphere, the bulk of which (75.6 wt.%) is free nitrogen (N 2). Complex organic derivatives of nitrogen are part of all living organisms. As a result of the death of these living organisms and the decay of their remains, simpler nitrogen compounds are formed, which, under favorable conditions (mainly the lack of moisture) can accumulate in the earth's crust.
Under normal conditions, nitrogen is a colorless, odorless gas. It is also colorless in liquid and solid states.
Free nitrogen is chemically very inert. There is a triple bond between the atoms in the nitrogen molecule (bond energy 940 kJ/mol). Under normal conditions, it practically does not react with either metals (except Li and Mg) or non-metals. Heating increases its chemical activity mainly towards metals, with some of which it combines to form nitrides. At a temperature of 3000 0 C it reacts with oxygen in the air.
14.2 Hydrogen nitrogen compounds (ammonia, hydrazine and hydroxylamine)
Formulas of hydrogen compounds, respectively:
NH 3, N 2 H 4, NH 2 OH, HN 3.
Ammonia is a colorless gas with a characteristic pungent odor (“ammonia”). Its solubility in water is greater than that of all other gases: one volume of water absorbs about 1200 volumes of NH 3 at 0ºC, and about 700 at 20ºC.
Hydrazine N2H4 is a colorless liquid that fumes in air and easily mixes with water, and hydroxylamine NH 2 OH It is colorless crystals, highly soluble in water.
For the chemical characterization of ammonia, hydrazine and hydroxylamine, reactions of three types are of primary importance: addition, hydrogen substitution and oxidation.
When dissolved in water, some of the ammonia molecules react chemically with water, forming a weak base (K d = 1.8 × 10 -5).
NH 3 + H 2 O ↔ NH 4 OH ↔ NH 4 + + OH¯
Hydrazine and hydroxylamine also partially react with water. Solutions of these substances are weaker bases compared to ammonia (K d = 8.5×10 -7 and K d = 2∙10 -8).
Hydronitric acid HN 3 is a colorless liquid with a pungent odor; its poisonous, corrosive mucous membranes, vapors explode with great force upon contact with heated objects.
The acid is stable in aqueous solutions. This is a weak (slightly weaker than acetic) acid (K = 1.2∙10-5), dissociating according to the following scheme:
HN 3 ↔ H + + N 3 -
The salts are called azides, explosives (detonators).
14.3 Oxygen compounds of nitrogen (nitrogen oxides, nitric and nitrous acids)
Nitrogen forms oxides: N 2 O, NO, N 2 O 3, NO 2, N 2 O 5. All oxides are gaseous substances under normal conditions, except N 2 O 5 (a colorless crystalline substance).
The first two are non-salt-forming, while the rest are acidic.
N 2 O 3 - nitrous acid anhydride (HNO 2).
NO 2 - nitrous anhydride (HNO 2). and nitric (HNO 3) acids.
N 2 O 5 – nitric acid anhydride.
Nitrogen forms several acids: H 2 N 2 O 2 - nitrous, HNO 2 - nitrous, HNO 3 - nitric.
Nitrous acid H 2 N 2 O 2 white crystalline substance, explosive, easily soluble in water. In an aqueous solution it is a weak, moderately stable, dibasic acid (K 1 d = 9 × 10 -8 and K 2 d = 10 -11).
Nitrous acid HNO 2 a weak and unstable monobasic acid (Kd = 5×10 -4), existing in aqueous solutions. Nitrite salts are stable. Nitrous acid and its salts exhibit redox duality because they contain nitrogen in an intermediate oxidation state (+3).
Clean nitric acid HNO 3-colorless liquid with a density of 1.51 g/cm at – 42°C, solidifying into a transparent crystalline mass
Nitric acid is one of the strongest acids; in dilute aqueous solutions it completely disintegrates into ions:
HNO 3 → H + + NO 3 ¯.
Nitric acid is a strong oxidizing agent. It oxidizes metals to salts, and non-metals to higher oxygen acids. At the same time, it is reduced in concentrated solutions to nitrogen dioxide, and in diluted solutions, the products of its reduction, depending on the activity of the metal, can contain N 2, NO, N 2 O, N 2 O 3, NH 4 NO 3.
Nitric acid has no effect on gold, platinum, rhodium and iridium. Some metals are passivated (coated with a protective film) in concentrated nitric acid. These are aluminum, iron and chrome.
Salts of nitric acid - nitrates. They dissolve well in water and are stable under normal conditions. When heated, they decompose releasing oxygen.
14.4 Phosphorus. Physical and chemical properties. Hydrogen and oxygen compounds
For solid phosphorus, several allotropic modifications are known, of which only two are practically encountered: white and red.
During storage, white phosphorus gradually (very slowly) turns into a more stable red form. The transition is accompanied by the release of heat (heat of transition):
P white = P red + 4 kcal
The chemical activity of phosphorus is much higher than that of nitrogen. Thus, it easily combines with oxygen, halogens, sulfur and many metals. In the latter case, phosphides similar to nitrides are formed (Mg 3 P 2, Ca 3 P 2, etc.).
Hydrogen compounds of phosphorus are phosphine (PH 3) and diphosphine (P 2 H 4).
Diphosphine (P 2 H 4) is liquid hydrogen phosphate, self-igniting in air (will-o-the-wisps in the cemetery are explained by the formation of this substance during the smoldering of remains).
Phosphorous hydrogen (“phosphine”) – PH 3 is a colorless gas with an unpleasant odor (“rotten fish”). Phosphine is a very strong reducing agent (phosphorus has an oxidation state of –3) and is highly toxic. In contrast to ammonia, addition reactions are not very common for phosphine. Phosphonium salts are known only for a few strong acids and are very unstable, and phosphine does not chemically interact with water (although it is quite soluble in it).
Oxygen compounds of phosphorus - oxides P 2 O 3 and P 2 O 5, existing in the form of dimers (P 2 O 3) 2 and (P 2 O 5) 2, as well as acids: H 3 PO 2 - hypophosphorous, H 3 PO 3 – phosphorous, H 3 PO 4 – phosphoric.
The combustion of phosphorus with a lack of air or slow oxidation produces mainly phosphorous anhydride (P 2 O 3). The latter is a white (wax-like) crystalline mass. When heated in air, it turns into P 2 O 5 (a white snow-like mass). Interacting with cold water, P 2 O 3 slowly forms phosphorous acid:
P 2 O 3 + 3H 2 O = 2H 3 PO 3
P 2 O 5 - higher oxide - phosphoric anhydride is obtained by combustion of phosphorus in excess oxygen (or air). Phosphoric anhydride (P 2 O 5) attracts moisture extremely vigorously and is therefore often used as a gas desiccant.
The interaction of P 2 O 5 with water, depending on the number of attached H 2 O molecules, leads to the formation of the following hydrate forms:
P 2 O 5 + H 2 O = 2HPO 3 (metaphosphoric)
P 2 O 5 + 2H 2 O = H 4 P 2 O 7 (pyrophosphoric acid)
P 2 O 5 + 3H 2 O = 2H 3 PO 4 (orthophosphoric acid)
H 3 PO 2 (phosphorous acid) - it is a colorless crystalline substance. In aqueous solution it is a strong monobasic acid. It is the strongest among phosphorus acids. The acid itself and its salts (hypophosphites) are reducing agents.
Free phosphorous acid (H 3 PO 3) are colorless crystals that diffuse in air and are easily soluble in water. It is a strong (but in most cases slow-acting) reducing agent. Despite the presence of three hydrogens in the molecule, H 3 PO 3 functions only as a dibasic acid of medium strength. Its salts (phosphorus or phosphites), as a rule, are colorless and poorly soluble in water. Of the derivatives of more commonly occurring metals, only Na, K, and Ca salts are highly soluble.
Of the pentavalent phosphorus acids, orthohydrate (H 3 PO 4) has the greatest practical importance.
Phosphoric acid It is a colorless crystal that diffuses in air. It is usually sold in the form of an 85% aqueous solution, approximately corresponding to the composition of 2H 3 PO 4 H 2 O and having the consistency of thick syrup. Unlike many other phosphorus derivatives, H 3 PO 4 is non-toxic. Oxidizing properties are not at all characteristic of it.
Being a tribasic acid of medium strength, H 3 PO 4 is capable of forming three series of salts, for example: acid salts Na 2 HPO 4 and Na 2 HPO 4, as well as the middle salt - Na 3 PO 4
NaH 2 PO 4 - sodium dihydrogen phosphate (primary sodium phosphate)
Na 2 HPO 4 - sodium hydrogen phosphate (secondary sodium phosphate)
Na 3 PO 4 – sodium phosphate (tertiary sodium phosphate).
14.5 Nitrogen and phosphorus fertilizers.
Nitrogen and phosphorus are macroelements that are needed by plant and animal organisms in large quantities. Nitrogen is part of protein. Phosphorus is part of bones. Organic derivatives of phosphoric acid are sources of energy for endothermic cell reactions.
Nitrogen fertilizers are salts of nitric acid: KNO 3 - potassium nitrate, NaNO 3 - sodium nitrate, NH 4 NO 3 - ammonium nitrate, Ca(NO 3) 2 - Norwegian nitrate. Solutions of ammonia in water are liquid nitrogen fertilizers.
Phosphorus fertilizers are salts of phosphoric acid: Ca(H 2 PO 4) 2 × 2CaSO 4 - simple superphosphate, Ca(H 2 PO 4) 2 - double superphosphate, CaHPO 4 × 2H 2 O - precipitate. Macrofertilizers are applied to the soil in large quantities (in centners per hectare).
Fertilizing with mineral fertilizers is the most important measure when caring for plants. Any mineral fertilizer is an artificially created concentrate that contains nutrients in the form of mineral salts. Usually the soil contains all the compounds necessary for the plant, but in certain phases of development the crop requires increased doses of any element. In such cases, you cannot do without mineral supplements. It allows you to get a high yield with a very modest investment of money and labor. Fertilizers can be simple or complex, depending on how much nutrients they include.
- ammonium nitrate and urea – 10-25 g/m2;
- sodium and calcium nitrate: up to 70 g/m2.
- nitrate;
- ammonium;
- ammonia;
- ammonium nitrate;
- amide.
- first grade – 16.4%;
- second grade – 16.3%;
- technical 15.5%.
- manganese;
- molybdenum;
- copper;
- cobalt.
- old leaves turn brown at the edges and take on a burnt appearance;
- leaves curl and corrugate;
- potato leaves acquire a characteristic bronze coating;
- Vegetable stems become tough and woody.
- potassium chloride - 20-40 g/m2;
- potassium sulfate - 10-15 g/m2;
- potassium nitrate - 15-20 g/m2.
- superphosphate - with ammonium nitrate, ammonium sulfate, potassium chloride;
- double superphosphate - with urea;
- all nitrogen fertilizers (except urea) - with manure.
Show all
Nitrogen
Soils in regions with a rainy climate and artificially irrigated, such as the lands of greenhouses, vegetable gardens, and household plots, are always poor in nitrogen. The element easily dissolves in water.
With high rainfall or frequent watering, nitrogen seeps from the top layer of soil, where the roots of crop plants are located, deeper and becomes unavailable. In such cases, nitrogen fertilizers provide a significant increase in yield, which can reach up to 50%.
In the Non-Black Earth Region, with an optimal dose of nitrogen fertilizer, each kilogram of nitrogen produces an additional 50-70 kg of potatoes, 20-30 kg of white cabbage, 6-7 kg of onions.
Average application rates of nitrogen fertilizers:
In Russia, the largest amount of precipitation falls on the Black Sea coast, in the northern part of the Urals, in the Irkutsk, Kemerovo regions, and Khanty-Mansiysk. The soil is heavily washed out in the Pskov, Smolensk, Vologda, and Leningrad regions. In these regions, it is impossible to get a good harvest without nitrogen fertilizers.
In one-component fertilizers, nitrogen can be in different forms:
Nitrate
Nitrogen in nitrate form is found in sodium and calcium nitrate. These fertilizers are a by-product of chemical production. There are only a few of them produced - less than 1% of all nitrogen fertilizers.
Sodium nitrate
Sodium or Chilean nitrate has the formula NaNO3. In addition to nitrogen, the product contains sodium – 26%.
Chilean saltpeter looks like small crystals of white or yellowish color. It dissolves well in water, giving it a bitter-salty taste. When properly stored, it practically does not cake, since it does not absorb moisture from the air.
After applying nitrate, the soil becomes slightly alkalized. In agriculture, the product is used to feed winter crops, perennial herbs, berries and vegetables. The fertilizer is especially useful for root crops: fodder and table beets, potatoes, carrots. This is explained by the fact that sodium accelerates the outflow of carbohydrates from the above-ground part to the underground part. As a result, the root vegetables grow larger and sweeter. Sodium nitrate can be mixed with superphosphate and potassium chloride.
Calcium nitrate
The fertilizer contains from 15 to 17% nitrogen. The fertilizer looks like small white crystals and quickly dissolves in water. The substance is capable of absorbing moisture from the air and, even under good storage conditions, quickly cakes, so it must be stored and transported in airtight packaging. To reduce hygroscopicity, some manufacturers press calcium nitrate into granules with a water-repellent shell, but even this helps little. The substance is mainly used on acidic soils, as it alkalizes.
The fertilizer is well suited for any vegetables except potatoes. This is the only composition containing calcium in a water-soluble form, so it is widely used in greenhouses and greenhouses for root and foliar feeding of cucumbers and tomatoes. Calcium nitrate, which rapidly absorbs water, is of little use for application to the soil. It is also not recommended to mix it with other fats, as the mixture will turn into a doughy mass.
The disadvantage of all saltpeter is its low nitrogen content. The costs of transportation and purchase may not be justified by the increase in yield.
Ammonium
Substances in this group contain nitrogen in the form of ammonium (NH4), which gives them good solubility in water. The main advantage of ammonium fertilizers is that nitrogen in the form of ammonium is readily available to plants. It is moderately mobile in the soil, that is, it is practically not washed out during rains and watering.
Ammonium fertilizers can be used in the fall - they will not be washed out of the soil by melt water in the spring, and will not turn into an inaccessible form during the winter. Experts recommend using ammonium fertilizers as basic fertilizers in the fall or spring, and nitrate fertilizers as top dressing.
Ammonium sulfate
Ammonium sulfate (ammonium sulfate) – formula (NH4)2SO4. The product contains two substances necessary for plants - nitrogen and sulfur. The fertilizer is of the highest grade (21% nitrogen) and technical (19% nitrogen).
Ammonium sulfate is produced synthetically and as a by-product of iron and steel industry. You can distinguish synthetic fertilizer from coke fertilizer by color. Synthetic is snow-white, while coke-chemical contains impurities, therefore it is colored greyish, bluish or reddish. The fertilizer almost does not absorb water from the air, so it cakes a little.
The product contains up to 24% sulfur. Onions, garlic, rapeseed, and mustard especially need this microelement. The characteristic smell of these plants is largely due to the sulfur they contain. When grown in soils with a high sulfur content or when ammonium sulfate is added, onions and garlic grow more fragrant and are less damaged by pests and diseases. After onions, cabbage, broccoli and canola have the highest sulfur requirements, followed by legumes and grains.
Sodium ammonium sulfate
The substance contains 17% nitrogen and 8% sodium. Externally, the fertilizer consists of white, dark gray or yellow crystals.
It is used in the same way as regular ammonium sulfate, but due to its sodium content it is more advisable to apply it under root vegetables.
Ammonium chloride
The chemical formula of the fertilizer is NH4Cl. It is a by-product of soda production. Contains 25% nitrogen. The composition contains up to 67% chlorine, which is harmful to plants, so it is not used for feeding crops that are sensitive to this element: grapes, tobacco, citrus fruits.
Ammonium chloride acidifies the soil. With a one-time application of fertilizer, the soil will not become worse, but with systematic use there is a risk of acidification of the beds.
Ammonia liquid fertilizers
Liquid fertilizers are readily available for plants. Recently, the production of liquid ammonia fertilizers has been increasing.
Chemical formula of liquid ammonia NH3. Fertilizer is obtained by exposing ammonia gas to high pressure. The result is a colorless liquid with a boiling point of 34 degrees. It cannot be stored in open containers as it evaporates quickly. Liquid ammonia is stored and transported in steel cylinders and tanks.
Ammonia water (ammonia aqueous) is ammonia dissolved in water. The fertilizer is available in two varieties. The first contains 20.5% nitrogen, the second - no less than 18%. Ammonia water is a colorless liquid with the smell of ammonia. It can only be stored and transported in sealed containers, since nitrogen easily evaporates.
Liquid nitrogen fertilizers are not for hobbyists. Their consumers are large agricultural enterprises.
Liquid fertilizers are much cheaper than solid ones, despite the fact that their transportation and storage require significant costs. At enterprises, only specially trained workers are allowed to work with liquid fertilizers. Ordinary summer residents and lovers of indoor flowers also use liquid nitrogen fertilizer - ammonia.
Ammonium-nitrate
Fertilizers of this type contain nitrogen in two forms at once: NO3 (nitrates) and NH4 (ammonium). Thus, in percentage terms they contain more nitrogen than the previous ones.
Ammonium nitrate
Ammonium nitrate is the main nitrogen fertilizer. Approximately 55-60% of all nitrogen compounds used in agriculture are ammonium nitrate. The fertilizer contains 34% nitrogen. It looks like white crystals or granules of various shapes. The substance absorbs water from the air, so it is stored in dry rooms in waterproof packaging.
The product is fire and explosive. It should be kept away from open flames and explosives. Ammonium nitrate does not contain ballast and dissolves without residue. Acts on the soil as an acidifier.
Calcium ammonium nitrate
The product is obtained by mixing ammonium nitrate with lime, chalk or dolomite. The fertilizer does not acidify the soil, is not explosive, and does not caking. Contains 22-26% nitrogen and 17-27% calcium carbonate, suitable for systematic use on soils requiring liming.
Amide - in these fertilizers nitrogen is in the form of (NH2)2. In Russia, only one fertilizer of this class is produced; even novice summer residents know it. This is urea (carbamide). Chemical formula products CO(NH2)2, nitrogen content 46%. Urea is produced from ammonia under high pressure. As a result, small white crystals are formed that are highly soluble in water. When stored correctly, urea does not cake.
Urea should not be spread over the surface of the soil, as the nitrogen will evaporate. It must be immediately embedded in the soil.
Urea is one of the best nitrogen compounds. It can be used on all soils and for any crops as the main fertilizer or top dressing, including foliar feeding. In addition, urea is used in livestock farming as a feed additive.
Phosphorus
Any plant needs phosphorus. When this element is deficient, the crop slows down and leaves turn green, purple or red. Then dark spots appear along the edges of the plates. Signs of phosphorus starvation appear primarily on the lower leaves. With acute phosphorus starvation, flowering and ripening are noticeably delayed. Plants especially urgently need phosphorus in the early stages of development, when their small root system cannot yet absorb a sufficient amount of the element from the soil.
Usually the soil contains a lot of phosphorus, but it is included in compounds that are inaccessible to plants. Therefore, phosphorus fertilizing is urgently needed for all agricultural crops. Russia has the world's richest deposit of apatite ore, a raw material for the production of phosphate fertilizers. Phosphorus-containing fertilizers listed in the table are produced from apatites.
Types of phosphate fertilizers:
The main phosphorus fertilizer for summer residents is superphosphates - simple and double. Superphosphate may contain additional beneficial microelements:
Gardeners believe that superphosphate is poorly soluble in water. In fact, the phosphorus contained in this fertilizer passes into water quite easily, and the gray insoluble granules are ordinary gypsum. The average application rate of double superphosphate is 40-50 g/m2.
There is more gypsum in simple superphosphate than in double superphosphate, so it is better to apply it to crops that react positively to calcium, for example, legumes. Superphosphate must be incorporated into the soil when planting, directly under the roots. In the top layer of soil it quickly dries out and becomes inaccessible to plants.
Potash
Potassium increases plant resistance to drought and cold. The element accelerates the flow of sugar from leaves into fruits and underground organs, so potassium fertilizers make fruits, berries, and root vegetables sweeter. After potassium feeding, the stems become resistant to lodging. Of the fruits and vegetables, potatoes need potassium the most - their tubers contain 2.4% potassium in terms of dry matter. For comparison, heads of cabbage contain 13 times less potassium – 0.18%.
Plants receiving potassium 3-5 times less than normal show signs of starvation:
Potassium usually accumulates in parts of plants that are not used for food: leaves, straw. It is enough to add unnecessary plant matter back into the soil, and next year the plants will be well supplied with potassium.
Types of potashfertilizers:
Chlorine in potash fertilizers is undesirable. Chlorine-free options are preferable. The most popular chlorine-free potassium fertilizer is potassium sulfate, a product of the processing of natural minerals. The fertilizer does not caking, is suitable for any soil, for all crops. The production of potassium sulfate is not cheap, so in stores it is more expensive than other potassium compounds.
Potassium magnesium contains potassium and magnesium in equal quantities. The fertilizer is ideal for crops that absorb a lot of magnesium (potatoes, clover). After feeding strawberries with potassium magnesia, the plantation suffers less from strawberry mites and other sucking insects, and the number of berries with rot is reduced. Fertilizing will be most beneficial on poor sandy and sandy loam soils.
Average application rates:
Complex
Complex fertilizers include several chemical elements necessary for the plant. Fertilizers of this variety are more concentrated, provide plants with several nutrients at once in the required ratio, and save time and labor costs.
Types of complex fertilizers:
Name | Nutrient content in percentage | Note |
||
| Nitrogen | Phosphorus | Potassium | ||
| 9-11 | – | Inexpensive nitrogen-phosphorus fertilizer, highly soluble in water, does not caking |
||
Diammofos | 19-21 | – | Highly concentrated, physiologically neutral fertilizer. Contains nitrogen and phosphorus in a readily available water-soluble form. One of the best complex nutritional compositions |
|
Nitroammofoska | 13-18 | 17-20 | ||
Diammofoska | 9-10 | 25-26 | ||
Azofoska | 16 | 16 | ||
Potassium nitrate | 13-15 | 39-45 | Chlorine-free nitrogen-potassium fertilizer, does not contain phosphorus. Used mainly for potatoes and grapes | |
Combined application of fertilizers
Do not mix mineral fertilizers randomly. Chemical reactions occur between them that can reduce the solubility of fats or lead to loss of nutrients.
It's better not to mix:
Mineral fertilizers can be used in any period, except winter, on any soil and for any crops. They provide a significant increase in yield, but do not improve its physical properties. Experienced gardeners use mineral fertilizers together with organic matter, which benefits both plants and soil.
Nitric acid HNO3 in its pure form is a colorless liquid with a pungent suffocating odor. It is formed in small quantities during lightning discharges and is present in rainwater. Under the influence of light, nitric acid partially decomposes with the release of NO2 and due to this acquires a light brown color: 4HNO3 = 4NO2 + 2H2O + O2. Nitric acid is one of the...
When solid nitrates are heated, they all decompose with the release of oxygen (ammonium nitrate is an exception), and they can be divided into four groups. The first group consists of alkali metal nitrates, which, when heated, decompose into nitrites and oxygen: 2KNO3 = 2KNO2 + O2. The second group consists of the majority of nitrates (from alkaline earth metals to copper inclusive), decomposing into metal oxide, NO2 and oxygen: 2Сu(NO3)2 = 2СuО + 4NO2 + O2. The third group consists of nitrates of the heaviest metals (AgNO3 and Нg(NO3)2 ), decomposing to free metal, NO2 and oxygen: Hg(NO3)2 = Hg + 2NO2 + O2. The fourth “group” is ammonium nitrate: NH4NO3 = N2O + 2H2O.
Nitrous acid HNO2 belongs to weak acids (K = 6.10-4 at 25 °C), is unstable and is known only in dilute solutions in which the equilibrium 2HNO2 NO + NO2 + H2O occurs. Nitrites, unlike the acid itself, are stable even when heated. The exception is crystalline ammonium nitrite, which when heated decomposes into free nitrogen and water.
Of the three phosphoric acids, orthophosphoric acid H3PO4 (often called simply phosphoric acid) is of greatest practical importance - a white solid, highly soluble in water. In an aqueous solution it dissociates stepwise. As a tribasic acid, phosphoric acid forms three types of salts: dihydrogen phosphates (NaH2PO4); hydrophosphates (Na2HPO4); phosphates (Na3PO4). All dihydrogen phosphates are soluble in water. Of the hydrophosphates and phosphates, only alkali metal and ammonium salts are soluble in water. Salts of phosphoric acid are valuable mineral fertilizers. The most common among them are superphosphate, precipitate and phosphate rock. Simple superphosphate is a mixture of calcium dihydrogen phosphate Ca(H2PO4)2 and “ballast” CaSO4. It is obtained by treating phosphorites and apatites with sulfuric acid. When mineral phosphates are treated with phosphoric acid, double superphosphate Ca(H2PO4)2 is obtained. When phosphoric acid is quenched with lime, the precipitate CaHPO4.2H2O is obtained. Complex fertilizers (i.e., containing both nitrogen and phosphorus; or nitrogen, phosphorus and potassium) are important. Of these, the most famous is ammophos - a mixture of NH4H2PO4 and (NH4)2HPO4.
Under normal conditions, it is a colorless gas with a pungent odor (the smell of “ammonia”); liquefies at -33.4 °C and solidifies at -77.7 °C. The ammonia molecule has the shape of a pyramid; in liquid ammonia, NH3 molecules are connected by hydrogen bonds, thereby causing an abnormally high boiling point. Polar NH3 molecules are very soluble in water (700 volumes of NH3 in one volume of H2O)…
Phosphorus forms two chlorides: phosphorus trichloride PCl3 and phosphorus pentachloride PCl5. Phosphorus trichloride is prepared by passing chlorine over the surface of white phosphorus. In this case, phosphorus burns with a pale green flame, and the resulting phosphorus chloride condenses as a colorless liquid. Phosphorus trichloride is hydrolyzed by water to form phosphorous acid and hydrogen chloride: PCl3 + 3H2O = H3PO3 + 3HCl. Phosphorus pentachloride can be obtained in the laboratory...
In oxides, the oxidation state of nitrogen varies from +1 to +5. The oxides N2O and NO are colorless gases, nitrogen oxide (IV) NO2 is a brown gas, called “fox tail” in the industry. Nitrogen oxide (III) N2O3 is a blue liquid, nitrogen oxide (V) N2O5 under normal conditions is transparent colorless crystals. The trivial name nitric oxide (I) is often used...
Phosphoric anhydride P2O5 (the “simplest” formula) is the most stable phosphorus oxide under normal conditions. It is a solid white substance with the composition P4O10. Phosphorous anhydride is described by the simplest formula P2O3 and the true formula P4O6. It has been shown that phosphorus in P4O6 is coordinatively unsaturated and therefore unstable. The interaction of P4O6 with hot water leads to disproportionation of P4O6 + 6H2O = PH3 + 3H3PO4; Gaseous HCl decomposes P4O6: P4O6 + 6HCl = 2H3PO3 + 2PCl3. P4O10 actively interacts with water and also takes it away from other compounds, forming, depending on the conditions, either metaphosphoric HPO3, orthophosphoric H3PO4, or pyrophosphoric H4P2O7 acids. That is why P4O10 is widely used as a desiccant for various substances from water vapor.