Presentation on nitrogen and phosphorus. Presentation on the topic "Nitrogen and phosphorus-p-elements of the VA-group"
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1. I warn you in advance: I am unbreathable! But everyone seems not to hear And they constantly breathe me. 2. I am a luminiferous element. I'll light a match for you in a moment. They will burn me - and under water my oxide will become acid.
The position of nitrogen and phosphorus in the Periodic system
Characteristics of nitrogen and phosphorus. properties of nitrogen.
Five famous chemists of the XVIII century. gave a certain non-metal, which in the form of a simple substance is a gas and consists of diatomic molecules, five different names. - "poisonous air" - "dephlogisticated air" - "spoiled air" - "suffocating air" - "lifeless air" In 1772, the Scottish chemist, botanist and physician Daniel Rutherford In 1772, the English chemist Joseph Priestley In 1773, the Swedish apothecary chemist Carl Scheele In 1774, the English chemist Henry Cavendish In 1776, the French chemist Antoine Lavoisier
FINDING OF NITROGEN IN NATURE: in a free state in the atmosphere
FINDING OF NITROGEN IN NATURE: in the form of inorganic compounds In small quantities in the soil: in the form of ammonium salts and nitrates. organic Nitrogen of plants and animals (Nucleic acids, proteins)
SIGNS OF COMPARISON NITROGEN PHOSPHORUS POSITION IN PSCE STRUCTURE OF THE ATOM Number of electrons in an atom 7, protons in the nucleus 7, number of neutrons in the nucleus 7 Electronic circuit: 1s 2 2s 2 2p 3 OXIDATION DEGREES 3 period V group main subgroup Ordinal number 15; relative atomic mass 31 2 period V group Main subgroup Ordinal number 7; relative atomic mass 14 P +15) 2) 8) 5 The number of electrons in an atom 15, protons in the nucleus 15, the number of neutrons in the nucleus 16 Electronic circuit: 1s 2 2s 2 2p 6 3s 2 2p 3 N + 7) 2) 5 + 3, +5 , -3 +1,+2,+3,+4, +5 , -3
Determine the oxidation states of nitrogen in the compounds: HNO 3, NH 3, NO, KNO 2, NO 2, N 2 O, HNO 2 s.o. -3 0 +1 +2 +3 +4 +5 compound NH 3 N 2 N 2 O NO N 2 O 3 NO 2 N 2 O 5
STRUCTURE OF THE MOLECULES N N N N BOND: - COVALENT NON-POLE - TRIPLE - STRONG MOLECULE: - VERY STABLE - LOW REACTIVITY 1 3 4 2
N 2 Physical properties: V, C, Z, M slightly lighter than air, t bale = -196 0 C, t pl = -210 0 C
In industry, nitrogen is obtained by distillation of air, in the laboratory - by thermal decomposition of compounds (most often NH 4 NO 2): NH 4 NO 2 → N 2 + 2 H 2 O Phosphorus is obtained by calcining calcium phosphate with coal and sand in electric furnaces at 1500 0 С : 2Ca 3 (PO 4) 2 + 10C + 6SiO 2 → 6CaSiO 3 + 10CO + P 4 Preparation.
Chemical properties of nitrogen Phosphorus with metals at room t reacts with Li 6 Li + N 2 = 2 Li 3 N at high t - with others Me 2Al + N 2 = 2AlN 3Mg + N 2 = Mg 3 N 2 reacts with Me 3 when heated Ca + 2 P \u003d Ca 3 P 2 with oxygen at very high t (about 3000 ° C) N 2 + O 2 \u003d 2 NO white phosphorus ignites spontaneously, and red burns when heated 4 P + 5 O 2 \u003d 2 P 2 O 5 with hydrogen in the presence of a catalyst at high pressure and t N 2 + 3 H 2 = 2 NH 3
Applications Ammonia production Creation of inert atmosphere Creation of low temperatures Saturation of steel surface to increase strength Liquid nitrogen in medicine Synthesis of ammonia Fertilizer industry Synthesis of nitric acid Creation of inert atmosphere N2
Questions for self-control The gas is colorless, tasteless and odorless The molecule is diatomic The content in air is 78% In the laboratory it is obtained by decomposition of KMnO 4 and H 2 O 2 In industry - from liquid air It is chemically inactive It interacts with almost all simple substances The processes of respiration and photosynthesis are associated with it Is an integral part of proteins Participates in the cycle of substances in nature
CHECK YOURSELF O 2 1, 2, 4, 5, 7, 8, 10. "5" N 2 1, 2, 3, 5, 6, 9, 10. "5" 1-2 errors "4" 3-4 errors « 3 » 5 errors and more « 2 » On the example of information about nitrogen, give arguments in favor of two points of view: 1. Nitrogen - "lifeless" 2. Nitrogen - the main element of life on Earth.
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In the VA-group of the periodic system, the non-metals nitrogenN and phosphorus P, the semi-metal arsenic As, as well as antimony Sb and bismuth Bi, which are classified as non-metals, are located.
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The atoms of the elements of the VA group have 5 electrons on the outer electron layer. The electronic configuration of their outer electron layer is ns2np3, for example: nitrogen - 2s2p3, phosphorus - 3s23p3.
In chemical compounds, nitrogen and phosphorus atoms can exhibit oxidation states from -3 to +5.
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nitrogen in nature
Nitrogen is denoted by the symbol N (lat. Nitrogenium, i.e. "giving birth to saltpeter"). The simple substance nitrogen (N2) is a rather inert gas under normal conditions, colorless, tasteless and odorless. Nitrogen, in the form of diatomic N2 molecules, makes up most of the atmosphere, where its content is 78.084% by volume (that is, about 3.87 1015 tons).
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nitrogen in space
Outside the Earth, nitrogen is found in gaseous nebulae, the solar atmosphere, on Uranus, Neptune, interstellar space, and others. Nitrogen is the 4th most abundant element in the solar system (after hydrogen, helium, and oxygen).
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Phosphorus in nature
Phosphorus occurs naturally in the form of phosphates. Thus, calcium phosphate Ca3(PO4)2 is the main component of the mineral apatite. Phosphorus is found in all parts of green plants, and even more in fruits and seeds. Contained in animal tissues, is part of proteins and other essential organic compounds (ATP, DNA), is an element of life. Apatite
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The simple substance nitrogen consists of diatomic N2 molecules. In the N2 molecule, the nitrogen atoms are linked by a triple covalent nonpolar bond. The triple bond energy is high and amounts to 946 kJ/mol. Therefore, bond breaking and the formation of nitrogen atoms and molecules occurs only at temperatures above 3000°C. The high bond strength in molecules determines the chemical inertness of nitrogen.
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In the free state, phosphorus forms several allotropic modifications, which are called white, red and black phosphorus.
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In the simplest P4 molecule, each of the four phosphorus atoms is covalently bonded to the other three. White phosphorus consists of such tetrahedral-shaped molecules. Cast in an inert atmosphere in the form of sticks (ingots), it is stored in the absence of air under a layer of purified water or in special inert media.
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Chemically, white phosphorus is extremely active! For example, it is slowly oxidized by air oxygen already at room temperature and glows (pale green glow). The phenomenon of this kind of glow due to chemical oxidation reactions is called chemiluminescence (sometimes erroneously phosphorescence). White phosphorus is highly toxic. The lethal dose of white phosphorus for an adult male is 0.05-0.1 g.
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Red phosphorus has an atomic polymer structure in which each phosphorus atom is bonded to three other atoms by covalent bonds. Red phosphorus is not volatile, insoluble in water, and non-toxic. It is used in the manufacture of matches.
In the light and when heated to 300 ° C without air, white phosphorus turns into red phosphorus.
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When heated under a pressure of about 1200 times greater than atmospheric pressure, white phosphorus turns into black phosphorus, which has an atomic layered crystal lattice. Black phosphorus is similar to metal in its physical properties: it conducts electricity and glistens. Outwardly, it is very similar to graphite. Black phosphorus is the chemically least active form of phosphorus.
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In 1830, the French chemist Charles Soria invented phosphorus matches, which consisted of a mixture of barthollet salt, white phosphorus and glue. These matches were very flammable, because they caught fire even from mutual friction in the box and when rubbing against any hard surface, for example, the sole of a boot. Because of white phosphorus, they were poisonous. In 1855, the Swedish chemist Johan Lundström applied red phosphorus to the surface of sandpaper and replaced white phosphorus in the match head with it. Such matches were no longer harmful to health, they easily ignited on a pre-prepared surface and practically did not ignite spontaneously. Johan Lundström patents the first "Swedish match", which has survived almost to this day. In 1855, Lundström's matches were awarded a medal at the World Exhibition in Paris. Later, phosphorus was completely removed from the composition of the match heads and remained only in the spread (grater). With the development of the production of "Swedish" matches, the production of matches using white phosphorus was banned in almost all countries.
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The simplest substance, nitrogen N2, is chemically inactive and, as a rule, enters into chemical reactions only at high temperatures. The oxidizing properties of nitrogen are manifested in the reaction with hydrogen and active metals. So, hydrogen and nitrogen combine in the presence of a catalyst at high temperature and high pressure, forming ammonia:
Of the metals, under normal conditions, nitrogen reacts only with lithium, forming lithium nitride:
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The oxidizing properties of phosphorus are manifested when it interacts with the most active metals:
The reducing properties of nitrogen and phosphorus are manifested when they interact with oxygen. So, nitrogen reacts with oxygen at a temperature of about 3000˚С, forming nitric oxide (II):
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Phosphorus is also oxidized by oxygen, thus exhibiting reducing properties. But different modifications of phosphorus have different chemical activity. For example, white phosphorus is easily oxidized in air at room temperature to form phosphorus(III) oxide:
Oxidation of white phosphorus is accompanied by luminescence. White and red phosphorus ignite when ignited and burn with a dazzlingly bright flame with the formation of white smoke of phosphorus (IV) oxide:
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Burning white phosphorus
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The most chemically active, toxic and combustible white phosphorus. Therefore, it is very often used in incendiary bombs. Unfortunately, phosphorus ammunition is also used in the 21st century!
During the siege of Sarajevo, phosphorus shells were used by Bosnian Serb artillery. In 1992, such shells burned down the building of the Institute of Oriental Studies, as a result of which many historical documents were destroyed. - in 2003-2004 they were used by British intelligence services in the vicinity of Basra in Iraq. - in 2004, the United States used against the guerrilla underground in Iraq in the battle for Fallujah. in the summer of 2006, during the Second Lebanon War, artillery shells with white phosphorus were used by the Israeli army. in 2009, during Operation Cast Lead in the Gaza Strip, the Israeli army used ammunition containing white phosphorus, which is allowed by international law. Since 2009 Palestinian terrorists have been loading their missiles with white phosphorus.
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The appearance of wandering lights in old cemeteries and swamps is caused by the ignition of phosphine PH3 and other compounds containing phosphorus in air. In air, the products of the combination of phosphorus with hydrogen spontaneously ignite with the formation of a luminous flame and droplets of phosphoric acid, a product of the interaction of phosphorus (V) oxide with water. These droplets create a blurry outline of the "ghost".
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The main application of nitrogen is the production of ammonia. Nitrogen is also used to create an inert environment in the drying of explosives and in the storage of valuable paintings and manuscripts. In addition, electric incandescent lamps are filled with nitrogen.
Application of simple substances Production of ammonia Most modern lamps are filled with chemically inert gases. Mixtures of nitrogen N2 with argon Ar are the most common due to their low cost.
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The presentation on the topic "Phosphorus" can be downloaded absolutely free of charge on our website. Project subject: Chemistry. Colorful slides and illustrations will help you keep your classmates or audience interested. To view the content, use the player, or if you want to download the report, click on the appropriate text under the player. The presentation contains 29 slide(s).
Presentation slides
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Material for repetition and preparation for the GIA Chemistry teacher of the Municipal Educational Institution "Gymnasium No. 1", Saratov Shishkina I.Yu.
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Introduction……………………………………………………………………………. The history of the development of phosphorus………………………………………………………... Natural compounds and the production of phosphorus……………………………………... Chemical properties ……………………………………………………………… Allotropic changes………………………………………………………….. a) white…………………………………………………………………………….. b) red………………………………… …………………………… c) black……………………………………………………………………………. Phosphorus oxides……………………………………………………………………… Orthophosphoric acid……………………………………………………… ……... Orthophosphates………………………………………………………………………. Phosphorus in the human body…………………………………………………….. Matches……………………………………………………………… …………………. Phosphate fertilizers…………………………………………………………….. Conclusion…………………………………………………………… ………………. 1. The value of phosphorus………………………………………………………………….. 2. The use of phosphorus…………………………………………… ………………… Bibliography………………………………………………..
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Introduction:
The fifth group of the Periodic system includes two typical elements nitrogen and phosphorus - and subgroups of arsenic and vanadium. There is a significant difference in properties between the first and second typical elements. In the state of simple substances, nitrogen is a gas, and phosphorus is a solid. These two substances received a wide range of applications, although when nitrogen was first isolated from the air it was considered a harmful gas, and a lot of money was made from the sale of phosphorus (phosphorus was valued for its ability to glow in the dark).
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The history of the discovery of phosphorus
Ironically, phosphorus has been discovered several times. And every time they got it from ... urine. There are references that the Arab alchemist Alhild Bekhil (XII century) discovered phosphorus during the distillation of urine mixed with clay, lime and coal. However, the date of discovery of phosphorus is considered to be 1669. The Hamburg amateur alchemist Henning Brand, a bankrupt merchant who dreamed of improving his affairs with the help of alchemy, processed a wide variety of products. Assuming that physiological products might contain the "primordial matter" thought to be the basis of the Philosopher's Stone, Brand became interested in human urine. He collected about a ton of urine from the soldiers' barracks and evaporated it to a syrupy liquid. This liquid he distilled again and obtained a heavy red "urinary oil", which was distilled to form a solid residue. Heating the latter, without access to air, he noticed the formation of white smoke, which settled on the walls of the vessel and shone brightly in the darkness. Brand named the substance he received phosphorus, which in Greek means "light-bearer". For several years, the "preparation recipe" for phosphorus was kept in the strictest confidence and was known only to a few alchemists. Phosphorus was discovered for the third time by R. Boyle in 1680. In a somewhat modified form, the old method of obtaining phosphorus was also used in the 18th century: a mixture of urine with lead oxide (PbO), common salt (NaCl), potash (K2CO3) and coal (C) was heated. Only by 1777, K.V. Scheele developed a method for obtaining phosphorus from animal horns and bones.
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Natural compounds and obtaining phosphorus
In terms of prevalence in the earth's crust, phosphorus is ahead of nitrogen, sulfur and chlorine. Unlike nitrogen, phosphorus, due to its high chemical activity, occurs in nature only in the form of compounds. The most important minerals of phosphorus are apatite Ca5 (PO4) 3X (X is fluorine, less often chlorine and a hydroxyl group) and phosphorite, the basis of which is Ca3 (PO4) 2. The largest apatite deposit is located on the Kola Peninsula, in the region of the Khibiny Mountains. Phosphorite deposits are located in the Karatau mountains, in the Moscow, Kaluga, Bryansk regions and in other places. Phosphorus is part of some protein substances contained in the generative organs of plants, in the nervous and bone tissues of animal and human organisms. Brain cells are especially rich in phosphorus. Today, phosphorus is produced in electric furnaces by reducing apatite with coal in the presence of silica: Ca3(PO4)2+3SiO2+5C 3CaSiO3+5CO+2P Phosphorus vapor at this temperature consists almost entirely of P2 molecules, which condense into P4 molecules upon cooling.
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Chemical properties
The electronic configuration of the phosphorus atom is 1s22s22p63s23p3 The outer electron layer contains 5 electrons. The presence of three unpaired electrons on the external energy level explains the fact that in the normal, unexcited state, the phosphorus valency is 3. But at the third energy level there are vacant cells of d-orbitals, therefore, upon transition to an excited state, 3S-electrons will separate, go to the d sublevel , which leads to the formation of 5 unpaired elements. Thus, the valency of phosphorus in the excited state is 5. In compounds, phosphorus usually exhibits an oxidation state of +5 (P2O5, H3PO4), less often +3 (P2O3, PF3), -3 (AlP, PH3, Na3P, Mg3P2).
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The transition of the phosphorus atom to an excited state
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White phosphorus
The white modification of phosphorus resulting from vapor condensation has a molecular crystal lattice, in the nodes of which P4 molecules are dislocated. Due to the weakness of intermolecular forces, white phosphorus is volatile, fusible, cut with a knife and dissolved in non-polar solvents, such as carbon disulfide. White phosphorus is a highly reactive substance. It reacts vigorously with oxygen, halogens, sulfur and metals. Oxidation of phosphorus in air is accompanied by heating and glow. Therefore, white phosphorus is stored under water, with which it does not react. White phosphorus is highly toxic. About 80% of the total production of white phosphorus goes to the synthesis of pure phosphoric acid. It, in turn, is used to produce sodium polyphosphates (they are used to reduce the hardness of drinking water) and food phosphates. The rest of the white phosphorus is used to create smoke-forming substances and incendiary mixtures. Safety engineering. In the production of phosphorus and its compounds, special precautions are required, because white phosphorus is a strong poison. Prolonged work in an atmosphere of white phosphorus can lead to disease of bone tissue, loss of teeth, necrosis of jaw areas. When ignited, white phosphorus causes painful burns that do not heal for a long time. White phosphorus should be stored under water, in airtight containers. Burning phosphorus is extinguished with carbon dioxide, CuSO4 solution or sand. Burnt skin should be washed with KMnO4 or CuSO4 solution. The antidote for phosphorus poisoning is a 2% solution of CuSO4. During long-term storage, as well as when heated, white phosphorus turns into a red modification (it was first received only in 1847). The 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, and compared to white phosphorus, they are less reactive and have a polymer structure: these are P4 tetrahedra connected to each other in endless chains.
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Red and black phosphorus
Red phosphorus is used in metallurgy, the production of semiconductor materials and incandescent lamps, and is used in match production. The most stable modification of phosphorus is black phosphorus. It is obtained by allotropic transformation of white phosphorus at t=2200C and high pressure. In appearance, it resembles graphite. The crystal structure of black phosphorus is layered, consisting of corrugated layers (Fig. 2). Black phosphorus is the least active modification of phosphorus. When heated without access to air, it, like red, passes into vapor, from which it condenses into white phosphorus.
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An experiment illustrating the transition of red phosphorus to white
1-molecules of white phosphorus; 2-crystalline. black phosphorus lattice
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Phosphorus (V) oxide - Р2О5
Phosphorus forms several oxides. The most important of them is phosphorus oxide (V) P4O10. Often its formula is written in a simplified form - P2O5. The structure of this oxide retains the tetrahedral arrangement of phosphorus atoms. White crystals, t melt = 5700°C, boil t = 6000°C, ρ = 2.7 g/cm3. Has several modifications. In vapor it consists of P4H10 molecules, it is very hygroscopic (used as a desiccant for gases and liquids). Preparation: 4P + 5O2 = 2P2O5 Chemical properties All chemical properties of acidic oxides: reacts with water, basic oxides and alkalis 1) P2O5 + H2O = 2HPO3 (metaphosphoric acid) P2O5 + 2H2O = H4P2O7 (pyrophosphoric acid) acid) 2) P2O5 + 3BaO =Ba3(PO4)2 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.
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Orthophosphoric acid.
Several acids containing phosphorus are known. The most important of them is orthophosphoric acid H3PO4. Anhydrous orthophosphoric acid is a light transparent crystals, deliquescent in air at room temperature. Melting point 42.35°C. With water, phosphoric acid forms solutions of any concentration.
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Physical properties of H3PO4
Orthophosphoric acid in its pure form under normal conditions is colorless rhombic crystals, melting at a temperature of 42.3 ° C. However, chemists rarely encounter such an acid. Much more often they deal with H3PO4 * 0.5 H2O hemihydrate, which precipitates in the form of colorless hexagonal prisms when concentrated aqueous solutions of phosphoric acid are cooled. The melting point of the hemihydrate is 29.3°C. Pure H3PO4 after melting forms a viscous oily liquid with low electrical conductivity and greatly reduced diffusivity. These properties, as well as a detailed study of the spectra, show that the H3PO4 molecules in this case are practically not dissociated and are united by strong hydrogen bonds into a single macromolecular structure. As a rule, molecules are connected to each other by one, rarely two, and very rarely three hydrogen bonds. If the acid is diluted with water, then its molecules are more likely to form hydrogen bonds with water than with each other. Because of such "sympathy" for water, acid mixes with it in any relationship. The hydration energy here is not as high as that of sulfuric acid; therefore, the heating of H3PO4 upon dilution is not so strong and dissociation is less pronounced. According to the first stage of dissociation, phosphoric acid is considered an electrolyte of medium strength (25 - 30%), according to the second - weak, according to the third - very weak.
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Chemical properties of H3PO4
When neutralizing phosphoric acid with alkalis, salts are formed: dihydrophosphates, hydrophosphates, and also phosphates, for example:
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Phosphorus in the human body
In a human body weighing 70 kg. Contains about 780 g of phosphorus. In the form of calcium phosphates, phosphorus is present in the bones of humans and animals. It is also included in the composition of proteins, phospholipids, nucleic acids; phosphorus compounds are involved in energy metabolism (adenisine triphosphoric acid, ATP). The daily need of the human body for phosphorus is 1.2 g. We consume the main amount of it with milk and bread (100 g of bread contains approximately 200 mg of phosphorus). Fish, beans and some types of cheese are the richest in phosphorus. Interestingly, for proper nutrition, it is necessary to maintain a balance between the amount of phosphorus and calcium consumed: the optimal ratio in these food elements is 1.5/1. An excess of phosphorus-rich food leads to leaching of calcium from the bones, and with an excess of calcium, urolithiasis develops.
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The incendiary surface of the matchbox is coated with a mixture of red phosphorus and glass powder. The composition of the match head includes oxidizing agents (PbO2, KClO3, BaCrO4) and reducing agents (S, Sb2S3). With friction from the incendiary surface, the mixture applied to the match ignites. The first phosphorus matches - with a white phosphorus head - were created only in 1827. 6P + 5KCLO3 = 5KCL + 3P2O5 Such matches caught fire when rubbed against any surface, which often led to fires. In addition, white phosphorus is highly toxic. Cases of poisoning with phosphorus matches are described, both due to careless handling and for the purpose of suicide: for this it was enough to eat a few match heads. That is why phosphorus matches were replaced by safe ones, which serve us faithfully to this day. Industrial production of safety matches began in Sweden in the 60s. XIX century.
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The value of phosphorus
Phosphoric acid is of great importance as one of the most important components of plant nutrition. Phosphorus is used by plants to build their most vital parts, seeds and fruits. Orthophosphoric acid derivatives are very necessary not only for plants, but also for animals. Bones, teeth, shells, claws, needles, spikes in most living organisms consist mainly of calcium orthophosphate. In addition, phosphoric acid, forming various compounds with organic substances, is actively involved in the metabolism of a living organism with the environment. As a result, phosphorus derivatives are found in bones, brain, blood, muscle and connective tissues of human and animal organisms. There is especially a lot of phosphoric acid in the composition of nerve (brain) cells, which allowed A.E. Fersman, a well-known geochemist, called phosphorus an "element of thought." Very negatively (animal disease rickets, anemia, etc.) affects the state of the body by lowering the content of phosphorus compounds in the diet or introducing them in an indigestible form.
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The use of phosphorus
Orthophosphoric acid is currently widely used. Its main consumer is the production of phosphate and combined fertilizers. For these purposes, about 100 million tons of phosphorus-containing ore are annually mined all over the world. Phosphorus fertilizers not only help to increase the yield of various crops, but also give plants winter hardiness and resistance to other adverse climatic conditions, create conditions for faster ripening of crops in areas with short vegetative period. They also have a beneficial effect on the soil, contributing to its structuring, the development of soil bacteria, changing the solubility of other substances contained in the soil and suppressing some of the resulting harmful organic substances. A lot of orthophosphoric acid is consumed by the food industry. The fact is that dilute phosphoric acid tastes very pleasant and its small additions to marmalades, lemonades and syrups significantly improve their taste. Some salts of phosphoric acid have the same property. Calcium hydrogen phosphates, for example, have long been included in baking powders, improving the taste of rolls and bread. Other industrial applications of phosphoric acid are also of interest. For example, it has been observed that the impregnation of wood with the acid itself and its salts makes the wood incombustible. On this basis, fire-retardant paints, non-combustible phospho-wood boards, non-combustible phosphate foam and other building materials are now being produced. Various salts of phosphoric acid are widely used in many industries, in construction, in various fields of technology, in public utilities and everyday life, for protection against radiation, for softening water, combating boiler scale and manufacturing various detergents. Phosphoric acid, condensed acids and dehydrogenated phosphates serve as catalysts in the processes of dehydration, alkylation and polymerization of hydrocarbons. A special place is occupied by organophosphorus compounds as extractants, plasticizers, lubricants, gunpowder additives and absorbents in refrigeration units. Acid alkyl phosphate salts are used as surfactants, antifreezes, special fertilizers, latex anticoagulants, etc. Acid alkyl phosphates are used for extraction processing of uranium ore liquors.
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Phosphorus 1. Write the electronic formula of the phosphorus atom. Explain what happens to the electronic configuration of an atom when it exhibits the highest oxidation state. 2. What oxidation states can phosphorus exhibit in compounds? Give examples of these compounds. Write the electronic formula of the phosphorus atom in the +3 oxidation state. 3. What are the main differences in the physical and chemical properties of red and white phosphorus. How can red phosphorus be separated from white impurities? 4. Calculate the relative density of phosphine from hydrogen and air. Is phosphine lighter or heavier than these gases? 5. How can the transition from red to white phosphorus and vice versa be made? Are these processes chemical phenomena? Explain the answer. 6. Calculate the mass of phosphorus that must be burned in oxygen to obtain phosphorus (V) oxide weighing 3.55 g? 7. A mixture of red and white phosphorus weighing 20 g was treated with carbon disulfide. The undissolved residue was separated and weighed, its mass was 12.6 g. Calculate the mass fraction of white phosphorus in the initial mixture. 8. What is the type of chemical bond in compounds: a) РН3; b) PCl5; c) Li3P. In polar substances, indicate the direction of displacement of common electron pairs. 9. Phosphine can be obtained by the action of hydrochloric acid on calcium phosphide. Calculate the volume of phosphine (normal conditions) that is formed from 9.1 g of calcium phosphide. The mass fraction of the product yield is 90%.
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Phosphoric acid and its salts
1. Write the reaction equations between phosphoric acid and the following substances: a) magnesium oxide; b) potassium carbonate; c) silver nitrate; d) iron sulfate (II). 2. Write the reaction equations between orthophosphoric acid and potassium hydroxide, as a result of which 3 types of salts are formed: medium and two acidic. 3. Which of the acids is a stronger oxidizing agent: nitric or orthophosphoric? Explain the answer. 4. Write the reaction equations by which the following transformations can be carried out: P → P205 → H3P04 → Na3P04 → Ca3(P04)2 P04)2→Ca(H2P04)2 Write the equations for these reactions. 6. Using the electronic balance method, select the coefficients in the schemes of the following redox reactions: a) RN3 + O2 → P2O5 + H2O b) Ca3 (PO4) 2 + C + SiO2 → CaSiO3 + P + CO acids 40% can be obtained from phosphorite weighing 100 kg with a mass fraction of Ca3 (PO4) 2 93%? 8. Phosphoric acid weighing 195 kg was obtained from natural phosphorite weighing 310 kg. Calculate the mass fraction of Ca3(PO4)2 in natural phosphorite. 9. An aqueous solution containing phosphoric acid weighing 19.6 g was neutralized with calcium hydroxide weighing 18.5 g. Determine the mass of the CaHPO4 2H2O precipitate formed. 10. There is a solution of phosphoric acid weighing 150 g (mass fraction of H3PO4 24.5%). Calculate the volume of ammonia (normal conditions) that must be passed through the solution to obtain ammonium dihydrogen phosphate. 11. What salt is formed if 2.8 g of potassium hydroxide is added to a solution containing H3PO4 weighing 4.9 g? Calculate the mass of the resulting salt
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Mineral fertilizers
1. What nitrogen and phosphate fertilizers do you know? Write the reaction equations for their production. Why do plants need nitrogen and phosphorus? 2. Determine the mass fraction of phosphorus (V) oxide in the CaHPO4 2H2O precipitate. 3. The mass fraction of phosphorus (V) oxide in superphosphate is 20%. Determine the mass of superphosphate to be introduced under a fruit tree if 15.5 g of phosphorus is required for the normal development of the tree. 4. The mass fraction of nitrogen in the fertilizer is 14%. All nitrogen is included in the fertilizer in the composition of urea CO(NH2)2. Calculate the mass fraction of urea in this fertilizer. 5. In superphosphate, the mass fraction of phosphorus (V) oxide is 25%. Calculate the mass fraction of Ca(H2PO4)2 in this fertilizer. 6. Calculate the mass of ammonium sulfate, which should be taken in order to introduce nitrogen weighing 2 tons into the soil on an area of 5 hectares. What mass of fertilizer should be applied to each square meter of soil? 7. Calculate the mass of ammonium nitrate to be applied to an area of 100 ha if the mass of nitrogen applied to an area of 1 ha is to be 60 kg. 8. Phosphorus (V) oxide weighing 0.4 kg must be introduced into the soil under the fruit tree. What mass of superphosphate should be taken in this case, if the mass fraction of assimilable phosphorus (V) oxide in it is 20%? 9. Under the fruit tree, it is necessary to add ammonium nitrate weighing 140 g (the mass fraction of nitrogen in nitrate is 35%). Determine the mass of ammonium sulfate, which can be used to add the same amount of nitrogen.
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References: 1. F.G. Feldman, G.E. Rudzitis. CHEMISTRY. Textbook for grade 9 educational institutions. - M., 5th edition, ENLIGHTENMENT, 1997. 2. CHEMISTRY. Reference materials. Under the editorship of Yu.D. Tretyakov, - M., EDUCATION, 1984. 3. CHEMISTRY. Schoolchildren's handbook, - M., 1995. 4. CHEMISTRY. Encyclopedia for children. Volume 17, AVANTA, 2000 5. Weser V.-J., Phosphorus and its compounds, trans. from English, - M., 1963. 6. Internet: http://school-sector.relarn.ru/nsm/chemistry/
Group V A subgroup The elements of this subgroup include: The elements of this subgroup include: N; P; As; Sb; Bi. N; P; As; Sb; Bi. Nitrogen and phosphorus are especially important Nitrogen and phosphorus are especially important Nitrogen is part of the air, is part of Nitrogen is part of the air, proteins, nucleic acids, proteins, nucleic acids, many rocks and minerals ( saltpeter) of many rocks and minerals (nitrate) Phosphorus is a constituent of proteins, nucleic acids, apatite and phosphorite minerals Phosphorus is a constituent of proteins, nucleic acids, apatite minerals and phosphorites
Characterization of nitrogen and phosphorus according to the periodic system Characteristic plan NitrogenPhosphorus
Electronic formulas of nitrogen and phosphorus N)) 1s²2s²2p³ N)) 1s²2s²2p³ 2 5 Compose the electron graphical formula of the graphical formula of nitrogen nitrogen +7 N highest oxidation state +7 N highest oxidation state lowest oxidation state lowest oxidation state -3 - 3
Did you know that ... Nitrogen was first discovered by scientists Nitrogen was first discovered by scientist D. Rutherford in 1772. The properties were investigated by K Scheele, G. Cavendish, D. Rutherford in 1772. The properties were investigated by K Scheele, G. Cavendish, J. Priestley. J. Priestley. A. Lavoisier proposed the term nitrogen, which is translated from Greek as "lifeless" A. Lavoisier proposed the term nitrogen, which is translated from Greek as "lifeless"
Nitrogen. Physical properties Molecular structure N2 Molecule structure N2 Structural formula N Ξ N Structural formula N Ξ N Electronic formula: N N: Electronic formula: N N: Bond covalent non-polar, very strong, triple 1σ(sigma) and 2π (pi) Bond covalent non-polar, very strong, triple 1σ (sigma) and 2π (pi) Nitrogen gas is colorless and odorless, poorly soluble in water, slightly lighter than air, Nitrogen gas is colorless and odorless, poorly soluble in water, slightly lighter than air, Тboil = ºС Тboil = ºС
Chemical properties of nitrogen Under normal conditions, low activity Under normal conditions, low activity At T=2000º it reacts with oxygen At T=2000º it reacts with oxygen N 2 + O 2 2 NO –Q N 2 + O 2 2 NO –Q At T=400ºC and p At T \u003d 400 ° C and p N 2 + 3H 2 2 NH 3 N 2 + 3H 2 2 NH 3 ammonia ammonia With some metals With some metals 3 Mg + N 2 Mg 3 N 2 3 Mg + N 2 Mg 3 N 2 magnesium nitride magnesium nitride
Ammonia Ammonia The compound of nitrogen with hydrogen is called ammonia NH 3 The compound of nitrogen with hydrogen is called ammonia NH 3 Molecule structure Molecule structure H – N – H H – N – H | H Covalent polar bond Covalent polar bond Shape of the molecule tetrahedron fig.11 page 47 Shape of the molecule tetrahedron fig.11 page 47
Obtaining in industry In 1913, the first plant for the catalytic synthesis of ammonia was established in Germany In 1913, the first plant for the catalytic synthesis of ammonia was established in Germany N2 + 3H2 2NH3 +Q N2 + 3H2 2NH3 +Q in the presence of a catalyst - The reaction is reversible, T = 300ºС, P = MPa, in the presence of a catalyst - porous iron porous iron
Obtaining in the laboratory By heating a mixture of ammonium salts with alkalis. When heating a mixture of ammonium salts with alkalis. 2NH4Cl +Ca(OH)2=CaCl2+2NH3 +2H2O ammonium chloride ammonia ammonium chloride ammonia Physical properties Physical properties Colorless gas with a characteristic pungent odor, almost 2 times lighter than air. Let's well dissolve in water. В 1V H2O – 700V NH3 Colorless gas with a characteristic pungent odor, almost 2 times lighter than air. Let's well dissolve in water. At 1V H2O - 700V NH3
Chemical properties Active substance Active substance Reacts with water Reacts with water NH3 + H2O NH4OH ammonium hydroxide NH3 + H2O NH4OH ammonium hydroxide With acids With acids NH3 + HCl = NH4Cl ammonium chloride NH3 + HCl = NH4Cl ammonium chloride 2NH3 + H2SO4 = (NH4)2SO4 ammonium sulfate 2NH3 + H2SO4 = (NH4)2SO4 ammonium sulfate
Chemical properties Weak compound decomposes when heated Weak compound decomposes when heated 2NH3 N2 + 3H2 2NH3 N2 + 3H2 Burns Burns? NH3 + ? O2? N2 + ?H2O ?NH3 + ? O2? N2 + ?H2O Oxidized in the presence of a Pt catalyst Oxidized in the presence of a Pt catalyst ? NH3+? O2? NO + ?H2O? NH3+? O2? NO + ?H2O check page 49 tab. 13 check p. 49 tab. 13 Reduces metals from their oxides Reduces metals from their oxides 2 NH3 + 3 CuO = N2 + 3Cu + 3 H2O 2 NH3 + 3 CuO = N2 + 3Cu + 3 H2O
Ammonium salts NH3 + HCl = NH4Cl ammonium chloride NH3 + HCl = NH4Cl ammonium chloride 2NH3 + H2SO4 = (NH4)2 SO4 ammonium sulfate 2NH3 + H2SO4 = (NH4)2 SO4 ammonium sulfate NH3 + H2SO4 = NH4HSO4 ammonium hydrosulfate NH3 + H2SO4 = NH4HSO4 ammonium hydrosulfate NH3 + HNO3 = ? Name NH3 + HNO3 = ? Name NH3 + H3PO4 = ? NH3 + H3PO4 = ? Qualitative reaction to ammonium ion Qualitative reaction to ammonium ion NH4 CL +NaOH =NaCl +NH3 +H2O smell of ammonia NH4 CL +NaOH =NaCl +NH3 +H2O smell of ammonia Decomposes when heated Decomposes when heated NH4NO3 = N2O +2 H2O NH4NO3 = N2O + 2 H2O NH4NO2 = N2 + 2H2O NH4NO2 = N2 + 2H2O
Questions and exercises What elements make up the VA group? What elements make up the VA group? What is the structure of the outer electron layer of nitrogen and phosphorus atoms? What is the structure of the outer electron layer of nitrogen and phosphorus atoms? What are the physical properties of nitrogen? What are the physical properties of nitrogen? Why is nitrogen chemically inactive? Why is nitrogen chemically inactive? How much nitrogen is in the air by volume? How much nitrogen is in the air by volume? What type of chemical bond is in a nitrogen molecule? What type of chemical bond is in a nitrogen molecule? Where is nitrogen found in nature? Where is nitrogen found in nature? How is nitrogen obtained? How is nitrogen obtained? Name the hydrogen compound of nitrogen, its physical properties. Name the hydrogen compound of nitrogen, its physical properties. How is ammonia obtained in the laboratory and industry? How is ammonia obtained in the laboratory and industry?
Questions and exercises What salts does ammonia form? What salt forms ammonia? What is a qualitative reaction for the ammonium cation? What is a qualitative reaction for the ammonium cation? Where are nitrogen, ammonia, ammonium salts used? Where are nitrogen, ammonia, ammonium salts used?
Exercise Solve transformation chain Solve transformation chain N2 NH3 NH4OH NH4Cl NH3 N2 NH3 NH4OH NH4Cl NH3 Solve OVR Solve OVR NH3 + O2 NO + H2O NH3 + O2 NO + H2O l of hydrogen? Calculate the volume of ammonia (N.O.) that is formed from 25 liters of nitrogen and 25 liters of hydrogen? What is the mass and volume of 5 moles of ammonia? What is the mass and volume of 5 moles of ammonia? Calculate relative density for hydrogen and ammonia for air? Calculate relative density for hydrogen and ammonia for air?
Nitric oxides Several nitrogen oxides are known Several nitrogen oxides are known in H 2 O "laughing gas" NO - nitric oxide II Colorless, odorless, slightly soluble N 2 O 3 nitric oxide III Dark blue liquid, sol. in water. NO 2 nitric oxide IV Brown gas, poisonous N 2 O 5 nitric oxide V Colorless
Obtaining In the laboratory In the laboratory NaNO3 + H2SO4 = NaHSO4 + HNO3 NaNO3 + H2SO4 = NaHSO4 + HNO3 sodium nitrate sodium hydrogen sulfate sodium nitrate sodium hydrogen sulfate In industry In industry 1. 4NH3 + O2 = 4NO + 6H2O +Q 2. 2NO + O2 = 2NO NO2 + H2O + O2 4 HNO3 +Q
Physical properties Colorless fuming liquid with a pungent odor. Well soluble in water. Concentrated is very dangerous. Decomposes under the influence of light. Store in a dark container. Strong oxidizer. Flammable. Colorless fuming liquid with a pungent odor. Well soluble in water. Concentrated is very dangerous. Decomposes under the influence of light. Store in a dark container. Strong oxidizer. Flammable.
Chemical properties Common with other acids Common with other acids 1. Strong acid, completely dissociates HNO3 H + NO3 ˉ HNO3 H + NO3 ˉ 2. React with basic oxides CuO + HNO3 = ? +? CuO + HNO3 = ? +? 3. React with bases Fe(OH)3 + HNO3 = ? +? Fe(OH)3 + HNO3 = ? +? 4 Reacts with salts of weaker acids Na2CO3 + HNO3 = ? +? +? Na2CO3 + HNO3 = ? +? +?
Specific properties Decomposes on heating and in the light Decomposes on heating and in the light 4HNO3 2 H2O + 4NO2 + O2 4HNO3 2 H2O + 4NO2 + O2 When interacting with proteins, a bright yellow substance is formed. When interacting with proteins, a bright yellow substance is formed. Reacts differently with metals, but hydrogen H2 is never released Reacts differently with metals, while hydrogen H2 is never released Me + HNO3 = Me NO3 + H2O + gas Me + HNO3 = Me NO3 + H2O + gas
Salts of nitric acid Nitric acid salts - nitrates Nitrogen salts - nitrates Potassium, sodium, ammonium and calcium nitrates are called saltpeters. Potassium, sodium, ammonium and calcium nitrates are called saltpeters. NaNO3 - sodium nitrate, sodium nitrate NaNO3 - sodium nitrate, sodium nitrate NH4NO3 - ammonium nitrate, ammonia NH4NO3 - ammonium nitrate, ammonium nitrate. saltpeter. All nitrates are soluble in water. All nitrates are soluble in water. They are strong oxidizing agents. They are strong oxidizing agents. When heated, all nitrates decompose with the release of oxygen O 2 When heated, all nitrates decompose with the release of oxygen O 2
Questions and exercises What nitrogen oxides do you know? What oxides of nitrogen do you know? What are the physical properties of nitric acid What are the physical properties of nitric acid Describe the chemical properties of nitric acid? Describe the chemical properties of nitric acid? What specific properties of nitric acid do you know? What specific properties of nitric acid do you know? How is nitric acid produced in the laboratory? How is nitric acid produced in the laboratory? How is nitric acid produced industrially? How is nitric acid produced industrially? Where is nitric acid used? Where is nitric acid used? What are nitric acid salts called and where are they used? What are nitric acid salts called and where are they used?
Exercises Write molecular and ionic reaction equations Write molecular and ionic reaction equations CaO + HNO3 = CaO + HNO3 = NaOH + HNO3 = NaOH + HNO3 = K2CO3 + HNO3 = K2CO3 + HNO3 = Write the reaction equation nitric acid with copper. Solve it as OVR Write the equation for the reaction of conc. nitric acid with copper. Solve it as OVR Cu + 4 HNO3 = Cu(NO3)2 + 2NO2 + 2H2O Cu + 4 HNO3 = Cu(NO3)2 + 2NO2 + 2H2O
Exercises Solve the chain of transformations Solve the chain of transformations N2 NO NO2 HNO3 N2 NO NO2 HNO3 KNO3 HNO3 Cu(NO3)2 NO2 KNO3 HNO3 Cu(NO3)2 NO2 Calculate the mass of magnesium nitrate, which was formed by the interaction of magnesium oxide with 120 g of nitric acid solution with 10% concentration. Calculate the mass of magnesium nitrate, which was formed by the interaction of magnesium oxide with 120 g of a solution of nitric acid with a 10% concentration. What volume of oxygen will be released during decomposition when 150 g of sodium nitrate is heated? What volume of oxygen will be released during decomposition when 150 g of sodium nitrate is heated? Calculate the mass fraction of nitrogen in aluminum nitrate. Calculate the mass fraction of nitrogen in aluminum nitrate.
Chemistry lesson in grade 10: "Nitrogen and phosphorus - p-elements of the VA group"
- prepared
- chemistry and biology teacher
- GUO secondary school №163 Minsk
- Kostyukevich Yury Mikhailovich
- In chemical compounds, nitrogen and phosphorus atoms can exhibit oxidation states from -3 to +5.
- Nitrogen is symbolized N
- (lat. Nitrogenium, i.e. "giving birth to saltpeter).
- The simple substance nitrogen (N2) is a rather inert gas under normal conditions, without color, taste and smell.
- Nitrogen, in the form of diatomic N2 molecules, makes up most of the atmosphere, where its content is 78.084% by volume (that is, about 3.87 1015 tons).
- Outside the Earth, nitrogen is found in gaseous nebulae, the solar atmosphere, on Uranus, Neptune, interstellar space, etc. Nitrogen is the 4th most abundant element in the solar system (after hydrogen, helium and oxygen).
- Phosphorus occurs naturally in the form of phosphates. Thus, calcium phosphate Ca3(PO4)2 is the main component of the mineral apatite.
- Phosphorus is found in all parts of green plants, and even more in fruits and seeds.
- Contained in animal tissues, is part of proteins and other essential organic compounds (ATP, DNA), is an element of life.
- Apatite
- Phosphorus
- In the light and when heated to 300 ° C without air, white phosphorus turns into red phosphorus.
- Of the metals, under normal conditions, nitrogen reacts only with lithium, forming lithium nitride:
- The reducing properties of nitrogen and phosphorus are manifested when they interact with oxygen. So, nitrogen reacts with oxygen at a temperature of about 3000˚С, forming nitric oxide (II):
- Oxidation of white phosphorus is accompanied by luminescence. White and red phosphorus ignite when ignited and burn with a dazzlingly bright flame with the formation of white smoke of phosphorus (IV) oxide:
- - During the siege of Sarajevo, phosphorus shells were used by Bosnian Serb artillery. In 1992, such shells burned down the building of the Institute of Oriental Studies, as a result of which many historical documents were destroyed.
- - in 2003-2004 they were used by British intelligence services in the vicinity of Basra in Iraq.
- - in 2004, the United States used against the guerrilla underground in Iraq in the battle for Fallujah.
- in the summer of 2006, during the Second Lebanon War, artillery shells with white phosphorus were used by the Israeli army.
- in 2009, during Operation Cast Lead in the Gaza Strip, the Israeli army used ammunition containing white phosphorus, which is allowed by international law.
- Since 2009 Palestinian terrorists have been loading their missiles with white phosphorus.
- The use of simple substances
- Production
- ammonia
- Most modern lamps are filled with chemically inert gases. Mixtures of nitrogen N2 with argon Ar are the most common due to their low cost.
- http://ru.wikipedia.org/wiki/Nitrogen
- http://en.wikipedia.org/wiki/Phosphorus
- http://distant-lessons.ru/ximiya/podgruppa-azota
- http://www.vredno.com.ua/2011/10/05/
- http://21region.org/sections/book/41869-istoriya-spichek.html
- http://x-ufo.ru/2008/08/19/fotografii_pjejjnobektov_s_kladbishha.html
- http://www.varson.ru/images/Himia_jpeg_big/7-04.jpg
- http://lols.ru/2010/11/09/
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