Four elements - "chalcogen" (i.e. "giving birth to copper") head the main subgroup of group VI (according to the new classification - the 16th group) of the periodic system. In addition to sulfur, tellurium and selenium, they also include oxygen. Let's take a closer look at the properties of this most common element on Earth, as well as the use and production of oxygen.
Element abundance
In a bound form, oxygen is included in the chemical composition of water - its percentage is about 89%, as well as in the composition of the cells of all living beings - plants and animals.
In the air, oxygen is in a free state in the form of O2, occupying a fifth of its composition, and in the form of ozone - O3.
Physical properties
Oxygen O2 is a colorless, tasteless and odorless gas. It is slightly soluble in water. The boiling point is 183 degrees below zero Celsius. In liquid form, oxygen has a blue color, and in solid form it forms blue crystals. The melting point of oxygen crystals is 218.7 degrees below zero Celsius.
Chemical properties
When heated, this element reacts with many simple substances, both metals and non-metals, while forming the so-called oxides - compounds of elements with oxygen. in which elements enter with oxygen is called oxidation.
For example,
4Na + O2= 2Na2O
2. Through the decomposition of hydrogen peroxide when it is heated in the presence of manganese oxide, which acts as a catalyst.
3. Through the decomposition of potassium permanganate.
The production of oxygen in industry is carried out in the following ways:
1. For technical purposes, oxygen is obtained from air, in which its usual content is about 20%, i.e. fifth part. To do this, the air is first burned, obtaining a mixture with a liquid oxygen content of about 54%, liquid nitrogen - 44% and liquid argon - 2%. These gases are then separated by a distillation process using a relatively small interval between the boiling points of liquid oxygen and liquid nitrogen - minus 183 and minus 198.5 degrees, respectively. It turns out that nitrogen evaporates before oxygen.
Modern equipment ensures the production of oxygen of any degree of purity. Nitrogen, which is obtained by separating liquid air, is used as a raw material in the synthesis of its derivatives.
2. also gives oxygen to a very pure degree. This method has become widespread in countries with rich resources and cheap electricity.
Application of oxygen
Oxygen is the most important element in the life of our entire planet. This gas, which is contained in the atmosphere, is consumed in the process by animals and humans.
Obtaining oxygen is very important for such areas of human activity as medicine, welding and cutting of metals, blasting, aviation (for breathing people and for the operation of engines), metallurgy.
In progress economic activity human oxygen is consumed in large quantities - for example, when burning various kinds fuels: natural gas, methane, coal, wood. In all these processes, it is formed. At the same time, nature has provided for the process of natural binding of this compound through photosynthesis, which takes place in green plants under the influence of sunlight. As a result of this process, glucose is formed, which the plant then uses to build its tissues.
Oxygen enters into compounds with almost all elements of the periodic system of Mendeleev.
The reaction of any substance with oxygen is called oxidation.
Most of these reactions involve the release of heat. When light is released during an oxidation reaction, it is called combustion. However, it is not always possible to notice the heat and light released, since in some cases the oxidation proceeds extremely slowly. It is possible to notice heat release when the oxidation reaction occurs quickly.
As a result of any oxidation - fast or slow - in most cases, oxides are formed: compounds of metals, carbon, sulfur, phosphorus and other elements with oxygen.
You probably have seen more than once how iron roofs are covered. Before covering them with new iron, the old is thrown down. Brown scales - rust - fall to the ground along with iron. This is iron oxide hydrate, which slowly, over several years, was formed on iron under the action of oxygen, moisture and carbon dioxide.
Rust can be thought of as a combination of iron oxide with a water molecule. It has a loose structure and does not protect iron from destruction.
To protect iron from destruction - corrosion - it is usually coated with paint or other corrosion-resistant materials: zinc, chromium, nickel and other metals. The protective properties of these metals, like aluminum, are based on the fact that they are covered with a thin stable film of their oxides, which protect the coating from further destruction.
Protective coatings significantly slow down the process of metal oxidation.
In nature, processes of slow oxidation, similar to combustion, constantly occur.
When rotting wood, straw, leaves and other organic matter there are processes of oxidation of carbon, which is part of these substances. Heat is released extremely slowly, and therefore usually goes unnoticed.
But sometimes this kind oxidative processes themselves are accelerated and go into combustion.
Spontaneous combustion can be observed in a wet haystack.
Rapid oxidation with the release of a large amount of heat and light can be observed not only during the combustion of wood, kerosene, candles, oil and other combustible materials containing carbon, but also during the combustion of iron.
Pour some water into the jar and fill it with oxygen. Then put an iron spiral into the jar, at the end of which a smoldering splinter is fixed. The splinter, and behind it the spiral, will light up with a bright flame, scattering star-shaped sparks in all directions.
This is the process of rapid oxidation of iron by oxygen. It began at a high temperature, which gave a burning splinter, and continues until the complete combustion of the spiral due to the heat released during the combustion of iron.
There is so much heat from this that the particles of oxidized iron formed during combustion glow white, brightly illuminating the jar.
The composition of the scale formed during the combustion of iron is somewhat different from the composition of the oxide formed in the form of rust during the slow oxidation of iron in air in the presence of moisture.
In the first case, the oxidation goes to ferrous oxide (Fe 3 O 4), which is part of the magnetic iron ore; in the second, an oxide is formed that closely resembles brown iron ore, which has the formula 2Fe 2 O 3 ∙ H 2 O.
Thus, depending on the conditions under which the oxidation proceeds, various oxides are formed, differing from each other in the content of oxygen.
So, for example, carbon in combination with oxygen gives two oxides - carbon monoxide and carbon dioxide. With a lack of oxygen, incomplete combustion of carbon occurs with the formation of carbon monoxide (CO), which in the hostel is called carbon monoxide. Complete combustion produces carbon dioxide, or carbon dioxide(CO2).
Phosphorus, burning under conditions of lack of oxygen, forms phosphorous anhydride (P 2 O 3), and with an excess - phosphorus anhydride (P 2 O 5). Sulfur under various combustion conditions can also give sulfurous (SO 2) or sulfuric (SO 3) anhydride.
In pure oxygen, combustion and other oxidation reactions proceed faster and reach completion.
Why does combustion proceed more vigorously in oxygen than in air?
Does pure oxygen have any special properties that the oxygen in the air does not have? Of course not. In both cases we have the same oxygen, with the same properties. Only air contains 5 times less oxygen than the same volume of pure oxygen, and, in addition, oxygen is mixed with oxygen in air. large quantities nitrogen, which not only does not burn itself, but also does not support combustion. Therefore, if the oxygen of the air is already used up directly near the flame, then another portion of it must break through nitrogen and combustion products. Consequently, more vigorous combustion in an oxygen atmosphere can be explained by its faster supply to the place of combustion. In this case, the process of combining oxygen with a burning substance is more energetic and more heat is released. The more oxygen is supplied to the burning substance per unit time, the brighter the flame, the higher the temperature and the stronger the combustion.
Does oxygen itself burn?
Take the cylinder and turn it upside down. Place a tube of hydrogen under the cylinder. Since hydrogen is lighter than air, it will completely fill the cylinder.
Ignite hydrogen near the open part of the cylinder and insert a glass tube through the flame into it, through which gaseous oxygen flows. Near the end of the tube, a fire will flare up, which will burn quietly inside a cylinder filled with hydrogen. It is not oxygen that is burning, but hydrogen in the presence of a small amount of oxygen coming out of the tube.
What is formed as a result of the combustion of hydrogen? What is the resulting oxide?
Hydrogen is oxidized to water. Indeed, droplets of condensed water vapor gradually begin to settle on the walls of the cylinder. 1 oxygen molecule goes to the oxidation of 2 hydrogen molecules, and 2 water molecules are formed (2H 2 + O 2 → 2H 2 O).
If the oxygen flows out of the tube slowly, it burns out completely in the hydrogen atmosphere, and the experiment goes smoothly.
One has only to increase the supply of oxygen so much that it does not have time to burn out completely, part of it will go beyond the flame, where pockets of a mixture of hydrogen and oxygen are formed, and separate small flashes will appear, similar to explosions.
A mixture of oxygen and hydrogen is an explosive gas. If you set fire to explosive gas, there will be a strong explosion: when oxygen combines with hydrogen, water is obtained and a high temperature develops. Water vapor and the surrounding gases expand greatly, creating a great pressure, at which not only a glass cylinder, but also a more durable vessel can easily burst. Therefore, working with an explosive mixture requires special care.
Oxygen has another interesting property. It enters into combination with some elements, forming peroxide compounds.
Let's bring characteristic example. Hydrogen, as you know, is monovalent, oxygen is bivalent: 2 hydrogen atoms can combine with 1 oxygen atom. This produces water. The structure of a water molecule is usually depicted as H - O - H. If 1 more oxygen atom is attached to a water molecule, then hydrogen peroxide is formed, the formula of which is H 2 O 2.
Where does the second oxygen atom enter in this compound and by what bonds is it held? The second oxygen atom, as it were, breaks the bond of the first one with one of the hydrogen atoms and becomes between them, thus forming H-O-O-N connection. The same structure has sodium peroxide (Na-O-O-Na), barium peroxide.
Characteristic of peroxide compounds is the presence of 2 oxygen atoms, interconnected by one valency. Therefore, 2 hydrogen atoms, 2 sodium atoms or 1 barium atom can attach to themselves not 1 oxygen atom with two valencies (-O-), but 2 atoms, which, as a result of the bond between themselves, also have only two free valences (-O- ABOUT-).
Hydrogen peroxide can be obtained by the action of dilute sulfuric acid on sodium peroxide (Na 2 O 2) or barium peroxide (BaO 2). It is more convenient to use barium peroxide, since when sulfuric acid acts on it, an insoluble precipitate of barium sulfate is formed, from which hydrogen peroxide is easily separated by filtration (BaO 2 + H 2 SO 4 → BaSO 4 + H 2 O 2).
Hydrogen peroxide, like ozone, is an unstable compound and decomposes into water and an oxygen atom, which at the time of release has a high oxidizing power. At low temperatures and in the dark, the decomposition of hydrogen peroxide is slow. And when heated and in the light, it happens much faster. Sand, powdered manganese dioxide, silver or platinum also accelerate the decomposition of hydrogen peroxide, while they themselves remain unchanged. Substances that only affect the rate of a chemical reaction, while themselves remain unchanged, are called catalysts.
If you pour a little hydrogen peroxide into a bottle, at the bottom of which there is a catalyst - manganese dioxide powder, the decomposition of hydrogen peroxide will proceed with such speed that you can notice the release of oxygen bubbles.
The ability to oxidize various compounds is possessed not only by gaseous oxygen, but also by some compounds in which it is included.
Hydrogen peroxide is a good oxidizing agent. It bleaches various dyes and is therefore used in technology for bleaching silk, fur and other products.
The ability of hydrogen peroxide to kill various microbes allows it to be used as a disinfectant. Hydrogen peroxide is used for washing wounds, gargling and in dental practice.
Has strong oxidizing properties Nitric acid(HNO3). If a drop of turpentine is added to nitric acid, a bright flash is formed: carbon and hydrogen, which are part of turpentine, are rapidly oxidized with the release of a large amount of heat.
Paper and fabrics moistened with nitric acid are rapidly destroyed. The organic substances from which these materials are made are oxidized by nitric acid and lose their properties. If paper or cloth soaked in nitric acid is heated, the oxidation process will accelerate so much that a flash may occur.
Nitric acid oxidizes not only organic compounds, but also some metals. Copper, when exposed to concentrated nitric acid, is oxidized first to copper oxide, releasing nitrogen dioxide from nitric acid, and then copper oxide turns into copper nitrate.
Not only nitric acid, but also some of its salts have strong oxidizing properties.
Nitric acid salts of potassium, sodium, calcium and ammonium, which are called saltpeter in technology, decompose when heated, releasing oxygen. At high temperatures in molten saltpeter, the ember burns so vigorously that a bright white light appears. If, however, a piece of sulfur is thrown into a test tube with molten saltpeter along with a smoldering coal, combustion will go on with such intensity and the temperature will rise so much that the glass will begin to melt. These properties of saltpeter have long been known to man; he took advantage of these properties to make gunpowder.
Black, or smoky, gunpowder is made from saltpeter, coal and sulfur. In this mixture, coal and sulfur are combustible materials. When burned, they turn into gaseous carbon dioxide (CO 2) and solid potassium sulfide (K 2 S). Saltpeter, decomposing, releases a large amount of oxygen and gaseous nitrogen. The released oxygen enhances the combustion of coal and sulfur.
As a result of combustion, such a high temperature develops that the gases formed could expand to a volume that is 2000 times the volume of the taken gunpowder. But the walls of a closed vessel, where gunpowder is usually burned, do not allow gases to expand easily and freely. An enormous pressure is created, which breaks the vessel at its weakest point. A deafening explosion is heard, gases burst out with noise, carrying with them crushed solid particles in the form of smoke.
So from potassium nitrate, coal and sulfur, a mixture is formed that has tremendous destructive power.
Compounds with strong oxidizing properties also include salts of oxygen-containing chlorine acids. Bertolet's salt, when heated, decomposes into potassium chloride and atomic oxygen.
Even easier than Bertolet's salt, chloride, or bleach, lime gives up its oxygen. White lime is used to bleach cotton, linen, paper and other materials. Chloric lime is also used as a remedy against poisonous substances: poisonous substances, like many other complex compounds, are destroyed by strong oxidizing agents.
The oxidizing properties of oxygen, its ability to easily combine with various elements and vigorously support combustion, while developing a high temperature, have long attracted the attention of scientists. various areas Sciences. Chemists and metallurgists were especially interested in this. But the use of oxygen was limited because there was no easy and cheap way to get it from air and water.
Physicists came to the aid of chemists and metallurgists. They found a very convenient way to extract oxygen from the air, and physical chemists learned how to get it in huge quantities from water.
Oxygen is an element of the 16th group (according to the outdated classification - the main subgroup of group VI), the second period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 8. It is denoted by the symbol O. Oxygen is a reactive non-metal and is the lightest element of the chalcogen group. The simple substance oxygen normal conditions- a gas without color, taste and smell, the molecule of which consists of two oxygen atoms (formula O2), in connection with which it is also called dioxygen]. Liquid oxygen has a light blue color, and solid oxygen is light blue crystals.
There are other allotropic forms of oxygen, for example - under normal conditions, gas blue color with a specific odor, the molecule of which consists of three oxygen atoms (formula O3).
Finding in nature. Natural oxygen consists of 3 stable isotopes o16, o17, o18.
Oxygen in the form of a simple substance o2 is part of the atmospheric air. = 21% In a bound form, the element of oxygen is an integral part of water of various minerals of many organic substances.
RECEIVING. At present, in industry, oxygen is obtained from the air. The main industrial method for obtaining oxygen is cryogenic distillation. Oxygen plants based on membrane technology are also well known and successfully used in industry.
In laboratories, industrial oxygen is used, supplied in steel cylinders under a pressure of about 15 MPa.
Small amounts of oxygen can be obtained by heating potassium permanganate KMnO4:
2KMNO4 = K2MnO4 + MnO2 + O2
The reaction of catalytic decomposition of hydrogen peroxide H2O2 in the presence of manganese(IV) oxide is also used:
2H2O2 =MnO2=2H2O + O2
Oxygen can be obtained by catalytic decomposition of potassium chlorate (bertolet salt) KClO3:
2KClO3 = 2KCl + 3O2
Laboratory methods for obtaining oxygen include the method of electrolysis of aqueous solutions of alkalis, as well as the decomposition of mercury (II) oxide (at t = 100 ° C):
On submarines, it is usually obtained by the reaction of sodium peroxide and carbon dioxide exhaled by a person:
2Na2O2 + 2CO2 = 2Na2CO3 + O2
CHEMICAL ST_VA. A strong oxidizing agent, interacts with almost all elements, forming oxides. The oxidation state is −2. As a rule, the oxidation reaction proceeds with the release of heat and accelerates with increasing temperature (see Combustion). An example of reactions occurring at room temperature:
4Li + O2 = 2Li2O
Oxidizes compounds that contain elements with a non-maximum oxidation state:
Oxidizes most organic compounds:
CH3CH2OH + 3O2 = 2CO2 + 3H2O
Under certain conditions, it is possible to carry out a mild oxidation of an organic compound:
CH3CH2OH +O2 = CH3COOH + H2O
Oxygen reacts directly (under normal conditions, when heated and/or in the presence of catalysts) with all simple substances, except for Au and inert gases (He, Ne, Ar, Kr, Xe, Rn); reactions with halogens occur under the influence of an electric discharge or ultraviolet radiation. Oxides of gold and heavy inert gases (Xe, Rn) were obtained indirectly. In all two-element compounds of oxygen with other elements, oxygen plays the role of an oxidizing agent, except for compounds with fluorine (see below #oxygen fluorides).
Oxygen forms peroxides with the oxidation state of the oxygen atom formally equal to −1.
For example, peroxides are produced by burning alkali metals in oxygen:
2Na + O2 = Na2O2
Some oxides absorb oxygen:
2BaO + O2 = 2BaO2
According to the combustion theory developed by A. N. Bach and K. O. Engler, oxidation occurs in two stages with the formation of an intermediate peroxide compound. This intermediate compound can be isolated, for example, when a flame of burning hydrogen is cooled with ice, along with water, hydrogen peroxide is formed:
In superoxides, oxygen formally has an oxidation state of −½, that is, one electron per two oxygen atoms (O−2 ion). Obtained by the interaction of peroxides with oxygen at elevated pressure and temperature:
Na2O2 + O2 = 2NaO2
Potassium K, rubidium Rb and cesium Cs react with oxygen to form superoxides:
Inorganic ozonides contain the O–3 ion with the oxidation state of oxygen formally equal to –1/3. Obtained by the action of ozone on alkali metal hydroxides:
2KOH + 3O3 = 2KO3 + H2O +2O2
In the dioxygenyl ion O2+, oxygen formally has an oxidation state of +½. Get by reaction:
PtF6 +O2 = O2PtF6
Oxygen fluorides Oxygen difluoride, OF2 oxygen oxidation state +2, is obtained by passing fluorine through an alkali solution:
2F2 + 2NaOH = 2NaF + H2O + OF2
Oxygen monofluoride (Dioxydifluoride), O2F2, is unstable, oxygen oxidation state is +1. Obtained from a mixture of fluorine and oxygen in a glow discharge at a temperature of −196 C:
Passing a glow discharge through a mixture of fluorine with oxygen at a certain pressure and temperature, mixtures of higher oxygen fluorides O3F2, O4F2, O5F2 and O6F2 are obtained.
Quantum mechanical calculations predict the stable existence of the trifluorohydroxonium ion (English) OF3+. If this ion really exists, then the oxidation state of oxygen in it will be +4.
Oxygen supports the processes of respiration, combustion, and decay.
In its free form, the element exists in two allotropic modifications: O2 and O3 (ozone). As established in 1899 by Pierre Curie and Maria Sklodowska-Curie, under the influence of ionizing radiation, O2 passes into O3 OZONE. Ozone is an allotropic modification of oxygen consisting of triatomic O3 molecules. Under normal conditions - blue gas. When liquefied, it turns into an indigo liquid. In solid form, it is dark blue, almost black crystals.
CHEM.CB-VA Ozone is a powerful oxidizing agent, much more reactive than diatomic oxygen. Oxidizes almost all metals (with the exception of gold, platinum and iridium) to their higher degrees oxidation. Oxidizes many non-metals. The reaction product is mainly oxygen.
2Cu2+ + 2H3O+ + O3 = 2Cu3+ + 3H2O + O2
Ozone increases the oxidation state of oxides:
NO + O3 = NO2 + O2
This reaction is accompanied by chemiluminescence. Nitrogen dioxide can be oxidized to nitric anhydride:
2NO2 + O3 = N2O5 + O2
Ozone reacts with carbon at normal temperature to form carbon dioxide:
2C +2O3 = 2CO2 + O2
Ozone does not react with ammonium salts, but reacts with ammonia to form ammonium nitrate:
2NH3 + 4O3 = NH4NO3 + 4O2 + H2O
Ozone reacts with hydrogen to form water and oxygen:
O3 + H2 = O2 + H2O
Ozone reacts with sulfides to form sulfates:
PbS + 4O3 = PbSO4 + 4O2
Ozone can be used to sulfuric acid from both elemental sulfur and sulfur dioxide:
S + H2O + O3 = H2SO4
3SO2 + 3H2O + O3 = 3H2SO4
All three oxygen atoms in ozone can react individually in the reaction of tin chloride with hydrochloric acid and ozone:
3SnCl2 + 6HCl + O3 = 3SnCl4 + 3H2O
In the gas phase, ozone reacts with hydrogen sulfide to form sulfur dioxide:
H2S + O3 = SO2 + H2O
IN aqueous solution two competing reactions take place with hydrogen sulfide, one with the formation of elemental sulfur, the other with the formation of sulfuric acid:
H2S + O3 = S + O2 + H2O
3H2S + 4O3 = 3H2SO4
By treating a solution of iodine in cold anhydrous perchloric acid with ozone, iodine(III) perchlorate can be obtained:
I2 + 6HClO4 +O3 = 2I(ClO4)3 + 3H2O
Solid nitrile perchlorate can be obtained by the reaction of gaseous NO2, ClO2 and O3:
2NO2 + 2ClO2 + 2O2 = 2NO2ClO4 + O2
Ozone can participate in combustion reactions, while the combustion temperatures are higher than with diatomic oxygen:
3C3N2 + 4O3 = 12CO + 3N2
Ozone can enter into chemical reactions and at low temperatures. At 77 K (-196 °C), atomic hydrogen reacts with ozone to form a superoxide radical with dimerization of the latter:
H + O3 = HO2. + O
2HO2 . = H2O2 +O2
Ozone can form inorganic ozonides containing the O3− anion. These compounds are explosive and can only be stored at low temperatures. Ozonides of all alkali metals (except France) are known. KO3, RbO3, and CsO3 can be obtained from the corresponding superoxides:
KO2 + O3 = KO3 + O2
Potassium ozonide can be obtained in another way from potassium hydroxide:
2KOH + 5O3 = 2KO3 + 5O2 + H2O
NaO3 and LiO3 can be obtained by the action of CsO3 in liquid ammonia NH3 on ion exchange resins containing Na+ or Li+ ions:
CsO3 + Na+ = Cs+ + NaO3
Treatment of a solution of calcium in ammonia with ozone leads to the formation of ammonium ozonide, and not calcium:
3Ca + 10NH3 + 7O3 = Ca * 6NH3 + Ca(OH)2 + Ca(NO3)2 + 2NH4O3 + 3O2 + 2H2O
Ozone can be used to remove manganese from water to form a precipitate that can be separated by filtration:
2Mn2+ + 2O3 + 4H2O = 2MnO(OH)2 + 2O2 + 4H+
Ozone converts toxic cyanides into less dangerous cyanates:
CN- + O3 = CNO- + O2
Ozone can completely decompose urea:
(NH2)2CO + O3 = N2 + CO2 + 2H2O
The interaction of ozone with organic compounds with an activated or tertiary carbon atom at low temperatures leads to the corresponding hydrotrioxides.
RECEIVING. Ozone is formed in many processes accompanied by the release of atomic oxygen, for example, during the decomposition of peroxides, the oxidation of phosphorus, etc.
In industry, it is obtained from air or oxygen in ozonizers by the action of an electric discharge. O3 liquefies more easily than O2 and is therefore easy to separate. Ozone for ozone therapy in medicine is obtained only from pure oxygen. When air is irradiated with hard ultraviolet radiation, ozone is formed. The same process takes place in the upper layers of the atmosphere, where, under the action of solar radiation the formation and maintenance of the ozone layer.
In the laboratory, ozone can be obtained by reacting chilled concentrated sulfuric acid with barium peroxide:
3H2SO4 + 3BaO2 = 3BaSO4 + O3 + 3H2O
Peroxides are complex substances in which oxygen atoms are connected to each other. Peroxides readily release oxygen. For inorganic substances, it is recommended to use the term peroxide; for organic substances, the term peroxide is often used in Russian today. Peroxides of many organic substances are explosive (acetone peroxide); in particular, they are easily formed photochemically when ethers are illuminated for a long time in the presence of oxygen. Therefore, before distillation, many ethers (diethyl ether, tetrahydrofuran) require testing for the absence of peroxides.
Peroxides slow down protein synthesis in the cell.
Depending on the structure, peroxides proper, superoxides, and inorganic ozonides are distinguished. Inorganic peroxides in the form of binary or complex compounds known for almost all elements. Alkaline peroxides and alkaline earth metals react with water to form the corresponding hydroxide and hydrogen peroxide.
Organic peroxides are subdivided into dialkyl peroxides, alkyl hydroperoxides, diacyl peroxides, acyl hydroperoxides (peroxocarboxylic acids), and cyclic peroxides. Organic peroxides are thermally unstable and often explosive. Used as sources of free radicals in organic synthesis and industry
Halides (halides) - compounds of halogens with other chemical elements or radicals. In this case, the halogen included in the compound must be electronegative; Thus, bromine oxide is not a halide.
According to the halogen involved in the compound, halides are also called fluorides, chlorides, bromides, iodides and astatides. Silver halides are best known by this name due to the mass distribution of film silver halide photography.
Compounds of halogens among themselves are called interhalides, or interhalogen compounds (for example, iodine pentafluoride IF5).
In halides, the halogen has negative power oxidation, and the element is positive.
A halide ion is a negatively charged halogen atom.
DEFINITION
Oxygen- the eighth element Periodic table. Designation - O from the Latin "oxygenium". Located in the second period, VIA group. Refers to non-metals. The nuclear charge is 8.
Oxygen is the most abundant element in the earth's crust. In a free state, it is found in atmospheric air, in a bound form it is part of water, minerals, rocks and all substances from which plant and animal organisms are built. Mass fraction of oxygen in earth's crust is about 47%.
As a simple substance, oxygen is a colorless, odorless gas. It is slightly heavier than air: the mass of 1 liter of oxygen under normal conditions is 1.43 g, and 1 liter of air is 1.293 g. Oxygen dissolves in water, although in small quantities: 100 volumes of water at 0 o C dissolve 4.9, and at 20 o C - 3.1 volumes of oxygen.
Atomic and molecular weight of oxygen
DEFINITION
Relative atomic mass A r is the molar mass of an atom of a substance, divided by 1/12 molar mass carbon atom-12 (12 C).
The relative atomic mass of atomic oxygen is 15.999 amu.
DEFINITION
Relative molecular weight M r is the molar mass of the molecule, referred to 1/12 of the molar mass of the carbon-12 atom (12 C).
This is a dimensionless quantity. It is known that the oxygen molecule is diatomic - O 2 . The relative molecular weight of an oxygen molecule will be equal to:
M r (O 2) \u003d 15.999 × 2 ≈32.
Allotropy and allotropic modifications of oxygen
Oxygen can exist in the form of two allotropic modifications - oxygen O 2 and ozone O 3 ( physical properties oxygen described above).
At normal conditions ozone is a gas. It can be separated from oxygen by strong cooling; ozone condenses into a blue liquid boiling at (-111.9 o C).
The solubility of ozone in water is much greater than that of oxygen: 100 volumes of water at 0 o C dissolve 49 volumes of ozone.
The formation of ozone from oxygen can be expressed by the equation:
3O 2 \u003d 2O 3 - 285 kJ.
Isotopes of oxygen
It is known that in nature oxygen can be in the form of three isotopes 16 O (99.76%), 17 O (0.04%) and 18 O (0.2%). Their mass numbers are 16, 17 and 18, respectively. The nucleus of an atom of the oxygen isotope 16 O contains eight protons and eight neutrons, and the isotopes 17 O and 18 O contain the same number of protons, nine and ten neutrons, respectively.
There are twelve radioactive isotopes of oxygen with mass numbers from 12 to 24, of which the most stable isotope is 15 O with a half-life of 120 s.
oxygen ions
On the outer energy level of the oxygen atom, there are six electrons that are valence:
1s 2 2s 2 2p 4 .
The structure of the oxygen atom is shown below:
As a result of chemical interaction, oxygen can lose its valence electrons, i.e. be their donor, and turn into positively charged ions or accept electrons from another atom, i.e. be their acceptor, and turn into negatively charged ions:
O 0 +2e → O 2-;
About 0 -1e → About 1+.
Molecule and atom of oxygen
The oxygen molecule consists of two atoms - O 2 . Here are some properties that characterize the oxygen atom and molecule:
Examples of problem solving
EXAMPLE 1
The earth's crust is 50% oxygen. The element is also present in minerals in the form of salts and oxides. Oxygen in a bound form is included in the composition (the percentage of the element is about 89%). Oxygen is also present in the cells of all living organisms and plants. Oxygen is in the air in a free state in the form of O₂ and its allotropic modification in the form of ozone O₃, and occupies a fifth of its composition,
Physical and chemical properties of oxygen
Oxygen O₂ is a colorless, tasteless and odorless gas. Slightly soluble in water, boils at a temperature of (-183) °C. Oxygen in the form of a liquid has a blue color, in the solid form the element forms blue crystals. Oxygen melts at a temperature of (-218.7) °C.
Liquid oxygen at room temperatureWhen heated, oxygen reacts with various simple substances (metals and non-metals), resulting in the formation of oxides - compounds of elements with oxygen. The interaction of chemical elements with oxygen is called an oxidation reaction. Examples of reaction equations:
4Na + О₂= 2Na₂O
S + O₂ = SO₂.
Some complex substances also interact with oxygen, forming oxides:
CH₄ + 2O₂ \u003d CO₂ + 2H₂O
2СО + О₂ = 2СО₂
oxygen as chemical element obtained in laboratories and industrial enterprises. in the laboratory in several ways:
- decomposition (potassium chlorate);
- decomposition of hydrogen peroxide when the substance is heated in the presence of manganese oxide as a catalyst;
- decomposition of potassium permanganate.
Chemical reaction of oxygen combustion
Pure oxygen does not have special properties that atmospheric oxygen does not have, that is, it has the same chemical and physical properties. The air contains five times less oxygen than the same volume of pure oxygen. In the air, oxygen is mixed with large amounts of nitrogen, a gas that does not burn itself and does not support combustion. Therefore, if the oxygen in the air near the flame has already been used up, then the next portion of oxygen will break through nitrogen and combustion products. Consequently, more vigorous combustion of oxygen in the atmosphere is explained by a faster supply of oxygen to the place of combustion. During the reaction, the process of combining oxygen with the burning substance is carried out more vigorously and more heat is released. The more oxygen is supplied to the burning substance per unit time, the brighter the flame burns, the higher the temperature and the stronger the combustion process.
How does the oxygen combustion reaction take place? This can be verified by experience. It is necessary to take the cylinder and turn it upside down, then bring a tube of hydrogen under the cylinder. Hydrogen, which is lighter than air, will completely fill the cylinder. It is necessary to ignite hydrogen near the open part of the cylinder and introduce a glass tube into it through the flame, through which gaseous oxygen flows. A fire will flare at the end of the tube, while the flame will burn quietly inside the hydrogen-filled cylinder. During the reaction, it is not oxygen that burns, but hydrogen in the presence of a small amount of oxygen escaping from the tube.
What results from the combustion of hydrogen and what kind of oxide is formed? Hydrogen is oxidized to water. Droplets of condensed water vapor are gradually deposited on the walls of the cylinder. Two molecules of hydrogen are oxidized by one molecule of oxygen, and two molecules of water are formed. Reaction equation:
2Н₂ + O₂ → 2Н₂O
If oxygen flows out of the tube slowly, it burns out completely in the hydrogen atmosphere, and the experiment goes smoothly.
As soon as the supply of oxygen increases so much that it does not have time to burn out completely, part of it goes beyond the flame, where pockets of a mixture of hydrogen and oxygen are formed, and separate, explosion-like, small flashes appear. A mixture of oxygen and hydrogen is an explosive gas.
When the explosive gas is ignited, a strong explosion occurs: when oxygen combines with hydrogen, water is formed and a high temperature develops. Vapors of water with surrounding gases expand greatly, a great pressure arises, at which not only a fragile cylinder, but also a more durable vessel can burst. Therefore, it is necessary to work with an explosive mixture with extreme caution.
Oxygen consumption during combustion
For the experiment, a glass crystallizer with a volume of 3 liters must be filled 2/3 with water and a tablespoon of caustic soda or caustic potassium should be added. Color the water with phenolphthalein or other suitable dye. Pour sand into a small flask and insert a wire vertically into it with cotton wool fixed at the end. The cone is placed in a crystallizer with water. The cotton wool remains 10 cm above the surface of the solution.
Lightly moisten a cotton ball with alcohol, oil, hexane, or other flammable liquid and set fire to it. Carefully cover the burning cotton wool with a 3-liter bottle and lower it below the surface of the alkali solution. In the process of combustion, oxygen passes into water and. As a result of the reaction, the alkali solution in the bottle rises. The cotton wool will soon go out. The bottle should be carefully placed on the bottom of the crystallizer. In theory, the bottle should be 1/5 full, since air contains 20.9% oxygen. During combustion, oxygen passes into water and carbon dioxide CO₂, absorbed by alkali. Reaction equation:
2NaOH + CO₂ = Na₂CO₃ + H₂O
In practice, combustion will stop before all the oxygen is used up; part of the oxygen passes into carbon monoxide, which is not absorbed by alkali, and part of the air leaves the bottle as a result of thermal expansion.
Attention! Do not try to repeat these experiments yourself!
