Oxygen formsperoxides with oxidation state −1.
— For example, peroxides are produced by the combustion of alkali metals in oxygen:
2Na + O 2 → Na 2 O 2

— Some oxides absorb oxygen:
2BaO + O 2 → 2BaO 2

— According to the principles of combustion 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, hydrogen peroxide is formed along with water:
H 2 + O 2 → H 2 O 2

Superoxides have an oxidation state of −1/2, that is, one electron per two oxygen atoms (O 2 - ion). Obtained by reacting peroxides with oxygen at elevated pressures and temperatures:
Na 2 O 2 + O 2 → 2NaO 2

Ozonides contain the O 3 - ion with an oxidation state of −1/3. Obtained by the action of ozone on alkali metal hydroxides:
KOH(tv) + O 3 → KO 3 + KOH + O 2

Ion dioxygenyl O 2 + has an oxidation state of +1/2. Obtained by the reaction:
PtF 6 + O 2 → O 2 PtF 6

Oxygen fluorides
Oxygen difluoride, OF 2 oxidation state +2, is obtained by passing fluorine through an alkali solution:
2F 2 + 2NaOH → OF 2 + 2NaF + H 2 O

Oxygen monofluoride (Dioxydifluoride), O 2 F 2, unstable, oxidation state +1. It is obtained from a mixture of fluorine and oxygen in a glow discharge at a temperature of −196 °C.

By passing a glow discharge through a mixture of fluorine and oxygen at a certain pressure and temperature, mixtures of higher oxygen fluorides O 3 F 2, O 4 F 2, O 5 F 2 and O 6 F 2 are obtained.
Oxygen supports the processes of respiration, combustion, and decay. In its free form, the element exists in two allotropic modifications: O 2 and O 3 (ozone).

Application of oxygen

The widespread industrial use of oxygen began in the middle of the 20th century, after the invention of turboexpanders - devices for liquefying and separating liquid air.

In metallurgy

The converter method of steel production involves the use of oxygen.

Welding and cutting of metals

Oxygen in cylinders is widely used for flame cutting and welding of metals.

Propellant

Liquid oxygen, hydrogen peroxide, nitric acid and other oxygen-rich compounds are used as oxidizers for rocket fuel. A mixture of liquid oxygen and liquid ozone is one of the most powerful oxidizers of rocket fuel (the specific impulse of the hydrogen-ozone mixture exceeds the specific impulse for the hydrogen-fluorine and hydrogen-oxygen fluoride pairs).

In medicine

Oxygen is used to enrich respiratory gas mixtures for breathing problems, for the treatment of asthma, in the form of oxygen cocktails, oxygen pillows, etc.

In the food industry

In the food industry, oxygen is registered as a food additive E948, as propellant and packaging gas.

Biological role of oxygen

Living things breathe oxygen from the air. Oxygen is widely used in medicine. In case of cardiovascular diseases, to improve metabolic processes, oxygen foam (“oxygen cocktail”) is injected into the stomach. Subcutaneous administration of oxygen is used for trophic ulcers, elephantiasis, gangrene and other serious diseases. Artificial ozone enrichment is used to disinfect and deodorize air and purify drinking water. The radioactive isotope of oxygen 15 O is used to study blood flow speed and pulmonary ventilation.

Toxic oxygen derivatives

Some oxygen derivatives (so-called reactive oxygen species), such as singlet oxygen, hydrogen peroxide, superoxide, ozone and hydroxyl radical, are highly toxic. They are formed during the process of activation or partial reduction of oxygen. Superoxide (superoxide radical), hydrogen peroxide and hydroxyl radical can form in cells and tissues of humans and animals and cause oxidative stress.

Isotopes of oxygen

Oxygen has three stable isotopes: 16 O, 17 O and 18 O, the average content of which is, respectively, 99.759%, 0.037% and 0.204% of the total number of oxygen atoms on Earth. The sharp predominance of the lightest of them, 16 O, in the mixture of isotopes is due to the fact that the nucleus of the 16 O atom consists of 8 protons and 8 neutrons. And such nuclei, as follows from the theory of the structure of the atomic nucleus, are particularly stable.

There are radioactive isotopes 11 O, 13 O, 14 O (half-life 74 sec), 15 O (T 1/2 = 2.1 min), 19 O (T 1/2 = 29.4 sec), 20 O (contradictory half-life data from 10 minutes to 150 years).

Additional information

Oxygen compounds
Liquid oxygen
Ozone

Oxygen, Oxygenium, O (8)
The discovery of oxygen (Oxygen, French Oxygene, German Sauerstoff) marked the beginning of the modern period in the development of chemistry. It has been known since ancient times that combustion requires air, but for many centuries the combustion process remained unclear. Only in the 17th century. Mayow and Boyle independently expressed the idea that the air contains some substance that supports combustion, but this completely rational hypothesis was not developed at that time, since the idea of ​​combustion as a process of combining a burning body with a certain component of the air seemed at that time contradicting such an obvious act as the fact that during combustion the decomposition of the burning body into elementary components takes place. It was on this basis that at the turn of the 17th century. The phlogiston theory arose, created by Becher and Stahl. With the advent of the chemical-analytical period in the development of chemistry (the second half of the 18th century) and the emergence of “pneumatic chemistry” - one of the main branches of the chemical-analytical direction - combustion, as well as respiration, again attracted the attention of researchers. The discovery of various gases and the establishment of their important role in chemical processes was one of the main incentives for the systematic studies of combustion processes undertaken by Lavoisier. Oxygen was discovered in the early 70s of the 18th century.

The first report of this discovery was made by Priestley at a meeting of the Royal Society of England in 1775. Priestley, by heating red mercury oxide with a large burning glass, obtained a gas in which the candle burned more brightly than in ordinary air, and the smoldering splinter flared up. Priestley determined some of the properties of the new gas and called it daphlogisticated air. However, two years earlier than Priestley (1772), Scheele also obtained oxygen by the decomposition of mercuric oxide and other methods. Scheele called this gas fire air (Feuerluft). Scheele was able to report his discovery only in 1777.

In 1775, Lavoisier spoke before the Paris Academy of Sciences with the message that he had succeeded in obtaining “the purest part of the air that surrounds us,” and described the properties of this part of the air. At first, Lavoisier called this “air” empyrean, vital (Air empireal, Air vital) the basis of vital air (Base de l'air vital). The almost simultaneous discovery of oxygen by several scientists in different countries gave rise to disputes about priority. Priestley was especially persistent in recognizing himself as a discoverer In essence, these disputes have not ended yet. A detailed study of the properties of oxygen and its role in the processes of combustion and the formation of oxides led Lavoisier to the incorrect conclusion that this gas is an acid-forming principle. In 1779, Lavoisier, in accordance with this conclusion. introduced a new name for oxygen - the acid-forming principle (principe acidifiant ou principe oxygine). Lavoisier derived the word oxygine, which appears in this complex name, from the Greek - acid and “I produce”.

Contents of the article

OXYGEN, O (oxygenium), a chemical element of the VIA subgroup of the periodic table of elements: O, S, Se, Te, Po - a member of the chalcogen family. This is the most common element in nature, its content in the Earth’s atmosphere is 21% (vol.), in the earth’s crust in the form of compounds of approx. 50% (wt.) and in the hydrosphere 88.8% (wt.).

Oxygen is necessary for the existence of life on earth: animals and plants consume oxygen during respiration, and plants release oxygen through photosynthesis. Living matter contains bound oxygen not only in body fluids (in blood cells, etc.), but also in carbohydrates (sugar, cellulose, starch, glycogen), fats and proteins. Clays, rocks, consist of silicates and other oxygen-containing inorganic compounds such as oxides, hydroxides, carbonates, sulfates and nitrates.

Historical information.

The first information about oxygen became known in Europe from Chinese manuscripts of the 8th century. At the beginning of the 16th century. Leonardo da Vinci published data related to the chemistry of oxygen, not yet knowing that oxygen was an element. Reactions of oxygen addition are described in the scientific works of S. Geils (1731) and P. Bayen (1774). K. Scheele's research in 1771–1773 on the interaction of metals and phosphorus with oxygen deserves special attention. J. Priestley reported the discovery of oxygen as an element in 1774, a few months after Bayen's report of reactions with air. The name oxygenium (“oxygen”) was given to this element shortly after its discovery by Priestley and comes from the Greek words meaning “acid-producing”; this is due to the misconception that oxygen is present in all acids. The explanation of the role of oxygen in the processes of respiration and combustion, however, belongs to A. Lavoisier (1777).

The structure of the atom.

Any naturally occurring oxygen atom contains 8 protons in the nucleus, but the number of neutrons can be 8, 9, or 10. The most common of the three isotopes of oxygen (99.76%) is 16 8 O (8 protons and 8 neutrons). The content of another isotope, 18 8 O (8 protons and 10 neutrons), is only 0.2%. This isotope is used as a label or for identifying certain molecules, as well as for conducting biochemical and medico-chemical studies (a method for studying non-radioactive traces). The third non-radioactive isotope of oxygen, 17 8 O (0.04%), contains 9 neutrons and has a mass number of 17. After the mass of the carbon isotope 12 6 C was adopted by the International Commission as the standard atomic mass in 1961, the weighted average atomic mass of oxygen became 15. 9994. Until 1961, chemists considered the standard unit of atomic mass to be the atomic mass of oxygen, assumed to be 16,000 for a mixture of three naturally occurring isotopes of oxygen. Physicists took the mass number of the oxygen isotope 16 8 O as the standard unit of atomic mass, so on the physical scale the average atomic mass of oxygen was 16.0044.

An oxygen atom has 8 electrons, with 2 electrons in the internal level and 6 electrons in the outer level. Therefore, in chemical reactions, oxygen can accept up to two electrons from donors, building up its outer shell to 8 electrons and forming an excess negative charge.

Molecular oxygen.

Like most other elements, the atoms of which lack 1-2 electrons to complete the outer shell of 8 electrons, oxygen forms a diatomic molecule. This process releases a lot of energy (~490 kJ/mol) and, accordingly, the same amount of energy must be spent for the reverse process of dissociation of the molecule into atoms. The strength of the O–O bond is so high that at 2300° C only 1% of oxygen molecules dissociate into atoms. (It is noteworthy that during the formation of the nitrogen molecule N2, the strength of the N–N bond is even higher, ~710 kJ/mol.)

Electronic structure.

In the electronic structure of the oxygen molecule, as might be expected, the distribution of electrons in an octet around each atom is not realized, but there are unpaired electrons, and oxygen exhibits properties typical of such a structure (for example, it interacts with a magnetic field, being paramagnetic).

Reactions.

Under appropriate conditions, molecular oxygen reacts with almost any element except the noble gases. However, under room conditions, only the most active elements react with oxygen quickly enough. It is likely that most reactions occur only after the dissociation of oxygen into atoms, and dissociation occurs only at very high temperatures. However, catalysts or other substances in the reacting system can promote the dissociation of O 2 . It is known that alkali (Li, Na, K) and alkaline earth (Ca, Sr, Ba) metals react with molecular oxygen to form peroxides:

Receipt and application.

Due to the presence of free oxygen in the atmosphere, the most effective method for its extraction is the liquefaction of air, from which impurities, CO 2, dust, etc. are removed. chemical and physical methods. The cyclic process includes compression, cooling and expansion, which leads to air liquefaction. With a slow rise in temperature (fractional distillation method), first noble gases (the most difficult to liquefy) evaporate from liquid air, then nitrogen, and liquid oxygen remains. As a result, liquid oxygen contains traces of noble gases and a relatively large percentage of nitrogen. For many applications these impurities are not a problem. However, to obtain oxygen of extreme purity, the distillation process must be repeated. Oxygen is stored in tanks and cylinders. It is used in large quantities as an oxidizer for kerosene and other fuels in rockets and spacecraft. The steel industry uses oxygen gas to blow through the molten iron using the Bessemer method to quickly and effectively remove C, S and P impurities. Oxygen blast produces steel faster and of higher quality than air blast. Oxygen is also used for welding and cutting metals (oxy-acetylene flame). Oxygen is also used in medicine, for example, to enrich the respiratory environment of patients with difficulty breathing. Oxygen can be produced by various chemical methods, and some of them are used to obtain small quantities of pure oxygen in laboratory practice.

Electrolysis.

One of the methods for producing oxygen is the electrolysis of water containing small additions of NaOH or H 2 SO 4 as a catalyst: 2H 2 O ® 2H 2 + O 2. In this case, small hydrogen impurities are formed. Using a discharge device, traces of hydrogen in the gas mixture are again converted into water, the vapors of which are removed by freezing or adsorption.

Thermal dissociation.

An important laboratory method for producing oxygen, proposed by J. Priestley, is the thermal decomposition of heavy metal oxides: 2HgO ® 2Hg + O 2 . To do this, Priestley focused the sun's rays on mercury oxide powder. A well-known laboratory method is also the thermal dissociation of oxo salts, for example potassium chlorate in the presence of a catalyst - manganese dioxide:

Manganese dioxide, added in small quantities before calcination, allows maintaining the required temperature and dissociation rate, and MnO 2 itself does not change during the process.

Methods for thermal decomposition of nitrates are also used:

as well as peroxides of some active metals, for example:

2BaO 2 ® 2BaO + O 2

The latter method was at one time widely used to extract oxygen from the atmosphere and consisted of heating BaO in air to form BaO 2 followed by thermal decomposition of the peroxide. The thermal decomposition method remains important for the production of hydrogen peroxide.

SOME PHYSICAL PROPERTIES OF OXYGEN
Atomic number	8
Atomic mass	15,9994
Melting point, °C	–218,4
Boiling point, °C	–183,0
Density
hard, g/cm 3 (at t pl)	1,27
liquid g/cm 3 (at t kip)	1,14
gaseous, g/dm 3 (at 0° C)	1,429
air relative	1,105
critical a, g/cm 3	0,430
Critical temperature a, °C	–118,8
Critical pressure a, atm	49,7
Solubility, cm 3 /100 ml of solvent
in water (0° C)	4,89
in water (100° C)	1,7
in alcohol (25° C)	2,78
Radius, Å	0,74
covalent	0,66
ionic (O 2–)	1,40
Ionization potential, V
first	13,614
second	35,146
Electronegativity (F=4)	3,5
a Temperature and pressure at which the densities of gas and liquid are the same.

Physical properties.

Oxygen under normal conditions is a colorless, odorless and tasteless gas. Liquid oxygen is pale blue in color. Solid oxygen exists in at least three crystalline modifications. Oxygen gas is soluble in water and probably forms weak compounds such as O2HH2O, and possibly O2H2H2O.

Chemical properties.

As already mentioned, the chemical activity of oxygen is determined by its ability to dissociate into O atoms, which are highly reactive. Only the most active metals and minerals react with O 2 at high rates at low temperatures. The most active alkali (IA subgroups) and some alkaline earth (IIA subgroups) metals form peroxides such as NaO 2 and BaO 2 with O 2 . Other elements and compounds react only with the dissociation product O2. Under suitable conditions, all elements, excluding the noble gases and the metals Pt, Ag, Au, react with oxygen. These metals also form oxides, but under special conditions.

The electronic structure of oxygen (1s 2 2s 2 2p 4) is such that the O atom accepts two electrons to the outer level to form a stable outer electron shell, forming an O 2– ion. In alkali metal oxides, predominantly ionic bonds are formed. It can be assumed that the electrons of these metals are almost entirely drawn to oxygen. In oxides of less active metals and non-metals, the electron transfer is incomplete, and the negative charge density on oxygen is less pronounced, so the bond is less ionic or more covalent.

When metals are oxidized with oxygen, heat is released, the magnitude of which correlates with the strength of the M–O bond. During the oxidation of some nonmetals, heat is absorbed, which indicates their weaker bonds with oxygen. Such oxides are thermally unstable (or less stable than oxides with ionic bonds) and are often highly reactive. The table shows for comparison the values ​​of the enthalpies of formation of oxides of the most typical metals, transition metals and nonmetals, elements of the A- and B-subgroups (the minus sign means the release of heat).

Several general conclusions can be drawn about the properties of oxides:

1. Melting temperatures of alkali metal oxides decrease with increasing atomic radius of the metal; So, t pl (Cs 2 O) t pl (Na 2 O). Oxides in which ionic bonding predominates have higher melting points than the melting points of covalent oxides: t pl (Na 2 O) > t pl (SO 2).

2. Oxides of reactive metals (IA–IIIA subgroups) are more thermally stable than oxides of transition metals and nonmetals. Oxides of heavy metals in the highest oxidation state upon thermal dissociation form oxides with lower oxidation states (for example, 2Hg 2+ O ® (Hg +) 2 O + 0.5O 2 ® 2Hg 0 + O 2). Such oxides in high oxidation states can be good oxidizing agents.

3. The most active metals react with molecular oxygen at elevated temperatures to form peroxides:

Sr + O 2 ® SrO 2 .

4. Oxides of active metals form colorless solutions, while the oxides of most transition metals are colored and practically insoluble. Aqueous solutions of metal oxides exhibit basic properties and are hydroxides containing OH groups, and non-metal oxides in aqueous solutions form acids containing the H + ion.

5. Metals and non-metals of A-subgroups form oxides with an oxidation state corresponding to the group number, for example, Na, Be and B form Na 1 2 O, Be II O and B 2 III O 3, and non-metals IVA–VIIA of subgroups C, N , S, Cl form C IV O 2, N V 2 O 5, S VI O 3, Cl VII 2 O 7. The group number of an element correlates only with the maximum oxidation state, since oxides with lower oxidation states of elements are possible. In combustion processes of compounds, typical products are oxides, for example:

2H 2 S + 3O 2 ® 2SO 2 + 2H 2 O

Carbon-containing substances and hydrocarbons, when heated slightly, oxidize (burn) to CO 2 and H 2 O. Examples of such substances are fuels - wood, oil, alcohols (as well as carbon - coal, coke and charcoal). The heat from the combustion process is utilized to produce steam (and then electricity or goes to power plants), as well as for heating houses. Typical equations for combustion processes are:

a) wood (cellulose):

(C6H10O5) n + 6n O 2 ® 6 n CO2+5 n H 2 O + thermal energy

b) oil or gas (gasoline C 8 H 18 or natural gas CH 4):

2C 8 H 18 + 25O 2 ® 16CO 2 + 18H 2 O + thermal energy

CH 4 + 2O 2 ® CO 2 + 2H 2 O + thermal energy

C 2 H 5 OH + 3O 2 ® 2CO 2 + 3H 2 O + thermal energy

d) carbon (coal or charcoal, coke):

2C + O 2 ® 2CO + thermal energy

2CO + O 2 ® 2CO 2 + thermal energy

A number of C-, H-, N-, O-containing compounds with a high energy reserve are also subject to combustion. Oxygen for oxidation can be used not only from the atmosphere (as in previous reactions), but also from the substance itself. To initiate a reaction, a small activation of the reaction, such as a blow or shake, is sufficient. In these reactions, combustion products are also oxides, but they are all gaseous and expand rapidly at the high final temperature of the process. Therefore, such substances are explosive. Examples of explosives are trinitroglycerin (or nitroglycerin) C 3 H 5 (NO 3) 3 and trinitrotoluene (or TNT) C 7 H 5 (NO 2) 3.

Oxides of metals or non-metals with lower oxidation states of an element react with oxygen to form oxides of high oxidation states of that element:

Natural oxides, obtained from ores or synthesized, serve as raw materials for the production of many important metals, for example, iron from Fe 2 O 3 (hematite) and Fe 3 O 4 (magnetite), aluminum from Al 2 O 3 (alumina), magnesium from MgO (magnesia). Light metal oxides are used in the chemical industry to produce alkalis or bases. Potassium peroxide KO 2 has an unusual use because in the presence of moisture and as a result of reaction with it, it releases oxygen. Therefore, KO 2 is used in respirators to produce oxygen. Moisture from the exhaled air releases oxygen in the respirator, and KOH absorbs CO 2. Production of CaO oxide and calcium hydroxide Ca(OH) 2 – large-scale production in ceramics and cement technology.

Water (hydrogen oxide).

The importance of water H 2 O both in laboratory practice for chemical reactions and in life processes requires special consideration of this substance WATER, ICE AND STEAM). As already mentioned, during the direct interaction of oxygen and hydrogen under conditions, for example, a spark discharge, an explosion and the formation of water occur, and 143 kJ/(mol H 2 O) is released.

The water molecule has an almost tetrahedral structure, the H–O–H angle is 104° 30°. The bonds in the molecule are partially ionic (30%) and partially covalent with a high density of negative charge on oxygen and, accordingly, positive charges on hydrogen:

Due to the high strength of H–O bonds, hydrogen is difficult to split off from oxygen and water exhibits very weak acidic properties. Many properties of water are determined by the distribution of charges. For example, a water molecule forms a hydrate with a metal ion:

Water gives one electron pair to an acceptor, which can be H +:

Oxoanions and oxocations

– oxygen-containing particles having a residual negative (oxoanions) or residual positive (oxocations) charge. The O 2– ion has high affinity (high reactivity) for positively charged particles such as H +. The simplest representative of stable oxoanions is the hydroxide ion OH –. This explains the instability of atoms with a high charge density and their partial stabilization as a result of the addition of a particle with a positive charge. Therefore, when an active metal (or its oxide) acts on water, OH– is formed, and not O 2–:

2Na + 2H 2 O ® 2Na + + 2OH – + H 2

Na 2 O + H 2 O ® 2Na + + 2OH –

More complex oxoanions are formed from oxygen with a metal ion or non-metallic particle that has a large positive charge, resulting in a low-charge particle that is more stable, for example:

°C a dark purple solid phase is formed. Liquid ozone is slightly soluble in liquid oxygen, and 49 cm 3 O 3 dissolves in 100 g of water at 0 ° C. In terms of chemical properties, ozone is much more active than oxygen and is second only to O, F 2 and OF 2 (oxygen difluoride) in oxidizing properties. During normal oxidation, oxide and molecular oxygen O 2 are formed. When ozone acts on active metals under special conditions, ozonides of the composition K + O 3 – are formed. Ozone is produced industrially for special purposes; it is a good disinfectant and is used to purify water and as a bleach, improves the condition of the atmosphere in closed systems, disinfects objects and food, and accelerates the ripening of grains and fruits. In a chemistry laboratory, an ozonizer is often used to produce ozone, which is necessary for some methods of chemical analysis and synthesis. Rubber is easily destroyed even when exposed to low concentrations of ozone. In some industrial cities, significant concentrations of ozone in the air lead to rapid deterioration of rubber products if they are not protected by antioxidants. Ozone is very toxic. Constant inhalation of air, even with very low concentrations of ozone, causes headaches, nausea and other unpleasant conditions.

DEFINITION

Oxygen- the eighth element of the Periodic Table. Designation - O from the Latin “oxygenium”. Located in the second period, group VIA. Refers to non-metals. The nuclear charge is 8.

Oxygen is the most common element in the earth's crust. In a free state, it is found in the atmospheric air; in a bound form, it is part of water, minerals, rocks and all substances from which the organisms of plants and animals are built. The mass fraction of oxygen in the earth's crust is about 47%.

In its simple form, 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 mass of oxygen

DEFINITION

Relative atomic mass A r is the molar mass of an atom of a substance divided by 1/12 of the molar mass of a carbon-12 atom (12 C).

The relative atomic mass of atomic oxygen is 15.999 amu.

DEFINITION

Relative molecular weight M r is the molar mass of a molecule divided by 1/12 the molar mass of a carbon-12 atom (12 C).

This is a dimensionless quantity. It is known that the oxygen molecule is diatomic - O 2. The relative molecular mass of an oxygen molecule will be equal to:

M r (O 2) = 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 (the physical properties of oxygen are described above).

Under 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 = 2O 3 - 285 kJ.

Isotopes of oxygen

It is known that in nature oxygen can be found 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 15 O with a half-life of 120 s.

Oxygen ions

The outer energy level of the oxygen atom has six electrons, which are valence electrons:

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- ;

O 0 -1e → O 1+ .

Oxygen molecule and atom

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 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 liquid form is blue; in solid form, the element forms blue crystals. Oxygen melts at a temperature of (-218.7) °C.

Liquid oxygen at room temperature

When 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= CO₂ + 2H₂O

2СО + О₂ = 2СО₂

Oxygen as a chemical element is obtained in laboratories and industrial plants. in the laboratory there are several ways:

• decomposition (potassium chlorate);
• decomposition of hydrogen peroxide when heating the substance 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 oxygen in the air does not have, that is, it has the same chemical and physical properties. The air contains 5 times less oxygen than the same volume of pure oxygen. In the air, oxygen is mixed with large quantities of nitrogen - a gas that does not burn itself and does not support combustion. Therefore, if air oxygen has already been consumed near the flame, then the next portion of oxygen will make its way through nitrogen and combustion products. Consequently, more energetic combustion of oxygen in the atmosphere is explained by a faster supply of oxygen to the combustion site. During the reaction, the process of combining oxygen with the burning substance is carried out more energetically 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 combustion reaction of oxygen occur? This can be verified experimentally. You need to take the cylinder and turn it upside down, then place a tube with 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 insert a glass tube into it through the flame, through which oxygen gas flows. A fire will break out 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 coming out of the tube.

What results from the combustion of hydrogen and what oxide is formed? Hydrogen is oxidized to water. Droplets of condensed water vapor are gradually deposited on the walls of the cylinder. The oxidation of two hydrogen molecules takes one oxygen molecule, and two water molecules are formed. Reaction equation:

2Н₂ + O₂ → 2Н₂O

If the oxygen flows out of the tube slowly, it burns completely in the hydrogen atmosphere, and the experiment proceeds calmly.

As soon as the supply of oxygen increases so much that it does not have time to burn completely, part of it goes beyond the flame, where pockets of a mixture of hydrogen and oxygen are formed, and individual small flashes similar to explosions appear. A mixture of oxygen and hydrogen is an explosive gas.

When detonating gas is ignited, a strong explosion occurs: when oxygen combines with hydrogen, water is formed and a high temperature develops. Water vapor with surrounding gases expands greatly, creating high pressure, at which not only a fragile cylinder, but also a more durable vessel can rupture. 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 must be added. Tint the water with phenolphthalein or another suitable dye. Pour sand into a small flask and vertically insert a wire with cotton wool attached to the end into it. The flask is placed in a crystallizer with water. The cotton wool remains 10 cm above the surface of the solution.

Lightly moisten the cotton wool with alcohol, oil, hexane or other flammable liquid and set it on fire. Carefully cover the burning cotton wool with a 3-liter bottle and lower it below the surface of the lye solution. During the combustion process, 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 the air contains 20.9% oxygen. During combustion, oxygen turns into water and carbon dioxide CO₂, which is absorbed by the alkali. Reaction equation:

2NaOH + CO₂ = Na₂CO₃ + H₂O

In practice, combustion will stop before all the oxygen is consumed; part of the oxygen turns into carbon monoxide, which is not absorbed by the alkali, and part of the air leaves the bottle as a result of thermal expansion.

Attention! Do not try to repeat these experiments yourself!

Ministry of Education and Science of the Russian Federation

"OXYGEN"

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General characteristics of oxygen.

OXYGEN (lat. Oxygenium), O (read “o”), chemical element with atomic number 8, atomic mass 15.9994. In Mendeleev's periodic table of elements, oxygen is located in the second period in group VIA.

Natural oxygen consists of a mixture of three stable nuclides with mass numbers 16 (dominates in the mixture, it contains 99.759% by mass), 17 (0.037%) and 18 (0.204%). The radius of a neutral oxygen atom is 0.066 nm. The configuration of the outer electronic layer of the neutral unexcited oxygen atom is 2s2р4. The energies of sequential ionization of the oxygen atom are 13.61819 and 35.118 eV, the electron affinity is 1.467 eV. The radius of the O 2 ion is at different coordination numbers from 0.121 nm (coordination number 2) to 0.128 nm (coordination number 8). In compounds it exhibits an oxidation state of –2 (valence II) and, less commonly, –1 (valency I). According to the Pauling scale, the electronegativity of oxygen is 3.5 (the second highest among non-metals after fluorine).

In its free form, oxygen is a colorless, odorless, and tasteless gas.

Features of the structure of the O 2 molecule: atmospheric oxygen consists of diatomic molecules. The interatomic distance in the O 2 molecule is 0.12074 nm. Molecular oxygen (gaseous and liquid) is a paramagnetic substance; each O2 molecule has 2 unpaired electrons. This fact can be explained by the fact that in the molecule there is one unpaired electron in each of the two antibonding orbitals.

The dissociation energy of the O 2 molecule into atoms is quite high and amounts to 493.57 kJ/mol.

Physical and chemical properties

Physical and chemical properties: in free form it is found in the form of two modifications O 2 (“ordinary” oxygen) and O 3 (ozone). O 2 is a colorless and odorless gas. Under normal conditions, the density of oxygen gas is 1.42897 kg/m3. The boiling point of liquid oxygen (the liquid is blue) is –182.9°C. At temperatures from –218.7°C to –229.4°C there is solid oxygen with a cubic lattice (modification), at temperatures from –229.4°C to –249.3°C there is a modification with a hexagonal lattice and at temperatures below –249.3°C - cubic modification. Other modifications of solid oxygen have been obtained at elevated pressure and low temperatures.

At 20°C, the solubility of O2 gas is: 3.1 ml per 100 ml of water, 22 ml per 100 ml of ethanol, 23.1 ml per 100 ml of acetone. There are organic fluorine-containing liquids (for example, perfluorobutyltetrahydrofuran), in which the solubility of oxygen is much higher.

The high strength of the chemical bond between the atoms in the O2 molecule leads to the fact that at room temperature oxygen gas is chemically quite inactive. In nature, it slowly undergoes transformation during decay processes. In addition, oxygen at room temperature is able to react with hemoglobin in the blood (more precisely with heme iron II), which ensures the transfer of oxygen from the respiratory organs to other organs.

Oxygen reacts with many substances without heating, for example, with alkali and alkaline earth metals (corresponding oxides such as Li 2 O, CaO, etc., peroxides such as Na 2 O2, BaO 2, etc., and superoxides such as KO 2, RbO 2 are formed etc.), causes the formation of rust on the surface of steel products. Without heating, oxygen reacts with white phosphorus, with some aldehydes and other organic substances.

When heated, even slightly, the chemical activity of oxygen increases sharply. When ignited, it reacts explosively with hydrogen, methane, other flammable gases, and a large number of simple and complex substances. It is known that when heated in an oxygen atmosphere or in air, many simple and complex substances burn, and various oxides are formed, for example:

S+O 2 = SO 2; C + O 2 = CO 2

4Fe + 3O 2 = 2Fe 2 O 3; 2Cu + O 2 = 2CuO

4NH 3 + 3O 2 = 2N 2 + 6H 2 O; 2H 2 S + 3O 2 = 2H 2 O + 2SO 2

If a mixture of oxygen and hydrogen is stored in a glass vessel at room temperature, then the exothermic reaction to form water

2H 2 + O 2 = 2H 2 O + 571 kJ

proceeds extremely slowly; According to calculations, the first drops of water should appear in the vessel in about a million years. But when platinum or palladium (playing the role of a catalyst) is introduced into a vessel with a mixture of these gases, as well as when ignited, the reaction proceeds with an explosion.

Oxygen reacts with nitrogen N2 either at high temperature (about 1500-2000°C), or by passing an electric discharge through a mixture of nitrogen and oxygen. Under these conditions, nitric oxide (II) is reversibly formed:

N2 + O2 = 2NO

The resulting NO then reacts with oxygen to form brown gas (nitrogen dioxide):

2NO + O 2 = 2NO2

Of non-metals, oxygen does not directly interact with halogens under any circumstances, and of metals - with noble metals - silver, gold, platinum, etc.

Binary oxygen compounds in which the oxidation state of oxygen atoms is –2 are called oxides (formerly called oxides). Examples of oxides: carbon monoxide (IV) CO 2, sulfur oxide (VI) SO 3, copper oxide (I) Cu 2 O, aluminum oxide Al 2 O 3, manganese oxide (VII) Mn 2 O 7.

Oxygen also forms compounds in which its oxidation state is –1. These are peroxides (the old name is peroxides), for example, hydrogen peroxide H 2 O 2, barium peroxide BaO 2, sodium peroxide Na 2 O 2 and others. These compounds contain a peroxide group - O - O -. With active alkali metals, for example, potassium, oxygen can also form superoxides, for example, KO 2 (potassium superoxide), RbO 2 (rubidium superoxide). In superoxides, the oxidation state of oxygen is –1/2. It may be noted that superoxide formulas are often written as K 2 O 4, Rb 2 O 4, etc.

With the most active nonmetal fluorine, oxygen forms compounds in positive oxidation states. So, in the compound O 2 F 2 the oxidation state of oxygen is +1, and in the compound O 2 F - +2. These compounds do not belong to oxides, but to fluorides. Oxygen fluorides can be synthesized only indirectly, for example, by the action of fluorine F2 on dilute aqueous solutions of KOH.

History of discovery

The history of the discovery of oxygen, like nitrogen, is connected with the study of atmospheric air that lasted several centuries. The fact that air by its nature is not homogeneous, but includes parts, one of which supports combustion and respiration, and the other does not, was known back in the 8th century by the Chinese alchemist Mao Hoa, and later in Europe by Leonardo da Vinci. In 1665, the English naturalist R. Hooke wrote that the air consists of the gas contained in nitrate, as well as inactive gas, which makes up most of the air. The fact that air contains a life-sustaining element was known to many chemists in the 18th century. The Swedish pharmacist and chemist Karl Scheele began studying the composition of air in 1768. For three years, he decomposed saltpeter (KNO 3, NaNO 3) and other substances by heating and obtained “fiery air” that supported respiration and combustion. But Scheele published the results of his experiments only in 1777 in the book “Chemical Treatise on Air and Fire.” In 1774, the English priest and naturalist J. Priestley obtained a gas that supports combustion by heating “burnt mercury” (mercuric oxide HgO). While in Paris, Priestley, who did not know that the gas he obtained was part of the air, reported his discovery to A. Lavoisier and other scientists. By this time, nitrogen had also been discovered. In 1775, Lavoisier came to the conclusion that ordinary air consists of two gases - a gas necessary for breathing and supporting combustion, and a gas of the “opposite nature” - nitrogen. Lavoisier called the combustion-supporting gas oxygene - “acid-forming” (from the Greek oxys - sour and gennao - I give birth; hence the Russian name “oxygen”), since he then believed that all acids contain oxygen. It has long been known that acids can be both oxygen-containing and oxygen-free, but the name given to Lavoisier’s element has remained unchanged. For almost a century and a half, 1/16 of the mass of an oxygen atom served as a unit for comparing the masses of different atoms with each other and was used to numerically characterize the masses of atoms of various elements (the so-called oxygen scale of atomic masses).

Occurrence in nature: oxygen is the most common element on Earth; its share (in various compounds, mainly silicates) accounts for about 47.4% of the mass of the solid earth's crust. Sea and fresh waters contain a huge amount of bound oxygen - 88.8% (by mass), in the atmosphere the content of free oxygen is 20.95% (by volume). The element oxygen is part of more than 1,500 compounds in the earth's crust.

Receipt:

Currently, oxygen is produced in industry by separating air at low temperatures. First, the air is compressed by a compressor, which heats up the air. The compressed gas is allowed to cool to room temperature and then allowed to expand freely. As it expands, the temperature of the gas drops sharply. Cooled air, the temperature of which is several tens of degrees lower than the ambient temperature, is again compressed to 10-15 MPa. Then the released heat is removed again. After several compression-expansion cycles, the temperature drops below the boiling point of both oxygen and nitrogen. Liquid air is formed, which is then subjected to distillation. The boiling point of oxygen (–182.9°C) is more than 10 degrees higher than the boiling point of nitrogen (–195.8°C). Therefore, nitrogen evaporates from the liquid first, and oxygen accumulates in the remainder. Due to slow (fractional) distillation, it is possible to obtain pure oxygen, in which the nitrogen impurity content is less than 0.1 percent by volume.