Oxygen in nature and its use. Report on the use of oxygen. Oxygen short message
The report on the topic “Uses of Oxygen”, summarized in this article, will tell you about the areas of industry in which this invisible substance brings incredible benefits.
Message about oxygen use
Oxygen is an integral part of the life of all living organisms and chemical processes on the planet. In this article we will look at the most common uses of oxygen:
Use of oxygen in medicine
In this area, it is extremely important: the chemical element is used to support the life of people suffering from difficulty breathing and to treat certain ailments. It is noteworthy that when normal pressure You cannot breathe pure oxygen for a long time. This is not safe for health.
Application of oxygen in the glass industry
This chemical element is used in glass melting furnaces as a component that improves combustion in them. Also, thanks to oxygen, the industry reduces nitrogen oxide emissions to a level that is safe for life.
Use of oxygen in the pulp and paper industry
This chemical element is used in alcoholization, delignification and other processes, such as:
- Whitening paper
- Cleaning Wastewater
- Preparation drinking water
- Intensification of combustion of waste incinerators
- Tire recycling
Application of oxygen in aviation
Since a person cannot breathe outside the atmosphere without oxygen, he needs to take a supply of this useful element with him. Artificially produced oxygen is used by people for breathing in an alien environment: in aviation during flights, in spacecraft.
Use of oxygen in nature
In nature, there is an oxygen cycle: during the process of photosynthesis, plants turn into carbon dioxide and water into organic compounds. This process is characterized by the release of oxygen. Like humans and animals, plants consume oxygen from the atmosphere at night. The oxygen cycle in nature is determined by the fact that humans and animals consume oxygen, and plants produce it during the day and consume it at night.
Application of oxygen in metallurgy
The chemical and metallurgical industries require pure oxygen, not atmospheric oxygen. Every year, enterprises around the world receive more than 80 million tons of this chemical element. It is used up in the process of producing steel from scrap metal and cast iron.
What is the use of oxygen in mechanical engineering?
In construction and mechanical engineering it is used for cutting and welding metals. These processes are carried out at high temperatures.
Use of oxygen in life
In life, a person uses oxygen in various areas, such as:
- Growing fish in pond farms (the water is saturated with oxygen).
- Water treatment during food production.
- Disinfection of storage facilities and production premises with oxygen.
- Development of oxygen cocktails for animals so that they gain weight.
Human use of oxygen in electricity
Thermal and power plants that run on oil, natural gas or coal use oxygen to burn the fuel. Without it, all industrial production plants simply would not work.
The abstract was completed by: 9th grade student “A” Vasilyeva N.
Ministry of Education of the Russian Federation
Secondary school No. 34.
Khabarovsk
I . Introduction.
If you look at the table of the periodic system D.I. Mendeleev and look at group VI, you can see that it contains elements whose atoms have 6 valence electrons and their highest oxidation state in compounds is +6. Group VI is divided into two subgroups – main and secondary. The main one includes elements of small and large periods: O (oxygen), S (sulfur), Se (selenium), Te (tellurium), Po (polonium); in the secondary - elements of only long periods: Cr (chromium), Mo (molybdenum), W (tungsten). Such a distribution indicates that within even one group there are elements that are closer in their properties to each other and less similar.
Indeed, in the main subgroup there are elements that are mainly non-metallic in nature. These properties are most pronounced in oxygen and sulfur. Selenium and tellurium occupy an intermediate position between metals and non-metals. In terms of chemical properties, they are closer to non-metals. In polonium, the heaviest element of the subgroup, radioactive and relatively short-lived, the metallic character is more pronounced, but in some properties it is close to tellurium. In accordance with this, during the transition from oxygen to polonium, a great diversity is observed in the structural types of crystal lattices, both in simple substances and in their compounds.
Oxygen, sulfur, selenium and tellurium are grouped together as “chalcogens,” which in Greek means “generating ores.” These elements are found in many ores. Thus, most metals in nature are found in bound state in the form of sulfides, oxides, selenides, etc. For example, the most important ores of iron and copper are red iron ore Fe2O3, magnetic iron ore Fe3O4, pyrite FeS2, red magnetic ore Cu2O, copper luster Cu2S. All of the above ores contain elements of group VI.
A side subgroup consists of metals: chromium, molybdenum and tungsten. In most physical and chemical properties, molybdenum and tungsten are similar to each other and slightly different from chromium.
II . Characteristics of elements VI subgroups.
The chemical properties of elements are determined primarily by the structure of their outer electronic layers (energy levels). The diagram shown (Fig. 1) shows the sequential filling of layers of atoms of group VI elements with electrons.
The maximum possible number of electrons in layers (Z) is determined by the formula: Z=2n2, where n is the layer number.
According to this dependence, the number of electrons should be equal: in the first layer - 2, in the second - 8, in the third - 18, in the fourth - 32, etc. However, more than 32 electrons in a layer of atoms of any currently known elements have not been found.
8 2 6 1 13 8 2 +24
16 2 8 6 1 13 18 8 2 +42
34 2 8 18 6 2 12 32 18 8 2 +74
84 2 8 18 32 18 6
Rice. 1. Scheme of the structure of atoms of elements of group VI.
The electronic structure of atoms of group VI elements can be presented as follows (Table 1).
Table 1
Electronic configurations of atoms of group VI elements
16S 1s2 2s2 2p6 3s2 3p4
34Se 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p4
52Te 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p4
84Po 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 6s2 6p4
24Cr 1s2 2s2 2p6 3s2 3p6 3d5 4s1
42Mo 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d5 5s1
74W 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d4 6s2
If you look closely at the structures depicted, you will notice that the sum of the electrons of the last two sublevels in the atoms of all these elements is equal to 6. This is the reason for the common chemical properties. But it is also visible a big difference in electronic configurations between atoms of elements of the main and secondary subgroups.
The atoms of the elements of the main subgroup on the outer electronic layer have the same number of electrons - 6. The latter are located on the s- and p-sublevels (s2 p4) and take part in the formation of chemical bonds.
Elements in the atoms of which the p-sublevel of the outer layer is filled with electrons are called p-elements. These are oxygen, sulfur, tellurium, selenium and polonium: in their atoms the s-sublevel is filled and the p-sublevel of the outer layer is filled with electrons. The atoms of these elements have a specific tendency to attract additional (two) electrons compared to neutral atoms. It manifests itself in their compounds with non-metals (CuS, Na2S, K2Te) and in the existence of negative ions in molten salts of the most active metals (S2-, Se2-, Te2-).
It should be noted that the penultimate layer of tellurium and polonium atoms is not completed, unlike oxygen, sulfur and selenium, where it is completely filled. But despite the common properties of p-elements of group VI, there are some differences between them.
Chromium and molybdenum atoms have 1 electron each in the outer electron layer and 13 electrons each in the penultimate one. For tungsten atoms, the number of electrons in the outer layer increases to 2, and in the penultimate layer decreases to 12. Elements in the atoms of which the d-sublevel of the layer adjacent to the outer layer is filled with electrons are called d-elements. These are chromium, molybdenum and tungsten.
Hence, outer layer elements of the secondary subgroup (d-elements) are represented only by the s-sublevel and in the formation of a chemical bond, in addition to 1-2 electrons from this sublevel, a certain number of electrons from the d-sublevel of the penultimate layer participate. These differences affect the chemical properties of d-elements. First of all, these are metals. Their specific properties are associated with a large number outer electrons in atoms. Under certain conditions, for example in aqueous solutions of acids, 2 or 3 electrons are completely transferred to other atoms, and the metal atoms are converted, respectively, into two- or three-charged hydrated cations. The ability of metal atoms to partially or completely displace their electrons to other atoms determines the formation of strong compounds with non-metals, the displacement of hydrogen from acids, the basic nature of oxides and hydroxides, etc.
So, the number and state of electrons at the outer levels of an atom is one of the most important signs of chemical nature. However, chemical individuality individual elements– their metallic and non-metallic activity – is determined not only by the outer electronic structures of the atoms, but also by the structure of their atoms as a whole: the charge of the nucleus, the number and state of electrons in individual layers, and the radii of the atoms.
The quantitative characteristics of the chemical properties of elements are determined by the structure of the outer electronic layer, which may include electrons from one layer of different sublevels or sometimes from adjacent sublevels of two adjacent layers (for example, in elements of side subgroups).
The oxygen atom has two unpaired p-electrons out of four, and therefore the formation of two electron pairs when interacting with a particular atom does not require excitation energy (a). Cells correspond to certain states (orbitals) of electrons at each sublevel; sublevels are characterized by different shapes of electron clouds. Electrons in the diagram are shown by arrows. In all compounds, oxygen has a typical oxidation state of –2, with the exception of O+2F2 and O+4O2 (ozone).
For analogues of oxygen (sulfur, selenium, tellurium and polonium) the situation is completely different. For example, in the outer electron layer of a sulfur atom there are also 6 electrons, but unlike oxygen there can be 18, i.e. there are vacancies (b). Therefore, in order for sulfur to react and acquire an oxidation state of +4 or +6 in compounds, a slight excitation of the atom is necessary, because electrons are transferred to the d-sublevel of the same energy layer, which undoubtedly requires a certain amount of energy (c and d).
The same explanation can be applied to selenium, tellurium, polonium and the chromium subgroup metals. These elements may show varying degrees oxidation: from -2 to +6.
table 2
Possible oxidation states of atoms of group VI elements
Table 2 shows the oxidation states of atoms of group VI elements.
The elements of the main subgroup have wide limits for changing the oxidation state: from the maximum possible negative -2 to the maximum positive, corresponding to the group number.
When going from oxygen to tellurium and from chromium to tungsten, the melting and boiling points increase. Oxygen has the lowest boiling and melting points, since the polarizability of its molecule is low. This can also explain the poor solubility of oxygen in water: 5 volumes of O2 in 100 volumes of H2O at 0°C.
Tungsten is the most refractory and high-boiling among all metals. Its boiling point is almost 6000°C, as on the surface of the Sun. Tungsten melts at 3380°C. At this temperature, most metals turn to steam.
High temperatures melting of group VI metals is explained by the fact that they have a high electron density, that is, a large number of free electrons per unit volume. As is known, the metallic bond is caused by the interaction of free electrons with ion atoms. In group VI metals, the number of free electrons reaches up to six for each atom ion, which is why they are refractory.
I will talk in more detail about oxygen.
III . History of the discovery of oxygen.
The discovery of oxygen 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 hundreds of years the combustion process remained unclear. Oxygen was discovered almost simultaneously by two outstanding chemists. half of the XVIII V. - Swede Karl Scheele and Englishman Joseph Priestley. K. Scheele was the first to obtain oxygen, but his work “On Air and Fire,” in which this gas was described, appeared somewhat later than D. Priestley’s message.
K. Scheele and D. Priestley discovered a new element, but did not understand its role in the processes of combustion and respiration. Until the end of their days, they remained defenders of the phlogiston theory: combustion was interpreted as the disintegration of a combustible body with the release of phlogiston, in which each combustible substance turned into non-combustible:
zinc = phlogiston + zinc scale
(flammable) (non-flammable)
Hence, metals, sulfur and other simple substances were considered complex and, conversely, complex substances were considered simple (lime, acids, etc.).
Proponents of the phlogiston theory explained the need for air for combustion by the fact that phlogiston does not simply disappear during combustion, but combines with air or some part of it. If there is no air, then combustion stops because phlogiston has nothing to connect with.
F. Engels wrote about the discovery of K. Scheele and D. Priestley: both “they did not know what was in their hands... The element that was destined to overthrow all phlogistic views and revolutionize chemistry disappeared in their hands completely fruitlessly.” Further, F. Engels wrote that the discovery of oxygen belongs to Lavoisier, since K. Scheele and D. Priestley did not even know what they were describing.
The liberation of chemistry from the phlogiston theory occurred as a result of the introduction into chemistry precise methods research, which began with the works of M.V. Lomonosov. In 1745-1748 M.V. Lomonosov experimentally proved that combustion is a reaction of substances combining with air particles.
Ten years (1771-1781) were spent by the French chemist Antoine Lavoisier to confirm the validity of the theory of combustion as a chemical interaction of various substances with oxygen. Starting to study the phenomena of combustion and “burning” of metals, he wrote: “I propose to repeat everything done by my predecessors, taking all possible precautions in order to combine what is already known about bound or liberated air with other facts and give a new theory. The works of the mentioned authors, if considered from this point of view, provide me with individual links in the chain... But many experiments must be done in order to obtain a complete sequence.” The corresponding experiments, begun in October 1772, were carried out by A. Lavoisier strictly quantitatively, with careful weighing of the initial and final reaction products. He heated mercury in a sealed retort and observed a decrease in the volume of air in it and the formation of red flakes of “mercury scale.” In another retort, he decomposed the “mercury scale” obtained in the previous experiment, obtained mercury and a small volume of that gas, which D. Priestley called “dephlogisticated air”, and concluded: how much air is consumed to convert mercury into scale, so much of it is released again during decomposition of scale.
The remaining air in the retort, which did not participate in the reaction, began to be called nitrogen, which meant lifeless (translated from the Greek “a” - negation, “zoe” - life). The gas formed as a result of the decomposition of “mercury scale” exhibited properties opposite to nitrogen - it supported respiration and combustion. Therefore, A. Lavoisier called it “vital”. Later he replaced this name with the Latin word “oxygenum”, borrowed from Greek language, where the word “oxys” means sour, and “gennao” - I give birth, produce (giving birth to acid). The name of the element is literally translated into Russian - “oxygen”.
So, in 1777, the essence of combustion was clarified. And the need for phlogiston - “fiery matter” - disappeared. The oxygen theory of combustion replaced the phlogiston theory.
IV . Biological role oxygen.
Oxygen is the most common element on Earth; its share (in various compounds, mainly silicates) accounts for about 47.4% of the mass of solids. 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.
Oxygen is the main biogenic element that is part of the molecules of all the most important substances that provide the structure and function of cells - proteins, nucleic acids, carbohydrates, lipids, as well as many low-molecular compounds. Every plant or animal contains much more oxygen than any other element (on average about 70%). Human muscle tissue contains 16% oxygen, bone tissue - 28.5%; In total, the body of an average person (body weight 70 kg) contains 43 kg of oxygen. Oxygen enters the body of animals and humans mainly through the respiratory organs (free oxygen) and with water (bound oxygen). The body's need for oxygen is determined by the level (intensity) of metabolism, which depends on the mass and surface of the body, age, gender, nutrition, external conditions etc. In ecology, the ratio of total respiration (that is, total oxidative processes) of a community of organisms to its total biomass is determined as an important energy characteristic.
Small amounts of oxygen are used in medicine: oxygen (from so-called oxygen pillows) is given to patients who have difficulty breathing for some time. However, it must be borne in mind that prolonged inhalation of air enriched with oxygen is dangerous to human health. High concentrations of oxygen cause the formation of free radicals in tissues, disrupting the structure and function of biopolymers. They also have a similar effect on the body. ionizing radiation. Therefore, a decrease in the oxygen content (hypoxia) in tissues and cells when the body is irradiated with ionizing radiation has a protective effect - the so-called oxygen effect. This effect is used in radiation therapy: increasing the oxygen content in the tumor and decreasing its content in the surrounding tissues increases radiation damage to tumor cells and reduces damage to healthy ones. For some diseases, saturation of the body with oxygen under high pressure is used - hyperbaric oxygenation.
V . Physical and chemical properties of oxygen.
The chemical element oxygen forms two simple substances - oxygen O2 and O3, which have different physical properties.
Oxygen O2 is a colorless and odorless gas. Its molecule is O2. It is paramagnetic (attracted by a magnet) because it contains two unpaired electrons. The structure of the oxygen molecule can be represented in the form of the following structural formulas:
O - O or O - O
Atmospheric oxygen consists of diatomic molecules. The interatomic distance in the O2 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 O2 molecule into atoms is quite high and amounts to 493.57 kJ/mol.
The oxygen molecule O2 is quite inert. The stability of the oxygen molecule and the high activation energy of most oxidation reactions mean that at low and room temperatures, many reactions involving oxygen proceed at a barely noticeable rate. Only when conditions are created for the appearance of radicals - O - or R-O-O-, which excite the chain process, does oxidation proceed quickly. In this case, for example, catalysts are used that can accelerate oxidative processes.
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 - modification with a hexagonal lattice and at temperatures below -249.3°C - cubic modification. At high blood pressure And low temperatures Other modifications of solid oxygen have also been obtained.
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, which ensures the transfer of oxygen from the respiratory organs to other organs.
Oxygen interacts with many substances without heating, for example, with alkali and alkaline earth metals (the corresponding oxides like Li2O, CaO, etc., peroxides like Na2O2, BaO2, etc., and superoxides like KO2, RbO2, etc. are formed), causing the formation 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.
Ordinary atmospheric oxygen consists of a mixture of three isotopes: 16O (99.7%), 17O (0.01%), 18O (0.2%). Due to the fact that the content of the 17O and 18O isotopes in oxygen is small compared to the 16O isotope, the atomic mass of oxygen is taken to be 15.9994 cu. e.
Depending on the natural conditions the isotopic composition of oxygen can change, sometimes becoming enriched in heavy isotopes, sometimes depleted in them. Thus, water molecules H216O pass into the vapor state relatively more easily than molecules H217O and H218O. Therefore, the composition of water vapor evaporating from the sea includes oxygen with a relatively lower content of heavy isotopes than the oxygen remaining in sea water.
With the help of atoms of the heavy oxygen isotope 18O, it was possible to determine the “origin” of oxygen released by plants during photosynthesis. It was previously thought that this was oxygen released from carbon monoxide molecules, not water. It has now become known that plants bind oxygen from carbon monoxide and return oxygen from water to the atmosphere.
Oxygen forms compounds with all elements except some noble gases (helium, neon, argon). Thus, oxygen reacts with most metals already at room temperature, for example:
2Na° + O2° = Na2+102-2
Na° -1(е) Na+1 2 reducing agent
O2° +2(е) 2 2O-2 oxidizing agent
2Zn° + O2° = 2Zn+2O-2
Zn° -2(е) Zn+2 reducing agent
O2° +2(е) 2 2O-2 oxidizing agent
Oxygen usually reacts with nonmetals when heated. Thus, oxygen reacts actively with phosphorus at a temperature of 60°C:
4Р° + 502° = 2Р2+505-2
P° -5(е) P+5 2 reducing agent
O2° +2(е) 2 2O-2 5 oxidizing agent
with sulfur - at a temperature of about 250°C:
S° + 02° = S+402-2
S° -4(е) S+4 reducing agent
O2° +2(е) 2 2O-2 2 oxidizing agent
with carbon (in the form of graphite) - at 700-800°C:
С° + О2° = С+4О2-2
C° -4(е) C+4 reducing agent
O2° +2(е) 2 2O-2 2 oxidizing agent
The interaction of oxygen with nitrogen begins only at 1200°C or in an electrical discharge:
N2 + O2 2NO - Q.
Oxygen also reacts with many complex compounds, for example, it reacts with nitrogen oxides already at room temperature:
2N+2O + O2° = 2N+4O2-2
N+2 -2(е) N+4 1 reducing agent
O2° +2(е) 2 2O-2 2 oxidizing agent
Hydrogen sulfide, reacting with oxygen when heated, gives sulfur:
2H2S-2 + O2° = 2S° + 2H2O-2
S-2 -2(е) S° reducing agent
O2° +2(е) 2 2O-2 oxidizing agent
or sulfur(IV) oxide
2H2S + 3О2 = 2SO2 + 2Н2О
depending on the ratio between oxygen and hydrogen sulfide.
In the above reactions, oxygen is the oxidizing agent. Most oxidation reactions involving oxygen release heat and light, a process called combustion.
Ozone is an allotropic modification of oxygen. Its molecule is triatomic - O3. Its structure can be represented by the following structural formula:
Any change in the number or arrangement of the same atoms in a molecule entails the appearance of a qualitatively new substance with different properties. Ozone has different properties from oxygen. Under normal conditions it is a gas of blue color, with a strong irritating odor. Its name comes from Greek word"ozane" which means smell. It's toxic. Unlike oxygen, the ozone molecule is characterized by a large molecular weight, polarizability and polarity. Therefore, ozone has a higher boiling point (-111.9°C) than oxygen (-182.9°C), intense color and better solubility in water.
Under natural conditions, ozone is formed from oxygen during lightning discharges, and at an altitude of 10-30 km - under the influence of ultraviolet sun rays. It blocks life-harming ultraviolet radiation from the Sun. In addition, ozone absorbs the Earth's infrared rays, preventing it from cooling. Consequently, the allotropic form of oxygen - ozone - plays a large role in preserving life on Earth.
The formation of ozone is accompanied by the release of atomic oxygen. These are basically chain reactions in which the appearance of an active particle (usually denoted by *) causes a large number (chain) of successive transformations of inactive molecules, for example O2. The chain reaction of ozone formation from oxygen can be expressed by the following diagram:
*O2 + O2 = O3 + O
O + O2 = O3,
or in total:
In technology, ozone is produced by electrical discharges in ozonizers.
The O3 molecule is unstable, and at high concentrations, ozone disintegrates explosively:
The oxidative activity of ozone is much higher than that of oxygen. For example, already under normal conditions, ozone oxidizes such low-active simple substances as silver and mercury with the formation of their oxides and oxygen:
8Ag + 2O3 = 4Ag2O + O2
As a strong oxidizing agent, ozone is used to purify drinking water and to disinfect air. Air coniferous forests It is considered useful because it contains a small amount of ozone, which is formed during the oxidation of the resin of coniferous trees.
An even stronger oxidizing agent than oxygen O2 is ozone O3 (allotropic modification of oxygen). It is formed in the atmosphere during lightning discharges, which explains the specific smell of freshness after a thunderstorm.
In laboratories, ozone is produced by passing a discharge through oxygen (endothermic reaction):
302 203 - 284 kJ.
When ozone reacts with a solution of potassium iodide, iodine is released, whereas this reaction does not occur with oxygen:
2KI + 03 + H20 = I2 + 2KON + 02.
The reaction is often used qualitatively for the detection of I- or ozone ions. To do this, starch is added to the solution, which gives a characteristic blue complex with released iodine. The reaction is also qualitative because ozone does not oxidize Cl- and Br- ions
There is another modification of oxygen - tetraatomic (O4):
This modification is formed by the weak interaction of two oxygen molecules. The content of tetraatomic molecules in gaseous oxygen under normal conditions is only 0.1% of total number molecules, in liquid and solid oxygen - up to 50%. There is a balance:
At low temperatures it is shifted to the right, i.e. towards the formation of O4 molecules. Structural changes in molecules cause differences in the properties of substances. Thus, liquid and solid oxygen, unlike gaseous oxygen, are colored blue.
When heated, oxygen reacts with hydrogen to form water. When a mixture of both gases is ignited in volumetric proportions of 2:1 (explosive gas), the reaction occurs explosively. But it can also proceed calmly if this mixture is brought into contact with a very small amount of finely divided platinum, which plays the role of a catalyst:
2H2 + O8 = 2H20 + 572.6 kJ/mol
Oxygen can directly oxidize all metals. If the metal is highly volatile, the oxidation process usually occurs in the form of combustion. Combustion of low-volatile metals in oxygen can be carried out under the condition of high volatility of the resulting oxide. The efficiency of this process depends on the reducing activity of the metal and is characterized by the heat of formation of the resulting product. The products of interaction of metals with oxygen (oxides) can be basic, acidic or amphoteric.
When some active metals burn in oxygen, sometimes not their oxides are formed, but superoxides and peroxides. Thus, when potassium and rubidium burn, superoxides of these metals are formed:
This is due to the fact that an oxygen molecule can gain or lose electrons to form molecular ions such as O2-2, O2- and O2+. The addition of one electron to oxygen causes the formation of superoxide ion O2:
O - O + e = [ O - O ] -
The presence of an unpaired electron in the O2- ion determines the paramagnetism of superoxides.
By adding two electrons, the oxygen molecule
rotates into the peroxide ion O2-2, in which the bond atoms
We have one two-electron bond, and therefore it is diamagnetic:
O - O + 2е = [ O - O ]-2
For example, the interaction of barium with oxygen leads to the formation of peroxide BaO2:
Ba + O2 = BaO2
VI. Obtaining oxygen.
The variety of chemical compounds containing oxygen and their availability make it possible to obtain oxygen different ways. All methods of producing oxygen can be divided into two groups: physical and chemical. Most of them are chemical, that is, the production of oxygen is based on certain reactions. For example, when especially pure oxygen is needed, it is obtained from water by decomposing it. Let's consider this method.
Electrodes, most often platinum, are lowered into a vessel filled with electrolytes (distilled water acidified with sulfuric acid), and electricity. Positively charged hydrogen ions move to the negatively charged electrode (cathode), and negatively charged hydroxide ions OH- and sulfate ions SO42- move to the positively charged electrode (anode). The ions are discharged at the electrodes. It should be noted that the discharge of H+ and OH- ions occurs much more easily than the sulfate ions SO42- Thus, hydrogen is released at the cathode, and oxygen at the anode:
4Н+ + 4е - 2Н2
4OH- - 4е - 2H2O + O2
The released gases are collected in different vessels or used directly.
In a school laboratory, it is more convenient to use an alkali solution as an electrolyte. Then the electrodes can be made from iron wire or sheet. In an alkaline environment, water molecules are directly subjected to discharge at the cathode:
H2O + e - H° + H-
Н° + Н° - H2
For the experiment, a laboratory electrolyzer is used. This is a U-shaped glass tube into which electrodes are soldered. The electrolytic method produces fairly pure oxygen (0.1% impurities).
Let's consider another chemical method for producing oxygen. If barium oxide BaO is heated to 540°C, it adds atmospheric oxygen to form barium peroxide BaO2. The latter decomposes when heated to 870°C, and oxygen is released:
2BaO + O2 = 2BaO2
2BaO2 = 2BaO + O2
Barium peroxide acts as an oxygen carrier.
In the last century, plants were developed to produce oxygen using this method. They included vertically located containers that had a heating system. A current of air was passed through barium oxide heated to 400 - 500°C. After the formation of barium peroxide, the air supply was stopped, and the containers were heated to 750°C (decomposition temperature of BaO2).
With the development of technology for obtaining low temperatures, it was developed physical method obtaining oxygen from atmospheric air. It is based on deep cooling of air and the use of differences in boiling points of gases that make up the air.
Liquid air produced in refrigeration units is a mixture consisting of 79% nitrogen and 21% oxygen by volume. Liquid nitrogen boils at a temperature of - 195.8°C, and liquid oxygen boils at a temperature of - 182.9°C. Their separation is based on the difference in boiling temperatures of nitrogen and oxygen. To completely separate liquid oxygen and gaseous nitrogen, repeated evaporation of liquid air is used, accompanied by condensation of its vapor. This process is called fractional distillation or rectification. Currently, this method has become the main way to obtain technical oxygen (cheap raw materials and high productivity of installations). Liquid oxygen is stored and transported in tanks and tanks specially adapted for this purpose, equipped with good thermal insulation.
Since the physical method of producing oxygen is widely used in industry, chemical methods of production have practically lost their technical significance and are used to obtain oxygen in the laboratory.
In connection with developing scientific and technological progress, people around the world are beginning to worry about the fate of oxygen and atmospheric pollution. In many cities it is already becoming difficult to breathe. According to world statistics, all cars emit up to 600 thousand tons of toxic carbon monoxide CO into the air in just one hour of operation. When 1 ton of gasoline is burned in a car, 600 kg of carbon monoxide CO is formed. Currently, the global automobile fleet numbers 190 million vehicles. According to experts, in 1980 their number will exceed 200 million. These figures make you think.
Air poisoning from automobile exhaust gases has become alarming in cities such as Tokyo, London, New York, Paris, Rome, and Moscow. In addition, the atmosphere is polluted by other harmful gases (SO2, H2S), ash, smoke emitted by many enterprises. As a result, over the past 100 years, the number of sunny days around industrial centers has decreased by a quarter: where there were 200, there were 150. In all major cities world, as a result of thick dirty fogs, solar illumination has decreased compared to the beginning of the 20th century. by 10-30%. In London in 1952, about 4,000 people died in a few days of dirty and unbreathable fog in the air. Therefore, the fight for clean air has become one of the current problems modern hygiene.
It is known that green plants are unsurpassed cleaners and sanitizers. earth's atmosphere. Photosynthesis is the only process that has maintained the oxygen cycle in the Earth's atmosphere for about 2 billion years. Green plants are a gigantic laboratory that produces oxygen and absorbs carbon monoxide CO2. Scientists have calculated that plants globe annually absorb about 86.5 billion tons of CO2 oxide. In this regard, the creation of green parks around large cities, the arrangement of gardens, the layout of squares and flower beds is an integral part of modern urban planning, as necessary as the installation of water supply and street lighting. It is estimated that in the green areas of Moscow, Leningrad, and Kharkov, air dust levels are 2-3 times less than on adjacent streets.
Over the past few years, Russia has faced an acute problem forest fires. Thousands of hectares forest plantations die in fire. I believe that if emergency measures are not taken to extinguish fires, restore forest areas waiting for us soon ecological catastrophy. Nature reserves and forests are burning and dying unique plants, animals. In the warm season, cities, villages... are shrouded in smoke. Harmful substances are found in large quantities in the air we breathe. In connection with which various chronic diseases arise or worsen in people, immunity decreases. Children are born with congenital malformations, immunodeficiency, damage to the central nervous system...
Nature conservation and reserves have existed for a long time. But probably on at this stage development of our country, this issue remained on last place. It is necessary for all people to come to their senses and take care of our nature. After all, 95% of all forest fires are caused by them.
VII . Use of oxygen.
The use of any substance is associated with their physical and chemical properties, as well as their distribution in nature.
The amount of metal produced per capita is one measure of the level of industrial development in each country. The smelting of ferrous and non-ferrous metals is impossible without oxygen.
Now in our country, only ferrous metallurgy absorbs over 60% of the oxygen produced. But oxygen is also used in non-ferrous metallurgy.
Oxygen intensifies not only pyrometallurgical processes, but also hydrometallurgical ones, where the main process of extracting metals from ores or their concentrates is based on the action of special reagents on aqueous solutions. Thus, currently the main method of extracting gold from ores is cyanidation. It allows you to extract up to 95% of gold from gold ores and is therefore used even when processing ores with low gold content. The process of dissolving gold contained in ores is a very labor-intensive operation. It turned out that the dissolution of this metal can be significantly accelerated if pure oxygen is used instead of air. Gold in cyanide solutions forms a complex compound Na, which is then treated with zinc, and as a result gold is released:
4Аu + 8NaCN + 2H2O + O2 = 4Na + 4NaOH
2Na [Аu(CN)2] + Zn = Na2 + 2Аu
This method of extracting gold from ores was developed by the Russian engineer P.R. Bagration, a relative of the hero Patriotic War 1812
Oxygen is widely used in the chemical industry. About 30% of the oxygen produced in our country is consumed for the needs of this industry. Replacing air with oxygen during the production of sulfuric acid by contact method increases the productivity of the installation by five to six times. But this is not the only benefit of using oxygen instead of air. Pure oxygen makes it possible to obtain 100 percent sulfur oxide without carrying out additional labor-intensive operations that are necessary when using air as an oxidizer.
When producing nitric acid by the catalytic oxidation of ammonia, oxygen is also used as an oxidizing agent. If its content in the air is increased to 25%, then the productivity of the installation doubles.
With the participation of oxygen in the process of thermal-oxidative cracking, acetylene is produced on a large scale, which is widely used for cutting and welding metals and for the synthesis of organic substances:
6CH4 + 4O2 = C2H2 + 8H2 + 3CO + CO2 + 3H2O
Oxygen is used to obtain high temperatures. If you burn hydrogen in a stream of oxygen, then when 1 mole of water is formed, 286.3 kJ is released, and 2 moles - 572.6 kJ. This is colossal energy! The high temperatures reached in the flame of such burners (up to 3000°C) are used for cutting and welding metals.
Oxygen also serves in space. So, in the engine of the second stage of the American space rocket“Centaur” used liquid oxygen as an oxidizer. Oxygen is also widely used in rockets for various high-altitude research.
Liquid oxygen is included in explosives. For a long time, ammonites and other nitrogen-containing explosives were used for various blasting operations. Their use presented certain difficulties, such as the complexity and danger of transportation, and the need to build warehouses. Currently, liquid oxygen explosives can be manufactured at the point of use. Any porous flammable substance (sawdust, peat, hay, straw), when saturated with liquid oxygen, becomes explosive. Such substances are called oxyliquits and, if necessary, can replace dynamite in development ore deposits. When an explosion occurs, an oxyliquit cartridge is used - a simple long bag filled with flammable material, into which an electronic fuse is inserted. It is charged immediately before insertion into the hole by immersion in liquid oxygen. A hole is a round hole that is usually drilled in rocks ah and filled with explosives. If for some reason an explosion of the oxyliquit cartridge in the hole does not occur, the cartridge discharges itself as a result of the evaporation of liquid oxygen from it. The action of oxyliquits is based on the extremely rapid combustion of organic substances in pure oxygen. The short-term combustion process is accompanied by intense release large quantities heat and gases, which determines the use of oxyliquits as powerful explosives with a blasting (crushing) effect.
Oxygen is used in medicine and aviation. In medical practice for pulmonary and cardiac diseases, when breathing is difficult, patients are given oxygen from oxygen pillows and placed in special rooms in which the required oxygen concentration is maintained. One breath of oxygen by a person is equivalent to five breaths of air. Thus, when inhaled, this gas not only enters the patient’s body in sufficient quantities, but also saves energy for the breathing process itself. In addition, subcutaneous administration of oxygen has proven effective in the treatment of certain diseases, such as gangrene, thrombophlebitis, elephantiasis and tropical ulcers.
The phenomenon of “oxygen starvation” in the body can also occur from a lack of oxygen in environment. For example, at an altitude of 10,000 m, barometric air pressure drops to 217 mm Hg. Art. and the absolute oxygen content in the air decreases fourfold. This amount of gas is too small for normal breathing. Therefore on high altitudes pilots use oxygen cylinders.
VIII. Ozone layer above the Earth.
Ozone - « brother» oxygen. Its molecule is formed by three atoms of this chemical element: O3. Where there is an electric spark, a peculiar smell of freshness appears, because an electric discharge is a condition for the conversion of air oxygen into ozone:
oxygen ozone
We smell ozone in the air after a thunderstorm. Ozone is present in coniferous forests, especially pine forests. During decomposition tree resin some ozone is formed.
Ozone in the lower layer of air is scattered and its content is low. This gas is short-lived because it turns back into oxygen:
ozone oxygen
Even in small quantities, ozone acts as an oxidizing agent for many substances. Ozone disinfects tap water and purifies the air from pathogenic bacteria. Due to its activity, ozone can become hazardous to human and animal health if its permissible content in the air is exceeded. However, this does not happen in nature.
High above the Earth, in the stratosphere at an altitude of up to 30 km (above sea level), there is a constant thin layer of ozone that protects life on our planet from the harmful effects of short-wave ultraviolet radiation from the Sun. Ozone absorbs solar ultraviolet radiation, and only a part of it penetrates to the Earth, without causing much harm to its inhabitants. Short waves, which are harmful to all living things, are blocked, and long ultraviolet waves, which are harmless, are transmitted to the Earth.
There is more ozone in the stratosphere than in surface air, however, this does not mean that the layer is formed only by ozone. There is only 1 molecule of ozone in the ozone layer for every 100,000 molecules of other gases. But this ozone is enough to protect life on the planet from ultraviolet radiation.
Long-wave ultraviolet rays affect human skin, causing a tan. But skin cells can react painfully to short-wave radiation, and various types of tumors will appear. Ultraviolet radiation is also harmful to vision.
This is why it is so important that there is a protective ozone layer above the Earth!
In the stratosphere, ozone exists for quite a long time; it does not often encounter reducing substances there, but if they penetrate there, the ozone reacts with them and its amount decreases. This phenomenon of decreasing ozone concentration in some parts of the stratosphere is called the formation of “ozone holes.” IN Lately recorded a decrease in ozone concentration in the stratosphere by almost 40% over Antarctica. This flat continent is surrounded by an ocean, above South Pole a funnel is formed from the winds circulating around the continent and bringing substances with which ozone reacts. What substances are these?
These are artificially obtained and very valuable substances in practical terms - chlorofluorocarbons of various compositions, for example the following:
These substances are obtained by substitution reactions of halogens for hydrogen atoms in hydrocarbons. Chlorofluorocarbons are stable substances, do not dissolve in water, are non-toxic, do not burn, do not cause corrosion, and are excellent insulators. They are used to make insulation for building walls, disposable tableware for hot drinks. Liquid substances from this group (freons) are good solvents, effective refrigerants in refrigerators and air conditioners. They are used in aerosol cans as harmless solvents of special substances, in automatic fire extinguishing systems (CBrF3).
The production of these substances developed at an accelerated pace until it was discovered that when they enter the stratosphere, they destroy ozone (Now they are trying to replace freons with less volatile substances. For example, fluorochloromethane is used as refrigerants, and liquefied gaseous saturated hydrocarbons are used for aerosol cans ).
These substances reach the stratosphere unchanged. After all, they are chemically stable. And in the stratosphere, where there is a lot of ultraviolet radiation, their molecules are destroyed, and active halogen atoms, in particular chlorine, are split off:
Monatomic chlorine radical reacts with ozone:
03 + Cl = O2 + ClO
ozone chlorine oxygen oxide
(radical) chlorine(II)
Under the influence of ultraviolet rays, oxygen is formed from ozone, which at the time of release is also in an active monatomic state:
ozone oxygen atomic
oxygen
Chlorine oxide (II) reacts with atomic oxygen, and then a chlorine radical is again formed, which again destroys ozone; a chain reaction occurs, repeating itself many times:
ClO + O = Cl + O2
oxide atomic chlorine oxygen
chlorine (II) oxygen (radical)
О3 + С1= О2 + СlO
One chlorine atom participates in a series of such reactions and can destroy up to 100,000 ozone molecules. Chlorine can “get out of the game” when it encounters a methane molecule. Then it, adding to itself one hydrogen atom from methane, forms hydrogen chloride, which, when dissolved in water, forms hydrochloric acid. This is how chlorine destroyer returns to Earth in the form of acid rain:
CH4 + 2C1 - CH3C1 + HC1
methane chlorine chlorine hydrogen chloride
(radical) methane (in solution - hydrochloric acid)
Even if the production of chlorofluorocarbons is reduced everywhere, the process of destruction of the ozone layer over the entire planet will continue. The ozone-depleted air gradually dissipates, the gases in the atmosphere mix, and the chlorofluorocarbons contained in the air will continue their work destroying ozone for a very long time, at least 100 years.
In 1990, representatives of the government of 92 countries in London signed an agreement to completely stop the production of chlorofluorocarbons by the year 2000. Compliance with this agreement will be a condition for the gradual restoration of the natural content of ozone in the atmosphere, because the concentration of chlorine already in the atmosphere should decrease over time, but this time - century.
IX . Conclusion.
So, we received various information from the field of chemistry of group VI elements and, to a greater extent, about oxygen, we learned about where and how oxygen is used and obtained, we also learned about the effect of oxygen on our lives, National economy and culture.
If, after reading my essay, you have a desire to take a closer look at the vast area of that science from which information on the elements of group VI of D. I. Mendeleev’s periodic system was obtained, then I have completed my task.
Bibliography
1. Chemistry. For schoolchildren Art. classes and entering universities: Proc. Manual / N. E. Kuzmenkoy, V. V. Eremin, V. A. Popkov - 4th ed., stereotype. - M.: Bustard, 2001. - 544 p.: ill.
2. A book to read on inorganic chemistry. Book for students. At 2 p.m. Part 1 / comp. V. A. Kritsman - 3rd ed. - M.: Education, 1993. - 192 pp., 8 l ill.: ill. - ISBN 5-09-002972-5
3. Chemistry. Textbook for 9th grade. avg. school / F. G. Feldman, G. E. Rudzitis - M.: Education, 1990. - 176 pp.: ill. ISBN 5-09-002624-6
4. Chemistry: Textbook. for 8-9 grades. general education Institutions / R. G. Ivanova. - 3rd ed., M.: Education, 2001. - 270 pp.: ill. - ISBN 5-09-010278-3
5. Traveling through the sixth group. Elements of group VI of the periodic system of D. I. Mendeleev. A manual for students. / G. L. Nemchaninova - M., “Enlightenment”, 1976 - 128 pp.: ill.
Oxygen O has atomic number 8, located in the main subgroup (subgroup a) VI group, in the second period. In oxygen atoms, valence electrons are located on the 2nd energy level, which has only s- And p-orbitals. This excludes the possibility of transition of O atoms to an excited state, therefore oxygen in all compounds exhibits a constant valency equal to II. Having high electronegativity, oxygen atoms in compounds are always negatively charged (c.d. = -2 or -1). An exception is the fluorides OF 2 and O 2 F 2 .
For oxygen, the oxidation states are known -2, -1, +1, +2
General characteristics of the element
Oxygen is the most abundant element on Earth, accounting for slightly less than half, 49%, of total mass earth's crust. Natural oxygen consists of 3 stable isotopes 16 O, 17 O and 18 O (16 O predominates). Oxygen is part of the atmosphere (20.9% by volume, 23.2 by mass), in the composition of water and more than 1,400 minerals: silica, silicates and aluminosilicates, marbles, basalts, hematite and other minerals and rocks. Oxygen makes up 50-85% of the mass of tissues of plants and animals, as it is contained in proteins, fats and carbohydrates that make up living organisms. The role of oxygen for respiration and oxidation processes is well known.
Oxygen is relatively slightly soluble in water - 5 volumes in 100 volumes of water. However, if all the oxygen dissolved in water passed into the atmosphere, it would occupy a huge volume - 10 million km 3 (n.s.). This is equal to approximately 1% of all oxygen in the atmosphere. The formation of an oxygen atmosphere on earth is due to the processes of photosynthesis.
It was discovered by the Swede K. Scheele (1771 - 1772) and the Englishman J. Priestley (1774). The first used heating of nitrate, the second – mercury oxide (+2). The name was given by A. Lavoisier (“oxygenium” - “giving birth to acids”).
In its free form, it exists in two allotropic modifications - “ordinary” oxygen O 2 and ozone O 3 .
The structure of the ozone molecule
3O 2 = 2O 3 – 285 kJ
Ozone in the stratosphere forms a thin layer that absorbs most of the biologically harmful ultraviolet radiation.
During storage, ozone spontaneously turns into oxygen. Chemically, oxygen O2 is less active than ozone. The electronegativity of oxygen is 3.5.
Physical properties of oxygen
O 2 – colorless, odorless and tasteless gas, m.p. –218.7 °C, bp. –182.96 °C, paramagnetic.
Liquid O2 is blue, solid O2 is blue. O 2 is soluble in water (better than nitrogen and hydrogen).
Obtaining oxygen
1. Industrial method - distillation of liquid air and electrolysis of water:
2H 2 O → 2H 2 + O 2 
2. In the laboratory oxygen is obtained:
1. Electrolysis of alkaline aqueous solutions or aqueous solutions of oxygen-containing salts (Na 2 SO 4, etc.)
2. Thermal decomposition of potassium permanganate KMnO 4:
2KMnO 4 = K 2 MnO4 + MnO 2 + O 2,
Berthollet salt KClO 3:
2KClO 3 = 2KCl + 3O 2 (MnO 2 catalyst)
Manganese oxide (+4) MnO 2:
4MnO 2 = 2Mn 2 O 3 + O 2 (700 o C),
3MnO 2 = 2Mn 3 O 4 + O 2 (1000 o C),
Barium peroxide BaO 2:
2BaO2 = 2BaO + O2
3. Decomposition of hydrogen peroxide:
2H 2 O 2 = H 2 O + O 2 (MnO 2 catalyst)
4. Decomposition of nitrates:
2KNO 3 → 2KNO 2 + O 2
On spaceships and submarines, oxygen is obtained from a mixture of K 2 O 2 and K 2 O 4:
2K 2 O 4 + 2H 2 O = 4KOH +3O 2
4KOH + 2CO 2 = 2K 2 CO 3 + 2H 2 O
Total:
2K 2 O 4 + 2CO 2 = 2K 2 CO 3 + 3O 2
When K 2 O 2 is used, the overall reaction looks like this:
2K 2 O 2 + 2CO 2 = 2K 2 CO 3 + O 2
If you mix K 2 O 2 and K 2 O 4 in equal molar (i.e. equimolar) quantities, then one mole of O 2 will be released per 1 mole of absorbed CO 2.
Chemical properties of oxygen
Oxygen supports combustion. Combustion - b
a rapid process of oxidation of a substance, accompanied by the release of a large amount of heat and light.
To prove that the bottle contains oxygen and not some other gas, you need to lower a smoldering splinter into the bottle. In oxygen, a smoldering splinter flashes brightly. The combustion of various substances in air is a redox process in which oxygen is the oxidizing agent. Oxidizing agents are substances that “take” electrons from reducing substances. The good oxidizing properties of oxygen can be easily explained by the structure of its outer electron shell.
The valence shell of oxygen is located at the 2nd level - relatively close to the core. Therefore, the nucleus strongly attracts electrons to itself. On the valence shell of oxygen 2s 2 2p 4 there are 6 electrons. Consequently, the octet is missing two electrons, which oxygen tends to accept from the electron shells of other elements, reacting with them as an oxidizing agent.
Oxygen has the second (after fluorine) electronegativity on the Pauling scale. Therefore, in the vast majority of its compounds with other elements, oxygen has negative degree of oxidation. The only stronger oxidizing agent than oxygen is its neighbor in the period, fluorine. Therefore, compounds of oxygen with fluorine are the only ones where oxygen has a positive oxidation state.
So, oxygen is the second most powerful oxidizing agent among all the elements of the Periodic Table. Most of its most important chemical properties are associated with this.
All elements react with oxygen except Au, Pt, He, Ne and Ar; in all reactions (except for the interaction with fluorine), oxygen is an oxidizing agent.
Oxygen easily reacts with alkali and alkaline earth metals:
4Li + O 2 → 2Li 2 O,
2K + O 2 → K 2 O 2,
2Ca + O 2 → 2CaO,
2Na + O 2 → Na 2 O 2,
2K + 2O 2 → K 2 O 4
Fine powder of iron (the so-called pyrophoric iron) spontaneously ignites in air, forming Fe 2 O 3, and steel wire burns in oxygen if it is heated in advance:
3 Fe + 2O 2 → Fe 3 O 4
2Mg + O 2 → 2MgO
2Cu + O 2 → 2CuO
Oxygen reacts with non-metals (sulfur, graphite, hydrogen, phosphorus, etc.) when heated:
S + O 2 → SO 2,
C + O 2 → CO 2,
2H 2 + O 2 → H 2 O,
4P + 5O 2 → 2P 2 O 5,
Si + O 2 → SiO 2, etc.
Almost all reactions involving oxygen O2 are exothermic, with rare exceptions, for example:
N2+O2 → 2NO–Q
This reaction occurs at temperatures above 1200 o C or in an electrical discharge.
Oxygen is capable of oxidizing complex substances, for example:
2H 2 S + 3O 2 → 2SO 2 + 2H 2 O (excess oxygen),
2H 2 S + O 2 → 2S + 2H 2 O (lack of oxygen),
4NH 3 + 3O 2 → 2N 2 + 6H 2 O (without catalyst),
4NH 3 + 5O 2 → 4NO + 6H 2 O (in the presence of a Pt catalyst),
CH 4 (methane) + 2O 2 → CO 2 + 2H 2 O,
4FeS 2 (pyrite) + 11O 2 → 2Fe 2 O 3 + 8SO 2.
Compounds containing the dioxygenyl cation O 2 + are known, for example, O 2 + - (the successful synthesis of this compound prompted N. Bartlett to try to obtain compounds of inert gases).
Ozone
Ozone is chemically more active than oxygen O2. Thus, ozone oxidizes iodide - I ions - in a Kl solution:
O 3 + 2Kl + H 2 O = I 2 + O 2 + 2KOH
Ozone is highly toxic and poisonous properties stronger than, for example, hydrogen sulfide. However, in nature, ozone contained in high layers of the atmosphere acts as a protector of all life on Earth from the harmful ultraviolet radiation of the sun. The thin ozone layer absorbs this radiation and it does not reach the Earth's surface. There are significant fluctuations in the thickness and extent of this layer over time (the so-called ozone hole); the reasons for such fluctuations have not yet been clarified.
Application of Oxygen O 2: to intensify the processes of producing cast iron and steel, in the smelting of non-ferrous metals, as an oxidizer in various chemical industries, for life support on submarines, as an oxidizer for rocket fuel (liquid oxygen), in medicine, in welding and cutting metals.
Application of ozone O 3: for disinfection of drinking water, waste water, air, for bleaching fabrics.
Plan:
History of discovery
Origin of name
Being in nature
Receipt
Physical properties
Chemical properties
Application
10. Isotopes
Oxygen
Oxygen- element of the 16th group (according to the outdated classification - the main subgroup of group VI), the second period of the periodic table chemical elements D.I. Mendeleev, with atomic number 8. Denoted by the symbol O (lat. Oxygenium). Oxygen is a chemically active non-metal and is the lightest element from the group of chalcogens. Simple substance oxygen(CAS number: 7782-44-7) under normal conditions is a colorless, tasteless and odorless gas, the molecule of which consists of two oxygen atoms (formula O 2), and therefore it is also called dioxygen. Liquid oxygen has a light blue color, and solid crystals are light blue in color.
There are other allotropic forms of oxygen, for example, ozone (CAS number: 10028-15-6) - under normal conditions, a blue gas with a specific odor, the molecule of which consists of three oxygen atoms (formula O 3).
History of discovery
It is officially believed that oxygen was discovered by the English chemist Joseph Priestley on August 1, 1774 by decomposing mercuric oxide in a hermetically sealed vessel (Priestley directed sunlight at this compound using a powerful lens).
However, Priestley initially did not realize that he had discovered a new simple substance; he believed that he had isolated one of the constituent parts of air (and called this gas “dephlogisticated air”). Priestley reported his discovery to the outstanding French chemist Antoine Lavoisier. In 1775, A. Lavoisier established that oxygen is a component of air, acids and is found in many substances.
A few years earlier (in 1771), oxygen was obtained by the Swedish chemist Karl Scheele. He calcined saltpeter with sulfuric acid and then decomposed the resulting nitric oxide. Scheele called this gas “fire air” and described his discovery in a book published in 1777 (precisely because the book was published later than Priestley announced his discovery, the latter is considered the discoverer of oxygen). Scheele also reported his experience to Lavoisier.
An important step that contributed to the discovery of oxygen was the work of the French chemist Pierre Bayen, who published works on the oxidation of mercury and the subsequent decomposition of its oxide.
Finally, A. Lavoisier finally figured out the nature of the resulting gas, using information from Priestley and Scheele. His work was of enormous importance because thanks to it, the phlogiston theory, which was dominant at that time and hampered the development of chemistry, was overthrown. Lavoisier conducted experiments on the combustion of various substances and disproved the theory of phlogiston, publishing results on the weight of the burned elements. The weight of the ash exceeded the original weight of the element, which gave Lavoisier the right to claim that during combustion a chemical reaction (oxidation) of the substance occurs, and therefore the mass of the original substance increases, which refutes the theory of phlogiston.
Thus, the credit for the discovery of oxygen is actually shared between Priestley, Scheele and Lavoisier.
Origin of name
The word oxygen (also called “acid solution” at the beginning of the 19th century) owes its appearance in the Russian language to some extent to M.V. Lomonosov, who introduced the word “acid”, along with other neologisms; Thus, the word “oxygen”, in turn, was a tracing of the term “oxygen” (French oxygène), proposed by A. Lavoisier (from ancient Greek ὀξύς - “sour” and γεννάω - “giving birth”), which is translated as “generating acid”, which is associated with its original meaning - “acid”, which previously meant substances called oxides according to modern international nomenclature.
Being 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 and 23.12% by mass. More than 1,500 compounds in the earth's crust contain oxygen.
Oxygen is part of many organic substances and is present in all living cells. In terms of the number of atoms in living cells, it is about 25%, and in terms of mass fraction - about 65%.
Receipt
Currently, in industry, oxygen is obtained from the air. The main industrial method for producing oxygen is cryogenic rectification. Oxygen plants operating on the basis of membrane technology are also well known and successfully used in industry.
Laboratories use industrially produced oxygen, supplied in steel cylinders under a pressure of about 15 MPa.
Small amounts of oxygen can be obtained by heating potassium permanganate KMnO 4:
The reaction of catalytic decomposition of hydrogen peroxide H2O2 in the presence of manganese(IV) oxide is also used:
Oxygen can be obtained by the catalytic decomposition of potassium chlorate (Berthollet salt) KClO 3:
Laboratory methods for producing oxygen include the method of electrolysis of aqueous solutions of alkalis, as well as the decomposition of mercury(II) oxide (at t = 100 °C):
In submarines it is usually obtained by the reaction of sodium peroxide and carbon dioxide exhaled by humans:
Physical properties
In the world's oceans, the content of dissolved O2 is greater in cold water and less in warm water.
Under normal conditions, oxygen is a gas without color, taste or smell.
1 liter of it has a mass of 1.429 g. Slightly heavier than air. Slightly soluble in water (4.9 ml/100 g at 0 °C, 2.09 ml/100 g at 50 °C) and alcohol (2.78 ml/100 g at 25 °C). It dissolves well in molten silver (22 volumes of O 2 in 1 volume of Ag at 961 ° C). Interatomic distance - 0.12074 nm. Is paramagnetic.
When gaseous oxygen is heated, its reversible dissociation into atoms occurs: at 2000 °C - 0.03%, at 2600 °C - 1%, 4000 °C - 59%, 6000 °C - 99.5%.
Liquid oxygen (boiling point −182.98 °C) is a pale blue liquid.
O2 phase diagram
Solid oxygen (melting point −218.35°C) - blue crystals. There are 6 known crystalline phases, three of which exist at a pressure of 1 atm:
α-O 2 - exists at temperatures below 23.65 K; bright blue crystals belong to the monoclinic system, cell parameters a=5.403 Å, b=3.429 Å, c=5.086 Å; β=132.53°.
β-O 2 - exists in the temperature range from 23.65 to 43.65 K; pale blue crystals (with increasing pressure the color turns pink) have a rhombohedral lattice, cell parameters a=4.21 Å, α=46.25°.
γ-O 2 - exists at temperatures from 43.65 to 54.21 K; pale blue crystals have cubic symmetry, lattice parameter a=6.83 Å.
Three more phases form at high pressures:
δ-O 2 temperature range 20-240 K and pressure 6-8 GPa, orange crystals;
ε-O 4 pressure from 10 to 96 GPa, crystal color from dark red to black, monoclinic system;
ζ-O n pressure more than 96 GPa, a metallic state with a characteristic metallic luster, at low temperatures it transforms into a superconducting state.
Chemical properties
A strong oxidizing agent, it interacts with almost all elements, forming oxides. Oxidation state −2. As a rule, the oxidation reaction proceeds with the release of heat and accelerates with increasing temperature (see Combustion). Example of reactions occurring at room temperature:
Oxidizes compounds that contain elements with less than the maximum oxidation state:
Oxidizes most organic compounds:
Under certain conditions, it is possible to carry out mild oxidation of an organic compound:
Oxygen reacts directly (under normal conditions, with heating and/or in the presence of catalysts) with all simple substances except Au and inert gases (He, Ne, Ar, Kr, Xe, Rn); reactions with halogens occur under the influence of an electrical 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
Oxygen forms peroxides with the oxidation state of the oxygen atom formally equal to −1.
For example, peroxides are produced by combustion alkali metals in oxygen:
Some oxides absorb oxygen:
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, hydrogen peroxide is formed along with water:
In superoxides, oxygen formally has an oxidation state of −½, that is, one electron per two oxygen atoms (O − 2 ion). Obtained by reacting peroxides with oxygen at elevated pressure and temperature:
Potassium K, rubidium Rb and cesium Cs react with oxygen to form superoxides:
In the dioxygenyl ion O 2 +, oxygen formally has an oxidation state of +½. Obtained by the reaction:
Oxygen fluorides
Oxygen difluoride, OF 2 oxidation state of oxygen +2, is prepared by passing fluorine through an alkali solution:
Oxygen monofluoride (dioxydifluoride), O 2 F 2, is unstable, the oxidation state of oxygen is +1. 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.
Quantum mechanical calculations predict the stable existence of the trifluorohydroxonium ion OF 3 +. If this ion really exists, then the oxidation state of oxygen in it will be equal to +4.
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). As Pierre Curie and Marie Skłodowska-Curie established in 1899, under the influence of ionizing radiation O 2 turns into O 3 .
Application
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.
INmetallurgy
The converter method of steel production or matte processing involves the use of oxygen. In many metallurgical units, for more efficient combustion of fuel, an oxygen-air mixture is used instead of air in the burners.
Welding and cutting of metals
Oxygen in blue cylinders is widely used for flame cutting and welding of metals.
Rocket fuel
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).
INmedicine
Medical oxygen is stored in high-pressure metal gas cylinders (for compressed or liquefied gases) of blue color of various capacities from 1.2 to 10.0 liters under pressure up to 15 MPa (150 atm) and is used to enrich respiratory gas mixtures in anesthesia equipment, when breathing disorders, to relieve an attack of bronchial asthma, to eliminate hypoxia of any origin, for decompression sickness, to treat pathologies of the gastrointestinal tract in the form of oxygen cocktails. For individual use, special rubberized containers - oxygen cushions - are filled from cylinders with medical oxygen. Oxygen inhalers of various models and modifications are used to supply oxygen or an oxygen-air mixture simultaneously to one or two victims in the field or in a hospital setting. The advantage of an oxygen inhaler is the presence of a condenser-humidifier of the gas mixture, which uses the moisture of the exhaled air. To calculate the amount of oxygen remaining in the cylinder in liters, the pressure in the cylinder in atmospheres (according to the pressure gauge of the reducer) is usually multiplied by the cylinder capacity in liters. For example, in a cylinder with a capacity of 2 liters, the pressure gauge shows an oxygen pressure of 100 atm. The volume of oxygen in this case is 100 × 2 = 200 liters.
INFood Industry
IN Food Industry oxygen is registered as a food additive E948, as a propellant and packaging gas.
INchemical industry
In the chemical industry, oxygen is used as an oxidizing agent in numerous syntheses, for example, the oxidation of hydrocarbons into oxygen-containing compounds (alcohols, aldehydes, acids), ammonia into nitrogen oxides in the production of nitric acid. Due to the high temperatures developing during oxidation, the latter are often carried out in combustion mode.
INagriculture
In greenhouse farming, for making oxygen cocktails, for weight gain in animals, for enriching the aquatic environment with oxygen in fish farming.
Biological role of oxygen
Emergency oxygen supply in a bomb shelter
Most living beings (aerobes) 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 oxygen isotope 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
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 (a double magic nucleus with filled neutron and proton shells). And such nuclei, as follows from the theory of the structure of the atomic nucleus, are particularly stable.
Radioactive isotopes of oxygen with mass numbers from 12 O to 24 O are also known. All radioactive isotopes of oxygen have a short half-life, the longest-lived of them is 15 O with a half-life of ~120 s. The shortest-lived isotope 12 O has a half-life of 5.8·10−22 s.
The discovery of oxygen happened twice, in the second half of the 18th century, several years apart. In 1771, the Swede Karl Scheele obtained oxygen by heating saltpeter and sulfuric acid. The resulting gas was called "fire air". In 1774, the English chemist Joseph Priestley carried out the process of decomposing mercuric oxide in a completely closed vessel and discovered oxygen, but mistook it for an ingredient in air. Only after Priestley shared his discovery with the Frenchman Antoine Lavoisier did it become clear that a new element (calorizator) had been discovered. Priestley takes the lead in this discovery because Scheele published his scientific work describing the discovery only in 1777.
Oxygen is an element of group XVI of period II of the periodic table of chemical elements by D.I. Mendeleev, has atomic number 8 and atomic mass 15.9994. It is customary to denote oxygen by the symbol ABOUT(from Latin Oxygenium- generating acid). In Russian the name oxygen became a derivative of acids, a term that was introduced by M.V. Lomonosov.
Being in nature
Oxygen is the most common element found in the earth's crust and the World Ocean. Oxygen compounds (mainly silicates) make up at least 47% of the mass of the earth's crust, oxygen is produced during photosynthesis by forests and all green plants, most of it comes from phytoplankton of marine and fresh water. Oxygen is an essential component of any living cells and is also found in most substances of organic origin.
Physical and chemical properties
Oxygen is a light non-metal, belongs to the group of chalcogens, and has high chemical activity. Oxygen, as a simple substance, is a colorless, odorless and tasteless gas; it has a liquid state - light blue transparent liquid and a solid state - light blue crystals. Consists of two oxygen atoms (denoted by the formula O₂).
Oxygen is involved in redox reactions. 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, and gangrene. Artificial ozone enrichment is used to disinfect and deodorize air and purify drinking water.
Oxygen is the basis of the life activity of all living organisms on Earth, it is the main biogenic element. It is found in the molecules of all the most important substances that are responsible for the structure and functions of cells (lipids, proteins, carbohydrates, nucleic acids). Every living organism contains much more oxygen than any element (up to 70%). For example, the body of an average adult human weighing 70 kg contains 43 kg of oxygen.
Oxygen enters living organisms (plants, animals and humans) through the respiratory system and the intake of water. Remembering that in the human body the most important respiratory organ is the skin, it becomes clear how much oxygen a person can receive, especially in the summer on the shore of a reservoir. Determining a person’s need for oxygen is quite difficult, because it depends on many factors - age, gender, body weight and surface area, nutrition system, external environment, etc.
Use of oxygen in life
Oxygen is used almost everywhere - from metallurgy to the production of rocket fuel and explosives used for road work in the mountains; from medicine to the food industry.
In the food industry, oxygen is registered as a food additive, as a propellant and packaging gas.