Oxygen is the most abundant element on Earth. Sea water contains 85.82% oxygen, atmospheric air 23.15% by weight or 20.93% by volume, and 47.2% by weight in the earth's crust. This concentration of oxygen in the atmosphere is maintained constant through the process of photosynthesis. In this process, green plants use sunlight to convert carbon dioxide and water into carbohydrates and oxygen. The main mass of oxygen is in a bound state; the amount of molecular oxygen in the atmosphere is only 0.01% of the total oxygen content in the earth's crust. In the life of nature, oxygen is of exceptional importance. Oxygen and its compounds are indispensable for sustaining life. They play an important role in metabolic processes and respiration. Oxygen is a part of proteins, fats, carbohydrates from which organisms are "built"; the human body, for example, contains about 65% oxygen. Most organisms obtain the energy they need to perform their vital functions by oxidizing certain substances with the help of oxygen. The decrease in oxygen in the atmosphere as a result of the processes of respiration, decay and combustion is compensated for by oxygen released during photosynthesis. Deforestation, soil erosion, various mine workings on the surface reduce the total mass of photosynthesis and reduce the circulation over large areas.

Oxygen has not always been part of the earth's atmosphere. It appeared as a result of the vital activity of photosynthetic organisms. Under the influence of ultraviolet rays, it turns into ozone. As ozone accumulated, an ozone layer formed in the upper atmosphere. The ozone layer, like a screen, reliably protects the Earth's surface from ultraviolet radiation, which is fatal to living organisms.

Geochemical oxygen cycle connects the gas and liquid shells with the earth's crust. Its main points are: the release of free oxygen during photosynthesis, the oxidation of chemical elements, the entry of extremely oxidized compounds into the deep zones of the earth's crust and their partial reduction, including due to carbon compounds, the removal of carbon monoxide and water to the surface of the earth's crust and their involvement in the reaction photosynthesis.

In addition to the oxygen cycle described above in an unbound form, this element also performs the most important cycle, entering the composition of water (Fig. 3). During the cycle, water evaporates from the surface of the ocean, water vapor moves along with air currents, condenses, and water returns in the form of precipitation to the surface of land and sea. There is a large water cycle, in which water that has fallen in the form of precipitation on land returns to the seas through surface and underground runoff; and the small water cycle, in which precipitation falls on the surface of the ocean.

The oxygen cycle is accompanied by its arrival and consumption.

The arrival of oxygen includes: 1) release during photosynthesis; 2) formation in the ozone layer under the influence of UV radiation (in a small amount); 3) dissociation of water molecules in the upper layers of the atmosphere under the influence of UV radiation; 4) the formation of ozone - O3.

Oxygen consumption includes: 1) consumption by animals during respiration; 2) oxidative processes in the earth's crust; 3) oxidation of carbon monoxide (CO) released during volcanic eruptions.

Oxygen is the main biogenic element that is part of the molecules of all the most important substances that provide the structure and functions of cells - proteins, nucleic acids, carbohydrates, lipids, as well as many low molecular weight compounds. In every plant or animal, there is much more oxygen than any other element (about 70% on average). 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 the community is determined as an important energy characteristic. organisms to its total biomass.

Small amounts of oxygen are used in medicine: oxygen (from the so-called oxygen pillows) is given some time to breathe for patients who have difficulty breathing. However, it must be borne in mind that prolonged inhalation of air enriched with oxygen is dangerous to human health. High oxygen concentrations cause the formation of free radicals in tissues that disrupt the structure and functions of biopolymers. Ionizing radiation has a similar effect on the body. 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: by increasing the oxygen content in the tumor and lowering its content in the surrounding tissues, they increase the radiation damage to tumor cells and reduce damage to healthy ones. In some diseases, saturation of the body with oxygen under high pressure is used - hyperbaric oxygenation.

The main (in fact, the only) function of oxygen is its participation as an oxidizing agent in redox reactions in the body. Due to the presence of oxygen, the organisms of all animals are able to utilize (actually “burn”) various substances (carbohydrates, fats, proteins) with the extraction of a certain “combustion” energy for their own needs. At rest, the body of an adult consumes 1.8-2.4 g of oxygen per minute.

Ozone(from other Greek ὄζω - I smell) - an allotropic modification of oxygen consisting of triatomic O 3 molecules. Under normal conditions - blue gas. When liquefied, it turns into an indigo liquid. In solid form, it is dark blue, almost black crystals.

Question

Sulfur- an element of the 16th group (according to the outdated classification - the main subgroup of group VI), the third period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 16. Shows non-metallic properties. Indicated by the symbol S(lat. sulfur). In hydrogen and oxygen compounds, it is part of various ions, forms many acids and salts. Many sulfur-containing salts are sparingly soluble in water.

In air, sulfur burns, forming sulfur dioxide - a colorless gas with a pungent odor:

Using spectral analysis, it was found that in fact the process of oxidation of sulfur to dioxide is a chain reaction and occurs with the formation of a number of intermediate products: sulfur monoxide S 2 O 2 , molecular sulfur S 2 , free sulfur atoms S and free radicals of sulfur monoxide SO .

The reducing properties of sulfur are manifested in the reactions of sulfur with other non-metals, however, at room temperature, sulfur reacts only with fluorine:

The sulfur melt reacts with chlorine, while the formation of two lower chlorides (sulfur dichloride and dithiodichloride) is possible:

With an excess of sulfur, various polyser dichlorides of the S n Cl 2 type are also formed.

When heated, sulfur also reacts with phosphorus, forming a mixture of phosphorus sulfides, among which is the highest sulfide P 2 S 5:

In addition, when heated, sulfur reacts with hydrogen, carbon, silicon:

(hydrogen sulfide)

(carbon disulfide)

When heated, sulfur interacts with many metals, often very violently. Sometimes a mixture of metal with sulfur ignites when ignited. In this interaction, sulfides are formed:

Solutions of alkali metal sulfides react with sulfur to form polysulfides:

Of the complex substances, first of all, the reaction of sulfur with molten alkali should be noted, in which sulfur disproportionates similarly to chlorine:

The resulting alloy is called sulfur liver.

With concentrated oxidizing acids (HNO 3, H 2 SO 4), sulfur reacts only with prolonged heating:

With an increase in temperature in sulfur vapor, changes occur in the quantitative molecular composition. The number of atoms in a molecule decreases:

At 800-1400 °C, vapors consist mainly of diatomic sulfur:

And at 1700 ° C, sulfur becomes atomic:

Biological role: Sulfur is constantly present in all living organisms, being an important biogenic element. Its content in plants is 0.3-1.2%, in animals 0.5-2% (marine organisms contain more sulfur than terrestrial ones). The biological significance of sulfur is determined primarily by the fact that it is part of the amino acids methionine and cysteine ​​and, consequently, in the composition of peptides and proteins. Disulfide bonds –S–S– in polypeptide chains are involved in the formation of the spatial structure of proteins, and sulfhydryl groups (–SH) play an important role in the active centers of enzymes. In addition, sulfur is included in the molecules of hormones, important substances. A lot of sulfur is found in the keratin of hair, bones, and nervous tissue. Inorganic sulfur compounds are essential for the mineral nutrition of plants. They serve as substrates for oxidative reactions carried out by naturally occurring sulfur bacteria.

The body of an average person (body weight 70 kg) contains about 1402 g of sulfur. The daily requirement of an adult for sulfur is about 4.

However, in terms of its negative impact on the environment and humans, sulfur (more precisely, its compounds) is one of the first places. The main source of sulfur pollution is the combustion of coal and other fuels containing sulfur. At the same time, about 96% of the sulfur contained in the fuel enters the atmosphere in the form of sulfur dioxide SO 2 .

In the atmosphere, sulfur dioxide is gradually oxidized to sulfur oxide (VI). Both oxides - both sulfur oxide (IV) and sulfur oxide (VI) - interact with water vapor to form an acid solution. These solutions then fall out as acid rain. Once in the soil, acidic waters inhibit the development of soil fauna and plants. As a result, unfavorable conditions are created for the development of vegetation, especially in the northern regions, where chemical pollution is added to the harsh climate. As a result, forests are dying, the grass cover is being disturbed, and the condition of water bodies is deteriorating. Acid rain destroys monuments made of marble and other materials, moreover, they cause the destruction of even stone buildings and metal products. Therefore, it is necessary to take various measures to prevent the ingress of sulfur compounds from the fuel into the atmosphere. To do this, oil and oil products are purified from sulfur compounds, and the gases formed during the combustion of fuel are purified.

By itself, sulfur in the form of dust irritates the mucous membranes, respiratory organs and can cause serious illness. MPC of sulfur in the air is 0.07 mg/m 3 .

Many sulfur compounds are toxic. Especially noteworthy is hydrogen sulfide, the inhalation of which quickly causes a dulling of the reaction to its unpleasant odor and can lead to severe poisoning, even with a fatal outcome. Maximum allowable concentration of hydrogen sulfide in the air of working premises is 10 mg/m 3 , in the atmospheric air 0.008 mg/m 3 .

Sulfur(II) oxide (sulfur monoxide, sulfur monoxide) is a binary inorganic compound. Under normal conditions, it is a colorless gas with a pungent, unpleasant odor. Reacts with water. It is extremely rare in the Earth's atmosphere. Thermodynamically unstable, exists as a dimer S 2 O 2 . It reacts very actively with oxygen, forming sulfur dioxide.

Receipt

The main method of obtaining is the combustion of sulfur:

Obtained by the decomposition of sulfur dioxide:

Chemical properties

It dissolves in water to form thiosulfuric acid:

Application

Due to its rarity and instability, sulfur monoxide has not been used.

Toxicity

Due to the instability of sulfur monoxide, it is difficult to determine its toxicity, but in concentrated form, sulfur monoxide turns into peroxide, which is toxic and corrosive.

Sulfur(IV) oxide (sulfur dioxide, sulfur dioxide, sulphur dioxide, sulfur dioxide) - a compound of sulfur with oxygen of the composition SO 2. Under normal conditions, it is a colorless gas with a characteristic pungent odor (the smell of a lighted match). It liquefies under pressure at room temperature. Dissolves in water to form unstable sulfurous acid; solubility 11.5 g/100 g water at 20 °C, decreases with increasing temperature. It also dissolves in ethanol and sulfuric acid. One of the main components of volcanic gases.

Receipt

The industrial method of obtaining is the burning of sulfur or the roasting of sulfides, mainly pyrite:

In laboratory conditions and in nature, SO 2 is obtained by the action of strong acids on sulfites and hydrosulfites. The resulting sulfurous acid H 2 SO 3 immediately decomposes into SO 2 and H 2 O:

Also, sulfur dioxide can be obtained by the action of concentrated sulfuric acid on low-active metals when heated:

Chemical properties

SO2 absorption spectrum in the ultraviolet range.

Refers to acidic oxides. It dissolves in water to form sulfurous acid (under normal conditions, the reaction is reversible):

Forms sulfites with alkalis:

The chemical activity of SO 2 is very high. The most pronounced reducing properties of SO 2, the degree of oxidation of sulfur in such reactions increases:

The penultimate reaction is a qualitative reaction to the sulfite ion SO 3 2− and to SO 2 (discoloration of the violet solution).

In the presence of strong reducing agents, SO 2 is able to exhibit oxidizing properties. For example, to extract sulfur from waste gases of the metallurgical industry, SO 2 reduction with carbon monoxide (II) is used:

Or to get hypophosphorous acid:

Application

Most of the sulfur(IV) oxide is used to produce sulfurous acid. It is also used in winemaking as a preservative (food additive E220). Since this gas kills microorganisms, vegetable stores and warehouses are fumigated with it. Sulfur(IV) oxide is used to bleach straw, silk and wool, materials that cannot be bleached with chlorine. It is also used as a solvent in laboratories. With this application, one should be aware of the possible content of impurities in SO 2 in the form of SO 3, H 2 O, and, as a result of the presence of water, H 2 SO 4 and H 2 SO 3. They are removed by passing concentrated H 2 SO 4 through a solvent; this is best done under vacuum or in other closed apparatus. Sulfur oxide (IV) is also used to obtain various salts of sulfurous acid.

Toxic action

SO 2 is very toxic. Symptoms of sulfur dioxide poisoning are a runny nose, cough, hoarseness, severe sore throat and a peculiar aftertaste. When inhaling sulfur dioxide at a higher concentration - suffocation, speech disorder, difficulty swallowing, vomiting, acute pulmonary edema is possible.

With short-term inhalation, it has a strong irritant effect, causes coughing and sore throat.

MPC (maximum permissible concentration):

· in the atmospheric air maximum one-time - 0.5 mg/m³, average daily - 0.05 mg/m³;

indoors (working area) - 10 mg/m³

Interestingly, the sensitivity to SO 2 is very different in individuals, animals and plants. Thus, among plants, birch and oak are the most resistant to sulfur dioxide, the least resistant are rose, pine and spruce.

Sulfur oxide (VI) (sulfuric anhydride, sulfur trioxide, sulfuric gas) SO 3 - higher sulfur oxide, type of chemical bond: covalent polar chemical bond. Under normal conditions, a highly volatile, colorless liquid with a suffocating odor. At temperatures below 16.9 ° C, it solidifies with the formation of a mixture of various crystalline modifications of solid SO 3.

Receipt

Obtained by oxidizing sulfur oxide (IV) with atmospheric oxygen when heated, in the presence of a catalyst (V 2 O 5 , Pt, NaVO 3 or iron oxide (III) Fe 2 O 3):

Can be obtained by thermal decomposition of sulfates:

or the interaction of SO 2 with ozone:

For the oxidation of SO 2, NO 2 is also used:

This reaction underlies the historically first, nitrous method for the production of sulfuric acid.

Chemical properties

1. Acid-base: SO 3 is a typical acid oxide, sulfuric anhydride. Its chemical activity is quite high. When reacted with water, it forms sulfuric acid:

However, in this reaction, sulfuric acid is formed in the form of an aerosol, and therefore, in industry, sulfur oxide (VI) is dissolved in sulfuric acid to form a moleum, which is then dissolved in water to form sulfuric acid of the desired concentration.

Interacts with bases:

and oxides:

SO 3 dissolves in 100% sulfuric acid, forming oleum.

"2" . Redox: SO 3 is characterized by strong oxidizing properties, it is usually reduced to sulfur dioxide:

3. When interacting with hydrogen chloride, chlorosulfonic acid is formed:

It also reacts with sulfur dichloride and chlorine to form thionyl chloride:

Application

Sulfuric anhydride is used mainly in the production of sulfuric acid.

Sulfuric anhydride is also released into the air when sulfur pellets are burned, which are used in the disinfection of premises. Upon contact with wet surfaces, sulfuric anhydride turns into sulfuric acid, which already destroys fungus and other harmful organisms.

SULFURIC ACID

H2S03H2S03, (S + 4S + 4) - sulfurous acid - an acid of medium strength, corresponds to the oxidation state of sulfur +4, a fragile compound, exists only in aqueous solutions (not isolated in a free state), oxidized by atmospheric oxygen, turning into sulfuric acid H2S04H2S04, good restorer. As a dibasic acid, it forms two series of salts: hydrosulfites (NaHSO3NaHSO3, in excess of alkali):

H2SO3+NaOH=NaHSO3+H2OH2SO3+NaOH=NaHSO3+H2O

and sulfites (Na2SO3Na2SO3 - with a lack of alkali):

H2SO3+2NaOH=Na2SO3+2H2OH2SO3+2NaOH=Na2SO3+2H2O

Like sulfur dioxide, sulfurous acid and its salts are strong reducing agents:

H2SO3+Br2+2O=H2SO4+2HBrH2SO3+Br2+2O=H2SO4+2HBr

When interacting with even stronger reducing agents, it can play the role of an oxidizing agent:

H2SO3+2H2S=3S+3H2OH2SO3+2H2S=3S+3H2O

A qualitative reaction to sulfite ions is the evolution of a gas with a pungent odor (SO2SO2) when interacting with acids:

SO2−3+2H+=SO2+H2OSO32−+2H+=SO2+H2O

In addition, a solution of sulfite ions discolors a solution of potassium permanganate:

5SO2−3+6H++2MnO−4=5SO2−4+2Mn2++3H2O5SO32−+6H++2MnO4−=5SO42−+2Mn2++3H2O

However, this reaction is rarely used for the qualitative detection of sulfite ions.

Sulfuric acid and its salts are used as reducing agents for bleaching wool, silk and other materials that cannot withstand bleaching with strong oxidizing agents (chlorine). Sulfuric acid is used in the preservation of fruits and vegetables. Calcium hydrosulfite (sulfite liquor, Ca (HSO3) 2Ca (HSO3) 2) is used to process wood into the so-called sulfite cellulose (calcium hydrosulfite solution dissolves lignin, a substance that binds cellulose fibers, as a result of which the fibers are separated from each other; treated in this way wood is used to make paper).

SULFURIC ACID

H2S04H2S04 (S + 6S + 6) - sulfuric acid - a colorless, odorless oily liquid, non-volatile, crystallizing at 10.3010.30С, heavy, actively absorbs water vapor, a strong oxidizing agent, dibasic acid, forms two series of salts: sulfates and hydrosulfates, of which only BaSO4BaSO4, PbSO4PbSO4, and SrSO4SrSO4 are practically insoluble.

The specific properties of sulfuric acid are discussed in detail in the topic "Interaction of sulfuric acid with metals and non-metals".

Due to the ability to replace hydrogen and sulfur atoms and the formation of oxygen "bridges", sulfur is able to form a number of oxygen-containing acids:

H2S207H2S207 (S + 6S + 6) - pyrosulfuric, or disulfuric acid.

When sulfuric anhydride S03S03 is dissolved in sulfuric acid, oleum is obtained, consisting mainly of pyrosulfuric acid. When the oleum is cooled, the acid separates out as colorless crystals. Pyrosulfuric acid forms salts - disulfates or pyrosulfates (Na2S2O7Na2S2O7), which, when heated above the melting point, decompose, turning into sulfates.

H2S02H2S02, ($S^(+2)) - (structural formula H-O-S-O-H) sulfoxylic acid; not isolated in the free state.

H2S208H2S208, (S + 6S + 6) - peroxysulfuric, or persulfuric, acid, has strong oxidizing properties, forms persulfate salts (see the structure in Figure 1).

H2S202H2S202 (S+4S+4) - thiosulfuric acid, is formed as an intermediate product in various reactions. Thiosulfuric acid can be considered as sulfurous acid in which the oxygen atom is replaced by sulfur. Neither the acid itself nor its salts have been isolated in the free state.

H2S203H2S203 (S + 4S + 4 - thiosulfuric acid - unstable, decomposes already at room temperature, forms salts - thiosulfates, which are much more stable than acid and are often used in industry as reducing agents

H2S204H2S204 (S+4S+4-dithionic or sulphurous acid, exists only in the form of salts.

There is a group of polythionic acids that correspond to the general formula H2Sx06H2Sx06 (S + 4S + 4, where x takes values ​​from 2 to 6. Polythionic acids are unstable and are known only in aqueous solutions. Their salts - polythionates - are more stable, some of them are obtained in the form of crystals .

Hydrogen sulfide (hydrogen sulfide, hydrogen sulfide, dihydrosulfide)- a colorless gas with a sweetish taste, having the smell of rotten chicken eggs. Binary chemical compound of hydrogen and sulfur. Chemical formula - H 2 S. Poorly soluble in water, well - in ethanol. Poisonous. At high concentrations, it interacts with many metals. Flammable. The concentration limits of ignition in a mixture with air are 4.5-45% hydrogen sulfide. It is used in the chemical industry for the synthesis of certain compounds, the production of elemental sulfur, sulfuric acid, and sulfides. Hydrogen sulfide is also used medicinally, such as in hydrogen sulfide baths.

The intrinsic ionization of liquid hydrogen sulfide is negligible.

Hydrogen sulfide is slightly soluble in water, an aqueous solution of H 2 S is a very weak acid:

K a \u003d 6.9 10 -7 mol / l; p K a = 6.89.l

Reacts with alkalis:

(medium salt, with excess NaOH)

(acid salt, at a ratio of 1:1)

Hydrogen sulfide is a strong reducing agent. Redox potentials:

In the air it burns with a blue flame:

with a lack of oxygen:

(The industrial method for producing sulfur is based on this reaction).

Hydrogen sulfide also reacts with many other oxidizing agents; when it is oxidized in solutions, free sulfur or an SO 4 2− ion is formed, for example:

A qualitative reaction to hydrogen sulfide, hydrosulfide acid and its salts is their interaction with lead salts, in which a black precipitate of lead sulfide is formed, for example:

When hydrogen sulfide is passed through human blood, it turns black, because hemoglobin is destroyed, and iron, which is part of its composition and gives the blood a red color, reacts with hydrogen sulfide and forms black iron sulfide.

Question

Halogens(from the Greek ἁλός - "salt" and γένος - "birth, origin"; sometimes an outdated name is used halides) - chemical elements of the 17th group of the periodic table of chemical elements of D. I. Mendeleev (according to the outdated classification - elements of the main subgroup of group VII).

They react with almost all simple substances, except for some non-metals. All halogens are energetic oxidizing agents, therefore they occur in nature only in the form of compounds. With an increase in the serial number, the chemical activity of halogens decreases, the chemical activity of halide ions F - , Cl - , Br - , I - , At - decreases.

Halogens include fluorine F, chlorine Cl, bromine Br, iodine I, astatine At, and (formally) the artificial element ununseptium Uus.

All halogens show high oxidative activity, which decreases when moving from fluorine to astatine. Fluorine is the most active of the halogens, it reacts with all metals without exception, many of them spontaneously ignite in an atmosphere of fluorine, releasing a large amount of heat, for example:

2Al + 3F 2 = 2AlF 3 + 2989 kJ,

2Fe + 3F 2 = 2FeF 3 + 1974 kJ.

Without heating, fluorine also reacts with many non-metals (H 2 , S, C, Si, P); all reactions are strongly exothermic, for example:

H 2 + F 2 = 2HF + 547 kJ,

Si + 2F 2 = SiF 4 (g) + 1615 kJ.

When heated, fluorine oxidizes all other halogens according to the scheme

Hal 2 + F 2 = 2HalF

where Hal = Cl, Br, I, At, and in HalF compounds, the oxidation states of chlorine, bromine, iodine, and astatine are +1.

Finally, when irradiated, fluorine reacts even with heavy inert (noble) gases:

Xe + F 2 = XeF 2 + 152 kJ.

The interaction of fluorine with complex substances also proceeds very vigorously. So, it oxidizes water, while the reaction is explosive:

3F 2 + ZN 2 O \u003d OF 2 + 4HF + H 2 O 2.

Free chlorine is also very reactive, although its activity is less than that of fluorine. It directly reacts with all simple substances except oxygen, nitrogen and noble gases. For comparison, we present the equations for the reactions of chlorine with the same simple substances as for fluorine:

2Al + 3Cl 2 = 2AlCl 3 (cr) + 1405 kJ,

2Fe + ZCl 2 = 2FeCl 3 (cr) + 804 kJ,

Si + 2Cl 2 = SiCl 4 (L) + 662 kJ,

H 2 + Cl 2 \u003d 2HCl (g) + 185 kJ.

Of particular interest is the reaction with hydrogen. So, at room temperature, without lighting, chlorine practically does not react with hydrogen, while when heated or illuminated (for example, in direct sunlight), this reaction proceeds with an explosion according to the following chain mechanism:

Cl2+ hν → 2Cl,

Cl + H 2 → HCl + H,

H + Cl 2 → HCl + Cl,

Cl + H 2 → HCl + H, etc.

The excitation of this reaction occurs under the action of photons ( hν), which cause the dissociation of Cl 2 molecules into atoms - in this case, a chain of successive reactions occurs, in each of which a particle appears, initiating the beginning of the next stage.

The reaction between H 2 and Cl 2 served as one of the first objects of study of chain photochemical reactions. The greatest contribution to the development of ideas about chain reactions was made by the Russian scientist, Nobel Prize winner (1956) N. N. Semyonov.

Chlorine reacts with many complex substances, such as substitution and addition with hydrocarbons:

CH 3 -CH 3 + Cl 2 → CH 3 -CH 2 Cl + HCl,

CH 2 \u003d CH 2 + Cl 2 → CH 2 Cl - CH 2 Cl.

Chlorine is capable of displacing bromine or iodine from their compounds with hydrogen or metals when heated:

Cl 2 + 2HBr \u003d 2HCl + Br 2,

Cl 2 + 2HI \u003d 2HCl + I 2,

Cl 2 + 2KBr \u003d 2KCl + Br 2,

and also reacts reversibly with water:

Cl 2 + H 2 O \u003d HCl + HClO - 25 kJ.

Chlorine, dissolving in water and partially reacting with it, as shown above, forms an equilibrium mixture of substances called chlorine water.

Note also that chlorine on the left side of the last equation has an oxidation state of 0. As a result of the reaction, the oxidation state of some chlorine atoms became −1 (in HCl), others +1 (in hypochlorous acid HOCl). Such a reaction is an example of a self-oxidation-self-healing, or disproportionation, reaction.

Chlorine can react (disproportionate) with alkalis in the same way:

Cl 2 + 2NaOH \u003d NaCl + NaClO + H 2 O (in the cold),

3Cl 2 + 6KOH \u003d 5KCl + KClO 3 + 3H 2 O (when heated).

The chemical activity of bromine is less than that of fluorine and chlorine, but still quite high due to the fact that bromine is usually used in a liquid state and therefore its initial concentrations, other things being equal, are greater than that of chlorine.

For example, we give the reactions of interaction of bromine with silicon and hydrogen:

Si + 2Br 2 \u003d SiBr 4 (l) + 433 kJ,

H 2 + Br 2 = 2HBr (g) + 73 kJ.

Being a "softer" reagent, bromine is widely used in organic chemistry.

Note that bromine, like chlorine, dissolves in water, and, partially reacting with it, forms the so-called "bromine water".

The solubility of iodine in water is 0.3395 grams per liter at 25 degrees Celsius, which is less than that of bromine. An aqueous solution of iodine is called "iodine water". Iodine is able to dissolve in iodide solutions with the formation of complex anions:

I 2 + I − → I − 3 .

The resulting solution is called Lugol's solution.

Iodine differs significantly in chemical activity from other halogens. It does not react with most non-metals, and reacts slowly with metals only when heated. The interaction of iodine with hydrogen occurs only with strong heating, the reaction is endothermic and highly reversible:

H 2 + I 2 \u003d 2HI - 53 kJ.

Thus, the chemical activity of halogens consistently decreases from fluorine to astatine. Each halogen in the F - At series can displace the next one from its compounds with hydrogen or metals, that is, each halogen in the form of a simple substance is able to oxidize the halide ion of any of the subsequent halogens.

Astatine is even less reactive than iodine. But it also reacts with metals (for example, with lithium):

2Li + At 2 = 2LiAt - lithium astatide.

During dissociation, not only anions are formed, but also cations At +: HAt dissociates into:

2HAt=H + +At - +H - +At + .

(hydrogen halides) - colorless gases with a pungent odor, fuming in moist air. They are highly soluble in water, their aqueous solutions are acids, bearing the common name - hydrohalic acids. Salts of hydrohalic acids (fluorides, chlorides, bromides and iodides) can be obtained by direct combination of metals with halogens. In composition, they are of the same type and have similar properties. So NaF, NaCl, NaBr, NaJ are white crystalline substances, readily soluble in water. Along with similarities, halogens also have certain differences in both physical and chemical properties. However, these properties change naturally with an increase in the atomic weight of the halogen.

- Hydrogen halides HF, HC1, HBr and HI are colorless gases that dissolve well in water. Of these, HF is a weak acid and the rest of the hydrogen halides are strong acids in aqueous solution.

So hydrogen compounds of halogens more stable than oxygen.

So hydrogen compounds of halogens more stable than oxygen. The redox properties and differences in the chemical behavior of halogens are easy to understand by comparing these properties as a function of the change in nuclear charge when going from fluorine to iodine. In the series F, C1, Br, I, iodine has the largest atomic radius (and, consequently, the lowest electron affinity), so it is characterized by less pronounced oxidizing properties than bromine, chlorine and fluorine.

For allowed to use the following names: hydrogen fluoride, hydrogen chloride, hydrogen bromide and hydrogen iodide. Type names hydrochloric acid refer to aqueous solutions of hydrogen halides.

Education hydrogen compounds of halogens goes with a greater release of heat than oxygen, so hydrogen compounds are more stable than oxygen. Of the oxygen compounds, the salts of oxygen acids are the most stable and the oxides are the least stable.

Oxygen compounds of halogens All oxygen compounds of halogens are obtained indirectly. Salts are the most stable, oxides and acids are the least stable. Halogens are characterized by the formation of a large number of oxides corresponding to different oxidation states. Most of all, BrO-2 and IO-2 ions are very unstable. stable oxides are formed by chlorine Cl, least of all - iodine I. Of the compounds of oxygen with fluorine, there is oxygen fluoride F-12O + 2: The bond between the atoms of fluorine and oxygen is covalent, very close to non-polar. It is a colorless gas with a pungent ozone odor, poorly soluble in water, boiling point = -145°C. It was opened in 1929. obtained by the interaction of fluorine with a 2% solution of sodium hydroxide: 2F2 + 2NaOH = 2NaF + H2O + F2O I Consider the most important of the oxygen compounds of the remaining halogens. All oxides are unstable, decompose with a large release of heat. Chlorine oxide (I) Сl2О is a brown-yellow gas with an unpleasant odor. It is characterized by a low boiling point, the relative density in air is 3. The bond in the oxide molecule is low-polar covalent. It has the following chemical properties: 1. When heated, it easily decomposes (with an explosion) into chlorine and oxygen: 2C12O=t2Cl2+O2 2. Being an acidic oxide, it hydrates to form hypochlorous acid: Cl2O+H2O=2HClO 3. It interacts with alkalis and basic oxides: Cl2O+2NaOH= 2NaClO + H2O Cl2O + K2O \u003d 2KClO Chlorine (I) oxide corresponds to hypochlorous acid. Hypochlorous acid HClO and its bromine and iodine counterparts are very weak acids, and their strength decreases when going from HClO to HIO. This is due to the fact that chlorine has a higher electronegativity and attracts the electron pair that binds it to oxygen more strongly than its counterparts. This, in turn, leads to a shift of the electron pair that binds hydrogen with oxygen to oxygen and an increase in the ability of hydrogen to split off. Hypochlorous acid is a yellow-green solution with a characteristic odor. She and her analogues have all the properties of weak volatile acids, are oxidizing acids. Moreover, the oxidative activity in the series HClO, HBrO, HIO decreases. 1. Hypochlorous acid decomposes in the light: HCl + 1O-2 \u003d hv HCl-1 + O0 2. Decomposes under the action of water-removing agents: 2HCl + 1O \u003d Cl + 12O + H2O 3. When hypochlorous acid is heated, hydrochloric and chloric acids are formed: 3HCl +1О=2НCl-1+НCl+5O3 oxidizing agent Сl++2е- Сl-reducing agent Сl+-4е- Сl+5 Salts of oxygen acids of chlorine are of the greatest importance. All of them can be obtained based on the reaction of the interaction of chlorine with water. HCl + HClO "Cl2 + H2O The equilibrium of this reaction can be easily shifted towards the reaction products by adding alkali to the solution, which reacts with two formed acids: HCl + HClO + 2KOH \u003d KCl + KClO + 2H2O I Summing up these two equations, we get: Сl2 + 2KOH \u003d KCl + KClO + H2O Cl2 + 2OH- \u003d Cl- + ClO- + H2O Salts of hypochlorous acid are called hypochlorites. An aqueous solution containing hypochlorite and potassium chloride is called javelin water. She, like chlorine (a solution of chlorine in water) water, is used for bleaching cotton fabric and paper. The mechanism of the oxidizing and disinfecting action of hypochlorous acid and its salts is explained by the presence of chlorine with an oxidation state of +1, which exhibits active oxidizing properties in these processes. Cl++1e-Cl° Cl++2e-Cl- Hypochlorites are very strong oxidizing agents. When chlorine is passed into an alkali solution heated to 100 ° C, the process proceeds with the formation of chlorates (salts of chloric acid HclO3) and chlorides: heating to 400 ° C in the absence of catalysts, perchlorates are formed from chlorates (salts of perchloric acid HclO4): with lye. In this case, hypochlorites are formed at room temperature, and chlorates are formed at 100°C. These are redox reactions. Chloric acid HClO2 - medium strength. It is unstable in aqueous solutions, and its analogues of bromine and iodine are even less durable. The strength of chlorine oxygen acids increases with an increase in the degree of oxidation of the central atom: HCl + 1O - weak; HCl + 3O2 - somewhat stronger; HCl + 5O3 is very strong and HClO + 74 is the strongest of all known acids. If chlorine interacts with calcium hydroxide, which is taken in the form of powder - fluff, then chloride, or whitewash, lime is formed - a loose white powder with the smell of chlorine. It consists mainly of calcium hypochlorite Ca(ClO)2, basic calcium salts and calcium chloride. Approximate equation: 2Cl2 + 2Ca (OH) 2 \u003d Ca (ClO) 2 + CaCl 2 + 2H2O Often, Ca (ClO) 2 is added to it to improve the quality of bleach. Chlorine is a strong oxidizing agent. She is very reactive. It is used for bleaching cotton fabrics, paper, for water chlorination, disinfection, and also for degassing areas contaminated with persistent toxic substances. The whitening and disinfecting properties of bleach are similar to the properties of javel and chlorine water: carbonic acid displaces hypochlorous acid from calcium hypochlorite; in the light, it decomposes with the release of atomic oxygen, which has an oxidizing effect.

Question

Functions of iodine in the body
Iodine is essential for the formation of thyroid hormones and for the functioning of macrophages. Macrophages are special cells that destroy various pathogenic microbes, viruses, fungi, etc.
What diseases are caused by iodine deficiency. Causes of iodine deficiency
Lack of iodine in the human body causes serious metabolic diseases (thyroid disease), mental retardation, and can also lead to chromosome damage and cancer. The concentration of cholesterol in the blood increases, all types of metabolism are disturbed. Perhaps the development of deafness, dumbness, paralysis, sterility, congenital malformations, miscarriage, drowsiness, edema, slowing of the heart rate.
Iodine deficiency develops due to inadequate intake with food and water, exposure to radiation, or due to the intake of certain drugs.

The norm of fluorine consumption. Role in the human body
Fluorine is an ambiguous element. Both excess and deficiency of fluorine are dangerous for human health. Fluorine is found in bones and teeth and is an essential element for building bone tissue. For a person, a sufficient amount of fluorine is 1-1.5 mg per 1 liter of water. We give data per liter of water because fluorine compounds are readily soluble. Fluoride is found in almost all foods and beverages. To date, it is impossible to talk about the development of fluorine deficiency, since almost all soils contain an excess of fluorine, which accumulates in excess in agricultural crops.
What causes excess and deficiency of fluoride?
The most well-known effect of a lack of fluoride in the body is the development of dental caries. An excess of fluorine causes osteochondrosis, changes in the shape and color of teeth (dental fluorosis), joint stiffness and the formation of bone growths. Marked loss of voice, dry choking cough, decreased pressure, hemorrhage. Contact with fluorine causes diseases of the skin (itching, irritation, desquamation) and mucous membranes, and also dramatically increases the risk of developing cancer of the gastrointestinal tract.
Causes of excess fluoride in modern products. What foods are high in fluoride
Fans of such a widespread drink as tea should know that the stronger the tea, and the longer you steep it, the more fluoride the drink contains. 1 liter of red wine contains 5 mg of fluorine - the maximum daily dose. Excess fluoride contains krill. In general, the excessive use of inorganic fertilizers in agricultural production has led to the accumulation of fluorine compounds in almost all plants.

Question

Iron- an element of the eighth group (according to the old classification - a side subgroup of the eighth group) of the fourth period of the periodic system of chemical elements D. I. Mendeleev with atomic number 26. Denoted by the symbol Fe(lat. Ferrum). One of the most common metals in the earth's crust (second place after aluminum).

A simple substance iron is a malleable silver-white metal with a high chemical reactivity: iron corrodes quickly at high temperatures or high humidity in the air. In pure oxygen, iron burns, and in a finely dispersed state, it ignites spontaneously in air.

Metabolism

oxygen exchange

Oxygen refers to organogenic elements. Its content is up to 65% of the human body weight, that is, more than 40 kg in an adult. Oxygen is the most common oxidizing agent on Earth, it is present in the environment in two forms - in the form of compounds (the earth's crust and water: oxides, peroxides, hydroxides, etc.) and in a free form (atmosphere).

The biological role of oxygen

The main (in fact, the only) function of oxygen is its participation as an oxidizing agent in redox reactions in the body. Due to the presence of oxygen, the organisms of all animals are able to utilize (actually “burn”) various substances ( , ) with the extraction of a certain “combustion” energy for their own needs. At rest, the body of an adult consumes 1.8-2.4 g of oxygen per minute.

Sources of oxygen

The main source of oxygen for humans is the Earth's atmosphere, from where, through breathing, the human body is able to extract the amount of oxygen necessary for life.

oxygen deficiency

With a deficiency in the human body, the so-called hypoxia develops.

Causes of oxygen deficiency

• absence or sharply reduced oxygen content in the atmosphere;
• reduced partial pressure of oxygen in the inhaled air (when climbing to high altitudes - in the mountains, aircraft);
• cessation or decrease in the supply of oxygen to the lungs during asphyxia;
• violations of oxygen transport (disturbances in the activity of the cardiovascular system, a significant decrease in hemoglobin in the blood during anemia, the inability of hemoglobin to perform its functions - to bind, transport or give oxygen to tissues, for example, in case of carbon monoxide poisoning);
• the inability of tissues to utilize oxygen due to a violation of redox processes in tissues (for example, with)

Consequences of oxygen deficiency

For acute hypoxia:

• loss of consciousness;
• disorder, irreversible damage and rapid death of the central nervous system (literally in minutes)

For chronic hypoxia:

• rapid physical and mental fatigue;
• disorders of the central nervous system;
• tachycardia and shortness of breath at rest or with little exertion

Excess oxygen

It is observed extremely rarely, as a rule, in artificial conditions (for example, hyperbaric chambers, improperly selected breathing mixtures when diving in water, etc.). In this case, prolonged inhalation of excessively oxygenated air is accompanied by oxygen poisoning - as a result of its excessive amount, a large amount of free radicals are formed in organs and tissues, the process of spontaneous oxidation of organic substances is initiated, including lipid peroxidation.

Plan:

Discovery history

Origin of name

Being in nature

Receipt

Physical Properties

Chemical properties

Application

The biological role of oxygen

Toxic oxygen derivatives

10. Isotopes

Oxygen

Oxygen- an element of the 16th group (according to the outdated classification - the main subgroup of group VI), the second period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 8. It is designated by the symbol O (lat. Oxygenium). Oxygen is a reactive non-metal and is the lightest element of the chalcogen group. simple substance oxygen(CAS number: 7782-44-7) under normal conditions - a colorless, tasteless and odorless gas, the molecule of which consists of two oxygen atoms (formula O 2), in connection with which it is also called dioxygen. Liquid oxygen has a light blue, and the solid is light blue crystals.

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).

1. Discovery history

It is officially believed that oxygen was discovered by the English chemist Joseph Priestley on August 1, 1774 by decomposing mercury oxide in a hermetically sealed vessel (Priestley directed the sun's rays at this compound using a powerful lens).

However, Priestley did not initially realize that he had discovered a new simple substance, he believed that he 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 an integral part of air, acids and is contained in many substances.

A few years earlier (in 1771), the Swedish chemist Carl Scheele had obtained oxygen. He calcined saltpeter with sulfuric acid and then decomposed the resulting nitric oxide. Scheele called this gas "fiery 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 stage that contributed to the discovery of oxygen was the work of the French chemist Pierre Bayen, who published work 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 great importance, because thanks to it, the phlogiston theory that dominated at that time and hindered the development of chemistry was overthrown. Lavoisier conducted an experiment on the combustion of various substances and refuted the theory of phlogiston by publishing the results on the weight of the burned elements. The weight of the ash exceeded the initial weight of the element, which gave Lavoisier the right to assert that during combustion a chemical reaction (oxidation) of the substance occurs, in connection with this, the mass of the original substance increases, which refutes the theory of phlogiston.

Thus, the credit for the discovery of oxygen is actually shared by Priestley, Scheele, and Lavoisier.

1. origin of name

The word oxygen (at the beginning of the 19th century it was still called "acid"), its appearance in the Russian language is to some extent due to M.V. Lomonosov, who introduced, along with other neologisms, the word "acid"; thus the word "oxygen", in turn, was a tracing-paper of the term "oxygen" (French oxygène), proposed by A. Lavoisier (from other Greek ὀξύς - "sour" and γεννάω - "I give birth"), which translates as “generating acid”, which is associated with its original meaning - “acid”, which previously meant substances called oxides according to modern international nomenclature.

1. Being in nature

Oxygen is the most common element on Earth; its share (as part of 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 1500 compounds of the earth's crust contain oxygen in their composition.

Oxygen is a constituent 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%, in terms of mass fraction - about 65%.