Iron number in the periodic table. Valency value and graphical formula

Iron, its position in the periodic table chemical elements D. I. Mendeleev.

In the periodic table of chemical elements of D.I. Mendeleev, iron Fe is located in the 4th period of group VIII of the secondary subgroup.

The distribution of electrons across the electron layers in an iron atom looks like this:

In basic condition.

In an excited state.

An iron atom has four electron layers. The d-sublevel of the third layer is filled with electrons; there are 6 electrons on it, and the s-sublevel on the fourth layer contains 2 electrons. In compounds, iron exhibits oxidation states +2 and +3.

Compounds with iron atoms in oxidation states +4, +6 and some others are also known.

Physical properties.

Iron is a typical metal, shiny silver white metal, its density is 7.87 g/cm3, m.p. 1539 C. Has good ductility. Iron is easily magnetized and demagnetized, and therefore is used as the cores of dynamos and electric motors. Iron consists of four stable isotopes with mass numbers 54,56,57 and 58. Iron is a moderate refractory metals. In the series of standard electrode potentials, iron is ranked before hydrogen and easily reacts with dilute acids.

Further, it is advisable to note that iron, after aluminum, is the most common metal in nature (the total content in the earth’s crust is 4.65% by weight). Known big number minerals that contain iron: magnetite ( magnetic iron ore) - Fe3O4, hematite (red iron ore) - Fe2O3, iron spar(siderite) - FeCO3, iron pyrite - FeS2, etc.

Chemical properties.

Iron is characterized by oxidation states of +2 and +3.

The oxidation state +2 corresponds to black oxide FeO and green hydroxide Fe(OH) 2. They are basic in nature. In salts, Fe(+2) is present as a cation. Fe(+2) is a weak reducing agent.

The oxidation state +3 corresponds to the red-brown oxide Fe 2 O 3 and the brown hydroxide Fe(OH) 3. They are amphoteric in nature, although acidic, and their basic properties are weakly expressed. Thus, Fe 3+ ions are completely hydrolyzed even in an acidic environment. Fe(OH) 3 dissolves (and even then not completely) only in concentrated alkalis. Fe 2 O 3 reacts with alkalis only upon fusion, giving ferrites (formal salts of the acid HFeO 2, which does not exist in free form):

Iron (+3) most often exhibits weak oxidizing properties.

Oxidation states +2 and +3 easily change between each other when redox conditions change.

In addition, there is the oxide Fe 3 O 4, the formal oxidation state of iron in which is +8/3. However, this oxide can also be considered as iron (II) ferrite Fe +2 (Fe +3 O 2) 2.

There is also an oxidation state of +6. The corresponding oxide and hydroxide do not exist in free form, but salts are obtained - ferrates (for example, K 2 FeO 4). Iron (+6) is present in them in the form of an anion. Ferrates are strong oxidizing agents.

Properties of a simple substance.

When stored in air at temperatures up to 200 °C, iron is gradually covered with a dense film of oxide, which prevents further oxidation of the metal. In humid air, iron becomes covered with a loose layer of rust, which does not prevent the access of oxygen and moisture to the metal and its destruction. Rust does not have a constant chemical composition; its approximate chemical formula can be written as Fe 2 O 3 xH 2 O.

Interacts with acids.

With hydrochloric acid:

With dilute sulfuric acid:

· Concentrated nitric and sulfuric acids passivate iron. It reacts with concentrated sulfuric acid only when heated:

· Interaction with oxygen:

Combustion of iron in air:

Combustion in pure oxygen:

Passing oxygen or air through molten iron:

· Interaction with sulfur powder when heated:

· Interaction with halogens when heated:

· Combustion in chlorine:

· At high blood pressure bromine vapor:

Interaction with iodine:

· Interaction with non-metals:

With nitrogen when heated:

With phosphorus when heated:

With carbon:

With silicon:

· Interaction of hot iron with water vapor:

· Iron reduces metals that are to the right of it in the activity series from salt solutions:

Iron reduces iron(III) compounds:

At elevated pressure, metallic iron reacts with carbon monoxide (II) CO, and a liquid is formed, at normal conditions highly volatile iron pentacarbonyl Fe(CO)5. Iron carbonyls of the compositions Fe2(CO)9 and Fe3(CO)12 are also known. Iron carbonyls serve as starting materials in the synthesis of organoiron compounds, including ferrocene of the composition (η5-C5H5)2Fe.

Pure metallic iron is stable in water and dilute alkali solutions. Iron does not dissolve in cold concentrated sulfuric and nitric acids due to passivation of the metal surface by a strong oxide film. Hot concentrated sulfuric acid, being a stronger oxidizing agent, interacts with iron.

Iron(II) compounds.

Iron(II) oxide FeO has basic properties; the base Fe(OH) 2 corresponds to it. Iron (II) salts have a light green color. When stored, especially in humid air, they turn brown due to oxidation to iron (III). The same process occurs when storing aqueous solutions of iron(II) salts:

Of the iron(II) salts in aqueous solutions, the most stable is Mohr's salt - double ammonium and iron(II) sulfate (NH 4) 2 Fe(SO 4) 2 6H 2 O.

Potassium hexacyanoferrate(III) K3 (red blood salt) can serve as a reagent for Fe 2+ ions in solution. When Fe 2+ and 3− ions interact, a precipitate of potassium iron (II) hexacyanoferrate (III) (Prussian blue) precipitates:

which rearranges intramolecularly into potassium iron(III) hexacyanoferrate(II):

Iron(III) compounds.

Iron(III) oxide Fe2O3 is the most stable natural oxygen-containing compound gland.

Iron(III) oxide Fe2O3 is weakly amphoteric; it is matched by an even weaker base than Fe (OH)2, Fe(OH)3, which reacts with acids:

Fe3+ salts are prone to the formation of crystalline hydrates. In them, the Fe3+ ion is usually surrounded by six water molecules. These salts are pink or purple in color.

The Fe3+ ion is completely hydrolyzed even in an acidic environment. At pH>4, this ion is almost completely precipitated in the form of Fe(OH)3:

With partial hydrolysis of the Fe3+ ion, polynuclear oxo- and hydroxocations are formed, which is why the solutions turn brown.

The basic properties of iron(III) hydroxide Fe(OH)3 are very weakly expressed. It is capable of reacting only with concentrated solutions of alkalis:

The resulting hydroxo complexes of iron(III) are stable only in strongly alkaline solutions. When solutions are diluted with water, they are destroyed, and Fe(OH)3 precipitates.

When alloyed with alkalis and oxides of other metals, Fe2O3 forms a variety of ferrites:

Iron(III) compounds in solutions are reduced by metallic iron:

Iron(III) is capable of forming double sulfates with singly charged cations such as alum, for example, KFe(SO4)2 - iron-potassium alum, (NH4)Fe(SO4)2 - iron-ammonium alum, etc.

For qualitative detection of iron(III) compounds in solution, a qualitative reaction of Fe3+ ions with inorganic thiocyanates SCN− is used. In this case, a mixture of bright red thiocyanate iron complexes 2+, +, Fe(SCN)3, - is formed. The composition of the mixture (and therefore the intensity of its color) depends on various factors Therefore, this method is not applicable for accurate qualitative determination of iron.

Another high-quality reagent for Fe3+ ions is potassium hexacyanoferrate (II) K4 (yellow blood salt). When Fe3+ and 4− ions interact, a bright blue precipitate of potassium iron (III) hexacyanoferrate (II) is formed:

Fe3+ ions are quantitatively determined by the formation of red (in a slightly acidic environment) or yellow (in a weakly acidic environment) alkaline environment) complexes with sulfosalicylic acid. This reaction requires proper selection of buffers, since some anions (in particular, acetate) form mixed complexes with iron and sulfosalicylic acid with their own optical characteristics.

Iron(VI) compounds.

Ferrates are salts of iron acid H2FeO4, which does not exist in free form. These are violet-colored compounds, reminiscent of permanganates in oxidative properties, and sulfates in solubility. Ferrates are obtained by the action of gaseous chlorine or ozone on a suspension of Fe(OH)3 in alkali:

Ferrates can also be obtained by electrolysis of a 30% alkali solution on an iron anode:

Ferrates are strong oxidizing agents. In an acidic environment they decompose with the release of oxygen:

Oxidative properties Ferrates are used to disinfect water.

Finding in nature: In the earth's crust, iron is quite widespread - it accounts for about 4.1% of the mass of the earth's crust (4th place among all elements, 2nd among metals). A large number of ores and minerals containing iron are known. Of greatest practical importance are red iron ores (hematite ore, Fe2O3; contains up to 70% Fe), magnetic iron ores (magnetite ore, Fe3O4; contains 72.4% Fe), brown iron ores (hydrogethite ore HFeO2 nH2O), as well as spar iron ores ( siderite ore, iron carbonate, FeCO3; contains about 48% Fe). In nature, there are also large deposits of pyrite FeS2 (other names are sulfur pyrite, iron pyrite, iron disulfide and others), but ores with high content sulfur for now practical significance Dont Have. Russia ranks first in the world in terms of iron ore reserves. In sea water there is 1·10–5 - 1·10–8% iron.

Biological role.

Iron is an essential component of hemoglobin, myoglobin, cytochromes, peroxidases and catalases. The iron and transferrin complex binds to specific receptors on the membranes of proliferating erythroid cells, and iron enters the cell. With iron deficiency, the body produces red blood cells with insufficient hemoglobin content, therefore the main manifestation of iron deficiency is hypochromic anemia. Treatment with iron preparations leads to a gradual regression of clinical (for example, weakness, fatigue, dizziness, tachycardia, soreness and dry skin) and laboratory symptoms.

Iron is present in the bodies of all plants and animals as a trace element, that is, in very small quantities (on average about 0.02%). However, iron bacteria, which use the energy of oxidation of iron (II) into iron (III) for chemosynthesis, can accumulate up to 17-20% iron in their cells. The main biological function of iron is participation in oxygen (O) transport and oxidative processes. Iron performs this function as part of complex proteins - hemoproteins, the prosthetic group of which is the iron porphyrin complex - heme. Among the most important hemoproteins are the respiratory pigments hemoglobin and myoglobin, universal electron carriers in the reactions of cellular respiration, oxidation and photosynthesis, cytochromes, catalose and peroxide enzymes, and others. In some invertebrates, the iron-containing respiratory pigments heloerythrin and chlorocruorin have a structure different from hemoglobins. During the biosynthesis of hemoproteins, iron is transferred to them from the protein ferritin, which stores and transports iron. This protein, one molecule of which contains about 4,500 iron atoms, is concentrated in the liver, spleen, bone marrow and intestinal mucosa of mammals and humans. A person’s daily need for iron (6-20 mg) is abundantly covered by food (meat, liver, eggs, bread, spinach, beets and others are rich in iron). The body of an average person (body weight 70 kg) contains 4.2 g of iron, 1 liter of blood contains about 450 mg. When there is a lack of iron in the body, glandular anemia develops, which is treated with drugs containing iron. Iron supplements are also used as general strengthening agents. An excessive dose of iron (200 mg or more) can have a toxic effect. Iron is also necessary for the normal development of plants, which is why there are microfertilizers based on iron preparations.

Iron supplements- a group of drugs containing salts or complexes of divalent and trivalent iron, as well as their combinations with other drugs. Mainly used for the treatment and prevention of iron deficiency anemia.

Iron supplements are indicated for:

Iron deficiency conditions (main indication);

· with intolerance to cow's milk;

· children who have suffered acute or long-term infectious diseases.

Iron deficiency can be caused by:

· insufficient intake of iron into the body of the fetus (during feto-fetal and feto-maternal transfusion), child or adult;

· impaired absorption from the intestinal lumen (malabsorption syndrome, inflammatory processes in the intestines, while taking tetracycline antibiotics and other drugs);

· acute massive or chronic blood loss (bleeding, helminthic infestations, nasal hemorrhages, juvenile uterine bleeding, prolonged hematuria and others);

· result of increased iron consumption (period of intensive growth, infectious diseases and others).

Side effects.

When taking iron supplements orally, dyspeptic effects (nausea, vomiting, diarrhea) may occur. The degree of their severity is higher, the more unabsorbed drug remains in the intestinal lumen. Reduced iron (only 0.5%) is absorbed worst of all (lowest bioavailability) from the gastrointestinal tract; it is these drugs that most often lead to intestinal dysfunction (should not be used in children).

By activating free radical reactions, iron preparations can damage cell membranes (including increasing the degree of hemolysis of red blood cells).

After parenteral administration of iron preparations, undesirable effects may occur: due to an increase in the concentration of free iron in the blood, the tone of small vessels - arterioles and venules - decreases - their permeability increases. Redness of the skin of the face, neck, and a rush of blood to the head and chest are observed. Further administration of the drug in this case is contraindicated. If the drug is not stopped, hemosiderosis develops in the future. internal organs and fabrics.

An overdose of iron taken orally results in bloody diarrhea and vomiting. With an overdose of any iron preparation, peripheral vascular resistance decreases, fluid transudation increases, and the volume of circulating blood decreases. As a result, blood pressure drops and tachycardia occurs.

In general, this category of medicines can be divided into several main groups: preparations based on divalent and trivalent iron salts, various complex iron compounds and combination products. Iron salt preparations are prescribed only orally.

Salts ferrous iron.

The absorption of iron by the cells of the gastrointestinal mucosa from salt compounds mainly occurs in divalent form, since apoferritin in enterocytes can only bind to Fe 2 ions. Therefore, preparations based on various iron (II) salts (sulfate, fumarate, gluconate, succinate, glutamate, lactate, etc.) have greater bioavailability, and general case more preferable than preparations containing iron(III) salts. In addition, they are the cheapest drugs compared to other iron preparations.

Despite these advantages, iron salt preparations also have significant disadvantages, in particular, a high level of gastrointestinal side effects (about 23%) when using high dosages. The bioavailability of iron(II) salts may decrease when interacting with various components food and other medications (phytins, oxalates, tannins, antacids, etc.), and therefore they are prescribed on an empty stomach, although this increases their negative effect on the intestinal mucosa. Any overdose of these drugs easily leads to acute poisoning(in the United States, between 1986 and 1996, there were 100 thousand reports of poisoning of children under 6 years of age with iron salts), which also somewhat limits their widespread use in children.

The main representatives of preparations made from divalent iron salts are products based on iron sulfate heptahydrate FeSO 4 7H 2 O(elemental iron content - 20% by weight of salt). Ferrous sulfate is highly soluble in water and, like other water-soluble salts, has relatively high bioavailability. It should be noted that iron(II) sulfate in a humid environment gradually oxidizes to iron(III) sulfate, which imposes some restrictions on its storage and use (cannot be used in the form of solutions, syrups and other liquid forms). Several trade names of medicines containing ferrous sulfate are registered in Russia: "Tardiferon", "Hemofer prolongatum", "Fenuls". Ferrous sulfate is also sometimes used in combination with stabilizing agents, such as ascorbic acid, which acts as an antioxidant (trade names "Sorbifer Durules", "Ferroplex").

Preparations based on ferric chloride tetrahydrate FeCl 2 4H 2 O(iron content 28%), unlike ferrous sulfate, do not oxidize in aqueous solutions, therefore they are produced in the form of drops for oral administration (trademark registered in Russia - “Hemofer”). When taking such drugs, it should be taken into account that solutions of iron salts can cause darkening of the teeth due to the deposition on their surface of insoluble iron sulfide, formed by the interaction of Fe 2+ ions with hydrogen sulfide, which may be contained in the oral cavity (for example, during dental caries).

Iron fumarate FeC 4 H 2 O 4(elemental iron content 33% by weight of the salt), unlike previous salts, is less soluble in water, but dissolves well in dilute solutions of acids, such as gastric juice. Therefore, preparations based on ferrous fumarate are more stable, do not have a characteristic iron taste, and do not bind to proteins in upper sections Gastrointestinal tract, but at the same time they dissolve well directly in the stomach and therefore are not inferior in bioavailability to water-soluble salts. Ferrous fumarate is registered in Russia as a medicine, but this moment did not receive distribution.

Ferric salts.

Preparations from ferric salts are traditionally less preferable compared to iron(II) salts, since in order to be absorbed by the body, Fe 3+ ions must first be reduced to Fe 2+, which is the reason for their lower bioavailability. In addition, iron(III) salts in the upper parts of the small intestine are easily hydrolyzed to form poorly soluble hydroxides, which also reduces their digestibility.

Manganese.

Manganese is an element of the side subgroup of the seventh group of the fourth period of the periodic system of chemical elements of D. I. Mendeleev with atomic number 25.

The electronic formula of manganese is:
1s2 2s2 2p6 3s2 3p6 4s2 3d5
Valence electrons are found in the 4s and 3d sublevels. There are 7 electrons in the valence orbitals of the manganese atom.

Distribution of Manganese in nature. The average content of manganese in the earth's crust is 0.1%, in most igneous rocks it is 0.06-0.2% by mass, where it is in a dispersed state in the form of Mn 2+ (an analogue of Fe 2+). On the earth's surface, Mn 2+ is easily oxidized; the minerals Mn 3+ and Mn 4+ are also known here. In the biosphere, Manganese migrates vigorously in reducing conditions and is inactive in an oxidizing environment. Manganese is most mobile in the acidic waters of the tundra and forest landscapes, where it is found in the form of Mn 2+. The Manganese content here is often elevated and cultivated plants in some places suffer from excess Manganese; Iron-manganese nodules, lake and swamp ores are formed in soils, lakes, and swamps. In dry steppes and deserts under conditions of an alkaline oxidizing environment, Manganese is inactive, organisms are poor in Manganese, and cultivated plants often need manganese microfertilizers. River waters are poor in Manganese (10 -6 -10 -5 g/l), however, the total removal of this element by rivers is enormous, and the bulk of it is deposited in the coastal zone. There is even less Manganese in the water of lakes, seas and oceans; In many places on the ocean floor, iron-manganese nodules formed in past geological periods are common.

Manganese minerals.

· pyrolusite MnO 2 · x H 2 O, the most common mineral (contains 63.2% manganese);

· manganite (brown manganese ore) MnO(OH) (62.5% manganese);

· brownite 3Mn 2 O 3 · MnSiO 3 (69.5% manganese);

· hausmannite (Mn II Mn 2 III) O 4 ;

· rhodochrosite (manganese spar, crimson spar) MnCO 3 (47.8% manganese);

· psilomelane m MnO MnO 2 n H 2 O (45-60% manganese);

· purpurite Mn 3+, (36.65% manganese).

Chemical properties.

Chemically, Manganese is quite active; when heated, it energetically interacts with non-metals - oxygen (a mixture of Manganese oxides of different valencies is formed), nitrogen, sulfur, carbon, phosphorus and others. At room temperature Manganese does not change in air: it reacts very slowly with water. It dissolves easily in acids (hydrochloric, dilute sulfuric), forming divalent manganese salts. When heated in a vacuum, Manganese easily evaporates even from alloys.

Manganese forms alloys with many chemical elements; most metals dissolve in their individual modifications and stabilize them. Thus, Cu, Fe, Co, Ni and others stabilize the γ-modification. Al, Ag and others expand the β- and σ-Mn regions in binary alloys. This is important for the production of Manganese-based alloys that are amenable to plastic deformation (forging, rolling, stamping).

In compounds, Manganese usually exhibits a valency from 2 to 7 (the most stable oxidation states are +2, +4 and +7). With an increase in the degree of oxidation, oxidative and acid properties Manganese compounds.

Mn(+2) compounds are reducing agents.

Oxide MnO- powder grey-green color; has basic properties. insoluble in water and alkalis, highly soluble in acids. Mn(OH)3 hydroxide is a white substance, insoluble in water. Mn(+4) compounds can act both as oxidizing agents (a) and reducing agents (b):

MnO 2 + 4HCl = MnCl 2 + Cl 2 + 2H 2 O (a)

(according to this edition, chlorine is produced in laboratories)

MnO 2 + KClO 3 + 6KOH = 3K 2 MnO 4 + KCl + 3H 2 O (b)

(the reaction occurs during fusion).

Manganese (II) oxide MnO2- black-brown in color, the corresponding Mn(OH)4 hydroxide is dark brown in color. Both compounds are insoluble in water, both are amphoteric with a slight predominance of acidic function. Salts of the K2MnO4 type are called manganites.

Of the Mn(+6) compounds, the most typical are permanganous acid and its manganate salts. Mn(+7) compounds are very important - manganese acid, manganese anhydride and permanganates.

Characteristic oxidation states of manganese: 0, +2, +3, +4, +6, +7 (+1, +5 are not very characteristic).

Passivates during oxidation in air. Powdered manganese burns in oxygen (Mn + O2 → MnO2). When heated, manganese decomposes water, displacing hydrogen (Mn + 2H2O → (t) Mn(OH)2 + H2), the resulting manganese hydroxide slows down the reaction.

Manganese absorbs hydrogen, and with increasing temperature its solubility in manganese increases. At temperatures above 1200 °C it reacts with nitrogen, forming nitrides of various compositions.

Carbon reacts with molten manganese to form carbides Mn3C and others. It also forms silicides, borides, and phosphides.

Reacts with hydrochloric and sulfuric acids according to the equation:

With concentrated sulfuric acid, the reaction proceeds according to the equation:

With dilute nitric acid, the reaction proceeds according to the equation:

Manganese is stable in alkaline solution.

Manganese forms the following oxides: MnO, Mn2O3, MnO2, MnO3 (not isolated in the free state) and manganese anhydride Mn2O7.

Mn2O7 under normal conditions is a liquid oily substance, dark Green colour, very unstable; when mixed with concentrated sulfuric acid, it ignites organic substances. At 90 °C Mn2O7 decomposes explosively. The most stable oxides are Mn2O3 and MnO2, as well as the combined oxide Mn3O4 (2MnO·MnO2, or salt Mn2MnO4).

When manganese (IV) oxide (pyrolusite) is fused with alkalis in the presence of oxygen, manganates are formed:

The manganate solution has a dark green color. When acidified, the reaction occurs:

The solution turns crimson due to the appearance of the MnO4− anion, and a brown precipitate of manganese (IV) oxide-hydroxide precipitates from it.

Manganese acid is very strong, but unstable, it cannot be concentrated to more than 20%. The acid itself and its salts (permanganates) are strong oxidizing agents. For example, potassium permanganate, depending on the pH of the solution, oxidizes various substances, being reduced to manganese compounds of varying degrees of oxidation. In an acidic environment - to manganese (II) compounds, in a neutral environment - to manganese (IV) compounds, in a strongly alkaline environment - to manganese (VI) compounds.

When heated, permanganates decompose with the release of oxygen (one of the laboratory methods for producing pure oxygen). The reaction proceeds according to the equation (using the example of potassium permanganate):

Under the influence of strong oxidizing agents, the Mn2+ ion transforms into the MnO4− ion:

This reaction is used for the qualitative determination of Mn2+.

When solutions of Mn(II) salts are alkalized, a precipitate of manganese(II) hydroxide precipitates out of them, which quickly turns brown in air as a result of oxidation.

Salts MnCl3, Mn2(SO4)3 are unstable. The hydroxides Mn(OH)2 and Mn(OH)3 are basic in nature, MnO(OH)2 is amphoteric. Manganese (IV) chloride MnCl4 is very unstable, decomposes when heated, which is used to produce chlorine:

The zero oxidation state of manganese manifests itself in compounds with σ-donor and π-acceptor ligands. Thus, carbonyl of the composition Mn2(CO)10 is known for manganese.

Other manganese compounds with σ-donor and π-acceptor ligands (PF3, NO, N2, P(C5H5)3) are also known.

Biological role.

Manganese in the body. Manganese is widely distributed in nature, being a constant integral part plant and animal organisms. The Manganese content in plants is ten-thousandths to hundredths, and in animals - hundredth-thousandths to thousandths of a percent. Invertebrate animals are richer in manganese than vertebrates. Among plants, significant amounts of Manganese accumulate in some rust fungi, water chestnut, duckweed, bacteria of the genera Leptothrix, Crenothrix and some diatoms (Cocconeis) (up to several percent in ash), among animals - red ants, some mollusks and crustaceans (up to hundredths of a percent ). Manganese is an activator of a number of enzymes, participates in the processes of respiration, photosynthesis, biosynthesis nucleic acids and others, enhances the effect of insulin and other hormones, affects hematopoiesis and mineral metabolism. Manganese deficiency in plants causes necrosis, chlorosis of apple and citrus fruits, spotting of cereals, burns of potatoes, barley, etc. Manganese is found in all human organs and tissues (the liver, skeleton and thyroid gland are the richest in it). The daily requirement of animals and humans for Manganese is several mg (a person receives 3-8 mg of Manganese daily from food). The need for Manganese increases with physical activity, with a lack sunlight; Children need more Manganese than adults. It has been shown that the lack of Manganese in animal food negatively affects their growth and development, causes anemia, so-called lactation tetany, and a violation of mineral metabolism of bone tissue. To prevent these diseases, Manganese salts are added to the feed.

Biological action manganese:
● antioxidant
● regulating blood glucose levels
● normalizing cholesterol levels and blood lipid composition
● antianemic
● antiallergic
● promoting maturation of germ cells, fetal development and full-term pregnancy
● restoring the structure of bone and cartilage tissue
● anticonvulsant, preventing PMS (premenstrual syndrome), etc.

Signs of manganese deficiency:

● Fatigue, weakness, dizziness, tinnitus
● Deterioration of brain activity, memory loss
vomit
● Spasms and cramps
● Pain in muscles and joints, movement disorders, tendency to sprains and sprains, arthritis, abnormal growth and development skeletal system
● Visual impairment
● Vitiligo, skin pigmentation disorders
● Delayed growth of nails and hair
● Diabetes, decreased glucose tolerance, overweight, high cholesterol, metabolic problems
● Risk of infertility, reproductive problems, early menopause, ovarian dysfunction, osteoporosis during menopause
● Decreased immunity, premature aging
● Allergies
● Risk of cancer
● Developmental delay in children, appearance of children with pathologies

Signs of manganese toxicity:

Excess manganese is toxic: it interferes with the absorption of iron and competes with copper in the process of hematopoiesis, causing anemia, and also causes other pathological changes.
● Weak appetite, apathy, depression
● General weakness, impotence
● Disturbed sleep
● Temporary insanity, dementia
● Neurological problems
● Parkinsonism or Parkinson's disease (muscle rigidity, tremors, monotone voice, frozen, mask-like face).

Cobalt

Co, chemical element with atomic number 27. Its atomic mass 58.9332. Natural cobalt consists of two stable nuclides: 59 Co (99.83% by weight) and 57 Co (0.17%). In the periodic system of elements of D.I. Mendeleev, cobalt is included in group VIII and, together with iron and nickel, forms in the 4th period in this group a triad of transition metals with similar properties. The configuration of the two outer electron layers of the cobalt atom is 3s 2 p 6 d 7 4s 2. It forms compounds most often in the +2 oxidation state, less often in the +3 oxidation state, and very rarely in the +1, +4 and +5 oxidation states.

Cobalt is a mineral that is part of vitamin B12. Usually measured in micrograms (mcg). Cobalt - essential for red blood cells. Must be obtained from food sources. There is no established daily value for cobalt, and only very small amounts of this mineral are needed in the diet (usually no more than 8 mcg).

Being in nature.

In the earth's crust, the cobalt content is 4·10 -3% by mass. Cobalt is a component of more than 30 minerals. These include carolite CuCo 2 S 4, linneite Co 3 S 4, cobaltine CoAsS, spherocobaltite CoCO 3, smaltite CoAs 2 and others. As a rule, cobalt in nature is accompanied by its neighbors in the 4th period - nickel, iron, copper, manganese. In sea water there is approximately (1-7)·10 -10% cobalt.

Cobalt is a relatively rare metal, and deposits rich in it are now almost exhausted. Therefore, cobalt-containing raw materials (often nickel ores containing cobalt as an impurity) are first enriched and a concentrate is obtained from it. Next, to extract cobalt, the concentrate is either treated with solutions of sulfuric acid or ammonia, or processed by pyrometallurgy into a sulfide or metal alloy. This alloy is then leached with sulfuric acid. Sometimes, to extract cobalt, sulfuric acid “heap” leaching of the original ore is carried out (crushed ore is placed in high heaps on special concrete platforms and these heaps are watered with a leaching solution on top).

Physical properties.

Cobalt is a hard metal that exists in two modifications. At temperatures from room temperature to 427 °C, the α-modification is stable. At temperatures from 427 °C to the melting point (1494 °C), the β-modification of cobalt (face-centered cubic lattice) is stable. Cobalt is a ferromagnet, Curie point 1121 °C. A thin layer of oxides gives it a yellowish tint.

Chemical properties.

Oxides.

· In air, cobalt oxidizes at temperatures above 300 °C.

· Cobalt oxide, stable at room temperature, is a complex oxide Co 3 O 4, having a spinel structure, in crystal structure of which one part of the nodes is occupied by Co 2+ ions, and the other by Co 3+ ions; decomposes to form CoO above 900 °C.

· At high temperatures, the α-form or β-form of CoO oxide can be obtained.

· All cobalt oxides are reduced with hydrogen:

Cobalt(III) oxide can be obtained by calcining cobalt(II) compounds, for example:

Other connections.

· When heated, cobalt reacts with halogens, and cobalt (III) compounds are formed only with fluorine.

· With sulfur, cobalt forms 2 different modifications of CoS. Silver-gray α-form (when powders are fused) and black β-form (precipitates from solutions).

When heating CoS in

Iron (lat. Ferrum) is a chemical element of group VIII of the periodic system of Mendeleev; atomic number 26, atomic mass 55.847.
Iron can be called the main metal of our time. This chemical element has been very well studied. Nevertheless, scientists do not know when and by whom iron was discovered: it was too long ago. Man began to use iron products at the beginning of the 1st millennium BC. The Bronze Age was replaced by the Iron Age. Iron metallurgy in Europe and Asia began to develop in the 9th-7th centuries. BC.
The first iron that fell into the hands of man was probably not earthly origin. Every year more than a thousand meteorites fall on our Earth, some of them are iron, consisting mainly of nickel iron. The largest iron meteorite discovered weighs about 60 tons. It was found in 1920 in southwestern Africa. “Heavenly” iron has one important technological feature: when heated, this metal cannot be forged; only cold meteorite iron can be forged. Weapons made from “heavenly” metal remained extremely rare and precious for many centuries.
Iron was also discovered on the Moon, and in the lunar soil it is present in a native, non-oxidized state, which is obviously explained by the absence of an atmosphere.
On Earth, iron is also sometimes found in a native state.
In ancient times, iron was highly valued. In the “Geography” of the ancient Greek scientist Strabo, written at the very beginning of our era, it is said that among the African peoples, iron was 10 times more expensive than gold... Perhaps this is fair if we consider the main criterion for high cost not chemical resistance and rarity, but value for technology, for the development of civilization. The main reasons that iron has become the most important metal for technology and production are the prevalence of compounds of this element and the relative ease of recovering metal from them.
The bulk of iron is found in deposits that can be developed industrially.
In terms of reserves in the earth's crust, iron ranks fourth among all elements, after oxygen, silicon and aluminum. There is much more iron in the planet's core, which, according to scientists, consists of nickel and iron. But this hardware is not available and is unlikely to become available in the foreseeable future. Therefore, the most important source of iron remains such minerals as magnetite Fe3O4, hydrogoethite FeO2-nH2O, hematite Fe2O3 and siderite FeCO3, located on the Earth’s surface or at shallow depths. They form the basis of the main iron ores - magnetic, brown, red and spar iron ore. The most iron, 72.4%, is in magnetite. The largest iron ore deposits in the USSR are the Kursk magnetic anomaly, the Krivoy Rog iron ore deposit, in the Urals (Mountains Magnitnaya, Vysokaya, Blagodat), in Kazakhstan the Sokolovskoye and Sarbaiskoye deposits.
Iron is a shiny silver-white metal that is easy to process: cutting, forging, rolling, stamping. It can be given greater strength and hardness using thermal and mechanical methods (hardening, rolling).

When talking about the properties of iron, it is necessary first of all to stipulate what kind of iron we are talking about - technically pure iron or iron of the highest purity. The difference in their properties - both physical and chemical - is quite large. Technically pure iron is called low-carbon electrical steel. This name reflects both the purpose of the material and the nature of the main impurities: carbon 0.02-0.04%, and oxygen, sulfur, nitrogen and phosphorus even less. Highest purity iron contains less than 0.001% impurities. Both materials have good magnetic properties, both weld well. However, if technically pure iron metal is of average chemical activity, then high purity is almost inert. The solubility of gases in it, especially oxygen, is also very low. The mechanical properties of high-purity iron are low, and the strength is much less than that of any steel or cast iron. Iron of the highest purity is unsuitable as a structural material. However, if high-purity iron is introduced into in a certain order alloying additives, it will be able to withstand loads of up to 600 kg/cm2 instead of the usual 17-21.
Iron in compounds can exhibit different oxidation states: + 2, +3, +6, rarely + 1, -r 4 and even 0 (in carbonyl Fe(CO)5). Of the divalent iron compounds, the best known are FeO, iron (II) oxide, as well as its sulfide and halides. Fe ions are formed when iron is dissolved in dilute acids. But in concentrated strong acids- nitrogen and sulfur - iron does not dissolve: it, as experts say, is passivated due to the formation of a thin and dense oxide film on the metal surface. Iron practically does not dissolve in alkalis (except for hot concentrated solutions).
Ferric iron salts Fe(III) are usually obtained by the oxidation of ferrous iron salts. Moreover, if a reaction occurs in a solution, the color of the solution changes; The light green color characteristic of Fe2+ changes to brown. Ferric salts are often prone to hydrolysis. Ferrous Н2FeО4 and ferrous НFeО2 acids were not obtained in a free state. However, their salts - ferrates and ferrites - are known and studied quite well.
Ferric oxide Fe2O3. The oxide of composition Fe3O4 is considered as a compound of FeO and Fe2O3. Di- and trivalent iron hydroxides Fe(OH)2 and Fe(OH)3 are poorly soluble in water and, unlike oxides, are of little practical importance. Oxides are important not only as a source of many iron compounds, but also as the most important raw material for ferrous metallurgy.
Like other transition metals, iron also forms many complex compounds.
Many iron compounds are practically important. Iron chloride FeCl3, for example, is used as a coagulant in water purification and as a catalyst in organic synthesis. Ferrites, especially divalent metals, are widely used in computer technology. It is important not to confuse two concepts: ferrites, salts of ferrous acid, and ferrite, a polymorphic modification of iron that is stable under normal conditions, otherwise called alpha iron.
For normal life, humans absolutely need iron-containing organic compounds. The most famous of them is the respiratory pigment hemoglobin. But in addition to hemoglobin, our body also contains iron in myolgobin, a protein that stores oxygen in the muscles.
It also contains iron-containing enzymes. Finally, there is the protein complex ferritin, from which all other iron-containing substances necessary for the body are formed.

Iron is a chemical element

1. The position of iron in the periodic table of chemical elements and the structure of its atom

Iron is a group VIII d element; serial number– 26; atomic mass Ar(Fe ) = 56; atomic composition: 26 protons; 30 – neutrons; 26 – electrons.

Atomic structure diagram:

Electronic formula: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2

Medium activity metal, reducing agent:

Fe 0 -2 e - → Fe +2 , the reducing agent is oxidized

Fe 0 -3 e - → Fe +3 , the reducing agent is oxidized

Main oxidation states: +2, +3

2. Iron prevalence

Iron is one of the most common elements in nature . In the earth's crust its mass fraction is 5.1%, according to this indicator it second only to oxygen, silicon and aluminum. A lot of iron is also found in celestial bodies, which is established from data spectral analysis. In samples of lunar soil delivered by the Luna automatic station, iron was found in an unoxidized state.

Iron ores are quite widespread on Earth. The names of the mountains in the Urals speak for themselves: Vysokaya, Magnitnaya, Zheleznaya. Agrochemists find iron compounds in soils.

Iron is a component of most rocks. To obtain iron, iron ores with an iron content of 30-70% or more are used.

The main iron ores are :

magnetite(magnetic iron ore) – Fe3O4 contains 72% iron, deposits are found in the Southern Urals, Kursk magnetic anomaly:

hematite(iron sheen, bloodstone)– Fe2O3 contains up to 65% iron, such deposits are found in the Krivoy Rog region:

limonite(brown iron ore) – Fe 2 O 3* nH 2 O contains up to 60% iron, deposits are found in Crimea:

pyrite(sulfur pyrite, iron pyrite, cat gold) – FeS 2 contains approximately 47% iron, deposits are found in the Urals.

3. The role of iron in the life of humans and plants

Biochemists have discovered the important role of iron in the life of plants, animals and humans. Being part of an extremely complex organic compound called hemoglobin, iron determines the red color of this substance, which in turn determines the color of human and animal blood. The body of an adult contains 3 g of pure iron, 75% of which is part of hemoglobin. The main role of hemoglobin is to transport oxygen from the lungs to the tissues, and in the opposite direction - CO 2.

Plants also need iron. It is part of the cytoplasm and participates in the process of photosynthesis. Plants grown on a substrate that does not contain iron have white leaves. A small addition of iron to the substrate and they turn green. Moreover, it is worth smearing a white sheet with a solution of salt containing iron, and soon the smeared area turns green.

So, for the same reason - the presence of iron in juices and tissues - the leaves of plants turn cheerfully green and a person’s cheeks brightly blush.

4. Physical properties of iron.

Iron is a silvery-white metal with a melting point of 1539 o C. It is very ductile, therefore it is easily processed, forged, rolled, stamped. Iron has the ability to be magnetized and demagnetized, therefore it is used as electromagnet cores in various electrical machines and devices. It can be given greater strength and hardness by thermal and mechanical methods, for example, by hardening and rolling.

There are chemically pure and commercially pure iron. Technically pure iron is essentially low-carbon steel; it contains 0.02-0.04% carbon, and even less oxygen, sulfur, nitrogen and phosphorus. Chemically pure iron contains less than 0.01% impurities. Chemically pure iron - silver-gray, shiny metal, very similar in appearance to platinum. Chemically pure iron is resistant to corrosion and has good resistance to acids. However insignificant shares impurities deprive it of these precious properties.

5. Getting iron

Reduction from oxides with coal or carbon monoxide (II), as well as hydrogen:

FeO + C = Fe + CO

Fe 2 O 3 + 3CO = 2Fe + 3CO 2

Fe 2 O 3 + 3H 2 = 2Fe + 3H 2 O

Experiment "Production of iron by aluminothermy"

6. Chemical properties of iron

As a secondary subgroup element, iron can exhibit several oxidation states. We will consider only compounds in which iron exhibits oxidation states +2 and +3. Thus, we can say that iron has two series of compounds, in which it is di- and trivalent.

1) In air, iron easily oxidizes in the presence of moisture (rusting):

4Fe + 3O 2 + 6H 2 O = 4Fe(OH) 3

2) Hot iron wire burns in oxygen, forming scale - iron oxide (II,III) - a black substance:

3Fe + 2O 2 = Fe 3 O 4

Coxygen in moist air is formed Fe 2 O 3 * nH 2 O

Experiment "Interaction of iron with oxygen"

3) At high temperatures (700–900°C), iron reacts with water vapor:

3Fe + 4H 2 O t˚C → Fe 3 O 4 + 4H 2

4) Iron reacts with non-metals when heated:

Fe + S t˚C → FeS

5) Iron easily dissolves in hydrochloric and dilute sulfuric acids under normal conditions:

Fe + 2HCl = FeCl 2 + H 2

Fe + H 2 SO 4 (dil.) = FeSO 4 + H 2

6) Iron dissolves in concentrated oxidizing acids only when heated

2Fe + 6H 2 SO 4 (conc. .) t˚C → Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O

Fe + 6HNO 3 (conc. .) t˚C → Fe(NO 3) 3 + 3NO 2 + 3H 2 OIron(III)

7. Use of iron.

The bulk of the iron produced in the world is used to produce cast iron and steel - alloys of iron with carbon and other metals. Cast irons contain about 4% carbon. Steels contain less than 1.4% carbon.

Cast irons are necessary for the production of various castings - heavy machine frames, etc.

Cast iron products

Steels are used to make machines, various building materials, beams, sheets, rolled products, rails, tools and many other products. To produce various grades of steel, so-called alloying additives are used, which are various metals: M

Simulator No. 2 - Genetic series Fe 3+

Simulator No. 3 - Equations of reactions of iron with simple and complex substances

Tasks for consolidation

No. 1. Write down reaction equations for the production of iron from its oxides Fe 2 O 3 and Fe 3 O 4, using as a reducing agent:
a) hydrogen;
b) aluminum;
c) carbon monoxide (II).
For each reaction, create an electronic balance.

No. 2. Carry out transformations according to the scheme:
Fe 2 O 3 -> Fe - +H2O, t -> X - +CO, t -> Y - +HCl ->Z
Name products X, Y, Z?

Iron English Iron, French Fer, German Eisen) is one of the seven metals of antiquity. It is very likely that man became acquainted with iron of meteorite origin earlier than with other metals. Meteoric iron is usually easy to distinguish from terrestrial iron, since it almost always contains from 5 to 30% nickel, most often 7-8%. Since ancient times, iron has been obtained from ores that occur almost everywhere. The most common ores are hematite (Fe 2 O 3), brown iron ore (2Fe 2 O 3, ZN 2 O) and its varieties (swamp ore, siderite, or spar iron FeCO,), magnetite (Fe 3 0 4) and some others . All these ores, when heated with coal, are easily reduced at a relatively low temperature, starting from 500 o C. The resulting metal had the appearance of a viscous spongy mass, which was then processed at 700-800 o With repeated forging.

The etymology of the names of iron in ancient languages quite clearly reflects the history of our ancestors’ acquaintance with this metal. Many ancient peoples undoubtedly became acquainted with it as a metal that fell from the sky, that is, as meteorite iron. Thus, in ancient Egypt, iron had the name bi-ni-pet (benipet, Coptic - benipe), which literally means heavenly ore, or heavenly metal. During the era of the first dynasties of Ur in Mesopotamia, iron was called an-bar (heavenly iron). The Ebers Papyrus (previously 1500 BC) contains two references to iron; in one case it is spoken of as a metal from the city of Kazi (Upper Egypt), in another - as a metal of heavenly manufacture (artpet). The ancient Greek name for iron, as well as the North Caucasian one - zido, is associated with the oldest word surviving in the Latin language - sidereus (stellar from Sidus - star, luminary). On ancient and modern Armenian language iron is called erkat, which means dripped (fell) from the sky. The fact that ancient people initially used iron of meteorite origin is also evidenced by the myths widespread among some peoples about gods or demons who dropped iron objects and tools from the sky - plows, axes, etc. It is also interesting that by the time of the discovery of America the Indians and Eskimos North America They were not familiar with the methods of obtaining iron from ores, but they knew how to process meteorite iron.

In ancient times and in the Middle Ages, the seven then known metals were compared with the seven planets, which symbolized the connection between metals and celestial bodies and the celestial origin of metals. This comparison became common more than 2000 years ago and is constantly found in literature until the 19th century. In the II century. n. e. iron was compared with Mercury and was called mercury, but later it began to be compared with Mars and called Mars, which, in particular, emphasized the external similarity of the reddish color of Mars with red iron ores.

However, some peoples did not connect the name of iron with the celestial origin of the metal. Thus, among the Slavic peoples, iron is called on a “functional” basis. Russian iron (South Slavic zalizo, Polish zelaso, Lithuanian gelesis, etc.) has the root “lez” or “rez” (from the word lezo - blade). This word formation directly indicates the function of objects made of iron - cutting tools and weapons. The prefix “zhe” is apparently a softening of the more ancient “ze” or “for”; it was preserved in its original form among many Slavic peoples (among the Czechs - zelezo). Old German philologists - representatives of the theory of Indo-European, or, as they called it, Indo-Germanic proto-language - sought to produce Slavic names from German and Sanskrit roots. For example, Fik compares the word iron with the Sanskrit ghalgha (molten metal, from ghal - to glow). But this is unlikely to correspond to reality: after all, iron smelting was inaccessible to ancient people. It is more likely that the Greek name for copper can be compared with the Sanskrit ghalgha, but not the Slavic word iron. Functional sign the names of iron are reflected in other languages. Thus, in Latin, along with the usual name for steel (chalybs), derived from the name of the Khalib tribe, who lived on the southern coast of the Black Sea, the name acies was used, literally meaning blade or point. This word corresponds exactly to the ancient Greek, which was used in the same sense. Let us mention in a few words the origin of the German and English names for iron. Philologists generally accept that the German word Eisen has Celtic origin, as well as the English Iron. Both terms reflect the Celtic names of rivers (Isarno, Isarkos, Eisack), which were then transformed) isarn, eisarn) and turned into Eisen. There are, however, other points of view. Some philologists derive the German Eisen from the Celtic isara, meaning "strong, strong." There are also theories that Eisen comes from ayas or aes (copper), and also from Eis (ice), etc. The Old English name for iron (before 1150) is iren; it was used along with isern and isen and passed into the Middle Ages. Modern Iron came into use after 1630. Note that in Ruland’s “Alchemical Lexicon” (1612) the word Iris is given as one of the old names for iron, meaning “rainbow” and consonant with Iron.

Became international Latin name Ferrum is accepted among the Romance peoples. It is probably related to the Greco-Latin fars (to be hard), which comes from the Sanskrit bhars (to harden). A comparison is also possible with ferreus, which among ancient writers meant “insensitive, unyielding, strong, hard, heavy,” as well as with ferre (to wear). Alchemists, along with Ferrum ynot, used many other names, for example Iris, Sarsar, Phaulec, Minera, etc.

Iron products made from meteorite iron were found in burials dating back to very ancient times (4th - 5th millennia BC) in Egypt and Mesopotamia. However iron age in Egypt began only in the 12th century. BC e., and in other countries even later. IN ancient Russian literature the word iron appears in ancient monuments(from the 11th century) under the names iron, iron, iron.

D.I. Mendeleev, interaction with sulfur, hydrochloric acid, salt solutions.

ANSWER PLAN:

position in p.s. and atomic structure physical properties chemical properties The chemical element iron is in the 4th period, 8th group, secondary subgroup. An iron atom has four electron layers. The d-sublevel of the third layer is filled with electrons; there are 6 electrons on it, and the s-sublevel on the fourth layer contains 2 electrons. In compounds, iron exhibits oxidation states +2 and +3.

IV period VIII group secondary subgroup	Fe))))	+2	+3
+26 2 8 8+6 2	4s	??
3d	??	?	?	?	?

The simple substance iron is a silvery-white metal with a melting point of 15390C, a density of 7.87 g/cm3, and has magnetic properties. Iron is a reactive metal. When heated, it reacts with sulfur to form iron(II) sulfide: Fe0 + S0 = Fe+2S-2. Iron displaces hydrogen from acid solutions, and iron(II) salts are formed, for example, when iron is exposed to hydrochloric acid, iron(II) chloride is formed: Fe0 + 2H+1Cl-1 = Fe+2Cl2-1 + H20. Iron can displace less active metals from solutions of their salts, for example, when iron acts on a solution of copper(II) sulfate, metallic copper and iron(II) sulfate are formed: Fe0 + Cu+2SO4 = Cu0 + Fe+2SO4.

In all reactions, iron exhibits the properties of a reducing agent. Stronger oxidizing agents - chlorine, oxygen, concentrated acids– oxidize iron to oxidation state +3.

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1. Iron: position in the periodic table of chemical elements by D.I. Mendeleev, atomic structure, possible oxidation states, physical properties, interaction with oxygen, halogens, solutions of acids and salts. The role of iron in modern technology. Iron alloys.

Iron is in the secondary subgroup of group VIII of the periodic table. Electronic formula of the iron atom:

Typical oxidation states of iron are +2 and +3. The +2 oxidation state occurs due to the loss of two 4s electrons. The oxidation state +3 also corresponds to the loss of one more Zd electron, and the Zd level is half filled; such electronic configurations are relatively stable.

Physical properties. Iron is a typical metal; it forms a metallic crystal lattice. Iron conducts electricity, is quite refractory, melting point 1539°C. Iron differs from most other metals in its ability to be magnetized.

Chemical properties . Iron reacts with many non-metals:

Iron scale is formed - mixed iron oxide. Its formula is also written as follows: FeO Fe2O3.

Reacts with acids to release hydrogen:

It enters into substitution reactions with metal salts located to the right of iron in the voltage series:

Iron compounds. FeO is a basic oxide that reacts with acid solutions to form iron (II) salts. Fe2O3 is an amphoteric oxide that also reacts with alkali solutions.

Iron hydroxides. Fe(OH)2 is a typical basic oxide; Fe(OH)3 has amphoteric properties and reacts not only with acids, but also with concentrated solutions of alkalis.

Iron (II) hydroxide is easily oxidized to iron (III) hydroxide by atmospheric oxygen:

When iron (II) and (III) salts react with alkalis, they precipitate insoluble hydroxides:

Iron alloys. The modern metallurgical industry produces iron alloys of various compositions.

All iron alloys are divided into two groups according to composition and properties. The first group includes various types of cast iron, the second group includes various types of steel.

Cast iron is brittle; steel is ductile, they can be forged, rolled, drawn, stamped. The difference in the mechanical properties of cast iron and steel depends primarily on their carbon content - cast iron contains about 4% carbon, and steel usually contains less than 1.4%.

In modern metallurgy, cast iron is first produced from iron ores, and then steel is obtained from cast iron. Pig iron is smelted in blast furnaces, steel is cooked in steel smelting furnaces. Up to 90% of all smelted iron is processed into steel.

Cast iron. Pig iron intended for processing into steel is called pig iron. It contains from 3.9 to 4.3% C, 0.3-1.5% Si, 1.5-3.5% Mn, no more than 0.3% P and no more than 0.07% S. Cast iron , intended "for producing castings, is called foundry iron. Ferroalloys are also smelted in blast furnaces, which are used primarily in the production of steel as additives. Ferroalloys have, compared to pig iron, an increased content of silicon (ferrosilicon), manganese (ferromanganese), chromium ( ferrochrome) and other elements.

Become. All steels are divided into carbon and alloy.

Carbon steels contain several times less carbon, silicon and manganese than cast iron, and very little phosphorus and sulfur. The properties of carbon steel depend primarily on the carbon content in it: the more carbon in the steel, the harder it is. The industry produces soft steels, medium-hard steels and hard steels. Soft steels and medium-hard steels are used for the manufacture of machine parts, pipes, bolts, nails, etc., and hard steels are used for the manufacture of tools.

Steels should contain as little sulfur and phosphorus as possible, since these impurities worsen the mechanical properties of steels. In increased quantities, sulfur causes red brittleness - the formation of cracks during hot machining of metal. Phosphorus causes cold brittleness in steel at ordinary temperatures. -

Alloy steels. The physical, chemical and mechanical properties of steels change significantly from the introduction into their composition of an increased amount of manganese and silicon, as well as chromium, nickel, tungsten and other elements. These elements are called alloying elements, and steels are called alloying elements [from the Latin word ligare - to bind, to connect].

Chromium is the most widely used alloying element. Especially great importance for the construction of machines, apparatus and many machine parts have chromium-nickel become. These steels have high ductility, strength, heat resistance and resistance to oxidizing agents. Nitric acid any concentration does not destroy them even at boiling temperatures. Chrome-nickel steels do not rust in atmospheric conditions and in water. Shiny, silver-colored sheets of chromium-nickel steel decorate the arches of the Mayakovskaya station of the Moscow metro. Stainless knives, spoons, forks and other household items are made from the same steel.

Molybdenum and vanadium increase the hardness and strength of steels at elevated temperatures and pressures. So, chrome molybdenum And chrome vanadium steels are used for the manufacture of pipelines and compressor parts in the production of synthetic ammonia and aircraft engines.

When cutting at high speeds, the tool becomes very hot and wears out quickly. By adding tungsten, the hardness of the steel is maintained even at elevated temperatures. Therefore, chrome-tungsten steels are used for the manufacture of cutting tools operating at high speeds."

Increasing the manganese content in steel increases its resistance to friction and impact. Manganese steels are used for the manufacture of railway ramps, switches, crosses, and stone crushing machines.

The use of alloy steels makes it possible to significantly reduce the weight of metal structures, increase their strength, durability and operational reliability.

2. Squirrels as biopolymers. Primary, secondary and tertiary structures of proteins. Properties and biological functions proteins.

Proteins (proteins, polypeptides) are high-molecular organic substances consisting of alpha-amino acids connected in a chain by a peptide bond.

Proteins, like polysaccharides, are biological polymers. Most protein molecules reach gigantic sizes compared to others organic compounds and has a very high molecular weight:

Molecular formula one of the proteins of the penicillin group - C 43 H 58 N 4 O 12; casein - cow's milk protein, - C 47 H 48 N 3 NaO 7 S 2; hemoglobin - C 3032 H 48I6 O 872 N 780 S 8 Fe 4;

Iron (symbol Fe)− chemical element of the eighth group, fourth period. Iron in the periodic table of chemical elements is located at number 26.

The Iron subgroup contains 4 elements: Fe iron, ruthenium Ru, osmium Os, Hs hasmium.

Characteristics of the chemical element Iron

Ferrum- Latin word, it means not only iron, but also hardness and weapons. From it came the names of iron in some European languages: French fer, Italian ferro, Spanish hierro and such terms as ferrites, ferromagnetism. Similar names for this metal in Slavic and Baltic languages: Lithuanian gelezis, Polish zelazo, Bulgarian zhelez, Ukrainian zalizo and Belarusian zalez. The English name Iron, German Eisen, Dutch ijzer are derived from the Sanskrit isira (strong, strong).

Distribution of Iron in Nature

Iron 26 element of the periodic table

Iron- first on globe and the second most abundant metal in the earth's crust, a very important metal for humans. Since time immemorial, people have encountered iron in the form of iron meteorites. Typically, meteorite iron contains from 5 to 30% nickel, almost 0.5% cobalt and up to 1% other elements. On the territory of Africa, 80 thousand years ago, the most large meteorite Goba, it weighed 66 tons. It contains 84% gland and 16% nickel. In the meteorite museum of the Russian Academy of Sciences, two fragments of an iron meteorite are stored, which weigh 256 kg, which fell on Far East. In 1947, in the Primorsky Territory, over an area of 35 km 2, thousands of fragments (weighing from 60 to 100 tons) of an iron meteorite fell like “iron rain”. A very rare mineral - native iron of terrestrial origin, occurs in the form of small grains and contains 2% nickel and tenths of a percent of other metals. Native iron was found on the Moon in a crushed state.

In the 13th-12th centuries BC. There is a collapse and change of cultures throughout the entire space of Eurasia from the Atlantic to the Pacific Ocean, and over the course of several centuries - until the 10th-8th centuries BC. migrations of peoples occur. This period was called the Bronze Age catastrophe and the beginning of the transition to the Iron Age.

There is a lot of iron in the earth's crust, but it is difficult to extract. This metal is tightly bound in ores with oxygen and sometimes with sulfur. Ancient furnaces could not produce the required temperature at which pure iron melted and iron was obtained in the form of a sponge with impurities from an ore called kritsa. When forging the kritsa, the iron was partially separated from the ore.

Many minerals contain iron. Magnetic iron ore, containing 72.3% iron, is the richest mineral in iron. The ancient Greek philosopher Thales of Miletus more than 2,500 years ago studied samples of ferrous metal that attract iron. He gave it the name magnetis lithos - a stone from Magnesia, which is how the name of the magnet came about. It is now known that it was magnetic iron ore - black iron oxide.

The role of iron in a living organism

The most important iron ore is hematite. It contains 69.9% iron. Hematite is also called red iron ore, and the ancient name is bloodstone. From the Greek haima, meaning blood. Other words related to blood also appeared, such as hemoglobin. Hemoglobin serves as a carrier of oxygen from the respiratory organs to the tissues of the body, and in the opposite direction carries carbon dioxide. Lack of iron in the body leads to a serious disease - iron deficiency anemia. With this disease, disorders of the skeleton, functions of the central nervous and vascular systems occur, and there is a lack of oxygen in the tissues. Iron is necessary for living organisms. It is also found in muscles, spleen and liver. An adult has about 4 g of iron; it is present in every cell of the body. A person should receive 15 milligrams of iron every day with food. If there is a lack of iron, doctors prescribe special medications that contain iron in an easily digestible form.

Applications of Iron

If the smelted iron contains more than 2% carbon, then cast iron is obtained; it is smelted hundreds of degrees lower than pure iron. Since cast iron is brittle, it can only be used to cast various products, it cannot be forged. Iron ore is smelted in blast furnaces a large number of cast iron, which is used for casting monuments, gratings and heavy machine beds. The bulk of cast iron is processed into steel. To do this, some of the carbon and other impurities are “burned out” of cast iron in converters or open-hearth furnaces.

All objects from rails to nails are made of steel with different carbon contents. If there is little carbon in the iron, soft low-carbon steel is obtained, and by introducing alloying impurities of other elements into the steel, different grades of special steels are obtained. A huge variety of steels are known and each has its own application.

The most famous is stainless steel, which contains nickel and chromium. Equipment for chemical plants and tableware are made from this steel. And if you add 18% tungsten, 1% vanadium and 4% chromium to steel, you get high-speed steel; drills and cutter tips are made from it. If you fuse iron with 1.5% carbon and 15% manganese, you get the kind of hard steel that is used to make bulldozer blades and excavator teeth. Steel containing 36% nickel, 0.5% carbon and 0.5% manganese is called invar; precision instruments and some watch parts are made from it. The steel, called platinite, contains 46% nickel and 15% carbon and expands when heated, just like glass. The junction of platinite with glass does not crack and therefore it is used in the manufacture of electric lamps.

Stainless steel is not magnetized and is not attracted to a magnet. Only carbon steel can be magnetized. Pure iron itself is not magnetized, but is attracted by a magnet; such iron is suitable for making electromagnet cores.

More than a billion tons of iron are smelted annually in the world. But corrosion, which is a terrible enemy of metal, not only destroys the metal itself, on the smelting of which enormous efforts were spent, but also disables finished products that are more expensive than the metal itself. It annually destroys tens of millions of tons of smelted metal. When iron corrodes, it reacts with oxygen and water, turning into rust.

METALS OF SUB-GROUPS

Characteristics of transition elements - copper, chromium, iron according to their position in the periodic system of chemical elements D.I. Mendeleev and the structural features of their atoms.

The term transition element is usually used to refer to any of the d- or f-elements. These elements occupy a transitional position between electropositive s-elements and electronegative p-elements. d-Elements form three transition series - in the 4th, 5th and 6th periods, respectively. The first transition series includes 10 elements, from scandium to zinc. It is characterized by the internal configuration of 3d orbitals. Chromium and copper have only one electron in their 4s orbitals. The fact is that half-filled or filled d-subshells are more stable than partially filled ones. The chromium atom has one electron in each of the five 3d orbitals that form the 3d subshell. This subshell is half-filled. In a copper atom, each of the five 3d orbitals contains a pair of electrons (the anomaly of silver is explained in a similar way). All d-elements are metals. Most of them have a characteristic metallic luster. Compared to s-metals, their strength is generally significantly higher. In particular, they are characterized by the following properties: high tensile strength; ductility; malleability (they can be flattened into sheets by blows). d-elements and their compounds have a number of characteristic properties: variable oxidation states; ability to form complex ions; formation of colored compounds. d-elements are also characterized by more high density compared to other metals. This is explained by the relatively small radii of their atoms. The atomic radii of these metals change little in this series. d-Elements are good conductors of electricity, especially those whose atoms have only one outer s-electron in addition to a half-filled or full d-shell. For example, copper.

Chemical properties.

The electronegativity of the metals of the first transition series increases in the direction from chromium to zinc. This means that the metallic properties of the elements of the first transition row gradually weaken in in the indicated direction. This change in their properties is also manifested in a consistent increase in redox potentials with a transition from negative to positive values.

Characteristics of chromium and its compounds

Chromium- hard, bluish-white metal.ρ = 7.2 g/cm 3, t melt = 1857 0 C CO: +1,+2,+3,+4,+5,+6

Chemical properties.

Under normal conditions, chromium reacts only with fluorine. At high temperatures (above 600 0 C) it interacts with oxygen, halogens, nitrogen, silicon, boron, sulfur, phosphorus.

4Cr + 3O 2 2Cr 2 O 3

2Cr + 3Cl 2 2CrCl 3

2Cr + 3S Cr 2 S 3

When heated, it reacts with water vapor:

2Cr + 3H 2 O Cr 2 O 3 + 3H 2

Chromium dissolves in dilute strong acids (HCl, H 2 SO 4). In the absence of air, Cr 2+ salts are formed, and in air, Cr 3+ salts are formed.

Cr + 2HCl → CrCl 2 + H 2 -

2Cr + 6HCl + O 2 → 2CrCl 3 + 2H 2 O + H 2 -

The presence of a protective oxide film on the surface of the metal explains its passivity towards cold concentrated acids - oxidizing agents. However, when heated strongly, these acids dissolve chromium:

2 Сr + 6 Н 2 SO 4 (conc) Сr 2 (SO 4) 3 + 3 SO 2 + 6 Н 2 О

Cr + 6 HNO 3 (conc) Cr(NO 3) 3 + 3 NO 2 + 3 H 2 O

Receipt.

Chromium compounds

Chromium oxide (II) CrO

Physical properties: a solid, water-insoluble substance of bright red or brownish-red color. Chemical properties. CrO is the main oxide.

Receipt.

Cr 2 O 3 + 3H 2 2Cr + 3H 2 O Chromium hydroxide (II) Cr(OH) 2 Physical properties: a yellow, water-insoluble solid. Chemical properties. Cr(OH) 2 is a weak base.

Interacts with acids: Cr(OH) 2 + 2HCl → CrCl 2 + 2H 2 O Easily oxidized in the presence of moisture by atmospheric oxygen in Cr(OH) 3:

4Cr(OH) 2 + O 2 + 2H 2 O → 4Cr(OH) 3

When heated, it decomposes: a) without air access: Cr(OH) 2 CrO + H 2 O b) in the presence of oxygen: 4Cr(OH) 2 2 Cr 2 O 3 + 4H 2 O Receipt.

The effect of alkali on solutions of Cr(II) salts: CrCl 2 + 2 NaOH = Cr(OH) 2 ↓ + 2 NaCl.

Trivalent chromium compounds

Chromium oxide (III) Cr 2 O 3 Physical properties: dark green, refractory substance, insoluble in water. Chemical properties. Cr 2 O 3 is an amphoteric oxide.

Sodium chromite

high temperature

Cr 2 O 3 + 3H 2 2Cr + 3H 2 O

Receipt.

Chromium hydroxide (III) Cr(OH) 3 Physical properties: green substance insoluble in water. Chemical properties. Cr(OH) 3 – amphoteric hydroxide

2Cr(OH) 3 + 3H 2 SO 4 →Cr 2 (SO 4) 3 + 6H 2 O

Cr(OH) 3 + KOH → KCrO 2 + 2H 2 O

(potassium chromite) Receipt.

When alkalis act on Cr 3+ salts, a gelatinous precipitate of green chromium (III) hydroxide precipitates:

Cr 2 (SO 4) 3 + 6NaOH → 2 Cr(OH) 3 ↓ + 3 Na 2 SO 4,

Hexavalent chromium compounds

Chromium oxide (VI) CrO 3 Physical properties: dark red solid, highly soluble in water. Poisonous! Chemical properties. CrO 3 is an acidic oxide.

Reacts with alkalis, forming yellow chromate salts:

CrO 3 + 2KOH → K 2 CrO 4 + H 2 O

Reacts with water to form acids: CrO 3 + H 2 O → H 2 CrO 4 chromic acid 2 CrO 3 + H 2 O → H 2 Cr 2 O 7 dichromic acid

Thermally unstable: 4 CrO 3 → 2Cr 2 O 3 + 3O 2Receipt.

Obtained from potassium chromate (or dichromate) by the action of H 2 SO 4 (conc.).

K 2 CrO 4 + H 2 SO 4 → CrO 3 + K 2 SO 4 + H 2 O

K 2 Cr 2 O 7 + H 2 SO 4 → 2CrO 3 + K 2 SO 4 + H 2 O

Hydroxideschromium(VI)H 2 CrO 4 - chromeacid, H 2 Cr 2 O 7 - dichromeacid Both acids are unstable, when trying to isolate them in pure form decompose into water and chromium(VI) oxide. However, their salts are quite stable. Salts of chromic acid are called chromates, they are colored yellow, and salts of dichromic acid are called dichromates, they are colored yellow. Orange color.

Iron and its compounds

Iron - a relatively soft malleable metal of silver color, ductile, magnetized. T melt = 1539 0 C. ρ = 7.87 g/cm 3. CO: +2 – with weak oxidizing agents – solutions of acids, salts, non-metals, except oxygen and halogens +3 – with strong oxidizing agents – concentrated acids, oxygen, halogens.

Chemical properties.

Burns in oxygen, forming scale - iron (II,III) oxide: 3Fe + 2O 2 → Fe 3 O 4 Iron reacts with non-metals when heated:

At high temperatures (700–900C), iron reacts with water vapor:

3Fe + 4H 2 O Fe 3 O 4 + 4H 2 -

In air in the presence of moisture it rusts: 4Fe + 3O 2 + 6H 2 O → 4Fe(OH) 3. Iron easily dissolves in hydrochloric and dilute sulfuric acids, exhibiting CO +2:

Fe + 2HCl → FeCl 2 + H 2 -

Fe + H 2 SO 4 (diluted) → FeSO 4 + H 2 -

In concentrated oxidizing acids, iron dissolves only when heated, exhibiting CO +3:

2Fe + 6H 2 SO 4 (conc.) Fe 2 (SO 4) 3 + 3SO 2 - + 6H 2 O

Fe + 6HNO 3 (conc.) Fe(NO 3) 3 + 3NO 2 - + 3H 2 O

(in the cold, concentrated nitric and sulfuric acids passivate iron).

Iron displaces metals that are to the right of it in the voltage series from solutions of their salts.

Fe + CuSO 4 → FeSO 4 + Cu ↓

Receipt.

Reduction from oxides with coal or carbon monoxide (II)

Fe 2 O 3 + 3CO 2Fe + 3CO 2

Ferrous compounds

ABOUTiron oxide (II) FeO

Physical properties: black solid, insoluble in water. Chemical properties: FeO – basic oxide 6 FeO + O 2 2Fe 3 O 4

Reduced by hydrogen, carbon, carbon monoxide (II) to iron:Receipt. Fe 3 O 4 + H 2 - 3 FeO + H 2 O

Iron hydroxide (II) Fe(OH) 2

Physical properties: powder white, insoluble in water. Chemical properties: Fe(OH) 2 is a weak base. Receipt.

Formed by the action of alkali solutions on iron (II) salts without air access:

FeCl 2 + 2KOH → 2KCl + Fe(OH) 2 ↓

Qualitative response to Fe 2+

When potassium hexacyanoferrate (III) K 3 (red blood salt) acts on solutions of ferrous iron salts, a blue precipitate (Turnboole blue) is formed:

3FeSO 4 + 2K 3  Fe 3 2  + 3K 2 SO 4

Ferric compounds

Iron oxide (III) Fe 2 O 3

Physical properties: red-brown solid. Chemical properties: Fe 2 O 3 is an amphoteric oxide. sodium ferrite Fe 2 O 3 + 3H 2 - 2 Fe + 3H 2 O Receipt.

Iron hydroxide (III) Fe(OH) 3

Physical properties: red-brown solid. Chemical properties: Fe(OH) 3 is an amphoteric hydroxide.

Reacts with acids as an insoluble base:

2Fe(OH) 3 + 3H 2 SO 4 →Fe 2 (SO 4) 3 + 6H 2 O

Reacts with alkalis as an insoluble acid:

Fe(OH) 3 + KOH (sol) → KFeO 2 + 2H 2 O

Fe(OH) 3 + 3KOH (conc) → K 3

Receipt.

Formed by the action of alkali solutions on ferric salts: it precipitates in the form of a red-brown precipitate:

Fe(NO 3) 3 + 3KOH  Fe(OH) 3  + 3KNO 3

Qualitative reactions to Fe 3+

When potassium hexacyanoferrate (II) K 4 (yellow blood salt) acts on solutions of ferric salts, a blue precipitate (Prussian blue) is formed:

4FeCl 3 +3K 4  Fe 4 3  + 12KCl

When potassium or ammonium thiocyanate is added to a solution containing Fe 3+ ions, an intense blood-red color of iron(III) thiocyanate appears:

FeCl 3 + 3KCNS  3КCl + Fe(CNS) 3

Copper and its compounds

Copper- quite soft metal red-yellow color, malleable, plastic, has high thermal and electrical conductivity. T melt = 1083 0 C. ρ = 8.96 g/cm 3. CO: 0,+1,+2

Chemical properties.

Interaction with simple substances.

Interaction with complex substances.

Copper is in the voltage series to the right of hydrogen, therefore it does not react with dilute hydrochloric and sulfuric acids, but dissolves in oxidizing acids:

3Cu + 8HNO 3 (dil.) → 3Cu(NO 3) 2 + 2NO- + 2H 2 O

Cu + 4HNO 3 (conc.) → Cu(NO 3) 2 + 2NO 2 -+ 2H 2 O

Cu + 2H 2 SO 4 (conc.) → CuSO 4 + SO 2 -+2H 2 O

Receipt.

CuO + CO Cu + CO 2

Cuprous compounds

Copper oxide(I) WITHu 2 O Physical properties: red solid, insoluble in water. Chemical properties: Cu 2 O is the main oxide. Receipt.

Obtained by reduction of copper (II) compounds, for example, glucose in an alkaline medium: 2CuSO 4 + C 6 H 12 O 6 + 5NaOH → Cu 2 O↓ + 2Na 2 SO 4 + C 6 H 11 O 7 Na + 3H 2 O Copper hydroxide(I) CuOH Physical properties: unstable, poorly soluble in water, yellow substance, not isolated in a free state. Chemical properties: CuOH is a weak base.

Interacts with acids: CuOH + HCl → CuCl + H 2 O In air, easily oxidizes to Cu(OH) 2: 4CuOH + O 2 + 2H 2 O → 4 Cu(OH) 2Receipt.

Cupric compounds

Iron number in the periodic table. Valency value and graphical formula

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Ferrous compounds

ABOUTiron oxide (II) FeO

Iron hydroxide (II) Fe(OH) 2

Qualitative response to Fe 2+

Ferric compounds

Iron oxide (III) Fe 2 O 3

Iron hydroxide (III) Fe(OH) 3

Qualitative reactions to Fe 3+