Iron is a metal of medium chemical activity. It is a constituent of many minerals: magnetite, hematite, limonite, siderite, pyrite.
Limonite sampleChemical and physical properties of iron
Under normal conditions and in its pure form, iron is a silvery-gray solid with a bright metallic luster. Iron is a good electrical and thermal conductor. This can be felt by touching an iron object in a cold room. Since metal conducts heat quickly, it takes most of the heat from human skin in a short period of time, so cold is felt when touched.
pure iron
The melting point of iron is 1538 °C, the boiling point is 2862 °C. The characteristic properties of iron are good ductility and fusibility.
Reacts with simple substances: oxygen, halogens (bromine, iodine, fluorine,), phosphorus, sulfur. When iron is burned, metal oxides are formed. Depending on the reaction conditions and the proportions between the two participants, iron oxides can be varied. Reaction equations:
2Fe + O₂ = 2FeO;
4Fe + 3O₂ = 2Fe₂O₃;
3Fe + 2O₂ = Fe₃O₄.
These reactions take place at high temperatures. you will learn what experiments to study the properties of iron can be done at home.
The reaction of iron with oxygen
For the reaction of iron with oxygen, preheating is necessary. Iron burns with a dazzling flame, scattering - red-hot particles of iron scale Fe₃O₄. The same reaction of iron and oxygen occurs in air, when it is strongly heated by friction during mechanical processing.
When iron is burned in oxygen (or in air), iron scale is formed. Reaction equation:
3Fe + 2O₂ = Fe₃O₄
3Fe + 2O₂ = FeO Fe₂O₃.
Iron oxide is a compound in which iron has different valence values.
Production of iron oxides
Iron oxides are products of the interaction of iron with oxygen. The most famous of them are FeO, Fe₂O₃ and Fe₃O₄.
Iron oxide (III) Fe₂O₃ is an orange-red powder formed during the oxidation of iron in air.
The substance is formed by the decomposition of a ferric salt in air at high temperature. A little iron (III) sulfate is poured into a porcelain crucible, and then it is calcined on the fire of a gas burner. Upon thermal decomposition, ferrous sulfate will decompose into sulfur oxide and iron oxide.
Iron oxide (II, III) Fe₃O₄ is formed by burning powdered iron in oxygen or in air. To obtain oxide, a little fine iron powder mixed with sodium or potassium nitrate is poured into a porcelain crucible. The mixture is ignited with a gas burner. When heated, potassium and sodium nitrates decompose with the release of oxygen. Iron in oxygen burns to form the oxide Fe₃O₄. After the end of combustion, the resulting oxide remains at the bottom of the porcelain cup in the form of iron scale.
Attention! Do not try to repeat these experiments yourself!
Iron(II) oxide FeO is a black powder that is formed by the decomposition of iron oxalate in an inert atmosphere.
The human body contains about 5 g of iron, most of it (70%) is part of the hemoglobin in the blood.
Physical properties
In the free state, iron is a silvery-white metal with a grayish tinge. Pure iron is ductile and has ferromagnetic properties. In practice, iron alloys are commonly used - cast irons and steels.
Fe is the most important and most common element of the nine d-metals of the secondary subgroup of group VIII. Together with cobalt and nickel, it forms the "iron family".
When forming compounds with other elements, it often uses 2 or 3 electrons (B \u003d II, III).
Iron, like almost all d-elements of group VIII, does not show a higher valency equal to the group number. Its maximum valency reaches VI and is extremely rare.
The most typical compounds are those in which the Fe atoms are in the +2 and +3 oxidation states.
Methods for obtaining iron
1. Commercial iron (in an alloy with carbon and other impurities) is obtained by carbothermal reduction of its natural compounds according to the scheme:
Recovery occurs gradually, in 3 stages:
1) 3Fe 2 O 3 + CO = 2Fe 3 O 4 + CO 2
2) Fe 3 O 4 + CO = 3FeO + CO 2
3) FeO + CO \u003d Fe + CO 2
The cast iron resulting from this process contains more than 2% carbon. In the future, steels are obtained from cast iron - iron alloys containing less than 1.5% carbon.
2. Very pure iron is obtained in one of the following ways:
a) decomposition of pentacarbonyl Fe
Fe(CO) 5 = Fe + 5CO
b) hydrogen reduction of pure FeO
FeO + H 2 \u003d Fe + H 2 O
c) electrolysis of aqueous solutions of Fe +2 salts
FeC 2 O 4 \u003d Fe + 2СO 2
iron(II) oxalate
Chemical properties
Fe - a metal of medium activity, exhibits general properties characteristic of metals.
A unique feature is the ability to "rust" in humid air:
In the absence of moisture with dry air, iron begins to noticeably react only at T > 150°C; when calcined, “iron scale” Fe 3 O 4 is formed:
3Fe + 2O 2 = Fe 3 O 4
Iron does not dissolve in water in the absence of oxygen. At very high temperatures, Fe reacts with water vapor, displacing hydrogen from water molecules:
3 Fe + 4H 2 O (g) \u003d 4H 2
The rusting process in its mechanism is electrochemical corrosion. The rust product is presented in a simplified form. In fact, a loose layer of a mixture of oxides and hydroxides of variable composition is formed. Unlike the Al 2 O 3 film, this layer does not protect the iron from further destruction.
Types of corrosion
Corrosion protection of iron
1. Interaction with halogens and sulfur at high temperature.
2Fe + 3Cl 2 = 2FeCl 3
2Fe + 3F 2 = 2FeF 3
Fe + I 2 \u003d FeI 2
Compounds are formed in which the ionic type of bond predominates.
2. Interaction with phosphorus, carbon, silicon (iron does not directly combine with N 2 and H 2, but dissolves them).
Fe + P = Fe x P y
Fe + C = Fe x C y
Fe + Si = FexSiy
Substances of variable composition are formed, since berthollides (the covalent nature of the bond prevails in the compounds)
3. Interaction with "non-oxidizing" acids (HCl, H 2 SO 4 dil.)
Fe 0 + 2H + → Fe 2+ + H 2
Since Fe is located in the activity series to the left of hydrogen (E ° Fe / Fe 2+ \u003d -0.44V), it is able to displace H 2 from ordinary acids.
Fe + 2HCl \u003d FeCl 2 + H 2
Fe + H 2 SO 4 \u003d FeSO 4 + H 2
4. Interaction with "oxidizing" acids (HNO 3 , H 2 SO 4 conc.)
Fe 0 - 3e - → Fe 3+
Concentrated HNO 3 and H 2 SO 4 "passivate" iron, so at ordinary temperatures the metal does not dissolve in them. With strong heating, slow dissolution occurs (without release of H 2).
In razb. HNO 3 iron dissolves, goes into solution in the form of Fe 3+ cations, and the acid anion is reduced to NO *:
Fe + 4HNO 3 \u003d Fe (NO 3) 3 + NO + 2H 2 O
It dissolves very well in a mixture of HCl and HNO 3
5. Attitude to alkalis
Fe does not dissolve in aqueous solutions of alkalis. It reacts with molten alkalis only at very high temperatures.
6. Interaction with salts of less active metals
Fe + CuSO 4 \u003d FeSO 4 + Cu
Fe 0 + Cu 2+ = Fe 2+ + Cu 0
7. Interaction with gaseous carbon monoxide (t = 200°C, P)
Fe (powder) + 5CO (g) \u003d Fe 0 (CO) 5 iron pentacarbonyl
Fe(III) compounds
Fe 2 O 3 - iron oxide (III).
Red-brown powder, n. R. in H 2 O. In nature - "red iron ore".
Ways to get:
1) decomposition of iron hydroxide (III)
2Fe(OH) 3 = Fe 2 O 3 + 3H 2 O
2) pyrite roasting
4FeS 2 + 11O 2 \u003d 8SO 2 + 2Fe 2 O 3
3) decomposition of nitrate
Chemical properties
Fe 2 O 3 is a basic oxide with signs of amphoterism.
I. The main properties are manifested in the ability to react with acids:
Fe 2 O 3 + 6H + = 2Fe 3+ + ZH 2 O
Fe 2 O 3 + 6HCI \u003d 2FeCI 3 + 3H 2 O
Fe 2 O 3 + 6HNO 3 \u003d 2Fe (NO 3) 3 + 3H 2 O
II. Weak acid properties. Fe 2 O 3 does not dissolve in aqueous solutions of alkalis, but when fused with solid oxides, alkalis and carbonates, ferrites are formed:
Fe 2 O 3 + CaO \u003d Ca (FeO 2) 2
Fe 2 O 3 + 2NaOH \u003d 2NaFeO 2 + H 2 O
Fe 2 O 3 + MgCO 3 \u003d Mg (FeO 2) 2 + CO 2
III. Fe 2 O 3 - feedstock for iron production in metallurgy:
Fe 2 O 3 + ZS \u003d 2Fe + ZSO or Fe 2 O 3 + ZSO \u003d 2Fe + ZSO 2
Fe (OH) 3 - iron (III) hydroxide
Ways to get:
Obtained by the action of alkalis on soluble salts Fe 3+:
FeCl 3 + 3NaOH \u003d Fe (OH) 3 + 3NaCl
At the time of receipt of Fe(OH) 3 - red-brown mucosamorphous precipitate.
Fe (III) hydroxide is also formed during the oxidation of Fe and Fe (OH) 2 in humid air:
4Fe + 6H 2 O + 3O 2 \u003d 4Fe (OH) 3
4Fe(OH) 2 + 2Н 2 O + O 2 = 4Fe(OH) 3
Fe(III) hydroxide is the end product of hydrolysis of Fe 3+ salts.
Chemical properties
Fe(OH) 3 is a very weak base (much weaker than Fe(OH) 2). Shows noticeable acidic properties. Thus, Fe (OH) 3 has an amphoteric character:
1) reactions with acids proceed easily:
2) a fresh precipitate of Fe(OH) 3 is dissolved in hot conc. solutions of KOH or NaOH with the formation of hydroxo complexes:
Fe (OH) 3 + 3KOH \u003d K 3
In an alkaline solution, Fe (OH) 3 can be oxidized to ferrates (salts of iron acid H 2 FeO 4 not isolated in the free state):
2Fe(OH) 3 + 10KOH + 3Br 2 = 2K 2 FeO 4 + 6KBr + 8H 2 O
Fe 3+ salts
The most practically important are: Fe 2 (SO 4) 3, FeCl 3, Fe (NO 3) 3, Fe (SCN) 3, K 3 4 - yellow blood salt \u003d Fe 4 3 Prussian blue (dark blue precipitate)
b) Fe 3+ + 3SCN - \u003d Fe (SCN) 3 Fe (III) thiocyanate (blood red solution)
Iron was known in prehistoric times, but it was widely used much later, since it is extremely rare in nature in the free state, and its production from ores became possible only at a certain level of technological development. Probably, for the first time, a person became acquainted with meteorite Iron, as evidenced by its names in the languages of ancient peoples: the ancient Egyptian "beni-pet" means "heavenly iron"; the ancient Greek sideros is associated with the Latin sidus (genus case sideris) - a star, a celestial body. In the Hittite texts of the 14th century BC. e. Iron is mentioned as a metal that fell from the sky. In the Romance languages, the root of the name given by the Romans has been preserved (for example, French fer, Italian ferro).
The method of obtaining Iron from ores was invented in the western part of Asia in the 2nd millennium BC. e.; after that, the use of Iron spread in Babylon, Egypt, Greece; The Bronze Age was replaced by the Iron Age. Homer (in the 23rd song of the Iliad) tells that Achilles awarded the winner of the discus throwing competition with an iron cry discus. In Europe and Ancient Russia for many centuries, iron was obtained by the cheese-making process. Iron ore was reduced with charcoal in a furnace built in a pit; air was pumped into the hearth with furs, the reduction product - kritsu was separated from the slag by hammer blows and various products were forged from it. As the methods of blowing were improved and the height of the hearth increased, the temperature of the process increased and part of the iron became carburized, that is, cast iron was obtained; this relatively fragile product was considered a waste product. Hence the name of cast iron "chushka", "pig iron" - English. pig iron. Later it was noticed that when not iron ore, but cast iron is loaded into the hearth, low-carbon iron bloom is also obtained, and such a two-stage process turned out to be more profitable than raw-dough. In the 12th-13th centuries, the screaming method was already widespread.
In the 14th century, cast iron began to be smelted not only as a semi-finished product for further processing, but also as a material for casting various products. The reconstruction of the hearth into a shaft furnace ("domnitsa"), and then into a blast furnace, also dates back to the same time. In the middle of the 18th century, the crucible process of obtaining steel began to be used in Europe, which was known in Syria in the early period of the Middle Ages, but later was forgotten. With this method, steel was obtained by melting a metal charge in small vessels (crucibles) from a highly refractory mass. In the last quarter of the 18th century, the puddling process of converting cast iron into iron began to develop on the hearth of a fiery reverberatory furnace. The industrial revolution of the 18th and early 19th centuries, the invention of the steam engine, the construction of railways, large bridges, and the steam fleet created an enormous demand for iron and its alloys. However, all existing methods of iron production could not meet the needs of the market. Mass production of steel began only in the middle of the 19th century, when the Bessemer, Thomas and open-hearth processes were developed. In the 20th century, the electric steelmaking process arose and became widespread, giving high quality steel.
Distribution of iron in nature. In terms of content in the lithosphere (4.65% by weight), iron ranks second among metals (aluminum is in first place). It migrates vigorously in the earth's crust, forming about 300 minerals (oxides, sulfides, silicates, carbonates, titanates, phosphates, etc.). Iron takes an active part in magmatic, hydrothermal and supergene processes, which are associated with the formation of various types of iron deposits. Iron is a metal of the earth's depths, it accumulates in the early stages of magma crystallization, in ultrabasic (9.85%) and basic (8.56%) rocks (it is only 2.7% in granites). In the biosphere, iron accumulates in many marine and continental sediments, forming sedimentary ores.
An important role in the geochemistry of iron is played by redox reactions - the transition of 2-valent iron to 3-valent and vice versa. In the biosphere, in the presence of organic substances, Fe 3+ is reduced to Fe 2+ and easily migrates, and when it encounters atmospheric oxygen, Fe 2+ is oxidized, forming accumulations of trivalent iron hydroxides. Widespread compounds of 3-valent Iron are red, yellow, brown. This determines the color of many sedimentary rocks and their name - "red-colored formation" (red and brown loams and clays, yellow sands, etc.).
Physical properties of iron. The importance of iron in modern technology is determined not only by its wide distribution in nature, but also by a combination of very valuable properties. It is plastic, easily forged both in a cold and heated state, can be rolled, stamped and drawn. The ability to dissolve carbon and other elements is the basis for obtaining a variety of iron alloys.
Iron can exist in the form of two crystal lattices: α- and γ-body-centered cubic (bcc) and face-centered cubic (fcc). Below 910°C, α-Fe with a bcc lattice is stable (a = 2.86645Å at 20°C). Between 910°C and 1400°C, the γ-modification with the fcc lattice is stable (a = 3.64Å). Above 1400°C, the δ-Fe bcc lattice (a = 2.94Å) is again formed, which is stable up to the melting point (1539°C). α-Fe is ferromagnetic up to 769 °C (Curie point). Modifications γ-Fe and δ-Fe are paramagnetic.
Polymorphic transformations of iron and steel during heating and cooling were discovered in 1868 by D.K. Chernov. Carbon forms interstitial solid solutions with Iron, in which C atoms having a small atomic radius (0.77 Å) are located at the interstices of the metal crystal lattice, which consists of larger atoms (Fe atomic radius 1.26 Å). A solid solution of carbon in γ-Fe is called austenite, and in α-Fe it is called ferrite. A saturated solid solution of carbon in γ-Fe contains 2.0% C by mass at 1130 °C; α-Fe dissolves only 0.02-0.04% C at 723 °C, and less than 0.01% at room temperature. Therefore, when austenite is quenched, martensite is formed - a supersaturated solid solution of carbon in α-Fe, which is very hard and brittle. The combination of quenching with tempering (heating to relatively low temperatures to reduce internal stresses) makes it possible to give the steel the required combination of hardness and ductility.
The physical properties of Iron depend on its purity. In industrial iron materials Iron is usually accompanied by impurities of carbon, nitrogen, oxygen, hydrogen, sulfur, and phosphorus. Even at very low concentrations, these impurities greatly change the properties of the metal. So, sulfur causes the so-called red brittleness, phosphorus (even 10 -2% P) - cold brittleness; carbon and nitrogen reduce plasticity, and hydrogen increases the brittleness of Iron (the so-called hydrogen brittleness). Reducing the content of impurities to 10 -7 - 10 -9% leads to significant changes in the properties of the metal, in particular to an increase in ductility.
The following are the physical properties of Iron, referring mainly to a metal with a total impurity content of less than 0.01% by mass:
Atomic radius 1.26Å
Ionic radii Fe 2+ 0.80Å, Fe 3+ 0.67Å
Density (20°C) 7.874 g/cm3
t bale about 3200°С
Temperature coefficient of linear expansion (20°C) 11.7 10 -6
Thermal conductivity (25°C) 74.04 W/(m K)
The heat capacity of Iron depends on its structure and changes in a complex way with temperature; average specific heat capacity (0-1000°C) 640.57 j/(kg K) .
Electrical resistivity (20°C) 9.7 10 -8 ohm m
Temperature coefficient of electrical resistance (0-100°C) 6.51 10 -3
Young's modulus 190-210 10 3 MN / m 2 (19-21 10 3 kgf / mm 2)
Temperature coefficient of Young's modulus 4 10 -6
Shear modulus 84.0 10 3 MN/m 2
Short-term tensile strength 170-210 MN/m2
Relative elongation 45-55%
Brinell hardness 350-450 MN/m2
Yield strength 100 MN/m2
Impact strength 300 MN/m2
Chemical properties of iron. The configuration of the outer electron shell of the atom is 3d 6 4s 2 . Iron exhibits a variable valency (the most stable compounds are 2- and 3-valent Iron). With oxygen, Iron forms oxide (II) FeO, oxide (III) Fe 2 O 3 and oxide (II,III) Fe 3 O 4 (compound of FeO with Fe 2 O 3 having a spinel structure). In humid air at ordinary temperatures, iron becomes covered with loose rust (Fe 2 O 3 nH 2 O). Due to its porosity, rust does not prevent the access of oxygen and moisture to the metal and therefore does not protect it from further oxidation. As a result of various types of corrosion, millions of tons of Iron are lost every year. When iron is heated in dry air above 200°C, it is covered with a very thin oxide film, which protects the metal from corrosion at ordinary temperatures; this is the basis of the technical method of protecting Iron - bluing. When heated in water vapor, iron is oxidized to form Fe 3 O 4 (below 570 °C) or FeO (above 570 °C) and release hydrogen.
Hydroxide Fe (OH) 2 is formed as a white precipitate by the action of caustic alkalis or ammonia on aqueous solutions of Fe 2+ salts in an atmosphere of hydrogen or nitrogen. Upon contact with air, Fe(OH) 2 first turns green, then turns black, and finally quickly turns into red-brown Fe(OH) 3 hydroxide. FeO oxide exhibits basic properties. Oxide Fe 2 O 3 is amphoteric and has a mildly acidic function; reacting with more basic oxides (for example, with MgO), it forms ferrites - compounds of the Fe 2 O 3 nMeO type, which have ferromagnetic properties and are widely used in radio electronics. Acidic properties are also expressed in 6-valent Iron, which exists in the form of ferrates, for example K 2 FeO 4 , salts of iron acid not isolated in the free state.
Iron easily reacts with halogens and hydrogen halides, giving salts, such as chlorides FeCl 2 and FeCl 3 . When iron is heated with sulfur, FeS and FeS 2 sulfides are formed. Iron carbides - Fe 3 C (cementite) and Fe 2 C (e-carbide) - precipitate from solid solutions of carbon in iron upon cooling. Fe 3 C is also released from solutions of carbon in liquid Iron at high concentrations of C. Nitrogen, like carbon, gives interstitial solid solutions with Iron; nitrides Fe 4 N and Fe 2 N are isolated from them. With hydrogen, iron gives only slightly stable hydrides, the composition of which has not been precisely determined. When heated, iron reacts vigorously with silicon and phosphorus to form silicides (eg Fe 3 Si and phosphides (eg Fe 3 P).
Iron compounds with many elements (O, S and others), which form a crystalline structure, have a variable composition (for example, the sulfur content in monosulfide can vary from 50 to 53.3 at.%). This is due to defects in the crystal structure. For example, in iron oxide (II), some of the Fe 2+ ions at the lattice sites are replaced by Fe 3+ ions; to maintain electrical neutrality, some lattice sites belonging to Fe 2+ ions remain empty.
The normal electrode potential of Iron in aqueous solutions of its salts for the reaction Fe = Fe 2+ + 2e is -0.44 V, and for the reaction Fe = Fe 3+ + 3e is -0.036 V. Thus, in the series of activities, iron is to the left of hydrogen. It readily dissolves in dilute acids with the release of H 2 and the formation of Fe 2+ ions. The interaction of iron with nitric acid is peculiar. Concentrated HNO 3 (density 1.45 g/cm 3) passivates Iron due to the formation of a protective oxide film on its surface; more dilute HNO 3 dissolves Iron with the formation of Fe 2+ or Fe 3+ ions, being reduced to NH 3 or N 2 and N 2 O. Solutions of salts of 2-valent Iron in air are unstable - Fe 2+ gradually oxidizes to Fe 3+. Aqueous solutions of iron salts are acidic due to hydrolysis. The addition of thiocyanate ions SCN- to solutions of Fe 3+ salts gives a bright blood-red color due to the appearance of Fe(SCN) 3, which makes it possible to reveal the presence of 1 part of Fe 3+ in about 10 6 parts of water. Iron is characterized by the formation of complex compounds.
Getting Iron. Pure iron is obtained in relatively small quantities by the electrolysis of aqueous solutions of its salts or by the reduction of its oxides with hydrogen. The production of sufficiently pure iron is gradually increasing by means of its direct reduction from ore concentrates with hydrogen, natural gas, or coal at relatively low temperatures.
The use of iron. Iron is the most important metal of modern technology. In its pure form, due to its low strength, iron is practically not used, although steel or cast iron products are often called "iron" in everyday life. The bulk of iron is used in the form of alloys with very different compositions and properties. Iron alloys account for approximately 95% of all metal products. Carbon-rich alloys (over 2% by weight) - cast iron, are smelted in blast furnaces from iron-rich ores. Steel of various grades (carbon content less than 2% by mass) is smelted from cast iron in open-hearth and electric furnaces and converters by oxidizing (burning out) excess carbon, removing harmful impurities (mainly S, P, O) and adding alloying elements. High-alloy steels (with a high content of nickel, chromium, tungsten and other elements) are smelted in electric arc and induction furnaces. New processes such as vacuum and electroslag remelting, plasma and electron-beam melting, and others are used for the production of steels and iron alloys for particularly important purposes. Methods are being developed for smelting steel in continuously operating units that ensure high quality of the metal and automation of the process.
Iron-based materials are created that can withstand the effects of high and low temperatures, vacuum and high pressures, aggressive media, high alternating voltages, nuclear radiation, etc. The production of iron and its alloys is constantly growing.
Iron as an art material has been used since ancient times in Egypt, Mesopotamia, and India. Since the Middle Ages, numerous highly artistic iron products have been preserved in European countries (England, France, Italy, Russia and others) - forged fences, door hinges, wall brackets, weather vanes, chest fittings, lights. Forged through products from rods and products from perforated sheet iron (often with a mica lining) are distinguished by planar shapes, a clear linear-graphic silhouette and are effectively visible against a light-air background. In the 20th century, iron is used for the manufacture of lattices, fences, openwork interior partitions, candlesticks, and monuments.
Iron in the body. Iron is present in the organisms of all animals and in plants (about 0.02% on average); it is necessary mainly for oxygen exchange and oxidative processes. There are organisms (the so-called concentrators) capable of accumulating it in large quantities (for example, iron bacteria - up to 17-20% of Iron). Almost all of the iron in animal and plant organisms is associated with proteins. Iron deficiency causes growth retardation and plant chlorosis associated with reduced chlorophyll production. An excess of iron also has a harmful effect on the development of plants, causing, for example, sterility of rice flowers and chlorosis. In alkaline soils, iron compounds that are inaccessible to plant roots are formed, and plants do not receive it in sufficient quantities; in acidic soils, iron passes into soluble compounds in excess. With a deficiency or excess of assimilable iron compounds in soils, plant diseases can be observed in large areas.
Iron enters the body of animals and humans with food (liver, meat, eggs, legumes, bread, cereals, spinach, and beets are the richest in iron). Normally, a person receives 60-110 mg of Iron with the diet, which significantly exceeds his daily requirement. The absorption of iron ingested with food occurs in the upper part of the small intestines, from where it enters the blood in a protein-bound form and is carried with the blood to various organs and tissues, where it is deposited in the form of an iron-protein complex - ferritin. The main depot of iron in the body is the liver and spleen. Due to ferritin, the synthesis of all iron-containing compounds of the body occurs: the respiratory pigment hemoglobin is synthesized in the bone marrow, myoglobin is synthesized in muscles, cytochromes and other iron-containing enzymes are synthesized in various tissues. Iron is excreted from the body mainly through the wall of the large intestine (in humans, about 6-10 mg per day) and to a small extent by the kidneys. The body's need for Iron varies with age and physical condition. For 1 kg of weight, children need - 0.6, adults - 0.1 and pregnant women - 0.3 mg of Iron per day. In animals, the need for Iron is approximately (per 1 kg of dry matter of the diet): for dairy cows - at least 50 mg, for young animals - 30-50 mg; for piglets - up to 200 mg, for pregnant pigs - 60 mg.
It is one of the most common elements in the earth's crust.
Physical properties of iron.
Iron- malleable silver-white metal with high chemical resistance. It tolerates high temperatures and humidity well. It quickly tarnishes (rusts) in air and in water. Very plastic, well gives in to forging and rolling. It has good thermal and electrical conductivity, an excellent ferromagnet.
Chemical properties of iron.
Iron transition metal. It can have an oxidation state of +2 and +3. Reacts with water vapor:
3 Fe + 4 H 2 O = Fe 3 O 4 + 4 H 2 .
But in the presence of moisture, iron rusts:
4 Fe + 3 O 2 + 6 H 2 O = 4 Fe(Oh) 3 .
2 Fe + 3 Cl 2 = 2 FeCl 3 .
Fe + H 2 SO 4 = FeSO 4 + H 2 .
Concentrated acids passivate iron in the cold, but dissolve when heated:
2Fe + 6H 2 SO 4 \u003d Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O.
Iron hydroxide (II) obtained by the action of alkali on salts of iron (II) without access to oxygen:
F 2 SO 4 + 2NaOH \u003d Fe (OH) 2 + Na 2 SO 4.
A white precipitate is formed, which quickly oxidizes in air:
4Fe(OH) 2 + O 2 + 2H 2 O = 4Fe(OH) 3 .
This hydroxide is amphoteric; when heated, it dissolves in alkalis with the formation of hexahydroferate:
Fe (OH) 3 + 3KOH \u003d K 3.
Iron forms two complex iron salts:
- yellow blood salt K 4 [ Fe(CN) 6 ];
- red blood salt K 3 [ Fe(CN) 6 ].
These compounds are qualitative for the determination of iron ions. Compound Prussian blue:
K 4 + Fe 2+ \u003d KFe III + 2K +.
The use of iron.
Iron is an essential component of the respiration process. It is part of the hemoglobin of the blood, is involved in the transfer of oxygen from the lungs to the tissues. In nature, iron is found in the composition of ores and minerals.
Iron is the eighth element of the fourth period in the periodic table. Its number in the table (also called atomic) is 26, which corresponds to the number of protons in the nucleus and electrons in the electron shell. It is designated by the first two letters of its Latin equivalent - Fe (lat. Ferrum - reads like "ferrum"). Iron is the second most common element in the earth's crust, the percentage is 4.65% (the most common is aluminum, Al). In its native form, this metal is quite rare, more often it is mined from mixed ore with nickel.
In contact with
What is the nature of this compound? Iron as an atom consists of a metal crystal lattice, which ensures the hardness of compounds containing this element and molecular stability. It is in connection with this that this metal is a typical solid body, unlike, for example, mercury.
Iron as a simple substance- silver-colored metal with properties typical for this group of elements: malleability, metallic luster and ductility. In addition, iron has a high reactivity. The latter property is evidenced by the fact that iron corrodes very quickly in the presence of high temperature and appropriate humidity. In pure oxygen, this metal burns well, and if it is crushed into very small particles, they will not only burn, but ignite spontaneously.
Often we call iron not a pure metal, but its alloys containing carbon ©, for example, steel (<2,14% C) и чугун (>2.14% C). Also of great industrial importance are alloys, to which alloying metals (nickel, manganese, chromium, and others) are added, due to which the steel becomes stainless, that is, alloyed. Thus, based on this, it becomes clear what an extensive industrial application this metal has.
Characteristic Fe
Chemical properties of iron
Let's take a closer look at the features of this element.
Properties of a simple substance
- Oxidation in air at high humidity (corrosive process):
4Fe + 3O2 + 6H2O \u003d 4Fe (OH) 3 - iron (III) hydroxide (hydroxide)
- Combustion of an iron wire in oxygen with the formation of a mixed oxide (it contains an element with both an oxidation state of +2 and an oxidation state of +3):
3Fe+2O2 = Fe3O4 (iron scale). The reaction is possible when heated to 160 ⁰C.
- Interaction with water at high temperature (600−700 ⁰C):
3Fe+4H2O = Fe3O4+4H2
- Reactions with non-metals:
a) Reaction with halogens (Important! With this interaction, it acquires the oxidation state of the element +3)
2Fe + 3Cl2 \u003d 2FeCl3 - ferric chloride
b) Reaction with sulfur (Important! In this interaction, the element has an oxidation state of +2)
Iron (III) sulfide - Fe2S3 can be obtained during another reaction:
Fe2O3+ 3H2S=Fe2S3+3H2O
c) Formation of pyrite
Fe + 2S \u003d FeS2 - pyrite. Pay attention to the degree of oxidation of the elements that make up this compound: Fe (+2), S (-1).
- Interaction with metal salts in the electrochemical series of metal activity to the right of Fe:
Fe + CuCl2 \u003d FeCl2 + Cu - iron (II) chloride
- Interaction with dilute acids (for example, hydrochloric and sulfuric):
Fe+HBr = FeBr2+H2
Fe+HCl = FeCl2+ H2
Note that these reactions produce iron with an oxidation state of +2.
- In undiluted acids, which are the strongest oxidizing agents, the reaction is possible only when heated; in cold acids, the metal is passivated:
Fe + H2SO4 (concentrated) = Fe2 (SO4) 3 + 3SO2 + 6H2O
Fe+6HNO3 = Fe(NO3)3+3NO2+3H2O
- The amphoteric properties of iron are manifested only when interacting with concentrated alkalis:
Fe + 2KOH + 2H2O \u003d K2 + H2 - potassium tetrahydroxyferrate (II) precipitates.
Iron making process in a blast furnace
- Roasting and subsequent decomposition of sulfide and carbonate ores (isolation of metal oxides):
FeS2 -> Fe2O3 (O2, 850 ⁰C, -SO2). This reaction is also the first step in the industrial synthesis of sulfuric acid.
FeCO3 -> Fe2O3 (O2, 550−600 ⁰C, -CO2).
- Burning coke (in excess):
С (coke) + O2 (air) —> CO2 (600−700 ⁰C)
CO2+С (coke) —> 2CO (750−1000 ⁰C)
- Recovery of ore containing oxide with carbon monoxide:
Fe2O3 —> Fe3O4 (CO, -CO2)
Fe3O4 —> FeO (CO, -CO2)
FeO —> Fe(CO, -CO2)
- Carburization of iron (up to 6.7%) and melting of cast iron (t⁰melting - 1145 ⁰C)
Fe (solid) + C (coke) -> cast iron. The reaction temperature is 900−1200 ⁰C.
In cast iron, cementite (Fe2C) and graphite are always present in the form of grains.
Characterization of compounds containing Fe
We will study the features of each connection separately.
Fe3O4
Mixed or double iron oxide, containing an element with an oxidation state of both +2 and +3. Also Fe3O4 is called iron oxide. This compound is resistant to high temperatures. Does not react with water, water vapor. Decomposed by mineral acids. Can be reduced with hydrogen or iron at high temperature. As you can understand from the above information, it is an intermediate product in the reaction chain of the industrial production of iron.
Directly iron oxide is used in the production of mineral-based paints, colored cement and ceramic products. Fe3O4 is what is obtained by blackening and bluing steel. A mixed oxide is obtained by burning iron in air (the reaction is given above). An ore containing oxides is magnetite.
Fe2O3
Iron(III) oxide, trivial name - hematite, red-brown compound. Resistant to high temperatures. In its pure form, it is not formed during the oxidation of iron with atmospheric oxygen. Does not react with water, forms hydrates that precipitate. Reacts poorly with dilute alkalis and acids. It can be alloyed with oxides of other metals, forming spinels - double oxides.
Red iron ore is used as a raw material in the industrial production of pig iron by the blast-furnace method. It also accelerates the reaction, that is, it is a catalyst in the ammonia industry. It is used in the same areas as iron oxide. Plus, it was used as a carrier of sound and pictures on magnetic tapes.
FeOH2
Iron(II) hydroxide, a compound that has both acidic and basic properties, the latter predominate, that is, it is amphoteric. A white substance that quickly oxidizes in air, "turns brown" to iron (III) hydroxide. Decomposes when exposed to temperature. It reacts with both weak solutions of acids and alkalis. We will not dissolve in water. In the reaction, it acts as a reducing agent. It is an intermediate product in the corrosion reaction.
Detection of Fe2+ and Fe3+ ions (“qualitative” reactions)
Recognition of Fe2+ and Fe3+ ions in aqueous solutions is carried out using complex complex compounds - K3, red blood salt, and K4, yellow blood salt, respectively. In both reactions, a precipitate of saturated blue color with the same quantitative composition, but a different position of iron with a valence of +2 and +3, is formed. This precipitate is also often referred to as Prussian blue or Turnbull blue.
Reaction written in ionic form
Fe2++K++3- K+1Fe+2
Fe3++K++4- K+1Fe+3
A good reagent for detecting Fe3+ is thiocyanate ion (NCS-)
Fe3++ NCS- 3- - these compounds have a bright red ("bloody") color.
This reagent, for example, potassium thiocyanate (formula - KNCS), allows you to determine even a negligible concentration of iron in solutions. So, he is able to determine if the pipes are rusty when examining tap water.
