Oxidative properties of iron. iron in nature. Chemical properties of these compounds

The first products made of iron and its alloys were found during excavations and date back to about the 4th millennium BC. That is, even the ancient Egyptians and Sumerians used meteorite deposits of this substance to make jewelry and household items, as well as weapons.

Today, various kinds of iron compounds, as well as pure metal, are the most common and used substances. No wonder the 20th century was considered iron. After all, before the advent and widespread use of plastic and related materials, it was this compound that was of decisive importance for humans. What is this element and what substances it forms, we will consider in this article.

Chemical element iron

If we consider the structure of the atom, then first of all we should indicate its location in the periodic system.

1. Ordinal number - 26.
2. The period is the fourth big one.
3. The eighth group, the secondary subgroup.
4. The atomic weight is 55.847.
5. The structure of the outer electron shell is denoted by the formula 3d 6 4s 2 .
6. - Fe.
7. The name is iron, the reading in the formula is "ferrum".
8. In nature, there are four stable isotopes of the element in question with mass numbers 54, 56, 57, 58.

The chemical element iron also has about 20 different isotopes that are not stable. The possible oxidation states that a given atom can exhibit are:

Not only the element itself is important, but also its various compounds and alloys.

Physical Properties

As a simple substance, iron has a pronounced metallicity. That is, it is a silvery-white metal with a gray tint, which has a high degree of ductility and ductility and a high melting and boiling point. If we consider the characteristics in more detail, then:

• melting point - 1539 0 С;
• boiling - 2862 0 С;
• activity - average;
• refractoriness - high;
• exhibits pronounced magnetic properties.

Depending on the conditions and different temperatures, there are several modifications that iron forms. Their physical properties differ from the fact that the crystal lattices differ.

All modifications have different types of structure of crystal lattices, and also differ in magnetic properties.

Chemical properties

As mentioned above, the simple substance iron exhibits medium chemical activity. However, in a finely dispersed state, it is capable of self-ignition in air, and the metal itself burns out in pure oxygen.

The corrosion ability is high, so the alloys of this substance are coated with alloying compounds. Iron is able to interact with:

• acids;
• oxygen (including air);
• gray;
• halogens;
• when heated - with nitrogen, phosphorus, carbon and silicon;
• with salts of less active metals, reducing them to simple substances;
• with sharp water vapor;
• with iron salts in the oxidation state +3.

It is obvious that, showing such activity, the metal is able to form various compounds, diverse and polar in properties. And so it happens. Iron and its compounds are extremely diverse and are used in various branches of science, technology, and industrial human activity.

Distribution in nature

Natural iron compounds are quite common, because it is the second most common element on our planet after aluminum. At the same time, in its pure form, the metal is extremely rare, as part of meteorites, which indicates its large accumulations in space. The main mass is contained in the composition of ores, rocks and minerals.

If we talk about the percentage of the element in question in nature, then the following figures can be given.

1. The cores of the terrestrial planets - 90%.
2. In the earth's crust - 5%.
3. In the Earth's mantle - 12%.
4. In the earth's core - 86%.
5. In river water - 2 mg/l.
6. In the sea and ocean - 0.02 mg / l.

The most common iron compounds form the following minerals:

• magnetite;
• limonite or brown iron ore;
• vivianite;
• pyrrhotite;
• pyrite;
• siderite;
• marcasite;
• lellingite;
• mispikel;
• milanterite and others.

This is still a long list, because there are really a lot of them. In addition, various alloys that are created by man are widespread. These are also such iron compounds, without which it is difficult to imagine the modern life of people. These include two main types:

• cast irons;
• become.

Iron is also a valuable addition to many nickel alloys.

Iron(II) compounds

These include those in which the oxidation state of the forming element is +2. They are quite numerous, because they include:

• oxide;
• hydroxide;
• binary compounds;
• complex salts;
• complex compounds.

The formulas of chemical compounds in which iron exhibits the indicated degree of oxidation are individual for each class. Consider the most important and common of them.

1. Iron(II) oxide. Black powder, insoluble in water. The nature of the connection is basic. It is able to quickly oxidize, however, it can also be easily reduced to a simple substance. It dissolves in acids to form the corresponding salts. Formula - FeO.
2. Iron(II) hydroxide. It is a white amorphous precipitate. Formed by the reaction of salts with bases (alkalis). It shows weak basic properties, is able to quickly oxidize in air to iron compounds +3. Formula - Fe (OH) 2.
3. The salts of an element in the specified oxidation state. As a rule, they have a pale green color of the solution, oxidize well even in air, acquiring and turning into iron salts 3. They dissolve in water. Examples of compounds: FeCL 2 , FeSO 4 , Fe(NO 3) 2 .

   

   Several compounds are of practical importance among the designated substances. First, (II). This is the main supplier of ions to the human body with anemia. When such an ailment is diagnosed in a patient, he is prescribed complex preparations, which are based on the compound in question. This is how iron deficiency in the body is replenished.

   Secondly, that is, iron (II) sulfate, together with copper, is used to destroy agricultural pests in crops. The method has been proving its effectiveness for more than a dozen years, therefore it is very much appreciated by gardeners and gardeners.

   Mora Salt

   This is a compound that is a crystalline hydrate of iron and ammonium sulfate. Its formula is written as FeSO 4 * (NH 4) 2 SO 4 * 6H 2 O. One of the iron (II) compounds, which has been widely used in practice. The main areas of human use are as follows.

   1. Pharmaceutics.
   2. Scientific research and laboratory titrimetric analyzes (to determine the content of chromium, potassium permanganate, vanadium).
   3. Medicine - as an additive to food with a lack of iron in the patient's body.
   4. For impregnation of wooden products, as Mora salt protects against decay processes.

   There are other areas in which this substance finds application. It got its name in honor of the German chemist who first discovered the manifested properties.

   Substances with an oxidation state of iron (III)

   The properties of iron compounds, in which it exhibits an oxidation state of +3, are somewhat different from those discussed above. Thus, the nature of the corresponding oxide and hydroxide is no longer basic, but pronounced amphoteric. We give a description of the main substances.

   
   

   Among the examples given, from a practical point of view, such a crystalline hydrate as FeCL 3 * 6H 2 O, or iron (III) chloride hexahydrate, is important. It is used in medicine to stop bleeding and replenish iron ions in the body with anemia.

   Iron(III) sulfate pentahydrate is used to purify drinking water, as it behaves as a coagulant.

   Iron(VI) compounds

   The formulas of the chemical compounds of iron, where it exhibits a special oxidation state of +6, can be written as follows:

   • K 2 FeO 4 ;
   • Na 2 FeO 4 ;
   • MgFeO 4 and others.

   All of them have a common name - ferrates - and have similar properties (strong reducing agents). They are also able to disinfect and have a bactericidal effect. This allows them to be used for the treatment of drinking water on an industrial scale.

   Complex compounds

   Special substances are very important in analytical chemistry and not only. Those that form in aqueous solutions of salts. These are complex compounds of iron. The most popular and well-studied of them are as follows.

   1. Potassium hexacyanoferrate (II) K 4 . Another name for the compound is yellow blood salt. It is used for qualitative determination of iron ion Fe 3+ in solution. As a result of exposure, the solution acquires a beautiful bright blue color, since another complex is formed - Prussian blue KFe 3+. Since ancient times it has been used as
   2. Potassium hexacyanoferrate (III) K 3 . Another name is red blood salt. It is used as a qualitative reagent for the determination of iron ions Fe 2+ . As a result, a blue precipitate is formed, which is called Turnbull blue. Also used as a dye for fabric.

   

   Iron in organic matter

   Iron and its compounds, as we have already seen, are of great practical importance in the economic life of man. However, in addition to this, its biological role in the body is no less great, on the contrary.

   There is one very important protein, which includes this element. This is hemoglobin. It is thanks to him that oxygen is transported and uniform and timely gas exchange is carried out. Therefore, the role of iron in the vital process - respiration - is simply enormous.

   

   In total, the human body contains about 4 grams of iron, which must be constantly replenished through the food consumed.

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, i.e., 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.

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 matter, 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. When in 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 pronounced 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 established. 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 crystal 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 weight) 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.

On the basis of iron, materials are created that can withstand 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 forms, 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, all the iron-containing compounds of the body are synthesized: the respiratory pigment hemoglobin is synthesized in the bone marrow, myoglobin is synthesized in the muscles, and 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.

Brazhnikova Alla Mikhailovna,

GBOU secondary school №332

Nevsky district of St. Petersburg

This manual considers questions on the topic "Chemistry of Iron". In addition to traditional theoretical issues, issues that go beyond the basic level are considered. It contains questions for self-control, which enable students to check the level of assimilation of the relevant educational material in preparation for the exam.

CHAPTER 1. IRON IS A SIMPLE SUBSTANCE.

The structure of the iron atom .

Iron is a d-element, located in a side subgroup of group VIII of the periodic system. The most common metal in nature after aluminum. It is part of many minerals: brown iron ore (hematite) Fe 2 O 3, magnetic iron ore (magnetite) Fe 3 O 4, pyrite FeS 2.

Electronic structure : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 .

Valence : II, III, (IV).

Oxidation states: 0, +2, +3, +6 (only in ferrates K 2 FeO 4).

physical properties.

Iron is a shiny, silvery-white metal, m.p. - 1539 0 C.

Receipt.

Pure iron can be obtained by reducing oxides with hydrogen when heated, as well as by electrolysis of solutions of its salts. Domain process - obtaining iron in the form of alloys with carbon (cast iron and steel):

1) 3Fe 2 O 3 + CO → 2Fe 3 O 4 + CO 2

2) Fe 3 O 4 + CO → 3FeO + CO 2

3) FeO + CO → Fe + CO 2

Chemical properties.

I. Interaction with simple substances - non-metals

1) With chlorine and sulfur (when heated). With a stronger oxidizing agent, chlorine oxidizes iron to Fe 3+, with a weaker one - sulfur - to Fe 2+:

2Fe 2 + 3Cl → 2FeCl 3

2) With coal, silicon and phosphorus (at high temperature).

3) In dry air, it is oxidized by oxygen, forming scale - a mixture of iron (II) and (III) oxides:

3Fe + 2O 2 → Fe 3 O 4 (FeO Fe 2 O 3)

II. Interaction with complex substances.

1) Corrosion (rusting) of iron proceeds in moist air:

4Fe + 3O 2 + 6H 2 O → 4Fe(OH) 3

At a high temperature (700 - 900 0 C) in the absence of oxygen, iron reacts with water vapor, displacing hydrogen from it:

3Fe+ 4H 2 O → Fe 3 O 4 + 4H 2

2) Displaces hydrogen from dilute hydrochloric and sulfuric acids:

Fe + 2HCl \u003d FeCl 2 + H 2

Fe + H 2 SO 4 (razb.) \u003d FeSO 4 + H 2

Highly concentrated sulfuric and nitric acids do not react with iron at ordinary temperatures due to its passivation.

With dilute nitric acid, iron is oxidized to Fe 3+, the products of HNO 3 reduction depend on its concentration and temperature:

8Fe + 30HNO 3 (very well dec.) → 8Fe(NO 3) 3 + 3NH 4 NO 3 + 9H 2 O

Fe + 4HNO 3 (diff.) → Fe (NO 3) 3 + NO + 2H 2 O

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

3) Reaction with solutions of metal salts to the right of iron in the electrochemical series of metal voltages:

Fe + CuSO 4 → FeSO 4 + Cu

CHAPTER2. IRON(II) COMPOUNDS.

Iron oxide(II) .

FeO oxide is a black powder, insoluble in water.

Receipt.

Recovery from iron oxide (III) at 500 0 C by the action of carbon monoxide (II):

Fe 2 O 3 + CO → 2FeO + CO 2

Chemical properties.

The main oxide, it corresponds to Fe (OH) 2 hydroxide: it dissolves in acids, forming iron (II) salts:

FeO+ 2HCl → FeCl 2 + H 2 O

Iron hydroxide (II).

Iron hydroxide Fe(OH) 2 is a water-insoluble base.

Receipt.

The action of alkalis on iron salts () without air access:

FeSO 4 + NaOH → Fe(OH) 2 ↓+ Na 2 SO 4

Chemical properties.

Hydroxide Fe(OH) 2 exhibits basic properties, dissolves well in mineral acids, forming salts.

Fe(OH) 2 + H 2 SO 4 → FeSO 4 + 2H 2 O

When heated, it decomposes:

Fe(OH) 2 → (temperature) FeO+ H 2 O

redox properties.

Iron (II) compounds exhibit sufficiently strong reducing properties, they are stable only in an inert atmosphere; in air (slowly) or in an aqueous solution under the action of oxidizing agents (quickly) they pass into iron (III) compounds:

4 Fe(OH) 2 (precipitate) + O 2 + 2H 2 O→ 4 Fe(OH) 3 ↓

2FeCl 2 + Cl 2 → 2FeCl 3

10FeSO 4 + 2KMnO 4 + 8H 2 SO 4 → 5 Fe 2 (SO 4) 3 + 2MnSO 4 + K 2 SO 4 + 8 H 2 O

Iron (II) compounds can also act as oxidizing agents:

FeO+ CO→ (temperature) Fe+ CO

CHAPTER 3. IRON COMPOUNDS (III).

Iron oxide(III)

Fe 2 O 3 oxide is the most stable natural oxygen-containing iron compound. It is an amphoteric oxide, insoluble in water. It is formed during the firing of pyrite FeS 2 (see 20.4 "Obtaining SO 2".

Chemical properties.

1) Dissolving in acids, it forms iron (III) salts:

Fe 2 O 3 + 6HCl → 2FeCl 3 + 3H 2 O

2) When fused with potassium carbonate, it forms potassium ferrite:

Fe 2 O 3 + K 2 CO 3 → (temperature) 2KFeO 2 + CO 2

3) Under the action of reducing agents, it acts as an oxidizing agent:

Fe 2 O 3 + 3H 2 → (temperature) 2Fe + 3H 2 O

Iron hydroxide (III)

Iron hydroxide Fe (OH) 3 is a red-brown substance, insoluble in water.

Receipt.

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

Chemical properties.

Fe (OH) 3 hydroxide is a weaker base than iron (II) hydroxide, has a weakly pronounced amphotericity.

1) Soluble in weak acids:

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

2) When boiled in 50% NaOH solution, it forms

Fe(OH) 3 + 3NaOH → Na 3

Iron salts (III).

Subject to strong hydrolysis in aqueous solution:

Fe 3+ + H 2 O ↔ Fe (OH) 2+ + H +

Fe 2 (SO 4) 3 + 2H 2 O ↔ Fe (OH) SO 4 + H 2 SO 4

Under the action of strong reducing agents in an aqueous solution, they exhibit oxidizing properties, turning into iron (II) salts:

2FeCl 3 + 2KI → 2FeCl 2 + I 2 + 2KCl

Fe 2 (SO 4) 3 + Fe → 3 Fe

CHAPTER4. QUALITATIVE REACTIONS.

Qualitative reactions to Fe 2+ and Fe 3+ ions.

1. The reagent for the Fe 2+ ion is potassium hexacyanoferrate (III) (red blood salt), which gives with it an intensely blue precipitate of a mixed salt - potassium-iron (II) hexacyanoferrate (III) or turnbull blue:

FeCl 2 + K 3 → KFe 2+ ↓ + 2KCl

1. The reagent for the Fe 3+ ion is the thiocyanate ion (thiocyanate ion) CNS -, when interacting with iron (III) salts, a blood-red substance is formed - iron (III) thiocyanate:

FeCl 3 + 3KCNS → Fe(CNS) 3 + 3KCl

3) Fe 3+ ions can also be detected using potassium hexacyanoferrate (II) (yellow blood salt). In this case, a water-insoluble substance of intense blue color is formed - potassium-iron (III) hexacyanoferrate (II) or Prussian blue:

FeCl 3 + K 4 → KFe 3+ ↓ + 3KCl

CHAPTER 5. MEDICAL AND BIOLOGICAL SIGNIFICANCE OF IRON.

The role of iron in the body.

Iron participates in the formation of hemoglobin in the blood, in the synthesis of thyroid hormones, in protecting the body from bacteria. It is necessary for the formation of immune protective cells, it is required for the "work" of B vitamins.

Iron is a part of more than 70 different enzymes, including respiratory ones, which ensure the processes of respiration in cells and tissues, and are involved in the neutralization of foreign substances entering the human body.

Hematopoiesis. Hemoglobin.

Gas exchange in the lungs and tissues.

Iron-deficiency anemia.

Iron deficiency in the body leads to diseases such as anemia, anemia.

Iron deficiency anemia (IDA) is a hematological syndrome characterized by impaired hemoglobin synthesis due to iron deficiency and manifested by anemia and sideropenia. The main causes of IDA are blood loss and lack of heme-rich food and drink.

The patient may be disturbed by fatigue, shortness of breath and palpitations, especially after physical exertion, often - dizziness and headaches, tinnitus, even fainting is possible. A person becomes irritable, sleep is disturbed, concentration of attention decreases. Because blood flow to the skin is reduced, increased sensitivity to cold may develop. There are also symptoms from the gastrointestinal tract - a sharp decrease in appetite, dyspeptic disorders (nausea, changes in the nature and frequency of stools).

Iron is an integral part of vital biological complexes, such as hemoglobin (oxygen and carbon dioxide transport), myoglobin (oxygen storage in muscles), cytochromes (enzymes). The body of an adult contains 4-5 g of iron.

LIST OF USED LITERATURE:

1. K.N. Zelenin, V.P. Sergutin, O.V. Malt "We pass the exam in chemistry perfectly." Elbl-SPb LLC, 2001.
2. K.A. Makarov "Medical chemistry". Publishing house of St. Petersburg State Medical University of St. Petersburg, 1996.
3. N.L. Glinka General Chemistry. Leningrad "Chemistry", 1985.
4. V.N. Doronkin, A.G. Berezhnaya, T.V. Sazhnev, V.A. Fevraleva "Chemistry. Thematic tests for preparing for the exam. Publishing house "Legion", Rostov-on-Don, 2012.

It has been known to people since antiquity: scientists attribute ancient household items made of this material to the 4th millennium BC.

Human life cannot be imagined without iron. It is believed that iron is used for industrial purposes more often than other metals. The most important structures are made from it. Iron is also found in small amounts in the blood. It is the content of the twenty-sixth element that colors the blood red.

Physical properties of iron

In oxygen, iron burns to form an oxide:

3Fe + 2O₂ = Fe₃O₄.

When heated, iron can react with non-metals:

Also, at a temperature of 700-900 ° C, it reacts with water vapor:

3Fe + 4H₂O = Fe₃O₄ + 4H₂.

Iron compounds

As you know, iron oxides have ions with two oxidation states: +2 and + 3. It is extremely important to know this, because completely different qualitative reactions will be carried out for different elements.

Qualitative reactions to iron

A qualitative reaction is needed in order to easily determine the presence of ions of one substance in solutions or impurities of another. Consider the qualitative reactions of ferrous and ferric iron.

Qualitative reactions for iron (III)

The content of ferric ions in a solution can be determined using alkali. With a positive result, a base is formed - iron (III) hydroxide Fe (OH) ₃.

Iron(III) hydroxide Fe(OH)₃

The resulting substance is insoluble in water and has a brown color. It is the brown precipitate that may indicate the presence of ferric ions in the solution:

FeCl₃ + 3NaOH = Fe(OH)₃↓+ 3NaCl.

Fe(III) ions can also be determined using K₃.

A solution of ferric chloride is mixed with a yellowish blood salt solution. As a result, you can see a beautiful bluish precipitate, which will indicate that ferric ions are present in the solution. you will find spectacular experiments on the study of the properties of iron.

Qualitative reactions for iron (II)

Fe²⁺ ions react with the red blood salt K₄. If a bluish precipitate forms when the salt is added, then these ions are present in the solution.