A side subgroup of the eighth group of the periodic table covers three triads of d-elements and three artificially obtained and little studied elements: hassium, Hs, meitnerium, Mt, darmstadtium Ds. The first triad is formed by the elements: iron, Fe, eobalt Co, nickel Ni; the second triad - ruthenium Ru, radium Ro, palladium Pd and the third triad - osmium Os, iridium Ir and platinum Pt. Artificially obtained hassium, matehrenium, darmstadtium with a short lifetime close the series of the heaviest elements known today.
Most of the Group VIIB elements under consideration have two valence electrons in the outer electron shell of the atom; they are all metals. In addition to external ns electrons, electrons from the penultimate electron shell (n-1)d take part in the formation of bonds.
Due to the increase in nuclear charge, the last element of each triad has a characteristic oxidation state lower than the first element. At the same time, an increase in the number of the period in which the element is located is accompanied by an increase in the characteristic degree of octlement (Table 9.1)
Table 9.1 Characteristic oxidation states of elements of the eighth secondary subgroup
The most common oxidation states of elements in their compounds are highlighted in Table. 41 in bold.
These elements are sometimes divided into three subgroups: the iron subgroup (Fe, Ru, Os), the cobalt subgroup (Co, Rh, Ir) and the nickel subgroup (Ni, Pd, Pt). This division is supported by the characteristic oxidation states of the elements (Table 42) and some other properties. For example, all elements of the iron subgroup are active catalysts for the synthesis of ammonia, and the nickel subgroup is active catalysts for the hydrogenation reactions of organic compounds. Elements of the cobalt subgroup are characterized by the formation of complex compounds [E(NH 3) 6 ]G 3, where G is a halogen ion
The redox properties of group VIIIB elements are determined by the following scheme:
Strengthening the oxidative properties of metal ions
All group VIIIB metals are catalytically active. All are more or less capable of absorbing hydrogen and activating it; they all form colored ions (compounds). All metals are prone to complex formation. A comparison of the physical and chemical properties of elements of subgroup VIII-B shows that Fe, Ni, Co are very similar to each other and at the same time very different from the elements of the other two triads, so they are classified into the iron family. The remaining six stable elements are united under a common name - the family of platinum metals.
Iron family metals
In the iron triad, the horizontal analogy, characteristic of d-elements in general, is most clearly manifested. The properties of the elements of the iron triad are given in table. 42.
Table 9.2 Properties of the elements of the iron triad
Natural resources. Iron is the fourth most abundant element in the earth's crust (after O 2 , Si, Al). It can be found in nature in a free state: it is iron of meteorite origin. Iron meteorites contain on average 90% Fe, 8.5% Ni, 0.5% Co. On average, there is one iron meteorite for every twenty stone meteorites. Sometimes native iron is found, carried out from the depths of the earth by molten magma.
To obtain iron, magnetic iron ore Fe 3 O 4 (magnetite mineral), red iron ore Fe 2 O 3 (hematite) and brown iron ore Fe 2 O 3 x H 2 O (limonite), FeS 2 - pyrite are used. In the human body, iron is present in hemoglobin.
Cobalt and nickel are found in the metallic state in meteorites. The most important minerals: cobaltine CoAsS (cobalt luster), iron-nickel pyrite (Fe,Ni) 9 S 8. These minerals are found in polymetallic ores.
Properties. Iron, cobalt, and nickel are silvery-white metals with grayish (Fe), pinkish (Co) and yellowish (Ni) tints. Pure metals are strong and ductile. All three metals are ferromagnetic. When heated to a certain temperature (Curie point), ferromagnetic properties disappear and metals become paramagnetic.
Iron and cobalt are characterized by polymorphism, while nickel is monomorphic and has an fcc structure up to the melting point.
The presence of impurities greatly reduces the resistance of these metals to aggressive atmospheres in the presence of moisture. This leads to the development of corrosion (iron rusting) due to the formation on the surface of a loose layer of a mixture of oxides and hydroxides of variable composition, which do not protect the surface from further destruction.
A comparison of the electrode potentials of the E 2+ /E systems for iron (-0.441 V), nickel (- 0.277 V) and cobalt (- 0.25 V), and the electrode potential of the Fe 3+ /Fe system (-0.036 V), shows that the most active element of this triad is iron. Dilute hydrochloric, sulfuric and nitric acids dissolve these metals to form E 2+ ions:
Fe + 2HC? =FeC? 2 +H 2 ;
Ni + H 2 SO 4 = NiSO 4 + H 2;
3Co + 8HNO 3 = 3Co(NO 3) 2 + 2NO + 4H 2 O;
4Fe + 10HNO 3 = 3Fe(NO 3) 2 + NH 4 No 3 + 3H 2 O.
More concentrated nitric acid and hot concentrated sulfuric acid (less than 70%) oxidize iron to Fe (III) with the formation of NO and SO2, for example:
Fe + 4HNO 3 = Fe(NO 3) 3 + No + 2H 2 O;
2Fe + 6H 2 SO 4 Fe 2 (SO 4) 3 + 3SO 2 +6H 2 O.
Very concentrated nitric acid (sp.v. 1.4) passivates iron, cobalt, nickel, forming oxide films on their surface.
Fe, Co, Ni are stable with respect to alkali solutions, but react with melts at high temperatures. All three metals do not react with water under normal conditions, but at a red-hot temperature, iron interacts with water vapor:
3Fe + 4H 2 o Fe 3 O 4 + 4H 2.
Cobalt and nickel are noticeably more resistant to corrosion than iron, which is consistent with their position in the series of standard electrode potentials.
Fine iron in oxygen burns when heated to form Fe 3 O 4, which is the most stable iron oxide and the same oxide forms cobalt. These oxides are derivatives of elements in oxidation states +2, +3 (EO E 2 O 3). Thermal oxidation of cobalt and nickel occurs at higher temperatures, resulting in the formation of NiO and CoO, which have a variable composition depending on the oxidation conditions.
For iron, nickel, cobalt, the oxides EO and E 2 O 3 are known (Table 9.3)
Table 9.3 Oxygen-containing compounds of elements of subgroup VIIIB
|
Item name |
Oxidation state |
Hydroxides |
|||||
|
Character |
Name |
Ion formula |
Name |
||||
|
Iron (Fe) |
Basic |
Iron(II) hydroxide |
Iron(II) salts |
||||
|
Amphoteric with a predominance of the main |
Iron(III) hydroxide |
Iron(III) salts |
|||||
|
Ferrous acid |
|||||||
|
Acid |
Iron acid |
||||||
|
Cobalt (Co) |
Basic |
Cobalt(II) hydroxide |
Cobalt(II) salts |
||||
|
Basic |
Cobalt(III) hydroxide |
Cobalt(III) salts |
|||||
|
Nickel (Ni) |
Basic |
Nickel(II) hydroxide |
Nickel(II) salts |
||||
|
Basic |
Nickel(III) hydroxide |
Nickel(III) salts |
Oxides EO and E 2 O 3 cannot be obtained in pure form by direct synthesis, since this produces a set of oxides, each of which is a phase of variable composition. They are obtained indirectly - by the decomposition of certain salts and hydroxides. Oxide E 2 O 3 is stable only for iron and is obtained by dehydrating the hydroxide.
EO oxides are insoluble in water and do not interact with it or with alkali solutions. The same is typical for the corresponding E(OH)2 hydroxides. E(OH)2 hydroxides easily react with acids to form salts. The acid-base properties of hydroxides of elements of the iron triad are given in Table. 42.
Iron (III) hydroxide Fe(OH) 3 is formed by the oxidation of Fe(OH) 2 with atmospheric oxygen:
4 Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3.
A similar reaction is typical for cobalt. Nickel (II) hydroxide is stable in relation to atmospheric oxygen. As a result, E(OH)3 hydroxides behave differently when interacting with acids. If Fe(OH) 3 forms iron (III) salts, then the reaction of Co(OH) 3 and Ni(OH) 3 with acids is accompanied by their reduction to E(+2):
Fe(OH) 3 + 3HC? =FeC? 3 + 3H 2 O;
2Ni(OH) 3 + 6HC? = 2NiC? 2+C? 2 + 6H 2 O.
Fe(OH)3 hydroxide also exhibits an acidic function, reacting with hot concentrated solutions of alkalis to form hydroxo complexes, for example, Na3. Derivatives of ferrous acid HFeO 2 (ferrites) are obtained by fusing alkalis or carbonates with Fe 2 O 3:
2NaOH + Fe 2 O 3 2NaFeO 2 + H 2 O;
MgCO 3 + Fe 2 O 3 MgFe 2 O 4 + CO 2.
Ferrites Me II Fe 2 O 4 belong to the spinel class. The oxides Fe 3 O 4 and Co 3 O 4 discussed above are formally spinels FeFe 2 O 4 and CoCo 2 O 4 .
Unlike cobalt and nickel, iron compounds are known in which its oxidation state is + 6. Ferrates are formed by the oxidation of Fe(OH) 3 in hot concentrated alkali in the presence of an oxidizing agent:
2Fe +3 (OH) 3 + 10KOH + 3Br 2 = 2K 2 Fe +6 O 4 + 6KBr + 2H 2 O.
Ferrates are thermally unstable and with slight heating (100-2000C) they turn into ferrites:
4K 2 FeO 4 4KfeO 2 + 2K 2 O + 3O 2.
In the free state, iron acid and its corresponding oxide FeO 3 are not isolated. In solubility and structural terms, ferrates are close to the corresponding chromates and sulfates. Potassium ferrate is formed by fusing Fe 2 O 3 with KNO 3 and KOH:
Fe 2 O 3 + 3KNO 3 + 4KOH = 2K 2 feO 4 + 3KNO 2 + 2H 2 O.
Ferrates are red-violet crystalline substances. When heated they decompose. The acid H 2 FeO 4 cannot be isolated; it instantly decomposes into Fe 2 O 3, H 2 O and O 2. Ferrates are strong oxidizing agents. In acidic and neutral environments, ferrates decompose, oxidizing water:
2Na 2 FeO 4 + 10 H 2 O 4Fe(OH) 3 + 4NaOH + O 2.
Compounds with non-metals. Fe, Ni, Co halides are relatively few in number and correspond to the most characteristic oxidation states +2 and +3. For iron, the halides FeG 2 and FeG 3 with fluorine, chlorine and bromine are known. During direct interaction, FeF 3, FeC? 3, FeBr 3. Dihalides are obtained indirectly by dissolving the metal (or its oxide) in the corresponding hydrohalic acid. Trifluoride CoF 3 and trichloride CoC were obtained for cobalt? 3. Nickel does not form trihalides. All dihalides of the iron triad are typical salt-like compounds with a noticeable ionic contribution to the chemical bond.
Iron, cobalt, nickel energetically interact with chalcogens and form chalcogenides: EC and EC 2. Monochalcogenides can be obtained by reacting the corresponding components in solutions:
CoC? 2 + (NH 4) 2 S = CoS + 2NH 4 C?.
All chalcogenides are phases of variable composition.
Compounds of metals of the iron triad with other nonmetals (pnictogens, carbon, silicon, boron) differ markedly from those discussed above. All of them do not obey the rules of formal valence and most of them have metallic properties.
Iron, cobalt, and nickel absorb hydrogen, but do not produce certain compounds with it. When heated, the solubility of hydrogen in metals increases. Hydrogen dissolved in them is in an atomic state.
Salts of oxygen-containing acids and complex compounds. All salts of hydrochloric, sulfuric and nitric acids are soluble in water.
Nickel (II) salts are green, cobalt (II) are blue, and their solutions and crystalline hydrates are pink (for example), iron (II) salts are greenish, and iron (III) are brown. The most important salts are: FeC? 3 6H 2 O; FeSO 4 7H 2 O - iron sulfate, (NH 4) 2 SO 4 FeSO 4 6H 2 O - Mohr's salt; NH 4 Fe(SO 4) 2 12H 2 O - ferroammonium alum; NiSO 4 6H 2 O, etc.
The ability of iron, cobalt and nickel salts to form crystalline hydrates indicates the tendency of these elements to form complexes. Crystal hydrates are a typical example of aqua complexes:
[E(H 2 O) 6 ](ClO 4) 2; [E(H 2 O) 6 ](NO 3) 2.
Anionic complexes are numerous for the elements of the iron triad: halide (Me I (EF 3), Me 2 I [EG 4], Me 3 [EG 4], etc.), thiocyanate (Me 2 I [E (CNS) 4] , Me 4 I [E(CNS) 6 ], Me 3 I [E(CNS) 6 ]), oxolate (Me 2 I [E(C 2 O 4) 2 ], Me 3 [E(C 2 O 4) 3 ]). Cyanide complexes are especially characteristic and stable: K 4 - potassium hexacyanoferrate (II) (yellow blood salt) and K 3 - potassium hexacyanoferrate (III) (red blood salt). These salts are good reagents for the detection of Fe+3 ions (yellow salt) and Fe2+ ions (red salt) at pH ??7:
4Fe 3+ + 4- = Fe 4 3;
Prussian blue
3Fe 2+ + 2 3- = Fe 3 2.
Turnbull blue
Prussian blue is used as a blue dye. When thiocyanate salts KCNS are added to a solution containing Fe 3+ ions, the solution turns blood red due to the formation of iron thiocyanate:
FeC? 3 + 3KCNS = Fe(CNS) 3 + 3KC?.
This reaction is very sensitive and is used to discover the Fe 3+ ion.
Cobalt (II) is characterized by stable simple salts and unstable complex compounds K2, K4, which transform into cobalt (III) compounds: K3, C? 3.
Characteristic complex compounds of iron, iron, cobalt and nickel are carbonyls. Similar compounds were discussed earlier for elements of the chromium and manganese subgroups. However, the most typical among carbonyls are: , , . Iron and nickel carbonyls are obtained in the form of liquids at normal pressure and 20-60 o C by passing a CO stream over metal powders. Cobalt carbonyl is obtained at 150-200 o C and a pressure of (2-3) 10 7 Pa. These are orange crystals. In addition, there are carbonyls of a more complex composition: Fe(CO) 9 and trinuclear carbonyls, which are cluster-type compounds.
All carbonyls are diamagnetic, since CO ligands (like CN?) create a strong field, as a result of which the valence d-electrons of the complexing agent form p-bonds with CO molecules according to the donor-acceptor mechanism. y-Bonds are formed due to lone electron pairs of CO molecules and the remaining vacant orbitals of the complexing agent:
Nickel (II), on the contrary, forms many stable complex compounds: (OH) 2, K 2; The 2+ ion is dark blue.
This reaction is widely used in qualitative and quantitative analysis for the determination of nickel. Nickel and especially cobalt compounds are poisonous.
Application. Iron and its alloys form the basis of modern technology. Nickel and cobalt are important alloying additives in steels. Heat-resistant nickel-based alloys (nichrome containing Ni and Cr, etc.) are widely used. Coins, jewelry, and household items are made from copper-nickel alloys (cupronickel, etc.). Many other nickel- and cobalt-containing alloys are of great practical importance. In particular, cobalt is used as a viscous component of the materials from which metal-cutting tools are made, in which particles of exclusively hard carbides MoC and WC are embedded. Galvanic nickel coatings of metals protect them from corrosion and give them a beautiful appearance.
Metals of the iron family and their compounds are widely used as catalysts. Sponge iron with additives is a catalyst for ammonia synthesis. Highly dispersed nickel (Raney nickel) is a very active catalyst for the hydrogenation of organic compounds, in particular fats. Raney nickel is obtained by reacting an alkali solution with the intermetallic compound NiA?, while aluminum forms a soluble aluminate, and nickel remains in the form of tiny particles. This catalyst is stored under a layer of organic liquid, since in a dry state it is instantly oxidized by atmospheric oxygen. Cobalt and manganese are part of the catalyst added to oil paints to speed up their "drying".
Fe 2 O 3 oxide and its derivatives (ferrites) are widely used in radio electronics as magnetic materials.
Located in the fourth period.
The atomic weight of iron is 55.84, nuclear charge +26. Distribution of electrons by energy levels (+26): 2, 8, 14, 2. Electronic configuration of the outer and pre-outer layer of iron 3s23p63d64s2.
Thus, the iron atom, in addition to two s-electrons of the fourth outer layer, there are six more d-electrons of the third pre-outer layer. Of these d-electrons are the most active 4 unpaired ones. Consequently, 6 electrons are especially actively involved in the formation of iron valence bonds - 2 from the outer and 4 from the pre-outer layers. The most common oxidation states of iron are Fe+2 and Fe+3. Iron is one of the most commonly found elements in nature. It ranks fourth in prevalence among other elements.
■ 57. Based on the structure of the iron atom, as well as the distribution of electrons in orbitals, indicate the possible oxidation states of this element.
Iron in the free state is a silvery-gray shiny metal with a density of 7.87, a melting point of 1535° and a boiling point of 2740°. Iron has pronounced ferromagnetic properties, i.e., under the influence of a magnetic field it becomes magnetized and when the field stops, it retains magnetic properties, becoming a magnet itself. All elements of the iron group have these properties.
In terms of its chemical properties, iron is a very active metal. In the absence of moisture, iron does not change in the air, but when exposed to moisture and oxygen in the air, it undergoes severe corrosion and becomes covered with a loose film of rust, which is iron, which does not protect it from further oxidation, and the iron gradually oxidizes in its entire mass:
4Fe + 2H2O + 3O2 = 2Fe2O3 2H2O
A number of methods have been developed to protect this valuable metal from corrosion.
In the voltage series, iron is located to the left of hydrogen. In this regard, it is easily exposed to dilute acids, turning into a ferrous iron salt, for example:
Fe + 2HCl = FeCl2 + H2
Iron does not react with concentrated sulfuric and nitric acids. These acids create such a strong and dense film of oxide on the surface of the metal that the metal becomes completely passive and no longer enters into other reactions. At the same time, when directly interacting with such strong oxidizing agents as iron, iron always exhibits an oxidation state of +3:
2Fe + 3Сl2 = 2FeCl3
Iron reacts with superheated steam; in this case, is displaced from the water, and the hot iron turns into oxide, and this is always either ferrous oxide FeO or iron oxide Fe3O4(Fe2O3 FeO):
Fe + H2O = FeO + H2
3Fe + 4H2O = Fe3O4 + 4H2
Iron heated in pure oxygen burns vigorously to form iron scale (see Fig. 40).
3Fe + 2O2 = Fe3O4
When calcined, iron forms an alloy with carbon and at the same time iron carbide Fe3C.
■ 58. List the physical properties of iron.
59. What are the chemical properties of iron? Give a reasoned answer.
Iron compounds
Iron forms two series of compounds - compounds Fe +2 and Fe +3. Iron is characterized by two oxides - oxide FeO and oxide Fe2O3. True, the mixed oxide Fe3O4 is known, the molecule of which is di- and trivalent iron: Fe2O3 · FeO. This oxide is also called iron scale, or iron oxide.
Ferrous iron compounds are less stable than iron oxide compounds, and in the presence of an oxidizing agent, even if it is only air, they usually turn into ferric iron compounds. For example, iron (II) hydroxide Fe(OH)2 is a white solid, but it can be obtained in pure form only when the solutions of the reacting substances do not contain dissolved oxygen and if the reaction is carried out in the absence of atmospheric oxygen:
FeSO4 + 2NaOH = Fe(OH)2 + Na2SO4
The salt from which iron (II) hydroxide is obtained, of course, should not contain the slightest admixture of oxide compounds. Since such conditions are very difficult to create in an ordinary educational laboratory, iron (II) hydroxide is obtained in the form of a more or less dark green precipitate of a gelatinous appearance, which indicates the oxidation of divalent iron compounds into ferric iron. If iron (II) hydroxide is kept in air for a long time, it gradually transforms into iron (III) hydroxide Fe(OH)3:
4Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3
iron are typical insoluble hydroxides. Iron (II) hydroxide has basic properties, while Fe(OH)3 has very weakly expressed amphoteric properties.
■ 60. List the properties of iron oxide as a typical basic oxide. Give a reasoned answer. Write all reaction equations in full and abbreviated ionic forms.
61. List the properties of iron (II) hydroxide. Support your answer with reaction equations.
Among iron (II) salts, the most important is iron sulfate FeSO4 · 7H2O, which contains 7 molecules of water of crystallization. Iron sulfate dissolves well in water. It is used to control agricultural pests, as well as in the manufacture of dyes.
Of the trivalent iron salts, the most important is ferric chloride FeCl3, which is very hygroscopic orange crystals that absorb water during storage and dissolve into a brown paste.
Iron (II) salts can easily transform into iron (III) salts, for example when heated with nitric acid or potassium permanganate in the presence of sulfuric acid:
6FeSO4 + 2HNO3 + 3H2SO4 = 3Fe2(SO4)3 + 2NO + 4H2O
Oxidation of Fe +2 salts into Fe +3 salts can also occur under the influence of atmospheric oxygen during storage of these compounds, but this process is longer. Very characteristic specific reagents are used to recognize Fe 2+ and Fe 3+ cations. For example, to recognize divalent iron, take the red blood salt K3, which, in the presence of divalent iron ions, gives with them a characteristic intense blue precipitate of Turnboule blue:
3FeSO4 + 2K3 = Fe32 + 3K2SO4
or in ionic form
3Fe 2+ + 2 3- = Fe32
To recognize Fe3+ salts, a reaction with yellow blood salt K4 is used:
4FeCl3 + 3K4 = Fe43 + 12KCl
4Fe 3+ + 3 4- = Fe43
In this case, an intense blue precipitate of Prussian blue appears. Prussian blue and Turnboule blue are used as dyes.
In addition, ferric iron can be recognized using soluble salts - potassium thiocyanate KCNS or ammonium thiocyanate NH4CNS. When these substances interact with Fe(III) salts, the solution acquires a blood-red color.
■ 62. List the properties of Fe +3 and Fe +2 salts. Which oxidation state is more stable?
63. How to convert Fe +2 salt into Fe +3 salt and vice versa? Give examples.
The reaction follows the equation:
FeCl3 + 3KCNS = Fe(CNS)3 + 3КCl
or in ionic form
Fe 3+ + 3CNS - = Fe(CNS),
Iron compounds play an important role in the life of organisms. For example, it is part of the main blood protein - hemoglobin, as well as green plants - chlorophyll. Iron enters the body mainly as part of organic matter in food products. Apples, eggs, spinach, and beets contain a lot of iron. As medicines, iron is used in the form of salts of organic acids. Ferric chloride serves as a hemostatic agent.
■ 64. Three test tubes contain: a) iron (II) sulfate, b) iron (III) sulfate and c) iron (III) chloride. How to determine which test tube contains which salt?
65. How to carry out a series of transformations:
Fe → FeCl2 → FeSO4 → Fe2(SO4)3 → Fe(OH)3 → Fe2O3.
66. The following are given: iron, caustic soda. How, using only these substances, can one obtain iron (II) hydroxide and iron (III) hydroxide?
67. A solution containing chromium (III) chloride and iron (III) chloride was treated with excess alkali. The resulting precipitate was filtered. What remained on the filter and what went into the filtrate? Give a reasoned answer using reaction equations in molecular, full ionic, and reduced ionic forms.
Iron alloys
Iron is the basis of ferrous metallurgy, so it is mined in huge quantities. The new program for the extensive construction of communism provides for the production of 250 million tons of steel in 1980. This is 3.8 times more than in 1960.
Iron is almost never used in its pure form, but only in the form of alloys. The most important alloys of iron are its with carbon - various cast irons and steels. The main difference between cast iron and steel is the carbon content: cast iron contains more than 1.7% carbon, and steel contains less than 1.7%.
Ferroalloys (an alloy of iron and silicon), ferrochrome (an alloy of iron and chromium), and ferromanganese (an alloy of iron and manganese) are of great practical importance. Ferroalloys are cast irons containing more than 10% iron and at least 10% of the corresponding component. In addition, they contain the same elements as cast iron. Ferroalloys are used mainly in the “deoxidation” of steel and as alloying impurities.
Among cast irons, a distinction is made between linear and pigment. Foundry cast iron is used for castings of various parts; pig iron is remelted into steel, as it has very high hardness and cannot be processed. Pipe iron is white, and foundry iron is gray. Pig iron contains more manganese.
Steels are carbon and alloyed. Carbon steels are usually an alloy of iron and carbon, while alloy steels contain alloying additives, i.e., admixtures of other metals that give the steel more valuable properties. gives steel ductility, elasticity, stability during hardening, and - hardness and heat resistance. Steels with zirconium additives are very elastic and ductile; they are used to make armor plates. Manganese impurities make steel resistant to impact and friction. Boron improves the cutting properties of steel in the manufacture of tool steels.
Sometimes even minor impurities of rare metals impart new properties to steel. If you keep a steel part in beryllium powder at a temperature of 900-1000°, the hardness of the steel and its wear resistance are greatly increased.
Chromium-nickel steel, or, as they are also called, stainless steel, is resistant to corrosion. Impurities of sulfur and phosphorus are very harmful to steel - they make the metal brittle.
■ 68. What important glands do you know?
69. What is the main difference between steel and cast iron?
70. What properties of cast iron and what types of cast iron do you know?
71. What are alloy steels and alloying additives?
Domain process
Cast iron is obtained by reduction smelting in blast furnaces. These are huge structures thirty meters high, producing more than 2000 tons of cast iron per day. A diagram of the blast furnace structure is shown in Fig. 83.
The upper part of the blast furnace, through which the charge is loaded, is called the top. Through the furnace the charge
Rice. 83. Scheme of a blast furnace.
falls into a long furnace shaft that widens downwards, which facilitates the movement of the loaded material from top to bottom. As the charge moves to the widest part of the furnace - the steam - a series of transformations occur with it, as a result of which cast iron is formed, flowing into the hearth - the hottest part of the furnace. This is where the slag collects. Pig iron and slag are discharged from the furnace through special holes in the forge, called tapholes. Air is blown into the blast furnace through the top of the furnace to keep the fuel burning in the furnace.
Let us consider the chemical processes that occur during the smelting of cast iron. The blast furnace charge, i.e., the complex of substances loaded into it, consists of iron ore, fuel and fluxes, or fluxes. There are many iron ores. The main ores are magnetic iron ore Fe3O4, red iron ore Fe2O3, brown iron ore 2Fe2O8 · 3H2O. In the blast furnace process, siderite FeCO3 and sometimes FeS2 are used as iron ore, which, after firing in pyrite furnaces, turns into cinder Fe2O3, which can be used in metallurgy. Such ore is less desirable due to its high sulfur content. Not only cast iron, but also ferroalloys are smelted in a blast furnace. The fuel loaded into the furnace serves both to maintain a high temperature in the furnace and to reduce iron from ore, and also takes part in the formation of an alloy with carbon. The fuel is usually coke.
During the smelting process of iron, coke is gasified, turning, as in a gas generator, first into dioxide and then into carbon monoxide:
C + O2 = CO3 CO2 + C = 2CO
The resulting carbon monoxide is a good gaseous reducing agent. With its help, iron ore is recovered:
Fe2O3 + 3СО = 3СО2 + 2Fe
Along with the ore containing iron, waste rock impurities necessarily enter the furnace. They can be very refractory and can clog a furnace that has been operating continuously for many years. In order for the waste rock to be easily removed from the furnace, it is converted into a low-melting compound, turning it into slag using fluxes (fluxes). To convert base rock containing, for example, limestone into slag, which decomposes in a furnace according to the equation
CaCO3 = CaO + CO2
add sand. Fusion with calcium oxide, sand forms silicate:
CaO + SiO3 = CaSiO3
This is a substance with an incomparably lower melting point. In a liquid state, it can be released from the oven.
If the rock is acidic, containing a large amount of silicon dioxide, then, on the contrary, limestone is loaded into the furnace, which converts the silicon dioxide into silicate, and the result is the same slag. Previously, slag was a waste, but now it is cooled with water and used as a building material.
To maintain fuel combustion, heated, oxygen-enriched air is continuously supplied to the blast furnace. It is heated in special air heaters - kiupers. Cowper is a high tower made of refractory bricks, where hot gases escaping from the blast furnace are diverted. Blast furnace gases contain carbon dioxide CO2, N2 and carbon monoxide CO. Carbon monoxide burns in the cowper, thereby increasing its temperature. Then the blast furnace gases are automatically sent to another cowper, and through the first one begins blowing air directed into the blast furnace. In a hot cowper, the air is heated, and thus fuel is saved, which in large quantities would be spent on heating the air entering the blast furnace. Each blast furnace has several cowpers.
■ 72. What is the composition of the blast furnace charge?
73. List the main chemical processes that occur during the smelting of cast iron.
74. What is the composition of blast furnace gas and how is it used in cowpers?
75. How much cast iron containing 4% carbon can be obtained from 519.1 kg of magnetic iron ore containing 10% impurities?
76. What amount of coke gives a volume of carbon monoxide sufficient to reduce 320 kg of iron oxide if the coke contains 97% pure carbon?
77. How should siderite be processed so that iron can be obtained from them?
Steelmaking
Steel is smelted in three types of furnaces - open-hearth regenerative furnaces, Bessemer converters and electric furnaces.
The open hearth furnace is the most modern furnace designed for smelting the bulk of steel (Fig. 84). An open hearth furnace, unlike a blast furnace, is not a continuously operating furnace.
Rice. 84. Diagram of an open-hearth furnace
Its main part is the bathtub, into which the necessary materials are loaded through the windows using a special machine. The bath is connected by special passages to regenerators, which serve to heat combustible gases and air supplied to the furnace. Heating occurs due to the heat of combustion products, which are passed through regenerators from time to time. Since there are several of them, they work in turn and heat up in turn. An open hearth furnace can produce up to 500 tons of steel per melt.
The charge of an open-hearth furnace is very diverse: the charge includes cast iron, scrap metal, ore, fluxes (fluxes) of the same nature as in the blast furnace process. As in the blast furnace process, during steel smelting, air and combustible gases are heated in regenerators using the heat of waste gases. The fuel in open-hearth furnaces is either fuel oil sprayed by nozzles or combustible gases, which are currently used especially widely. The fuel here serves only to maintain a high temperature in the furnace.
The process of steel smelting is fundamentally different from the blast furnace process, since the blast furnace process is a reducing process, and steel smelting is an oxidative process, the purpose of which is to reduce the carbon content by oxidizing it in the metal mass. The processes that take place are quite complex.
Contained in the ore and supplied with air to the furnace for burning gaseous fuel, it oxidizes, as well as a significant amount of iron, converting it mainly into iron (II) oxide: 2Fe + O2 = 2FeO
Contained in cast iron, or any impurities of other metals at high temperatures reduce the resulting iron (II) oxide again to metallic iron according to the equation: Si + 2FeO = SiO2 + 2Fe Mn + FeO = MnO + Fe
Reacts similarly with iron (II) oxide and: C + FeO = Fe + CO
At the end of the process, “deoxidizers” - ferroalloys - are added to restore the remaining iron (II) oxide (or, as they say, to “deoxidize” it). The additives of manganese and silicon present in them reduce the remaining iron (II) oxide according to the above equations. After this, the melting ends. Melting in open hearth furnaces lasts 8-10 hours.
Rice. 85. Bessemer converter design diagram
The Bessemer converter (Fig. 85) is an older type furnace, but with very high productivity. Since the converter operates without fuel consumption, this method of steel production occupies a significant place in metallurgy. The converter is a pear-shaped steel vessel with a capacity of 20-30 tons, lined on the inside with refractory bricks. Each melting in the converter lasts 12-15 minutes. The converter has a number of disadvantages: it can only operate on liquid cast iron. This is due to the fact that carbon oxidation is carried out by air passed from below through the entire mass of liquid cast iron, which significantly speeds up melting and increases the intensity of oxidation. Naturally, the “waste” of iron in this case is especially great. At the same time, the short melting time does not allow it to be regulated or alloyed to be added, so mainly carbon steels are smelted in converters. At the end of the melting, the air supply is stopped and, as in the open-hearth process, “deoxidizers” are added.
In electric furnaces (Fig. 86) alloy steel of special grades is smelted, mainly with a high melting point, containing and other additives. The finished steel is sent to rolling. There, on huge rolling mills - blooming and slab mills - hot steel ingots are compressed using rolls, which make it possible to produce various shapes from the steel ingot.
Figure 86. Diagram of an electric arc furnace. 1 - electrodes, 2 - loading window, 3 - chute for steel release, 4 - rotary mechanism
Iron in the form of alloys is widely used in the national economy. Not a single sector of the national economy can do without it. In order to save ferrous metals, currently, whenever possible, they are trying to replace them with synthetic materials.
Ferrous metals are used to make machine tools and cars, airplanes and tools, reinforcement for reinforced concrete structures, tin for cans and roofing sheets, ships and bridges, agricultural machines and beams, pipes and a whole range of household products.
■ 78. What is the fundamental difference between the steel smelting process and the blast furnace process?
79. What furnaces are used to smelt steel?
80. What are regenerators in an open hearth furnace?
81. Indicate the composition of the open-hearth furnace charge and its difference from the composition of the blast furnace charge?
82. What are “deoxidizers”?
83. Why is steel smelting called oxidative smelting?
84. How much steel containing 1% carbon can be produced from 116.7 kg of cast iron containing 4% carbon?
85. How much ferromanganese containing 80% manganese is needed to “deoxidize” 36 kg of ferrous oxide?
Article on the topic Iron, a secondary subgroup of group VIII
IRON AND ELECTRICITY The properties of steels are varied. There are steels designed to last long in sea water, steels that can withstand high temperatures and...
The side subgroup of the eighth group covers three triads of d-elements.
The first triad is formed by the elements iron, cobalt and nickel, second – ruthenium, rhodium, palladium, and the third triad - osmium, iridium and platinum.
Most elements of the subgroup under consideration have two electrons in the outer electron shell of the atom; they are all metals.
In addition to outer electrons, electrons from the previous unfinished electron shell also take part in the formation of chemical bonds.
The iron family includes iron, cobalt and nickel. The increase in electronegativity in the series Fe (1.83) – Co (1.88) – Ni (1.91) shows that from iron to nickel there should be a decrease in basic and reducing properties. In the electrochemical voltage series, these elements come before hydrogen.
In terms of its prevalence in nature, the use of compounds in medicine and technology, and its role in the body, iron ranks first in this group.
Elements of the iron family in compounds exhibit oxidation states +2,
Iron(II) compounds. Ferrous salts are formed when iron dissolves in dilute acids. The most important of them is iron (II) sulfate, or ferrous sulfate, FeSO 4 . 7H 2 O, forming light green
crystals, highly soluble in water. In air, iron sulfate gradually erodes and at the same time oxidizes from the surface, turning into a yellow-brown basic salt of iron (III).
Iron(II) sulfate is prepared by dissolving steel scraps in 20-30% sulfuric acid:
Fe + H 2 SO 4 = FeSO 4 + H 2
Iron (II) sulfate is used to control plant pests, in the production of inks and mineral paints, and in textile dyeing. When a solution of an iron (II) salt reacts with an alkali, a white precipitate of iron (II) hydroxide Fe(OH) 2 precipitates, which in air due to oxidation quickly takes on a greenish and then brown color, turning into iron (III) hydroxide Fe(OH) 3 :
4Fe(OH) 2 + O 2 + 2H 2 O = 4Fe(OH) 3
Divalent iron compounds are reducing agents and can easily be converted to ferric iron compounds:
6FeSO 4 + 2HNO 3 + 3H 2 SO 4 = 3Fe 2 (SO 4) 3 + 2NO + 4H 2 O
10FeSO 4 + 2KMnO 4 + 8H 2 SO 4 = 5Fe 2 (SO 4) 3 + K 2 SO 4 + 2MnSO 4 + 8H 2 O
Ferric oxide and hydroxide have amphoteric properties. Iron (III) hydroxide is a weaker base than iron (II) hydroxide, this is expressed in the fact that ferric iron salts are strongly hydrolyzed, and Fe(OH) 3 does not form salts with weak acids (for example, carbonic acid, hydrogen sulfide).
The acidic properties of ferric iron oxide and hydroxide are manifested in the fusion reaction with alkali metal carbonates, as a result of which ferrites are formed - salts of ferrous acid HFeO 2 not obtained in a free state:
Fe 2 O 3 + Na 2 CO 3 = 2NaFeO 2 + CO
If you heat steel filings or iron (III) oxide with potassium nitrate and hydroxide, an alloy is formed containing potassium ferrate K 2 FeO 4 - a salt of iron acid H 2 FeO 4 not released in the free state:
Fe 2 O 3 + 4KOH + 3KNO 3 = 2K 2 FeO 4 + 3KNO 2 + 2H 2 O
In biogenic compounds, iron is complexed with organic ligands (myoglobin, hemoglobin). The degree of iron oxidation in these complexes is debated. Some authors believe that the oxidation state is +2, others suggest that it varies from +2 to +3 depending on the degree of interaction with oxygen.
Application
Dissociation constants of some acids and bases /at 25 0 C/
| Compound | K 1 | K 2 | K 3 |
| HF | 6,8 . 10 -4 | ||
| HClO | 5,0 . 10 -8 | ||
| HBrO | 2,5 . 10 -9 | ||
| H2S | 9,5 . 10 -8 | 1.0 . 10 -14 | |
| H2SO3 | 1,7 . 10 -2 | 6,2 . 10 -8 | |
| HNO2 | 5,1 . 10 -4 | ||
| H3PO4 | 7,6 . 10 -3 | 6,2 . 10 -8 | 4,2 . 10 -13 |
| H2CO3 | 4,5 . 10 -7 | 4,8 . 10 -11 | |
| CH3COOH | 1,8 . 10 -5 | ||
| HCN | 6,2 . 10 -10 | ||
| NH4OH | 1,8 . 10 -5 |
These elements: helium ( Not), neon ( Ne), argon ( Ar), krypton ( Kr), xenon
(Heh) and radon ( Rn) are called inert gases because they have very low chemical reactivity. The outer energy level of helium has two electrons, while the other elements have eight electrons each, which corresponds to an energetically favorable electron configuration.
Neon and argon are used to fill incandescent lamps. Argon welding of stainless steels, titanium, aluminum and aluminum alloys produces exceptionally clean and strong welds.
Krypton, xenon and radon are capable of combining with other elements and, above all, with fluorine. These compounds (XeF 2, XeF 6, XeO 3, etc.) have strong oxidizing properties. Radon is a radioactive element with a half-life of 3.8 days. However, in nature it is constantly formed. In terms of molar mass, it is 7.65 times heavier than air, so it accumulates in basement, unventilated areas. During the day, the concentration of radon in an unventilated room increases 6 times, and when using a shower in the bathroom, 40 times. Most of the radiation a person receives comes from the radioactive decay of radon.
18 Complex connections
Complex connections – these are compounds containing a complex ion capable of independent existence in solution. The large number of possible complexing agents and ligands, as well as the phenomenon of isomerism, lead to the diversity of these compounds.
18.1 Composition of complex compounds
To answer this question, we will conduct a comparative analysis of the dissociation of an ordinary salt, a double salt and a complex compound:
1) dissociation of medium salts - potassium and aluminum sulfates
K 2 SO 4 → 2K + + SO 4 2– ,
Al 2 (SO 4) 3) → 2Al 3+ + 3SO 4 2– ;
2) dissociation of double salt – potassium alum
KAl(SO 4) 2 → K + + Al 3+ + 2SO 4 2– ;
3) dissociation of a complex compound - potassium hexacyanoferrate (III)
K 3 → 3K + + 3– .
From the given equations of electrolytic dissociation it is clear that the dissociation products of the double salt completely coincide with the dissociation products of potassium and aluminum sulfates. In the case of a complex salt, the dissociation products contain a complex particle (complex ion), enclosed in square brackets, and simple ions neutralizing its charge.
The complex ion, in turn, dissociates like a weak electrolyte, that is, reversibly and stepwise:
3– ↔ Fe 3+ + 6CN – .
For complex ions, it is allowed to write dissociation products in all steps in one equation.
Dissociation products of a complex ion:
1) Fe 3+ – complexing agent,
2) 6СN – – ligands.
Thus, the complex compound includes:
1) complexing agent – central atom,
2) ligands – particles coordinated around the complexing agent,
3) particles that neutralize the charge of the complex ion. If the charge of a complex ion is zero, then it accordingly consists only of a complexing agent and ligands.
Typical complexing agents are metal cations of side subgroups: Ag +, Cu 2+, Fe 3+ and others.
Typical ligands: NH 3, H 2 O, CN –, NO 2 –, halide ions and others.
The complexing agent forms a strong bond with the ligands through covalent bonds and/or electrostatic interactions.
The coordination number is the number of monodentate ligands coordinated around the complexing agent. The coordination number is usually equal to twice the charge of the complexing agent.
The number of bonds formed by each ligand with the central atom is called the denticity of the ligand. For example:
1) monodentate ligands: F – , Br – , I – , CN – , OH – , NH 3 , H 2 O etc.;
2) bidentate ligands: H 2 N–CH 2 –CH 2 –NH 2 – ethylenediamine, oxalate ion, carbonate ion, etc.;
3) polydentate ligands - an example is ethylenediaminetetraacetate ion (EDTA). Complexes with polydentate ligands are called chelate complexes. They are widespread in nature and play an important role in biological processes, for example, blood hemoglobin (Fe 2+ complexing agent), chlorophyll (Mg 2+ complexing agent).
Name of the complex compound consists of the names of an anion and cation. The name of the compound is read from right to left, with the anion being named in the nominative case and the cation in the genitive case.
The number of ligands is indicated by Greek numerals: 1 - mono, 2 - di, 3 - tri, 4 - tetra, 5 - penta, 6 - hexa, 7 - hepta, 8 - octa. Names of the most common ligands: F - - fluoro, Cl - - chloro, Br - - bromo, I - - iodo, OH - - hydroxo, SO 3 2- - sulfito, NO 2 - - nitro, CN - - cyano,
CNS – rhodano, NH 3 – ammine, en – ethylenediamine, H 2 O – aqua.
The name of the complexing agent depends on the charge of the ion it is included in. In the case of a complex cation or a complex particle without an outer sphere, the Russian name of the complexing agent is used, and in the case of a complex anion, the root of the Latin name of the complexing element and the ending “at” are added after the name of the ligands.
If the inner sphere of the complex includes molecules and anions as ligands, then the anions (with the ending in “o”) are called first, and then the molecules. If several oxidation states are possible for a complexing agent, it is indicated in parentheses with a Roman numeral.
Examples of names of complex compounds:
1) anion type:
Na – sodium tetrahydroxyaluminate,
K 4 – potassium hexacyanoferrate(II);
2) cationic type:
SO 4 - tetraammine copper(II) sulfate,
Cl 2 – dichlorotetraammineplatinum(IV) chloride;
3) electrically neutral complexes:
– trifluorotriaquachrome,
– iron pentacarbonyl.
Dissociation of complex compounds occurs like a strong electrolyte into a complex ion and ions of the outer sphere. In turn, the complex ion or electrically neutral complex dissociates like a weak electrolyte into a complexing agent and ligands.
Quantitatively, the state of equilibrium is characterized by the corresponding value of Kr. In relation to the dissociation of a complex ion, the equilibrium constant (Kp) is called the instability constant (Kn). The smaller the Kn, the more stable the complex. For example:
1) Dissociation of the anionic complex
K 3 → 3K + + 3– ,
3- ↔ Fe 3+ + 6CN – ,
2) Dissociation of the cationic complex
SO 4 → 2+ + SO 4 2– ,
2+ ↔ Cu 2+ + 4NH 3 ,
The ferric cyanide complex is more stable.
18.2 Reactions involving complex compounds
Examples of reactions for the formation of complex compounds with a complex cation (1), a complex anion (2) and a neutral complex (3):
1) Ni(NO 3) 2 + 6NH 3 → (NO 3) 2,
Ni 2+ + 6NH 3 → 2+ ;
2) Cr(OH) 3 + 3KOH conc. ↔ K 3,
Cr(OH) 3 + 3OH – ↔ 3– ;
3) Fe + 5СО = .
Conclusion: complex compounds are formed if complexing ions and ligands are present in the solution.
As an example of the transition from one complex compound to
For another, let’s look at the reaction of converting the ammonia complex of copper into a cyanide complex:
SO 4 + 4KSN ↔ K 2 + K 2 SO 4 + 4NH 3,
2+ + SO 4 2– + 4К + + 4СN – ↔ 2– + 4К + + 4NH 3 + SO 4 2– ,
Cu(NH 3) 4 ] 2+ + 4СN – ↔ 2– + 4NH 3 .
Kn(2+) = 5.0·10 – 4, and Kn(2–) = 5.0·10 – 28,
that is, a stronger complex ion is formed.
Let us analyze the destruction of a complex compound using the example of the ammonia complex of silver:
NO 3 + KI ↔ AgI¯ + 2NH 3 + KNO 3,
NO 3 – + K + + I – ↔ AgI¯ + 2NH 3 + NO 3 – + K + ,
I – ↔ AgI¯ + 2NH 3 .
The equilibrium of this reaction is shifted to the right, since
Kn(+) = 6.8·10 – 8, and PR(AgI) = 1.5·10 – 16,
that is, a compound poorly soluble in water is formed - silver iodide.
The above reactions characterize the participation of complex compounds in ion exchange reactions.
As an example of a redox reaction, let us examine the reaction of converting the cyanide complex of ferrous iron into the cyanide complex of ferric iron:
K 4 + O 2 + H 2 O → K 3 + KOH,
Fe +2 – 1е = Fe +3 | ×4,
О 2 + 4е = 2О –2 | × 1.
4K 4 + O 2 + 2H 2 O → 4K 3 + 4KOH.
The elements of the eighth (iron, ruthenium, osmium, hassium), ninth (cobalt, rhodium, iridium, meitnerium) and tenth (nickel, palladium, platinum, darmstadtium) groups are historically considered together in connection with their unification into a single eighth group of the short-period version of the periodic table . The elements of the fifth and sixth periods (ruthenium, osmium, rhodium, iridium, palladium, platinum) included in its composition are noble, often found together in the form of alloys in which platinum predominates, therefore they are usually combined into the family of platinum metals (platinoids). Likewise, iron, cobalt and nickel are sometimes considered a separate triad (iron triad). While there is certainly some similarity between the platinum metals, the chemistry of elements belonging to different groups, for example, osmium, rhodium and palladium, differs significantly, but at the same time, there is a noticeable similarity between similar compounds of elements within the group, for example, cobalt(III) ammoniaates, rhodium(III) and iridium(III). Therefore, the chemical properties of oxygen-containing and complex compounds are described in the textbook by group. The seventh period elements hassium, meitnerium and darmstadtium are radioactive with a short half-life and are obtained only in quantities of several tens of atoms.
Iron is one of the seven metals of antiquity, that is, it has been known to mankind since the earliest periods of the history of society. Although the ability of cobalt compounds to give glass a bright blue color was already known to the Egyptians and Phoenicians, the element itself in the form of a simple substance was obtained only in 1735 by the German chemist G. Brandt, and a few years later by the Swedish metallurgist A.F. Kronstedt isolated nickel from copper ore. Platinum is traditionally considered the metal of the Ecuadorian Indians, as it was used by them to make jewelry and ritual masks before the arrival of the conquistadors. The infusible metal, which looks like silver, was given the name platina by the Spaniards, a diminutive derogatory word for “silver.” For a long time, the metal did not find any use due to its high hardness and refractoriness. The English chemist W. Wollaston was the first to obtain malleable platinum in 1805, after improving the hot forging process. He is also credited with the discovery of palladium (named after the asteroid Pallas, discovered in 1802) and rhodium, named after the pink-red color of the salts. From the powder remaining after treating raw platinum with aqua regia, iridium (from the Latin iris - rainbow, from compounds that have bright colors of various colors) and osmium (from the Greek οσμη - smell, from the sharp unpleasant odor of volatile tetroxide) were soon isolated. In 1844, Klaus, a professor of chemistry at Kazan University, isolated ruthenium from Ural ore sent to him for analysis, which he named in honor of Russia.
Superheavy platinum metals are radioactive hassium, meitnerium and darmstadtium. These elements were obtained in the 1980s - 1990s. at the super-powerful nuclear accelerator in Darmstadt (Germany) on reactions
208 Pb + 58 Fe 265 Hs + 1 n τ 1/2 (265 Hs) = 2×10 –3 s
209 Bi + 58 Fe 266 Mt + 1 n τ 1/2 (266 Mt) = 3.4×10 –3 s
208 Pb + 62 Ni 269 Ds + 1 n τ 1/2 (269 Ds) = 2.7×10 –4 s
Hassium was named in honor of the state of Hesse, where the city of Darmstadt is located, meitnerium - in honor of the Australian scientist Lise Meitner, who studied the fission reactions of uranium nuclei, and darmschatttium in honor of Darmstadt. The name of the last element was approved by the IUPAC commission in 2003.
Group 8 elements have a common ground state electronic configuration (n – 1)d 6 ns 2 is violated in ruthenium due to “electron leakage”. Similar phenomena occur in the rhodium atom, which is part of the ninth group, with a common electron configuration (n – 1)d 7 ns 2 . Among the elements of the tenth group is the configuration (n – 1)d 8 ns 2 observed only in the nickel atom: for platinum in the ground state there is a “leap” of one electron, and for palladium – two, which leads to the complete completion of the d-shell (Table 6.1).
Table 6.1.
Some properties of elements of the eighth - tenth groups.
| Group | Eighth | Ninth | Tenth | ||||||||
| Core charge | 26 Fe | 44 Ru | 76 Os | 27Co | 45 Rh | 77Ir | 28 Ni | 46 Pd | 78 Pt | ||
| Number of natural isotopes | |||||||||||
| Electronic configuration | 3d 6 4s 2 | [Kr] 4d 7 5s 1 | [Xe]4f 14 5d 6 6s 2 | 3d 7 4s 2 | [Kr]4d 8 5s 1 | [Xe]4f 14 5d 7 6s 2 | 3d 8 4s 2 | [Kr]4d 10 | [Xe]4f 14 5d 9 6s 1 | ||
| Metal radius, nm | 0.126 | 0.134 | 0.135 | 0.125 | 0.134 | 0.136 | 0.124 | 0.137 | 0.139 | ||
| Ionization energy, kJ/mol, | I 1 I 2 I 3 I 4 I 5 | (4500) (6100) | (1600) (2400) (3900) (5200) | (4400) (6500) | (1680) (2600) (3800) (5500) | (4700) (6300) | (2800) (3900) (5300) | ||||
| Ionic radius, nm (CN = 6) | E 2+ E 3+ E 4+ E 5+ E 6+ E 7+ | 0.061* 0.065* 0.059 | - 0.068 0.062 0.057 | - - 0.063 0.058 0.055 0.053 | 0.065* 0.054* 0.053 | - 0.067 0.060 0.055 | - 0.068 0.063 0.057 | 0.069 0.056* 0.048 | 0.086 0.076 0.062 | 0.080 ‘ 0.063 0.057 | |
| Electronegativity according to Pauling | 1.83 | 2.2 | 2.2 | 1.88 | 2.28 | 2.20 | 1.91 | 2.20 | 2.28 | ||
| Electronegativity according to Allred-Rochow | 1.64 | 1.42 | 1.52 | 1.70 | 1.45 | 1.55 | 1.75 | 1.35 | 1.44 | ||
| Oxidation states | (–2), (–1), 0, +2, +3, (+4), (+5), +6 | (–2), 0, (+2), (+3), +4, (+5), +6, +7, +8 | (–2), 0, (+2), +3, +4, (+5), +6, +7, +8 | (–1), 0, (+1) (+2), (+3), +4, (+5), (+6), (+7), +8 | (–1), 0, +1, +2, +3, (+4), (+5), (+6) | (–1), 0, +1, (+2), +3, +4, (+5), (+6) | (–1), 0, (+1), +2, (+3), (+4) | 0, (+1), +2, (+3), (+4) | 0, (+1), +2, (+3), +4, (+5), (+6) | ||
* in low spin state
The patterns of changes in the properties of elements of groups 8 - 10 when moving across a period and group obey the general patterns discussed in Chapter 1. The first ionization energies in the eighth and ninth groups decrease during the transition from a 3d metal to a 4d metal (Table 6.1.), which is related with increasing atomic radius and removal of valence electrons from the nucleus. A further increase in E 1 during the transition to d-metals of the sixth period is explained by screening effects associated with the filling of the 4f sublevel. The general pattern does not apply to elements of the tenth group due to the significant stabilization of the d-orbitals of the nickel atom caused by the double “breakthrough” of electrons.
The metals of the iron triad, like other elements of the 3d series, having a small atomic radius and relatively small d-orbitals with an insignificant degree of overlap, have much higher chemical reactivity compared to platinum metals. In contrast, iron, cobalt and nickel displace hydrogen from acid solutions and oxidize in air. They are not characterized by cluster compounds, which, even if formed, often turn out to be unstable in air and in aqueous solution. Platinum metals in general can be considered as the least chemically active metals, due to their relatively low (compared to the d-elements at the beginning of the transition series) atomic radius and the high degree of overlap of d-orbitals. Of these, only osmium is able to directly interact with oxygen, and only palladium reacts with concentrated nitric acid. Platinum metals in general are characterized by complex compounds, including complexes with π-acceptor ligands (carbon monoxide, alkenes, alkadienes), hydrides, which are often stable even in aqueous solution, and clusters. Like other heavy transition metals, platinoids exhibit high oxidation states, up to +8 (OsO 4). The stability of higher oxidation states increases down the groups (Footnote: For a review of the chemistry of platinum metals in oxidation states from +4 to +8, see D.J. Gulliver, W. Levason, Coord. Chem. Rev., 1982, 46, 1).
When moving through the period, as the number of valence electrons increases and their pairing occurs, the d-sublevel stabilizes and the stability of higher oxidation states decreases. Thus, iron can be oxidized in an aqueous solution to ferrate FeO 4 2–, containing a metal atom in the oxidation state +6, cobalt and nickel under these conditions acquire an oxidation state of +3. The highest oxidation states are most stable for elements of the eighth group - iron (+6), ruthenium (+8) and osmium (+8) (Footnote: There is information about obtaining an iron compound in the oxidation state +8: See Kiselev Yu. M., Kopelev N. S., Spitsyn V. I., Martynenko L. I. Reports of the USSR Academy of Sciences, 1987, vol. 3, p. 628). These metals exhibit lower oxidation states in compounds with π-acceptor ligands, for example, in carbonyls: K 2 , K. The value of the most stable oxidation state decreases monotonically when moving through the period: for iron the oxidation state is most characteristic +3, cobalt exists in aqueous solutions predominantly in the +2 oxidation state, and in +3 complexes, nickel - exclusively in the +2 oxidation state. This is consistent with the increase in third ionization energies in the series Fe – Co – Ni (Table 6.1.). Ni 2+ ions are resistant to oxidation by atmospheric oxygen at any pH, cobalt(II) salts are stable in acidic and neutral environments, and in the presence of OH ions they are oxidized, iron(II) is converted into iron(III) under the influence of oxygen (E 0 ( O 2 /H 2 O) = 1.229 V, pH = 0, and 0.401 V, pH = 14) at any pH. The reduction activity of the triad metals also decreases when moving along the 3d row (Table 6.2.).
Table 6.2. Standard electrode potentials M(III)/M(II) and M(III)/M(0) for elements of the iron triad
The change in oxidation states that are stable in aqueous solutions can be represented in the form of a diagram:
Examples of compounds of elements of groups 8 – 10 with different oxidation states are given in table. 6.3. Ions with electronic configurations d 3 (Ru +5), d 5 (Fe +3,) and d 6 (Fe +2, Co +3, Rh +3, Ir +3) are characterized by octahedral complexes, for configurations d 4 ( Ru +4, Os +4) and d 7 (Co +2) - tetragonally distorted octahedral, arising due to the Jahn-Teller effect, for d 8 - octahedral (Ni +2 with weak and medium field ligands) - or planar square ( Pd +2, Pt +2, as well as Ni +2 with high field ligands). Molecules and ions with tetrahedral geometry arise from the interaction of metal ions with bulky ligands (PR3, Cl–, Br–, I–) or when the d-sublevel is completely filled (d10, Pd0, Rh–1, Ru–2).
A consistent decrease in atomic and ionic radii as one moves through the period leads to a gradual decrease in the maximum coordination numbers from 10 for iron (in ferrocene) to 8 for cobalt (in 2–) and 7 for nickel (in complexes with macrocyclic ligands). Heavy analogues of iron - ruthenium and osmium - also rarely increase the coordination number above six. For platinum(II) and palladium(II), which have an electronic configuration of d 8, square-planar complexes with a coordination number of 4 are most characteristic.
Another consequence of the decrease in ionic radii is a slight decrease in the values of the solubility product of hydroxides M(OH) 2, and, consequently, their basicity constants when moving along the 3d series:
Mn(OH) 2 Fe(OH) 2 Co(OH) 2 Ni(OH) 2
PR, 20 °C 1.9×10 –13 7.1×10 –16 2.0×10 –16 6.3×10 –18
The degree of hydrolysis of salts with anions of the same name also increases in the same direction. This leads to the fact that when manganese(II) and iron(II) salts are exposed to a solution of average sodium carbonate, average carbonates precipitate, and cobalt and nickel ions under these conditions give basic salts. An increase in the Pearson softness of 3d metal cations as they move across a period as the d sublevel is filled and the ionic radii decrease causes the M-S bond to strengthen compared to M-O. This clearly illustrates the monotonic change in the solubility products of sulfides:
MnS FeS CoS NiS CuS
PR, 20 °C 2.5×10 –13 5.0×10 –1 8 2.0×10 – 25 2.0×10 – 26 6.3×10 – 36
Thus, manganese and iron are found in nature mainly in the form of oxygen compounds, followed by iron, cobalt, nickel and copper in polysulfide ores.
Table 6.3. Oxidation states, electronic configurations, coordination numbers (CN) and geometry of molecules and ions
| Electronic Configuration | K.Ch. | Geometry | Eighth group | Ninth group | Tenth group | ||||
| Oxidation state | Examples | Oxidation state | Examples | Oxidation state | Examples | ||||
| d 10 | tetrahedron | –2 | 2– , M = Fe, Ru, Os | –1 | – , M = Co, Rh | Ni(CO) 4 , M(PF 3) 4 , M = Pd, Pt | |||
| d 9 | trigonal bipyramid | –1 | 2– | – | +1 | – | |||
| – | Co 2 (CO) 8 , M 4 (CO) 12 , M = Rh, Ir | – | |||||||
| octahedron | |||||||||
| d 8 | octahedron | – | +1 | – | +2 | 2+ , 3+ | |||
| trigonal bipyramid | , | 3– | |||||||
| – | – | 2– | |||||||
| tetrahedron | |||||||||
| – | RhCl(PPh 3) 2 | 2– . | |||||||
| 2– , M = Pd, Pt | |||||||||
| square | octahedron | +1 | + | +2 | d 7 | +3 | 2+, Rh 2 (CH 3 COO) 4 (H 2 O) 2 | ||
| tetrahedron | – | 2– | – | ||||||
| 3– , M = Ni, Pd | tetrahedron | +2 | 2– | +3 | 5– | +4 | d 6 | ||
| octahedron | 2+ , 4– | 3+ | – | ||||||
| 2– , 2– , M = Pd, Pt | tetrahedron | +3 | – | +4 | – | +5 | – | ||
| octahedron | 3+ , 3– | d 5 | – | ||||||
| 2– , 2– , M = Co, Rh | tetrahedron | +4 | +5 | – | +6 | d 4 | |||
| octahedron | PtF 6 | 2– , M = Ru, Os | – | ||||||
| – , M = Rh, Ir | tetrahedron | +5 | d 3 | +6 | 3– , – , M = Ru, Os | ||||
| MF 6, M = Rh, Ir | tetrahedron | +6 | 2– , 2– , | – | |||||
| d 2 | tetrahedron | +7 | d 1 | – | |||||
| octahedron | – , M = Ru, Os | – | |||||||
| OsOF 5 | pentagonal bipyramid | – | |||||||
| OsF 7 | tetrahedron | +8 | d 0 | – |
MO 4, M = Ru, Os
Although the adult human body contains only about 4 g of iron, it plays a critical role in the processes of transporting oxygen to tissues and cells, removing carbon dioxide, and oxidative phosphorylation. Three-quarters of the iron atoms in the body are in the form of hemoglobin, which consists of a porphyrin complex of iron called heme and the globin protein. Hemoglobin ensures the transport of oxygen to the tissues of the body, and its related protein, myoglobin, which has a simpler structure and, unlike hemoglobin, does not have a quaternary structure, determines the ability of tissues to store oxygen. Hemoglobin is found in red blood cells, and myoglobin is present in muscle tissue. Both compounds have a red color due to the presence of an iron atom in them in the +2 oxidation state, and the oxidation of iron leads to the loss of their biological activity! In the protein structure, the heme is located in the gap between two helices formed by the polypeptide chain. The porphyrin complex ensures the square-planar coordination of the iron atom by the four nitrogen atoms of the porphyrin ring. The nitrogen atom of the imidazole ring of the amino acid histidine, which belongs to the nearest polypeptide chain, complements the coordination number of iron to five. Thus, in the non-oxygenated form of hemoglobin, the sixth position in the coordination sphere of the iron atom remains vacant. This is where the oxygen molecule attaches. When oxygen is added, the iron atom moves out of the plane of the porphyrin ring by 0.02 nm compared to the deoxy form. This leads to conformational changes in the arrangement of polypeptide chains. In this case, the complex becomes diamagnetic due to the transition of the iron atom to a low-spin state:
Arterial blood contains predominantly oxyhemoglobin, and as the oxygen molecules contained in it are converted into myoglobin, the color of the blood becomes darker - this indicates the return of heme to its previous deoxy form. Hemoglobin not only carries oxygen from the lungs to peripheral tissues, but also accelerates the transport of carbon dioxide from the tissues to the lungs. Immediately after oxygen is released, it binds approximately 15% of the CO 2 dissolved in the blood.
The CO molecule is capable of forming a stronger complex with heme than an oxygen molecule, thereby preventing its transport from the lungs to the tissues. This is why inhalation of carbon monoxide leads to death from lack of oxygen. Cyanide ion also plays a similar role, although its toxicity is mainly due to interaction with other iron-containing hemoproteins - cytochromes. Cytochromes are involved in oxidative phosphorylation - the oxidation of pyruvate that occurs in mitochondria, which is formed during the primary oxidation of carbohydrates. The energy released in this case accumulates in the form of high-energy bonds of the ATP molecule. In the complex chain of oxidative phosphorylation, cytochromes a, b and c are the carriers of electrons from one enzyme to another and, ultimately, to oxygen. In this case, the iron atom constantly changes its oxidation state.
The most studied is cytochrome P 450, which is a heme that differs from the heme in hemoglobin in a set of substituents and contains iron +3, coordinated by a water molecule and a sulfur atom belonging to the amino acid cysteine (Fig. 6.1. Model of the active center of cytochrome P 450, surrounded by the protein part of the molecule) . Its role is to hydroxylate lipophilic compounds foreign to the body, formed as by-products or entering the body from the outside:
R–H + O 2 + 2e – + 2H + ¾® ROH + H 2 O
At the first stage (Fig. 6.2. Catalytic cycle of cytochrome P 450). Cytochrome attaches a substrate molecule, which then (step 2) undergoes reduction by another enzyme. The third stage is the addition of oxygen, similar to that described above for hemoglobin. In the low-spin iron complex formed in this case, the coordinated O2 molecule is reduced to a peroxide ion (stage 4), which, as a result of intramolecular electron transfer, leads to an oxoferryl complex containing iron in the +5 oxidation state (stage 5). When it is reduced, the oxidized substrate is separated, and the cytochrome returns to its original state (stage 6).
Heme also underlies catalases and peroxidases, enzymes that catalyze oxidation reactions with hydrogen peroxide. One molecule of catalase per second can cause the decomposition of 44,000 molecules of H 2 O 2.
Along with cytochromes, ferredoxins participate in oxidative phosphorylation - iron-sulfur proteins, the active center of which is a cluster containing an iron atom, sulfide bridges and cysteine amino acid residues (Fig. 6.3. Structure of bacterial ferredoxin (a), active center of ferredoxin (b)). Ferredoxins found in bacteria, containing eight atoms of iron and sulfur, play a key role in the processes of atmospheric nitrogen fixation. In the molecule of bacterial ferredoxin, two identical groups of Fe 4 S 4 were found, having the shape of a cube and located at a distance of 1.2 nm from each other. These two clusters are located inside a cavity formed by chains of amino acids linked to each other. The composition of nitrogenase (see page 169, volume 2) also includes proteins with a molecular weight of about 220 thousand, containing two molybdenum atoms and up to 32 iron atoms. (R. Murray, D. Grenner, P. Mayes, V. Rodwell, Human Biochemistry, M., Mir, 1993).
END OF ADDENDUM
6.2. Prevalence in nature, production and use of simple substances of 8 - 10 groups.
In terms of prevalence in nature, among elements of groups 8–10, the undisputed leader is iron, or more precisely, its isotope 56 Fe, whose nuclei have the highest binding energy of protons and neutrons, and, therefore, have high stability.
Indeed, the number of iron atoms in the Universe significantly exceeds the number of atoms of any of the neighboring elements in the Periodic Table and is close in order to hydrogen and helium. For example, on the Sun, the hydrogen content is estimated at 1 × 10 12 conventional units, helium at 6.31 × 10 10, and iron at 3.16 × 10 17. This is explained by the fact that the nucleus of the 56 Fe nuclide is one of the magic ones, that is, having completely filled nuclear shells. As the number of nucleons in a nucleus increases, the binding energy per nucleon initially increases rapidly, reaching a maximum just at the iron nucleus, and then gradually decreases (Fig. 6.4. Binding energy per nucleon as a function of the element’s atomic number). (R. J. Theiler, Origin chemical elements, M., Mir, 1975).
In terms of content in the earth’s crust, iron is in fourth place (4.1%), second only to oxygen, silicon and aluminum, nickel (8×10–3%) is in the second ten, cobalt (2×10–3%) is in third, and platinum metals are rare (Ru 10–7%, Pt 10–7%, Pd 6×10–8%, Rh 2×10–8%, Os 10–8%, Ir 3×10–10%) . In the earth's crust, iron is represented mainly by hematite Fe 2 O 3 (red iron ore), magnetite Fe 3 O 4 (magnetic iron ore), limonite Fe 2 O 3 ×xH 2 O (brown iron ore), siderite FeCO 3 (iron spar, spar iron ore ), ilmenite FeTiO 3 and sulfur-containing mineral pyrite FeS 2 (iron pyrite). In total, more than 300 iron-containing minerals are known. A significant amount of iron is included in the composition of various silicates and aluminosilicates that make up rocks. When they weather, iron compounds, mainly iron(III) oxide and oxohydroxide, enter quartz sand, clays and soil, giving them a yellow-brown, earthy color. Iron of meteorite origin is found in free form on earth, often in the form of an alloy with nickel. Native iron is also known in the form of flakes or small leaves embedded in basalts. Only occasionally does it form separate pieces. Such finds are so rare that in the Stone and Bronze Ages, tools made from it were valued much more than gold. The Earth's mantle contains significant amounts of iron in the form of spinels, silicates, and oxides. It is believed that iron, mixed with nickel and sulfur, is the main part of the earth's core. In the surface layer of the Moon, the iron content reaches 0.5%.
The development of iron production from iron ore marked the beginning of the Iron Age. To reduce iron oxides with coal, a temperature above 1400 °C is required, which a conventional fire could not provide. That is why, in the early stages of the development of society, iron ores were not available as raw materials for metal production. People had to limit themselves to only random finds of meteorite iron. At the beginning of the first millennium BC. the cheese-blowing method of ore recovery was mastered, based on the use of a forge - a structure made of stones coated with clay. Holes were left in the walls of the forge, into which air was pumped through special clay tubes - nozzles - using leather bags called bellows. Charcoal and iron ore were poured into the forge, and a fire was lit on top. The resulting metal was welded into kritsa - a porous mass from which products were forged. The cheese-making method was replaced by blast furnace production. This occurred as a result of an increase in the height of the furnace, which also required the introduction of fluxes - special additives that form low-melting slags with the waste rock contained in the ore. Since in a blast furnace, unlike a forge, the molten metal is in contact with coal for a long time, it carbonizes, turning into cast iron. This requires an unnecessary operation to “convert” cast iron into steel and iron. The first blast furnaces appeared in the Netherlands at the end of the 14th - beginning of the 15th centuries; in the 16th century they reached a height of 4 - 5 m. In Russia, blast furnace production arose in the 17th century, and in the next century it developed in the Urals.
Addition. State diagram of the iron-carbon system.
State diagram of the Fe-C system in the region up to 6.5 wt. % C, shown in Fig. 6.5 a, is important in metallurgy for the targeted production of various types of steels and cast irons. Pure iron crystallizes in three modifications, α, γ and δ, each of which dissolves a certain amount of carbon and is stable over a certain temperature range. Solid solutions of carbon in these modifications, α-Fe, γ-Fe and δ-Fe-C, are called α-ferrite, γ-austenite and δ-ferrite, respectively. α-Fe and δ-Fe have cubic body-centered lattices and γ-Fe has cubic face-centered lattices. The solubility of carbon is greatest in austenite (γ-Fe).
Melts containing up to 1.75 wt. % C, after rapid cooling to 1150 o C, they form a homogeneous solid solution - austenite. These alloys make steel. In melts containing more than 1.75% C after cooling to 1150 o C, in addition to solid austenite, there is also a liquid eutectic of composition point A (Fig. 6.5.a). When cooled below 1150 o C, it crystallizes and fills the space between with a thin mixture of crystals. austenite crystals. The resulting solid systems are cast iron. Depending on the conditions, the eutectic can crystallize in two ways. When cooled rapidly, the solidified ectectic consists of austenite crystals and unstable Fe 3 C crystals called cementite. With slow cooling, a mixture of austenite crystals and stable graphite is formed. Cast iron containing cementite is called white, and cast iron containing graphite is called gray. The solidified eutectic of austenite and cementite is called ledeburite, and only ledeburite is released from a melt containing 4.3% C.
When austenite is cooled below 1150 o C, it recrystallizes. From solid solutions containing less than 0.9 wt. % C, α-Fe ferrite is released first (see inset in Fig. 6.5.a), and from solutions containing more than 0.9 wt. % C, cementite is primarily released, which is called secondary cementite. In both cases, the composition of the remaining solid solution approaches eutectoid point B. At this point, ferrite and cementite crystals simultaneously precipitate in the form of a thin layered mixture called pearlite. A melt containing 0.9% C, when cooled, can form pure pearlite, which does not contain previously released large crystals of ferrite or Fe 3 C.
By adjusting the composition of the initial melt, the cooling rate and heating time at temperatures selected from the diagram, it is possible to obtain alloys with different microstructures, compositions, orientations and stresses in the crystals. If the resulting system is then cooled (hardened) very quickly, then all further transformations are greatly inhibited, and the created structure is preserved, although it turns out to be thermodynamically unstable. This is the way to obtain different grades of steel.
Rice. 6.5. Phase diagram of the iron-carbon system
END OF ADDENDUM.
Currently, iron ore is reduced with coke in blast furnaces, while molten iron partially reacts with carbon, forming iron carbide Fe 3 C (cementite), and partially dissolves it. When the melt solidifies, it forms cast iron. Pig iron used to produce steel is called pig iron. Steel, unlike cast iron, contains less carbon. The excess carbon contained in the cast iron must be burned off. This is achieved by passing air enriched with oxygen over the molten cast iron. There is also a direct method for obtaining iron, based on the reduction of magnetic iron ore pellets with natural gas or hydrogen:
Fe 3 O 4 + CH 4 = 3Fe + CO 2 + 2H 2 O.
Very pure iron in powder form is obtained by decomposition of carbonyl Fe(CO) 5 .
ADDITION. Iron alloys.
Iron-based alloys are divided into cast iron and steel.
Cast iron– an alloy of iron with carbon (contains from 2 to 6% C), containing carbon in the form of a solid solution, as well as crystals of graphite and cementite Fe 3 C. There are several types of cast iron, differing in properties and fracture color. White cast iron contains carbon in the form of cementite. It is highly fragile and does not find direct use. All white cast iron is processed into steel (pig iron). Gray cast iron contains graphite inclusions - they are clearly visible at the fracture. It is less fragile than white and is used for the manufacture of flywheels and water heating radiators. Adding a small amount of magnesium to the melt causes graphite to separate out not in the form of plates, but in the form of spherical inclusions. This modified cast iron has high strength and is used to make engine crankshafts. Mirror cast iron, containing 10–20% manganese and about 4% carbon, is used as a deoxidizing agent in steel production.
Fig.6.6. Gray cast iron (a) and high-strength cast iron (b) under a microscope.
The raw materials for the production of cast iron are iron ore and coke. Cast iron is smelted in blast furnaces - large furnaces up to 80 m high, lined with refractory bricks on the inside and covered with a steel casing on top. The upper part of the blast furnace is called the shaft, the lower part is called the hearth, and the upper hole, which serves to load the charge, is called the top. Hot air enriched with oxygen is supplied to the furnace from below. In the upper part of the hearth, coal is burned to form carbon dioxide. The heat released in this case is sufficient for the process to proceed. Carbon dioxide, passing through layers of coke, is reduced to carbon monoxide (II) CO, which, reacting with iron ore, reduces it to metal. To remove impurities contained in the ore, for example, quartz sand SiO 2, fluxes are added to the furnace - limestone or dolomite, which decompose to oxides CaO, MgO, which bind the slag into low-melting fluxes (CaSiO 3, MgSiO 3). In addition to iron, coke also reduces impurities contained in the ore, for example, phosphorus, sulfur, manganese, and partially silicon:
Ca 3 (PO 4) 2 + 5C = 3CaO + 5CO + 2P,
CaSO 4 + 4C = CaS + 4CO,
MnO + C = Mn + CO,
SiO2 + 2C = Si + 2CO.
In the molten metal, sulfur is present in the form of FeS sulfide, phosphorus in the form of Fe 3 P phosphide, silicon in the form of SiC silicide, and excess carbon in the form of Fe 3 C carbide (cementite). The gases coming out of the blast furnace are called blast furnace or top gases. They consist of approximately one-third by volume carbon monoxide, so they are used as fuel to heat the air entering the blast furnace.
RICE. 6.7 Blast furnace diagram
Steel– an alloy of iron and carbon (contains from 0.5 to 2% C), containing carbon only in the form of a solid solution. Steel is harder than iron, harder to bend, more elastic, easier to break, although not as brittle as cast iron. The higher the carbon content, the harder it is. In conventional steel grades, no more than 0.05% sulfur and 0.08% phosphorus are allowed. Even a slight admixture of sulfur makes steel brittle when heated; in metallurgy, this property of steel is called red brittleness. The phosphorus content in steel causes cold brittleness - brittleness at low temperatures. Hardened steel is formed by sudden cooling of steel heated to a red-hot temperature. This steel has high hardness, but is brittle. Cutting tools are made from hardened steel. With slow cooling, tempered steel is obtained - it is soft and ductile. By introducing alloying additives into the melt ( doping) – chromium, manganese, vanadium, etc., special grades of steel are obtained. Steel containing more than 13% chromium loses its ability to corrode in air and becomes stainless. It is used in the chemical industry, in everyday life, and in construction. Particularly strong steels containing vanadium are used to cast armor.
The raw material for steel production is cast iron, and the essence of the processes occurring during smelting is to remove excess carbon from the alloy. To do this, oxygen is passed through molten cast iron, which oxidizes the carbon contained in the cast iron in the form of graphite or cementite to carbon monoxide CO. However, in this case, part of the iron is also oxidized by oxygen to oxide:
2Fe + O 2 = 2FeO.
To reduce FeO back to iron, deoxidizers are introduced into the melt; as a rule, these are active metals - manganese, barium, calcium, lanthanum. They reduce oxidized iron to metal:
Mn + FeO = MnO + Fe,
and then separated from the melt, floating to its surface in the form of fusible slag, interacting either with the furnace lining or with specially added fluxes:
MnO + SiO 2 = MnSiO 3.
Steel smelting is carried out in special furnaces. Depending on the type of furnace, there are several methods of steelmaking. In an open hearth furnace, the melting space is a bath covered with a vault of refractory bricks (Fig. 6.8. Steel production: (a) Open hearth furnace, Oxygen converter). Fuel is injected into the upper part of the furnace - it is natural gas or fuel oil. The heat released during its combustion heats the charge and causes it to melt. Over the 6–8 hours during which the molten cast iron is in the open-hearth furnace, the carbon in it gradually burns out. After this, the molten steel is poured out and after some time the cast iron is loaded again. The open-hearth process is periodic. Its main advantage is that the resulting steel can be cast into large molds. In terms of productivity, the open-hearth process is inferior to the oxygen-converter process, which is carried out not in large furnaces, but in small converters - pear-shaped devices, welded from steel and lined with refractory bricks on the inside. Oxygen-enriched air is blown from above through a converter mounted on a horizontal axis. The resulting oxides of manganese and iron react with the silicate lining of the converter, forming slag. The process lasts about 40 minutes, after which the converter is moved to an inclined position and molten steel and slag are poured sequentially (Fig. 6.8. b). Converters lined with sand-lime brick, called Bessemer converters after the English inventor Henry Bessemer, are not suitable for smelting steel from cast iron containing iron phosphides. For the processing of phosphorus-rich cast iron, Thomas converters are used, which are lined with limestone or dolomite on the inside. Steel smelting is carried out in the presence of lime, which binds the phosphorus contained in cast iron into phosphates, forming slag (Thomas slag), which is used as fertilizer. Alloy steels are smelted in electric furnaces at temperatures above 3000 °C. This makes it possible to obtain steels with special properties, including ultra-strong and refractory.
END OF ADDENDUM
Cobalt occurs in nature, mainly in the form of compounds with arsenic, smaltite CoAs 2 (cobalt speis) and cobaltite CoAsS (cobalt luster), but these minerals are too rare and do not form independent deposits. It is also part of complex copper-cobalt-nickel and copper-cobalt sulfide ores; it is found in small quantities in clays and shales that formed under conditions of lack of oxygen.
Nickel, like cobalt, has a high affinity for post-transition elements of the fifth period - arsenic and sulfur, and due to the proximity of ionic radii, it is often isomorphic to compounds of cobalt, iron, and copper. Due to this, large amounts of nickel in the lithosphere are bound into polysulfide copper-nickel ores. Among the sulfide minerals, the most important are millerite NiS (yellow nickel pyrite), pentlandite (Fe, Ni) 9 S 8, chloantite NiAs 2 (white nickel pyrite). Another important nickel raw material is serpentine rocks, which are basic silicates, for example, garnierite (Ni, Mg) 6 × 4H 2 O. Nickel compounds are found in small quantities in fossil coals, shale, and oil.
The main raw materials for the production of cobalt and nickel are polysulfide ores (footnote: silicates and other oxygen-containing nickel ores are first converted into sulfides by fusion with dehydrated gypsum and coal at 1500 °C: CaSO 4 + 4C = CaS + 4CO; 3NiO + 3CaS = Ni 3 S 2 + 3CaO + S). The agglomerated ore is mixed with sulfuric acid and smelted in a shaft furnace into a matte consisting of iron, cobalt, nickel and copper sulfides. This allows it to be separated from silicates that form slags. When the molten matte is cooled, sulfides are released in crystalline form. They are crushed and then heated to 1300 °C in a stream of air. The ability of sulfides to oxidize decreases in the order FeS > CoS > Ni 3 S 2, so iron sulfide first reacts with oxygen, which is converted into slag by adding silica. Further oxidation leads to the formation of cobalt and nickel oxides
2Ni2S3 + 7O2 = 6NiO + 4SO2.
They are transferred into solution by treatment with sulfuric acid or by resorting to anodic oxidation. The copper impurity is removed by introducing nickel powder, which reduces it to a simple substance. Cobalt and nickel have similar chemical properties. To separate them, the solution is alkalized and treated with sodium chlorate, which oxidizes only cobalt ions:
2CoSO 4 + Cl 2 + 3Na 2 CO 3 + 3H 2 O = 2Co(OH) 3 ¯ + 2NaCl + 3CO 2 + 2Na 2 SO 4.
In a slightly acidic environment, cobalt remains in the precipitate in the form of hydroxide, and nickel goes into solution in the form of a salt, which is converted into hydroxide. Oxides obtained by calcination of hydroxides are reduced with coal:
Co 3 O 4 + 4C = 3CO + 4CO,
NiO + C = Ni + CO.
During reduction, carbides Co 3 C and Ni 3 C are also formed; to remove them, the oxide is taken in excess:
Ni 3 C + NiO = 4Ni + CO.
To obtain purer metals, electrolytic refining is used. It also makes it possible to isolate platinum metals contained in the matte.
More than half of the cobalt and nickel produced is spent on the production of alloys. Magnetic alloys based on cobalt (Fe-Co-Mo, Fe-Ni-Co-Al, Sm-Co) are able to retain magnetic properties at high temperatures. Metal-ceramic alloys, which are carbides of titanium, tungsten, molybdenum, vanadium and tantalum, cemented with cobalt, are used for the manufacture of cutting tools. Steels with a high content of nickel and chromium do not corrode in air; surgical instruments and equipment for the chemical industry are made from them. The heat-resistant chromium-nickel alloy nichrome, containing 20–30% chromium, has high electrical resistance; coils of electric heaters are made from it. Copper-nickel alloys constantan (40% Ni, 60% Cu) and nickel (30% Ni, 56% Cu, 14% Zn) and monel (68%Ni, 28% Cu, 2.5) are also used as heating elements. % Fe, 1.5 % Mn) mint a coin.
Are important superalloys– materials based on iron, cobalt or nickel, specially designed for use at high temperatures. They have high corrosion resistance, maintain strength in the temperature range at which gas turbines operate, and are characterized by a high modulus of elasticity and a low coefficient of thermal expansion. The combination of oxidation resistance and strength makes these materials unmatched. Many superalloys have a face-centered cubic lattice, which, being the densest of all crystalline structures, provides exceptional thermomechanical properties of the material. The alloy consists of a base (Fe, Co, Ni), contains metal additives that increase surface resistance (Cr) and elements (Al) that form a cubic γ'-phase (γ'-Ni 3 Al), which has high strength and oxidation resistance . The introduction of small amounts of carbon (0.05 - 0.2%) into superalloys leads to the formation of carbides, for example, TiC, which, during the operation of the alloy at high temperatures, gradually transform into carbides of the composition M 23 C 6 and M 6 C, which are easily exposed to heat treatment. The carbon formed in this case goes into the form of a solid solution. Thus, the structure of a superalloy can be represented as a solid solution with fine-crystalline inclusions of intermetallic compounds and carbides, which ensure its hardness and strength. Additional alloying helps to slow down diffusion processes and increase the stability of the structure at high temperatures. One of the first superalloys was Rex-78, developed in 1935, consisting of 60% iron, 18% Ni, 14% Cr, and also containing small amounts of molybdenum, titanium, copper, boron, and carbon. It is used for the manufacture of turbine blades and nozzles (Superalloys II. Heat-resistant materials for aerospace and industrial power plants, M., Metallurgy, 1995)
Finely dispersed cobalt and nickel have high catalytic activity. Fine cobalt powder deposited on a support serves as an active catalyst for Fischer-Tropsch hydrocarbonylation. Nickel often replaces platinum in hydrogenation processes, for example, of vegetable fats. In the laboratory, catalytically active fine nickel powder (skeletal nickel, Raney nickel) is obtained by treating a nickel-aluminum alloy with alkali in an inert or reducing atmosphere. Nickel is used to produce alkaline batteries.
Many cobalt compounds are brightly colored and have been used since ancient times as pigments for the preparation of paints: cobalt aluminate CoAl 2 O 4 (“cobalt blue”, “Gzhel blue”) has a blue color, stannate Co 2 SnO 4 (“ceruleum”, “sky blue”) blue") - blue with a bluish tint, phosphates Co 3 (PO 4) 2 ("cobalt violet dark") and CoNH 4 PO 4 × H 2 O ("cobalt violet light") - reddish-violet, mixed cobalt(II) oxide ) and zinc CoO×xZnO (“green cobalt”) – bright green, cobalt silicates (“schmalt”, “cobalt glass”) – dark blue (E.F. Belenkiy, I.V. Riskin, Chemistry and technology of pigments, L ., Chemistry, 1974). Adding cobalt oxide to glass gives it a blue color.
Iron pigments are usually yellow-brown or red-brown in various shades. Among natural pigments, the most famous are ocher - crystalline oxohydroxide FeOOH and sienna containing clay. When heated, they dehydrate, turning red. Brown umber is formed by the weathering of iron ores containing manganese. The black pigment is magnetite.
Platinum metals are found in nature mainly in native form - in the form of simple substances, alloys with each other and with other noble metals. In very small quantities they are part of some polysulfide ores; finds of their own sulfide minerals, for example, laurite RuS 2, cooperite PtS, are extremely rare. The average total content of platinum metals in the Ural sulfide rads is 2–5 grams per ton. In nature, platinum grains are often found in the same placers as gold, therefore, in the form of separate inclusions, they are sometimes visible on the surface of ancient gold products, mainly of Egyptian origin. Large reserves of native platinum are concentrated in the South American Andes. In their constituent rocks, grains of platinum, together with particles of gold, are often included in pyroxenes and other basic silicates, from which, as a result of erosion, they turn into river sands. The gold washed from them contains small crystals of platinum, which are extremely difficult to separate. In the Middle Ages they did not strive for this: the admixture of heavy grains only increased the mass of the precious metal. Large platinum nuggets, up to nine kilograms, are also occasionally found. They necessarily contain impurities of iron, copper, platinum iodes, and sometimes gold and silver. For example, metal from the Choco deposit in Colombia, which was mined by the ancient Incas, has an approximate composition of Pt 86.2%, Pd 0.4%, Rh 2.2%, Ir 1.2%, Os 1.2%, Cu 0, 40%, Fe 8.0%, Si 0.5%. Native iridium contains 80–95% Ir, up to 2.7% Ru, up to 6.1% Pt; osmium – 82 – 98.9% Os, 0.9 – 19.8% Ir, up to 10% Ru, 0.1 – 3.0% Pt, up to 1.3% Rh, up to 1% Fe.
In Russia, the first platinum placer was discovered in 1824 in the Northern Urals, and soon mining began in the Nizhny Tagil region. From that time until 1934, Russia was the leader in the market of global platinum suppliers, giving way first to Canada, and since 1954 to South Africa, which has the largest deposits of the metal.
ADDITION. Refining.
Refining is the process of obtaining high-purity precious metals. Refining of platinum metals is based on the separation of chemical compounds of these elements, due to the difference in some of their properties - solubility, volatility, reactivity. The raw materials are enriched sludge left over from copper and nickel production, obtained by dissolving scrap technical products containing precious metals, including spent catalysts. The sludge contains platinum metals, as well as gold, silver, copper, and iron. To remove silica and base metals, most technological schemes resort to melting the sludge with lead litharge and charcoal. In this case, the base metals contained in the sludge are oxidized by lead litharge to oxides, and the resulting lead concentrates silver, gold and platinum group metals. The resulting lead bead, also called werkbley, is subjected to cupellation - oxidative melting on a droplet - in a porous vessel made of bone ash, magnesite and Portland cement. In this case, most of the lead is oxidized and absorbed by the droplet material. After cupellation, the alloy is treated with sulfuric acid to remove silver. It now contains noble metals. The most important refining operation is the interaction with aqua regia (Fig. 6.9. Simplified scheme for refining precious metals), in which most of the gold, palladium and platinum are dissolved, and ruthenium, osmium, rhodium and iridium mainly remain in the sediment. To separate gold from platinum and palladium, apply iron sulfate to the solution, which leads to the release of gold in free form. Palladium and platinum, present in solution in the form of chlorides and chloride complexes, are separated based on the different solubilities of the salts. Boiling the sludge for many hours in aqua regia leads to a partial transition of other platinum metals into solution, so the platinum obtained according to this scheme contains impurities of rhodium and iridium. From the residue, insoluble in aqua regia, rhodium is isolated by fusion with sodium hydrogen sulfate. When the melt is leached, it goes into solution in the form of complex sulfates. Ruthenium, osmium and iridium, which are resistant to acid attack, are subjected to oxidative fusion with alkali. The solution obtained by leaching the melt contains ruthenates and osmates, and most of the iridium precipitates in the form of dioxide. The separation of ruthenium from osmium is based on the sublimation of their higher oxides with their capture in a solution of hydrochloric acid. In this case, ruthenium oxide is reduced and goes into solution, and osmium anhydride goes into the gas phase and partially escapes into the atmosphere. This is not surprising, since osmium is the least popular of the platinum metals. The exact refining scheme is selected for a specific raw material, depending on the percentage of various metals in it.
END OF ADDENDUM.
Due to its high melting point, platinum, unlike gold and silver, did not melt in a forge and could not be forged either cold or hot. Therefore, the metal did not find practical use for a long time; it was in demand only among counterfeiters, who mixed it with gold to increase its mass. It got to the point that the King of Spain in 1755 issued a decree according to which all platinum mined during the development of Colombian placers in Choco had to be carefully separated from gold and drowned in rivers. During the 43 years that the decree was in effect, up to four tons of precious metal were destroyed.
Russian engineers were first able to obtain an ingot of metal in 1826. To do this, grains of native platinum were dissolved in aqua regia and then precipitated in the form of a porous spongy mass, which was molded under a press at 1000 °C. At the same time, the metal acquired malleability and ductility. In Russia, from 1828 to 1845, platinum coins were minted, as well as medals and jewelry. Settings for diamonds and many other precious stones made of platinum look much more impressive than silver ones. Adding platinum to silver jewelry makes it heavier and more durable. “White gold” is widely used in jewelry - a silver-white alloy of palladium and gold in a ratio of 1: 5. Interestingly, gold does not mix with platinum in solid form; such an alloy is a mixture of solid solutions of platinum in gold and gold in platinum . As the percentage of platinum increases, the color of gold changes to grayish-yellow and silver-gray. Such alloys were used by Faberge jewelers.
Annual global consumption of platinum metals is estimated at 200 tons. Platinum is slightly more expensive than gold, while rhodium, iridium, ruthenium and osmium are several times more expensive than platinum. The cheapest of the platinum metals is palladium. It costs less than $4 per gram.
The most important areas of use of platinum metals are presented in the table
Table 6.4. Structure of consumption of platinum metals in%
It does not include osmium, the global annual production of which is only a few kilograms. Although hydrogenation catalysts developed on its basis are even more effective than platinum ones, and adding it to alloys greatly increases their wear resistance, osmium and its compounds have not yet found practical use due to their high cost.
Among the consumers of platinum, rhodium and palladium, the automotive industry is in first place, which widely introduces catalysts made on their basis that improve the afterburning of exhaust gases. The effectiveness of their use directly depends on the quality of gasoline - the high content of organic sulfur compounds in it leads to rapid poisoning of the catalyst and nullifies its effect. In reforming processes, platinum-rhenium alloys are used, in hydrogenation, as well as in the oxidation of ammonia to nitrogen oxide (II) and sulfur dioxide to sulfuric anhydride - platinized asbestos, in the production of synthetic acetaldehyde (Wacker process) - palladium (II) chloride. Rhodium compounds are used mainly in homogeneous catalysis. Among them, the best known is triphenylphosphine rhodium(I) chloride Rh(PPh 3) 3 Cl, often called Wilkinson's catalyst. In its presence, many hydrogenation processes occur already at room temperature.
Due to their high heat resistance and high thermo-EMF values, platinum metal alloys are used in the production of thermocouples for measuring high temperatures: platinum-rhodium thermocouples operate when heated to 1300 °C, and rhodium-iridium thermocouples - 2300 °C.
Chemical inertness and refractoriness make platinum and platinoids convenient materials for the manufacture of electrodes, laboratory glassware, and chemical reactors, for example, glass melting apparatus. Palladium is the main material for multilayer ceramic capacitors used in computers and mobile phones. In electrical engineering, platinum and palladium are used to apply protective coatings to electrical contacts and resistances, so they can be recovered from used electrical devices. Platinum drugs are used in chemotherapy of oncological tumor diseases.