The only substance remains liquid at temperatures up to 0 K. It crystallizes only under a pressure of 25 atm. has the lowest boiling point. at temperatures below 2.2 K, liquid helium exists as a mixture of two liquids, one of which has anomalous properties - in particular, superfluidity (viscosity is 10 billion times lower than that of water).
Helium is the second most abundant (after hydrogen) element in the universe. About 10% of it is the Sun (discovered in 1868). On earth, helium was found in 1895 in reaction gases when the mineral kleveite was dissolved in acids. The remaining noble gases were isolated from the air.
Neon is a light gas: it is 1.44 times lighter than air, almost 2 times lighter than argon, but 5 times heavier than helium. According to the complex of properties, it is closer to helium than to argon. The spectrum of neon is rich: more than 900 lines are distinguished in it. The brightest lines form a beam in the red, orange and yellow parts of the spectrum at wavelengths from 6599 to 5400 Ǻ. These rays are much less absorbed and scattered by air and particles suspended in it than the rays of short waves - blue, blue, violet.
In 1898, in the Old World, when using a spectroscope to study the first portions of gas evaporating from liquid air, the Scottish chemist William Ramsay (Ramsey), together with Morris William Traver, discovered in them a new gas, Neon (Ne 6), an inert gas contained in air in microscopic quantities.
Argon is a monatomic gas with a boiling point (at normal pressure) of -185.9°C (slightly lower than that of oxygen, but slightly higher than that of nitrogen), melting point of -189.3°C In 100 ml of water at 20°C 3.3 ml of argon dissolves; in some organic solvents, argon dissolves much better than in water.
Discovered by J. Rayleigh and the English physicist W. Ramsay in 1894 from the air. The gas was distinguished by a monatomic composition of molecules and almost complete chemical inactivity (argon does not enter into any chemical reactions). new gas and got its name (Greek argos inactive).
Krypton is an inert monatomic gas without color, taste or smell. 3 times heavier than air. equal to 3.74 g/l. Opened in 1898 by W. Ramsay (England) Application: for filling incandescent lamps. Krypton compounds are oxidizers and fluorinating agents in chemical synthesis reactions.
Xenon is an inert monatomic gas without color, taste or smell. Tm 112 °C, Tk t 108 °C, glow in the discharge in purple. In 1889, the English scientist Wu Ramsay isolated a mixture from liquid air, in which two gases were discovered by the spectral method: krypton (“hidden”, “secret”) and xenon (“alien”, “unusual”).
Radon is a radioactive monatomic gas, colorless and odorless. Solubility in water 460 ml/l; in organic solvents, in human adipose tissue, the solubility of radon is ten times higher than in water. Radon's own radioactivity causes it to fluoresce. Gaseous and liquid radon fluoresces with blue light. The color of the glow in the gas discharge of radon is blue.
Colorless crystals, soluble in water. The molecule is linear. A solution in water is a very strong oxidizing agent, especially in an acidic environment, where it oxidizes bromine and manganese to the highest oxidation states of +7. In an alkaline environment, it hydrolyzes according to the equation: XeF 2 + 4KOH \u003d 2Xe + 4KF + O 2 + 2H 2 O
When interacting with water, XeF 4 disproportionates: 6XeF H 2 O \u003d 2XeO HF + 4Xe + 3O 2
It is formed during the hydrolysis of XeF 4. This is a white, non-volatile, highly explosive substance, highly soluble in water, and the solution has a slightly alkaline reaction. Under the action of ozone on such a solution, a salt of xenonic acid is formed, in which xenon has an oxidation state of +8: XeO 3 + O 3 + 4NaOH \u003d Na 4 XeO 6 + O H 2 O
It can be obtained by reacting barium perxenate with anhydrous sulfuric acid at low temperatures: Ba 2 XeO 6 + 2H 2 SO 4 \u003d 2 BaSO 4 + XeO H 2 O XeO 4 is a colorless gas that is very explosive and decomposes at temperatures above 0 ° C : 3XeО 4 = 2XeО 3 + Xe + 3О 2
The subgroup consists of 9 elements and is unique in this sense in the Periodic Table. Another unique property of this group is that the elements of this subgroup do not reach the highest oxidation state (with the exception of Ru and Os). It is generally accepted to divide 9 elements into 4 families: the triad of iron and the dyads Ru-Os, Rh-Ir, Pd-Pt. This division is justified by the kainosymmetry of the 3d sublevel of the elements Fe, Co, and Ni, as well as by the lanthanide contraction of Os, Ir, and Pt.
Chemistry of the elements of the iron triad Simple substances
Iron is the fourth most abundant on Earth, but most of it is in a state unsuitable for industrial use (aluminosilicates). Only ores based on iron oxides FeO and Fe 2 O 3 are of industrial importance. Cobalt and nickel are rare elements, which, although they form their own minerals, are industrially extracted from polymetallic ores.
Obtaining elements is reduced to restoring them from oxides. Derivatives of carbon (coke, CO) are used as a reducing agent, so the resulting metal contains up to several percent carbon. Iron containing more than 2% carbon is called cast iron; this material is well suited for casting massive products, but its mechanical strength is low. By burning carbon in open-hearth furnaces or converters, steel is obtained, from which mechanically strong products can be obtained. The dependence of material properties on the method of its production and processing is especially clear for iron: the combination of quenching and tempering makes it possible to obtain materials with different properties.
Obtaining Co and Ni is a complex process. At the final stage, metal oxides (CoO, Co 2 O 3 , NiO) are reduced with coal, and the resulting metal is purified by electrolysis.
The properties of simple substances strongly depend on the presence of impurities of other elements in them. Pure compact metals are stable in air at ordinary temperatures due to the formation of a strong oxide film, especially Ni. However, in the highly dispersed state, these metals are pyrophoric; self-ignite.
When heated, Fe, Co, Ni react with basic non-metals, and the interaction of iron with chlorine is especially intense due to the volatility of the resulting FeCl 3, which does not protect the metal surface from oxidation. On the contrary, the interaction of Ni with fluorine practically does not occur due to the formation of a strong fluoride film; therefore, nickel equipment is used when working with fluorine.
Fe, Co, Ni do not form certain compounds with hydrogen, but they are able to absorb it in noticeable amounts, especially in a highly dispersed state. Therefore, iron family metals are good catalysts for hydrogenation processes.
Metals react well with non-oxidizing acids:
E + 2HCl ECl 2 + H 2
Oxidizing acids passivate metals, but with alkalis, the reaction does not proceed due to the basic nature of metal oxides.
e(0) compounds
This oxidation state is characteristic of carbonyls. Iron forms a carbonyl of the composition Fe(CO) 5 , cobalt forms Co 2 (CO) 8 , and nickel forms Ni(CO) 4 . Nickel carbonyl is especially easy to form (50 °C, atmospheric pressure), so it is used to obtain pure nickel.
E(+2) compounds
The stability of compounds in this oxidation state increases from Fe to Ni. This is due to the fact that an increase in the charge of the nucleus with the same size of the atom strengthens the bond between the nucleus and d-electrons, so the latter are more difficult to break off.
E(+2) compounds are obtained by dissolving metals in acids. Hydroxides E (OH) 2 precipitate when an alkali solution is added to aqueous solutions of salts:
ECl 2 + 2NaOH \u003d E (OH) 2 + 2NaCl
From this we can conclude that the salts of the metals under consideration are susceptible to hydrolysis by the cation. As a result of hydrolysis, various products are obtained, including polynuclear complexes, such as NiOH + ,.
By calcining E (OH) 2 without access to air, oxides can be obtained. Oxides and hydroxides are predominantly basic; ferrates (+2), cobaltates (+2) and nickelates (+2) are obtained only under harsh conditions, for example, by fusion:
Na 2 O + NiO = Na 2 NiO 2
E(+2) sulfides can be precipitated from aqueous solutions with Na 2 S or even H 2 S (unlike MnS, which is not precipitated with H 2 S), but these sulfides dissolve in strong acids, which is used in chemical analysis :
E 2+ + S 2– E 2 S, E 2 S + 2H + (ex.) E 2+ + H 2 S
Of the E(+2) compounds, only Fe(+2) exhibits noticeable reducing properties. So, all simple (non-complex) Fe(+2) compounds are oxidized by atmospheric oxygen and other strong oxidizing agents:
4Fe(OH) 2 + 2H 2 O + O 2 4Fe(OH) 3
10FeSO 4 + 2KMnO 4 + 8H 2 SO 4 5Fe 2 (SO 4) 3 + K 2 SO 4 + 2MnSO 4 + 8H 2 O
Compounds of cobalt (+2) and nickel (+2) are oxidized only by strong oxidizing agents, such as NaOCl:
E(OH) 2 + NaOCl + x H 2 O E 2 O 3 x H 2 O + NaCl
E(+3) compounds
Stable compounds in this oxidation state are produced by iron and, to some extent, cobalt. Of the Ni(+3) derivatives, only complex compounds are stable.
Hydroxides E (OH) 3 are obtained by the action of alkali on salt solutions or by oxidation of E (OH) 2:
FeCl 3 + 3NaOH = Fe(OH) 3 ↓ + 3NaCl
2Co(OH) 2 + H 2 O 2 = 2Co(OH) 3
This results in products containing a variable amount of water (not having a constant composition). Oxides are the final products of dehydration of hydroxides, however, it is not possible to obtain pure Co 2 O 3 and Ni 2 O 3 due to their decomposition into oxygen and lower oxide. For iron and cobalt, it is possible to obtain oxides of the composition E 3 O 4 , which can be considered as mixed oxides EOE 2 O 3 . On the other hand, E 3 O 4 are salts corresponding to the acidic function of E(OH) 3 hydroxides.
Fe 2 O 3 + Na 2 O 2NaFeO 2
The main functions of Fe (OH) 3 are much better expressed:
Fe(OH) 3 + 3HCl FeCl 3 + 3H 2 O
Due to the fact that Fe (OH) 3 is a weak electrolyte, Fe (+3) salts are subject to hydrolysis. The hydrolysis products color the solution in a characteristic brown color, and when the solution is boiled, Fe (OH) 3 precipitates:
Fe 3+ + 3H 2 O Fe (OH) 3 + 3H +
It is not possible to obtain simple salts Co (+3) and Ni (+3), which correspond to the main function of E (OH) 3 hydroxide: in an acidic environment, redox reactions occur with the formation of E (+2):
2Co 3 O 4 + 12HCl 6CoCl 2 + O 2 + 6H 2 O
Compounds Co(+3) and Ni(+3) can only be oxidizing agents, and rather strong ones, while iron(+3) is not among the strong oxidizers. Nevertheless, it is not always possible to obtain E(+3) salts with a reducing anion (I – , S 2–). For instance:
2Fe(OH) 3 + 6HI 2FeI 2 + 6H 2 O + I 2
Unlike cobalt and nickel, iron gives Fe (+6) derivatives, which are obtained by hard oxidation of Fe (OH) 3 in an alkaline medium:
2Fe(OH) 3 + 3Br 2 +10KOH 2K 2 FeO 4 + 6KBr + 8H 2 O
Ferrates(+6) are stronger oxidizers than permanganates.
The elements of the eighth (iron, ruthenium, osmium, gassium), 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 included in it (ruthenium, osmium, rhodium, iridium, palladium, platinum) are noble, often found together in the form of alloys in which platinum predominates, so they are usually combined into a family of platinum metals (platinoids). Likewise, iron, cobalt and nickel are sometimes treated as a separate triad (iron triad). With some unconditional similarity of platinum metals, the chemistry of elements included in 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) ammoniates, rhodium(III) and iridium(III). Therefore, the chemical properties of oxygen-containing and complex compounds are described in the textbook by groups. Elements of the seventh period gassium, meitnerium and darmstadtium are radioactive with a short half-life and are obtained only in the amount of several tens of atoms.
Iron is one of the seven metals of antiquity, that is, known to mankind from the earliest periods of the history of society. Although the Egyptians and Phoenicians already knew the ability of cobalt compounds to give glasses a bright blue color, 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 the Swedish metallurgist A.F. Cronstedt isolated nickel from copper ore. Platinum is traditionally considered the metal of the Indians of Ecuador, as it was used by them to make jewelry and ritual masks before the arrival of the conquistadors. The infusible metal, outwardly similar to silver, received from the Spaniards the name platina, a diminutive of the word "silver". For a long time, the metal did not find any use due to its high hardness and refractoriness. For the first time, the English chemist W. Wollaston managed to obtain malleable platinum in 1805, who improved the process of hot forging. He is 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. Iridium (from the Latin iris - rainbow, according to compounds that have a bright color of various colors) and osmium (from the Greek οσμη - smell, according to the sharp unpleasant smell of volatile tetroxide) were soon isolated from the powder remaining after processing raw platinum with aqua regia. In 1844, Klaus, professor of chemistry at Kazan University, isolated ruthenium, which he named after Russia, from the Ural ore sent to him for analysis.
The superheavy platinum metals are the radioactive gassium, 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
Gassium was named after the land of Hesse, where the city of Darmstadt is located, meitnerium - in honor of the Australian scientist Lise Meitner, who studied the reactions of fission of uranium nuclei, and darmshadtium in honor of Darmstadt. The name of the last element was approved by the IUPAC commission in 2003.
The elements of the eighth group have a common electronic configuration in the ground state (n – 1)d 6 ns 2 is broken in ruthenium due to "electron slip". 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, the configuration (n – 1)d 8 ns 2 is observed only in the nickel atom: in platinum in the ground state, one electron "breakthrough" occurs, and in 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 | 77 Ir | 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 change in the properties of elements of groups 8–10 when moving along a period and across a group obey the general laws discussed in Chapter 1. The first ionization energies in the eighth and ninth groups decrease when moving from 3d metal to 4d (Table 6.1.), which is due with an increase in the atomic radius and the removal of valence electrons from the nucleus. The further increase in E 1 on going to d-metals of the sixth period is explained by the screening effects associated with the filling of the 4f sublevel. The general pattern does not apply to the elements of the tenth group due to the significant stabilization of the d-orbitals of the nickel atom, caused by a double "breakthrough" of electrons.
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 a much higher chemical activity compared to platinum metals. Unlike them, iron, cobalt and nickel displace hydrogen from acid solutions and oxidize in air. They are not characterized by cluster compounds, which, if formed, often turn out to be unstable in air and in aqueous solution. In general, platinum metals can be considered as the least chemically active metals, due to the relatively low (compared to the d-elements at the beginning of the transition rows) atomic radius and the high degree of overlapping of d-orbitals. Of these, only osmium is able to directly interact with oxygen, and only palladium reacts with concentrated nitric acid. In general, platinum metals 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 along the period, as the number of valence electrons increases and their pairing, the d-sublevel stabilizes, and the stability of higher oxidation states decreases. So, 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 the oxidation state +3. The highest oxidation states are most stable in the elements of the eighth group - iron (+6), ruthenium (+8) and osmium (+8) (Footnote: There is information on the preparation of iron compounds in the oxidation state +8: See Kiselev Yu. Kopelev N. S., Spitsyn V. I., Martynenko L. I. Doklady AN SSSR, 1987, v. 292, no. 3, p. 628). These metals show the lowest oxidation states with compounds with π-acceptor ligands, for example, in carbonyls: K 2 , K. The value of the most stable oxidation state decreases monotonically as you move along the period: for iron, the most characteristic oxidation state is +3, cobalt in aqueous solutions exists predominantly in the +2 oxidation state, and in +3 complexes, nickel exclusively in the +2 oxidation state. This is consistent with the increase in the third ionization energies in the series Fe - Co - Ni (Table 6.1.). Ni 2+ ions are resistant to air oxygen oxidation at any pH, cobalt(II) salts are stable in acidic and neutral media, and are oxidized in the presence of OH ions, iron(II) is converted into iron(III) under the action of oxygen (E 0 ( O 2 /H 2 O) = 1.229 B, pH = 0, and 0.401 B, pH = 14) at any pH. The reducing activity of the metals of the triad also decreases when moving along the 3d-series (Table 6.2.).
Table 6.2. Standard electrode potentials M(III)/M(II) and M(III)/M(0) for iron triad elements
The change in oxidation states that are stable in aqueous solutions can be represented as a diagram:
Examples of compounds of elements of 8 - 10 groups with different degrees of oxidation are given in table. 6.3. For ions with electronic configurations d 3 (Ru +5), d 5 (Fe +3,) and d 6 (Fe +2, Co +3, Rh +3, Ir +3), octahedral complexes are characteristic, for configurations d 4 ( Ru +4, Os +4) and d 7 (Co +2) are tetragonally distorted octahedral, arising due to the Jahn-Teller effect, for d 8 - octahedral (Ni +2 with weak and medium field ligands) - or square planar ( Pd +2 , Pt +2 , and also Ni +2 with strong field ligands). Molecules and ions with tetrahedral geometry arise when metal ions interact with bulk ligands (PR3, Cl–, Br–, I–) or when the d sublevel is completely filled (d10, Pd0, Rh–1, Ru–2).
A successive decrease in atomic and ionic radii as one moves along 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 analogs of iron - ruthenium and osmium also rarely increase the coordination number more than six. For platinum(II) and palladium(II), which have the d 8 electronic configuration, square planar complexes with a coordination number of 4 are most characteristic.
Another consequence of a decrease in ionic radii is a certain 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 give basic salts under these conditions. An increase in the Pearson softness of 3d-metal cations as they move along the period as the d-sublevel is filled and the ionic radii decrease causes the strengthening of the M-S bond in comparison with 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 occur 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 (C.N.) 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 | |||
| d9 | trigonal bipyramid | –1 | 2– | – | +1 | – | |||
| – | Co 2 (CO) 8, M 4 (CO) 12, M = Rh, Ir | – | |||||||
| octahedron | |||||||||
| d8 | octahedron | – | +1 | – | +2 | 2+ , 3+ | |||
| trigonal bipyramid | , | 3– | |||||||
| – | – | 2– | |||||||
| tetrahedron | |||||||||
| – | RhCl(PPh 3) 2 | 2–. 2– , M = Pd, Pt | |||||||
| square | |||||||||
| d7 | octahedron | +1 | + | +2 | 2+, Rh 2 (CH 3 COO) 4 (H 2 O) 2 | +3 | 3– , M = Ni, Pd | ||
| tetrahedron | – | 2– | – | ||||||
| d6 | tetrahedron | +2 | 2– | +3 | 5– | +4 | 2– , 2– , M = Pd, Pt | ||
| octahedron | 2+ , 4– | 3+ | – | ||||||
| d5 | tetrahedron | +3 | – | +4 | – | +5 | – | ||
| octahedron | 3+ , 3– | 2– , 2– , M = Co, Rh | – | ||||||
| d4 | tetrahedron | +4 | +5 | – | +6 | PTF 6 | |||
| octahedron | 2– , M = Ru, Os | – , M = Rh, Ir | – | ||||||
| d3 | tetrahedron | +5 | 3– , – , M = Ru, Os | +6 | MF 6 , M = Rh, Ir | ||||
| d2 | tetrahedron | +6 | 2– , 2– , | – | |||||
| d1 | tetrahedron | +7 | – , M = Ru, Os | – | |||||
| octahedron | OsOF 5 | – | |||||||
| pentagonal bipyramid | OSF 7 | – | |||||||
| d0 | tetrahedron | +8 | MO 4 , M = Ru, Os | – |
ADDITION. Biochemistry of iron.
Although the body of an adult contains only about 4 g of iron, it plays an important role in the processes of oxygen transfer to tissues and cells, the removal of carbon dioxide, and oxidative phosphorylation. Three-quarters of the iron atoms in the body are in the form of hemoglobin, which is made up of a porphyrin complex of iron called heme and the protein globin. Hemoglobin provides oxygen transport to the tissues of the body, and the 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 found 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, 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 cycle. The nitrogen atom of the imidazole ring of the histidine amino acid belonging 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 comes in. When oxygen is added, the iron atom leaves the plane of the porphyrin cycle 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 pass 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 tissues to the lungs. Immediately after the release of oxygen, it binds approximately 15% of the CO 2 dissolved in the blood.
The CO molecule is able to form a stronger complex with heme than the oxygen molecule, thereby preventing its transport from the lungs to the tissues. That is why inhalation of carbon monoxide leads to death from lack of oxygen. The 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 occurring in mitochondria, which is formed during the primary oxidation of carbohydrates. The energy released in this process is stored in the form of high-energy bonds of the ATP molecule. In a complex chain of oxidative phosphorylation, cytochromes a, b, and c are electron carriers 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 heme in hemoglobin by 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 450surrounded by the protein part of the molecule) . Its role is to hydroxylate lipophilic compounds alien to the body, which are formed as by-products or enter 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 is then (step 2) subjected to reduction by another enzyme. The third stage is the addition of oxygen, similar to that described above for hemoglobin. In the resulting low-spin iron complex, 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 restored, the oxidized substrate is separated, and the cytochrome passes into its original state (stage 6).
Heme also forms the basis of catalases and peroxidases, enzymes that catalyze oxidation reactions with hydrogen peroxide. One catalase molecule per second is capable of causing the decomposition of 44,000 H 2 O 2 molecules.
In oxidative phosphorylation, along with cytochromes, ferredoxins are involved - 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. The structure of bacterial ferredoxin (a), the active center of ferredoxin (b)). Ferredoxins found in bacteria, containing eight iron and sulfur atoms each, 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 p. 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. Meyes, W. Rodwell, Human Biochemistry, M., Mir, 1993).
END OF SUPPLEMENT
6.2. Distribution in nature, production and use of simple substances of 8-10 groups.
In terms of prevalence in nature among the elements of groups 8-10, the undisputed leader is iron, more precisely, its isotope 56 Fe, the nuclei of which have the highest binding energy of protons and neutrons, and, therefore, are highly stable.
Indeed, the number of iron atoms in the Universe significantly exceeds the number of atoms of any of the neighboring elements in the Periodic system and is close in order to hydrogen and helium. For example, on the Sun, the content of hydrogen is estimated at 1 × 10 12 conventional units, helium - at 6.31 × 10 10, and iron - at 3.16 × 10 17. This is due to the fact that the nucleus of the nuclide 56 Fe belongs to the number of magic, that is, having completely filled nuclear shells. As the number of nucleons in the nucleus increases, the binding energy per nucleon first 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 atomic number of the element). (R.J. Theiler, Origin chemical elements, M., Mir, 1975).
According to the 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%) - in the 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-bearing mineral pyrite FeS 2 (iron pyrites). In general, more than 300 iron-containing minerals are known. A significant amount of iron is part of various silicates and aluminosilicates that make up rocks. When weathered, iron compounds, mainly iron(III) oxide and oxohydroxide, enter quartz sand, clay, and soil, giving them a yellow-brown, earthy color. Iron of meteoric 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 interspersed in basalts. Only occasionally does it form separate pieces. Such finds are so rare that in the Stone and Bronze Age 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 with an admixture of 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 obtaining iron from iron ore was the beginning of the Iron Age. To reduce iron oxides with coal, a temperature of over 1400 °C is required, which an ordinary fire could not provide. That is why, in the early stages of the development of society, iron ores were not available as a raw material for the production of metal. People had to limit themselves to only random finds of meteoric iron. At the beginning of the first millennium BC. In the 18th century, a raw-working 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 injected through special clay tubes - nozzles - with the help of leather bags called furs. Charcoal and iron ore were poured into the furnace, and a fire was made on top. The resulting metal was welded into a kritsa - a porous mass, from which products were obtained by forging. Blast-furnace production has come to replace the raw-dough method. This happened 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 waste rock contained in the ore. Since in a blast furnace, unlike a hearth, the molten metal is in contact with coal for a long time, it carburizes, turning into cast iron. This requires an extra operation for the "redistribution" of 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 was developed in the Urals.
Addition. Diagram of the state 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 grades 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 in 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 a cubic body-centered and γ-Fe have a cubic face-centered lattice. The solubility of carbon is greatest in austenite (γ-Fe).
Melts containing up to 1.75 wt. % C, after rapid cooling to 1150 ° C, they are a homogeneous solid solution - austenite. Steel is made from these alloys. In melts containing more than 1.75% C after cooling to 1150 ° C, in addition to solid austenite, there is also a liquid eutectic of the composition of point A (Fig. 6.5.a) When cooled below 1150 ° C, it crystallizes and fills the space between austenite crystals. The resulting solid systems are cast iron. Depending on the conditions, the eutectic can crystallize in two ways. Upon rapid cooling, the solidified ectectics 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 graphite containing is called gray. The solidified eutectic from austenite and cementite is called ledeburite, and only ledeburite is released from the 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 the eutectoid point B. At this point, ferrite and cementite crystals precipitate simultaneously in a thin layered mixture called pearlite. A melt containing 0.9% C, upon cooling, can form pure pearlite, which does not contain large crystals of ferrite or Fe 3 C that have precipitated earlier.
By adjusting the composition of the initial melt, the cooling rate, and the heating time at the temperatures selected from the diagram, it is possible to obtain alloys with different microstructure, composition, orientation, and stresses in the crystals. If then the resulting system is cooled very quickly (quenched), then all further transformations are strongly 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
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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, cast iron. Cast iron used to make steel is called pig iron. Steel, unlike cast iron, contains less carbon. Excess carbon contained in cast iron must be burned off. This is achieved by passing oxygen-enriched air over the molten iron. There is also a direct method for producing iron, based on the reduction of magnetic iron ore pellets with natural gas or hydrogen:
Fe 3 O 4 + CH 4 \u003d 3Fe + CO 2 + 2H 2 O.
Very pure iron in powder form is obtained by decomposition of Fe(CO) 5 carbonyl.
ADDITION. Alloys of iron.
Iron-based alloys are divided into cast irons and steels.
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 that differ in properties and fracture color. White cast iron contains carbon in the form of cementite. It is highly brittle and does not find direct application. All white cast iron is converted into steel (pig iron). Gray cast iron contains graphite inclusions - they are clearly visible at the fracture. It is less brittle than white and is used to make flywheels and water heating radiators. The addition of a small amount of magnesium to the melt causes the precipitation of graphite 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 deoxidizer in steel production.
Fig.6.6. Gray cast iron (a) and heavy duty cast iron (b) under a microscope.
Iron ore and coke are the raw materials for iron production. Pig iron is smelted in blast furnaces - large furnaces, up to 80 m high, lined with refractory bricks from 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 mountain, and the upper hole, which serves to load the charge, is called the top. From below, hot air enriched with oxygen is fed into the furnace. In the upper part of the hearth, coal is burned with the formation of carbon dioxide. The heat released in this case is sufficient for the process to proceed. Carbon dioxide, passing through the 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, binding slag into low-melting fluxes (CaSiO 3, MgSiO 3). In addition to iron, coke also reduces the impurities contained in the ore, for example, phosphorus, sulfur, manganese, and partly silicon:
Ca 3 (PO 4) 2 + 5C = 3CaO + 5CO + 2P,
CaSO 4 + 4C \u003d CaS + 4CO,
MnO + C = Mn + CO
SiO2 + 2C = Si + 2CO.
In the molten metal, sulfur is present in the form of FeS sulfide, phosphorus is in the form of Fe 3 P phosphide, silicon is in the form of SiC silicide, and excess carbon is in the form of Fe 3 C carbide (cementite). The gases leaving the blast furnace are called blast furnace or blast furnace gases. About one third by volume they consist of carbon monoxide, so they are used as fuel to heat the air entering the blast furnace.
RICE. 6.7 Scheme of a blast furnace
Steel– an alloy of iron with 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 resilient, easier to break, although not as brittle as cast iron. The more carbon it contains, the harder it is. In ordinary 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 content of phosphorus in the steel causes cold brittleness - brittleness at low temperatures. Hardened steel is formed during the sharp cooling of steel heated to a red heat temperature. Such 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., receive special grades of steel. 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, in construction. Especially strong steels containing vanadium are used for armor casting.
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 the molten iron, which oxidizes the carbon contained in the 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 an oxide:
2Fe + O 2 \u003d 2FeO.
For the reverse reduction of FeO 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 \u003d MnSiO 3.
Steel is smelted in special furnaces. Depending on the type of furnaces, there are several methods of steelmaking. In an open hearth furnace, the melting space is a bath covered with a vault of refractory bricks (Figure 6.8. Steelmaking: (a) Open-hearth furnace, Oxygen converter). Fuel is injected into the upper part of the furnace - they are natural gas or fuel oil. The heat released during its combustion heats the mixture and causes it to melt. For 6 - 8 hours, during which the molten cast iron is in the open-hearth furnace, carbon gradually burns out in it. After that, the molten steel is poured out and after a while the cast iron is loaded again. The open-hearth process is periodic. Its main advantage is that the resulting steel can be poured into large molds. In terms of performance, 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 apparatus welded from steel and lined with refractory bricks from the inside. From above, oxygen-enriched air is blown through a converter mounted on a horizontal axis. The formed oxides of manganese and iron react with the silicate lining of the converter, forming slags. The process lasts about 40 minutes, after which the converter is moved to an inclined position and molten steel and slag are poured out sequentially (Fig. 6.8. b). Sand-lime brick-lined converters, called Bessemer converters after the English inventor Henry Bessemer, are not suitable for making steel from irons containing iron phosphides. For the redistribution of cast iron rich in phosphorus, Thomas converters are used, which are lined with limestone or dolomite from the inside. Steel smelting is carried out in the presence of lime, which binds the phosphorus contained in the cast iron into phosphates, which form slag (Thomas slag), which is used as a fertilizer. Alloy steels are smelted in electric furnaces at temperatures above 3000 °C. This makes it possible to obtain steels with special properties, including superstrong and refractory ones.
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Cobalt occurs in nature mainly in the form of compounds with arsenic, smaltite CoAs 2 (cobalt spice) and cobaltite CoAsS (cobalt luster), however, these minerals are too rare and do not form independent deposits. It is also a part of complex copper-cobalt-nickel and copper-cobalt sulfide ores; it is found in small amounts in clays and shales, which were formed under conditions of oxygen deficiency.
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, millerite NiS (yellow nickel pyrite), pentlandite (Fe, Ni) 9 S 8 , and chloantite NiAs 2 (white nickel pyrite) are of the greatest importance. Another important nickel raw material is serpentine rocks, which are basic silicates, for example, garnierite (Ni, Mg) 6 × 4H 2 O. Small amounts of nickel compounds are found 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 melted in a shaft furnace into a matte consisting of iron, cobalt, nickel and copper sulfides. This allows you to separate it from the 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 series FeS > CoS > Ni 3 S 2, therefore, iron sulfide first reacts with oxygen, which is converted into slag by the addition of silica. Further oxidation leads to the formation of oxides of cobalt and nickel
2Ni 2 S 3 + 7O 2 \u003d 6NiO + 4SO 2.
They are brought 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 \u003d 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 passes into solution in the form of a salt, which is converted into hydroxide. Oxides obtained by calcining hydroxides are reduced with coal:
Co 3 O 4 + 4C \u003d 3CO + 4CO,
NiO + C = Ni + CO.
During the reduction, carbides Co 3 C, Ni 3 C are also formed, to remove them, the oxide is taken in excess:
Ni 3 C + NiO = 4Ni + CO.
Electrolytic refining is used to obtain purer metals. It also makes it possible to isolate the platinum metals contained in the matte.
More than half of the produced cobalt and nickel is spent on the production of alloys. Cobalt-based magnetic alloys (Fe-Co-Mo, Fe-Ni-Co-Al, Sm-Co) are able to retain magnetic properties at high temperatures. Metal-ceramic alloys, which are titanium, tungsten, molybdenum, vanadium and tantalum carbides cemented with cobalt, are used to manufacture cutting tools. Steels with a high content of nickel and chromium do not corrode in air; they are used to make surgical instruments and equipment for the chemical industry. Heat-resistant chromium-nickel alloy nichrome, containing 20 - 30% chromium, has a high electrical resistance, it is used to make electric heater coils. Copper-nickel alloys constantan (40% Ni, 60% Cu) and nickelin (30% Ni, 56% Cu, 14% Zn), from monel (68% Ni, 28% Cu, 2.5 % Fe, 1.5% Mn) mint a coin.
Are important superalloys– materials based on iron, cobalt or nickel, specially designed for high temperature service. They have high corrosion resistance, retain strength in the temperature range at which gas turbines operate, are characterized by a high modulus of elasticity and a low coefficient of thermal expansion. The combination of oxidation resistance and strength of these materials is unrivaled. Many superalloys have a face-centered cubic lattice, which, being the densest of all crystal 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), which 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 turn into carbides of the composition M 23 C 6 and M 6 C, which are easily affected heat treatment. The resulting carbon passes into the form of a solid solution. Thus, the structure of a superalloy can be represented as a solid solution with finely crystalline inclusions of intermetallic compounds and carbides, which provide its hardness and strength. Additional doping contributes to slowing down diffusion processes, increasing the stability of the structure at high temperatures. One of the first superalloys was developed in 1935, Rex-78, consisting of 60% iron, 18% Ni, 14% Cr, and also containing small amounts of molybdenum, titanium, copper, boron, 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. The fine cobalt powder deposited on the support serves as an active catalyst for the Fischer-Tropsch hydrocarbonylation. Nickel often replaces platinum in hydrogenation processes, such as 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 goes to the production of 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 (“blue cobalt”, “Gzhel blue”) has a blue color, stannate Co 2 SnO 4 (“ceruleum”, “sky- blue") - blue with a bluish tint, phosphates Co 3 (PO 4) 2 ("cobalt purple dark") and CoNH 4 PO 4 × H 2 O ("cobalt purple light") - reddish-violet, mixed oxide of cobalt (II ) and zinc CoO × xZnO (“green cobalt”) - bright green, cobalt silicates (“schmalt”, “cobalt glass”) - dark blue (E.F. Belenky, 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 best known are ocher - crystalline oxohydroxide FeOOH and sienna containing clay. When calcined, they dehydrate, acquiring a red color. Brown umber is formed by the weathering of iron ores containing manganese. The black pigment is magnetite.
Platinum metals occur in nature mainly in their 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, RuS 2 laurite, PtS cooperite, are extremely rare. The average total content of platinum metals in the Ural sulphide 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 items, 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 gold particles, often turn out to be included in pyroxenes and other basic silicates, from which, as a result of erosion, they pass into river sands. The gold washed out of 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. Occasionally there are also large nuggets of platinum, up to nine kilograms. They necessarily contain impurities of iron, copper, platinum iodides, and sometimes gold and silver. For example, the metal from the Choco deposit in Colombia, which was developed 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 world suppliers of platinum, giving way first to Canada, and from 1954 to South Africa, which has the largest deposits of the metal.
ADDITION. Refining.
Refining is the production of high purity precious metals. The 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 material is enriched sludge left from copper and nickel production, obtained by dissolving scrap of technical products containing precious metals, including spent catalysts. Sludges contain platinum metals, as well as gold, silver, copper, and iron. To remove silica and base metals, most technological schemes resort to melting sludge with lead litharge and charcoal. In this case, 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 werkble, is subjected to cupellation - oxidative melting on a drop - 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. Now it contains noble metals. The most important refining operation is the interaction with aqua regia (Fig. 6.9. Simplified refining of noble metals), in which most of the gold, palladium and platinum are dissolved, while ruthenium, osmium, rhodium and iridium mainly remain in the sediment. To separate gold from platinum and palladium, iron sulfate is applied to the solution, which leads to the release of gold in a 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. Many hours of boiling the sludge 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 hydrosulfate. 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 trapping in a solution of hydrochloric acid. In this case, ruthenium oxide is reduced and goes into solution, while osmium anhydride goes into the gas phase and partially escapes into the atmosphere. This is not surprising, since osmium is the least sought after of the platinum metals. The exact refining scheme is selected for a specific raw material, depending on the percentage of various metals in it.
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Due to its high melting point, platinum, unlike gold and silver, did not melt in a furnace, could not be forged either cold or hot. Therefore, the metal did not find practical application for a long time, it was in demand only among counterfeiters, who mixed it with gold to increase its mass. Things 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 was to be carefully separated from gold and drowned in rivers. For 43 years of the decree, up to four tons of precious metal were destroyed.
For the first time, Russian engineers managed 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 pressure at 1000 °C. In this case, the metal acquired malleability and ductility. In Russia, from 1828 to 1845, platinum coins were minted, as well as medals and jewelry. Setting for diamonds and many other precious stones made of platinum looks much more impressive than silver ones. The addition of platinum to silver jewelry makes it heavier and more durable. Widespread use in jewelry is "white gold" - a silver-white alloy of palladium and gold in a ratio of 1: 5. Interestingly, gold does not mix with platinum in solid form, this alloy is a mixture of solid solutions of platinum in gold and gold in platinum . With an increase in the percentage of platinum, the color of gold changes to grayish yellow and silvery gray. Such alloys were used by Faberge jewelers.
The annual world 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 world's annual production of which is only a few kilograms. Although the hydrogenation catalysts developed on its basis are even more efficient than platinum ones, and its addition to alloys greatly increases their wear resistance, osmium and its compounds have not yet found practical application due to their high cost.
Among the consumers of platinum, rhodium and palladium, the automotive industry is in the first place, which widely introduces catalysts made on their basis that improve afterburning of exhaust gases. The efficiency of their use directly depends on the quality of gasoline - a high content of organic sulfur compounds in it leads to a rapid poisoning of the catalyst and reduces its effect to nothing. In reforming processes, platinum-rhenium alloys are used, in hydrogenation, as well as in the oxidation of ammonia to nitric oxide (II) and sulfur dioxide to sulfuric anhydride - platinized asbestos, in the production of synthetic acetaldehyde (Wacker process) - palladium (II) chloride. Rhodium compounds find application mainly in homogeneous catalysis. Among them, triphenylphosphinerhodium(I) chloride Rh(PPh 3) 3 Cl, often referred to as Wilkinson's catalyst, is best known. In its presence, many hydrogenation processes proceed even at room temperature.
Due to their high thermal stability and high thermal EMF values, platinum metal alloys are used in the production of thermocouples for measuring high temperatures: platinum-rhodium thermocouples operate at temperatures up to 1300 °C, and rhodium-iridium - 2300 °C.
The chemical inertness and refractoriness makes platinum and platinoids convenient materials for the manufacture of electrodes, laboratory glassware, chemical reactors, for example, glass melters. 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 on electrical contacts and resistances, so they can be removed from used electrical devices. Platinum preparations are used in chemotherapy of oncological tumor diseases.
Located in the fourth period.
The atomic weight of iron is 55.84, the nuclear charge is +26. Distribution of electrons by energy levels (+26): 2, 8, 14, 2. Electronic configuration of the outer and pre-outer iron layer 3s23p63d64s2.
Thus, for an iron atom, in addition to two s-electrons of the fourth outer layer, there are six more d-electrons of the third preexternal layer. Of these d-electrons are most active 4 unpaired. 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. In terms of prevalence among other elements, it ranks fourth.
■ 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 lustrous 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 ceases, retains its magnetic properties, becoming itself a magnet. All elements of the iron group have these properties.
According to its chemical properties, iron is a very active metal. In the absence of moisture, iron does not change in 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 iron gradually oxidizes in its entire mass:
4Fe + 2Н2О + 3О2 = 2Fe2O3 2H2O
A number of methods have been developed to protect this most valuable metal from corrosion.
In the series of voltages, iron is located to the left of hydrogen. In this regard, it is easily exposed to the action of dilute acids, turning into a ferrous 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 oxide film on the surface of the metal that the metal becomes completely passive and no longer enters into other reactions. At the same time, in direct interaction with such strong oxidizing agents as iron always exhibits an oxidation state of +3:
2Fe + 3Cl2 = 2FeCl3
Iron reacts with superheated steam; at the same time, it is displaced from the water, and the red-hot iron turns into oxide, and it is always either ferrous oxide FeO or ferrous oxide Fe3O4 (Fe2O3 FeO):
Fe + H2O = FeO + H2
3Fe + 4H2O = Fe3O4 + 4H2
Iron heated in pure oxygen burns vigorously with the formation of 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, a mixed oxide Fe3O4 is known, the molecule of which is two- and ferric iron: Fe2O3 · FeO. This oxide is also called iron oxide, or ferrous oxide.
Ferrous iron compounds are less stable than oxide-o, 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 impurity of oxidic compounds. Since it is very difficult to create such conditions in an ordinary educational laboratory, iron (II) hydroxide is obtained in the form of a more or less dark green gelatinous precipitate, which indicates the ongoing oxidation of ferrous compounds to ferric. If iron (II) hydroxide is kept in air for a long time, it gradually turns into iron (III) hydroxide Fe (OH) 3:
4Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3
iron are typical insoluble hydroxides. Iron hydroxide (II) 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 hydroxide (II). Support your answer with reaction equations.
Among the salts of iron (II), iron sulfate FeSO4 7H2O, which contains 7 molecules of water of crystallization, is of the greatest importance. Iron sulphate is highly soluble in water. It is used to control pests in agriculture, as well as in the manufacture of dyes.
Of the ferric salts, the most important is ferric chloride FeCl3, which is a very hygroscopic orange crystals that absorb water during storage and spread into a brown slurry.
Salts of iron (II) can easily turn into salts of iron (III), 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 in Fe +3 salts can also occur under the action of atmospheric oxygen during storage of these compounds, but only this process is longer. For the recognition of Fe 2+ and Fe 3+ cations, very characteristic specific reagents are used. For example, to recognize ferrous iron, they take red blood salt K3, which, in the presence of ferrous ions, gives them a characteristic intense blue precipitate of turnbull 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 precipitates. Prussian blue and turnbull 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 salts Fe +3 and Fe +2. Which oxidation state is more stable?
63. How to convert Fe +2 salt into Fe +3 salt and vice versa? Give examples.
The reaction goes according to the equation:
FeCl3 + 3KCNS = Fe(CNS)3 + 3KCl
or in ionic form
Fe 3+ + 3CNS - \u003d 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 in the organic matter of food products. A lot of iron contain apples, eggs, spinach, beets. 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, to obtain iron (II) hydroxide and iron (III) hydroxide?
67. A solution containing chromium (III) chloride and iron (III) chloride was treated with an excess of alkali. The resulting precipitate was filtered off. What remains on the filter and what passes into the filtrate? Give an informed answer using the reaction equations in molecular, full ionic, and abbreviated ionic forms.
iron alloys
Iron is the basis of ferrous metallurgy, so it is mined in huge quantities. The new program for the full-scale 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 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, while steel contains less than 1.7%.
Of great practical importance are ferroalloys (an alloy of iron with silicon), ferrochrome (an alloy of iron with chromium), and ferromanganese (an alloy of iron with manganese). 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 in cast iron. Ferroalloys are mainly used in the "deoxidation" of steel and as alloying impurities.
Among cast irons, linear and pig irons are distinguished. Foundry iron is used for casting various parts, pig iron is melted down into steel, as it has a very high hardness and cannot be processed. Pig 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., impurities of other metals that give the steel more valuable properties. gives steel malleability, 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 give steel new properties. If a steel part is kept in beryllium powder at a temperature of 900-1000 °, the hardness of the steel and its wear resistance increase greatly.
Chrome-nickel or, as they are also called, stainless steels are resistant to corrosion. Sulfur and phosphorus impurities are very harmful to steel - they make the metal brittle.
■ 68. What are the most important iron known to you?
69. What is the main difference between steel and cast iron?
70. What are the properties of cast iron and what types of cast iron do you know?
71. What are alloy steels and alloying additives?
domain process
Pig iron is produced by reduction smelting in blast furnaces. These are huge structures thirty meters high, producing more than 2,000 tons of cast iron per day. The scheme of the blast furnace device is shown in fig. 83.
The upper part of the blast furnace, through which the charge is loaded, is called the top. Through the top of the charge
Rice. 83. Scheme of the device of a blast furnace.
falls into a long shaft of the furnace, expanding from top to bottom, which facilitates the movement of the loaded material from top to bottom. As the charge moves to the widest part of the furnace - steam - a series of transformations occurs with it, as a result of which cast iron is formed, flowing into the hearth - the hottest part of the furnace. This is where slag collects. Cast iron and slag are released from the furnace through special holes in the hearth, called tapholes. Through the upper part of the furnace, air is blown into the blast furnace, which supports the combustion of fuel in the furnace.
Consider the chemical processes that occur during the smelting of iron. The charge of a blast furnace, 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, FeCO3 siderite is used as iron ore, and sometimes FeS2, which, after firing in pyrite furnaces, turns into Fe2O3 cinder, which can be used in metallurgy. Such ore is less desirable due to the large admixture of sulfur. Not only cast iron, but also ferroalloys are smelted in a blast furnace. The fuel loaded into the furnace serves simultaneously to maintain a high temperature in the furnace and to recover iron from ore, and also takes part in the formation of an alloy with carbon. The fuel is usually coke.
In the process of iron smelting, 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 restored:
Fe2O3 + 3CO = 3CO2 + 2Fe
Together with iron-containing ore, waste rock impurities are sure to enter the furnace. They are very refractory and can clog a furnace that has been running continuously for many years. In order for the waste rock to be easily removed from the furnace, it is converted into a fusible compound, turning it into slag with fluxes (fluxes). To transfer to slag the base rock containing, for example, limestone, which decomposes in a furnace according to the equation
CaCO3 = CaO + CO2
add sand. Fusing with calcium oxide, sand forms a silicate:
CaO + SiO3 = CaSiO3
This is a substance with an incomparably lower melting point. In the liquid state, it can be released from the furnace.
If the rock is acidic, containing a large amount of silicon dioxide, then, on the contrary, limestone is loaded into the furnace, which converts silicon dioxide into silicate, and the same slag is obtained as a result. Previously, slag was a waste, but now it is cooled with water and used as a building material.
To maintain the combustion of fuel, heated, oxygen-enriched air is continuously supplied to the blast furnace. It is heated in special air heaters - kiupers. Cowper is a high tower built of refractory bricks, where hot gases from the blast furnace are discharged. Blast-furnace gases contain carbon dioxide CO2, N2 and carbon monoxide CO. Carbon monoxide burns in the cowper, thereby raising its temperature. Then the blast-furnace gases are automatically sent to another cowper, and through the first one, the blowing of air directed to the blast furnace begins. In a hot cowper, the air is heated, and thus fuel is saved, which would be spent in large quantities 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 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?
steel smelting
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 main mass of steel (Fig. 84). An open-hearth furnace, unlike a blast furnace, is not a continuously operating furnace.
Rice. 84. Scheme of the device of an open-hearth furnace
Its main part is a bath, where the necessary materials are loaded through the windows by a special machine. The bath is connected by special passages to regenerators, which serve to heat the combustible gases and air supplied to the furnace. Heating occurs due to the heat of the combustion products, which are passed through the 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 composition of the charge includes pig iron, scrap metal, ore, fluxes (fluxes) of the same nature as in the blast furnace process. As in the blast furnace process, during steelmaking, air and combustible gases are heated in regenerators due to the heat of the exhaust 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 reduction 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 involved in this 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, turning it mainly into iron oxide (II): 2Fe + O2 \u003d 2FeO
Contained in cast iron, or any impurities of other metals at high temperature, reduce the resulting iron (II) oxide again to metallic iron according to the equation: Si + 2FeO \u003d SiO2 + 2Fe Mn + FeO \u003d MnO + Fe
Reacts similarly with iron oxide (II) and: C + FeO = Fe + CO
At the end of the process, to restore the remaining iron oxide (II) (or, as they say, to "deoxidize" it), "deoxidizers" - ferroalloys - are added. The additives of manganese and silicon present in them reduce the remaining iron oxide (II) according to the above equations. After that, the melting ends. Melting in open-hearth furnaces takes 8-10 hours.
Rice. 85. Bessemer converter device diagram
Bessemer converter (Fig. 85) - an older type furnace, but with very high performance. Since the converter works 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 from the inside with refractory bricks. Each heat in the converter lasts 12-15 minutes. The converter has a number of disadvantages: it can only work on liquid iron. This is due to the fact that the oxidation of carbon is carried out by air passing from below through the entire mass of liquid iron, which significantly accelerates 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 to regulate it, to add alloying impurities, therefore, mainly carbon steels are smelted in converters. At the end of the melt, the air supply is stopped and, as in the open-hearth process, “deoxidizers” are added.
In electric furnaces (Fig. 86), alloyed steel of special grades is smelted, mainly with a high melting temperature, containing and other additives. Finished steel is sent to rolling. There, on huge rolling mills - bloomings and slabs - they compress red-hot steel ingots with the help of rolls, which make it possible to produce various shapes from a steel ingot.
Fig 86. Scheme 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 branch of the national economy can do without it. In order to save ferrous metals, at present, as far as possible, they are trying to replace them with synthetic materials.
Machine tools and automobiles, airplanes and tools, fittings for reinforced concrete structures, tin for tin boxes and roofing sheets, ships and bridges, agricultural machines and beams, pipes and a whole range of household products are made from ferrous metals.
■ 78. What is the fundamental difference between the steelmaking process and the blast furnace process?
79. What furnaces are used for steel smelting?
80. What are regenerators in an open hearth furnace?
81. Specify the composition of the charge of an open-hearth furnace and its difference from the composition of the charge of a blast furnace?
82. What are "deoxidizers"?
83. Why is steel smelting called oxidative smelting?
84. How much steel containing 1% carbon can be obtained from 116.7 kg of cast iron containing 4% carbon?
85. How much ferromanganese containing 80% manganese will be required to “deoxidize” 36 kg of ferrous oxide?
Article on the topic of Iron, a secondary subgroup of group VIII
IRON AND ELECTRICITY The properties of steels are varied. There are steels designed for a long stay in sea water, steels that can withstand high temperatures and...
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Ministry of Health and Social Development of the Russian Federation
State educational institution of higher professional education
First Moscow State Medical University
named after I.M. Sechenov
Department of General Chemistry
Topic: “Chemistry of elements of group VIII B of the Periodic system of D.I. Mendeleev"
Completed by a student
Kirillova Anastasia Alexandrovna
2 courses of pediatric faculty
Lecturer: Garnova N.Yu.
Moscow 2015
General characteristics of the chemical elements of group VIIIB
Group VIII of the secondary subgroup includes elements of the triads of iron (iron, cobalt, nickel), ruthenium (ruthenium, rhodium, palladium) and osmium (osmium, iridium, platinum). Most of the elements of the subgroup under consideration have two electrons in the outer electron layer of the atom; they are all metals. In addition to outer electrons, electrons from the previous unfinished layer also take part in the formation of chemical bonds. These elements are characterized by oxidation states equal to 2, 3, 4. Higher oxidation states are less common.
A comparison of the physical and chemical properties of the elements of the eighth group shows that iron, cobalt and nickel, which are in the first large period, are very similar to each other and at the same time are very different from the elements of the other two triads. Therefore, they are usually isolated in the iron family. The remaining six elements of the eighth group are combined under the general name of platinum metals.
The atoms of the iron triad at the last level contain 2 electrons, however, the number of electrons at the 3d sublevel is different for them. Fe - 3d 6.4s 2; Co-3d 7.4s 2; Ni -3d 8.4s 2 . The atoms of the elements of the iron family, in contrast to the atoms of platinum metals, do not have a free f-sublevel, therefore their properties are very different from the properties of elements of other triads (Ru, Rh, Pd and Os, Ir, Pt) In their compounds, Fe, Co, Ni exhibit the degree +2 and +3 oxidations. Iron can also exhibit an oxidation state of +6, cobalt +5, nickel +4. All elements of the three triads are strong complexing agents.
The history of the discovery of chemical elements of group VIIIB
subgroup of iron.
Iron as a tool material has been known since ancient times. The history of the production and use of iron dates back to the prehistoric era, most likely with the use of meteoric iron. Smelting in a cheese-blast furnace was used in the 12th century BC. e. in India, Anatolia and the Caucasus. The use of iron in the smelting and manufacture of tools and tools is also noted in 1200 BC. e. In Africa south of the Sahara.
Ruthenium was discovered by Kazan University professor Karl Klaus in 1844. Klaus isolated it from the Ural platinum ore in its pure form and pointed out the similarity between the triads of ruthenium - rhodium - palladium and osmium - iridium - platinum. He named the new element ruthenium in honor of Russia (Ruthenia is the Latin name for Russia).
Osmium was discovered in 1804 by the English chemist Smithson Tennant in the sediment left after dissolving platinum in aqua regia.
subgroup of cobalt.
Cobalt compounds have been known to man since ancient times, blue cobalt glasses, enamels, paints are found in the tombs of Ancient Egypt. In 1735, the Swedish mineralogist Georg Brand managed to isolate a previously unknown metal from this mineral, which he called cobalt. He also found out that the compounds of this particular element turn glass blue - this property was used even in ancient Assyria and Babylon.
Rhodium was discovered in England in 1803 by William Hyde Wollaston. The name comes from ancient Greek - rose, typical compounds of rhodium have a deep dark red color.
Iridium was discovered in 1803 by the English chemist S. Tennant simultaneously with osmium, which were present as impurities in natural platinum delivered from South America. The name (ancient Greek - rainbow) was due to the various colors of its salts.
Nickel subgroup.
Nickel was discovered in 1751. However, long before that, Saxon miners were well aware of the ore, which outwardly resembled copper ore and was used in glass making to color glass green.
Palladium was discovered by the English chemist William Wollaston in 1803. Wollaston isolated it from platinum ore brought from South America.
Platinum was unknown in Europe until the 18th century. For the first time, platinum was obtained in pure form from ores by the English chemist W. Wollaston in 1803. In Russia, back in 1819, in alluvial gold mined in the Urals, a “new Siberian metal” was discovered. At first it was called white gold, platinum was found on the Verkh-Isetsky, and then on the Nevyansk and Bilimbaevsky mines. Rich placers of platinum were discovered in the second half of 1824, and the following year, its mining began in Russia.
Distribution in nature
subgroup of iron.
Iron is the most common metal on the globe after aluminum: it makes up 4% (mass) of the earth's crust. Iron occurs in the form of various compounds: oxides, sulfides, silicates. Iron is found in the free state only in meteorites.
The most important iron ores include magnetic iron ore FeO·Fe2O3, red iron ore Fe2O3, brown iron ore 2Fe2O3* 3H2O and spar iron ore FeCO3.
Pyrite, or iron pyrite FeS2, which occurs in large quantities, is rarely used in metallurgy, since cast iron is obtained from it of very low quality due to the high sulfur content. Nevertheless, iron pyrite has the most important use - it serves as a feedstock for the production of sulfuric acid.
Ruthenium is the most common platinum metal in humans, but almost the rarest of all. Does not play a biological role. It is concentrated mainly in muscle tissue. Higher ruthenium oxide is extremely toxic and, being a strong oxidizing agent, can ignite flammable substances. Osmium may also exist in humans in imperceptibly small amounts.
In its native state, osmium occurs in the form of solid solutions with iridium, containing from 10 to 50% osmium. Osmium is found in polymetallic ores containing also platinum and palladium (sulfide copper-nickel and copper-molybdenum ores), in platinum minerals and waste from the processing of gold-bearing ores. The main minerals of osmium are natural alloys of osmium and iridium (nevyanskite and sysertskite) belonging to the class of solid solutions. Nevyanskite forms dense (c = 17000–22000 kg/m3) white or light gray lamellar hexagonal crystals with a hardness of 6–7 points on the Mohs scale. The content of osmium in nevyanskite can reach 21–49.3%.
subgroup of cobalt.
The mass fraction of cobalt in the earth's crust is 4·10?3%. In total, about 30 cobalt-containing minerals are known. The content in sea water is approximately (1.7) 10-10%. Rhodium is found in platinum ores, in some golden sands of South America and other countries. The content of rhodium and iridium in the earth's crust is 10–11%.
Cobalt is one of the trace elements vital to the body. It is part of vitamin B12 (cobalamin). Cobalt is involved in hematopoiesis, the functions of the nervous system and liver, enzymatic reactions. The human need for cobalt is 0.007-0.015 mg daily. The human body contains 0.2 mg of cobalt for every kilogram of human weight. In the absence of cobalt, acobaltosis develops.
Excess cobalt is also harmful to humans.
Rhodium and iridium are perhaps the rarest elements in the human body (until it is fully proven that they exist there at all).
Rhodium is a very rare and trace element. Only the isotope 103Rh occurs in nature. The average content of rhodium in the earth's crust is 1·10?7% by mass, in stony meteorites 4.8·10?5%. The rhodium content is elevated in ultramafic igneous rocks. It does not have its own minerals. Found in some of the golden sands of South America. It is found in nickel and platinum ores as a simple compound. Up to 43% of rhodium comes from Mexican gold deposits. It is also found in an isomorphic admixture of minerals of the osmic iridium group (up to 3.3%), in copper-nickel ores. A rare variety of osmic iridium, rhodium nevyanskite, is the richest mineral in rhodium (up to 11.3%).
Iridium is relatively common in meteorites. It is possible that the actual content of the metal on the planet is much higher: its high density and high affinity for iron (siderophilicity) could lead to the displacement of iridium deep into the Earth, into the core of the planet, in the process of its formation from the protoplanetary disk. A small amount of iridium has been found in the solar photosphere.
Nickel subgroup.
Nickel is quite common in nature - its content in the earth's crust is approx. 0.01% (wt.). It is found in the earth's crust only in bound form; iron meteorites contain native nickel (up to 8%). Its content in ultrabasic rocks is approximately 200 times higher than in acidic ones (1.2 kg/t and 8 g/t). In ultramafic rocks, the predominant amount of nickel is associated with olivines containing 0.13-0.41% Ni. It replaces iron and magnesium isomorphically. A small part of nickel is present in the form of sulfides. Nickel exhibits siderophilic and chalcophilic properties. With an increased content of sulfur in the magma, nickel sulfides appear along with copper, cobalt, iron, and platinoids. In a hydrothermal process, together with cobalt, arsenic and sulfur, and sometimes with bismuth, uranium and silver, nickel forms elevated concentrations in the form of nickel arsenides and sulfides. Nickel is commonly found in sulfide and arsenic-bearing copper-nickel ores.
nickeline (red nickel pyrite, kupfernickel) NiAs
chloantite (white nickel pyrite) (Ni, Co, Fe)As2
garnierite(Mg, Ni)6(Si4O11)(OH)6*H2O and other silicates
magnetic pyrites (Fe, Ni, Cu)S
arsenic-nickel gloss (gersdorfite) NiAsS,
pentlandite(Fe,Ni)9S8
In plants, on average, 5 × 10?5 weight percent of nickel, in marine animals - 1.6 × 10?4, in terrestrial animals - 1 × 10?6, in the human body - 1 ... 2 × 10?6. Much is known about nickel in organisms. It has been established, for example, that its content in human blood changes with age, that in animals the amount of nickel in the body is increased, and finally, that there are some plants and microorganisms - "concentrators" of nickel, containing thousands and even hundreds of thousands of times more nickel, than the environment.
Palladium and platinum in imperceptibly small amounts and without performing any role, according to some sources, exist in living organisms.
Natural platinum occurs as a mixture of six isotopes: 190Pt (0.014%), 192Pt (0.782%), 194Pt (32.967%), 195Pt (33.832%), 196Pt (25.242%), 198Pt (7.163%). One of them is weakly radioactive. The existence of very weak radioactivity of two more natural isotopes of platinum is predicted: alpha decay 192Pt? 188Os and double beta decay 198Pt? it has only been established that the half-lives exceed 4.7×1016 years and 3.2×1014 years, respectively.
General properties of the iron subgroup.
Change the size of the radii of atoms and ions.
Atomic radius Fe=0.126nm, Radius of Fe2+ ions =0.80E, Fe3+ =0.67E
Ru=0.134nm Ru=0.68E
Os=0.135nm Os=0.63E
Atomic radius Co=0.125nm Radius of ions Co2+=0.78E and Co3+=0.64E
Rh=0.1342nm Rh=0.68E
Ir=0.136nm Ir=0.625Е
Atomic radius Ni=0.124nm Ion radius Ni=0.69Е
Pd=0.137nm Pd=0.86E
Pt=0.138nm Pt=0.625E
Change in the ionization potential.
In periods, as a rule, the ionization potential increases from left to right.
Ionization potential at consecutive. transition from Fe0 to Fe5+ = 7.893; 16.183; 30.65; 57.79 eV respectively
Ionization potential of Ru0:Ru1+:Ru2+:Ru3+ resp. = 7.366; 16.763 and 28.46 eV
Ionization potential Os0 to Os2+ = 8.5 eV, 17 eV
Potential last. ionization of the cobalt atom 7.865, 17.06, 33.50, 53.2 and 82.2 eV
Ionization potential Rh0:Rh+ : Rh2+ : Rh3+ resp. 7.46, 18.077 and 31.04 eV
Ionization potential at last. transition from Ir0 to Ir5+ are equal respectively. 9.1, 17.0, (27), (39), (57) eV
Nickel ionization potential 7.635; 18.15; 35.17; 56.0 and 79 eV
Ionization potential of palladium Pd0 ? Pd+? Pd2+? Pd3+ resp. equal to 8.336, 19.428 and 32.92 eV
Ionization potential of platinum Pt0 ? Pt+? Pt2+? Pt3+ resp. equal to 9.0, 18.56 and 23.6 eV
Chemical element of group VIII of the periodic system, atomic number 26, atomic mass 55.847. The configuration of the two outer electron layers is 3s2p6d64s2. Iron exhibits a variable valency (the most stable compounds are 2- and 3-valent Iron). With oxygen, iron forms oxide (II) FeO, oxide (III) Fe2O3, and oxide (II,III) Fe3O4.
At high temperatures (700-900) it reacts with water vapor:
3Fe+4H2O? Fe3O4+4H2
Reacts with dilute acids HCl and H2SO4 to form iron salt two and hydrogen:
Fe+2HCl= FeCl2+H2
Fe+H2SO4(razb.)=FeSO4+H2
In concentrated oxidizing acids, iron dissolves only when heated and immediately passes into the Fe3 + cation:
2Fe+6H2SO4(conc)=Fe2(SO4)3+3SO2+6H2O
Fe+6HNO3(conc)=Fe(NO3)3+3NO2+3H2O
(in the cold, conc. nitric and sulfuric acids passivate iron)
In razb. With nitric acid, iron is oxidized both in the cold and when heated, deeply reducing it.
10Fe + 36HNO3diff. = 10Fe(NO3)3 + 3N2 + 18H20
In air in the presence of water, it easily oxidizes (rusts)
4Fe + 3O2 + 6H2O = 4Fe(OH)3
Iron reacts with halogens when heated. When iron and iodine 1 react, iodide Fe3I8 is formed.
2Fe+3Cl2=2FeCl3 when heated
When heated, iron reacts with nitrogen, forming iron nitride Fe3N, with phosphorus, forming phosphides FeP, Fe2P and Fe3P, with carbon, forming Fe3C carbide, with silicon, forming several silicides, for example, FeSi.
Qualitative reaction to iron (II) ion - reaction with red blood salt.
In the presence of iron(II) ions, a dark blue precipitate is formed. It is the Turnbull blue complex iron salt KFe).
2 K3 + 3 Fe SO4 \u003d KFe)? + 3K2SO4
A qualitative reaction to an iron (II) ion is a reaction with alkali.
A gray-green precipitate forms.
FeSO4+2 NaOH = Fe(OH)2? + Na2SO4
A qualitative reaction to the iron (III) ion is a reaction with alkali.
A brown precipitate indicates the presence of iron (III) ions in the initial solution.
FeCl3 + 3 NaOH = Fe(OH)3? + 3 NaCl
Qualitative reaction to iron (III) ion - reaction with yellow blood salt.
Yellow blood salt is potassium hexacyanoferrate K4.
3K4 +4 FeCl3 \u003d KFe)? + 12 KCl
A qualitative reaction to the iron (III) ion is a reaction with potassium thiocyanate.
A red substance is formed. It is iron(III) thiocyanate.
FeCl3 + 3 KCNS= Fe(CNS)3 + 3KCl
When interacting ions 4? with Fe3 + cations, a dark blue precipitate is formed - Prussian blue:
3Fe(CN)6]4? + 4Fe3+ = Fe43?
OXIDATION PROPERTIES OF IRON(III) COMPOUNDS
1. 2FeCl3 + H2S = 2FeCl2 + S + 2HCl
2. CuS + Fe2(SO4)3 = CuSO4 + 2FeSO4 + S
3. 2FeCl3 + 2KI = 2FeCl2 + I2 + 2KCl
4. 2FeCl3 + Na2SO3 + H2O = Na2SO4 + 2HCl + 2FeCl2
5. 2CuI + 2Fe2(SO4)3 = I2 + 2CuSO4 + 4FeSO4
6. SnCl2 + 2FeCl3 = SnCl4 + 2FeCl2
7. 2FeCl3 + Fe = 3FeCl2 8. Cu2S + 2Fe2(SO4)3 = 4FeSO4 + 2CuSO4 + S
iron compounds.
Iron pentacarbonyl is an inorganic compound, iron carbonyl complex of Fe(CO)5 composition. Light yellow liquid, immiscible with water. At elevated pressure, metallic iron reacts with carbon monoxide CO, and liquid, under normal conditions, easily volatile iron pentacarbonyl Fe (CO) 5 is formed. Iron carbonyls of compositions Fe2(CO)9 and Fe3(CO)12 are also known. Iron carbonyls serve as starting materials in the synthesis of organo-iron compounds, including ferrocene composition.
Receipt:
The action of pressurized carbon monoxide on iron powder:
Action of carbon monoxide under pressure on iron(II) and copper iodide:
Decomposition upon heating of nonacarbonyl iron:
Decomposes on heating:
Reacts with hot water:
Reacts with acids (in diethyl ether):
Oxidized by oxygen:
Reacts with hydroiodic acid:
When a solution in acetic acid is irradiated with ultraviolet light, higher carbonyls are formed:
Under the action of a sodium alcoholate catalyst on a solution in ethanol, higher carbonyls are formed:
Reacts with bases in methanol:
Reacts with sodium in liquid ammonia:
Reacts with nitrogen monoxide under pressure:
Iron oxide (II) FeO has basic properties, it corresponds to the base Fe (OH) 2. Iron oxide (III) Fe2O3 is weakly amphoteric, it corresponds to an even weaker than Fe (OH) 2 base Fe (OH) 3, which reacts with acids:
2Fe(OH)3+ 3H2SO4= Fe2(SO4)3 + 6H2O
Iron hydroxide (III) Fe(OH)3 exhibits slightly amphoteric properties; it is able to react only with concentrated alkali solutions:
Fe(OH)3 + KOH = K
The resulting iron(III) hydroxocomplexes are stable in strongly alkaline solutions. When solutions are diluted with water, they are destroyed, and iron (III) Fe(OH)3 hydroxide precipitates.
Of the salts of iron (II) in aqueous solutions, Mohr's salt is stable - double ammonium sulfate and iron (II) (NH4) 2Fe (SO4) 2 6H2O.
Iron (III) is able to form double sulfates with singly charged alum-type cations, for example, Kfe (SO4) 2 - iron-potassium alum, (NH4) Fe (SO4) 2 - iron-ammonium alum, etc.
Obtained by cooling an acidified mixture of saturated solutions of ferrous sulfate and ammonium:
In medicine, they are used as an astringent, cauterizing, hemostatic agent; as an antiperspirant.
Ferrous compounds.
Iron oxide (II) FeO. Diamagnetic black unstable crystalline powder. Transforms into when heated in air. Slightly soluble in water and alkalis. Soluble in acids. Decomposes water when heated. Obtained by oxidation of metallic iron, reduction of iron (III) oxide with CO or hydrogen, calcination of a mixture of Fe2O3 and iron powder.
Iron hydroxide (II) Fe (OH) 2. It is formed in the form of a flaky yellowish-white precipitate when solutions of iron (II) salts are treated with alkalis without air access. Slightly soluble in alkalis. Soluble in acids. Shows basic properties. In the presence of oxidizing agents, it instantly turns into Fe(OH)3.
Iron sulfate (II) FeSO4. Toxic, highly hygroscopic, paramagnetic, white, orthorhombic crystals. When heated in air, it turns into Fe2O3. Obtained by calcining pyrite, heating PbSO4 with iron, dehydration of FeSO4.7H2O crystalline hydrate.
In medicine, it is used as a drug for the treatment and prevention of iron deficiency anemia.
Iron (II) orthophosphate Fe3(PO4)2.8H2O. It occurs naturally as the mineral vivianite. Bluish-white monoclinic crystals. Soluble in mineral acids.
Iron (II) carbonate FeCO3. It occurs naturally as the mineral siderite or iron spar. Slightly soluble in water, soluble in mineral acids and sodium bicarbonate solutions. Oxidizes in moist air. Decomposes on heating into FeO and CO2. It is reduced by hydrogen when heated. Obtained by treating solutions of iron (II) salts with solutions of sodium carbonate or bicarbonate.
Iron hexacyanoferrate (II) K4.3H2O. Diamagnetic yellow monoclinic crystals, non-toxic, salty and bitter taste. Soluble in water, ethylamine, acetone. Obtained by the action of KCN on Fe(CN)2. It is used for the manufacture of photographic paper, as a chemical reagent for the determination of iron, zinc, copper, uranium, methylene blue, and in the production of mineral dyes.
IRON(III) COMPOUNDS
Iron oxide (III) - Fe2O3 - a poorly soluble red-brown substance, exhibits amphoteric properties. Reacts with acids and fuses with alkalis and carbonates to form ferrites. Ferrites have strong magnetic properties and are used in electromagnets.
Fe2O3 + 6HCl = FeCl 3 + 3H2O;
Fe2O3 + 2NaOH = 2NaFeO2 + H 2 O;
Fe2O3 + Na2CO3 = 2NaFeO2 + CO2:
Fe2 O 3 + CaCO3 \u003d Ca (FeO2) 2 + CO2
Iron oxide (III) is obtained by calcination of iron hydroxide or salts.
2Fe(OH)3 = Fe2O3 + 3H2O
4Fe(NO3)3 = 2Fe2O3 + 8NO2 + O2
4FeCO3 + O2 = 2Fe2O3 + 4CO2
Iron (III) hydroxide - Fe (OH) 3 - a slightly soluble brown substance, obtained by the action of alkalis or carbonates on iron (III) salts, since Fe2O3 does not interact with water. Generally speaking, this hydroxide is a polymeric compound with a variable composition Fe2O3 .nH2O
FeCl3 + 3KOH = Fe(OH)3 + 3KCl
Fe(OH)3 is an unstable compound and gradually loses water, turning into Fe2O3. It has weak amphoteric properties (reacts with acids, reacts with alkalis in solutions or during fusion).
Fe(OH)3 = FeO(OH) + H2O;
2FeO(OH) = Fe2O3 + H2O
Fe(OH)3 + 3HCl = FeCl3 + 3H2O;
Fe(OH)3 + NaOH = NaFeO2 + 2H2O;
Fe(OH)3 + NaOH = Na - sodium tetrahydroxyferrite
Salts of iron (III)
FeCl3 .6H2O - etching of radio circuit boards, water treatment Fe2(SO4)3 *9H2O - coagulant in water treatment
Fe(NO3)3 - mordant for dyeing fabrics
Phosphates and sulfides of iron (III) are insoluble in water. Salts of iron (III) are hydrolyzed and their solutions are acidic:
FeCl3 + Na2HPO4 + CH3COONa = FePO4 + 2NaCl + CH3COOH
2FeCl3 + 3(NH4)2S \u003d Fe2S3 + 6NH4Cl (elemental sulfur is released when heated or acidified)
FeCl3 + H2O = Fe(OH)Cl2 + Hcl;
Fe(OH)Cl2 + H2O = Fe(OH)2Cl + Hcl
Fe(OH)2Cl + H2O = Fe(OH)3 + Hcl;
2FeCl3 + 3Na2CO3 + 3H2O = 2Fe(OH)3 + 6NaCl + 3CO2
IRON(VI) COMPOUNDS
Salts of iron acid - ferrates are obtained by oxidation of iron (III) compounds.
2Fe(OH)3 + 10KOH + Br2 = 2K2FeO4 + 6KBr + 8H2O
Fe2O3 + 3KNO3 + 2K2CO3 = 2K2FeO4 + 3KNO2 + 2CO2
Fe2O3 + 3KNO3 + 4KOH = 2K2FeO4 + 3KNO2 + 2H2O
Fe2O3 + KClO3 + 4KOH = 2K2FeO4 + KCl + 2H2O
Ferrates are strong oxidizers: CrCl3 + K2FeO4 = K2CrO4 + FeCl3
2NH4OH + 2K2FeO4 = N2 + 2Fe(OH)3 + 4KOH
Iron is also included in the heme structure. Heme is a prosthetic group of many proteins: hemoglobin, myoglobin, cytochromes of mitochondrial CPE, cytochrome P450 involved in microsomal oxidation. The enzymes catalase, peroxidase, cytochrome oxidase contain heme as a coenzyme.
Heme consists of ferrous ion and porphyrin. The structure of porphyrins is based on porphin. Porphin consists of four pyrrole rings linked by methene bridges. Depending on the structure of the substituents in the pyrrole rings, several types of porphyrins are distinguished: protoporphyrins, etioporphyrins, mesoporphyrins, and coproporphyrins. Protoporphyrins are the precursors of all other types of porphyrins.
The hemes of different proteins can contain different types of porphyrins. The iron in heme is in the reduced state (Fe+2) and is bound by two covalent and two coordination bonds to the nitrogen atoms of the pyrrole rings. When iron is oxidized, heme is converted to hematin (Fe3+). The largest amount of heme is contained in erythrocytes filled with hemoglobin, muscle cells with myoglobin, and liver cells due to the high content of cytochrome P450 in them.
The structure of porphin (A), protoporphyrin IX (B), and hemoglobin heme (C). Porphin is a cyclic structure consisting of four pyrrole rings linked by methane bridges.
Myoglobin is an oxygen-binding protein in skeletal and cardiac muscles. Myoglobin contains a non-protein part (heme) and a protein part (apomyoglobin).
Apomyoglobin is the protein part of myoglobin; the primary structure is represented by a sequence of 153 amino acids, which are arranged in 8 α-helices in the secondary structure. ?-Helices are designated by Latin letters from A to H, starting from the N-terminus of the polypeptide chain, and contain from 7 to 23 amino acids. To designate individual amino acids in the primary structure of apomyoglobin, either writing their serial number from the N-terminus (for example, His64, Phen138), or the letter?-helix and the serial number of this amino acid in this helix, starting from the N-terminus (for example, His F8 ).
The tertiary structure has the form of a compact globule (there is practically no free space inside), formed due to loops and turns in the region of non-coiled sections of the protein. The inner part of the molecule consists almost entirely of hydrophobic radicals, with the exception of two His residues located in the active center.
Cytochrome P450-dependent monooxygenases catalyze the breakdown of various substances through hydroxylation involving the electron donor NADH and molecular oxygen. In this reaction, one oxygen atom is attached to the substrate, and the second is reduced to water.
Enzymes of the cytochrome P450 family, unlike other hemoproteins, as a rule, having one type of activity and a strictly defined function, are quite diverse in functions, types of enzymatic activity, and often have low substrate specificity. P450s can exhibit both monooxygenase and oxygenase activity, and therefore are sometimes referred to as mixed-function oxidases.
* functional (as part of hemoglobin, myoglobin, enzymes and coenzymes);
* transport (transferrin, lactoferrin, mobilferrin);
*deposited (ferritin, hemosiderin);
* iron forming a free pool.
Electronic formula 4s24p64d75s1
Oxidation state: +2, +3, +4, +5, +6, +7, +8; valency: 2, 3, 4, 5, 6, 7, 8
Physical properties: silver-white brittle metal, tmelt=2250оС, tboil=4200оС, density 12.4 g/cm3
Prevalence in nature: content in the earth's crust 5.0.10-6% (mass.)
Main Mineral: Laurite RuS2
Obtaining: as a result of complex processing of ores, RuO4 is obtained, which is reduced with hydrogen
Chemical properties: inactive metal. In the series of voltages, it is after hydrogen. Does not react with acids. Reacts when heated with halogens. Reacts with oxidizing-alkaline mixtures.
Compounds of divalent ruthenium.
Ruthenium (II) hydroxide Ru(OH)2. A brown precipitate formed when a solution of ruthenium(II) chloride is treated with alkali. It is not very stable and easily oxidized, turning into Ru(OH)3.
Ruthenium(II) chloride RuCl2. Brown powder, slightly soluble in cold water, acids, alkalis and soluble in alcohol. Obtained by the action of chlorine on powdered ruthenium metal heated to 250°C.
Compounds of trivalent ruthenium.
Ruthenium (III) hydroxide Ru(OH)3. It is formed as a black precipitate during the treatment of solutions of ruthenium salts with alkalis or oxidation of ruthenium (II) hydroxide.
Ruthenium (III) chloride RuCl3. Brilliant brown-black crystalline powder. Slightly soluble in water and acids.
RuCl3+ 2KCl (conc.) + H2O = K2
With concentrated hydrochloric acid forms chlorocomplexes:
With concentrated solutions of alkali metal chlorides it forms complex salts:
When heated, it is oxidized by atmospheric oxygen:
Recovered with hydrogen:
Ruthenium (III) bromide RuBr3. Black crystalline powder. Obtained by direct interaction of the elements or by treatment of ruthenium (III) hydroxide with hydrogen bromide.
Ruthenium(III) iodide RuI3. A black crystalline powder that decomposes into its elements when heated to 127°C. Obtained by direct interaction of the elements when heated or by treating ruthenium (III) hydroxide with hydrogen iodide.
The reaction of ruthenium (III) chloride and potassium iodide:
Compounds of tetravalent ruthenium.
Ruthenium (IV) oxide RuO2. Very stable. Slightly soluble in water and alcohol. Soluble in acids. It is obtained by heating powdered ruthenium in oxygen or by calcining ruthenium (IV) sulfide in a stream of oxygen.
Decomposition on heating of ruthenium(VIII) oxide:
Reaction of ruthenium(VIII) oxide and hydrogen peroxide:
Ruthenium(IV) sulfide RuS2. It occurs naturally as the mineral laurite. It is obtained by heating a mixture of powdered ruthenium with sulfur or by the action of hydrogen sulfide on solutions of ruthenium (IV) salts.
Compounds of pentavalent ruthenium.
Ruthenium(V) fluoride RuF5. Transparent dark green crystals. Corrodes glass. Decomposed by water and reduced by iodine to ruthenium(III) fluoride. Obtained by the action of fluorine on ruthenium metal heated to 300°C.
4RuF5+10H2O =3RuO2+ RuO4 + 20HF
Compounds of hexavalent ruthenium.
Potassium ruthenate K2RuO4. Dark green tetrahedral crystals. Reduced with hydrogen to ruthenium(IV) oxide or ruthenium metal. Obtained by the action of an oxidizing-alkaline mixture on ruthenium metal powder.
Reacts with dilute acids:
Reacts with concentrated acids:
Reacts with chlorine:
Compounds of heptavalent ruthenium.
Potassium perruthenate KruO4. Black tetragonal crystals. Obtained by oxidation of potassium ruthenate with gaseous chlorine or liquid bromine.
Octavalent ruthenium compounds.
Ruthenium (VIII) oxide RuO4. Decomposes explosively into ruthenium (IV) oxide and oxygen when heated below the boiling point. Transforms into ruthenates under the action of alkalis. The vapors have a strong ozone smell and are toxic. Has oxidizing properties. Obtained by calcining ruthenium in oxygen at temperatures above 1000°C.
Electronic configuration ext. electron shells 5s25p65d66s2
Osmium is isolated from the enriched raw material of platinum metals by calcining this concentrate in air at temperatures of 800–900 °C. In this case, vapors of highly volatile osmium tetroxide OsO4 quantitatively sublimate, which are then absorbed by the NaOH solution.
By evaporating the solution, a salt is isolated - sodium perosmate, which is then reduced with hydrogen at 120 ° C to osmium:
Compounds of divalent osmium.
Osmium(II) oxide OsO. Grayish-black powder, slightly soluble in water and acids. Obtained by heating a mixture of osmium, osmium (II) sulfite and sodium carbonate in a stream of carbon dioxide.
Osmium(II) chloride OsCl2. Soluble in alcohol, ether and nitric acid. Obtained by heating osmium (III) chloride at 500°C and reduced pressure. chemical element iron cobalt nickel
Osmium(II) iodide OsI2. A green solid that forms when acid solutions of osmium(IV) salts are treated with potassium iodide.
Trivalent osmium compounds.
Osmium(III) oxide Os2O3. Dark brown powder (or copper-red flakes), slightly soluble in water. Obtained by reduction of OsO4 with metallic osmium by heating or by heating salts of osmium (III) with sodium carbonate in a stream of carbon dioxide.
Osmium(III) chloride OsCl3. Hygroscopic brown cubic crystals. Easily soluble in water and alcohol. It decomposes into OsCl2 and Cl2 above 500°C. The crystalline hydrate OsCl3.3H2O is known.
Compounds of tetravalent osmium.
Osmium (IV) oxide OsO2. Slightly soluble in water and acids. Obtained by heating finely dispersed metallic osmium in OsO4 vapor.
Osmium(IV) fluoride OsF4. Brown powder that decomposes with water or heat. It is obtained by passing fluorine over metallic osmium heated to 280 ° C.
Osmium(IV) chloride OsCl4. Brown-red needles, which are slowly converted to OsO2.2H2O by water. It is obtained by treating OsO4 with concentrated HCl or by cooling the brown-yellow vapors formed by passing chlorine over metallic osmium heated to 650-700 ° C.
Osmium(IV) sulfide OsS2. Black cubic crystals. Slightly soluble in water and alcohol. When heated in air, it turns into OsO4. Reduced with hydrogen to metal.
Compounds of valent osmium.
Osmium (VI) fluoride OsF6. Corrodes glass. Hydrolyzes with water. Vapors of osmium(VI) fluoride are colorless and toxic. Obtained together with OsF4 and OsF8 by heating osmium in a fluorine environment.
Potassium osmate K2OsO4. Violet octahedral crystals. When heated, it turns into OsO4. Decomposed by acids. Obtained by reduction with potassium nitrite or OsO4 alcohol in a solution of caustic potash.
Octovalent osmium compounds.
Osmium oxide (VIII) OsO4. It has a pungent odor. Soluble in water, alcohol, ether. The vapors are very toxic. Shows oxidizing properties. Used in many reactions as a catalyst.
Osmium(VIII) fluoride OsF8. Yellow crystals. Decomposes when heated above 225°C. Has oxidizing properties. Causes burns on contact with skin. Reacts with water to form OsO4 and hydrofluoric acid.
Osmium(VIII) sulfide OsS4. Dark brown powder. Obtained by passing hydrogen sulfide through an acidified solution of OsO4 or by the action of ammonium or sodium sulfides on alkaline solutions of OsO4.
Cobalt (lat. Сobaltum, Co) is a chemical element with atomic number 27, atomic mass 58.9332. The configuration of the two outer electron layers of the cobalt atom is 3s2 3p6 3d7 4s2 . Cobalt forms compounds most often in the oxidation state +2 (valency II), less often in the oxidation state +3 (valence III) and very rarely in the oxidation states +1, +4 and +5 (valencies, respectively, I, IV, V).
Simple substances of Co in powder form exhibit a fairly high activity with respect to acids. As a result of their interaction with acids, salts with an oxidation state of +2 are formed. Cobalt salts are colored pink due to the formation of the 2- aqua complex.
Co + 2HCl = CoCl2 + H2?
Co + H2SO4 = CoSO4 + H2?
3 Co + 8HNO3(sc.) = 3Co(NO3)2 + 2NO + 4H2O
Cold concentrated nitric acid passivates Co. When heated, the protective film is destroyed and the metal reacts with concentrated nitric acid:
Co + 4HNO3(conc.) = Co(NO3)2 + 2NO2? + 2H2O
With oxygen, cobalt forms EO oxides with basic properties. These oxides do not dissolve in water, do not interact with alkalis, but readily react with acids to form Co(II) salts. Co(II) salts are most often used for the synthesis of the corresponding hydroxides, for example: CoCl2 + NaOH = Co(OH)2? + NaCl
Upon receipt of cobalt (II) hydroxide from salts, a blue precipitate of poorly soluble basic salts Co(OH)nX2-n? xH2O and then pink hydroxide Co(OH)2. The appearance of blue coloration can also be explained by the formation of cobalt hydroxide with the composition 3Co(OH)2?2H2O, which is formed together with basic salts. With further addition of alkali as a result of dehydration and aging, it changes color from blue to pink. Cobalt(II) hydroxide shows little evidence of amphoterism with predominantly basic properties. It dissolves easily in acids (with the formation of Co (II) salts), and dissolution in alkali is very difficult. However, the presence of acidic properties of Co (OH) 2 is confirmed by the existence of a 2- hydroxo complex. Co (II) hydroxide is very slowly oxidized by atmospheric oxygen and turns into Co(III) hydroxide, colored brown: 4Co(OH)2 + O2 + 2H2O = 4Co(OH)3
In the presence of stronger oxidizing agents, such as hydrogen peroxide, the Co(II) oxidation process is much faster: 2Co(OH)2 + H2O2 = 2Co(OH)3
A qualitative reaction to the Co(II) ion is the formation of its yellow nitro complex.
CoCl2 + 7KNO2 + 2CH3COOH = K3 ? + NO + 2CH3COOK + 2KCl
Qualitative reaction for cobalt (II) cations Co2+
The peculiarity of these cations is the formation of complex salts with ammonia molecules - ammoniates: Co2 + + 4NH3 \u003d 2+ Ammonias color solutions in bright colors. Yellow-brown cobalt(II) ammonia is gradually oxidized by atmospheric oxygen to cherry-red cobalt(III) ammonia. In the presence of oxidizing agents, this reaction proceeds instantaneously.
The oxidation state (III) is unstable for cobalt, so Co(III) hydroxide exhibits oxidizing properties, even under the influence of such a weak reducing agent as the Cl- ion: 2Co(OH)3 + 6HCl = 2CoCl2 + Cl2? + 6H2O
Cobalt forms a large number of insoluble salts, many of which, such as phosphates, can be synthesized using exchange reactions in aqueous solutions:
3 CoCl2 + 4Na2HPO4 = 2 Co3(PO)4 + 8NaCl + HCl
Medium Co(II) carbonates cannot be obtained by adding alkali metal carbonate to solutions of their salts. Due to increased hydrolysis in the presence of carbonate ions, the formation of poorly soluble basic rather than intermediate carbonates occurs:
2CoCl2 + Na2CO3 + 2H2O ? (CoOH)2CO3? + 2NaCl + 2HCl
In the middle of the 20th century, cobalamin, vitamin B12, was isolated from the liver of animals, and only this vitamin contains cobalt - it contains only 4%, and it is there in active form.
Cobalt takes part in many processes in the body: it activates the process of hematopoiesis - thanks to it, red blood cells are produced in the bone marrow, iron is better absorbed, and the composition of the blood constantly remains normal. The intestinal microflora, "responsible" for the absorption of iron, needs cobalt - for these bacteria it is food, therefore, if cobalt is not enough, various types of anemia often develop; the process of blood circulation with a lack of cobalt also cannot proceed normally.
Taking part in metabolic processes, cobalt normalizes the activity of the endocrine system, activates the production of enzymes, and participates in the synthesis of proteins, carbohydrates and fats. Interacting with other substances, cobalt starts the process of renewal of all body cells, also participating in the synthesis of RNA and DNA.
For the normal development and preservation of the structure of bone tissue, it is also important that there is enough cobalt in the body, so products with cobalt are especially necessary for children, women and the elderly.
Cobalt is important for maintaining a healthy state of blood vessels - it prevents the development of atherosclerosis, as it not only reduces the amount of "bad" cholesterol in the blood, but also helps the body to remove it, so it does not have time to be deposited in the vessels.
The immunostimulatory effect of cobalt is manifested by its ability to increase the phagocytic activity of leukocytes - this means that leukocytes more actively bind, absorb and digest pathogens that enter the body. The activity of the pancreas also depends on the amount of cobalt in the body - if it is lacking, this organ cannot work normally.
Cobalt, together with other substances, helps the body to maintain youth - for example, together with copper and manganese, it prevents early graying of hair and speeds up recovery after serious illnesses.
The electronic formula is 4s24p64d85s1.
Oxidation state: +1, +2, +3, +4, +6; valency: 1, 2, 3, 4, 6
Chemical properties: inactive metal. It dissolves in aqua regia and in concentrated sulfuric and hydrochloric acids in the presence of oxygen. Oxidized by oxidizing-alkaline mixtures.
Rhodium(II) oxide RhO. Black-brown substance, slightly soluble in water and acids.
Rhodium (II) chloride RhCl2. Powder that can be dyed in various colors - from dark brown to purple-red. Obtained by heating rhodium (III) chloride at 950°C.
Rhodium(II) sulfide RhS. Dark gray crystals. Slightly soluble in water and aqua regia. Obtained by heating metallic rhodium to a red heat in sulfur vapor.
Trivalent rhodium compounds.
Rhodium (III) oxide Rh2O3. Slightly soluble in water, acids and aqua regia. It is reduced to metallic rhodium by hydrogen when heated. Obtained by heating powdered rhodium, rhodium (III) nitrate or rhodium (III) chloride in air at 800°C.
Rhodium hydroxide (III) Rh(OH)3. Yellow gelatinous precipitate obtained by treating rhodium(III) salts with alkalis or alkali metal carbonates. Soluble in acids or excess alkali. When dehydrated, it turns into Rh2O3.
Rhodium(III) fluoride RhF3. Red rhombic crystals. Obtained by passing fluorine over metallic rhodium heated to 500-600°C.
Rhodium (III) chloride RhCl3. Reddish-brown powder deliquescent in air. Poorly soluble in water and acids. Above 948°C it decomposes into elements. Known crystal hydrate RhCl3.4H2O. It is obtained by the action of chlorine on rhodium heated to 250-300 ° C or by dehydration of RhCl3.4H2O.
Rhodium (III) iodide RhI3. A black substance, poorly soluble in water. Decomposes at 327°C. It is obtained by the action of potassium iodide on solutions of rhodium (III) salts during boiling.
Rhodium(III) sulfide Rh2S3. Greyish-black powder, stable up to 500°C. Above this temperature, in air or oxygen, it ignites and burns to form rhodium metal. Obtained by the action of hydrogen sulfide on rhodium (III) chloride.
Compounds of tetravalent rhodium.
Rhodium (IV) oxide RhO2. Solid black substance. Obtained by fusing metallic rhodium with hydroxide and potassium nitrate.
Hydrated oxide of rhodium (IV) RhO2.nH2O. Olive green solid. Soluble in acids. Transforms into Rh2O3 when heated. It is obtained by electrolytic oxidation of Rh (OH) 3 in an excess of alkali or by oxidation of solutions of rhodium (III) salts with chlorine in an alkaline medium.
Rhodium (IV) bromide RhBr4. Brown powder. Decomposes into elements when heated to 527°C. Obtained by direct interaction of elements when heated.
Electronic formula 5s25p65d76s2
Oxidation state: +1, +2, +3, +4, +5, +6; valency: 1, 2, 3, 4, 5, 6
Physical Properties: Silvery white hard metal
Obtaining: as a result of complex processing of ores, (NH4) 2 is obtained, during the thermal decomposition of which iridium is obtained.
Chemical properties: inert metal. Does not interact with acids; does not react even with aqua regia. Reacts with oxidizing-alkaline mixtures. When heated, it interacts with halogens.
Compounds of divalent iridium.
Iridium(II) chloride IrCl2. Brilliant dark green crystals. Poorly soluble in acids and alkalis. Obtained by heating metallic iridium or IrCl3 in a stream of chlorine at 763°C.
Iridium(II) sulfide IrS. Brilliant dark blue solid. Soluble in potassium sulfide. Obtained by heating metallic iridium in sulfur vapor.
Compounds of trivalent iridium.
Iridium(III) oxide Ir2O3. Dark blue solid. Slightly soluble in water and alcohol. Dissolves in sulfuric acid. Obtained by light calcination of iridium (III) sulfide.
Iridium(III) chloride IrCl3. The volatile compound is olive green in color. Slightly soluble in water, alkalis and acids. Obtained by the action of chlorine on iridium heated to 600 ° C.
Iridium(III) bromide IrBr3. Olive green crystals. Soluble in water, slightly soluble in alcohol. Obtained by the interaction of IrO2 with hydrobromic acid.
Iridium(III) sulfide Ir2S3. Brown solid. Slightly soluble in water. Soluble in nitric acid and potassium sulfide solution. Obtained by the action of hydrogen sulfide on iridium (III) chloride.
Compounds of tetravalent iridium.
Iridium(IV) oxide IrO2. Black tetragonal crystals with a rutile lattice. Slightly soluble in water, alcohol and acids. It is reduced to metal with hydrogen. Thermally dissociates into elements when heated. Obtained by heating powdered iridium in air or oxygen at 700°C, heating IrO2.nH2O.
Iridium(IV) fluoride IrF4. Yellow oily liquid that decomposes in air and hydrolyzes with water. Obtained by heating IrF6 with iridium powder at 150°C.
Iridium(IV) chloride IrCl4. Hygroscopic brown solid. It dissolves in cold water and decomposes in warm water. Obtained by heating (600-700 ° C) metallic iridium with chlorine at elevated pressure.
Iridium(IV) bromide IrBr4. A blue substance that floats in the air. Soluble in alcohol. Obtained by the interaction of IrO2 with hydrobromic acid at low temperature.
Iridium(IV) sulfide IrS2. Brown solid. Slightly soluble in water. Obtained by passing hydrogen sulfide through solutions of iridium (IV) salts.
Compounds of hexavalent iridium.
Iridium(VI) fluoride IrF6. Yellow tetragonal crystals. Under the action of metallic iridium, it turns into IrF4, is reduced by hydrogen to metallic iridium. Corrodes wet glass. Obtained by heating iridium in a fluorine atmosphere in a fluorite tube.
Iridium(VI) sulfide IrS3. Gray powder, slightly soluble in water. Obtained by heating powdered metallic iridium with excess sulfur in a vacuum.
Electronic formula 3s23p63d84s2
Oxidation state: (+1), +2, (+3, +4); valency: (1), 2, (3, 4)
Physical properties: gray hard metal
Basic minerals: iron-nickel pyrite (Fe,Ni)9S8, nickeline NiAs
Chemical properties: inactive metal. Resistant to water and humid air. Reacts slowly with dilute acids. When heated, it reacts with oxygen, halogens, nitrogen, sulfur and other non-metals. Resistant at normal temperature to the action of fluorine, which is stored in cylinders made of nickel.
A qualitative reaction for nickel is the reaction of nickel with dimethylglyoxime. Dimethylglyoxime was proposed in 1905 by L.A. Chugaev, so dimethylglyoxime is sometimes called "Chugaev's reagent".
When reacting with nickel ions, dimethylglyoxime forms a red complex, which can be easily precipitated and determined gravimetrically.
Simple Ni substances in powdered form show a rather high activity with respect to acids. As a result of their interaction with acids, salts with an oxidation state of +2 are formed. Aqueous solutions of Ni salts are colored green due to the presence of the 2- ion.
Ni + 2HCl = NiСl2 + H2?
Ni + H2SO4 = NiSO4 + H2?
3Ni + 8HNO3(diff.) = 3Ni(NO3)2 + 2NO + 4H2O
Cold concentrated nitric acid passivates Ni. When heated, the protective film is destroyed and the metal reacts with concentrated nitric acid:
Ni+ 4HNO3(conc.) = Ni(NO3)2 + 2NO2? + 2H2O
With oxygen, nickel forms EO oxides with basic properties. These oxides do not dissolve in water, do not interact with alkalis, but readily react with acids to form Ni(II) salts. Ni(II) salts are most often used for the synthesis of the corresponding hydroxides, for example:
NiCl2 + NaOH = Ni(OH)2? + NaCl
Nickel(II) hydroxide, green in color, is similar in acid-base properties to Co(II) hydroxide. It is easily soluble in acids and practically insoluble in alkalis. Prolonged exposure of alkalis to the Ni(OH)2 precipitate leads to the formation of a hydroxocomplex of indeterminate composition with the conditional formula 2? . Nickel(II) hydroxide is not oxidized to Ni(OH)3 neither by atmospheric oxygen nor by hydrogen peroxide. To oxidize it, a stronger oxidizing agent is needed, for example, bromine:
2Ni(OH)2 + 2NaOH + Br2 = 2Ni(OH)3 + 2NaBr
Nickel in oxidation states +2 and +3 forms a large number of complex compounds. Their most stable cationic complexes are aqua complexes and ammoniates, as well as complexes where the ligands are polydentate organic molecules, for example, dimethylglyoximate. Nickel forms a large number of insoluble salts, many of which, such as phosphates, can be synthesized using exchange reactions in aqueous solutions:
3NiCl2 + 4Na2HPO4 = 2Ni3(PO)4 + 8NaCl + HCl
Medium Ni(II) carbonates cannot be obtained by adding alkali metal carbonate to solutions of their salts. Due to increased hydrolysis in the presence of carbonate ions, the formation of poorly soluble basic rather than intermediate carbonates occurs:
2NiCl2 + Na2CO3 + 2H2O ? (NiOH)2CO3? + 2NaCl + 2HCl
Nickel is a trace element that affects the processes of hematopoiesis and is involved in many redox processes in the body.
The body of an adult contains only about 5-14 ml of nickel. The element accumulates in muscle tissue, liver, lungs, kidneys, pancreas and thyroid glands, pituitary gland, brain and epithelium. It is noted that with age, the concentration of nickel in the lungs increases. Nickel is excreted from the body mainly with feces (up to 95%).
in combination with cobalt, iron, copper participates in the processes of hematopoiesis (affects the maturation of young red blood cells and increases the level of hemoglobin)
increases hypoglycemic activity (increases the effectiveness of insulin)
participates in the structural organization and functioning of DNA, RNA and proteins
enhances the passage of redox processes in tissues (provides cells with oxygen)
enhances the antidiuretic action of the pituitary gland
activates a number of enzymes (including arginase)
important for hormonal regulation of the body
involved in fat metabolism
oxidizes vitamin C
lowers blood pressure
daily requirement
The daily requirement for nickel, depending on age, gender and weight, is about 100-300 mcg.
Deficiency and overdose symptoms
Overdose of nickel is an infrequent phenomenon, which is explained by the high level of toxicity of the element - about 20-40 mg per day. With its excess, the following symptoms are observed:
tachycardia
dermatitis
decreased resistance to infectious diseases
irritation of the mucous membranes of the upper respiratory tract
increased excitability of the nervous system
decreased immunity
magnesium deficiency in the body
accumulation of iron or zinc
retardation of bone growth
edema of the lungs and brain
cancer risk
Palladium.
Electronic formula 4s24p64d10
Oxidation state: +1, +2, +3, +4; valency: 1, 2, 3, 4
Physical Properties: Silvery white soft metal
Main minerals: palladite PdO; also found in native form
Chemical properties: inactive metal. Stable in dry and humid air at normal temperatures. Soluble in concentrated nitric acid and aqua regia. When heated, it dissolves in concentrated sulfuric acid. When heated, it reacts with non-metals.
Monovalent palladium compounds.
Palladium(I) sulfide Pd2S. Greenish-gray amorphous substance. Slightly soluble in water, acids and aqua regia. Obtained by heating Cl2 with sulfur under a layer of borax.
Compounds of divalent palladium.
Palladium(II) oxide PdO. The most stable palladium oxide. Black powder. Slightly soluble in water and acids. Has oxidizing properties. Obtained by strong heating of metallic palladium in oxygen or by calcination of palladium (II) nitrate.
Palladium(II) hydroxide Pd(OH)2. Brown-red powder. Slightly soluble in water. Soluble in acids. Possesses weakly expressed oxidizing properties. It is obtained by hydrolysis of palladium (II) nitrate with hot water or by the action of alkalis on palladium (II) salts.
Palladium(II) fluoride PdF2. Brown crystals. Poorly soluble in water. It dissolves in hydrofluoric acid to form tetrafluoropalladic acid H2. Obtained by direct interaction of the elements or by heating PdF3 with palladium.
Palladium(II) bromide PdBr2. Brown-red substance. Slightly soluble in water. It dissolves in an aqueous solution of hydrogen bromide (with the formation of tetrabromopalladic acid H2) or alkali metal bromides. Obtained by the action of bromine water on palladium metal or by the action of potassium bromide on PdCl2.
Palladium(II) iodide PdI2. Dark red powder. Slightly soluble in water and alcohols. It dissolves in hydroiodic acid (with the formation of tetraiodopalladic acid H2) or solutions of alkali metal iodides. Obtained by treating a solution of palladium (II) chloride with potassium iodide.
Palladium(II) sulfide PdS. Dark brown solid metal-like substance. Slightly soluble in water, hydrochloric acid, ammonium sulfide. Soluble in nitric acid and aqua regia. Obtained by the interaction of elements during heating, thermal decomposition of PdS2, or by passing hydrogen sulfide through an aqueous solution of palladium (II) salts.
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