General characteristics of the elements of the 8th group of the secondary subgroup. Elements of subgroup VIIIB

MINISTRY OF EDUCATION AND SCIENCE OF THE RUSSIAN FEDERATION VORONEZH STATE UNIVERSITY

CHEMISTRY OF THE ELEMENTS OF GROUP VIII OF THE PERIODIC SYSTEM

Tutorial

Publishing and Printing Center of Voronezh State University

Approved by the Scientific and Methodological Council of the Faculty of Chemistry on December 12, 2012, Protocol No. 9

Compiled by: I.Ya. Mittova, E.V. Tomina, B.V. Sladkopevtsev, D.O. Solodukhin

Reviewer Dr. chem. Sciences, Professor V.N. Semenov

The textbook was prepared at the Department of Materials Science and Industry of Nanosystems, Faculty of Chemistry, Voronezh State University.

For directions: 020300 - Chemistry, physics and mechanics of materials, 020100 - Chemistry

FOREWORD

This textbook is a continuation of the first three parts, in which the Periodic Law was considered as the basis of inorganic chemistry, the chemistry of elements of groups I–VI of the Periodic system. The fourth part deals with the chemistry of elements of group VIII of the Periodic Table of Chemical Elements by D.I. Mendeleev.

The manual is designed to help the first-year student in studying the discipline "Inorganic Chemistry" and, in fact, is a summary of the lecture course, which displays all the main key points that need to be taken into account when studying it.

As a continuation of the cycle of manuals for the course "Inorganic Chemistry", this edition generally retains the structure and sequence of presentation of the material. The description begins with a general description of simple substances, their prevalence in nature, methods of preparation and chemical properties, in separate subsections the properties of the compounds of the elements of the group are considered. Particular attention is paid to the use of chemical elements and their compounds as a variety of modern materials.

To implement the principle of clarity, the manual contains a large amount of illustrative material and tables that allow you to present vast amounts of material in a compact form and reflect the main patterns in the change in the properties of chemical elements and their compounds.

When writing, modern literary sources were used, a list of which is given at the end of the manual. The illustrative material is mostly taken from the textbooks "Inorganic Chemistry" (under the editorship of Yu.D. Tretyakov, M.: Asademia, 2004) and "Chemistry of the Elements" (N. Greenwood, A. Earnshaw. M.: BINOM. Lab. Knowledge , 2008).

This manual is primarily intended for first-year students of the Faculty of Chemistry, but it can also be useful for senior students, in particular masters studying the disciplines "Modern Inorganic Chemistry" and "Modern Problems of Inorganic Chemistry", to update the knowledge gained earlier.

CHAPTER 1. VIII-A GROUP

1.1. Simple substances

1.1.1. Element Properties

Group VIII-A elements: helium 2 He, neon 10 Ne, argon 18 Ar, krypton 36 Kr, xenon 54 Xe and radon 86 Rn - are called noble gases. Electron-

naya configuration of the first representative of the group, helium - 1s 2 . The atoms of other noble gases have eight valence electrons at the outer level (Table 1), which corresponds to a stable electronic configuration.

Table 1

Properties of Group VIII-A Elements

Property

Nuclear charge Z

Electronic config-					4f 14
walkie-talkie	[Not] 2s 2p	3s 3p
walkie-talkie
			4 s 4 p	5 s 5 p	5p6
Atomic radius, nm

The first ion energy
zation I 1, kJ/mol
Excitation energy
ns2 np6 →ns2 np5 (n + 1) s1 ,

Electronegative-

A completely completed configuration of the outer electron layer (in the case of helium and neon) or the presence of an octet of electrons determines the high ionization energies of noble gas atoms and, as a result, their low chemical activity. The ability of atoms of these elements to enter into chemical reactions increases with increasing atomic radius due to the weakening of the attraction of valence electrons to the nucleus. To date, chemical compounds have been obtained only for heavy noble gases: krypton, xenon and radon.

1.1.2. Being in nature, getting

Helium is the second (after hydrogen) element in abundance in the Universe. At the same time, the mass of "terrestrial" helium is only one millionth of the mass of the earth's crust. On the Sun, a significant number of helium nuclei are formed during the nuclear "burning" of hydrogen, so the content of this element in the Universe is gradually increasing. Helium is also formed during the α-decay of radionuclides. It fills voids in radioactive rocks and minerals, and from there it enters the atmosphere. As an impurity, helium accompanies methane. The main source of helium is natural gas.

All noble gases contained in the air, which is the raw material for their industrial production.

Radon is a radioactive element. The longest-lived isotope 222 Rn, formed by the α-decay of 226 Ra, has a half-life of 3.82 days. One gram of radium-226 releases 6.6 x 10–4 ml of radon per day. Thorium minerals contain a certain amount of the 220 Rn isotope.

1.1.3. Physical properties

All noble gases are colorless, tasteless and odorless, have low melting and boiling points. Their molecules are monoatomic. Argon, krypton, and xenon form clathrates based on water and hydroquinone, for example, Xe 3C6 H4 (OH)2, in which noble gas atoms are located in the cavities of the “host” substance structure. Smaller atoms of helium and argon are not able to stay in cavities. The main physical properties of simple substances are given in Table. 2.

	Properties of simple substances				table 2
	Properties of simple substances
Property

Standard en-
talpia evaporate-
ion, kJ/mol
ion, kJ/mol


t pl , ° С
t bale, ° С
	5.2 10–4	1.8 10–3	1.1 10–3	8.7 10–6	6.0 10–18
in the air, %
Solubility in
water at 20 °C,

1.2. Chemical properties

True chemical compounds have been obtained only for krypton, xenon and radon. The chemistry of xenon is best studied, since krypton compounds are extremely unstable, and radon is radioactive.

The interaction of xenon with fluorine leads to the formation of a mixture of fluorides. A convenient method for the synthesis of difluoride, which makes it possible to avoid direct fluorination, is the oxidation of xenon with silver (II) fluoride in the presence of a Lewis acid:

2AgF2 + 2BF3 + Xe = XeF2 + 2AgBF4.

Xenon fluorides are colorless volatile crystalline substances that are easily hydrolyzed. Xenon difluoride forms stable solutions that decompose within a few hours:

2XeF2 + 2H2 O = 2Xe + 4HF + O2.

Xenon tetra- and hexafluoride are much more sensitive to air moisture - when they enter water, they instantly hydrolyze to form XeO3:

6XeF4 + 12H2 O = 2XeO3 + 4Xe + 3O2 + 24HF, XeF6 + 3H2 O = XeO3 + 6HF.

Xenon fluorides have a molecular structure (Fig. 1). XeF2 is a linear molecule with three unshared electron pairs lying in the equatorial plane (AB2 E3 type); XeF4 has the shape of a square with two lone pairs (type AB4 E2), and XeF6 is a distorted octahedron with one lone pair of electrons (type AB6 E). Free XeF6 molecules are known in pairs.

Rice. 1. Structure of XeF2 (a), XeF4 (b), XeF6 molecules (dynamic model with a migrating electron pair) (c)

The molecular orbital method describes the formation of xenon fluorides in terms of three-center four-electron bonds. For example, p x orbitals of a xenon atom and two fluorine atoms are involved in the formation of the XeF2 molecule (Fig. 2). Their interaction leads to the appearance of three molecular σ orbitals: bonding, nonbonding, and loosening, the first two of which are filled with electrons. The bond order is thus equal to one. Compounds containing three-center four-electron bonds are called hypervalent.

Rice. 2. Scheme of molecular orbitals of the XeF2 molecule. The combinations of atomic orbitals involved in the formation of each of the molecular orbitals of the molecule are shown on the right.

Xenon fluorides are strong oxidizing agents. They convert bromates to perbromates, iodates to periodates, sulfur to hexafluoride, manganese (II) salts to permanganates:

3XeF2 + S = 3Xe + SF6,

5XeF2 + 2Mn(NO3 )2 + 16KOH = 2KMnO4 + 10KF + 4KNO3 + 8H2O + 5Xe.

This is the basis for the use of xenon fluorides in the synthesis of higher transition metal fluorides:

XeF2 + 2CeF3 → Xe + 2CeF4 .

Another important property of xenon fluorides is their ability to act as both donors and acceptors of fluoride ions. Donor properties decrease in the series XeF2 > XeF6 > XeF4 . With typical Lewis acids PF5 , AsF5 , SbF5 , PtF5 and others, xenon difluoride most easily interacts, forming salts + - , + - :

XeF2 + AsF5 = + – .

The interaction of XeF2 with an excess of antimony pentafluoride at a pressure of 3 atm made it possible to obtain dark green crystals containing the paramagnetic dixenon cation Xe2 +:

4XeF2 + 8SbF5 = 2Xe2 + – + 3F2 .

The Xe–Xe distance in the cation is 0.309 nm, which indicates only a very weak interaction.

The acceptor properties decrease in the series XeF6 > XeF4 > XeF2 . They are most typical for xenon hexafluoride, which easily reacts with fluorides of heavy alkali metals (rubidium and cesium):

XeF6 + CsF = Cs+ – .

For krypton, only compounds with fluorine in the +2 oxidation state are known. Fluoride KrF2 is formed from simple substances at liquid temperature

whom nitrogen. It is usually produced by passing an electrical discharge through a mixture of krypton and fluorine in a reactor cooled with liquid nitrogen. In structure and properties, KrF2 resembles xenon difluoride, being an even stronger oxidizing agent in comparison with it. KrF2 oxidizes gold trifluoride to pentafluoride and chlorine pentafluoride to + ion, turns metallic gold into gold (V):

7KrF2 + 2Au = 2KrF+ – + 5Kr.

Interestingly, free fluorine, unlike krypton difluoride, is not capable of oxidizing gold to AuF5.

Oxygen compounds are known only for xenon. Xenon forms two oxides: XeO3 and XeO4 (Fig. 3), both are extremely unstable and easily explode from the slightest shock. XeO3 oxide is formed by the hydrolysis of tetra- and hexafluorides or by the action of hexafluoride on silicon oxide:

2XeF6 + 3SiO2 = 2XeO3 + 3SiF4 .

In free form, it is a colorless crystals, highly soluble in water.

Rice. 3. Structure of XeO3 (a) and XeO4 (b) molecules

It was possible to isolate only acid xenates of alkali metals (M) of the composition МНХеО4, which, when an excess of alkali is added, disproportionate:

2NaHXeO4 + 2NaOH = Na4 XeO6 + Xe + O2 + 2H2 O.

This is how perxenates are obtained - salts of perxenonic acid H4 XeO6. They contain the [XeO6 ]4– ion, which has an octahedral structure.

By acting on perxenates with 100% sulfuric acid, the highest xenon oxide XeO4 is obtained:

Na4 XeO6 + 2H2 SO4 = 2Na2 SO4 + XeO4 + 2H2 O.

It is a colorless gas, spontaneously exploding, its solutions in donor solvents (BrF5, HF) are more stable, they can be stored at –33 °C. Xenon tetroxide and perxenates are among the strongest oxidizers.

1.3. Application

The initial use of helium as a non-combustible gas for filling balloons (its lifting force is approximately 1 kg / m3) loss-

lo its importance, although it is still used for meteorological probes. Helium is used as a cryogenic liquid to maintain temperatures of the order of 4.2 K and below (30% of the produced He is used for these purposes); 2/3 are spent on spectrometers and NMR tomographs. Other important applications are arc welding (21%), sealing and cleaning (11%). The choice between Ar and He for this purpose is determined by the cost of the gas, and everywhere except the United States, argon is generally preferred. Small in volume, but important areas of application of helium are as follows:

a) to replace N2 in artificial gas mixtures when breathing at great depths (the low solubility of helium in the blood minimizes the outgassing that occurs in the case of nitrogen - when the diver goes through decompression - and sometimes leads to death);

b) as a working medium in gas leak detectors; c) as a coolant in the cooling system of high-temperature

nuclear reactors; d) as a carrier gas in gas-liquid chromatography;

e) for deaeration of solutions and generally as an inert diluent or inert atmosphere.

Ar is used mainly as an inert gaseous medium in high-temperature metallurgical processes and, to a lesser extent, for filling incandescent lamps. Together with Ne, Kr and Xe, which are obtained in much smaller quantities, Ar is also used in discharge tubes - the tube color obtained depends on the composition of the gas mixture. Noble gases are also used in fluorescent tubes, although in this case the color does not depend on the gas filling the tube, but on the phosphorus coating the inside of the tube walls. Another important application is lasers, although compared to other applications, the amount of gas used here is negligible.

Other noble gases are significantly more expensive, so their use is limited only to highly specialized areas. Radon has been used in the treatment of cancer and as a source of radioactivity in metal casting flaw detection, but due to its short half-life (3.824 days) it has been supplanted by other materials. That small amount of radon that is required in practice is obtained as a decay product of 226 Ra (1 g of which gives 0.64 cm3 of radon within 30 days).

CHAPTER 2. VIII-B GROUP

2.1. Simple substances

2.1.1. Electronic structure

Group VIII-B includes nine elements at once: iron 26 Fe, ruthenium 44 Ru, osmium 76 Os, cobalt 27 Co, rhodium 45 Rh, iridium 77 Ir, nickel 28 Ni, palladium 46 Pd and platinum 78 Pt.

The properties of the chemical elements of group VIII-B do not differ too much, which was the reason for their combination into triads. The similarity in properties is due to the preservation of the composition and structure of the outer electron shell of atoms with a consistent increase in the atomic number of the element and, accordingly, the total number of electrons in an isolated atom. For elements of triads, with an unchanged structure of the outer electron shell (principal quantum number n = 4, 5, 6), the corresponding d-sublevel (electronic n - 1-layer) is completed (with an increase in the atomic number), the degree of filling of which does not have a determining effect on the dimensions atoms and ions, as well as the properties of compounds - at least if the chemical bond in them is predominantly ionic in nature.

At the same time, the properties of the compounds of the elements of the iron triad differ from the properties of the compounds of the elements of the triads of palladium and platinum (a family of platinum elements) that are similar in composition to a very significant extent.

One of the reasons for the greater similarity between the compounds of platinum elements (PE) in comparison with the compounds of the iron triad is the effect of lanthanide compression. Thus, the atomic radii of the elements of the palladium and platinum triads are almost the same, but differ significantly from the atomic radii of the elements of the iron triad.

When moving from top to bottom in a group, the stability of compounds containing an element in the highest oxidation state increases (see the diagram below). If for iron the oxidation states +2 and +3 are the most characteristic, and the states +6 and especially +8 are unstable, then for osmium compounds containing the element in the highest possible oxidation state +8 are quite stable. A similar regularity is observed upon passing from Co and Ni to their heavy analogs. So, for nickel, the most stable compounds are those where it has an oxidation state of +2, while for palladium, and especially for platinum, an oxidation state of +4 is characteristic.

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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"

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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 Rus' (Ruthenia is the Latin name for Rus').

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=0.625E

Atomic radius Ni=0.124nm Ion radius Ni=0.69E

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 is easily soluble in acids (with the formation of Co(II) salts, and dissolution in alkali is very difficult. However, the presence of the acidic properties of Co(OH)2 is confirmed by the existence of the 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 immunostimulating 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 crystalline 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. As a selective reagent for the determination of nickel in an alkaline environment in the presence of oxidizing agents, 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.

...

Metals of the iron family

In the triad of iron, the horizontal analogy, characteristic of the d-elements as a whole, is most clearly manifested. The properties of the elements of the iron triad are given in Table. 42.

Table 9.2 Properties of the elements of the iron triad

Natural resources. Iron is the fourth (after O 2 , Si, Al) most abundant element in the earth's crust. It can occur in nature in a free state: it is iron of meteoric origin. Iron meteorites contain on average 90% Fe, 8.5% Ni, 0.5% Co. There is an average of one iron meteorites for twenty stone meteorites. Sometimes native iron is found, taken out of the bowels of the earth by molten magma.

To obtain iron, magnetic iron ore Fe 3 O 4 (magnetite mineral), red iron ore Fe 2 O 3 (hematite) and brown iron ore Fe 2 O 3 x H 2 O (limonite), FeS 2 - pyrite are used. In the human body, iron is present in hemoglobin.

Cobalt and nickel in the metallic state are found in meteorites. The most important minerals: cobaltite CoAsS (cobalt luster), iron-nickel pyrite (Fe,Ni) 9 S 8 . These minerals are found in polymetallic ores.

Properties. Iron, cobalt, nickel are silvery-white metals with a grayish (Fe), pinkish (Co) and yellowish (Ni) tint. Pure metals are strong and ductile. All three metals are ferromagnetic. When heated to a certain temperature (Curie point), the ferromagnetic properties disappear and the metals become paramagnetic.

Iron and cobalt are characterized by polymorphism, while nickel is monomorphic and has an fcc structure up to the melting temperature.

The presence of impurities greatly reduces the resistance of these metals to an aggressive atmosphere in the presence of moisture. This leads to the development of corrosion (iron rusting) due to the formation of a loose layer of a mixture of oxides and hydroxides of variable composition on the surface, which do not protect the surface from further destruction.

Comparison of the electrode potentials of the E 2+ /E systems for iron (-0.441 V), nickel (-0.277 V) and cobalt (-0.25 V), and the electrode potential of the Fe 3+ /Fe system (-0.036 V), shows that that the most active element of this triad is iron. Dilute hydrochloric, sulfuric and nitric acids dissolve these metals with the formation of E 2+ ions:

Fe+2HC? = FeC? 2+H2;

Ni + H 2 SO 4 \u003d NiSO 4 + H 2;

3Co + 8HNO 3 \u003d 3Co (NO 3) 2 + 2NO + 4H 2 O;

4Fe + 10HNO 3 \u003d 3Fe (NO 3) 2 + NH 4 No 3 + 3H 2 O.

More concentrated nitric acid and hot concentrated sulfuric acid (less than 70%) oxidize iron to Fe (III) with the formation of NO and SO2, for example:

Fe + 4HNO 3 \u003d Fe (NO 3) 3 + No + 2H 2 O;

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

Very concentrated nitric acid (sp.v. 1.4) passivates iron, cobalt, nickel, forming oxide films on their surface.

With respect to alkali solutions, Fe, Co, Ni are stable, but react with melts at high temperatures. All three metals do not react with water under normal conditions, but at a red-hot temperature, iron interacts with water vapor:

3Fe + 4H 2 o Fe 3 O 4 + 4H 2 .

Cobalt and nickel are noticeably more resistant to corrosion than iron, which is consistent with their position in the series of standard electrode potentials.

Finely dispersed iron in oxygen burns when heated to form Fe 3 O 4 , which is the most stable iron oxide and cobalt forms the same oxide. These oxides are derivatives of elements in oxidation states +2, +3 (EO E 2 O 3). Thermal oxidation of cobalt and nickel proceeds at higher temperatures, with the formation of NiO and CoO, which have a variable composition depending on the oxidation conditions.

For iron, nickel, cobalt, oxides EO and E 2 O 3 are known (Table 9.3)

Table 9.3 Oxygen-containing compounds of elements of subgroup VIIIB

Element name	Oxidation state		Hydroxides
	Character		Name	Ion formula	Name
Iron (Fe)			Basic		Iron(II) hydroxide		Salts of iron (II)
		Amphoteric with a predominance of basic		Iron(III) hydroxide		Salts of iron (III)
	ferrous acid
		Acid		iron acid
Cobalt (Co)			Basic		Cobalt(II) hydroxide		Salts of cobalt (II)
		Basic		Cobalt(III) hydroxide		Salts of cobalt (III)
Nickel (Ni)			Basic		Nickel(II) hydroxide		Nickel(II) salts
		Basic		Nickel(III) hydroxide		Nickel(III) salts

Oxides EO and E 2 O 3 cannot be obtained in pure form by direct synthesis, since in this case a set of oxides is formed, each of which is a phase of variable composition. They are obtained indirectly - by the decomposition of certain salts and hydroxides. Oxide E 2 O 3 is stable only for iron and is obtained by dehydration of the hydroxide.

EO oxides are insoluble in water and do not interact with it or with alkali solutions. The same is true for the corresponding E(OH) 2 hydroxides. E(OH) 2 hydroxides readily react with acids to form salts. The acid-base properties of the hydroxides of the elements of the iron triad are given in Table. 42.

Iron hydroxide (III) Fe (OH) 3 is formed during the oxidation of Fe (OH) 2 with atmospheric oxygen:

4 Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3.

A similar reaction is typical for cobalt. Nickel hydroxide (II) is stable with respect to atmospheric oxygen. As a result, E(OH) 3 hydroxides behave differently when interacting with acids. If Fe (OH) 3 forms iron (III) salts, then the reaction of Co (OH) 3 and Ni (OH) 3 with acids is accompanied by their reduction to E (+2):

Fe(OH)3 + 3HC? = FeC? 3+3H2O;

2Ni(OH)3 + 6HC? = 2 NiC? 2+C? 2+6H2O.

Fe(OH) 3 hydroxide also exhibits an acid function, reacting with hot concentrated alkali solutions to form hydroxo complexes, for example, Na 3 . Ferrous acid derivatives HFeO 2 (ferrites) are obtained by fusing alkalis or carbonates with Fe 2 O 3:

2NaOH + Fe 2 O 3 2NaFeO 2 + H 2 O;

MgCO 3 + Fe 2 O 3 MgFe 2 O 4 + CO 2.

Ferrites Me II Fe 2 O 4 belong to the class of spinels. The Fe 3 O 4 and Co 3 O 4 oxides discussed above are formally FeFe 2 O 4 and CoCo 2 O 4 spinels.

Unlike cobalt and nickel, iron compounds are known in which its oxidation state is + 6. Ferrates are formed during the oxidation of Fe (OH) 3 in hot concentrated alkali in the presence of an oxidizing agent:

2Fe +3 (OH) 3 + 10KOH + 3Br 2 = 2K 2 Fe +6 O 4 + 6KBr + 2H 2 O.

Ferrates are thermally unstable and with slight heating (100-2000C) turn into ferrites:

4K 2 FeO 4 4KfeO 2 + 2K 2 O + 3O 2 .

In the free state, iron acid and the corresponding FeO 3 oxide have not been isolated. By solubility and structurally, ferrates are close to the corresponding chromates and sulfates. Potassium ferrate is formed by fusing Fe 2 O 3 with KNO 3 and KOH:

Fe 2 O 3 + 3KNO 3 + 4KOH \u003d 2K 2 feO 4 + 3KNO 2 + 2H 2 O.

Ferrates are red-violet crystalline substances. When heated, they decompose. The acid H 2 FeO 4 cannot be isolated, it instantly decomposes into Fe 2 O 3 , H 2 O and O 2 . Ferrates are strong oxidizing agents. In acidic and neutral environments, ferrates decompose, oxidizing water:

2Na 2 FeO 4 + 10 H 2 O 4Fe (OH) 3 + 4NaOH + O 2.

Compounds with non-metals. Fe, Ni, Co halides are relatively few and correspond to the most typical oxidation states +2 and +3. For iron, the halides FeG 2 and FeG 3 with fluorine, chlorine and bromine are known. With direct interaction, FeF 3 , FeC? 3, FeBr3. Dihalides are obtained indirectly - by dissolving the metal (or its oxide) in the corresponding hydrohalic acid. Cobalt trifluoride CoF 3 and trichloride CoC? 3 . Nickel does not form trihalides. All dihalides of the iron triad are typical salt-like compounds with a significant ionic contribution to the chemical bond.

Iron, cobalt, nickel interact vigorously with chalcogens and form chalcogenides: EX and EX 2 . Monochalcogenides can be obtained by the interaction of the corresponding components in solutions:

CoC? 2 + (NH 4) 2 S \u003d CoS + 2NH 4 C?.

All chalcogenides are phases of variable composition.

The compounds of the metals of the iron triad with other non-metals (pnictogens, carbon, silicon, boron) differ markedly from those considered above. All of them do not obey the rules of formal valency and most of them have metallic properties.

Iron, cobalt, nickel absorb hydrogen, but do not give certain compounds with it. When heated, the solubility of hydrogen in metals increases. Hydrogen dissolved in them is in the atomic state.

Salts of oxygen-containing acids and complex compounds. All salts of hydrochloric, sulfuric and nitric acids are soluble in water.

Nickel (II) salts are green, cobalt (II) - blue, and their solutions and crystalline hydrates - pink (for example,), iron (II) salts - greenish, and iron (III) - brown. The most important salts are: FeC? 36H2O; FeSO 4 7H 2 O - iron sulfate, (NH 4) 2 SO 4 FeSO 4 6H 2 O - Mohr's salt; NH 4 Fe (SO 4) 2 12H 2 O - iron ammonium alum; NiSO 4 6H 2 O, etc.

The ability of iron, cobalt and nickel salts to form crystalline hydrates indicates the tendency of these elements to complex formation. Crystalline hydrates are a typical example of aquacomplexes:

[E (H 2 O) 6] (ClO 4) 2; [E (H 2 O) 6] (NO 3) 2.

Anionic complexes are numerous for the elements of the iron triad: halide (Me I (EF 3), Me 2 I [EG 4], Me 3 [EG 4], etc.), thiocyanate (Me 2 I [E (CNS) 4] , Me 4 I [E (CNS) 6 ], Me 3 I [E (CNS) 6 ]), oxolate (Me 2 I [E (C 2 O 4) 2 ], Me 3 [E (C 2 O 4) 3]). Cyanide complexes are especially characteristic and stable: K 4 - potassium hexacyanoferrate (II) (yellow blood salt) and K 3 - potassium hexacyanoferrate (III) (red blood salt). These salts are good reagents for detecting Fe +3 ions (yellow salt) and Fe 2+ ions (red salt) at pH ??7:

4Fe 3+ + 4- = Fe 4 3;

Prussian blue

3Fe 2+ + 2 3- = Fe 3 2 .

turnbull blue

Prussian blue is used as a blue dye. When thiocyanate salts KCNS are added to a solution containing Fe 3+ ions, the solution turns blood red due to the formation of iron thiocyanate:

FeC? 3 + 3KCNS = Fe(CNS) 3 + 3KC?.

This reaction is very sensitive and is used to open the Fe 3+ ion.

Cobalt (II) is characterized by stable simple salts and unstable complex compounds K 2, K 4, which turn into cobalt (III) compounds: K 3, C? 3 .

Characteristic complex compounds of iron, iron, cobalt and nickel are carbonyls. Similar compounds were considered earlier for the elements of the chromium and manganese subgroups. However, the most typical among carbonyls are: , , . Iron and nickel carbonyls are obtained as liquids at ordinary pressure and 20-60° C. by passing a CO stream over metal powders. Cobalt carbonyl is obtained at 150-200 o C and pressure (2-3) 10 7 Pa. These are orange crystals. In addition, there are carbonyls of a more complex composition: Fe (CO) 9 and trinuclear carbonyls, which are compounds of a cluster type.

All carbonyls are diamagnetic, since CO ligands (as well as CN?) create a strong field, as a result of which the valence d-electrons of the complexing agent form p-bonds with CO molecules by the donor-acceptor mechanism. y-bonds are formed due to lone electron pairs of CO molecules and the remaining vacant orbitals of the complexing agent:

Nickel (II), on the contrary, forms many stable complex compounds: (OH) 2 , K 2 ; the 2+ ion is dark blue.

This reaction is widely used in qualitative and quantitative analysis for the determination of nickel. Compounds of nickel and especially cobalt are poisonous.

Application. Iron and its alloys form the basis of modern technology. Nickel and cobalt are important alloying additions in steels. Nickel-based heat-resistant alloys (nichrome containing Ni and Cr, etc.) are widely used. Copper-nickel alloys (melchior, etc.) are used to make coins, jewelry, and household items. Many other nickel- and cobalt-containing alloys are of great practical importance. In particular, cobalt is used as a viscous component of the materials from which metal-cutting tools are made, in which particles of exceptionally hard carbides MoC and WC are interspersed. Electroplated nickel coatings of metals protect the latter from corrosion and give them a beautiful appearance.

Iron family metals and their compounds are widely used as catalysts. Sponge iron with additives - a catalyst for the synthesis of ammonia. Highly dispersed nickel (Raney nickel) is a very active catalyst for the hydrogenation of organic compounds, in particular fats. Raney nickel is obtained by acting with an alkali solution on the NiA? intermetallic compound, while aluminum forms a soluble aluminate, and nickel remains in the form of tiny particles. This catalyst is stored under a layer of organic liquid, since in a dry state it is instantly oxidized by atmospheric oxygen. Cobalt and manganese are part of the catalyst added to oil paints to speed up their "drying".

Fe 2 O 3 oxide and its derivatives (ferrites) are widely used in radio electronics as magnetic materials.

Name	Refers, at. weight	Electronic formula	Radius, pm	Main isotopes (%)
Helium Helium [from Greek. helios - sun]			atomic 128	3 He* (0.000138) 4 He* (99.99986)
Neon Neon [from Greek. neos - new]			Vanderwaals 160
Argon Argon [from Greek. argos - inactive]		ls 2 2s 2 2p 6 3s 2 3p 6	atomic 174
Krypton Krypton [from Greek. Kryptos - hidden]		3d 10 4s 2 4p 6	Covalent 189
Xenon Xenon [from Greek. xenos - stranger]		4d 10 5s 2 5p 6	atomic 218, covalent 209	129 He* (26.4)
Radon Radon [named after radium]		4f 14 5d 10 6s 2 6p 6		219*,220,222 Rn (traces)

In the earth's crust (%)	In the ocean (%)	In the human body
In the earth's crust (%)	In the ocean (%)	Average (with a body weight of 70 kg)	Blood (mg/l)
				non-toxic, but may cause asphyxiation



				non-toxic
				toxic due to radioactivity

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General characteristics of the elements of the 8th group of the secondary subgroup. Elements of subgroup VIIIB

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