Characteristics of d-elements of group VI. General characteristics of the elements of the VI A subgroup Elements 6a of the subgroup
Elements of the main subgroup of group VI bear the common group name "chalcogens". Their atoms are electronic counterparts, since they have the same structure of the outer electronic layer (ns 2 np 4).
p-elements, non-metals (except polonium)
Valency II, IV; VI
Oxidation states -2, +2, +4, +6 (oxygen is an exception)
Valence states of atoms of the sulfur subgroup
In unexcited atoms, there are 2 unpaired electrons that participate in the formation of ionic or covalent bonds with other atoms (B \u003d II).
Interacting with more EO atoms, sulfur, selenium, and tellurium (as well as Po) can pass into excited states, which is accompanied by the transition of electrons to vacant d-orbitals.
In this case, the number of unpaired electrons increases to 4 or 6, as a result of which the atoms can exhibit B equal to IV and VI.
The difference between oxygen and other elements of the subgroup
In O atoms, valence electrons are located at the 2nd energy level, which has only s- and p-orbitals. This excludes the possibility of the transition of O atoms to excited states, therefore oxygen in all compounds exhibits a constant B = II.
Having a high EO (second only to fluorine), oxygen atoms are always negatively charged in compounds (s.o. = -2 or -1). Exception - fluorides OF 2 and O 2 F 2
Simple substances
Simple substances formed by the elements of this subgroup exist in the form of various allotropic modifications:
O 2 - oxygen, O 3 - ozone
S - plastic, monoclinic, rhombic
Se - red, glassy, gray
Te - crystalline, amorphous
In addition to gaseous oxygen and ozone, all other simple substances at ordinary temperatures are solids.
Compounds with hydrogen (chalcogen hydrogens) H 2 E
H 2 S - hydrogen sulfide, H 2 Se - hydrogen selenide, H 2 Te - hydrogen telluride
Colorless gases with an unpleasant odor. Very poisonous. Strong reducing agents. Aqueous solutions exhibit the properties of weak acids.
The most important connections with C.O. +4
SO 2 (g.) Se) 2 (tv.), TeO 2 (tv.) - acid oxides
H 2 SO 3 - sulfurous acid, H 2 SeO 3 - selenious acid, H 2 TeO 3 - telluric - weak acids, reducing agents.
Acid properties are weakening. Recovery ability is reduced.
The most important connections with C.O. +6
SO 3 (l.) Se) 3 (tv.), TeO 3 (tv.) - acid oxides
H 2 SO 4 - sulfuric acid, H 2 SeO 4 - selenic acid - strong acids, H 2 TeO 4 - orthotelluric - weak acid.
Salts: sulfates, selenates, tellurates.
The elements of the 6-A group include: oxygen (8 O), sulfur (16 S), selenium (34 Se), tellurium (52 Te) and polonium (84 Po). The name of the group "Chalcogenes" literally translates as "giving birth to salts" from the Greek. "Chalkos" - copper and "genos" - genus, origin. In nature, chalcogens are indeed found most often in the form of copper compounds (except oxygen) - these are sulfides, copper selenides. Copper(II) sulfide Copper(I) selenide
Ø When moving from oxygen to polonium, the size of the atom increases, the non-metallic properties weaken, and the metallic ones increase: oxygen and sulfur are typical non-metals, selenium and tellurium are metalloids with non-metallic properties, polonium is a metal. Ø The ionization energy of atoms in the same series decreases E 0 → E +, which means an increase in reducing properties (the ability to donate an electron). Ø Due to the high electron density and strong interelectronic repulsion, the electron affinity energy for oxygen is less than for all other elements of the 6-A group. As a result, the S 2 anion is much more stable than similar anions of selenium and tellurium, and the O 2 anion practically does not exist in a free form. Ø The decrease in electronegativity in the series from oxygen to polonium means a decrease in the degree of polarity of covalent bonds in the series of chalcogens.
The higher the bond energy, the stronger the bond and the more stable the connection. Due to the high electron density and interelectronic repulsion forces, a single bond between oxygen atoms (O - O) is the least stable. On the contrary, the formation of a double bond for oxygen is much more favorable, since the energy of a double bond is much higher than the energy of two single bonds.
Forms of existence of compounds 6 -A of group E Prost. in-in H 2 E EO 3 H 2 EO 4 CO O O 2 H 2 O - - -2, -1, 0, +1, +2 S S H 2 S SO 3 H 2 SO 4 -2, 0, + 2, +4, +6 Se Se H 2 Se Se. O 3 H 2 Se. O 4 -2, 0, +2, +4, +6 Te Te H 2 Te Te. O 3 H 2 Te. O 4 -2, 0, +2, +4, +6 Po Po H 2 Po Po. O 3 -2, 0, +2, +4, +6 -
The electronic configuration of the unexcited atom is 1 s 22 p 4. In most compounds, it exhibits an oxidation state of -2, but compounds with oxidation states of -1 are known; 0; +1; +2; +4. The most common element in the earth's crust (58%). 3 stable isotopes were found: 168 O (99.759%), 178 O (0.037%), 18 O (0.204%). 8 More than 1400 minerals containing oxygen are known.
Features of the structure of the oxygen atom 1) There is no low-energy d-sublevel; 2) High electronegativity (place after fluorine), unable to donate more than 2 electrons (all other elements of the subgroup exhibit the highest oxidation state +6) 3) Small atomic radius
The structure of the oxygen molecule 1) 12 electrons of the outer energy level are located in 8 molecular orbitals (MO); 2) The orbitals of oxygen atoms are close in energy, so no non-bonding orbitals are formed; 3) 8 electrons are located in bonding orbitals and 4 electrons in loosening ones, so the bond order is (8 – 4)/2 = 2; 4) The presence of 2 unpaired electrons on loosening π * orbitals gives the oxygen molecule paramagnetic properties.
Physical properties of oxygen 1) The diatomic molecule is the most stable, since the dissociation energy O = O is 494 k. J / mol, while the O - O bonds are only 210 k. J / mol; 2) Oxygen molecules are weakly polarized, so the intermolecular bonds between them are weak. Tm. = -218 °С Тboil. \u003d -183 ° С. 3) Poorly soluble in water (5 volumes in 100 volumes of water at 0 °C). 4) Liquid and solid oxygen are attracted by a magnet, since its molecules have paramagnetic properties. 5) Solid oxygen is blue, and liquid oxygen is blue.
There are 3 allotropic modifications of oxygen O 2, O 3 (ozone) and O 4 - unstable tetraoxygen.
Allotropic modifications of oxygen A bluish gas with a characteristic pungent odor. The ozone molecule is polar. Slightly soluble in water Tm. = -193 °С Тboil. = -112 °С
Chemical properties of oxygen 1) A strong oxidizing agent. 2) Directly does not interact only with inert gases, halogens, silver, gold and platinum group metals (except osmium). Gold (III) oxide Au - yellow O - red
Obtaining oxygen 1) Biological origin of oxygen: 6 H 2 O + 6 CO 2 hν C 6 H 12 O 6 + 6 O 2 400 -500 ° C 2) Decomposition of potassium chlorate (bertolet salt): KCLO 3 3 O 2 + 2 KCl 3) Decomposition of potassium permanganate: KMn. O 4 210 -240 ° С K 2 Mn. O4 + Mn. O 2 + O 2
Oxygen compounds (-2) The water molecule has a tetrahedral (not flat) structure. The oxygen atom is in the state. sp 3 hybridization The VS method does not explain why one of the two lone electron pairs (LEP) is more active than the other
The structure of the water molecule (MO method) v Two 1 s atomic orbitals (one from each hydrogen atom), one 2 s orbital and three 2 p orbitals of the oxygen atom take part in the formation of the water molecule. In total, 8 electrons are located in 6 molecular orbitals of a water molecule. v Interaction (overlapping) of 1 s orbitals of two hydrogen atoms with 2 s and 2 px orbitals of an oxygen atom leads to the formation of 4 MO, of which two are bonding (2 a 1, 1 b 1) and 2 loosening (4 a 1, 2 b one). v The energy of bonding orbitals is dominated by lower energy AOs (oxygen), while the energy of antibonding orbitals is dominated by higher energy AOs (hydrogen). v 2 pz-orbital of oxygen and 1 s-orbital of one of the hydrogens slightly overlap and form a weakly binding orbital (3 a 1). v Orbital 2 py of the oxygen atom is perpendicular to the plane of overlapping orbitals, does not overlap with 1 s-orbitals at. H and, therefore, forms a non-bonding 1 b 2 orbital.
The structure of the water molecule (MO method) In total, 2 binding, 2 non-binding and 2 loosening MOs are formed in the water molecule. 8 electrons are arranged in pairs in 2 bonding and 2 non-bonding orbitals.
The structure of water (MO method) MO 2 a 1, 1 b 1 electrons form O-H bonds, and ē, located on non-bonding orbitals (3 a 1, 1 b 2), correspond to the free EP of the water molecule. However, MO 3 a 1 and 1 b 2 differ in localization and energy. 1 b 2 is localized on the oxygen atom and has a purely p-character. 3 a 1 has a lower energy and is delocalized, since it is formed with the participation of AO at. H and O. Localization of 1 b 2 on the oxygen atom leads to the fact that the negative charge in the water molecule is concentrated near the oxygen atom, and near the hydrogen atom - positive. Hence the polarity of the water molecule (μ = 1.84 D).
Hydrogen bonds in a water molecule Consequences of the polarity of water molecules: 1) The ability to form intermolecular hydrogen bonds; 2) High melting and boiling points (0 °C and 100 °C) 3) Strong surface tension
Boiling point of chalcogens H 2 O Tbp. , ° С H 2 S 100 -60 H 2 Se H 2 Te -42 -2
Hydrogen bonds in an ice molecule Ice crystals have a hexagonal structure. Each water molecule is connected to three neighboring molecules through hydrogen bonds. Each water molecule forms 4 hydrogen bonds using both NEPs. During melting, one hydrogen bond is broken (at 0 °C, 15% hydrogen is broken. St.). In this case, some of the molecules get inside the framework, which explains the fact: at 4 °C, the density of water is maximum.
Oxides In accordance with the nature of the element in a positive oxidation state, the nature of oxides in the periods and groups of the periodic system naturally changes. In periods, the effective negative charge on oxygen atoms decreases and a posteriori transition from basic through amphoteric oxides to acidic ones takes place.
Interaction of oxides with water In a series of acidic oxides, acidic properties and the ability to interact with water increase in the period, which is confirmed by a decrease in the potential ∆G 0298.
Superoxides The attachment of one electron to an oxygen molecule causes the formation of a superoxide ion O 2 -. O 2 derivatives are called superoxides, which are known for the most active alkali metals (potassium, rubidium, cesium). Superoxides are formed by direct interaction of simple substances: K + O 2 \u003d KO 2. The unpaired electron of the O 2 ion - determines the paramagnetism of superoxides and the presence of color in them. Superoxides are very strong oxidizing agents. They react violently with water to release hydrogen.
The V and A groups of the periodic system include: oxygen 8O, Sulfur 16S, Selenium 34Se, tellurium 52Te and radioactive Polonium 84Ro. The common name of these elements “chalcogens” (translated from Greek - those that give birth to copper ores) is due to the fact that in nature the elements of this subgroup (except oxygen) are most often found in the form of copper compounds: sulfides, selenides, and the like.
With a decrease in the ionization energy in the series O - S - Se - Te - Po, the properties of chalcogens change from non-metallic to metallic: oxygen and sulfur are typical non-metals, Selenium and tellurium are metalloids, Polonium is a metal.
In the ground state, chalcogen atoms have the outer layer electronic configuration ns2np4, with an even number of valence electrons, of which two are unpaired. And this already in simple compounds leads to the alternative possibility of the formation of either a multiple bond between two atoms in the E2 molecule, or a single bond in chain structures. Due to the high electron density and strong interelectron repulsion, the electron affinity of oxygen and the strength of the E–E single bond are lower than for sulfur, selenium, and Tellurium. Oxygen is able to form strong pπ-pπ bonds with other atoms, for example, with oxygen (O2, O3), carbon, nitrogen, phosphorus. For sulfur and its analogues, single bonds are energetically favorable. They are characterized by the phenomenon of catenation (the ability of atoms of elements to combine into rings or chains).
Oxygen, like other elements of the second period, differs in properties from the elements of its subgroup. The properties of sulfur are more similar to those of Selenium and Tellurium than to oxygen and Polonium.
Oxygen, which does not have a vacant d-orbital, is bivalent in most compounds, while other chalcogens are capable of forming up to six valence bonds. When passing from oxygen to Polonium, the size of atoms and their possible coordination numbers increase, while the value of the ionization energy and electronegativity decrease. The electronegativity of oxygen is second only to that of fluorine. Compounds in which oxygen exhibits an oxidation state of +2 are strong oxidizing agents and are very unstable. All other chalcogens show a high oxidation state (6). The stability of E + 6 decreases from sulfur to Polonium, for which compounds with an oxidation state of +4, +2 are stable (which is explained by an increase in the strength of the bond between 6s2 electrons and the nucleus).
3.2 Being in nature
Oxygen is the most abundant element in the earth's crust (its content is 49% of its total mass). Oxygen is found in water, silica, limestone, marble, basalt, bauxite, hematite, and many other minerals and rocks. The earth's atmosphere contains about 21% (by volume) of oxygen in the form of a simple substance - oxygen O2. Atmospheric oxygen is of biological origin and is formed in green plants during photosynthesis.
Other chalcogens are much less common. Their content (wt.%) Decreases with increasing serial number: S - 0.0048; Se—8 10-5; That is 1 10-6; Ro - 2 10-14. Sulfur, selenium and tellurium are concentrated in ore deposits, where they are predominantly combined with metals. A significant part of sulfur is in its native state (volcanic sulfur), or in the form of sulfides and sulfates. The most important sulfur minerals are: FeS2 (pyrite or iron pyrites) ZnS (zinc blende) HgS (cinnabar) PbS (lead luster); CuFeS2 (chalcopyrite), CaSO4 · 2 H2O (gypsum), Na2SO4 · 10H2O (mirabilite), etc. Many sulfur is found in oil and petroleum gases and creates man-made (pipeline corrosion) and environmental problems. Sulfur is part of proteins. Selenium and tellurium, rare elements in nature, are found as impurities in similar natural sulfur compounds. Polonium, as a decay product of uranium, is contained in uranium ore (the half-life for 210Ро is 138.4 days).
3.3 Physical properties and allotropy
The element oxygen exists in the free form of two allotropic modifications - oxygen O2 and less stable ozone O3. Oxygen is a colorless gas, odorless and tasteless, slightly heavier than air. The melting and boiling points are -183 ° C and -219 ° C, respectively. The solubility in water at 0 ° C is 4.89 volumes per 100 volumes of water. Ozone is formed in the upper atmosphere under the influence of ultraviolet radiation and during lightning discharges. Ozone is a colorless gas, Tm. = -193 ° С, Тb.p. \u003d -112 ° C, soluble in water (at 0 ° C - 1.82 volumes per 100 volumes of water. Unlike oxygen, ozone has a characteristic pungent odor.
For sulfur, several allotropic modifications are known. The most stable are rhombic (α-sulfur) and monoclinic (β-sulfur), which consist of cyclic S8 molecules located at the sites of rhombic and monoclinic crystal lattices. Rhombic sulfur is stable at room temperature - a solid, low-melting crystalline substance, light yellow in color, practically insoluble in water, soluble in organic solvents. When heated to about 96°C, it becomes monoclinic (long light yellow crystals). If molten sulfur is poured into cold water, dark brown plastic sulfur is formed (molecules with open chains, closed molecules S4, S6).
The polymorphism of selenium and tellurium is associated with the way in which molecular chains and cycles are packed in crystals. The red monoclinic modifications of selenium (α, β, γ) formed during the crystallization of selenium solutions in carbon disulfide consist of cyclic Se8 molecules. Amorphous red selenium precipitates from its aqueous solutions under the action of reducing agents. The most thermodynamically stable gray hexagonal modification of selenium has a metallic luster (Tm. = 200 ° C). It is formed by heating all other modifications and consists of unbranched helical polymer chains Sen.
For tellurium, only one hexagonal modification is known (Tm. = 452 ° C), similar to selenium. There are weak intermolecular bonds between the chains, so selenium and tellurium melt at relatively low temperatures. Due to this structure, selenium and tellurium form a continuous series of solid solutions with a random alternation of Se atoms and then in helical chains.
Metallic polonium exists in the form of two crystalline modifications: low-temperature cubic and high-temperature hexagonal. For both modifications, an increase in electrical resistance, typical for metals, is observed upon heating.
In the O - S - Se - Te - Po series, with an increase in the radius of atoms, the intermolecular interaction increases, which leads to an increase in the melting and boiling points. Oxygen and sulfur are typical dielectrics, selenium and tellurium are semiconductors, and polonium is a conductor.
3.4 Methods for the extraction of simple substances
In industry, oxygen is produced by electrolysis of water, as well as by multi-stage distillation of liquefied air (since the boiling point of oxygen (-183 ° C) is higher than the boiling point of nitrogen (-195.8 ° C).
Oxygen of very high purity is obtained by the reactions: 2BaO + O2 2BaO2; 2BaO22BaO + О2
Under laboratory conditions, oxygen is obtained: 2KMnO4 K2MnO4 + MnO2 + O2 2KClO3 2KCl + 3O2 2NaNO3 2NaNO2 + O2 2 H2O2 2 H2O + O2 2 HgO 2 Hg + O2.
Sulfur in industry is easier to extract from underground deposits by the mine method or by smelting it from the rock under the action of hot water vapor.
From natural gases containing hydrogen sulfide, sulfur is produced by oxidizing it to sulfur dioxide SO2 and then reacting SO2 with H2S in the presence of iron and aluminum oxide catalysts:
SO2 + H2S → 3S + H2O.
Pyrite is also an important source of sulfur extraction:
FeS2 → FeS + S.
The main source of selenium and tellurium are residues (sludge) after the electrolytic purification of copper, waste from sulfate and pulp and paper production, ores, in which these elements are in the form of chalcogenides.
In industry, sludge containing selenium and tellurium is subjected to oxidative roasting with soda at 650 ° C:
Ag2Se + Na2CO3 + O2 → 2Ag + Na2SeO3 + CO2;
Cu2Te + Na2CO3 + 2O2 → 2CuO + Na2TeO3 + CO2.
Separation of selenium and tellurium is achieved by treatment with a solution of sulfuric acid, while tellurium precipitates in the form of hydrated oxide (which is dissolved in meadows and electrolytically reduced to tellurium), and selenious acid H2SeO3 remains in solution, from which red selenium is precipitated under the action of SO2.
Metallic polonium is obtained by thermal decomposition of polonium sulfide or polonium dioxide in vacuum, followed by sublimation of the metal, as well as the reduction of PoO2 with hydrogen (or the reduction of PoBr2 with dry ammonia at 2000C).
3.5 Chemical properties of oxygen and its compounds
Most metals and non-metals combine with oxygen to form oxides:
4Fe + 3O2 → 2Fe2O3;
4P + 5O2 → P4O10;
S + O2 → SO2.
Oxygen does not directly react only with inert gases, halogens (except fluorine), silver, gold and platinum metals (with the exception of osmium). The reactivity of oxygen is highly dependent on temperature and the presence of water. In some cases, the rate of interaction is so high that an explosion occurs (for example, H2, CH4, CO). Explosive are mixtures of air with coal dust, flour and other finely dispersed substances.
Ozone is a strong oxidizing agent due to the atomic oxygen formed during its decomposition: O3 → O2 + O ∙.
In the atmosphere of ozone, many organic substances ignite spontaneously; it easily oxidizes copper, silver, and mercury; Under the influence of ozone, sulfides are converted to sulfates. In the laboratory, it is produced by passing quiet electrical discharges through oxygen. A characteristic reaction to ozone is blue starch iodine paper:
O3 + 2KI + → I2 + O2 + 2KOH.
In compounds, oxygen exhibits oxidation states +2, +1, -1, -2.
Characteristic oxidation state: 2. Important oxygen compounds with an oxidation state of -2 are oxides. Oxides of all elements are known except for three inert gases - not, Ne, Ar.
In periods, there is a gradual transition from basic oxides to acid ones. So for elements of the third period: Na2O, MgO Al2O3 SiO2, P2O5, SO3, Cl2O7 Basic oxides Amphoteric oxide Acid oxides Basic oxides form bases with water: MgO + H2O → Mg (OH) 2; acidic - acids: SO3 + H2O → H2SO4. Basic and acidic oxides react with each other to form salts: MgO + SO3 → MgSO4.
When alkali metals interact with oxygen, peroxides are formed, in which oxygen has an oxidation state of 1: 2Na + O2 → Na2O2. Among peroxides, hydrogen peroxide H2O2 is of the greatest practical importance. H2O2 is miscible with water in any ratio. In practice, 3% and 30% H2O2 solutions are used (30% hydrogen peroxide is called perhydrol). 80% hydrogen peroxide is obtained by electrolysis of a solution of H2SO4 with a concentration of at least 50% or by the action of dilute sulfuric acid on barium peroxide at 0°C. H2O2 in aqueous solutions exists in the form of a dihydrate, it is a weak acid, the compound is unstable, decomposes when heated and in the light with the release of oxygen: 2 H2O2 → 2 H2O + O2. Hydrogen peroxide is characterized by redox properties, but oxidizing properties predominate.
For example, KNO2 + H2O2-1 → KNO3 + H2O-2 oxidizer Ag2O + H2O2-1 → 2Ag + O20 + H2O reducing agent The property of peroxides to interact with carbon dioxide is used in gas masks and submarines:
2Na2O2 + 2CO2 → 2Na2CO3 + O2
In addition, when interacting with alkali metals, namely: potassium, rubidium, cesium, superoxide is formed (for example, K2O4), where oxygen has an oxidation state of -1/2. When alkali metals interact with ozone, ozonides (for example, KO3) are formed, where oxygen has an oxidation state of -1/3. .
Oxygen exhibits positive oxidation states +1 and +2 only when interacting with fluorine and fluorides (F2O2, OF2).
3.6 Chemical properties of sulfur, selenium, tellurium and polonium
In the series from S to Po, metallic properties are enhanced, reducing activity increases.
Interaction with simple substances:
At room temperature, sulfur reacts only with mercury, but its chemical activity increases significantly when heated (the breaking of S-S bonds is facilitated). Under these conditions, it directly reacts with many simple substances. Sulfur is a fairly strong oxidizing agent, but it can also exhibit reducing properties (in relation to elements with high electronegativity - oxygen, halogens).
Spilled mercury can be collected with finely ground sulfur: Hg + S → HgS.
When heated, selenium, tellurium and polonium easily combine with oxygen, hydrogen, halogens, and also with metals, similarly to sulfur: Se + H2 → H2Se Se + Na → Na2Se
Te + O2 → TeO2 Te 2 Cl2 → TeCl4
In a fluorine atmosphere, S, Se, To burn to form hexafluoride EF6. As a result of the interaction of sulfur with chlorine, which accelerates significantly when heated, S2Cl2, SCl2 are formed. For bromine, only S2Br2 is known; sulfur iodides are unstable. Selenium and tellurium form ECl4 under these conditions.
Interaction with water:
Sulfur and selenium do not interact with water, tellurium reacts with water at a temperature of 100-1600C: Te + 2H2O → TeO2 + 2H2
Interaction with alkalis:
Sulfur, selenium, tellurium interact with alkalis during boiling (disproportionation reaction):
3E + 6KOH → 2K2E + K2EO3 + 3H2O.
In alkali melts, they are oxidized to the highest oxidation states:
2Se + 4KOH + 3O2 → 2K2SeO4 + 2H2O.
Interaction with acids:
Sulfur and selenium do not react with aqueous solutions of acids, although they dissolve in concentrated sulfuric and nitric acids:
3 S + 4 HNO3 (conc.) → 3 SO2 + 4 NO + 2H2O;
Se + 4 HNO3 (conc.) → 2 H2SeO3 + 4NO2 + 2H2O.
Tellurium is also inert to non-oxidizing acids. Diluted nitric acid oxidizes it to teluric acid H2TeO3, concentrated nitric acid to oxohydroxonitrate, and hypochlorous acid to orthoteluric acid H6TeO6.
Polonium readily dissolves in acids to form divalent salts (pink), which rapidly oxidize to tetravalent polonium compounds (yellow).
Characteristic oxidation states in compounds: 2, +2, +4, +6.
The oxidation state -2 is typical for hydrogen compounds of chalcogens, as well as their oxygen-free salts.
Hydrides of all chalcogens are known, with the exception of polonium hydride, which was obtained only in residual quantities and has not been practically characterized.
In the H2S-H2Se-H2Te series, due to an increase in the size of chalcogen atoms, the H-E bond length increases, the binding energy, formation energy and thermodynamic stability of H2E molecules decrease, intermolecular interaction increases, and, accordingly, the melting and boiling points increase.
All H2E hydrides are poisonous gases with an unpleasant odor.
In practice, H2S is produced by the action of dilute acids on metal sulfides. For instance:
FeS + HCl → FeCl2 + H2S.
H2S is a very strong reducing agent.
Silver and copper things in the air or in water containing hydrogen sulfide tarnish due to the formation of the corresponding sulfides: 4Ag + 2H2S + O2 → 2Ag2S + 2H2O
Selenide and teluride (H2Se, H2Te) of hydrogen can be obtained by the action of water or acids on the selenides and telurides of some metals:
Al2Te3 + 6H2O → 2Al(OH) 3 + 3H2Te.
In aqueous solutions, they are rapidly oxidized by atmospheric oxygen:
2H2E + O2 → 2H2O + 2E (where E is Se or To).
Chalcogen waters burn in air with the formation of dioxides, and with a lack of an oxidizing agent, the formation of simple substances is also possible.
2H2E + O2 → 2H2O + 2 EO2.
When dissolved in water, H2E forms the corresponding weak dibasic acids. The ability to dissociate, the strength of acids and reducing properties increase in the series H2S, H2Se, H2Te.
Hydrosulfide acid and its salts behave as reducing agents, oxidizing to free sulfur, sulfur dioxide, polythionates, sulfuric acid. I2 + H2S → 2HI + S; H2S + 4 Cl2 + 4 H2O → H2SO4 + 8 HCl.
Sulfur is characterized by the formation of polysulfanes H2Sn (n = 2-8), they are unstable and easily oxidized and disproportionate to H2S and S.
Among sulfides, only salts of alkali, alkaline earth metals and ammonium are soluble. Sulfides MnS, FeS, CoS, NiS, ZnS are poorly soluble in water, soluble in strong non-oxidizing acids; CdS, CuS, PbS, Sb2S3, As2S3, SnS, SnS2 (gold leaf) are completely insoluble. Many sulfides are colored. Sulfides are obtained by direct interaction of simple substances, exposure to hydrogen sulfide metals or by the reduction of hydroxy acids. Insoluble sulfides can be obtained by the exchange reaction: СdSO4 + H2S → CdS ↓ + H2SO4. Sulfides, especially soluble ones, are strong reducing agents.
In addition, sulfur, selenium and tellurium are capable of forming polysulfides, polyselenides and polytelurides.
3.7 oxygenated chalcogen compounds
+4 oxidation state
Sulfur (IV) oxide SO2 - sulfur dioxide, colorless, with a very stuffy odor. SeO2, TeO2, PoO2 are solids. At the same time, SO2, SeO2 are acidic oxides, TeO2 is an amphoteric oxide, PoO2 is a basic oxide.
SO2, TeO2, SeO2 exhibit redox properties:
2SO2 + O2 → 2 SO3 (SO2 is a reducing agent)
2H2S + SO2 → 2H2O + S (SO2 is an oxidizing agent).
SO2 is of the greatest industrial importance, which is produced by the reaction: 4FeS2 + 11O2 → 2Fe2O3 + 8SO2.
SO2 can be obtained by burning sulfur S + O2 → SO2.
In the laboratory - under the action of concentrated H2SO4 on copper filings:
SO2 is a typical acidic oxide: it reacts with basic oxides and hydroxides to form sulfites: Na2O + SO2. → Na2SO3.
When SO2 dissolves in water, sulfite (or sulfurous) acid is formed.
SO2 + H2O → H2SO3.
This is a dibasic acid of medium strength, which exists only in dilute solutions, oxidizes to H2SO4 in air. 2H2SO3 + O2 → 2H2SO4. Sulfurous acid exhibits redox properties (but reducing properties are more characteristic of it): reducing properties: H2SO3 + Cl2 + H2O → H2SO4 + 2HCl; oxidizing properties: H2SO3 + 2 H2S → 3S + 3H2O.
H2SO3 forms medium salts - sulfites and acidic - hydrosulfites. Sulfites of active metals decompose upon heating to form sulfides and sulfates: 4Na2SO3 → Na2S + 3Na2SO4.
Sulfites are able to restore free sulfur:
Na2SO3 + S → Na2S2O3. sodium thiosulfate
In the series H2SO3 → H2SeO3 → H2TeO3, the strength of acids decreases, and the reducing properties also decrease. Oxidizing properties are most pronounced in H2SeO3.
+6 oxidation state
Among the trioxide, sulfur (VI) oxide SO3 - sulfuric anhydride - is of the greatest importance - a transparent liquid, which is obtained by oxidizing SO2 at a temperature of 400 - 6000C in the presence of a catalyst (V2O5):
2SO2 + O2 → 2SO3.
It is also formed during the thermal decomposition of sulfates and pyrosulfate:
Na2S2O7 → SO3 + Na2SO4.
SO3 is one of the most reactive sulfur compounds. Shows oxidizing properties, for example:
SO3 + C → 2SO2 + CO2.
When SO3 dissolves in water, a strong sulfate (sulfuric) acid is formed.
Anhydrous sulfuric acid H2SO4 is one of the most important products of the chemical industry. It is a colorless oily substance with a density of 1.84 g/cm3. Sulfuric acid is mixed with water in any ratio, this process is accompanied by the release of a large amount of heat, the mixture can even boil, splash (therefore, it is necessary to add acid to water, and not vice versa). Concentrated sulfuric acid is used as a dehumidifier gas desiccant.
It is mined by contact and nitrous methods. In the production of sulfuric acid, SO3 is dissolved not in water, but in a concentrated solution of sulfuric acid, since gaseous SO3, when dissolved in water, reacts with water vapor above the surface of the water, forming a significant part of H2SO4 in the form of a mist. A solution of SO3 in H2SO4 is technically called oleum. In industry, oleum is mined containing 20-65% SO3. To obtain concentrated H2SO4, oleum is mixed with sulfuric acid and contains some water.
In the contact method for the production of sulfuric acid, the following processes occur:
Stage I: 4FeS2 + 11O2 → 2Fe2O3 + 8SO2;
II stage: 2SO2 + O2 → 2SO3 (V2O5 catalyst)
Stage III: SO3 + H2SO4 → H2S2O7 (oleum)
Stage IV: H2S2O7 + H2O → 2H2SO4.
Nitrous way:
SO2 + NO2 → SO3 + NO;
SO3 + H2O → H2SO4 (≈75%)
2NO + O2 → 2NO2 (return to the reaction).
Concentrated sulfuric acid is a fairly strong oxidizing agent.
Active metals (Mg, Zn) are able to reduce concentrated H2SO4 to H2S, S or SO2.
4Zn + 5H2SO4 → 4ZnSO4 + H2S + 4H2O
Cu, Ag, Hg reduce H2SO4 to SO2:
Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O.
Under the action of concentrated acid on Fe, Mn, Te, V, Cr, Al, etc., passivation occurs - oxide films of metals with a high degree of oxidation are formed.
Concentrated H2SO4 is capable of oxidizing C, P, S:
2 H2SO4 + C → CO2 + 2SO2 + 2H2O;
5 H2SO4 + 2P → 2H3PO4 + 5SO2 + 2H2O.
Hydrogen bromide and hydrogen iodide under the action of concentrated H2SO4 are oxidized to free halogens:
8HI + H2SO4 → 4I2 + H2S + 4H2O.
Dilute sulfuric acid does not show noticeable oxidizing properties.
Unlike concentrated H2SO4, dilute H2SO4 reacts only with metals that are in the voltage series up to hydrogen:
H2SO4 + Zn → ZnSO4 + H2 .
Salts of sulfuric acid - sulfates, dissolve well in water. They tend to form hydrates. BaSO4, SrSO4, PbSO4 are poorly soluble. Sulfates are characterized by high thermal stability, decompose upon strong heating with the release of metal oxide, SO2 and O2.
Hydrogen sulfates, when heated, form pyrosulfate or disulfate, salts of disulfate acid:
2NaHSO4 → Na2S2O7 + H2O;
Na2S2O7 → Na2SO4 + SO3.
Selenic anhydride SeO3 is a white hygroscopic substance similar in many properties to SO3. SeO3 is a very strong oxidizing agent. When dissolved in water, it forms selenic acid H2SeO4. It is obtained by oxidizing selenous acid with a concentrated solution of hydrogen peroxide:
H2SeO3 + H2O2 → H2SeO4 + H2O.
TeO3 is a white unstable solid compound, unlike SO3 and SeO3, it practically does not hydrate, but quickly interacts with alkalis to form teluratives. The reducing properties of TeO3 are much less pronounced than in SeO3.
Telluric acid is synthesized by the oxidation of tellurium dioxide or tellurium, as well as by exchange reactions. Orthoteluric acid H6TeO6 is a colorless crystalline solid, highly soluble in water.
In the series SO42- → SeO42- → H5TeO6-, a nonmonotonic change in thermodynamic stability and oxidizing ability is observed: selenic acid and its salts are thermodynamically less stable. Oxidizing properties are most pronounced in H2SeO4: it releases chlorine with concentrated HCl, dissolves copper and even gold without heating 2 Au + 6 H2SeO3 → Au2 (SeO4) 3 + H2SeO3 + 3 H2O.
Sulfur forms a large amount of oxygenated acids, where the formal oxidation state of sulfur is +2, +5 and the like. In addition to the acids discussed above, it is necessary to note thiosulfate acid H2S2O3, dithionic and polythionic acids H2S2O6 and H2SnOn, peroxodisyrcane H2S2O8, peroxomonosyrchane H2SO5, halosulfonic acids HSO3X. It is convenient to consider all these acids as the result of a formal substitution in H2SO4 of the final oxygen atom or hydroxyl group by an isoelectronic group.
For example, thiosulfate acid is an unstable compound that decomposes by the reaction:
H2S2O3 → S ↓ + SO2 + H2O
Its salts - thiosulfates - are strong reducing agents:
Na2S2O3 + 4Cl2 + 5H2O → 2NaHSO4 + 8HCl.
This reaction is the basis for the use of thiosulfate in industry to reduce excess chlorine in the process of bleaching fabrics (“antichlor”). Sodium thiosulfate is used in iodometric titration:
I2 + 2Na2SO3 → 2NaI + Na2S4O6 sodium tetrathionate.
Thiosulfate is decomposed by acids:
Na2S2O3 + H2SO4 → Na2SO4 + S ↓ + SO2 + H2O
Lecture 2
TOPIC : GROUP ITEMS VI B
Questions studied at the lecture:
- general characteristics d elements of group VI.
- Finding in nature and obtaining chromium, molybdenum and tungsten.
- Physical properties of metals.
- Chemical properties of chromium, molybdenum and tungsten.
- The most important compounds of the elements of the chromium subgroup: a) compounds
E (P); b) connections E (Sh); c) compounds E ( VI).
- Chromium peroxide.
side subgroup VI group is represented by the following elements: C r, Mo and W . All of them are d -elements, since they are built up with electrons d - sublevel of the pre-external level. The valence electrons of these elements are the electrons of the outer S -sublevel and preexternal d - sublevel - only 6 electrons.
Electronic configuration of external level and pre-external d-sublevel: С r 3 d 5 4 S 1 ; Mo 4 d 5 5 S 1 ; W 5 d 4 6 S 2 .
d elements of the 6th group take 4th place in their decade d elements, so d the sublevel must contain 4 electrons, and the outer level must contain two s electron, as is observed for tungsten. For chromium and molybdenum, there is a "breakthrough" of one s electron from the outer level to the pre-outer d sublevel, as a result of which each d the orbital will be occupied by one electron, which corresponds to the most stable state of the atom.
│││││ │ (n 1) d → ││││││ (n 1) d
nS │↓│ nS ││
Table 3
Basic parameters of atoms of elements VI B group
|
Atomic radius r a , nm |
Ion radiusr E 6+ , nm |
E E o → E + , eV |
Ar |
chemical activity |
|
|
C r |
0,127 |
0,035 |
6,76 |
│reduce- │ staggers |
|
|
0,137 |
0,065 |
7,10 |
|||
|
0,140 |
0,065 |
7,98 |
Analyzing these data, we can say that there is a common for all d -elements pattern: the radii of atoms from top to bottom in the subgroup increase, but only slightly. Since the mass of atoms in the same row increases greatly, this leads to a densification of the electron shells in molybdenum and especially in tungsten. It is more difficult to extract an electron from such a compacted structure, so the ionization energy increases when passing from chromium to tungsten, as a result of which the chemical activity of elements decreases from top to bottom in the subgroup. Due to the fact that molybdenum and tungsten have approximately the same atomic and ionic radii, their properties are closer to each other than to chromium.
In compounds, chromium and its analogues exhibit oxidation states (CO) of 0, +1, +2, +3, +4, +5 and +6. Maximum S.O. corresponds to the number of valence electrons. Characteristic S.O. chromium +3 and to a lesser extent +6 and +2. Molybdenum and tungsten, like the other 4 d- and 5d - elements, the highest S.O. is most characteristic, that is, +6. Thus, for elements of the subgroup Cr there is a common d pattern elements: increase in the group from top to bottom steady S.O. Therefore, the oxidizing power of the compounds where the elements exhibit the highest S.O. of +6 decreases from top to bottom in the subgroup, as the stability of the compounds in this series increases. For example, in the series of acids:
H 2 C rO 4 │ stability Cr +6 │ oxidizing capacity
H 2 MoO 4 │mo increases+6 │ decreases
H 2 W O 4 │ W +6 │
↓ ↓
For Cr, Mo, W the most typical coordination numbers are 6 and 4. Derivatives are also known, in which k.ch. My W reaches 8.
Examples: [ Cr (OH ) 4 ] - [ Cr (H 2 O ) 6 ] 3+
3- 2-
At the same time, in the formation of connections can participate d -orbitals of the pre-outer level, as well as S - and p-orbitals of the outer level.
The nature of the relationship of elements of subgroup C r in compounds is determined largely by S.O. element. For Cr, Mo, W at low S.O. (+1, +2) ionic bonds are characteristic, and at high S.O. covalent bonds. In accordance with this r+2 O basic oxide, C r 2 +3 O 3 amphoteric, and C r +6 O 3 acidic. Similar to C r (OH) 2 base, C r (OH) 3 amphoteric hydroxide, N 2 C r O 4 acid.
Finding in nature and obtaining chromium, molybdenum and tungsten
The content of chromium in the earth's crust is 0.02% (mass), molybdenum 10-3 % (mass), tungsten 7 ∙ 10-3 % (mass). The main ore of chromium is chromium iron ore. Fe (CrO 2 ) 2 (chromite). Molybdenum occurs as the mineral molybdenite Mo S2 (molybdenum luster), as well as molybdates: PvMoO 4 (wulfenite) and M gMoO 4 . The most important tungsten ores wolframite (mixture FeWO 4 and M nWO 4 ), scheelite Ca WO 4 and stolcite Pv WO 4 .
To obtain pure chromium, an oxide is first obtained Cr2O3 , which is then restored by the aluminothermic method:
Cr 2 O 3 + 2Al → Al 2 O 3 + 2Cr.
For the purposes of metallurgy, chromium is obtained in the form of an alloy with iron (ferrochromium). To do this, chromium iron ore is reduced with coal in an electric furnace.
Fe(CrO 2 ) 2 + 4C → Fe + 2Cr + 4CO.
Molybdenum and tungsten are obtained by converting the minerals listed above into oxides, from which the metal is reduced with hydrogen at high temperatures:
2 Mo S 2 + 7O 2 → 2MoO 3 + 4SO 2
MoO 3 + 3H 2 → Mo + 3H 2 O.
Physical properties of metals
In the form of simple substances, chromium, molybdenum and tungsten are grayish-white shiny metals. All of them are refractory, and tungsten is the most refractory of metals (T sq. = 3380 about C).
The electrical conductivity of metals in the transition from chromium to tungsten generally increases and for molybdenum and tungsten is approximately 30% of the electrical conductivity of silver. The properties of metals are greatly affected by impurities. Thus, technical chromium is one of the hardest metals, while pure chromium is ductile.
Chemical properties of chromium, molybdenum and tungsten
Reactivity in a row Cr Mo W noticeably decreases. Under normal conditions, all three metals noticeably interact only with fluorine:
Me + 3F 2 → MeF 6 (CrF 3 ).
Under normal conditions, these metals are resistant to atmospheric oxygen and water.
In a series of standard electrode potentials of metals, chromium is before hydrogen between zinc and iron, molybdenum is also before hydrogen, but not far from it, and tungsten is after hydrogen. Therefore, chromium displaces hydrogen from dilute HC l and H 2 SO 4.
C r + 2 HCl → CrCl 2 + H 2
Cr + H 2 SO 4 → CrSO 4 + H 2
In concentrated H 2 SO 4 and H NO 3 chromium is passivated in the cold. When heated, chromium slowly dissolves in these acids.
2Cr + 6H 2 SO 4 conc. → Cr 2 (SO 4) 3 + 3SO 2 + 6H 2 O.
Hydrochloric acid and dilute H 2 SO 4 on Mo and W do not work. Molybdenum dissolves only in hot conc. H 2 SO 4 . Tungsten dissolves only in a hot mixture of hydrofluoric and nitric acids
E o + 2H N +5 O 3 + 8 HF → H 2 [E +6 F 8] + 2 N +2 O + 4 H 2 O, where E \u003d Mo, W.
At high temperature, especially in a finely divided state, C r , Mo , W quite easily oxidized by many non-metals:
│ O 2 → Cr 2 O 3
│to
│ S → CrS
│to
Cr + │ Cl 2 → CrCl 3
│to
│ N 2 → c melts
│to
│ C → c fusion
In the case of chromium, compounds with the most stable S.O. are most often formed. chromium (+3). In the interaction of Mo and W with non-metals, as a rule, compounds are formed in which S.O. element is +6.
Common to the elements of the chromium subgroup is the absence of interaction with hydrogen.
The most important compounds of the elements of the chromium subgroup
I .Compounds E (P), that is, S.O. = +2.
1. Black chromium oxide (P) C rO very hard to get. It is formed during the oxidation of chromium amalgam (that is, there is no oxide film) with air under normal conditions: C r + ½ O 2 → CrO.
When heated, oxidation continues to C r2O3.
CrO unstable connection of the main character:
With rO + 2HCl → CrCl 2 + H 2 O.
2.Chromium hydroxide (P) C r(OH)2 a yellow substance insoluble in water, which is obtained by alkalizing solutions of chromium (P) salts:
C rCl 2 + 2NaOH → Cr(OH) 2 ↓ + 2NaCl.
Chromium hydroxide (P) has a basic character, that is, it interacts only with acids and does not dissolve in alkali solutions:
C r (OH) 2 + 2 HCl ↔ CrCl 2 + 2 H 2 O.
C r (OH) 2 is a weak base.
C r (P) forms a number of complexes. For chromium in S.O. +2 is characterized by a coordination number of 6. For example, in aqueous solutions, the ion C r2+ hydrates to form blue aquacomplexes [ Cr (H 2 O ) 6 ] 2+ . Chromium (P) halides absorb gaseous ammonia, forming ammonia:
CrCl 2 + 6NH 3 → Cl 2.
P. Compounds E (W), that is, S.O. = +3
Chromium S.O. +3 in compounds is the most stable.
- Chromium oxide (W) C r 2 O 3 get:
a) when chromium metal powder is heated in air:
4Cr + 3O 2 → 2Cr 2 O 3
b) calcination of chromium oxide ( VI ) or ammonium dichromate:
t o t o
4CrO 3 → 2Cr 2 O 3 + 3O 2 (NH 4) 2 Cr 2 O 7 → Cr 2 O 3 + N 2 + 4H 2 O;
c) when heating chromium hydroxide (Ш):
2 C r (OH) 3 → Cr 2 O 3 + 3H 2 O.
Amorphous oxide Cr 2 O 3 dark green powder. Crystal modification Cr2O3 black powder. It has high refractoriness and is chemically inert.In water, acids and alkali solutions dissolves . However, when melting the oxide Cr (III) with alkalis and basic oxides, salts of metachromous acid are formed:
t o t o
Cr 2 O 3 + 2KOH → 2K Cr O 2 + H 2 O; Cr 2 O 3 + CaO → Ca (Cr O 2) 2.
When alloying Cr 2 O 3 with potassium disulfate, chromium sulfate (III) is formed.
3K 2 S 2 O 7 \u003d 3K 2 SO 4 + 3SO 3;
Cr 2 O 3 + 3SO 3 \u003d Cr 2 (SO 4) 3
─────────────────────────
Cr 2 O 3 + 3K 2 S 2 O 7 \u003d Cr 2 (SO 4) 3 + 3K 2 SO 4.
These reactions show an amphoteric character Cr 2 O 3 .
- Chromium hydroxide (SH) Cr(OH)3 precipitated from solutions of chromium (III) salts with alkalis in the form of a voluminous gelatinous grayish-greenish precipitate, insoluble in water.
C r +3 + 3 OH - → Cr (OH) 3 ↓;
CrCl 3 + 3NaOH → Cr(OH) 3 ↓ + 3NaCl.
Chromium hydroxide Cr(OH) 3 has an amphoteric character and freshly obtained chromium (III) hydroxide is easily soluble in acids and alkali solutions.
Cr (OH) 3 + 3HC l ↔ CrCl 3 + 3 H 2 O
Cr(OH) 3 + NaOH ↔ Na.
The basic and especially acidic properties of chromium (III) hydroxide are weakly expressed. Therefore salt Cr+3 undergo significant hydrolysis in solutions, and soluble chromites, in the absence of an excess of alkali, are almost completely hydrolyzed.
Cr 3+ + HOH ↔ Cr(OH) 2+ + H + ;
3- + HOH ↔ 2- + OH-.
Alum. Cr (W), like A l (Sh), forms with active metals and NH4+ double salts alum. Example: KS r (SO 4 ) 2 ∙12 H 2 O and (NH 4 )Cr (SO 4 ) 2 ∙12 H 2 O. They are formed by the interaction of solutions M 2 +1 SO 4 and Cr 2 (SO 4 ) 3 . In solution, these salts dissociate:
K Cr(SO 4 ) 2 ↔ K + + Cr 3+ + 2 SO 4 2-.
Cr 3+ + H 2 O ↔ Cr (OH) 2+ + H + - acidic.
Cr (W) as well as Cr (P) - active complexing agent. coordination number Cr (W) is equal to 6 and 4.
Examples: aquacomplex [ Cr (H 2 O ) 6 ] 3+ - blue-violet color;
hydroxocomplex [ Cr(OH)6]3- - emerald green color;
amino complex [ Cr (NH 3 ) 6 ] 3+ - purple color.
Sh. Compounds E (VI), that is, S.O. = +6
Compounds in which S.O. element is +6, most characteristic of Mo, W and to a lesser extent for cr.
- oxides EO 3 (Cr O 3 , MoO3 andWO3 ).
MoE3 andWO3 formed when metals are heated in air:
to
2E + 3O2 → 2EO3.
CrO3 can only be obtained indirectly, since when heatedCrformed in the airCr2 O3 .
CrO3 precipitated when an excess of concentratedH2 SO4 to a saturated solution of chromate:
To2 CrO4 + H2 SO4 conc. = CrO3 ↓ + K2 SO4 + H2 O
MoE3 colorless crystals;
WO3 light yellow crystals;
CrO3 dark red crystals.
MoE3 andWO3 are stable and, when heated, pass into the gas phase without decomposition. When heatedCrO3 easily decomposes, releasing O2 .
to
4 CrO3 → 2 Cr2 O3 + 3O2 .
CrO3 easily soluble in water, forming chromic acid
CrO3 + H2 O → H2 CrO4 .
MoE3 andWO3 do not dissolve in water. The acidic nature of these oxides manifests itself when dissolved in alkali solutions:
EO3 + 2KOH → K2 EO4 + H2 O.
In this case, salts of chromic, molybdic and tungstic acids are formed, respectively.
- Hydroxides E (VI) - H2 EO4
H2 CrO4 , N2 MoE4 , N2 WO4 .
Chromic acid is obtained by dissolvingCrO3 in H2 A. Molybdic and tungstic acids are obtained indirectly - by acidifying solutions of their salts:
(NH4 ) 2 MaboutO4 + 2HNO3 → H2 MaboutO4 ↓+2NH4 NO3 .
The strength of acids in the series H2 FROMrO4 N2 MoE4 - H2 WO4 is decreasing.
Chromic acid H2 FROMrO4 medium strength acid (K1 = 2∙10 -1 , TO2 =3∙10 -7 ), is not isolated in the free state.
H2 MoO4 released in free form. It is a white powder, almost insoluble in water. Constants of the first stage of acidic and basic dissociationH2 MoO4 have order respectively 10-2 and 10-13 .
3. Salt.
Most salts of acids H2 EO4 sparingly soluble in H2 A. Only salts dissolve well.Na+ and K+ . Chromates are stained yellow with ion CrO4 2- , molybdates and tungstates are colorless. All salts of chromic acids are poisonous.
When acidifying a chromate solution, hydrochromate is formed, which is very unstable and, releasing water, turns into dichromate:
2 CrO4 2- + 2H+ ↔ 2Н СrO4 - ↔ Cr2 O7 2- + H2 O.
In this case, the yellow color of the solution changes to orange, characteristic of the C ion.r2 O7 2- . This balance is very fluid. It can be shifted by changing the pH of the medium: adding acids (H ions) to the solution+ ) shifts the equilibrium towards the formation of dichromate, and the addition of alkali - to the left (due to the binding of H ions+ ). Thus, in the presence of an excess of OH ions- only C ions exist in solutionrO4 2- , and with an excess of hydrogen ions - ions Cr2 O7 2- .
2K2 FROMrO4 + H2 SO4 → K2 FROMr2 O7 + K2 SO4 + H2 O;
To2 FROMr2 O7 + 2KOH → 2K2 FROMrO4 + H2 O.
Bichromic acid H2 FROMr2 O7 much stronger than chromium, K2 = 2∙10 -2 . It is also not isolated in free form.
ConnectionsCr (VI) - strong oxidizers, pass in redox reactions to derivatives of Cr(III). Most strongly oxidizing propertiesCr(VI) are expressed in acid medium.
To2 FROMr2 O7 + 6KJ + 7H2 SO4 →Cr2 (SO4 ) 3 + 3J2 + 4K2 SO4 + 7H2 Oh
In this case, the orange color of the potassium bichromate solution is replaced by a green or greenish-violet color of the solutions.Cr3+ .
In contrast to chromium, the oxidizing properties of Mo(VI) andW (VI) even in an acidic environment appear only when interacting with the strongest reducing agents, for example, with hydrogen at the time of isolation.
Chromium peroxide
When an acidic solution of chromate or dichromate is treated with hydrogen peroxide, chromium peroxide C is formedrO(O2 ) 2 or withrO5 .
FROMr2 O7 2- + 4H2 O2 + 2H+ = 2 CrO (O2 ) 2 + 5 H2 O.
CrO (O2 ) 2 blue, unstable in aqueous solution and decomposes into oxygen and aqua complexes [Cr(H2 O) 6 ] 3+ .
Chromium peroxide is stable in ether and forms a peroxo complex
CrO (O2 ) 2 ∙ L, whereLether, pyridine, etc. These complexes have the form of a pentagonal pyramid with an oxygen atom at the top:
Chromium peroxide in its composition contains two peroxide groups (-O-O-), due to which it exhibits oxidizing properties.
Group VI p-elements include oxygen ( O), sulfur ( S), selenium ( Se), tellurium ( Te), polonium ( Ro).
The general electronic formula of the valence band of atoms has the form ns 2 np 4, from which it follows that there are six electrons on the outer electron layer of the atoms of the elements under consideration and they can exhibit even valencies 2, 4, 6. Since the oxygen atom does not have a d-sublevel, therefore, excited states are impossible and the oxygen valence is only 2.
All elements of this subgroup, with the exception of polonium, are non-metals.
Oxygen is the most abundant element in the earth's crust. The oxygen molecule is diatomic (O 2). Under normal conditions, it is a colorless and odorless gas, poorly soluble in water. The Earth's atmosphere contains 21% (by volume) of oxygen. In natural compounds, oxygen occurs in the form of oxides (H 2 O, SiO 2) and salts of oxygen acids. An industrial method for obtaining oxygen is the rectification of liquid air. Air, nitrogen and oxygen are stored in a liquid state in Dewar flasks.
Oxygen plays an important role in nature. It is involved in a vital process - breathing. Its application is diverse: the production of sulfuric and nitric acids, the smelting of metals, etc.
The allotropic modification of oxygen is ozone (O 3). Ozone is one of the strongest oxidizers; in oxidative activity, it is second only to fluorine. It oxidizes all metals except gold and platinum metals, as well as most non-metals. The Earth's stratosphere contains the ozone layer, which absorbs most of the ultraviolet radiation.
The following reaction is used to detect ozone:
2KI + O 3 + H 2 O \u003d I 2 + 2KOH + O 2.
An important oxygen compound is H 2 O 2 (H–O–O–H), hydrogen peroxide.
The oxygen atoms in H 2 O 2 are in an intermediate oxidation state of -1 and therefore can exhibit both oxidizing and reducing
properties. For example:
1) H 2 O 2 + 2KI \u003d I 2 + 2KOH
2O - 1 + 2 e \u003d 2O - 2, 2I - – 2e = I 2 ;
2) 5H 2 O 2 + 2KMnO 4 + 3H 2 SO 4 \u003d 5O 2 + 2MnSO 4 + K 2 SO 4 + 8H 2 O
2O - 1 – 2e \u003d O 2, Mn +7 + 5e \u003d Mn +2.
Sulfur exists in several allotropic modifications: rhombic, monoclinic, plastic. Under normal conditions, sulfur is a solid yellow substance, insoluble in water, but highly soluble in organic solvents.
Sulfur interacts directly with many metals (Zn, Al, Fe, Сu, alkali and alkaline earth metals). For example,
2Al + 3S → Al 2 S 3 .
At high temperatures, sulfur reacts with hydrogen to form hydrogen sulfide (H 2 S) - a colorless gas with a characteristic odor (rotten eggs)
H 2 + S → H 2 S.
Hydrogen sulfide is very toxic and can cause severe poisoning.
Hydrosulfuric acid is a weak dibasic acid:
H 2 S ↔ H + + NS -, K 1 \u003d 6 ∙ 10 - 8;
HS - ↔ H + + S 2 -, K 2 \u003d 1 ∙ 10 - 14.
Hydrosulfuric acid forms salts - sulfides, many of which are characterized by low solubility. For example:
CuSO 4 + H 2 S ↔ CuS ↓ + H 2 SO 4,
Cu 2+ + SO 4 2 – + H 2 S ↔ CuS↓ + 2H + + SO 4 2 – ,
Cu 2+ + H 2 S ↔ CuS ↓ + 2H + .
When ignited in air, hydrogen sulfide burns with a bluish flame.
2H 2 S + 3O 2 → 2SO 2 + 2H 2 O (in excess oxygen).
Sulfur oxide (IV) is formed during the combustion of sulfur in air. It is highly soluble in water with the formation of sulfurous acid:
SO 2 + H 2 O ↔ H 2 SO 3.
Sulfurous acid is a weak dibasic acid. It is a good reducing agent and oxidizes to sulfuric acid:
2H 2 SO 3 + O 2 → 2H 2 SO 4.
At high temperature, in the presence of a catalyst (V 2 O 5 , platinum-based alloys), sulfur dioxide is oxidized by oxygen to trioxide, which
in turn used to produce sulfuric acid
SO 3 + H 2 O → H 2 SO 4.
H 2 SO 4 is a strong dibasic acid. In dilute aqueous solutions, it dissociates almost completely H 2 SO 4 → 2H + + SO 4 2 -. When concentrated sulfuric acid is dissolved in water, a large amount of heat is released.
Concentrated sulfuric acid, especially when hot, is a vigorous oxidizing agent. It is reduced by metals to SO 2, S or H 2 S. The more active the metal, the more deeply the acid is reduced:
Cu + 2H 2 SO 4 (conc.) → CuSO 4 + SO 2 + 2H 2 O,
3Zn + 4Н 2 SO 4 (conc.) → 3ZnSO 4 + S↓ + 4Н 2 O,
Salts of sulfuric acid, sulfates, as a rule, are highly soluble. They are isolated from aqueous solutions in the form of crystalline hydrates, called vitriols: CuSO 4 5H 2 O, FeSO 4 7H 2 O, etc. Sulfuric acid also forms double salts - alum, existing in a crystalline state:
K 2 SO 4 Al 2 (SO 4) 3 24H 2 O or KAl (SO 4) 2 12H 2 O, etc.
In the series H 2 O - H 2 S - H 2 Se - H 2 Te with an increase in molecular weights, an increase in boiling points should be observed. As can be seen from Figure 17.1, this dependence is observed, with the exception of H 2 O.
Figure 17.1 - Dependence of the boiling points of hydrogen compounds of p-elements of group VI on the molecular weight of the compound
Previously, it was shown that the abnormally high boiling point of H 2 O is a consequence of the formation of hydrogen bonds between individual water molecules.
Selenium properties close to sulfur. Selenic acid (H 2 SeO 4) is also a strong acid. Selenium is an important biological trace element.
Tellurium forms a very weak orthotelluric acid H 6 TeO 6 . Selenium and tellurium are semiconductors. Tellurium serves as an alloying addition to lead, improving its mechanical properties. All compounds of selenium and tellurium are poisonous.
