Carbon - a characteristic of the element and chemical properties. Oxides of carbon (II) and (IV). Transition metal carbonyls The effect of carbon monoxide on the human body

Carbon (C) is a typical non-metal; in the periodic system is in the 2nd period of the IV group, the main subgroup. Ordinal number 6, Ar = 12.011 amu, nuclear charge +6.

Physical properties: carbon forms many allotropic modifications: diamond one of the hardest substances graphite, coal, soot.

A carbon atom has 6 electrons: 1s 2 2s 2 2p 2 . The last two electrons are located in separate p-orbitals and are unpaired. In principle, this pair could occupy one orbital, but in this case the interelectron repulsion strongly increases. For this reason, one of them takes 2p x, and the other, either 2p y , or 2p z-orbitals.

The difference between the energies of the s- and p-sublevels of the outer layer is small, therefore, the atom quite easily passes into an excited state, in which one of the two electrons from the 2s-orbital passes to a free 2r. A valence state arises having the configuration 1s 2 2s 1 2p x 1 2p y 1 2p z 1 . It is this state of the carbon atom that is characteristic of the diamond lattice - the tetrahedral spatial arrangement of hybrid orbitals, the same length and energy of bonds.

This phenomenon is known to be called sp 3 -hybridization, and the resulting functions are sp 3 -hybrid . The formation of four sp 3 bonds provides the carbon atom with a more stable state than three rr- and one s-s-bond. In addition to sp 3 hybridization, sp 2 and sp hybridization are also observed at the carbon atom . In the first case, there is a mutual overlap s- and two p-orbitals. Three equivalent sp 2 - hybrid orbitals are formed, located in the same plane at an angle of 120 ° to each other. The third orbital p is unchanged and directed perpendicular to the plane sp2.

In sp hybridization, the s and p orbitals overlap. An angle of 180° arises between the two equivalent hybrid orbitals formed, while the two p-orbitals of each of the atoms remain unchanged.

Allotropy of carbon. diamond and graphite

In a graphite crystal, carbon atoms are located in parallel planes, occupying the vertices of regular hexagons in them. Each of the carbon atoms is linked to three adjacent sp 2 hybrid bonds. Between parallel planes, the connection is carried out due to van der Waals forces. Free p-orbitals of each of the atoms are directed perpendicular to the planes of covalent bonds. Their overlap explains the additional π-bond between carbon atoms. So from the valence state in which carbon atoms are in a substance, the properties of this substance depend.

Chemical properties of carbon

The most characteristic oxidation states: +4, +2.

At low temperatures, carbon is inert, but when heated, its activity increases.

Carbon as a reducing agent:

- with oxygen
C 0 + O 2 - t ° \u003d CO 2 carbon dioxide
with a lack of oxygen - incomplete combustion:
2C 0 + O 2 - t° = 2C +2 O carbon monoxide

- with fluorine
C + 2F 2 = CF 4

- with steam
C 0 + H 2 O - 1200 ° \u003d C + 2 O + H 2 water gas

— with metal oxides. In this way metal is smelted from ore.
C 0 + 2CuO - t ° \u003d 2Cu + C +4 O 2

- with acids - oxidizing agents:
C 0 + 2H 2 SO 4 (conc.) \u003d C +4 O 2 + 2SO 2 + 2H 2 O
С 0 + 4HNO 3 (conc.) = С +4 O 2 + 4NO 2 + 2H 2 O

- forms carbon disulfide with sulfur:
C + 2S 2 \u003d CS 2.

Carbon as an oxidizing agent:

- forms carbides with some metals

4Al + 3C 0 \u003d Al 4 C 3

Ca + 2C 0 \u003d CaC 2 -4

- with hydrogen - methane (as well as a huge amount organic compounds)

C 0 + 2H 2 \u003d CH 4

- with silicon, forms carborundum (at 2000 ° C in an electric furnace):

Finding carbon in nature

Free carbon occurs as diamond and graphite. In the form of compounds, carbon is found in minerals: chalk, marble, limestone - CaCO 3, dolomite - MgCO 3 * CaCO 3; bicarbonates - Mg (HCO 3) 2 and Ca (HCO 3) 2, CO 2 is part of the air; carbon is the main integral part natural organic compounds - gas, oil, hard coal, peat, is part of organic matter, proteins, fats, carbohydrates, amino acids that are part of living organisms.

Inorganic carbon compounds

Neither C 4+ ions, nor C 4- - under any normal chemical processes are not formed: in carbon compounds there are covalent bonds different polarity.

Carbon monoxide (II) SO

Carbon monoxide; colorless, odorless, sparingly soluble in water, soluble in organic solvents, poisonous, bp = -192°C; t sq. = -205°C.

Receipt
1) In industry (in gas generators):
C + O 2 = CO 2

2) In the laboratory - thermal decomposition formic or oxalic acid in the presence of H 2 SO 4 (conc.):
HCOOH = H2O + CO

H 2 C 2 O 4 \u003d CO + CO 2 + H 2 O

Chemical properties

Under ordinary conditions, CO is inert; when heated - reducing agent; non-salt-forming oxide.

1) with oxygen

2C +2 O + O 2 \u003d 2C +4 O 2

2) with metal oxides

C +2 O + CuO \u003d Cu + C +4 O 2

3) with chlorine (in the light)

CO + Cl 2 - hn \u003d COCl 2 (phosgene)

4) reacts with alkali melts (under pressure)

CO + NaOH = HCOONa (sodium formate)

5) forms carbonyls with transition metals

Ni + 4CO - t° = Ni(CO) 4

Fe + 5CO - t° = Fe(CO) 5

Carbon monoxide (IV) CO2

Carbon dioxide, colorless, odorless, solubility in water - 0.9V CO 2 dissolves in 1V H 2 O (under normal conditions); heavier than air; t°pl.= -78.5°C (solid CO 2 is called "dry ice"); does not support combustion.

Receipt

1. Thermal decomposition of salts of carbonic acid (carbonates). Limestone firing:

CaCO 3 - t ° \u003d CaO + CO 2

1. The action of strong acids on carbonates and bicarbonates:

CaCO 3 + 2HCl \u003d CaCl 2 + H 2 O + CO 2

NaHCO 3 + HCl \u003d NaCl + H 2 O + CO 2

ChemicalpropertiesCO2
Acid oxide: reacts with basic oxides and bases to form carbonic acid salts

Na 2 O + CO 2 \u003d Na 2 CO 3

2NaOH + CO 2 \u003d Na 2 CO 3 + H 2 O

NaOH + CO 2 \u003d NaHCO 3

May exhibit oxidizing properties at elevated temperatures

C +4 O 2 + 2Mg - t ° \u003d 2Mg +2 O + C 0

Qualitative reaction

Turbidity of lime water:

Ca (OH) 2 + CO 2 \u003d CaCO 3 ¯ (white precipitate) + H 2 O

It disappears when CO 2 is passed through lime water for a long time, because. insoluble calcium carbonate is converted to soluble bicarbonate:

CaCO 3 + H 2 O + CO 2 \u003d Ca (HCO 3) 2

carbonic acid and itssalt

H2CO3 — Weak acid, exists only in aqueous solution:

CO 2 + H 2 O ↔ H 2 CO 3

Dual base:
H 2 CO 3 ↔ H + + HCO 3 - Acid salts - bicarbonates, bicarbonates
HCO 3 - ↔ H + + CO 3 2- Medium salts - carbonates

All properties of acids are characteristic.

Carbonates and bicarbonates can be converted into each other:

2NaHCO 3 - t ° \u003d Na 2 CO 3 + H 2 O + CO 2

Na 2 CO 3 + H 2 O + CO 2 \u003d 2NaHCO 3

Metal carbonates (except alkali metals) decarboxylate when heated to form an oxide:

CuCO 3 - t ° \u003d CuO + CO 2

Qualitative reaction- "boiling" under the action of a strong acid:

Na 2 CO 3 + 2HCl \u003d 2NaCl + H 2 O + CO 2

CO 3 2- + 2H + = H 2 O + CO 2

Carbides

calcium carbide:

CaO + 3 C = CaC 2 + CO

CaC 2 + 2 H 2 O \u003d Ca (OH) 2 + C 2 H 2.

Acetylene is released when zinc, cadmium, lanthanum and cerium carbides react with water:

2 LaC 2 + 6 H 2 O \u003d 2La (OH) 3 + 2 C 2 H 2 + H 2.

Be 2 C and Al 4 C 3 are decomposed by water to form methane:

Al 4 C 3 + 12 H 2 O \u003d 4 Al (OH) 3 \u003d 3 CH 4.

Titanium carbides TiC, tungsten W 2 C (hard alloys), silicon SiC (carborundum - as an abrasive and material for heaters) are used in technology.

cyanides

obtained by heating soda in an atmosphere of ammonia and carbon monoxide:

Na 2 CO 3 + 2 NH 3 + 3 CO \u003d 2 NaCN + 2 H 2 O + H 2 + 2 CO 2

Hydrocyanic acid HCN is an important product chemical industry, widely used in organic synthesis. Its world production reaches 200 thousand tons per year. The electronic structure of the cyanide anion is similar to carbon monoxide (II), such particles are called isoelectronic:

C = O:[:C = N:]-

Cyanides (0.1-0.2% water solution) are used in gold mining:

2 Au + 4 KCN + H 2 O + 0.5 O 2 \u003d 2 K + 2 KOH.

When cyanide solutions are boiled with sulfur or when solids are fused, thiocyanates:
KCN + S = KSCN.

When cyanides of low-active metals are heated, cyanide is obtained: Hg (CN) 2 \u003d Hg + (CN) 2. cyanide solutions are oxidized to cyanates:

2KCN + O2 = 2KOCN.

Cyanic acid exists in two forms:

H-N=C=O; H-O-C = N:

In 1828, Friedrich Wöhler (1800-1882) obtained urea from ammonium cyanate: NH 4 OCN \u003d CO (NH 2) 2 by evaporating an aqueous solution.

This event is usually seen as the victory of synthetic chemistry over "vitalistic theory".

There is an isomer of cyanic acid - fulminic acid

H-O-N=C.
Its salts (mercury fulminate Hg(ONC) 2) are used in impact igniters.

Synthesis urea(carbamide):

CO 2 + 2 NH 3 \u003d CO (NH 2) 2 + H 2 O. At 130 0 C and 100 atm.

Urea is an amide of carbonic acid, there is also its "nitrogen analogue" - guanidine.

Carbonates

The most important inorganic compounds of carbon are salts of carbonic acid (carbonates). H 2 CO 3 is a weak acid (K 1 \u003d 1.3 10 -4; K 2 \u003d 5 10 -11). Carbonate buffer supports carbon dioxide balance in the atmosphere. The oceans have a huge buffer capacity because they are an open system. The main buffer reaction is the equilibrium during the dissociation of carbonic acid:

H 2 CO 3 ↔ H + + HCO 3 -.

When acidity decreases, additional absorption occurs carbon dioxide from the atmosphere to form acid:
CO 2 + H 2 O ↔ H 2 CO 3.

With an increase in acidity, carbonate rocks (shells, chalk and limestone deposits in the ocean) dissolve; this compensates for the loss of hydrocarbonate ions:

H + + CO 3 2- ↔ HCO 3 -

CaCO 3 (tv.) ↔ Ca 2+ + CO 3 2-

Solid carbonates are converted into soluble hydrocarbons. It is this process of chemical dissolution of excess carbon dioxide that counteracts the "greenhouse effect" - global warming due to the absorption of the Earth's thermal radiation by carbon dioxide. Approximately one third of the world's production of soda (sodium carbonate Na 2 CO 3) is used in the manufacture of glass.

Carbon forms two extremely stable oxides (CO and CO 2), three much less stable oxides (C 3 O 2 , C 5 O 2 and C 12 O 9), a number of unstable or poorly studied oxides (C 2 O, C 2 O 3 etc.) and non-stoichiometric graphite oxide. Among the listed oxides, CO and CO 2 play a special role.

DEFINITION

carbon monoxide under normal conditions, a combustible gas, colorless and odorless.

It is quite toxic due to its ability to form a complex with hemoglobin, which is about 300 times more stable than the oxygen-hemoglobin complex.

DEFINITION

Carbon dioxide under normal conditions - colorless gas, about 1.5 times heavier than air, due to which it can be poured, like a liquid, from one vessel to another.

The mass of 1 liter of CO 2 under normal conditions is 1.98 g. The solubility of carbon dioxide in water is low: 1 volume of water at 20 o C dissolves 0.88 volumes of CO 2 , and at 0 o C - 1.7 volumes.

Direct oxidation of carbon with a lack of oxygen or air leads to the formation of CO, with a sufficient amount of them, CO 2 is formed. Some properties of these oxides are presented in table. one.

Table 1. Physical properties of carbon oxides.

Obtaining carbon monoxide

Pure CO can be obtained in the laboratory by dehydration formic acid(HCOOH) concentrated sulfuric acid at ~140 °C:

HCOOH \u003d CO + H 2 O.

In small quantities, carbon dioxide can be easily obtained by the action of acids on carbonates:

CaCO 3 + 2HCl \u003d CaCl 2 + H 2 O + CO 2.

On an industrial scale, CO 2 is produced mainly as a by-product in the ammonia synthesis process:

CH 4 + 2H 2 O \u003d CO 2 + 4H 2;

CO + H 2 O \u003d CO 2 + H 2.

Large quantities carbon dioxide is produced by burning limestone:

CaCO 3 \u003d CaO + CO 2.

Chemical properties of carbon monoxide

Carbon monoxide is chemically active at high temperatures. It manifests itself as a strong reducing agent. Reacts with oxygen, chlorine, sulfur, ammonia, alkalis, metals.

CO + NaOH = Na(HCOO) (t = 120 - 130 o C, p);

CO + H 2 \u003d CH 4 + H 2 O (t \u003d 150 - 200 o C, kat. Ni);

CO + 2H 2 \u003d CH 3 OH (t \u003d 250 - 300 o C, kat. CuO / Cr 2 O 3);

2CO + O 2 \u003d 2CO 2 (kat. MnO 2 / CuO);

CO + Cl 2 \u003d CCl 2 O (t \u003d 125 - 150 o C, cat. C);

4CO + Ni = (t = 50 - 100 o C);

5CO + Fe = (t = 100 - 200 o C, p).

Carbon dioxide exhibits acid properties: reacts with alkalis, ammonia hydrate. Recovering active metals, hydrogen, carbon.

CO 2 + NaOH dilute = NaHCO 3 ;

CO 2 + 2NaOH conc \u003d Na 2 CO 3 + H 2 O;

CO 2 + Ba(OH) 2 = BaCO 3 + H 2 O;

CO 2 + BaCO 3 + H 2 O \u003d Ba (HCO 3) 2;

CO 2 + NH 3 × H 2 O \u003d NH 4 HCO 3;

CO 2 + 4H 2 \u003d CH 4 + 2H 2 O (t \u003d 200 o C, kat. Cu 2 O);

CO 2 + C \u003d 2CO (t\u003e 1000 o C);

CO 2 + 2Mg \u003d C + 2MgO;

2CO 2 + 5Ca = CaC 2 + 4CaO (t = 500 o C);

2CO 2 + 2Na 2 O 2 \u003d 2Na 2 CO 3 + O 2.

Application of carbon monoxide

Carbon monoxide is widely used as a fuel in the form of producer gas or water gas, and is also formed during the separation of many metals from their oxides by reduction with coal. Generator gas is obtained by passing air through hot coal. It contains about 25% CO, 4% CO2 and 70% N 2 with traces of H 2 and CH 4 62.

The use of carbon dioxide is most often due to its physical properties. It is used as a cooling agent, for carbonating beverages, for producing lightweight (foamed) plastics, and as a gas for creating an inert atmosphere.

Examples of problem solving

EXAMPLE 1

EXAMPLE 2

Exercise	Determine how many times heavier than air is carbon monoxide (IV)CO 2.
Solution	The ratio of the mass of a given gas to the mass of another gas taken in the same volume, at the same temperature and the same pressure, is called the relative density of the first gas over the second. This value shows how many times the first gas is heavier or lighter than the second gas. The relative molecular weight of air is taken equal to 29 (taking into account the content of nitrogen, oxygen and other gases in the air). It should be noted that the concept of "relative molecular weight of air" is used conditionally, since air is a mixture of gases. D air (CO 2) \u003d M r (CO 2) / M r (air); D air (CO 2) \u003d 44 / 29 \u003d 1.517. M r (CO 2) \u003d A r (C) + 2 × A r (O) \u003d 12 + 2 × 16 \u003d 12 + 32 \u003d 44.
Answer	Carbon monoxide (IV)CO 2 is 1.517 times heavier than air.

Chemical properties: At ordinary temperatures, carbon is chemically inert, at sufficiently high temperatures it combines with many elements, and exhibits strong reducing properties. The chemical activity of different forms of carbon decreases in the series: amorphous carbon, graphite, diamond, in air they ignite at temperatures above 300-500 °C, 600-700 °C and 850-1000 °C, respectively. Oxidation states +4 (for example, CO 2), -4 (e.g., CH 4), rarely +2 (CO, metal carbonyls), +3 (C 2 N 2); electron affinity 1.27 eV; the ionization energy during the successive transition from C 0 to C 4+ is 11.2604, 24.383, 47.871 and 64.19 eV, respectively.

The three best known carbon monoxide:

1) Carbon monoxide CO(It is a colorless, odorless, tasteless gas. Combustible. The so-called "carbon monoxide smell" is actually the smell of organic impurities.)

2) Carbon dioxide CO 2 (Non-toxic, but does not support breathing. A high concentration in the air causes suffocation. A lack of carbon dioxide is also dangerous. Carbon dioxide in animal organisms also has physiological significance, for example, it is involved in the regulation of vascular tone)

3) Tricarbon dioxide C 3 O 2 (colored poisonous gas with a pungent, suffocating odor, easily polymerized under normal conditions to form a product that is insoluble in water, yellow, red or purple.)

Compounds with non-metals have their own names - methane, tetrafluoromethane.

Products burning carbon in oxygen are CO and CO 2 (carbon monoxide and carbon dioxide, respectively). Also known to be unstable underoxide carbon C 3 O 2 (melting point −111 ° C, boiling point 7 ° C) and some other oxides (for example C 12 O 9, C 5 O 2, C 12 O 12). Graphite and amorphous carbon begin to react with hydrogen at 1200 °C, with fluorine at 900 °C.

carbon dioxide reacts with water, forming a weak carbonic acid - H 2 CO 3, which forms salts - carbonates. Calcium carbonates (mineral forms - chalk, marble, calcite, limestones, etc.) and magnesium are the most widespread on Earth.

43 Question. Silicon

Silicon (Si) - stands in period 3, group IV of the main subgroup periodic. systems.

Phys. saints: silicon exists in two modifications: amorphous and crystalline. Amorphous silicon - brown powder is dissolved in metal melts. Crystalline silicon is dark gray crystals with a steely luster, hard and brittle. Silicon is made up of three isotopes.

Chem. saints: electronic configuration: 1s 2 2s 2 2p 6 3 s 2 3p 2 . Silicon is a non-metal. On the external energy ur-not silicon has 4 e, which determines its oxidation states: +4, -4, -2. Valence - 2, 4. Amorphous silicon has a greater reactivity than crystalline. Under normal conditions, it interacts with fluorine: Si + 2F 2 = SiF 4.

From to-t silicon interacts only with a mixture of nitric and hydrofluoric acids:

With respect to metals, it behaves differently: it dissolves well in molten Zn, Al, Sn, Pb, but does not react with them; with other melts of metals - with Mg, Cu, Fe, silicon interacts with the formation of silicides: Si + 2Mg = Mg2Si. Silicon burns in oxygen: Si + O2 = SiO2 (sand).

Receipt: Free silicon can be obtained by calcining fine white sand with magnesium, which, according to chem. composition is almost pure silicon oxide, SiO2 + 2Mg \u003d 2MgO + Si.

Silicon(II) oxide SiO- Resin-like amorphous in-in, under normal conditions, resistant to oxygen. Refers to non-salt-forming oxides. SiO does not occur in nature. Gaseous silicon monoxide has been found in gas and dust clouds of interstellar media and on sunspots. Receipt: Silicon monoxide can be obtained by heating silicon in a lack of oxygen at a temperature of 2Si + O 2 weeks → 2SiO. When heated in an excess of oxygen, silicon oxide (IV) SiO2 is formed: Si + O 2 wt → SiO 2.

SiO is also formed during the reduction of SiO2 with silicon at high temperatures: SiO 2 + Si → 2SiO.

Silicon oxide (IV) SiO2 - colorless crystals have high hardness and strength. Saints: Belongs to the acid group. oxides. When heated, it interacts with the main. oxides and alkalis. P-ryatsya in hydrofluoric acid. SiO2 belongs to the group of glass-forming oxides, i.e. prone to the formation of a supercooled melt - glass. One of the best dielectrics (does not conduct electric current). It has an atomic crystal lattice.

Nitride - binary inorganic. chemical compound, which is a compound of silicon and nitrogen Si 3 N 4 . Saints: Silicon nitride has good mech.and fiz.-chem. St. you. Thanks to the silicon nitride bond means. the operational properties of refractories based on silicon carbide, periclase, forsterite, etc. are improved. Nitride-bonded refractories have high thermal and wear resistance, have excellent cracking resistance, and also to the impact of, alkalis, aggressive melts and vapors of metals.

Silicon(IV) chloride silicon - colorless in-in, chem. cat formula. SiCl 4. It is used in the production of silicon-organic. connections; used to create smoke screens. Techn. silicon tetrachloride is intended for the production of ethyl silicates, aerosil.

Silicon carbide- binary inorganic. chem. compound of silicon with carbon SiC. In nature, it occurs in the form of an extremely rare mineral - moissanite.

Silicon dioxide or silica- stable connection Si, is widely distributed in nature. It reacts with its fusion with alkalis, basic oxides, forming salts of silicic acid - silicates. Receipt: in industry, pure silicon is obtained by reducing silicon dioxide with coke in electric furnaces: SiO 2 + 2C \u003d Si + 2CO 2.

In the laboratory, silicon is obtained by calcining white sand with magnesium or aluminum:

SiO 2 + 2Mg \u003d 2MgO + Si.

3SiO 2 + 4Al \u003d Al 2 O 3 + 3Si.

Silicon forms to-you: H 2 SiO 3 – meta-silicon acid; H 2 Si 2 O 5 - two-methacrylic acid.

Finding in nature: quartz mineral - SiO2. Quartz crystals are in the form of a hexagonal prism, colorless and transparent, called rock crystal. Amethyst - rhinestone, stained with impurities in purple; smoky topaz is painted brownish; agate and jasper - crystalline. varieties of quartz. Amorphous silica is less common and exists as the mineral opal. Diatomite, tripolite or diatomaceous earth (diatomaceous earth) are earthy forms of amorphous silicon. formula silicon to - t - n SiO2?m H2O. In nature, nah-Xia is mainly in the form of salts, in free. A few have been identified, for example, HSiO (orthosilicon) and H 2 SiO 3 (silicon or metasilicon).

Obtaining silicic acid:

1) the interaction of silicates alkali. metals with to-tami: Na 2 SiO 3 + 2HCl = H 2 SiO 3 + 2NaCl;

2) flint to-that yavl. thermally unstable: H 2 SiO 3 \u003d H 2 O + SiO 2.

H 2 SiO 3 forms supersaturated p-ry, in the cat. in the result of those polymerization forms colloids. Using stabilizers, stable colloids (sols) can be obtained. They are used in production. Without stabilizers, a gel is formed from the silicon to-you solution, after drying which you can get silica gel (used as an adsorbent).

silicates- silicon salts to-you. silicates are common in nature, Earth's crust consists mostly of silica and silicates (feldspars, mica, clay, talc, etc.). Granite, basalt and other rocks contain silicates. Emerald, topaz, aquamarine - silicate crystals. Only sodium and potassium silicates are soluble, the rest are insoluble. The silicates are complex. chem. compound: Kaolin Al 2 O 3 ; 2SiO 2 ; 2H 2 O or H 4 Al 2 SiO 9 .

Asbestos CaO; 3MgO; 4SiO 2 or CaMgSi 4 O 12 .

Receipt: fusion of silicon oxide with alkalis or carbonates.

Soluble glass- sodium and potassium silicates. Liquid glass- water solutions of silicates of potassium and sodium. Its use for the manufacture of acid-resistant cement and concrete, kerosene-proof plasters, fire-retardant paints. Aluminosilicates- silicates containing aluminum ( feldspar, mica). feldspars In addition to oxides of silicon and aluminum, they consist of oxides of potassium, sodium, and calcium. micas have in their composition, in addition to silicon and aluminum, also hydrogen, sodium or potassium, less often - calcium, magnesium, iron. Granites and gneisses (rocks)- comp. from quartz, feldspar and mica. Horn. rocks and minerals, being on the Earth, interact with water and air, which causes their change and destruction. This process is called weathering.

Application: silicate rocks (granite) how construction material, silicates - as a raw material in the production of cement, glass, ceramics, fillers; mica and asbestos - as electrical and thermal insulation.

Let's talk about how to determine the nature of the oxide. Let's start with the fact that all substances are usually divided into two groups: simple and complex. Elements are divided into metals and non-metals. Complex compounds are divided into four classes: bases, oxides, salts, acids.

Definition

Since the nature of oxides depends on their composition, first let's define this class inorganic substances. Oxides are which consist of two elements. Their peculiarity is that oxygen is always located in the formula as the second (last) element.

The most common option is the interaction with oxygen of simple substances (metals, non-metals). For example, when magnesium reacts with oxygen, a mineral that exhibits basic properties is formed.

Nomenclature

The nature of oxides depends on their composition. There are certain rules by which such substances are called.

If the oxide is formed by metals of the main subgroups, the valence is not indicated. For example, calcium oxide CaO. If in the compound the first metal of a similar subgroup is located, which has a variable valence, then it must be indicated by a Roman numeral. Placed after the connection name in parentheses. For example, there are oxides of iron (2) and (3). When compiling the formulas of oxides, one must remember that the sum of the oxidation states in it must be equal to zero.

Classification

Consider how the nature of oxides depends on the degree of oxidation. Metals having an oxidation state of +1 and +2 form basic oxides with oxygen. A specific feature of such compounds is the basic nature of the oxides. Such compounds enter into chemical interaction with salt-forming oxides of non-metals, forming salts with them. In addition, they react with acids. The product of the interaction depends on the amount in which the starting substances were taken.

Non-metals, as well as metals with oxidation states from +4 to +7, form acidic oxides with oxygen. The nature of oxides suggests interaction with bases (alkalis). The result of the interaction depends on the amount in which the initial alkali was taken. With its deficiency, as a product of interaction, acid salt. For example, in the reaction of carbon monoxide (4) with sodium hydroxide, sodium bicarbonate (acid salt) is formed.

In the case of interaction of an acid oxide with an excess amount of alkali, the reaction product will be an average salt (sodium carbonate). The nature of acid oxides depends on the degree of oxidation.

They are divided into salt-forming oxides (in which the oxidation state of the element is equal to the group number), as well as indifferent oxides that are not capable of forming salts.

Amphoteric oxides

There is also an amphoteric nature of the properties of oxides. Its essence lies in the interaction of these compounds with both acids and alkalis. Which oxides exhibit dual (amphoteric) properties? These include binary compounds of metals with an oxidation state of +3, as well as oxides of beryllium, zinc.

How to get

There are various ways The most common option is the interaction with oxygen simple substances(metals, non-metals). For example, when magnesium reacts with oxygen, a mineral that exhibits basic properties is formed.

In addition, oxides can also be obtained by the interaction complex substances with molecular oxygen. For example, when burning pyrite (iron sulfide 2), two oxides can be obtained at once: sulfur and iron.

Another option for obtaining oxides is the reaction of decomposition of salts of oxygen-containing acids. For example, when calcium carbonate is decomposed, carbon dioxide and calcium oxide can be obtained.

Basic and amphoteric oxides are also formed during the decomposition of insoluble bases. For example, when iron (3) hydroxide is calcined, iron (3) oxide is formed, as well as water vapor.

Conclusion

Oxides are a class of inorganic substances with wide industrial applications. They are used in the construction industry, pharmaceutical industry, medicine.

In addition, amphoteric oxides are often used in organic synthesis as catalysts (accelerators of chemical processes).

Two oxides of carbon are known: CO and CO 2 .

Carbon monoxide (II) CO (carbon monoxide). In the molecule of this oxide, the carbon atom is in an unexcited state. Due to the two p-electrons, it forms two bonds with the oxygen atom. The third bond is formed according to the donor-acceptor mechanism, and oxygen is the donor of an electron pair, which the carbon atom accepts into a free 2p orbital.

Carbon monoxide (II) CO is formed during the combustion of coal with a lack of oxygen. In industry, it is obtained by passing carbon dioxide over hot coal:

CO 2 + C \u003d 2CO

Under laboratory conditions, CO is obtained by the action of concentrated sulfuric acid on formic acid when heated (H 2 SO 4 takes away water):

HCOOH®H 2 O+CO

Carbon monoxide (II) CO is a colorless gas, odorless. Highly

slightly soluble in water. Poisonous. Permissible CO content in

industrial premises is 0.03 mg per 1 liter of air. In quantities dangerous to life, it is contained in the exhaust gases of cars. The toxic effect is

that it irreversibly interacts with blood hemoglobin,

as a result of which the transfer of oxygen from the lungs to the

Chemically, CO is an inert compound (at low temperature). With an increase in temperature to 200 ° C and a pressure of 15 10 5 Pa, carbon monoxide (II) reacts with NaOH, forming the sodium salt of formic acid:

Oxidation to CO 2 occurs at a temperature of 700 ° C: 2CO + O 2 \u003d 2CO 2

When interacting with water vapor, CO 2 and H 2 are formed: CO + H 2 O®CO 2 + H 2

CO is an energetic reducing agent. It restores many metals from their oxides, which is used in metallurgy to obtain metals from ores:

Fe 2 O 3 + 3CO \u003d 2Fe + 3CO 2

In the presence of catalysts (platinum or activated carbon) or under the influence of direct sunlight, carbon monoxide combines with chlorine, forming an extremely poisonous gas - phosgene:

CO + Cl 2 ®COCl 2

Unique is the ability of carbon monoxide (II) at elevated temperatures and pressures to form with some metals unusual (complex) compounds called carbonyls:

Under normal conditions, liquids are carbonyls Ni(CO) 4 , Fe(CO) 5 , Ru(CO) 5 , Os(CO) 5 . All the rest are crystalline substances. Metal carbonyls are diamagnetic, indicating the presence of paired electrons. All of them are highly resistant to various chemical reagents. Relative independence in the interpretation of the behavior of s - and p-electrons allows us to understand the features of the electronic structure of carbonyl complexes. If the metal, connecting with the ligand, exhibits low valence values, then in s-bonds the charge is transferred from the ligand to the metal, and in p-bonds, on the contrary, from the metal to the ligand. As a result, the metal atom goes into a state close to neutral. This is exactly how the CO molecule acts as an acceptor in p - connections.

When heated, metal carbonyls decompose into CO and metal, which is used to obtain high-purity metals.

Carbon monoxide (IV) CO 2 (carbon dioxide) It is formed in nature during combustion and decay of organic substances. Contained in the air (volume fraction 0.03%), as well as in many mineral springs (Narzan, Borjomi). It is released during the respiration of animals and plants.

In the laboratory, it can be obtained by the action of dilute acids on carbonates:

CaCO 3 + 2HCl \u003d CaCl 2 + CO 2 + H 2 O

In industry, it is obtained by burning limestone:

CaCO 3 \u003d CaO + CO 2

Structural formula CO 2 molecules: O=C=O. It has a linear shape. The bond between carbon and oxygen is polar. However, due to the symmetrical arrangement of bonds, the CO 2 molecule itself is non-polar.

Under normal conditions, CO 2 is a colorless gas, 1.5 times heavier than air. Soluble in water (at 0°C 1.7 l of CO 2 in 1 l of H 2 O). It does not support combustion and respiration, but serves as a source of nutrition for green plants. With strong cooling, CO 2 crystallizes in the form of a white snow-like mass, which, in a compressed state, evaporates very slowly, lowering the temperature environment. This explains its use as "dry ice".