Lecture 8

THEME : Group elements IVA.

Carbon

Questions studied at the lecture:

1. IVA groups.
2. Carbon. General characteristics of carbon.
3. Chemical properties of carbon.
4. The most important carbon compounds.

General characteristics of the elements IVA groups

To the elements of the main subgroup IV groups include C, Si, Ge, Sn, P v. Electronic formula of external valence level nS 2 np 2 , that is, they have 4 valence electrons and these are p - elements, therefore they are in the main subgroup IV group.

││││

│ ↓ │ np

In the ground state of an atom, two electrons are paired and two are unpaired. The pre-outer electron shell of carbon has 2 electrons, silicon has 8, and Ge, Sn, P c - 18 electrons each. That's why Ge, Sn, P in united in the subgroup germanium (these are full electronic analogs).

In this subgroup of p - elements, as in other subgroups of p - elements, the properties of the atoms of the elements change periodically:

Table 9

Element	Covalent atom radius, nm	Metallic radius of an atom, nm	Conditional ion radius, nm		Energy ionization E E o → E +, ev.	Relative electronegativity
								E 2+	E 4+
	0,077				11,26
	0,117	0,134		0,034	8,15
	0,122	0,139	0,065	0,044	7,90
	0,140	0,158	0,102	0,067	7,34
P in		0,175	0,126	0,076	7,42

Thus, from top to bottom in the subgroup, the radius of the atom increases, therefore the ionization energy decreases, therefore the ability to donate electrons increases, and the tendency to supplement the outer electron shell to an octet sharply decreases, therefore, the reducing properties and metallic properties increase from C to PB, and the non-metallic properties decrease. ... Carbon and silicon are typical non-metals, y Ge already there are metallic properties and in appearance it looks like a metal, although it is a semiconductor. Tin already has metallic properties, and lead is a typical metal.

Having 4 valence electrons, atoms in their compounds can exhibit oxidation states from minimum (-4) to maximum (+4), and they are characterized by even S.О .: -4, 0, +2, +4; S.O. = -4 is characteristic for C and Si with metals.

The nature of the relationship with other elements.Carbon forms only covalent bonds, silicon also predominantly forms covalent bonds. For tin and lead, especially in S.O. = +2, the ionic nature of the bond is more characteristic (for example, Рв ( NO 3) 2).

Covalence determined by the valence structure of the atom. The carbon atom has 4 valence orbitals and the maximum covalence is 4. For the rest of the elements, covalence can be more than four, since there is a valence d -sublevel (for example, H 2 [SiF 6]).

Hybridization ... The type of hybridization is determined by the type and number of valence orbitals. Carbon only has S - and p-valence orbitals, so it can be Sp (carbyne, CO 2, CS 2), Sp 2 (graphite, benzene, COCl 2), Sp 3 -hybridization (CH 4, diamond, CCl 4 ). For silicon, the most characteristic Sp 3 - hybridization (SiO 2, SiCl 4 ), but it has a valence d -sublevel, therefore there is also Sp 3 d 2 -hybridization, for example, H 2 [SiF 6].

IV the PSE group is the middle of the Mendeleev table. A sharp change in properties from non-metals to metals is clearly seen here. Consider carbon separately, then silicon, then elements of the germanium subgroup.

Carbon. General characteristics of carbon

The carbon content in the earth's crust is low (approximately 0.1% of the mass). Most of it is contained in the composition of sparingly soluble carbonates (CaCO 3, MgCO 3 ), oil, coal, natural gas. CO content 2 in the air is small (0.03%), but its total mass is about 600 million tons. Carbon is part of the tissues of all living organisms (the main constituent of flora and fauna). Carbon is also found in a free state, mainly in the form of graphite and diamond.

In nature, carbon is known in the form of two stable isotopes: 12 C (98.892%) and 13 C (1.108%). Under the action of cosmic rays, a certain amount of β-radioactive isotope is also formed in the atmosphere 14 WITH: . By content 14 C in plant residues is judged on their age. Radioactive isotopes with mass numbers from 10 to 16 have also been obtained.

Unlike F 2, N 2, O 2 simple carbon substances have a polymeric structure. In accordance with the characteristic types of hybridization of valence orbitals, C atoms can combine into polymer formations of three-dimensional modification (diamond, Sp 3 ), two-dimensional or layered modification (graphite, Sp 2 ) and linear polymer (carbyne, Sp).

Chemical properties of carbon

Chemically, carbon is very inert. But when heated, it is able to interact with many metals and non-metals, while exhibiting both oxidizing and reducing properties.

Diamond + 2 F 2 → CF 4 and graphite forms graphite fluoride CF

(and further + F 2 → CF 4 ). One of the methods of separating diamond from graphite is based on a different attitude to fluorine. Carbon does not react with other halogens. With oxygen (O 2 ) carbon with a lack of oxygen forms CO, with an excess of oxygen forms CO 2 .

2C + O 2 → 2CO; C + O 2 → CO 2.

At high temperatures, carbon reacts with metals to form metal carbides:

Ca + 2C = CaC 2.

When heated, it reacts with hydrogen, sulfur, silicon:

t o t o

С + 2 Н 2 = СН 4 С + 2S ↔ CS 2

C + Si = SiC.

Carbon also reacts with complex substances. If steam is passed through the heated coal, a mixture of CO and H is formed 2 - water gas (at a temperature of more than 1200 o C):

C + HOH = CO + H 2.

This mixture is widely used as a gaseous fuel.

At high temperatures, carbon is able to reduce many metals from their oxides, which is widely used in metallurgy.

ZnO + C → Zn + CO

The most important carbon compounds

1. Metal carbides.

Since it is typical for carbon to form homochains, the composition of most carbides does not correspond to the oxidation state of carbon equal to (-4). By the type of chemical bond, covalent, ionic - covalent and metal carbides are distinguished. In most cases, carbides are obtained by strong heating of the corresponding simple substances or their oxides with carbon.

T o t o

V 2 O 5 + 7C → 2VC + 5CO; Ca + 2 C → CaC 2.

In this case, carbides of different composition are obtained.

Salt-like or ionic-covalent carbides are compounds of active and some other metals: Be 2 C, CaC 2, Al 4 C 3, Mn 3 C ... In these compounds, the chemical bond is intermediate between ionic and covalent. When exposed to water or dilute acids, they are hydrolyzed and hydroxides and corresponding hydrocarbons are obtained:

CaC 2 + 2HOH → Ca (OH) 2 + C 2 H 2;

Al 4 C 3 + 12HOH → 4Al (OH) 3 + 3CH 4.

In metal carbides, carbon atoms occupy octahedral voids in the structures of metals (side subgroups IV - VIII groups). These are very hard, refractory and heat-resistant substances, many of them exhibit metallic properties: high electrical conductivity, metallic luster. The composition of such carbides varies widely. So, titanium carbides have the composition TiC 0.6 - 1.0.

Covalent carbides - SiC and B 4 C. They are polymeric. The chemical bond in them approaches purely covalent, since boron and silicon are neighbors of carbon in the PES and are close to it along the radius of the atom and OEO. They are very hard and chemically inert. Methane CH 4 .

1. Carbon halides

Carbon forms many compounds with halogens, the simplest of which have the formula C H al 4 , that is, carbon tetrahalides. In them S.O. carbon is +4, Sp 3 -hybridization of the C atom, therefore the molecules CH al 4 - tetrahedrons. CF 4 - gas, CCl 4 - liquid, CBr 4 and CJ 4 - solids. Only CF 4 obtained directly from F 2 and C, carbon does not react with other halogens. Carbon tetrachloride is obtained by chlorination of carbon disulfide:

CS 2 + 3Cl 2 = CCl 4 + S 2 Cl 2.

All C H al 4 do not dissolve in water, but dissolve in organic solvents.

t o, Kat

C H al 4 (g) + 2HOH (g) = CO 2 + 4HHa l (d) (hydrolysis occurs with strong heating and in the presence of a catalyst). Are of practical importance CF 4, CC l 4.

CF 4 , like other fluorine-containing carbon compounds, for example CF 2 Cl 2 (difluorodichloromethane) is used as freons - working substances of refrigeration machines.

CCl 4 It is used as a non-flammable solvent for organic substances (fats, oils, resins), as well as a liquid for fire extinguishers.

1. Carbon monoxide (P).

Carbon monoxide (P) CO is a colorless, odorless, slightly water-soluble gas. Very toxic (carbon monoxide): blood hemoglobin associated with CO, loses its ability to combine with O 2 and be its carrier.

Carbon monoxide (P) is obtained:

• with incomplete oxidation of carbon 2C + O 2 = 2CO;

• in industry is obtained by the reaction: CO 2 + C = 2CO;
• when passing superheated steam over hot coal:

C + HOH = CO + H 2 t o

• decomposition of carbonyls Fe (CO) 5 → Fe + 5 CO;
• in the laboratory CRM is obtained by acting on formic acid with dehydrating substances ( H 2 SO 4, P 2 O 5):

НСООН → СО + НОН.

However, CO is not formic acid anhydride, since in CO carbon is trivalent, and in HCOOH it is tetravalent. Thus, CO is a non-salt-forming oxide.

The solubility of CO in water is low and no chemical reaction occurs. In the CO molecule, as in the molecule N 2 - triple bond. According to the method of valence bonds, 2 bonds are formed due to the pairing of two unpaired p - electrons C and O (of each atom), and the third - according to the donor-acceptor mechanism due to the free 2p - orbital of the C atom and 2p - the electron pair of the oxygen atom: C ≡ O . The CO triple bond is very strong and its energy is very high (1066 kJ / mol) - more than in N 2 ... Carbon monoxide (P) is characterized by the following three types of reactions:

1. oxidation reactions... CO is a strong reducing agent, however, due to the strong triple bond in the molecule, redox reactions involving CO proceed quickly only at high temperatures. Reduction of oxides with CO during heating is of great importance in metallurgy.

Fe 2 O 3 + 3CO = 3CO 2 + 2Fe.

CO can be oxidized by oxygen: t o

2CO + O 2 = 2CO 2.

1. another characteristic chemical property of CO is the tendency toaddition reactions, which is due to the valence unsaturation of carbon in CO (in these reactions, carbon goes into a tetravalent state, which is more characteristic for it than the trivalent state of carbon in CO).

So, CO reacts with chlorine to form phosgene COC l 2:

CO + Cl 2 = COCl 2 (CO is also a reducing agent in this reaction). The reaction is accelerated by the action of light and a catalyst. Phosgene is a brown gas, very poisonous - a strong poisonous substance. Slowly hydrolyzed COCl 2 + 2 HOH → 2 HCl + H 2 CO 3.

Phosgene is used in the synthesis of various substances and was used in the First World War as a chemical warfare agent.

When heated, CO reacts with sulfur to form carbon sulfoxide COS:

CO + S = COS (gas).

When heated under pressure, CO forms methanol upon interaction with hydrogen

t o, p

CO + 2H 2 ↔ CH 3 OH.

Synthesis of methanol from CO and H 2 - one of the most important chemical industries.

1. Unlike most other carbon compounds, the CO molecule has a lone electron pair at the C atom. Therefore, the CO molecule can act ligand in various complexes. The products of CO addition to metal atoms, which are called carbonyls, are especially numerous. About 1000 carbonyls are known, including carbonyls containing other ligands besides CO. Carbonyls (complexes) get:

T, p t, p

Fe + 5CO → Ni + 4CO →.

There are gaseous, liquid and solid carbonyls, in which the metal has an oxidation state of 0. When heated, carbonyls decompose and powder metals of a very high purity are obtained:

t o

Ni (CO) 4 → Ni + 4CO.

Carbonyls are used in syntheses and for the production of highly pure metals. All carbonyls, like CO, are extremely toxic.

1. Carbon monoxide (IV).

CO 2 molecule has a linear structure (O = C = O), Sp - hybridization of the carbon atom. Two bonds of the σ - type arise due to the overlap of two Sp - hybrid orbitals of the C atom and two 2p NS - orbitals of two oxygen atoms, on which there are unpaired electrons. Two other bonds of the π - type arise when 2p overlaps y - and 2p z - orbitals of the C atom (non-hybrid) with the corresponding 2p y - and 2p z - orbitals of oxygen atoms.

Obtaining CO 2:

- in industryobtained by calcining limestone

CaCO 3 → CaO + CO 2;

In the laboratory are obtained in the Kipp apparatus by the reaction

CaCO 3 + 2HCl → CaCl 2 + CO 2 + HOH.

Physical properties of CO 2 : it is a gas, heavier than air, solubility in water is low (at 0 O C in 1 liter of water dissolves 1.7 liters of CO 2, and at 15 о C dissolves 1 l of CO 2 ), while some of the dissolved CO 2 interacts with water to form carbonic acid:

HOH + CO 2 ↔ H 2 CO 3 ... Equilibrium is shifted to the left (←), so most of the dissolved CO 2 as CO 2, not acid.

V chemically CO 2 exhibits: a) the properties of an acidic oxide and when interacting with alkali solutions, carbonates are formed, and with an excess of CO 2 - hydrocarbons:

2NaOH + CO 2 → Na 2 CO 3 + H 2 O NaOH + CO 2 → NaHCO 3.

b) oxidizing properties, but oxidizing properties CO 2 very weak, since S.O. = +4 is the most characteristic oxidation state of carbon. In this case, CO 2 reduced to CO or C:

C + CO 2 ↔ 2CO.

C О 2 used in the production of soda, for extinguishing fires, preparation of mineral water, as an inert medium in syntheses.

1. Carbonic acid and its salts

Carbonic acid is known only in dilute aqueous solutions. Formed by the interaction of CO 2 with water. In an aqueous solution, most of the dissolved CO 2 in a hydrated state and only a small part in the form of H 2 CO 3, HCO 3 -, CO 3 2- , that is, equilibrium is established in an aqueous solution:

CO 2 + HOH ↔ H 2 CO 3 ↔ H + + HCO 3 - ↔ 2H + + CO 3 2-.

Equilibrium is strongly shifted to the left (←) and its position depends on temperature, environment, etc.

Carbonic acid is considered a weak acid (K 1 = 4,2 ∙ 10 -7 ). This is the apparent ionization constant K and he. , it is related to the total amount of CO dissolved in water 2 , and not to the true concentration of carbonic acid, which is not known exactly. But since the molecules H 2 CO 3 is small in solution, then the true K and he. carbonic acid is much more than indicated above. So, apparently, the true value of K 1 ≈ 10 -4 , that is, carbonic acid is an acid of medium strength.

Salts (carbonates) are usually slightly soluble in water. Carbonates K dissolve well+, Na +, R в +, Cs +, Tl +1, NH 4 + ... Hydrocarbonates, unlike carbonates, are mostly soluble in water.

Salt hydrolysis: Na 2 CO 3 + HOH ↔ NaHCO 3 + NaOH (pH> 7).

When heated, carbonates decompose to form metal oxide and CO 2 The more pronounced the metallic properties of the cation-forming element, the more stable the carbonate. So, Na 2 CO 3 melts without decomposition; CaCO 3 decomposes at 825о С, а Ag 2 CO 3 decomposes at 100 O C. Bicarbonates decompose on low heating:

2NaHCO 3 → Na 2 CO 3 + CO 2 + H 2 O.

1. Urea and carbon disulfide.

Urea or urea is obtained by the action of CO 2 for aqueous solution H 3 N at 130 о С and 1 ∙ 10 7 Pa.

CO 2 + 2H 3 N = CO (NH 2) 2 + H 2 O.

Urea is a white crystalline substance. It is used as a nitrogen fertilizer, for feeding livestock, for obtaining plastics, pharmaceuticals (veronal, luminal).

Carbon disulfide (carbon disulfide) - CS 2 under normal conditions - a volatile colorless liquid, poisonous. Clean CS 2 has a faint pleasant odor, but upon contact with air - a disgusting smell of its oxidation products. Carbon disulfide does not dissolve in water; when heated (150 O C) hydrolyzed to CO 2 and H 2 S:

CS 2 + 2HOH = CO 2 + 2H 2 S.

Carbon disulfide is easily oxidized and flammable in air with slight heating: CS 2 + 3 O 2 = CO 2 + 2 SO 2.

Get carbon disulfide by the interaction of sulfur vapors with hot coal. Carbon disulfide is used as a good solvent for organic substances, phosphorus, sulfur, iodine. The bulk CS 2 it is used to obtain viscose silk and as a means for combating agricultural pests.

1. Hydrocyanic, thiocyanic and cyanic acids.

Hydrocyanic acid HCN (or hydrocyanic acid) has a linear structure, consists of molecules of 2 types, which are in tautomeric equilibrium, which is shifted to the left at room temperature:

H - C ≡ N ↔ H - N ≡ C

cyanide isocyanide

hydrogen hydrogen

HCN Is a volatile liquid with the smell of almonds, one of the strongest poisons, mixes up with water in any ratio. In aqueous solution HCN - weak acid (K = 7.9 ∙ 10-10 ), that is, much weaker than carbonic acid.

In industry HCN obtained by catalytic reaction:

t o, kat

CO + NH 3 → HCN + HOH.

Salts (cyanides) are obtained by reduction of carbonates with carbon by heating:

Na 2 CO 3 + C + 2NH 3 = 2NaCN + 3H 2 O.

Hydrogen cyanide is used in organic synthesis, and NaCN and KCN - in gold mining, to obtain complex cyanides, etc.

Cyanides are basic ( NaCN) and acidic (JCN ). Hydrolysis of basic cyanide:

NaCN + HOH ↔ NaOH + HCN (pH> 7).

When acid cyanide is hydrolyzed, two acids are formed:

JCN + HOH = HJO + HCN.

Cyanide d -elements do not dissolve in water, but due to complexation they easily dissolve in the presence of basic cyanides:

4KCN + Mn (CN) 2 = K 4.

Complex cyanides are very stable.

Hydrogen thiocyanate HSCN or HNCS has a linear structure and consists of two types of molecules: H - S - C ≡ NorH – N = C = S... In crystalline thiocyanatesNaNCS, Ba(NCS) 2 the metal ion is near the nitrogen atom; vAgSCN, Hg(SCN) 2 the metal ion is near the sulfur atom.

Rhodanides or thiocyanates are obtained by the action of sulfur on alkali metal cyanides (boiling solutions with sulfur):

to

KCN + S = KNCS.

Anhydrous hydrogen thiocyanate is obtained by heating lead (or mercury) thiocyanate in a currentH2 S:

to

PB(SCN)2 + H2 S →PBS ↓ + 2HNCS.

HNCS- colorless oily liquid with a pungent odor, easily decomposes. It dissolves well in water, in aqueous solutionHNCSforms strong thiocyanate acid (K = 0.14). Rodanides are mainly used for dyeing fabrics, andNH4 CNSused as a reagent for ionsFe3+ .

Also known are tautomeric cyanic (HOCN) and isocyanic (HNCO) acids:

.

This equilibrium at room temperature is shifted to the left.

Salts - cyanates and isocyanates are obtained by oxidation of cyanides: 2KCN + O2 = 2 KOCN... Cyanic acid in aqueous solution is a medium strength acid.

IVA group of chemical elements of the periodic table D.I. Mendeleev includes non-metals (carbon and silicon), as well as metals (germanium, tin, lead). The atoms of these elements contain four electrons at the external energy level (ns 2 np 2), two of which are not paired. Therefore, the atoms of these elements in compounds can exhibit valence II. Atoms of group IVA elements can pass into an excited state and increase the number of unpaired electrons to 4 and, accordingly, in compounds exhibit a higher valence equal to the number of group IV. Carbon in compounds exhibits oxidation states from –4 to +4, for the rest, oxidation states are stabilized: –4, 0, +2, +4.

In a carbon atom, unlike all other elements, the number of valence electrons is equal to the number of valence orbitals. This is one of the main reasons for the stability of the C – C bond and the exceptional tendency of carbon to form homochains, as well as the existence of a large number of carbon compounds.

In the change in the properties of atoms and compounds in the series C – Si – Ge – Sn – Pb, secondary peridicity is manifested (Table 5).

Table 5 - Characteristics of atoms of group IV elements

	6 C	1 4 Si	3 2 Ge	50 Sn	82 Pb
Atomic mass	12,01115	28,086	72,59	118,69	207,19
Valence electrons	2s 2 2p 2	3s 2 3p 2	4s 2 4p 2	5s 2 5p 2	6s 2 6p 2
Atom covalent radius, Ǻ	0,077	0,117	0,122	0,140	–
Metallic radius of an atom, Ǻ	–	0,134	0,139	0,158	0,175
Conditional ion radius, E 2+, nm	–	–	0,065	0,102	0,126
Conditional ion radius E 4+, nm	–	0,034	0,044	0,067	0,076
Ionization energy E 0 - E +, ev	11,26	8,15	7,90	7,34	7,42
Content in the earth's crust, at. %	0,15	20,0	2∙10 –4	7∙10 – 4	1,6∙10 – 4

Secondary periodicity (non-monotonic change in the properties of elements in groups) is due to the nature of the penetration of external electrons to the nucleus. Thus, the nonmonotonicity of the change in atomic radii on going from silicon to germanium and from tin to lead is due to the penetration of s electrons under the 3d 10 electron shield in germanium and the double shield of 4f 14 and 5d 10 electrons in lead, respectively. Since the penetrating power decreases in the series s> p> d, the internal periodicity in the change in properties is most clearly manifested in the properties of elements determined by s-electrons. Therefore, it is most typical for compounds of the elements of the A-groups of the periodic table, corresponding to the highest oxidation state of the elements.

Carbon differs significantly from other p-elements of the group by its high value of ionization energy.

Carbon and silicon have polymorphic modifications with different crystal lattice structures. Germanium belongs to metals, silvery-white in color with a yellowish tint, but has a diamond-like atomic crystal lattice with strong covalent bonds. Tin has two polymorphic modifications: a metal modification with a metal crystal lattice and a metal bond; non-metallic modification with an atomic crystal lattice, which is stable at temperatures below 13.8 C. Lead is a dark gray metal with a metallic face-centered cubic crystal lattice. A change in the structure of simple substances in the germanium – tin – lead series corresponds to a change in their physical properties. So germanium and non-metallic tin are semiconductors, metallic tin and lead are conductors. The change in the type of chemical bond from predominantly covalent to metallic is accompanied by a decrease in the hardness of simple substances. So, germanium is quite hard, while lead is easily rolled into thin sheets.

Compounds of elements with hydrogen have the formula EN 4: CH 4 - methane, SiH 4 - silane, GeH 4 - german, SnH 4 - stannane, PbH 4 - plumbane. They are insoluble in water. From top to bottom, in the series of hydrogen compounds, their stability decreases (plumban is so unstable that its existence can only be judged by indirect signs).

Compounds of elements with oxygen have the general formulas: EO and EO 2. The oxides CO and SiO are non-salt-forming; GeO, SnO, PbO - amphoteric oxides; CO 2, SiO 2 GeO 2 - acidic, SnO 2, PbO 2 - amphoteric. With an increase in the oxidation state, the acidic properties of the oxides increase, the basic properties weaken. The properties of the corresponding hydroxides change in a similar way.

| | | | | | | |

16.1. General characteristics of the elements of IIIA, IVA and VA groups

B Boron 0,776	C Carbon 0,620	N Nitrogen 0,521
Al Aluminum 1,312	Si Silicon 1,068	P Phosphorus 0,919
Ga Gallium 1,254	Ge Germanium 1,090	As Arsenic 1,001
In Indium 1,382	Sn Tin 1,240	Sb Antimony 1,193
Tl Thallium 1,319	Pb Lead 1,215	Bi Bismuth 1,295

The composition of these three groups of the natural system of elements is shown in Figure 16.1. The values ​​of the orbital radii of atoms (in angstroms) are also given here. It is in these groups that the boundary between the elements forming metals (the orbital radius is greater than 1.1 angstroms) and the elements forming non-metals (the orbital radius is less than 1.1 angstroms) is most clearly traced. In the figure, this border is shown with a double line. It should not be forgotten that this border is still conditional: aluminum, gallium, tin, lead and antimony are certainly amphoteric metals, but boron, germanium, and arsenic show some signs of amphotericity.
Of the atoms of elements of these three groups in the earth's crust, the following are most often found: Si (w = 25.8%), Al (w = 7.57%), P (w = 0.090%), C (w = 0.087%) and N (w = 0.030%). It is with them that you will get acquainted in this chapter.
The general valence electronic formulas of the atoms of the IIIA group elements are ns 2 np 1, IVA groups - ns 2 np 2, VA groups - ns 2 np 3. The highest oxidation states are equal to the group number. Intermediate 2 less.
All simple substances formed by the atoms of these elements (with the exception of nitrogen) are solid. Many elements are characterized by allotropy (B, C, Sn, P, As). There are only three stable molecular substances: nitrogen N 2, white phosphorus P 4 and yellow arsenic As 4.

Non-metallic elements of these three groups tend to form molecular hydrogen compounds with covalent bonds. Moreover, carbon has so many of them that hydrocarbons and their derivatives are studied by a separate science - organic chemistry. Boron is the second most hydrogen compound among these elements. Borohydrides (boranes) are very numerous and complex in structure, therefore, the chemistry of borohydrides has also emerged as a separate section of chemistry. Silicon forms a total of 8 hydrogen compounds (silanes), nitrogen and phosphorus - two each, the rest - one hydrogen compound each. Molecular formulas of the simplest hydrogen compounds and their names:

The composition of the higher oxides corresponds to the highest oxidation state equal to the group number. The type of higher oxides in each of the groups with an increase in the serial number gradually changes from acidic to amphoteric or basic.

The acid-base character of hydroxides is very diverse. So, HNO 3 is a strong acid, and TlOH is an alkali.

1. Make abbreviated electronic formulas and energy diagrams of atoms of elements of IIIA, IVA and VA groups. Indicate the outer and valence electrons.

The nitrogen atom has three unpaired electrons, therefore, according to the exchange mechanism, it can form three covalent bonds. It can form another covalent bond by the donor-acceptor mechanism, while the nitrogen atom acquires a positive formal charge +1 e... Thus, the maximum nitrogen is pentavalent, but its maximum covalent is equal to 4. (This explains the often met statement that nitrogen cannot be pentavalent)
Almost all of the terrestrial nitrogen is in the atmosphere of our planet. A significantly smaller part of nitrogen is present in the lithosphere in the form of nitrates. Nitrogen is a part of organic compounds contained in all organisms and in their decomposition products.
Nitrogen forms the only simple molecular substance N 2 with a triple diatomic bond in the molecule (Fig. 16.2). The energy of this bond is equal to 945 kJ / mol, which exceeds the values ​​of other binding energies (see table 21). This explains the inertness of nitrogen at ordinary temperatures. According to its physical characteristics, nitrogen is a colorless, odorless gas, familiar to us from birth (the earth's atmosphere is three quarters of nitrogen). Nitrogen is slightly soluble in water.

Nitrogen forms two hydrogen compounds: ammonia NH 3 and hydrazine N 2 H 6:

Ammonia is a colorless gas with a pungent, suffocating odor. Careless inhalation of concentrated ammonia fumes can lead to cramping and suffocation. Ammonia is very soluble in water, which is explained by the formation of four hydrogen bonds with water molecules by each ammonia molecule.

The ammonia molecule is a base particle (see Appendix 14). Taking a proton, it turns into an ammonium ion. The reaction can take place both in an aqueous solution and in a gas phase:

NH 3 + H 2 O NH 4 + OH (in solution);
NH 3 + H 3 O B = NH 4 + H 2 O (in solution);
NH 3g + HCl g = NH 4 Cl cr (in the gas phase).

Aqueous solutions of ammonia are alkaline enough to precipitate insoluble hydroxides, but not alkaline enough for amphoteric hydroxides to dissolve in them to form hydroxo complexes. Therefore, an ammonia solution is convenient to use for the preparation of amphoteric hydroxides p-elements: Al (OH) 3, Be (OH) 2, Pb (OH) 2, etc., for example:

Pb 2 + 2NH 3 + 2H 2 O = Pb (OH) 2 + 2NH 4.

When ignited in air, ammonia burns out, forming nitrogen and water; when interacting with oxygen in the presence of a catalyst (Pt), it is reversibly oxidized to nitrogen monoxide:

4NH 3 + 3O 2 = 2N 2 + 6H 2 O (without catalyst),
4NH 3 + 5O 2 4NO + 6H 2 O (with catalyst).

When heated, ammonia can reduce oxides of not very active metals, for example, copper:

3CuO + 2NH 3 = 3Cu + N 2 + 3H 2 O

Ammonium salts in their properties (except for thermal stability) are similar to alkali metal salts. like the latter, almost all of them are soluble in water, but since the ammonium ion is a weak acid, they are cationically hydrolyzed. When heated, ammonium salts decompose:

NH 4 Cl = NH 3 + HCl;
(NH 4) 2 SO 4 = NH 4 HSO 4 + NH 3;
(NH 4) 2 CO 3 = 2NH 3 + CO 2 + H 2 O;
NH 4 HS = NH 3 + H 2 S;
NH 4 NO 3 = N 2 O + 2H 2 O;
NH 4 NO 2 = N 2 + 2H 2 O;
(NH 4) 2 HPO 4 = NH 3 + (NH 4) H 2 PO 4;
(NH 4) H 2 PO 4 = NH 4 PO 3 + H 2 O.

Nitrogen in various oxidation states forms five oxides: N 2 O, NO, N 2 O 3, NO 2 and N 2 O 5.
The most stable of these is nitrogen dioxide. It is a poisonous brown gas with an unpleasant odor. Reacts with water:

2NO 2 + H 2 O = HNO 2 + HNO 3.

With an alkali solution, the reaction proceeds with the formation of nitrate and nitrite.
N 2 O and NO are non-salt-forming oxides.
N 2 O 3 and N 2 O 5 are acidic oxides. Reacting with water, they respectively form solutions of nitrous and nitric acids.

Nitrogen oxoacid in oxidation state + III - nitrous acid HNO 2. It is a weak acid whose molecules exist only in aqueous solution. Its salts are nitrites. Nitrogen in nitrous acid and nitrites is readily oxidized to the + V oxidation state.

Unlike nitric acid, nitric acid HNO 3 is a strong acid. The structure of its molecule can be expressed in two ways:

Nitric acid mixes with water in all respects, reacting completely with it in dilute solutions:

HNO 3 + H 2 O = H 3 O + NO 3

Nitric acid and its solutions are strong oxidizing agents. When nitric acid is diluted, its oxidative activity decreases. In solutions of nitric acid of any concentration, the oxidizing atoms are primarily nitrogen atoms, not hydrogen. Therefore, when various substances are oxidized with nitric acid, hydrogen, if it is released, is only as a by-product. Depending on the concentration of the acid and the reducing activity of another reagent, the reaction products can be NO 2, NO, N 2 O, N 2, and even NH 4. Most often, a mixture of gases is formed, but in the case of concentrated nitric acid, only nitrogen dioxide is released:

Cu + 4HNO 3 = Cu (NO 3) 2 + 2NO 2 + 2H 2 O
3FeS + 30HNO 3 = Fe 2 (SO 4) 3 + Fe (NO 3) 3 + 27NO 2 + 15H 2 O

In the case of dilute nitric acid, nitrogen monoxide is most often released:

Fe + 4HNO 3 = Fe (NO 3) 3 + NO + 2H 2 O
3H 2 S + 2HNO 3 = 2NO + 4H 2 O + 3S

In the case of very dilute nitric acid reacting with a strong reducing agent (Mg, Al, Zn), ammonium ions are formed:

4Mg + 10HNO 3 = 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O

Those metals that are passivated with concentrated sulfuric acid are also passivated with concentrated nitric acid.
Nitric acid salts - nitrates - thermally unstable compounds. When heated, they decompose:
2KNO 3 = 2KNO 2 + O 2;
2Zn (NO 3) 2 = 2ZnO + 4NO 2 + O 2;
2AgNO 3 = 2Ag + 2NO 2 + O 2.

1. Make the equations of the reactions given in the text of the paragraph descriptively.
2. Make up the equations of reactions characterizing the chemical properties of a) ammonia, b) nitric acid, c) zinc nitrate.
Chemical properties of ammonia and nitric acid.

16.3. Phosphorus

Unlike the nitrogen atom, atom phosphorus can form five covalent bonds by the exchange mechanism. The traditional explanation for this boils down to the possibility of arousal of one of the 3 s-electrons and its transition to 3 d-sublevel.
The element phosphorus forms quite a lot allotropic modifications... Of these, three modifications are the most stable: white phosphorus, red phosphorus, and black phosphorus. White phosphorus is a waxy poisonous substance prone to spontaneous combustion in air, consisting of P 4 molecules. Red phosphorus is a dark red, non-molecular, less active substance with a rather complex structure. Usually, red phosphorus always contains an admixture of white, therefore, both white and red phosphorus are always stored under a layer of water. Black phosphorus is also a non-molecular substance with a complex framework structure.
Molecules of white phosphorus are tetrahedral, the phosphorus atom in them is trivalent. Ball-and-stick model and structural formula of the white phosphorus molecule:

The structure of red phosphorus can be expressed by the structural formula:

Phosphorus is obtained from calcium phosphate by heating with sand and coke:

Ca 3 (PO 4) 2 + 3SiO 2 + 5C = 3CaSiO 3 + 2P + 5CO.

For phosphorus, compounds with the + V oxidation state are most characteristic. When combined with an excess of chlorine, phosphorus forms pentachloride. When any allotropic modification of phosphorus burns in excess of oxygen, oxide phosphorus (V):

4P + 5O 2 = 2P 2 O 5.

There are two modifications of phosphorus (V) oxide: non-molecular (with the simplest formula P 2 O 5) and molecular (with the molecular formula P 4 O 10). Usually phosphorus oxide is a mixture of these substances.

This highly hygroscopic acidic oxide reacts with water to form metaphosphoric, diphosphoric and orthophosphoric acids in succession:

P 2 O 5 + H 2 O = 2HPO 3, 2HPO 3 + H 2 O = H 4 P 2 O 7, H 4 P 2 O 7 + H 2 O = 2H 3 PO 4.

Orthophosphoric acid(usually referred to simply as phosphoric) is a weak tribasic acid (see Appendix 13). It is a colorless crystalline substance, very soluble in water. When reacting with strong bases, depending on the ratio of reagents, forms three rows salts(orthophosphates, hydrogen phosphates and dihydrogen orthophosphates - usually in their names the prefix "ortho" is omitted):

H 3 PO 4 + OH = H 2 PO 4 + H 2 O,
H 3 PO 4 + 2OH = HPO 4 2 + 2H 2 O,
H 3 PO 4 + 3OH = PO 4 3 + 3H 2 O.

Most of the average phosphates (except for the salts of alkaline elements other than lithium) are insoluble in water. There are significantly more soluble acid phosphates.
Phosphoric acid is obtained from natural calcium phosphate by treating it with excess sulfuric acid. With a different ratio of calcium phosphate and sulfuric acid, a mixture of dihydrogen phosphate and calcium sulfate is formed, which is used in agriculture as a mineral fertilizer called "simple superphosphate":
Ca 3 (PO 4) 2 + 3H 2 SO 4 = 2H 3 PO 4 + 3CaSO 4;
Ca 3 (PO 4) 2 + 2H 2 SO 4 = Ca (H 2 PO 4) 2 + 2CaSO 4.

The more valuable "double superphosphate" is obtained by the reaction

Ca 3 (PO 4) 2 + 4H 3 PO 4 = 3Ca (H 2 PO 4) 3.

The main substance of this mineral fertilizer is calcium dihydrogen phosphate.

1. Make the molecular equations of reactions for which the ionic equations are given in the text of the paragraph.
2. Make the equations of the reactions given in the text of the paragraph descriptively.
3. Make the equations of reactions characterizing the chemical properties of a) phosphorus, b) phosphorus (V) oxide, c) orthophosphoric acid, d) sodium dihydrogen phosphate.
Chemical properties of phosphoric acid.

16.4. Carbon

Carbon is the main constituent of all organisms. In nature, there are both simple substances formed by carbon (diamond, graphite) and compounds (carbon dioxide, various carbonates, methane and other hydrocarbons in the composition of natural gas and oil). The mass fraction of carbon in coal reaches 97%.
Atom carbon in the ground state can form two covalent bonds by the exchange mechanism, but under normal conditions such compounds are not formed. A carbon atom, passing into an excited state, uses all four valence electrons.
Carbon forms quite a lot allotropic modifications(see fig.16.2). These are diamond, graphite, carbyne, and various fullerenes.

Diamond is a very hard, colorless, transparent crystalline substance. Diamond crystals are composed of carbon atoms in sp 3 -hybridized state, forming a space frame.
Graphite is a rather soft crystalline substance of gray-black color. Graphite crystals consist of flat layers in which carbon atoms are located in sp 2 -hybrid state and form a grid with hexagonal cells.
Carbyne is a colorless fibrous substance, consisting of linear molecules in which carbon atoms are located in sp-hybrid state (= С = С = С = С = or –С С – С С–).
Fullerenes are molecular allotropic modifications of carbon with molecules C 60, C 80, etc. The molecules of these substances are hollow reticulated spheres.
All carbon modifications exhibit reducing properties to a greater extent than oxidizing ones, for example, coke (a product of coal processing; contains up to 98% carbon) is used to reduce iron from oxide ores and a number of other metals from their oxides:

Fe 2 O 3 + 3C = 2Fe + 3CO (at high temperature).

Most of the carbon compounds are studied in organic chemistry, which you will learn in 10th and 11th grades.
In inorganic substances, the oxidation state of carbon is + II and + IV. With these oxidation states of carbon, there are two oxide.
Carbon monoxide (II) is a colorless poisonous gas, odorless. The trivial name is carbon monoxide. Formed by incomplete combustion of carbonaceous fuel. For the electronic structure of its molecule, see page 121. According to the chemical properties of CO, a non-salt-forming oxide, when heated, exhibits reducing properties (it reduces many oxides of not very active metals to metal).
Carbon monoxide (IV) is a colorless, odorless gas. The trivial name is carbon dioxide. Acidic oxide. It is slightly soluble in water (physically), partially reacts with it, forming coal acid H 2 CO 3 (molecules of this substance exist only in very dilute aqueous solutions).
Carbonic acid - very weak acid (see Appendix 13), dibasic, forms two rows salts(carbonates and hydrocarbons). Most carbonates are insoluble in water. Of bicarbonates, only alkali metal and ammonium bicarbonates exist as individual substances. Both the carbonate ion and the bicarbonate ion are base particles; therefore, both carbonates and bicarbonates in aqueous solutions undergo hydrolysis at the anion.
Of the carbonates, the most important are sodium carbonate Na 2 CO 3 (soda, soda ash, washing soda), sodium bicarbonate NaHCO 3 (baking soda, baking soda), potassium carbonate K 2 CO 3 (potash) and calcium carbonate CaCO 3 (chalk, marble, limestone).
Qualitative reaction for the presence of carbon dioxide in the gas mixture: the formation of a calcium carbonate precipitate when the test gas is passed through lime water (saturated solution of calcium hydroxide) and the subsequent dissolution of the precipitate with further gas passing. Proceeding reactions: The silicon element forms one simple substance with the same name. It is a non-molecular substance with a diamond structure, to which silicon is only slightly inferior in hardness. Over the past half century, silicon has become an absolutely essential material for our civilization, since its single crystals are used in almost all electronic equipment.
Silicon is a fairly inert substance. at room temperature, it practically does not react with anything except fluorine and hydrogen fluoride:
Si + 2F 2 = SiF 4;
Si + 4HF = SiF 4 + 2H 2.
When heated in the form of a finely divided powder, it burns in oxygen, forming dioxide (SiO 2). When fusion with alkali or when boiled with concentrated solutions of alkalis forms silicates:

Si + 4NaOH = Na 4 SiO 4 + 2H 2;
Si + 2NaOH + H 2 O = Na 2 SiO 3 + 2H 2.

Silicon monoxide SiO - non-salt-forming oxide; easily oxidized to dioxide.
Silicon dioxide SiO 2 is a non-molecular substance of the frame structure. Does not react with water. acidic oxide - when fusion with alkalis forms silicates, for example:
SiO 2 + 2NaOH = Na 2 SiO 3 + H 2 O. Aluminum is the next most abundant element in the Earth's lithosphere after silicon. On its own and together with silicon, it forms many minerals: feldspars, micas, corundum Al 2 O 3 and its precious varieties (colorless leucosapphire containing chromium impurities, ruby, containing titanium impurities sapphire).
The simple substance aluminum is a silvery-white shiny light metal. Pure aluminum is very soft, it can be rolled into thin foil, wire drawn out of it. Aluminum has good electrical conductivity. It is weatherproof. Aluminum alloys are hard enough but work well. Aluminum is not poisonous. All this makes it possible to use aluminum in a wide variety of industries: in the aviation, electrical, food, and construction industries. Aluminum is widely used in everyday life. Aluminum is obtained by electrolysis of the melt of its compounds.
The chemical inertness of aluminum is caused by the presence of a dense oxide film on its surface, which prevents the metal from contacting the reagent. When this film is removed chemically or mechanically, aluminum becomes highly reactive. So, devoid of an oxide film, aluminum ignites spontaneously and burns in air without additional heating.
The reducing properties of aluminum are especially pronounced when heated. Under these conditions, it reduces many metals from oxides: not only iron, titanium, zirconium, but even calcium and barium.
Aluminum oxide Al 2 O 3 (trivial names are alumina, corundum) is a non-molecular substance, the bond in which is poorly described as both ionic and covalent. As always in these cases, it is amphoteric oxide. It is obtained by calcining aluminum hydroxide, which also has amphoteric properties.
The hydrated aluminum ion is a cationic acid, so the soluble aluminum salts are highly hydrolyzed.
Of the aluminum salts, the most commonly used potassium alum KAl (SO 4) 2 · 12H 2 O - dodecahydrate of potassium-aluminum sulfate. It is a non-hygroscopic, excellent crystallizing substance. Its solution behaves like a mixture of solutions of two different sulfates: potassium sulfate and aluminum sulfate. The structure of alum can be expressed by the formula: (SO 4) 2.

1. Make the equations of the reactions given in the text of the paragraph descriptively.
2. Make up the equations of reactions characterizing the chemical properties of a) aluminum, b) aluminum hydroxide, i) potassium alum ..
Chemical properties of aluminum salts

The IVA group contains the most important elements, without which there would be neither us, nor the Earth on which we live. This carbon is the basis of all organic life, and silicon is the "monarch" of the mineral kingdom.

If carbon and silicon are typical non-metals, and tin and lead are metals, then germanium occupies an intermediate position. Some textbooks classify it as a non-metal, while others refer to it as a metal. It is silvery white in color and looks like a metal, but has a diamond-like crystal lattice and is a semiconductor, like silicon.

From carbon to lead (with a decrease in non-metallic properties):

w decreases the stability of the negative oxidation state (-4)

w decreases the stability of the highest positive oxidation state (+4)

w increased stability of low positive oxidation state (+2)

Carbon is the main constituent of all organisms. In nature, there are both simple substances formed by carbon (diamond, graphite) and compounds (carbon dioxide, various carbonates, methane and other hydrocarbons in the composition of natural gas and oil). The mass fraction of carbon in coal reaches 97%.
A carbon atom in the ground state can form two covalent bonds by the exchange mechanism, but such compounds are not formed under normal conditions. A carbon atom, passing into an excited state, uses all four valence electrons.
Carbon forms quite a few allotropic modifications (see Figure 16.2). These are diamond, graphite, carbyne, and various fullerenes.

In inorganic substances, the oxidation state of carbon is + II and + IV. With these oxidation states of carbon, there are two oxides.
Carbon monoxide (II) is a colorless poisonous gas, odorless. The trivial name is carbon monoxide. Formed by incomplete combustion of carbonaceous fuel. For the electronic structure of its molecule, see page 121. According to the chemical properties of CO, a non-salt-forming oxide, when heated, exhibits reducing properties (it reduces many oxides of not very active metals to metal).
Carbon monoxide (IV) is a colorless, odorless gas. The trivial name is carbon dioxide. Acidic oxide. It is slightly soluble in water (physically), partially reacts with it, forming carbonic acid H2CO3 (molecules of this substance exist only in very dilute aqueous solutions).
Carbonic acid is a very weak, dibasic acid, forms two series of salts (carbonates and hydrocarbonates). Most carbonates are insoluble in water. Of bicarbonates, only alkali metal and ammonium bicarbonates exist as individual substances. Both the carbonate ion and the bicarbonate ion are base particles; therefore, both carbonates and bicarbonates in aqueous solutions undergo hydrolysis at the anion.
Of the carbonates, the most important are sodium carbonate Na2CO3 (soda, soda ash, washing soda), sodium bicarbonate NaHCO3 (baking soda, baking soda), potassium carbonate K2CO3 (potash) and calcium carbonate CaCO3 (chalk, marble, limestone).
Qualitative reaction to the presence of carbon dioxide in the gas mixture: the formation of a precipitate of calcium carbonate when the test gas is passed through lime water (saturated solution of calcium hydroxide) and the subsequent dissolution of the precipitate with further gas passing. Proceeding reactions:

Ca2 + 2OH + CO2 = CaCO3 + H2O;
CaCO3 + CO2 + H2O = Ca2 + 2HCO3.

In pharmacology and medicine, various carbon compounds are widely used - derivatives of carbonic acid and carboxylic acids, various heterocycles, polymers and other compounds. Thus, carbolene (activated carbon) is used to absorb and remove various toxins from the body; graphite (in the form of ointments) - for the treatment of skin diseases; radioactive isotopes of carbon - for scientific research (radiocarbon analysis).

Carbon is the basis of all organic matter. Any living organism is made up largely of carbon. Carbon is the basis of life. The carbon source for living organisms is usually CO 2 from the atmosphere or water. As a result of photosynthesis, it enters biological food chains, in which living things eat each other or each other's remains and thereby extract carbon to build their own bodies. The biological carbon cycle ends with either oxidation and re-entry into the atmosphere, or disposal in the form of coal or oil.

Analytical reactions of carbonate ion CO 3 2-

Carbonates are salts of an unstable, very weak carbonic acid H 2 CO 3, which is unstable in a free state in aqueous solutions and decomposes with the release of CO 2: H 2 CO 3 - CO 2 + H 2 O

Carbonates of ammonium, sodium, rubidium, cesium are soluble in water. Lithium carbonate is slightly soluble in water. Carbonates of other metals are slightly soluble in water. Hydrocarbonates dissolve in water. Carbonate - ions in aqueous solutions are colorless and undergo hydrolysis. Aqueous solutions of alkali metal bicarbonates do not stain when a drop of phenolphthalein solution is added to them, which makes it possible to distinguish carbonate solutions from bicarbonate solutions (pharmacopoeial test).

1.Reaction with barium chloride.

Ва 2+ + СОЗ 2 - -> ВаСО 3 (white fine crystalline)

Similar precipitates of carbonates are produced by cations of calcium (CaCO 3) and strontium (SrCO 3). The precipitate dissolves in mineral acids and acetic acid. In the H 2 SO 4 solution, a white precipitate of BaSO 4 is formed.

A solution of HC1 is slowly added dropwise to the precipitate until the precipitate is completely dissolved: BaCO3 + 2 HC1 -> BaC1 2 + CO 2 + H 2 O

2. Reaction with magnesium sulfate (pharmacopoeial).

Mg 2+ + СОЗ 2 - -> MgCO 3 (white)

Bicarbonate - ion HCO 3 - forms a precipitate of MgCO 3 with magnesium sulfate only when boiling: Mg 2+ + 2 HCO3- -> MgCO 3 + CO 2 + H 2 O

The precipitate MgCO 3 dissolves in acids.

3. Reaction with mineral acids (pharmacopoeial).

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

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

H 2 CO 3 - CO 2 + H 2 O

The evolved gaseous CO 2 is detected by the turbidity of baritone or lime water in the device for detecting gases, gas bubbles (CO 2), in the test tube - the receiver - the turbidity of the solution.

4. Reaction with uranyl hexacyanoferrate (II).

2CO 3 2 - + (UО 2) 2 (brown) -> 2 UO 2 CO 3 (colorless) + 4 -

A brown solution of uranyl hexacyanoferrate (II) is obtained by mixing a solution of uranyl acetate (CH 3 COO) 2 UO 2 with a solution of potassium hexacyanoferrate (II):

2 (CH 3 COO) 2 GO 2 + K 4 -> (UO 2) 2 + 4 CH 3 COOK

To the resulting solution is added dropwise a solution of Na 2 CO 3 or K 2 CO 3 with stirring until the brown color disappears.

5. Separate discovery of carbonate - ions and bicarbonate - ions by reactions with calcium cations and with ammonia.

If carbonate - ions and bicarbonate - ions are simultaneously present in the solution, then each of them can be opened separately.

To do this, first, an excess of CaCl 2 solution is added to the analyzed solution. In this case, СОz 2 - are precipitated in the form of CaCO 3:

COz 2 - + Ca 2+ = CaCO 3

Bicarbonate - ions remain in solution, as Ca (HCO 3) 2 solutions in water. The precipitate is separated from the solution and an ammonia solution is added to the latter. HCO 2 - -anions with ammonia and calcium cations again give a precipitate of CaCO 3: HCO 3 - + Ca 2+ + NH 3 -> CaCO 3 + NH 4 +

6. Other reactions of carbonate - ion.

Carbonate ions, upon reaction with iron (III) chloride FeCl 3, form a brown precipitate Fe (OH) CO 3, with silver nitrate - a white precipitate of silver carbonate Ag 2 CO3, soluble in HbTO3 and decomposing upon boiling in water to a dark precipitate Ag 2 O iCO 2: Ag 2 CO 3 -> Ag 2 O + CO 2

Analytical reactions of acetate - ion CH 3 COO "

Acetate - ion CH 3 COO - - anion of weak monobasic acetic acid CH 3 COOH: colorless in aqueous solutions, subject to hydrolysis, does not have redox properties; rather effective ligand and forms stable acetate complexes with cations of many metals. When reacting with alcohols in an acidic medium, it gives esters.

Acetates of ammonium, alkali and most other metals are readily soluble in water. Silver acetates CH 3 COOAg and mercury (I) are less soluble in water than acetates of other metals.

1.Reaction with iron (III) chloride (pharmacopoeial).

At pH = 5-8 acetate - ion with Fe (III) cations forms a soluble dark red (strong tea color) acetate or iron (III) oxyacetate.

In aqueous solution, it is partially hydrolyzed; acidification of the solution with mineral acids suppresses hydrolysis and leads to the disappearance of the red color of the solution.

3 CH3COOH + Fe -> (CH 3 COO) 3 Fe + 3 H +

When boiling, a red-brown precipitate of basic iron (III) acetate precipitates from the solution:

(CH 3 COO) 3 Fe + 2 H 2 O<- Fe(OH) 2 CH 3 COO + 2 СН 3 СООН

Depending on the ratios of the concentrations of iron (III) and acetate ions, the composition of the sediment can change and respond, for example, to the formulas: Fe OH (CH 3 COO) 2, Fe 3 (OH) 2 O 3 (CH 3 COO), Fe 3 O (OH) (CH 3 COO) 6 or Fe 3 (OH) 2 (CH 3 COO) 7.

The reaction is interfered with by the anions CO 3 2 -, SO 3 "-, PO 4 3 -, 4, which form precipitates with iron (III), as well as SCN- anions (giving red complexes with Fe 3+ cations), iodide is the G ion, oxidized to iodine 1 2, which gives the solution a yellow color.

2. Reaction with sulfuric acid.

Acetate - an ion in a strongly acidic medium turns into weak acetic acid, the vapors of which have a characteristic vinegar smell:

CH 3 COO- + H +<- СН 3 СООН

The reaction is interfered with by the anions NO 2 \ S 2 -, SO 3 2 -, S 2 O 3 2 -, which also emit gaseous products with a characteristic odor in a concentrated H 2 SO4 environment.

3. Reaction of ethyl acetate formation (pharmacopoeial).

The reaction is carried out in a sulfuric acid medium. With ethanol:

CH 3 COO- + H + - CH 3 COOH CH 3 COOH + C 2 H 5 OH = CH 3 COOC 2 H 4 + H 2 O

Released ethyl acetate is detected by its characteristic pleasant odor. Silver salts catalyze this reaction, therefore it is recommended to add a small amount of AgNO 3 during this reaction.

Similarly, when reacting with amyl alcohol С 5 НцОН, a pleasant smelling amyl acetate СН 3 СООС 5 Н (- pearl) is also formed. A characteristic smell of ethyl acetate is felt, which intensifies with careful heating of the mixture.

Analytical reactions of tartrate - ion POC - CH (OH) - CH (OH) - COMPOSITION. Tartrate ion - the anion of a weak dibasic tartaric acid:

NO-CH-COOH

HO-CH-COOH

Tartrate - the ion is highly soluble in water. In aqueous solutions, tartrate ions are colorless, undergo hydrolysis, are prone to complexation, giving stable tartrate complexes with cations of many metals. Tartaric acid forms two rows of salts - medium tartrates containing two charged tartrate - the COCH (OH) CH (OH) COO - ion, and acidic tartrates - hydrotartrates containing a singly charged hydrotartrate - HOOOCH (OH) CH (OH) COO - ion. Potassium hydrogen tartrate (-tartar) KNS 4 H 4 O 6 is practically not a solution in water, which is used to open potassium cations. Medium calcium salt is also slightly soluble in water. Medium potassium salt K 2 C 4 H 4 O 6 is readily soluble in water.

I. Reaction with potassium chloride (pharmacopoeial).

С 4 Н 4 О 6 2 - + К + + Н + -> KNS 4 Н 4 О 6 1 (white)

2. Reaction with resorcinol in an acidic medium (pharmacopoeial).

Tartrates when heated with resorcinol meta - C 6 H 4 (OH) 2 in concentrated sulfuric acid form the reaction products of cherry red color.

14) Reactions with the ammonia complex of silver. A black precipitate of metallic silver falls out.

15) Reaction with iron (II) sulfate and hydrogen peroxide.

The addition of a dilute aqueous solution of FeSO 4 and H 2 O 2 to a solution containing tartrates. leads to the formation of an unstable iron complex of a crinkled color. Subsequent treatment with an alkali solution of NaOH leads to the formation of a blue complex.

Analytical reactions of oxalate ion С 2 О 4 2-

Oxalate - ion С 2 О 4 2 - - anion of dibasic oxalic acid Н 2 С 2 О 4 of medium strength, relatively well soluble in water. Oxalate ion in aqueous solutions is colorless, partially hydrolyzed, a strong reducing agent, an effective ligand - forms stable oxalate complexes with cations of many metals. Oxalates of alkali metals, magnesium and ammonium dissolve in water, while other metals are slightly soluble in water.

1Reaction with barium chloride Ba 2+ + C 2 O 4 2- = BaC 2 O 4 (white) The precipitate dissolves in mineral acids and in acetic acid (during boiling). 2. Reaction with calcium chloride (pharmacopoeial): Ca 2+ + C 2 O 4 2 - = CaC 2 O 4 (white)

The precipitate dissolves in mineral acids, but does not dissolve in acetic acid.

3. Reaction with silver nitrate.

2 Ag + + С 2 О 4 2 - -> Ag2C2O 4. |. (Curd) Test for solubility. The sediment is divided into 3 parts:

a). In the first test tube with the precipitate, add dropwise with stirring a solution of HNO 3 until the precipitate dissolves;

b). A concentrated ammonia solution is added dropwise to the second test tube with the precipitate with stirring until the precipitate dissolves; v). Add 4-5 drops of HC1 solution to the third test tube with sediment; a white precipitate of silver chloride remains in the test tube:

Ag 2 C 2 O 4 + 2 HC1 -> 2 AC1 (white) + H 2 C 2 O 4

4. Reaction with potassium permanganate. Oxalate ions with KMnO 4 in an acidic medium are oxidized with the release of CO 2; the KMnO 4 solution is discolored due to the reduction of manganese (VII) to manganese (II):

5 C 2 O 4 2 - + 2 MnO 4 "+ 16 H + -> 10 CO 2 + 2 Mn 2+ + 8 H 2 O

Diluted solution of KMnO 4. The latter is discolored; there is a release of gas bubbles - CO 2.

38 Elements of group VA

General characteristics of VA group of the Periodic system. in the form of s x p y is the electronic configuration of the external energy level of the elements of the VA group.

Arsenic and antimony have different allotropic modifications: both with molecular and metal crystal lattices. However, based on a comparison of the stability of cationic forms (As 3+, Sb 3+), arsenic is referred to as non-metals, and antimony to metals.

oxidation states stable for group VA elements

From nitrogen to bismuth (with a decrease in non-metallic properties):

w decreases the stability of the negative oxidation state (-3) (m. properties of hydrogen compounds)

w decreases the stability of the highest positive oxidation state (+5)

w increased stability of low positive oxidation state (+3)