Chemical properties of non-metals. Interaction of metals with non-metals
INTERACTION OF METALS WITH NONMETALS
Non-metals exhibit oxidizing properties in reactions with metals, accepting electrons from them and recovering.
Interaction with halogens
Halogens (F 2, Cl 2, Br 2, I 2 ) are strong oxidizing agents, therefore, all metals interact with them under normal conditions:
2Me + n Hal 2 → 2 MeHal n
The product of this reaction is a metal halide salt ( MeF n -fluoride, MeCl n -chloride, MeBr n -bromide, MeI n -iodide). When interacting with a metal, the halogen is reduced to the lowest oxidation state (-1), andnequal to the oxidation state of the metal.
The reaction rate depends on the chemical activity of the metal and halogen. The oxidative activity of halogens decreases in the group from top to bottom (from F to I ).
Interaction with oxygen
Oxygen oxidizes almost all metals (except Ag, Au, Pt ), resulting in the formation of oxides Me 2 O n .
active metals easily interact with atmospheric oxygen under normal conditions.
2 Mg + O 2 → 2 MgO (with flash)
Intermediate activity metals also react with oxygen at ordinary temperature. But the rate of such a reaction is significantly lower than with the participation of active metals.
Inactive metals oxidized by oxygen when heated (combustion in oxygen).
oxides Chemical properties of metals can be divided into three groups:
1. Basic oxides ( Na 2 O, CaO, Fe II O, Mn II O, Cu I O etc.) are formed by metals in low oxidation states (+1, +2, as a rule, below +4). Basic oxides interact with acidic oxides and acids to form salts:
CaO + CO 2 → CaCO 3
CuO + H 2 SO 4 → CuSO 4 + H 2 O
2. Acid oxides ( Cr VI O 3 , Fe VI O 3 , Mn VI O 3 , Mn 2 VII O 7 etc.) are formed by metals in high oxidation states (as a rule, above +4). Acid oxides interact with basic oxides and bases to form salts:
FeO 3 + K 2 O → K 2 FeO 4
CrO 3 + 2KOH → K 2 CrO 4 + H 2 O
3. Amphoteric oxides ( BeO, Al 2 O 3, ZnO, SnO, MnO 2, Cr 2 O 3, PbO, PbO 2 etc.) have a dual nature and can interact with both acids and bases:
Cr 2 O 3 + 3H 2 SO 4 → Cr 2 (SO 4) + 3H 2 O
Cr 2 O 3 + 6NaOH → 2Na 3
Interaction with sulfur
All metals interact with sulfur (except Au ), forming salts - sulfides Me 2 S n . In this case, sulfur is reduced to the oxidation state of "-2". Platinum ( Pt ) interacts with sulfur only in a finely divided state. alkali metals, and Ca and Mg react with sulfur when heated with an explosion. Zn, Al (powder) and Mg in reaction with sulfur give a flash. In the direction from left to right in the activity series, the rate of interaction of metals with sulfur decreases.
Interaction with hydrogen
With hydrogen, some active metals form compounds - hydrides:
2 Na + H 2 → 2 NaH
In these compounds, hydrogen is in its rare oxidation state "-1".
E.A. Nudnova, M.V. Andriukhova
CHEMICAL PROPERTIES OF HYDROGEN
1. WITH METALS
(Li, Na, K, Rb, Cs, Ca, Sr, Ba) → with alkali and alkaline earth metals, when heated, forms solid unstable substances hydrides, other metals do not react.
2K + H₂ = 2KH (potassium hydride)
Ca + H₂ = CaH₂
2. WITH NONMETALS
with oxygen, halogens under normal conditions, when heated, it reacts with phosphorus, silicon and carbon, with nitrogen under pressure and a catalyst.
2Н₂ + O₂ = 2Н₂O Н₂ + Cl₂ = 2HCl
3Н₂ + N₂↔ 2NH₃ H₂ + S = H₂S
3. INTERACTION WITH WATER
Does not react with water
4. INTERACTION WITH OXIDES
Reduces oxides of metals (inactive) and non-metals to simple substances:
CuO + H₂ = Cu + H₂O 2NO + 2H₂ = N₂ + 2H₂O
SiO₂ + H₂ = Si + H₂O
5. INTERACTION WITH ACIDS
Does not react with acids
6. INTERACTION WITH ALKALI
Does not react with alkalis
7. INTERACTION WITH SALT
Restores inactive metals from salts
CuCl₂ + H₂ = Cu + 2HCl
CHEMICAL PROPERTIES OF OXYGEN
1. INTERACTION WITH METALS
With alkali metals under normal conditions - oxides and peroxides (lithium - oxide, sodium - peroxide, potassium, cesium, rubidium - superoxide
4Li + O2 = 2Li2O (oxide)
2Na + O2 = Na2O2 (peroxide)
K+O2=KO2 (superoxide)
With the rest of the metals of the main subgroups, under normal conditions, it forms oxides with an oxidation state equal to the group number
2 FROMa+O2=2FROMaO
4Al + O2 = 2Al2O3
1. INTERACTION WITH METALS
With metals of secondary subgroups, under normal conditions and when heated, it forms oxides of various degrees of oxidation, and with iron, iron scaleFe3 O4 ( FeO∙ Fe2 O3)
3Fe + 2O2 = Fe3O4 4Cu + O₂ = 2Cu₂⁺¹O (red);
2Cu + O₂ = 2Cu⁺²O (black); 2Zn + O₂ = ZnO
4Cr + 3О2 = 2Cr2⁺³О3
forms oxides - often of an intermediate oxidation state
C + O₂(ex)=CO₂; C+ O₂ (week) =CO
S + O₂ = SO₂N₂ + O₂ = 2NO - Q
3. INTERACTION WITH WATER
Does not react with water
4. INTERACTION WITH OXIDES
Oxidizes lower oxides to oxides with a higher oxidation state
Fe⁺²O + O2 = Fe2⁺³O3; C⁺²O + O2 = C⁺⁴O2
5. INTERACTION WITH ACIDS
Anhydrous anoxic acids (binary compounds) burn in an oxygen atmosphere
2H2S + O2 = 2S + 2H2O 2H2S + 3O2 = 2SO2 + 2H2O
In oxygen-containing, it increases the degree of oxidation of the non-metal.
2HN⁺³O2 + O2 = 2HN⁺⁵O3
6. INTERACTION WITH BASES
Oxidizes unstable hydroxides in aqueous solutions to a higher oxidation state
4Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3
7. INTERACTION WITH SALT AND BINARY COMPOUNDS
Enters into combustion reactions.
4FeS2 +11O2 = 2Fe2O3 + 8SO2
CH4 + 2O2 = CO2 + 2H2O
4NH3 + 3O2 = 2N2 + 6H2O
catalytic oxidation
NH3 + O2 = NO + H2O
CHEMICAL PROPERTIES OF THE HALOGENS
1. INTERACTION WITH METALS
With alkaline under normal conditions, withF, Cl, Brignite:
2 Na + Cl2 = 2 NaCl(chloride)
Alkaline earth and aluminum react under normal conditions:
FROMa+Cl2=FROMaCl2 2Al+3Cl2 = 2AlCl3
Metals of secondary subgroups at elevated temperatures
Cu + Cl₂ = Cu⁺²Cl₂
2Cu + I₂ = 2Cu⁺¹I (there is no copper (II) iodide!)
2Fe + ЗС12 = 2Fe⁺³Cl3 iron (III) chloride
Fluorine reacts with metals (often explosively), including gold and platinum.
2Au + 3F₂ = 2AuF
2. INTERACTION WITH NON-METALS
They do not directly interact with oxygen (except for F₂), they react with sulfur, phosphorus, silicon. The chemical activity of bromine and iodine is less pronounced than that of fluorine and chlorine:
H2 +F2 = 2NF ; Si + 2 F2 = SiF4.; 2 P + 3 Cl2 = 2 P⁺³ Cl3; 2 P + 5 Cl2 = 2 P⁺⁵ Cl5; S + 3 F2 = S⁺⁶ F6;
S + Cl2 = S⁺²Cl2
F₂
Reacts with oxygen:F2 + O2 = O⁺² F2
Reacts with other halogens:Cl₂ + F₂ = 2 Cl⁺¹ F¯¹
Reacts even with inert gases 2F₂ + Xe= Xe⁺⁸ F₄¯¹.
3. INTERACTION WITH WATER
Fluorine under normal conditions forms hydrofluoric acid + + O₂
2F2 + 2H2O → 4HF + O2
Chlorine, when the temperature rises, forms hydrochloric acid + O₂,
2Сl₂ + 2H₂O → 4HCl + O₂
at n.o. - "chlorine water"
Сl2 + Н2О ↔ НCl + НClO (hydrochloric and hypochlorous acids)
Bromine under normal conditions forms "bromine water"
Br2 + H2O ↔ HBr + HBrO (hydrobromic and hypobromous acids
Iodine → no reaction
I2 + H₂O ≠
5. INTERACTION WITH OXIDES
Only fluorine F₂ REACTS, displacing oxygen from the oxide, forming fluorides
SiO2‾² + 2F2⁰ = SiF4‾¹ + O2⁰
6. INTERACTION WITH ACIDS.
react with oxygen-free acids, displacing less active non-metals.
H2S‾² + I2⁰ → S⁰↓+ 2HI‾
7. INTERACTION WITH ALKALI
Fluorine forms fluoride + oxygen and water
2F2 + 4NaOH = 4NaF¯¹ + O2 + 2H2O
Chlorine, when heated, forms chloride, chlorate and water.
3 Cl₂ + 6 KOH = 5 KCl¯¹ + KCl⁺⁵ O3 + 3 H2 O
In the cold, chloride, hypochlorate and water, with calcium hydroxide bleach and water
Cl2 + 2KOH-(cold)= KCl¯¹ + KCl⁺¹O + H2O
Cl2 + Ca(OH) 2 = CaOCl2 (bleach - mixture of chloride, hypochlorite and hydroxide) + H2O
Bromine when heated → bromide, bromate and and water
3Br2 + 6KOH =5KBr¯¹ + KBr ⁺⁵O3 + 3H2O
Iodine when heated → iodide, iodate and water
3I2 + 6NaOH = 5NaI¯¹ + NaI ⁺⁵O3 + 3H2O
9. INTERACTION WITH SALT
Displacement of less active halogens from salts
2KBr + Cl2 → 2KCl + Br2
2KCl + Br2 ≠
2KCl + F2 → 2KF + Cl2
2KBr + J2≠
Oxidize non-metals in salts to a higher oxidation state
2Fe⁺²Cl2 + Cl2⁰ → 2Fe⁺³Cl 3 ‾¹
Na2S⁺⁴O3 + Br2⁰ + 2H2O →Na2S⁺⁶O4 + 2HBr‾
CHEMICAL PROPERTIES OF SULFUR
1. INTERACTION WITH METALSreacts when heated even with alkali metals, with mercury under normal conditions: with sulfur - sulfides:
2K + S = K2S
2Cr + 3S = Cr2⁺³S3 Fe + S = Fe⁺²S
2. INTERACTION WITH NON-METALS
When heated with hydrogen,coxygen (sulfur dioxide)chalogens (except iodine), with carbon, nitrogen and silicon and does not react
S + Cl₂ = S⁺²Cl₂ ; S + O₂ =S⁺⁴O₂
H₂ + S = H₂S¯² ; 2P + 3S = P₂S₃¯²
FROM+ 3S = CS₂¯²
WITH WATER, OXIDES, SALT
DOES NOT REACT
3. INTERACTION WITH ACIDS
Oxidized by sulfuric acid when heated to sulfur dioxide and water
2H2SO4 (conc) = 2H2O + 3S⁺⁴O2
Nitric acid when heated to sulfuric acid, nitric oxide (+4) and water
S + 6HNO3(conc) =H2SO4 + 6N⁺⁴O2 + 2H2O
4. INTERACTION WITH ALKALI
Forms sulfite when heated, sulfide + water
3S + 6KOH = K2SO3 + 2K2S + 3H2O
CHEMICAL PROPERTIES OF NITROGEN
1. INTERACTION WITH METALSreactions proceed when heated (exception: lithium with nitrogen under normal conditions):
With nitrogen - nitrides
6Li + N2 = 3Li2N (lithium nitride) (n.o.) 3Mg + N2 = Mg3N2 (magnesium nitride) 2Cr + N2 = 2CrN
Iron in these compounds has an oxidation state of +2
2. INTERACTION WITH NON-METALS
(due to the triple bond, nitrogen is very inactive). Under normal conditions, it does not react with oxygen. Reacts with oxygen only when high temperature(electric arc), in nature - during a thunderstorm
N2+O2=2NO (email. arc, 3000 0C)
With hydrogen at high pressure, elevated temperature and in the presence of a catalyst:
t,p,kat
3N2+3H2 ↔ 2NH3
WITH WATER, OXIDES, ACIDS, ALKALS AND SALT
DOES NOT REACT
CHEMICAL PROPERTIES OF PHOSPHORUS
1. INTERACTION WITH METALS
reactions proceed when heated with phosphorus - phosphides
3Ca + 2P = K3P2, Iron in these compounds has an oxidation state of +2
2. INTERACTION WITH NON-METALS
Combustion in oxygen
4P + 5O₂ = 2P₂⁺⁵O₅ 4P + 3O₂ = 2P₂⁺³O₃
With halogens and sulfur when heated
2P + 3Cl₂ = 2P⁺³Cl₃ 2P + 5Cl₂ = 2P⁺⁵Cl₅; 2P + 5S = P₂⁺⁵S₅
Does not interact directly with hydrogen, carbon, silicon
WITH WATER AND OXIDES
DOES NOT REACT
3. INTERACTION WITH ACIDS
With concentrated nitric acid nitric oxide (+4), with dilute nitric oxide (+2) and phosphoric acid
3P + 5HNO₃(conc) =3H₃PO₄ + 5N⁺⁴O₂
3P + 5HNO₃ + 2H₂O = 3H₃PO₄ + 5N⁺²O
With concentrated sulfuric acid, phosphoric acid, sulfur oxide (+4) and water are formed
3P + 5H₂SO₄(conc.) =3H₃PO₄ + 5S⁺⁴O₂+ 2H₂O
4. INTERACTION WITH ALKALI
Forms phosphine and hypophosphite with alkali solutions
4P⁰ + 3NaOH + 3H2O = P¯³H 3 + 3NaH 2 P ⁺1O 2
5. INTERACTION WITH SALT
5. INTERACTION WITH SALT
With strong oxidizing agents, exhibiting reducing properties
3P⁰ + 5NaN⁺⁵O₃ = 5NaN⁺³O₂ + P₂⁺⁵O₅
CHEMICAL PROPERTIES OF CARBON
1. INTERACTION WITH METALS
reactions take place when heated
Metals - d-elements form with carbon compounds of non-stoichiometric composition such as solid solutions: WC, ZnC, TiC - are used to obtain superhard steels
with carbon carbides 2Li + 2C = Li2C2,
Ca + 2C = CaC2
2. INTERACTION WITH NON-METALS
Of the halogens, it directly reacts only with fluorine, with the rest when heated.
С + 2F₂ = CF₄.
Interaction with oxygen:
2C + O₂ (lack) \u003d 2C⁺²O (carbon monoxide),
С + О₂(ex) = С⁺⁴О₂(carbon dioxide).
Interaction with other non-metals at elevated temperature, does not interact with phosphorus
C + Si = SiC¯⁴ ; C + N₂ = C₂⁺⁴N₂ ;
C + 2H₂ = C¯⁴H₄ ; C + 2S = C⁺⁴S₂;
3. INTERACTION WITH WATER
The passage of water vapor through hot coal - carbon monoxide and hydrogen are formed (synthesis gas
C + H₂O = CO + H₂
4. INTERACTION WITH OXIDES
CARBON REDUCES METALS AND NON-METALS FROM OXIDES TO A SIMPLE SUBSTANCE WHEN HEATED (CARBOTHERMY), reduces the degree of oxidation in carbon dioxide
2ZnO + C = 2Zn + CO; fourFROM+ Fe₃O₄ = 3Fe + 4CO;
P₂O₅ + C = 2P + 5CO; 2FROM+ SiO₂ = Si + 2CO;
FROM+ C⁺⁴O₂ = 2C⁺²O
5. INTERACTION WITH ACIDS
Oxidized by concentrated nitric and sulfuric acids to carbon dioxide
C +2H2SO4(conc)=C⁺⁴O2+ 2S⁺⁴O2+ 2H2O; C+4HNO3 (conc) = C⁺⁴O2 + 4N⁺⁴O2 + 2H2O.
WITH ALKALI AND SALT
DOES NOT REACT
CHEMICAL PROPERTIES OF SILICON
1. INTERACTION WITH METALS
reactions proceed when heated: active metals react with silicon - silicides
4Cs + Si = Cs4Si,
1. INTERACTION WITH NON-METALS
From halogens directly only with fluorine.
Reacts with chlorine when heated
Si + 2F2 = SiF4; Si + 2Cl2 = SiCl4;
Si + O₂ = SiO₂; Si+C=SiC; 3Si + 2N₂ = Si₃N;
Does not interact with hydrogen
3. INTERACTION WITH ACIDS
interacts only with a mixture of hydrofluoric and nitric acids, forming hexafluorosilicic acid
3Si + 4HNO₃ + 18HF = 3H₂ + 4NO + 8H₂O
Interaction with hydrogen halides (these are not acids) - displaces hydrogen, silicon halides and hydrogen are formed
Reacts with hydrogen fluoride under normal conditions.
Si + 4HF = SiF₄ + 2H₂
4. INTERACTION WITH ALKALI
It dissolves when heated in alkalis, forming silicate and hydrogen:
Si + 2NaOH + H₂O = Na₂SiO₃ + 2H₂
Lecture 24
Nonmetals.
Lecture plan:
Non-metals are simple substances
The position of non-metals in the periodic system
The number of non-metal elements is much less than that of metal elements. Ten have typical non-metallic properties. chemical elements(H, C, N, P, O, S, F, Cl, Br, I). Six elements, which are usually referred to as non-metals, exhibit dual (both metallic and non-metallic) properties (B, Si, As, Se, Te, At). And 6 more items in recent times began to be included in the list of non-metals. These are the so-called noble (or inert) gases (He, Ne, Ar, Kg, Xe, Rn). So, 22 of the known chemical elements are usually classified as non-metals.
Elements that exhibit non-metallic properties in the periodic system are located above the boron-astat diagonal (Fig. 26).
The atoms of most non-metals, unlike metal atoms, have a large number of electrons on the outer electron layer - from 4 to 8. The exceptions are the atoms of hydrogen, helium, boron, which have 1, 2 and 3 electrons at the outer level, respectively.
Among non-metals, only two elements - hydrogen (1s 1) and helium (1s 2) belong to the s-family, all the rest belong to R-family .
Atoms of typical non-metals (A) are characterized by high electronegativity and high electron affinity, which determines their ability to form negatively charged ions with the electronic configurations of the corresponding inert gases:
A 0 + nê → A n -
These ions are part of ionic compounds of non-metals with typical metals. Non-metals also have negative oxidation states in covalent compounds with other less electronegative non-metals (in particular, with hydrogen).
Atoms of non-metals in covalent compounds with more electronegative non-metals (in particular, with oxygen) have positive oxidation states. Highest positive oxidation state of a non-metal, usually, equal to group number in which it is located.
Non-metals are simple substances
Despite not big number non-metal elements, their role and significance both on Earth and in space are enormous. 99% of the mass of the Sun and other stars are non-metals hydrogen and helium. air shell Earth consists of non-metal atoms - nitrogen, oxygen and noble gases. The Earth's hydrosphere is formed by one of the most important substances for life - water, the molecules of which consist of non-metals hydrogen and oxygen. In living matter, 6 non-metals predominate - carbon, oxygen, hydrogen, nitrogen, phosphorus, sulfur.
Under normal conditions, non-metal substances exist in different states of aggregation:
1) gases: hydrogen H 2, oxygen O 2, nitrogen N 2, fluorine F 2, chlorine C1 2, inert gases: He, Ne, Ar, Kg, Xe, Rn
2) liquid: bromine Br 2
3) solid substances iodine I 2, carbon C, silicon Si, sulfur S, phosphorus P, etc.
Seven non-metal elements form simple substances that exist in the form of diatomic molecules E 2 (hydrogen H 2, oxygen O 2, nitrogen N 2, fluorine F 2, chlorine C1 2, bromine Br 2, iodine I 2).
Since there are no free electrons between atoms in the crystal lattice of non-metals, they differ in physical properties from metals:
¾ do not have gloss;
¾ fragile, have different hardness;
¾ conduct heat and electricity poorly.
Non-metal solids are practically insoluble in water; gaseous O 2 , N 2 , H 2 and halogens have a very low solubility in water.
A number of non-metals are characterized allotropy- the phenomenon of the existence of one element in the form of several simple substances. Allotropic modifications known for oxygen (oxygen O 2 and ozone O 3), sulfur (rhombic, monoclinic and plastic), phosphorus (white, red and black), carbon (graphite, diamond and carbine, etc.), silicon (crystalline and amorphous).
Chemical properties of non-metals
According to the chemical activity of non-metals differ significantly from each other. So, nitrogen and noble gases enter into chemical reactions only under very harsh conditions ( high pressure and temperature, the presence of a catalyst).
The most reactive non-metals are halogens, hydrogen and oxygen. Sulfur, phosphorus, and especially carbon and silicon, are reactive only at elevated temperatures.
Nonmetals in chemical reactions exhibit both oxidizing and reducing properties. The highest oxidizing capacity is characteristic of halogens and oxygen. In such non-metals as hydrogen, carbon, silicon, reducing properties predominate.
I. Oxidizing properties of non-metals:
1. Interaction with metals. In this case, binary compounds are formed: with oxygen - oxides, with hydrogen - hydrides, nitrogen - nitrides, halogens - halides, etc.:
2Cu + O 2 → 2CuO
2Fe + 3Cl 2 → 2FeCl 3
2. Interaction with hydrogen. Non-metals also act as oxidizing agents in reactions with hydrogen, forming volatile hydrogen compounds:
H 2 + C1 2 → 2HC1
N 2 + 3H 2 → t, p, cat. 2NH3
3. Interaction with non-metals. Non-metals also exhibit oxidizing properties in reactions with less electronegative non-metals:
2P + 5C1 2 → 2PC1 5 ;
C + 2S → CS 2 .
4. Interaction with complex substances. The oxidizing properties of nonmetals can also manifest themselves in reactions with complex substances. For example, water burns in an atmosphere of fluorine:
2F 2 + 2H 2 O → 4HF + O 2.
II. Reducing properties of non-metals
1. Interaction with non-metals. Non-metals can exhibit reducing properties in relation to non-metals with higher electronegativity, and primarily in relation to fluorine and oxygen:
4P + 5O 2 → 2P 2 O 5;
N 2 + O 2 → 2NO
2. Interaction with complex substances. Some non-metals can be reducing agents, which allows them to be used in metallurgical production:
C + ZnO → Zn + CO;
5H 2 + V 2 O 5 → 2V + 5H 2 O.
SiO 2 + 2C → Si + 2CO.
Non-metals exhibit reducing properties when interacting with complex substances - strong oxidizing agents, for example:
3S + 2KSlO 3 → 3SO 2 + 2KS1;
6P + 5KSlO 3 → ZR 2 O 5 + 5KS1.
C + 2H 2 SO 4 → CO 2 + 2SO 2 + 2H 2 O;
3P + 5HNO 3 + 2H 2 O → ZH 3 RO 4 + 5NO.
General methods for obtaining non-metals
Some non-metals are found in nature in a free state: these are sulfur, oxygen, nitrogen, noble gases. First of all, simple substances - non-metals are part of the air.
Large amounts of gaseous oxygen and nitrogen are obtained by rectification of air (separation).
The most active non-metals - halogens - are obtained by electrolysis of melts or solutions from compounds. In industry, with the help of electrolysis, three most important products are simultaneously obtained in large quantities: the closest analogue of fluorine is chlorine, hydrogen, and sodium hydroxide. The electrolyte used is a sodium chloride solution fed into the cell from above.
In more detail, methods for obtaining non-metals will be discussed later in the relevant lectures.
If we draw a diagonal from beryllium to astatine in the periodic table of elements of D.I. Mendeleev, then there will be metal elements on the diagonal at the bottom left (they also include elements of secondary subgroups, highlighted in blue), and non-metal elements at the top right (highlighted yellow). Elements located near the diagonal - semimetals or metalloids (B, Si, Ge, Sb, etc.) have a dual character (highlighted in pink).
As can be seen from the figure, the vast majority of elements are metals.
In its own way chemical nature Metals are chemical elements whose atoms donate electrons from the outer or pre-outer energy levels, thus forming positively charged ions.
Almost all metals have relatively large radii and a small number of electrons (from 1 to 3) at the external energy level. Metals are characterized by low electronegativity values and reducing properties.
The most typical metals are located at the beginning of periods (starting from the second), further from left to right, the metallic properties weaken. In a group from top to bottom, metallic properties are enhanced, because the radius of atoms increases (due to an increase in the number of energy levels). This leads to a decrease in the electronegativity (the ability to attract electrons) of the elements and an increase in the reducing properties (the ability to donate electrons to other atoms in chemical reactions).
typical metals are s-elements (elements of the IA group from Li to Fr. elements of the PA group from Mg to Ra). The general electronic formula of their atoms is ns 1-2. They are characterized by oxidation states + I and + II, respectively.
The small number of electrons (1-2) in the outer energy level of typical metal atoms suggests that these electrons are easily lost and exhibit strong reducing properties, which reflect low electronegativity values. This implies the limited chemical properties and methods for obtaining typical metals.
A characteristic feature of typical metals is the tendency of their atoms to form cations and ionic chemical bonds with non-metal atoms. Compounds of typical metals with non-metals are ionic crystals "metal cation anion of non-metal", for example, K + Br -, Ca 2+ O 2-. Typical metal cations are also included in compounds with complex anions - hydroxides and salts, for example, Mg 2+ (OH -) 2, (Li +) 2CO 3 2-.
The A-group metals forming the amphoteric diagonal in the Be-Al-Ge-Sb-Po Periodic Table, as well as the metals adjacent to them (Ga, In, Tl, Sn, Pb, Bi) do not exhibit typically metallic properties. The general electronic formula of their atoms ns 2 np 0-4 suggests a greater variety of oxidation states, a greater ability to retain their own electrons, a gradual decrease in their reducing ability and the appearance of an oxidizing ability, especially in high oxidation states ( characteristic examples- compounds Tl III, Pb IV, Bi v). A similar chemical behavior is also characteristic of most (d-elements, i.e., elements of the B-groups of the Periodic Table (typical examples are the amphoteric elements Cr and Zn).
This manifestation of duality (amphoteric) properties, both metallic (basic) and non-metallic, is due to the nature of the chemical bond. In the solid state, compounds of atypical metals with non-metals contain predominantly covalent bonds (but less strong than bonds between non-metals). In solution, these bonds are easily broken, and the compounds dissociate into ions (completely or partially). For example, gallium metal consists of Ga 2 molecules, in the solid state aluminum and mercury (II) chlorides AlCl 3 and HgCl 2 contain strongly covalent bonds, but in a solution AlCl 3 dissociates almost completely, and HgCl 2 - to a very small extent (and then into HgCl + and Cl - ions).
General physical properties of metals
Due to the presence of free electrons ("electron gas") in the crystal lattice, all metals exhibit the following characteristic general properties:
1) Plastic- the ability to easily change shape, stretch into a wire, roll into thin sheets.
2) metallic luster and opacity. This is due to the interaction of free electrons with light incident on the metal.
3) Electrical conductivity. It is explained by the directed movement of free electrons from the negative to the positive pole under the influence of a small potential difference. When heated, the electrical conductivity decreases, because. with an increase in temperature, vibrations of atoms and ions in the nodes increase crystal lattice, which hinders the directed movement of the "electron gas".
4) Thermal conductivity. It is due to the high mobility of free electrons, due to which the temperature is quickly equalized by the mass of the metal. The highest thermal conductivity is in bismuth and mercury.
5) Hardness. The hardest is chrome (cuts glass); the softest - alkali metals - potassium, sodium, rubidium and cesium - are cut with a knife.
6) Density. It is less the less atomic mass metal and larger atomic radius. The lightest is lithium (ρ=0.53 g/cm3); the heaviest is osmium (ρ=22.6 g/cm3). Metals having a density less than 5 g/cm3 are considered "light metals".
7) Melting and boiling points. The most fusible metal is mercury (m.p. = -39°C), the most refractory metal is tungsten (t°m. = 3390°C). Metals with t°pl. above 1000°C are considered refractory, below - low melting point.
General chemical properties of metals
Strong reducing agents: Me 0 – nē → Me n +
A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions. 
I. Reactions of metals with non-metals
1) With oxygen:
2Mg + O 2 → 2MgO
2) With sulfur:
Hg + S → HgS
3) With halogens:
Ni + Cl 2 – t° → NiCl 2
4) With nitrogen:
3Ca + N 2 – t° → Ca 3 N 2
5) With phosphorus:
3Ca + 2P – t° → Ca 3 P 2
6) With hydrogen (only alkali and alkaline earth metals react):
2Li + H 2 → 2LiH
Ca + H 2 → CaH 2
II. Reactions of metals with acids
1) Metals standing in the electrochemical series of voltages up to H reduce non-oxidizing acids to hydrogen:
Mg + 2HCl → MgCl 2 + H 2
2Al+ 6HCl → 2AlCl 3 + 3H 2
6Na + 2H 3 PO 4 → 2Na 3 PO 4 + 3H 2
2) With oxidizing acids:
In the interaction of nitric acid of any concentration and concentrated sulfuric acid with metals hydrogen is never released! 
Zn + 2H 2 SO 4 (K) → ZnSO 4 + SO 2 + 2H 2 O
4Zn + 5H 2 SO 4(K) → 4ZnSO 4 + H 2 S + 4H 2 O
3Zn + 4H 2 SO 4(K) → 3ZnSO 4 + S + 4H 2 O
2H 2 SO 4 (c) + Cu → Cu SO 4 + SO 2 + 2H 2 O
10HNO 3 + 4Mg → 4Mg(NO 3) 2 + NH 4 NO 3 + 3H 2 O
4HNO 3 (c) + Сu → Сu (NO 3) 2 + 2NO 2 + 2H 2 O
III. Interaction of metals with water
1) Active (alkali and alkaline earth metals) form a soluble base (alkali) and hydrogen:
2Na + 2H 2 O → 2NaOH + H 2
Ca+ 2H 2 O → Ca(OH) 2 + H 2
2) Metals of medium activity are oxidized by water when heated to oxide:
Zn + H 2 O – t° → ZnO + H 2
3) Inactive (Au, Ag, Pt) - do not react.
IV. Displacement by more active metals of less active metals from solutions of their salts:
Cu + HgCl 2 → Hg + CuCl 2
Fe+ CuSO 4 → Cu+ FeSO 4
In industry, not pure metals are often used, but their mixtures - alloys in which the beneficial properties of one metal are complemented by the beneficial properties of another. So, copper has a low hardness and is of little use for the manufacture of machine parts, while alloys of copper with zinc ( brass) are already quite hard and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but is too soft. On its basis, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing useful properties aluminum, acquires high hardness and becomes suitable in the aircraft industry. Alloys of iron with carbon (and additions of other metals) are widely known cast iron and steel.
Metals in free form are reducing agents. However, the reactivity of some metals is low due to the fact that they are covered with surface oxide film, to varying degrees resistant to the action of such chemical reagents as water, solutions of acids and alkalis.
For example, lead is always covered with an oxide film; its transition into solution requires not only exposure to a reagent (for example, dilute nitric acid), but also heating. The oxide film on aluminum prevents its reaction with water, but is destroyed under the action of acids and alkalis. Loose oxide film (rust), formed on the surface of iron in moist air, does not interfere with the further oxidation of iron.
Under the influence concentrated acids are formed on metals sustainable oxide film. This phenomenon is called passivation. So, in concentrated sulfuric acid passivated (and then do not react with acid) such metals as Be, Bi, Co, Fe, Mg and Nb, and in concentrated nitric acid - metals A1, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb , Th and U.
When interacting with oxidizing agents in acidic solutions, most metals turn into cations, the charge of which is determined by the stable oxidation state of a given element in compounds (Na +, Ca 2+, A1 3+, Fe 2+ and Fe 3+)
The reducing activity of metals in an acidic solution is transmitted by a series of stresses. Most metals are converted into a solution of hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only sulfuric (concentrated) and nitric acids, and Pt and Au - "aqua regia".
Corrosion of metals
unwanted chemical property metals is theirs, i.e. active destruction (oxidation) upon contact with water and under the influence of oxygen dissolved in it (oxygen corrosion). For example, the corrosion of iron products in water is widely known, as a result of which rust is formed, and the products crumble into powder.
Corrosion of metals proceeds in water also due to the presence of dissolved CO 2 and SO 2 gases; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 ( hydrogen corrosion).
The point of contact between two dissimilar metals can be especially corrosive ( contact corrosion). Between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water, a galvanic couple occurs. The flow of electrons goes from the more active metal, which is to the left in the series of voltages (Re), to the less active metal (Sn, Cu), and the more active metal is destroyed (corrodes).
It is because of this that the tinned surface of cans (tin-plated iron) rusts when stored in a humid atmosphere and carelessly handled (iron quickly collapses after even a small scratch appears, allowing contact of iron with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, because even if there are scratches, it is not iron that corrodes, but zinc (a more active metal than iron).
Corrosion resistance for a given metal is enhanced when it is coated with a more active metal or when they are fused; for example, coating iron with chromium or making an alloy of iron with chromium eliminates the corrosion of iron. Chrome-plated iron and steel containing chromium ( stainless steel) have high corrosion resistance.
electrometallurgy, i.e., obtaining metals by electrolysis of melts (for the most active metals) or salt solutions;
pyrometallurgy, i.e., the recovery of metals from ores at high temperature (for example, the production of iron in the blast furnace process);
hydrometallurgy, i.e., the isolation of metals from solutions of their salts by more active metals (for example, the production of copper from a CuSO 4 solution by the action of zinc, iron or aluminum).
Native metals are sometimes found in nature (typical examples are Ag, Au, Pt, Hg), but more often metals are in the form of compounds ( metal ores ). In terms of prevalence in earth's crust metals are different: from the most common - Al, Na, Ca, Fe, Mg, K, Ti) to the rarest - Bi, In, Ag, Au, Pt, Re.
Chemical properties of non-metals
In accordance with the numerical values of the relative electronegativity the oxidizing power of non-metals increases in the following order: Si, B, H, P, C, S, I, N, Cl, O, F.
Nonmetals as oxidizers
The oxidizing properties of non-metals are manifested when they interact:
· with metals: 2Na + Cl 2 = 2NaCl;
· with hydrogen: H 2 + F 2 = 2HF;
· with non-metals that have a lower electronegativity: 2P + 5S = P 2 S 5;
· with some complex substances: 4NH 3 + 5O 2 = 4NO + 6H 2 O,
2FeCl 2 + Cl 2 \u003d 2 FeCl 3.
Nonmetals as reducing agents1. All non-metals (except fluorine) exhibit reducing properties when interacting with oxygen:
S + O 2 \u003d SO 2, 2H 2 + O 2 \u003d 2H 2 O.
Oxygen in combination with fluorine can also exhibit a positive oxidation state, i.e., be a reducing agent. All other non-metals exhibit reducing properties. So, for example, chlorine does not combine directly with oxygen, but its oxides (Cl 2 O, ClO 2, Cl 2 O 2) can be obtained indirectly, in which chlorine exhibits a positive oxidation state. Nitrogen at high temperatures directly combines with oxygen and exhibits reducing properties. Sulfur reacts even more easily with oxygen.
2. Many non-metals exhibit reducing properties when interacting with complex substances:
ZnO + C \u003d Zn + CO, S + 6HNO 3 conc \u003d H 2 SO 4 + 6NO 2 + 2H 2 O.
3. There are also such reactions in which the same non-metal is both an oxidizing agent and a reducing agent:
Cl 2 + H 2 O \u003d HCl + HClO.
4. Fluorine is the most typical non-metal, which is not characterized by reducing properties, i.e., the ability to donate electrons in chemical reactions.
Compounds of non-metalsNonmetals can form compounds with different intramolecular bonds.
Types of non-metal compounds
General formulas of hydrogen compounds by groups of the periodic system of chemical elements are given in the table:
|
RH 2 |
RH 3 |
RH4 |
RH 3 |
H2R |
||
|
Non-volatile hydrogen compounds |
Volatile hydrogen compounds |
|||||
With oxygen, non-metals form acidic oxides. In some oxides, they exhibit a maximum oxidation state equal to the group number (for example, SO 2 , N 2 O 5 ), while in others, a lower one (for example, SO 2 , N 2 O 3 ). Acid oxides correspond to acids, and of the two oxygen acids of one non-metal, the one in which it exhibits a higher degree of oxidation is stronger. For example, nitric acid HNO 3 is stronger than nitrous HNO 2, and sulphuric acid H 2 SO 4 is stronger than sulfurous H 2 SO 3 .
Characteristics of oxygen compounds of non-metals
1. Properties of higher oxides (i.e. oxides, which include an element of this group with the highest degree oxidation) in periods from left to right gradually change from basic to acidic.
2. In groups from top to bottom, the acidic properties of higher oxides gradually weaken. This can be judged by the properties of the acids corresponding to these oxides.
3. The increase in the acidic properties of higher oxides of the corresponding elements in periods from left to right is explained by a gradual increase in the positive charge of the ions of these elements.
4. In the main subgroups of the periodic system of chemical elements in the direction from top to bottom, the acidic properties of higher oxides of non-metals decrease.