Metals occupy Periodic table lower left corner. Metals belong to the families of s-elements, d-elements, f-elements and partially p-elements.
The most typical property of metals is their ability to donate electrons and transform into positively charged ions. Moreover, metals can only show a positive oxidation state.
Me - ne = Me n +
1. Interaction of metals with non-metals.
a ) Interaction of metals with hydrogen.
Alkali and alkaline earth metals react directly with hydrogen to form hydrides.
For example:
Ca + H 2 = CaH 2
Non-stoichiometric compounds with an ionic crystal structure are formed.
b) Interaction of metals with oxygen.
All metals with the exception of Au, Ag, Pt are oxidized by atmospheric oxygen.
Example:
2Na + O 2 = Na 2 O 2 (peroxide)
4K + O 2 = 2K 2 O
2Mg + O 2 = 2MgO
2Cu + O 2 = 2CuO
c) Interaction of metals with halogens.
All metals react with halogens to form halides.
Example:
2Al + 3Br 2 = 2AlBr 3
These are mainly ionic compounds: MeHal n
d) Interaction of metals with nitrogen.
Alkali and alkaline earth metals interact with nitrogen.
Example:
3Ca + N 2 = Ca 3 N 2
Mg + N 2 = Mg 3 N 2 - nitride.
e) Interaction of metals with carbon.
Compounds of metals and carbon - carbides. They are formed by the interaction of melts with carbon. Active metals form stoichiometric compounds with carbon:
4Al + 3C = Al 4 C 3
Metals - d-elements form compounds of non-stoichiometric composition such as solid solutions: WC, ZnC, TiC - are used to obtain superhard steels.
2. Interaction of metals with water.
Metals react with water that have a more negative potential than the redox potential of water.
Active metals react more actively with water, decomposing water with the release of hydrogen.
Na + 2H 2 O = H 2 + 2NaOH
Less active metals slowly decompose water and the process is inhibited due to the formation of insoluble substances.
3. Interaction of metals with salt solutions.
Such a reaction is possible if the reacting metal is more active than that in the salt:
Zn + CuSO 4 = Cu 0 ↓ + ZnSO 4
0.76 B., = + 0.34 B.
A metal with a more negative or less positive standard electrode potential displaces another metal from its salt solution.
4. Interaction of metals with alkali solutions.
Metals that give amphoteric hydroxides or have high oxidation states in the presence of strong oxidants can interact with alkalis. When metals interact with alkali solutions, water is the oxidizing agent.
Example:
Zn + 2NaOH + 2H 2 O = Na 2 + H 2
1 Zn 0 + 4OH - - 2e = 2- oxidation
Zn 0 - reducing agent
1 2H 2 O + 2e = H 2 + 2OH - reduction
H 2 O - oxidizing agent
Zn + 4OH - + 2H 2 O = 2- + 2OH - + H 2
Metals with high oxidation states can interact with alkalis during fusion:
4Nb + 5O 2 + 12KOH = 4K 3 NbO 4 + 6H 2 O
5. Interaction of metals with acids.
it complex reactions, the interaction products depend on the activity of the metal, on the type and concentration of the acid, and on the temperature.
According to their activity, metals are conventionally divided into active, medium-active and low-activity.
Acids are conventionally divided into 2 groups:
Group I - acids with low oxidizing ability: HCl, HI, HBr, H 2 SO 4 (dil.), H 3 PO 4, H 2 S, the oxidizing agent here is H +. When interacting with metals, oxygen (H 2) is released. Metals with a negative electrode potential react with acids of the first group.
Group II - acids with high oxidizing ability: H 2 SO 4 (conc.), HNO 3 (diluted), HNO 3 (conc.). In these acids, acid anions are oxidizing agents:. Anion reduction products can be very diverse and depend on the activity of the metal.
H 2 S - with active metals
H 2 SO 4 + 6е S 0 ↓ - with metals of medium activity
SO 2 - with low-activity metals
NH 3 (NH 4 NO 3) - with active metals
HNO 3 + 4,5e N 2 O, N 2 - with metals of medium activity
NO - with low active metals
HNO 3 (conc.) - NO 2 - with metals of any activity.
If metals have variable valence, then with Group I acids the metals acquire the lowest positive oxidation state: Fe → Fe 2+, Cr → Cr 2+. When interacting with group II acids, the oxidation state is +3: Fe → Fe 3+, Cr → Cr 3+, while hydrogen is never released.
Some metals (Fe, Cr, Al, Ti, Ni, etc.) in solutions strong acids being oxidized, they are covered with a dense oxide film, which protects the metal from further dissolution (passivation), but when heated, the oxide film dissolves, and the reaction proceeds.
Poorly soluble metals with a positive electrode potential can dissolve in Group I acids in the presence of strong oxidants.
1. Metals react with non-metals.
2 Me + n Hal 2 → 2 MeHal n
4Li + O2 = 2Li2O
Alkali metals, with the exception of lithium, form peroxides:
2Na + O 2 = Na 2 O 2
2. Metals standing up to hydrogen react with acids (except for nitric and sulfuric conc.) With the release of hydrogen
Me + HCl → salt + H2
2 Al + 6 HCl → 2 AlCl3 + 3 H2
Pb + 2 HCl → PbCl2 ↓ + H2
3. Active metals react with water to form alkali and release hydrogen.
2Me + 2n H 2 O → 2Me (OH) n + n H 2
The product of metal oxidation is its hydroxide - Me (OH) n (where n is the oxidation state of the metal).
For example:
Ca + 2H 2 O → Ca (OH) 2 + H 2
4. Metals of medium activity react with water when heated to form metal oxide and hydrogen.
2Me + nH 2 O → Me 2 O n + nH 2
The oxidation product in such reactions is the metal oxide Me 2 O n (where n is the oxidation state of the metal).
3Fe + 4H 2 O → Fe 2 O 3 FeO + 4H 2
5. Metals behind hydrogen do not react with water and acid solutions (except for nitric and sulfuric conc.)
6. More active metals displace less active metals from solutions of their salts.
CuSO 4 + Zn = Zn SO 4 + Cu
CuSO 4 + Fe = Fe SO 4 + Cu
Active metals - zinc and iron replaced copper in sulfate and formed salts. Zinc and iron were oxidized, and copper was reduced.
7. Halogens react with water and alkali solution.
Fluorine, unlike other halogens, oxidizes water:
2H 2 O + 2F 2 = 4HF + O 2 .
in the cold: Cl2 + 2KOH = KClO + KCl + H2OCl2 + 2KOH = KClO + KCl + H2O chloride and hypochlorite are formed
when heated: 3Cl2 + 6KOH− → KClO3 + 5KCl + 3H2O3Cl2 + 6KOH → t, ∘CKClO3 + 5KCl + 3H2O loride and chlorate are formed
8 Active halogens (except for fluorine) displace less active halogens from solutions of their salts.
9. Halogens do not react with oxygen.
10. Amphoteric metals (Al, Be, Zn) react with solutions of alkalis and acids.
3Zn + 4H2SO4 = 3 ZnSO4 + S + 4H2O
11. Magnesium reacts with carbon dioxide and silicon oxide.
2Mg + CO2 = C + 2MgO
SiO2 + 2Mg = Si + 2MgO
12. Alkali metals (except lithium) form peroxides with oxygen.
2Na + O 2 = Na 2 O 2
3. Classification of inorganic compounds
Simple substances - substances whose molecules consist of atoms of one type (atoms of one element). V chemical reactions cannot decompose to form other substances.
Complex substances (or chemical compounds) - substances whose molecules consist of atoms of different types (atoms of various chemical elements). In chemical reactions, they decompose to form several other substances.
Simple substances are divided into two large groups: metals and non-metals.
Metals - a group of elements with characteristic metallic properties: solid substances (with the exception of mercury) have a metallic luster, are good guides heat and electricity, malleable (iron (Fe), copper (Cu), aluminum (Al), mercury (Hg), gold (Au), silver (Ag), etc.).
Nonmetals - a group of elements: solid, liquid (bromine) and gaseous substances that do not have a metallic luster are insulators that are brittle.
A complex substances in turn are subdivided into four groups, or classes: oxides, bases, acids and salts.
Oxides - these are complex substances, the composition of the molecules of which includes atoms of oxygen and some other substance.
Foundations Are complex substances in which metal atoms are combined with one or more hydroxyl groups.
From the point of view of the theory of electrolytic dissociation, bases are complex substances, the dissociation of which in an aqueous solution forms metal cations (or NH4 +) and hydroxide - OH- anions.
Acids - These are complex substances, the molecules of which include hydrogen atoms that can be replaced or exchanged for metal atoms.
Salt Are complex substances, the molecules of which are composed of metal atoms and acidic residues. Salt is a product of partial or complete substitution of a metal for hydrogen atoms of an acid.
Metal atoms relatively easily donate valence electrons and transform into positively charged ions. Therefore, metals are reducing agents. Metals interact with simple substances: Ca + C12 - CaC12, Active metals react with water: 2Na + 2H20 = 2NaOH + H2f. Metals in the series of standard electrode potentials up to hydrogen interact with dilute acid solutions (except for HNO3) with the evolution of hydrogen: Zn + 2HC1 = ZnCl2 + H2f. Metals react with aqueous solutions of salts of less active metals: Ni + CuSO4 = NiS04 + Cu J. Metals react with oxidizing acids: C. Methods of obtaining metals Modern metallurgy receives more than 75 metals and numerous alloys based on them. Depending on the methods of obtaining metals, pyrohydrometallurgy and electrometallurgy are distinguished. GG) Pyrometallurgy covers methods of obtaining metals from ores using reduction reactions held at high temperatures... Coal, active metals, carbon monoxide (II), hydrogen, methane are used as reducing agents. Cu20 + C - 2Cu + CO, t ° Cu20 + CO - 2Cu + C02, t ° Cg203 + 2A1 - 2Cg + A1203, (alumothermy) t ° TiCl2 + 2Mg - Ti + 2MgCl2, (magnesium heat) t ° W03 + 3H2 = W + 3H20. (hydrogenothermy) | C Hydrometallurgy is the production of metals from solutions of their salts. For example, when treating copper ore containing copper (I) oxide with dilute sulfuric acid, copper goes into solution in the form of sulfate: CuO + H2SO4 = CuSO4 + H20. Then copper is removed from the solution either by electrolysis or by displacement with iron powder: CuSO4 + Fe = FeSO4 + Cu. [h] Electrometallurgy is a method of obtaining metals from their molten oxides or salts using electrolysis: electrolysis 2NaCl - 2Na + Cl2. Questions and tasks for independent solution 1. Indicate the position of metals in periodic system D.I. Mendeleev. 2. Show the physical and chemical properties of metals. 3. Explain the reason for the generality of the properties of metals. 4. Show the change in the chemical activity of metals of the main subgroups I and II of the periodic system. 5. How do the metallic properties of the elements of the II and III periods change? Name the most refractory and the lowest melting metals. 7. Indicate which metals occur naturally in the native state and which - only in the form of compounds. How can this be explained? 8. What is the nature of the alloys? How the composition of the alloy affects its properties. Show with specific examples. Please indicate the most important ways obtaining metals from ores. 10l Name the types of pyrometallurgy. What reducing agents are used in each specific method? Why? 11. Name the metals that are obtained using hydrometallurgy. What is the essence and what are the advantages this method in front of others? 12. Give examples of obtaining metals using electrometallurgy. In what case is this method used? 13. What are modern ways obtaining metals high degree purity? 14. What is "electrode potential"? Which of the metals has the highest and which is the lowest electrode potential in an aqueous solution? 15. Describe a number of standard electrode potentials? 16. Is it possible to displace metallic iron from an aqueous solution of its sulfate using metallic zinc, nickel, sodium? Why? 17. What is the working principle of galvanic cells? What metals can be used in them? 18. What processes are corrosive? What types of corrosion are you aware of? 19. What is called electrochemical corrosion? What methods of protection against it do you know? 20. How does its contact with other metals affect the corrosion of iron? Which metal will break first on the damaged surface of tinned, galvanized and nickel-plated iron? 21. What process is called electrolysis? Write the reactions reflecting the processes occurring at the cathode and anode during the electrolysis of sodium chloride melt, aqueous solutions sodium chloride, copper sulfate, sodium sulfate, sulfuric acid. 22. What is the role of the electrode material in the course of electrolysis processes? Give examples of electrolysis processes with soluble and insoluble electrodes. 23. The alloy used to make copper coins contains 95% copper. Determine the second metal included in the alloy if there is an excess of of hydrochloric acid 62.2 ml of hydrogen (n.u.) was released. aluminum. 24. A sample of metal carbide weighing 6 g is burned in oxygen. This formed 2.24 liters of carbon monoxide (IV) (n.u.). Determine which metal was included in the carbide. 25. Show what products will be released during the electrolysis of an aqueous solution of nickel sulfate, if the process proceeds: a) with coal; b) with nickel electrodes? 26. When electrolysis of an aqueous solution copper sulfate 2.8 liters of gas (n.u.) were released at the anode. What kind of gas is it? What and in what quantity was released at the cathode? 27. Make a diagram of the electrolysis of an aqueous solution of potassium nitrate flowing on the electrodes. What is the amount of passed electricity if 280 ml of gas (n.u.) was released at the anode? What and in what quantity was released at the cathode?
The structure of metal atoms determines not only the characteristic physical properties simple substances- metals, but also their general chemical properties.
With a large variety, all chemical reactions of metals are redox reactions and can be of only two types: compounds and substitutions. Metals are capable of donating electrons during chemical reactions, that is, being reducing agents, showing only a positive oxidation state in the resulting compounds.
V general view this can be expressed by the scheme:
Ме 0 - ne → Me + n,
where Me is a metal - a simple substance, and Me 0 + n is a metal chemical element in conjunction.
Metals are able to donate their valence electrons to atoms of non-metals, hydrogen ions, ions of other metals, and therefore will react with non-metals - simple substances, water, acids, salts. However, the reducing ability of metals is different. The composition of the reaction products of metals with various substances also depends on the oxidizing ability of substances and the conditions under which the reaction proceeds.
At high temperatures, most metals burn out in oxygen:
2Mg + O 2 = 2MgO
Only gold, silver, platinum and some other metals are not oxidized under these conditions.
Many metals react with halogens without heating. For example, aluminum powder, when mixed with bromine, ignites:
2Al + 3Br 2 = 2AlBr 3
When metals interact with water, hydroxides are formed in some cases. Under normal conditions, alkali metals, as well as calcium, strontium, barium, interact very actively with water. The scheme of this reaction in general looks like this:
Ме + HOH → Me (OH) n + H 2
Other metals react with water when heated: magnesium when it boils, iron in water vapor when it boils red. In these cases, metal oxides are obtained.
If the metal reacts with an acid, then it is part of the resulting salt. When the metal interacts with acid solutions, it can be oxidized by the hydrogen ions present in this solution. Abbreviated ionic equation in general, it can be written like this:
Me + nH + → Me n + + H 2
Anions of oxygen-containing acids such as concentrated sulfuric and nitric acids have stronger oxidizing properties than hydrogen ions. Therefore, those metals that are not capable of being oxidized by hydrogen ions react with these acids, for example, copper and silver.
When metals interact with salts, a substitution reaction occurs: electrons from the atoms of the substituting - more active metal pass to the ions of the substituted - less active metal. Then the network is the replacement of the metal with the metal in the salts. These reactions are not reversible: if metal A displaces metal B from the salt solution, then metal B will not displace metal A from the salt solution.
In decreasing order of chemical activity manifested in the reactions of displacing metals from each other from aqueous solutions of their salts, metals are located in the electrochemical series of voltages (activities) of metals:
Li → Rb → K → Ba → Sr → Ca → Na → Mg → Al → Mn → Zn → Cr → → Fe → Cd → Co → Ni → Sn → Pb → H → Sb → Bi → Cu → Hg → Ag → Pd → Pt → Au
The metals located to the left in this row are more active and are able to displace the following metals from salt solutions.
The electrochemical series of voltages of metals includes hydrogen, as the only non-metal that separates with metals common property- to form positively charged ions. Therefore, hydrogen replaces some metals in their salts and itself can be replaced by many metals in acids, for example:
Zn + 2 HCl = ZnCl 2 + H 2 + Q
Metals in the electrochemical series of voltages up to hydrogen displace it from solutions of many acids (hydrochloric, sulfuric, etc.), and all those following it, for example, copper, do not displace.
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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, be drawn into wire, rolled 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 directional 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 increasing temperature, the vibrations of atoms and ions in the sites crystal lattice, which complicates the directional movement of the "electron gas".
4) Thermal conductivity. It is caused by the high mobility of free electrons, due to which there is a rapid equalization of temperature over the mass of the metal. Bismuth and mercury have the highest thermal conductivity.
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 the smaller, the less the atomic mass of the metal and the greater the radius of the atom. The lightest is lithium (ρ = 0.53 g / cm3); the heaviest is osmium (ρ = 22.6 g / cm3). Metals with a density of less than 5 g / cm3 are considered “light metals”.
7) Melting and boiling points. The lowest-melting metal is mercury (mp = -39 ° C), the most refractory metal- tungsten (t ° pl. = 3390 ° C). Metals with t ° pl. above 1000 ° C are considered refractory, below - low melting.
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.
1. Reactions of metals with non-metals
1) With oxygen:
2Mg + O 2 → 2MgO
2) With gray:
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
2. Reactions of metals with acids
1) Metals 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:
With the interaction of nitric acid of any concentration and concentrated sulfuric with metals hydrogen is never released!
Zn + 2H 2 SO 4 (К) → ZnSO 4 + SO 2 + 2H 2 O
4Zn + 5H 2 SO 4 (К) → 4ZnSO 4 + H 2 S + 4H 2 O
3Zn + 4H 2 SO 4 (К) → 3ZnSO 4 + S + 4H 2 O
2H 2 SO 4 (k) + 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) + Cu → Cu (NO 3) 2 + 2NO 2 + 2H 2 O
3. 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.
4. 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 copper-zinc alloys ( brass) are already quite solid and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but 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 aircraft construction. Alloys of iron with carbon (and additives of other metals) are widely known cast iron and steel.
Free metals are reducing agents. However, the reactivity of some metals is low due to the fact that they are coated surface oxide film, v varying degrees resistant to the action of chemicals such as water, solutions of acids and alkalis.
For example, lead is always covered with an oxide film; for its transition into solution, not only the action of a reagent (for example, dilute nitric acid) is required, but also heating. The oxide film on aluminum prevents it from reacting with water, but is destroyed by acids and alkalis. Loose oxide film (rust), formed on the surface of iron in humid air, does not interfere with further oxidation of iron.
Under the influence concentrated acids on metals are formed steady oxide film. This phenomenon is called passivation... So, in concentrated sulfuric acid metals such as Be, Bi, Co, Fe, Mg and Nb are passivated (and then do not react with acid), and in concentrated nitric acid - metals A1, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb , Th and U.
When interacting with oxidants in acidic solutions, most metals are converted into cations, the charge of which is determined by the stable oxidation state of this 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 voltages. 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
An undesirable chemical property of metals is their corrosion, i.e., active destruction (oxidation) upon contact with water and under the influence of oxygen dissolved in it. (oxygen corrosion). For example, corrosion of iron products in water is widely known, as a result of which rust is formed and the products are crumbled into powder.
Corrosion of metals occurs in water also due to the presence of dissolved gases CO 2 and SO 2; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 ( hydrogen corrosion).
The place of contact of two dissimilar metals ( contact corrosion). A galvanic pair arises between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the series of voltages (Pe), to the less active metal (Sn, Cu), and the more active metal is destroyed (corroded).
It is because of this that the tinned surface rusts. cans(tin-plated iron) when stored in a humid atmosphere and carelessly handling them (iron is quickly destroyed after the appearance of even a small scratch that allows the iron to come into contact with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, because even in the presence of 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, plating iron with chromium or making an iron-chromium alloy eliminates iron corrosion. Chromium-plated iron and steel containing chromium ( stainless steel ), have high corrosion resistance.
