Chemical properties metals: interaction with oxygen, halogens, sulfur and relation to water, acids, salts.

The chemical properties of metals are due to the ability of their atoms to easily give up electrons from an external energy level, turning into positively charged ions. Thus, in chemical reactions, metals act as energetic reducing agents. This is their main common chemical property.

The ability to donate electrons in atoms of individual metallic elements is different. The more easily a metal gives up its electrons, the more active it is, and the more vigorously it reacts with other substances. Based on the research, all metals were arranged in a row according to their decreasing activity. This series was first proposed by the outstanding scientist N. N. Beketov. Such a series of activity of metals is also called the displacement series of metals or the electrochemical series of metal voltages. It looks like this:

Li, K, Ba, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, H2, Cu, Hg, Ag, Рt, Au

Using this series, you can find out which metal is the active of the other. This series contains hydrogen, which is not a metal. Its visible properties are taken for comparison as a kind of zero.

Having the properties of reducing agents, metals react with various oxidizing agents, primarily with non-metals. Metals react with oxygen under normal conditions or when heated to form oxides, for example:

2Mg0 + O02 = 2Mg+2O-2

In this reaction, magnesium atoms are oxidized and oxygen atoms are reduced. The noble metals at the end of the row react with oxygen. Reactions with halogens actively occur, for example, the combustion of copper in chlorine:

Cu0 + Cl02 = Cu+2Cl-2

Reactions with sulfur most often occur when heated, for example:

Fe0 + S0 = Fe+2S-2

Active metals in the activity series of metals in Mg react with water to form alkalis and hydrogen:

2Na0 + 2H+2O → 2Na+OH + H02

Metals of medium activity from Al to H2 react with water under more severe conditions and form oxides and hydrogen:

Pb0 + H+2O Chemical properties of metals: interaction with oxygen Pb+2O + H02.

The ability of a metal to react with acids and salts in solution also depends on its position in the displacement series of metals. Metals to the left of hydrogen in the displacement series of metals usually displace (reduce) hydrogen from dilute acids, and metals to the right of hydrogen do not displace it. So, zinc and magnesium react with acid solutions, releasing hydrogen and forming salts, while copper does not react.

Mg0 + 2H+Cl → Mg+2Cl2 + H02

Zn0 + H+2SO4 → Zn+2SO4 + H02.

Metal atoms in these reactions are reducing agents, and hydrogen ions are oxidizing agents.

Metals react with salts in aqueous solutions. Active metals displace less active metals from the composition of salts. This can be determined from the activity series of metals. The reaction products are a new salt and a new metal. So, if an iron plate is immersed in a solution of copper (II) sulfate, after a while copper will stand out on it in the form of a red coating:

Fe0 + Cu+2SO4 → Fe+2SO4 + Cu0 .

But if a silver plate is immersed in a solution of copper (II) sulfate, then no reaction will occur:

Ag + CuSO4 ≠ .

To carry out such reactions, one should not take too active metals (from lithium to sodium), which are capable of reacting with water.

Therefore, metals are able to react with non-metals, water, acids and salts. In all these cases, the metals are oxidized and are reducing agents. To predict the course of chemical reactions involving metals, a displacement series of metals should be used.

Moscow State Industrial University

Faculty of Applied Mathematics and Technical Physics

Department of Chemistry

Laboratory work

Chemical properties of metals

Moscow 2012

Objective. Exploring properties s-, p-, d- metal elements (Mg, Al, Fe, Zn) and their compounds.

1. Theoretical part

All metals are reducing agents in terms of their chemical properties; they donate electrons during a chemical reaction. Metal atoms donate valence electrons relatively easily and become positively charged ions.

1.1. Interaction of metals with simple substances

When metals interact with simple substances, non-metals usually act as oxidizing agents. Metals react with non-metals to form binary compounds.

1. When interacting with oxygen metals form oxides:

2Mg + O 2 2MgO,

2Cu + O2 2CuO.

2. Metals react with halogens(F 2, Cl 2, Br 2, I 2) with the formation of salts of hydrohalic acids:

2Na + Br 2 \u003d 2NaBr,

Ba + Cl 2 \u003d BaCl 2,

2Fe + 3Cl 2 2FeCl3.

3. When metals interact with gray sulfides are formed (salts of hydrosulfide acid H 2 S):

4. C hydrogen active metals interact with the formation of metal hydrides, which are salt-like substances:

2Na + H2 2NaH,

Ca+H2 CaH2.

In metal hydrides, hydrogen has an oxidation state (-1).

Metals can also interact with other nonmetals: nitrogen, phosphorus, silicon, carbon to form nitrides, phosphides, silicides, and carbides, respectively. For example:

3Mg + N2 Mg3N2,

3Ca + 2P Ca 3 P 2 ,

2Mg + Si Mg 2 Si,

4Al + 3C Al 4 C 3 .

5. Metals can also interact with each other to form intermetallic compounds:

2Mg + Cu \u003d Mg 2 Cu,

2Na + Sb = Na 2 Sb.

Intermetallic compounds(or intermetallics) are the compounds formed between elements, which usually belong to metals.

1.2. Interaction of metals with water

The interaction of metals with water is a redox process in which the metal is a reducing agent and water acts as an oxidizing agent. The reaction proceeds according to the scheme:

Me + n H 2 O \u003d Me (OH) n + n/2H2.

Under normal conditions, alkali and alkaline earth metals interact with water to form soluble bases and hydrogen:

2Na + 2H 2 O \u003d 2NaOH + H 2,

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

Magnesium reacts with water when heated:

Mg + 2H 2 O Mg (OH) 2 + H 2.

Iron and some other active metals interact with hot water vapor:

3Fe + 4H 2 O Fe 3 O 4 + 4H 2.

Metals with positive electrode potentials do not interact with water.

Do not interact with water 4 d-elements (except Cd), 5 d-elements and Cu (3 d-element).

1.3. The interaction of metals with acids

According to the nature of the action on metals, the most common acids can be divided into two groups.

1. Non-oxidizing acids: hydrochloric (hydrochloric, HCl), hydrobromic (HBr), hydroiodic (HI), hydrofluoric (HF), acetic (CH 3 COOH), dilute sulfuric (H 2 SO 4 (dil.)), dilute orthophosphoric (H 3 PO 4 (diff.)).

2. Oxidizing acids: nitric (HNO 3) in any concentration, concentrated sulfuric (H 2 SO 4 (conc.)), concentrated selenic (H 2 SeO 4 (conc.)).

Interaction of metals with non-oxidizing acids. Oxidation of metals by hydrogen ions H + in solutions of non-oxidizing acids occurs more vigorously than in water.

All metals that have a negative value of the standard electrode potential, i.e. which are in the electrochemical series of voltages up to hydrogen, displace hydrogen from non-oxidizing acids. The reaction proceeds according to the scheme:

Me+ n H+=Me n + + n/2H2.

For example:

2Al + 6HCl \u003d 2AlCl 3 + 3H 2,

Mg + 2CH 3 COOH \u003d Mg (CH 3 COO) 2 + H 2,

2Ti + 6HCl \u003d 2TiCl 3 + 3H 2.

Metals with a variable oxidation state (Fe, Co, Ni, etc.) form ions in their lowest oxidation state (Fe 2+, Co 2+, Ni 2+ and others):

Fe + H 2 SO 4 (razb) \u003d FeSO 4 + H 2.

When some metals interact with non-oxidizing acids: HCl, HF, H 2 SO 4 (diff.), HCN, insoluble products are formed that protect the metal from further oxidation. Thus, the surface of lead in HCl (diff) and H 2 SO 4 (diff) is passivated by poorly soluble salts PbCl 2 and PbSO 4, respectively.

Interaction of metals with oxidizing acids. Sulfuric acid in a dilute solution is a weak oxidizing agent, but in a concentrated solution it is a very strong one. The oxidizing ability of concentrated sulfuric acid H 2 SO 4 (conc.) is determined by the anion SO 4 2 , the oxidizing potential of which is much higher than that of the H + ion. Concentrated sulfuric acid is a strong oxidizing agent due to the sulfur atoms in the oxidation state (+6). In addition, a concentrated solution of H 2 SO 4 contains few H + ions, since it is weakly ionized in a concentrated solution. Therefore, when metals interact with H 2 SO 4 (conc.), hydrogen is not released.

Reacting with metals as an oxidizing agent, H 2 SO 4 (conc.) Most often passes into sulfur oxide (IV) (SO 2), and when interacting with strong reducing agents - into S or H 2 S:

Me + H 2 SO 4 (conc)  Me 2 (SO 4) n + H 2 O + SO 2 (S, H 2 S).

For ease of remembering, consider the electrochemical series of voltages, which looks like this:

Li, Rb, K, Cs, Ba, Sr, Ca, Na, Mg, Be, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, (H), Cu, Hg, Ag, Pt, Au.

In table. 1. shows the products of the reduction of concentrated sulfuric acid when interacting with metals of various activity.

Table 1.

Products of the interaction of metals with concentrated

sulfuric acid

Cu + 2H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O,

4Mg + 5H 2 SO 4 (conc) = 4MgSO 4 + H 2 S + 4H 2 O.

For metals of medium activity (Mn, Cr, Zn, Fe), the ratio of reduction products depends on the acid concentration.

The general trend is: the higher the concentration H2SO4, the deeper the recovery goes.

This means that formally each sulfur atom from H 2 SO 4 molecules can take not only two electrons from the metal (and go to ), but also six electrons (and go to) and even eight (and go to ):

Zn + 2H 2 SO 4 (conc) = ZnSO 4 + SO 2 + 2H 2 O,

3Zn + 4H 2 SO 4 (conc) = 3ZnSO 4 + S + 4H 2 O,

4Zn + 5H 2 SO 4 (conc) = 4ZnSO 4 + H 2 S + 4H 2 O.

Lead with concentrated sulfuric acid interacts with the formation of soluble lead (II) hydrosulfate, sulfur oxide (IV) and water:

Pb + 3H 2 SO 4 \u003d Pb (HSO 4) 2 + SO 2 + 2H 2 O.

Cold H 2 SO 4 (conc) passivates some metals (for example, iron, chromium, aluminum), which makes it possible to transport acid in steel containers. With strong heating, concentrated sulfuric acid interacts with these metals:

2Fe + 6H 2 SO 4 (conc) Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O.

Interaction of metals with nitric acid. The oxidizing ability of nitric acid is determined by the NO 3 - anion, the oxidizing potential of which is much higher than that of H + ions. Therefore, when metals interact with HNO 3, hydrogen is not released. Nitrate ion NO 3 , which has in its composition nitrogen in the oxidation state (+ 5), depending on the conditions (acid concentration, nature of the reducing agent, temperature), can accept from one to eight electrons. Reduction of the anion NO 3  can proceed with the formation of various substances according to the following schemes:

NO 3  + 2H + + e \u003d NO 2 + H 2 O,

NO 3  + 4H + + 3e \u003d NO + 2H 2 O,

2NO 3  + 10H + + 8e = N 2 O + 5H 2 O,

2NO 3  + 12H + + 10e = N 2 + 6H 2 O,

NO 3  + 10H + + 8e = NH 4 + + 3H 2 O.

Nitric acid has an oxidizing power at any concentration. Other things being equal, the following tendencies appear: the more active the metal that reacts with the acid, and the lower the concentration of the nitric acid solution,the more deeply it recovers.

This can be explained by the following diagram:

, ,
,
,

Acid concentration

metal activity

Oxidation of substances with nitric acid is accompanied by the formation of a mixture of products of its reduction (NO 2, NO, N 2 O, N 2, NH 4 +), the composition of which is determined by the nature of the reducing agent, the temperature and concentration of the acid. Oxides NO 2 and NO predominate among the products. Moreover, when interacting with a concentrated solution of HNO 3, NO 2 is more often released, and with a dilute solution - NO.

The equations of redox reactions involving HNO 3 are compiled conditionally, with the inclusion of only one reduction product, which is formed in a larger amount:

Me + HNO 3  Me (NO 3) n + H 2 O + NO 2 (NO, N 2 O, N 2, NH 4 +).

For example, in a gas mixture formed by the action of zinc on a sufficiently active metal (
= - 0.76 B) concentrated (68%) nitric acid, NO 2 prevails, 40% - NO; 20% - N 2 O; 6% - N 2. Very dilute (0.5%) nitric acid is reduced to ammonium ions:

Zn + 4HNO 3 (conc.) \u003d Zn (NO 3) 2 + 2NO 2 + 2H 2 O,

3Zn + 8HNO 3 (40%) = 3Zn(NO 3) 2 + 2NO + 4H 2 O,

4Zn + 10HNO 3 (20%) = 4Zn(NO 3) 2 + N 2 O + 5H 2 O,

5Zn + 12HNO 3 (6%) = 5Zn(NO 3) 2 + N 2 + 6H 2 O,

4Zn + 10HNO 3 (0.5%) = 4Zn(NO 3) 2 + NH 4 NO 3 + 3H 2 O.

With inactive metal copper (
= + 0.34B) reactions proceed according to the following schemes:

Cu + 4HNO 3 (conc) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O,

3Cu + 8HNO 3 (razb) \u003d 3 Cu (NO 3) 2 + 2NO + 4H 2 O.

Almost all metals are dissolved in concentrated HNO 3, except for Au, Ir, Pt, Rh, Ta, W, Zr. And metals such as Al, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb, Th, U, as well as stainless steels, are passivated with acid to form stable oxide films that adhere tightly to the metal surface and protect it from further oxidation. However, Al and Fe begin to dissolve when heated, and Cr is resistant to even hot HNO 3:

Fe + 6HNO 3 Fe(NO 3) 3 + 3NO 2 + 3H 2 O.

Metals, which are characterized by high oxidation states (+6, +7, +8), form oxygen-containing acids with concentrated nitric acid. In this case, HNO 3 is reduced to NO, for example:

3Re + 7HNO 3 (conc) = 3HReO 4 + 7NO + 2H 2 O.

Very dilute HNO 3 already lacks HNO 3 molecules, only H + and NO 3 - ions exist. Therefore, a very dilute acid (~ 3-5%) interacts with Al and does not transfer Cu and other low-active metals into solution:

8Al + 30HNO 3 (very dilute) = 8Al(NO 3) 3 + 3NH 4 NO 3 + 9H 2 O.

A mixture of concentrated nitric and hydrochloric acids (1:3) is called aqua regia. It dissolves Au and platinum metals (Pd, Pt, Os, Ru). For example:

Au + HNO 3 (conc.) + 4HCl = H + NO + 2H 2 O.

These metals dissolve in HNO 3 and in the presence of other complexing agents, but the process is very slow.

Reaction equations for the ratio of metals:

a) to simple substances: oxygen, hydrogen, halogens, sulfur, nitrogen, carbon;
b) to complex substances: water, acids, alkalis, salts.

Metals include s-elements of groups I and II, all s-elements, p-elements of group III (except boron), as well as tin and lead (group IV), bismuth (group V) and polonium (group VI). Most metals have 1-3 electrons in their outer energy level. For atoms of d-elements inside the periods, from left to right, the d-sublevels of the pre-outer layer are filled.
The chemical properties of metals are due to the characteristic structure of their outer electron shells.

Within a period, with an increase in the charge of the nucleus, the radii of atoms with the same number of electron shells decrease. Alkali metal atoms have the largest radii. The smaller the atomic radius, the greater the ionization energy, and the larger the atomic radius, the lower the ionization energy. Since metal atoms have the largest atomic radii, they are characterized mainly by low values of ionization energy and electron affinity. Free metals exhibit exclusively reducing properties.

3) Metals form oxides, for example:

Only alkali and alkaline earth metals react with hydrogen, forming hydrides:

Metals react with halogens to form halides, with sulfur - sulfides, with nitrogen - nitrides, with carbon - carbides.

With an increase in the algebraic value of the standard electrode potential of the metal E 0 in a series of voltages, the ability of the metal to react with water decreases. So, iron reacts with water only at very high temperature:

Metals with a positive value of the standard electrode potential, that is, those standing after hydrogen in a series of voltages, do not react with water.

Typical reactions of metals with acids. Metals with a negative value of E 0 displace hydrogen from solutions of Hcl, H 2 S0 4, H 3 P0 4, etc.

A metal with a lower value of E 0 displaces a metal with great value E 0 from salt solutions:

The most important calcium compounds obtained in industry, their chemical properties and methods of preparation.

Calcium oxide CaO is called quicklime. It is obtained by roasting limestone CaCO 3 --> CaO + CO, at a temperature of 2000 ° C. Calcium oxide has the properties of a basic oxide:

a) reacts with water with the release of a large amount of heat:

CaO + H 2 0 \u003d Ca (OH) 2 (slaked lime).

b) reacts with acids to form salt and water:

CaO + 2HCl \u003d CaCl 2 + H 2 O

CaO + 2H + = Ca 2+ + H 2 O

c) reacts with acid oxides to form a salt:

CaO + C0 2 \u003d CaC0 3

Calcium hydroxide Ca(OH) 2 is used in the form of slaked lime, milk of lime and lime water.

Lime milk is a suspension formed by mixing excess slaked lime with water.

Lime water is a clear solution obtained by filtering milk of lime. Used in the laboratory to detect carbon monoxide (IV).

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

With prolonged transmission of carbon monoxide (IV), the solution becomes transparent, since an acid salt is formed that is soluble in water:

CaC0 3 + C0 2 + H 2 O \u003d Ca (HCO 3) 2

If the resulting transparent solution of calcium bicarbonate is heated, then turbidity occurs again, since CaCO 3 precipitates:

1. Metals react with non-metals.

2Me + n Hal 2 → 2 MeHal n

4Li + O2 = 2Li2O

Alkali metals, with the exception of lithium, form peroxides:

2Na + O 2 \u003d Na 2 O 2

2. Metals standing up to hydrogen react with acids (except 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 H2

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. Intermediate activity metals 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 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 standing after hydrogen do not react with water and acid solutions (except for nitric and sulfuric conc.)

6. More active metals displace less active ones from solutions of their salts.

CuSO 4 + Zn \u003d ZnSO 4 + Cu

CuSO 4 + Fe \u003d FeSO 4 + Cu

Active metals - zinc and iron replaced copper in sulfate and formed salts. Zinc and iron are oxidized, and copper is restored.

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

heating: 3Cl2+6KOH−→KClO3+5KCl+3H2O3Cl2+6KOH→t,∘CKClO3+5KCl+3H2O forms loride and chlorate

8 Active halogens (except 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 \u003d Na 2 O 2

3. Classification of inorganic compounds

Simple substances - substances whose molecules consist of atoms of the same type (atoms of the same element). In chemical reactions, they 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: solids (with the exception of mercury) have a metallic luster, are good conductors of heat and electricity, malleable (iron (Fe), copper (Cu), aluminum (Al), mercury (Hg), gold (Au), silver (Ag), etc.).

non-metals - a group of elements: solid, liquid (bromine) and gaseous substances that do not have a metallic sheen, are insulators, brittle.

And complex substances, in turn, are divided 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 - These are complex substances in which metal atoms are connected to 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 produces metal cations (or NH4 +) and hydroxide - anions OH-.

acids - these are complex substances whose molecules include hydrogen atoms that can be replaced or exchanged for metal atoms.

salt - These are complex substances, the molecules of which consist of metal atoms and acid residues. Salt is a product of partial or complete replacement of hydrogen atoms of an acid by a metal.