Chemical properties of metals with examples. General characteristics of metals
The structure of metal atoms determines not only the characteristic physical properties of 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 chemical reactions donate electrons, that is, to be reducing agents, to show only a positive oxidation state in the resulting compounds.
In general terms, this can be expressed by the following scheme:
Ме 0 - ne → Me + n,
where Me is a metal - a simple substance, and Me 0 + n is a metal - a chemical element in a compound.
Metals are able to donate their valence electrons to non-metal atoms, 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 the substances and the conditions under which the reaction proceeds.
At high temperatures, most metals burn 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. The abbreviated ionic equation in general form can be written as follows:
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 react with these acids that are not capable of being oxidized by hydrogen ions, 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 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.
Hydrogen is included in the electrochemical series of voltages of metals, as the only non-metal that shares a common property with metals - 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, do not displace copper.
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Metals occupy the lower left corner of the Periodic Table. 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.
These are complex reactions, the products of interaction 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-active 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 of strong acids, being oxidized, become 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.
The first material that people have learned to use for their needs is stone. However, later, when a person became aware of the properties of metals, the stone moved far back. It is these substances and their alloys that have become the most important and main material in the hands of people. They were used to make household items, tools, premises were built. Therefore, in this article we will consider what metals are, the general characteristics, properties and application of which are so relevant to this day. After all, literally immediately after the Stone Age, a whole galaxy of metal followed: copper, bronze and iron.
Metals: general characteristics
What unites all representatives of these simple substances? Of course, this is the structure of their crystal lattice, types of chemical bonds and features of the electronic structure of the atom. After all, hence the characteristic physical properties that underlie the use of these materials by humans.
First of all, consider metals as chemical elements the periodic system. In it, they are located quite freely, occupying 95 cells out of 115 known today. There are several features of their location in the general system:
- Form the main subgroups of groups I and II, as well as III, starting with aluminum.
- All side subgroups are composed only of metals.
- They are located below the conventional diagonal from boron to astatine.
Based on such data, it is easy to trace that non-metals are collected in the upper right side of the system, and all the rest of the space belongs to the elements we are considering.
All of them have several features of the electronic structure of the atom:
general characteristics metals and non-metals allows you to identify patterns in their structure. So, the crystal lattice of the former is metallic, special. In its nodes there are several types of particles at once:
- ions;
- atoms;
- electrons.
A common cloud accumulates inside, called an electron gas, which explains all the physical properties of these substances. The type of chemical bond in metals is the same name with them.
Physical properties
There are a number of parameters that all metals have in common. General characteristics of their physical properties looks like that.
The listed parameters are the general characteristics of metals, that is, everything that unites them into one large family. However, it should be understood that there are exceptions to every rule. Moreover, there are too many elements of this kind. Therefore, within the family itself there are also subdivisions into various groups, which we will consider below and for which we will indicate the characteristic features.
Chemical properties
From the point of view of the science of chemistry, all metals are reducing agents. Moreover, they are very strong. The fewer electrons at the outer level and the larger the atomic radius, the stronger the metal in this parameter.
As a result, metals are able to react with:
This is just a general overview of the chemical properties. Indeed, for each group of elements, they are purely individual.
Alkaline earth metals
The general characteristics of alkaline earth metals are as follows:
Thus, alkaline earth metals are common elements of the s-family, exhibiting high chemical activity and are strong reducing agents and important participants in biological processes in the body.
Alkali metals
The general description begins with their name. They received it for the ability to dissolve in water, forming alkalis - caustic hydroxides. Reactions with water are very violent, sometimes with inflammation. These substances do not occur in free form in nature, since their chemical activity is too high. They react with air, water vapor, non-metals, acids, oxides and salts, that is, with almost everything.
This is due to their electronic structure. On the outer level, there is only one electron, which they easily donate. These are the strongest reducing agents, which is why it took a long time to get them in their pure form. This was first done by Humphrey Davy as early as the 18th century by electrolysis of sodium hydroxide. Now all representatives of this group are mined using this method.
The general characteristic of alkali metals also lies in the fact that they constitute the first group of the main subgroup of the periodic system. All of them are important elements that form many valuable natural compounds used by humans.
General characteristics of metals of the d- and f-families
This group of elements includes all those whose oxidation state can vary. This means that, depending on the conditions, the metal can act as both an oxidizing agent and a reducing agent. Such elements have a great ability to react. Among them are a large number of amphoteric substances.
The common name for all these atoms is transition elements. They received it for the fact that, in terms of the manifested properties, they really stand in the middle, as it were, between typical metals of the s-family and non-metals of the p-family.
The general characteristic of transition metals implies the designation of their similar properties. They are as follows:
- a large number of electrons at the outer level;
- large atomic radius;
- several oxidation states (from +3 to +7);
- are on the d- or f-sublevel;
- form 4-6 large periods of the system.
As simple substances, metals of this group are very strong, ductile and malleable, therefore they are of great industrial importance.
Side subgroups of the periodic system
The general characteristics of the metals of the secondary subgroups completely coincide with that of the transitional ones. And this is not surprising, because, in fact, they are exactly the same thing. It's just that the side subgroups of the system are formed precisely by representatives of the d- and f-families, that is, transition metals. Therefore, we can say that these concepts are synonyms.
The most active and important of them is the first row of 10 representatives from scandium to zinc. All of them are of great industrial importance and are often used by humans, especially for smelting.
Alloys
The general characteristics of metals and alloys makes it possible to understand where and how it is possible to use these substances. Such compounds have undergone great transformations in the past decades, because more and more additives are being discovered and synthesized to improve their quality.
The most famous alloys today are:
- brass;
- duralumin;
- cast iron;
- steel;
- bronze;
- will win;
- nichrome and others.
What is an alloy? This is a mixture of metals obtained by melting the latter in special furnace devices. This is done in order to obtain a product that is superior in properties to the pure substances that form it.
Comparison of the properties of metals and non-metals
If we talk about general properties, then the characteristics of metals and non-metals will differ in one very important point: for the latter, it is impossible to distinguish similar features, since they are very different in terms of the manifested properties, both physical and chemical.
Therefore, it is impossible to create such a characteristic for non-metals. You can only consider separately the representatives of each group and describe their properties.
The 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 for boron), as well as tin and lead (group IV), bismuth (group V) and polonium (group VI). Most metals have 1-3 electrons at the external energy level. In d-element atoms, 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 the period, with an increase in the nuclear charge, the radii of the atoms with the same number of electron shells decrease. Atoms of alkali metals have the largest radii. The smaller the radius of the atom, the greater the ionization energy, and the larger the radius of the atom, the lower the ionization energy. Since metal atoms have the largest radii of atoms, they are characterized mainly by low values of ionization energy and electron affinity. Free metals exhibit extremely reducing properties.
3) Metals form oxides, for example:
Only alkali and alkaline earth metals react with hydrogen, forming hydrides:
Metals react with halogens, forming 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 the series of voltages, the ability of the metal to react with water decreases. So, iron reacts with water only when very high temperature:
Metals with a positive value of the standard electrode potential, that is, standing after hydrogen in the series of voltages, do not react with water.
Reactions of metals with acids are characteristic. 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 E 0 value displaces a metal with a large E 0 value from salt solutions:
The most important calcium compounds obtained in industry, their chemical properties and methods of production.
Calcium oxide CaO is called quicklime. It is obtained by burning limestone CaCO 3 -> CaO + CO, at a temperature of 2000 ° C. Calcium oxide has the properties of a basic oxide:
a) reacts with water to release a large number warmth:
CaO + H 2 0 = Ca (OH) 2 (slaked lime).
b) reacts with acids, forming salt and water:
CaO + 2HCl = CaCl 2 + H 2 O
CaO + 2H + = Ca 2+ + H 2 O
c) reacts with acidic oxides to form a salt:
CaO + C0 2 = CaCO 3
Calcium hydroxide Ca (OH) 2 is used in the form of slaked lime, milk of lime and lime water.
Lime milk is a slurry formed by mixing excess hydrated 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 = CaCO 3 + H 2 O
With prolonged transmission of carbon monoxide (IV), the solution becomes transparent, since an acidic salt is formed, soluble in water:
CaCO 3 + CO 2 + H 2 O = Ca (HCO 3) 2
If the resulting transparent solution of calcium bicarbonate is heated, then turbidity occurs again, since CaCO 3 precipitates.
CHEMICAL PROPERTIES OF METALS
By chemical properties metals are subdivided into:1 ) Active (alkali and alkaline earth metals, Mg, Al, Zn, etc.)
2) Metalsaverage activity (Fe, Cr, Mn, etc.);
3 ) Inactive (Cu, Ag)
4) Noble metals - Au, Pt, Pd, etc.
The reactions contain only reducing agents. Metal atoms easily donate electrons of the outer (and some of the pre-outer) electron layer, turning into positive ions. Possible oxidation states of Ме Low 0, + 1, + 2, + 3 High + 4, + 5, + 6, + 7, + 8
1. INTERACTION WITH NON-METALS
1.WITH HYDROGEN
When heated, metals of groups IA and IIA, except beryllium, react. Solid unstable hydrides are formed, the rest of the metals do not react.
2K + H₂ = 2KH (potassium hydride)
Ca + H₂ = CaH₂
2.With oxygen
All metals react, except for gold and platinum. The reaction with silver occurs at high temperatures, but silver (II) oxide is practically not formed, since it is thermally unstable. Alkali metals under normal conditions form oxides, peroxides, superoxides (lithium - oxide, sodium - peroxide, potassium, cesium, rubidium - superoxide
4Li + O2 = 2Li2O (oxide)
2Na + O2 = Na2O2 (peroxide)
K + O2 = KO2 (superoxide)
The remaining metals of the main subgroups under normal conditions form oxides with an oxidation state equal to the group number 2Ca + O2 = 2CaO
2Сa + O2 = 2СaO
Metals of side subgroups form oxides under normal conditions and when heated, oxides of different oxidation states, and iron iron oxide Fe3O4 (Fe⁺²O ∙ Fe2⁺³O3)
3Fe + 2O2 = Fe3O4
4Cu + O₂ = 2Cu₂⁺¹O (red) 2Cu + O₂ = 2Cu⁺²O (black);
2Zn + O₂ = ZnO 4Cr + 3О2 = 2Cr2О3
3.WITH HALOGENS
halides (fluorides, chlorides, bromides, iodides). Alkaline under normal conditions with F, Cl, Br ignite:
2Na + Cl2 = 2NaCl (chloride)
Alkaline earth and aluminum react under normal conditions:
WITHa + Cl2 =WITHaCl2
2Al + 3Cl2 = 2AlCl3
Side subgroup metals at elevated temperatures
Cu + Cl₂ = Cu⁺²Cl₂ Zn + Cl₂ = ZnCl₂
2Fe + ЗС12 = 2Fe⁺³Cl3 ferric chloride (+3) 2Cr + 3Br2 = 2Cr⁺³Br3
2Cu + I₂ = 2Cu⁺¹I(there is no copper iodide (+2)!)
4. INTERACTION WITH SULFUR
when heated even with alkali metals, with mercury under normal conditions. All metals react except gold and platinum
withgray – sulfides: 2K + S = K2S 2Li + S = Li2S (sulfide)
WITHa + S =WITHaS (sulfide) 2Al + 3S = Al2S3 Cu + S = Cu⁺²S (black)
Zn + S = ZnS 2Cr + 3S = Cr2⁺³S3 Fe + S = Fe⁺²S
5. INTERACTION WITH PHOSPHORUS AND NITROGEN
proceeds when heated (exception: lithium with nitrogen under normal conditions):
with phosphorus - phosphides: 3Ca + 2 P= Ca3P2,
With nitrogen - nitrides 6Li + N2 = 3Li2N (lithium nitride) (n.o.) 3Mg + N2 = Mg3N2 (magnesium nitride) 2Al + N2 = 2A1N 2Cr + N2 = 2CrN 3Fe + N2 = Fe₃⁺²N₂¯³
6. INTERACTION WITH CARBON AND SILICON
proceeds when heated:
Carbides are formed with carbon. Only the most active metals react with carbon. Of alkali metals, carbides form lithium and sodium, potassium, rubidium, cesium do not interact with carbon:
2Li + 2C = Li2C2, Ca + 2C = CaC2
Metals - d-elements form compounds of non-stoichiometric composition with carbon, such as solid solutions: WC, ZnC, TiC - are used to obtain superhard steels.
with silicon - silicides: 4Cs + Si = Cs4Si,
7. INTERACTION OF METALS WITH WATER:
Metals that stand up to hydrogen in the electrochemical series of voltages react with water Alkaline and alkaline earth metals react with water without heating, forming soluble hydroxides (alkalis) and hydrogen, aluminum (after the destruction of the oxide film - amalgamation), magnesium when heated, form insoluble bases and hydrogen ...
2Na + 2HOH = 2NaOH + H2
WITHa + 2HOH = Ca (OH) 2 + H2
2Аl + 6Н2O = 2Аl (OH) 3 + 3H2
The rest of the metals react with water only in a red-hot state, forming oxides (iron - iron scale)
Zn + H2O = ZnO + H2 3Fe + 4HOH = Fe3O4 + 4H2 2Cr + 3H₂O = Cr₂O₃ + 3H₂
8 WITH OXYGEN AND WATER
In air, iron and chromium are easily oxidized in the presence of moisture (rusting)
4Fe + 3O2 + 6H2O = 4Fe (OH) 3
4Cr + 3O2 + 6H2O = 4Cr (OH) 3
9. INTERACTION OF METALS WITH OXIDES
Metals (Al, Mg, Ca), reduce non-metals or less active metals from their oxides at high temperatures → non-metal or low-activity metal and oxide (calcium-thermal, magnesium-thermal, aluminothermy)
2Al + Cr2O3 = 2Cr + Al2O3 ЗСа + Cr₂O₃ = ЗСаО + 2Cr (800 ° C) 8Al + 3Fe3O4 = 4Al2O3 + 9Fe (thermite) 2Mg + CО2 = 2MgO + С Mg + N2O = MgO + N2 ZnO + CO2 = Z + 2NO = 2CuO + N2 3Zn + SO2 = ZnS + 2ZnO
10. WITH OXIDES
The metals iron and chromium react with oxides, reducing the oxidation state
Cr + Cr2⁺³O3 = 3Cr⁺²O Fe + Fe2⁺³O3 = 3Fe⁺²O
11. INTERACTION OF METALS WITH ALKALI
Alkalis interact only with those metals, oxides and hydroxides of which have amphoteric properties ((Zn, Al, Cr (III), Fe (III), etc.) MELT → metal salt + hydrogen.
2NaOH + Zn → Na2ZnO2 + H2 (sodium zincate)
2Al + 2 (NaOH H2O) = 2NaAlO2 + 3H2
SOLUTION → complex metal salt + hydrogen.
2NaOH + Zn0 + 2H2O = Na2 + H2 (sodium tetrahydroxozincate) 2Al + 2NaOH + 6H2O = 2Na + 3H2
12. REACTION WITH ACIDS (EXCEPT HNO3 and H2SO4 (conc.)
Metals standing in the electrochemical series of metal voltages to the left of hydrogen displace it from dilute acids → salt and hydrogen
Remember! Nitric acid never releases hydrogen when it interacts with metals.
Mg + 2HC1 = MgCl2 + H2
Al + 2HC1 = Al⁺³Сl₃ + Н2
13. REACTIONS WITH SALTS
Active metals displace less active metals from salts. Recovery from solutions:
CuSO4 + Zn = Zn SO4 + Cu
FeSO4 + Cu =REACTIONSNO
Mg + CuCl2 (pp) = MgCl2 +WITHu
Recovery of metals from molten salts
3Na + AlCl₃ = 3NaCl + Al
TiCl2 + 2Mg = MgCl2 + Ti
Group B metals react with salts, lowering the oxidation state
2Fe⁺³Cl3 + Fe = 3Fe⁺²Cl2