Cu is an active metal. Metals. Interaction of metals with alkali solutions
Metals mean a group of elements, which are presented in the form of the simplest substances. They have characteristic properties, namely high electrical and thermal conductivity, positive temperature coefficient of resistance, high ductility and metallic luster.
Note that of the 118 chemical elements that have been discovered so far, the following should be classified as metals:
- among the group of alkaline earth metals there are 6 elements;
- among alkali metals there are 6 elements;
- among transition metals 38;
- in the group of light metals 11;
- There are 7 elements among semimetals,
- 14 among lanthanides and lanthanum,
- 14 in the group of actinides and sea anemones,
- Beryllium and magnesium are outside the definition.
Based on this, 96 elements are classified as metals. Let's take a closer look at what metals react with. Since most metals have a small number of electrons from 1 to 3 at the outer electronic level, in most of their reactions they can act as reducing agents (that is, they give up their electrons to other elements).
Reactions with the simplest elements
- Except for gold and platinum, absolutely all metals react with oxygen. Note also that the reaction occurs with silver at high temperatures, but silver(II) oxide is not formed at normal temperatures. Depending on the properties of the metal, oxides, superoxides and peroxides are formed as a result of reaction with oxygen.
Here are examples of each chemical education:
- lithium oxide – 4Li+O 2 =2Li 2 O;
- potassium superoxide – K+O 2 =KO 2;
- sodium peroxide – 2Na+O 2 =Na 2 O 2.
In order to obtain an oxide from a peroxide, it must be reduced with the same metal. For example, Na 2 O 2 +2Na=2Na 2 O. With low- and medium-active metals, a similar reaction will occur only when heated, for example: 3Fe+2O 2 =Fe 3 O 4.
- Metals can only react with nitrogen with active metals, however, at room temperature only lithium can react, forming nitrides - 6Li+N 2 = 2Li 3 N, however, when heated, the following chemical reaction occurs: 2Al+N 2 = 2AlN, 3Ca+N 2 =Ca 3 N 2.
- Absolutely all metals react with sulfur, as with oxygen, with the exception of gold and platinum. Note that iron can only react when heated with sulfur, forming sulfide: Fe+S=FeS
- Only active metals can react with hydrogen. These include metals of groups IA and IIA, except beryllium. Such reactions can only occur when heated, forming hydrides.
Since the oxidation state of hydrogen is considered? 1, the metals in this case act as reducing agents: 2Na + H 2 = 2NaH.
- The most active metals also react with carbon. As a result of this reaction, acetylenides or methanides are formed.
Let's consider what metals react with water and what do they produce as a result of this reaction? Acetylenes, when reacting with water, will give acetylene, and methane will be obtained as a result of the reaction of water with methanides. Here are examples of these reactions:
- Acetylene – 2Na+2C= Na 2 C 2 ;
- Methane - Na 2 C 2 +2H 2 O=2NaOH+C 2 H 2.
Reaction of acids with metals
Metals can also react differently with acids. Only those metals that are in the series of electrochemical activity of metals up to hydrogen react with all acids.
Let's give an example of a substitution reaction that shows what metals react with. In another way, this reaction is called redox: Mg+2HCl=MgCl 2 +H 2 ^.
Some acids can also interact with metals that come after hydrogen: Cu+2H 2 SO 4 =CuSO 4 +SO 2 ^+2H 2 O.
Note that such a dilute acid can react with a metal according to the classical scheme shown: Mg + H 2 SO 4 = MgSO 4 + H 2 ^.
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 wire, and roll into thin sheets.
2) Metallic shine 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 pole to the positive one under the influence of a small potential difference. When heated, electrical conductivity decreases, because As the temperature increases, vibrations of atoms and ions in the nodes of the crystal lattice intensify, 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 the temperature quickly equalizes over the mass of the metal. The highest thermal conductivity is found 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. The smaller the atomic mass of the metal and the larger the radius of the atom, the smaller it is. 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 most fusible metal is mercury (mp = -39°C), the most refractory metal is tungsten (mp = 3390°C). Metals with melting temperature 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 voltages 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 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
2. Reactions of metals with acids
1) Metals in the electrochemical voltage series 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:
When nitric acid of any concentration and concentrated sulfuric acid interact 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 (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 (k) + 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 an oxide:
Zn + H 2 O – t° → ZnO + H 2
3) Inactive (Au, Ag, Pt) - do not react.
4. Displacement of less active metals by more active metals from solutions of their salts:
Cu + HgCl 2 → Hg+ CuCl 2
Fe+ CuSO 4 → Cu+ FeSO 4
In industry, they often use not pure metals, but mixtures of them - alloys, in which the beneficial properties of one metal are complemented by the beneficial properties of another. Thus, copper has low hardness and is unsuitable for the manufacture of machine parts, while alloys of copper and 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. Based on it, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing the beneficial properties of 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 restorers. However, some metals have low reactivity due to the fact that they are coated surface oxide film, to varying degrees, resistant to chemical reagents such 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 by acids and alkalis. Loose oxide film (rust), formed on the surface of iron in moist air, does not interfere with further oxidation of iron.
Under the influence concentrated acids form on metals sustainable 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 oxidizing agents in acidic solutions, most metals transform 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 transferred into solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only with sulfuric (concentrated) and nitric acids, and Pt and Au - with “regia vodka”.
Metal corrosion
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, the corrosion of iron products in water is widely known, as a result of which rust forms and the products crumble into powder.
Corrosion of metals also occurs in water 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 area of contact between two dissimilar metals can be especially corrosive ( contact corrosion). A galvanic couple occurs between one metal, for example Fe, and another metal, for example Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the voltage series (Re), to the less active metal (Sn, Cu), and the more active metal is destroyed (corroded).
It is because of this that the tinned surface of cans (iron coated with tin) rusts when stored in a humid atmosphere and handled carelessly (iron quickly collapses after even a small scratch appears, allowing 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, since even if there are scratches, it is not the iron that corrodes, but the zinc (a more active metal than iron).
Corrosion resistance for a given metal increases when it is coated with a more active metal or when they are fused; Thus, coating iron with chromium or making an alloy of iron and chromium eliminates corrosion of iron. Chromed iron and steel containing chromium ( stainless steel), have high corrosion resistance.
REMEMBER!!!
Alkali metals – this is group I, A is the main subgroup – Li, Na, K, Rb, Cs, Fr
Alkaline earth metals – this is group II, A – the main subgroup (Be, Mg do not belong) – Ca, Sr, Ba, Ra
n I
Grounds Me(OH) n
OH – hydroxyl group, with valency (I)
Alkalis – these are water-soluble bases (see SOLUBILITY TABLE)
I n
Acids - these are complex substances with the general formula N n (KO)
(KO) – acidic residue
V - VII
Acid oxide – neMe x O y And Fur x O y
I, II
Basic oxides – Fur x O y
I. Interaction of water with metals.
Depending on the activity of the metal, the reaction occurs under different conditions and different products are formed.
1). Interaction with the most active metals , standing in the periodic table at I A and I I A groups (alkali and alkaline earth metals) and aluminum . In the activity series, these metals are located up to aluminum (inclusive)
The reaction proceeds under normal conditions, producing alkali and hydrogen.
I I
2Li + 2 H 2 O = 2 Li OH + H 2
HOH hydroxide
lithium
I II
Ba + 2 H 2 O= Ba (OH) 2 + H 2
2 Al + 6 H 2 O = 2Al (OH) 3 + 3 H 2
hydroxide
aluminum
OH is a hydroxo group, it is always monovalent
CONCLUSION – active metals - Li, Na, K, Rb, Cs, Fr, Ca, Sr, Ba, Ra+ Al - react like this
|
Me + H 2 O =Me(OH) n + H 2( R. substitution) Base |
2) Interaction with less active metals, which are located in the activity series from aluminum to hydrogen.
The reaction occurs only with steam water, i.e. when heated.
In this case, the following are formed: the oxide of this metal and hydrogen.
I II I
Fe + H 2 O = FeO + H 2 (substitution reaction occurs)
oxide
gland
Ni + H 2 O = NiO + H 2
(The valence of a metal can be easily determined by the activity series of metals; above their symbol there is a value, for example +2, this means that the valency of this metal is 2).
CONCLUSION – metals of medium activity, standing in the activity series up to (H 2) – Be, Mg, Fe, Pb, Cr, Ni, Mn, Zn - react like this
3) Metals in the activity series after hydrogen do not react with water.
Cu + H 2 O = no reaction
I I. Interaction with oxides (basic and acidic)
Only those oxides that react with water react with water to produce a water-soluble product (acid or alkali).
1). Interaction with basic oxides.
Only the main oxides of active metals, which are located in the I A and I I A groups, interact with water, except for Be and Mg (aluminum oxide does not react, because it is amphoteric). The reaction proceeds under normal conditions, and only an alkali is formed.
I II
Na 2 O + H 2 O = 2 NaOHBaO + H 2 O =Ba (OH) 2 (a compound reaction occurs)
2) Interaction of acid oxides with water.
Acidic oxides all react with water. The only exception is SiO 2.
This produces acids. In all acids, hydrogen comes first, so the reaction equation is written as follows:
SO 3 + H 2 O = H 2 SO 4 P 2 O 5 + H 2 O = 2 HPO 3
SO 3 cold
+H2O P2O5
H2SO4 + H2O
H2P2O6
P 2 O 5 +3 H 2 O=2 H 3 PO 4
hot
P2O5
+ H 6 O 3
H6P2O8
note that depending on the temperature of water, different products are formed when interacting with P 2 O 5.
IVWater interaction cnon-metals
Examples: Cl 2 +H 2 O =HCl +HClO
C +H 2 O =CO +H 2
carbon dioxide
Si +2H 2 O =SiO 2 +2H 2.
Metals vary greatly in their chemical activity. The chemical activity of a metal can be approximately judged by its position in.
The most active metals are located at the beginning of this row (on the left), the least active are at the end (on the right).
Reactions with simple substances. Metals react with nonmetals to form binary compounds. The reaction conditions, and sometimes their products, vary greatly for different metals.
For example, alkali metals actively react with oxygen (including in air) at room temperature to form oxides and peroxides
4Li + O 2 = 2Li 2 O;
2Na + O 2 = Na 2 O 2
Medium activity metals react with oxygen when heated. In this case, oxides are formed:
2Mg + O 2 = t 2MgO.
Low-active metals (for example, gold, platinum) do not react with oxygen and therefore practically do not change their luster in air.
Most metals, when heated with sulfur powder, form the corresponding sulfides:
Reactions with complex substances. Compounds of all classes react with metals - oxides (including water), acids, bases and salts.
Active metals react violently with water at room temperature:
2Li + 2H 2 O = 2LiOH + H 2;
Ba + 2H 2 O = Ba(OH) 2 + H 2.
The surface of metals such as magnesium and aluminum is protected by a dense film of the corresponding oxide. This prevents the reaction from occurring with water. However, if this film is removed or its integrity is disrupted, then these metals also actively react. For example, powdered magnesium reacts with hot water:
Mg + 2H 2 O = 100 °C Mg(OH) 2 + H 2.
At elevated temperatures, less active metals also react with water: Zn, Fe, Mil, etc. In this case, the corresponding oxides are formed. For example, when passing water vapor over hot iron filings, the following reaction occurs:
3Fe + 4H 2 O = t Fe 3 O 4 + 4H 2.
Metals in the activity series up to hydrogen react with acids (except HNO 3) to form salts and hydrogen. Active metals (K, Na, Ca, Mg) react with acid solutions very violently (at high speed):
Ca + 2HCl = CaCl 2 + H 2;
2Al + 3H 2 SO 4 = Al 2 (SO 4) 3 + 3H 2.
Low-active metals are often practically insoluble in acids. This is due to the formation of a film of insoluble salt on their surface. For example, lead, which is in the activity series before hydrogen, is practically insoluble in dilute sulfuric and hydrochloric acids due to the formation of a film of insoluble salts (PbSO 4 and PbCl 2) on its surface.
You need to enable JavaScript to voteCharacteristic chemical properties of simple substances - metals
Most chemical elements are classified as metals - 92 out of 114 known elements. Metals- these are chemical elements whose atoms give up electrons from the outer (and some from the outer) electron layer, turning into positive ions. This property of metal atoms is determined by that they have relatively large radii and a small number of electrons(mostly 1 to 3 on the outer layer). The only exceptions are 6 metals: germanium, tin, and lead atoms on the outer layer have 4 electrons, antimony and bismuth atoms - 5, polonium atoms - 6. For metal atoms characterized by small electronegativity values(from 0.7 to 1.9) and exclusively restorative properties, i.e. the ability to donate electrons. In the Periodic Table of Chemical Elements of D.I. Mendeleev, metals are located below the boron - astatine diagonal, as well as above it, in secondary subgroups. In the periods and main subgroups, there are known patterns in the changes in the metallic, and therefore the reducing properties of the atoms of the elements.
Chemical elements located near the boron-astatine diagonal (Be, Al, Ti, Ge, Nb, Sb, etc.) have dual properties: in some of their compounds they behave like metals, in others they exhibit the properties of non-metals. In secondary subgroups, the reducing properties of metals most often decrease with increasing atomic number.
Compare the activity of the metals of group I of the secondary subgroup known to you: Cu, Ag, Au; Group II of the secondary subgroup: Zn, Cd, Hg - and you will see this for yourself. This can be explained by the fact that the strength of the bond between the valence electrons and the nucleus in the atoms of these metals is largely influenced by the magnitude of the nuclear charge, and not by the radius of the atom. The nuclear charge increases significantly, and the attraction of electrons to the nucleus increases. In this case, although the atomic radius increases, it is not as significant as for the metals of the main subgroups.
Simple substances formed by chemical elements - metals, and complex metal-containing substances play a vital role in the mineral and organic “life” of the Earth. Suffice it to remember that atoms (ions) of metal elements are an integral part of compounds that determine metabolism in the body of humans and animals. For example, 76 elements are found in human blood, and only 14 of them are not metals.
In the human body, some metal elements (calcium, potassium, sodium, magnesium) are present in large quantities, i.e. they are macroelements. And metals such as chromium, manganese, iron, cobalt, copper, zinc, molybdenum are present in small quantities, i.e. these are trace elements. If a person weighs 70 kg, then his body contains (in grams): calcium - 1700, potassium - 250, sodium - 70, magnesium - 42, iron - 5, zinc - 3. All metals are extremely important, health problems arise and with their deficiency, and with their excess.
For example, sodium ions regulate water content in the body and the transmission of nerve impulses. Its deficiency leads to headaches, weakness, poor memory, loss of appetite, and its excess leads to increased blood pressure, hypertension, and heart disease.
Simple substances - metals
The emergence of civilization (Bronze Age, Iron Age) is associated with the development of the production of metals (simple substances) and alloys. The scientific and technological revolution that began about 100 years ago, affecting both industry and the social sphere, is also closely related to the production of metals. Based on tungsten, molybdenum, titanium and other metals, they began to create corrosion-resistant, super-hard, refractory alloys, the use of which greatly expanded the capabilities of mechanical engineering. In nuclear and space technology, tungsten and rhenium alloys are used to make parts that operate at temperatures up to 3000 °C; In medicine, surgical instruments made of tantalum and platinum alloys and unique ceramics based on titanium and zirconium oxides are used.
And, of course, we must not forget that most alloys use the long-known metal iron, and the basis of many light alloys is made up of relatively “young” metals - aluminum and magnesium. Composite materials have become supernovae, representing, for example, polymer or ceramics, which inside (like concrete with iron rods) are strengthened with metal fibers from tungsten, molybdenum, steel and other metals and alloys - it all depends on the goal set and the properties of the material necessary to achieve it. The figure shows a diagram of the crystal lattice of sodium metal. In it, each sodium atom is surrounded by eight neighbors. The sodium atom, like all metals, has many empty valence orbitals and few valence electrons. Electronic formula of the sodium atom: 1s 2 2s 2 2p 6 3s 1 3p 0 3d 0, where 3s, 3p, 3d - valence orbitals.
Single valence electron of sodium atom 3s 1 can occupy any of the nine free orbitals - 3s (one), 3p (three) and 3d (five), because they do not differ much in energy level. When atoms approach each other, when a crystal lattice is formed, the valence orbitals of neighboring atoms overlap, due to which electrons move freely from one orbital to another, establishing bonds between all atoms of the metal crystal. Such a chemical bond is called metallic.
A metallic bond is formed by elements whose atoms in the outer layer have few valence electrons compared to a large number of outer orbitals that are energetically close. Their valence electrons are weakly held in the atom. The electrons that carry out the communication are socialized and move throughout the crystal lattice of the generally neutral metal. Substances with a metallic bond are characterized by metallic crystal lattices, which are usually depicted schematically as shown in the figure. Cations and metal atoms located at the sites of the crystal lattice provide its stability and strength (socialized electrons are depicted as small black balls).
Metal connection- this is a bond in metals and alloys between metal atoms located at the nodes of the crystal lattice, carried out by shared valence electrons. Some metals crystallize in two or more crystalline forms. This property of substances - to exist in several crystalline modifications - is called polymorphism. Polymorphism of simple substances is known as allotropy. For example, iron has four crystalline modifications, each of which is stable in a certain temperature range:
α - stable up to 768 °C, ferromagnetic;
β - stable from 768 to 910 °C, non-ferromagnetic, i.e. paramagnetic;
γ - stable from 910 to 1390 °C, non-ferromagnetic, i.e. paramagnetic;
δ - stable from 1390 to 1539 °C (£° pl iron), non-ferromagnetic.
Tin has two crystalline modifications:
α - stable below 13.2 °C (p = 5.75 g/cm3). This is gray tin. It has a diamond-type crystal lattice (atomic);
β - stable above 13.2 °C (p = 6.55 g/cm3). This is white tin.
White tin is a silvery-white, very soft metal. When cooled below 13.2 °C, it crumbles into gray powder, since during the transition its specific volume increases significantly. This phenomenon was called the “tin plague.”
Of course, a special type of chemical bond and the type of crystal lattice of metals must determine and explain their physical properties. What are they? These are metallic luster, ductility, high electrical and thermal conductivity, an increase in electrical resistance with increasing temperature, as well as such significant properties as density, high melting and boiling points, hardness, and magnetic properties. A mechanical effect on a crystal with a metal crystal lattice causes a displacement of layers of ion-atoms relative to each other (Fig. 17), and since electrons move throughout the crystal, bond breaking does not occur, therefore metals are characterized by greater plasticity. A similar effect on a solid with covalent bonds (an atomic crystal lattice) leads to the breaking of covalent bonds. Breaking bonds in the ionic lattice leads to mutual repulsion of like-charged ions. Therefore, substances with atomic and ionic crystal lattices are fragile. The most ductile metals are Au, Ag, Sn, Pb, Zn. They are easily drawn into wire, can be forged, pressed, or rolled into sheets. For example, gold foil 0.003 mm thick can be made from gold, and a thread 1 km long can be drawn from 0.5 g of this metal. Even mercury, which is liquid at room temperature, becomes malleable in its solid state at low temperatures, like lead. Only Bi and Mn do not have plasticity; they are brittle.
Why do metals have a characteristic shine and are also opaque?
Electrons filling the interatomic space reflect light rays (rather than transmit them like glass), and most metals equally scatter all rays of the visible part of the spectrum. Therefore they are silvery white or gray in color. Strontium, gold and copper absorb short wavelengths (close to violet) to a greater extent and reflect long wavelengths of the light spectrum, and therefore have light yellow, yellow and “copper” colors. Although in practice, metal does not always seem like a “light body” to us. Firstly, its surface can oxidize and lose its shine. Therefore, native copper appears as a greenish stone. And secondly, even pure metal may not shine. Very thin sheets of silver and gold have a completely unexpected appearance - they have a bluish-green color. And fine metal powders appear dark gray, even black. Silver, aluminum, and palladium have the greatest reflectivity. They are used in the manufacture of mirrors, including spotlights.
Why do metals have high electrical conductivity and conduct heat?
Chaotically moving electrons in a metal, under the influence of an applied electrical voltage, acquire directional movement, i.e., they conduct electric current. As the temperature of the metal increases, the vibration amplitudes of the atoms and ions located at the nodes of the crystal lattice increase. This makes it difficult for electrons to move, and the electrical conductivity of the metal drops. At low temperatures, the oscillatory motion, on the contrary, is greatly reduced and the electrical conductivity of metals increases sharply. Near absolute zero, metals have virtually no resistance; most metals exhibit superconductivity.
It should be noted that non-metals that have electrical conductivity (for example, graphite), at low temperatures, on the contrary, do not conduct electric current due to the lack of free electrons. And only with increasing temperature and the destruction of some covalent bonds does their electrical conductivity begin to increase. Silver, copper, as well as gold and aluminum have the highest electrical conductivity; manganese, lead, and mercury have the lowest.
Most often, the thermal conductivity of metals changes with the same pattern as electrical conductivity. It is due to the high mobility of free electrons, which, colliding with vibrating ions and atoms, exchange energy with them. The temperature is equalized throughout the entire piece of metal.
Mechanical strength, density, melting point of metals are very different. Moreover, with an increase in the number of electrons connecting ion-atoms and a decrease in the interatomic distance in crystals, the indicators of these properties increase.
So, alkali metals(Li, K, Na, Rb, Cs), the atoms of which have one valence electron, soft (cut with a knife), with low density (lithium is the lightest metal with p = 0.53 g/cm 3) and melts at low temperatures (for example, the melting point of cesium is 29 ° C). The only metal that is liquid under normal conditions is mercury, which has a melting point of -38.9 °C. Calcium, which has two electrons in the outer energy level of its atoms, is much harder and melts at a higher temperature (842 °C). Even more durable is the crystal lattice formed by scandium ions, which have three valence electrons. But the strongest crystal lattices, high densities and melting temperatures are observed in metals of secondary subgroups V, VI, VII, VIII. This is explained by the fact that metals of side subgroups, which have unpaired valence electrons at the d-sublevel, are characterized by the formation of very strong covalent bonds between atoms, in addition to the metallic one, carried out by electrons of the outer layer from the s-orbitals.
The heaviest metal- this is osmium (Os) with p = 22.5 g/cm 3 (a component of super-hard and wear-resistant alloys), the most refractory metal is tungsten W with t = 3420 ° C (used for the manufacture of incandescent lamp filaments), the hardest metal is - This is Cr chrome (scratch glass). They are part of the materials from which metal-cutting tools, brake pads of heavy machines, etc. are made. Metals interact with the magnetic field in different ways. Metals such as iron, cobalt, nickel and gadolinium stand out for their ability to be highly magnetized. They are called ferromagnets. Most metals (alkali and alkaline earth metals and a significant part of transition metals) are weakly magnetized and do not retain this state outside a magnetic field - they are paramagnetic. Metals pushed out by a magnetic field are diamagnetic (copper, silver, gold, bismuth).
When considering the electronic structure of metals, we divided metals into metals of the main subgroups (s- and p-elements) and metals of secondary subgroups (transition d- and f-elements).
In technology, it is customary to classify metals according to various physical properties:
1. Density - light (p< 5 г/см 3) и тяжелые (все остальные).
2. Melting point - low-melting and refractory.
There are classifications of metals based on their chemical properties. Metals with low chemical activity are called noble(silver, gold, platinum and its analogues - osmium, iridium, ruthenium, palladium, rhodium). Based on the similarity of chemical properties, they distinguish alkaline(metals of the main subgroup of group I), alkaline earth(calcium, strontium, barium, radium), as well as rare earth metals(scandium, yttrium, lanthanum and lanthanides, actinium and actinides).
General chemical properties of metals
Metal atoms are relatively easy donate valence electrons and turn into positively charged ions, that is, they are oxidized. This is the main common property of both atoms and simple substances - metals. Metals are always reducing agents in chemical reactions. The reducing ability of atoms of simple substances - metals formed by chemical elements of one period or one main subgroup of D. I. Mendeleev's Periodic Table changes naturally.
The reduction activity of a metal in chemical reactions that occur in aqueous solutions is reflected by its position in the electrochemical voltage series of metals.
Based on this series of voltages, the following important conclusions can be drawn about the chemical activity of metals in reactions occurring in aqueous solutions under standard conditions (t = 25 °C, p = 1 atm).
· The further to the left a metal is in this row, the more powerful a reducing agent it is.
· Each metal is capable of displacing (reducing) from salts in solution those metals that are located after it in the series of stresses (to the right).
· Metals located in the voltage series to the left of hydrogen are capable of displacing it from acids in solution
· Metals that are the strongest reducing agents (alkali and alkaline earth) react primarily with water in any aqueous solution.
The reduction activity of a metal, determined from the electrochemical series, does not always correspond to its position in the periodic table. This is explained by the fact that when determining the position of a metal in a series of stresses, not only the energy of electron abstraction from individual atoms is taken into account, but also the energy expended on the destruction of the crystal lattice, as well as the energy released during the hydration of ions. For example, lithium is more active in aqueous solutions than sodium (although Na is a more active metal by position in the periodic table). The fact is that the hydration energy of Li + ions is much greater than the hydration energy of Na +, so the first process is energetically more favorable. Having examined the general provisions characterizing the reducing properties of metals, let us move on to specific chemical reactions.
Interaction of metals with non-metals
· Most metals form oxides with oxygen- basic and amphoteric. Acidic transition metal oxides, such as chromium (VI) oxide CrOg or manganese (VII) oxide Mn 2 O 7, are not formed by direct oxidation of the metal with oxygen. They are obtained indirectly.
Alkali metals Na, K react actively with oxygen in the air, forming peroxides:
Sodium oxide is obtained indirectly by calcining peroxides with the corresponding metals:
Lithium and alkaline earth metals react with atmospheric oxygen, forming basic oxides:
Other metals, except gold and platinum metals, which are not oxidized by atmospheric oxygen at all, interact with it less actively or when heated:
· With halogens, metals form salts of hydrohalic acids, For example:
· The most active metals form hydrides with hydrogen- ionic salt-like substances in which hydrogen has an oxidation state of -1, for example:
Many transition metals form hydrides of a special type with hydrogen - it is as if hydrogen is dissolved or introduced into the crystal lattice of metals between atoms and ions, while the metal retains its appearance, but increases in volume. The absorbed hydrogen is in the metal, apparently in atomic form.
There are also intermediate metal hydrides.
· Gray metals form salts - sulfides, For example:
· Metals react somewhat more difficultly with nitrogen, because the chemical bond in the nitrogen molecule N2 is very strong; In this case, nitrides are formed. At ordinary temperatures, only lithium reacts with nitrogen:
Interaction of metals with complex substances
· With water. Under normal conditions, alkali and alkaline earth metals displace hydrogen from water and form soluble bases - alkalis, for example:
Other metals that are in the voltage series before hydrogen can also, under certain conditions, displace hydrogen from water. But aluminum reacts violently with water only if the oxide film is removed from its surface:
Magnesium reacts with water only when boiled, and hydrogen is also released:
If burning magnesium is added to water, the combustion continues because the reaction occurs:
Iron reacts with water only when it is hot:
· With acids in solution (HCl, H 2 SO 4 ), CH 3 COOH and others, except HNO 3 ) metals that are in the voltage series up to hydrogen interact. This produces salt and hydrogen.
But lead (and some other metals), despite its position in the voltage series (to the left of hydrogen), is almost insoluble in dilute sulfuric acid, since the resulting lead sulfate PbSO 4 is insoluble and creates a protective film on the metal surface.
· With salts of less active metals in solution. As a result of this reaction, a salt of a more active metal is formed and a less active metal is released in free form.
It must be remembered that the reaction occurs in cases where the resulting salt is soluble. The displacement of metals from their compounds by other metals was first studied in detail by N. N. Beketov, a great Russian scientist in the field of physical chemistry. He arranged metals according to their chemical activity into a “displacement series,” which became the prototype of a series of metal stresses.
· With organic substances. Interaction with organic acids is similar to reactions with mineral acids. Alcohols can exhibit weak acidic properties when interacting with alkali metals:
Phenol reacts similarly:
Metals participate in reactions with haloalkanes, which are used to obtain lower cycloalkanes and for syntheses during which the carbon skeleton of the molecule becomes more complex (A. Wurtz reaction):
· Metals whose hydroxides are amphoteric react with alkalis in solution. For example:
· Metals can form chemical compounds with each other, which are collectively called intermetallic compounds. They most often do not exhibit oxidation states of atoms, which are characteristic of compounds of metals with non-metals. For example:
Cu 3 Au, LaNi 5, Na 2 Sb, Ca 3 Sb 2, etc.
Intermetallic compounds usually do not have a constant composition; the chemical bond in them is mainly metallic. The formation of these compounds is more typical for metals of secondary subgroups.
Metals of the main subgroups of groups I-III of the Periodic Table of Chemical Elements by D. I. Mendeleev
general characteristics
These are metals of the main subgroup of group I. Their atoms at the outer energy level have one electron each. Alkali metals - strong reducing agents. Their reducing power and chemical activity increase with increasing atomic number of the element (i.e., from top to bottom in the Periodic Table). All of them have electronic conductivity. The strength of the bond between alkali metal atoms decreases with increasing atomic number of the element. Their melting and boiling points also decrease. Alkali metals react with many simple substances - oxidizing agents. In reactions with water they form water-soluble bases (alkalis). Alkaline earth elements are called the elements of the main subgroup of group II. The atoms of these elements contain at the outer energy level two electrons each. They are the strongest reducing agents, have an oxidation state of +2. In this main subgroup, general patterns in changes in physical and chemical properties are observed, associated with an increase in the size of atoms in the group from top to bottom, and the chemical bond between atoms also weakens. As the size of the ion increases, the acidic properties of oxides and hydroxides become weaker and the basic ones increase.
The main subgroup of group III consists of the elements boron, aluminum, gallium, indium and thallium. All elements are p-elements. At the external energy level they have three(s) 2 p 1 ) electron, which explains the similarity of properties. Oxidation state +3. Within a group, as the nuclear charge increases, the metallic properties increase. Boron is a non-metallic element, while aluminum already has metallic properties. All elements form oxides and hydroxides.
Most metals are found in subgroups of the Periodic Table. Unlike the elements of the main subgroups, where the outer level of atomic orbitals is gradually filled with electrons, the d-orbitals of the penultimate energy level and the s-orbitals of the last one are filled in the elements of the secondary subgroups. The number of electrons corresponds to the group number. Elements with an equal number of valence electrons are grouped under the same number. All elements of subgroups are metals.
Simple substances formed by subgroup metals have strong crystal lattices that are resistant to heat. These metals are the strongest and most refractory among other metals. In d-elements, a transition with an increase in their valency from basic properties through amphoteric to acidic is clearly visible.
Alkali metals (Na, K)
At the external energy level, the alkali metal atoms of the elements contain one electron each, located at a great distance from the core. They easily give up this electron, so they are strong reducing agents. In all compounds, alkali metals exhibit an oxidation state of +1. Their reducing properties increase with increasing atomic radius from Li to Cs. All of them are typical metals, have a silvery-white color, are soft (can be cut with a knife), light and fusible. Actively interact with everyone non-metals:
All alkali metals, when reacting with oxygen (with the exception of Li), form peroxides. Alkali metals are not found in free form due to their high chemical reactivity.
Oxides- solids with basic properties. They are obtained by calcining peroxides with the corresponding metals:
Hydroxides NaOH, KOH- solid white substances, hygroscopic, soluble in water with the release of heat, they are classified as alkalis:
Alkali metal salts are almost all soluble in water. The most important of them: Na 2 CO 3 - sodium carbonate; Na 2 CO 3 10H 2 O - crystalline soda; NaHCO 3 - sodium bicarbonate, baking soda; K 2 CO 3 - potassium carbonate, potash; Na 2 SO 4 10H 2 O - Glauber's salt; NaCl - sodium chloride, table salt.
Group I elements in tables
Alkaline earth metals (Ca, Mg)
Calcium (Ca) is a representative alkaline earth metals, which are the names of the elements of the main subgroup of group II, but not all, but only starting from calcium and down the group. These are the chemical elements that, when interacting with water, form alkalis. Calcium at the external energy level contains two electrons, oxidation state +2.
The physical and chemical properties of calcium and its compounds are presented in the table.
Magnesium (Mg) has the same atomic structure as calcium, its oxidation state is also +2. It is a soft metal, but its surface is covered with a protective film in air, which slightly reduces chemical reactivity. Its combustion is accompanied by a blinding flash. MgO and Mg(OH) 2 exhibit basic properties. Although Mg(OH) 2 is slightly soluble, it colors the phenolphthalein solution crimson.
Mg + O 2 = MgO 2
MO oxides are hard, white, refractory substances. In engineering, CaO is called quicklime, and MgO is called burnt magnesia; these oxides are used in the production of building materials. The reaction of calcium oxide with water is accompanied by the release of heat and is called slaking of lime, and the resulting Ca(OH) 2 is called slaked lime. A transparent solution of calcium hydroxide is called lime water, and a white suspension of Ca(OH) 2 in water is called milk of lime.
Magnesium and calcium salts are obtained by reacting them with acids.
CaCO 3 - calcium carbonate, chalk, marble, limestone. Used in construction. MgCO 3 - magnesium carbonate - is used in metallurgy to remove slag.
CaSO 4 2H 2 O - gypsum. MgSO 4 - magnesium sulfate - called bitter, or English, salt, found in sea water. BaSO 4 - barium sulfate - due to its insolubility and ability to block X-rays, it is used in diagnostics (“barite porridge”) of the gastrointestinal tract.
Calcium accounts for 1.5% of the human body weight, 98% of calcium is found in bones. Magnesium is a bioelement; there is about 40 g of it in the human body; it is involved in the formation of protein molecules.
Alkaline earth metals in tables
Aluminum
Aluminum (Al)- element of the main subgroup of group III of the periodic system of D.I. Mendeleev. The aluminum atom contains at the outer energy level three electrons, which it easily releases during chemical interactions. The ancestor of the subgroup and the upper neighbor of aluminum - boron - has a smaller atomic radius (for boron it is 0.080 nm, for aluminum - 0.143 nm). In addition, the aluminum atom has one intermediate eight-electron layer (2e; 8e; 3e), which prevents the outer electrons from reaching the nucleus. Therefore, the reducing properties of aluminum atoms are quite pronounced.
In almost all of its compounds, aluminum has oxidation state +3.
Aluminum is a simple substance
Silver-white light metal. Melts at 660 °C. It is very plastic, easily drawn into wire and rolled into foil up to 0.01 mm thick. It has very high electrical and thermal conductivity. They form light and strong alloys with other metals. Aluminum is a very active metal. If aluminum powder or thin aluminum foil is heated strongly, they ignite and burn with a blinding flame:
This reaction can be observed when sparklers and fireworks burn. Aluminum, like all metals, Reacts easily with non-metals, especially in powder form. In order for the reaction to begin, initial heating is necessary, with the exception of reactions with halogens - chlorine and bromine, but then all reactions of aluminum with non-metals proceed very violently and are accompanied by the release of a large amount of heat:
Aluminum dissolves well in dilute sulfuric and hydrochloric acids:
And here concentrated sulfuric and nitric acids passivate aluminum, forming on the metal surface dense durable oxide film, which prevents the further progress of the reaction. Therefore, these acids are transported in aluminum tanks.
Aluminum oxide and hydroxide have amphoteric properties, therefore aluminum dissolves in aqueous solutions of alkalis, forming salts - aluminates:
Aluminum is widely used in metallurgy to produce metals - chromium, manganese, vanadium, titanium, zirconium from their oxides. This method is called aluminothermy. In practice, thermite is often used - a mixture of Fe 3 O 4 with aluminum powder. If this mixture is set on fire, for example, using a magnesium tape, then a vigorous reaction occurs, releasing a large amount of heat:
The heat released is quite sufficient to completely melt the resulting iron, so this process is used for welding steel products.
Aluminum can be obtained by electrolysis - the decomposition of the melt of its oxide Al 2 O 3 into its component parts using an electric current. But the melting point of aluminum oxide is about 2050 °C, so electrolysis requires large amounts of energy.
Aluminum connections
Aluminosilicates. These compounds can be considered as salts formed by the oxide of aluminum, silicon, alkali and alkaline earth metals. They make up the bulk of the earth's crust. In particular, aluminosilicates are part of feldspars, the most common minerals and clays.
Bauxite- a rock from which aluminum is obtained. It contains aluminum oxide Al 2 O 3.
Corundum- a mineral of the composition Al 2 O 3, has very high hardness, its fine-grained variety containing impurities - emery, is used as an abrasive (grinding) material. Another natural compound, alumina, has the same formula.
Transparent, colored with impurities, corundum crystals are well known: red - rubies and blue - sapphires, which are used as precious stones. Currently, they are obtained artificially and are used not only for jewelry, but also for technical purposes, for example, for the manufacture of parts for watches and other precision instruments. Ruby crystals are used in lasers.
Aluminum oxide Al 2 O 3 - a white substance with a very high melting point. Can be obtained by decomposing aluminum hydroxide by heating:
Aluminum hydroxide Al(OH) 3 precipitates in the form of a gelatinous precipitate under the action of alkalis on solutions of aluminum salts:
How amphoteric hydroxide it dissolves easily in acids and alkali solutions:
Aluminates are called salts of unstable aluminum acids - orthoaluminum H 2 AlO 3, meta-aluminum HAlO 2 (it can be considered as orthoaluminum acid, from the molecule of which a water molecule has been removed). Natural aluminates include noble spinel and precious chrysoberyl. Aluminum salts, except phosphates, are highly soluble in water. Some salts (sulfides, sulfites) are decomposed by water. Aluminum chloride AlCl 3 is used as a catalyst in the production of many organic substances.
Group III elements in tables
Characteristics of transition elements - copper, zinc, chromium, iron
Copper (Cu)- element of a secondary subgroup of the first group. Electronic formula: (…3d 10 4s 1). Its tenth d-electron is mobile, because it has moved from the 4S sublevel. Copper in compounds exhibits oxidation states +1 (Cu 2 O) and +2 (CuO). Copper is a light pink metal, malleable, viscous, and an excellent conductor of electricity. Melting point 1083 °C.
Like other metals of subgroup I of group I of the periodic system, copper stands to the right of hydrogen in the activity series and does not displace it from acids, but reacts with oxidizing acids:
Under the influence of alkalis on solutions of copper salts, a precipitate of a weak base of blue color precipitates.- copper (II) hydroxide, which when heated decomposes into basic black oxide CuO and water:
Chemical properties of copper in tables
Zinc (Zn)- element of a secondary subgroup of group II. Its electronic formula is as follows: (…3d 10 4s 2). Since the penultimate d-sublevel in zinc atoms is completely complete, zinc in compounds exhibits an oxidation state of +2.
Zinc is a silver-white metal that practically does not change in air. It is corrosion resistant due to the presence of an oxide film on its surface. Zinc is one of the most active metals at elevated temperatures reacts with simple substances:
Zinc, like other metals, displaces less active metals from their salts:
Zn + 2AgNO 3 = 2Ag + Zn(NO 3) 2
Zinc hydroxide is amphoteric, i.e., exhibits the properties of both acids and bases. When a solution of alkali is gradually added to a solution of zinc salt, the precipitate that initially formed dissolves (the same happens with aluminum):
Chemical properties of zinc in tables
For example chromium (Cr) it can be shown that properties of transition elements do not change significantly along the period: A quantitative change occurs due to a change in the number of electrons in the valence orbitals. The maximum oxidation state of chromium is +6. The metal in the activity series is to the left of hydrogen and displaces it from acids:
When an alkali solution is added to such a solution, a precipitate of Me(OH) is formed 2 , which is quickly oxidized by atmospheric oxygen:
It corresponds to the amphoteric oxide Cr 2 O 3. Chromium oxide and hydroxide (in the highest oxidation state) exhibit the properties of acidic oxides and acids, respectively. Chromic acid salts (H 2 CrO 4 ) in an acidic environment transform into dichromates- salts of dichromic acid (H 2 Cr 2 O 7). Chromium compounds have a high oxidizing ability.
Chemical properties of chromium in tables
Iron Fe- an element of the secondary subgroup of group VIII and the 4th period of the periodic table of D. I. Mendeleev. Iron atoms are structured somewhat differently from the atoms of the elements of the main subgroups. As befits an element of the 4th period, iron atoms have four energy levels, but it is not the last one that is filled, but the penultimate level, the third from the nucleus. At the last level, iron atoms contain two electrons. At the penultimate level, which can accommodate 18 electrons, the iron atom has 14 electrons. Consequently, the distribution of electrons across levels in iron atoms is as follows: 2e; 8e ; 14e; 2e. Like all metals, iron atoms exhibit reducing properties, giving away during chemical interactions not only two electrons from the last level, and acquiring an oxidation state of +2, but also an electron from the penultimate level, while the oxidation state of the atom increases to +3.
Iron is a simple substance
It is a silvery-white shiny metal with a melting point of 1539 °C. It is very plastic, therefore it is easy to process, forge, roll, stamp. Iron has the ability to be magnetized and demagnetized. It can be given greater strength and hardness using thermal and mechanical methods. There are technically pure and chemically pure iron. Technically pure iron is essentially low-carbon steel; it contains 0.02-0.04% carbon, and even less oxygen, sulfur, nitrogen and phosphorus. Chemically pure iron contains less than 0.01% impurities. For example, paper clips and buttons are made from technically pure iron. Such iron easily corrodes, while chemically pure iron is almost not subject to corrosion. Currently, iron is the basis of modern technology and agricultural engineering, transport and communications, spaceships and, in general, all modern civilization. Most products, from a sewing needle to spacecraft, cannot be made without the use of iron.
Chemical properties of iron
Iron can exhibit oxidation states +2 and +3, accordingly, iron gives two series of compounds. The number of electrons that an iron atom gives up during chemical reactions depends on the oxidizing ability of the substances reacting with it.
For example, with halogens, iron forms halides, in which it has an oxidation state of +3:
and with sulfur - iron (II) sulfide:
Hot iron burns in oxygen with the formation of iron scale:
At high temperatures (700-900 °C) iron reacts with water vapor:
In accordance with the position of iron in the electrochemical voltage series, it can displace metals to the right of it from aqueous solutions of their salts, for example:
Iron dissolves in dilute hydrochloric and sulfuric acids, i.e., it is oxidized by hydrogen ions:
Iron also dissolves in dilute nitric acid., this produces iron (III) nitrate, water and the products of reduction of nitric acid - N 2, NO or NH 3 (NH 4 NO 3) depending on the concentration of the acid.
Iron compounds
In nature, iron forms a number of minerals. This is magnetic iron ore (magnetite) Fe 3 O 4, red iron ore (hematite) Fe 2 O 3, brown iron ore (limonite) 2Fe 2 O 3 3H 2 O. Another natural iron compound is iron, or sulfur, pyrite (pyrite) FeS 2, does not serve as iron ore for metal production, but is used for the production of sulfuric acid.
Iron is characterized by two series of compounds: iron(II) and iron(III) compounds. Iron (II) oxide FeO and its corresponding iron (II) hydroxide Fe(OH) 2 are obtained indirectly, in particular, through the following chain of transformations:
Both compounds have distinct basic properties.
Iron(II) cations Fe 2 + easily oxidized by atmospheric oxygen to iron (III) cations Fe 3 + . Therefore, the white precipitate of iron (II) hydroxide turns green and then turns brown, turning into iron (III) hydroxide:
Iron(III) oxide Fe 2 O 3 and the corresponding iron (III) hydroxide Fe(OH) 3 is also obtained indirectly, for example, along the chain:
Of the iron salts, sulfates and chlorides are of greatest technical importance.
Crystal hydrate of iron (II) sulfate FeSO 4 7H 2 O, known as iron sulfate, is used to control plant pests, to prepare mineral paints and for other purposes. Iron (III) chloride FeCl 3 is used as a mordant when dyeing fabrics. Iron (III) sulfate Fe 2 (SO 4) 3 9H 2 O is used for water purification and other purposes.
The physical and chemical properties of iron and its compounds are summarized in the table:
Chemical properties of iron in tables
Qualitative reactions to Fe 2+ and Fe 3+ ions
For recognition of iron (II) and (III) compounds carry out qualitative reactions to Fe ions 2+ and Fe 3+ . A qualitative reaction to Fe 2+ ions is the reaction of iron (II) salts with the K 3 compound, called red blood salt. This is a special group of salts called complex salts, which you will become familiar with later. In the meantime, you need to understand how such salts dissociate:
The reagent for Fe 3+ ions is another complex compound - yellow blood salt - K 4, which dissociates in solution in a similar way:
If solutions containing Fe 2+ and Fe 3+ ions are added, respectively, to solutions of red blood salt (reagent for Fe 2+) and yellow blood salt (reagent for Fe 3+), then in both cases the same blue precipitate precipitates:
To detect Fe 3+ ions, the interaction of iron (III) salts with potassium thiocyanate KNCS or ammonium thiocyanate NH 4 NCS is also used. In this case, a brightly colored FeNCNS 2+ ion is formed, as a result of which the entire solution acquires an intense red color:
Solubility table
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