General characteristics of metals. Metals: general characteristics of metals and alloys
Typical chemical properties of simple substances - metals
Most of the chemical elements are classified as metals - 92 out of 114 known elements. Metals- these are chemical elements, the atoms of which donate electrons of the outer (and some - and pre-outer) electronic layer, turning into positive ions. This property of metal atoms is determined by the fact 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 small values of electronegativity are characteristic(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 below the boron - astatine diagonal, as well as above it, in side subgroups. In the periods and main subgroups, the regularities known to you operate in the change of 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 side subgroups, the reducing properties of metals with an increase in the serial number most often decrease.
Compare the activity of the metals of the I group of the secondary subgroup known to you: Cu, Ag, Au; II group of a side subgroup: Zn, Cd, Hg - and you will see for yourself. This can be explained by the fact that the strength of the bond of valence electrons with the nucleus of atoms of these metals is more influenced by the magnitude of the nuclear charge, and not the radius of the atom. The magnitude of the nuclear charge increases significantly, the attraction of electrons to the nucleus increases. At the same time, the radius of the atom increases, but not as significantly as for metals of the main subgroups.
Simple substances formed by chemical elements - metals, and complex metal-containing substances play an important role in the mineral and organic "life" of the Earth. Suffice it to recall that the atoms (ions) of metal elements are an integral part of the compounds that determine the metabolism in the human body and animals. For example, 76 elements have been found in human blood, and only 14 of them are not metals.
In the human body, some elements, metals (calcium, potassium, sodium, magnesium) are present in large quantities, that is, they are macroelements. And metals such as chromium, manganese, iron, cobalt, copper, zinc, molybdenum are present in small quantities, that is, 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 lack, and with an excess.
For example, sodium ions regulate the water content in the body, the transmission of nerve impulses. Its deficiency leads to headaches, weakness, poor memory, loss of appetite, and excess leads to increased blood pressure, hypertension, heart disease.
Simple substances - metals
The development of the production of metals (simple substances) and alloys is associated with the emergence of civilization (Bronze Age, Iron Age). The scientific and technological revolution that began about 100 years ago, which affected both industry and the social sphere, is also closely related to the production of metals. On the basis of tungsten, molybdenum, titanium and other metals, they began to create corrosion-resistant, superhard, refractory alloys, the use of which greatly expanded the possibilities 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 are used from tantalum and platinum alloys, unique ceramics based on titanium and zirconium oxides.
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. Supernovae have become composite materials, representing, for example, a polymer or ceramics, which inside (like concrete with iron rods) are reinforced with metal fibers from tungsten, molybdenum, steel and other metals, and alloys - it all depends on the goal, the properties of the material necessary to achieve it. The figure shows a diagram of the crystal lattice of metallic sodium. In it, each sodium atom is surrounded by eight neighbors. The sodium atom, like all metals, has many free valence orbitals and few valence electrons. Electronic formula of sodium atom: 1s 2 2s 2 2p 6 3s 1 3p 0 3d 0, where 3s, 3p, 3d - valence orbitals.
The only valence electron of the sodium atom is 3s 1 can occupy any of the nine free orbitals - 3s (one), 3p (three) and 3d (five), because they are not very different in energy level. When atoms approach each other, when a crystal lattice is formed, the valence orbitals of neighboring atoms overlap, due to which the electrons move freely from one orbital to another, making a bond between all the atoms of the metal crystal. This chemical bond is called metallic.
A metal bond is formed by elements whose atoms on the outer layer have few valence electrons in comparison with a large number of outer energetically close orbitals. Their valence electrons are weakly held in the atom. The electrons that make the connection are socialized and move throughout the entire crystal lattice of the neutral metal as a whole. Metallic-bonded substances have metallic crystal lattices, which are usually depicted schematically as shown in the figure. Cations and metal atoms located in the nodes of the crystal lattice ensure its stability and strength (the socialized electrons are depicted as black small balls).
Metal bond- this is a bond in metals and alloys between metal atom-ions located in 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. The 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 ° С, ferromagnetic;
β - stable from 768 to 910 ° С, non-ferromagnetic, i.e. paramagnetic;
γ - stable from 910 to 1390 ° С, non-ferromagnetic, i.e. paramagnetic;
δ - stable from 1390 to 1539 ° C (£ ° pl of iron), non-ferromagnetic.
Tin has two crystalline modifications:
α - stable below 13.2 ° C (p = 5.75 g / cm 3). This is gray tin. It has a diamond-type crystal lattice (atomic);
β - stable above 13.2 ° C (p = 6.55 g / cm 3). This is white tin.
White tin is a silvery white very soft metal. When cooled below 13.2 ° C, it disintegrates into a gray powder, since during the transition, its specific volume increases significantly. This phenomenon has received the name "tin plague".
Of course, the special type of chemical bond and the type of crystal lattice of metals should determine and explain them. physical properties... What are they? These are metallic luster, plasticity, high electrical conductivity 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. Mechanical action on a crystal with a metal crystal lattice causes a displacement of the layers of ion-atoms relative to each other (Fig. 17), and since electrons move throughout the crystal, bond breakage does not occur, therefore, metals are characterized by high plasticity. A similar effect on a solid with covalent bonds (atomic crystal lattice) leads to the breaking of covalent bonds. The breaking of 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 plastic metals are Au, Ag, Sn, Pb, Zn. They are easily drawn into wire, amenable to forging, pressing, rolling into sheets. For example, gold foil can be made of 0.003 mm thick, and 0.5 g of this metal can be used to draw a thread 1 km long. Even mercury, which is liquid at room temperature, at low temperatures in the solid state it becomes malleable, like lead. Only Bi and Mn do not have plasticity, they are brittle.
Why do metals have a characteristic luster and are also opaque?
The electrons that fill the interatomic space reflect light rays (rather than transmit them, like glass), with most metals equally scattering all rays of the visible part of the spectrum. Therefore, they are silvery white or gray in color. Strontium, gold and copper absorb to a greater extent short wavelengths (close to violet) and reflect long waves of the light spectrum, therefore they have light yellow, yellow and "copper" colors. Although in practice the metal does not always seem to us a "light body". First, its surface can oxidize and lose its luster. Therefore, native copper looks like a greenish stone. A Secondly, and 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, palladium have the highest reflectivity. They are used in the manufacture of mirrors, including spotlights.
Why do metals have high electrical and thermal conductivity?
Chaotically moving electrons in a metal, under the influence of an applied electric voltage, acquire directional motion, that is, they conduct an electric current. With an increase in the temperature of the metal, the amplitudes of vibrations of atoms and ions located in the nodes of the crystal lattice increase. This makes it difficult for electrons to move, the electrical conductivity of the metal drops. At low temperatures, the vibrational motion, on the contrary, is greatly reduced and the electrical conductivity of metals increases sharply. Near absolute zero, there is practically no resistance in metals, and superconductivity appears in most metals.
It should be noted that non-metals with electrical conductivity (for example, graphite), on the contrary, do not conduct electric current at low temperatures due to the absence of free electrons. And only with an increase in temperature and the destruction of some covalent bonds, their electrical conductivity begins to increase. Silver, copper, as well as gold, aluminum have the highest electrical conductivity, manganese, lead, and mercury have the lowest electrical conductivity.
Most often, with the same regularity as electrical conductivity, the thermal conductivity of metals changes. 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 piece of metal.
Mechanical strength, density, melting point of metals are very different... Moreover, with an increase in the number of electrons that bind ion-atoms and a decrease in the interatomic distance in crystals, the indices of these properties increase.
So, alkali metals(Li, K, Na, Rb, Cs), whose atoms have one valence electron, soft (cut with a knife), with a low density (lithium is the lightest metal with p = 0.53 g / cm 3) and melt at low temperatures (for example, the melting point of cesium is 29 ° C). The only metal that is liquid under normal conditions, mercury, has a melting point of -38.9 ° C. Calcium, which has two electrons at the outer energy level of atoms, is much harder and melts at more high temperature(842 ° C). Even stronger is the crystal lattice formed by scandium ions, which has three valence electrons. But the strongest crystal lattices, high densities and melting points are observed in metals of side subgroups V, VI, VII, VIII of groups. This is due to the fact that metals of side subgroups with unpaired valence electrons on the d-sublevel are characterized by the formation of very strong covalent bonds between atoms, in addition to the metallic one, carried out by the electrons of the outer layer from the s-orbitals.
Heaviest metal- this is osmium (Os) with p = 22.5 g / cm 3 (a component of superhard and wear-resistant alloys), the most refractory metal is tungsten W with t = 3420 ° C (used to make lamp filaments), the hardest metal is it is chrome Cr (scratches glass). They are part of the materials from which metal-cutting tools, brake pads of heavy machines, etc. are made. Metals interact differently with the magnetic field. 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 the magnetic field - these are paramagnets. Metals pushed out magnetic field, - diamagnets (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 - lungs (p< 5 г/см 3) и тяжелые (все остальные).
2. Melting point - fusible and refractory.
There are chemical classifications of metals. Metals with low reactivity are called noble(silver, gold, platinum and its analogues - osmium, iridium, ruthenium, palladium, rhodium). By the proximity of chemical properties, they are distinguished alkaline(metals of the main subgroup of group I), alkaline earth(calcium, strontium, barium, radium) and rare earth metals(scandium, yttrium, lanthanum and lanthanides, anemones and actinides).
General chemical properties of metals
Metal atoms are relatively easy donate valence electrons and pass into positively charged ions, that is, they are oxidized. This is the main thing common property and atoms, and simple substances - metals. Metals in chemical reactions always restorers. The reducing ability of atoms of simple substances - metals, formed by chemical elements of one period or one main subgroup of DI Mendeleev's Periodic Table, changes naturally.
The reducing activity of a metal in chemical reactions that take place in aqueous solutions reflects its position in the electrochemical series of metal voltages.
Based on this series of stresses, the following important conclusions can be drawn about the chemical activity of metals in reactions taking place in aqueous solutions under standard conditions (t = 25 ° C, p = 1 atm).
· The more to the left the 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 in the series of voltages after it (to the right).
· Metals located in the series of voltages to the left of hydrogen are able to displace it from acids in solution.
· Metals, which are the strongest reducing agents (alkaline and alkaline earth), in any aqueous solution interact primarily with water.
The reducing activity of a metal, determined by the electrochemical series, does not always correspond to its position in the periodic table. This is due to the fact that when determining the position of a metal in a series of stresses, not only the energy of electron detachment from individual atoms is taken into account, but also the energy spent 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, according to its position in the periodic table, Na is a more active metal). The fact is that the hydration energy of Li + ions is much higher than the hydration energy of Na +, so the first process is energetically more favorable. Having considered the general provisions characterizing the reducing properties of metals, we turn to specific chemical reactions.
Interaction of metals with non-metals
· With oxygen, most metals form oxides- basic and amphoteric. Acid oxides of transition metals, for example, chromium (VI) oxide CrO g 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 actively react with atmospheric oxygen forming peroxides:
Sodium oxide is obtained indirectly by calcining peroxides with the corresponding metals:
Lithium and alkaline earth metals interact with atmospheric oxygen to form basic oxides:
Metals other than gold and platinum metals, which are not oxidized at all by atmospheric oxygen, interact with it less actively or when heated:
· With halogens, metals form salts of hydrohalic acids, For example:
· With hydrogen, the most active metals form hydrides- 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 - there is, as it were, dissolution or introduction of hydrogen 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.
· With gray metals form salts - sulfides, For example:
· Metals react with nitrogen a little more difficult. because the chemical bond in the nitrogen molecule N 2 is very strong; in this case, nitrides are formed. At ordinary temperatures, only lithium interacts 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, which stand in the series of voltages before hydrogen, can also displace hydrogen from water under certain conditions. But aluminum reacts violently with water only if the oxide film is removed from its surface:
Magnesium interacts with water only when boiling, while hydrogen is also released:
If the burning magnesium is added to water, then combustion continues, since the reaction proceeds:
Iron interacts with water only in a red-hot state:
· With acids in solution (HCl, H 2 SO 4 ), CH 3 COOH etc., except for HNO 3 ) interact metals, standing in the series of stresses up to hydrogen. This produces salt and hydrogen.
But lead (and some other metals), despite its position in the series of voltages (to the left of hydrogen), hardly dissolves 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 liberated in a free form.
It must be remembered that the reaction takes place in cases where the resulting salt is soluble. The displacement of metals from their compounds by other metals was first studied in detail by NN Beketov, a great Russian scientist in the field of physical chemistry. He arranged the metals according to their chemical activity in the "displacement series", which became the prototype of a number of metal stresses.
· With organic substances. Interaction with organic acids is similar to reactions with mineral acids. Alcohols, on the other hand, 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 (reaction of A. Würz):
· Metals, the hydroxides of which are amphoteric, interact with alkalis in solution. For instance:
· Metals can form chemical compounds with each other, which are collectively called intermetallic compounds. They most often do not show the oxidation states of atoms, which are characteristic of compounds of metals with non-metals. For instance:
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 side subgroups.
Metals of the main subgroups of groups I-III of the Periodic table of chemical elements of D. I. Mendeleev
general characteristics
These are metals of the main subgroup of group I. Their atoms on the outer energy level each have one electron. Alkali metals - strong reducing agents... Their reductive ability and reactivity increase with an increase in the ordinal number of the element (i.e., from top to bottom in the Periodic Table). All of them are electronically conductive. The bond strength between alkali metal atoms decreases with an increase in the ordinal number of the element. Their melting and boiling points are also reduced. Alkali metals interact with many simple substances - oxidants... In reactions with water, they form water-soluble bases (alkalis). Alkaline earth elements the elements of the main subgroup of group II are called. The atoms of these elements contain on the external energy level two electrons... They are the strongest restorers, have an oxidation state of +2. In this main subgroup, general laws are observed in the change in physical and chemical properties associated with an increase in the size of atoms in the group from top to bottom, and the chemical bond between atoms also weakens. With an increase in the size of the ion, the acidic properties weaken and the basic properties of oxides and hydroxides increase.
The main subgroup of group III is made up 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, with an increase in the charge of the nucleus, the metallic properties increase. Boron is a non-metallic element, while aluminum already has metallic properties. All elements form oxides and hydroxides.
Most of the metals are found in subgroups of the Periodic Table. Unlike the elements of the main subgroups, where there is a gradual filling of the outer level of atomic orbitals with electrons, the elements of the subsidiary subgroups are filled with the d-orbitals of the penultimate energy level and the s-orbitals of the latter. The number of electrons corresponds to the group number. Elements with an equal number of valence electrons are grouped under one number. All elements of the subgroups are metals.
Simple substances formed by metals of subgroups 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 valence from basic properties through amphoteric to acidic is clearly manifested.
Alkali metals (Na, K)
At the external energy level, the alkali metal atoms of the elements contain one electron at a time located at a great distance from the nucleus. They easily donate this electron, therefore 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, soft (cut with a knife), light and fusible. Actively interact with everyone non-metals:
All alkali metals react with oxygen (excluding Li) to form peroxides. Alkali metals do not occur in free form due to their high chemical activity.
Oxides- solids, have basic properties. They are obtained by calcining peroxides with the corresponding metals:
Hydroxides NaOH, KOH- solid white substances, hygroscopic, readily soluble in water with the release of heat, they are referred to as alkalis:
Almost all alkali metal salts are 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, edible salt.
Elements of group I in tables
Alkaline earth metals (Ca, Mg)
Calcium (Ca) is a representative alkaline earth metals, which are called the elements of the main subgroup of group II, but not all, but only starting with calcium and down the group. These are the chemical elements that, interacting with water, form alkalis. Calcium at the external energy level contains two electrons, oxidation state +2.
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. Soft metal, but its surface is covered in air protective film, which slightly reduces the chemical activity. Its burning is accompanied by a blinding flash. MgO and Mg (OH) 2 exhibit basic properties. Although Mg (OH) 2 is slightly soluble, it stains the phenolphthalein solution in a crimson color.
Mg + O 2 = MgO 2
MO oxides are solid white refractory substances. In the technique, 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 lime slaking, and the resulting Ca (OH) 2 is called slaked lime. A clear solution of calcium hydroxide is called lime water, and a white suspension of Ca (OH) 2 in water is called lime milk.
Magnesium and calcium salts are obtained by their interaction with acids.
CaCO 3 - calcium carbonate, chalk, marble, limestone. It is used in construction. MgCO 3 - magnesium carbonate - is used in metallurgy to remove slags.
CaSO 4 2H 2 O - gypsum. MgSO 4 - magnesium sulfate - is called bitter, or Epsom, salt, found in sea water. BaSO 4 - barium sulfate - due to its insolubility and ability to block X-rays, it is used in the diagnosis ("barite porridge") of the gastrointestinal tract.
Calcium accounts for 1.5% of human body weight, 98% of calcium is found in bones. Magnesium is a bioelement, there are 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)- an element of the main subgroup of the III group of the periodic system of D. I. Mendeleev. The aluminum atom contains on the external energy level three electrons, which he easily gives up during chemical interactions. In the ancestor of the subgroup and the upper neighbor of aluminum - boron - the radius of the atom is smaller (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 extension of the outer electrons to 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 a 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 fine aluminum foil heat up strongly, then they ignite and burn with a blinding flame:
This reaction can be observed when burning sparklers and fireworks. Aluminum, like all metals, reacts easily with non-metals, especially in a powdery state. In order for the reaction to start, initial heating is necessary, with the exception of reactions with halogens - chlorine and bromine, but then all reactions of aluminum with non-metals go 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 course 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 obtain metals - chromium, manganese, vanadium, titanium, zirconium from their oxides. This method is called alumothermy. In practice, thermite is often used - a mixture of Fe 3 O 4 with aluminum powder. If this mixture is ignited, for example, with a magnesium tape, then a vigorous reaction occurs with the release a large number warmth:
The released heat is quite enough for the complete melting of the formed iron, therefore 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 constituent parts using an electric current. But the melting point of aluminum oxide is about 2050 ° C, therefore, high energy costs are required for electrolysis.
Aluminum compounds
Aluminosilicates... These compounds can be considered as salts formed by oxide of aluminum, silicon, alkali and alkaline earth metals. They constitute the bulk of the earth's crust. In particular, aluminosilicates are found in feldspars, the most common minerals and clays.
Bauxite- the rock from which aluminum is obtained. It contains aluminum oxide Al 2 O 3.
Corundum- a mineral of composition Al 2 O 3, has a 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.
Well-known transparent, colored by impurities, crystals of corundum: 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 decomposition on heating aluminum hydroxide:
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 called salts of unstable aluminum acids - ortho-aluminum H 2 AlO 3, meta-aluminum HAlO 2 (it can be considered as ortho-aluminum acid, from the molecule of which the water molecule was taken away). Natural aluminates include noble spinel and precious chrysoberyl. Aluminum salts, except for phosphates, are readily 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.
Elements of group III in tables
Characterization of transition elements - copper, zinc, chromium, iron
Copper (Cu)- an element of a secondary subgroup of the first group. Electronic formula: (… 3d 10 4s 1). Its tenth d-electron is mobile, since 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, viscous, viscous, an excellent conductor of electricity. Melting point 1083 ° C.
Like other metals of subgroup I of group I of the periodic system, copper stands in the row of activity to the right of hydrogen and does not displace it from acids, but reacts with oxidizing acids:
Under the action of alkalis on solutions of copper salts, a precipitate of a weak base of blue color- copper (II) hydroxide, which, when heated, decomposes into a basic black oxide CuO and water:
Chemical properties copper in tables
Zinc (Zn)- an 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 completed, zinc in the compounds exhibits an oxidation state of +2.
Zinc is a silvery-white metal that practically does not change in air. It has corrosion resistance 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 amphoterine, that is, it exhibits the properties of both acid and base. With the gradual addition of the alkali solution to the zinc salt solution, the initially precipitated precipitate dissolves (the same happens with aluminum):
Zinc chemical properties in tables
For example chromium (Cr) it can be shown that properties of transition elements change along the period not fundamentally: there is a quantitative change associated with a change in the number of electrons in the valence orbitals. The maximum oxidation state of chromium is +6. The metal in the row of activity 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 rapidly oxidized by atmospheric oxygen:
It corresponds to 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 Cr O 4 ) in an acidic medium turn into dichromates- dichromic acid salts (H 2 Cr 2 O 7). Chromium compounds are highly oxidizing.
Chemical properties of chromium in tables
Iron Fe- an element of a side subgroup of the VIII group and the 4th period of the periodic system of D. I. Mendeleev. Iron atoms are arranged 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 not the last one is filled, but the penultimate, third from the nucleus, level. At the last level, iron atoms contain two electrons. At the penultimate level, which can hold 18 electrons, the iron atom has 14 electrons. Consequently, the distribution of electrons over levels in iron atoms is as follows: 2e; 8e; 14e; 2e. Like all metals iron atoms exhibit reducing properties, donating during chemical interactions not only two electrons from the last level, and acquiring the oxidation state +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 easily processed, forged, rolled, stamped. Iron has the ability to magnetize and demagnetize. It can be given greater strength and hardness by methods of thermal and mechanical action. Distinguish between technically pure and chemically pure iron. Technically pure iron, in fact, is a 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 of technically pure iron. Such iron corrodes easily, while chemically pure iron hardly corrodes. At present, iron is the basis of modern technology and agricultural engineering, transport and communications, spacecraft and, in general, all modern civilization. Most products, from sewing needles to spacecraft, cannot be made without the use of iron.
Iron chemical properties
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 in 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 series of voltages, it can displace the metals that are to the right of it from aqueous solutions their salts, for example:
Iron dissolves in dilute hydrochloric and sulfuric acids, i.e. it is oxidized by hydrogen ions:
Iron dissolves in dilute nitric acid, in this case, iron (III) nitrate, water and nitric acid reduction products - N 2, NO or NH 3 (NH 4 NO 3) are formed, depending on the concentration of the acid.
Iron compounds
In nature, iron forms a number of minerals. These are 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 an 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 the corresponding iron (II) hydroxide Fe (OH) 2 are obtained indirectly, in particular, through the following transformation chain:
Both compounds have pronounced basic properties.
Iron (II) cations Fe 2 + easily oxidized by atmospheric oxygen to cations of iron (III) Fe 3 + ... Therefore, the white precipitate of iron (II) hydroxide becomes 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 the greatest technical importance.
Crystalline hydrate of iron sulfate (II) FeSO 4 7H 2 O, known as ferrous sulfate, is used to combat plant pests, to prepare mineral paints and for other purposes. Iron (III) chloride FeCl 3 is used as a mordant for dyeing fabrics. Iron (III) sulfate Fe 2 (SO 4) 3 9H 2 O is used for water purification and for other purposes.
The physical and chemical properties of iron and its compounds are summarized in the table:
Iron chemical properties in tables
Qualitative reactions for Fe 2+ and Fe 3+ ions
For the recognition of iron (II) and (III) compounds carry out qualitative reactions for Fe ions 2+ and Fe 3+ ... A qualitative reaction to Fe 2+ ions is the reaction of iron (II) salts with a compound K 3 called red blood salt. This is a special group of salts, which are called complex salts, you will get acquainted with them in the future. In the meantime, you need to learn how such salts dissociate:
The reagent for Fe 3+ ions is another complex compound - yellow blood salt - K 4, which dissociates in solution in the same way:
If, respectively, solutions of red blood salt (reagent for Fe 2+) and yellow blood salt (reagent for Fe 3+) are added to solutions containing ions Fe 2+ and Fe 3+, then in both cases the same blue precipitate will form:
To detect Fe 3+ ions, the interaction of iron (III) salts with potassium thiocyanate KNCS or ammonium 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
Due to the presence of free electrons ("electron gas") in the crystal lattice, all metals exhibit the following characteristic general properties:
1) Plastic- the ability to easily change shape, be drawn into wire, rolled into thin sheets.
2) Metallic luster and opacity. This is due to the interaction of free electrons with light incident on the metal.
3) Electrical conductivity... It is explained by the directional movement of free electrons from the negative to the positive pole under the influence of a small potential difference. When heated, the electrical conductivity decreases, because with an increase in temperature, the 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 there is a rapid equalization of temperature over the mass of the metal. Bismuth and mercury have the highest thermal conductivity.
5) Hardness. The hardest is chrome (cuts glass); the softest - alkali metals - potassium, sodium, rubidium and cesium - are cut with a knife.
6) Density. It is the less, the less atomic mass metal and larger radius of the atom. The lightest is lithium (ρ = 0.53 g / cm3); the heaviest is osmium (ρ = 22.6 g / cm3). Metals with a density of less than 5 g / cm3 are considered “light metals”.
7) Melting and boiling points. The lowest-melting metal is mercury (melting point = -39 ° C), the most refractory metal is tungsten (melting point = 3390 ° C). Metals with t ° pl. above 1000 ° C are considered refractory, below - low melting.
General chemical properties of metals
Strong reducing agents: Me 0 - nē → Me n +
A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions.
1. Reactions of metals with non-metals
1) With oxygen:
2Mg + O 2 → 2MgO
2) With gray:
Hg + S → HgS
3) With halogens:
Ni + Cl 2 - t ° → NiCl 2
4) With nitrogen:
3Ca + N 2 - t ° → Ca 3 N 2
5) With phosphorus:
3Ca + 2P - t ° → Ca 3 P 2
6) With hydrogen (only alkali and alkaline earth metals react):
2Li + H 2 → 2LiH
Ca + H 2 → CaH 2
2. Reactions of metals with acids
1) Metals in the electrochemical series of voltages up to H reduce non-oxidizing acids to hydrogen:
Mg + 2HCl → MgCl 2 + H 2
2Al + 6HCl → 2AlCl 3 + 3H 2
6Na + 2H 3 PO 4 → 2Na 3 PO 4 + 3H 2
2) With oxidizing acids:
With the interaction of nitric acid of any concentration and concentrated sulfuric with metals hydrogen is never released!
Zn + 2H 2 SO 4 (К) → ZnSO 4 + SO 2 + 2H 2 O
4Zn + 5H 2 SO 4 (К) → 4ZnSO 4 + H 2 S + 4H 2 O
3Zn + 4H 2 SO 4 (К) → 3ZnSO 4 + S + 4H 2 O
2H 2 SO 4 (k) + Cu → Cu SO 4 + SO 2 + 2H 2 O
10HNO 3 + 4Mg → 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O
4HNO 3 (c) + Cu → Cu (NO 3) 2 + 2NO 2 + 2H 2 O
3. Interaction of metals with water
1) Active (alkali and alkaline earth metals) form a soluble base (alkali) and hydrogen:
2Na + 2H 2 O → 2NaOH + H 2
Ca + 2H 2 O → Ca (OH) 2 + H 2
2) Metals of medium activity are oxidized by water when heated to oxide:
Zn + H 2 O - t ° → ZnO + H 2
3) Inactive (Au, Ag, Pt) - do not react.
4. Displacement of less active metals from solutions of their salts by more active metals:
Cu + HgCl 2 → Hg + CuCl 2
Fe + CuSO 4 → Cu + FeSO 4
In industry, not pure metals are often used, but their mixtures - alloys, in which the beneficial properties of one metal are complemented by the beneficial properties of another. So, copper has a low hardness and is of little use for the manufacture of machine parts, while copper-zinc alloys ( brass) are already quite solid and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but too soft. On its basis, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing useful properties aluminum, acquires high hardness and becomes suitable in aircraft construction. Alloys of iron with carbon (and additives of other metals) are widely known cast iron and steel.
Free metals are reducing agents. However, the reactivity of some metals is low due to the fact that they are coated surface oxide film, in varying degrees, resistant to the action of chemicals such as water, solutions of acids and alkalis.
For example, lead is always covered with an oxide film; for its transition into solution, not only the action of a reagent (for example, dilute nitric acid) is required, but also heating. The oxide film on aluminum prevents it from reacting with water, but is destroyed by acids and alkalis. Loose oxide film (rust), formed on the surface of iron in humid air, does not interfere with further oxidation of iron.
Under the influence concentrated acids on metals are formed steady oxide film. This phenomenon is called passivation... So, in concentrated sulfuric acid metals such as Be, Bi, Co, Fe, Mg and Nb are passivated (and then do not react with acid), and metals A1, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb in concentrated nitric acid , Th and U.
When interacting with oxidants in acidic solutions, most metals are converted into cations, the charge of which is determined by the stable oxidation state of 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 voltages. Most of the metals are converted into a solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only sulfuric (concentrated) and nitric acids, and Pt and Au - "aqua regia".
Corrosion of metals
An undesirable chemical property of metals is their corrosion, i.e., active destruction (oxidation) upon contact with water and under the influence of oxygen dissolved in it. (oxygen corrosion). For example, corrosion of iron products in water is widely known, as a result of which rust is formed and the products are crumbled into powder.
Corrosion of metals occurs in water also due to the presence of dissolved gases CO 2 and SO 2; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 ( hydrogen corrosion).
The place of contact of two dissimilar metals ( contact corrosion). A galvanic pair arises between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the series of voltages (Fe), 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 (tin-coated iron) rusts when stored in a humid atmosphere and carelessly handling them (iron quickly collapses after the appearance of at least a small scratch that allows iron to come into contact with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, because even in the presence of scratches, it is not iron that corrodes, but zinc (a more active metal than iron).
Corrosion resistance for a given metal is enhanced when it is coated with a more active metal or when they are fused; thus, plating iron with chromium or making an iron-chromium alloy eliminates the corrosion of iron. Chromium-plated iron and steel containing chromium ( stainless steel), have high corrosion resistance.
1. Metals react with non-metals.
2 Me + n Hal 2 → 2 MeHal n
4Li + O2 = 2Li2O
Alkali metals, with the exception of lithium, form peroxides:
2Na + O 2 = Na 2 O 2
2. Metals standing up to hydrogen react with acids (except for nitric and sulfuric conc.) With the release of hydrogen
Me + HCl → salt + H2
2 Al + 6 HCl → 2 AlCl3 + 3 H2
Pb + 2 HCl → PbCl2 ↓ + H2
3. Active metals react with water to form alkali and release hydrogen.
2Me + 2n H 2 O → 2Me (OH) n + n H 2
The product of metal oxidation is its hydroxide - Me (OH) n (where n is the oxidation state of the metal).
For example:
Ca + 2H 2 O → Ca (OH) 2 + H 2
4. Metals of medium activity react with water when heated to form metal oxide and hydrogen.
2Me + nH 2 O → Me 2 O n + nH 2
The oxidation product in such reactions is the metal oxide Me 2 O n (where n is the oxidation state of the metal).
3Fe + 4H 2 O → Fe 2 O 3 FeO + 4H 2
5. Metals after hydrogen do not react with water and acid solutions (except for nitric and sulfuric conc.)
6. More active metals displace less active metals from solutions of their salts.
CuSO 4 + Zn = Zn SO 4 + Cu
CuSO 4 + Fe = Fe SO 4 + Cu
Active metals - zinc and iron replaced copper in sulfate and formed salts. Zinc and iron were oxidized, and copper was reduced.
7. Halogens react with water and alkali solution.
Fluorine, unlike other halogens, oxidizes water:
2H 2 O + 2F 2 = 4HF + O 2 .
in the cold: Cl2 + 2KOH = KClO + KCl + H2OCl2 + 2KOH = KClO + KCl + H2O chloride and hypochlorite are formed
when heated: 3Cl2 + 6KOH− → KClO3 + 5KCl + 3H2O3Cl2 + 6KOH → t, ∘CKClO3 + 5KCl + 3H2O loride and chlorate are formed
8 Active halogens (except for fluorine) displace less active halogens from solutions of their salts.
9. Halogens do not react with oxygen.
10. Amphoteric metals (Al, Be, Zn) react with solutions of alkalis and acids.
3Zn + 4H2SO4 = 3 ZnSO4 + S + 4H2O
11. Magnesium reacts with carbon dioxide and silicon oxide.
2Mg + CO2 = C + 2MgO
SiO2 + 2Mg = Si + 2MgO
12. Alkali metals (except lithium) form peroxides with oxygen.
2Na + O 2 = Na 2 O 2
3. Classification of inorganic compounds
Simple substances - substances whose molecules consist of atoms of one type (atoms of one element). In chemical reactions, they cannot decompose with the formation of other substances.
Complex substances (or chemical compounds) - substances whose molecules consist of atoms of different types (atoms of various chemical elements). In chemical reactions, they decompose to form several other substances.
Simple substances are divided into two large groups: metals and non-metals.
Metals - a group of elements with characteristic metallic properties: solid substances (with the exception of mercury) have a metallic luster, are good conductors of heat and electricity, malleable (iron (Fe), copper (Cu), aluminum (Al), mercury (Hg), gold (Au), silver (Ag), etc.).
Nonmetals - a group of elements: solid, liquid (bromine) and gaseous substances that do not have a metallic luster, are insulators, fragile.
And complex substances, in turn, are subdivided into four groups, or classes: oxides, bases, acids and salts.
Oxides - these are complex substances, the composition of the molecules of which includes atoms of oxygen and some other substance.
Foundations Are complex substances in which metal atoms are combined with one or more hydroxyl groups.
From the point of view of the theory of electrolytic dissociation, bases are complex substances, the dissociation of which in an aqueous solution forms metal cations (or NH4 +) and hydroxide - OH- anions.
Acid - These are complex substances, the molecules of which include hydrogen atoms that can be replaced or exchanged for metal atoms.
Salt Are complex substances, the molecules of which are composed of metal atoms and acidic residues. Salt is a product of partial or complete substitution of a metal for hydrogen atoms of an acid.
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 donating electrons during chemical reactions, that is, being reducing agents, showing 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 chemical element in conjunction.
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 of their salts, metals are located in the electrochemical series of voltages (activities) of metals:
Li → Rb → K → Ba → Sr → Ca → Na → Mg → Al → Mn → Zn → Cr → → Fe → Cd → Co → Ni → Sn → Pb → H → Sb → Bi → Cu → Hg → Ag → Pd → Pt → Au
The metals located to the left in this row are more active and are able to displace the following metals from salt solutions.
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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CHEMICAL PROPERTIES OF METALS
According to their chemical properties, metals are divided 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