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 voteIron and oxygen form three oxides
2Fe + O 2 ↔ 2FeO (nitrous oxide containing 22.7% O 2);
6FeO + O 2 ↔ 2Fe 3 O 4 (oxide-oxide containing 27.64% O 2);
4Fe 3 O 4 + O 2 ↔ 6Fe 2 O 3 (oxide containing 30.06% O 2).
Of these three oxides, only FeO oxide is soluble in iron and therefore most strongly affects its properties. The remaining oxides do not dissolve in iron; they can be present in it in the form of separate inclusions and easily decompose at high temperatures. At the melting point of iron, the maximum solubility of oxygen in it is 0.17%, and at room temperature - thousandths of a percent.
Sources of metal oxidation during welding are:
1. Free oxygen in the gas phase (air oxygen with insufficient protection; the presence of complex gases CO 2, H 2 O, capable of releasing oxygen upon dissociation).
2. oxides located on the molten edges of the metal being welded and on the filler material.
3.Oxides present in the slag and soluble in the metal.
4. Chemically active slags that release oxygen to the metal as a result of exchange redox reactions.
Oxidation of metal with free gas phase oxygen occurs according to the reaction nMe + O 2 = mMe n / m O 2/ m. If the metal and oxide are condensed phases (solid or liquid), then the equilibrium constant K p of such an oxidation reaction is determined only by the partial pressure of oxygen p O2, corresponding to the elasticity of oxide dissociation at a given temperature and pressure, i.e. K p = p o2. The dependence log p o 2 = f(T) for different temperatures is given in Fig. 4.5.
The lower the elasticity of dissociation, the greater the strength of the oxide. Since the elasticity of dissociation of the oxides located in the upper part of Fig. 4.5, at a specific temperature there are more than those located below, their affinity for oxygen is less (affinity for oxygen is judged by the amount of work that must be spent on destroying the oxide of this element). Therefore, when the metal Me 1 with a high affinity for oxygen comes into contact with the oxide of another metal Me 2, which has a lower affinity for oxygen, a reaction is possible
Me 1 + Me 2 O = Me 1 O + Me 2.
On this basis, deoxidation is carried out in welding processes, and the element Me 1 in relation to the element Me 2 is a deoxidizer.
Rice. 4.5. Dependence of the elasticity of oxide dissociation on temperature.
Let us rank the metals in order of decreasing affinity for oxygen:
Cu, Ni, Fe, Mo, Cr, Mn, Si, Ti, Al, Mg, Ca, C (at high temperatures).
The possibility of oxidation or reduction of any element as a result of its interaction with a gas phase containing free oxygen under specific external conditions (when welding at different temperatures) is determined by comparing the elasticity of oxide dissociation p O2 (MeO) = p O2 and the partial pressure of free oxygen in gas phase [p O2]. If the pressure [p O2 ] p O2 is greater than the elasticity of oxide dissociation, then oxidation will occur; at [p O2]< р О2 – восстановление.
Elasticity of dissociation of oxides in solution, p! O2 differs from the elasticity of dissociation of free oxides p O2. Wherein
The denominator is the maximum saturation of the metal solution with oxide.
The air contains approximately 20% oxygen i.e. [p O2] = 0.2 kgf/cm 2 and it is a strong oxidizing agent for iron under welding conditions [p O2] p O2.
Oxidation of metal in the melting zone by surface oxides carried out by remelting the oxides located on the edges and on the surface of the filler metal. When the edges of the base metal, the surface of which is covered with oxides, melt, the additional amounts of oxygen they introduce enter the weld pool, leading to greater oxidation of the weld pool. Oxygen is introduced similarly in the case of the presence of oxides on the filler metal.
In order to minimize the increase in oxidation of the pool and weld metal by such oxides, the surfaces of the base metal that are subject to melting during welding must first be cleaned mechanically. The filler wire is cleaned of oxides either mechanically or by etching.
Oxidation of metal by oxides present in the slag and soluble in the metal, occurs due to the redistribution of such oxides between the slag and metal. Such free oxide tends to be distributed between the metal and slag phases, determined by the distribution constant L MeO = (MeO)/, where [MeO] and (MeO) are the concentrations of these oxides in the metal and slag, respectively. This constant changes with temperature. For FeO this dependence is expressed by the formula
Lg1/L FeO = log/(FeO) = -(6300/T) + 1.386.
When the temperature changes from Tm of iron to 2500 0 C, the value of the constant increases from 0.011 to 0.125.
Oxidation with oxygen-reactive slags occurs in connection with exchange reactions such as
(Meh!! x O y) + y[Meh! ] = y(Me! O) + x[Me!! ].
When welding carbon steels according to this scheme, silicon and manganese reduction processes occur in the presence of significant quantities of SiO 2 and MnO in the slag. If the metal contains elements with a stronger affinity for oxygen, their oxidation by silica and manganese oxide can be very intense. When welding steels containing elements with a very high affinity for oxygen (Al, Ti), their burnout can occur almost completely.
The oxidation of liquid metal in the welding zone depends on the content of elements in it - deoxidizers. Deoxidizers are elements with a greater affinity for oxygen than the metal that is the basis of the alloy. Nickel cannot be a deoxidizer for iron and its burnout when welding iron-based alloys should be negligible. Manganese, even at concentrations of more than 0.5% at 2300 0 C and about 0.1% at 1540 0 C, has a lower elasticity of oxide dissociation than oxygen-saturated iron, and can act as a deoxidizer that removes oxygen from the iron base. Chromium is weaker than manganese. Stronger iron deoxidizers are silicon, titanium, and aluminum. At high temperatures, carbon becomes the most powerful deoxidizing agent.
Due to the fact that the affinity of elements for oxygen decreases with increasing temperature, the concentration of oxygen in the molten metal at high temperatures can be significant. As the temperature in the tail part of the bath decreases, the deoxidizing ability of the deoxidizing elements increases and the reactions shift towards the binding of oxygen by these elements. The products of these reactions, being practically insoluble in the metal, precipitate as a separate phase. This deoxidation is called besieger.
Particles of precipitated oxides can be removed by floating or being pushed out by growing crystals or remain in the solidified metal in the form of slag inclusions.
The products of carbon deoxidation are gaseous. When they are released, the tail part of the bath bubbles (boils), and the bubbles, which have not had time to leave the solidified metal, form pores filled with gas in the metal.
Restorative properties- these are the main chemical properties characteristic of all metals. They manifest themselves in interaction with a wide variety of oxidizing agents, including oxidizing agents from the environment. In general, the interaction of a metal with oxidizing agents can be expressed by the following scheme:
Me + Oxidizing agent" Me(+X),
Where (+X) is the positive oxidation state of Me.
Examples of metal oxidation.
Fe + O 2 → Fe(+3) 4Fe + 3O 2 = 2 Fe 2 O 3
Ti + I 2 → Ti(+4) Ti + 2I 2 = TiI 4
Zn + H + → Zn(+2) Zn + 2H + = Zn 2+ + H 2
Metal activity series
The reducing properties of metals differ from each other. Electrode potentials E are used as a quantitative characteristic of the reduction properties of metals.
The more active the metal, the more negative its standard electrode potential E o.
Metals arranged in a row as their oxidative activity decreases form an activity series.
Metal activity series
| Me | Li | K | Ca | Na | Mg | Al | Mn | Zn | Cr | Fe | Ni | Sn | Pb | H 2 | Cu | Ag | Au |
| Me z+ | Li+ | K+ | Ca2+ | Na+ | Mg 2+ | Al 3+ | Mn 2+ | Zn 2+ | Cr 3+ | Fe 2+ | Ni 2+ | Sn 2+ | Pb 2+ | H+ | Cu 2+ | Ag+ | Au 3+ |
| E o ,B | -3,0 | -2,9 | -2,87 | -2,71 | -2,36 | -1,66 | -1,18 | -0,76 | -0,74 | -0,44 | -0,25 | -0,14 | -0,13 | 0 | +0,34 | +0,80 | +1,50 |
The reduction of a metal from a solution of its salt with another metal with higher reducing activity is called cementation. Cementation is used in metallurgical technologies.
In particular, Cd is obtained by reducing it from a solution of its salt with zinc.
Zn + Cd 2+ = Cd + Zn 2+
3.3. 1. Interaction of metals with oxygen
Oxygen is a strong oxidizing agent. It can oxidize the vast majority of metals exceptAuAndPt . Metals exposed to air come into contact with oxygen, so when studying the chemistry of metals, one always pays attention to the peculiarities of the interaction of the metal with oxygen.
Everyone knows that iron in humid air becomes covered with rust - hydrated iron oxide. But many metals in a compact state at not too high temperatures exhibit resistance to oxidation, since they form thin protective films on their surface. These films of oxidation products prevent the oxidizing agent from contacting the metal. The phenomenon of formation of protective layers on the surface of a metal that prevent oxidation of the metal is called passivation of the metal.
An increase in temperature promotes the oxidation of metals with oxygen. The activity of metals increases in a finely crushed state. Most metals in powder form burn in oxygen.
s-metals
Show the greatest reducing activitys-metals. Metals Na, K, Rb Cs can ignite in air, and they are stored in sealed vessels or under a layer of kerosene. Be and Mg are passivated at low temperatures in air. But when ignited, the Mg tape burns with a blinding flame.
MetalsIIA-subgroups and Li, when interacting with oxygen, form oxides.
2Ca + O2 = 2CaO
4 Li + O 2 = 2 Li 2 O
Alkali metals, exceptLi, when interacting with oxygen, they form not oxides, but peroxidesMe 2 O 2 and superoxidesMeO 2 .
2Na + O 2 = Na 2 O 2
K + O 2 = KO 2
p-metals
Metals belonging top- the block is passivated in air.
When burning in oxygen
- metals of the IIIA subgroup form oxides of the type Me 2 O 3,
- Sn is oxidized to SnO 2 , and Pb - up to PbO
- Bi goes to Bi2O3.
d-metals
Alld-period 4 metals are oxidized by oxygen. Sc, Mn, Fe are most easily oxidized. Particularly resistant to corrosion are Ti, V, Cr.
When burned in oxygen of alld
When burned in oxygen of alld-of period 4 elements, only scandium, titanium and vanadium form oxides in which Me is in the highest oxidation state, equal to group number. The remaining period 4 d-metals, when burned in oxygen, form oxides in which Me is in intermediate but stable oxidation states.
Types of oxides formed by period 4 d-metals upon combustion in oxygen:
- MeO form Zn, Cu, Ni, Co. (at T>1000°C Cu forms Cu 2 O),
- Me 2 O 3, form Cr, Fe and Sc,
- MeO 2 - Mn, and Ti,
- V forms a higher oxide - V 2 O 5 .
When burned in oxygend-metals of periods 5 and 6, as a rule, form higher oxides, the exceptions are the metals Ag, Pd, Rh, Ru.
Types of oxides formed by d-metals of periods 5 and 6 during combustion in oxygen:
- Me 2 O 3- form Y, La; Rh;
- MeO 2- Zr, Hf; Ir:
- Me 2 O 5- Nb, Ta;
- MeO 3- Mo, W
- Me 2 O 7- Tc, Re
- MeO 4 - Os
- MeO- Cd, Hg, Pd;
- Me 2 O- Ag;
Interaction of metals with acids
In acid solutions, the hydrogen cation is an oxidizing agent. The H+ cation can oxidize metals in the activity series up to hydrogen, i.e. having negative electrode potentials.
Many metals, when oxidized, transform into cations in acidic aqueous solutionsMe z + .
Anions of a number of acids are capable of exhibiting oxidizing properties that are stronger than H +. Such oxidizing agents include anions and the most common acids H 2 SO 4 AndHNO 3 .
NO 3 - anions exhibit oxidizing properties at any concentration in solution, but the reduction products depend on the concentration of the acid and the nature of the metal being oxidized.
SO 4 2- anions exhibit oxidizing properties only in concentrated H 2 SO 4.
Reduction products of oxidizing agents: H + , NO 3 - , SO 4 2 -
2Н + + 2е - =H 2
SO 4
2-
from concentrated H 2 SO 4 SO 4
2-
+ 2e -
+ 4
H +
=
SO 2
+ 2
H 2
O
(formation of S, H 2 S is also possible)
NO 3 - from concentrated HNO 3 NO 3 - + e -
+ 2H + =
NO 2 + H 2 O
NO 3 - from dilute HNO 3 NO 3 - + 3e -
+4H+=NO+2H2O
(formation of N 2 O, N 2, NH 4 + is also possible)
Examples of reactions between metals and acids
Zn + H 2 SO 4 (diluted) " ZnSO 4 + H 2
8Al + 15H 2 SO 4 (k.) " 4Al 2 (SO 4) 3 + 3H 2 S + 12H 2 O
3Ni + 8HNO 3 (dil.) " 3Ni(NO 3) 2 + 2NO + 4H 2 O
Cu + 4HNO 3 (k.) " Cu(NO 3) 2 + 2NO 2 + 2H 2 O
Products of metal oxidation in acidic solutions
Alkali metals form a Me + type cation, s-metals of the second group form cations Me 2+.
When dissolved in acids, p-block metals form the cations indicated in the table.
The metals Pb and Bi are dissolved only in nitric acid.
| Me | Al | Ga | In | Tl | Sn | Pb | Bi |
| Mez+ | Al 3+ | Ga 3+ | In 3+ | Tl+ | Sn 2+ | Pb 2+ | Bi 3+ |
| Eo,B | -1,68 | -0,55 | -0,34 | -0,34 | -0,14 | -0,13 | +0,317 |
All d-metals of 4 periods, except Cu , can be oxidized by ionsH+ in acidic solutions.
Types of cations formed by period 4 d-metals:
- Me 2+(form d-metals ranging from Mn to Cu)
- Me 3+ ( form Sc, Ti, V, Cr and Fe in nitric acid).
- Ti and V also form cations MeO 2+
In acidic solutions, H + can oxidize: Y, La, Cd.
The following can dissolve in HNO 3: Cd, Hg, Ag. Pd, Tc, Re dissolve in hot HNO 3.
The following dissolve in hot H 2 SO 4: Ti, Zr, V, Nb, Tc, Re, Rh, Ag, Hg.
Metals: Ti, Zr, Hf, Nb, Ta, Mo, W are usually dissolved in a mixture of HNO 3 + HF.
In aqua regia (a mixture of HNO 3 + HCl) Zr, Hf, Mo, Tc, Rh, Ir, Pt, Au and Os can be dissolved with difficulty). The reason for the dissolution of metals in aqua regia or in a mixture of HNO 3 + HF is the formation of complex compounds.
Example. The dissolution of gold in aqua regia becomes possible due to the formation of a complex -
Au + HNO 3 + 4HCl = H + NO + 2H 2 O
Interaction of metals with water
The oxidizing properties of water are due to H(+1).
2H 2 O + 2e -" N 2 + 2OH -
Since the concentration of H + in water is low, its oxidizing properties are low. Metals can dissolve in water E< - 0,413 B. Число металлов, удовлетворяющих этому условию, значительно больше, чем число металлов, реально растворяющихся в воде. Причиной этого является образование на поверхности большинства металлов плотного слоя оксида, нерастворимого в воде. Если оксиды и гидроксиды металла растворимы в воде, то этого препятствия нет, поэтому щелочные и щелочноземельные металлы энергично растворяются в воде. Alls-metals, except Be and Mg easily dissolve in water.
2 Na + 2 HOH = H 2 + 2 OH -
Na reacts vigorously with water, releasing heat. The released H2 may ignite.
2H 2 +O 2 =2H 2 O
Mg dissolves only in boiling water, Be is protected from oxidation by an inert insoluble oxide
P-block metals are less powerful reducing agents thans.
Among p-metals, the reducing activity is higher in metals of the IIIA subgroup, Sn and Pb are weak reducing agents, Bi has Eo > 0.
p-metals do not dissolve in water under normal conditions. When the protective oxide is dissolved from the surface in alkaline solutions with water, Al, Ga and Sn are oxidized.
Among d-metals, they are oxidized by water when Sc and Mn, La, Y are heated. Iron reacts with water vapor.
Interaction of metals with alkali solutions
In alkaline solutions, water acts as an oxidizing agent..
2H 2 O + 2e - =H 2 + 2OH - Eo = - 0.826 B (pH = 14)
The oxidizing properties of water decrease with increasing pH, due to a decrease in H + concentration. Nevertheless, some metals that do not dissolve in water dissolve in alkali solutions, for example, Al, Zn and some others. The main reason for the dissolution of such metals in alkaline solutions is that the oxides and hydroxides of these metals exhibit amphotericity and dissolve in alkali, eliminating the barrier between the oxidizing agent and the reducing agent.
Example. Dissolution of Al in NaOH solution.
2Al + 3H 2 O + 2NaOH + 3H 2 O = 2Na + 3H 2
Oxygen atoms can form two types of molecules: O 2 - oxygen and O 3 - ozone.
The phenomenon of the existence of several simple substances formed by atoms of one chemical element is called alotropy. And simple substances formed by one element are called alotropic modifications.
Consequently, ozone and oxygen are allotropic modifications of the element Oxygen.
|
Properties |
Oxygen |
Ozone |
|
Compound Formula |
O2 |
O 3 |
|
Appearance under normal conditions |
Gas |
Gas |
|
Color |
Oxygen in vapor is colorless. Liquid is pale blue and solid is blue |
Ozone vapor is light blue. The liquid is blue, and the solid is dark purple crystals. |
|
Smell and taste |
Odorless and tasteless |
Pungent characteristic odor (in small concentrations gives the air a fresh smell) |
|
Melting temperature |
219 °C |
192 °C |
|
Boiling temperature |
183 °C |
112 °C |
|
Density at n. u. |
1.43 g/l |
2.14 g/l |
|
Solubility |
4 volumes of oxygen in 100 volumes of water |
45 volumes of ozone in 100 volumes of water |
|
Magnetic properties |
Liquid and solid oxygen are paramagnetic substances, i.e. are drawn into a magnetic field |
It has diamagnetic properties, that is, it does not interact with the magnetic field |
|
Biological role |
Necessary for the respiration of plants and animals (mixed with nitrogen or inert gas). Inhalation of pure oxygen leads to severe poisoning |
In the atmosphere it forms the so-called ozone layer, which protects the biosphere from the harmful effects of ultraviolet radiation. Poisonous |
Chemical properties of oxygen and ozone
Interaction of oxygen with metals
Molecular oxygen is a fairly strong oxidizing agent. It oxidizes almost all metals (except gold and platinum). Many metals oxidize slowly in air, but in an atmosphere of pure oxygen they burn very quickly, forming an oxide:
However, when some metals burn, they form not oxides, but peroxides (in such compounds the oxidation state of Oxygen is -1) or superoxides (the oxidation state of the Oxygen atom is fractional). Examples of such metals are barium, sodium and potassium:
Interaction of oxygen with nonmetals
Oxygen exhibits an oxidation state of -2 in compounds that are formed with all nonmetals except Fluorine, Helium, Neon and Argon. When heated, oxygen molecules directly interact with all non-metals, except halogens and inert gases. In an oxygen atmosphere, phosphorus and some other non-metals spontaneously ignite:
When oxygen interacts with fluorine, oxygen fluoride is formed, and not fluorine oxide, since the Fluorine atom has a higher electronegativity than the Oxygen atom. Oxygen fluoride is a pale yellow gas. It is used as very strongoxidizing agent and fluorovalent agent. In this compound, the oxidation state of Oxygen is +2.
In an excess of fluorine, dioxygen difluoride can be formed, in which the oxidation state of Oxygen is +1. The structure of such a molecule is similar to the hydrogen peroxide molecule.
Application of oxygen and ozone. Meaning ozone layer
Oxygen is used by all aerobic living things for respiration. During photosynthesis, plants release oxygen and absorb carbon dioxide.
Molecular oxygen is used for so-called intensification, that is, acceleration of oxidative processes in the metallurgical industry. Oxygen is also used to produce a high-temperature flame. When acetylene (C 2 H 2) burns in oxygen, the flame temperature reaches 3500 ° C. In medicine, oxygen is used to facilitate the breathing of patients. It is also used in breathing apparatus for people working in difficult-to-breathe atmospheres. Liquid oxygen is used as a rocket fuel oxidizer.
Ozone is used in laboratory practice as a very strong oxidizing agent. In industry, it is used to disinfect water, since it has a strong oxidizing effect that destroys various microorganisms.
Peroxides, superoxides and ozonides of alkali metals are used for oxygen regeneration in spacecraft and submarines. This application is based on the reaction of these substances with carbon dioxide CO 2:
In nature, ozone is found in high layers of the atmosphere at an altitude of about 20-25 km, in the so-called ozone layer, which protects the Earth from harsh solar radiation. A decrease in ozone concentration in the stratosphere by even 1 can lead to serious consequences, such as an increase in the number of skin cancers in humans and animals, an increase in the number of diseases associated with suppression of the human immune system, a slowdown in the growth of terrestrial plants, a decrease in the growth rate of phytoplankton, etc. .
Without the ozone layer, life on the planet would be impossible. Meanwhile, atmospheric pollution from various industrial emissions leads to the destruction of the ozone layer. The most dangerous substances for ozone are freons (they are used as refrigerants in refrigeration machines, as well as fillers for deodorant cans) and waste rocket fuel.
The world community is very concerned about the formation of a hole in the ozone layer at the poles of our planet, and therefore, in 1987, the Montreal Protocol on Substances that Deplete the Ozone Layer was adopted, which limited the use of substances harmful to the ozone layer.
Physical properties of substances formed by the element Sulfur
Sulfur atoms, as well as Oxygen, can form various allotropic modifications ( S∞; S 12; S 8; S 6; S 2 and others). At room temperature, sulfur is in the formα -sulfur (or rhombic sulfur), which is yellow, brittle crystals, odorless, insoluble in water. At temperatures above +96 °C a slow transition occursα-sulfur to β -sulfur (or monoclinic sulfur), which is almost white plates. If molten sulfur is poured into water, the liquid sulfur supercools and yellow-brown, rubber-like plastic sulfur is formed, which later turns back into a-sulfur. Sulfur boils at a temperature of +445 ° C, forming dark brown vapors.
All modifications of sulfur are insoluble in water, but dissolve quite well in carbon disulfide(CS 2) and some other non-polar solvents.
Application of sulfur
The main product of the sulfur industry is sulfate acid. Its production accounts for about 60% of the sulfur that is mined. In the gum industry, sulfur is used to convert rubber into high-quality rubber, that is, to vulcanize rubber. Sulfur is the most important component of any pyrotechnic mixtures. For example, match heads contain about 5%, and the spread on a box contains about 20% sulfur by weight. In agriculture, sulfur is used to control pests in vineyards. In medicine, sulfur is used in the manufacture of various ointments to treat skin diseases.
Metals are active reducing agents with a positive oxidation state. Due to their chemical properties, metals are widely used in industry, metallurgy, medicine, and construction.
Metal activity
In reactions, metal atoms give up valence electrons and become oxidized. The more energy levels and fewer electrons a metal atom has, the easier it is for it to give up electrons and undergo reactions. Therefore, metallic properties increase from top to bottom and from right to left in the periodic table.
Rice. 1. Changes in metallic properties in the periodic table.
The activity of simple substances is shown in the electrochemical voltage series of metals. To the left of hydrogen are active metals (activity increases towards the left), to the right are inactive metals.
The greatest activity is exhibited by alkali metals that are in group I of the periodic table and are to the left of hydrogen in the electrochemical voltage series. They react with many substances already at room temperature. They are followed by alkaline earth metals, which are included in group II. They react with most substances when heated. Metals in the electrochemical series from aluminum to hydrogen (medium activity) require additional conditions to enter into reactions.
Rice. 2. Electrochemical series of voltages of metals.
Some metals exhibit amphoteric properties or duality. Metals, their oxides and hydroxides react with acids and bases. Most metals react only with certain acids, displacing hydrogen and forming a salt. The most pronounced dual properties are exhibited by:
- aluminum;
- lead;
- zinc;
- iron;
- copper;
- beryllium;
- chromium.
Each metal is capable of displacing another metal standing to the right of it in the electrochemical series from salts. Metals to the left of hydrogen displace it from dilute acids.
Properties
Features of the interaction of metals with different substances are presented in the table of chemical properties of metals.
|
Reaction |
Peculiarities |
The equation |
|
With oxygen |
Most metals form oxide films. Alkali metals spontaneously ignite in the presence of oxygen. In this case, sodium forms peroxide (Na 2 O 2), the remaining metals of group I form superoxides (RO 2). When heated, alkaline earth metals spontaneously ignite, while metals of intermediate activity oxidize. Gold and platinum do not interact with oxygen |
4Li + O 2 → 2Li 2 O; 2Na + O 2 → Na 2 O 2 ; K + O 2 → KO 2 ; 4Al + 3O 2 → 2Al 2 O 3; 2Cu + O 2 → 2CuO |
|
With hydrogen |
At room temperature alkaline compounds react, and when heated, alkaline earth compounds react. Beryllium does not react. Magnesium additionally requires high blood pressure |
Sr + H 2 → SrH 2 ; 2Na + H 2 → 2NaH; Mg + H 2 → MgH 2 |
|
Only active metals. Lithium reacts at room temperature. Other metals - when heated |
6Li + N 2 → 2Li 3 N; 3Ca + N 2 → Ca 3 N 2 |
|
|
With carbon |
Lithium and sodium, the rest - when heated |
4Al + 3C → Al 3 C4; 2Li+2C → Li 2 C 2 |
|
Gold and platinum do not interact |
2K + S → K 2 S; Fe + S → FeS; Zn + S → ZnS |
|
|
With phosphorus |
When heated |
3Ca + 2P → Ca 3 P 2 |
|
With halogens |
Only low-active metals do not react, copper - when heated |
Cu + Cl 2 → CuCl 2 |
|
Alkali and some alkaline earth metals. When heated, in acidic or alkaline conditions, metals of medium activity react |
2Na + 2H 2 O → 2NaOH + H 2 ; Ca + 2H 2 O → Ca(OH) 2 + H 2; Pb + H 2 O → PbO + H 2 |
|
|
With acids |
Metals to the left of hydrogen. Copper dissolves in concentrated acids |
Zn + 2HCl → ZnCl 2 + 2H 2 ; Fe + H 2 SO 4 → FeSO 4 + H 2; Cu + 2H 2 SO 4 → CuSO 4 + SO 2 +2H 2 O |
|
With alkalis |
Only amphoteric metals |
2Al + 2KOH + 6H 2 O → 2K + 3H 2 |
|
Reactive metals replace less reactive metals |
3Na + AlCl 3 → 3NaCl + Al |
Metals interact with each other and form intermetallic compounds - 3Cu + Au → Cu 3 Au, 2Na + Sb → Na 2 Sb.
Application
The general chemical properties of metals are used to create alloys, detergents, and are used in catalytic reactions. Metals are present in batteries, electronics, and supporting structures.
The main areas of application are listed in the table.
Rice. 3. Bismuth.
What have we learned?
From the 9th grade chemistry lesson we learned about the basic chemical properties of metals. The ability to interact with simple and complex substances determines the activity of metals. The more active a metal is, the more easily it reacts under normal conditions. Active metals react with halogens, nonmetals, water, acids, and salts. Amphoteric metals react with alkalis. Low-active metals do not react with water, halogens, and most non-metals. We briefly reviewed the areas of application. Metals are used in medicine, industry, metallurgy, and electronics.
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