The reaction of the compound with metals. General characteristics of metals. − dilute nitric acid

These are the elements of group I of the periodic system: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr); very soft, ductile, fusible and light, usually silvery white; chemically very active; react violently with water to form alkalis(whence the name).

All alkali metals are extremely active, exhibit reducing properties in all chemical reactions, donate their only valence electron, turning into a positively charged cation, and exhibit a single oxidation state +1.

The reducing ability increases in the series ––Li–Na–K–Rb–Cs.

All alkali metal compounds are ionic in nature.

Almost all salts are soluble in water.

low melting points,

Small values ​​of density,

Soft, cut with a knife

Due to their activity, alkali metals are stored under a layer of kerosene to block the access of air and moisture. Lithium is very light and floats to the surface in kerosene, so it is stored under a layer of petroleum jelly.

Chemical properties of alkali metals

1. Alkali metals actively interact with water:

2Na + 2H 2 O → 2NaOH + H 2

2Li + 2H 2 O → 2LiOH + H 2

2. Reaction of alkali metals with oxygen:

4Li + O 2 → 2Li 2 O (lithium oxide)

2Na + O 2 → Na 2 O 2 (sodium peroxide)

K + O 2 → KO 2 (potassium superoxide)

In air, alkali metals instantly oxidize. Therefore, they are stored under a layer of organic solvents (kerosene, etc.).

3. In the reactions of alkali metals with other non-metals, binary compounds are formed:

2Li + Cl 2 → 2LiCl (halides)

2Na + S → Na 2 S (sulfides)

2Na + H 2 → 2NaH (hydrides)

6Li + N 2 → 2Li 3 N (nitrides)

2Li + 2C → Li 2 C 2 (carbides)

4. Reaction of alkali metals with acids

(rarely carried out, there is a competing reaction with water):

2Na + 2HCl → 2NaCl + H 2

5. Interaction of alkali metals with ammonia

(sodium amide is formed):

2Li + 2NH 3 = 2LiNH 2 + H 2

6. The interaction of alkali metals with alcohols and phenols, which in this case exhibit acidic properties:

2Na + 2C 2 H 5 OH \u003d 2C 2 H 5 ONa + H 2;

2K + 2C 6 H 5 OH = 2C 6 H 5 OK + H 2 ;

7. Qualitative reaction to alkali metal cations - coloring of the flame in the following colors:

Li + - carmine red

Na + - yellow

K + , Rb + and Cs + - violet

Obtaining alkali metals

Lithium, sodium and potassium metal receive electrolysis of molten salts (chlorides), and rubidium and cesium - reduction in vacuum when their chlorides are heated with calcium: 2CsCl + Ca \u003d 2Cs + CaCl 2
On a small scale, vacuum thermal production of sodium and potassium is also used:

2NaCl + CaC 2 \u003d 2Na + CaCl 2 + 2C;
4KCl + 4CaO + Si \u003d 4K + 2CaCl 2 + Ca 2 SiO 4.

Active alkali metals are released in vacuum thermal processes due to their high volatility (their vapors are removed from the reaction zone).

Features of the chemical properties of s-elements of group I and their physiological effect

The electronic configuration of the lithium atom is 1s 2 2s 1 . It has the largest atomic radius in the 2nd period, which facilitates the detachment of the valence electron and the emergence of the Li + ion with a stable inert gas (helium) configuration. Therefore, its compounds are formed with the transfer of an electron from lithium to another atom and the occurrence of an ionic bond with a small amount of covalence. Lithium is a typical metallic element. In substance form, it is an alkali metal. It differs from other members of group I in its small size and the smallest, in comparison with them, activity. In this respect, it resembles the group II element, magnesium, located diagonally from Li. In solutions, the Li + ion is highly solvated; it is surrounded by several tens of water molecules. Lithium, in terms of solvation energy - the addition of solvent molecules, is closer to a proton than to alkali metal cations.

The small size of the Li + ion, the high nuclear charge and only two electrons create conditions for the emergence of a rather significant positive charge field around this particle, therefore, in solutions, a significant number of polar solvent molecules are attracted to it and its coordination number is large, the metal is able to form a significant number of organolithium compounds .

Sodium begins the 3rd period, so it has only 1e at the external level - , occupying the 3s orbital. The radius of the Na atom is the largest in the 3rd period. These two features determine the nature of the element. Its electronic configuration is 1s 2 2s 2 2p 6 3s 1 . The only oxidation state of sodium is +1. Its electronegativity is very low, therefore sodium is present in compounds only in the form of a positively charged ion and gives the chemical bond an ionic character. The size of the Na + ion is much larger than Li +, and its solvation is not so great. However, it does not exist in free form in solution.

The physiological significance of K + and Na + ions is associated with their different adsorbability on the surface of the components that make up the earth's crust. Sodium compounds are only slightly adsorbed, while potassium compounds are strongly retained by clay and other substances. Cell membranes, being the cell-environment interface, are permeable to K + ions, as a result of which the intracellular concentration of K + is much higher than that of Na + ions. At the same time, the concentration of Na + in the blood plasma exceeds the content of potassium in it. This circumstance is associated with the emergence of the membrane potential of cells. Ions K + and Na + - one of the main components of the liquid phase of the body. Their ratio with Ca 2+ ions is strictly defined, and its violation leads to pathology. The introduction of Na + ions into the body does not have a noticeable harmful effect. An increase in the content of K + ions is harmful, but under normal conditions, an increase in its concentration never reaches dangerous values. The effect of Rb + , Cs + , Li + ions has not yet been sufficiently studied.

Of the various lesions associated with the use of alkali metal compounds, burns with hydroxide solutions are most common. The action of alkalis is associated with the dissolution of skin proteins in them and the formation of alkaline albuminates. Alkali is released again as a result of their hydrolysis and acts on the deeper layers of the body, causing the appearance of ulcers. Nails under the influence of alkalis become dull and brittle. Eye damage, even with very dilute alkali solutions, is accompanied not only by superficial destruction, but by violations of deeper parts of the eye (iris) and leads to blindness. During the hydrolysis of alkali metal amides, alkali and ammonia are simultaneously formed, causing fibrinous-type tracheobronchitis and pneumonia.

Potassium was obtained by G. Davy almost simultaneously with sodium in 1807 during the electrolysis of wet potassium hydroxide. From the name of this compound - "caustic potash" and the element got its name. The properties of potassium differ markedly from the properties of sodium, due to the difference in the radii of their atoms and ions. In potassium compounds, the bond is more ionic, and in the form of the K + ion, it has a lesser polarizing effect than sodium, due to its large size. The natural mixture consists of three isotopes 39 K, 40 K, 41 K. One of them is 40 K ‑ is radioactive and a certain proportion of the radioactivity of minerals and soil is associated with the presence of this isotope. Its half-life is long - 1.32 billion years. Determining the presence of potassium in a sample is quite easy: vapors of the metal and its compounds turn the flame purple-red. The spectrum of the element is quite simple and proves the presence of 1e - on the 4s orbital. The study of it served as one of the grounds for finding general patterns in the structure of the spectra.

In 1861 Robert Bunsen discovered a new element while studying the salt of mineral springs by spectral analysis. Its presence was proved by dark red lines in the spectrum, which other elements did not give. By the color of these lines, the element was named rubidium (rubidus-dark red). In 1863, R. Bunsen obtained this metal in its pure form by reducing rubidium tartrate (tartar salt) with soot. A feature of the element is the slight excitability of its atoms. Electron emission from it appears under the action of red rays of the visible spectrum. This is due to a small difference in the energies of the atomic 4d and 5s orbitals. Of all the alkaline elements with stable isotopes, rubidium (like cesium) has one of the largest atomic radii and a low ionization potential. Such parameters determine the nature of the element: high electropositivity, extreme chemical activity, low melting point (39 0 C) and low resistance to external influences.

The discovery of cesium, like rubidium, is associated with spectral analysis. In 1860, R. Bunsen discovered two bright blue lines in the spectrum that did not belong to any element known at that time. Hence the name "caesius" (caesius), which means sky blue. It is the last element of the alkali metal subgroup still found in measurable amounts. The largest atomic radius and the smallest first ionization potentials determine the nature and behavior of this element. It has a pronounced electropositivity and pronounced metallic qualities. The desire to donate the outer 6s-electron leads to the fact that all its reactions proceed extremely violently. A small difference in the energies of the atomic 5d and 6s orbitals is responsible for the slight excitability of the atoms. Electronic emission in cesium is observed under the action of invisible infrared rays (thermal). This feature of the atomic structure determines the good electrical conductivity of the current. All this makes cesium indispensable in electronic devices. Recently, more and more attention has been paid to cesium plasma as a fuel of the future and in connection with the solution of the problem of thermonuclear fusion.

In air, lithium actively reacts not only with oxygen, but also with nitrogen and is covered with a film consisting of Li 3 N (up to 75%) and Li 2 O. The remaining alkali metals form peroxides (Na 2 O 2) and superoxides (K 2 O 4 or KO 2).

The following substances react with water:

Li 3 N + 3 H 2 O \u003d 3 LiOH + NH 3;

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

K 2 O 4 + 2 H 2 O \u003d 2 KOH + H 2 O 2 + O 2.

For air regeneration in submarines and spaceships, in insulating gas masks and breathing apparatus of combat swimmers (submarine saboteurs), a mixture of "oxon" was used:

Na 2 O 2 + CO 2 \u003d Na 2 CO 3 + 0.5 O 2;

K 2 O 4 + CO 2 \u003d K 2 CO 3 + 1.5 O 2.

This is currently the standard filling of regenerating cartridges for insulating gas masks for firefighters.
Alkali metals react when heated with hydrogen to form hydrides:

Lithium hydride is used as a strong reducing agent.

Hydroxides alkali metals corrode glass and porcelain dishes, they can not be heated in quartz dishes:

SiO 2 + 2NaOH \u003d Na 2 SiO 3 + H 2 O.

Sodium and potassium hydroxides do not split off water when heated up to their boiling point (more than 1300 0 C). Some sodium compounds are called soda:

a) soda ash, anhydrous soda, laundry soda or just soda - sodium carbonate Na 2 CO 3;
b) crystalline soda - sodium carbonate crystal hydrate Na 2 CO 3. 10H2O;
c) bicarbonate or drinking - sodium bicarbonate NaHCO 3;
d) sodium hydroxide NaOH is called caustic soda or caustic.

This lesson is devoted to the study of the topic “General properties of metals. Metal connection. During the lesson, the general chemical properties of metals, the features of the metallic chemical bond will be considered. The teacher will explain the similarities between the chemical and physical properties of metals using a model of their internal structure.

Topic: Chemistry of metals

Lesson: General properties of metals. metal connection

Metals are characterized by common physical properties: they have a special metallic luster, high thermal and electrical conductivity, and ductility.

Metals also share some common chemical properties. It is important to remember that in chemical reactions, metals act as reducing agents: they donate electrons and increase their oxidation state. Consider some reactions in which metals participate.

INTERACTION WITH OXYGEN

Many metals can react with oxygen. Usually the products of these reactions are oxides, but there are exceptions, which you will learn about in the next lesson. Consider the interaction of magnesium with oxygen.

Magnesium burns in oxygen to form magnesium oxide:

2Mg + O 2 \u003d 2MgO

Rice. 1. Combustion of magnesium in oxygen

Magnesium atoms donate their outer electrons to oxygen atoms: two magnesium atoms donate two electrons each to two oxygen atoms. In this case, magnesium acts as a reducing agent, and oxygen acts as an oxidizing agent.

Metals react with halogens. The product of this reaction is a metal halide, such as chloride.

Rice. 2. Potassium combustion in chlorine

Potassium burns in chlorine to form potassium chloride:

2K + Cl 2 \u003d 2KCl

Two potassium atoms donate one electron to the chlorine molecule. Potassium, increasing the oxidation state, plays the role of a reducing agent, and chlorine, lowering the oxidation state, plays the role of an oxidizing agent.

Many metals react with sulfur to form sulfides. In these reactions, metals also act as reducing agents, while sulfur will act as an oxidizing agent. Sulfur in sulfides is in the -2 oxidation state, i.e. it lowers its oxidation state from 0 to -2. For example, when heated, iron reacts with sulfur to form iron (II) sulfide:

Rice. 3. Interaction of iron with sulfur

Metals can also react with hydrogen, nitrogen, and other non-metals under certain conditions.

Only active metals, such as alkali and alkaline earth, react with water without heating. During these reactions, alkali is formed and hydrogen gas is released. For example, calcium reacts with water to form calcium hydroxide and hydrogen, and a large amount of heat is released:

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

Less active metals, such as iron and zinc, react with water only when heated to form metal oxide and hydrogen. For example:

Zn + H 2 O \u003d ZnO + H 2

In these reactions, the oxidizing agent is the hydrogen atom, which is part of the water.

Metals to the right of hydrogen in the voltage series do not react with water.

You already know that metals that are in the series of voltages to the left of hydrogen react with acids. In these reactions, metals donate electrons and act as a reducing agent. The oxidizing agent is hydrogen cations formed in acid solutions. For example, zinc reacts with hydrochloric acid:

Zn + 2HCl \u003d ZnCl 2 + H 2

Otherwise, the reactions of metals with nitric and concentrated sulfuric acids proceed. Almost no hydrogen is released in these reactions. We will talk about such interactions in the next lessons.

A metal can react with a salt solution if it is more active than the metal in the salt. For example, iron replaces copper from copper (II) sulfate:

Fe + CuSO 4 \u003d FeSO 4 + Cu

Iron is a reducing agent, copper cations are an oxidizing agent.

Let's try to explain why metals have common physical and chemical properties. To do this, consider a model of the internal structure of the metal.

Metal atoms have relatively large radii and a small number of external electrons. These electrons are weakly attracted to the nucleus, therefore, in chemical reactions, metals act as reducing agents, donating electrons from the external energy level.

In the nodes of the crystal lattice of metals there are not only neutral atoms, but also metal cations, because outer electrons move freely in the crystal lattice. At the same time, atoms donating electrons become cations, and cations, accepting electrons, turn into electrically neutral atoms.

Rice. 4. Model of the internal structure of the metal

A chemical bond that is formed as a result of the attraction of metal cations to freely moving electrons is called metallic.

The electrical and thermal conductivity of metals is explained by the presence of free electrons, which can be carriers of electric current and heat carriers. The plasticity of the metal is explained by the fact that under mechanical action the chemical bond does not break, because. a chemical bond is established not between specific atoms and cations, but between all metal cations with all free electrons in a metal crystal.

1. Mikityuk A.D. Collection of tasks and exercises in chemistry. Grades 8-11 / A.D. Mikityuk. - M.: Ed. "Exam", 2009.

2. Orzhekovsky P.A. Chemistry: 9th grade: textbook. for general inst. / P.A. Orzhekovsky, L.M. Meshcheryakova, L.S. Pontak. - M.: AST: Astrel, 2007. (§23)

3. Orzhekovsky P.A. Chemistry: 9th grade: textbook for general education. inst. / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashova. - M.: Astrel, 2013. (§6)

4. Rudzitis G.E. Chemistry: inorgan. chemistry. Organ. chemistry: textbook. for 9 cells. / G.E. Rudzitis, F.G. Feldman. - M .: Education, JSC "Moscow textbooks", 2009.

5. Khomchenko I.D. Collection of problems and exercises in chemistry for high school. - M.: RIA "New Wave": Publisher Umerenkov, 2008.

6. Encyclopedia for children. Volume 17. Chemistry / Chapter. ed. V.A. Volodin, leading. scientific ed. I. Leenson. - M.: Avanta +, 2003.

Additional web resources

1. A single collection of digital educational resources (video experiences on the topic) ().

2. Electronic version of the journal "Chemistry and Life" ().

Homework

p.41 Nos. A1, A2 from P.A. Orzhekovsky’s Textbook. "Chemistry: 9th grade" (M.: Astrel, 2013).

General properties of metals.

The presence of valence electrons weakly bound to the nucleus determines the general chemical properties of metals. In chemical reactions, they always act as a reducing agent; simple substances, metals, never exhibit oxidizing properties.

Getting metals:
- recovery from oxides with carbon (C), carbon monoxide (CO), hydrogen (H2) or more active metal (Al, Ca, Mg);
- recovery from salt solutions with a more active metal;
- electrolysis of solutions or melts of metal compounds - recovery of the most active metals (alkali, alkaline earth metals and aluminum) using electric current.

In nature, metals are found mainly in the form of compounds, only low-active metals are found in the form of simple substances (native metals).

Chemical properties of metals.
1. Interaction with simple substances non-metals:
Most metals can be oxidized with non-metals such as halogens, oxygen, sulfur, nitrogen. But most of these reactions require preheating to start. In the future, the reaction can proceed with the release of a large amount of heat, which leads to the ignition of the metal.
At room temperature, reactions are possible only between the most active metals (alkali and alkaline earth) and the most active non-metals (halogens, oxygen). Alkali metals (Na, K) react with oxygen to form peroxides and superoxides (Na2O2, KO2).

a) interaction of metals with water.
At room temperature, alkali and alkaline earth metals interact with water. As a result of the substitution reaction, an alkali (soluble base) and hydrogen are formed: Metal + H2O \u003d Me (OH) + H2
When heated, other metals interact with water, standing in the activity series to the left of hydrogen. Magnesium reacts with boiling water, aluminum - after a special surface treatment, resulting in the formation of insoluble bases - magnesium hydroxide or aluminum hydroxide - and hydrogen is released. Metals in the activity range from zinc (inclusive) to lead (inclusive) interact with water vapor (i.e. above 100 C), while oxides of the corresponding metals and hydrogen are formed.
Metals to the right of hydrogen in the activity series do not interact with water.
b) interaction with oxides:
active metals interact in a substitution reaction with oxides of other metals or non-metals, reducing them to simple substances.
c) interaction with acids:
Metals located to the left of hydrogen in the activity series react with acids to release hydrogen and form the corresponding salt. Metals to the right of hydrogen in the activity series do not interact with acid solutions.
A special place is occupied by the reactions of metals with nitric and concentrated sulfuric acids. All metals except noble ones (gold, platinum) can be oxidized by these oxidizing acids. As a result of these reactions, the corresponding salts will always be formed, water and the product of nitrogen or sulfur reduction, respectively.
d) with alkalis
Metals that form amphoteric compounds (aluminum, beryllium, zinc) are capable of reacting with melts (with the formation of medium salts of aluminates, beryllates or zincates) or alkali solutions (with the formation of the corresponding complex salts). All reactions will release hydrogen.
e) In accordance with the position of the metal in the activity series, reactions of reduction (displacement) of a less active metal from a solution of its salt by another more active metal are possible. As a result of the reaction, a salt of a more active and simple substance is formed - a less active metal.

General properties of nonmetals.

There are much fewer non-metals than metals (22 elements). However, the chemistry of non-metals is much more complicated due to the greater filling of the external energy level of their atoms.
The physical properties of non-metals are more diverse: among them are gaseous (fluorine, chlorine, oxygen, nitrogen, hydrogen), liquids (bromine) and solids, which differ greatly from each other in melting point. Most non-metals do not conduct electricity, but silicon, graphite, germanium have semiconductor properties.
Gaseous, liquid and some solid non-metals (iodine) have a molecular structure of the crystal lattice, the rest of the non-metals have an atomic crystal lattice.
Fluorine, chlorine, bromine, iodine, oxygen, nitrogen and hydrogen under normal conditions exist in the form of diatomic molecules.
Many non-metal elements form several allotropic modifications of simple substances. So oxygen has two allotropic modifications - oxygen O2 and ozone O3, sulfur has three allotropic modifications - rhombic, plastic and monoclinic sulfur, phosphorus has three allotropic modifications - red, white and black phosphorus, carbon - six allotropic modifications - soot, graphite, diamond , carbine, fullerene, graphene.

Unlike metals, which exhibit only reducing properties, non-metals in reactions with simple and complex substances can act both as a reducing agent and as an oxidizing agent. According to their activity, non-metals occupy a certain place in the series of electronegativity. Fluorine is considered the most active non-metal. It exhibits only oxidizing properties. Oxygen is in second place in terms of activity, nitrogen is in third, then halogens and other non-metals. Hydrogen has the lowest electronegativity among non-metals.

Chemical properties of non-metals.

1. Interaction with simple substances:
Nonmetals interact with metals. In such a reaction, metals act as a reducing agent, non-metals as an oxidizing agent. As a result of the reaction of the compound, binary compounds are formed - oxides, peroxides, nitrides, hydrides, salts of oxygen-free acids.
In the reactions of non-metals with each other, a more electronegative non-metal exhibits the properties of an oxidizing agent, a less electronegative one - the properties of a reducing agent. As a result of the compound reaction, binary compounds are formed. It must be remembered that non-metals can exhibit variable oxidation states in their compounds.
2. Interaction with complex substances:
a) with water:
Under normal conditions, only halogens interact with water.
b) with oxides of metals and non-metals:
Many non-metals can react at high temperatures with oxides of other non-metals, reducing them to simple substances. Non-metals to the left of sulfur in the electronegativity series can also interact with metal oxides, reducing metals to simple substances.
c) with acids:
Some non-metals can be oxidized with concentrated sulfuric or nitric acids.
d) with alkalis:
Under the action of alkalis, some non-metals can undergo dismutation, being both an oxidizing agent and a reducing agent.
For example, in the reaction of halogens with alkali solutions without heating: Cl2 + 2NaOH = NaCl + NaClO + H2O or when heated: 3Cl2 + 6NaOH = 5NaCl + NaClO3 + 3H2O.
e) with salts:
When interacting, being strong oxidizing agents, they exhibit reducing properties.
Halogens (except fluorine) enter into substitution reactions with solutions of salts of hydrohalic acids: a more active halogen displaces a less active halogen from a salt solution.

Group IIA contains only metals - Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium) and Ra (radium). The chemical properties of the first representative of this group, beryllium, differ most strongly from the chemical properties of the other elements of this group. Its chemical properties are in many ways even more similar to aluminum than to other group IIA metals (the so-called "diagonal similarity"). Magnesium, in terms of chemical properties, also differs markedly from Ca, Sr, Ba, and Ra, but still has much more similar chemical properties with them than with beryllium. Due to the significant similarity of the chemical properties of calcium, strontium, barium and radium, they are combined into one family, called alkaline earth metals.

All elements of group IIA belong to s-elements, i.e. contain all of their valence electrons s-sublevel. Thus, the electronic configuration of the outer electron layer of all chemical elements of this group has the form ns 2 , where n– number of the period in which the element is located.

Due to the peculiarities of the electronic structure of group IIA metals, these elements, in addition to zero, are capable of having only one single oxidation state, equal to +2. Simple substances formed by elements of group IIA, when participating in any chemical reactions, can only be oxidized, i.e. donate electrons:

Me 0 - 2e - → Me +2

Calcium, strontium, barium and radium are extremely reactive. The simple substances formed by them are very strong reducing agents. Magnesium is also a strong reducing agent. The reducing activity of metals obeys the general laws of the periodic law of D.I. Mendeleev and increases down the subgroup.

Interaction with simple substances

with oxygen

Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of a liquid inert to them, most often kerosene.

Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO 2):

2Mg + O 2 \u003d 2MgO

2Ca + O 2 \u003d 2CaO

2Ba + O 2 \u003d 2BaO

Ba + O 2 \u003d BaO 2

It should be noted that during the combustion of alkaline earth metals and magnesium in air, the side reaction of these metals with atmospheric nitrogen also proceeds, as a result of which, in addition to compounds of metals with oxygen, nitrides with the general formula Me 3 N 2 are also formed.

with halogens

Beryllium reacts with halogens only at high temperatures, while the rest of the Group IIA metals already at room temperature:

Mg + I 2 \u003d MgI 2 - magnesium iodide

Ca + Br 2 \u003d CaBr 2 - calcium bromide

Ba + Cl 2 \u003d BaCl 2 - barium chloride

with non-metals of IV–VI groups

All metals of group IIA react when heated with all non-metals of groups IV-VI, but depending on the position of the metal in the group, as well as the activity of non-metals, a different degree of heating is required. Since beryllium is the most chemically inert among all metals of group IIA, its reactions with nonmetals require significantly more about high temperature.

It should be noted that the reaction of metals with carbon can form carbides of various nature. There are carbides related to methanides and conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by a metal. They, like methane, contain carbon in the -4 oxidation state, and during their hydrolysis or interaction with non-oxidizing acids, methane is one of the products. There is also another type of carbides - acetylenides, which contain the C 2 2- ion, which is actually a fragment of the acetylene molecule. Carbides of the acetylenide type upon hydrolysis or interaction with non-oxidizing acids form acetylene as one of the reaction products. What type of carbide - methanide or acetylenide - will be obtained by the interaction of one or another metal with carbon depends on the size of the metal cation. As a rule, methanides are formed with metal ions having a small radius, and acetylides with larger ions. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:

The remaining metals of group II A form acetylenides with carbon:

With silicon, group IIA metals form silicides - compounds of the Me 2 Si type, with nitrogen - nitrides (Me 3 N 2), phosphorus - phosphides (Me 3 P 2):

with hydrogen

All alkaline earth metals react when heated with hydrogen. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, an increased pressure of hydrogen is also required. Beryllium does not react with hydrogen under any conditions.

Interaction with complex substances

with water

All alkaline earth metals actively react with water to form alkalis (soluble metal hydroxides) and hydrogen. Magnesium reacts with water only during boiling, due to the fact that when heated, the protective oxide film of MgO dissolves in water. In the case of beryllium, the protective oxide film is very resistant: water does not react with it either when boiling or even at a red heat temperature:

with non-oxidizing acids

All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the activity series to the left of hydrogen. In this case, a salt of the corresponding acid and hydrogen are formed. Reaction examples:

Be + H 2 SO 4 (razb.) \u003d BeSO 4 + H 2

Mg + 2HBr \u003d MgBr 2 + H 2

Ca + 2CH 3 COOH = (CH 3 COO) 2 Ca + H 2

with oxidizing acids

− dilute nitric acid

All Group IIA metals react with dilute nitric acid. In this case, the reduction products instead of hydrogen (as in the case of non-oxidizing acids) are nitrogen oxides, mainly nitric oxide (I) (N 2 O), and in the case of highly diluted nitric acid, ammonium nitrate (NH 4 NO 3):

4Ca + 10HNO 3 ( razb .) \u003d 4Ca (NO 3) 2 + N 2 O + 5H 2 O

4Mg + 10HNO3 (very disaggregated)\u003d 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O

− concentrated nitric acid

Concentrated nitric acid at ordinary (or low) temperature passivates beryllium, i.e. does not react with it. When boiling, the reaction is possible and proceeds mainly in accordance with the equation:

Magnesium and alkaline earth metals react with concentrated nitric acid to form a wide range of different nitrogen reduction products.

− concentrated sulfuric acid

Beryllium is passivated with concentrated sulfuric acid, i.e. does not react with it under normal conditions, however, the reaction proceeds during boiling and leads to the formation of beryllium sulfate, sulfur dioxide and water:

Be + 2H 2 SO 4 → BeSO 4 + SO 2 + 2H 2 O

Barium is also passivated by concentrated sulfuric acid due to the formation of insoluble barium sulfate, but reacts with it when heated, barium sulfate dissolves when heated in concentrated sulfuric acid due to its conversion to barium hydrogen sulfate.

The remaining metals of the main group IIA react with concentrated sulfuric acid under any conditions, including in the cold. Sulfur reduction can occur to SO 2, H 2 S and S, depending on the activity of the metal, the reaction temperature and the concentration of the acid:

Mg + H 2 SO 4 ( conc .) \u003d MgSO 4 + SO 2 + H 2 O

3Mg + 4H2SO4 ( conc .) \u003d 3MgSO 4 + S↓ + 4H 2 O

4Ca + 5H2SO4 ( conc .) \u003d 4CaSO 4 + H 2 S + 4H 2 O

with alkalis

Magnesium and alkaline earth metals do not interact with alkalis, and beryllium easily reacts both with alkali solutions and with anhydrous alkalis during fusion. Moreover, when the reaction is carried out in an aqueous solution, water is also involved in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and gaseous hydrogen:

Be + 2KOH + 2H 2 O \u003d H 2 + K 2 - potassium tetrahydroxoberyllate

When carrying out the reaction with solid alkali during fusion, beryllates of alkali or alkaline earth metals and hydrogen are formed.

Be + 2KOH \u003d H 2 + K 2 BeO 2 - potassium beryllate

with oxides

Alkaline earth metals, as well as magnesium, can reduce less active metals and some non-metals from their oxides when heated, for example:

The method of restoring metals from their oxides with magnesium is called magnesiumthermy.

CHEMICAL PROPERTIES OF METALS

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.

In reactions - only reducing agents. Metal atoms easily donate electrons from the outer (and some of them from the pre-outer) electron layer, turning into positive ions. Possible oxidation states Me Lower 0,+1,+2,+3 Higher +4,+5,+6,+7,+8

1. INTERACTION WITH NON-METALS

1. WITH HYDROGEN

Metals of groups IA and IIA react when heated, except for beryllium. Solid unstable substances hydrides are formed, other metals do not react.

2K + H₂ = 2KH (potassium hydride)

Ca + H₂ = CaH₂

2. WITH OXYGEN

All metals react except 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 2Сa + O2 = 2СaO

2Сa+O2=2СaO

The metals of the secondary subgroups form oxides under normal conditions and when heated, oxides of various degrees of oxidation, and iron iron scale Fe3O4 (Fe⁺²O∙Fe2⁺³O3)

3Fe + 2O2 = Fe3O4

4Cu + O₂ = 2Cu₂⁺¹O (red) 2Cu + O₂ = 2Cu⁺²O (black);

2Zn + O₂ = ZnO 4Cr + 3O2 = 2Cr2O3

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:

FROMa+Cl2=FROMaCl2

2Al+3Cl2 = 2AlCl3

Metals of secondary subgroups at elevated temperatures

Cu + Cl₂ = Cu⁺²Cl₂ Zn + Cl₂ = ZnCl₂

2Fe + ЗС12 = 2Fe⁺³Cl3 iron 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

fromgray – sulfides: 2K + S = K2S 2Li+S = Li2S (sulfide)

FROMa+S=FROMaS(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

leaks 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

flows when heated:

Carbides are formed with carbon. Only the most active metals react with carbon. From 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 such as solid solutions with carbon: WC, ZnC, TiC - are used to obtain superhard steels.

with silicon - silicides: 4Cs + Si = Cs4Si,

7. INTERACTION OF METALS WITH WATER:

Metals that reach hydrogen in the electrochemical series of voltages react with water. Alkali and alkaline earth metals react with water without heating, forming soluble hydroxides (alkalis) and hydrogen, aluminum (after the destruction of the oxide film - amalgation), magnesium when heated, form insoluble bases and hydrogen .

2Na + 2HOH = 2NaOH + H2
FROMa + 2HOH = Ca(OH)2 + H2

2Al + 6H2O = 2Al(OH)3 + ZH2

The remaining metals react with water only in a 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 easily oxidize 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 temperature → non-metal or low-active metal and oxide (calciumthermy, magnesiumthermy, aluminothermy)

2Al + Cr2O3 = 2Cr + Al2O3 3Са + Cr₂O₃ = 3СаО + 2Cr (800 °C) 8Al + 3Fe3O4 = 4Al2O3 + 9Fe (thermite) 2Mg + CO2 = 2MgO + С Mg + N2O = MgO + N2 Zn + CO2 = ZnO + CO 2Cu + 2NO = 2CuO + N2 3Zn + SO2 = ZnS + 2ZnO

10. WITH OXIDES

Metals iron and chromium react with oxides, reducing the degree of oxidation

Cr + Cr2⁺³O3 = 3Cr⁺²O Fe+ Fe2⁺³O3 = 3Fe⁺²O

11. INTERACTION OF METALS WITH ALKALI

Only those metals interact with alkalis, the 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. INTERACTION WITH ACIDS (EXCEPT HNO3 and H2SO4 (conc.)

Metals standing in the electrochemical series of voltages of metals to the left of hydrogen displace it from dilute acids → salt and hydrogen

Remember! Nitric acid never releases hydrogen when interacting with metals.

Mg + 2HC1 = MgCl2 + H2
Al + 2HC1 = Al⁺³Cl₃ + H2

13. REACTIONS WITH SALT

Active metals displace less active metals from salts. Recovery from solutions:

CuSO4 + Zn = ZnSO4 + Cu

FeSO4 + Cu =REACTIONSNO

Mg + CuCl2(pp) = MgCl2 +FROMu

Recovery of metals from melts of their salts

3Na+ AlCl₃ = 3NaCl + Al

TiCl2 + 2Mg = MgCl2 + Ti

Group B metals react with salts, lowering their oxidation state.

2Fe⁺³Cl3 + Fe = 3Fe⁺²Cl2