Compound reaction with metals. General characteristics of metals. - diluted 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, plastic, low-melting and light, usually silvery-white; chemically very active; react violently with water, forming alkalis(where the name comes from).

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

The regenerative capacity increases in the order –– Li – Na – K – Rb – Cs.

All alkali metal compounds are ionic.

Almost all salts are water soluble.

Low melting points,

Low density values,

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 are instantly oxidized. 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. Interaction of alkali metals with alcohols and phenols, which in this case exhibit acidic properties:

2Na + 2C 2 H 5 OH = 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 + - purple

Obtaining alkali metals

Metallic lithium, sodium and potassium get by electrolysis of molten salts (chlorides), and rubidium and cesium - by reduction in vacuum when their chlorides are heated with calcium: 2CsCl + Ca = 2Cs + CaCl 2
The vacuum-thermal production of sodium and potassium is also used on a small scale:

2NaCl + CaC 2 = 2Na + CaCl 2 + 2C;
4KCl + 4CaO + Si = 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 group I s-elements and their physiological action

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 a valence electron and the appearance of a Li + ion with a stable configuration of an inert gas (helium). Consequently, its compounds are formed with the transfer of an electron from lithium to another atom and the appearance of an ionic bond with a small fraction of covalence. Lithium is a typical metallic element. As a substance, it is an alkali metal. It differs from other members of group I by its small size and the smallest, in comparison with them, activity. In this respect, it resembles the Group II element, which is located diagonally from Li, magnesium. In solutions, the Li + ion is highly solvated; it is surrounded by several tens of water molecules. Lithium in terms of the energy of solvation - the addition of solvent molecules, is closer to the proton than to the cations of alkali metals.

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

The third period begins with sodium, therefore, at the external level, it has only 1e - , 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 small, therefore, sodium is present in compounds only in the form of a positively charged ion and gives the chemical bond an ionic character. The Na + ion is much larger in size 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 adsorption capacity on the surface of the components that make up the earth's crust. Sodium compounds are only slightly subject to adsorption, while potassium compounds are firmly held by clay and other substances. Cell membranes, being the cell - medium 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 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 + are 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 influence 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 ulcers. Under the influence of alkalis, nails become dull and brittle. Damage to the eyes, even with very dilute alkali solutions, is accompanied not only by superficial damage, but by disturbances in the deeper parts of the eye (iris) and leads to blindness. During the hydrolysis of alkali metal amides, alkali and ammonia are simultaneously formed, causing tracheobronchitis of the fibrinous type and pneumonia.

Potassium was obtained by G. Davy almost simultaneously with sodium in 1807 by electrolysis of wet potassium hydroxide. From the name of this compound - "caustic potassium" and the element got its name. The properties of potassium differ markedly from those of sodium, which is 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 lower 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 ‑ 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. It is quite easy to determine the presence of potassium in a sample: vapors of the metal and its compounds color the flame in a violet-red color. The spectrum of the element is quite simple and proves the presence of 1e - in the 4s orbital. Its study served as one of the grounds for finding general patterns in the structure of the spectra.

In 1861, while studying the salt of mineral springs by spectral analysis, Robert Bunsen discovered a new element. Its presence was proved by dark red lines in the spectrum, which were not given by other elements. According to 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 reduction of rubidium tartrate (tartrate salt) with soot. A feature of the element is the easy excitability of its atoms. Electronic emission from it appears under the influence of red rays of the visible spectrum. This is due to the small difference in the energies of the atomic 4d and 5s orbitals. Of all alkaline elements with stable isotopes, rubidium (like cesium) has one of the largest atomic radii and a small ionization potential. Such parameters determine the nature of the element: high electropositiveness, 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 by that time. Hence the name "cesius" (caesius), which means sky blue. It is the last element in the alkali metal subgroup that is still found in measurable amounts. The largest atomic radius and the smallest first ionization potentials determine the character and behavior of this element. It has a pronounced electropositiveness and pronounced metallic qualities. The desire to donate the outer 6s electron leads to the fact that all its reactions are extremely violent. The small difference in the energies of the atomic 5d and 6s orbitals is responsible for the slight excitability of atoms. Electronic emission from cesium is observed under the influence of invisible infrared rays (heat). The specified 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 the 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 rest of the alkali metals form peroxides (Na 2 O 2) and superoxides (K 2 O 4 or KO 2).

The listed substances react with water:

Li 3 N + 3 H 2 O = 3 LiOH + NH 3;

Na 2 O 2 + 2 H 2 O = 2 NaOH + H 2 O 2;

K 2 O 4 + 2 H 2 O = 2 KOH + H 2 O 2 + O 2.

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

Na 2 O 2 + CO 2 = Na 2 CO 3 + 0.5O 2;

K 2 O 4 + CO 2 = K 2 CO 3 + 1.5 O 2.

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

Lithium hydride is used as a strong reducing agent.

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

SiO 2 + 2NaOH = Na 2 SiO 3 + H 2 O.

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

a) soda ash, anhydrous soda, laundry soda or just soda - sodium carbonate Na 2 CO 3;
b) crystalline soda - sodium carbonate crystalline hydrate Na 2 CO 3. 10H 2 O;
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 bond ". During the lesson, the general chemical properties of metals, 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. Metallic bond

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

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. Let's look at some of the reactions involving metals.

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 = 2MgO

Rice. 1. Combustion of magnesium in oxygen

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

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

Rice. 2. Combustion of potassium in chlorine

Potassium burns in chlorine to form potassium chloride:

2K + Cl 2 = 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 be 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, for example, alkali and alkaline earth metals, 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 = 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 = ZnO + H 2

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

Metals in the voltage row to the right of hydrogen 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 = ZnCl 2 + H 2

Otherwise, the reactions of metals with nitric and concentrated sulfuric acids proceed. In these reactions, hydrogen is practically not evolved. We will burn 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 = 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 an external energy level.

In the nodes of the crystal lattice of metals there are not only neutral atoms, but also metal cations, since external electrons move freely along the crystal lattice. In this case, atoms, donating electrons, become cations, and cations, accepting electrons, turn into electrically neutral atoms.

Rice. 4. Model of the internal structure of metal

The chemical bond that forms as a result of the attraction of metal cations to freely moving electrons is called metal.

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 the chemical bond does not break under mechanical action, because a chemical bond is established not between specific atoms and cations, but between all metal cations with all free electrons in the 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. institutions / 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. institutions / 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 cl. / 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 / Chap. ed. V.A. Volodin, led. 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 the Textbook of P.A. Orzhekovsky. "Chemistry: 9th grade" (Moscow: 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 show oxidizing properties.

Obtaining metals:
- reduction from oxides with carbon (C), carbon monoxide (CO), hydrogen (H2) or a more active metal (Al, Ca, Mg);
- recovery from salt solutions with a more active metal;
- electrolysis of solutions or melts of metal compounds - reduction 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-activity 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 earths) and the most active non-metals (halogens, oxygen). Alkali metals (Na, K) react with oxygen to form peroxides and superoxides (Na2O2, KO2).

a) the interaction of metals with water.
At room temperature, alkali and alkaline earth metals interact with water. As a result of the substitution reaction, alkali (soluble base) and hydrogen are formed: Metal + H2O = Me (OH) + H2
When heated, the rest of the metals in the activity row to the left of hydrogen interact with water. Magnesium reacts with boiling water, aluminum - after special surface treatment, as a result, insoluble bases - magnesium hydroxide or aluminum hydroxide - are formed and hydrogen is released. Metals in the range of activity from zinc (inclusive) to lead (inclusive) interact with water vapor (i.e., above 100 C), thus forming oxides of the corresponding metals and hydrogen.
Metals in the line of activity to the right of hydrogen do not interact with water.
b) interaction with oxides:
active metals interact by a substitution reaction with oxides of other metals or non-metals, reducing them to simple substances.
c) interaction with acids:
Metals located in the row of activity to the left of hydrogen react with acids with the release of hydrogen and the formation of the corresponding salt. Metals in the activity row to the right of hydrogen 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 for noble metals (gold, platinum), can be oxidized by these oxidizing acids. As a result of these reactions, the corresponding salts, water and the product of nitrogen or sulfur reduction, respectively, will always be formed.
d) with alkalis
Metals that form amphoteric compounds (aluminum, beryllium, zinc) are capable of reacting with melts (in this case, medium salts of aluminates, beryllates or zincates are formed) or with alkali solutions (in this case, the corresponding complex salts are formed). In all reactions, hydrogen will be released.
e) In accordance with the position of the metal in the line of activity, reduction (displacement) reactions 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 non-metals.

There are much less 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 there are gaseous (fluorine, chlorine, oxygen, nitrogen, hydrogen), liquids (bromine) and solids, which differ greatly from each other in terms of their melting point. Most non-metals do not conduct electric current, but silicon, graphite, germanium have semiconducting 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-metallic 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 , carbyne, 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 electronegativity series. The most active non-metal is fluorine. It only exhibits 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:
Non-metals interact with metals. In such reactions, 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 anoxic 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 react 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 in the electronegativity series to the left of sulfur 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 with heating: 3Cl2 + 6NaOH = 5NaCl + NaClO3 + 3H2O.
e) with salts:
When interacting, they are strong oxidants and exhibit reducing properties.
Halogens (except for fluorine) enter into substitution reactions with solutions of salts of hydrohalic acids: the more active halogen displaces the less active halogen from the 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, are most different from the chemical properties of the rest of the elements of this group. Its chemical properties are in many ways even more similar to aluminum than to other metals of IIA group (the so-called "diagonal similarity"). Magnesium also differs markedly from Ca, Sr, Ba and Ra in chemical properties, but it still has much more similar chemical properties with them than with beryllium. Due to the significant similarity in the chemical properties of calcium, strontium, barium and radium, they are combined into one family, called alkaline earth metals.

All elements of the IIA group belong to s-elements, i.e. contain all their valence electrons on s-sub-level. Thus, the electronic configuration of the outer electron layer of all chemical elements of a given 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 the elements of group IIA, with participation in any chemical reactions, can only be oxidized, i.e. donate electrons:

Ме 0 - 2e - → Ме +2

Calcium, strontium, barium and radium are extremely reactive. The simple substances formed by them are very strong reducing agents. Magnesium is also a powerful 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 either with atmospheric oxygen or with 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 liquid inert to them, most often kerosene.

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

2Mg + O 2 = 2MgO

2Ca + O 2 = 2CaO

2Ba + O 2 = 2BaO

Ba + O 2 = BaO 2

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

with halogens

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

Mg + I 2 = MgI 2 - magnesium iodide

Ca + Br 2 = CaBr 2 - calcium bromide

Ba + Cl 2 = BaCl 2 - barium chloride

with non-metals of IV-VI groups

All metals of group IIA react when heated with all non-metals of IV-VI groups, 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 IIA metals, when carrying out its reactions with non-metals, it is required to significantly O higher temperature.

It should be noted that the reaction of metals with carbon can form carbides of different nature. Distinguish between carbides belonging to methanides and conditionally considered derivatives of methane, in which all hydrogen atoms are replaced by metal. They, like methane, contain carbon in the oxidation state -4, and during their hydrolysis or interaction with non-oxidizing acids, one of the products is methane. 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 - is obtained by the interaction of a particular metal with carbon depends on the size of the metal cation. With metal ions with a small radius, methanides are formed, as a rule, with ions of a larger size, acetylenides. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:

The rest of the II A group metals form acetylenides with carbon:

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

with hydrogen

All alkaline earth metals react with hydrogen when heated. In order for magnesium to react with hydrogen, heating alone, as is the case with alkaline earth metals, is not enough; in addition to a 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 when 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 during boiling, or even at red heat:

with non-oxidizing acids

All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the line of activity to the left of hydrogen. This forms the salt of the corresponding acid and hydrogen. Examples of reactions:

Be + H 2 SO 4 (dil.) = BeSO 4 + H 2

Mg + 2HBr = MgBr 2 + H 2

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

with oxidizing acids

- diluted nitric acid

All metals of group IIA 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 nitrogen oxide (I) (N 2 O), and in the case of highly dilute nitric acid, ammonium nitrate (NH 4 NO 3):

4Ca + 10HNO 3 ( smashed .) = 4Ca (NO 3) 2 + N 2 O + 5H 2 O

4Mg + 10HNO 3 (badly broken)= 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O

- concentrated nitric acid

Concentrated nitric acid passivates beryllium at ordinary (or low) temperatures, 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 rest of the metals of the main IIA group react with concentrated sulfuric acid under any conditions, including 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 acid concentration:

Mg + H 2 SO 4 ( end .) = MgSO 4 + SO 2 + H 2 O

3Mg + 4H 2 SO 4 ( end .) = 3MgSO 4 + S ↓ + 4H 2 O

4Ca + 5H 2 SO 4 ( end .) = 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. In this case, when the reaction is carried out in an aqueous solution, water also participates in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and gaseous hydrogen:

Be + 2KOH + 2H 2 O = H 2 + K 2 - potassium tetrahydroxoberyllate

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

Be + 2KOH = 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 reducing metals from their oxides with magnesium is called magnesium heat.

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.

The reactions contain only reducing agents. Metal atoms easily donate electrons to the outer (and some of the pre-outer) electron layer, turning into positive ions. Possible oxidation states of Ме Low 0, + 1, + 2, + 3 Highest + 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 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 GRAY

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. 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 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 + 3Н2

The rest of the metals react with water only in an incandescent 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-active 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