Features of lithium compounds in comparison with compounds of other alkali metals.

Hydrides, oxides, peroxides, hydroxides of alkali metals: chemical bond in compounds, getting and properties.

Obtaining sodium, sodium hydroxide and sodium carbonate in industry.

Interaction with alkali solutions: a) amphoteric metals; b) non-metals; c) acid oxides; d) amphoteric oxides.

Subgroup IA metals periodic system elements of I. I. Mendeleev Li, Na, K, Rb, Cs and Fr are called alkaline.

Alkali, alkaline earth metals, Be and Mg are among the most electropositive elements. In compounds with other elements, the oxidation state of + 1 is typical for metals of the IA subgroup, and +2 for the metals of the PA subgroup. With an increase in the number of electron layers and an increase in radii, the ionization energy of atoms decreases. As a result, the chemical activity of elements in subgroups increases with the growth of their serial number. The photoelectric effect characteristic of them is associated with low ionization energy, as well as their coloring with flame salts of a gas burner. Due to the easy return of external electrons, alkali and alkaline earth metals form compounds predominantly with an ionic bond.

Alkali and alkaline earth metals exhibit a high

chemical activity. When heated in hydrogen, they form

hydrides are salt-like compounds containing hydrogen

as a negatively charged ion. Air alkaline

metals are rapidly oxidized, forming, depending on their activity

oxides, peroxides, superoxides or ozonides.

At the same time, Ci, Na and K "light up in air or in an atmosphere of dry oxygen only

when heated, a, Rb and Cs spontaneously ignite without heating.

The formation during combustion of an oxide of the composition M 2 O is characteristic only

for lithium. Sodium forms a peroxide of the composition M 2 O 2, potassium, rubidium

and cesium - superoxides composition MO 2 .

Alkali metals react vigorously with water, displacing hydrogen from it and forming the corresponding hydroxides. The activity of the interaction of these metals with water increases as the atomic number of the element increases. So, - lithium reacts with water without melting, sodium - melts, potassium - ignites spontaneously, the interaction of rubidium and cesium proceeds even more vigorously. Alkali metals react vigorously with halogens, and when heated, with sulfur.

Alkali metal hydroxides - compounds with predominantly

but ionic. In aqueous solutions, they completely dissociate

The characteristic nature of the bond also explains their high thermal

stability: they do not split off water even when heated to a boiling point (above 1300 ° C) The exception is lithium hydroxide, which, when heated, decomposes with splitting of water. The behavior of lithium also differs in other respects from that of the other alkali metals. This is explained by its incomplete electronic analogy with the other elements of the group.

Of the alkali metals, only lithium, with relatively small when heated, it interacts with nitrogen, carbon and silicon, forming, respectively, Li 3 N nitride, Li 2 C 2 carbide and Li 6 Si 2 silicide. In the presence of moisture, nitride formation proceeds already at room temperature.

Unlike alkali metals, almost all salts of which are highly soluble in water, lithium forms poorly soluble fluoride LiF carbonate Li 2 CO 3 and phosphate Li 3 PO 4 .

Calcium, strontium and barium behave like alkali metals in relation to oxygen and water. They decompose water with the release of hydrogen and the formation of hydroxides M(OH) 2 . Interacting with oxygen, they form oxides (CaO) and peroxides (SrO 2, BaO 2), which react with water like similar alkali metal compounds.

Magnesium also differs significantly from the alkaline earth metals. For example, due to the low solubility of its hydroxide, it does not interact with cold water. When heated, the process is facilitated.

In general, the metals of the PA subgroup are chemically active: when heated, they interact with halogens and sulfur to form the corresponding salts, and combine with molecular nitrogen.

Alkaline earth metal salts, like alkali metal salts, are composed of ions. The salts of these metals color the flame of the burner. V characteristic colors, this is not observed for Be and Mg compounds.

Unlike alkali metal salts, many metal salts of the PA subgroup are sparingly soluble, in particular fluorides (except for BeF 2). sulfates (except for BeSO 4 and MgSO 4), carbonates. From aqueous solutions, Be 2+ precipitates in the form of basic carbonates of variable composition, Mg 2+ - in the form of 4MgCO 3 -Mg (OH) 2 -5H 2 O, and Ca 2+, Sr 2 + and Ba 2+ precipitate in in the form of medium carbonates MCO3.

A) Be+2NaOH= Na2BeO2+H2

Al+NaOH+H2O=NaAlO2+H2

B) Non-metals, with the exception of halogens, do not react with alkali solutions.

Cl2+NaOH=NaClO3+NaCl+H2O

IN) acidic oxides dissolve only in alkalis to form salt and water

SO3+2NaOH=Na2So4+H2o

G) Amf me react with strong alkalis, thereby showing their acidic properties, for example:

ZnO + 2NaOH → Na 2 ZnO 2 + H 2 O Amphoteric oxides can react with alkalis in two ways: in solution and in melt.

When reacted with an alkali in the melt, an ordinary medium salt is formed (as shown in the example above).

When reacting with alkali in solution, a complex salt is formed.

Al 2 O 3 + 2NaOH + 3H 2 O → 2Na (In this case, sodium tetrahydroxoalluminate is formed)

ALKALI METALS
SUB-GROUP IA. ALKALI METALS
LITHIUM, SODIUM, POTASSIUM, RUBIDIUM, CESIUM, FRANCE
The electronic structure of alkali metals is characterized by the presence of one electron on the outer electron shell, which is relatively weakly bound to the nucleus. Each alkali metal starts a new period in the periodic table. The alkali metal is able to donate its outer electron more easily than any other element of this period. The cut of an alkali metal in an inert medium has a bright silvery sheen. Alkali metals are characterized by low density, good electrical conductivity, and melt at relatively low temperatures (Table 2).
Due to their high activity, alkali metals do not exist in pure form, but occur in nature only in the form of compounds (excluding francium), for example, with oxygen (clays and silicates) or with halogens (sodium chloride). Chlorides are raw materials for obtaining alkali metals in a free state. Sea water contains ALKALINE METALS 3% NaCl and trace amounts of other salts. It is clear that the lakes inland seas, as well as underground salt deposits and brines contain alkali metal halides in higher concentrations than sea water. For example, the salt content in the waters of the Great Salt Lake (Utah, USA) is 13.827.7%, and in the Dead Sea (Israel) up to 31%, depending on the area of ​​the water surface, which varies with the season. It can be assumed that the insignificant content of KCl in sea ​​water compared to NaCl is explained by the assimilation of the K+ ion by marine plants.
In the free form, alkali metals are obtained by electrolysis of melts of salts such as NaCl, CaCl2, CaF2 or hydroxides (NaOH), since there is no more active metal capable of displacing the alkali metal from the halide. During the electrolysis of halides, it is necessary to isolate the metal released at the cathode, since at the same time gaseous halogen is released at the anode, which actively reacts with the released metal.
See also ALKALI PRODUCTION
Since alkali metals have only one electron on the outer layer, each of them is the most active in its period, so Li is the most active metal in the first period of eight elements, Na, respectively, in the second, and K is the most active metal of the third period, containing 18 elements (first transition period). In the alkali metal subgroup (IA), the ability to donate an electron increases from top to bottom.
Chemical properties. All alkali metals actively react with oxygen, forming oxides or peroxides, differing from each other in this: Li turns into Li2O, and other alkali metals into a mixture of M2O2 and MO2, while Rb and Cs ignite. All alkali metals form with hydrogen salt-like, thermally stable at high temperatures, hydrides of composition M + H, which are active reducing agents; hydrides are decomposed by water with the formation of alkalis and hydrogen and the release of heat, causing ignition of the gas, and the rate of this reaction for lithium is higher than for Na and K.
See also HYDROGEN; OXYGEN.
In liquid ammonia, alkali metals dissolve to form blue solutions, and (unlike the reaction with water) can be isolated again by evaporating ammonia or adding an appropriate salt (for example, NaCl from its ammonia solution). When reacting with gaseous ammonia, the reaction proceeds similarly to the reaction with water:

Alkali metal amides exhibit basic properties similar to hydroxides. Most alkali metal compounds, except for some lithium compounds, are highly soluble in water. In terms of atomic size and charge density, lithium is close to magnesium, so the properties of the compounds of these elements are similar. In terms of solubility and thermal stability, lithium carbonate is similar to magnesium and beryllium carbonates of subgroup IIA elements; these carbonates decompose at relatively low temperatures due to the stronger binding of MO. Lithium salts are better soluble in organic solvents (alcohols, ethers, petroleum solvents) than other alkali metal salts. Lithium (like magnesium) reacts directly with nitrogen to form Li3N (magnesium forms Mg3N2), while sodium and other alkali metals can only form nitrides under harsh conditions. The metals of subgroup IA react with carbon, but the most easy reaction is with lithium (apparently due to its small radius) and the least easy with cesium. Conversely, active alkali metals directly react with CO, forming carbonyls (for example, K(CO)x), while less active Li and Na only under certain conditions.
Application. Alkali metals are used both in industry and in chemical laboratories, for example, for syntheses. Lithium is used to produce hard light alloys, which differ, however, in brittleness. Large quantities sodium are consumed to obtain an Na4Pb alloy, from which tetraethyl lead Pb(C2H5)4 is obtained, an antiknock gasoline fuel. Lithium, sodium and calcium are used as components of soft bearing alloys. The only and therefore mobile electron on the outer layer makes alkali metals excellent conductors of heat and electricity. Potassium and sodium alloys, which remain liquid over a wide temperature range, are used as a heat exchange fluid in some types of nuclear reactors and, due to the high temperatures in a nuclear reactor, are used to produce steam. Sodium metal in the form of supply busbars is used in electrochemical technology to transmit high power currents. Lithium hydride LiH is a convenient source of hydrogen released as a result of the reaction of the hydride with water. Lithium aluminum hydride LiAlH4 and lithium hydride are used as reducing agents in organic and inorganic synthesis. Due to the small ionic radius and correspondingly high charge density, lithium is active in reactions with water, therefore lithium compounds are highly hygroscopic, and lithium chloride LiCl is used to dry the air during the operation of devices. Alkali metal hydroxides are strong bases, highly soluble in water; they are used to create alkaline environment. Sodium hydroxide, as the cheapest alkali, is widely used (in the USA alone, more than 2.26 million tons of it are consumed per year).
Lithium. The lightest metal, has two stable isotopes with atomic masses 6 and 7; the heavy isotope is more common, its content is 92.6% of all lithium atoms. Lithium was discovered by A. Arfvedson in 1817 and isolated by R. Bunsen and A. Mathisen in 1855. It is used in the production of thermonuclear weapons ( H-bomb), to increase the hardness of alloys and in pharmaceuticals. Lithium salts are used to increase the hardness and chemical resistance of glass, in the technology of alkaline batteries, and to bind oxygen during welding.
Sodium. Known since antiquity, it was isolated by H. Davy in 1807. It is a soft metal, its compounds such as alkali (sodium hydroxide NaOH), baking soda (sodium bicarbonate NaHCO3) and soda ash (sodium carbonate Na2CO3) are widely used. Metal is also used in the form of vapors in dim gas-discharge lamps for street lighting.
Potassium. Known since antiquity, it was also identified by H. Davy in 1807. Potassium salts are well known: potassium nitrate (potassium nitrate KNO3), potash (potassium carbonate K2CO3), caustic potash (potassium hydroxide KOH), etc. Potassium metal also finds various applications in technologies of heat exchange alloys.
Rubidium was discovered by spectroscopy by R. Bunsen in 1861; contains 27.85% radioactive rubidium Rb-87. Rubidium, like other metals of subgroup IA, is highly reactive and must be stored under a layer of oil or kerosene to avoid oxidation by atmospheric oxygen. Rubidium finds a variety of applications, including photovoltaic technology, radio vacuum devices and pharmaceuticals.
Cesium. Cesium compounds are widely distributed in nature, usually in small quantities together with compounds of other alkali metals. The mineral pollucite silicate contains 34% cesium oxide Cs2O. The element was discovered by R. Bunsen by spectroscopy in 1860. The main application of cesium is the production of photocells and electronic lamps, one of the radioactive isotopes of cesium Cs-137 is used in radiation therapy and scientific research.
France. The last member of the alkali metal family, francium, is so radioactive that it is not found in earth's crust in more than trace amounts. Information about francium and its compounds is based on the study of its insignificant amount, artificially obtained (at a high-energy accelerator) during the a-decay of actinium-227. The longest-lived isotope 22387Fr decays in 21 min into 22388Ra and b-particles. According to a rough estimate, the metallic radius of francium is 2.7 . Francium has most of the properties of other alkali metals and is highly electron-donating. It forms soluble salts and hydroxide. Francium exhibits oxidation state I in all compounds.

Collier Encyclopedia. - Open Society. 2000 .

alkali metals- these are elements of the 1st group of the periodic table of chemical elements (according to the outdated classification - elements of the main subgroup of group I): lithium Li, sodium Na, potassium K, rubidium rb, cesium cs, francium Fr, and ununenniy Uue. When alkali metals are dissolved in water, soluble hydroxides are formed, called alkalis.

Chemical properties of alkali metals

Due to the high chemical activity of alkali metals with respect to water, oxygen, and sometimes even nitrogen (Li, Cs), they are stored under a layer of kerosene. To carry out the reaction with an alkali metal, a piece of the required size is carefully cut off with a scalpel under a layer of kerosene, the metal surface is thoroughly cleaned from the products of its interaction with air in an argon atmosphere, and only then the sample is placed in the reaction vessel.

1. Interaction with water. An important property of alkali metals is their high activity with respect to water. Lithium reacts most calmly (without explosion) with water:

When carrying out a similar reaction, sodium burns with a yellow flame and a small explosion occurs. Potassium is even more active: in this case, the explosion is much stronger, and the flame is colored purple.

2. Interaction with oxygen. The combustion products of alkali metals in air have a different composition depending on the activity of the metal.

· Only lithium burns in air to form an oxide of stoichiometric composition:

・When burning sodium Na 2 O 2 peroxide is mainly formed with a small admixture of NaO 2 superoxide:

In combustion products potassium, rubidium And cesium contains mainly superoxides:

To obtain oxides of sodium and potassium, mixtures of hydroxide, peroxide or superoxide are heated with an excess of metal in the absence of oxygen:

For oxygen compounds of alkali metals, the following regularity is characteristic: as the radius of the alkali metal cation increases, the stability of oxygen compounds containing peroxide ion O 2 2− and superoxide ion O 2 − increases.

Heavy alkali metals are characterized by the formation of fairly stable ozonides composition of EO 3 . All oxygen compounds have different colors, the intensity of which deepens in the series from Li to Cs:

Alkali metal oxides have all the properties of basic oxides: they react with water, acidic oxides and acids:

Peroxides And superoxides exhibit the properties of strong oxidizers:

Peroxides and superoxides interact intensively with water, forming hydroxides:

3. Interaction with other substances. Alkali metals react with many non-metals. When heated, they combine with hydrogen to form hydrides, with halogens, sulfur, nitrogen, phosphorus, carbon and silicon to form, respectively, halides, sulfides, nitrides, phosphides, carbides And silicides:

When heated, alkali metals are able to react with other metals, forming intermetallics. Alkali metals react actively (with an explosion) with acids.

Alkali metals dissolve in liquid ammonia and its derivatives - amines and amides:

When dissolved in liquid ammonia, an alkali metal loses an electron, which is solvated by ammonia molecules and gives the solution a blue color. The resulting amides are easily decomposed by water with the formation of alkali and ammonia:

Alkali metals interact with organic matter alcohols (with the formation of alcoholates) and carboxylic acids (with the formation of salts):

4. Qualitative determination of alkali metals. Since the ionization potentials of alkali metals are low, when a metal or its compounds is heated in a flame, an atom is ionized, coloring the flame in certain color:

Coloring the flame with alkali metals
and their compounds

alkaline earth metals.

alkaline earth metals- chemical elements of the II group of the periodic table of elements: beryllium, magnesium, calcium, strontium, barium and radium.

Physical properties

All alkaline earth metals are gray, solid substances at room temperature. Unlike alkali metals, they are much harder, and they are mostly not cut with a knife (the exception is strontium). Density alkaline earth metals with a serial number grows, although growth is clearly observed only starting with calcium, which has the lowest density among them (ρ = 1.55 g / cm³), the heaviest is radium, whose density is approximately equal to the density of iron.

Chemical properties

Alkaline earth metals have an electronic configuration of the outer energy level ns², and are s-elements, along with alkali metals. Having two valence electrons, alkaline earth metals easily donate them, and in all compounds they have an oxidation state of +2 (very rarely +1).

The chemical activity of alkaline earth metals increases with increasing serial number. Beryllium in a compact form does not react with either oxygen or halogens, even at a red heat temperature (up to 600 ° C, an even higher temperature is needed to react with oxygen and other chalcogens, fluorine is an exception). Magnesium is protected by an oxide film at room temperature and higher (up to 650 °C) temperatures and does not oxidize further. Calcium oxidizes slowly and at room temperature in depth (in the presence of water vapor), and burns out with slight heating in oxygen, but is stable in dry air at room temperature. Strontium, barium, and radium rapidly oxidize in air to give a mixture of oxides and nitrides, so they, like the alkali metals (and calcium), are stored under a layer of kerosene.

Oxides and hydroxides of alkaline earth metals tend to increase in basic properties with increasing serial number: Be (OH) 2 - amphoteric, water-insoluble hydroxide, but soluble in acids (and also exhibits acidic properties in the presence of strong alkalis), Mg (OH) 2 - weak base, insoluble in water, Ca (OH) 2 - strong, but slightly soluble in water base, Sr (OH) 2 - more soluble in water than calcium hydroxide, strong base (alkali) at high temperatures close to the boiling point water (100 ° C), Ba (OH) 2 - a strong base (alkali), not inferior in strength to KOH or NaOH, and Ra (OH) 2 - one of the strongest alkalis, a very corrosive substance

Being in nature

All alkaline earth metals are present (in different quantities) in nature. Due to their high chemical activity, all of them are not found in the free state. The most common alkaline earth metal is calcium, the amount of which is 3.38% (of the mass of the earth's crust). Magnesium is slightly inferior to it, the amount of which is 2.35% (of the mass of the earth's crust). Barium and strontium are also common in nature, which, respectively, are 0.05 and 0.034% of the mass of the earth's crust. Beryllium is a rare element, the amount of which is 6·10 −4% of the mass of the earth's crust. As for radium, which is radioactive, it is the rarest of all alkaline earth metals, but it is always found in small quantities in uranium ores. In particular, it can be separated from there by chemical means. Its content is 1 10 −10% (of the mass of the earth's crust)

Aluminum.

Aluminum- an element of the main subgroup of the third group of the third period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 13. It is indicated by the symbol Al(lat. Aluminum). Belongs to the group of light metals. Most common metal and third most common chemical element in the earth's crust (after oxygen and silicon).

simple substance aluminum- light, paramagnetic silver-white metal, easily molded, cast, machined. Aluminum has a high thermal and electrical conductivity, resistance to corrosion due to the rapid formation of strong oxide films that protect the surface from further interaction.

Aluminum was first obtained by the Danish physicist Hans Oersted in 1825 by the action of potassium amalgam on aluminum chloride, followed by distillation of mercury. The modern method of obtaining was developed independently by the American Charles Hall and the Frenchman Paul Héroux in 1886. It consists in dissolving aluminum oxide Al 2 O 3 in a melt of Na 3 AlF 6 cryolite followed by electrolysis using consumable coke or graphite electrodes. This method of obtaining requires large amounts of electricity, and therefore was in demand only in the 20th century.

The production of 1000 kg of crude aluminum requires 1920 kg of alumina, 65 kg of cryolite, 35 kg of aluminum fluoride, 600 kg of anode mass and 17 thousand kWh of DC electricity

Alkali metals - francium, cesium, rubidium, potassium, sodium, lithium - are so called because they form alkalis when interacting with water. Because of their high reactivity, these elements should be stored under mineral oil or kerosene. The most active of all these substances is francium (possesses radioactivity).

Alkali metals are soft, silvery substances. Their freshly cut surface has a characteristic luster. Alkali metals boil and melt at low temperatures, have high thermal and electrical conductivity. They also have a low density.

Chemical properties of alkali metals

Substances are strong reducing agents, exhibit in compounds the oxidation state (single) +1. With the increase atomic mass alkali metals increases and reducing ability. Almost all compounds are soluble in water, all of them are ionic in nature.

When heated moderately, alkali metals ignite in air. In combination with hydrogen, substances form salt-like hydrides. Combustion products are usually peroxides.

Alkali metal oxides are solids yellow (rubidium and potassium oxides), white and lithium), and orange (cesium oxide) colors. These oxides are capable of reacting with water, acids, oxygen, acid and amphoteric oxides. These basic properties are inherent in all of them and are pronounced.

Alkali metal peroxides are yellowish-white powders. They are able to react with carbon dioxide and carbon monoxide, acids, non-metals, water.

Alkali metal hydroxides are white, water-soluble solids. In these compounds, the basic properties of alkalis are manifested (quite clearly). From lithium to francium, the strength of the bases and the degree of solubility in water increase. Hydroxides are considered fairly strong electrolytes. They react with salts, and oxides, individual non-metals. With the exception of the compound with lithium, all the others exhibit thermal stability. When calcined, it decomposes into water and oxide. These compounds are obtained by electrolysis of chloride aqueous solutions, a series of exchange reactions. Hydroxides are also obtained by the interaction of elements (or oxides) with water.

Almost all salts of the described metals (with the exception of individual lithium salts) are well soluble in water. Formed by weak acids, salt solutions have a medium reaction (alkaline) due to hydrolysis, formed by the same strong acids salts are not hydrolyzed. Common salts are stone silicate glue (soluble glass), Bertolet salt, potassium permanganate, drinking soda, soda ash and others.

All alkali metal compounds have the ability to change the color of the flame. This is applied in chemical analysis. So, the flame is colored with lithium ions, purple with potassium ions, yellow with sodium, whitish-pink with rubidium, violet-red with cesium.

Due to the fact that all alkaline elements are the strongest reducing agents, they can be obtained by electrolysis of molten salts.

Application of alkali metals

Elements are used in various fields of human activity. For example, cesium is used in solar cells. Lithium is used as a catalyst in bearing alloys. Sodium is present in gas discharge lamps, nuclear reactors as a heat carrier. Rubidium is used in research activities.

The most active among metals are alkali metals. They actively react with simple and complex substances.

General information

Alkali metals are in group I of the periodic table of Mendeleev. These are soft monovalent metals of a gray-silver color with a low melting point and low density. They show a single oxidation state +1, being reducing agents. Electronic configuration - ns 1 .

Rice. 1. Sodium and lithium.

The general characteristics of the metals of group I are given in the table.

List of alkali metals	Formula	Number	Period	t° sq. , °C	t° b.p. , °C	ρ, g/cm 3

Active metals quickly react with other substances, therefore, in nature they are found only in the composition of minerals.

Receipt

To obtain pure alkali metal, several methods are used:

electrolysis of melts, most often chlorides or hydroxides -

2NaCl → 2Na + Cl 2, 4NaOH → 4Na + 2H 2 O + O 2;

calcination of soda (sodium carbonate) with coal to obtain sodium -

Na 2 CO 3 + 2C → 2Na + 3CO;

calcium reduction of rubidium from chloride at high temperatures -

2RbCl + Ca → 2Rb + CaCl 2 ;

• reduction of cesium from carbonate using zirconium -

  2Cs 2 CO 3 + Zr → 4Cs + ZrO 2 + 2CO 2 .

Interaction

The properties of alkali metals are due to their structure. Being in the first group of the periodic table, they have only one valence electron in the outer energy level. The only electron easily passes to the oxidizing atom, which contributes to the rapid entry into the reaction.

Metallic properties increase in the table from top to bottom, so lithium parted with a valence electron more difficult than francium. Lithium is the hardest element among all alkali metals. The reaction of lithium with oxygen takes place only under the influence of high temperature. Lithium reacts with water much more slowly than other metals of the group.

Are common Chemical properties presented in the table.

Reaction	Products	The equation
With oxygen	Oxide (R 2 O) forms only lithium. Sodium forms a mixture of oxide and peroxide (R 2 O 2). The remaining metals form superoxides (RO 2)	4Li + O 2 → 2Li 2 O; 6Na + 2O 2 → 2Na 2 O + Na 2 O 2; K + O 2 → KO 2
With hydrogen		2Na + H 2 → 2NaH
	Hydroxides	2Na + 2H 2 O → 2NaOH + H 2
With acids		2Na + 2HCl → 2NaCl + H 2-
With halogens	Halides	2Li + Cl 2 → 2LiCl
With nitrogen (only lithium reacts at room temperature)		6Li + N 2 → 2Li 3 N
	Sulfides	2Na + S → Na 2 S
With carbon (only lithium and sodium react)		2Li + 2C → Li 2 C 2 ; 2Na + 2C → Na 2 C 2
with phosphorus		3K + P → K 3 P
With silicon	Silicides	4Cs + Si → Cs 4 Si
With ammonia		2Li + 2NH 3 → 2LiNH 2 + H 2

With a qualitative reaction, they have a different color of the flame. Lithium burns crimson, sodium yellow, and cesium a pink-violet flame. Alkali metal oxides also have different colors. Sodium turns white, rubidium and potassium turn yellow.

Rice. 2. Qualitative reaction of alkali metals.

Application

Simple metals and their compounds are used to make light alloys, metal parts, fertilizers, soda and other substances. Rubidium and potassium are used as catalysts. Sodium vapor is used in fluorescent lamps. Doesn't have practical application only francium due to its radioactive properties. How the elements of group I are used is briefly described in the table of the use of alkali metals.

Application area	Application
Chemical industry	Sodium speeds up the reaction in rubber production; Potassium and sodium hydroxide - soap production; Sodium and potassium carbonate - manufacture of glass, soap; Sodium hydroxide - making paper, soap, fabric; Potassium nitrate - fertilizer production
food industry	Sodium chloride - table salt; Sodium bicarbonate - baking soda
Metallurgy	Potassium and sodium are reducing agents in the production of titanium, zirconium, and uranium
Energy	Melts of potassium and sodium are used in nuclear reactors and aircraft engines; Lithium is used to make batteries
Electronics	Cesium - production of solar cells
Aviation and astronautics	Aluminum and lithium alloys are used for car bodies and rockets

Rice. 3. Drinking soda.

What have we learned?

From the 9th grade lesson, we learned about the features of alkali metals. They are in group I of the periodic table and give up one valence electron during reactions. These are soft metals that easily enter into chemical reactions with simple and complex substances - halogens, non-metals, acids, water. In nature, they are found only in the composition of other substances, therefore, electrolysis or a reduction reaction is used to extract them. Are applied in the industry, construction, metallurgy, power.

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