Examples of alkali metal salts. Alkali metals

Alkali metals - francium, cesium, rubidium, potassium, sodium, lithium - are so called because they form alkalis when interacting with water. Due to their high reactivity, these elements should be stored under a layer of mineral oil or kerosene. Francium is considered the most active of all these substances (it is radioactive).

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

Chemical properties of alkali metals

The substances are strong reducing agents and exhibit a (single) oxidation state of +1 in their compounds. As the atomic mass of alkali metals increases, the reducing ability also increases. Almost all compounds are soluble in water, all of them are ionic in nature.

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

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

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

Alkaline 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 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 compounds with lithium, all others exhibit thermal stability. When calcined, it decomposes into water and oxide. These compounds are obtained by electrolysis of aqueous chloride solutions and a series of exchange reactions. Hydroxides are also obtained by reacting elements (or oxides) with water.

Almost all salts of the described metals (with the exception of individual lithium salts) are highly soluble in water. Salt solutions formed by weak acids have a medium reaction (alkaline) due to hydrolysis, while salts formed by strong acids do not hydrolyze. Common salts are rock silicate glue (soluble glass), Berthollet salt, potassium permanganate, baking soda, soda ash and others.

All alkali metal compounds have the ability to change the color of the flame. This is used in chemical analysis. Thus, the flame is colored by lithium ions, violet by potassium ions, yellow by sodium, whitish-pink by rubidium, and violet-red by 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

The elements are used in various fields of human activity. For example, cesium is used in solar cells. Bearing alloys use lithium as a catalyst. Sodium is present in gas-discharge lamps and nuclear reactors as a coolant. Rubidium is used in scientific 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. These are soft monovalent metals of gray-silver color with a low melting point and low density. They exhibit a single oxidation state of +1, being reducing agents. Electronic configuration - ns 1.

Rice. 1. Sodium and lithium.

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

Active metals react quickly with other substances, so they are found in nature only in minerals.

Receipt

Several methods are used to obtain pure alkali metal:

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;

reduction of rubidium from chloride by calcium 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 determined by their structure. Being in the first group of the periodic table, they have only one valence electron in the outer energy level. A single electron easily goes 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 loses its valence electron more difficult than francium. Lithium is the hardest element among all alkali metals. The reaction of lithium with oxygen occurs only under the influence of high temperature. Lithium reacts with water much more slowly than the other metals in the group.

General chemical properties are presented in the table.

With a high-quality reaction, they have different flame colors. Lithium burns with a crimson flame, sodium with a yellow flame, and cesium with 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. Only francium has no practical use due to its radioactive properties. How group I elements are used is briefly described in the table on the use of alkali metals.

Rice. 3. Baking soda.

What have we learned?

From the 9th grade lesson we learned about the characteristics 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. They are found in nature only as part of other substances, so electrolysis or a reduction reaction is used to extract them. They are used in industry, construction, metallurgy, and energy.

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Alkali metals are the common name for elements of group 1 of the periodic table of chemical elements. Its composition is: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr), and a hypothetical element - ununennium (Uue). The name of the group comes from the name of soluble sodium and potassium hydroxides, which have an alkaline reaction and taste. Let's consider the general features of the structure of atoms of elements, properties, preparation and use of simple substances.

Outdated and new group numbering

According to the outdated numbering system, the alkali metals occupying the leftmost vertical column of the periodic table belong to group I-A. In 1989, the International Chemical Union (IUPAC) proposed a different option (long-period) as the main one. Alkali metals, in accordance with the new classification and continuous numbering, belong to group 1. This complex is opened by a representative of the 2nd period - lithium, and is completed by the radioactive element of the 7th period - francium. All metals of group 1 contain one s-electron in the outer shell of their atoms, which they easily give up (recover).

Structure of alkali metal atoms

Elements of group 1 are characterized by the presence of a second energy level, repeating the structure of the previous inert gas. Lithium has 2 electrons in the penultimate layer, and 8 electrons in the rest. In chemical reactions, atoms easily give up their outer s electron, acquiring an energetically favorable noble gas configuration. Group 1 elements have low ionization energies and electronegativity (EO). They easily form singly charged positive ions. When moving from lithium to francium, the number of protons and electrons and the radius of the atom increase. Rubidium, cesium and francium give up their outer electron more easily than the elements preceding them in the group. Consequently, in the group from top to bottom, the regenerative capacity increases.

The easy oxidation of alkali metals leads to the fact that elements of group 1 exist in nature in the form of compounds of their singly charged cations. The content of sodium in the earth's crust is 2.0%, potassium - 1.1%. Other elements are found in small quantities, for example, francium reserves - 340 g. Sodium chloride is dissolved in sea water, brine of salt lakes and estuaries, and forms deposits of rock or table salt. Along with halite, sylvinite NaCl occurs. KCl and sylvite KCl. Feldspar is formed by potassium aluminosilicate K2. Sodium carbonate is dissolved in the water of a number of lakes, and the reserves of sulfate of the element are concentrated in the waters of the Caspian Sea (Kara-Bogaz-Gol). There are deposits of sodium nitrate in Chile (Chilean saltpeter). There are a limited number of naturally occurring lithium compounds. Rubidium and cesium are found as impurities in compounds of group 1 elements, and francium is found in uranium ores.

Sequence of discovery of alkali metals

The British chemist and physicist G. Davy in 1807 carried out the electrolysis of alkali melts, obtaining sodium and potassium in free form for the first time. In 1817, the Swedish scientist Johann Arfvedson discovered the element lithium in minerals, and in 1825 G. Davy isolated the pure metal. Rubidium was first discovered in 1861 by R. Bunsen and G. Kirchhoff. German researchers analyzed the composition of aluminosilicates and obtained a red line in the spectrum corresponding to the new element. In 1939, Margarita Pere, an employee of the Paris Institute of Radioactivity, established the existence of the francium isotope. She named the element in honor of her homeland. Ununennium (eka-francium) is the tentative name of a new type of atom with atomic number 119. The chemical symbol Uue is temporarily used. Since 1985, researchers have been attempting to synthesize a new element, which will be the first in the 8th period, the seventh in the 1st group.

Physical properties of alkali metals

Almost all alkali metals have a silvery-white color and a metallic luster when freshly cut (cesium has a golden-yellow color). In air the luster fades and a gray film appears; on lithium it turns greenish-black. This metal has the greatest hardness among its group neighbors, but is inferior to talc, the softest mineral on the Mohs scale. Sodium and potassium are easy to bend and can be cut. Rubidium, cesium and francium in their pure form are a dough-like mass. Melting of alkali metals occurs at relatively low temperatures. For lithium it reaches 180.54 °C. Sodium melts at a temperature of 97.86 °C, potassium - at 63.51 °C, rubidium - at 39.32 °C, cesium - at 28.44 °C. The density of alkali metals is less than that of their related substances. Lithium floats in kerosene, rises to the surface of the water, potassium and sodium also float in it.

Crystalline state

Crystallization of alkali metals occurs in the cubic system (body-centered). The atoms in its composition have a conduction band, to the free levels of which electrons can move. It is these active particles that carry out a special chemical bond—metallic. The common structure of energy levels and the nature of crystal lattices explain the similarity of the elements of group 1. When moving from lithium to cesium, the masses of atoms of elements increase, which leads to a natural increase in density, as well as to a change in other properties.

Chemical properties of alkali metals

The only outer electron in alkali metal atoms is weakly attracted to the nucleus, so they are characterized by low ionization energy and negative or close to zero electron affinity. Elements of group 1, having reducing activity, are practically incapable of oxidizing. In the group from top to bottom, activity in chemical reactions increases:

Preparation and use of alkali metals

Metals belonging to group 1 are produced industrially by electrolysis of melts of their halides and other natural compounds. When decomposed by an electric current, positive ions at the cathode gain electrons and are reduced to free metal. At the opposite electrode, the anion is oxidized.

During the electrolysis of hydroxide melts at the anode, OH - particles are oxidized, oxygen is released and water is obtained. Another method is the thermal reduction of alkali metals from molten salts with calcium. Simple substances and compounds of elements of group 1 are of practical importance. Lithium serves as a raw material in nuclear energy and is used in rocket technology. In metallurgy it is used to remove residual hydrogen, nitrogen, oxygen, and sulfur. Hydroxide is used to supplement the electrolyte in alkaline batteries.

Sodium is necessary for nuclear energy, metallurgy, and organic synthesis. Cesium and rubidium are used in the manufacture of solar cells. Hydroxides and salts are widely used, especially chlorides, nitrates, sulfates, and carbonates of alkali metals. Cations have biological activity; sodium and potassium ions are especially important for the human body.

These are elements of group I of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr); very soft, ductile, fusible and light, usually silver-white in color; chemically very active; react violently with water, forming alkalis(hence the name).

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

Low densities,

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 Vaseline.

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 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 + H2

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

Preparation of alkali metals

Metal 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
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 effects

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 removal 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 by transferring an electron from lithium to another atom and forming an ionic bond with a small amount of covalency. Lithium is a typical metal element. In the form of a substance it is an alkali metal. It differs from other members of group I in its small size and the least activity compared to them. 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 dozen water molecules. In terms of the energy of solvation - the addition of solvent molecules, lithium is closer to a proton than to alkali metal cations.

The small size of the Li + ion, the high charge of the nucleus and only two electrons create conditions for the appearance of a fairly 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 high, the metal is capable of forming 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 greatest in the 3rd period. These two features determine the character 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, in compounds, sodium is present 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 adsorbability on the surface of the components that make up the earth's crust. Sodium compounds are only slightly susceptible to adsorption, while potassium compounds are firmly held by clay and other substances. Cell membranes, being the interface between the cell and the environment, are permeable to K + ions, as a result of which the intracellular concentration of K + is significantly 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. The emergence of cell membrane potential is associated with this circumstance. K + and Na + ions are one of the main components of the liquid phase of the body. Their relationship 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 the increase in its concentration never reaches dangerous values. The influence of Rb + , Cs + , Li + ions has not yet been sufficiently studied.

Of the various injuries associated with the use of alkali metal compounds, the most common are burns with hydroxide solutions. The effect of alkalis is associated with the dissolution of skin proteins in them and the formation of alkaline albuminates. The 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. Damage to the eyes, even with very dilute alkali solutions, is accompanied not only by superficial destruction, but also by damage to 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 fibrinous tracheobronchitis and pneumonia.

Potassium was obtained by G. Davy almost simultaneously with sodium in 1807 through the electrolysis of wet potassium hydroxide. The element got its name from the name of this compound – “caustic potassium”. The properties of potassium differ markedly from the properties 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 less 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. It is quite easy to determine the presence of potassium in a sample: vapors of the metal and its compounds color the flame violet-red. 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 spectra.

In 1861, while studying the salt of mineral springs by spectral analysis, Robert Bunsen discovered a new element. Its presence was proven by dark red lines in the spectrum, which were not produced by other elements. Based on 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 (tartrate) with soot. A feature of the element is the easy excitability of its atoms. Its electron emission appears under the influence of red rays of the visible spectrum. This is due to the slight difference in the energies of the atomic 4d and 5s orbitals. Of all the alkali elements that have 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 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. This is where the name “caesius” comes from, which means sky blue. It is the last element of the alkali metal subgroup that still occurs in measurable quantities. The largest atomic radius and the smallest first ionization potentials determine the character and behavior of this element. It has 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. The small difference in the energies of the atomic 5d and 6s orbitals causes the slight excitability of atoms. Electron emission from cesium is observed under the influence of invisible infrared rays (heat). This feature of the atomic structure determines good electrical conductivity of 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 solving the problem of thermonuclear fusion.

In air, lithium reacts actively not only with oxygen, but also with nitrogen and becomes 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 = 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.

To regenerate air in submarines and spaceships, in isolating gas masks and breathing apparatus of combat swimmers (underwater saboteurs), the Oxon mixture 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.

This is currently the standard filling for regenerating gas mask cartridges for firefighters.
Alkali metals react with hydrogen when heated, forming 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 temperatures (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 - crystalline hydrate of sodium carbonate 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.