Alkali metals physical properties. Alkali metals - chemical and biogenic elements

Metals Common iron Iron is one of the most common elements in nature. Its content in the earth's crust is about 4.7% by weight, therefore iron, from the point of view of its occurrence in nature, is usually called a macroelement. In natural water, iron

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. With an increase in the atomic mass of alkali metals, 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 capable of reacting 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 structure of the outer electronic layers in the atoms of group I elements allows us, first of all, to assume that they do not have a tendency to add electrons. On the other hand, the donation of a single external electron, it would seem, should occur very easily and lead to the formation of stable monovalent cations of the elements in question.

As experience shows, these assumptions are fully justified only in relation to the elements of the left column (Li, Na, K and analogues). For copper and its analogues, they are only half true: in the sense of their lack of tendency to add electrons. At the same time, their 18-electron layer, which is farthest from the nucleus, turns out to be not yet completely fixed and, under certain conditions, is capable of partial loss of electrons. The latter makes it possible to exist, along with monovalent Cu, Agand Aualso compounds of the elements under consideration, corresponding to their higher valency.

Such a discrepancy between assumptions derived from atomic models and experimental results shows that consideration of the properties of elements based ononlythe electronic structures of atoms without taking into account other features are not always sufficient for the chemical characterization of these elements even in the roughest terms.

Alkali metals.

The name alkali metals applied to elements of the Li-Cs series is due to the fact that their hydroxides are strong alkalis. Sodium And potassium are among the most common elements, accounting for 2.0 and 1.1%, respectively, of the total number of atoms in the earth’s crust. Contents in it lithium (0,02%), rubidium (0.004%) and cesium (0.00009%) is already significantly less, and France - negligible. Elementary Na and K were isolated only in 1807. Lithium was discovered in 1817, cesium and rubidium - in 1860 and 1861, respectively. Element No. 87 - francium - was discovered in 1939, and received its name in 1946. Natural sodium and cesium are “pure” elements (23 Na and 133 Cs), lithium is composed of the isotopes 6 Li (7.4%) and 7 Li (92.6%), potassium is made of the isotopes 39 K (93.22%) .
40 K (0.01%) and 41 K (6.77%), rubidium - from the isotopes 85 Rb (72.2%) and 87 Rb (27.8%). Of the isotopes of francium, the most important is the naturally occurring 223 Fr (the average lifespan of an atom is 32 minutes).

Prevalence:

Only compounds of alkali metals are found in nature. Sodium and potassium are permanent constituents of many silicates. Of the individual minerals, sodium is the most important - salt (NaCl) is part of sea water and in certain areas of the earth's surface forms huge deposits of rock salt under a layer of alluvial rocks. The upper layers of such deposits sometimes contain accumulations of potassium salts in the form of layers sylvinite (mKCl∙nNaCl), ka rnallite (KCl MgCl 2 6H 2 O), etc., which serve as the main source for obtaining compounds of this element. Only a few natural accumulations of potassium salts of industrial importance are known. A number of minerals are known for lithium, but their accumulations are rare. Rubidium and cesium occur almost exclusively as impurities in potassium. Traces of France are always contained in uranium ores . Lithium minerals are, for example, spodumene And lepidolite (Li 2 KAl). Part of the potassium in the latter of them is sometimes replaced by rubidium. The same applies to carnallite, which can serve as a good source of rubidium. The relatively rare mineral is most important for cesium technology pollucite - CsAI(SiO 3) 2.

Receipt:

In their free state, alkali metals can be isolated by electrolysis of their molten chloride salts. Sodium is of primary practical importance, the annual world production of which is more than 200 thousand tons. The installation diagram for its production by electrolysis of molten NaCl is shown below. The bath consists of a steel casing with fireclay lining, a graphite anode (A) and an annular iron cathode (K), between which a mesh diaphragm is located. The electrolyte is usually not pure NaCl (mp 800 ℃), but a more fusible mixture of approximately 40% NaCl and 60% CaCl 2, which makes it possible to work at temperatures of about 580 ° C. Metallic sodium, which collects in the upper part of the annular cathode space and passes into the collector, contains a small (up to 5%) admixture of calcium, which is then almost completely released (the solubility of Ca in liquid sodium at its melting point is only 0.01%). As electrolysis progresses, NaCl is added to the bath. Electricity consumption is about 15 kWh per 1 kg Na.

2NaCl→ 2Na+Cl 2

This is interesting:

Before the introduction of the electrolytic method into practice, metallic sodium was obtained by heating soda with coal according to the reaction:

Na 2 CO 3 +2C+244kcal→2Na+3CO

The production of metallic K and Li is incomparably less than that of sodium. Lithium is obtained by electrolysis of the LiCl + KCl melt, and potassium is obtained by the action of sodium vapor on the KCl melt, which flows countercurrently to them in special distillation columns (from the upper part of which potassium vapor comes out). Rubidium and cesium are almost never mined on a large scale. To obtain small quantities of these metals, it is convenient to use heating of their chlorides with metallic calcium in a vacuum.

2LiCl→2Li+Cl 2

Physical properties:

In the absence of air, lithium and its analogs are silvery-white (with the exception of yellowish cesium) substances with a more or less strong metallic luster. All alkali metals are characterized by low densities, low hardness, low melting and boiling points, and good electrical conductivity. Their most important constants are compared below:

Due to their low density, Li, Na and K float on water (Li even on kerosene). Alkali metals are easily cut with a knife, and the hardness of the softest of them - cesium - does not exceed the hardness of wax. The non-luminous flame of a gas burner is colored by alkali metals and their volatile compounds in characteristic colors, of which the bright yellow inherent in sodium is the most intense.

This is interesting:

Externally manifested in the form of coloring of the flame, the emission of light rays by heated atoms of alkali metals is caused by the jump of electrons from higher to lower energy levels. For example, the characteristic yellow line in the spectrum of sodium appears when an electron jumps from the 3p level to the 3s level. Obviously, for such a jump to be possible, a preliminary excitation of the atom is necessary, that is, the transfer of one or more of its electrons to a higher energy level. In the case under consideration, excitation is achieved due to the heat of the flame (and requires an expenditure of 48 kcal/g-atom); in general, it can result from the imparting of energy of various types to the atom. Other alkali metals cause the appearance of the following flame colors: Li - carmine-red, K-violet, Rb - bluish-red, Cs - blue.

The luminescence spectrum of the night sky shows the constant presence of yellow sodium radiation. The altitude of the place of its origin is estimated at 200-300 km.T. That is, the atmosphere at these altitudes contains sodium atoms (of course, in negligible quantities). The occurrence of radiation is described by a number of elementary processes (the asterisk indicates the excited state; M is any third particle - O 2, O 0, N 2, etc.): Na + O 0 + M = NaO + M*, then NaO + O=O 2 + Na* and finally Na*= Na +λν.

Sodium and potassium should be stored in tightly closed containers under a layer of dry and neutral kerosene. Their contact with acids, water, chlorinated organic compounds and solid carbon dioxide is unacceptable. Do not accumulate small potassium scraps, which oxidize especially easily (due to their relatively large surface). Unused potassium and sodium residues in small quantities are destroyed by interaction with excess alcohol, in large quantities - by burning on the coals of a fire. Alkali metals that catch fire in a room are best extinguished by covering them with dry soda ash powder.

Chemical properties:

From the chemical point of view, lithium and its analogues are extremely reactive metals (and their activity usually increases in the direction from Li to Cs). In all compounds, alkali metals are monovalent. Located at the extreme left of the voltage series, they energetically interact with water according to the following scheme:

2E + 2H 2 O = 2EON + H 2

When reacting with Li and Na, the release of hydrogen is not accompanied by its ignition; for K it already occurs, and for Rb and Cs the interaction proceeds with an explosion.

· In contact with air, fresh sections of Na and K (to a lesser extent, Li) are immediately covered with a loose film of oxidation products. In view of this, Na and K are usually stored under kerosene. Na and K heated in air easily ignite, while rubidium and cesium spontaneously ignite even at ordinary temperatures.

4E+O 2 →2E 2 O (for lithium)

2E+O 2 →E 2 O 2 (for sodium)

E+O 2 →EO 2(for potassium, rubidium and cesium)

Practical application is mainly found in sodium peroxide (Na 2 0 2). Technically, it is obtained by oxidation at 350°C of atomized sodium metal:

2Na+O 2 →Na 2 O 2 +122kcal

· Melts of simple substances are capable of combining with ammonia to form amides and imides, solvates:

2Na melt +2NH 3 →2NaNH 2 +H 2 (sodium amide)

2Na melt +NH 3 →Na 2 NH+H 2 (sodium imide)

Na melt +6NH 3 → (sodium solvate)

When peroxides interact with water, the following reaction occurs:

2E 2 O 2 +2H 2 O=4EOH+O 2

The interaction of Na 2 O 2 with water is accompanied by hydrolysis:

Na 2 O 2 +2H 2 O→2NaOH + H 2 O 2 +34 kcal

This is interesting:

InteractionNa 2 O 2 with carbon dioxide according to the scheme

2Na 2 O 2 + 2CO 2 =2Na 2 CO 3 +O 2 +111 kcal

serves as the basis for the use of sodium peroxide as a source of oxygen in insulating gas masks and on submarines. Pure or containing various additives (for example, bleach mixed with Ni or C saltsu) sodium peroxide has the technical name "oxylitol". Mixed oxylit preparations are especially convenient for obtaining oxygen, which they release under the influence of water. Oxylitol compressed into cubes can be used to obtain a uniform flow of oxygen in a conventional apparatus for producing gases.

Na 2 O 2 +H 2 O=2NaOH+O 0 (atomic oxygen is released due to the decomposition of hydrogen peroxide).

Potassium superoxide ( KO 2) is often included in oxylitol. Its interaction with carbon dioxide in this case follows the overall equation:

Na 2 O 2 + 2KO 2 + 2CO 2 = Na 2 CO 3 + K 2 CO 3 + 2O 2 + 100 kcal, i.e. carbon dioxide is replaced by an equal volume of oxygen.

· Capable of forming ozonides. The formation of potassium ozonide-KO 3 follows the equation:

4KOH+3O 3 = 4KO 3 + O 2 +2H 2 O

It is a red crystalline substance and is a strong oxidizing agent. During storage, KO 3 decomposes slowly according to the equation 2NaO 3 →2NaO 2 +O 2 +11 kcal already under normal conditions. It instantly decomposes with water according to the overall scheme 4 KO 3 +2 H 2 O=4 KOH +5 O 2

· Capable of reacting with hydrogen to form ionic hydrides, according to the general scheme:

The interaction of hydrogen with heated alkali metals is slower than with alkaline earth metals. In the case of Li, heating to 700-800 °C is required, while its analogs interact already at 350-400 °C. Alkali metal hydrides are very strong reducing agents. Their oxidation by atmospheric oxygen in a dry state is relatively slow, but in the presence of moisture the process accelerates so much that it can lead to spontaneous ignition of the hydride. This especially applies to hydrides K, Rb and Cs. A violent reaction occurs with water according to the following scheme:

EN+ H 2 O= H 2 +EON

EH+O 2 →2EOH

When NaH or KH reacts with carbon dioxide, the corresponding salt of formic acid is formed:

NaH+CO 2 →HCOONa

Capable of forming complexes:

NaH+AlCl 3 →NaAlH 4 +3NaCl (sodium allanate)

NaAlH 4 → NaH+AlH 3

Normal alkali metal oxides (with the exception of Li 2 0) can be prepared only indirectly . They are solids of the following colors:

Na 2 O+2HCl=2NaCl+H 2 O

Alkali metal hydroxides (EOH) are colorless, very hygroscopic substances that corrode most materials that come into contact with them. Hence their sometimes used name in practice - caustic alkalis. When exposed to alkalis, the skin of the human body swells greatly and becomes slippery; with longer action, a very painful deep burn is formed. Caustic alkalis are especially dangerous for the eyes (it is recommended to wear safety glasses when working). Any alkali that gets on your hands or dress should be immediately washed off with water, then the affected area should be moistened with a very dilute solution of any acid and rinsed again with water.

All of them are relatively fusible and volatile without decomposition (except for LiOH, which eliminates water). To obtain hydroxide-alkali metals Electrolytic methods are mainly used. The most large-scale production is sodium hydroxide electrolysis concentrated aqueous table salt solution:

2NaCl+2H 2 O→2NaOH+Cl 2 +H 2

Ø Are typical grounds:

NaOH+HCl=NaCl+H2O

2NaOH+CO 2 =Na 2 CO 3 +H 2 O

2NaOH+2NO 2 =NaNO 3 +NaNO 2 +H 2 O

Ø Capable of forming complexes:

NaOH+ZnCl 2 = (ZnOH)Cl+NaCl

2Al+2NaOH+6H 2 O=2Na+3H 2

Al 2 O 3 + 6NaOH = 2Na 3 AlO 3 + 3H 2 O

Al(OH) 3 +NaOH=Na

Ø Capable of reacting with non-metals:

Cl 2 +2KOH=KCl+KClO+H 2 O (reaction occurs without heating)

Cl 2 +6KOH=5KCl+KClO 3 +3H 2 O (the reaction occurs with heating)

3S+6NaOH=2Na 2 S+Na 2 SO 3 +3H 2 O

Ø Used in organic synthesis (in particular, potassium and sodium hydroxide, sodium hydroxide is indicated in the examples):

NaOH+C 2 H 5 Cl=NaCl+C 2 H 4 (method for producing alkenes, ethylene (ethene) in this case), an alcohol solution of sodium hydroxide was used.

NaOH+C 2 H 5 Cl=NaCl+C 2 H 5 OH(a method for producing alcohols, ethanol in this case), an aqueous solution of sodium hydroxide was used.

2NaOH+C 2 H 5 Cl=2NaCl+C 2 H 2 +H 2 O (method for producing alkynes, acetylene (ethyne) in this case), an alcohol solution of sodium hydroxide was used.

C 6 H 5 OH (phenol) +NaOH= C 6 H 5 ONa+H 2 O

NaOH(+CaO)+CH 3 COONa→Na 2 CO 3 CH 4 (one of the methods for producing methane)

Ø You need to know the decomposition of several salts:

2KNO 3 →2KNO 2 +O 2

4KClO 3→ KCl+3KClO 4

2KClO 3→ KCl+3O 2

4Na 2 SO 3 →Na 2 S+3Na 2 SO 4

It is noteworthy that the decomposition of nitrates occurs approximately in the range of 450-600 ℃, then they melt without decomposition, but when reaching approximately 1000-1500 ℃, decomposition occurs according to the following scheme:

4LiNO 2 →2Li 2 O+4NO+O 2

This is interesting:

K 4 [ Fe(CN) 6 ]+ FeCl 3 = KFe[ Fe(CN) 6 ]+3 KCl(qualitative reaction toFe3+)

3K 4 +4FeCl 3 =Fe 4 3 +12KCl

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

4NaO 2 +2H 2 O=4NaOH+ 3O 2

4NaO 3 +2H 2 O=4NaOH+5O 2 (reaction of sodium ozonide with water )

2NaO 3 → 2NaO 2 +O 2(Decomposition occurs at different temperatures, for example: decomposition of sodium ozonide at -10 °C, cesium ozonide at +100°C)

NaNH 2 +H 2 O→ NaOH+NH 3

Na 2 NH+2H 2 O→ 2NaOH+NH 3

Na 3 N+3H 2 O→3NaOH+NH 3

KNO 2 +2Al+KOH+5H 2 O→2K+NH 3

2NaI + Na 2 O 2 + 2H 2 SO 4 →I 2 ↓+ 2Na 2 SO 4 + 2H 2 O

Fe 3 O 4 +4NaH=4NaOH+3Fe

5NaN 3 +NaNO 3 →8N 2 +3Na 2 O

Application:

Sodium is widely used in the synthesis of organic compounds and, partly, for the preparation of some of its derivatives. In nuclear technology it is used as a coolant.

Lithium is of absolutely exceptional importance for thermonuclear technology. In the rubber industry it is used in the production of artificial rubber (as a polymerization catalyst), in metallurgy - as a valuable additive to some other metals and alloys. For example, adding only hundredths of a percent of lithium greatly increases the hardness of aluminum and its alloys, and adding 0.4% lithium to lead almost triples its hardness without compromising bending resistance. There are indications that a similar cesium additive greatly improves the mechanical properties of magnesium and protects it from corrosion, but this is not the case with its use. Sodium hydride is sometimes used in metallurgy to isolate rare metals from their compounds. Its 2% solution in molten NaOH is used to remove scale from steel products (after a minute of holding in it, the hot product is immersed in water, which is reduced according to the equation

Fe 3 O 4 + 4NaH = 4NaOH + 3Fe (scale disappears).

Schematic diagram of a factory installation for producing soda by ammoniamethod (Solvay, 1863).

Limestone is fired in the furnace (L), and the resulting CO 2 enters the carbonization tower (B), and CaO is quenched with water (C), after which Ca(OH) 2 is pumped into the mixer (D), where it meets NH 4 Cl , this releases ammonia. The latter enters the absorber (D) and saturates a strong NaCl solution there, which is then pumped into the carbonization tower, where, when interacting with CO 2, NaHCO 3 and NH 4 Cl are formed. The first salt is almost completely precipitated and retained on the vacuum filter (E), and the second is pumped back into the mixer (D). Thus, NaCl and limestone are constantly consumed, and NaHCO 3 and CaCl 2 are obtained (the latter in the form of production waste). Sodium bicarbonate is then transferred by heating into soda.

Editor: Galina Nikolaevna Kharlamova