Chemistry tests for grade 9 Topic: “Alkali metals and their compounds.”

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The most active among the metal group are alkali and alkaline earth metals. These are soft light metals that react with simple and complex substances.

general description

Active metals occupy the first and second groups of the periodic table. Complete list of alkali and alkaline earth metals:

• lithium (Li);
• sodium (Na);
• potassium (K);
• rubidium (Rb);
• cesium (Cs);
• francium (Fr);
• beryllium (Be);
• magnesium (Mg);
• calcium (Ca);
• strontium (Sr);
• barium (Ba);
• radium (Ra).

Rice.
1. Alkali and alkaline earth metals in the periodic table. The electronic configuration of alkali metals is ns1, alkaline earth metals are ns2.

Accordingly, the constant valency of alkali metals is I, alkaline earth metals are II. Due to the small number of valence electrons at the outer energy level, active metals exhibit powerful reducing properties, donating outer electrons in reactions. The more energy levels, the less connection from the outer electrons with the nucleus of the atom. Therefore, metallic properties increase in groups from top to bottom.

Due to their activity, metals of groups I and II are found in nature only in rocks. Pure metals are isolated using electrolysis, calcination, and substitution reactions.

As presented in nature

The alkaline group is represented on the planet in different ways:

• The most common elements are sodium and potassium.
• Lithium, rubidium, cesium are classified as rare and dispersed.
• The rarest alkali metal is francium. In terms of rarity, this radiogenic material is the second on the planet: the total volume in the earth’s crust does not exceed a third of a kilogram.

Due to their increased activity, alkali metals do not occur in nature. Only as connections with other elements.

Suppliers of sodium with potassium:

• Sodium chloride is common rock salt. It contains the water of the seas and oceans.
• Glauber's salt.
• Potassium salts are found in soils.

Metals are included in the structure of a number of minerals. These are mainly aluminosilicates: albite - sodium; orthoclase (feldspar) – potassium.

Physical properties

Alkali metals have a silvery-white color with a metallic luster. Cesium is a silvery-yellow metal. These are the most active and soft metals. Sodium, potassium, rubidium, cesium are cut with a knife. They resemble wax in their softness.

Rice. 2. Cutting sodium with a knife.

Alkaline earth metals are gray in color. Compared to alkali metals, they are harder, denser substances. Only strontium can be cut with a knife. The densest metal is radium (5.5 g/cm3).

The lightest metals are lithium, sodium and potassium. They float on the surface of the water.

Where are they used?

Industrialists appreciated the utilitarian properties of alkali metals. They are fusible, ductile (rolled out like foil), forged well, and transmit heat and electricity.

The most famous product is table salt (formula NaCl). It is supplemented with soda ash with caustic soda (carbonate, sodium hydroxide), potassium permanganate (potassium permanganate).

They are produced in millions of tons:

• Caustic soda (colloquially caustic soda) is an ingredient in soap making, aluminum production, and artificial fibers.
• Soda ash is a raw material for producing soap, glass, and a substitute for laundry soap.

Caustic and soda ash, Glauber's salt are inedible. Only table soda and table salt.

• Baking soda has found use as a home remedy for sore throats, an essential ingredient in baking, and a natural preservative.

Albite and orthoclase are classified as collection and decorative materials.

Chemical properties

Alkali and alkaline earth metals react with simple substances and complex compounds, forming salts, oxides, and alkalis. The main properties of active metals are described in the table.

Rice.
3. Reaction of potassium with water. Alkali and alkaline earth metals can be detected using a qualitative reaction. When burning, metals are painted a certain color. For example, sodium burns with a yellow flame, potassium with a violet flame, barium with a light green flame, and calcium with a dark orange flame.

§ 1. 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. The content of lithium (0.02%), rubidium (0.004%) and cesium (0.00009) in it is already significantly less, and france is negligible.

In nature, alkali metals are found exclusively in the form of compounds. Sodium and potassium are permanent constituents of many very common silicates. Of the individual minerals, sodium, the most important—table salt (NaCl)—is part of sea water and in certain areas of the earth’s surface forms huge deposits of so-called rock salt under a layer of alluvial rocks (in the USSR - Solikamsk, Artemoven, Iletsk, etc.). The upper layers of such deposits sometimes contain accumulations of potassium salts [in the form of minerals sylvinite (KCl–NaCl), carnallite (KCl MgCl2 6H2 O), etc.], which serve as the main source of compounds of this element. Only a few natural accumulations of potassium salts of industrial importance are known. The most important of them is the Solikamsk field in the USSR.

A number of minerals are known for lithium [for example, spodumene - LiAl (SiO3) 2 ], but their accumulations are rare. Rubidium and cesium occur almost exclusively as impurities in other alkali metals. Traces of francium are always found in uranium ores.

Sodium and potassium compounds are of great importance for life. Suffice it to recall that a person consumes 5–10 kg of NaCl annually. In the same way, plants need potassium salts. In this regard, about 90% of all mined potassium compounds are used to fertilize soils. The remaining 10%, as well as huge quantities of various sodium compounds, are used in industry. Only relatively small applications have so far been found for lithium derivatives and very limited applications for Rb and Cs compounds.

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 50 thousand tons.

1) The installation diagram for producing sodium metal by electrolysis of molten NaCl is shown in Fig. 184. The bath itself consists of a steel casing with a 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°C), but a more fusible mixture of it with 25% NaF and 12% KCl (which allows the process to be carried out at 610–650°C). Metallic sodium is collected in the upper part of the annular cathode space, from where it passes into the collector. As electrolysis progresses, NaCl is added to the bath.

In the absence of air, lithium and its analogs are silvery-white substances with a metallic luster. All alkali metals are characterized by low specific gravities, low hardness, low melting and boiling points, and good electrical conductivity. Their most important constants are compared in the table below. The course of change of constants in the Li–Cs series can be seen from Fig. 185.

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 softest of them, cesium, does not exceed wax in hardness.

All of them are soluble in mercury (the worst is lithium). The non-luminous flame of a gas burner is colored by alkali metals and their volatile compounds in characteristic colors: Li - carmine red, Na - bright yellow, K, Rb and Cs - pink-violet.

Lithium and its analogues are extremely reactive metals (and their activity usually increases in the direction from Li to Cs). In all their 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 + 2H2O = 2EON + H2

During the reaction with Li and Na, the release of hydrogen is not accompanied by its ignition; with K it already occurs, and with 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.

When Na and K are heated in air, they easily ignite, while rubidium and cesium spontaneously ignite even at ordinary temperatures.

In the presence of traces of moisture, alkali metals ignite in a chlorine atmosphere. The interaction of Cs, Rb and K with liquid bromine is accompanied by a strong explosion, while Na and Li react only superficially at ordinary temperatures. Reactions with iodine proceed vigorously only when heated. In all cases of interaction with halogens, the reaction product is the corresponding salt (EG). The formation of sulfide (E2 S) when rubbing an alkali metal with powder, sulfur, is accompanied by an explosion. When heated in a hydrogen atmosphere, lithium and its analogues form hydrides (H), which have the character of typical salts in which the negative ion is hydrogen (H–). Only lithium combines directly with nitrogen and carbon. The formation of its nitride (Li3 N) occurs slowly in a nitrogen atmosphere even at ordinary temperatures. In contrast, lithium carbide (Li2C2) can be obtained from elements only by heating. The heat generated by alkali metal compounds is compared in Fig. 186.

2) In appearance and most physical properties, alkali metal hydrides are similar to the corresponding halide salts. Thus, the best studied LiH forms solid colorless crystals (like NaCl), which melt without decomposition at 680 °C in the absence of air.

The salt-like nature of the hydrides in question was also directly proven by the evolution of hydrogen during the electrolysis of molten LiH at the anode.

Alkali metal hydrides are strong reducing agents, and when they interact with water, a violent reaction occurs according to the following scheme:

EN + H2O = H2 + EON

or in ions:

H– + H+ = H2

When alkali metals burn in excess oxygen, compounds of the following composition are formed:

Li2 O – white Na2 O2 – white KO2 – yellow RbO2 – yellow CsO2 – yellow

Of all these substances, only Li2O is a normal oxide, and the rest are peroxide compounds.

Practical application is mainly found in sodium peroxide (Na2 O2). It is obtained by burning sodium metal in aluminum vessels:

2Na + O2 = Na2 O2

The resulting product is usually a yellowish powder or grains.

The interaction of Na2 O2 with water is accompanied by hydrolysis:

Na2 O2 + 2H2 O < = >2NaOH + H2 O2

The release of H2O2 during this reaction is the basis for the use of sodium peroxide for bleaching various materials. Interaction of Na2 O2 with carbon dioxide according to the scheme

2Na2O2 + 2СО2 = 2Na2CO3 + O2

serves as the basis for the use of sodium peroxide as a source of oxygen in insulating gas masks and on submarines. Sodium peroxide reacts so vigorously with easily oxidized substances that an explosion can sometimes occur even with simple contact.

3) Pure sodium peroxide or containing various additives (for example, bleach mixed with Ni or Cu salts) has the technical name oxylit. Mixed oxylitol preparations are especially convenient for obtaining oxygen, which is released when interacting with water.

4) Peroxides K, Rb and Cs type EO2 are solids that melt at about 400 °C. All of them diffuse in air and are extremely energetic oxidizing agents.

Their interaction with dilute acids proceeds according to the following scheme:

2EO2 + H2 SO4 = E2 SO4 + H2 O2 + O2

The lattice of the EO2 peroxide crystal is similar to the CaC2 lattice (Fig. 138), i.e., it is formed by E+ and O2 ions. Potassium peroxide (KO2) is often included in xylitol. The interaction of the latter with carbon dioxide in this case follows the overall equation

Na2 O2 + 2KO2 + 2CO2 = Na2 CO3 + K2CO3 + 2O2

that is, carbon dioxide is replaced by an equal volume of oxygen.

Normal alkali metal oxides (with the exception of Li2O) can only be obtained indirectly. All of them are solids of the following colors:

Li2O – white

Na2O – white

K2 O – white

Rb2 O – yellow

Cs2 O – orange

Lithium oxide hydrates relatively slowly. On the contrary, the oxides of other alkali metals react with water very vigorously. The interaction proceeds according to the scheme

E2 O + H2 O = 2EON

and is accompanied by a large release of heat. Alkali metal oxides have not yet found direct use.

Alkali metal hydroxides (EOH) are colorless, very hygroscopic substances that corrode most materials that come into contact with them. Hence their name, often used in practice, is caustic alkalis.

Alkali metal hydroxides are highly soluble in water (and alcohol). Since they dissociate into ions to a greater extent than the hydroxides of all other metals, caustic alkalis are the strongest bases.

5) The properties of alkali metal hydroxides are compared below:

In contrast to the hydroxides of other metals, NaOH and its analogues do not melt water even when heated to their boiling temperatures (NaOH - 1388 °C, KOH - 1324 °C).

The exception is LiOH, which begins to lose water at about 600 °C. Since in the molten state caustic alkalis strongly corrode glass, porcelain and (with access to air) platinum dishes, vessels made of silver, nickel or iron are used to melt them.

6) Under the influence of caustic alkalis, the skin of the human body swells greatly and becomes slippery; In the future, a very painful deep burn will form. 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. Fabrics made from animal fibers are quickly destroyed by the action of alkalis, while plant fibers are quite resistant to them (on the contrary, acids corrode plant fibers faster than animal fibers).

Sodium hydroxide (otherwise known as caustic soda or caustic soda) is consumed by a wide variety of industries. Its annual global production amounts to several million tons, most of which is produced by electrolysis of NaCl solutions. Currently, the exchange decomposition of soda with slaked lime is used less frequently:

Na2 CO3 + Ca(OH)2 = CaCO3 + 2NaOH

The production of NaOH in the USSR amounted to 961 thousand tons in 1962 (against 190 thousand tons in 1940 and 55 thousand tons in 1913). Of the hydrates of oxides of other alkali metals, only KOH (caustic potash) finds significant practical use. It is produced by conventional electrolysis of KCl solutions.

7) Electrolysis of NaCl (and KCl) solutions is one of the main processes of the chemical industry, as it leads to the simultaneous production of two very important substances for technology - caustic alkali and free chlorine (cf. VII § 2). Hydrogen is also produced as a by-product.

The most important condition for the correct operation of an electrolytic installation is the absence of interaction between the resulting products (alkali and chlorine), which can be achieved by making it extremely difficult to mix the anodic and cathodic liquids. With the especially frequently used diaphragm method (see Fig. 96), the anode and cathode spaces are separated from each other by a diaphragm made of asbestos cardboard, which is highly permeable to liquids. The anode is made of graphite, the cathode is made of iron. During the electrolysis process, an alkali chloride solution is continuously fed into the anode space, and a solution of a mixture of alkali chloride and alkali continuously flows out of the cathode space. When it is evaporated, the chloride crystallizes. The resulting NaOH solution is evaporated in iron vats, after which the dry residue is melted down.

An example of an installation for the electrolysis of a NaCl solution using the so-called mercury method is the bath shown schematically in Fig. 187. The cathode space (K) in it is completely separated from the anode space (A) by blind side partitions and a layer of mercury lying on a thin plate of finely porous material (3). The anode is graphite electrodes (D), the cathode is the layer of mercury itself and an auxiliary iron electrode (G). During the electrolysis of a NaCl solution located in the anode space, chlorine is released at the anodes (D), and cations are discharged on mercury, and sodium amalgam is obtained. Since the latter has a lower density than mercury itself, it floats to the surface of the layer and interacts with the water filling the cathode space to form a caustic alkali and release hydrogen (at electrode G). The alkali obtained by this method is distinguished by its high purity.

Alkali metal ions are colorless. Almost all the salts they form with ordinary acids are highly soluble in water. In contrast to the K, Rb and Cs salts that are usually released without crystallization water, the formation of crystalline hydrates is very typical for lithium salts. Sodium occupies an intermediate position. Salts of alkali metals and weak acids show an alkaline reaction in solution due to hydrolysis. Complex compounds with an alkali metal ion as a complexing agent are known only for lithium. On the contrary, complex derivatives in which the alkali metal ions are located in the outer sphere are quite common. Many of these complexes are distinguished by their stability, which increases in the Li–Cs series.

The halide salts of the elements under consideration are rather refractory crystalline substances, with the exception of LiF (and partly NaF), highly soluble in water (Fig. 188). Of these, NaCl has the greatest practical importance. In addition to consumption with food (hence the name table salt), huge quantities of sodium chloride are used by industry. Its annual global consumption amounts to tens of millions of tons.

The sources of industrial production of NaCl are, on the one hand, natural deposits of rock salt, and on the other, seas and salt lakes (in the USSR - Baskunchak, etc.). The latter is simply broken out of rock salt deposits and then crushed. This salt is often so pure that it requires no further purification. NaCl is obtained from seas and salt lakes by evaporating brines under the influence of the sun or freezing water. The salt obtained in this way is often contaminated with impurities (mainly Ca2+, Mg2+ and SO42– ions) and becomes damp in humid air. On the contrary, completely pure table salt is non-hygroscopic. Of other alkali metal halides, KCl, the basis of potassium fertilizers, is of great importance.

Alkali metal nitrates are relatively fusible and highly soluble in water (Fig. 189). Of these, almost exclusively NaNO3 and KNO3 are of practical importance. Both salts are used mainly as mineral fertilizers. The first of them is usually obtained as a by-product of nitric acid production (due to the absorption of nitrogen oxides from exhaust gases by alkalis), and the second is obtained by the exchange decomposition of KCl and NaNO3.

Due to the dibasicity of carbonic acid, it forms two types of salts with alkali metals - acidic (ENCO3) and medium (E2 CO3). Acid carbonates (bicarbonates) are known for all alkali metals except Li. They are isolated from solutions without water of crystallization. When heated, bicarbonates quite easily transform into medium salts of carbonic acid:

2EHCO3 = E2 CO3 + CO2 + H2 O

Along the Na-Cs series, their thermal stability increases (Fig. 190).

With the exception of NaHCO3, the bicarbonates in question are highly soluble. Due to hydrolysis, their solutions show a very slightly alkaline reaction. When these solutions are heated, CO2 is partially released from them (in accordance with the decomposition equation given above), and the reaction becomes highly alkaline. In contact with air, this release of bicarbonates by CO2 solutions occurs very slowly even at ordinary temperatures. Practical application is found mainly in NaHCO3 (“baking soda”), used in medicine, the confectionery industry, etc.

Normal alkali metal carbonates, with the exception of Li2 CO3, are highly soluble in water (Fig. 191), and as a result of hydrolysis, their solutions show a strongly alkaline reaction. The most important is soda (Na2 CO3). It is produced either in an anhydrous state. (“soda ash”) or in the form of crystalline hydrate Na2 CO3 ·10H2 O (“crystalline soda”), which erodes in air.

Soda consumers are many industries. In addition, it is used in washing clothes, etc. The production of soda ash in the USSR amounted to 2.3 million tons in 1962 (compared to 536 thousand tons in 1940 and 160 thousand tons in 1913). Potash – K2 CO3 – is also consumed in significant quantities by industry (mainly glass).

The ammonia method, based on the reaction, is of primary importance for the production of soda

NaCl + NH4 HCO3 = NaHCO3 + NH4 Cl

The equilibrium of which, due to the relatively low solubility of NaHCO3, is shifted to the right. A concentrated NaCl solution is first saturated with ammonia and then treated with carbon dioxide (obtained by calcining CaCO3). The released NaHCO3 is filtered off and the wood products are heated in Na2 CO3, and the resulting carbon dioxide is returned to production. The mother liquor containing NH4 Cl is treated with slaked lime and the ammonia released during this process is also returned to production. Thus, the only waste is the CaCl2 remaining in solution.

9) A schematic diagram of a factory installation for producing soda using the ammonia method is shown in Fig. 192. In the furnace (A), limestone is fired, in which the resulting carbon dioxide 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 NH4 Cl and releases ammonia from it. The latter enters the absorber (D) and saturates a strong NaCl solution there, which is then pumped into the carbonization tower, where it interacts with CO2, forming NaHCO3 and NH4 Cl. Of both of these salts, the first is retained on the vacuum filter (E), and the second is again pumped into the mixer (D). Thus, NaCl and limestone are constantly consumed, and NaHCO3 and CaCl2 are obtained (the latter as production waste). Sodium bicarbonate is then transferred by heating into soda.

10) If there are natural sources of Na2 SO4, the older, sulfate method of soda production may also be profitable. The latter is carried out by alloying a mixture of Na2 SO4, limestone and coal at 1000 °C. In this case the reaction occurs:

Na2 SO4 + 2C + CaCO3 = 2СО2 + Na2 CO3 + CaS

Soda is separated from sparingly soluble CaS by treating the alloy with water. Production waste – CaS – can serve as a starting product for obtaining H2 S and then sulfur from it.

Like carbonates, alkali metal sulfates are also known to be acidic (EHSO4) and moderate (E2SO4). Both are highly soluble in water (Fig. 193). Of practical importance are mainly Na2 SO4 (often called simply sulfate in technology) and K2 SO4, especially the first of them. Their most important consumer is the glass industry. Na2 SO4 ·10H2 O crystalline hydrate (“mirabilite” or “Glauber’s salt”) is used in medicine as a laxative.

The industrial production of Na2 SO4 and K2 SO4 is based either on their isolation from natural minerals or on the treatment of the corresponding chlorides with sulfuric acid. In the latter case, sulfates are by-products of hydrochloric acid production. The waters of the Kara-Bogaz-Gol contain enormous quantities of mirabilite.

From a comparison of the properties of alkali metals and their compounds, it is clear that in the Li–Cs series they generally change quite regularly. Just as was observed in B and Be, the first element of the subgroup, lithium, occupies a somewhat special position. The low solubility of its salts with the anions CO32–, PO43– and F–, as well as hydroxide, the comparative ease of elimination from the latter water when heated, and some other properties bring lithium closer to magnesium and calcium. However, in general, lithium is still a typical alkali metal.

More on the topic:

• ALKALI METALS - chemical encyclopedia

What have we learned?

Alkali and alkaline earth metals are the most active metals. These are soft simple substances of gray or silver color with low density. Lithium, sodium, potassium float on the surface of the water. Alkaline earth metals are harder and denser than alkali metals. They oxidize quickly in air. Alkali metals form superoxides and peroxides; only lithium forms an oxide. Reacts violently with water at room temperature. They react with non-metals when heated. Alkaline earth metals react with oxides, displacing less active metals. Only beryllium reacts with alkalis.

Receiving technology

Alkali metals are obtained in several ways:

• Electrolysis. The material is melts of their chlorides (or other halides) or hydroxides. Salt solutions are not suitable as a source: hydrogen and alkalis become the final product.
• Reduction from bromide, chromate or chloride . The reducing agents are magnesium, zirconium, calcium, and silicon. The production process takes place in a vacuum at temperatures below 1000°C, and the resulting metal is periodically distilled off.

Sodium is recovered from carbonate. Ingredients: coal, limestone, temperature 990°C. For industrial needs, hydroxide is synthesized from a strong solution of table salt.

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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.

Qualitative reactions

Detection of lithium cations

Metal ions in salt can be detected using a qualitative reaction.

Phosphoric acid is used to detect lithium cations. The resulting white lithium phosphate is soluble only in concentrated nitric acid and ammonium salts:

3Li+ + PO43- = Li 3PO 4↓

In soluble salts, lithium can be detected using ammonium fluoride. The reaction produces a white insoluble precipitate of lithium fluoride:

Li+ + F- = LiF↓

Detection of sodium cation

Sodium ions can be detected by reaction with the complex salt hexahydroxostibate (V). At low temperatures and in a neutral environment, a finely crystalline white sodium salt is formed:

Na+ + – = Na[Sb(OH)6] ↓

This reaction is carried out on a glass slide. The formation and shape of crystals is observed under a microscope or binocular.

Detection of potassium cation

In acidic and neutral environments, potassium ions form a double complex salt with sodium hexanitrocabaltate (III). This salt is insoluble in water and has a yellow color:

2K+ + Na3[Co(NO2)6] = NaK2[Co(NO2)6] ↓+ 2Na+

The peculiarity of the reaction is its slow flow. The reaction mixture is left for several hours. To speed up the reaction, you can rub a glass rod along the wall of the test tube. This leads to the formation of static electricity, which increases the attraction of the ions to each other.

Another way to detect potassium cations is the reaction with sodium hydrogen tartrate NaHC4H4O6. In this case, a white salt is formed. The reaction can be accelerated by rubbing a glass rod against the surface of the test tube:

K+ + NaHC4H406= KHC4H4O6↓ + Na+.

Detection of rubidium and cesium cations

Rubidium and cesium cations are highly active ions and therefore are not detected by qualitative reactions. All compounds are transparent and highly soluble in water. The main methods for their detection in salt composition are precipitation with organic aromatic solutions. There is currently no method for separating cesium and rubidium from a mixture.

France ion detection

As a radioactive metal, francium is not included in ordinary salts. Its detection is carried out using spectrometry and radiation instruments.