Interaction of alkali and alkaline earth metals with water. Alkaline earth metals

The elements of the calcium subgroup are called alkaline earth metals. The origin of this name is due to the fact that their oxides (“earths” of alchemists) impart an alkaline reaction to water. Alkaline earth metals are most often classified as onlycalcium , strontium, barium and radium , less often magnesium . The first element of this subgroup, beryllium , in most properties it is much closer to aluminum.

Prevalence:

Calcium accounts for 1.5% of the total number of atoms in the earth's crust, while the content of radium in it is very small (8-10-12%). Intermediate elements - strontium (0.008) and barium (0.005%) - are closer to calcium. Barium was discovered in 1774, strontium in 1792. Elementary Ca, Sr and Ba were first obtained in 1808. Natural calcium th is composed of isotopes with mass numbers 40 (96.97%), 42 (0.64), 43 (0.14), 44 (2.06), 46 (0.003), 48 (0.19); strontium - 84 (0,56%), 86 (9,86), 87 (7,02), 88 (82,56); barium -130 (0.10%), 132 (0.10), 134 (2.42), 135 (6.59), 136 (7.81), 137 (11.32), 138 (71.66) . From isotopes radium Of primary importance is the naturally occurring 226 Ra (the average lifespan of an atom is 2340 years).

Calcium compounds (limestone, gypsum) were known and practically used in ancient times. In addition to various silicate rocks, Ca, Sr and Ba are found mainly in the form of their sparingly soluble carbon dioxide and sulfate salts, which are the minerals:

CaC0 3 - calcite CaS0 4 - en hydrite

SrC0 3 - strontianite SrS0 4 - celestine

ВаС0 3 - witherite BaS0 4 - heavy spar

CaMg(CO 3) 2 - dolomite MgCO 3 - magnesite

Calcium carbonate in the form of limestone and chalk sometimes forms entire mountain ranges. Much less common is the crystallized form of CaCO 3 - marble. For calcium sulfate, it is most typical to be found in the form of the mineral gypsum (CaSO 4 2H 2 0), deposits of which are often extremely powerful. In addition to those listed above, an important calcium mineral is fluorite -CaF 2, used to obtain hydrofluoric acid according to the equation:

CaF 2 +H 2 SO 4 (conc.) →CaSO 4 +HF

For strontium and barium, sulfate minerals are more common than carbon dioxide minerals. Primary deposits of radium are associated with uranium ores (and per 1000 kg of uranium, the ore contains only 0.3 g of radium).

Receipt:

Aluminothermic production of free alkaline earth metals is carried out at temperatures of about 1200 °C according to the following scheme:

ZE0 + 2Al=Al 2 O 3 +ZE

by heating their oxides with aluminum metal in a high vacuum. In this case, the alkaline earth metal is distilled off and deposited on the cooler parts of the installation. On a large scale (about thousands of tons annually), only calcium is produced, which is also obtained by electrolysis of molten CaCl 2. The process of aluminothermy is complicated in that it involves partial fusion with Al 2 O 3. For example, in the case of calcium, the reaction proceeds according to the equation:

3CaO + Al 2 O 3 →Ca 3 (AlO 3) 2

Partial fusion of the resulting alkaline earth metal with aluminum may also occur.

Electrolyzer for the production of metallic calcium is a furnace with an internal graphite lining, cooled from below by running water. Anhydrous CaCl 2 is loaded into the furnace, and the electrodes are an iron cathode and graphite anodes. The process is carried out at a voltage of 20-30V, current up to 10 thousand amperes, low temperature (about 800 °C). Thanks to the latter circumstance, the graphite lining of the furnace remains always covered with a protective layer of solid salt. Since calcium is deposited well only at a sufficiently high current density on the cathode (about 100 A/cm 3), the latter is gradually raised upward as the electrolysis progresses, so that only its end remains immersed in the melt. Thus, in fact, the cathode is the metallic calcium itself (which is isolated from the air by a solidified salt crust). Its purification is usually carried out by distillation in a vacuum or in an argon atmosphere.

Physical properties:

Calcium and its analogs are malleable, silvery-white metals. Of these, calcium itself is quite hard, strontium and especially barium are much softer. Some constants of alkaline earth metals are compared below:

Density, g/cm 3
Melting point, °C
Boiling point, °C

Volatile compounds of alkaline earth metals color the flame in characteristic colors: Ca - orange-red (brick), Sr and Ra - carmine-red, Ba - yellowish-green. This is used in chemical analyzes to discover the elements in question.

Chemical properties :

In air, calcium and its analogues are covered with a film, along with normal oxides (EO), partially also containing peroxides (E0 2) and nitrides (E 3 N 2). In the voltage series, alkaline earth metals are located to the left of magnesium and therefore easily displace hydrogen not only from dilute acids, but also from water. When going from Ca to Ra, the energy of interaction increases. In their compounds, the elements in question are divalent. Alkaline earth metals combine with metalloids very energetically and with a significant release of heat.

· Usually, when alkali earth metals interact with oxygen, the formation of an oxide is indicated:

2E +O 2 →2EO

It is important to know the trivial names of several compounds:

bleach, bleach (bleach) – CaCl 2 ∙ Ca(ClO) 2

slaked (fluff) – Ca(OH) 2

lime - a mixture of Ca(OH) 2, sand and water

milk of lime – suspension of Ca(OH) 2 in lime water

soda - a mixture of solid NaOH and Ca(OH) 2 or CaO

quicklime (boiling liquid) – CaO

· Interaction with water, using the example of calcium and its oxide:

Ca+2H 2 O→Ca(OH) 2 +H 2

CaO+H 2 O→Ca(OH) 2 +16 kcal (“slaking” lime)

When interacting with acids, oxides and hydroxides of alkaline earth metals easily form the corresponding salts, which are usually colorless.

This is interesting:

When slaking lime, if you replace water with a NaOH solution, you get so-called soda lime. In practice, during its production, crushed CaO is added to a concentrated solution of sodium hydroxide (in a weight ratio of 2:1 to NaOH). After mixing the resulting mass, it is evaporated to dryness in iron vessels, lightly calcined and then crushed. Soda lime is a close mixture Ca(OH) 2 with NaOH and is widely used in laboratories to absorb carbon dioxide.

Along with normal oxides, white peroxides of the E0 2 type are known for elements of the calcium subgroup. Of these, barium peroxide (Ba02) is of practical importance, used, in particular, as a starting product for the production of hydrogen peroxide:

BaO 2 + H 2 SO 4 = BaSO 4 + H 2 O 2

Technically, Ba0 2 is obtained by heating BaO in a stream of air to 500 °C. In this case, oxygen is added according to the reaction

2BaO + O 2 = 2BaO 2 + 34 kcal

Further heating leads, on the contrary, to the decomposition of Ba0 2 into barium oxide and oxygen. Therefore, the combustion of barium metal is accompanied by the formation of only its oxide.

· Interaction with hydrogen to form hydrides:

EN 2 hydrides do not dissolve (without decomposition) in any of the usual solvents. They react vigorously with water (even traces of it) according to the following scheme:

EH 2 + 2H 2 O = E(OH) 2 + 2H 2

This reaction can serve as a convenient method for producing hydrogen, since for its implementation it requires, in addition to CaH 2 (1 kg of which gives approximately 1 m 3 H 2), only water. It is accompanied by such a significant release of heat that CaH 2 moistened with a small amount of water spontaneously ignites in air. The interaction of EN 2 hydrides with dilute acids occurs even more vigorously. On the contrary, they react more calmly with alcohols than with water:

CaH 2 +2HCl→CaCl 2 +2H 2

CaH 2 +2ROH→2RH+Ca(OH) 2

3CaH 2 +N 2 → Ca 3 N 2 +ЗH 2

CaH 2 +O 2 →CaO+H 2 O

Calcium hydride is used as an effective drying agent for liquids and gases. It is also successfully used for the quantitative determination of water content in organic liquids, crystalline hydrates, etc.

· Can interact directly with non-metals:

Ca+Cl 2 →CaCl 2

· Interaction with nitrogen. E 3 N 2 white refractory bodies. They form very slowly even under normal conditions:

3E+N 2 →E 3 N 2

They decompose with water according to the following scheme:

E 3 N 2 +6H 2 O→3Ca(OH) 2 +2NH 3

4E 3 N 2 →N 2 +3E 4 N 2 (for Ba and Sr subnitrides)

E 4 N 2 +8H 2 O→4E(OH) 2 +2NH 3 +H 2

Ba 3 N 2 +2N 2 →3 Ba N 2 (barium pernitride)

When interacting with dilute acids, these pernitrides, along with two molecules of ammonia, also split off a molecule of free nitrogen:

E 4 N 2 +8HCl→4ESl 2 +2NH 3 +H 2

E 3 N 2 + ZSO = 3EO + N 2 + ZS

The reaction is different in the case of barium:

B a 3 N 2 + 2CO = 2BaO + Ba(CN) 2

This is interesting :

E+NH 3(liquid) →(E(NH 2) 2 +H 2 +ENH+H 2)

4E(NH 2) 2 → EN 2 +2H 2

I wonder whatE(NH 3) 6 - ammonia compounds are formed by the interaction of elements with gaseous ammonia, and are capable of decomposition according to the following scheme:

E(NH 3) 6 →E(NH 2) 2 +4NH 3 +H 2

Further heating:

E(NH 2) 2 →ENH+NH 3

3ENH→NH 3 +E 3 N 2

But the interaction of metal with ammonia at high temperatures proceeds according to the following scheme:

6E+2N.H. 3 →EH 2 +E 3N 2

Nitrides are capable of adding halides:

E 3 N 2 + EHal 2 → 2 E 2 NHal

· Alkali metal oxides and hydroxides exhibit basic properties, with the exception of beryllium:

CaO+2 HCl→СаСl 2 +H2O

Ca(OH) 2 +2HCl→SaSl 2 +2H 2 O

Be+2NaOH+2H 2 O→Na 2 +H 2

BeO+2HCl→BeWITHl 2 +H 2 O

BeO+2NaOH→Na 2 BeO 2 +H 2 O

· Qualitative reactions to alkaline metal cations. Most publications indicate only qualitative reactions to Ca 2+ and Ba 2+. Let us consider them immediately in ionic form:

Ca 2+ +CO 3 2- →CaCO 3 ↓ (white precipitate)

Ca 2+ +SO 4 2- →CaSO 4 ↓ (white flocculent precipitate)

CaCl 2 + (NH 4) 2 C 2 O 4 →2NH 4 Cl + CaC 2 O 4 ↓

Ca 2+ +C 2 O 4 2- → CaC 2 O 4 ↓ (white precipitate)

Ca 2+ - coloring of the flame brick color

Ba 2+ +CO 3 2- →BaCO 3 ↓ (white precipitate)

Ba 2+ +SO 4 2- →BaSO 4 ↓ (white precipitate)

Ba 2+ +CrO 4 2- →BaCrO 4 ↓ (yellow precipitate, similar for strontium)

Ba 2+ +Cr 2 O 7 2- +H 2 O→2BaCrO 4 +2H + (yellow precipitate, similar for strontium)

Ba 2+ - coloring the flame green.

Application:

Industrial application is found almost exclusively in compounds of the elements under consideration, the characteristic properties of which determine the areas of their use. An exception is radium salts, the practical significance of which is associated with their general property - radioactivity. Practical use (mainly in metallurgy) is almost exclusively calcium. Calcium nitrate is widely used as a nitrogen-containing mineral fertilizer. Strontium and barium nitrates are used in pyrotechnics for the manufacture of compositions that burn with red (Sr) or green (Ba) flames. The use of individual natural varieties of CaCO 3 is different. Limestone is directly used in construction work, and also serves as the raw material for the production of the most important building materials - lime and cement. Chalk is used as a mineral paint, as a base for polishing compounds, etc. Marble is an excellent material for sculpting, making electrical distribution boards, etc. Practical application is found mainly in natural CaF 2, which is widely used in the ceramic industry and serves as the starting material for the production of HF.

Due to its hygroscopicity, anhydrous CaCl 2 is often used as a drying agent. The medical uses of calcium chloride solutions (intravenously and intravenously) are very diverse. Barium chloride is used to control agricultural pests and as an important reagent (for SO 4 2- ion) in chemical laboratories.

This is interesting:

If 1 wt. quickly pour a saturated solution of Ca(CH 3 COO) 2 into a vessel containing 17 wt. parts of ethyl alcohol, then the entire liquid immediately solidifies. The “dry alcohol” obtained in this way, after ignition, burns slowly with a non-smoking flame. This fuel is especially convenient for tourists.

Water hardness.

The content of calcium and magnesium salts in natural water is often assessed in terms of its “hardness”. In this case, a distinction is made between carbonate (“temporary”) and non-carbonate (“permanent”) hardness. The first is due to the presence of Ca(HC0 3) 2, less often Mg(HC0 3) 2. It is called temporary because it can be eliminated by simply boiling water: the bicarbonates are destroyed, and the insoluble products of their decomposition (Ca and Mg carbonates) settle on the walls of the vessel in the form of scale:

Ca(HCO 3) 2 →CaCO 3 ↓+CO 2 +H 2 O

Mg(HCO 3) 2 →MgCO 3 ↓+CO 2 +H 2 O

The constant hardness of water is due to the presence of calcium and magnesium salts in it, which do not produce sediment when boiled. The most common are sulfates and chlorides. Of these, poorly soluble CaS0 4 is of particular importance, which settles in the form of a very dense scale.

When a steam boiler operates on hard water, its heated surface becomes covered with scale. Since the latter conducts heat poorly, first of all, the operation of the boiler itself becomes uneconomical: already a layer of scale 1 mm thick increases fuel consumption by approximately 5%. On the other hand, the boiler walls, insulated from water by a layer of scale, can heat up to very high temperatures. In this case, the iron gradually oxidizes and the walls lose strength, which can lead to an explosion of the boiler. Since steam power systems exist in many industrial enterprises, the issue of water hardness is very practically important.

Since purifying water from dissolved salts by distillation is too expensive, in areas with hard water they use chemical methods to “soften” it. Carbonate hardness is usually eliminated by adding Ca(OH) 2 to water in an amount that strictly corresponds to the bicarbonate content determined by analysis. At the same time, according to the reaction

Ca(HCO 3) 2 + Ca(OH) 2 = 2CaCO 3 ↓ + 2H 2 O

all bicarbonate turns into normal carbonate and precipitates. Non-carbonate hardness is most often removed by adding soda to water, which causes the formation of a precipitate by the reaction:

СaSO 4 + Na 2 CO 3 = CaCO 3 ↓ + Na 2 SO 4

The water is then allowed to settle and only after that is it used to power boilers or in production. To soften small amounts of hard water (in laundries, etc.), you usually add a little soda to it and let it sit. In this case, calcium and magnesium are completely precipitated in the form of carbonates, and the sodium salts remaining in the solution do not interfere.

From the above it follows that soda can be used to eliminate both carbonate and non-carbonate hardness. Nevertheless, in technology they still try to use Ca(OH) 2 whenever possible, which is due to the much lower cost of this product compared to soda

Both carbonate and non-carbonate hardness of water is estimated by the total number of milligram equivalents of Ca and Mg contained in one liter (mg-eq/l). The sum of temporary and permanent hardness determines the total hardness of the water. The latter is characterized by this characteristic by the following names: soft (<4), средне жёсткая (4-8), жесткая (8-12), очень жесткая (>12 mEq/l). The hardness of individual natural waters varies within very wide limits. For open reservoirs, it often depends on the time of year and even the weather. The “softest” natural water is atmospheric (rain, snow), which contains almost no dissolved salts. Interestingly, there is evidence that heart disease is more common in areas with soft water.

To completely soften water, instead of soda, Na 3 PO 4 is often used, which precipitates calcium and magnesium in the form of their sparingly soluble phosphates:

2Na 3 PO 4 +3Ca(HCO 3) 2 →Ca 3 (PO 4) 2 ↓+6NaHCO 3

2Na 3 PO 4 +3Mg(HCO 3) 2 →Mg 3 (PO 4) 2 ↓+6NaHCO 3

There is a special formula for calculating water hardness:

Where 20.04 and 12.16 are the equivalent masses of calcium and magnesium, respectively.

Editor: Galina Nikolaevna Kharlamova

Properties of alkaline earth metals

Physical properties

Alkaline earth metals (compared to alkali metals) have higher t╟pl. and boiling point, ionization potentials, densities and hardness.

Chemical properties

1. Very reactive.

2. They have a positive valence of +2.

3. React with water at room temperature (except Be) to release hydrogen.

4. They have a high affinity for oxygen (reducing agents).

5. With hydrogen they form salt-like hydrides EH 2.

6. Oxides have the general formula EO. The tendency to form peroxides is less pronounced than for alkali metals.

Being in nature

3BeO ∙ Al 2 O 3 ∙ 6SiO 2 beryl

Mg

MgCO 3 magnesite

CaCO 3 ∙ MgCO 3 dolomite

KCl ∙ MgSO 4 ∙ 3H 2 O kainite

KCl ∙ MgCl 2 ∙ 6H 2 O carnallite

CaCO 3 calcite (limestone, marble, etc.)

Ca 3 (PO 4) 2 apatite, phosphorite

CaSO 4 ∙ 2H 2 O gypsum

CaSO 4 anhydrite

CaF 2 fluorspar (fluorite)

SrSO 4 celestine

SrCO 3 strontianite

BaSO 4 barite

BaCO 3 witherite

Receipt

Beryllium is obtained by reduction of fluoride:

BeF 2 + Mg═ t ═ Be + MgF 2

Barium is obtained by reduction of the oxide:

3BaO + 2Al═ t ═ 3Ba + Al 2 O 3

The remaining metals are obtained by electrolysis of chloride melts:

CaCl 2 = Ca + Cl 2 ╜

cathode: Ca 2+ + 2ē = Ca 0

anode: 2Cl - - 2ē = Cl 0 2

MgO + C = Mg + CO

Metals of the main subgroup of group II are strong reducing agents; compounds exhibit only the +2 oxidation state. The activity of metals and their reducing ability increases in the series: Be Mg Ca Sr Ba╝

1. Reaction with water.

Under normal conditions, the surface of Be and Mg is covered with an inert oxide film, so they are resistant to water. In contrast, Ca, Sr and Ba dissolve in water to form hydroxides, which are strong bases:

Mg + 2H 2 O═ t ═ Mg(OH) 2 + H 2

Ca + 2H 2 O = Ca(OH) 2 + H 2 ╜

2. Reaction with oxygen.

All metals form oxides RO, barium peroxide BaO 2:

2Mg + O2 = 2MgO

Ba + O 2 = BaO 2

3. Binary compounds are formed with other non-metals:

Be + Cl 2 = BeCl 2 (halides)

Ba + S = BaS (sulfides)

3Mg + N 2 = Mg 3 N 2 (nitrides)

Ca + H 2 = CaH 2 (hydrides)

Ca + 2C = CaC 2 (carbides)

3Ba + 2P = Ba 3 P 2 (phosphides)

Beryllium and magnesium react relatively slowly with nonmetals.

4. All metals dissolve in acids:

Ca + 2HCl = CaCl 2 + H 2 ╜

Mg + H 2 SO 4 (diluted) = MgSO 4 + H 2 ╜

Beryllium also dissolves in aqueous solutions of alkalis:

Be + 2NaOH + 2H 2 O = Na 2 + H 2 ╜

5. Qualitative reaction to cations of alkaline earth metals - coloring of the flame in the following colors:

Ca 2+ - dark orange

Sr 2+ - dark red

Ba 2+ - light green

The Ba 2+ cation is usually discovered by an exchange reaction with sulfuric acid or its salts:

Barium sulfate is a white precipitate, insoluble in mineral acids.

Alkaline earth metal oxides

Receipt

1) Oxidation of metals (except Ba, which forms peroxide)

2) Thermal decomposition of nitrates or carbonates

CaCO 3 ═ t ═ CaO + CO 2 ╜

2Mg(NO 3) 2 ═ t ═ 2MgO + 4NO 2 ╜ + O 2 ╜

Chemical properties

Typical basic oxides. Reacts with water (except BeO), acid oxides and acids

MgO + H 2 O = Mg(OH) 2

3CaO + P 2 O 5 = Ca 3 (PO 4) 2

BeO + 2HNO 3 = Be(NO 3) 2 + H 2 O

BeO is an amphoteric oxide, soluble in alkalis:

BeO + 2NaOH + H 2 O = Na 2

Alkaline earth metal hydroxides R(OH) 2

Receipt

Reactions of alkaline earth metals or their oxides with water: Ba + 2H 2 O = Ba(OH) 2 + H 2

CaO(quicklime) + H 2 O = Ca(OH) 2 (slaked lime)

Chemical properties

Hydroxides R(OH) 2 are white crystalline substances, less soluble in water than hydroxides of alkali metals (the solubility of hydroxides decreases with decreasing atomic number; Be(OH) 2 is insoluble in water, soluble in alkalis). The basicity of R(OH) 2 increases with increasing atomic number:

Be(OH) 2 - amphoteric hydroxide

Mg(OH) 2 - weak base

the remaining hydroxides are strong bases (alkalis).

1) Reactions with acid oxides:

Ca(OH) 2 + SO 2 = CaSO 3 ¯ + H 2 O

Ba(OH) 2 + CO 2 = BaCO 3 ¯ + H 2 O

2) Reactions with acids:

Mg(OH) 2 + 2CH 3 COOH = (CH 3 COO) 2 Mg + 2H 2 O

Ba(OH) 2 + 2HNO 3 = Ba(NO 3) 2 + 2H 2 O

3) Exchange reactions with salts:

Ba(OH) 2 + K 2 SO 4 = BaSO 4 ¯+ 2KOH

4) Reaction of beryllium hydroxide with alkalis:

Be(OH) 2 + 2NaOH = Na 2

Water hardness

Natural water containing Ca 2+ and Mg 2+ ions is called hard water. Hard water forms scale when boiled and food products cannot be cooked in it; Detergents do not produce foam.

Carbonate (temporary) hardness is caused by the presence of calcium and magnesium bicarbonates in water, non-carbonate (permanent) hardness is caused by chlorides and sulfates.

The total hardness of water is considered as the sum of carbonate and non-carbonate.

Water hardness is removed by precipitation of Ca 2+ and Mg 2+ ions from solution:

1) boiling:

Сa(HCO 3) 2 ═ t ═ CaCO 3 ¯ + CO 2 + H 2 O

Mg(HCO 3) 2 ═ t═ MgCO 3 ¯ + CO 2 + H 2 O

2) adding lime milk:

Ca(HCO 3) 2 + Ca(OH) 2 = 2CaCO 3 ¯ + 2H 2 O

3) adding soda:

Ca(HCO 3) 2 + Na 2 CO 3 = CaCO 3 ¯+ 2NaHCO 3

CaSO 4 + Na 2 CO 3 = CaCO 3 ¯ + Na 2 SO 4

MgCl 2 + Na 2 CO 3 = MgCO 3 ¯ + 2NaCl

To remove temporary hardness, all four methods are used, and for permanent hardness, only the last two are used.

Thermal decomposition of nitrates.

E(NO3)2 =t= EO + 2NO2 + 1/2O2

Features of the chemistry of beryllium.

Be(OH)2 + 2NaOH (g) = Na2

Al(OH)3 + 3NaOH (g) = Na3

Be + 2NaOH + 2H2O = Na2 + H2

Al + 3NaOH + 3H2O = Na3 + 3/2H2

Be, Al + HNO3 (Conc) = passivation

The lesson will cover the topic “Metals and their properties. Alkali metals. Alkaline earth metals. Aluminum". You will learn the general properties and patterns of alkali and alkaline earth elements, study separately the chemical properties of alkali and alkaline earth metals and their compounds. Using chemical equations, we will consider such a concept as water hardness. Get acquainted with aluminum, its properties and alloys. You will learn about oxygen regenerating mixtures, ozonides, barium peroxide and oxygen production.

Topic: Basic metals and non-metals

Lesson: Metals and their properties. Alkali metals. Alkaline earth metals. Aluminum

The main subgroup of group I of the Periodic System D.I. Mendeleev's elements are lithium Li, sodium Na, potassium K, rubidium Rb, cesium Cs and francium Fr. Elements of this subgroup belong to. Their common name is alkali metals.

Alkaline earth metals are in the main subgroup of group II of the D.I. Periodic Table. Mendeleev. These are magnesium Mg, calcium Ca, strontium Sr, barium Ba and radium Ra.

Alkali and alkaline earth metals, as typical metals, exhibit pronounced reducing properties. For elements of the main subgroups, the metallic properties increase with increasing radius. Alkali metals exhibit particularly strong reducing properties. So strong that it is almost impossible to carry out their reactions with dilute aqueous solutions, since the reaction of their interaction with water will occur first. The situation is similar for alkaline earth metals. They also interact with water, but much less intensely than alkali metals.

Electronic configurations valence layer of alkali metals - ns 1 , where n is the number of the electronic layer. They are classified as s-elements. For alkaline earth metals - ns 2 (s-elements). Aluminum has valence electrons …3 s 2 3р 1(p-element). These elements form compounds with an ionic bond type. When compounds are formed, their oxidation state corresponds to the group number.

Detection of metal ions in salts

Metal ions can be easily identified by changes in flame color. Rice. 1.

Lithium salts - carmine-red color of the flame. Sodium salts - yellow. Potassium salts - purple through cobalt glass. Rubidium is red, cesium is violet-blue.

Rice. 1

Salts of alkaline earth metals: calcium - brick-red, strontium - carmine-red and barium - yellowish-green. Aluminum salts do not change the color of the flame. Salts of alkali and alkaline earth metals are used to create fireworks. And you can easily determine by the color which metal salts were used.

Properties of metals

Alkali metals- These are silvery-white substances with a characteristic metallic luster. They quickly fade in air due to oxidation. These are soft metals; the softness of Na, K, Rb, Cs is similar to wax. They are easy to cut with a knife. They are light. Lithium is the lightest metal with a density of 0.5 g/cm 3 .

Chemical properties of alkali metals

1. Interaction with non-metals

Due to their high reducing properties, alkali metals react violently with halogens to form the corresponding halide. When heated, they react with sulfur, phosphorus and hydrogen to form sulfides, hydrides, and phosphides.

2Na + Cl 2 → 2NaCl

Lithium is the only metal that reacts with nitrogen already at room temperature.

6Li + N 2 = 2Li 3 N, the resulting lithium nitride undergoes irreversible hydrolysis.

Li 3 N + 3H 2 O → 3LiOH + NH 3

2. Interaction with oxygen

Only with lithium does lithium oxide form immediately.

4Li + O 2 = 2Li 2 O, and when oxygen reacts with sodium, sodium peroxide is formed.

2Na + O 2 = Na 2 O 2. When all other metals burn, superoxides are formed.

K + O 2 = KO 2

3. Interaction with water

By reaction with water, you can clearly see how the activity of these metals changes in the group from top to bottom. Lithium and sodium react calmly with water, potassium with a flash, and cesium with an explosion.

2Li + 2H 2 O → 2LiOH + H 2

4.

8K + 10HNO 3 (conc) → 8KNO 3 + N 2 O +5 H 2 O

8Na + 5H 2 SO 4 (conc) → 4Na 2 SO 4 + H 2 S + 4H 2 O

Preparation of alkali metals

Due to the high activity of metals, they can be obtained by electrolysis of salts, most often chlorides.

Alkali metal compounds are widely used in various industries. See Table. 1.

COMMON ALKALI METAL COMPOUNDS
	Caustic soda (caustic soda)
	Table salt
	Chilean saltpeter
Na 2 SO 4 ∙10H 2 O	Glauber's salt
Na 2 CO 3 ∙10H 2 O	Soda crystal
	Caustic potassium
	Potassium chloride (sylvine)
	Indian saltpeter

Their name is due to the fact that the hydroxides of these metals are alkalis, and the oxides were previously called “earths”. For example, barium oxide BaO is barium earth. Beryllium and magnesium are most often not classified as alkaline earth metals. We will not consider radium either, since it is radioactive.

Chemical properties of alkaline earth metals.

1. Interaction withnon-metals

Сa + Cl 2 → 2СaCl 2

Ca + H 2 CaH 2

3Ca + 2P Ca 3 P 2-

2. Interaction with oxygen

2Ca + O 2 → 2CaO

3. Interaction with water

Sr + 2H 2 O → Sr(OH) 2 + H 2, but the interaction is calmer than with alkali metals.

4. Interaction with acids - strong oxidizing agents

4Sr + 5HNO 3 (conc) → 4Sr(NO 3) 2 + N 2 O +4H 2 O

4Ca + 10H 2 SO 4 (conc) → 4CaSO 4 + H 2 S + 5H 2 O

Preparation of alkaline earth metals

Metallic calcium and strontium are obtained by electrolysis of molten salts, most often chlorides.

CaCl 2 Ca + Cl 2

High purity barium can be obtained aluminothermally from barium oxide

3BaO +2Al 3Ba + Al 2 O 3

COMMON ALKALINE EARTH METALS COMPOUNDS

The most famous compounds of alkaline earth metals are: CaO - quicklime. Ca(OH) 2 - slaked lime, or lime water. When carbon dioxide is passed through lime water, turbidity occurs, as insoluble calcium carbonate CaCO 3 is formed. But we must remember that with further passage of carbon dioxide, soluble bicarbonate is formed and the precipitate disappears.

Rice. 2

СaO + H 2 O → Ca(OH) 2

Ca(OH) 2 + CO 2 → CaCO 3 ↓+ H 2 O

CaCO 3 ↓+ H 2 O + CO 2 → Ca(HCO 3) 2

Plaster - this is CaSO 4 ∙2H 2 O, alabaster is CaSO 4 ∙0.5H 2 O. Gypsum and alabaster are used in construction, in medicine and for the manufacture of decorative items. Rice. 2.

Calcium carbonate CaCO 3 forms many different minerals. Rice. 3.

Rice. 3

Calcium phosphate Ca 3 (PO 4) 2 - phosphorite, phosphorus flour is used as a mineral fertilizer.

Pure anhydrous calcium chloride CaCl 2 is a hygroscopic substance, therefore it is widely used in laboratories as a desiccant.

Calcium carbide- CaC2. You can get it like this:

CaO + 2C →CaC 2 +CO. One of its uses is to produce acetylene.

CaC 2 + 2H 2 O →Ca(OH) 2 + C 2 H 2

Barium sulfate BaSO 4 - barite. Rice. 4. Used as a white standard in some studies.

Rice. 4

Water hardness

Natural water contains calcium and magnesium salts. If they are contained in noticeable concentrations, then soap does not lather in such water due to the formation of insoluble stearates. When it boils, scale forms.

Temporary hardness due to the presence of calcium and magnesium hydrocarbonates Ca(HCO 3) 2 and Mg(HCO 3) 2. This type of water hardness can be removed by boiling.

Ca(HCO 3) 2 CaCO 3 ↓ + CO 2 + H 2 O

Constant water hardness is caused by the presence of cations Ca 2+, Mg 2+ and anions H 2 PO 4 -, Cl -, NO 3 -, etc. Constant water hardness is eliminated only due to ion exchange reactions, as a result of which magnesium and calcium ions will be transferred to sediment.

Homework

1. No. 3, 4, 5-a (p. 173) Gabrielyan O.S. Chemistry. 11th grade. Basic level. 2nd ed., erased. - M.: Bustard, 2007. - 220 p.

2. What reaction of the medium does an aqueous solution of potassium sulfide have? Confirm your answer with the equation of the hydrolysis reaction.

3. Determine the mass fraction of sodium in sea water, which contains 1.5% sodium chloride.

The chemical properties of alkali and alkaline earth metals are similar. The outer energy level of alkali metals has one electron, and that of alkaline earth metals has two. During reactions, metals easily part with valence electrons, exhibiting the properties of a strong reducing agent.

Alkaline

Group I of the periodic table includes alkali metals:

• lithium;
• sodium;
• potassium;
• rubidium;
• cesium;
• French

Rice. 1. Alkali metals.

They are distinguished by their softness (can be cut with a knife), low melting and boiling points. These are the most active metals.

The chemical properties of alkali metals are presented in the table.

Reaction	Peculiarities	Equation
With oxygen	They oxidize quickly in air. Lithium forms oxide at temperatures above 200°C. Sodium forms a mixture of 80% peroxide (R 2 O 2) and 20% oxide. The remaining metals form superoxides (RO 2)	4Li + O 2 → 2Li 2 O; 2Na + O 2 → Na 2 O 2 ; Rb + O 2 → RbO 2
	Only lithium reacts at room temperature	6Li + N 2 → 2Li 3 N
With halogens	The reaction is vigorous	2Na + Cl 2 → 2NaCl
With non-metals	When heated. They form sulfides, hydrides, phosphides, and silicides. Only lithium and sodium react with carbon, forming carbides	2K + S → K 2 S; 2Na + H 2 → 2NaH; 2Cs + 5P → Cs 2 P 5 ; Rb + Si → RbSi; 2Li + 2C → Li 2 C 2
	Only lithium reacts calmly. Sodium burns with a yellow flame. Potassium reacts with a flash. Cesium and rubidium explode	2Na + 2H 2 O → 2NaOH + H 2 -
With acids	They react explosively with hydrochloric, phosphoric, and dilute sulfuric acids. When reacting with concentrated sulfuric acid, hydrogen sulfide is released, with concentrated nitric acid it forms nitric oxide (I), with dilute nitric acid - nitrogen	2Na + 2HCl → 2NaCl + H 2 ; 8Na + 5H 2 SO 4 (conc) → 4Na 2 SO 4 + H 2 S + 4H 2 O; 8K + 10HNO 3 (conc) → 8KNO 3 + N 2 O + 5H 2 O; 10Na + 12HNO 3 (diluted) → N 2 + 10NaNO 3 + 6H 2 O
With ammonia	Form amines	2Li + 2NH 3 → 2LiNH 2 + H 2

Can react with organic acids and alcohols.

Alkaline earth

In group II of the periodic table there are alkaline earth metals:

• beryllium;
• magnesium;
• calcium;
• strontium;
• barium;
• radium.

Rice. 2. Alkaline earth metals.

Unlike alkali metals, they are harder. Only strontium can be cut with a knife. The densest metal is radium (5.5 g/cm3).

Beryllium reacts with oxygen only when heated to 900°C. Does not react with hydrogen and water under any conditions. Magnesium oxidizes at a temperature of 650°C and reacts with hydrogen under high pressure.

The table shows the main chemical properties of alkaline earth metals.

Reaction	Peculiarities	Equation
With oxygen	Form oxide films. Self-ignites when heated to 500°C	2Mg + O 2 → 2MgO
With hydrogen	At high temperatures they form hydrides	Sr + H 2 → SrH 2
With halogens and non-metals	React when heated	Be + Cl 2 → BeCl 2 ; Mg + S → MgS; 3Ca + 2P → Ca 3 P 2 ; 3Ca + N 2 → Ca 3 N 2; Ba + 2C → BaC 2
	At room temperature	Mg + 2H 2 O → Mg(OH) 2 + H 2
With acids	All metals react to form salts	4Ca + 10HNO 3 (conc.) → 4Ca(NO 3) 2 + N 2 O + 5H 2 O
With alkalis	Only beryllium reacts	Be + 2NaOH + 2H 2 O → Na 2 + H 2
Substitution	Replace less active metals in oxides. The exception is beryllium.	2Mg + ZrO 2 → Zr + 2MgO

Alkali and alkaline earth metal ions in salts are easily detected by changes in flame color. Sodium salts burn with a yellow flame, potassium - violet, rubidium - red, calcium - brick-red, barium - yellow-green. Salts of these metals are used to create fireworks.

Rice. 3. Qualitative reaction.

What have we learned?

Alkali and alkaline earth metals are active elements of the periodic table that react with simple and complex substances. Alkali metals are softer, react violently with water and halogens, easily oxidize in air, forming oxides, peroxides, superoxides, and interact with acids and ammonia. When heated, they react with non-metals. Alkaline earth metals react with nonmetals, acids, and water. Beryllium does not react with hydrogen and water, but reacts with alkalis and oxygen at high temperatures.

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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 ns 1, alkaline earth metals are ns 2.

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.

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.

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.

Interaction	Alkali metals	Alkaline earth metals
With oxygen	Self-ignites in air. They form superoxides (RO 2), except lithium and sodium. Lithium forms an oxide when heated above 200°C. Sodium forms a mixture of peroxide and oxide. 4Li + O 2 → 2Li 2 O; 2Na + O 2 → Na 2 O 2 ; Rb + O 2 → RbO 2	In air, protective oxide films quickly form. When heated to 500°C, they ignite spontaneously. 2Mg + O 2 → 2MgO; 2Ca + O 2 → 2CaO
With non-metals	React when heated with sulfur, hydrogen, phosphorus: 2K + S → K 2 S; 2Na + H 2 → 2NaH; 2Cs + 5P → Cs 2 P 5 . Only lithium reacts with nitrogen, and lithium and sodium react with carbon: 6Li + N 2 → 2Li 3 N; 2Na + 2C → Li 2 C 2	React when heated: Ca + Br 2 → CaBr 2; Be + Cl 2 → BeCl 2 ; Mg + S → MgS; 3Ca + 2P → Ca 3 P 2 ; Sr + H 2 → SrH 2
With halogens	React violently to form halides: 2Na + Cl 2 → 2NaCl
	Alkalies are formed. The lower the metal is located in the group, the more active the reaction occurs. Lithium reacts calmly, sodium burns with a yellow flame, potassium with a flash, cesium and rubidium explode. 2Na + 2H 2 O → 2NaOH + H 2 -; 2Li + 2H 2 O → 2LiOH + H 2	Less active than alkali metals, they react at room temperature: Mg + 2H 2 O → Mg(OH) 2 + H 2 ; Ca + 2H 2 O → Ca(OH) 2 + H 2
With acids	They react explosively with weak and dilute acids. They form salts with organic acids. 8K + 10HNO 3 (conc) → 8KNO 3 + N 2 O + 5H 2 O; 8Na + 5H 2 SO 4 (conc) → 4Na 2 SO 4 + H 2 S + 4H 2 O; 10Na + 12HNO 3 (diluted) → N 2 + 10NaNO 3 + 6H 2 O; 2Na + 2CH 3 COOH → 2CH 3 COONa + H 2	Salts are formed: 4Sr + 5HNO 3 (conc) → 4Sr(NO 3) 2 + N 2 O + 4H 2 O; 4Ca + 10H 2 SO 4 (conc) → 4CaSO 4 + H 2 S + 5H 2 O
With alkalis		Of all the metals, only beryllium reacts: Be + 2NaOH + 2H 2 O → Na 2 + H 2
With oxides		All metals react except beryllium. Replaces less active metals: 2Mg + ZrO 2 → Zr + 2MgO

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.

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.

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