General characteristics of group IIA of the Periodic Table of elements.

The following elements are located in this group: Be, Mg, Ca, Sr, Ba, Ra. They have a common electronic configuration: (n-1)p6ns2, except Be 1s22s2. Due to the latter, the properties of Be are slightly different from the properties of the subgroup as a whole. The properties of magnesium also differ from the properties of the subgroup, but to a lesser extent. In the series Ca – Sr – Ba – Ra, the properties change sequentially. The relative electronegativity in the Be – Ra series decreases because As the size of the atom increases, valence electrons are given up more readily. The properties of subgroup IIA elements are determined by the ease of losing two ns electrons. In this case, E2+ ions are formed. When studying X-ray diffraction, it turned out that in some compounds elements of the IIA subgroup exhibit monovalency. An example of such compounds are EG, which are obtained by adding E to the EG2 melt. All elements of this series are not found in nature in a free state due to their high activity.

Alkaline earth metals.

Calcium, strontium, barium and radium are called alkaline earth metals. They are named so because their oxides give the water an alkaline environment.

History of alkaline earth metals
Limestone, marble and gypsum were used by the Egyptians in construction already in ancient times (5000 years ago). Until the end of the 18th century, chemists considered lime to be a simple substance. In 1746, I. Pott obtained and described fairly pure calcium oxide. In 1789, Lavoisier suggested that lime, magnesia, and barite were complex substances. Long before the discovery of strontium and barium, their “undeciphered” compounds were used in pyrotechnics to produce red and green lights, respectively. Until the mid-40s of the last century, strontium was primarily a metal of “amusing fires”. In 1787, a new mineral was found in a lead mine near the Scottish village of Strontian, which was called strontianite SrCO3. A. Crawford suggested the existence of a still unknown “earth”. In 1792, T. Hop proved that the composition of the found mineral included a new element - strontium. At that time, with the help of Sr(OH)2, insoluble strontium disaccharate (C12H22O4.2SrO) was isolated to obtain sugar from molasses. Sr production increased. However, it was soon noticed that the similar calcium saccharate was also insoluble, and calcium oxide was undoubtedly cheaper. Interest in strontium immediately disappeared and increased again only in the 40s of the last century. Heavy spar was the first known compound of barium. It was discovered at the beginning of the 17th century. Italian alchemist Casciarolo. He also established that this mineral, after strong heating with coal, glows in the dark with red light and gave it the name “lapis solaris” (sun stone). In 1808, Davy, subjecting a mixture of wet slaked lime with mercuric oxide to electrolysis with a mercury cathode, prepared calcium amalgam, and by distilling mercury from it, he obtained a metal called “calcium” (from the Latin Calx, genus calcis - lime). Ba and Sr were obtained by the same Davy method. An industrial method for producing calcium was developed by Suter and Redlich in 1896 at the Rathenau plant (Germany). In 1904, the first calcium production plant began operating.
Radium was predicted by Mendeleev in 1871 and discovered in 1898 by the spouses Marie and Pierre Curie. They discovered that uranium ores are more radioactive than uranium itself. The cause was radium compounds. They treated the remaining uranium ore with alkali, and what did not dissolve with hydrochloric acid. The residue after the second procedure had more radioactivity than the ore. Radium was discovered in this fraction. The Curies reported their discovery in a report in 1898.

Alkaline earth metal abundance
The calcium content in the lithosphere is 2.96% of the total mass of the earth's crust, strontium - 0.034%, barium - 0.065%, radium - 1.10-10%. In nature, calcium consists 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(0.1%), 132(0.1%), 134(2.42%), 135(6.59%), 136(7.81), 137(11, 32%), 138 (71.66). Radium is radioactive. The most stable natural isotope is 226Ra. The main minerals of alkaline earth elements are carbon and sulphate salts: CaCO3 - calcite, CaSO4 - andidrite, SrCO3 - strontianite, SrSO4 - celestine, BaCO3 - witherite. BaSO4 is a heavy spar. Fluorite CaF2 is also a useful mineral.
Ca plays an important role in life processes. The human body contains 0.7-1.4 wt.% calcium, 99% of which is in bone and dental tissue. Plants also contain large amounts of calcium. Calcium compounds are found in natural waters and soil. Barium, strontium and radium are found in the human body in negligible quantities.

Preparation of alkaline earth metals
First, oxides or chlorides of E are obtained. EO is obtained by calcination of ESO3, and ES12 by the action of hydrochloric acid on ESO3. All alkaline earth metals can be obtained by aluminothermic reduction of their oxides at a temperature of 1200 °C according to the approximate scheme: 3EO + 2Al = Al2O3 + 3E. The process is carried out in a vacuum to avoid oxidation of calcium. Calcium (like all other elements) can be obtained by electrolysis of a CaCl2 melt followed by distillation in a vacuum or thermal dissociation of CaCl2. Ba and Sr can be obtained by pyrolysis of E2N3, E(NH3)6, EN2. Radium is mined as a by-product from uranium ores.

Physical properties of alkaline earth metals
Ca and its analogues are silvery-white metals. Calcium is the hardest of them all. Strontium and especially barium are much softer than calcium. All alkaline earth metals are ductile and amenable to forging, cutting and rolling. Calcium under normal conditions crystallizes in the fcc structure with a period a = 0.556 nm (CN = 12), and at temperatures above 464 ° C in the bcc structure. Ca forms alloys with Li, Mg, Pb, Cu, Cd, Al, Ag, Hg. Strontium has an fcc structure; at a temperature of 488 °C, strontium undergoes a polymorphic transformation and crystallizes in a hexagonal structure. It is paramagnetic. Barium crystallizes in a bcc structure. Ca and Sr are capable of forming a continuous series of solid solutions between themselves, and areas of separation appear in the Ca-Ba and Sr-Ba systems. In the liquid state, strontium mixes with Be, Hg, Ga, In, Sb, Bi, Tl, Al, Mg, Zn, Sn, Pb. With the last four, Sr forms intermetallic compounds. The electrical conductivity of alkaline earth metals decreases with increasing pressure, contrary to the reverse process for other typical metals. Below are some constants for alkaline earth metals:

Ca Sr Ba Ra
Atomic radius, nm 0.197 0.215 0.221 0.235
Radius of E2+ ion, nm 0.104 0.127 0.138 0.144
Energy cr. lattices, µJkmol 194.1 164.3 175.8 130
, gcm3 1.54 2.63 3.5 5.5-6
Melt.,оС 852 770 710 800
Boil point, oC 1484 1380 1640 1500
Electrical conductivity (Hg=1) 22 4 2
Heat of fusion kcalg-atom 2.1 2.2 1.8
Heat of vaporization kcalg-atom 36 33 36
Specific heat capacity, J(kg.K) 624 737 191.93 136
Liquefiability Pa-1.10-11 5.92 8.36

Chemical properties of alkaline earth metals and their compounds
The fresh surface of E quickly darkens due to the formation of an oxide film. This film is relatively dense - over time, all the metal slowly oxidizes. The film consists of EO, as well as EO2 and E3N2. The normal electrode potentials of the reactions E-2e = E2+ are equal to  = -2.84 V (Ca),  = -2.89 (Sr). These are very active elements: they dissolve in water and acids, displace most metals from their oxides, halides, and sulfides. Primary (200-300°C) calcium interacts with water vapor according to the following scheme: 2Ca + H2O = CaO + CaH2. Secondary reactions have the form: CaH2 + 2H2O = Ca(OH)2 + 2H2 and CaO + H2O = Ca(OH)2. ESO4 are almost insoluble in strong sulfuric acid due to the formation of a film of poorly soluble ESO4. E react violently with dilute mineral acids, releasing hydrogen. When heated above 800°C, calcium reacts with methane according to the following scheme: 3Ca + CH4 = CaH2 + CaC2. When heated, they react with hydrogen, sulfur and ammonia gas. In terms of chemical properties, radium is closest to Ba, but it is more active. At room temperature, it noticeably combines with oxygen and nitrogen in the air. In general, its chemical properties are slightly more pronounced than those of its analogues. All radium compounds slowly decompose under the influence of their own radiation, acquiring a yellowish or brown color. Radium compounds have the property of autoluminescence. As a result of radioactive decay, 1 g of Ra releases 553.7 J of heat every hour. Therefore, the temperature of radium and its compounds is always 1.5 degrees higher than the ambient temperature. It is also known that 1 g of radium per day releases 1 mm3 of radon (226Ra = 222Rn + 4He), on which its use as a source of radon for radon baths is based.
E hydrides are white, crystalline salt-like substances. They are obtained directly from the elements by heating. The starting temperatures of the reaction E + H2 = EN2 are 250 oC (Ca), 200 oC (Sr), 150 oC (Ba). Thermal dissociation of EN2 begins at 600°C. In a hydrogen atmosphere, CaH2 does not decompose at the melting point (816°C). In the absence of moisture, alkaline earth metal hydrides are stable in air at ordinary temperatures. They do not react with halogens. However, when heated, the chemical activity of EN2 increases. They are capable of reducing oxides to metals (W, Nb, Ti, Ce, Zr, Ta), for example 2CaH2 + TiO2 = 2CaO + 2H2 + Ti. The reaction of CaH2 with Al2O3 occurs at 750°C: 3CaH2 + Al2O3 = 3CaO + 3H2 + 2Al, and then: CaH2 + 2Al = CaAl2 + H2. CaH2 reacts with nitrogen at 600°C according to the following scheme: 3CaH2 + N2 = Ca3N2 + 3H2. When EN2 is ignited, they burn slowly: EN2 + O2 = H2O + CaO. Explosive when mixed with solid oxidizing agents. When water acts on EN2, hydroxide and hydrogen are released. This reaction is highly exothermic: EN2 moistened with water in air spontaneously ignites. EN2 reacts with acids, for example, according to the scheme: 2HCl + CaH2 = CaCl2 + 2H2. EN2 is used to obtain pure hydrogen, as well as to determine traces of water in organic solvents. E nitrides are colorless, refractory substances. They are obtained directly from elements at elevated temperatures. They decompose with water according to the following scheme: E3N2 + 6H2O = 3E(OH)2 + 2NH3. E3N2 reacts with CO when heated according to the following scheme: E3N2 + 3CO = 3EO + N2 + 3C. The processes that occur when heating E3N2 with coal look like this:
E3N2 + 5C = ECN2 + 2ES2; (E = Ca, Sr); Ba3N2 + 6C = Ba(CN)2 + 2BaC2;
Strontium nitride reacts with HCl to produce Sr and ammonium chlorides. Phosphides E3P2 are formed directly from elements or by calcination of trisubstituted phosphates with coal:
Ca3(PO4)2 + 4C = Ca3P2 + 4CO
They are hydrolyzed by water according to the scheme: E3P2 + 6H2O = 2PH3 + 3E(OH)2. With acids, phosphides of alkaline earth metals give the corresponding salt and phosphine. This is the basis for their use for obtaining phosphine in the laboratory.
Complex ammonia compounds E(NH3)6 are solid substances with a metallic luster and high electrical conductivity. They are obtained by the action of liquid ammonia on E. They spontaneously ignite in air. Without air access, they decompose into the corresponding amides: E(NH3)6 = E(NH2)2 + 4NH3 + H2. When heated, they vigorously decompose according to the same pattern.
Carbides of alkaline earth metals, which are obtained by calcination of ethylene with coal, decompose with water, releasing acetylene: ES2 + 2H2O = E(OH)2 + C2H2. The reaction with BaC2 is so violent that it ignites in contact with water. The heats of formation of ES2 from elements for Ca and Ba are 14 and 12 kcalmol. When heated with nitrogen, ES2 gives CaCN2, Ba(CN)2, SrCN2. Silicides are known (ESi and ESi2). They can be obtained by heating directly from the elements. They are hydrolyzed by water and react with acids, giving H2Si2O5, SiH4, the corresponding compound E and hydrogen. EV6 borides are known that are obtained from elements by heating.
Calcium oxides and its analogues are white, refractory (TbpCaO = 2850°C) substances that energetically absorb water. This is the basis for the use of BaO to obtain absolute alcohol. They react violently with water, releasing a lot of heat (except for SrO, the dissolution of which is endothermic). EOs dissolve in acids and ammonium chloride: EO + 2NH4Cl = SrCl2 + 2NH3 + H2O. EO is obtained by calcination of carbonates, nitrates, peroxides or hydroxides of the corresponding metals. The effective charges of barium and oxygen in BaO are 0.86. SrO at 700 °C reacts with potassium cyanide:
KCN + SrO = Sr + KCNO.
Strontium oxide dissolves in methanol to form Sr(OSH3)2. During the magnesium-thermal reduction of BaO, the intermediate oxide Ba2O can be obtained, which is unstable and disproportionate.
Alkaline earth metal hydroxides are white, water-soluble substances. They are strong bases. In the Ca-Sr-Ba series, the basic nature and solubility of hydroxides increase. pPR(Ca(OH)2) = 5.26, pPR(Sr(OH)2) = 3.5, pPR(Ba(OH)2) = 2.3. Ba(OH)2.8H2O, Sr(OH)2.8H2O, Ca(OH)2.H2O are usually isolated from hydroxide solutions. EOs add water to form hydroxides. This is the basis for the use of CaO in construction. A close mixture of Ca(OH)2 and NaOH in a 2:1 weight ratio is called soda lime, and is widely used as a CO2 absorber. Ca(OH)2, when standing in air, absorbs CO2 according to the following scheme: Ca(OH)2 + CO2 = CaCO3 + H2O. At about 400°C, Ca(OH)2 reacts with carbon monoxide: CO + Ca(OH)2 = CaCO3 + H2. Barite water reacts with CS2 at 100 °C: CS2 + 2Ba(OH)2 = BaCO3 + Ba(HS)2 + H2O. Aluminum reacts with barite water: 2Al + Ba(OH)2 + 10H2O = Ba2 + 3H2. E(OH)2 is used to discover carbonic anhydride.
E form white peroxides. They are significantly less stable, unlike oxides, and are strong oxidizing agents. Of practical importance is the most stable BaO2, which is a white, paramagnetic powder with a density of 4.96 g1cm3, etc. 450°. BaO2 is stable at ordinary temperatures (can be stored for years), is poorly soluble in water, alcohol and ether, and dissolves in dilute acids with the release of salt and hydrogen peroxide. The thermal decomposition of barium peroxide is accelerated by the oxides Cr2O3, Fe2O3 and CuO. Barium peroxide reacts when heated with hydrogen, sulfur, carbon, ammonia, ammonium salts, potassium ferricyanide, etc. Barium peroxide reacts with concentrated hydrochloric acid, releasing chlorine: BaO2 + 4HCl = BaCl2 + Cl2 + 2H2O. It oxidizes water to hydrogen peroxide: H2O + BaO2 = Ba(OH)2 + H2O2. This reaction is reversible and in the presence of even carbonic acid the equilibrium is shifted to the right. BaO2 is used as a starting product for the production of H2O2, and also as an oxidizing agent in pyrotechnic compositions. However, BaO2 can also act as a reducing agent: HgCl2 + BaO2 = Hg + BaCl2 + O2. BaO2 is obtained by heating BaO in a stream of air to 500°C according to the scheme: 2BaO + O2 = 2BaO2. As the temperature increases, the reverse process occurs. Therefore, when Ba burns, only oxide is released. SrO2 and CaO2 are less stable. The general method for obtaining EO2 is the interaction of E(OH)2 with H2O2, which releases EO2.8H2O. Thermal decomposition of EO2 begins at 380 °C (Ca), 480 °C (Sr), 790 °C (Ba). When heating EO2 with concentrated hydrogen peroxide, yellow unstable substances can be obtained - EO4 superperoxides.
E salts are usually colorless. Chlorides, bromides, iodides and nitrates are highly soluble in water. Fluorides, sulfates, carbonates and phosphates are poorly soluble. The Ba2+ ion is toxic. E halides are divided into two groups: fluorides and all others. Fluorides are almost insoluble in water and acids, and do not form crystalline hydrates. On the contrary, chlorides, bromides, and iodides are highly soluble in water and are released from solutions in the form of crystalline hydrates. Some properties of EG2 are presented below:

CaF2 CaCl2 CaBr2 CaI2 SrF2 SrCl2 SrBr2 SrI2 BaF2 BaCl2 BaBr2 BaI2
Warm. arr., kcalmol. 290 191 164 128 189 198 171 134 286 205 181 145
Ecr. lattices, kcalmol. 617 525 508 487 580 504 489 467 547 468 463 440
Heat, °C 1423 782 760 575 1473 872 643 515 1353 962 853 740
Boil point, °C 2500 2000 1800 718 2460 2030 2260 1830
D(EG) in pairs, nm. 2.1 2.51 2.67 2.88 2.20 2.67 2.82 3.03 2.32 2.82 2.99 3.20

When obtained by exchange decomposition in solution, fluorides are released in the form of voluminous mucous deposits, which quite easily form colloidal solutions. EG2 can be obtained by acting with the corresponding halogens on the corresponding E. EG2 melts are capable of dissolving up to 30% E. When studying the electrical conductivity of melts of chlorides of elements of the second group of the main subgroup, it was found that their molecular ionic composition is very different. The degrees of dissociation according to the scheme ESl2 = E2+ + 2Cl– are equal to: BeCl2 - 0.009%, MgCl2 - 14.6%, CaCl2 - 43.3%, SrCl2 - 60.6%, BaCl2 - 80.2%. E halides (except fluorides) contain water of crystallization: CaCl2.6H2O, SrCl2.6H2O and BaCl2.2H2O. X-ray diffraction analysis established the structure of E[(OH2)6]G2 for crystalline hydrates of Ca and Sr. By slowly heating EG2 crystalline hydrates, anhydrous salts can be obtained. CaCl2 easily forms supersaturated solutions. Natural CaF2 (fluorite) is used in the ceramics industry, and it is also used to produce HF and is a fluorine mineral. Anhydrous CaCl2 is used as a desiccant due to its hygroscopicity. Calcium chloride crystalline hydrate is used to prepare refrigeration mixtures. BaCl2 – used in cx and to open SO42- (Ba2+ + SO42- = BaSO4). By fusing EG2 and EN2, hydrohalides can be obtained: EG2 + EN2 = 2ENH. These substances melt without decomposition but are hydrolyzed by water: 2ENH + 2H2O = EG2 + 2H2 + E(OH)2. The solubility of chlorates, bromates and iodates in water decreases in the order of Ca – Sr – Ba and Cl – Br – I. Ba(ClO3)2 – used in pyrotechnics. E perchlorates are highly soluble not only in water but also in organic solvents. The most important of E(ClO4)2 is Ba(ClO4)2.3H2O. Anhydrous barium perchlorate is a good desiccant. Its thermal decomposition begins only at 400 °C. Calcium hypochlorite Ca(ClO)2.nH2O (n=2,3,4) is obtained by the action of chlorine on lime milk. It is an oxidizing agent and is highly soluble in water. Bleached lime can be produced by treating solid slaked lime with chlorine. It decomposes with water and smells of chlorine in the presence of moisture. Reacts with CO2 in air:
CO2 + 2CaOCl2 = CaСO3 + CaCl2 + Cl2O.
Bleach is used as an oxidizing agent, bleaching agent and as a disinfectant.
Azides E(N3)2 and thiocyanates E(CNS)2.3H2O are known for alkaline earth metals. Azides are much less explosive than lead azide. Rodanides easily lose water when heated. They are highly soluble in water and organic solvents. Ba(N3)2 and Ba(CNS)2 can be used to obtain azides and thiocyanates of other metals from sulfates by exchange reaction.
Calcium and strontium nitrates usually exist in the form of crystalline hydrates Ca(NO3)2.4H2O and Sr(NO3)2.4H2O. Barium nitrate is not characterized by the formation of crystalline hydrate. When Ca(NO3)2.4H2O and Sr(NO3)2.4H2O are heated, I easily lose water. In an inert atmosphere, E nitrates are thermally stable up to 455 oC (Ca), 480 oC (Sr), 495 oC (Ba). The melt of calcium nitrate crystalline hydrate has an acidic environment at 75 °C. A special feature of barium nitrate is the low rate of dissolution of its crystals in water. Only barium nitrate, for which the unstable K2 complex is known, exhibits a tendency to form complexes. Calcium nitrate is soluble in alcohols, methyl acetate, and acetone. Strontium and barium nitrates are almost insoluble there. The melting point of E nitrates is estimated at 600°C, but at this same temperature decomposition begins: E(NO3)2 = E(NO2)2 + O2. Further decomposition occurs at a higher temperature: E(NO2)2 = EO + NO2 + NO. E nitrates have long been used in pyrotechnics. Highly volatile E salts color the flame in the corresponding colors: Ca – orange-yellow, Sr – red-carmine, Ba – yellow-green. Let's understand the essence of this using the example of Sr: Sr2+ has two VAOs: 5s and 5p or 5s and 4d. Let's impart energy to this system - heat it up. Electrons from orbitals closer to the nucleus will move to these VAOs. But such a system is not stable and will release energy in the form of a light quantum. It is Sr2+ that emits quanta with a frequency corresponding to the red wavelengths. When preparing pyrotechnic compositions, it is convenient to use saltpeter, because It not only colors the flame, but is also an oxidizing agent, releasing oxygen when heated. Pyrotechnic compositions consist of a solid oxidizer, a solid reducing agent and some organic substances that decolorize the flame of the reducing agent and act as a binding agent. Calcium nitrate is used as fertilizer.
All phosphates and hydrogen phosphates of E are poorly soluble in water. They can be obtained by dissolving the appropriate amount of CaO or CaCO3 in orthophosphoric acid. They are also deposited during exchange reactions of the type: (3-x)Ca2+ + 2HxPO4-(3-x) = Ca(3-x)(HxPO4)2. Of practical importance (as a fertilizer) is monosubstituted calcium orthophosphate, which, along with Ca(SO4), is part of superphosphate. It is obtained according to the scheme:
Ca3(PO4)2 + 2H2SO4 = Ca(H2PO4)2 + 2CaSO4
Oxalates are also slightly soluble in water. Of practical importance is calcium oxalate, which is dehydrated at 200 °C and decomposes at 430 °C according to the following scheme: CaC2O4 = CaCO3 + CO. E acetates are isolated in the form of crystalline hydrates and are highly soluble in water.
E sulfates are white substances that are poorly soluble in water. The solubility of CaSO4.2H2O per 1000 g of water at ordinary temperature is 8.10-3 mol, SrSO4 - 5.10-4 mol, BaSO4 - 1.10-5 mol, RaSO4 - 6.10-6 mol. In the Ca–Ra series, the solubility of sulfates decreases rapidly. Ba2+ is a reagent for sulfate ion. Calcium sulfate contains water of crystallization. Above 66 oC, anhydrous calcium sulfate is released from the solution, below - gypsum CaSO4.2H2O. Heating of gypsum above 170 °C is accompanied by the release of hydrate water. When gypsum is mixed with water, this mass quickly hardens due to the formation of crystalline hydrate. This property of gypsum is used in construction. The Egyptians used this knowledge 2000 years ago. The solubility of ESO4 in strong sulfuric acid is much higher than in water (BaSO4 up to 10%), which indicates complex formation. The corresponding ESO4.H2SO4 complexes can be obtained in the free state. Double salts with alkali metal and ammonium sulfates are known only for Ca and Sr. (NH4)2 is soluble in water and is used in analytical chemistry to separate Ca from Sr, because (NH4)2 is slightly soluble. Gypsum is used for the combined production of sulfuric acid and cement, because when heated with a reducing agent (coal), gypsum decomposes: CaSO4 + C = CaO + SO2 + CO. At a higher temperature (900 oC), sulfur is reduced even more according to the following scheme: CaSO4 + 3C = CaS + CO2 + 2CO. A similar decomposition of Sr and Ba sulfates begins at higher temperatures. BaSO4 is non-toxic and is used in medicine and in the production of mineral paints.
E sulfides are white solids that crystallize like NaCl. The heats of their formation and the energies of the crystal lattices are equal (kcalmol): 110 and 722 (Ca), 108 and 687 (Sr), 106 and 656 (Ba). They can be obtained by synthesis from elements by heating or by calcination of sulfates with coal: ESO4 + 3C = ES + CO2 + 2CO. The least soluble is CaS (0.2 hl). ES enters into the following reactions when heated:
ES + H2O = EO + H2S; ES + G2 = S + EG2; ES + 2O2 = ESO4; ES + xS = ESx+1 (x=2.3).
Sulfides of alkaline earth metals in a neutral solution are completely hydrolyzed according to the scheme: 2ES + 2H2O = E(HS)2 + E(OH)2. Acid sulfides can also be obtained in a free state by evaporating a solution of sulfides. They react with sulfur:
E(HS)2 + xS = ESx+1 + H2S (x=2,3,4).
Among crystalline hydrates, BaS.6H2O and Ca(HS)2.6H2O, Ba(HS)2.4H2O are known. Ca(HS)2 is used for hair removal. ES are subject to the phenomenon of phosphorescence. The following polysulfides of E are known: ES2, ES3, ES4, ES5. They are obtained by boiling a suspension of ES in water with sulfur. In air, ES are oxidized: 2ES + 3O2 = 2ESO3. By passing air through a CaS suspension, you can obtain Ca thiosulfate according to the scheme: 2CaS + 2O2 + H2O = Ca(OH)2 + CaS2O3. It is highly soluble in water. In the series Ca – Sr – Ba, the solubility of thiosulfates decreases. E tellurides are slightly soluble in water and are also subject to hydrolysis, but to a lesser extent than sulfides.
The solubility of E chromates in the Ca–Ba series drops just as sharply as in the case of sulfates. These yellow substances are obtained by the interaction of soluble salts of E with chromates (or dichromates) of alkali metals: E2+ + CrO42- = ECrO4. Calcium chromate is released in the form of crystalline hydrate – CaCrO4.2H2O (pPR CaCrO4 = 3.15). Even before the melting point, it loses water. SrCrO4 and BaCrO4 do not form crystalline hydrates. pPR SrCrO4 = 4.44, pPR BaCrO4 = 9.93.
E carbonates are white, poorly soluble substances in water. When heated, ECO3 transforms into EO, splitting off CO2. In the Ca–Ba series, the thermal stability of carbonates increases. The most practically important of these is calcium carbonate (limestone). It is directly used in construction, and also serves as a raw material for the production of lime and cement. The annual world production of lime from limestone amounts to tens of millions of tons. Thermal dissociation of CaCO3 is endothermic: CaCO3 = CaO + CO2 and requires a cost of 43 kcal per mole of limestone. Calcination of CaCO3 is carried out in shaft furnaces. A by-product of roasting is valuable carbon dioxide. CaO is an important building material. When mixed with water, crystallization occurs due to the formation of hydroxide and then carbonate according to the following schemes:
CaO + H2O = Ca(OH)2 and Ca(OH)2 + CO2 = CaCO3 + H2O.
A hugely important practical role is played by cement - a greenish-gray powder consisting of a mixture of various silicates and calcium aluminates. When mixed with water it hardens due to hydration. During its production, a mixture of CaCO3 and clay is fired until sintering begins (1400-1500 oC). Then the mixture is ground. The composition of cement can be expressed as the percentage of the components CaO, SiO2, Al2O3, Fe2O3, with CaO representing the base, and the rest being acid anhydrides. The composition of silicate (Portlad) cement consists mainly of Ca3SiO5, Ca2SiO4, Ca3(AlO3)2 and Ca(FeO2)2. Its setting proceeds according to the following schemes:
Ca3SiO5 + 3H2O = Ca2SiO4.2H2O + Ca(OH)2
Ca2SiO4 + 2H2O = Ca2SiO4.2H2O
Ca3(AlO3)2 + 6H2O = Ca3(AlO3)2.6H2O
Ca(FeO2)2 + nH2O = Ca(FeO2)2.nH2O.
Natural chalk is added to various putties. Fine-crystalline CaCO3, precipitated from a solution, is part of tooth powders. BaO is obtained from BaCO3 by calcination with coal according to the following scheme: BaCO3 + C = BaO + 2CO. If the process is carried out at a higher temperature in a stream of nitrogen, barium cyanide is formed: BaCO3 + 4C + N2 = 3CO + Ba(CN)2. Ba(CN)2 is highly soluble in water. Ba(CN)2 can be used to produce cyanides of other metals by exchange decomposition with sulfates. Hydrogen carbonates are soluble in water and can be obtained only in solution, for example, by passing carbon dioxide into a suspension of CaCO3 in water: CO2 + CaCO3 + H2O = Ca(HCO3)2. This reaction is reversible and shifts to the left when heated. The presence of calcium and magnesium bicarbonates in natural waters causes water hardness.

Water hardness and ways to eliminate it
Soluble calcium and magnesium salts cause overall water hardness. If they are present in water in small quantities, then the water is called soft. If the content of these salts is high (100–200 mg of calcium salts per liter in terms of ions), the water is considered hard. In such water, soap does not foam well, since calcium and magnesium salts form insoluble compounds with it. Hard water does not cook food well, and when boiled, it forms scale on the walls of household utensils and steam boilers. Scale has low thermal conductivity, causes an increase in fuel consumption or power consumption of an electrical appliance and accelerates the wear of the walls of the vessel for boiling water. When heated, acidic calcium and magnesium carbonates decompose and turn into insoluble basic carbonates: Ca(HCO3) = H2O + CO2 + CaCO3↓ The solubility of calcium sulfate CaSO4 also decreases when heated, so it is part of the scale. Hardness caused by the presence of acidic calcium and magnesium carbonates in water is called carbonate or temporary hardness, since it can be eliminated. In addition to carbonate hardness, there is also non-carbonate hardness, which depends on the content of ECl2 and ESO4 in water, where E is Ca, Mg. These salts are not removed by boiling, and therefore non-carbonate hardness is also called permanent hardness. Carbonate and non-carbonate hardness add up to total hardness. To completely eliminate it, the water is sometimes distilled. But it's expensive. To remove carbonate hardness, water can be boiled, but this is also expensive and scale forms. Hardness is removed by adding the appropriate amount of Ca(OH)2: Ca(OH)2 + Ca(HCO3)2 = CaCO3↓ + 2H2O. General hardness is eliminated either by adding Na2CO3, or using so-called cation exchangers. When using sodium carbonate, soluble calcium and magnesium salts are also converted into insoluble carbonates: Ca2+ + Na2CO3 = 2Na+ + CaCO3↓. Removing hardness using cation exchange resins is a more advanced process. Cation exchangers are high-molecular sodium-containing organic compounds, the composition of which can be expressed by the formula Na2R, where R is a complex acid residue. When water is filtered through a layer of cation exchange resin, the Na+ cations of the crystal lattice are exchanged for Ca2+ and Mg2+ cations from the solution according to the scheme: Ca2+ + Na2R = 2Na+ + CaR. Consequently, Ca ions pass from the solution into the cation exchanger, and Na+ ions pass from the cation exchanger into the solution. To restore the used cation exchanger, it is washed with a concentrated solution of table salt. In this case, the reverse process occurs: Ca2+ ions in the crystal lattice of the cation exchanger are replaced by Na+ ions from the solution. The regenerated cation exchanger is again used for water purification. Filters based on permutite work in a similar way:
Na2 + Ca2+ = 2Na+ + Ca

Being in nature
Due to its high chemical activity, calcium does not occur in free form in nature.
Most of the calcium is contained in silicates and aluminosilicates of various rocks (granites, gneisses, etc.), especially in feldspar - Ca anorthite.
In the form of sedimentary rocks, calcium compounds are represented by chalk and limestones, consisting mainly of the mineral calcite (CaCO3). The crystalline form of calcite - marble - is found in nature much less frequently.
Calcium minerals such as calcite CaCO3, anhydrite CaSO4, alabaster CaSO4 0.5H2O and gypsum CaSO4 2H2O, fluorite CaF2, apatite Ca5(PO4)3(F,Cl,OH), dolomite MgCO3 CaCO3 are quite widespread. The presence of calcium and magnesium salts in natural water determines its hardness.
Calcium, vigorously migrating in the earth's crust and accumulating in various geochemical systems, forms 385 minerals (the fourth largest number of minerals).
Calcium accounts for 3.38% of the mass of the earth's crust (5th most abundant after oxygen, silicon, aluminum and iron). The content of the element in sea water is 400 mg/l.
Isotopes
Calcium occurs in nature as a mixture of six isotopes: 40Ca, 42Ca, 43Ca, 44Ca, 46Ca and 48Ca, among which the most common - 40Ca - accounts for 96.97%.
Of the six natural isotopes of calcium, five are stable. The sixth isotope, 48Ca, the heaviest of the six and very rare (its isotopic abundance is only 0.187%), was recently discovered to undergo double beta decay with a half-life of 5.3 x 1019 years.
Receipt
Free metallic calcium is obtained by electrolysis of a melt consisting of CaCl2 (75-80%) and KCl or from CaCl2 and CaF2, as well as aluminothermic reduction of CaO at 1170-1200 °C:
4CaO + 2Al → CaAl2O4 + 3Ca.
Chemical properties
Calcium is a typical alkaline earth metal. The chemical activity of calcium is high, but lower than that of all other alkaline earth metals. It easily reacts with oxygen, carbon dioxide and moisture in the air, which is why the surface of calcium metal is usually dull gray, so in the laboratory calcium is usually stored, like other alkaline earth metals, in a tightly closed jar under a layer of kerosene or liquid paraffin .
In the series of standard potentials, calcium is located to the left of hydrogen. The standard electrode potential of the Ca2+/Ca0 pair is −2.84 V, so calcium reacts actively with water, but without ignition:
Ca + 2H2O → Ca(OH)2 + H2 + Q.
Calcium reacts with active non-metals (oxygen, chlorine, bromine) under normal conditions:
2Ca + O2 → 2CaO
Ca + Br2 → CaBr2.
When heated in air or oxygen, calcium ignites. Calcium interacts with less active non-metals (hydrogen, boron, carbon, silicon, nitrogen, phosphorus and others) when heated, for example:
Ca + H2 → CaH2, Ca + 6B = CaB6,
3Ca + N2 → Ca3N2, Ca + 2C → CaC2,
3Ca + 2P → Ca3P2 (calcium phosphide), calcium phosphides of the compositions CaP and CaP5 are also known;
2Ca + Si → Ca2Si (calcium silicide); calcium silicides of the compositions CaSi, Ca3Si4 and CaSi2 are also known.
The occurrence of the above reactions, as a rule, is accompanied by the release of a large amount of heat (that is, these reactions are exothermic). In all compounds with non-metals, the oxidation state of calcium is +2. Most of the calcium compounds with non-metals are easily decomposed by water, for example:
CaH2 + 2H2O → Ca(OH)2 + 2H2,
Ca3N2 + 6H2O → 3Ca(OH)2 + 2NH3.
The Ca2+ ion is colorless. When soluble calcium salts are added to the flame, the flame turns brick-red.
Calcium salts such as CaCl2 chloride, CaBr2 bromide, CaI2 iodide and Ca(NO3)2 nitrate are highly soluble in water. Fluoride CaF2, carbonate CaCO3, sulfate CaSO4, orthophosphate Ca3(PO4)2, oxalate CaC2O4 and some others are insoluble in water.
It is important that, unlike calcium carbonate CaCO3, acidic calcium carbonate (bicarbonate) Ca(HCO3)2 is soluble in water. In nature, this leads to the following processes. When cold rain or river water, saturated with carbon dioxide, penetrates underground and falls on limestone, their dissolution is observed:
CaCO3 + CO2 + H2O → Ca(HCO3)2.
In the same places where water saturated with calcium bicarbonate comes to the surface of the earth and is heated by the sun's rays, a reverse reaction occurs:
Ca(HCO3)2 → CaCO3 + CO2 + H2O.
This is how large masses of substances are transferred in nature. As a result, huge gaps can form underground, and beautiful stone “icicles” - stalactites and stalagmites - form in caves.
The presence of dissolved calcium bicarbonate in water largely determines the temporary hardness of water. It is called temporary because when water boils, bicarbonate decomposes and CaCO3 precipitates. This phenomenon leads, for example, to the fact that scale forms in the kettle over time.
Applications of calcium metal
The main use of calcium metal is as a reducing agent in the production of metals, especially nickel, copper and stainless steel. Calcium and its hydride are also used to obtain difficult-to-reduce metals such as chromium, thorium and uranium. Alloys of calcium with lead are used in batteries and bearing alloys. Calcium granules are also used to remove traces of air from vacuum devices.
Biological role
Calcium is a common macronutrient in the body of plants, animals and humans. In humans and other vertebrates, most of it is contained in the skeleton and teeth in the form of phosphates. The skeletons of most groups of invertebrates (sponges, coral polyps, mollusks, etc.) are made from various forms of calcium carbonate (lime). Calcium ions are involved in blood clotting processes, as well as in ensuring constant osmotic pressure of the blood. Calcium ions also serve as one of the universal second messengers and regulate a variety of intracellular processes - muscle contraction, exocytosis, including the secretion of hormones and neurotransmitters, etc. The calcium concentration in the cytoplasm of human cells is about 10−7 mol, in intercellular fluids about 10− 3 mol.
STRONTIUM
Being in nature
Strontium is not found in free form. It is part of about 40 minerals. Of these, the most important is celestine SrSO4 (51.2% Sr). Strontianite SrCO3 (64.4% Sr) is also mined. These two minerals are of industrial importance. Most often, strontium is present as an impurity in various calcium minerals.
Other strontium minerals include:
SrAl3(AsO4)SO4(OH)6 - kemmlicite;
Sr2Al(CO3)F5 - stenonite;
SrAl2(CO3)2(OH)4 H2O - strontiodresserite;
SrAl3(PO4)2(OH)5 H2O - goyasite;
Sr2Al(PO4)2OH - goodkenite;
SrAl3(PO4)SO4(OH)6 - svanbergite;
Sr(AlSiO4)2 - slosonite;
Sr(AlSi3O8)2 5H2O - brewsterite;
Sr5(AsO4)3F - fermorite;
Sr2(B14O23) 8H2O - strontioginorite;
Sr2(B5O9)Cl Н2О - strontiohilgardite;
SrFe3(PO4)2(OH)5 H2O - lyusunite;
SrMn2(VO4)2 4H2O - santafeite;
Sr5(PO4)3OH - belovite;
SrV(Si2O7) - Haradaite.
In terms of physical abundance in the earth's crust, strontium ranks 23rd - its mass fraction is 0.014% (in the lithosphere - 0.045%). The mole fraction of the metal in the earth's crust is 0.0029%. Strontium is found in sea water (8 mg/l).
In nature, strontium occurs in the form of a mixture of 4 stable isotopes 84Sr (0.56%), 86Sr (9.86%), 87Sr (7.02%), 88Sr (82.56%).

Receipt
There are 3 ways to obtain strontium metal:
thermal decomposition of some compounds
electrolysis
reduction of oxide or chloride.
Chemical properties
Strontium in its compounds always exhibits a valence of +2. In terms of properties, strontium is close to calcium and barium, occupying an intermediate position between them.
In the electrochemical voltage series, strontium is among the most active metals (its normal electrode potential is −2.89 V. It reacts vigorously with water, forming hydroxide:
Sr + 2H2O = Sr(OH)2 + H2
Interacts with acids, displaces heavy metals from their salts. It reacts weakly with concentrated acids (H2SO4, HNO3).
Metal strontium quickly oxidizes in air, forming a yellowish film, in which, in addition to SrO oxide, SrO2 peroxide and Sr3N2 nitride are always present. When heated in air, it ignites; powdered strontium in air is prone to self-ignition.
Reacts vigorously with non-metals - sulfur, phosphorus, halogens. Interacts with hydrogen (above 200°C), nitrogen (above 400°C). Practically does not react with alkalis.
At high temperatures it reacts with CO2 to form carbide:
5Sr + 2CO2 = SrC2 + 4SrO
Easily soluble strontium salts with the anions Cl−, I−, NO3−. Salts with anions F−, SO42−, CO32−, PO43− are slightly soluble.
Application
The main areas of application of strontium and its chemical compounds are the radio-electronic industry, pyrotechnics, metallurgy, and the food industry.
Metallurgy
Strontium is used for alloying copper and some of its alloys, for introduction into battery lead alloys, for desulfurization of cast iron, copper and steels.
Metallothermy
Strontium with a purity of 99.99-99.999% is used for the reduction of uranium.
Magnetic materials
Hard magnetic strontium ferrites are widely used as materials for the production of permanent magnets.
Pyrotechnics
In pyrotechnics, strontium carbonate, nitrate, and perchlorate are used to color the flame carmine red. The magnesium-strontium alloy has strong pyrophoric properties and is used in pyrotechnics for incendiary and signal compositions.
Nuclear energy
Strontium uranate plays an important role in the production of hydrogen (strontium-uranate cycle, Los Alamos, USA) by thermochemical method (atomic-hydrogen energy), and in particular, methods are being developed for the direct fission of uranium nuclei in the composition of strontium uranate to produce heat from the decomposition of water to hydrogen and oxygen.

Strontium oxide is used as a component of superconducting ceramics.
Chemical current sources
Strontium fluoride is used as a component of solid-state fluorine batteries with enormous energy capacity and energy density.
Strontium alloys with tin and lead are used for casting battery current leads. Strontium-cadmium alloys for galvanic cell anodes.
Biological role
Effect on the human body
The effect on the human body of natural (non-radioactive, low-toxic and, moreover, widely used for the treatment of osteoporosis) and radioactive isotopes of strontium should not be confused.
Natural strontium is a component of microorganisms, plants and animals. Strontium is an analogue of calcium, so it is most efficiently deposited in bone tissue. Less than 1% is retained in soft tissues. Strontium accumulates at a high rate in the body of children up to the age of four, when bone tissue is actively formed. Strontium metabolism changes in certain diseases of the digestive system and cardiovascular system.
BARIUM
Being in nature
The barium content in the earth's crust is 0.05% by weight; in sea water the average barium content is 0.02 mg/liter. Barium is active; it belongs to the subgroup of alkaline earth metals and is bound quite tightly in minerals. Main minerals: barite (BaSO4) and witherite (BaCO3).
Rare barium minerals: celsian or barium feldspar (barium aluminosilicate), hya-lophane (mixed barium and potassium aluminosilicate), nitrobarite (barium nitrate), etc.

Isotopes
Natural barium consists of a mixture of seven stable isotopes: 130Ba, 132Ba, 134Ba, 135Ba, 136Ba, 137Ba, 138Ba. The latter is the most common (71.66%). Radioactive isotopes of barium are also known, the most important of which is 140Ba. It is formed by the decay of uranium, thorium and plutonium.
Receipt
The main raw material for barium production is barite concentrate (80-95% BaSO4), which in turn is obtained by barite flotation. Barium sulfate is subsequently reduced with coke or natural gas:
BaSO4 + 4C = BaS + 4CO
BaSO4 + 2CH4 = BaS + 2C + 4H2O.
Next, the sulfide, when heated, is hydrolyzed to barium hydroxide Ba(OH)2 or, under the influence of CO2, is converted into insoluble barium carbonate BaCO3, which is then converted into barium oxide BaO (calcination at 800 °C for Ba(OH)2 and over 1000 °C for BaCO3):
BaS + 2H2O = Ba(OH)2 + H2S
BaS + H2O + CO2 = BaCO3 + H2S
Ba(OH)2 = BaO + H2O
BaCO3 = BaO + CO2
Barium metal is obtained from the oxide by reduction with aluminum in a vacuum at 1200-1250 °C:
4BaO + 2Al = 3Ba + BaAl2O4.
Barium is purified by vacuum distillation or zone smelting.
Chemical properties
Barium is an alkaline earth metal. In air, barium quickly oxidizes, forming a mixture of barium oxide BaO and barium nitride Ba3N2, and with slight heating it ignites. Reacts vigorously with water, forming barium hydroxide Ba(OH)2:
Ba + 2H2O = Ba(OH)2 + H2
Actively interacts with dilute acids. Many barium salts are insoluble or slightly soluble in water: barium sulfate BaSO4, barium sulfite BaSO3, barium carbonate BaCO3, barium phosphate Ba3(PO4)2. Barium sulfide BaS, unlike calcium sulfide CaS, is highly soluble in water.
Reacts easily with halogens to form halides.
When heated with hydrogen, it forms barium hydride BaH2, which in turn with lithium hydride LiH gives the Li complex.
Reacts when heated with ammonia:
6Ba + 2NH3 = 3BaH2 + Ba3N2
When heated, barium nitride Ba3N2 reacts with CO, forming cyanide:
Ba3N2 + 2CO = Ba(CN)2 + 2BaO
With liquid ammonia it gives a dark blue solution, from which ammonia can be isolated, which has a golden sheen and easily decomposes with the elimination of NH3. In the presence of a platinum catalyst, ammonia decomposes to form barium amide:
= Ba(NH2)2 + 4NH3 + H2
Barium carbide BaC2 can be obtained by heating BaO with coal in an arc furnace.
With phosphorus it forms phosphide Ba3P2.
Barium reduces the oxides, halides and sulfides of many metals to the corresponding metal.
Application
Anti-corrosion material
Barium is added together with zirconium to liquid metal coolants (alloys of sodium, potassium, rubidium, lithium, cesium) to reduce the aggressiveness of the latter to pipelines and in metallurgy.
Ferro- and piezoelectric
Barium titanate is used as a dielectric in the manufacture of ceramic capacitors, and as a material for piezoelectric microphones and piezoceramic emitters.
Optics
Barium fluoride is used in the form of single crystals in optics (lenses, prisms).
Pyrotechnics
Barium peroxide is used for pyrotechnics and as an oxidizing agent. Barium nitrate and barium chlorate are used in pyrotechnics to color flames (green fire).
Nuclear-hydrogen energy
Barium chromate is used in the production of hydrogen and oxygen by thermochemical method (Oak Ridge cycle, USA).
High temperature superconductivity
Barium oxide, together with oxides of copper and rare earth metals, is used to synthesize superconducting ceramics operating at liquid nitrogen temperatures and above.
Nuclear energy
Barium oxide is used to melt a special type of glass - used for coating uranium rods. One of the widespread types of such glasses has the following composition - (phosphorus oxide - 61%, BaO - 32%, aluminum oxide - 1.5%, sodium oxide - 5.5%). In glass melting for the nuclear industry, barium phosphate is also used.
Chemical current sources
Barium fluoride is used in solid-state fluorion batteries as a component of the fluoride electrolyte.
Barium oxide is used in high-power copper oxide batteries as a component of the active mass (barium oxide-copper oxide).
Barium sulfate is used as an expander of the active mass of the negative electrode in the production of lead-acid batteries.

Prices
Prices for barium metal in ingots with a purity of 99.9% fluctuate around $30 per 1 kg.
Biological role
The biological role of barium has not been sufficiently studied. It is not included in the list of vital microelements. All soluble barium salts are highly poisonous.
RADIUM
Radium (lat. Radium), Ra, radioactive chemical element of group II of the periodic system of Mendeleev, atomic number 88. Isotopes of Ra are known with mass numbers 213, 215, 219-230. The longest-lived is a-radioactive 226Ra with a half-life of about 1600 years. In nature, 222Ra (the special name of the isotope is actinium-X, symbol AcX), 224Ra (thorium-X, ThX), 226Ra and 228Ra (mesothorium-I, MsThI) are found as members of natural radioactive series.
STORY
The discovery of Ra was reported in 1898 by the spouses P. and M. Curie together with J. Bemont, shortly after A. Becquerel first (in 1896) discovered the phenomenon of radioactivity in uranium salts. In 1897, M. Sklodowska-Curie, who worked in Paris, found that the intensity of radiation emitted by uranium tar (the mineral uraninite) was much higher than could be expected, given the uranium content in the tar. Sklodowska-Curie suggested that this was caused by the presence of still unknown highly radioactive substances in the mineral. A thorough chemical study of uranium tar made it possible to discover two new elements - first polonium, and a little later - R. During the isolation of R., the behavior of the new element was monitored by its radiation, which is why the element was named from the Latin. radius - ray. To isolate the pure compound R., the Curies in the laboratory processed about 1 ton of factory waste remaining after extracting uranium from uranium tar. In particular, no less than 10,000 recrystallizations were performed from aqueous solutions of a mixture of BaCl2 and RaCl2 (barium compounds serve as so-called isomorphic carriers in the extraction of R). As a result, we managed to obtain 90 mg of pure RaCI2.
Ra is an extremely rare element. In uranium ores, which are its main source, there is no more than 0.34 g of Ra per 1 ton of U. R. belongs to highly dispersed elements and is found in very small concentrations in a wide variety of objects.
All Ra compounds exhibit a pale bluish glow in air. Due to the self-absorption of a- and b-particles emitted during the radioactive decay of 226Ra and its daughter products, each gram of 226Ra releases about 550 J (130 cal) of heat per hour, so the temperature of Ra preparations is always slightly higher than the ambient temperature.
PHYSICAL PROPERTIES
Ra is a silvery-white shiny metal that quickly tarnishes in air. Body-centered cubic lattice, estimated density 5.5 g/cm3. According to various sources, tpl. is 700-960 °C, tkip is about 1140 °C. The outer electron shell of the R atom contains 2 electrons (7s2 configuration). In accordance with this, R. has only one oxidation state +2 (valency II). In terms of chemical properties, R. is most similar to barium, but is more active. At room temperature, R. combines with oxygen, giving the oxide RaO, and with nitrogen, giving the nitride Ra3N2. R. reacts violently with water, releasing H2, and the strong base Ra (OH)2 is formed. R. chloride, bromide, iodide, nitrate and sulfide are highly soluble in water; carbonate, sulfate, chromate, and oxalate are poorly soluble.
CHEMICAL PROPERTIES
According to chemistry Holy shit, radium is similar to Va. Almost all radium compounds are isomorphic to the corresponding compounds. Va. In air, metallic radium quickly becomes covered with a dark film, which is a mixture of radium nitride and radium oxide. Metallic radium reacts violently with water to form the water-soluble hydroxide Ra(OH)2 and release H2. The electrode potential for radium release from water solutions is -1.718V (relative to a normal calomel electrode).

Radium compounds have the property of autoluminescence - glow in the dark due to their own properties. radiation. Mn. radium salts are colorless, but when decomposed under the action. own radiation acquires a yellow or brown color. Well sol. in water RaCl2 (mp 900 °C, density 4.91 g/cm3; see also table), RaBr2 (mp 728 °C, density 5.79 g/cm3), RaI2 and Ra(NO3)2. Better than other solutions. in water RaBr2 (70 g in 100 g at 20 ° C). Radium chloride and bromide crystallize from water in the form of crystalline hydrates with two or six molecules of H2O. Slightly soluble compounds are RaSO4 sulfate (approx. 2 10-4 g in 100 g of water at 20°C), Ra(IO3)2 iodate, RaF2 fluoride, RaCrO4 chromate, RaCO3 carbonate and RaC2O4 oxalate. Complexes of radium with lemon, tartar, apple, milk, ethylenediaminetetraacetic and other ligands are known. Compared to other alkali-earth. With metals, radium has a weaker tendency to form complexes.
Radium is isolated in the form of RaCl2 or other salts as a by-product of the processing of uranium ores (after extraction of U from them), using the methods of precipitation, fractional crystallization, and ion exchange; metallic radium is obtained by electrolysis of RaCl2 solution on a mercury cathode, reduction of RaO by aluminum upon heating. in a vacuum.

APPLICATION
The study of the properties of Ra played a huge role in the development of scientific knowledge, because made it possible to clarify many issues related to the phenomenon of radioactivity. For a long time, Ra was the only element whose radioactive properties found practical application (in medicine, for the preparation of luminous compounds, etc.). However, now in most cases it is more profitable to use not Ra, but cheaper artificial radioactive isotopes of other elements. Ra has retained some importance in medicine as a source of radon in the treatment of radon baths. Rum is used in small quantities for the preparation of neutron sources (in a mixture with beryllium) and in the production of light compositions (in a mixture with zinc sulfide).

BIOLOGICAL ROLE
Radium in the body. Of the natural radioactive isotopes, the long-lived 226Ra has the greatest biological significance. R. is unevenly distributed in different parts of the biosphere. There are geochemical provinces with a high content of phosphorus. The accumulation of phosphorus in the organs and tissues of plants obeys the general laws of absorption of mineral substances and depends on the type of plant and its growing conditions. As a rule, there is more R. in the roots and leaves of herbaceous plants than in the stems and reproductive organs; R. is most abundant in bark and wood. The average content of R. in flowering plants is 0.3-9.0 × 10-11 curie/kg, in sea. algae 0.2-3.2×10-11 curies/kg.
It enters the body of animals and humans with food, in which it is constantly present (in wheat 20-26×10-15 g/g, in potatoes 67-125×10-15 g/g, in meat 8×10-15 g/g) , as well as with drinking water. The daily intake of 226Ra into the human body with food and water is 2.3×10-12 curies, and losses with urine and feces are 0.8×10-13 and 2.2×10-12 curies. About 80% of R. that enters the body (it is similar in chemical properties to Ca) accumulates in bone tissue. The content of R. in the human body depends on the area of ​​residence and the nature of nutrition. Large concentrations of R. in the body have a harmful effect on animals and humans, causing painful changes in the form of osteoporosis, spontaneous fractures, and tumors. The content of R. in the soil above 1×10-7-10-8 curie/kg noticeably inhibits the growth and development of plants.

Let us consider the chemical properties of alkaline earth metals. Let us determine the features of their structure, production, occurrence in nature, and application.

Position in the PS

First, let's determine the location of these elements in Mendeleev. They are located in the second group of the main subgroup. These include calcium, strontium, radium, barium, magnesium, and beryllium. All of them do not contain two valence electrons. In general, beryllium, magnesium and alkaline earth metals have ns2 electrons in their outer shell. In chemical compounds they exhibit an oxidation state of +2. When interacting with other substances, they exhibit reducing properties, donating electrons from the external energy level.

Changing Properties

As the nucleus of an atom grows, beryllium and magnesium increase their metallic properties, as the radius of their atoms increases. Let's consider the physical properties of alkaline earth metals. Beryllium in its normal state is a gray metal with a steely luster. It has a dense hexagonal crystal lattice. Upon contact with atmospheric oxygen, beryllium immediately forms an oxide film, as a result of which its chemical activity decreases and a matte coating is formed.

Physical properties

Magnesium as a simple substance is a white metal that forms an oxide coating in air. It has a hexagonal crystal lattice.

The physical properties of the alkaline earth metals calcium, barium, and strontium are similar. They are metals with a characteristic silvery luster, which become covered with a yellowish film under the influence of atmospheric oxygen. Calcium and strontium have a face-centered cubic lattice, while barium has a body-centered structure.

The chemistry of alkaline earth metals is based on the fact that they have a metallic bond. That is why they are characterized by high electrical and thermal conductivity. Their melting and boiling points are higher than those of alkali metals.

Methods of obtaining

Beryllium is produced on an industrial scale by recovering the metal from fluoride. The condition for this chemical reaction to occur is preheating.

Considering that alkaline earth metals occur in nature in the form of compounds, to obtain magnesium, strontium, and calcium, electrolysis of molten salts is carried out.

Chemical properties

The chemical properties of alkaline earth metals are associated with the need to first remove the oxide film layer from their surface. It is this that determines the inertness of these metals to water. Calcium, barium, and strontium, when dissolved in water, form hydroxides that have pronounced basic properties.

The chemical properties of alkaline earth metals suggest their interaction with oxygen. For barium, the reaction product is peroxide; for all others, oxides are formed after the reaction. All representatives of this class of oxides exhibit basic properties; only beryllium oxide is characterized by amphoteric properties.

The chemical properties of alkaline earth metals also manifest themselves in reactions with sulfur, halogens, and nitrogen. When reacting with acids, dissolution of these elements is observed. Considering that beryllium is an amphoteric element, it is capable of chemical interaction with alkali solutions.

Qualitative reactions

The basic formulas of alkaline earth metals discussed in the course of inorganic chemistry are associated with salts. To identify representatives of this class in a mixture with other elements, a qualitative definition can be used. When salts of alkaline earth metals are added to the flame of an alcohol lamp, coloring of the flame by cations is observed. The strontium cation produces a dark red tint, the calcium cation produces an orange tint, and the barium cation produces a green tint.

To identify the barium cation in qualitative analysis, sulfate anions are used. As a result of this reaction, white barium sulfate is formed, which is insoluble in inorganic acids.

Radium is a radioactive element that occurs in nature in trace quantities. When magnesium interacts with oxygen, a blinding flash is observed. This process has been used for some time when photographing in dark rooms. Magnesium flares have now been replaced by electrical systems. Beryllium is a member of the alkaline earth metal family, which reacts with many chemicals. Calcium and magnesium, like aluminum, can reduce such rare metals as titanium, tungsten, molybdenum, niobium. The data are called calcithermia and magnesothermia.

Features of application

What are the uses of alkaline earth metals? Calcium and magnesium are used to make light alloys and rare metals.

For example, magnesium is contained in duralumin, and calcium is a component of lead alloys used to produce cable sheaths and create bearings. Alkaline earth metals are widely used in technology in the form of oxides. (calcium oxide) and burnt magnesium (magnesium oxide) are required for the construction industry.

When calcium oxide interacts with water, a significant amount of heat is released. (calcium hydroxide) is used for construction. The white suspension of this substance (lime milk) is used in the sugar industry for the process of purifying beet juice.

Salts of metals of the second group

Salts of magnesium, beryllium, and alkaline earth metals can be obtained by reacting with acids of their oxides. Chlorides, fluorides, and iodides of these elements are white crystalline substances, generally highly soluble in water. Among sulfates, only magnesium and beryllium compounds are soluble. Its decrease is observed from beryllium salts to barium sulfates. Carbonates are practically insoluble in water or have minimal solubility.

Sulfides of alkaline earth elements are found in small quantities in heavy metals. If you shine light on them, you can get different colors. Sulfides are included in luminous compounds called phosphors. Similar paints are used to create luminous dials and road signs.

Common alkaline earth metal compounds

Calcium carbonate is the most common element on the earth's surface. It is an integral part of compounds such as limestone, marble, chalk. Among them, limestone has the main use. This mineral is indispensable in construction and is considered an excellent building stone. In addition, quicklime and slaked lime, glass, and cement are obtained from this inorganic compound.

The use of limestone helps strengthen roads, and thanks to the powder, soil acidity can be reduced. represents the shells of ancient animals. This compound is used to make rubber, paper, and school crayons.

Marble is in demand among architects and sculptors. Many of Michelangelo's unique creations were created from marble. Some Moscow metro stations are lined with marble tiles. Magnesium carbonate is used in large quantities in the manufacture of brick, cement, and glass. It is needed in the metallurgical industry to remove waste rock.

Calcium sulfate, found naturally in the form of gypsum (calcium sulfate crystalline hydrate), is used in the construction industry. In medicine, this compound is used to make impressions, as well as to create plaster casts.

Alabaster (semi-hydrous gypsum) releases a huge amount of heat when interacting with water. This is also applied in industry.

Epsom salts (magnesium sulfate) are used medicinally as a laxative. This substance has a bitter taste and is found in sea water.

“Barite porridge” (barium sulfate) does not dissolve in water. That is why this salt is used in x-ray diagnostics. Salt blocks X-rays, which makes it possible to detect diseases of the gastrointestinal tract.

Phosphorites (rock) and apatites contain calcium phosphate. They are needed to obtain calcium compounds: oxides, hydroxides.

Calcium plays a special role in living organisms. It is this metal that is necessary to build the bone skeleton. Calcium ions are necessary to regulate heart function and increase blood clotting. Its deficiency causes disturbances in the functioning of the nervous system, loss of coagulability, and loss of the ability of the hands to hold various objects normally.

In order to avoid health problems, a person should consume approximately 1.5 grams of calcium every day. The main problem is that in order for the body to absorb 0.06 grams of calcium, you need to eat 1 gram of fat. The maximum amount of this metal is found in lettuce, parsley, cottage cheese, and cheese.

Conclusion

All representatives of the second group of the main subgroup of the periodic table are necessary for the life and activity of modern man. For example, magnesium is a stimulator of metabolic processes in the body. It must be present in nervous tissue, blood, bones, and liver. Magnesium is an active participant in photosynthesis in plants, as it is a component of chlorophyll. Human bones make up about a fifth of the total weight. They contain calcium and magnesium. Oxides and salts of alkaline earth metals have found various applications in the construction industry, pharmaceuticals and medicine.

The main subgroup of the second group of the periodic table covers the elements: beryllium, magnesium, calcium, strontium, barium and radium. Based on the main representatives of this subgroup - calcium, strontium and barium - known collectively as alkaline earth metals, the entire main subgroup of the second group is also called the subgroup alkaline earth metals.

These metals (sometimes magnesium is also added to them) received the name “alkaline earth” because their oxides in their chemical properties are intermediate, on the one hand, between alkalis, i.e. oxides or hydroxides of alkali metals and, on the other hand, “ earths,” that is, oxides of such elements, a typical representative of which is aluminum, the main component of clays. Due to this intermediate position, the oxides of calcium, strontium and barium were given the name “alkaline earths”.

The first element of this subgroup, beryllium (if you do not take into account its valency), is much closer in its properties to aluminum than to the higher analogs of the top group to which it belongs. The second element of this group, magnesium, also differs in some respects significantly from the alkaline earth metals in the narrow sense of the term. Some reactions bring it closer to elements of the secondary subgroup of the second group, especially zinc; Thus, magnesium and zinc sulfates, in contrast to sulfates of alkaline earth metals, are easily soluble, isomorphic to each other and form double salts of similar composition. Previously, a rule was stated according to which the first element exhibits properties that are transitional to the next main subgroup, the second - to a secondary subgroup of the same group; and usually only the third element has properties characteristic of the group; This rule is especially clearly manifested in the group of alkaline earth metals.

The heaviest of the elements of the second group - radium - in its chemical properties, of course, corresponds to typical representatives of the alkaline earth metals. However, it is usually not customary to include it in the group of alkaline earth metals in a narrower sense. Due to the peculiarities of its distribution in nature, as well as due to its most characteristic property - radioactivity, it is more appropriate to give it a special place. In the discussion of the general properties of the elements of this subgroup, radium will not be considered, since the corresponding physicochemical properties have not yet been sufficiently studied.

With the exception of radium, all elements of the alkaline earth subgroup are light metals. Light metals are those whose specific gravity does not exceed 5. In terms of their hardness, the metals of the main subgroup of group II are significantly superior to alkaline metals. The softest of them, barium (whose properties are closest to the alkali metals) has approximately the hardness of lead. The melting points of metals in this group are significantly higher than those of the alkali metals.

What is common to all elements of the main subgroup of group II is their property of exhibiting positive valency 2 in their compounds, and only in very exceptional cases are they positively monovalent. Their typical valency 2+, as well as the atomic numbers of the elements, undoubtedly force these metals to be classified as the main subgroup of the second group. In addition, they all exhibit a strongly electropositive character, which is determined by their position on the left side of the electrochemical voltage series, as well as a strong affinity for electronegative elements.

In accordance with the value of the normal potentials of the elements of the main subgroup of the second group, all of the listed metals decompose water; however, the effect of beryllium and magnesium on water occurs very slowly due to the low solubility of the hydroxides resulting from this reaction, for example for magnesium:

Mg + 2НН = Mg(OH) 2 + H 2

Having formed on the metal surface, Be and Mg hydroxides hinder the further course of the reaction. Therefore, even small errors of magnesium must be kept at normal temperature in contact with water for several days before they are completely converted into magnesium hydroxide. The remaining alkaline earth metals react with water much more vigorously, which is explained by the better solubility of their hydroxides. Barium hydroxide is the easiest to dissolve; Ba's normal potential is the lowest compared to other elements in the group, so it reacts very vigorously with water, as well as with alcohol. The resistance of alkaline earth metals to air decreases in the direction from magnesium to barium. In accordance with their position in the stress series, the named metals displace all heavy metals from solutions of their salts.

Normal oxides M II O are always obtained as combustion products of alkaline earth metals. Peroxides of alkaline earth metals are much less stable than in the series of alkali metals.

Oxides of alkaline earth metals combine with water to form hydroxides, Moreover, the energy of this reaction increases very noticeably in the direction from BeO to BaO. The solubility of hydroxides also increases greatly from beryllium hydroxide to barium hydroxide; But even the solubility of the latter at normal temperature is very low. The basic character of these compounds increases in the same order - from amphoteric beryllium hydroxide to strongly basic caustic barium.

It is interesting to note the strong affinity of the elements of the main subgroup of the second group for nitrogen. The tendency to form compounds with nitrogen increases in these elements with increasing atomic weight (despite the fact that the heat of formation of nitrides in this direction decreases); In the alkali earth metals themselves, the tendency to form nitrides is so great that the latter slowly combine with nitrogen even at normal temperatures.

Alkaline earth metals like alkali metals, they combine with hydrogen to form hydrides, for example:

Ca+H2 = CaH2.

Ethn hydrides also have a salt-like character, and therefore it should be assumed that in them, as in alkali metal hydrides, hydrogen is an electronegative component.

It is more difficult to obtain MgH 2 directly from the elements, but it was not possible to synthesize BeH 2 in this way at all. MgH 2 and BeH 2 are solid and non-volatile compounds, like hydrides of alkaline earth metals, but unlike the latter they do not have a pronounced salt-like character.

All elements of the main subgroup of the second group form colorless ions with a positive charge 2: Be 2+, Mg 2+, Ca 2+, Sr 2+, Ba 2+, Ra 2+. Beryllium also forms colorless anions [BeO 2 ] 2+ and [Be(OH) 4 ] 2+. All salts M II X 2 of these elements are also colorless, unless they are derivatives of colored anions.

Radium salts themselves are also colorless. However, some of them, such as radium chloride and bromide, are gradually colored by the radiation of the radium contained in them and finally acquire a color from brown to black. When recrystallized they become white again.

Many alkaline earth metal salts are poorly soluble in water. A certain pattern is often revealed in the change in solubility of these salts: for example, for sulfates, the solubility quickly decreases with increasing atomic weight of the alkaline earth metal. The solubility of chromites changes in approximately the same way. Most of the salts formed by alkaline earth metals with weak acids and with acids of medium strength are difficult to dissolve, for example phosphates, oxalates and carbonates; some of them, however, are easily soluble; the latter include sulfides, cyanides, thiocyanates and acetates. Due to the weakening of the basic character of hydroxides during the transition from Ba to Be, the degree of hydrolysis of their carbonates increases in the same sequence. Their thermal stability also changes in the same direction: while barium carbonate, even at white-hot temperatures, is far from completely decomposed, calcium carbonate can be completely decomposed into CaO and CO 2 even with relatively weak calcination, and magnesium carbonate decomposes even more easily.

From the point of view of Kossel's theory, the reason for the divalency of the elements of the alkaline earth group is the fact that in the periodic table they are all removed from the corresponding inert gases with: 2 elements, therefore each of them has 2 more electrons than the previous inert gas. Due to the tendency of atoms to adopt the configuration of inert gases in the elements of the alkaline earth group, a slight abstraction of two electrons occurs, but no more, since further abstraction would cause the destruction of the configuration of the inert gases.

Alkaline earth metals include metals of group IIA of the Periodic Table D.I. Mendeleev - calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). In addition to them, the main subgroup of group II includes beryllium (Be) and magnesium (Mg). The outermost energy level of alkaline earth metals contains two valence electrons. The electronic configuration of the outer energy level of alkaline earth metals is ns 2. In their compounds they exhibit a single oxidation state of +2. In OVR they are reducing agents, i.e. give up an electron.

With an increase in the charge of the nucleus of atoms of elements included in the group of alkaline earth metals, the ionization energy of atoms decreases, and the radii of atoms and ions increase, the metallic characteristics of chemical elements increase.

Physical properties of alkaline earth metals

In the free state, Be is a steel-gray metal with a dense hexagonal crystal lattice, quite hard and brittle. In air, Be becomes covered with an oxide film, which gives it a matte tint and reduces its chemical reactivity.

Magnesium in the form of a simple substance is a white metal, which, like Be, when exposed to air acquires a matte tint due to the formation of an oxide film. Mg is softer and more ductile than beryllium. The Mg crystal lattice is hexagonal.

Ca, Ba and Sr in free form are silvery-white metals. When exposed to air, they instantly become covered with a yellowish film, which is the product of their interaction with the components of the air. Calcium is a fairly hard metal, Ba and Sr are softer.

Ca and Sr have a face-centered cubic crystal lattice, barium has a body-centered cubic crystal lattice.

All alkaline earth metals are characterized by the presence of a metallic type of chemical bond, which determines their high thermal and electrical conductivity. The boiling and melting points of alkaline earth metals are higher than those of alkali metals.

Preparation of alkaline earth metals

Be is produced by the reduction reaction of its fluoride. The reaction occurs when heated:

BeF 2 + Mg = Be + MgF 2

Magnesium, calcium and strontium are obtained by electrolysis of molten salts, most often chlorides:

CaCl 2 = Ca + Cl 2

Moreover, when producing Mg by electrolysis of a dichloride melt, NaCl is added to the reaction mixture to lower the melting point.

To obtain Mg in industry, metal- and carbon-thermal methods are used:

2(CaO×MgO) (dolomite) + Si = Ca 2 SiO 4 + Mg

The main method of obtaining Ba is the reduction of the oxide:

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

Chemical properties of alkaline earth metals

Since in no. the surface of Be and Mg is covered with an oxide film - these metals are inert towards water. Ca, Sr and Ba dissolve in water to form hydroxides exhibiting strong basic properties:

Ba + H 2 O = Ba(OH) 2 + H 2

Alkaline earth metals are capable of reacting with oxygen, and all of them, with the exception of barium, as a result of this interaction form oxides, barium - peroxide:

2Ca + O2 = 2CaO

Ba + O 2 = BaO 2

Oxides of alkaline earth metals, with the exception of beryllium, exhibit basic properties, Be - amphoteric properties.

When heated, alkaline earth metals are capable of interacting with nonmetals (halogens, sulfur, nitrogen, etc.):

Mg + Br 2 =2MgBr

3Sr + N 2 = Sr 3 N 2

2Mg + 2C = Mg 2 C 2

2Ba + 2P = Ba 3 P 2

Ba + H 2 = BaH 2

Alkaline earth metals react with acids and dissolve in them:

Ca + 2HCl = CaCl 2 + H 2

Mg + H 2 SO 4 = MgSO 4 + H 2

Beryllium reacts with aqueous solutions of alkalis - dissolves in them:

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

Qualitative reactions

A qualitative reaction to alkaline earth metals is the coloring of the flame by their cations: Ca 2+ colors the flame dark orange, Sr 2+ - dark red, Ba 2+ - light green.

A qualitative reaction to the barium cation Ba 2+ is SO 4 2- anions, resulting in the formation of a white precipitate of barium sulfate (BaSO 4), insoluble in inorganic acids.

Ba 2+ + SO 4 2- = BaSO 4 ↓

Examples of problem solving

EXAMPLE 1

The second group of D.I. Mendeleev’s periodic table contains a group of elements that are very similar in their properties to alkali metals, but are inferior to them in activity. It includes beryllium and magnesium, as well as calcium, strontium, barium and radium. They are known collectively as alkaline earth elements. In our article we will get acquainted with their distribution in nature and use in industry, and also study the most important chemical properties of alkaline earth metals.

general characteristics

All atoms of the above elements contain two electrons in their outer energy layer. When interacting with other substances, they always give up their negative particles, turning into the state of cations with a charge of 2+. In redox reactions, elements behave as strong reducing agents. As the nuclear charge increases, the chemical properties of alkaline earth metals and their activity increase. In air they quickly oxidize, forming an oxide film on their surface. The general formula of all oxides is RO. They correspond to hydroxides with the formula R(OH) 2. Their basic properties and solubility in water also increase with increasing atomic number of the element.

Special properties of beryllium and magnesium

In some of their properties, the first two representatives of the main subgroup of the second group are somewhat different from other alkaline earth elements. This manifests itself, in particular, during their interaction with water. For example, the chemical properties of beryllium are such that it does not react at all with H 2 O. Magnesium reacts with water only when heated. But all alkaline earth elements easily react with it at ordinary temperatures. What substances are formed in this case?

Alkaline earth metal bases

Being active elements, calcium, barium and other representatives of the group quickly displace hydrogen from water, resulting in their hydroxides. The interaction of alkaline earth metals with water proceeds violently, with the release of heat. Solutions of calcium, barium, and strontium bases are soapy to the touch and cause severe burns if they come into contact with the skin and mucous membrane of the eyes. The first aid in such cases will be to treat the wound surface with a weak solution of acetic acid. It will neutralize alkali and reduce the risk of necrosis of damaged tissue.

Chemical properties of alkaline earth metals

Interaction with oxygen, water and non-metals is the main list of properties of metals included in the second group of the periodic system of chemical elements. For example, calcium, even under normal conditions, reacts with halogens: fluorine, chlorine, bromine and iodine. When heated, it combines with sulfur, carbon and nitrogen. Hard oxidation - combustion, ends with the formation of calcium oxide: 2Ca + O 2 = 2 CaO. The interaction of metals with hydrogen leads to the appearance of hydrides. They are white, refractory substances with ionic crystal lattices. Important chemical properties of alkaline earth metals include their interaction with water. As stated earlier, the product of this displacement reaction will be a metal hydroxide. We also note that in the main subgroup of the second group, calcium occupies the most significant place. Therefore, let us dwell on its characteristics in more detail.

Calcium and its compounds

The content of the element in the earth's crust is up to 3.5%, which indicates its widespread occurrence in minerals such as limestone, chalk, marble and calcite. Natural calcium contains six types of isotopes. It is also found in natural water sources. Alkali metal compounds are studied in detail in the course of inorganic chemistry. For example, in 9th grade lessons, students learn that calcium is a light but strong silvery-white metal. Its melting and boiling points are higher than those of alkaline elements. The main method of production is electrolysis of a mixture of molten salts of calcium chloride and calcium fluoride. The main chemical properties include its reactions with oxygen, water and non-metals. Of the alkali metal compounds, calcium oxide and base are the most important for industry. The first compound is obtained from chalk or limestone by burning them.

Next, calcium hydroxide is formed from calcium oxide and water. Its mixture with sand and water is called mortar. It continues to be used as plaster and for joining bricks when laying walls. A calcium hydroxide solution called limewater is used as a reagent to detect carbon dioxide. When carbon dioxide is passed through a clear aqueous solution of Ca(OH) 2, it becomes cloudy due to the formation of an insoluble precipitate of calcium carbonate.

Magnesium and its characteristics

The chemistry of alkaline earth metals studies the properties of magnesium, focusing on some of its features. It is a very light, silvery-white metal. Magnesium, molten in an atmosphere with high humidity, actively absorbs hydrogen molecules from water vapor. As the metal cools, it almost completely releases them back into the air. It reacts very slowly with water due to the formation of a slightly soluble compound - magnesium hydroxide. Alkalis have no effect on magnesium at all. The metal does not react with some acids: concentrated sulfate and hydrofluoric acids, due to its passivation and the formation of a protective film on the surface. Most mineral acids dissolve the metal, which is accompanied by the rapid release of hydrogen. Magnesium is a strong reducing agent; it replaces many metals from their oxides or salts:

BeO + Mg = MgO + Be.

The metal, together with beryllium, manganese, and aluminum, is used as an alloying additive to steel. Magnesium-containing alloys - electrons - have especially valuable properties. They are used in aircraft and automobile production, as well as in parts of optical equipment.

The role of elements in the life of organisms

Let us give examples of alkaline earth metals, the compounds of which are common in living nature. Magnesium is the central atom in chlorophyll molecules in plants. It is involved in the process of photosynthesis and is part of the active centers of green pigment. Magnesium atoms capture light energy, then converting it into the energy of chemical bonds of organic compounds: glucose, amino acids, glycerol and fatty acids. The element plays an important role as a necessary component of enzymes that regulate metabolism in the human body. Calcium is a macroelement that ensures the effective passage of electrical impulses through nerve tissue. The presence of its phosphoric acid salts in bones and tooth enamel gives them hardness and strength.

Beryllium and its properties

Alkaline earth metals also include beryllium, barium and strontium. Consider beryllium. The element is not very common in nature; it is mainly found in minerals, such as beryl. Its varieties containing multi-colored impurities form precious stones: emeralds and aquamarines. The peculiarity of the physical properties is fragility and high hardness. A distinctive feature of the element’s atom is the presence on the second outer energy level of not eight, like all other alkaline earth metals, but only two electrons.

Therefore, the radius of the atom and ion is disproportionately small, and the ionization energy is high. This determines the high strength of the metal crystal lattice. The chemical properties of beryllium also distinguish it from other elements of the second group. It reacts not only with acids, but also with alkali solutions, displacing hydrogen and forming hydroxoberyllates:

Be + 2NaOH + 2H 2 O = Na 2 + H 2.

The metal has a number of unique characteristics. Due to its ability to transmit x-rays, it is used to make windows for x-ray tubes. In the nuclear industry, the element is considered the best moderator and reflector of neutrons. In metallurgy, it is used as a valuable alloying additive that increases the anti-corrosion properties of alloys.

Strontium and barium

The elements are quite common in nature and, like the alkaline earth metal magnesium, are found in minerals. Let's call them: barite, celestine, strontianite. Barium has the appearance of a ductile metal with a silvery-white color. Like calcium, it is represented by several isotopes. In air, it actively interacts with its components - oxygen and nitrogen, forming barium oxide and nitride. For this reason, the metal is stored under a layer of paraffin or mineral oil, avoiding its contact with air. Both metals form peroxides when heated to 500°C.

Of these, barium peroxide has practical application, used as a fabric bleach. The chemical properties of the alkaline earth metals barium and strontium are similar to those of calcium. However, their interaction with water is much more active, and the resulting bases are stronger than calcium hydroxide. Barium is used as an additive to liquid metal coolants that reduces corrosion, in optics, and in the manufacture of vacuum electronic devices. Strontium is in demand in the production of photocells and phosphors.

Qualitative reactions using alkaline earth metal ions

Barium and strontium compounds are examples of alkaline earth metals widely used in pyrotechnics due to the bright coloring of the flame by their ions. Thus, strontium sulfate or carbonate gives a carmine-red glow of the flame, and the corresponding barium compounds give a yellow-green glow. To detect calcium ions in the laboratory, several grains of calcium chloride are poured onto the burner flame; the flame turns brick-red.

A solution of barium chloride is used in analytical chemistry to identify ions of the acidic residue of sulfate acid in a solution. If, when the solutions are drained, a white precipitate of barium sulfate is formed, it means that there were SO 4 2- particles in it.

In our article, we studied the properties of alkaline earth metals and gave examples of their use in various industries.