General characteristics of the IIA group of the Periodic Table of elements.
This group contains the following elements: Be, Mg, Ca, Sr, Ba, Ra. They have a common electronic configuration: (n-1)p6ns2, except for 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 Ca - Sr - Ba - Ra series, the properties change sequentially. The relative electronegativity in the series Be - Ra falls because. as the size of the atom increases, valence electrons are given away more readily. The properties of the elements of the IIA subgroup are determined by the ease of giving off two ns-electrons. In this case, E2+ ions are formed. When studying X-ray diffraction, it turned out that in some compounds the elements of the IIA subgroup exhibit monovalence. 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 already used in ancient times (5000 years ago) by the Egyptians in the construction business. 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, barite are 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 “funny lights”. In 1787, a new mineral was found in a lead mine near the Scottish village of Strontian, which they called SrCO3 strontianite. A. Crawford suggested the existence of the still unknown "earth". In 1792, T. Hop proved that the composition of the found mineral includes a new element - strontium. At the same time, insoluble strontium disaccharate (С12Н22О4.2SrO) was isolated using Sr (OH) 2, when obtaining sugar from molasses. Sr production increased. However, it was soon noticed that the analogous calcium sucrose was also insoluble, and calcium oxide was undoubtedly cheaper. Interest in strontium immediately disappeared and only rose again in the 1940s. Heavy spar was the first known barium compound. 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 to electrolysis with a mercury cathode a mixture of wet slaked lime with mercury oxide, prepared a calcium amalgam, and after driving mercury out of it, he obtained a metal called "calcium" (from Latin Calx, genus case calcis - lime). In the same way, Devi received Ba and Sr. An industrial method for obtaining calcium was developed by Suter and Redlich in 1896 at the Rathenau plant (Germany). In 1904, the first plant for the production of calcium began to operate.
Radium was predicted by Mendeleev in 1871 and discovered in 1898 by Marie and Pierre Curie. They found that uranium ores are more radioactive than uranium itself. The reason was radium compounds. They treated the remains of uranium ore with alkali, and what did not dissolve with hydrochloric acid. The residue after the second treatment was more radioactive than the ore. Radium was found in this fraction. The Curie spouses reported their discovery in a report for 1898.
Prevalence of alkaline earth metals
The content of calcium 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 sulfate salts: CaCO3 - calcite, CaSO4 - andidrite, SrCO3 - strontianite, SrSO4 - celestine, BaCO3 - witherite. BaSO4 - heavy spar. Fluorite CaF2 is also a useful mineral.
Sa plays 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 contained in the human body in negligible amounts.
Obtaining alkaline earth metals
First, oxides or chlorides are obtained by E. EO is obtained by calcining ESO3, and ES12 by the action of 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 vacuum in order to avoid oxidation of E. Calcium (like all other E) can be obtained by electrolysis of the CaCl2 melt, followed by vacuum distillation or thermal dissociation of CaC2. Ba and Sr can be obtained by pyrolysis of E2N3, E(NH3)6, EN2. Radium is mined incidentally from uranium ores.
Physical properties alkaline earth metals
Ca and its analogs are silver-white metals. Calcium is the hardest of them all. Strontium and especially barium are much softer than calcium. All alkaline earth metals are ductile, lend themselves well 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 a temperature 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. He is paramagnetic. Barium crystallizes in the bcc structure. Ca and Sr are able to form a continuous series of solid solutions with each other, 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
E2+ ion radius, 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
Tm.,оС 852 770 710 800
Boiling point, оС 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 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, the entire 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 =-2.84V(Ca), =-2.89(Sr). E are very active elements: they dissolve in water and acids, displace most metals from their oxides, halides, sulfides. Primarily (200-300°C) calcium interacts with water vapor according to the scheme: 2Ca + H2O = CaO + CaH2. Secondary reactions have the form: CaH2 + 2H2O = Ca(OH)2 + 2H2 and CaO + H2O = Ca(OH)2. In strong sulfuric acid, E almost do not dissolve due to the formation of a film of poorly soluble ESO4. With dilute mineral acids, E react violently with the release of hydrogen. When heated above 800°C, calcium reacts with methane according to the scheme: 3Ca + CH4 = CaH2 + CaC2. When heated, they react with hydrogen, with sulfur and with gaseous ammonia. 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 counterparts. All radium compounds slowly decompose under the action of their own radiation, while 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 higher than the temperature environment by 1.5 deg. It is also known that 1 g of radium per day releases 1 mm3 of radon (226Ra = 222Rn + 4He), which is the basis for its use as a source of radon for radon baths.
Hydrides E are white, crystalline salt-like substances. They are obtained directly from the elements when heated. The temperatures of the start of the reaction E + H2 = EN2 are 250 °C (Ca), 200 °C (Sr), 150 °C (Ba). Thermal dissociation of EN2 starts 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 able to reduce oxides to metals (W, Nb, Ti, Ce, Zr, Ta), for example, 2CaH2 + TiO2 = 2CaO + 2H2 + Ti. The reaction of CaH2 with Al2O3 proceeds at 750°C: 3CaH2 + Al2O3 = 3CaO + 3H2 + 2Al, and then: CaH2 + 2Al = CaAl2 + H2. CaH2 reacts with nitrogen at 600°C according to the scheme: 3CaH2 + N2 = Ca3N2 + 3H2. When EN2 is ignited, they slowly burn out: EN2 + O2 = H2O + CaO. Explosive when mixed with solid oxidizers. Under the action of water on EN2, hydroxide and hydrogen are released. This reaction is highly exothermic: EN2 wetted with water in air ignites spontaneously. 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 the elements at elevated temperatures. They decompose with water according to the scheme: E3N2 + 6H2O = 3E (OH) 2 + 2NH3. E3N2 react when heated with CO according to the scheme: E3N2 + 3CO = 3EO + N2 + 3C. The processes that occur when E3N2 is heated with coal look like this:
E3N2 + 5C = ECN2 + 2ES2; (E = Ca, Sr); Ba3N2 + 6C = Ba(CN)2 + 2BaC2;
Strontium nitride reacts with HCl to give Sr and ammonium chlorides. Phosphides E3P2 are formed directly from the elements or by calcining 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, alkaline earth metal phosphides give the corresponding salt and phosphine. This is the basis for their application to obtain phosphine in the laboratory.
Complex ammonia composition E(NH3)6 - solids with a metallic luster and high electrical conductivity. They are obtained by the action of liquid ammonia on E. They ignite spontaneously in air. Without air access, they decompose into the corresponding amides: E (NH3) 6 \u003d E (NH2) 2 + 4NH3 + H2. When heated, they vigorously decompose according to the same pattern.
Alkaline earth metal carbides, which are obtained by calcining E with coal, are decomposed by water with the release of acetylene: ES2 + 2H2O \u003d E (OH) 2 + C2H2. The reaction with BaC2 is so violent that it ignites on contact with water. The heats of formation of ES2 from the elements for Ca and Ba are 14 and 12 kcalmoles. When heated with nitrogen, ES2 gives CaCN2, Ba(CN)2, SrCN2. Known silicides (ESi and ESi2). They can be obtained by heating directly from the elements. They hydrolyze with water and react with acids to give H2Si2O5, SiH4, the corresponding E compound, and hydrogen. Known EV6 borides obtained from the elements by heating.
Calcium oxides and its analogues are white refractory (TboilCaO = 2850 ° C) substances that absorb water vigorously. 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). EO dissolve in acids and ammonium chloride: EO + 2NH4Cl = SrCl2 + 2NH3 + H2O. EO is obtained by calcining 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(OCH3)2. During magnesium-thermal reduction of BaO, an intermediate oxide Ba2O can be obtained, which is unstable and disproportionate.
Hydroxides of alkaline earth metals are white substances soluble in water. 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, and Ca(OH)2.H2O are usually isolated from hydroxide solutions. EOs add water to form hydroxides. The use of CaO in construction is based on this. 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 scavenger. Ca(OH)2, when standing in air, absorbs CO2 according to the 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 are used to open carbonic anhydride.
E form white peroxides. They are significantly less stable than oxides and are strong oxidizers. Of practical importance is the most stable BaO2, which is a white, paramagnetic powder with a density of 4.96 g1cm3, and so on. 450°. BaO2 is stable at ordinary temperature (it can be stored for years), it is poorly soluble in water, alcohol and ether, it dissolves in dilute acids with the release of salt and hydrogen peroxide. Thermal decomposition of barium peroxide is accelerated by 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 balance 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 air flow to 500°C according to the scheme: 2BaO + O2 = 2BaO2. As the temperature rises, the reverse process takes place. Therefore, when Ba burns, only oxide is released. SrO2 and CaO2 are less stable. General Method to obtain EO2 is the interaction of E(OH)2 with H2O2, with the release of EO2.8H2O. Thermal decomposition of EO2 starts at 380°C (Ca), 480°C (Sr), 790°C (Ba). When EO2 is heated with concentrated hydrogen peroxide, yellow unstable substances, EO4 superoxides, can be obtained.
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. Halides E are divided into two groups: fluorides and all the rest. 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 isolated 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
Temp. arr-I, 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
Tm., оС 1423 782 760 575 1473 872 643 515 1353 962 853 740
Boiling point, оС 2500 2000 1800 718 2460 2030 2260 1830
D(EG) in vapors, 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 precipitates, 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 chloride melts 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: BeCl2 - 0.009%, MgCl2 - 14.6%, CaCl2 - 43.3%, SrCl2 - 60.6%, BaCl2 - 80.2%. Halides (except fluorides) E contain water of crystallization: CaCl2.6H2O, SrCl2.6H2O and BaCl2.2H2O. X-ray diffraction analysis established the structure of E[(OH2)6]G2 for Ca and Sr crystalline hydrates. With slow heating of EG2 crystalline hydrates, anhydrous salts can be obtained. CaCl2 readily forms supersaturated solutions. Natural CaF2 (fluorite) is used in the ceramic industry and is also used to produce HF and is a fluorine mineral. Anhydrous CaCl2 is used as a desiccant due to its hygroscopicity. Calcium chloride hydrate is used for the preparation of refrigeration mixtures. ВаСl2 - used in cx and to open SO42-(Ва2+ + SO42- = ВаSO4). Fusion of EG2 and EN2 hydrohalides can be obtained: EG2 + EN2 = 2ENG. These substances melt without decomposition but are hydrolyzed by water: 2ENG + 2H2O = EG2 + 2H2 + E(OH)2. The water solubility of chlorates, bromates and iodates in water decreases in the series Ca - Sr - Ba and Cl - Br - I. Ba (ClO3) 2 - is used in pyrotechnics. E perchlorates are highly soluble not only in water but also in organic solvents. The most important of the 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 milk of lime. It is an oxidizing agent and is highly soluble in water. Bleach can be obtained by acting with chlorine on solid slaked lime. It decomposes with water and smells like chlorine in the presence of moisture. Reacts with CO2 in the air:
CO2 + 2CaOCl2 = CaCO3 + CaCl2 + Cl2O.
Chlorine lime is used as an oxidizing agent, bleach and as a disinfectant.
Azides E(N3)2 and thiocyanides E(CNS)2.3H2O are known for alkaline earth metals. Azides compared to lead azide are much less explosive. The thiocyanates 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 an exchange reaction.
Calcium and strontium nitrates usually exist in the form of Ca(NO3)2.4H2O and Sr(NO3)2.4H2O crystalline hydrates. 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). Melt of calcium nitrate crystalline hydrate has an acidic environment at 75 °C. A feature of barium nitrate is the low rate of dissolution of its crystals in water. Only barium nitrate exhibits a tendency to complex formation, for which an unstable complex K2 is known. Calcium nitrate is soluble in alcohols, methyl acetate, acetone. Strontium and barium nitrates are almost insoluble there. The melting points of nitrates E are estimated at 600 ° C, however, at the same temperature, decomposition begins: E (NO3) 2 \u003d 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 salts E 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 HAOs: 5s and 5p or 5s and 4d. We will inform the energy of this system - we will heat it. Electrons from orbitals closer to the nucleus will move to these HAOs. But such a system is not stable and will release energy in the form of a quantum of light. It is Sr2+ that emits quanta with a frequency corresponding to red wavelengths. When obtaining 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 matter, bleaching the flame of the reducing agent, and being a binding agent. Calcium nitrate is used as a fertilizer.
All phosphates and hydrophosphates E are poorly soluble in water. They can be obtained by dissolving an appropriate amount of CaO or CaCO3 in orthophosphoric acid. They are also precipitated during exchange reactions of the type: (3-x)Ca2+ + 2HxPO4-(3-x) = Ca(3-x)(HxPO4)2. Practical value (as a fertilizer) is monosubstituted calcium orthophosphate, which, along with Ca (SO4), is part of superphosphate. It is received 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 dehydrates at 200 °C, and decomposes at 430 °C according to the scheme: CaC2O4 = CaCO3 + CO. E acetates are isolated in the form of crystalline hydrates and are highly soluble in water.
Sulfates E are white, poorly soluble substances 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 sulfate ion reagent. Calcium sulfate contains water of crystallization. Above 66 °C, anhydrous calcium sulfate is released from the solution, below - gypsum CaSO4.2H2O. Heating of gypsum above 170°C is accompanied by the release of water of hydration. When mixing gypsum 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 as early as 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 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.
Sulfides E are white solids that crystallize as NaCl. The heats of their formation and the energies of the crystal lattices are (kcalmol): 110 and 722 (Ca), 108 and 687 (Sr), 106 and 656 (Ba). They can be obtained by synthesis from elements during heating or by calcining sulfates with coal: ES04 + 3C = ES + CO2 + 2CO. Less soluble CaS (0.2 hl). ES enters into the following reactions when heated:
ES + H2O = EO + H2S; ES + G2 = S + EG2; ES + 2O2 = ES04; ES + xS = ESx+1 (x=2.3).
Alkaline earth metal sulfides 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(NS)2 + xS = ESx+1 + H2S (x=2,3,4).
Of the crystalline hydrates, BaS.6H2O and Ca(HS)2.6H2O, Ba(HS)2.4H2O are known. Ca(HS)2 is used to remove hair. ES are subject to the phenomenon of phosphorescence. Polysulfides 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, Ca thiosulfate can be obtained according to the scheme: 2CaS + 2O2 + H2O \u003d Ca (OH) 2 + CaS2O3. It is highly soluble in water. In the Ca - Sr - Ba series, the solubility of thiosulfates decreases. Tellurides E 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 as sharply as in the case of sulfates. These substances yellow color are obtained by the interaction of soluble salts E with chromates (or dichromates) of alkali metals: E2 + + CrO42- \u003d ECrO4. Calcium chromate is isolated in the form of a crystalline hydrate - CaCrO4.2H2O (rPR 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, ESO3 pass into EO, splitting off CO2. In the Ca-Ba series, the thermal stability of carbonates increases. The most practically important of them 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 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 schemes:
CaO + H2O = Ca(OH)2 and Ca(OH)2 + CO2 = CaCO3 + H2O.
A colossally 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. In its production, a mixture of CaCO3 with clay is fired before the start of sintering (1400-1500 ° C). The mixture is then ground. The composition of cement can be expressed as a percentage of the components CaO, SiO2, Al2O3, Fe2O3, with CaO being a base, and everything else being acid anhydrides. The composition of silicate (Portland) cement consists mainly of Ca3SiO5, Ca2SiO4, Ca3(AlO3)2 and Ca(FeO2)2. His grasp goes according to the schemes:
Ca3SiO5 + 3H2O = Ca2SiO4.2H2O + Ca(OH)2
Ca2SiO4 + 2Н2О = Ca2SiO4.2Н2О
Ca3(AlO3)2 + 6Н2О = Ca3(AlO3)2.6Н2О
Ca(FeO2)2 + nH2O = Ca(FeO2)2.nH2O.
Natural chalk is introduced into the composition of various putties. Fine-crystalline, precipitated from a solution of CaCO3 is part of the tooth powders. From BaCO3 by calcination with coal, BaO is obtained according to the scheme: BaCO3 + C \u003d BaO + 2CO. If the process is carried out at a higher temperature in a stream of nitrogen, barium cyanide is formed: ВаСО3 + 4С + 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. Bicarbonates E are soluble in water and can only be obtained 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 how to fix it
Soluble calcium and magnesium salts determine the overall hardness of water. If they are present in water in small quantities, then the water is called soft. With a high content of these salts (100 - 200 mg of calcium salts - in 1 liter in terms of ions), water is considered hard. In such water, soap foams poorly, since calcium and magnesium salts form insoluble compounds with it. In hard water, food products are poorly boiled, and when boiled, it gives scale on the walls of household utensils and steam boilers. Scale has a low thermal conductivity, causes an increase in fuel consumption or power consumption of an electrical appliance and accelerates the wear of the walls of a 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. The hardness caused by the presence of acidic calcium and magnesium carbonates in water is called carbonate or temporary hardness, as it can be eliminated. In addition to carbonate hardness, non-carbonate hardness is also distinguished, 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 constant hardness. Carbonate and non-carbonate hardness add up to total hardness. To completely eliminate it, water is sometimes distilled. But it's expensive. To eliminate carbonate hardness, water can be boiled, but this is also expensive and scale forms. Hardness is removed by adding an appropriate amount of Ca(OH)2: Ca(OH)2 + Ca(HCO3)2 = CaCO3↓ + 2H2O. The overall hardness is eliminated either by adding Na2CO3 or by using the so-called cation exchangers. When sodium carbonate is used, soluble calcium and magnesium salts are also converted into insoluble carbonates: Ca2+ + Na2CO3 = 2Na+ + CaCO3↓. Removing stiffness with cation exchangers is a more advanced process. Cation exchangers are high-molecular sodium-containing organic compounds, whose composition can be expressed by the formula Na2R, where R is a complex acid residue. When water is filtered through a layer of cation exchanger, 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 table salt. In this case, the reverse process occurs: the Ca2+ ions in the crystal lattice in the cation exchanger are replaced by Na+ ions from the solution. The regenerated cation exchanger is again used for water purification. Permutite-based filters work in a similar way:
Na2 + Ca2+ = 2Na+ + Ca
Being in nature
Due to the high chemical activity of calcium in the free form in nature is not found.
Most of the calcium is contained in the composition of silicates and aluminosilicates of various rocks (granites, gneisses, etc.), especially in feldspar - anorthite Ca.
In the form of sedimentary rocks, calcium compounds are represented by chalk and limestone, 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, apatites 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, which migrates vigorously in the earth's crust and accumulates in various geochemical systems, forms 385 minerals (fourth in terms of the number of minerals).
Calcium accounts for 3.38% of the mass of the earth's crust (5th place in abundance after oxygen, silicon, aluminum and iron). Element content in sea water- 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 - is 96.97%.
Of the six naturally occurring calcium isotopes, five are stable. The sixth 48Ca isotope, 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 5.3×1019 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 by 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 interacts with oxygen, carbon dioxide and moisture in the air, due to which the surface of metallic calcium is usually dull gray, therefore, 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 that calcium actively reacts with water, but without ignition:
Ca + 2H2O → Ca(OH)2 + H2 + Q.
With active non-metals (oxygen, chlorine, bromine), calcium reacts at normal conditions:
2Ca + O2 → 2CaO
Ca + Br2 → CaBr2.
When heated in air or oxygen, calcium ignites. With less active non-metals (hydrogen, boron, carbon, silicon, nitrogen, phosphorus and others), calcium interacts when heated, for example:
Ca + H2 → CaH2, Ca + 6B = CaB6,
3Ca + N2 → Ca3N2, Ca + 2C → CaC2,
3Ca + 2P → Ca3P2 (calcium phosphide), calcium phosphides of CaP and CaP5 compositions are also known;
2Ca + Si → Ca2Si (calcium silicide), calcium silicides of compositions CaSi, Ca3Si4 and CaSi2 are also known.
The course 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. CaF2 fluoride, CaCO3 carbonate, CaSO4 sulfate, Ca3(PO4)2 orthophosphate, CaC2O4 oxalate and some others are insoluble in water.
Of great importance is the fact that, unlike calcium carbonate CaCO3, acidic calcium carbonate (hydrocarbonate) Ca(HCO3)2 is soluble in water. In nature, this leads to following processes. When cold rain or river water, saturated with carbon dioxide, penetrates underground and falls on limestones, 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 heats up sunbeams, the reverse reaction takes place:
Ca(HCO3)2 → CaCO3 + CO2 + H2O.
So in nature there is a transfer of large masses of substances. As a result, huge gaps can form underground, and beautiful stone "icicles" - stalactites and stalagmites - form in the caves.
The presence of dissolved calcium bicarbonate in water largely determines the temporary hardness of water. It is called temporary because when boiling water, the bicarbonate decomposes, and CaCO3 precipitates. This phenomenon leads, for example, to the fact that scale forms in the teapot over time.
Applications of metallic calcium
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 hard-to-recover metals such as chromium, thorium and uranium. Alloys of calcium with lead are used in storage batteries and bearing alloys. Calcium granules are also used to remove traces of air from electrovacuum devices.
Biological role
Calcium is a common macronutrient in plants, animals and humans. In humans and other vertebrates, most of it is found in the skeleton and teeth in the form of phosphates. From various forms calcium carbonate (lime) consists of the skeletons of most groups of invertebrates (sponges, coral polyps, mollusks, etc.). Calcium ions are involved in the processes of blood coagulation, as well as in maintaining a 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 does not occur 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 - gudkenite;
SrAl3(PO4)SO4(OH)6 - svanbergite;
Sr(AlSiO4)2 - slosonite;
Sr(AlSi3O8)2 5H2O - brusterite;
Sr5(AsO4)3F - fermorite;
Sr2(B14O23) 8H2O - strontium-ginorite;
Sr2(B5O9)Cl Н2О - strontium chilhardite;
SrFe3(PO4)2(OH)5 H2O - lusunite;
SrMn2(VO4)2 4Н2О - santafeite;
Sr5(PO4)3OH - whitish;
SrV(Si2O7) - charadaite.
In terms of the level of physical abundance in the earth's crust, strontium occupies the 23rd place - its mass fraction is 0.014% (in the lithosphere - 0.045%). The mole fraction of metal in the earth's crust is 0.0029%. Strontium is found in sea water (8 mg/l).
In nature, strontium occurs as 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 metallic strontium:
thermal decomposition of some compounds
electrolysis
reduction of oxide or chloride.
Chemical properties
Strontium in its compounds always exhibits a +2 valency. By properties, strontium is close to calcium and barium, occupying an intermediate position between them.
In the electrochemical series of voltages, 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
Reacts with acids heavy metals from their salts. It reacts weakly with concentrated acids (H2SO4, HNO3).
Strontium metal rapidly 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.
Vigorously reacts 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 salts of strontium with anions Cl–, I–, NO3–. Salts with anions F–, SO42–, CO32–, PO43– are sparingly soluble.
Application
The main areas of application of strontium and its chemical compounds- this is the radio-electronic industry, pyrotechnics, metallurgy, food industry.
Metallurgy
Strontium is used for alloying copper and some of its alloys, for introducing into battery lead alloys, for the desulfurization of cast iron, copper and steels.
Metalthermy
Strontium with a purity of 99.99-99.999% is used to reduce uranium.
Magnetic materials
Magnetically hard strontium ferrites are widely used as materials for the production of permanent magnets.
Pyrotechnics
In pyrotechnics, strontium carbonate, nitrate, perchlorate are used to color the flame in carmine red. The magnesium-strontium alloy has the strongest 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 the 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 during 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.
Alloys of strontium with tin and lead are used for casting battery down conductors. Strontium-cadmium alloys for anodes of galvanic cells.
Biological role
Impact on the human body
One should not confuse 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.
Natural strontium is an integral part of microorganisms, plants and animals. Strontium is an analog of calcium, so it is most effectively 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 there is an active formation of bone tissue. The exchange of strontium changes in some diseases of the digestive system and the cardiovascular system.
BARIUM
Being in nature
The content of barium in the earth's crust is 0.05% by weight; in sea water, the average content of barium is 0.02 mg / liter. Barium is active, it belongs to the subgroup of alkaline earth metals and is quite strongly bound in minerals. The main minerals are barite (BaSO4) and witherite (BaCO3).
Rare barium minerals: Celsian or barium feldspar (barium aluminosilicate), hyalofan (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 during the decay of uranium, thorium and plutonium.
Receipt
The main raw material for obtaining barium is barite concentrate (80-95% BaSO4), which in turn is obtained by barite flotation. Barium sulfate is further reduced with coke or natural gas:
BaSO4 + 4С = BaS + 4CO
BaSO4 + 2CH4 = BaS + 2С + 4H2O.
Further, the sulfide, when heated, is hydrolyzed to barium hydroxide Ba (OH) 2 or under the action 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 oxide by aluminum reduction in vacuum at 1200-1250 °C:
4BaO + 2Al = 3Ba + BaAl2O4.
Barium is purified by vacuum distillation or zone melting.
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 ignites when heated slightly. Vigorously reacts with water, forming barium hydroxide Ba (OH) 2:
Ba + 2H2O \u003d 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 well soluble in water.
Easily reacts with halogens to form halides.
When heated with hydrogen, it forms barium hydride BaH2, which, in turn, with lithium hydride LiH gives a Li complex.
Reacts on heating with ammonia:
6Ba + 2NH3 = 3BaH2 + Ba3N2
When heated, barium nitride Ba3N2 reacts with CO to form 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 luster 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 the phosphide Ba3P2.
Barium reduces 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.
Ferroelectric 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).
Atomic hydrogen energy
Barium chromate is used in the production of hydrogen and oxygen by a 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 to coat 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 glassmaking for the nuclear industry, barium phosphate is also used.
Chemical current sources
Barium fluoride is used in solid-state fluorine batteries as a component of the fluoride electrolyte.
Barium oxide is used in powerful copper oxide batteries as a component of the active mass (barium oxide-copper oxide).
Barium sulfate is used as a negative electrode active mass expander in the production of lead-acid batteries.
Prices
Prices for metal barium in ingots with a purity of 99.9% fluctuate around $30 per 1 kg.
Biological role
The biological role of barium has not been studied enough. It is not included in the number of vital trace elements. All soluble barium salts are highly toxic.
RADIUM
Radium (lat. Radium), Ra, a radioactive chemical element of group II of the periodic system of Mendeleev, atomic number 88. Ra isotopes with mass numbers 213, 215, 219-230 are known. The longest-lived is a-radioactive 226Ra with a half-life of about 1600 years. In nature, as members of the natural radioactive series, there are 222Ra (the special name of the isotope is actinium-X, symbol AcX), 224Ra (Thorium-X, ThX), 226Ra and 228Ra (mesothorium-I, MsThI).
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 on uranium salts. In 1897, M. Sklodowska-Curie, who worked in Paris, established that the intensity of radiation emitted by uranium resin (the mineral uraninite) is much higher than could be expected, given the content of uranium in the resin. Sklodowska-Curie suggested that this was due to the presence in the mineral of yet unknown strongly radioactive substances. A thorough chemical study of uranium pitch 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, and therefore the element was named from lat. radius - ray. In order to isolate the pure compound R., the Curie spouses processed about 1 ton of factory waste in laboratory conditions, which remained after the extraction of uranium from uranium pitch. In particular, at least 10,000 recrystallizations were carried out 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, 90 mg of pure RaCI2 was obtained.
Ra is an extremely rare element. In uranium ores, which are its main source, 1 ton of U accounts for no more than 0.34 g of Ra. R. belongs to highly scattered elements and is found in very small concentrations in a wide variety of objects.
All Ra compounds exhibit a pale bluish luminescence 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.
PHYSICAL PROPERTIES
Ra is a silvery white lustrous metal that quickly tarnishes in air. The cubic lattice is body-centered, the calculated density is 5.5 g/cm3. According to various sources, tpl. is 700-960 °С, tkip is about 1140 °С. On the outer electron shell of the R. atom, there are 2 electrons (configuration 7s2). In accordance with this, R. has only one oxidation state, +2 (valency II). By 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, liberating H2, and the strong base Ra(OH)2 is formed. Chloride, bromide, iodide, nitrate, and R. sulfide are readily soluble in water; carbonate, sulfate, chromate, and oxalate are poorly soluble.
CHEMICAL PROPERTIES
According to chem. St-you radium is similar to Va. Almost all compounds of radium are isomorphic to the corresponding compounds. Wa. In air, metallic radium is quickly covered with a dark film, which is a mixture of nitride and radium oxide. Metallic radium reacts violently with water to form hydroxide Ra(OH)2, which is soluble in water, and release H2. The electrode potential of radium release from aqueous solutions is -1.718V (relative to the normal calomel electrode).
Radium compounds have St-tion autoluminescence-glow in the dark due to their own. radiation. Mn. salts of radium colorless., but when decomposed under the action. own radiation becomes yellow or brown. 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 solvents 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 H2O molecules. Slightly soluble compounds - sulfate RaSO4 (approx. 2 10-4 g in 100 g of water at 20 ° C), iodate Ra (IO3) 2, fluoride RaF2, chromate RaCrO4, carbonate RaCO3 and oxalate RaC2O4. Radium complexes with citric, tartaric, malic, lactic, ethylenediaminetetraacetic acids and other ligands are known. Compared with other alkaline earth. metals, radium has a weaker tendency to complex formation.
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, ion exchange; metallic radium is obtained by electrolysis of the RaCl2 solution on a mercury cathode, by the reduction of RaO with aluminum during heating. in a vacuum.
APPLICATION
The study of the properties of Ra has played a huge role in the development of scientific knowledge, because. made it possible to elucidate many questions connected with 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 compositions, 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. In small quantities, R. is spent on the preparation of neutron sources (mixed with beryllium) and in the production of light compositions (mixed 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 high content R. The accumulation of R. in the organs and tissues of plants obeys the general patterns of absorption of mineral substances and depends on the type of plant and the conditions of its growth. Usually found in roots and leaves herbaceous plants R. more than in the stems and organs of reproduction; most R. in the bark and wood. The average content of R. in flowering plants is 0.3-9.0 × 10-11 curie / kg, in the sea. algae 0.2-3.2×10-11 curie/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 the losses with urine and feces are 0.8×10-13 and 2.2×10-12 curies. About 80% of the R. that enters the body (it is close in chemical properties to Ca) accumulates in bone tissue. R.'s maintenance in a human body depends on the area of residence and character of food. 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 over 1 × 10-7-10-8 curie / kg significantly inhibits the growth and development of plants.
Class: 9
Lesson type: learning new material.
Type of lesson: combined lesson
Lesson objectives:
Tutorials: the formation of students' knowledge about alkaline earth elements as typical metals, the concept of the relationship between the structure of atoms and properties (physical and chemical).
Developing: skill development research activities, the ability to extract information from various sources, compare, generalize, draw conclusions.
Educators: nurturing a steady interest in the subject, cultivating such moral qualities as accuracy, discipline, independence, responsible attitude to the task assigned.
Methods: problematic, search, laboratory work, independent work students.
Equipment: computer, safety table, disk “Virtual laboratory in chemistry”, presentation .
During the classes
1. Organizational moment.
2. Introductory word of the teacher.
We study the section, metals, and you know what metals have great importance in life modern man. In previous lessons, we got acquainted with the elements of group I of the main subgroup - alkali metals. Today we are starting to study the metals of group II of the main subgroup - alkaline earth metals. In order to assimilate the material of the lesson, we need to remember the most important questions that were considered in the previous lessons.
3. Actualization of knowledge.
Conversation.
Where are the alkali metals in the periodic system of D.I. Mendeleev?
Student:
In the periodic system, alkali metals are located in group I of the main subgroup, on the outer level 1 electron, which alkali metals easily give away, therefore, in all compounds they exhibit an oxidation state of +1. With an increase in the size of atoms from lithium to francium, the ionization energy of atoms decreases and, as a rule, their chemical activity increases.
Teacher:
Physical properties of alkali metals?
Student:
All alkali metals are silvery-white in color with slight tints, light, soft and fusible. Their hardness and melting point naturally decrease from lithium to cesium.
Teacher:
We will check the knowledge of the chemical properties of alkali metals in the form of a small test work on the options:
- Ioption: Write the reaction equations for the interaction of sodium with oxygen, chlorine, hydrogen, water. Specify the oxidizing agent and reducing agent.
- I option: Write the reaction equations for the interaction of lithium with oxygen, chlorine, hydrogen, water. Specify the oxidizing agent and reducing agent.
- I I I option: Write the reaction equations for the interaction of potassium with oxygen, chlorine, hydrogen, water. Specify the oxidizing agent and reducing agent.
Teacher: The topic of our lesson is “Alkaline earth metals”
Lesson objectives: Give general characteristics alkaline earth metals.
Consider their electronic structure, compare physical and chemical properties.
Learn about the most important compounds of these metals.
Determine the scope of these compounds.
Our lesson plan is written on the board, we will work according to the plan, look at the presentation.
- The position of metals in the periodic system D.I. Mendeleev.
- The structure of the alkali metal atom.
- physical properties.
- Chemical properties.
- The use of alkaline earth metals.
Conversation.
Teacher:
Based on the previous knowledge, we will answer the following questions: next questions: Let's use the periodic table to answer chemical elements DI. Mendeleev.
1. List the alkaline earth metals
Student:
These are magnesium, calcium, strontium, barium, radium.
Teacher:
2. Why are these metals called alkaline earth?
Student:
The origin of this name is due to the fact that their hydroxides are alkalis, and their oxides are similar in refractoriness to oxides of aluminum and iron, which were previously worn. common name"lands"
Teacher:
3. Location of alkaline earth metals in PSCE D.I. Mendeleev.
Student:
Group II is the main subgroup. Metals of group II of the main subgroup have 2 electrons at the external energy level, located at a smaller distance from the nucleus than alkali metals. Therefore, their reducing properties, although great, are still less than those of the elements of group I. Strengthening of the reducing properties is also observed in the transition from Mg to Ba, which is associated with an increase in the radii of their atoms, in all compounds they exhibit an oxidation state of +2.
Teacher: Physical properties of alkaline earth metals?
Student:
Metals of group II of the main subgroup are silvery-white substances that conduct heat well and electricity. Their density increases from Be to Ba, while the melting point, on the contrary, decreases. They are much harder than alkali metals. All, except beryllium, have the ability to color the flame in different colors.
Problem: How are alkaline earth metals found in nature?
Why do alkaline earth metals mostly exist in nature in the form of compounds?
Answer: In nature, alkaline earth metals are in the form of compounds, because they have high chemical activity, which in turn depends on the characteristics electronic structure atoms (the presence of two unpaired electrons at the outer energy level)
Fizkultminutka - rest for the eyes.
Teacher:
Knowing the general physical properties, the activity of metals, assume the chemical properties of alkaline earth metals. What substances do alkali metals interact with?
Student:
Alkaline earth metals interact with both simple and complex substances. They actively interact with almost all non-metals (with halogens, hydrogen, forming hydrides). From complex substances with water - forming water-soluble bases - alkalis and with acids.
Teacher:
And now, in experiments, we will verify the correctness of our assumptions about the chemical properties of alkaline earth metals.
4. Laboratory work on the virtual laboratory.
Target: carry out reactions confirming the chemical properties of alkaline earth metals.
We repeat the safety rules for working with alkaline earth metals.
- work in a fume hood
- on a tray
- with dry hands
- take in small quantities
We work with the text that we read in the virtual laboratory.
Experience No. 1. Interaction of calcium with water.
Experience number 2. Combustion of magnesium, calcium, strontium, barium
Write down the reaction and observation equations in a notebook.
5. Summing up the lesson, grading.
5. Reflection.
What do you remember about the lesson, what did you like?
6. Homework.
§ 12 exercise 1(b) exercise 4
Literature.
- Rudzitis G.E., Feldman F.G. Chemistry 9.- Moscow.: Education, 2001
- Gabrielyan O.S. Chemistry 9.-Moscow.: Bustard, 2008
- Gabrielyan O.S., Ostroumov I.G. Handbook of the teacher. Chemistry 9.-Moscow.: Bustard 2002
- Gabrielyan O.S. Control and verification work. Chemistry 9.-Moscow.: Bustard, 2005.
- Collection of the Virtual Laboratory. Educational electronic edition
Group IIA contains only metals - Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium) and Ra (radium). The chemical properties of the first representative of this group, beryllium, differ most strongly from the chemical properties of the other elements of this group. Its chemical properties are in many ways even more similar to aluminum than to other group IIA metals (the so-called "diagonal similarity"). Magnesium, in terms of chemical properties, also differs markedly from Ca, Sr, Ba, and Ra, but still has much more similar chemical properties with them than with beryllium. Due to the significant similarity of the chemical properties of calcium, strontium, barium and radium, they are combined into one family, called alkaline earth metals.
All elements of group IIA belong to s-elements, i.e. contain all of their valence electrons s-sublevel. Thus, the electronic configuration of the outer electron layer of all chemical elements of this group has the form ns 2 , Where n– number of the period in which the element is located.
Due to the peculiarities of the electronic structure of group IIA metals, these elements, in addition to zero, are capable of having only one single oxidation state, equal to +2. Simple substances formed by elements of group IIA, with the participation of any chemical reactions can only oxidize, i.e. donate electrons:
Me 0 - 2e - → Me +2
Calcium, strontium, barium and radium are extremely reactive. The simple substances formed by them are very strong reducing agents. Magnesium is also a strong reducing agent. The reducing activity of metals obeys the general laws of the periodic law of D.I. Mendeleev and increases down the subgroup.
Interaction with simple substances
with oxygen
Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of a liquid inert to them, most often kerosene.
Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO 2):
2Mg + O 2 \u003d 2MgO
2Ca + O 2 \u003d 2CaO
2Ba + O 2 \u003d 2BaO
Ba + O 2 \u003d BaO 2
It should be noted that during the combustion of alkaline earth metals and magnesium in air, the reaction of these metals with atmospheric nitrogen also proceeds side by side, as a result of which, in addition to compounds of metals with oxygen, nitrides with the general formula Me 3 N 2 are also formed.
with halogens
Beryllium reacts with halogens only at high temperatures, while the rest of the Group IIA metals already at room temperature:
Mg + I 2 \u003d MgI 2 - magnesium iodide
Ca + Br 2 \u003d CaBr 2 - calcium bromide
Ba + Cl 2 \u003d BaCl 2 - barium chloride
with non-metals of IV–VI groups
All metals of group IIA react when heated with all non-metals of groups IV-VI, but depending on the position of the metal in the group, as well as the activity of non-metals, a different degree of heating is required. Since beryllium is the most chemically inert among all metals of group IIA, its reactions with nonmetals require significantly more O high temperature.
It should be noted that the reaction of metals with carbon can form carbides of various nature. There are carbides related to methanides and conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by a metal. They, like methane, contain carbon in the -4 oxidation state, and during their hydrolysis or interaction with non-oxidizing acids, methane is one of the products. There is also another type of carbides - acetylenides, which contain the C 2 2- ion, which is actually a fragment of the acetylene molecule. Carbides of the acetylenide type upon hydrolysis or interaction with non-oxidizing acids form acetylene as one of the reaction products. What type of carbide - methanide or acetylenide - will be obtained by the interaction of one or another metal with carbon depends on the size of the metal cation. As a rule, methanides are formed with metal ions having a small radius, with ions more large size- acetylenides. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:
The remaining metals of group II A form acetylenides with carbon:
With silicon, group IIA metals form silicides - compounds of the Me 2 Si type, with nitrogen - nitrides (Me 3 N 2), phosphorus - phosphides (Me 3 P 2):
with hydrogen
All alkaline earth metals react when heated with hydrogen. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, an increased pressure of hydrogen is also required. Beryllium does not react with hydrogen under any conditions.
Interaction with complex substances
with water
All alkaline earth metals actively react with water to form alkalis (soluble metal hydroxides) and hydrogen. Magnesium reacts with water only during boiling, due to the fact that when heated, the protective oxide film of MgO dissolves in water. In the case of beryllium, the protective oxide film is very resistant: water does not react with it either when boiling or even at a red heat temperature:
with non-oxidizing acids
All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the activity series to the left of hydrogen. In this case, a salt of the corresponding acid and hydrogen are formed. Reaction examples:
Be + H 2 SO 4 (razb.) \u003d BeSO 4 + H 2
Mg + 2HBr \u003d MgBr 2 + H 2
Ca + 2CH 3 COOH = (CH 3 COO) 2 Ca + H 2
with oxidizing acids
− dilute nitric acid
With diluted nitric acid all metals of group IIA react. In this case, the reduction products instead of hydrogen (as in the case of non-oxidizing acids) are nitrogen oxides, mainly nitrogen oxide (I) (N 2 O), and in the case of highly dilute nitric acid, ammonium nitrate (NH 4 NO 3):
4Ca + 10HNO 3 ( razb .) \u003d 4Ca (NO 3) 2 + N 2 O + 5H 2 O
4Mg + 10HNO3 (very disaggregated)\u003d 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O
− concentrated nitric acid
Concentrated nitric acid at ordinary (or low) temperature passivates beryllium, i.e. does not react with it. When boiling, the reaction is possible and proceeds mainly in accordance with the equation:
Magnesium and alkaline earth metals react with concentrated nitric acid to form a wide range of different nitrogen reduction products.
− concentrated sulfuric acid
Beryllium is passivated with concentrated sulfuric acid, i.e. does not react with it under normal conditions, however, the reaction proceeds during boiling and leads to the formation of beryllium sulfate, sulfur dioxide and water:
Be + 2H 2 SO 4 → BeSO 4 + SO 2 + 2H 2 O
Barium is also passivated by concentrated sulfuric acid due to the formation of insoluble barium sulfate, but reacts with it when heated, barium sulfate dissolves when heated in concentrated sulfuric acid due to its conversion to barium hydrogen sulfate.
The remaining metals of the main group IIA react with concentrated sulfuric acid under any conditions, including in the cold. Sulfur reduction can occur to SO 2, H 2 S and S, depending on the activity of the metal, the reaction temperature and the concentration of the acid:
Mg + H 2 SO 4 ( conc .) \u003d MgSO 4 + SO 2 + H 2 O
3Mg + 4H2SO4 ( conc .) \u003d 3MgSO 4 + S↓ + 4H 2 O
4Ca + 5H2SO4 ( conc .) \u003d 4CaSO 4 + H 2 S + 4H 2 O
with alkalis
Magnesium and alkaline earth metals do not interact with alkalis, and beryllium easily reacts both with alkali solutions and with anhydrous alkalis during fusion. At the same time, when the reaction is carried out in aqueous solution water is also involved in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and hydrogen gas:
Be + 2KOH + 2H 2 O \u003d H 2 + K 2 - potassium tetrahydroxoberyllate
When carrying out the reaction with solid alkali during fusion, beryllates of alkali or alkaline earth metals and hydrogen are formed.
Be + 2KOH \u003d H 2 + K 2 BeO 2 - potassium beryllate
with oxides
Alkaline earth metals, as well as magnesium, can reduce less active metals and some non-metals from their oxides when heated, for example:
The method of restoring metals from their oxides with magnesium is called magnesiumthermy.
Alkaline earth metals are elements that belong to the second group periodic table. These include substances such as calcium, magnesium, barium, beryllium, strontium and radium. The name of this group indicates that in water they give an alkaline reaction.
Alkali and alkaline earth metals, or rather their salts, are widely distributed in nature. They are represented by minerals. The exception is radium, which is considered a fairly rare element.
All of the above metals have some common qualities, which made it possible to combine them into one group.
Alkaline earth metals and their physical properties
Almost all of these elements are grayish solids (at least at normal conditions and By the way, the physical properties are slightly different - these substances, although quite persistent, are easily affected.
Interestingly, with the serial number in the table, such an indicator of the metal as density also grows. For example, in this group, calcium has the lowest index, while radium is similar in density to iron.
Alkaline earth metals: chemical properties
To begin with, it is worth noting that the chemical activity increases according to the serial number of the periodic table. For example, beryllium is a fairly stable element. It reacts with oxygen and halogens only when heated strongly. The same goes for magnesium. But calcium is able to slowly oxidize even at room temperature. The remaining three representatives of the group (radium, barium and strontium) quickly react with atmospheric oxygen even at room temperature. That is why these elements are stored, covering with a layer of kerosene.
The activity of oxides and hydroxides of these metals increases in the same way. For example, beryllium hydroxide does not dissolve in water and is considered an amphoteric substance, but is considered a fairly strong alkali.
Alkaline earth metals and their a brief description of
Beryllium is a light gray hard metal with high toxicity. The element was first discovered in 1798 by the chemist Vauquelin. There are several minerals of beryllium in nature, of which the following are considered the most famous: beryl, phenakite, danalite and chrysoberyl. By the way, some beryllium isotopes are highly radioactive.
Interestingly, some forms of beryl are valuable gemstones. These include emerald, aquamarine and heliodor.
Beryllium is used to make some alloys. This element is used to slow down neutrons.
Calcium is one of the best known alkaline earth metals. In its pure form, it is a soft white substance with a silvery tint. Pure calcium was first isolated in 1808. In nature, this element is present in the form of minerals such as marble, limestone and gypsum. Calcium is widely used in modern technologies. It is used as a chemical fuel source and also as a fire retardant material. It's no secret that calcium compounds are used in the production building materials and medicines.
This element is also found in every living organism. Basically, he is responsible for the operation of the motor apparatus.
Magnesium is a light and fairly malleable metal with a characteristic grayish color. It was isolated in its pure form in 1808, but its salts became known much earlier. Magnesium is found in minerals such as magnesite, dolomite, carnallite, kieserite. By the way, magnesium salt provides a huge number of compounds of this substance can be found in sea water.
Metals of the main subgroups of groups I and II. Hardness of water
In the periodic system of elements, metals are mainly located in the main subgroups of groups I-Ill, as well as in secondary subgroups.
In the IA group, the atoms of the elements on the external energy level have 1 electron in the s 1 state, in the IIA group, the atoms on the external EU have 2 electrons in the s 2 state. These elements are s-elements. In group IIIA, all elements have 3 electrons in the s 2 p 1 state on the external EC. They refer to p-elements.
The IA group includes alkali metals Li, Na, K, Rb, Cs, Fr, the activity of which increases when moving from top to bottom due to an increase in the radius of atoms, the metallic properties increase in the same way as in alkaline earth metals of group IIA Be, Mg, Ca, Sr, Ba, Ra and Group IIIA metals Al, Ga, In, Tl.
Oxides of the R 2 O type are characteristic only for Li, for all other alkali metals, peroxides R 2 O 2 are characteristic, which are strong oxidizing agents.
All metals of these groups form basic oxides and hydroxides, except for Be and Al, which exhibit amphoteric properties.
Physical properties
In the free state, all metals are silvery-white substances. Magnesium and alkaline earth metals are malleable and ductile, rather soft, although harder than alkali metals. Beryllium is notable for its considerable hardness and brittleness, while barium splits upon a sharp blow.
In the crystalline state under normal conditions, beryllium and magnesium have a hexagonal crystal lattice, calcium, strontium have a cubic face-centered crystal lattice, barium has a cubic body-centered crystal lattice with a metallic type chemical bond, which causes their high thermal and electrical conductivity.
Metals have melting and boiling points higher than those of alkali metals, and with an increase in the ordinal number of the element, the melting point of the metal changes nonmonotonically, which is associated with a change in the type of crystal lattice.
Beryllium and magnesium are covered with a strong oxide film and do not change in air. Alkaline earth metals are very active, they are stored in sealed ampoules, under a layer of vaseline oil or kerosene.
Some physical properties of beryllium, magnesium and alkaline earth metals are given in the table.
alkali metals- These are silvery-white substances with a characteristic metallic luster. They quickly tarnish in air due to oxidation. These are soft metals, Na, K, Rb, Cs are similar in softness to wax. They are easily cut with a knife. They are light. Lithium is the lightest metal with a density of 0.5 g/cm3.
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
2Na + S → Na 2 S
2Na + H 2 → 2NaH
3Na + P → Na 3 P
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
Lithium oxide is formed immediately with lithium.
4Li + O 2 \u003d 2Li 2 O, and when oxygen reacts with sodium, sodium peroxide is formed.
2Na + O 2 \u003d Na 2 O 2. When all other metals are burned, superoxides are formed.
K + O 2 \u003d KO 2
By reacting with water, one can clearly see how the activity of these metals in the group changes from top to bottom. Lithium and sodium calmly interact 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
Obtaining 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.
alkaline earth metals
Their name is due to the fact that the hydroxides of these metals are alkalis, and the oxides used to be 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 with non-metals
Ca + Cl 2 → 2CaCl 2
Ca + S → CaS
Ca + H 2 → CaH 2
3Сa + 2P → Сa 3 P 2-
2. Interaction with oxygen
2Сa + 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
Obtaining 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 by aluminothermic process from barium oxide
