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Transcript
• ALKALI EARTH METALS Introduction Properties Earth`s crust abundance Beryllium 1.9x10 Electron configuration o Melting point ( C) o Boiling point ( C) Density Common oxidation num. Physical State Color • • • • • • • • • • • • -4 2 Magnesium Calcium 2.9 5 2 2 [He]2s [Ne]3s [Ar]4s 1287 650 842 • 2469 1.85 1090 1.74 1484 1.55 • 2 solid grey 2 solid white 2 solid White The word “earth” was used for the oxides of group 2A by the ancient scientists. Alkaline earth metals, after the alkali metals, are secondary metals with strong metallic properties. The group 2A elements are less active than those of 1A, but more active than those of group 3A. Except Be, all form ionic compounds. Be forms mostly covalent compounds. The atomic radius increases from top to bottom. The melting point, boiling point and the hardness of the elements decrease top to bottom. Beryllium is the hardest alkaline earth metal Barium is the softest alkaline earth metal. -2 -2 Compounds of CO3 and O are insoluble in water. Malleable and ductile. +2 +2 Excess Mg and Ca in water means that it is hard. 1. Occurrence Minerals of Some Alkaline Earth Metals Beryllium Calcium Beryl, Limestone (marble), Be3Al2(SiO3)6 CaCO3 Gypsum, Magnesium CaSO4.2H2O Dolomite, Fluorite, CaCO3. MgCO3 CaF2 Carnallite, Strontium KCI.MgCI2.6H2O Epsom salt, Celestite, MgSO4.7H2O SrSO4 Strontianite, SrCO3 Barium Barytes, BaSO4 Witherite, BaCO3 Radium All isotopes are radioactive 2. Chemical Properties • They are Covered by their oxides when they are exposed to air. • Burn in air with bright flame. • Alkaline earth metals form MSO4, MCO3, MO, M(OH)2, compounds where M alkaline earth metal. • Activity of the metals increases from Be to Ba. • Reactions of Ca, Sr and Ba are easy with water. Reactions of Be and Mg with water are difficult. +2 Ca(s) + 2H2O(l) → Ca (aq) + 2OH (aq) + H2(g) Mg(s) + 2H2O(l) → Mg(OH)2(s) + H2(g) 2Mg(s) + O2(g) → 2MgO(s) Mg (s) + H2(g) → MgH2(s) Ca (s) + 2HCI(aq) → CaCl2 + H2(g) Be (s) + 2HCI(aq) → BeCI2 + H2(g) All oxides and hydroxides of alkaline earth metals show basic properties. Oxides and its hydroxides of beryllium show amphoteric properties. 3. Compounds • Ionic • Have high melting and boiling points. • Thermally stable. • Less soluble in water than those of 1A metals. • Solubility of salts increase from Be to Ba. Calcium Oxide, CaO • Obtained by heating calcium carbonate (limestone) at 900°C. • CaCO3(s) → CaO(s) + CO2(g) Calcium Carbonate, CaCO3 • It forms by evolving CO2 gas from Ca(OH)2 solution Calcium Carbide, CaC2 • Small gray solid pieces • Used widely in industry • Prepared by the reaction of calcium oxide (lime) and coke. CaO(s) + 3C(s) ↔ CaC2(s) + CO (g) CaC2(s) + 2H2O(l) → C2H2(g) + Ca(OH)2(s) Calcium Sulfate, CaSO4 CaSO4 : Gypsum CaSO4.2H2O : Alabaster CaSO4.1/2H2O : Plaster of Paris • Gypsum is a white substance and used for fabrication into wallboard. Uses Beryllium • Very light element • To prepare some special alloys. • 2% Be in Cu alloy in springs. • Making glass for X-ray tubes. • In the aerospace industry. Magnesium • Light metal • Production of planes, missiles • Some light household items. • Flash in photography Magnesium • To increase the hardness, durability and resistance to corrosion of alloys. • • Additive in rocket fuels and signal rockets. Mg(OH)2 is used as a stomach antacid. Calcium Compounds • Construction materials, bleaching agents. • CaO in the production of some chemicals, steel, glass, paper and sugar. • CaC2is used to produce acetylene gas for welding. • In fertilizers and toothpaste Barium and Its Compounds • Ba is used as a gas absorbent in vacuum tubes. • The alloys of Ba with Ni donate electrons so used in vacuum tubes and in ignitors and spark plugs. • Barium nitrate and barium chlorate are used to obtain green light in fireworks. Barium and Its Compounds • BaSO4 is used to take the photos of stomach and intestines. • Ba does not allow X-rays to pass like all elements with large atomic number. +2 • The Ba ion is toxic, but since the solubility of BaSO4 in water is very low, its poisonous effect disappears Strontium • The salts of strontium in signal rockets and fireworks to produce red light. • Nuclear batteries. Radium • In self luminous paints and in radiotherapy. • As a neutron source. • On luminous watch dials.