\nGroup 1 of the periodic table consists of the elements: Lithium, Sodium, Potassium, Rubidium, Caesium and Francium. They are collectively known as alkali metals.<\/p>\n
Group 2 consists of Beryllium, Mgnesium, Calcium, Strontium, Barium and Radium. These elements except of beryllium are known as the alkaline earth metals. The general electronic configuration of s-block elements is [noble gasjns1 for alkali metals and [noble gas] ns\u00b2 for alkaline earth metals. The first elements of Group 1 and Group 2 respectively exhibit diagonal similarity, which is commonly referred to as diagonal relationship in the periodic table. The diagonal relationship is due to the similarity in ionic sizes and \/or charge\/radius ratio of the elements.<\/p>\n
Group 1 Elements: Alkali Metals<\/span><\/p>\n 1) Electronic Configuration: 2) Atomic And Ionic Radii: 3) Hydration Enthalpy: Physical Properties<\/span> Chemical Properties<\/span> They also react with proton donors such as alcohol, gaseous ammonia and alkynes.AII the alkali metal hydrides are ionic solids with high melting points. The alkali metals readily react vigorously with halogens to form ionic halides, M+<\/sup>X–<\/sup>. However, lithium halides are somewhat covalent. It is because of the high polarisation capability of lithium-ion. The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful. The alkali metals dissolve in liquid ammonia giving deep blue solutions. The solutions are paramagnetic and on standing slowly liberate hydrogen.<\/p>\n General Characteristics Of The Compounds Of The Alkali Metals<\/span><\/p>\n Oxides And Hydroxides<\/span> The normal oxides dissolve in water to form hydroxides (MOH) which are strong bases. However, LiOH is only slightly soluble in water and it decomposes on heating. The peroxides and superoxides also dis-solve in water to form basic hydroxides. The basic character of alkali metal hydroxides increases down the group.<\/p>\n Halides<\/span> Anomalous Properties Of Lithium<\/span> As a result, there is increased covalent character of lithium compounds which is responsible for their solubility in organic solvents.<\/p>\n Points Of Similarities Between Lithium And Magnesium<\/span> Some Important Compounds Of Sodium Sodium Carbonate (Washing Soda), Na2<\/sub>CO3<\/sub>.10H2<\/sub>O<\/span> In this process, NH3<\/sub> is recovered when the solution containing NH4<\/sub>Cl is treated with Ca(OH)2<\/sub>. On heating washing soda becomes monohydrate and then completely anhydrous i.e., soda ash.<\/p>\n Sodium Chloride, NaCl<\/span> Uses:<\/p>\n Sodium Hydroxide (Caustic Soda), NaOH<\/span> Uses: Biological Importance Of Sodium And Potassium<\/span> Group 2 Elements: Alkaline Earth Metals<\/span> 2) Atomic And Ionic Radii: 3) Hydration Enthalpy: Physical Properties<\/span> Chemical Properties<\/span> Reactivity towards hydrogen<\/span> Reactivity towards acids: General Characteristics Of Compounds Of The Alkaline Earth Metals<\/span> BeO is amphoteric in character, while the oxides of the rest of the elements in group 2 are basic. The oxides of Ca, Sr and Ba react with water to form their corresponding hydroxides.<\/p>\n The hydroxides of alkaline earth metals are bases except Li(OH)2<\/sub> which is amphoteric. The basic strength increases from Mg(OH)2<\/sub> to Ba(OH)2<\/sub>. The solubility and thermal stability of hydroxides increase downward in the group. Be(OH)2<\/sub> and Mg(OH)2<\/sub> are almost insoluble. Ca(OH)2<\/sub> is sparingly soluble, while Sr(OH)2<\/sub> and Ba(OH)2<\/sub> are increasingly more soluble.<\/p>\n ii) Halides: Group 2 metals directly combine with halogen to form divalent halides of the formula<\/p>\n The s-Block Elements<\/span> iii) Salts of Oxoacids: Anomalous Behaviour Of Beryllium<\/span> Diagonal Relationship Between Beryllium And Aluminium<\/span> Some Important Compounds Of Calcium<\/span> Calcium Oxide Or Quick Lime, CaO<\/span> Uses:<\/p>\n Calcium Hydroxide (Slaked Lime), Ca(OH)2<\/sub><\/span> Uses:<\/p>\n Calcium Carbonate, CaCO2<\/sub><\/span> Uses:<\/p>\n Calcium Sulphate (Plaster Of Paris), CaSO4<\/sub>.\u00bdH2<\/sub>O<\/span> Used:<\/p>\n
\nAll the alkali metals have one valence electron, ns\u00b9 outside the noble gas core. The loosely held s-electron readily lose electron to give monovalent M+<\/sup> ions.<\/p>\n
\nThe atomic and ionic radii of alkali metals increase on moving down the group. Hence, ionization enthalpies of the alkali metals are considerably low and decrease down the group.<\/p>\n
\nThe hydration enthalpies of alkali metal ions decrease with increase in ionic sizes. Li+<\/sup> > Na+<\/sup> > K+<\/sup> > Rb+<\/sup> > Cs+<\/sup> Li+<\/sup> has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., LiCl- 2H2<\/sub>O<\/p>\n
\nWhen heat is supplied to alkali metal or its salt the electrons are excited to higher energy levels. As these electrons return to their original level; radiations are emitted which fall in the visible region of electromagnetic spectrum. Thus they appear coloured. Li imparts crimson red colour, K gives violet colour and Na gives golden yellow colour to the flame.<\/p>\n
\nThe reactivity of these metals increases with their size. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide O2<\/sub> ion is stable only in the presence of large cations such as K, Rb, Cs.
\n4Li + O2<\/sub> \u2192 2LizO(oxide)
\n2Na + O2<\/sub> \u2192 Na2<\/sub>O2<\/sub> (peroxide)
\nM + O2<\/sub> \u2192 MO2<\/sub>(superoxide)
\n(M=K, Rb, Cs)
\nBecause of their high reactivity towards air and water, they are normally kept in kerosene oil.lt may be noted that although lithium has most negative E\u00b0 value.
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<\/p>\n
\n2M + H2<\/sub> \u2192 2M+<\/sup>H–<\/sup>.<\/p>\n
\nReactivity of alkali metals with oxygen increases down the group. Lithium, when heated in air, forms the normal oxide (Li2<\/sub>O) while sodium forms the per-oxide (Na2<\/sub>O2<\/sub>). Potassium, Rubidium and caesium form superoxides (MO2<\/sub>).
\n4Li + O2<\/sub> \u2192 2Li2<\/sub>O; 2Na+ O2<\/sub> \u2192 Na2<\/sub>O2<\/sub>; K + O2<\/sub> \u2192 KO2<\/sub><\/p>\n
\nAlkali metals react vigorously with halogens to form metal halides of the general formula MX. 2M+X2<\/sub> \u2192 2MX X=F, Cl, Br or l and M= alkali metal Reactivity of alkali metal towards halogen increases from Li to Cs. Halides of alkali metals are ionic compounds readily soluble in water. But LiF is almost insoluble due to high lattice energy.<\/p>\n
\nThe anomalous behaviour of lithium is due to the :<\/p>\n\n
\nThe similarity between lithium and magnesium is particularly striking and arises because of their similar sizes: atomic radii, Li = 152 pm, Mg= 160 pm; ionic radii: Li+<\/sup> = 76 pm, Mg2+<\/sup> = 72 pm. The main points of similarity are:<\/p>\n\n
\nSodium carbonate is generally prepared by Solvay Process.
\nThe equations for the complete process may be written as:
\n2NH3<\/sub> + H2<\/sub>O + CO2<\/sub> \u2192 (NH4<\/sub>)2<\/sub>CO3<\/sub>
\n(NH4<\/sub>)2<\/sub>CO3<\/sub> + H2<\/sub>O + CO2<\/sub> \u2192 2NH4<\/sub>HCO3<\/sub>
\nNH4<\/sub>HCO3<\/sub> +NaCl \u2192 NH4<\/sub>Cl + NaHCO3<\/sub>
\n2NaHCO3<\/sub> \u2192 Na2<\/sub>CO3<\/sub> +CO2<\/sub> +H2<\/sub>O<\/p>\n
\nThe most abundant source of sodium chloride is seawater. Common salt is generally obtained by evaporation of seawater. Crude sodium chloride, generally obtained by crystallisation of brine solution, contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. Calcium chloride, CaCl2<\/sub>, and magnesium chloride, MgCl2<\/sub> are impurities because they are deliquescent (absorb moisture easily from the atmosphere). To obtain pure sodium chloride, the crude salt is dissolved in minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas. Crystals of pure sodium chloride separate out. Calcium and magnesium chloride, being more soluble than sodium chloride, remain in solution.<\/p>\n\n
\nSodium hydroxide is generally prepared commercially by the electrolysis of sodium chloride in Castner-Kellner cell. A brine solution is electrolysed using a mercury cathode and a carbon anode.
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\nThe amalgam is treated with water to give sodium hydroxide and hydrogen gas.
\n2 Na – amalgam + 2H2<\/sub>O \u2192 2NaOH + 2Hg +H2
\n<\/sub>The sodium hydroxide solution at the surface reacts with the C02 in the atmosphere to form Na2<\/sub>CO3<\/sub>.<\/p>\n
\nIt is used in (i)the manufacture of soap, paper, artificial silk and a number of chemicals,(ii) in petroleum refining, (iii) in the purification of bauxite, (iv) in the textile industries for mercerising cotton fabrics and (v) for the preparation of pure fats and oils.<\/p>\n
\nSodium ions participate in the transmission of nerve signals. The concentration gradient of Na+<\/sub> and K+<\/sub> demonstrates that-a discriminatory mechanism called sodium-potassium pump, operates across the cell membranes.<\/p>\n
\nThe group 2 elements (except beryllium) are known as alkaline earth metals. The first element beryllium differs from the rest of the members and shows diagonal relationship to aluminium.
\n1) Electronic Configuration:
\nThese elements have two electrons in the s-orbital of the valence shell. Their general electronic configuration may be represented as [noble gas] ns\u00b2.<\/p>\n
\nWithin the group, the atomic and ionic radii increase with increase in atomic number due to the increased nuclear charge in these elements. They have low ionisation enthalpy and it decreases down the group with increase in size.<\/p>\n
\nHydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group. Be2+<\/sup> > Mg2+<\/sup> > Ca2+<\/sup> > Sr2+<\/sup> > Ba2+<\/sup> The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.<\/p>\n
\nCalcium, Strontium and Barium impart characteristic brick red, crimson and apple green colours respectively to the flame. Inflame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light. The electrons in Be and Mg are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame.<\/p>\n
\nThe alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.
\nReactivity towards air and water Beryllium and Magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. However, powdered beryllium burns brilliantly on ignition in air to give BeO and Be3<\/sub>N2<\/sub>.
\nReactivity towards halogen
\nM + X2<\/sub> \u2192 MX2<\/sub> (X = F, Cl, Br, I)<\/p>\n
\nAll the elements except beryllium form their hydrides, MH2<\/sub>.BeH2<\/sub>, however, can be prepared by the reaction of BeCl2<\/sub> with LiAlH4<\/sub>.
\n2BeCl2<\/sub> +LiAlH4<\/sub> \u2192 2BeH2<\/sub> +LiCl + AlCl3<\/sub><\/p>\n
\nThe alkaline earth metals readily react with acids liberating dihydrogen.<\/p>\n
\ni) Oxides and Hydroxides: Alkaline earth metals burn in air or oxygen to form their oxides. (Oxides are also prepared by the thermal decomposition of their carbonates). Be, Mg and Ca form monoxides (MO). The tendency to form peroxide increases as the size of the metal ion increases. Strontium and barium form peroxides (MO2<\/sub>)
\n2M + O2<\/sub> \u2192 MO (M = Be, Mg or Ca)
\nM+O2<\/sub> \u2192 MO2<\/sub> (M = Sr or Ba)<\/p>\n
\nMX2<\/sub> where X is the halogen. The metal halides are also formed by the action of halogen acids on metals, their oxides, carbonates and hydroxides. BeCl2<\/sub> is, however, prepared by passing Cl2<\/sub> over a hot mixture of BeO and coke.
\nIn contrast to the halides of other alkaline earth metals, beryllium halides are covalent. In the solid-state BeCl2<\/sub> has a polymeric chain structure involving Be-CI-Be bridges. The anhydrous halides are hygroscopic and form hydrates such as MgCl2<\/sub>.6H2<\/sub>O, CaCl2<\/sub>.6H2<\/sub>O etc. Due to this reason, anhydrous calcium chloride is widely used as a dehydrating agent. Fluorides are relatively less soluble due to high lattice energies,<\/p>\n
\nThe alkaline earth metals also form salts of oxoacids. Some of these are : Carbonates, Sulphates and Nitrates.<\/p>\n
\nBeryllium differs from the rest element in many of its properties. These are<\/p>\n\n
\nThe ionic radius of Be2+<\/sup> is estimated to be same as that of the Al3+<\/sup> ion. Hence Be resembles Al in some ways. Some of the similarities are:<\/p>\n\n
\nImportant compounds of calcium and their preparations are given below.<\/p>\n
\nIt is prepared by the following reaction.
\nCaCO3<\/sub> \\(\\rightleftharpoons \\) Ca0 + CO2<\/sub>
\nCO2<\/sub> is removed as soon as it is produced to enable the reaction to proceed to completion.
\nCaO + H2<\/sub>O \u2192 Ca(OH)2<\/sub>
\nThis process is called slaking of lime. CaO is a basis oxide.<\/p>\n\n
\nIt is prepared by adding water to CaO. The aqueous solution of Ca(OH)2<\/sub> is known as lime water and the suspension of slaked lime is known as milk of lime. When CO2<\/sub> is passed through lime water it turns milky due to the formation of CaCO3<\/sub>
\nCa(OH)2<\/sub> + CO2<\/sub> \u2192 CaCO3<\/sub> +H2<\/sub>O<\/p>\n\n
\nIt occurs in limestone, chalk, marble etc.
\nIt can be prepared by the following reactions.
\nCa(OH)2<\/sub> + CO2<\/sub> \u2192 CaCO3<\/sub> + H2<\/sub>O
\nCaCl2<\/sub> + Na2<\/sub>CO3<\/sub> \u2192 CaCO3<\/sub> + 2NaCl
\nCaCO3<\/sub> reacts with dilute acids to liberate carbon dioxide.<\/p>\n\n
\nIt is obtained by heating gypsum (CaSO2<\/sub>.2H2<\/sub>O)
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\nAbove 393K anhydrous calcium sulphate is formed. This is known as \u2018dead burnt plaster\u2019<\/p>\n