THE s- BLOCK ELEMENTS General electronic configuration- [ noble gas] ns 1-2 GROUP 1 ELEMENTS : ALKALI METALS General electronic configuration- [ noble gas] ns 1 Members- Li, Na, K, Rb, Cs, Fr Atomic and Ionic Radii- Largest in a particular period Monovalent ions ( M + ) are smaller than the parent atom. On moving down the group, it increases.( increase in no. of shells) Ionisation Enthalpy- Low ( loss of one electron gives stable configuration) On moving down the group, it decreases ( effect of increasing size outweighs the increasing nuclear charge) Hydration Enthalpy- Li + > Na + > K + > Rb + > Cs + It decreases with increase in ionic sizes because as the size increases, charge density decreases. So, Lithium salts are mostly hydrated. E.g. LiCl.2H 2 O PHYSICAL PROPERTIES- Due to large size, they have low density which increases down the group. Low melting and boiling points due to weak metallic bonding ( presence of single valence electron)
Alkali metals and their salts impart characteristic colour in the flame ( Heat from the flame excites the outermost electron to a higher energy level. When the excited electron comes back to the ground state, it emits radiation in the visible range) Li - Crimson red, Na- Yellow, K- Violet, Rb Red violet,cs- Blue CHEMICAL PROPERTIES- Highly reactive due to large size and low Ionisation Enthalpy Reactivity Towards Air- They burn in oxygen forming oxides. Li forms monoxide, Na forms peroxide, the other metals form superoxides. ( Superoxide ion is stable with large cations only such as K, Rb, Cs. It is so because large anions are stabilized by large cations through lattice energy effects.) 4Li + O 2 2Li 2 O 2Na + O 2 Na 2 O 2 K + O 2 KO 2 Reactivity Towards Water- Except Li, other members of the group react explosively with water. 2M + 2H 2 O 2M + + 2OH - + H 2 Although Li has most negative E 0 value, its reaction with water is less vigorous than Na ( due to small size and high hydration enthalpy of Li) Reactivity Towards Dihydrogen- 2M + H 2 2MH Metal hydrides are ionic solids.
Reactivity Towards Halogens- 2M + X 2 2MX Alkali metal halides are ionic in nature except LiX ( due to small size, charge density is high and so have high polarization capability. LiI is most covalent alkali metal halide) Reducing Nature- Strong reducing agents. The standard electrode potential depends on sublimation enthalpy, ionization enthalpy and hydration enthalpy. M (s) M (g) sublimation enthalpy M (g) M + (g) +e - ionization enthalpy M + (g) + H 2 O M + (aq) hydration enthalpy On moving down the group, E o value becomes more and more negative from Na to Cs. Li has most negative E 0 value and hence most powerful reducing agent. (Due to small size, Li has the highest hydration enthalpy which overcomes the high ionization enthalpy and so has high negative E o.) Solutions in liquid ammonia- The alkali metals dissolve in liquid ammonia giving deep blue conducting solutions. M + (x+y)nh 3 [M(NH 3 )x] + + [e(nh 3 )y] - The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region. ( absorbs from red region, so transmitted light is blue in colour)
The solutions are paramagnetic and on standing slowly liberate H 2 and so amide is formed. M + (am) + e - + NH 3 (l) MNH 2(am) + ½ H 2 (g) In conc. Solution, the blue colour changes to bronze colour and becomes diamagnetic. GENERAL CHARACTERISTICS OF THE COMPOUNDS OF THE ALKALI METALS I OXIDES AND HYDROXIDES They burn in oxygen forming oxides. Li forms monoxide, Na forms peroxide, the other metals form superoxides. ( Superoxide ion is stable with large cations only such as K, Rb, Cs. It is so because large anions are stabilized by large cations through lattice energy effects.) 4Li + O 2 2Li 2 O [M 2 O] 2Na + O 2 Na 2 O 2 [M 2 O 2 ] K + O 2 KO 2 [MO 2 ] Superoxides are paramagnetic. ( O 2 - is paramagnetic because of one unpaired electron in π* 2p molecular orbital.) Oxides are hydrolysed by water to form hydroxides( white crystalline solids) M 2 O + H 2 O 2M + + 2OH - M 2 O 2 + 2H 2 O 2M + + 2OH - + H 2 O 2 2MO 2 + 2H 2 O 2M + + 2OH - + H 2 O 2 + O 2 The alkali metal hydroxides are the strongest of all bases and dissolve freely in water with evolution of heat.
II HALIDES They are prepared by reaction of HX with oxides, hydroxides or carbonates of alkali metals. M 2 O + 2HX MOH + HX M 2 CO 3 + 2HX 2MX + H 2 O MX + H 2 O 2MX + CO 2 +H 2 O They have high negative values of Δ f H o. On moving down the group, Δ f H o value for fluorides become less negative. On moving down the group, Δ f H o value for chlorides, bromides and iodides become more negative. For a given metal, Δ f H o value always become less negative from fluoride to iodide. Trend of m.p. : For the same alkali metal ( due to decrease in lattice enthalpy) the order is : F - > Cl - > Br - > I - For the same halide ion, m.p. decreases from Na to Cs. But m.p. of Lithium halides is low because they are covalent in nature. Solubility in water- Low solubility of LiF due to high lattice enthalpy Low solubility of CsI due to smaller hydration enthalpy of ions. III SALTS OF OXOACIDS- Oxoacids are those in which the acidic proton is on a hydroxyl group with an oxo group attached to the same atom e.g. H 2 CO 3, H 2 SO 4. The alkali metals form salts with all the oxoacids.
Carbonates- ( M 2 CO 3 ) They are stable to heat but Li 2 CO 3 is not so stable to heat. Li 2 CO 3 Li 2 O + CO 2 Hydrogen carbonates ( MHCO 3 ) They are also stable to heat. Li does not form solid MHCO 3. Solubility in water- All carbonates and hydrogencarbonates are soluble in water which increases down the group. POINTS OF DIFFERENCE BETWEEN Lithium and other Alkali Metals- 1. Lithium salts are mostly hydrated. E.g. LiCl.2H 2 O whereas other metal chlorides do not form hydrates. 2. Li does not form solid MHCO 3 whereas other alkali metals form. 3. Li does not form ethynide on reaction with ethyne whereas other alkali metals form. 4. Li is least reactive but the strongest reducing agents among all the alkali metals. Points of Similarities between Lithium and Magnesium 1. They react slowly with water. 2. They form nitrides by direct combination with nitrogen. 3. Their carbonates decompose easily on heating. 4. They do not form solid hydrogencarbonates. 5. Their chlorides are deliquescent. LiCl.2H 2 O, MgCl 2.8H 2 O
GROUP 2 ELEMENTS : ALKALINE EARTH METALS General electronic configuration- [ noble gas] ns 2 Members- Be 4, Mg 12, Ca 20, Sr 38, Ba 56, Ra 88 Atomic and Ionic Radii- On moving down the group, it increases.( increase in no. of shells) Ionisation Enthalpy- Low ( due to large size of atoms) IE 1 of alkaline earth metals are higher than IE 1 of alkali metals. IE 2 of alkaline earth metals are smaller than IE 2 alkali metals. On moving down the group, it decreases ( effect of increasing size outweighs the increasing nuclear charge) Hydration Enthalpy- Be +2 > Mg +2 > Ca +2 > Sr +2 > Ba +2 It decreases with increase in ionic sizes because as the size increases, charge density decreases. The hydration enthalpies of Gp 2 metal ions are larger than those of Gp.1 ions. So, they are hydrated. E.g. MgCl 2.8H 2 O, CaCl 2.6H 2 O, SrCl 2.6H 2 O, BaCl 2.2H 2 O PHYSICAL PROPERTIES- On moving down the group electropositive character increases. Higher melting and boiling points than the corresponding alkali metals due to smaller sizes.( but not systematic trend) Ca, Sr, Ba metals and their salts impart characteristic colour in the flame.
Ca- Brick red, Sr - Crimson red, Ba- Grassy green or Apple green Be and Mg do not impart any colour their electrons are too strongly bound to get excited by flame) CHEMICAL PROPERTIES- Less reactive than alkali metals Reactivity Towards Air- Be and Mg are kinetically inert to oxygen. Powdered Be gives BeO and Be 3 N 2, when burnt in air. Mg, Ca, Sr and Ba give oxides and nitrides. E.g. MgO and Mg 3 N 2. Reactivity Towards Water- Be and Mg are kinetically inert to water. Ca, Sr and Ba react with wate rand form hydroxides. Reactivity Towards Dihydrogen- All the elements except Be form hydrides on heating. M + H 2 MH 2 Reactivity Towards Halogens- M + X 2 MX 2 Reducing Nature- Strong reducing agents. Their reducing power is less than the corresponding alkali meatls. Solutions in liquid ammonia-
The alkali metals dissolve in liquid ammonia giving deep blue black solutions. M + (x+y)nh 3 [M(NH 3 )x] +2 + 2[e(NH 3 )y] - The blue blackcolour of the solution is due to the ammoniated electron which absorbs energy in the visible region GENERAL CHARACTERISTICS OF THE COMPOUNDS OF THE ALKALINE EARTH METALS I OXIDES AND HYDROXIDES They burn in oxygen forming monoxides. BeO is covalent and amphoteric. All other oxides are ionic and basic. Oxides react with water to form basic hydroxides except Be which forms amphoteric hydroxide. MO + H 2 O M(OH) 2 These hydroxides are sparingly soluble in water. II HALIDES Be halides are covalent while other halides are ionic in nature. In the vapour phase BeCl2 forms dimer which dissociates at higher temperature. The tendency to form hydrates gradually decreases on going down the group e.g. MgCl 2.8H 2 O, CaCl 2.6H 2 O, SrCl 2.6H 2 O, BaCl 2.2H 2 O
III SALTS OF OXOACIDS- Oxoacids are those in which the acidic proton is on a hydroxyl group with an oxo group attached to the same atom e.g. H 2 CO 3, H 2 SO 4. The alkaline earth metals form salts of oxoacids. Carbonates- ( MCO 3 ) They decompose on heating to give oxide and CO 2. MCO 3 MO + CO 2 They are insoluble in water. Sulphates : MSO 4 White solids Stable to heat BeSO 4 and MgSO 4 are readily soluble in water ; The solubility decreases from CaSO 4 to BaSO 4.( high hydration enthalpies of Be 2+ and Mg 2+ overcomes the lattice enthalpy) Nitrates : M(NO 3 ) 2 On moving down the group, tendency to form hydrates decreases with increasing size and decreasing hydration enthalpy. E.g. Mg(NO 3 ) 2 crystallises with 6 H 2 O molecules whereas Ba(NO 3 ) 2 is anhydrous. They decompose on heating 2 M(NO 3 ) 2 2MO + 4NO 2 + O 2 ANOMALOUS BEHAVIOUR OF BERYLLIUM 1. Due to small size and high electronegativity it forms compounds which are covalent and get easily hydrolysed. 2. Max. covalency is 4 due to absence of d orbitals.
3.Oxide and hydroxide of Be are amphoteric in nature. DIAGONAL RELATIONSHIP BETWEEN BERYLLIUM AND ALUMINIUM 1. Their chlorides form dimer. 2. Their chlorides are strong Lewis acids and are used as Friedel Craft catalysts. 3. They are not readily attacked by acids because their oxide form a protective film on the surface. 4. They have strong tendency to form complexes. BeF 4 2-, AlF 6 3-