JEE Chemistry S Block Elements Notes Set 05

Group 1 and Group 2 Elements: General Features

1. The s-block contains group 1 (alkali metals Li, Na, K, Rb, Cs, Fr) with outer configuration ns¹ and group 2 (alkaline earth metals Be, Mg, Ca, Sr, Ba, Ra) with outer configuration ns². Fr and Ra are radioactive.
2. Alkali metals are named so because their hydroxides are strongly alkaline. Oxides and hydroxides of group 2 are also alkaline in nature, and the oxides occur in the earth's crust.
3. These metals are silvery white, light, soft (group 1) or harder (group 2), and are malleable and ductile with metallic lustre.

Physical Properties and Trends

1. Atomic and ionic radii increase down each group. A group 2 atom is smaller than the group 1 atom in the same period because of its larger nuclear charge. Cations are smaller than their parent atoms.
2. Melting and boiling points of group 1 fall down the group (Li highest) because metallic bonding weakens with size and each atom contributes only one electron. Group 2 metals melt higher because two electrons take part in bonding; Be is highest and the order is irregular (Mg is the lowest of Be, Mg, Ca, Sr, Ba).
3. Ionisation energy falls down the group. The first ionisation energy of group 2 is higher than that of group 1, but the second ionisation energy of group 1 is very high (a noble gas core is broken), whereas it is only moderate for group 2. This is why group 1 shows only +1 and group 2 shows only +2.
4. Electropositive and metallic character increase down the group. Density rises down the group with a break at K (lighter than Na) and Ca (lighter than Mg).
5. Hydration enthalpy of an ion falls as its size rises, so Li+ is hydrated most and its hydrated radius is the largest: hydrated radii \( \ce{Li+} \) > \( \ce{Na+} \) > \( \ce{K+} \) > \( \ce{Rb+} \) > \( \ce{Cs+} \), while ionic mobility and conductance rise from Li+ to Cs+.
6. All the metals are good conductors of heat and electricity because of free outer electrons.
7. Flame test: Li crimson red, Na golden yellow, K violet, Rb red-violet, Cs blue; Ca brick red, Sr crimson, Ba apple green. Be and Mg give no colour because their electrons are too tightly held to be excited in a Bunsen flame.
8. Photoelectric effect: K, Rb and Cs, with large size and low ionisation energy, emit electrons in visible light, and Cs is used in photocells. Be and Mg do not show it.
9. Standard oxidation potential: group 1 metals are strong reducing agents. In aqueous solution Li has the highest oxidation potential (+3.05 V) because of the very high hydration energy of Li+; order Li > Cs, Rb, K > Na. In group 2 the reducing power rises from Be to Ba.
10. Complex formation needs small size, high charge density and vacant orbitals, so it is weak in the s-block. Li+, \( \ce{Be2} \)+ and \( \ce{Mg2} \)+ are the exceptions, for example \( \ce{[BeF4]^2-} \) and chlorophyll (a Mg complex).

Chemical Properties of the Alkali Metals and Alkaline Earth Metals

1. Reactivity rises with falling ionisation energy: Li < Na < K < Rb < Cs for group 1, and Be < Mg < Ca < Sr < Ba for group 2. None occurs free in nature.
2. In dry air, alkali metals form oxides on the surface and tarnish, and are stored in kerosene. In moist air the oxide forms hydroxide and then carbonate. Group 2 metals (except Be) tarnish easily, and in dry air they give oxide and nitride together.
3. With oxygen, Li gives the normal oxide \( \ce{Li2O} \), Na gives mainly the peroxide \( \ce{Na2O2} \), and K, Rb and Cs give superoxides \( \ce{MO2} \) in excess oxygen. Superoxides are paramagnetic and coloured. Stability of the oxide types falls: normal oxide > peroxide > superoxide. In group 2 the normal oxide MO is formed, and only Ca, Sr and Ba also form peroxides.
4. Oxide basicity: BeO is amphoteric, MgO weakly basic, and CaO, SrO and BaO strongly basic; basic strength increases down the group.
5. With hydrogen, alkali metals form ionic hydrides MH, which give \( \ce{H2} \) and the hydroxide with water. Stability of the hydrides falls down the group. In group 2 all except Be form \( \ce{MH2} \); \( \ce{BeH2} \) and \( \ce{MgH2} \) are covalent. \( \ce{BeH2} \) is made from \( \ce{BeCl2} \) and \( \ce{LiAlH4} \).
6. With water, alkali metals give the hydroxide and \( \ce{H2} \), increasing in vigour from Li (slow) through Na (vigorous) and K (burns) to Rb and Cs (explosive). In group 2, Be does not react, Mg reacts with hot water or steam, and Ca, Sr and Ba react with cold water, less vigorously than alkali metals. Hydroxide basicity rises from \( \ce{Be(OH)2} \) to \( \ce{Ba(OH)2} \).
7. Halides: alkali metal halides are ionic, and ionic character rises from LiCl to CsCl. LiCl is covalent because of the polarising Li+, so it dissolves in organic solvents and hydrolyses more. K, Rb and Cs halides add halogen to give polyhalides such as \( \ce{KI3} \). Group 2 halides are ionic except \( \ce{BeCl2} \) and \( \ce{MgCl2} \) (covalent); ionic character rises from \( \ce{BeCl2} \) to \( \ce{BaCl2} \).
8. Nitrides: only Li reacts directly with \( \ce{N2} \) to give \( \ce{Li3N} \), which gives \( \ce{NH3} \) on hydrolysis. All group 2 metals form nitrides \( \ce{M3N2} \) that also give ammonia with water.
9. Solutions in liquid ammonia: all alkali metals, and Ca, Sr and Ba, dissolve to give deep blue conducting solutions containing ammoniated metal ions and ammoniated electrons. The dilute solution is paramagnetic and a strong reducing agent. At high concentration it becomes bronze coloured and diamagnetic due to metal-ion clusters. On standing, or with a catalyst, the metal amide forms and the blue colour fades: \( \ce{2Na + 2NH3 -> 2NaNH2 + H2} \). Li gives the imide \( \ce{Li2NH} \). Be and Mg do not dissolve.
10. Alkali metals form amalgams with mercury and alloys with other metals.

Anomalous Behaviour of Lithium and Beryllium

1. Lithium differs from the rest of group 1 because of its very small size, high polarising power and high ionisation energy. It is harder, forms only the normal oxide, reacts slowly with water, forms a nitride directly, and its LiOH and \( \ce{Li2CO3} \) decompose on heating. Li forms no solid bicarbonate, its \( \ce{Li2SO4} \) forms no alum, and its halides are covalent and soluble in ethanol.
2. Li resembles Mg (diagonal relationship): both are harder than their group members, form nitrides, give only normal oxides, have carbonates and hydroxides that decompose on heating, no solid bicarbonates, and chlorides that are soluble in ethanol, deliquescent and crystallise as hydrates.
3. Beryllium is the hardest of group 2, has the highest melting point and ionisation energy, forms covalent compounds, and dissolves in alkali to give beryllate and hydrogen. BeO and \( \ce{Be(OH)2} \) are amphoteric. \( \ce{Be2C} \) gives methane on hydrolysis, and Be forms complexes readily.
4. Be resembles Al: similar charge-to-radius ratio and electronegativity, passivity in conc. \( \ce{HNO3} \) due to an oxide film, amphoteric oxides and hydroxides, soluble beryllates and aluminates, chlorides with bridged structures and Lewis acid character, and tetrahedral and octahedral complexes such as \( \ce{[BeF4]^2-} \) and \( \ce{[AlF6]^3-} \).

Sodium Chloride, Sodium Hydroxide and Sodium Carbonates

1. Sodium chloride: obtained from sea water by solar evaporation; impurities \( \ce{Na2SO4} \), \( \ce{MgCl2} \) and \( \ce{CaCl2} \) are removed by filtering and by passing HCl gas through the saturated solution to crystallise pure NaCl (common ion effect). Table salt gets damp because of \( \ce{MgCl2} \) and \( \ce{CaCl2} \). It gives a white precipitate of AgCl with \( \ce{AgNO3} \).
2. Sodium hydroxide is made by electrolysis of brine in the diaphragm (Nelson) cell, with \( \ce{H2} \) at the steel cathode and \( \ce{Cl2} \) at the carbon anode, or in the mercury cathode (Castner-Kellner) cell where sodium forms an amalgam that gives NaOH and \( \ce{H2} \) with water.
3. NaOH is a deliquescent white solid and a strong base that absorbs \( \ce{CO2} \) from air to give \( \ce{Na2CO3} \). It reacts with \( \ce{SiO2} \), \( \ce{Al2O3} \), B and Si, and with Al, Zn, Be, Sn and Pb to give hydrogen and the metal oxoanions (aluminate, zincate, beryllate, stannite, plumbite). It does not react with \( \ce{H2} \), \( \ce{N2} \) or C. It precipitates \( \ce{Zn(OH)2} \), which dissolves in excess alkali. Uses: soap, rayon, dyes, paper, drugs and petroleum refining.
4. Sodium carbonate (washing soda, \( \ce{Na2CO3.10H2O} \)) and sodium hydrogencarbonate are made by the Solvay (ammonia-soda) process: \( \ce{CaCO3 -> CaO + CO2} \); \( \ce{NH3 + H2O + CO2 -> NH4HCO3} \); \( \ce{NaCl + NH4HCO3 -> NaHCO3 + NH4Cl} \); \( \ce{2NaHCO3 -> Na2CO3 + CO2 + H2O} \); and \( \ce{NH3} \) is recovered by \( \ce{2NH4Cl + CaO -> CaCl2 + 2NH3 + H2O} \). Calcium chloride is the by-product. \( \ce{KHCO3} \) is too soluble, so \( \ce{K2CO3} \) cannot be made this way.
5. \( \ce{NaHCO3} \) hydrolyses to a weakly alkaline solution, decomposes above 373 K to \( \ce{Na2CO3} \), and reacts with acids (giving \( \ce{CO2} \)) and with bases, so it shows amphoteric behaviour. Uses: baking powder, effervescent drinks, fire extinguishers and antacids.
6. \( \ce{Na2CO3.10H2O} \) effloresces to the monohydrate in air, and its aqueous solution is alkaline by anionic hydrolysis. Uses: glass, caustic soda, soap powders, laundries and softening of water.

Compounds of Calcium and Cement

1. Quick lime CaO: made by heating limestone at about 1070-1270 K. It reacts with water to give slaked lime, combines with \( \ce{SiO2} \) and \( \ce{P4O10} \) (basic oxide), and with carbon at 2273 K to give \( \ce{CaC2} \), which with nitrogen gives nitrolim \( \ce{CaCN2} \). Uses: bleaching powder, cement, glass, carbide, sugar purification, drying \( \ce{NH3} \), furnace lining and soda lime.
2. Slaked lime \( \ce{Ca(OH)2} \): lime water turns milky with \( \ce{CO2} \) owing to \( \ce{CaCO3} \), and the milkiness disappears in excess \( \ce{CO2} \) as soluble \( \ce{Ca(HCO3)2} \) forms, which gives \( \ce{CaCO3} \) on boiling. Dry \( \ce{Ca(OH)2} \) with \( \ce{Cl2} \) below 308 K gives bleaching powder \( \ce{CaOCl2} \). Uses: softening water, sugar and coal gas purification, bleaching powder, mortar, whitewash.
3. Gypsum \( \ce{CaSO4.2H2O} \) gives plaster of Paris (\( \ce{CaSO4} \) with half a water molecule per unit) at about 393 K and anhydrous \( \ce{CaSO4} \) (dead burnt plaster) at about 473 K. Plaster of Paris is a white powder that sets to a hard mass on mixing with water; it is used for casts, statues, moulds and chalk.
4. Portland cement is a mixture of calcium silicates and aluminates: tricalcium silicate, dicalcium silicate (slowest setting), tricalcium aluminate (fastest setting) and tetracalcium aluminoferrite. A typical composition is CaO about 61.5 percent and \( \ce{SiO2} \) about 22.5 percent, with \( \ce{Al2O3} \), \( \ce{Fe2O3} \), MgO, \( \ce{SO3} \), \( \ce{K2O} \) and \( \ce{Na2O} \) in small amounts. Raw materials are limestone, clay and gypsum (added to slow the setting). Setting is hydration and is exothermic. Mortar is cement with sand and water, concrete adds gravel, and reinforced concrete has iron bars.

Biological Importance

1. Na+ and K+ carry nerve signals, regulate water balance and transport sugars and amino acids; K+ is the main cation inside cells and activates enzymes. \( \ce{Mg2} \)+ is the central atom of chlorophyll and a cofactor of ATP-using enzymes. \( \ce{Ca2} \)+ takes part in nerve and muscle function, cell membrane integrity and blood clotting. By abundance in the human body K > Na > Fe > Cu.