JEE Chemistry S Block Elements Notes Set 07

General Features and Flame Test

The s-block consists of groups 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) and 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra), in which the last electron enters an s-orbital: ns\(^1\) and ns\(^2\). In flame tests the heat excites the outer electron to a higher level and light of a characteristic colour is emitted when it falls back.
1. Li crimson red, Na golden yellow, K violet (lilac), Rb red violet, Cs blue.
2. Ca brick red, Sr crimson red, Ba apple green.
3. Be and Mg give no colour, because their electrons are too tightly held to be excited by a Bunsen flame.

Periodic Trends

1. Group 1 forms only M\(^+\) because removing a second electron means breaking a noble gas core. Group 2 forms M\(^{2+}\) because the extra lattice or hydration energy outweighs the second ionisation energy.
2. The ionic mobility in water runs opposite to ionic size: the small Li\(^+\) is the most heavily hydrated and moves slowest, Li\(^+\) < Na\(^+\) < K\(^+\) < Rb\(^+\) < Cs\(^+\).
3. In water Li has the most negative reduction potential (-3.04 V) because of its very high hydration energy, so it is the strongest reducing agent in aqueous solution even though its ionisation energy is the highest. In the gas phase the order is reversed (Cs strongest).
4. Low ionisation energy gives the alkali metals the photoelectric effect: Cs and K are used in photocells.

Reaction with Air, Water, Hydrogen and Halogens

Oxygen. The main product of burning in excess air depends on the cation size: Li gives the oxide \( \ce{Li2O} \); Na gives the peroxide \( \ce{Na2O2} \); K, Rb, Cs give the superoxides \( \ce{MO2} \). Larger cations stabilise larger anions through lattice energy. Peroxides are diamagnetic, superoxides are paramagnetic (one unpaired electron in a \( \pi^* \) orbital). Peroxides and superoxides give \( \ce{H2O2} \) with water (and \( \ce{O2} \) from superoxides), which makes them bleaching agents; \( \ce{KO2} \) is used to absorb CO\(_2\) and supply oxygen in submarines. Among group 2, Be and Mg are protected by an oxide film; Mg burns brilliantly in air giving MgO and \( \ce{Mg3N2} \); Ba forms \( \ce{BaO2} \).
Nitrogen. Only lithium combines directly at room temperature giving \( \ce{Li3N} \); group 2 metals form nitrides on heating (\( \ce{Mg3N2}, \ce{Ca3N2} \)). Nitrides hydrolyse to ammonia: \( \ce{Mg3N2 + 6H2O -> 3Mg(OH)2 + 2NH3} \).
Water. Alkali metals give \( \ce{2M + 2H2O -> 2MOH + H2} \) with violence increasing down the group. Lithium is slower than sodium although its potential is more negative: this is a matter of kinetics (it does not melt, and its small ion is slow). Ca, Sr and Ba react readily with cold water, Mg only with hot water or steam, Be not at all.
Hydrogen. The metals form ionic hydrides MH and MH\(_2\) (stable solids that react with water giving H\(_2\)); BeH\(_2\) and MgH\(_2\) are covalent and polymeric. Hydrides form on heating the metal in hydrogen; \( \ce{BeH2} \) is made from \( \ce{BeCl2} \) and \( \ce{LiAlH4} \).
Halogens. All form halides. Group 1 halides are ionic and soluble (LiX has some covalent character because of the polarising power of Li\(^+\)). Among group 2, \( \ce{BeCl2} \) is essentially covalent: a chain polymer in the solid, a chloro-bridged dimer in the vapour, and the linear monomer at about 1200 K.

Oxides, Hydroxides and Amphoteric Beryllium

All alkali metal oxides are strongly basic and dissolve in water to give hydroxides. The basic strength and solubility of the hydroxides increase down both groups (LiOH < NaOH < KOH < RbOH < CsOH; Be(OH)\(_2\) < Mg(OH)\(_2\) < Ca(OH)\(_2\) < Sr(OH)\(_2\) < Ba(OH)\(_2\)). BeO and Be(OH)\(_2\) are amphoteric, dissolving in strong alkali: \( \ce{Be(OH)2 + 2NaOH -> Na2[Be(OH)4]} \) (compare Al). BeO and MgO have very high melting points and are used as refractories.
Halogens disproportionate in alkali: cold dilute NaOH with \( \ce{Cl2} \) gives NaCl + NaOCl; hot concentrated NaOH gives NaCl + \( \ce{NaClO3} \). White phosphorus and sulphur also disproportionate in hot alkali (\( \ce{P4} \) gives \( \ce{PH3} \) and \( \ce{NaH2PO2} \); S gives \( \ce{Na2S} \) and \( \ce{Na2S2O3} \)).

Carbonates, Sulphates and Nitrates

1. Thermal stability of carbonates increases down each group, since larger cations polarise the carbonate ion less. Li\(_2\)CO\(_3\) decomposes to Li\(_2\)O + CO\(_2\); the other alkali carbonates do not decompose at ordinary flame temperatures. Group 2: BeCO\(_3\) < MgCO\(_3\) < CaCO\(_3\) < SrCO\(_3\) < BaCO\(_3\).
2. Bicarbonates: only the alkali metals (except Li) form solid bicarbonates; they decompose on heating, \( \ce{2MHCO3 -> M2CO3 + H2O + CO2} \). Group 2 bicarbonates exist only in solution. NaHCO3 is much less soluble than KHCO3.
3. Solubility of group 2 carbonates and sulphates decreases down the group. BeSO\(_4\) and MgSO\(_4\) are freely soluble because the hydration energy of the small ions exceeds the lattice energy; BaSO\(_4\) is almost insoluble (this is why barium is detected, and barium poisoning is treated with sodium sulphate).
4. Nitrates: group 1 nitrates (except Li) give the nitrite and O\(_2\); LiNO\(_3\) and all group 2 nitrates give the oxide + NO\(_2\) + O\(_2\). The lithium behaviour is another example of its small size and strong polarising power.
5. Fluorides and hydroxides of group 2 follow the trend in solubility: fluorides decrease down the group, hydroxides increase.

Solutions in Liquid Ammonia

Alkali metals dissolve in liquid ammonia giving blue solutions that conduct electricity: \( \ce{M + (x+y)NH3 -> [M(NH3)_x]^+ + [e(NH3)_y]^-} \). The ammoniated electron gives the blue colour and the strong reducing power. Dilute solutions are paramagnetic (free electrons); in concentrated solutions (above about 3 M) the electrons pair up, the solution becomes diamagnetic and bronze coloured. On standing, or in the presence of a catalyst, the solution slowly gives the amide and hydrogen: \( \ce{2M + 2NH3 -> 2MNH2 + H2} \).

Anomalous Behaviour of Lithium and Beryllium

The first member of each group differs from the rest because of its very small size and high polarising power. Lithium resembles magnesium (diagonal relationship): both form normal oxides, nitrides directly with N\(_2\), have carbonates that decompose on heating, soluble chlorides that are deliquescent, no solid bicarbonates, and nitrates that decompose to oxide. Beryllium resembles aluminium: amphoteric oxide and hydroxide, covalent chloride that is hydrolysed in water, passivation by nitric acid, and complex formation such as \( \ce{[BeF4]^{2-}} \). Its covalency is limited to 4 because it has no d-orbitals, whereas Al can reach 6.

Important Compounds of Sodium and Calcium

1. NaOH is made by electrolysis of brine in the Castner-Kellner mercury cathode cell: sodium amalgam forms at the cathode and is decomposed by water to NaOH and H\(_2\), with Cl\(_2\) at the anode.
2. Na\(_2\)CO\(_3\) (washing soda is Na\(_2\)CO\(_3\).10H\(_2\)O) is made by the Solvay process: \( \ce{NH3 + CO2 + H2O -> NH4HCO3} \); \( \ce{NH4HCO3 + NaCl -> NaHCO3 + NH4Cl} \); the bicarbonate is heated to give the carbonate, CO\(_2\) is recycled from limestone, and NH\(_3\) is recovered from \( \ce{NH4Cl + Ca(OH)2} \) with CaCl\(_2\) as the by-product. K\(_2\)CO\(_3\) cannot be made this way because KHCO\(_3\) is too soluble. Washing soda loses water on exposure to air (efflorescence).
3. NaHCO3 is used in fire extinguishers and in baking powder (with a solid acid such as calcium dihydrogen phosphate).
4. CaO (quicklime) from heating limestone; with water it gives slaked lime Ca(OH)\(_2\) (milk of lime is a suspension, lime water is the clear solution, which turns milky with CO\(_2\) and clear again with excess CO\(_2\)). Ca(OH)\(_2\) with Cl\(_2\) gives bleaching powder, CaOCl\(_2\).
5. Gypsum CaSO\(_4\).2H\(_2\)O on heating at 393 K gives plaster of Paris, 2CaSO\(_4\).H\(_2\)O (that is, CaSO\(_4\).\(\frac12\)H\(_2\)O), which sets back to gypsum on mixing with water.
6. Portland cement is a mixture of calcium silicates and aluminates; gypsum is added to retard setting.
7. Carbides: CaC\(_2\) gives ethyne with water; heated with N\(_2\) it gives CaNCN (calcium cyanamide), which with steam gives NH\(_3\) and CaCO\(_3\); Be\(_2\)C gives methane and Mg\(_2\)C\(_3\) gives propyne.

Worked Example 1. Why is KO\(_2\) paramagnetic while Na\(_2\)O\(_2\) is diamagnetic?
Answer:
The superoxide ion \( \ce{O2^-} \) has 17 electrons, so one electron is unpaired in a \( \pi^* \) molecular orbital: paramagnetic.
The peroxide ion \( \ce{O2^{2-}} \) has 18 electrons, all paired: diamagnetic.

Worked Example 2. Arrange MgCl\(_2\), MgBr\(_2\), MgI\(_2\) in decreasing ionic character.
Answer:
The large I- ion is polarised the most (Fajans' rules), so covalent character is highest in MgI2.
Ionic character: MgCl2 > MgBr2 > MgI2.