s and p block elements: Class 12 Chemistry

Key points

The s block contains groups 1 and 2, where the last electron enters an s orbital. Group 1 (alkali metals: Li, Na, K, Rb, Cs) has the configuration ns¹. Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba) has ns². The p block contains groups 13 to 18, with configuration ns²np¹ to ns²np⁶ (helium is 1s²).

Alkali metals are soft, silvery, low in density and low in melting point, and they are stored under kerosene. They have large atomic size and the lowest ionisation energies in their periods, so they form M⁺ ions and are strong reducing agents. Atomic radius, reactivity and the ease of losing electrons increase down the group, while ionisation energy and melting point decrease. They react with water to give the hydroxide and hydrogen: 2Na + 2H₂O → 2NaOH + H₂. The reaction becomes more violent from Li to Cs.

Flame tests identify these metals. Li gives crimson red, Na golden yellow and K lilac. Among the alkaline earths, Ca gives brick red, Sr crimson and Ba apple green. The colour arises when excited electrons fall back to lower energy levels and emit light. Be and Mg give no colour because their electrons need too much energy to excite.

The alkaline earth metals are harder and denser than alkali metals and form M²⁺ ions. Their hydroxides become more soluble and more basic down the group, while their sulfates become less soluble (BaSO₄ is almost insoluble, MgSO₄ is soluble). Thermal stability of carbonates increases down the group. Ca(OH)₂ is slaked lime, CaO is quicklime and CaSO₄·½H₂O is plaster of Paris.

Lithium behaves differently from the rest of group 1 because of its small size and high charge density. It forms the nitride Li₃N and the normal oxide Li₂O (not peroxides), its hydroxide and carbonate decompose on heating, and it resembles magnesium. This is a diagonal relationship. Beryllium similarly differs from group 2 and resembles aluminium. Its compounds are covalent, and BeO is amphoteric.

Industrial compounds of sodium: Sodium is made by electrolysis of molten NaCl in the Down's cell. Sodium carbonate is made by the Solvay process: NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ + NH₄Cl. The sparingly soluble NaHCO₃ is filtered and heated: 2NaHCO₃ → Na₂CO₃ + H₂O + CO₂. The ammonia is recovered from NH₄Cl with Ca(OH)₂, and CaCO₃ supplies the CO₂. Na₂CO₃ is used in glass, soap and paper. NaOH is made from brine electrolysis.

Group 13 (B, Al, Ga, In, Tl) has configuration ns²np¹ and the common oxidation state +3. Boron is a non-metal and a metalloid, and the rest are metals. Boron forms covalent compounds, and the main source is borax, Na₂B₄O₇·10H₂O. Aluminium is the most abundant metal in the earth's crust. It is extracted from purified bauxite (Al₂O₃) by electrolysis in molten cryolite (Na₃AlF₆) in the Hall-Heroult process. Al₂O₃ is amphoteric. Tl shows +1 because of the inert pair effect.

Group 14 (C, Si, Ge, Sn, Pb) has ns²np². Carbon shows catenation and forms allotropes: diamond, graphite and fullerenes. CO₂ is acidic and SiO₂ is a giant covalent solid, in glass and sand. Silicones are polymers with Si-O-Si chains. Going down the group, the +2 state becomes more stable than +4 (inert pair effect), so Pb²⁺ is more stable than Pb⁴⁺ and PbO₂ is a strong oxidising agent.

Group 15 (N, P, As, Sb, Bi) has ns²np³. Nitrogen exists as N₂ with a strong triple bond, so it is unreactive. Ammonia is made by the Haber process, N₂ + 3H₂ ⇌ 2NH₃, at about 450 °C and 200 atm with iron catalyst. Nitric acid is made by the Ostwald process: NH₃ is oxidised over platinum to NO, then NO₂, and absorbed in water. Phosphorus exists as white phosphorus (reactive, P₄, stored under water) and red phosphorus (stable). Metallic character increases down the group.

Group 16 (O, S, Se, Te, Po) has ns²np⁴. Oxygen forms O₂ and ozone O₃. Sulfur shows allotropes rhombic and monoclinic. Sulfuric acid is made by the Contact process: S burns to SO₂, SO₂ is oxidised to SO₃ with V₂O₅ catalyst at about 450 °C, SO₃ is absorbed in concentrated H₂SO₄ to give oleum, and oleum is diluted. H₂SO₄ is a dehydrating agent and an oxidising agent when hot and concentrated.

Group 17 halogens (F, Cl, Br, I, At) have ns²np⁵ and exist as diatomic molecules. Fluorine is a gas, chlorine a gas, bromine a liquid and iodine a solid. Electronegativity and oxidising power decrease down the group, so F₂ > Cl₂ > Br₂ > I₂, and a more reactive halogen displaces a less reactive one from its salt. Acid strength of the hydrogen halides rises from HF (weak) to HI (strongest) as bond strength falls. Chlorine is used in bleaching and water treatment.

Group 18 noble gases (He, Ne, Ar, Kr, Xe, Rn) have complete outer shells (ns²np⁶) and are mostly unreactive and monatomic. Their boiling points increase down the group with the increase in dispersion forces. Xenon forms compounds with fluorine and oxygen such as XeF₂, XeF₄, XeF₆ and XeO₃. Uses: helium in balloons and for cooling, neon in lighting signs, argon in electric bulbs and welding.

Study tip: Study this chapter group by group using a table: configuration, trend, key compounds, uses and one industrial process. Learn each process with its equation, catalyst and conditions. For trends always give the reason, such as atomic size or ionisation energy, because marks are given for the explanation.

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Common mistakes

• Saying that ionisation energy increases down a group. It decreases because the outer electron is farther from the nucleus.
• Writing that Li reacts with nitrogen like other alkali metals. It is the only one in group 1 that forms a nitride directly.
• Forgetting that the Contact process uses V₂O₅ and that SO₃ is absorbed in H₂SO₄, not directly in water.
• Writing the Solvay equation without ammonia, or forgetting that NaHCO₃ is heated to give Na₂CO₃.
• Giving Pb⁴⁺ as more stable than Pb²⁺, ignoring the inert pair effect.

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