Samacheer Kalvi 12th Chemistry Chapter 2 Exercise Solutions

(a) Silicon (b) Germanium (c) Carbon (d) Tin

(a) Alkali metals (b) Transition metal cations (c) Halide ions (d) Sulphate ions

(a) CH₄ (b) C₂H₆ (c) B₂H₆ (d) SiH₄

(a) Donates a proton (b) Accepts an electron pair from water (c) Donates an electron pair (d) Releases H⁺ directly

(a) Boron (b) Aluminium (c) Gallium (d) Thallium

(a) Inorganic polymers with Si-Si backbone (b) Organic polymers with C-C backbone (c) Polymers with Si-O backbone and organic side groups (d) Ionic compounds of silicon

(a) Pure molten alumina (b) Alumina dissolved in molten cryolite (c) Alumina dissolved in water (d) Aluminium chloride solution

(a) Graphite (b) Fullerene (c) Diamond (d) Carbon black

(a) Boron and oxygen (b) Sodium metaborate and boric anhydride (c) Boric acid and sodium oxide (d) Diborane and water

(a) Boron shows properties different from the rest of group 13 (b) Aluminium is more metallic than boron (c) Gallium has a low melting point (d) Indium shows +3 oxidation state

(a) Silicates with chain structure (b) Aluminosilicates with a framework of pores (c) Simple silicon oxides (d) Borosilicate glass

(a) K₂SO₄·Al₂(SO₄)₃·24H₂O (b) Na₂SO₄·Al₂(SO₄)₃·24H₂O (c) KAl(SO₄)₂·12H₂O (d) Both (a) and (c) are equivalent formulae

The inert pair effect is the tendency of the ns² electrons in heavier p-block elements to remain paired and not participate in bonding, due to poor shielding by d and f electrons. In group 13, boron shows +3, but as we go down, the +1 oxidation state becomes more stable. Thallium (Tl) prefers +1 because the 6s² electrons act as an "inert pair," making Tl⁺ more stable than Tl³⁺.

Boron differs from other group 13 elements because: (1) It is a non-metal while others are metals. (2) It has a very small atomic size and high ionisation energy. (3) It forms only covalent compounds, unlike the ionic compounds of aluminium. (4) It shows a maximum coordination number of 4 (not 6 like Al) due to its small size. (5) It forms electron-deficient compounds like diborane with 3-centre 2-electron bonds.

Aluminium oxide reacts with both acids and bases, proving its amphoteric nature:

Silicones are synthetic polymers with a backbone of alternating silicon and oxygen atoms, with organic groups (methyl or phenyl) attached to silicon. Their general formula is (R₂SiO)ₙ. They are thermally stable, water-repellent, and chemically inert. Uses: (1) As sealants and water-repellent coatings, and (2) As insulating materials in electrical equipment.

Diamond has sp³ hybridised carbon atoms in a 3-dimensional tetrahedral network, making it the hardest natural substance, a poor conductor of electricity, and transparent. Graphite has sp² hybridised carbon atoms arranged in hexagonal layers held by weak van der Waals forces, making it soft, slippery, a good conductor of electricity (due to delocalised π electrons), and opaque.

The borax bead test is a qualitative test for identifying transition metal cations. When borax (Na₂B₄O₇·10H₂O) is heated on a platinum wire loop, it forms a colourless glassy bead of sodium metaborate and boric anhydride (B₂O₃). When this bead is heated with a metal salt, a coloured metaborate is formed. For example, cobalt gives a blue bead (Co(BO₂)₂), and chromium gives a green bead.

Aluminium is extracted from bauxite (Al₂O₃·2H₂O) in two main stages:

Stage 1 — Concentration (Bayer's Process): Bauxite ore is crushed and digested with concentrated NaOH solution under pressure at 473–523 K. Alumina dissolves as sodium aluminate, while impurities (Fe₂O₃, SiO₂, TiO₂) remain as red mud. The solution is filtered, cooled, and seeded with Al(OH)₃ crystals.

Stage 2 — Hall-Heroult Process (Electrolytic Reduction): Pure alumina is dissolved in molten cryolite (Na₃AlF₆) at about 950 °C in an iron cell lined with carbon (cathode). Carbon rods serve as the anode. On passing current, Al³⁺ ions are reduced at the cathode.

The oxygen liberated reacts with the carbon anode to form CO and CO₂, so anodes must be replaced periodically. The aluminium obtained is about 99.5% pure and is further refined by Hoopes process if needed.

(a) Borax (Na₂B₄O₇·10H₂O): Also called sodium tetraborate decahydrate, borax is a white crystalline solid. On heating, it first loses water molecules, then decomposes to sodium metaborate and boric anhydride, forming a glassy bead used in the borax bead test.

Borax is used as a flux in welding, in the borax bead test, in making optical and borosilicate glass, and as a mild antiseptic.

(b) Orthoboric Acid (H₃BO₃): Orthoboric acid is obtained by treating borax with sulphuric acid:

It is a weak monobasic Lewis acid — it does not donate a proton but accepts OH⁻ from water:

On heating, it forms metaboric acid and then boric anhydride:

(c) Diborane (B₂H₆): Diborane is the simplest boron hydride. It is an electron-deficient compound with 12 valence electrons. It has two types of B-H bonds: four terminal B-H bonds (normal 2-centre 2-electron bonds) and two bridging B-H-B bonds (3-centre 2-electron or "banana" bonds). Diborane burns in air to give boron trioxide and water:

It hydrolyses with water:

Group 14 elements (C, Si, Ge, Sn, Pb) show the following trends:

(a) Atomic and Ionic Radii: Atomic radii increase down the group from carbon to lead as additional shells are added. However, the increase is not uniform due to d- and f-orbital filling (lanthanide contraction effect on Pb).

(b) Ionisation Energy: Ionisation energy decreases down the group as atomic size increases. Carbon has the highest first ionisation energy. There is a slight irregularity at Ge-Sn due to poor shielding by d-electrons.

(c) Metallic Character: Increases down the group. Carbon and silicon are non-metals (silicon is a metalloid), germanium is a metalloid, while tin and lead are metals.

(d) Oxidation States: The common oxidation states are +4 and +2. The +4 state is more stable for lighter elements (C, Si), while the +2 state becomes more stable for heavier elements (Sn, Pb) due to the inert pair effect. Carbon also shows -4 in carbides.

(e) Catenation: The ability to form self-linked chains decreases down the group. Carbon shows maximum catenation (forms millions of organic compounds). Silicon shows limited catenation, and tin and lead show almost none.

(f) Allotropy: All group 14 elements show allotropy. Carbon has diamond and graphite (and fullerene). Silicon has amorphous and crystalline forms. Tin has grey (α) and white (β) forms.

(a) Silicates: Silicates are minerals composed of SiO₄⁴⁻ tetrahedra linked in various patterns through shared oxygen atoms. The classification depends on how the tetrahedra are linked:

1. Orthosilicates (Nesosilicates): Independent SiO₄⁴⁻ units. Example: Olivine (Mg₂SiO₄). 2. Pyrosilicates (Sorosilicates): Two tetrahedra share one oxygen → Si₂O₇⁶⁻. Example: Thortveitite. 3. Chain silicates (Inosilicates): Tetrahedra share two oxygens forming chains — single chain (SiO₃²⁻)ₙ or double chain (Si₄O₁₁⁶⁻)ₙ. Example: Asbestos. 4. Sheet silicates (Phyllosilicates): Each tetrahedron shares three oxygens → (Si₂O₅²⁻)ₙ. Example: Talc, mica. 5. 3D framework silicates (Tectosilicates): All four oxygens are shared → SiO₂ (quartz). When Al replaces some Si, aluminosilicates form.

(b) Zeolites: Zeolites are microporous aluminosilicates with a 3D framework of SiO₄ and AlO₄ tetrahedra. They have regularly spaced pores of molecular dimensions, making them excellent molecular sieves. Uses: (1) As molecular sieves to separate molecules by size. (2) As ion exchangers in water softening. (3) As catalysts in petrochemical cracking (e.g., ZSM-5). (4) In removing hardness of water by exchanging Ca²⁺/Mg²⁺ with Na⁺.

Carbon shows anomalous behaviour compared to other group 14 elements due to its small size, high electronegativity, and high ionisation energy:

1. Maximum Catenation: Carbon forms millions of compounds by linking with itself (C-C bonds) and with other elements. This is because the C-C bond energy (348 kJ/mol) is very high, comparable to C-H and C-O bonds. Silicon, in contrast, has a much lower Si-Si bond energy (222 kJ/mol).

2. Multiple Bond Formation: Carbon readily forms double and triple bonds (C=C, C≡C, C=O, C≡N) due to its small size and effective p-orbital overlap. Other group 14 elements rarely form stable multiple bonds due to poor p-orbital overlap.

3. Allotropy: Carbon has two important allotropic forms — diamond (sp³, 3D network, hardest substance) and graphite (sp², layered, conductor) — plus fullerene (C₆₀). Other elements show limited allotropy.

4. Tetravalency with Small Size: Carbon's small size allows four bond pairs to be accommodated around it without repulsion, giving stable tetrahedral geometry.

5. Hybridisation: Carbon shows sp³, sp², and sp hybridisation, enabling formation of single, double, and triple bonds with directional properties.

6. Formation of Extremely Stable Compounds: CO₂ is a gas (molecular), while SiO₂ is a solid (network). This is because CO₂ forms discrete molecules with C=O double bonds, whereas SiO₂ forms a 3D network of single bonds.