Silicones are synthetic polymers with Si-O-Si backbones and organic side groups. Repeat unit R2SiO. They are thermally stable, hydrophobic and chemically inert, used as lubricants, sealants, waterproofing agents, insulators and medical implants.

Group 14 Elements: The Carbon Family

Q: What is catenation and why does carbon show it more than silicon?

Catenation is the ability of atoms of an element to link into long chains or rings. Carbon shows extreme catenation because the C-C bond enthalpy is very high, small size gives excellent orbital overlap and there are no lone pairs to repel neighbours. Silicon's catenation is limited to about 8 atoms.

Q: Why does the +2 oxidation state become more stable down Group 14?

Due to the inert pair effect. The ns2 pair of electrons in the heavier atoms Sn and Pb is reluctant to participate in bonding because of poor d and f shielding. So Pb prefers Pb2+, while C and Si dominate in the +4 state.

Q: What are the three main allotropes of carbon?

Diamond (sp3, 3D tetrahedral network), graphite (sp2, 2D hexagonal layers with delocalised pi system) and fullerene C60 (sp2 cage of 12 pentagons and 20 hexagons). Diamond is hard and insulating, graphite is soft and conducting, fullerene is a molecular solid.

Q: Why is CO2 a gas but SiO2 is a solid?

Carbon is small enough to form strong pi bonds with oxygen giving discrete O=C=O molecules held only by weak van der Waals forces, so CO2 is a gas. Silicon is larger and cannot form effective pi bonds with oxygen; instead each Si single bonds to four O and each O bridges two Si atoms, giving a 3D covalent network solid.

Q: Why is CO a good ligand for transition metals?

CO donates a lone pair from carbon to an empty metal orbital (sigma donation) and its empty pi-star molecular orbitals accept electron density back from filled metal d-orbitals (pi back-donation). This synergic bonding produces very stable metal carbonyls like Ni(CO)4 and Fe(CO)5.

Q: What is a silicate and how are silicates classified?

Silicates are compounds built from SiO4 4- tetrahedra. They are classified by the number of corner oxygens shared with neighbouring tetrahedra: orthosilicates (0 shared), pyrosilicates (1), cyclic (2 in a ring), chain (2 in a line), sheet (3) and three-dimensional (4).

Q: What are zeolites and why are they called molecular sieves?

Zeolites are 3D aluminosilicates with pores of well-defined size that let small molecules enter and exclude larger ones. Uses include water softening by ion exchange, petroleum catalysis such as ZSM-5, and drying of organic solvents.

Q: Why does CCl4 not hydrolyse in water while SiCl4 does?

Hydrolysis needs the central atom to accept a lone pair from water. Si has empty 3d orbitals to accept the lone pair; carbon has no accessible d orbitals in its valence shell so cannot expand its octet. Hence SiCl4 hydrolyses but CCl4 does not.

ChemistryThe p-Block Elements: Group 13 & 14For JEE aspirants

The Carbon family or Group 14 consists of five elements: carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). All have the valence-shell configuration $n s^{2} n p^{2}$ – four valence electrons that allow four covalent bonds. Carbon is a non-metal, Si and Ge are metalloids, and Sn and Pb are true metals. Carbon shows a unique ability to form long chains and rings (catenation) that underlies the whole of organic chemistry, while silicon dominates the mineral world through its silicates. Down the group, the $+ 2$ oxidation state becomes progressively more stable due to the inert pair effect, so that $P b^{2 +}$ is more stable than $P b^{4 +}$ .

Key Formulas & Facts – Quick Reference

General valence configuration: $n s^{2} n p^{2}$ (n = 2 to 6)
Common oxidation states: $+ 4$ (all), $+ 2$ (stability increases down group; $P b^{2 +} > P b^{4 +}$ )
Catenation order: $C ≫ S i > G e \approx S n ≫ P b$
Allotropes of carbon: diamond ( $s p^{3}$ ), graphite ( $s p^{2}$ , layered), fullerene C $_{60}$ ( $s p^{2}$ , cage)
Producer gas: $C + ai r \to C O + N_{2}$ ; Water gas: $C + H_{2} O \to C O + H_{2}$
CO $_{2}$ is linear, monomeric, gas at RT ( $O = C = O$ ); SiO $_{2}$ is 3D covalent network, solid
Basic silicate unit: $S i O_{4}^{4 -}$ tetrahedron; silicones are $(R_{2} S i O)_{n}$ polymers
Zeolites are aluminosilicate cage structures used as molecular sieves

1. Introduction and Electronic Configuration

Group 14 sits in the middle of the $p$ -block. The four valence electrons split as two $s$ + two $p$ , allowing $s p^{3}$ , $s p^{2}$ or $s p$ hybridisation and rich covalent chemistry.

Element	Symbol	Z	Electronic Configuration	Character
Carbon	C	6	$[H e] 2 s^{2} 2 p^{2}$	Non-metal
Silicon	Si	14	$[N e] 3 s^{2} 3 p^{2}$	Metalloid
Germanium	Ge	32	$[A r] 3 d^{10} 4 s^{2} 4 p^{2}$	Metalloid
Tin	Sn	50	$[K r] 4 d^{10} 5 s^{2} 5 p^{2}$	Metal
Lead	Pb	82	$[X e] 4 f^{14} 5 d^{10} 6 s^{2} 6 p^{2}$	Metal

Occurrence

Carbon: in free state as diamond and graphite; in combined state as fossil fuels (coal, petroleum, natural gas), carbonate minerals (limestone $C a C O_{3}$ , dolomite $C a C O_{3} \cdot M g C O_{3}$ ) and $C O_{2}$ in air ( $\sim 0.04%$ ).
Silicon: second most abundant element in earth's crust ( $\sim 27.7%$ ); occurs as silica ( $S i O_{2}$ , quartz sand) and silicate minerals.
Ge, Sn, Pb: Ge from zinc ores; Sn as cassiterite $S n O_{2}$ ; Pb chiefly as galena $P b S$ .

2. Trends in Physical Properties

Figure 1: Atomic radius (solid) increases sharply from C to Si, then slowly. Ionization enthalpy (dashed) falls sharply C→Si, dips at Sn but rises at Pb because of the poor

4 f^{14}

shielding (lanthanoid contraction – same reason Tl in Group 13 has higher IE than In).

Atomic size, IE and metallic character

Atomic radii increase down the group, but the C → Si jump is huge because the second shell has no $d$ -orbitals.
First IE generally decreases; Ge and Pb show small anomalies due to $d$ -block and lanthanoid contractions.
Metallic character increases: C (non-metal) → Si, Ge (metalloids) → Sn, Pb (metals).

Oxidation states and inert pair effect

Group 14 shows both $+ 4$ and $+ 2$ states. Down the group the $+ 2$ state becomes more stable:

Stability of

+ 4

C > S i > G e > S n > P b

Stability of

+ 2

C < S i < G e < S n < P b

So $S n C l_{2}$ is a good reducing agent (Sn wants to go to $S n^{4 +}$ ), while $P b C l_{4}$ is an oxidising agent (Pb wants to go to $P b^{2 +}$ ). This is again the inert pair effect: the $n s^{2}$ pair becomes reluctant to participate in bonding down the group.

Figure 2: The inert pair effect. Down Group 14 the

n s^{2}

pair becomes increasingly reluctant to take part in bonding, so the

+ 4

state is progressively destabilised and the

+ 2

state stabilised. This is why

S n C l_{2}

acts as a reducing agent while

P b C l_{4}

acts as an oxidising agent.

3. Catenation and Anomalous Behaviour of Carbon

Catenation

Catenation is the tendency of an element's atoms to link to each other and form chains and rings. Carbon shows this to an extreme:

Catenation order:

C ≫ S i > G e \approx S n ≫ P b

Reasons for carbon's exceptional catenation:

Very high $C - C$ bond enthalpy ( $\sim 348$ kJ mol $^{- 1}$ ; the $S i - S i$ bond is only $\sim 297$ kJ mol $^{- 1}$ ).
Small size gives strong overlap between $s p^{3}$ hybrid orbitals.
No lone pairs of electrons that would repel each other in long chains.

Figure 3: Element-element single bond enthalpies. The

C - C

bond is far stronger than

S i - S i

G e - G e

S n - S n

, and the bond weakens steadily down the group. This single trend is the main reason catenation is essentially unlimited for carbon but dies out after silicon.

Carbon's catenation is the entire foundation of organic chemistry (millions of $C$ - $C$ compounds). Silicon's catenation is limited to about 8 atoms (in silanes $S i_{n} H_{2 n + 2}$ , $n \leq 8$ ).

Anomalous behaviour of carbon

Like B in Group 13, carbon differs from the rest of its group due to:

Very small atomic size and high electronegativity ( $2.55$ ).
Absence of $d$ -orbitals in valence shell – maximum covalency is $4$ .
Ability to form strong $p π$ - $p π$ multiple bonds ( $C = C$ , $C \equiv C$ , $C = O$ , $C \equiv N$ ).

Consequences: $C O_{2}$ is a discrete linear molecule (with $C = O$ multiple bonds), but $S i O_{2}$ is a giant 3D covalent network of $S i - O$ single bonds (Si cannot form effective $p π$ - $p π$ bonds with O because of size mismatch).

4. Allotropes of Carbon

Carbon exists in several crystalline and amorphous forms. The three most important crystalline allotropes are diamond, graphite and fullerene C $_{60}$ .

4.1 Diamond

Each carbon is $s p^{3}$ hybridised and bonded to four other carbons at $109.5°$ in a rigid three-dimensional tetrahedral network.

Figure 4: Diamond has each carbon

s p^{3}

hybridised, bonded to four others at

109.5°

in an infinite three-dimensional network. Every valence electron is in a strong

C - C

σ

bond, so diamond has no free electrons – it is a poor electrical conductor but excellent thermal conductor and the hardest known natural material.

Properties: extremely hard, transparent, high melting point ( $4373$ K), very high refractive index ( $2.42$ – hence the "fire" of a cut diamond), poor electrical conductor (no free electrons), excellent thermal conductor.

Uses: gemstones, cutting/drilling/grinding tools, glass cutting, high-precision bearings.

4.2 Graphite

Each carbon is $s p^{2}$ hybridised and bonded to three others in a flat hexagonal sheet. The fourth valence electron sits in an unhybridised $p$ -orbital and joins a delocalised $π$ -system across the whole sheet. These sheets stack in parallel layers held only by weak van der Waals forces.

Figure 5: Graphite has

s p^{2}

hybridised carbons in hexagonal sheets. The delocalised

π

system (shaded lobes above and below each sheet) makes graphite a good electrical conductor along the sheets. The sheets are held together only by weak van der Waals forces, so they slide over one another – giving graphite its lubricating and marking (pencil) properties.

Properties: soft, greasy, black, opaque; good conductor of electricity parallel to sheets (poor across sheets); very high melting point ( $3873$ K under pressure).

Uses: pencil leads (mixed with clay), electrodes (dry cells, electric arcs, Hall-Heroult anodes), high-temperature crucibles, dry lubricant, moderator in nuclear reactors.

JEE Advanced

Artificial preparation of graphite (Acheson process): a mixture of coke (or petroleum coke) and a little silica or iron oxide catalyst is heated to $\sim 3273$ K in an electric furnace. The coke rearranges into graphite; volatile silicon monoxide escapes.

$S i O_{2} + 3 C 3273 K S i C + 2 C O$ ; then $S i C Δ S i ↑ + C (graphite)$

Acheson graphite is used for electrodes, crucibles and high-performance carbon brushes.

4.3 Fullerenes

Fullerenes are cage-like molecular allotropes of carbon. The most famous, Buckminsterfullerene $C_{60}$ , was discovered in $1985$ by Kroto, Curl and Smalley (Nobel Prize $1996$ ).

Figure 6 (JEE Advanced):

C_{60}

has the shape of a truncated icosahedron – the same pattern as a football. It contains exactly

12

pentagonal and

20

hexagonal faces, with

60

s p^{2}

-hybridised carbons at the vertices. Each C is bonded to three neighbours; the remaining

p

-electron gives a partially delocalised

π

system. The molecule is aromatic in a limited sense.

JEE Advanced

C $_{60}$ geometry – the Euler check: for any closed polyhedron, $V - E + F = 2$ (Euler's formula). For a truncated icosahedron built from pentagons ( $P$ ) and hexagons ( $H$ ): $F = P + H$ ; each edge is shared by two faces, so $E = \frac{5 P + 6 H}{2}$ ; each vertex is shared by three faces, so $V = \frac{5 P + 6 H}{3}$ . Setting up Euler and requiring integer solutions gives $P = 12$ exactly, and hexagons can be any number. In $C_{60}$ , $H = 20$ , giving $V = 60$ , $E = 90$ , $F = 32$ . Two important rules: (i) no two pentagons share an edge (isolated pentagon rule); (ii) each pentagon is surrounded by five hexagons.

Properties: $C_{60}$ is dark solid, sublimes at $\sim 800$ K; soluble in benzene and toluene giving deep purple solutions. Other members: $C_{70}$ , $C_{76}$ , $C_{84}$ , and elongated cage structures called carbon nanotubes and sheet-like graphene.

Uses: superconductors (doped fullerides like $K_{3} C_{60}$ ), drug delivery (endohedral fullerenes), lubricants and organic electronics.

Figure 7: The three crystalline allotropes side-by-side. Diamond (3D

s p^{3}

) is hard and insulating; graphite (2D

s p^{2}

) is soft and conducting; fullerene

C_{60}

(0D

s p^{2}

cage) is neither – it is a discrete molecule that dissolves in organic solvents.

5. Chemical Properties of Group 14

Reactivity with water

C and Si do not react with water under ordinary conditions. Sn reacts slowly with steam; Pb has a protective oxide film that resists water.

$S n + 2 H_{2} O steam S n O_{2} + 2 H_{2}$

Very hot coke reacts with steam to give water gas (used industrially):

$C + H_{2} O 1273 K water gas C O + H_{2}$

Reactivity with halogens

All Group 14 elements form tetrahalides $M X_{4}$ ; the heavier ones also form $M X_{2}$ .

$M + 2 X_{2} \to M X_{4}$ ( $M = C, S i, G e, S n, P b$ )

Except $C X_{4}$ , all tetrahalides can be hydrolysed by water because M has vacant $d$ -orbitals to accept lone pairs from water:

$S i C l_{4} + 4 H_{2} O \to S i (O H)_{4} + 4 H C l$

$C C l_{4} + H_{2} O \to$ no reaction (C has no $d$ -orbitals available in valence shell)

Oxides

All members form monoxides $M O$ (except Si) and dioxides $M O_{2}$ .

Acidity of dioxides:

C O_{2}, S i O_{2}

(acidic)

>

G e O_{2}

(amphoteric)

>

S n O_{2}, P b O_{2}

(amphoteric)
Acidity of monoxides:

C O

(neutral);

G e O, S n O, P b O

(amphoteric)

6. Important Oxides of Carbon and Silicon

6.1 Carbon monoxide, $C O$

Colourless, odourless, highly poisonous gas.

Preparation:

Laboratory: $H C O O H conc. H_{2} S O_{4} Δ C O + H_{2} O$
Industrial: $C + \frac{1}{2} O_{2} \to C O$ (limited air, in blast furnace).
Water gas (also called synthesis gas): $C + H_{2} O 1273 K C O + H_{2}$
Producer gas: $2 C + O_{2} + 4 N_{2} Δ 2 C O + 4 N_{2}$ (with air, so mostly N $_{2}$ ).

Figure 8 (JEE Advanced): CO has a triple bond (

1 σ + 2 π

) between C and O, giving bond order

3

. One of the

π

bonds is a dative bond (both electrons from O). Both atoms retain a lone pair. This is why CO acts as a strong

σ

-donor and

π

-acceptor to transition metals, forming metal carbonyls like

N i (C O)_{4}

and

F e (C O)_{5}

Properties & reactions:

Combustion: $2 C O + O_{2} \to 2 C O_{2}$ (blue flame; used as fuel gas).
Powerful reducing agent (used in blast furnace to reduce iron ore): $F e_{2} O_{3} + 3 C O \to 2 F e + 3 C O_{2}$
With transition metals: $N i + 4 C O \to N i (C O)_{4}$ (Mond process); $F e + 5 C O \to F e (C O)_{5}$
Poisonous nature: CO binds to haemoglobin $\sim 300$ times more strongly than O $_{2}$ , forming carboxyhaemoglobin and preventing oxygen transport – hence its lethal effect.

6.2 Carbon dioxide, $C O_{2}$

Colourless, odourless gas; small $\sim 0.04%$ component of air. Denser than air. Slightly soluble in water giving weakly acidic carbonic acid.

Preparation:

Laboratory: $C a C O_{3} + 2 H C l \to C a C l_{2} + H_{2} O + C O_{2}$
Industrial: combustion of coke ( $C + O_{2} \to C O_{2}$ ); calcination of limestone ( $C a C O_{3} Δ C a O + C O_{2}$ ); fermentation ( $C_{6} H_{12} O_{6} \to 2 C_{2} H_{5} O H + 2 C O_{2}$ ).

Reactions:

With water: $C O_{2} + H_{2} O ⇌ H_{2} C O_{3}$ (weak carbonic acid).
With lime water: $C a (O H)_{2} + C O_{2} \to C a C O_{3} ↓ + H_{2} O$ (milky). Excess CO $_{2}$ : $C a C O_{3} + C O_{2} + H_{2} O \to C a (H C O_{3})_{2}$ (milkiness disappears).
With Mg (only strong reductant that can beat it): $2 M g + C O_{2} \to 2 M g O + C$ (magnesium burns in CO $_{2}$ ).
Photosynthesis: $6 C O_{2} + 6 H_{2} O chlorophyll, sunlight C_{6} H_{12} O_{6} + 6 O_{2}$

Greenhouse effect:

C O_{2}

absorbs infrared radiation strongly (asymmetric stretch and bending modes match IR). Rising atmospheric

C O_{2}

from fossil fuel burning traps more IR near the surface, warming the planet.

6.3 Why $C O_{2}$ is a gas but $S i O_{2}$ is a solid

This is a classic JEE contrast question, and the answer lies in the ability (or not) of the central atom to form $p π$ - $p π$ multiple bonds with oxygen.

Figure 9: CO

_{2}

vs SiO

_{2}

. Carbon is small enough to form strong

p π

p π

multiple bonds with oxygen, so each O is doubly bonded to C giving a discrete linear molecule (a gas). Silicon is much larger; the

S i 3 p

O 2 p

overlap is poor. Instead each Si single-bonds to four O atoms and each O bridges two Si atoms, giving a giant 3D covalent solid.

7. Silicon Dioxide (Silica), $S i O_{2}$

Silica is one of the commonest solids on earth: it makes up quartz sand, quartz crystal, agate and flint. Structure was covered in Figure 9.

Physical forms:

Quartz – hexagonal crystalline; the purest form is rock crystal. Used in optics, oscillators (quartz watches), lenses.
Cristobalite and tridymite – other crystalline polymorphs stable at higher temperatures.
Amorphous forms – silica gel (drying agent), fumed silica, opal.

Chemical reactions:

With alkalis (acidic oxide behaviour): $S i O_{2} + 2 N a O H \to N a_{2} S i O_{3} + H_{2} O$
With hydrofluoric acid (only acid that attacks silica): $S i O_{2} + 4 H F \to S i F_{4} + 2 H_{2} O$ ; then $S i F_{4} + 2 H F \to H_{2} [S i F_{6}]$
With carbonates on fusion: $S i O_{2} + N a_{2} C O_{3} \to N a_{2} S i O_{3} + C O_{2}$

Uses: optics, silica gel (drying agent), semiconductor manufacture (as ultra-pure SiO $_{2}$ layer), glass making, cement.

8. Silicates – the Structural Chemistry

Silicates are salts (or compounds) built from $S i O_{4}^{4 -}$ tetrahedral units. The tetrahedra can share $0$ , $1$ , $2$ , $3$ or $4$ corner oxygens with neighbours, giving five structural classes.

Figure 10: The building block of all silicates is a

[S i O_{4}]^{4 -}

tetrahedron. Silicon at the centre is

s p^{3}

hybridised, bonded to four oxygens with tetrahedral geometry. The overall

- 4

charge is balanced by cations (

N a^{+}

M g^{2 +}

C a^{2 +}

A l^{3 +}

, etc.) in the crystal.

JEE Advanced

Classification of silicates by corner sharing:

Type	Corner O shared	Unit formula	Example mineral
Ortho- (island)	0	$[S i O_{4}]^{4 -}$	Zircon $Z r S i O_{4}$ , olivine $M g_{2} S i O_{4}$
Pyro- (dimer)	1	$[S i_{2} O_{7}]^{6 -}$	Thortveitite $S c_{2} S i_{2} O_{7}$
Cyclic (ring)	2	$[S i_{n} O_{3 n}]^{2 n -}$	Beryl $B e_{3} A l_{2} [S i_{6} O_{18}]$
Chain (single)	2	$[S i O_{3}]_{n}^{2 n -}$	Pyroxenes like spodumene $L i A l (S i O_{3})_{2}$
Sheet	3	$[S i_{2} O_{5}]_{n}^{2 n -}$	Mica, kaolin, talc
Three-dimensional	4	$[S i O_{2}]_{n}$	Quartz, feldspars, zeolites

Figure 11 (JEE Advanced): Three important silicate structural types shown from above (each triangle is a

[S i O_{4}]^{4 -}

tetrahedron looking down its apex). Left: single chain (pyroxene) – each Si shares

2

O. Middle: cyclic six-ring (beryl) – each Si also shares

2

O but in a closed loop. Right: 2D sheet (mica) – each Si shares

3

O with three neighbours.

9. Silicones

Silicones are synthetic polymers with the general repeat unit $- (R_{2} S i O) -$ , where R is usually methyl or phenyl. The name "silicone" was coined because the empirical formula $R_{2} S i O$ looks like a ketone $R_{2} C O$ , though the structures are completely different (silicones have $S i - O - S i$ backbones, not $C = O$ ).

Preparation (Rochow-Muller direct process): alkyl or aryl halides react with silicon in the presence of a Cu catalyst at $\sim 570$ K to give alkyl chlorosilanes, which are then hydrolysed and polymerised:

$2 C H_{3} C l + S i C u 570 K (C H_{3})_{2} S i C l_{2}$

$(C H_{3})_{2} S i C l_{2} + 2 H_{2} O \to (C H_{3})_{2} S i (O H)_{2} + 2 H C l$

The silanediol then polymerises with loss of water to give a linear silicone:

$n (C H_{3})_{2} S i (O H)_{2} \to [- (C H_{3})_{2} S i - O -]_{n} + n H_{2} O$

Figure 12: Linear silicone chain. The

S i - O - S i

backbone is very strong and thermally stable; the outward-pointing organic groups (methyl in this case) give the material a hydrophobic surface. Silicones are used as high-temperature lubricants, waterproofing agents, sealants, medical implants and cosmetics.

JEE Advanced

Types of silicone polymers depend on the starting chlorosilane:

$R_{3} S i C l$ (mono-functional) → only chain ends → low molecular weight; used as chain-terminators.
$R_{2} S i C l_{2}$ (di-functional) → linear chains and rings (silicone oils, greases).
$R S i C l_{3}$ (tri-functional) → cross-linked 3D networks (silicone rubbers, resins).
$S i C l_{4}$ (tetra-functional) → hard cross-linked resins.

By blending different chlorosilanes, silicones with any desired viscosity, flexibility or hardness can be tailored.

Uses: high-temperature lubricants and greases, waterproofing agents for textiles and buildings, electrical insulators, sealants, silicone rubber tubing, breast implants and other medical prostheses.

10. Zeolites

Zeolites are three-dimensional aluminosilicates in which some $S i^{4 +}$ ions in the $S i O_{2}$ framework are replaced by $A l^{3 +}$ ions. The charge deficit is balanced by exchangeable cations (Na $^{+}$ , Ca $^{2 +}$ ) sitting in cavities of the structure.

General formula: $M_{x / n} [(A l O_{2})_{x} (S i O_{2})_{y}] \cdot m H_{2} O$ , where $M$ is a cation of charge $n$ .

Figure 13: Zeolite cage schematic. The 3D aluminosilicate framework has pores of well-defined size (typically

300 - 1000

pm). Molecules smaller than the pore diameter pass through and are adsorbed inside the cavities; larger molecules are excluded. This is the "molecular sieve" action.

Uses:

Water softeners (ion exchange: $C a^{2 +}$ in hard water swaps in, $N a^{+}$ swaps out).
Selective catalysts in petroleum industry: ZSM-5 converts alcohols to petrol; other zeolites crack heavy hydrocarbons.
Molecular sieves for drying solvents ( $3$ -Å sieve absorbs H $_{2}$ O; $4$ -Å absorbs H $_{2}$ O, CO $_{2}$ ; $5$ -Å absorbs straight-chain hydrocarbons but excludes branched isomers).
Deodorisers and gas separation membranes.

11. Compounds of Tin and Lead (brief)

The inert pair effect makes $P b^{2 +}$ far more stable than $P b^{4 +}$ (opposite for Sn).

$S n C l_{2}$ – strong reducing agent: $2 F e C l_{3} + S n C l_{2} \to 2 F e C l_{2} + S n C l_{4}$ ; $S n^{2 +}$ → $S n^{4 +}$ .
$P b O_{2}$ – strong oxidising agent: $P b O_{2} + 4 H C l \to P b C l_{2} + C l_{2} + 2 H_{2} O$ ; $P b^{4 +}$ → $P b^{2 +}$ .
$P b_{3} O_{4}$ (red lead) – a mixed oxide, $2 P b O \cdot P b O_{2}$ , used in anticorrosion primer paints.

Solved Example 1

Explain why

C O_{2}

is a gas at room temperature but

S i O_{2}

is a hard high-melting solid, though both have the same

M O_{2}

formula.

Solution:

Carbon is small ( $77$ pm) and its $2 p$ orbitals overlap effectively with the $2 p$ orbitals of oxygen, so C forms two strong $p π$ - $p π$ double bonds with two oxygens, giving a discrete linear $O = C = O$ molecule. Weak van der Waals forces between these small non-polar molecules mean CO $_{2}$ is a gas at room temperature. Silicon is much bigger ( $118$ pm) and $S i 3 p$ overlaps very poorly with $O 2 p$ – $p π$ - $p π$ bonding fails. To satisfy its octet, each Si single-bonds to four O atoms and each O bridges two Si atoms, generating a giant 3D covalent network. Breaking such a network requires breaking many strong $σ$ bonds, so SiO $_{2}$ is a hard solid melting near $1983$ K.

Solved Example 2

Diamond is one of the hardest natural substances yet graphite is soft enough to leave a mark on paper. Explain the difference in terms of structure and bonding.

Solution:

In diamond, every carbon is $s p^{3}$ hybridised and bonded to four other carbons in an isotropic 3D covalent network. To scratch or fracture diamond you have to break a huge number of strong $C - C$ $σ$ bonds simultaneously in any direction – hence its extreme hardness. In graphite, each carbon is $s p^{2}$ hybridised and bonded to three others in a 2D hexagonal sheet. Adjacent sheets are held only by weak van der Waals forces ( $\sim 335$ pm apart). Under mild shear (say, dragging graphite on paper), the sheets slide past one another easily – this leaves a smear on the paper (the "lead" of a pencil) and makes graphite an excellent solid lubricant.

Solved Example 3

Why does silicon form silicones

(R_{2} S i O)_{n}

readily, but carbon does not form similar

(R_{2} C O)_{n}

polymers?

Solution:

$R_{2} S i O$ actually exists as a polymer because Si does not form $S i = O$ multiple bonds effectively (poor $3 p$ - $2 p$ overlap). Each Si prefers to make four $σ$ bonds instead, so it uses two $σ$ bonds to R groups and two $σ$ bonds to bridging O atoms, giving the long $\dots - O - S i R_{2} - O - S i R_{2} - \dots$ chain. Carbon, being small, forms strong $C = O$ $π$ bonds; the analogous unit $R_{2} C = O$ is a stable ketone, a discrete monomer. Ketones do not polymerise via $C = O$ because that would require breaking the strong $C = O$ $π$ bond, which is not thermodynamically favoured. This is a beautiful illustration of the anomalous behaviour of carbon (multiple bonding).

Common Mistakes to Avoid

Watch out

Do not write $C O_{2}$ and $S i O_{2}$ as similar because both are "dioxides". They differ fundamentally: CO $_{2}$ is a discrete linear molecule with $C = O$ double bonds; SiO $_{2}$ is a 3D network with Si-O single bonds. Different physical states (gas vs solid) follow directly from this.
Never write PbO $_{2}$ as a base – it is amphoteric (dissolves in both acids and alkalis) and is a strong oxidising agent because $P b^{4 +}$ is unstable.
$C C l_{4}$ does not hydrolyse in water; $S i C l_{4}$ does. The reason is availability of $d$ -orbitals on Si to accept the incoming water lone pair – C has no such $d$ -orbitals in its valence shell.
Silicones are not "silicon ketones". The name is analogy only – the actual repeat unit $R_{2} S i O$ contains Si-O-Si single bonds, not $S i = O$ double bonds.
In fullerene $C_{60}$ , all $60$ carbons are $s p^{2}$ hybridised, but only the $12$ pentagons are strained. The bonds shared between two hexagons (double bonds) are shorter ( $139$ pm) than those between a hexagon and a pentagon (single bonds, $145$ pm). Do not mark $C_{60}$ as fully aromatic like benzene.
Zeolites are not just "silicates". They are aluminosilicates – some Si is replaced by Al, and cations balance the charge deficit. Without the Al substitution, there would be no exchangeable cations and no ion-exchange or catalytic action.

Frequently Asked Questions

Q1. What is catenation and why does carbon show it more than silicon?

Catenation is the ability of atoms of the same element to link together in long chains or rings. Carbon shows exceptional catenation because the $C - C$ bond enthalpy ( $\sim 348$ kJ mol $^{- 1}$ ) is much higher than $S i - S i$ ( $\sim 297$ kJ mol $^{- 1}$ ), small size gives excellent orbital overlap, and there are no valence lone pairs to repel neighbouring atoms. As a result, carbon forms millions of stable chain and ring compounds – the entire field of organic chemistry. Silicon's catenation is limited to $\sim 8$ atoms in silanes.

Q2. Why does the $+ 2$ oxidation state become more stable down Group 14?

Because of the inert pair effect – the $n s^{2}$ pair of electrons becomes increasingly reluctant to participate in bonding in the heavier atoms (Sn, Pb) due to poor $d$ and $f$ shielding. So Pb prefers to keep its $6 s^{2}$ pair inert and use only the $6 p^{2}$ electrons, giving $P b^{2 +}$ . Conversely, carbon and silicon readily use all four valence electrons and dominate in the $+ 4$ state.

Q3. What are the three main allotropes of carbon and how do they differ?

Diamond ( $s p^{3}$ , 3D tetrahedral network), graphite ( $s p^{2}$ , 2D hexagonal layers with delocalised $π$ system), and fullerene $C_{60}$ ( $s p^{2}$ , 0D cage of $12$ pentagons + $20$ hexagons). Diamond is the hardest natural material and an electrical insulator. Graphite is soft, greasy and conducts electricity along its sheets. Fullerene is a molecular solid that sublimes readily and dissolves in benzene.

Q4. Why is $C O_{2}$ a gas but $S i O_{2}$ is a solid?

Carbon's small size lets it form strong $p π$ - $p π$ double bonds with oxygen ( $C = O$ ), giving the discrete linear molecule $O = C = O$ . Weak van der Waals attractions between these molecules keep $C O_{2}$ gaseous at room temperature. Silicon is much larger; $S i 3 p$ overlaps poorly with $O 2 p$ , so $S i = O$ double bonds don't form. Each Si single-bonds to four O and each O bridges two Si, generating a giant 3D covalent network that only melts near $1983$ K.

Q5. Why is CO a good ligand for transition metals?

CO has a lone pair on carbon that donates to an empty metal orbital ( $σ$ donation), and empty $π^{*}$ molecular orbitals that accept electron density back from filled metal $d$ -orbitals ( $π$ acceptance). This synergic $σ$ -donation/ $π$ -back-donation locks CO tightly to transition metals, giving rise to stable metal carbonyls like $N i (C O)_{4}$ , $F e (C O)_{5}$ and $C r (C O)_{6}$ . This bonding is also why CO is so toxic – it binds haemoglobin more strongly than $O_{2}$ .

Q6. What is a silicate and how are silicates classified?

Silicates are compounds built from $[S i O_{4}]^{4 -}$ tetrahedral units. They are classified by how many corner oxygens each tetrahedron shares with its neighbours: (i) orthosilicates (0 shared, isolated units, e.g. zircon $Z r S i O_{4}$ ); (ii) pyrosilicates (1 shared, dimers); (iii) cyclic silicates (2 shared in a ring, e.g. beryl); (iv) chain silicates (2 shared in an infinite line, e.g. pyroxene); (v) sheet silicates (3 shared, 2D layers, e.g. mica, talc); (vi) 3D silicates (4 shared, e.g. quartz, feldspar, zeolites).

Q7. What are silicones and what makes them useful?

Silicones are synthetic polymers with $S i - O - S i$ backbones and organic side groups (usually methyl or phenyl), general repeat unit $(R_{2} S i O)_{n}$ . The $S i - O$ bond is very strong and thermally stable, and the outward-pointing organic groups make the surface hydrophobic. This combination gives silicones exceptional properties: they resist heat, water and chemicals, so they are used as high-temperature lubricants, waterproofing agents, sealants, electrical insulators, and biocompatible medical implants.

Q8. What are zeolites and why are they called molecular sieves?

Zeolites are three-dimensional aluminosilicates in which some Si $^{4 +}$ ions are replaced by Al $^{3 +}$ ; exchangeable cations (Na $^{+}$ , Ca $^{2 +}$ ) sit in cavities to balance the charge. The framework has pores of very well-defined size ( $300 - 1000$ pm) that let small molecules enter and be adsorbed, but exclude larger molecules – hence "molecular sieve". Uses include water softening (ion exchange), petroleum cracking catalysis (e.g. ZSM-5), and drying of organic solvents.

Q9. Why does CCl $_{4}$ not hydrolyse in water while SiCl $_{4}$ does?

Hydrolysis requires the central atom to accept a lone pair from an incoming water molecule. Si has empty $3 d$ -orbitals in its valence shell that can accept this lone pair, forming a transition state that leads to $S i (O H)_{4} + 4 H C l$ . Carbon has no accessible $d$ -orbitals in its valence shell ( $2 s$ and $2 p$ only), so it cannot expand its octet, and $C C l_{4}$ resists attack by water. This is why $C C l_{4}$ was once used as a fire extinguisher (it doesn't react with water or air) while $S i C l_{4}$ fumes in moist air.

Previous year questions on Group 14 Elements: The Carbon Family

8 questions from past papers, each with a step-by-step solution.

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1. Introduction and Electronic Configuration

Occurrence

2. Trends in Physical Properties

Atomic size, IE and metallic character

Oxidation states and inert pair effect

3. Catenation and Anomalous Behaviour of Carbon

Catenation

Anomalous behaviour of carbon

4. Allotropes of Carbon

4.1 Diamond

4.2 Graphite

4.3 Fullerenes

5. Chemical Properties of Group 14

Reactivity with water

Reactivity with halogens

Oxides

6. Important Oxides of Carbon and Silicon

6.1 Carbon monoxide, CO

6.2 Carbon dioxide, CO2​

6.3 Why CO2​ is a gas but SiO2​ is a solid

7. Silicon Dioxide (Silica), SiO2​

8. Silicates – the Structural Chemistry

9. Silicones

10. Zeolites

11. Compounds of Tin and Lead (brief)

Common Mistakes to Avoid

Frequently Asked Questions

Q1. What is catenation and why does carbon show it more than silicon?

Q2. Why does the +2 oxidation state become more stable down Group 14?

Q3. What are the three main allotropes of carbon and how do they differ?

Q4. Why is CO2​ a gas but SiO2​ is a solid?

Q5. Why is CO a good ligand for transition metals?

Q6. What is a silicate and how are silicates classified?

Q7. What are silicones and what makes them useful?

Q8. What are zeolites and why are they called molecular sieves?

Q9. Why does CCl4​ not hydrolyse in water while SiCl4​ does?

Previous year questions on Group 14 Elements: The Carbon Family

Ready to master The p-Block Elements: Group 13 & 14?

6.1 Carbon monoxide, $C O$

6.2 Carbon dioxide, $C O_{2}$

6.3 Why $C O_{2}$ is a gas but $S i O_{2}$ is a solid

7. Silicon Dioxide (Silica), $S i O_{2}$

Q2. Why does the $+ 2$ oxidation state become more stable down Group 14?

Q4. Why is $C O_{2}$ a gas but $S i O_{2}$ is a solid?

Q9. Why does CCl $_{4}$ not hydrolyse in water while SiCl $_{4}$ does?