Some p-Block Elements (Groups 13 and 14)NEET Chemistry · Class 11 · NCERT Chapter 12

Group 13: Boron Family — Overview

Group 13 contains five elements: boron (B), aluminium (Al), gallium (Ga), indium (In), and thallium (Tl). Together they are called the boron family or the p-block Group 13 elements. Their valence shell configuration is $n{s}^{2}n{p}^1$, so they have three valence electrons and can show a maximum oxidation state of +3.

Electronic Configurations

Because Ga, In, and Tl have filled d (and f for Tl) orbitals just below their valence shell, their nuclear charges are poorly shielded. This leads to some unexpected irregularities in trends, particularly in atomic radii and ionisation enthalpies.

Trends in Properties of Group 13

Atomic and Ionic Radius

Atomic radii generally increase down the group: B < Al < In > Tl. But notice the anomaly: Ga has a smaller atomic radius than Al. This happens because Ga has 10 d-electrons that shield the nucleus poorly, so the effective nuclear charge felt by the outer electrons is higher, pulling them in. Similarly, Tl has a smaller radius than expected because of the contraction caused by filled 4f and 5d subshells.

Ionisation Enthalpy

The first ionisation enthalpy decreases from B to Al (as expected with increasing size), but then increases from Al to Tl (again due to poor shielding by d and f electrons). The sum of the first three ionisation enthalpies determines whether the +3 state is energetically accessible. For Tl, this energy is too high, so it prefers to lose only one electron and exist as Tl⁺ (inert pair effect).

Oxidation States

All Group 13 elements show a +3 oxidation state. However, the +1 oxidation state becomes progressively more stable down the group due to the inert pair effect. In thallium, Tl⁺ is more stable than Tl³⁺. The ns² electrons (the "inert pair") become difficult to remove because relativistic effects and poor d/f shielding stabilise them.

• B: only +3 (covalent compounds; B³⁺ ion does not exist due to tiny size and high charge density)
• Al: predominantly +3 (ionic in Al²O₃, AlCl₃ which is covalent)
• Ga: +1 and +3
• In: +1 and +3
• Tl: +1 is more stable than +3

Nature of Oxides and Hydroxides

The oxides and hydroxides change character from acidic to amphoteric to basic down the group:

• B₂O₃: acidic oxide (forms boric acid with water)
• Al₂O₃, Al(OH)₃: amphoteric (reacts with both acids and alkalis)
• Ga₂O₃, In₂O₃, Tl₂O₃: increasingly basic in character

Aluminium hydroxide dissolving in NaOH:

Electron Deficiency

Group 13 elements have only 3 valence electrons but 4 available orbitals (one s + three p). After forming three bonds, one vacant orbital remains. This makes their compounds Lewis acids (electron-pair acceptors). BF₃, BCl₃, and AlCl₃ all readily accept electron pairs.

Anomalous Properties of Boron

Boron differs significantly from the rest of Group 13 because of its very small atomic size and very high ionisation enthalpy.

• Covalent character: B³⁺ is too small and has too high a charge density to exist as a free ion. All boron compounds are covalent. In contrast, Al forms partially ionic compounds (e.g., Al₂O₃ is ionic).
• Non-metal: Boron is the only non-metal in Group 13. Al, Ga, In, and Tl are all metals.
• Network structure: Boron has a very high melting point (2300 °C) due to a covalent network solid structure. The other Group 13 elements are soft metals with low melting points (Ga melts just above room temperature at 30 °C).
• No reaction with water or dilute HCl: Boron does not react with water or dilute HCl under normal conditions. Aluminium reacts with dilute HCl:
  
  $$2\text{Al}+6\text{HCl}\to 2{\text{AlCl}}_3+3{\text{H}}_2\uparrow$$
• Diagonal relationship with Si: Boron shows a diagonal relationship with silicon (Group 14). Both form covalent hydrides (B₂H₆ and SiH₄), both form acidic oxides (B₂O₃ and SiO₂), both are semiconductors, and both form halides that fume in moist air.
• No d-orbitals in the valence shell: Boron is restricted to a maximum coordination number of 4 (forming species like BF₄⁻). Heavier Group 13 elements can expand their coordination number by using d-orbitals.

Important Compounds of Boron

Diborane, B₂H₆

Diborane is the simplest stable boron hydride. It is a colourless, toxic gas at room temperature (bp −92.5 °C). Its structure is peculiar because it has fewer electrons than needed for conventional bonds.

Structure: Diborane has two types of B-H bonds:

• 4 terminal B-H bonds (normal 2c-2e bonds, two on each B)
• 2 bridging B-H-B bonds (3-centre 2-electron bonds, also called banana bonds or bridge bonds)

In each bridge bond, one electron pair is shared across three atoms (B, H, B). The four terminal H atoms and both B atoms lie in one plane. The two bridging H atoms are above and below that plane. The B-H-B angle in the bridge is about 97°. There are only 12 valence electrons in diborane (3 from each B + 1 from each of 6 H = 12), which is enough for 6 bonds but the molecule needs 8 bonds (4 terminal + 4 bridging = 8 bond positions). The 3c-2e bonding resolves this.

Preparation:

Reactions:

• With water (hydrolysis): ${\text{B}}_2{\text{H}}_6+6{\text{H}}_2\text{O}\to 2{\text{B(OH)}}_3+6{\text{H}}_2$
• With ammonia (at low temperature): forms B₂H₆·2NH₃ and ultimately borazine (B₃N₃H₆), the "inorganic benzene" at higher temperatures
• Combustion: ${\text{B}}_2{\text{H}}_6+3{\text{O}}_2\to {\text{B}}_2{\text{O}}_3+3{\text{H}}_2\text{O}$

Borax, Na₂B₄O₇·10H₂O

Borax (sodium tetraborate decahydrate) is the most important compound of boron. It occurs naturally in dry lake beds.

Structure: In borax, the tetraborate unit is ${\text{[B}}_4{\text{O}}_5{\text{(OH)}}_4{]}^{2-}$. It contains two sp²-hybridised B atoms (trigonal planar) and two sp³-hybridised B atoms (tetrahedral).

Uses and reactions:

• Borax dissolves in hot water to give an alkaline solution (pH ≈ 9) because it hydrolyses:
  
  $${\text{Na}}_2{\text{B}}_4{\text{O}}_7+7{\text{H}}_2\text{O}\to 2\text{NaOH}+4{\text{H}}_3{\text{BO}}_3$$
• Borax bead test: When heated on a platinum wire in a flame, borax first swells (loses water of crystallisation), then forms a transparent glassy bead of boron oxide (B₂O₃). This bead dissolves metal oxides forming characteristic coloured metal metaborates. Used to identify metal ions in qualitative analysis.
• Used as a cleaning agent (laundry), pH buffer, in glass and enamel manufacturing, and as a laboratory reagent.

Orthoboric Acid, H₃BO₃ (or B(OH)₃)

Boric acid is a white, flaky, crystalline solid. It is a weak monobasic acid.

Structure: Each B(OH)₃ unit is planar (sp² hybridised at B). The units form infinite sheets in the crystal through hydrogen bonding between O-H groups of adjacent molecules. This layered structure is why boric acid has a soapy feel (layers slide).

Important point: Boric acid does NOT donate a proton like a normal acid. Instead, it acts as a Lewis acid by accepting OH⁻ from water:

This is a Lewis acid mechanism (not Bronsted-Lowry). The boron expands its coordination from 3 to 4 by accepting OH⁻.

Preparation from borax:

Uses: Antiseptic (dilute solution), fireproofing, in the glass industry (borosilicate glass), and in the manufacture of cosmetics.

Important Compounds of Aluminium

Aluminium Oxide, Al₂O₃ (Alumina)

Alumina is an amphoteric oxide. It reacts with both acids and bases:

It exists in several crystalline forms (polymorphs). α-Al₂O₃ (corundum) is very hard (used as abrasive). Ruby and sapphire are forms of corundum coloured by trace metal impurities (Cr³⁺ gives red, Ti⁴⁺/Fe²⁺ gives blue).

Aluminium Chloride, AlCl₃

Anhydrous AlCl₃ is a covalent compound (not ionic, despite Al being a metal). It exists as a dimer Al₂Cl₆ in the vapour phase. It is a strong Lewis acid used as a catalyst in Friedel-Crafts reactions.

When AlCl₃ is dissolved in water, it hydrolyses to give an acidic solution:

Alums

Alums are double sulphates of the formula ${\text{M}}^{+}{\text{M}}^{3+}({\text{SO}}_4{)}_2\cdot 12{\text{H}}_2\text{O}$, where M⁺ is a univalent metal ion (K⁺, Na⁺, NH₄⁺) and M³⁺ is a trivalent metal ion (Al³⁺, Cr³⁺, Fe³⁺).

The most common alum is potash alum: KAl(SO₄)₂·12H₂O. Uses include water purification (the Al³⁺ ion hydrolyses to form Al(OH)₃, a colloid that flocculates suspended impurities), as a mordant in dyeing, in leather tanning, and in baking powder.

Reactions of Aluminium Metal

• With dilute HCl: $2\text{Al}+6\text{HCl}\to 2{\text{AlCl}}_3+3{\text{H}}_2\uparrow$
• With hot concentrated NaOH: $2\text{Al}+2\text{NaOH}+2{\text{H}}_2\text{O}\to 2{\text{NaAlO}}_2+3{\text{H}}_2\uparrow$
• Al is passive (unreactive) with concentrated HNO₃ because a thin, hard oxide layer (Al₂O₃) forms on the surface and prevents further reaction. This is called passivation.
• Thermite reaction: 2Al + Fe₂O₃ → Al₂O₃ + 2Fe (large amount of heat released; used in welding railway tracks)

Group 14: Carbon Family — Overview

Group 14 contains: carbon (C), silicon (Si), germanium (Ge), tin (Sn), and lead (Pb). Their valence shell configuration is $n{s}^{2}n{p}^2$, giving them four valence electrons and a maximum oxidation state of +4.

Carbon is unique in its ability to form long chains (catenation) due to the high C-C bond energy (347 kJ/mol). Silicon can also catenate but to a much lesser extent (Si-Si bond is weaker). This is the basis of organic chemistry.

Electronic Configurations

Trends in Properties of Group 14

Atomic Radius and Ionisation Enthalpy

Atomic radius increases down the group: C < Si < Ge < Sn < Pb (but the increase from Si to Ge is small due to d-electron contraction; from Sn to Pb is also small due to f-electron and relativistic effects).

Ionisation enthalpy decreases down the group, but with irregularities at Ge (because of the 3d¹⁰ electrons increasing effective nuclear charge slightly).

Oxidation States and Inert Pair Effect

All Group 14 elements show +4 and +2 oxidation states. The +2 state becomes more stable as you go down the group (inert pair effect). Lead strongly prefers the +2 state:

• Carbon and Silicon: +4 is the stable state; +2 exists in CO (carbon monoxide)
• Germanium: GeO (GeO₂ is more stable)
• Tin: both SnO and SnO₂ exist; SnCl₂ is a reducing agent
• Lead: PbO (Pb²⁺) is more stable than PbO₂ (Pb⁴⁺). PbO₂ is a strong oxidising agent precisely because it wants to become Pb²⁺ by gaining electrons.

Catenation

Catenation is the ability of an atom to bond with atoms of the same element to form long chains or rings. Carbon shows the highest degree of catenation among all elements because:

• C-C bond energy (347 kJ/mol) is high
• Carbon atoms are small so each bond is strong
• Carbon does not have d-orbitals in its valence shell, so it does not react with water

Silicon also catenates, but Si-Si bond energy is lower (226 kJ/mol) and Si has d-orbitals that make it susceptible to attack by water and nucleophiles. The maximum silane chain is about 6-8 Si atoms.

Nature of Hydrides

Hydrides of Group 14 (methane CH₄, silane SiH₄, germane GeH₄, stannane SnH₄, plumbane PbH₄) are all covalent. Their stability decreases down the group (C-H bond is stronger than Si-H, Ge-H, etc.). Methane is very stable; plumbane barely exists.

Nature of Oxides

Select a p-block element to see its highest oxide formula, nature (acidic/basic/amphoteric), and reactions with acid and base.

The character of the oxides changes from acidic (CO₂, SiO₂) to amphoteric (GeO₂, SnO₂) to basic (PbO). PbO₂ is amphoteric. Lead(II) oxide (PbO, litharge) is the most commonly encountered form.

Allotropes of Carbon

Allotropy is the existence of an element in two or more physical forms in the same physical state. Carbon has three main crystalline allotropes: diamond, graphite, and fullerenes. Amorphous forms (coal, charcoal, coke) also exist.

Diamond

• Each carbon is sp³ hybridised
• Forms a 3D tetrahedral network — each C bonded to 4 other C atoms (coordination number = 4). All four bonds are equivalent sigma bonds.
• Hardest natural substance (Mohs hardness = 10). Used as a cutting and abrasive tool.
• Non-conductor of electricity (all 4 valence electrons are used in C-C bonds; no free electrons)
• Very high melting point (3550 °C) due to strong covalent network
• Good conductor of heat (due to strong covalent bonds that transmit vibrations efficiently)
• Transparent to light; high refractive index (total internal reflection makes diamonds sparkle)

Graphite

• Each carbon is sp² hybridised
• Forms 2D hexagonal layers (honeycomb lattice). Each C is bonded to 3 other C atoms in the layer. The fourth valence electron of each carbon is delocalised in a $\pi$ system across the entire layer.
• Good conductor of electricity (due to delocalised $\pi$ electrons). Used as electrodes in electrolytic cells and in batteries.
• Soft and slippery: the van der Waals forces between layers are weak, so layers slide easily. Used as a lubricant and in pencils.
• Black and opaque (absorbs light)
• Less dense than diamond (2.2 g/cm³ vs 3.5 g/cm³ for diamond)
• The interlayer distance in graphite is 3.35 Å (too large for covalent bonding; held by weak van der Waals forces)

Fullerenes

• Discovered in 1985 (Kroto, Curl, Smalley — Nobel Prize in Chemistry 1996). The most well-known fullerene is C₆₀ (Buckminsterfullerene), nicknamed "buckyball."
• C₆₀ has a structure resembling a football (soccer ball): 20 hexagonal + 12 pentagonal faces. Each carbon is bonded to 2 others via C-C single bonds (within hexagons and pentagons) and to one via a C=C double bond (between hexagons). All C atoms are sp² hybridised.
• No dangling bonds; completely closed cage structure
• Semiconductors: can be doped with metals to become superconductors
• Higher fullerenes include C₇₀, C₈₀, etc. Carbon nanotubes are rolled-up graphene sheets.
• All fullerenes can be solubilised in organic solvents (unlike diamond or graphite), making them processable in solution.

Important Compounds of Carbon

Carbon Monoxide, CO

CO is a colourless, odourless, and highly poisonous gas. It is produced by incomplete combustion of carbon-containing fuels.

Structure: The C-O bond in CO has a bond order of 3 (one sigma + one pi + one dative pi from O to C). It is isoelectronic with N₂.

Toxicity: CO binds to haemoglobin (Hb) to form carboxyhaemoglobin (HbCO) with an affinity about 200 times greater than O₂. This prevents O₂ from binding, causing suffocation at the cellular level.

Chemical properties:

• Strong reducing agent: used in the extraction of metals (blast furnace — reduces iron ore):
  
  $${\text{Fe}}_2{\text{O}}_3+3\text{CO}\to 2\text{Fe}+3{\text{CO}}_2$$
• Reacts with Cl₂ to form phosgene (COCl₂), a war gas:
  
  $$\text{CO}+{\text{Cl}}_2\stackrel{\text{UV}}{\to }{\text{COCl}}_2$$
• With NaOH (under pressure): forms sodium formate HCOONa (Kolbe's synthesis of formic acid)
• Combines with transition metals to form metal carbonyls, e.g., Ni(CO)₄ (Mond process for purifying nickel)

CO is neutral to litmus and does not react with water directly.

Carbon Dioxide, CO₂

CO₂ is a colourless, odourless, non-toxic gas. It is a linear molecule (O=C=O) where carbon is sp hybridised.

Properties:

• Acidic oxide: dissolves in water to form carbonic acid (H₂CO₃), a diprotic weak acid
• Reaction with lime water: ${\text{Ca(OH)}}_2+{\text{CO}}_2\to {\text{CaCO}}_3\downarrow +{\text{H}}_2\text{O}$ (milky precipitate; excess CO₂ clears it by forming soluble Ca(HCO₃)₂)
• Heavier than air; does not support combustion (used in fire extinguishers)
• CO₂ can act as an oxidising agent at very high temperatures:
  
  $${\text{CO}}_2+\text{C}\to 2\text{CO (Boudouard reaction)}$$
• CO₂ is a greenhouse gas: it absorbs infrared radiation from Earth's surface and radiates it back, contributing to global warming.
• Solid CO₂ (dry ice, −78.5 °C) sublimes directly without melting.

Silicon and Silicates

Silicon Dioxide, SiO₂

SiO₂ (silica, quartz) is a covalent network solid in which each Si is bonded to 4 O atoms in a tetrahedral arrangement (Si is sp³ hybridised). Unlike CO₂ (molecular gas), SiO₂ is a macromolecular solid. This is because Si uses d-orbitals to form Si-O-Si bridges and cannot form strong Si=O double bonds (unlike C which forms C=O).

SiO₂ reacts with HF (only) among acids:

SiO₂ reacts with alkalis and basic oxides:

SiO₂ does not react with water directly (unlike CO₂). It is used in making glass, ceramics, and in electronics (as semiconductor substrate).

Silicates

Silicates are salts of silicic acid. The basic structural unit is the ${\text{SiO}}_4^{4-}$ tetrahedron. Different silicates arise by sharing different numbers of oxygen atoms between adjacent tetrahedra:

You only need to know the general principle: more O-sharing = lower O:Si ratio = more compact structure. For NEET, remember the two extremes: orthosilicate (SiO₄⁴⁻, no sharing) and framework silicate (SiO₂, complete sharing, 4 O per Si).

Silicones

Silicones are synthetic organosilicon polymers containing Si-O-Si linkages (siloxane bonds). The general repeating unit is ${\text{(R}}_2{\text{SiO)}}_n$, where R is an organic group (methyl, phenyl, etc.).

Properties: heat-resistant, water-repellent, electrical insulators, good lubricants, chemically inert.

Uses: sealants, lubricants, surgical implants (breast implants), electrical insulation, waterproofing textiles, cosmetics.

Sodium Silicate, Na₂SiO₃ (Water Glass)

Made by fusing SiO₂ with Na₂CO₃. Dissolves in hot water to give a viscous solution. Used in preserving eggs (forms impermeable silica gel coating), making silica gel (for drying/adsorption), and in fire-retardant treatments.

Worked NEET Problems

Summary Cheat Sheet

Group 13 Quick Reference

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