The p-Block Elements Groups 13 and 14 covers boron, aluminium, carbon, silicon and their heavier family members, studying how their ns2 np1 and ns2 np2 configurations drive periodic trends, the inert pair effect, and key compounds like borax, diborane, alums, carbon allotropes, silica and silicones. It explains why lower oxidation states win down the group (Tl plus and Pb 2 plus) and why only carbon shows strong catenation and multiple bonding. It is a high-yield inorganic chemistry chapter for NEET, with reliable questions on the inert pair effect, diborane bonding, boron and aluminium compounds, and carbon group trends.
Table of Contents
- Key Concepts - general introduction to p-block, Group 13, Group 14, important compounds
- Weightage in Board & Entrance Exams
- Important Definitions
- Solved Examples
- Important Questions for Board Exams
- Quick Revision Points
Key Concepts
1. General Introduction to the p-Block
The p-block contains elements in which the last electron enters one of the three p-orbitals of their outermost shell. It spans Groups 13 to 18 of the periodic table, and its general outer electronic configuration is ns²np¹⁻⁶.
This block is special because it holds metals, non-metals, and metalloids all together - the only block where all three appear. The non-metallic character is strongest at the top-right and metallic character increases as you go down and to the left.
The Inert Pair Effect
As we move down a group, the tendency of the two ns electrons to remain paired and not participate in bonding increases. This is called the inert pair effect, and it explains why the lower oxidation state becomes more stable down a group (e.g., Tl⁺ is more stable than Tl³⁺, and Pb²⁺ more stable than Pb⁴⁺).
2. Group 13 - The Boron Family
Group 13 consists of boron (B), aluminium (Al), gallium (Ga), indium (In), and thallium (Tl). Their general outer electronic configuration is ns²np¹.
Boron is a non-metal (a metalloid); all the others are metals. The common oxidation state is +3, but due to the inert pair effect the +1 state becomes increasingly stable down the group.
Trends in Group 13
- Atomic radius: increases down the group, but Ga is slightly smaller than Al (poor shielding by 3d electrons).
- Ionisation enthalpy: decreases down the group overall (with minor irregularities).
- Metallic character: increases down the group - B is a metalloid, Tl is a soft metal.
- Oxidation state: +3 dominates at the top; +1 stability rises down the group due to inert pair effect.
- Nature of oxides: B₂O₃ is acidic, Al₂O₃ and Ga₂O₃ are amphoteric, In₂O₃ and Tl₂O₃ are basic.
3. Anomalous Behaviour of Boron
Boron, the first member of Group 13, differs sharply from the rest of its family. This is due to its small size, high ionisation enthalpy, high electronegativity, and the absence of d-orbitals in its valence shell.
- Boron is a non-metal (metalloid); the rest are metals.
- Boron is never found as a B³⁺ ion - it forms only covalent compounds, while others can form ionic compounds.
- The maximum covalency of boron is 4 (no d-orbitals), whereas heavier members can expand their covalency beyond 4.
- Boron forms electron-deficient compounds like BF₃ and diborane (B₂H₆).
Diagonal relationship: Boron resembles silicon (the diagonally placed element of Group 14) more than it resembles aluminium - for example, both B and Si form covalent, polymeric, acidic oxides.
4. Important Compounds of Boron
Borax (Na₂B₄O₇·10H₂O)
Borax is sodium tetraborate decahydrate, a white crystalline solid. Its correct structural formula is Na₂[B₄O₅(OH)₄]·8H₂O, containing two triangular (BO₃) and two tetrahedral (BO₄) units.
- An aqueous solution of borax is alkaline (it hydrolyses), so it acts as a buffer.
- Borax bead test: on heating, borax swells and then melts to a transparent glassy bead of sodium metaborate and boric anhydride (NaBO₂ + B₂O₃), used to identify coloured metal ions.
- Reaction on heating: Na₂B₄O₇·10H₂O → 2NaBO₂ + B₂O₃ + 10H₂O.
Orthoboric Acid (H₃BO₃)
Boric acid is a white, soft, soapy solid with a layered structure in which planar BO₃ units are joined by hydrogen bonds.
- It is a weak monobasic acid - but not a protonic acid. It acts as a Lewis acid by accepting OH⁻ from water: B(OH)₃ + 2H₂O → [B(OH)₄]⁻ + H₃O⁺.
- On heating, it loses water in steps: H₃BO₃ → HBO₂ (metaboric acid) → B₂O₃ (boric anhydride).
- Used as a mild antiseptic and in the glass/ceramics industry.
Diborane (B₂H₆)
Diborane is the simplest boron hydride, a colourless, toxic gas that catches fire spontaneously in air. It is an electron-deficient molecule.
[DIAGRAM: B₂H₆ structure - two boron atoms each bonded to two terminal H atoms by normal covalent bonds, and bridged by two H atoms through three-centre two-electron (banana) bonds.]
- It contains four terminal B–H bonds (normal 2-centre 2-electron bonds) and two bridging B–H–B bonds (3-centre 2-electron “banana” bonds).
- Combustion: B₂H₆ + 3O₂ → B₂O₃ + 3H₂O (highly exothermic).
- With ammonia it gives borazine (B₃N₆H₆), called “inorganic benzene”.
5. Aluminium
Aluminium is the most abundant metal in the Earth’s crust. It is a silvery-white, light, malleable metal showing the +3 oxidation state, and it is the most important member of Group 13 commercially.
- Amphoteric nature: aluminium and its oxide react with both acids and alkalis.
2Al + 6HCl → 2AlCl₃ + 3H₂; 2Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂. - Passivity: a thin, tough oxide layer protects aluminium from further corrosion, so concentrated HNO₃ renders it passive.
- AlCl₃ exists as a dimer (Al₂Cl₆) in the vapour and non-polar solvents, completing aluminium’s octet.
- Used in alloys (duralumin), electrical cables, packaging, and as a reducing agent in the thermite process.
6. Group 14 - The Carbon Family
Group 14 consists of carbon (C), silicon (Si), germanium (Ge), tin (Sn), and lead (Pb). Their general outer electronic configuration is ns²np².
The group shows a clear shift from non-metal to metal: carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are metals. The common oxidation states are +4 and +2.
Trends in Group 14
- Atomic radius: increases down the group; the increase from C to Si is large, then small thereafter.
- Metallic character: increases down the group (C → non-metal, Sn, Pb → metals).
- Oxidation state: +4 is stable at the top; +2 stability rises down the group (Pb²⁺ > Pb⁴⁺) due to the inert pair effect.
- Catenation: the self-linking ability decreases down the group: C ≫ Si > Ge ≈ Sn ≫ Pb. The strong C–C bond explains the millions of carbon compounds.
- Nature of oxides: CO₂ and SiO₂ are acidic, GeO₂ is weakly acidic, SnO₂ and PbO₂ are amphoteric.
7. Anomalous Behaviour of Carbon & Allotropes
Carbon, like boron, differs from the rest of its group because of its small size, high electronegativity, high ionisation enthalpy, and the absence of d-orbitals. Two unique features stand out: its maximum covalency is 4, and it has an exceptional ability for catenation and pπ–pπ multiple bonding.
Allotropes of Carbon
- Diamond: each carbon is sp³ hybridised and bonded tetrahedrally to four others, forming a rigid 3-D network. It is the hardest natural substance and does not conduct electricity.
- Graphite: each carbon is sp² hybridised, forming planar hexagonal layers held by weak van der Waals forces. The delocalised electrons make it a good conductor and a lubricant.
- Fullerenes (e.g., C₆₀): cage-like molecules (“buckyballs”) with both 5- and 6-membered rings; the only pure, neat allotrope.
[TABLE: Diamond is sp³, 3-D, hard, insulator; Graphite is sp², 2-D layers, soft, conductor; Fullerene is sp², spherical cage, molecular solid.]
8. Oxides of Carbon - CO and CO₂
Carbon Monoxide (CO)
Carbon monoxide is a colourless, odourless, highly poisonous gas formed by the incomplete combustion of carbon. It is neutral and a powerful reducing agent.
- It is toxic because it binds to haemoglobin about 300 times more strongly than oxygen, forming carboxyhaemoglobin and blocking oxygen transport.
- It is a good reducing agent, used in metallurgy: Fe₂O₃ + 3CO → 2Fe + 3CO₂.
Carbon Dioxide (CO₂)
Carbon dioxide is a colourless, odourless acidic gas. It is a linear, non-polar molecule (O=C=O) and the main greenhouse gas responsible for global warming.
- It dissolves in water to form weak carbonic acid: CO₂ + H₂O ⇌ H₂CO₃.
- Solid CO₂ (“dry ice”) sublimes directly and is used as a refrigerant.
- Essential for photosynthesis, which keeps the carbon cycle balanced.
9. Silicones
Silicones are synthetic organosilicon polymers containing repeating R₂SiO units, with a backbone of alternating silicon and oxygen atoms (–Si–O–Si–O–) and organic groups attached to silicon.
- They are prepared by the hydrolysis of dialkyl/diaryl dichlorosilanes (R₂SiCl₂) followed by polymerisation.
- They are water-repellent (hydrophobic), heat-resistant, and chemically inert.
- Uses: water-proofing fabrics, lubricants, sealants, electrical insulators, and biomedical implants.
10. Silicates and Zeolites
Silicates
Silicates are compounds in which the basic structural unit is the SiO₄⁴⁻ tetrahedron. These tetrahedra link by sharing oxygen corners to form chains, sheets, or three-dimensional networks.
- When all four corner oxygens are shared, the neutral, giant covalent solid silica (SiO₂) results.
- Examples include feldspar, asbestos, mica, and zeolites.
Zeolites
Zeolites are three-dimensional aluminosilicates with a porous, cage-like structure, formed when some silicon atoms in SiO₄ are replaced by aluminium (giving an AlO₄ unit and a negative charge balanced by cations).
- Used as ion-exchangers to soften hard water (Permutit process).
- ZSM-5 is used as a catalyst to convert alcohols directly into petrol (gasoline).
- Used as molecular sieves and as catalysts in the petrochemical industry.
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Weightage in Board & Entrance Exams
| Exam | Typical Weightage | Most-Tested Areas |
|---|---|---|
| CBSE Board (Class 11) | 6–8 marks | Anomalous behaviour, borax/boric acid, diborane structure, allotropes, silicones |
| JEE Main / Advanced | 1–2 questions | Inert pair effect, structure of diborane & borax, catenation trend |
| NEET | 1–2 questions | Properties of CO/CO₂, oxides’ acidic-basic nature, uses of zeolites & silicones |
[TABLE: Question-type split - VSA (1 mark): definitions & formulae; SA (2–3 marks): anomalous behaviour, diborane bonding, CO toxicity; LA (5 marks): borax structure & bead test, allotropes comparison, silicones preparation.]
Important Definitions
| Term | Definition |
|---|---|
| p-Block element | Element in which the last electron enters a p-orbital; configuration ns²np¹⁻⁶ |
| Inert pair effect | Reluctance of the ns² electron pair to bond, stabilising lower oxidation states down a group |
| Diagonal relationship | Similarity between an element and the one placed diagonally to it (e.g., B and Si) |
| Electron-deficient compound | A molecule with fewer electrons than needed for normal bonds, e.g., B₂H₆, BF₃ |
| Three-centre two-electron bond | A “banana” bond where two electrons bind three atoms, as in the B–H–B bridge of diborane |
| Catenation | Self-linking of like atoms into chains/rings; strongest in carbon |
| Allotropy | Existence of an element in two or more forms differing in physical properties |
| Silicones | Organosilicon polymers with an –Si–O–Si– backbone and organic side groups |
| Zeolite | Porous three-dimensional aluminosilicate used as ion-exchanger and catalyst |
| Amphoteric oxide | An oxide that reacts with both acids and bases, e.g., Al₂O₃ |
Solved Examples
Example 1
Why does boron form only covalent compounds and never the B³⁺ ion?
Answer: Boron has a very small size and very high ionisation enthalpy. Removing three electrons to form B³⁺ needs more energy than is released by lattice/hydration, so it is energetically unfavourable. Hence boron shares electrons and forms only covalent compounds.
Example 2
Explain why diborane (B₂H₆) is called an electron-deficient molecule.
Answer: B₂H₆ has 12 valence electrons but needs more to form eight normal 2-electron bonds. It has four normal terminal B–H bonds and two three-centre two-electron B–H–B bridge bonds. Since there are not enough electrons for all conventional bonds, it is electron-deficient.
Example 3
Why does the +1 oxidation state become more stable than +3 down Group 13?
Answer: Down the group the inert pair effect increases - the ns² electrons become reluctant to participate in bonding. So Tl prefers the +1 state, making Tl⁺ more stable than Tl³⁺.
Example 4
Why is carbon monoxide poisonous?
Answer: CO binds to the haemoglobin of blood about 300 times more strongly than O₂, forming carboxyhaemoglobin. This blocks oxygen transport to the tissues, which can be fatal.
Example 5
Why does carbon show the maximum tendency for catenation in Group 14?
Answer: The C–C bond is exceptionally strong (about 348 kJ/mol) because of carbon’s small size and effective orbital overlap. Bond strength falls down the group (Si–Si, Ge–Ge weaker), so catenation is greatest in carbon - explaining its millions of compounds.
Example 6
What happens when borax is heated strongly? Write the reaction.
Answer: Borax first loses water of crystallisation, swells, and then melts to a clear glassy bead: Na₂B₄O₇·10H₂O → 2NaBO₂ + B₂O₃ + 10H₂O. The bead (sodium metaborate + boric anhydride) is the basis of the borax bead test for coloured metal ions.
Important Questions for Board Exams
1-Mark Questions (VSA)
- Write the general outer electronic configuration of p-block elements.
- What is the inert pair effect?
- Why is boric acid considered a weak monobasic acid?
- Name the hardest allotrope of carbon and give the hybridisation of its carbon atoms.
- Give one important use of zeolites.
2–3-Mark Questions (SA)
- Explain the anomalous behaviour of boron with any three points.
- Describe the structure of diborane and explain the bonding in it.
- Compare the structures and properties of diamond and graphite.
- What are silicones? How are they prepared, and give two uses.
5-Mark Questions (LA)
- Discuss the trends in oxidation state and the nature of oxides down Group 13, explaining them in terms of the inert pair effect.
- Explain the structure of borax and describe the borax bead test with relevant reactions.
- Discuss catenation in Group 14 and explain why carbon shows it to the greatest extent. Compare the acidic/basic nature of the oxides of the group.
Quick Revision Points
- p-block: Groups 13–18; general configuration ns²np¹⁻⁶; holds metals, non-metals, metalloids
- Inert pair effect → lower oxidation state more stable down a group (Tl⁺, Pb²⁺)
- Group 13 (ns²np¹): +3 common; oxide nature B₂O₃ acidic → Al₂O₃ amphoteric → Tl₂O₃ basic
- Boron is anomalous - small size, no d-orbitals, max covalency 4, only covalent compounds; diagonal with Si
- Borax: Na₂B₄O₇·10H₂O → 2NaBO₂ + B₂O₃ on heating; basis of bead test
- Boric acid H₃BO₃ - weak monobasic Lewis acid; layered, H-bonded structure
- Diborane B₂H₆ - electron-deficient; 4 terminal B–H + 2 bridging 3c–2e bonds
- Aluminium - amphoteric, made passive by conc. HNO₃; AlCl₃ dimerises to Al₂Cl₆
- Group 14 (ns²np²): +4 and +2; catenation C ≫ Si > Ge ≈ Sn ≫ Pb
- Carbon allotropes: diamond (sp³, hard, insulator), graphite (sp², conductor, lubricant), fullerene (C₆₀)
- CO - neutral, poisonous (carboxyhaemoglobin), reducing agent; CO₂ - acidic, greenhouse gas
- Silicones - –Si–O–Si– polymers, water-repellent; zeolites - porous aluminosilicates, ion-exchangers/catalysts
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Meet Groups 13 & 14
Two p-block families that share an outer-shell pattern but turn more metallic as you go down.
First members B and C are anomalous: small, electronegative, no d-orbitals.
- Group 13 has 3 valence electrons, Group 14 has 4.
- Metallic character rises down each group (B metalloid → Tl metal; C non-metal → Pb metal).
- B and C have a maximum covalency of 4 (no valence d-orbitals).
Size, IE & the Ga Anomaly
Atoms grow larger down a group, but poor 3d shielding makes gallium smaller than aluminium.
Group 13 IE1 is lower than the Group 2 neighbour (easy np1 vs stable ns2).
- Radius increases down a group, but Ga < Al due to poor 3d-shielding.
- Ionisation enthalpy generally falls down a group, with d-/f-shielding wobbles.
- Tl sits high in the IE1 order because bad 4f/5d shielding grips the 6s pair tightly.
Inert Pair Effect
The outer ns2 pair becomes reluctant to bond, so the lower oxidation state wins down the group.
It is the ns2 pair (not np or d) that turns inert; effect grows down the group.
- Heavy ns2 pair is held tightly and costs too much energy to unpair for weak bonds.
- PbO2 and Tl3⁺ are strong oxidisers; SnCl2 and Ge2⁺ are strong reducers.
- Stability: Tl⁺ > Tl3⁺, Pb2⁺ > Pb4⁺, but Sn4⁺ > Pb4⁺ and Pb2⁺ > Sn2⁺.
Borax & Boric Acid
Borax is an alkaline borate; orthoboric acid is a weak Lewis acid, not a proton donor.
Borax has 2 sp2 + 2 sp3 boron atoms; its aqueous solution is alkaline.
- Bead test: Na2B4O7 → 2NaBO2 + B2O3 gives coloured metaborates with metal salts.
- H3BO3 is weak, monobasic and a Lewis acid (it accepts OH⁻, not donates H⁺).
- Solid H3BO3 is a layered, hydrogen-bonded sheet structure.
Diborane: Banana Bonds
B2H6 has too few electrons for normal bonds, so it bridges with three-centre, two-electron bonds.
Electron-deficient: 8 ordinary B–H bonds would need 16 e⁻ but only 12 exist.
- Two bridge bonds use 2 pairs; the remaining 4 pairs make the 4 terminal B–H bonds.
- Each banana bond glues three atoms (B–H–B) with a single electron pair.
- Boron never exceeds a covalency of 4 (no d-orbitals).
Aluminium & Its Compounds
Aluminium is amphoteric, goes passive in conc. HNO3, and its chloride dimerises to complete the octet.
Both acid and alkali liberate H2; conc. HNO3 forms a passivating Al2O3 layer.
- Anhydrous AlCl3 exists as Cl-bridged dimer Al2Cl6; strong Lewis acid (Friedel–Crafts).
- Boron halide Lewis acidity: BF3 < BCl3 < BBr3 (back-bonding weakest in BBr3).
- Alums are M⁺M3⁺(SO4)2·12H2O, e.g. potash alum KAl(SO4)2·12H2O for water purification.
Carbon’s Three Faces
Pure carbon takes three forms whose properties come entirely from how the atoms are bonded.
Graphite is the thermodynamically most stable allotrope of carbon.
- Diamond is sp3, the hardest natural substance and an insulator.
- Graphite is sp2, conducts electricity (delocalised e⁻) and is a lubricant (sliding sheets).
- Fullerene C60 is the only pure allotrope with no dangling bonds.
Catenation & pπ–pπ Bonds
Carbon’s strong C–C bond gives it unmatched self-linking, and only carbon makes good multiple bonds.
Heavier Group 14 atoms cannot form effective pπ–pπ multiple bonds.
- Catenation order C > Si > Ge > Sn falls as element–element bonds weaken.
- Carbon forms strong C=C, C≡C and C=O via pπ–pπ overlap.
- Why CO2 is a gas (O=C=O) but SiO2 is a giant 3D solid (no Si=O).
Oxides, Silicates & Silicones
Group 14 oxides shift acidic→amphoteric down the group, and silicon builds networks and polymers.
CO is neutral and toxic (binds haemoglobin ~300× stronger than O2); CO2 is acidic, linear.
- SiO2 is a giant covalent network reacting only with NaOH and HF.
- Silicates are corner-sharing SiO44⁻ tetrahedra; all four shared gives quartz.
- Silicones come from hydrolysing R2SiCl2; producer gas = CO + N2, water gas = CO + H2.
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Frequently Asked Questions
The inert pair effect is the reluctance of the outer ns2 electron pair to take part in bonding for heavier p-block elements, so the lower oxidation state becomes more stable down the group. This is why Tl plus is more stable than Tl 3 plus in Group 13 and Pb 2 plus is more stable than Pb 4 plus in Group 14.
Diborane has only 12 valence electrons but would need 16 to form eight ordinary two-centre B-H bonds, so it is electron deficient. It uses two three-centre two-electron bonds (banana bonds) where one electron pair bridges three atoms in a B-H-B bridge, alongside four normal terminal B-H bonds.
Yes, it is a high-yield part of inorganic chemistry and almost always contributes questions to NEET. Common favourites are the inert pair effect, diborane and boric acid, properties of aluminium and AlCl3, and carbon group trends like catenation and oxide acidity.
In diamond each carbon is sp3 hybridised in a rigid 3D tetrahedral network, making it the hardest natural substance and an electrical insulator. In graphite each carbon is sp2 in layered hexagonal sheets with delocalised electrons, so it conducts electricity and acts as a lubricant, and it is the thermodynamically most stable allotrope of carbon.
Orthoboric acid does not donate its own protons; instead it acts as a Lewis acid by accepting a hydroxide ion from water, B(OH)3 plus 2H2O gives [B(OH)4] minus plus H3O plus, releasing only one H3O plus. Because it supplies just one hydronium ion and does so weakly, it is classed as a weak monobasic acid.