p-Block Elements · TOPIC LESSON
Teach both group families through Lewis acidity, multicentre bonds, carbon networks and silicon–oxygen chemistry.
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BEFORE YOU START
Original educational lesson. Work through the examples before revealing the practice answers. Check the current official syllabus separately; historical question frequency does not establish eligibility or predict future questions.
CONNECT THE CONCEPT TO THE EXAM
These selected examples are separate from the original practice below. Each identifies its exam, year and wording. They do not represent every paper or every selected year.
No source-verified exam questions have been published for this lesson yet. The worked examples and practice below are original teaching material, not past-exam questions.
01 / UNDERSTAND & APPLY
Group 13 has valence configuration ns²np¹ and group 14 ns²np². These patterns motivate common +3 and +4 oxidation states, but heavier members increasingly stabilize states two units lower: +1 in group 13 and +2 in group 14. The relative reluctance of the ns² pair to participate in some bonding is commonly called the inert-pair effect.
Boron is a small covalent metalloid, unlike the more metallic heavier group-13 elements. Carbon is a nonmetal with exceptional catenation and multiple-bond chemistry; silicon and germanium have extended covalent structures, while tin and lead show more metallic behaviour. A family comparison must teach these changes rather than transfer every property of its first element to the whole group.
Small size, high ionization energy, orbital overlap and bond strengths explain why first members are often anomalous. Boron commonly forms electron-deficient compounds rather than simple aqueous B³⁺ ions. Carbon forms strong C–C and effective 2p–2p π bonds, enabling chains, rings and multiple bonds.
Oxidation state is formal bookkeeping and does not establish whether a bond is ionic. BCl₃, CCl₄ and SiCl₄ all have positive central-atom oxidation states but are molecular covalent compounds. Bonding, structure and reaction conditions must be examined separately.
02 / UNDERSTAND & APPLY
BCl₃ has three B–Cl bonds arranged approximately trigonal-planar around boron. In a simple Lewis picture, boron has six electrons in its valence shell and an available acceptor orbital. It can accept a lone pair from a donor such as NH₃ to form an adduct, changing the local boron coordination toward tetrahedral.
A Lewis acid accepts an electron pair; a Lewis base donates one. The adduct BCl₃←NH₃ is therefore formed without a required transfer of a proton. Do not confuse Lewis acidity with the presence of hydrogen in a formula.
Boron trihalide acidity depends on more than the electronegativity of the halogen. Ligand-to-boron π donation, bond strength, steric effects and the chosen donor or medium all affect an observed affinity. The common school-level BF₃ versus BCl₃ comparison is often rationalized by stronger F-to-B π interaction in BF₃, but it should not become an unconditional ranking for every measurement.
BCl₃ hydrolyzes overall as BCl₃+3H₂O → B(OH)₃+3HCl. The reaction changes ligand identity and acid composition, not boron oxidation state, which remains +3. These are explanatory equations for reactive chemicals, not laboratory directions.
ORIGINAL WORKED EXAMPLE
For BCl₃+NH₃ → Cl₃B←NH₃, identify acid/base and the coordination change at boron.
Answer BCl₃ is acid, NH₃ base; boron becomes four-coordinate.
ORIGINAL PRACTICE · 1
No.
Its acidity is electron-pair acceptance, not proton donation.
03 / UNDERSTAND & APPLY
Boric acid is conveniently written B(OH)₃, emphasizing three hydroxyl groups around trigonal boron. In water its weak acidity is commonly represented as acceptance of hydroxide from water: B(OH)₃+2H₂O ⇌ B(OH)₄⁻+H₃O⁺. It is therefore treated as a monobasic Lewis acid in this context, not a tribasic acid merely because its formula contains three H atoms.
The structural change to tetrahedral B(OH)₄⁻ accompanies that equilibrium. The number of hydrogens in a molecular formula does not alone determine acid basicity; the actual aqueous reaction does. Complex formation with suitable polyols can change the apparent acidity by stabilizing borate species.
Borax is commonly written Na₂B₄O₇·10H₂O. This composition notation counts crystal water but is not a complete structural drawing of the borate ion. On acidification, tetraborate can be converted to boric acid: B₄O₇²⁻+2H⁺+5H₂O → 4B(OH)₃. Borax solutions are alkaline because of borate hydrolysis.
Heating hydrated borax removes water and eventually gives a glassy borate mixture. The familiar net anhydrous balance Na₂B₄O₇ → 2NaBO₂+B₂O₃ is useful school-level bookkeeping. The borax-bead test involves coloured metal-containing borate species; the colour is not an intrinsic universal property of pure borax.
B(OH)₃+2H₂O ⇌ B(OH)₄⁻+H₃O⁺
Aqueous Lewis-acid description; one hydronium formed per boric-acid unit.
ORIGINAL WORKED EXAMPLE
How many moles H⁺ are required by the displayed tetraborate balance for 0.030 mol B₄O₇²⁻, and how much B(OH)₃ forms?
Answer 0.060 mol H⁺ consumed; 0.120 mol B(OH)₃ produced.
ORIGINAL PRACTICE · 1
Its relevant equilibrium generates one H₃O⁺ while forming B(OH)₄⁻.
Acid behaviour depends on the reaction and structure, not the count of all hydrogen atoms.
04 / UNDERSTAND & APPLY
B₂H₆ has twelve valence electrons: two borons contribute six and six hydrogens contribute six. Four terminal B–H bonds use eight electrons. The remaining four electrons form two B–H–B bridge bonds, each a three-centre two-electron interaction.
Each bridging hydrogen participates in one delocalized bonding interaction spanning both borons; it does not make two independent ordinary two-electron bonds. The term banana bond is a visual nickname for the curved bridge interaction, not a different electron-count rule.
In the familiar structural picture, four terminal hydrogens and the two borons lie in one plane, while two bridging hydrogens lie above and below it. A flat connectivity sketch should state that it is a projection. The terminal and bridging B–H contacts are not equivalent.
Hydrolysis is represented overall by B₂H₆+6H₂O → 2B(OH)₃+6H₂. This strongly reactive chemistry must not be confused with a safe handling procedure. Diborane is hazardous; here the equation is used only for atom and mole accounting.
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ORIGINAL WORKED EXAMPLE
If a student draws eight separate two-centre two-electron B–H bonds in B₂H₆, how many electrons does the drawing require, and what correct accounting replaces it?
Answer The incorrect drawing needs 16 electrons; the correct terminal-plus-bridge model uses 12.
ORIGINAL PRACTICE · 1
Two.
Each B–H–B bridge contains two bonding electrons distributed over three centres; four terminal B–H bonds are separate.
05 / UNDERSTAND & APPLY
Aluminium usually forms +3 compounds. A thin oxide layer can passivate the metal under suitable conditions; reactivity is therefore not predicted from its electrochemical tendency alone. An oxide film is a kinetic barrier, not evidence that aluminium is a noble metal.
Al₂O₃ and Al(OH)₃ are amphoteric. Al(OH)₃ reacts with acid to yield hydrated Al³⁺ and with excess aqueous hydroxide to form aluminate, commonly represented as [Al(OH)₄]⁻. Amphoterism means reaction with both acids and bases; it does not mean an oxide is neutral in every medium.
Anhydrous aluminium chloride can form Al₂Cl₆ through chloride bridges, reducing the electron deficiency at aluminium. Its bridging description should not be copied from diborane: chloride supplies lone-pair donation, whereas the borane bridges are electron-deficient three-centre two-electron bonds. Conditions affect the molecular and ionic species present.
Down group 13 the +1 state becomes more important, with Tl(I) especially stable relative to Tl(III) in many settings. Tl(III) can consequently be oxidizing. The trend is relative stability, not a prohibition on compounds in the higher state.
Al(OH)₃+3H⁺ → Al³⁺+3H₂O
Acid dissolves aluminium hydroxide; aqueous Al³⁺ is hydrated.
Al(OH)₃+OH⁻ → [Al(OH)₄]⁻
Excess aqueous base forms a soluble hydroxo complex.
ORIGINAL PRACTICE · 1
It forms soluble [Al(OH)₄]⁻ under suitable aqueous conditions.
The dissolution demonstrates amphoteric/hydroxo-complex chemistry, not reduction of Al(III).
06 / UNDERSTAND & APPLY
Allotropes are different structural forms of the same element. Diamond has an extended three-dimensional network with each carbon bonded tetrahedrally to four neighbours. Strong covalent connectivity in all directions contributes to hardness. Its valence electrons are largely tied into bonding states, so pure diamond is an electrical insulator, although defects or doping can change behaviour.
Graphite consists of sheets in which each carbon forms three approximately trigonal-planar σ bonds. Remaining p-derived electrons form delocalized π states within the layers, supporting electrical conduction especially along them. Relatively weak interactions between layers permit sliding; softness does not mean the in-layer C–C bonds are weak.
Graphene is one such extended carbon layer. Fullerenes are closed cage molecules, with C₆₀ containing pentagons and hexagons; they are not infinite diamond networks. Carbon nanotubes have rolled-sheet-like structures whose electronic behaviour depends on structure. Do not assign every carbon allotrope the same conductivity merely because all contain carbon.
The empirical formula alone cannot predict hardness or electrical response: both diamond and graphite are elemental carbon. Explain coordination, dimensionality and electronic delocalization. A drawing of one hexagon is a local fragment, not a complete graphite molecule with a fixed molecular mass.
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ORIGINAL WORKED EXAMPLE
A solid has three σ bonds per carbon in layers and mobile π electrons. Another has four σ bonds per carbon in a 3D network. Assign graphite/diamond and explain conductivity.
Answer Graphite is the layered conductor; diamond is the tetrahedral-network insulator.
ORIGINAL PRACTICE · 1
Layers can slide because interlayer interactions are much weaker than the in-layer covalent network.
Mechanical response must distinguish directions and bonding types.
ORIGINAL PRACTICE · 2
No.
C₆₀ is a discrete fullerene cage molecule; its molecular solid has different connectivity.
07 / UNDERSTAND & APPLY
Carbon monoxide is a neutral oxide in ordinary acid–base classification, but it is not chemically inert. It can act as a reducing agent and as a ligand through carbon. CO₂ is an acidic oxide that reacts with bases and participates in aqueous carbonic-acid/bicarbonate/carbonate equilibria. Neither oxidation-state assignment alone predicts the entire chemistry.
CO has carbon in +2 and CO₂ carbon in +4. Oxidation of CO to CO₂ transfers two electrons per carbon. Both gases are molecular; CO₂ is linear and its bond dipoles cancel, whereas its dissolution and reaction with water still produce acid–base chemistry. Nonpolar does not mean unable to react with water.
CO₂ gives CaCO₃ precipitation with limewater: CO₂+Ca(OH)₂ → CaCO₃↓+H₂O. Further CO₂ in water can dissolve that carbonate by forming bicarbonate: CaCO₃+CO₂+H₂O → Ca²⁺+2HCO₃⁻. The changing observation reflects sequential equilibria and reagent amount, not a contradiction.
Many metal carbonates decompose on heating to oxide and CO₂, but stability varies with the cation. Do not universalize calcium carbonate behaviour to every group-1 carbonate. For acid reaction, CO₃²⁻+2H⁺ → CO₂+H₂O provides a widely useful net balance.
ORIGINAL WORKED EXAMPLE
What amount of CO₂ forms when 0.045 mol CO₃²⁻ reacts completely with sufficient acid? How much H⁺ is consumed?
Answer 0.045 mol CO₂ and 0.090 mol H⁺.
ORIGINAL PRACTICE · 1
It converts carbonate into soluble bicarbonate species under the stated aqueous conditions.
The second reaction consumes CaCO₃; the initial precipitation and later dissolution have different reagent ratios.
08 / UNDERSTAND & APPLY
SiO₂ is an extended covalent network in common crystalline forms: each silicon is surrounded tetrahedrally by four oxygens, and each bridging oxygen connects two silicons. Counting shared oxygen atoms gives SiO₂, not isolated SiO₄ molecules with no links. Strong Si–O bonding and network connectivity account for properties very different from molecular CO₂.
Silicates are built from SiO₄ tetrahedral units with different degrees of oxygen sharing. Isolated SiO₄⁴⁻ units, chains, sheets and 3D frameworks have different O:Si ratios and charges. Shared oxygen must be counted only once for the overall formula.
Silicones have Si–O–Si backbones bearing organic groups on silicon. Their flexibility and useful thermal/water-repellent properties depend on polymer structure; they are neither elemental silicon nor ordinary silica. Zeolites are porous aluminosilicate frameworks; substitution of Al for Si creates negative framework charge balanced by cations that may undergo ion exchange.
SiCl₄ hydrolyzes readily, with Si–O bond formation and accessible reaction pathways contributing to the contrast with the ordinary hydrolytic resistance of CCl₄. Avoid the obsolete claim that strong participation of empty silicon 3d orbitals is required to explain the chemistry. Bond polarity, energetics and coordination pathways provide a more defensible account.
SiCl₄+2H₂O → SiO₂+4HCl
Overall idealized hydrolysis balance; actual aqueous silica may be hydrated or polymeric.
ORIGINAL PRACTICE · 1
2: 1.
Four contacts per Si divided by two Si sharing each oxygen gives two oxygen atoms per silicon.
09 / UNDERSTAND & APPLY
The relative stability of +2 increases down group 14. Sn(II) is often reducing because oxidation to Sn(IV) is accessible. Pb(IV) is often oxidizing because reduction to Pb(II) is favourable in many media. These are useful representative tendencies, not a replacement for the actual reaction potential and ligands.
PbCl₄ is less stable than the lighter tetrachlorides and can decompose to PbCl₂+Cl₂. The central atom is reduced from+4 to+2 while chloride is oxidized from−1 to 0. This example ties the inert-pair trend to a balanced redox process.
Acid–base behaviour of group 14 oxides changes down the group: carbon and silicon dioxides are acidic in suitable reactions, while heavier oxides can be amphoteric or more basic. CO is a neutral-oxide exception that should be named explicitly. “All group 14 oxides are acidic” is too broad.
Catenation is strongest for carbon in the familiar family chemistry because robust C–C bonding supports extensive chains and rings. Silicon also catenates, but Si–Si bonding and reactivity give a different range of stability. Neither the existence of silicon chains nor carbon’s dominant range should be turned into an all-or-none rule.
10 / UNDERSTAND & APPLY
For group 13, connect three valence electrons to electron deficiency, Lewis acidity, donor adducts and borane bridges, then include aluminium amphoterism and the heavier +1 trend. For group 14, connect four valence electrons to carbon networks, silicon–oxygen frameworks and the increasing importance of +2 toward lead.
For structural questions, count electrons in diborane and shared atoms in silicates. For allotropes, explain coordination and dimensionality rather than listing properties alone. For reaction questions, preserve conditions and distinguish acid–base changes, hydrolysis and oxidation-state changes.
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