The p-Block Elements (Class 11) — NEET notes

Group 13: The Boron Family

Group 13 elements (B, Al, Ga, In, Tl) have ns²np¹ configuration with 3 valence electrons. Boron is anomalous (non-metal, no d-orbitals, forms covalent compounds). Oxidation states are +1 and +3 — the +1 state becomes more stable down the group due to the inert pair effect. Aluminum is the most abundant metal in Earth's crust.

General config: ns^2np^1

Electronic Configuration & Oxidation States

Group 13 elements have one electron in the p-orbital. They show +3 (most common) and +1 oxidation states. The +1 state gains importance down the group due to the inert pair effect.

All Group 13 elements have the general electronic configuration ns^2np^1 . With 3 valence electrons, the expected oxidation state is +3 — and that's indeed the dominant state for B and Al. But things get interesting as you go down.

• Elements: B ( Z=5 ), Al ( Z=13 ), Ga ( Z=31 ), In ( Z=49 ), Tl ( Z=81 )
• B: [He] 2s^2 2p^1 , Al: [Ne] 3s^2 3p^1 , Ga: [Ar] 3d^ 10 4s^2 4p^1
• In: [Kr] 4d^ 10 5s^2 5p^1 , Tl: [Xe] 4f^ 14 5d^ 10 6s^2 6p^1
• Common oxidation states: +3 and +1
• +3 is dominant for B, Al, Ga — but +1 becomes increasingly stable for In and Tl

Trends in Physical & Chemical Properties

Atomic radii generally increase down the group but Ga is smaller than Al due to poor shielding by 3d electrons. IE₁ shows irregularities for the same reason. Boron is a non-metal; Al is a metal; Ga, In, Tl are metals with some peculiarities.

Group 13 trends are full of surprises because of the d-orbital and f-orbital contraction effects. The 3d electrons in Ga and 4f electrons in Tl provide poor shielding, pulling outer electrons closer to the nucleus.

• Boron is the only non-metal in Group 13 — it forms only covalent compounds
• Al, Ga, In, Tl are metals with increasing metallic character down the group
• Gallium has an unusually low melting point ( 29.8°C ) — it melts in your hand!
• Boron exists in several allotropic forms — all contain icosahedral B_ 12 units

Anomalous Behavior of Boron

Boron differs from other Group 13 elements: it's a non-metal, forms only covalent compounds, has no d-orbitals in its valence shell, shows extremely high melting point, and acts as a Lewis acid (electron-deficient) rather than forming B³⁺ ions.

Boron is the black sheep of Group 13. While Al, Ga, In, Tl are metals that form ionic compounds and conduct electricity, boron is a non-metal that forms covalent bonds and acts as a semiconductor.

• Sum of first three IEs of B is very high → B^ 3+ never forms
• Boron forms only covalent compounds — its compounds are electron-deficient
• B has maximum covalence of 4 (e.g., BF_4^- , BH_4^- ) — no d-orbitals to expand
• Boron oxides and hydroxides are acidic (unlike amphoteric Al_2O_3 )
• B_2O_3 is acidic; Al_2O_3 is amphoteric — a key NCERT distinction

Boron & Its Compounds

Borax (Na₂B₄O₇·10H₂O) is the most important boron compound — its bead test identifies colored metal ions. Orthoboric acid (H₃BO₃) has a layered structure held by H-bonds and is a monobasic Lewis acid. Diborane (B₂H₆) contains unique 3-centre-2-electron 'banana bonds'. BN exists in two forms mimicking diamond and graphite.

Na_2B_4O_7 + 7H_2O arrow 2NaOH + 4H_3BO_3

Borax

Borax (Na₂B₄O₇·10H₂O) is the sodium salt of tetraboric acid. On heating it swells (loses water), then fuses to a transparent glass bead of sodium metaborate + boric anhydride. The bead test identifies transition metals by characteristic colors.

Borax ( Na_2B_4O_7 10H_2O ) is the most commercially important boron compound. Its formula can be written as Na_2[B_4O_5(OH)_4] 8H_2O , revealing that it contains two types of boron atoms — two tetrahedral ( sp^3 ) and two trigonal planar ( sp^2 ).

Na_2B_4O_7 10H_2O Na_2B_4O_7 2NaBO_2 + B_2O_3

When heated, borax first swells up (loses water of crystallization), then melts to form a clear glass bead of sodium metaborate ( NaBO_2 ) and boric anhydride ( B_2O_3 ). This bead dissolves metal oxides to give characteristic colored beads — the famous borax bead test.

• Borax solution is basic (hydrolysis of strong base + weak acid salt): pH > 7
• Na_2B_4O_7 + 7H_2O arrow 2NaOH + 4H_3BO_3
• Bead test: B_2O_3 fuses with metal oxide → colored metaborate bead
• Example: CuO + borax bead → blue (copper metaborate); CoO → deep blue; Cr₂O₃ → green
• Borax is used in making glass, ceramics, detergents, and as a flux in metallurgy

Quick recall

How many boron atoms in borax are sp² hybridized and how many are sp³?

2 boron atoms are sp^2 (trigonal planar, bonded to 3 oxygen atoms) and 2 boron atoms are sp^3 (tetrahedral, bonded to 4 oxygen atoms including one −OH). Total 4 boron atoms per formula unit.

Orthoboric Acid

H₃BO₃ is a unique monobasic acid — it doesn't donate a proton but accepts OH⁻ from water (Lewis acid behavior). It has a layered structure with trigonal planar BO₃ units held together by hydrogen bonds.

Orthoboric acid ( H_3BO_3 ) is prepared from borax by treating with HCl. It looks like a simple triprotic acid — three OH groups! But here's the twist: it's actually monobasic, and it doesn't donate protons at all.

H_3BO_3 + H_2O leftharpoons [B(OH)_4]^- + H^+

• Preparation: Na_2B_4O_7 + 2HCl + 5H_2O arrow 2NaCl + 4H_3BO_3
• Structure: Trigonal planar BO_3 units with sp^2 boron
• Layered structure: BO_3 units linked by H-bonds between layers → soft, soapy feel
• Monobasic Lewis acid ( K_a = 5.8 10^ -10 ) — very weak
• On heating: H_3BO_3 373K HBO_2 red hot B_2O_3

The thermal decomposition sequence is important: H_3BO_3 373K HBO_2 (metaboric acid) red hot B_2O_3 (boric anhydride, glassy solid)

Diborane

B₂H₆ is the simplest borohydride with 12 valence electrons but 8 bonds to explain. It uses 3-centre-2-electron (3c-2e) bonds — called 'banana bonds' — for the bridging hydrogens. It's prepared from NaBH₄ or BF₃ and reacts with Lewis bases to form adducts.

Diborane ( B_2H_6 ) is the poster child for electron-deficient bonding. With only 12 valence electrons (3 from each B + 1 from each of 6 H), there aren't enough electrons for 8 conventional 2-centre-2-electron bonds. Nature's solution? Banana bonds!

• Structure: two types of H atoms — 4 terminal H ( H_t ) and 2 bridging H ( H_b )
• Terminal B−H bonds: normal 2-centre-2-electron ( 2c - 2e ) bonds
• Bridging B−H−B bonds: 3-centre-2-electron ( 3c - 2e ) banana bonds
• B−H _b bond is longer than B−H _t (because electrons spread over 3 atoms)
• Each boron is approximately sp^3 hybridized
• Total electron count: 4(2c-2e) terminal + 2(3c-2e) bridging = 8 + 4 = 12 electrons ✓

Preparation: 2BF_3 + 6NaH arrow B_2H_6 + 6NaF

Other preparations: - 2NaBH_4 + I_2 arrow B_2H_6 + 2NaI + H_2 - Industrial: 2BF_3 + 6NaH arrow B_2H_6 + 6NaF

• Diborane reacts with O_2 : B_2H_6 + 3O_2 arrow B_2O_3 + 3H_2O (highly exothermic — used as rocket fuel!)
• With Lewis bases (e.g., NH_3 ): B_2H_6 + 2NH_3 arrow 2[H_3B arrowNH_3] (adduct, then on heating → borazine B_3N_3H_6 )
• Borazine ( B_3N_3H_6 ) is called 'inorganic benzene' — isoelectronic with benzene

Quick recall

How many 3c-2e bonds are present in B₂H₆?

Two. Each bridging hydrogen forms one 3c-2e (banana) bond: B−H _b −B. There are 2 bridging H atoms, so 2 banana bonds. The remaining 4 terminal B−H bonds are normal 2c-2e bonds.

Boron Nitride

BN is isoelectronic with carbon and exists in two structural forms: hexagonal BN (soft, lubricant, like graphite) and cubic BN (extremely hard, like diamond). It's thermally stable and electrically insulating unlike graphite.

Since B and N together have the same number of electrons as two C atoms, BN is isoelectronic with C. This means BN can mimic both allotropes of carbon!

• BN is isoelectronic and isostructural with carbon allotropes
• Hexagonal BN: layers held by van der Waals forces — but no conductivity
• Cubic BN (borazon): synthesis requires high temperature and pressure
• BN is thermally and chemically very stable — used in high-temperature crucibles
• Borazine ( B_3N_3H_6 ) is called 'inorganic benzene' — planar, 6-membered ring

Aluminum & Its Compounds

Aluminum is amphoteric — it dissolves in both acids and bases. AlCl₃ is a powerful Lewis acid that exists as a dimer (Al₂Cl₆) due to electron deficiency. Alums are double salts with the formula M⁺M³⁺(SO₄)₂·12H₂O used in water purification. Aluminum is extracted from bauxite by the Hall-Héroult process.

Al_2O_3 + 2NaOH arrow 2NaAlO_2 + H_2O

Amphoteric Nature of Aluminum

Aluminum and its oxide Al₂O₃ are amphoteric — they react with both acids and bases. With acids, Al forms Al³⁺ salts; with bases, it forms aluminates [Al(OH)₄]⁻. This dual behavior is a defining characteristic of Al.

Aluminum is the textbook example of an amphoteric metal. Its oxide ( Al_2O_3 ) and hydroxide ( Al(OH)_3 ) react with both acids and bases — making Al versatile but also tricky in qualitative analysis.

With acid: Al_2O_3 + 6HCl arrow 2AlCl_3 + 3H_2O

With base: Al_2O_3 + 2NaOH arrow 2NaAlO_2 + H_2O

• Al_2O_3 reacts with acids → Al^ 3+ salts (behaves as base)
• Al_2O_3 reacts with bases → aluminates AlO_2^- (behaves as acid)
• Al(OH)_3 is similarly amphoteric: dissolves in HCl AND NaOH
• In Group 13: B₂O₃ = acidic, Al₂O₃ = amphoteric, Ga₂O₃ = amphoteric, In₂O₃/Tl₂O₃ = basic
• Trend: oxide character changes from acidic → amphoteric → basic down the group

Aluminum Chloride

AlCl₃ exists as a dimer Al₂Cl₆ in the vapor phase due to electron deficiency. Each Al achieves an octet by accepting a lone pair from bridging Cl atoms. AlCl₃ is a powerful Lewis acid and Friedel-Crafts catalyst.

Aluminum chloride ( AlCl_3 ) is one of the most important Lewis acids in chemistry. In the monomer, Al has only 6 electrons — it's electron-deficient. To complete its octet, it forms a dimer ( Al_2Cl_6 ) with two bridging chlorine atoms.

• Monomeric AlCl_3 : Al is sp^2 hybridized, trigonal planar, electron-deficient
• Dimer Al_2Cl_6 : Al becomes sp^3 (tetrahedral) with 4 Cl around each Al
• 2 bridging Cl atoms + 4 terminal Cl atoms in the dimer
• At high temperature ( 400°C ), the dimer breaks into monomers
• AlCl_3 is the catalyst for Friedel-Crafts alkylation and acylation reactions

Alums

Alums have the general formula M⁺M³⁺(SO₄)₂·12H₂O. The most common is potash alum K₂SO₄·Al₂(SO₄)₃·24H₂O. They are used in water purification (coagulation), dyeing, and as styptic to stop bleeding.

Alums are double salts — they contain two different cations and the sulfate anion. The general formula is: M^+_2SO_4 M^ 3+ _2(SO_4)_3 24H_2O or equivalently: M^+M^ 3+ (SO_4)_2 12H_2O per formula unit.

• M^+ can be: Na⁺, K⁺, NH₄⁺, etc.
• M^ 3+ can be: Al³⁺, Fe³⁺, Cr³⁺, etc.
• Most common alum: Potash alum K_2SO_4 Al_2(SO_4)_3 24H_2O
• Alums crystallize in the octahedral crystal system
• Each M^ 3+ ion is surrounded by 6 water molecules: [M(H_2O)_6]^ 3+

Quick recall

Why is alum effective in water purification?

The Al^ 3+ ions (high charge density) neutralize the negative charge on colloidal impurity particles. Once neutralized, the particles aggregate (coagulate) and settle down. This is an application of coagulation from surface chemistry.

Group 14: The Carbon Family

Group 14 elements (C, Si, Ge, Sn, Pb) have ns²np² configuration. Carbon is anomalous — small size, high electronegativity, no d-orbitals, and unique ability to catenate. Oxidation states are +2 and +4; the +2 state becomes more stable down the group (inert pair effect). Carbon and silicon are non-metals; germanium is a metalloid; tin and lead are metals.

General config: ns^2np^2

Electronic Configuration & Oxidation States

Group 14 elements have 4 valence electrons (ns²np²). Both +2 and +4 oxidation states exist, but +4 dominates for C and Si while +2 becomes preferred for Sn and Pb due to inert pair effect.

Group 14 is the heart of the periodic table — containing carbon (basis of life), silicon (basis of electronics), and four other elements that beautifully illustrate the transition from non-metal to metal.

• Elements: C ( Z=6 ), Si ( Z=14 ), Ge ( Z=32 ), Sn ( Z=50 ), Pb ( Z=82 )
• C: [He] 2s^2 2p^2 , Si: [Ne] 3s^2 3p^2 , Ge: [Ar] 3d^ 10 4s^2 4p^2
• Sn: [Kr] 4d^ 10 5s^2 5p^2 , Pb: [Xe] 4f^ 14 5d^ 10 6s^2 6p^2
• Oxidation states: +4 (dominant for C, Si) and +2 (dominant for Pb)
• Ge and Sn show both +2 and +4 states — transition zone

Catenation

Catenation is the ability of an element to form bonds with its own atoms, creating chains, branches, and rings. Carbon is the champion of catenation due to its small size and strong C−C bond (356 kJ/mol). This ability decreases sharply down Group 14.

Catenation is the reason organic chemistry exists. Carbon's ability to link with itself in chains, branches, and rings — with single, double, and triple bonds — creates millions of compounds. No other element comes close.

• C−C bond energy: 356 kJ/mol (very strong)
• Si−Si bond energy: 226 kJ/mol (much weaker → poor catenation)
• Catenation order in Group 14: C >> Si > Ge > Sn >> Pb
• Carbon forms: chains ( C_n ), branches, rings, double bonds ( C=C ), triple bonds ( C )
• Silicon can catenate slightly (silanes Si_nH_ 2n+2 ) but chains rarely exceed 7-8 atoms

Anomalous Behavior of Carbon

Carbon differs from other Group 14 members: highest electronegativity in the group, no d-orbitals (max covalence = 4), exceptional catenation, and ability to form strong pπ-pπ multiple bonds. These make carbon chemistry (organic chemistry) a separate discipline.

Carbon is as anomalous in Group 14 as boron is in Group 13 — and for the same fundamental reasons. Its uniqueness literally created an entire branch of chemistry (organic chemistry).

Carbon Allotropes

Carbon exhibits remarkable allotropy — diamond (sp³, 3D tetrahedral, hardest substance), graphite (sp², layered hexagonal, conductor, lubricant), and fullerenes (C₆₀, sp², cage-like, soccer ball shape). The drastically different properties arise from different bonding arrangements of the same element.

Diamond: sp^3 (3D) | Graphite: sp^2 (2D layers)

Diamond

In diamond, each carbon is sp³ hybridized and bonded to 4 other carbons in a rigid 3D tetrahedral network. This makes diamond the hardest known natural substance, an electrical insulator, and transparent with a very high refractive index.

Diamond is pure carbon arranged in a 3D tetrahedral network. Every carbon atom is sp^3 hybridized and bonded to 4 neighboring carbons by strong bonds — forming a giant covalent structure that extends in all three dimensions.

• Each C is sp^3 hybridized → tetrahedral geometry (bond angle 109.5° )
• C−C bond length: 1.54 Å (same as in alkanes)
• Hardest natural substance (10 on Mohs scale) — rigid 3D network with no weak points
• Electrical insulator — all electrons are localized in bonds, no free electrons
• Very high refractive index (2.42) → 'fire' and brilliance of cut diamonds
• High melting point ( 3550°C ) — strong covalent bonds in all directions

Graphite

In graphite, each carbon is sp² hybridized, forming hexagonal planar layers. The unhybridized p-orbital creates delocalized π-electrons between layers — making graphite an electrical conductor. Weak van der Waals forces between layers make it soft and a good lubricant.

Graphite is the other common allotrope of carbon — and it couldn't be more different from diamond. Each carbon is sp^2 hybridized, forming flat hexagonal layers (like chicken wire). The fourth electron (in the unhybridized 2p orbital) is delocalized across the entire layer.

• Each C is sp^2 → trigonal planar within the layer (bond angle 120° )
• C−C bond length within layers: 1.42 Å (shorter than diamond's 1.54 Å due to partial double bond character)
• Interlayer distance: 3.35 Å — held by weak van der Waals forces
• Electrical conductor along the layers — delocalized -electrons move freely
• Soft and slippery — layers slide over each other → excellent lubricant
• Thermodynamically more stable than diamond under normal conditions

Quick recall

Why is the C−C bond length in graphite (1.42 Å) shorter than in diamond (1.54 Å)?

In graphite, each C−C bond has partial double bond character due to delocalized -electrons (resonance). It's intermediate between a single bond (1.54 Å) and a double bond (1.34 Å). In diamond, all bonds are pure single bonds.

Fullerenes

Fullerenes (C₆₀, C₇₀, etc.) are cage-like allotropes of carbon. C₆₀ (buckminsterfullerene) has a soccer ball shape with 20 hexagons and 12 pentagons. Carbon atoms are sp² hybridized. Fullerenes are soluble in organic solvents, unlike diamond and graphite.

Fullerenes are the third major allotrope of carbon, discovered in 1985. The most famous is C_ 60 (buckminsterfullerene), shaped like a soccer ball — 60 carbon atoms arranged in 20 hexagons and 12 pentagons.

• C_ 60 : 20 hexagons + 12 pentagons = 32 faces, 60 vertices, 90 edges
• Each carbon is sp^2 hybridized (like graphite) — but curved into a cage
• All carbons are equivalent — each is at the junction of 2 hexagons and 1 pentagon
• Soluble in organic solvents (toluene, CS₂) — unlike diamond and graphite
• Made by heating graphite in electric arc in inert atmosphere (He/Ar)
• C₇₀ (rugby ball shape), C₇₆, C₈₄ also exist — larger fullerenes

Carbon nanotubes (CNTs) are related structures — imagine rolling a graphene sheet into a tube. They have extraordinary strength (100× stronger than steel) and are excellent conductors. CNTs are a frontier material in nanotechnology.

Important Compounds of Carbon

Carbon monoxide (CO) is a deadly, odorless, colorless gas that's a powerful reducing agent and forms metal carbonyls. Carbon dioxide (CO₂) is a linear molecule, the main greenhouse gas, and dissolves in water to form carbonic acid. Carbonates and bicarbonates of sodium are industrially vital — Na₂CO₃ (soda ash) and NaHCO₃ (baking soda).

CO_2 + H_2O leftharpoons H_2CO_3 leftharpoons H^+ + HCO_3^-

Carbon Monoxide

CO is formed by incomplete combustion. It has a triple bond (:C≡O:), is isoelectronic with N₂, acts as a powerful reducing agent in metallurgy, and binds to hemoglobin 200× more strongly than O₂ — making it fatally poisonous.

Carbon monoxide ( CO ) is colorless, odorless, and deadly. It has a triple bond ( :C : ) — one and one bond from sharing, plus one coordinate bond from O to C. CO is isoelectronic with N_2 (both have 14 electrons).

Preparation: 2C + O_2 limited 2CO

• Lab prep: HCOOH H_2SO_4, CO + H_2O (dehydration of formic acid)
• CO + O_2 → CO_2 (burns with blue flame)
• Reducing agent in metallurgy: Fe_2O_3 + 3CO arrow 2Fe + 3CO_2 (blast furnace)
• Water gas: CO + H_2 (made by passing steam over hot coke)
• Producer gas: CO + N_2 (made by passing air over hot coke)

Carbon Dioxide

CO₂ is a linear molecule (O=C=O, sp hybridized carbon, bond angle 180°). It's the main greenhouse gas, dissolves in water to form weak carbonic acid (H₂CO₃), and doesn't support combustion (except burning Mg). Solid CO₂ is 'dry ice' — it sublimes.

Carbon dioxide ( CO_2 ) is a linear molecule ( O=C=O , bond angle 180° ) with carbon sp hybridized. Despite having two polar C=O bonds, the molecule is non-polar overall because the dipole moments cancel (linear symmetry).

CO_2(g) + H_2O(l) leftharpoons H_2CO_3(aq) leftharpoons H^+ + HCO_3^-

• Structure: linear, sp hybridized C, bond angle = 180° , non-polar
• Dissolves in water → weak dibasic acid ( H_2CO_3 )
• Does NOT support combustion (exceptions: Mg and Na burn in CO₂)
• Turns lime water milky: Ca(OH)_2 + CO_2 arrow CaCO_3 + H_2O
• Excess CO₂ turns it clear again: CaCO_3 + CO_2 + H_2O arrow Ca(HCO_3)_2 (soluble)
• Solid CO₂ = dry ice — sublimes at -78°C (no liquid phase at 1 atm)

Quick recall

Why does CO₂ turn lime water milky, then clear on passing excess?

First: Ca(OH)_2 + CO_2 arrow CaCO_3 (insoluble white ppt → milky). Then with excess CO₂: CaCO_3 + CO_2 + H_2O arrow Ca(HCO_3)_2 (soluble → clear). This is the standard test for CO₂.

Carbonates & Bicarbonates

Na₂CO₃ (soda ash/washing soda) and NaHCO₃ (baking soda) are industrially vital. Na₂CO₃ is manufactured by the Solvay process. NaHCO₃ decomposes on heating to give CO₂ (used in baking). Carbonates of alkali metals are thermally stable; bicarbonates decompose easily.

The carbonate ion ( CO_3^ 2- ) has a trigonal planar structure with sp^2 carbon. All three C−O bonds are equivalent due to resonance — each bond has a bond order of 1 1 3 . This ion forms two major classes of salts: carbonates and bicarbonates.

• NaHCO_3 (baking soda): mild base, used in antacids, fire extinguishers, baking
• Na_2CO_3 10H_2O (washing soda): laundry, water softening, glass making
• Solvay process: NaCl + NH_3 + CO_2 + H_2O arrow NaHCO_3 + NH_4Cl
• Heating NaHCO₃: 2NaHCO_3 Na_2CO_3 + H_2O + CO_2 (why bread rises!)
• Thermal stability: metal carbonates (Li₂CO₃ least stable) — stability increases down the group

Silicon & Its Compounds

Silicon forms unique compounds: silicones (R₂SiO polymers used in sealants and implants), silicates (SiO₄⁴⁻ tetrahedra linked in chains/sheets/3D networks — basis of rocks and minerals), zeolites (aluminosilicates used as molecular sieves and catalysts), and glass (amorphous SiO₂ with Na₂O/CaO). Silicon is the semiconductor powering all modern electronics.

SiO_4^ 4- (basic tetrahedral unit of all silicates)

Silicones

Silicones are synthetic organosilicon polymers with alternating Si−O−Si backbone and organic groups (R) attached to Si. They are water-repellent, heat-resistant, electrical insulators, and chemically inert — used in sealants, lubricants, implants, and cookware.

Silicones are organosilicon polymers with the general formula (R_2SiO)_n . The backbone is an alternating Si−O−Si chain (siloxane linkage), with organic groups (usually CH_3 ) attached to each silicon. They're made by hydrolyzing chlorosilanes like (CH_3)_2SiCl_2 .

n(CH_3)_2SiCl_2 + nH_2O arrow [(CH_3)_2SiO]_n + 2nHCl

• Backbone: −Si−O−Si−O− (siloxane linkage) — very stable
• Organic groups (R = CH₃, C₆H₅, etc.) attached to Si → hydrophobic surface
• Properties: water-repellent, heat-resistant ( -60°C to +250°C ), flexible, insulating
• Chain length determined by R_3SiCl (chain stopper) vs R_2SiCl_2 (chain extender)
• Cross-linking possible with RSiCl_3 → rigid 3D networks

Silicates

Silicates are built from SiO₄⁴⁻ tetrahedral units that can link by sharing corners (oxygen atoms). Different linking patterns create different silicate types: isolated, chain, sheet, 3D framework. About 95% of Earth's crust is made of silicates.

Silicates are the most abundant minerals on Earth — they make up ~95% of the crust. The fundamental building block is the SiO_4^ 4- tetrahedron — Si at the center bonded to 4 oxygen atoms at the corners. These tetrahedra can link by sharing corner oxygens.

• Basic unit: SiO_4^ 4- tetrahedron ( sp^3 Si at center)
• Linking occurs by sharing corner oxygens only — never edges or faces
• More sharing → lower Si:O ratio → lower charge → harder, more rigid structure
• If some Si atoms are replaced by Al → aluminosilicates (feldspars, zeolites)
• Asbestos is a double-chain silicate — fibrous structure causes health hazards

Zeolites

Zeolites are hydrated aluminosilicates with 3D framework structures containing channels and cavities. They act as molecular sieves (separating molecules by size), ion exchangers (water softening), and shape-selective catalysts (petroleum cracking).

Zeolites are 3D framework aluminosilicates where some Si^ 4+ atoms are replaced by Al^ 3+ . This substitution creates a negative charge on the framework, which is balanced by cations (Na⁺, Ca²⁺, K⁺) sitting inside channels and cavities.

• General formula: M_ x/n [(AlO_2)_x(SiO_2)_y] zH_2O
• 3D cage-like structure with channels of specific sizes
• Molecular sieve: channels allow only molecules of certain size to enter
• Ion exchange: Na⁺/Ca²⁺ in zeolite can be exchanged → water softening
• Shape-selective catalysis: ZSM-5 catalyst converts alcohols to gasoline
• Used in petroleum cracking — the specific pore size selects for desired hydrocarbons

Glass

Glass is an amorphous (non-crystalline) solid made by fusing SiO₂ with Na₂O and CaO (soda-lime glass). Different additives create specialized glasses: borosilicate (Pyrex — heat-resistant), lead crystal (PbO — high refractive index), and colored glasses (metal oxides).

Glass is a supercooled liquid — an amorphous solid with no long-range crystalline order. Common soda-lime glass is made by fusing SiO_2 (sand) with Na_2CO_3 (soda ash) and CaCO_3 (limestone) at ~ 1500°C .

Na_2SiO_3 CaSiO_3 4SiO_2 (approximate soda-lime glass composition)

• Glass is amorphous — no sharp melting point, softens gradually over a range
• Soda-lime glass: ~75% SiO₂, 15% Na₂O, 10% CaO (cheapest, most common)
• Pyrex glass has low coefficient of thermal expansion → doesn't crack with temperature change
• Glass can be tempered (rapidly cooled) for extra strength → car windshields
• Glass is 100% recyclable without loss of quality — infinitely!