General Trends · Anomalous Properties · Interhalogen Structures · Chemical Reactions & Industrial Uses
P-block elements are those in which the last differentiating electron enters the p-orbital of the outermost shell. They occupy Groups 13 to 18 of the periodic table, comprising a total of 35 elements (excluding He, which is placed in Group 18 but belongs to s-block by electronic configuration).
The general outer electronic configuration is ns²np¹⁻⁶, where n ranges from 2 to 6 (periods 2–6). The s and p electrons together form the valence shell, determining chemical behavior, oxidation states, and bonding patterns.
P-Block Element Grid (Period 2–6, Groups 13–18)
| Property | Trend Down a Group | Trend Across a Period | Key Reason |
|---|---|---|---|
| Atomic Radius | Increases ↓ | Decreases → | More shells added going down; increasing nuclear charge going across |
| Ionisation Enthalpy | Generally decreases ↓ | Generally increases → | Larger atoms hold outer electrons less tightly; higher Zeff across |
| Electronegativity | Decreases ↓ | Increases → | Atomic size increases down, Zeff increases across |
| Electron Gain Enthalpy | Becomes less negative ↓ | Becomes more negative → | Larger atoms less effectively attract extra electron |
| Metallic Character | Increases ↓ | Decreases → | Lower ionisation enthalpy; tendency to lose electrons increases |
| Melting/Boiling Point | Variable (depends on structure) | Generally increases from Grp 13→14, then decreases | C, Si have giant covalent → very high mp; rest molecular |
Elements show oxidation states from –(8–n) to +n, where n is the group number. The lower oxidation state becomes more stable on going down a group due to the inert pair effect (reluctance of ns² electrons to participate in bonding).
The ability to form chains/rings with themselves. Carbon has maximum catenation due to its strong C–C bond and small size. It decreases down the group: C > Si > Ge > Sn > Pb.
Many p-block elements show allotropy (different structural forms in the same physical state). Notable: Carbon (diamond, graphite, fullerene), Sulphur (rhombic, monoclinic), Phosphorus (white, red, black).
Down the group, the ns² electrons become increasingly reluctant to participate in bond formation due to poor shielding by d and f electrons. This stabilizes the lower oxidation state. Examples: Tl(I) > Tl(III); Pb(II) > Pb(IV); Bi(III) > Bi(V).
2nd period elements (C, N, O) readily form multiple bonds via pπ–pπ overlap due to their small size. Heavier congeners prefer dπ–pπ or single bonds — their larger size reduces effective orbital overlap for pπ–pπ bonding.
The first member of each p-block group differs significantly from the rest of the group because of:
Interhalogen compounds are formed between two different halogen atoms with the general formula XX'ₙ, where X is the larger (less electronegative) halogen and X' is the smaller (more electronegative) halogen, and n = 1, 3, 5, or 7. They are generally more reactive than the parent halogens because the X–X' bond is weaker than X–X.
The outer halogen X' always has an odd number of lone pairs on the central X after bonding. The central halogen X uses its d orbitals to accommodate extra bond pairs. The number of X' atoms increases in odd steps based on available vacant d orbitals and steric factors.
T-shaped structure; 2 lone pairs in equatorial positions reduce F–Cl–F angle slightly below 90°
1 axial lone pair distorts octahedral geometry to square pyramidal; F–Br–F angle ≈ 84.8°
Only interhalogen with n=7; IF₇ has no lone pairs, giving ideal pentagonal bipyramidal geometry
ICl is polar due to electronegativity difference (Cl > I); acts as a source of Cl⁺ or I⁺ in reactions
Boron trihalides are strong Lewis acids (electron-deficient). The Lewis acidity order is BF₃ < BCl₃ < BBr₃ (contrary to expected electronegativity trend) because back-donation of lone pair from F into empty p-orbital of B reduces Lewis acidity of BF₃ more.
Diborane has 3c-2e (banana) bonds. It hydrolyses readily and is a reducing agent.
Boric acid is a weak monobasic acid (Lewis acid, not Brønsted). It acts by accepting OH⁻ from water rather than donating a proton.
Borax dissolves in water to give an alkaline solution. When heated, it gives a transparent bead of borax that can dissolve metallic oxides to give characteristic coloured beads — used for identification of metal ions (borax bead test).
Carbon forms two important oxides: CO (reducing agent, ligand) and CO₂ (acidic oxide). CO is a product of incomplete combustion and is extremely toxic (binds to haemoglobin).
Unlike CO₂, SiO₂ is a giant covalent network solid (each Si bonded to 4 oxygens tetrahedrally). It dissolves only in HF acid and reacts with NaOH at high temperature.
Silicones are synthetic polymers with Si–O–Si backbone and organic groups (R = CH₃, C₆H₅) on silicon. They are heat resistant, water repellent, and chemically inert.
Pb(II) is more stable than Pb(IV); Sn(IV) is more stable than Sn(II) (Sn²⁺ is a good reducing agent). PbO₂ is a strong oxidising agent.
Nitrogen forms a series of oxides from N₂O (+1) to N₂O₅ (+5). Key ones: NO (colourless, paramagnetic, odd electron), NO₂ (brown, paramagnetic, toxic), N₂O₄ (dimer of NO₂).
Manufactured by the Ostwald process (catalytic oxidation of ammonia). HNO₃ is a strong oxidising acid; it dissolves most metals except Au and Pt (dissolve in aqua regia). Fe, Al, Cr are passivated by conc. HNO₃ (form oxide layer).
Phosphorus forms many oxoacids. Basicity (number of ionisable H) is determined by the number of P–OH groups, NOT P–H bonds (P–H bonds are non-ionisable).
| Acid | Formula | Oxidation State | Basicity | P–H bonds |
|---|---|---|---|---|
| Hypophosphorous acid | H₃PO₂ | +1 | Monobasic | 2 (strong reducer) |
| Phosphorous acid | H₃PO₃ | +3 | Dibasic | 1 (reducer) |
| Orthophosphoric acid | H₃PO₄ | +5 | Tribasic | 0 |
| Pyrophosphoric acid | H₄P₂O₇ | +5 | Tetrabasic | 0 |
| Meta-phosphoric acid | (HPO₃)ₙ | +5 | Monobasic/unit | 0 |
Ozone is an allotrope of oxygen. Its structure has a bent geometry (O–O–O angle = 116.8°) with resonance; one O has a +1 formal charge. It is a powerful oxidising agent (decomposes to give nascent oxygen).
Sulphur forms many oxoacids. Key ones: H₂SO₄ (king of chemicals), H₂SO₃, H₂S₂O₇ (oleum), and thiosulphate H₂S₂O₃.
Halogens are the strongest oxidising agents in the p-block. Oxidising power decreases: F₂ > Cl₂ > Br₂ > I₂. Each halogen can oxidise the halide ions below it in the group (displacement reactions).
Halogens (except F) form oxoacids. Acidic strength increases with oxidation state of halogen: HOCl < HOClO < HOClO₂ < HOClO₃. Stability and oxidising power of oxoacids varies.
| Acid | Oxidation State of Cl | Nature | Strength |
|---|---|---|---|
| HOCl (Hypochlorous acid) | +1 | Weak acid | Weak |
| HOClO (Chlorous acid) | +3 | Unstable | Moderate |
| HOClO₂ (Chloric acid) | +5 | Strong acid | Strong |
| HOClO₃ (Perchloric acid) | +7 | Strongest halogen acid | Very Strong |
Chlorine is manufactured by the Chlor-alkali (electrolytic) process from brine (NaCl solution). It reacts as an oxidiser, bleaching agent, and in substitution/addition reactions.
Thermal stability decreases: HF > HCl > HBr > HI. Reducing power increases: HF < HCl < HBr < HI. Acid strength increases: HF << HCl < HBr < HI (HF is weak acid due to very strong H–F bond and F⁻ solvation).
Noble gases have completely filled s and p orbitals (ns²np⁶, except He: 1s²) — the most stable electronic configuration. This gives them very high ionisation enthalpies and negligible electron affinity, making bond formation very unfavourable under normal conditions.
Neil Bartlett (1962) prepared the first noble gas compound Xe⁺[PtF₆]⁻ by reacting Xe with PtF₆, because IE of Xe (1170 kJ/mol) ≈ IE of O₂ (1165 kJ/mol), and PtF₆ had already oxidised O₂. This proved noble gases could form compounds.
| Compound | Geometry | Hybridisation | Lone Pairs on Xe |
|---|---|---|---|
| XeF₂ | Linear | sp³d | 3 |
| XeF₄ | Square Planar | sp³d² | 2 |
| XeF₆ | Distorted Octahedral | sp³d³ | 1 |
| XeOF₄ | Square Pyramidal | sp³d² | 1 |
| XeO₂F₂ | See-Saw | sp³d | 1 |
| XeO₃ | Pyramidal | sp³ | 1 |
| XeO₄ | Tetrahedral | sp³ | 0 |
| Group | Common Positive States | Common Negative States | Most Stable (lower members) |
|---|---|---|---|
| 13 | +1, +3 | – | +1 (Tl) |
| 14 | +2, +4 | –4 (C in CaC₂) | +2 (Pb, Sn to some extent) |
| 15 | +3, +5 | –3 (NH₃) | +3 (Bi) |
| 16 | +2, +4, +6 | –2 (H₂S, H₂O) | +4 (Te, Po) |
| 17 | +1, +3, +5, +7 | –1 (HCl) | –1 for F only |
| 18 | 0 (mostly); +2,+4,+6 for Xe | – | 0 (all except Xe compounds) |