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Chemistryp-Block Elements
Match the xenon compounds with their hybridisation and geometry:
A. $XeO_3$ → I. $sp^3d$, linear
B. $XeF_2$ → II. $sp^3$, pyramidal
C. $XeOF_4$ → III. $sp^3d^3$, distorted octahedral
D. $XeF_6$ → IV. $sp^3d^2$, square pyramidal
Options
1
A-II, B-I, C-IV, D-III
2
A-I, B-II, C-III, D-IV
3
A-III, B-IV, C-I, D-II
4
A-IV, B-III, C-II, D-I
Correct Answer
A-II, B-I, C-IV, D-III
Solution
1

Count electron pairs around Xe:

XeO₃: 3 BP + 1 LP = 4 → sp³, pyramidal → A-II ✓

XeF₂: 2 BP + 3 LP = 5 → sp³d, linear → B-I ✓

2

XeOF₄: 5 BP + 1 LP = 6 → sp³d², square pyramidal → C-IV ✓

XeF₆: 6 BP + 1 LP = 7 → sp³d³, distorted octahedral → D-III ✓

Answer: A-II, B-I, C-IV, D-III

XeO₃: sp³(pyramidal) | XeF₂: sp³d(linear) | XeOF₄: sp³d²(sq. pyramidal) | XeF₆: sp³d³(dist. oct.)
Theory: p-Block Elements
1. Noble Gas Chemistry — Xenon Compounds

Noble gases (Group 18): He, Ne, Ar, Kr, Xe, Rn. Generally very stable due to filled electron shells. Xenon is the most reactive noble gas (large size, low ionisation energy → easier to oxidise). Neil Bartlett (1962) first synthesised O₂⁺[PtF₆]⁻, then by analogy (Xe has similar IE to O₂) synthesised Xe[PtF₆]. This demolished the "inert gas" concept. Xenon fluorides: XeF₂, XeF₄, XeF₆ (formed by direct reaction Xe + F₂ under varying conditions). Xenon oxides: XeO₃ (explosive!), XeO₄ (unstable), XeO₂F₂, XeOF₄. Xenon fluorides are powerful fluorinating and oxidising agents. XeF₂ used: selective fluorination of organic compounds, mild oxidising agent in synthesis. Krypton fluoride: KrF₂ (less stable than XeF₂). No stable compounds of He, Ne, Ar (ionisation energies too high for any ligand to oxidise them at reasonable conditions).

2. Hybridisation and VSEPR for Xenon Compounds

The key: count total electron pairs (bonding + lone pairs) on Xe. 4 pairs: sp³ hybridisation. 5 pairs: sp³d. 6 pairs: sp³d². 7 pairs: sp³d³. Then determine geometry using VSEPR (lone pairs occupy more space, repel more). XeF₂ (5 pairs: 2 BP + 3 LP): sp³d. Geometry: linear (LPs equatorial in trigonal bipyramid). XeF₄ (6 pairs: 4 BP + 2 LP): sp³d². LPs at opposite axial positions → square planar. XeO₃ (4 pairs: 3 BP + 1 LP): sp³. Pyramidal (like NH₃). XeOF₄ (6 pairs: 5 BP + 1 LP): sp³d². LP in axial position → square pyramidal. XeF₆ (7 pairs: 6 BP + 1 LP): sp³d³. Lone pair distorts octahedral → distorted octahedral. XeO₄ (4 pairs, all bonding): sp³, tetrahedral. Note: unlike Cl, Xe forms fewer compounds with O (harder to oxidise).

3. Hybridisation Table Summary

Number of electron pairs → hybridisation → geometry (if all BP) → geometry with LP: 2 pairs: sp → linear. 3: sp² → trigonal planar. With 1 LP: bent. 4: sp³ → tetrahedral. 1 LP: pyramidal (like NH₃). 2 LP: bent (like H₂O). 5: sp³d → trigonal bipyramidal. 1 LP: seesaw (like SF₄). 2 LP: T-shaped (like ClF₃). 3 LP: linear (like XeF₂, I₃⁻). 6: sp³d² → octahedral. 1 LP: square pyramidal (like BrF₅, XeOF₄). 2 LP: square planar (like XeF₄). 7: sp³d³ → pentagonal bipyramidal. 1 LP: distorted octahedral (like XeF₆). Important examples: SF₆ (sp³d², octahedral). PCl₅ (sp³d, trig. bipyramidal). SF₄ (sp³d, seesaw). ClF₃ (sp³d, T-shaped). ICl₄⁻ (sp³d², square planar).

4. Reactions of Xenon Fluorides

XeF₂ + H₂O → Xe + HF + O (slowly hydrolyses). XeF₂ is a selective fluorinating agent: RCH₂R' + XeF₂ → RCHFR' (fluorination of organic compounds). XeF₄ + H₂O → XeO₃ + HF (fast, explosive XeO₃ formed). XeF₄ powerful oxidiser: 2Ce³⁺ → 2Ce⁴⁺ (can oxidise Ce³⁺, Pt, Au). XeF₆ + H₂O → XeOF₄ (partial hydrolysis). XeF₆ + 3H₂O → XeO₃ + 6HF (complete hydrolysis). XeO₃ is explosive when dry. XeO₃ + NaOH → NaHXeO₄ (sodium hydrogen xenate). XeO₄ (xenon tetroxide): sp³, tetrahedral, unstable gas (decomposes explosively). XeF₂ is commercially available (expensive) and used in VLSI chip manufacturing for silicon etching (more selective than F₂).

5. Structure and Bonding in Noble Gas Compounds

Noble gas-F bonds: highly polar (Xe δ+ — F δ-). Xe-F bond is weak (Xe not very electronegative). Bond energies: Xe-F ≈ 130 kJ/mol (weak). Compare: C-F ≈ 486 kJ/mol, H-F ≈ 570 kJ/mol. Xe-O bond somewhat stronger. The stability of xenon compounds results from very electronegative F creating enough favourable energy to overcome the unfavourable disruption of xenon's filled shell. For lighter nobles (He, Ne, Ar): IE too high, no compound is thermodynamically stable. Kr: only KrF₂ stable (barely). Rn: radioactive, difficult to study, RnF₂ probably exists. Bonding model: Xe has available d-orbitals for expanded octet. Alternative: 3-centre 4-electron (3c4e) bonding in XeF₂ (delocalized molecular orbital model). MO approach: three p-orbitals on F-Xe-F combine to give bonding, non-bonding, and antibonding MOs. 4 electrons fill bonding and non-bonding → net bond order of ½ for each Xe-F (2 bonds shared among 4 electrons).

6. Shapes of Important Molecules — Summary

Linear (180°): BeCl₂ (sp), CO₂ (sp), HCN (sp), CS₂ (sp), C₂H₂ (sp on each C), XeF₂ (sp³d). Trigonal planar (120°): BF₃ (sp²), SO₃ (sp²), AlCl₃ (sp²), C₂H₄ (sp² on each C). Bent: SO₂ (sp²), O₃ (sp²), SnCl₂ (sp²), H₂O (sp³, 104.5°). Tetrahedral (109.5°): CH₄, NH₄⁺, SO₄²⁻, XeO₄. Pyramidal: NH₃ (sp³, 107°), PH₃ (sp³, 93°), PCl₃, XeO₃. Seesaw: SF₄ (sp³d, 4 BP + 1 LP). T-shaped: ClF₃ (sp³d, 3 BP + 2 LP). Square planar: XeF₄ (sp³d², 4 BP + 2 LP), [PtCl₄]²⁻, [Ni(CN)₄]²⁻. Square pyramidal: IF₅ (sp³d², 5 BP + 1 LP), BrF₅, XeOF₄. Trigonal bipyramidal: PCl₅ (sp³d), SOF₄. Octahedral: SF₆ (sp³d²), [Co(NH₃)₆]³⁺. Distorted octahedral: XeF₆ (sp³d³).

7. Periodic Trends in Group 18

Atomic radii: He < Ne < Ar < Kr < Xe < Rn. Boiling points: He < Ne < Ar < Kr < Xe < Rn (-269, -246, -186, -152, -107, -62°C). Increases with molecular mass and polarisability → stronger London dispersion forces. Ionisation energies: He (2372) > Ne (2081) > Ar (1521) > Kr (1351) > Xe (1170) > Rn (~1036) kJ/mol. Decreases down group (outer electrons farther, shielded). He has highest IE of all elements. Uses: He: balloons (lighter than air, non-flammable unlike H₂), cryogenics (liquid He for MRI magnets and low-temperature physics), diving (HeO₂ mix avoids N₂ narcosis). Ne: neon lights (red-orange glow). Ar: welding shielding gas (prevents oxidation), incandescent bulbs. Kr: flash photography, high-intensity lights. Xe: anaesthesia (no long-term toxicity), ion propulsion in spacecraft, high-intensity discharge lamps. Rn: radioactive hazard in buildings (from uranium-containing rocks/soil).

8. d-block and Expanded Octet

Second and third period elements (P, S, Cl, Xe, etc.) can have more than 8 electrons in valence shell (expanded octet) by using d-orbitals. First period (N, O, F) and second period up to Ne: no d-orbitals in valence shell → maximum 4 bonds (octet). PCl₅: 10 electrons around P (5 bonds). SF₆: 12 electrons around S (6 bonds). XeF₂: 10 electrons around Xe (2 bonds + 3 lone pairs). The d-orbitals used are the empty d-orbitals of the same principal quantum shell (3d for P, S, Cl; 5d for Xe). Controversy: some chemists argue that d-orbitals don't significantly contribute to bonding in these molecules (hypervalent bonding explained by resonance or ionic models). But for NEET: the sp³d, sp³d², sp³d³ hybridisation model with d-orbital participation is accepted and tested.

Frequently Asked Questions
1. How do you determine hybridisation of xenon in XeF2? ⌄
XeF₂: Xe has 8 valence electrons. 2 electrons used for 2 bonds with F. Remaining 6 electrons = 3 lone pairs. Total electron pairs = 2 (bonding) + 3 (lone pairs) = 5. 5 electron pairs → sp³d hybridisation. Arrangement: trigonal bipyramidal. In trigonal bipyramidal, lone pairs prefer equatorial positions (more space) → 3 lone pairs go equatorial. 2 bond pairs go axial → linear geometry for the F-Xe-F part. Bond angle: F-Xe-F = 180° (linear). Compare with similar: I₃⁻ (3 LP + 2 BP = sp³d, linear), ClF₂⁻ (3 LP + 2 BP = sp³d, linear).
2. Why do lone pairs prefer equatorial positions in trigonal bipyramidal? ⌄
In a trigonal bipyramidal arrangement: equatorial positions have 2 adjacent positions at 90° and 2 at 120°. Axial positions have 3 adjacent positions at 90°. Lone pairs repel more strongly than bond pairs (lone pairs are "fatter" — not directed away in a bond, so more electron density close to central atom). At 90° angles, repulsion is strongest. Equatorial lone pairs have fewer 90° interactions (2 axial bonds at 90°) vs axial lone pairs (3 equatorial bonds at 90°). So lone pairs go equatorial to minimise 90° LP-BP repulsion. This principle explains: SF₄ (1 LP equatorial → seesaw), ClF₃ (2 LP equatorial → T-shape), XeF₂ (3 LP equatorial → linear).
3. What is the structure of XeF4? ⌄
XeF₄: Xe has 8 valence electrons. 4 bonds to F (4 electron pairs). Remaining 4 electrons = 2 lone pairs. Total = 6 electron pairs → sp³d². Octahedral arrangement. 2 lone pairs prefer to be opposite each other (trans, at 180°) to minimise mutual repulsion (two big lone pairs as far apart as possible). 4 F atoms in equatorial plane → square planar geometry. Bond angles: F-Xe-F = 90° (in the square) and 180° (opposite F atoms). Confirmed by experimental electron diffraction. Compare: IF₄⁻ (same: sp³d², square planar). XeF₄ is a powerful fluorinating agent, used in chemical synthesis.
4. Why is XeO3 explosive but XeF2 is stable? ⌄
XeF₂: Xe-F bond has some stability because F is very electronegative → the polar Xe⁺-F⁻ bond is stabilised by electrostatic attraction. Also, F is the only element that can "force" Xe to form bonds (highest electronegativity). XeO₃: Xe-O bonds are weaker than Xe-F bonds. Also, XeO₃ has a thermodynamic driving force to decompose: Xe(s) + 3/2 O₂(g) is much more stable. XeO₃ is a strong oxidising agent — it readily releases its oxygen atoms → oxidises almost anything → explosive reaction with organic matter or reducing agents. Dry XeO₃ is extremely shock-sensitive (detonates easily). Xenon-oxygen bonds are inherently less stable than xenon-fluorine bonds because O has two bonding slots (forms two bonds in most compounds) while F has one. The Xe=O double bonds in XeO₃ are relatively weak and highly reactive.
5. What applications do xenon compounds have in synthesis? ⌄
XeF₂ is used in organic synthesis as a mild, selective fluorinating agent: (1) Fluorination of aromatic compounds (electrophilic aromatic fluorination without the harsh conditions of F₂). (2) Dehydrofluorination. (3) Oxidation of organic sulfides to sulfoxides/sulfones. (4) Silicon etching in semiconductor manufacturing (more selective than F₂ gas). XeF₂ is commercially available but expensive (about 40x more expensive than equivalent fluorine). Advantages over F₂ (extremely reactive gas): safer to handle (solid), more selective (gentle oxidising agent), can be stored in glass. XeO₃ (if handled carefully, extremely hazardous) is a very powerful oxidant. XeOF₄ as fluorinating/oxidising agent. These compounds demonstrate that noble gas chemistry has practical applications beyond just curiosity.
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