Ozone (O3) is an angular sp2 molecule (bond angle ~117°) that detects itself via the blue-black KI-starch paper test, working through nascent oxygen release upon decomposition.

The most stable sulfur allotrope at room temperature is rhombic sulfur (S8 rings), which converts to monoclinic sulfur above the transition temperature of 95.6°C.

The Contact process manufactures H2SO4 through four steps: S → SO2 → SO3 (V2O5, 450°C) → oleum (SO3 + H2SO4) → H2SO4; SO3 must never be dissolved in water directly.

Concentrated H2SO4 is simultaneously a diprotic strong acid, a dehydrating agent (chars sugar), an oxidizing agent when hot, and a sulfonating agent.

Oxidizing power of halogens decreases down Group 17: F2 > Cl2 > Br2 > I2, while acidic strength of HX increases: HF < HCl < HBr < HI.

HF is anomalously the weakest HX acid because the H-F bond $\frac{568 kJ}{mol}$ is the strongest among H-X bonds, and intermolecular H-bonding stabilizes undissociated HF.

Chlorine oxoacid strength increases with oxidation state: HOCl (+1) < HClO2 (+3) < HClO3 (+5) < HClO4 (+7) because more oxygen atoms better stabilize the conjugate base.

Interhalogen compounds have shapes: AB (linear), AB3 (T-shaped, sp3d), AB5 (square pyramidal, sp3d2), AB7 (pentagonal bipyramidal, sp3d3).

Xenon fluorides have predictable VSEPR geometries: XeF2 (linear, 3 lone pairs), XeF4 (square planar, 2 lone pairs), XeF6 (distorted octahedral, 1 lone pair), with the decisive rule that lone pairs go equatorial in TBP and trans in octahedral.

XeF6 hydrolysis produces the explosive XeO3 and HF, while XeF4 undergoes disproportionation, and clathrate compounds physically trap noble gases without any chemical bonding.