Statement A — TRUE ✓
Fluorine (F, electronegativity = 4.0 on Pauling scale) is the most electronegative element in the entire periodic table. This makes F the strongest oxidising agent among halogens and explains many of its unique properties.
Statement B — FALSE ✗
Fluorine is a stronger oxidising agent than chlorine — NOT the other way around. Oxidising power of halogens decreases down the group: F₂ > Cl₂ > Br₂ > I₂. F₂ can oxidise water, Cl₂ cannot oxidise F⁻. Statement B is wrong.
Statement C — TRUE ✓
Physical states at room temperature: F₂ = pale yellow gas, Cl₂ = yellow-green gas, Br₂ = reddish-brown liquid, I₂ = shiny grey-black solid. Iodine is indeed a solid at room temperature (25°C). It sublimes directly from solid to purple vapour when heated.
Statement D — TRUE ✓
Bond dissociation energy: F₂ = 158 kJ/mol, Cl₂ = 242 kJ/mol. F₂ has LOWER bond dissociation enthalpy than Cl₂ — unusual because normally bond energy decreases as atoms get larger. F₂ is weak because lone pair–lone pair repulsion is severe between small F atoms bonded together.
Halogens (Group 17: F, Cl, Br, I, At) are the most reactive non-metals. They have 7 valence electrons (ns²np⁵) and need one more electron to complete their octet — making them powerful oxidising agents. Down the group, reactivity decreases as electronegativity decreases and atomic size increases. All halogens exist as diatomic molecules (X₂) in their elemental form.
📌 F₂: pale yellow gas, bp = −188°C, most reactive halogen
📌 Cl₂: yellow-green gas, bp = −34°C, pungent smell, used in water treatment
📌 Br₂: reddish-brown liquid, bp = 59°C, dense toxic vapour
📌 I₂: shiny grey-black solid, bp = 184°C, sublimes (purple vapour)
📌 At (astatine): radioactive, very short half-life, not practically available
Fluorine shows several anomalous properties compared to other halogens due to its small size, highest electronegativity, and absence of d-orbitals. Key anomalies:
📌 F₂ has lower bond dissociation energy (158 kJ/mol) than Cl₂ (242 kJ/mol) — due to lone pair repulsion in small F atoms
📌 F shows only −1 oxidation state (no d-orbitals, cannot expand octet)
📌 Other halogens show −1, +1, +3, +5, +7 oxidation states
📌 HF is a weak acid (extensive H-bonding), HCl/HBr/HI are strong acids
📌 F₂ reacts with water to give OF₂ (fluorine displaces oxygen): 2F₂ + 2H₂O → O₂ + 4HF. Actually: F₂ + H₂O → HF + HOF
📌 F cannot be made by electrolysis of aqueous solution (oxidises water) — made by electrolysis of molten KHF₂
Oxidising power follows the trend F₂ > Cl₂ > Br₂ > I₂, determined by the reduction potential (E°). F₂ has the highest reduction potential (E° = +2.87 V) — it can oxidise all other halide ions. Cl₂ can displace Br⁻ and I⁻ but not F⁻. Br₂ can displace I⁻ but not Cl⁻ or F⁻. I₂ cannot displace any other halide ion from solution. This is the basis of the "halogen displacement reactions" used in analytical chemistry.
📌 Acid strength in aqueous solution: HF < HCl < HBr < HI
📌 HF is weak acid (Ka ≈ 6.8×10⁻⁴) — extensive H-bonding, strong H-F bond
📌 HI is strongest acid — weakest H-I bond, easiest to donate H⁺
📌 Thermal stability: HF > HCl > HBr > HI (bond strength decreases)
📌 HI decomposes at 300°C; HF is stable even at 1000°C
📌 Reducing power: HF < HCl < HBr < HI (HI is a good reducing agent)
Chlorine forms a series of oxoacids: HOCl (hypochlorous acid, +1), HOClO (chlorous acid, +3), HOClO₂ (chloric acid, +5), HOClO₃ (perchloric acid, +7). Acid strength increases with oxidation state of Cl: HClO₄ > HClO₃ > HClO₂ > HClO. More oxygen atoms pull more electron density from H-O bond, making H⁺ easier to release. HClO₄ (perchloric acid) is one of the strongest acids known. Fluorine forms no oxoacids (cannot show positive oxidation state — most electronegative).
Interhalogens are compounds formed between two different halogens: XY, XY₃, XY₅, XY₇. Examples: ClF, BrF₃, IF₅, IF₇. The central halogen (larger) is the one with positive formal charge; fluorine always acts as the smaller terminal halogen (most electronegative). Properties are intermediate between parent halogens. ClF is used as fluorinating agent. IF₇ is the most complex known interhalogen. All are reactive and used as fluorinating/chlorinating agents in industry.
Noble gases (He, Ne, Ar, Kr, Xe, Rn) have completely filled outer shells (ns²np⁶, except He which is 1s²). Traditionally considered inert — until Bartlett (1962) prepared XePtF₆, the first noble gas compound. Xenon forms XeF₂, XeF₄, XeF₆, XeO₃, XeO₄. Radon is radioactive. Noble gases have the highest ionisation enthalpies in their periods. Uses: He in balloons, He-O₂ for deep sea diving, Ne in signs, Ar in light bulbs, Kr in photography flashes, Xe in high-intensity lamps.
📌 Atomic radius: increases F < Cl < Br < I
📌 Electronegativity: decreases F > Cl > Br > I
📌 Electron affinity: Cl > F > Br > I (F anomalously low — small size, high repulsion)
📌 Ionisation enthalpy: decreases down group
📌 Melting and boiling points: increase (F → I) due to stronger van der Waals forces with larger molecules
📌 Oxidising power: decreases F > Cl > Br > I
📌 Reducing power of halide ions: increases F⁻ < Cl⁻ < Br⁻ < I⁻