Given below are two statements about reactions of NaCl and NaI with H2SO4:
Statement I: Both NaCl and NaI react with concentrated H2SO4 to give the respective hydrogen halides HCl and HI as the only products.
Statement II: With NaI, concentrated H2SO4 acts as an oxidising agent, converting I- to I2.
Choose the most appropriate answer:

The halogens (Group VIIA/17) — fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) — form one of the most chemically coherent and extensively studied groups in the periodic table. They are the most electronegative and most reactive non-metals, all existing as diatomic molecules (X2) in their standard states. Physical state at room temperature: F2 (pale yellow gas), Cl2 (greenish-yellow gas), Br2 (reddish-brown liquid, only liquid non-metal at room temperature), I2 (shiny black-purple solid that sublimes to violet vapour), At (radioactive, short-lived, solid, properties largely predicted by extrapolation). Trends down the group: atomic radius increases, electronegativity decreases (F=4.0, Cl=3.0, Br=2.8, I=2.5), oxidising power decreases, melting and boiling points increase (London dispersion forces increase with molecular size), X-X bond energy decreases (F2 < Cl2 due to lone pair repulsions in very small F2), HX acid strength increases (HI > HBr > HCl >> HF), reducing power of X- increases (I- > Br- > Cl- > F-).

The different reactions of chloride, bromide, and iodide ions with concentrated sulfuric acid beautifully illustrate the trend in reducing power across the halide series and provide a practical test to distinguish between these ions. Chloride (NaCl + conc. H2SO4): initially forms NaHSO4 and HCl gas (which can be identified by its characteristic pungent smell and white fumes with NH3). Even with excess concentrated H2SO4, HCl is not oxidised to Cl2 because the standard reduction potential of Cl2/Cl- (+1.36 V) is higher than that of SO4^2-/SO2 (+0.17 V), so thermodynamically H2SO4 cannot oxidise Cl- under normal conditions. Product: only HCl gas and no Cl2. Bromide (NaBr + conc. H2SO4): initially forms HBr, but some HBr is oxidised to Br2 by the concentrated H2SO4 (the S(VI) in H2SO4 is reduced to S(IV) as SO2): 2HBr + H2SO4 → Br2 + SO2 + 2H2O. Products: HBr gas, Br2 vapour (orange-brown), and SO2. Iodide (NaI + conc. H2SO4): HI is a powerful reducing agent; it readily reduces concentrated H2SO4 in a stepwise manner, producing progressively more reduced sulfur species (SO2, then S, then H2S in the presence of excess HI). Overall: 8HI + H2SO4 → 4I2 + H2S + 4H2O. Products: I2 (violet vapour and dark solid), H2S (distinctive rotten egg smell), and sometimes S (yellow) and SO2. The I2 can be confirmed by the starch test (blue-black colour with starch paper).

The methods for preparing hydrogen halides in the laboratory differ based on their reducing power and the stability of HX toward oxidation. HF preparation: CaF2 (fluorite/fluorspar) + conc. H2SO4 → CaSO4 + 2HF↑ (heating in lead/platinum apparatus). This works because HF is not oxidised by H2SO4 (F- is the weakest reducing agent among halides) and fluorite is easily available. HCl preparation: NaCl + conc. H2SO4 (gently warm) → NaHSO4 + HCl↑. At higher temperatures: NaCl + H2SO4 → Na2SO4 + 2HCl↑. This is a practical and commonly used laboratory preparation. HBr preparation: cannot use H2SO4 (would oxidise HBr). Instead: PBr3 + 3H2O → H3PO3 + 3HBr. OR: red P + Br2 → PBr3 → react with H2O. OR: bromobenzene/alkyl bromide hydrolysis. Alternative: NaBr + H3PO4(conc.) → NaH2PO4 + HBr (H3PO4 is non-oxidising). HI preparation: same problem as HBr, compounded by stronger reducing power of HI. Methods: red P + I2 + H2O → H3PO3 + 3HI (same as HBr method). OR: NaI + H3PO4 → HI + NaH2PO4. OR: H2S + I2 → 2HI + S (H2S reduces I2, but this requires anhydrous conditions). Industrial HCl: by-product of chlorination of organic compounds (e.g., in manufacture of chlorinated solvents, vinyl chloride): R-H + Cl2 → R-Cl + HCl. Also: Mannheim process: 2NaCl + H2SO4 → Na2SO4 + 2HCl (at 600°C).

The decreasing oxidising power of halogens down the group (F2 > Cl2 > Br2 > I2) means that a halogen higher in the group can displace a halide ion lower in the group from its salt solution. This is observed in halogen displacement reactions: Cl2 + 2NaBr → 2NaCl + Br2 (Cl2 displaces Br- from solution). Cl2 + 2NaI → 2NaCl + I2 (Cl2 displaces I- from solution). Br2 + 2NaI → 2NaBr + I2 (Br2 displaces I- from solution). F2 displaces all three: Cl-, Br-, I-. I2 cannot displace any of the others. These displacement reactions form the basis of the "halogen tests" used to identify halide ions in qualitative analysis, and are also used in industrial processes for extracting bromine from sea water (reaction of Cl2 with Br- in sea water) and in the extraction of iodine from brine (Cl2 oxidises I- to I2).

The halogens and their compounds are among the most commercially important chemical substances, with applications spanning medicine, agriculture, industry, and everyday consumer products. Chlorine: manufactured by electrolysis of brine (NaCl solution) in chlor-alkali process, producing NaOH and H2 as co-products. Uses: water purification (kills microorganisms), manufacture of PVC (polyvinyl chloride), paper bleaching, production of pesticides, disinfectants (bleaching powder = CaOCl2, sodium hypochlorite NaOCl). Chlorinated organic compounds: chloroform (CHCl3), carbon tetrachloride (CCl4), DDT (now banned), chlorobenzene. Bromine: extracted from sea water by Cl2 displacement; major uses include 1,2-dibromoethane (antiknock additive in leaded petrol, now phased out), silver bromide (AgBr, used in photographic film), methyl bromide (CH3Br, soil fumigant, being phased out under Montreal Protocol), flame retardants (polybrominated diphenyl ethers, PBDEs, in electronics and textiles). Iodine: obtained from brine associated with oil wells (high iodide content) by Cl2 oxidation; essential micronutrient (thyroid hormone synthesis, iodised salt prevents goitre), iodine/iodophor antiseptics, iodoform (CHI3) antiseptic, contrast agents for X-ray imaging (iodinated compounds). Fluorine: the most reactive of all elements; used in manufacture of fluoropolymers (Teflon, PVDF), fluorinated pharmaceuticals and agrochemicals, uranium enrichment (UF6 used in gaseous diffusion), and etching of silicon in semiconductor manufacturing.

Interhalogen compounds are formed by combination of two different halogen elements, with general formula XYn (n = 1, 3, 5, or 7) where X is the larger, less electronegative halogen and Y is the smaller, more electronegative halogen. Examples: ClF (chlorine monofluoride), ClF3 (chlorine trifluoride), ClF5 (chlorine pentafluoride), BrF (bromine monofluoride), BrF3 (bromine trifluoride), BrF5 (bromine pentafluoride), IF (iodine monofluoride), IF3 (iodine trifluoride), IF5 (iodine pentafluoride), IF7 (iodine heptafluoride), IBr (iodine monobromide), ICl (iodine monochloride). The structures are predicted by VSEPR theory: ClF3 has a T-shaped geometry (3 bonding pairs + 2 lone pairs in trigonal bipyramidal arrangement), BrF5 has a square pyramidal geometry, IF7 has a pentagonal bipyramidal geometry. Interhalogen compounds are generally more reactive than their parent halogens (the weaker X-Y bond compared to X-X makes them easier to dissociate) and are powerful fluorinating and oxidising agents used in industrial fluorination reactions and as rocket propellants (ClF3 and ClF5 are hypergolic with most organic and inorganic fuels).