Chapter 12: Group 17

Physically, halogens exist as diatomic molecules and exhibit clear trends down the group: atomic radius increases, while volatility decreases (making iodine the least volatile) and color deepens (from pale yellow fluorine to purple iodine vapor). The decrease in volatility is attributed to increasing instantaneous dipole-induced dipole forces (van der Waals' forces) as molecular size and electron count rise. Chemically, the halogens are powerful oxidizing agents because their atoms readily accept a single electron to achieve a stable noble gas configuration; however, their oxidizing power decreases down the group, correlating with reduced electronegativity and increasing atomic size. This trend is demonstrated by halogen displacement reactions, where a more reactive halogen (like chlorine) can displace a less reactive halogen (like bromine) from its halide salt solution. Reactions with hydrogen form hydrogen halides (HX), with reaction vigor decreasing significantly from explosive fluorine to equilibrium-forming iodine. Correspondingly, the thermal stability of these hydrogen halides decreases down the group, explained by the reduction in bond energy as the hydrogen-halogen bond length increases. Conversely, the halide ions (Cl-, Br-, I-) function as increasingly effective reducing agents going down Group 17. This reducing capability is evident in their reactions with concentrated sulfuric acid: while chloride ions are only protonated to form hydrogen chloride gas, bromide and iodide ions are sufficiently strong reducing agents to be oxidized, leading to the reduction of the sulfuric acid into products like sulfur dioxide, sulfur, or hydrogen sulfide, alongside elemental bromine or iodine. Halide ions can be analytically identified using aqueous silver ions, yielding distinct silver halide precipitates—silver chloride (white), silver bromide (cream), and silver iodide (pale yellow)—which are then differentiated based on their solubility in dilute versus concentrated aqueous ammonia. Finally, the chapter details disproportionation reactions, where chlorine is both oxidized and reduced simultaneously, such as its reactions with cold or hot aqueous sodium hydroxide. A critical practical application involves the chlorination of water for purification, where chlorine disproportionates in water to produce chloric(I) acid and hypochlorite ions, both effective sterilizing agents against bacteria.