Q.Fluorine exhibits only –1 oxidation state whereas other halogens exhibit +1, +3, +5 and +7 oxidation states also. Explain.

Group 17 elements (F, Cl, Br, I) have the outer configuration $ns^2np^5$, one electron short of a stable noble-gas octet, which makes them the most electronegative elements in their respective periods and gives them the most negative electron gain enthalpies of any group, since adding one electron completes a stable octet and releases a large amount of energy. Although fluorine's electron gain enthalpy is numerically less negative than chlorine's — because the very small fluorine atom packs high inter-electronic repulsion into its compact 2p subshell, making it slightly harder to add one more electron — fluorine remains the strongest oxidising agent of the group, since overall oxidising power depends on the full thermodynamic cycle of bond dissociation, electron gain and hydration enthalpy together, and fluorine's very low X-X bond dissociation enthalpy together with the exceptionally large hydration enthalpy of the small $F^-$ ion outweigh its modest electron gain enthalpy. Down the group, the bond dissociation enthalpy of $X_2$ generally decreases (fluorine itself is anomalously weakly bonded because of strong lone-pair-lone-pair repulsion between its very small atoms), and the acid strength of the hydrogen halides increases from HF to HI as the H-X bond weakens with increasing atomic size; a similar trend runs the other way for the Group 15 hydride bases, whose base strength decreases from $NH_3$ to $BiH_3$. Halogens are coloured because their molecules absorb visible light, promoting electrons to higher energy levels, with the absorbed wavelength shifting to longer wavelength (lower energy gap) down the group, from pale yellow $F_2$ to deep violet $I_2$. Fluorine shows markedly anomalous behaviour compared with the rest of the group — for instance it reacts vigo …