`=>` All the halogens exhibit `-1` oxidation state.

`=>` However, chlorine, bromine and iodine exhibit `+ 1`, `+ 3`, `+ 5` and `+ 7` oxidation states also as explained in Fig.

● The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms. e.g., in interhalogens, oxides and oxoacids.

● The oxidation states of `+4` and `+6` occur in the oxides and oxoacids of chlorine and bromine.

● The fluorine atom has no `d`-orbitals in its valence shell and therefore cannot expand its octet.

● Being the most electronegative, it exhibits only `–1` oxidation state.

`=>` All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

`=>` The ready acceptance of an electron is the reason for the strong oxidising nature of halogens.

● `color{red}(F_2)` is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the solid phase.

● In general, a halogen oxidises halide ions of higher atomic number.

`color{red}(F_2 + 2 X^(-) → 2F^(-) + X_2 ( X = Cl , Br text(or) I))`

`color{red}(Cl_2 +2X^(-) → 2Cl^(-) +X_2 ( X = Br text(or) I))`

`color{red}(Br_2 + 2I^(-) → 2Br^(-) +I_2)`

`=>` The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials (Table 7.8) which are dependent on the parameters indicated below:

`color{red}(1/2 X_2 (g) overset ( 1/2 Delta_(diss) H^(⊖) ) → X (g) overset (Delta(_eg)H^(⊖))→ X^(-) (s) overset (Delta _(hyd) H^(⊖)) →X^(-) (aq))`

`=>` The relative oxidising power of halogens can further be illustrated by their reactions with water.

● Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids.

● The reaction of iodine with water is nonspontaneous. In fact, `color{red}(I^–)` can be oxidised by oxygen in acidic medium; just the reverse of the reaction observed with fluorine.

`color{red}(2F_2 (g) +2H_2O(l) → 4H^(+) (aq) +4F^(-) (aq) +O_2 (g))`

`color{red}(X_2 (g) +H_2O (l) → HX (aq) +HOX (aq) \ \ \ \ ( text(where ) X = Cl text(or) Br))`

`color{red}(4I^(-) (aq) +4H^(+) (aq) +O_2 (g) → 2I_2 (s) +2H_2O(l))`

`=>` Like other elements of p-block present in second period of the periodic table, fluorine is anomalous in many properties.

`=>` `color{red}(E"xample")` : (i) Ionisation enthalpy, electronegativity, enthalpy of bond dissociation and electrode potentials are all higher for fluorine than expected from the trends set by other halogens.

(ii) Ionic and covalent radii, m.p. and b.p. and electron gain enthalpy are quite lower than expected.

`=>` The anomalous behaviour of fluorine is due to its small size, highest electronegativity, low `color{red}(F-F)` bond dissociation enthalpy, and non availability of `color{red}(d)`-orbitals in valence shell.

`=>` Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements).

`=>` It forms only one oxoacid while other halogens form a number of oxoacids.

`=>` Hydrogen fluoride is a liquid (b.p. `293 K`) due to strong hydrogen bonding. Other hydrogen halides are gases.

(i) `color{green}("Reactivity towards Hydrogen ")` They all react with hydrogen to give hydrogen halides but affinity for hydrogen decreases from fluorine to iodine.

● They dissolve in water to form hydrohalic acids.

● Some of the properties of hydrogen halides are given in Table 7.9.

● The acidic strength of these acids varies in the order: `color{red}(HF < HCl < HBr < HI)`.

● The stability of these halides decreases down the group due to decrease in bond `color{red}(H–X)` dissociation enthalpy in the order : `color{red}(H–F > H–Cl > H–Br > H–I)`.

(ii) `color{green}("Reactivity towards Oxygen ")`: Halogens form many oxides with oxygen but most of them are unstable.

● Fluorine forms two oxides `color{red}(OF_2)` and `color{red}(O_2F_2)`.

● But, only `color{red}(OF_2)` is thermally stable at `298 K`.

● These oxides are essentially oxygen fluorides because of the higher electronegativity of fluorine than oxygen.

● Both are strong fluorinating agents.

● `color{red}(O_2F_2)` oxidises plutonium to `color{red}(PuF_6)` and the reaction is used in removing plutonium as `color{red}(PuF_6)` from spent nuclear fuel.

`=>` Chlorine, bromine and iodine form oxides in which the oxidation states of these halogens range from `+1` to `+7`.

`=>` A combination of kinetic and thermodynamic factors lead to the generally decreasing order of stability of oxides formed by halogens, `color{red}(I > Cl > Br)`.

`=>` The higher oxides of halogens tend to be more stable than the lower ones.

`=>` Chlorine oxides, `color{red}(Cl_2O, ClO_2, Cl_2O_6)` and `color{red}(Cl_2O_7)` are highly reactive oxidising agents and tend to explode. `color{red}(ClO_2)` is used as a bleaching agent for paper pulp and textiles and in water treatment.

`=>` The bromine oxides, `color{red}(Br_2O, BrO_2, BrO_3)` are the least stable halogen oxides (middle row anomally) and exist only at low temperatures. They are very powerful oxidising agents.

`=>` The iodine oxides, `color{red}(I_2O_4, I_2O_5, I_2O_7)` are insoluble solids and decompose on heating. `color{red}(I_2O_5)` is a very good oxidising agent and is used in the estimation of carbon monoxide.

(iii) `color{green}("Reactivity towards Metals ")` : Halogens react with metals to form metal halides. For e.g., bromine reacts with magnesium to give magnesium bromide.

`color{red}(Mg (s) +Br_2 (l) → Mg Br_2 (s))`

● The ionic character of the halides decreases in the order `color{red}(MF > MCl > MBr > MI)` where `color{red}(M)` is a monovalent metal. If a metal exhibits
more than one oxidation state, the halides in higher oxidation state will be more covalent than the one in lower oxidation state.

● For e.g., `color{red}(SnCl_4, PbCl_4, SbCl_5)` and `color{red}(UF_6)` are more covalent than `color{red}(SnCl_2,PbCl_2, SbCl_3)` and `color{red}(UF_4)` respectively.

(iv) `color{green}("Reactivity of Halogens towards other Halogens ")` : Halogens combine amongst themselves to form a number of compounds known as interhalogens of the types `color{red}(XX' , XX_3′, XX_5′)` and `color{red}(XX_7′)` where `color{red}(X)` is a larger size halogen and `color{red}(X′)` is smaller size halogen.