The acidic strength of HX (X = F, $\mathrm{Cl}$, Br and $\mathrm{I}$) follows the order : $\mathrm{HF}>\mathrm{HCl}>\mathrm{HBr}>\mathrm{HI}$.

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

The enthalpy of dissociation of ${\mathrm{F}}_2$ is smaller than that of ${\mathrm{Cl}}_2$.

Fluorine is stronger oxidising agent than chlorine.

Statement I is incorrect but statement II is true.

Both statement I and statement II are true.

Both statement I and statement II are false.

Statement I is correct but statement II is false.

${\mathrm{CO}}_2<{\mathrm{SiO}}_2<{\mathrm{SnO}}_2<{\mathrm{PbO}}_2$ : Increasing oxidizing power

$\mathrm{HF}<\mathrm{HCl}<\mathrm{HBr}<\mathrm{HI}$ : Increasing acidic strength

${\mathrm{H}}_2\mathrm{O}<{\mathrm{H}}_2\mathrm{S}<{\mathrm{H}}_2\mathrm{Se}<{\mathrm{H}}_2\mathrm{Te}$ : Increasing $p{K}_a$ values

${\mathrm{NH}}_3<{\mathrm{PH}}_3<{\mathrm{AsH}}_3<{\mathrm{SbH}}_3$ : Increasing acidic character

All form monobasic oxyacids.

All are oxidizing agents.

All but fluorine show positive oxidation states.

Chlorine has the highest electron-gain enthalpy.

${\mathrm{Br}}_2>{\mathrm{I}}_2>{\mathrm{F}}_2>{\mathrm{Cl}}_2$

${\mathrm{F}}_2>{\mathrm{Cl}}_2>{\mathrm{Br}}_2>{\mathrm{I}}_2$

${\mathrm{I}}_2>{\mathrm{Br}}_2>{\mathrm{Cl}}_2>{\mathrm{F}}_2$

${\mathrm{Cl}}_2>{\mathrm{Br}}_2>{\mathrm{F}}_2>{\mathrm{I}}_2$

There is strong hydrogen bonding between HF molecules.

The bond energy of HF molecules is greater than in other hydrogen halides.

The effect of nuclear shielding is much reduced in fluorine which polarises the HF molecule.

The electronegativity of fluorine is much higher than for other elements in the group.