(i) The given statement is true.
(ii) The relative acidic strength of HF, HCI, HBr and HI depends upon their bond dissociation enthalpies.
(iii) Since, the bond dissociation enthalpy of H-X bond from H-F to H-I as the size of atom increases from F to l. Therefore, the acidic strength increases in the opposite order as:
H - F < H - CI < H - Br < H - I

(i) SO₂ present in the atmosphere dissolves in rain water and causes acid rain.
(ii) Acid rain destroys buildings, inhibits plant growth and also causes eye irritation in human.

(i) Ozone acts as a powerful oxidizing agent because ozone liberates nascent oxygen very easily on decomposition.
\({ O }_{ 3 }\longrightarrow { O }_{ 2 }+\left[ O \right] \)
(ii) This nascent or atomic oxygen brings about the oxidations it is very reactive.

(i) Sulphur has great catenation property and forms large molecule having high weight, which results higher melting point.
(ii) While oxygen exists as diatomic molecule of less molecular weight, hence has very low melting point.

(i) The electron gain enthalpy for oxygen is less negative because of its small size due to which the electron repulsions in the relatively small 2p-subshell are comparatively large.
(ii) Hence the Incoming electrons are not accepted with the same ease as in case of sulphur as it has relatively large size.

(i) HCl < HBr < HI < HF
(ii) Fluorine being the most electronegative element cannot have positive oxidation state. Other halogens have d-orbitals, therefore can expand their octet.

This is due to,
(i) Lowbond dissociation enthalpy of F-F-bond
(ii) High hydration enthalpy of F^-.

Element of group 15 have stable half-filled p-orbitals hence large amount of energy is required to remove electrons as compared to group 16, which has an incomplete p subshell.

(i) All the halogens have strong oxidising property due to their high electron affinity.
(ii) Fluorine is the strongest oxidising agent. It oxidises other halide ions into halogens in solution or when dry.
\({ F }_{ 2 }+{ 2X }^{ - }\longrightarrow { 2F }^{ - }+{ X }_{ 2 }\) (where X^- = Cl^-, Br^-, I^-)

Yellow phosphorus reacts with alkali on boiling in an inert atmosphere liberating phosphine. Phosphorus acts as a reducing agent.
\({ P }_{ 4 }+NaOH+{ H }_{ 2 }O\longrightarrow \underset { Sodiumhypophosphite }{ { 3NaH }_{ 2 }{ PO }_{ 2 } } +\underset { Phosphine }{ { PH }_{ 3 } } \uparrow \)

White Phosphorous glows in the dark due to oxidation which is called as phosphorescence.

(i) Nitric acid is used as a oxidising agent and in the preparation of aqua regia.
(ii) Salts of nitric acid are used in photography (AgNO₃) and gun powder for fire arms (NaNO₃)·

Ammonia is prepared in the laboratory by heating an ammonium salt with a base.
\({ 2NH }_{ 4 }^{ + }+{ OH }^{ - }\longrightarrow { 2NH }_{ 3 }+{ H }_{ 2 }O\)
\({ 2NH }_{ 4 }Cl+CaO\longrightarrow { CaCl }_{ 2 }+{ 2NH }_{ 3 }+{ H }_{ 2 }O\)

(i) Nitrogen is used for the manufacture of ammonia, nitric acid and calcium Cyanamide etc.
(ii) Liquid nitrogen is used for producing low temperature required in cryosurgery and so in biological preservation.

(i) Nitrogen is separated industrially from liquid air by fractional distillation.
(ii) Pure nitrogen gas can be obtained by the thermal decomposition of sodium azide about 575K.
\(2Na{ N }_{ 3 }\longrightarrow 2Na+{ 3N }_{ 2 }\)