Solutions-Class-12-Chemistry-Chapter-7-Elements of Groups 16, 17, and 18-Maharashtra Board

(B) Chlorine

Reason : Chlorine is larger in size and has a lower electronegativity value than fluorine. As a result, it readily accepts an electron and has the highest electron gain enthalpy value.

(C) H₂Te > H₂Se > H₂S > H₂O

(A) O

(C) I₂

(A) NO

(D) HF

(D) Only +6

(C) ClF > ICl > IBr > BrCl > BrF

(D) HOBr + HCl

(B) ClF₃

(D) F

(C) IF₅

(D) [BrF₄] : square planar

(D) 8

The decreasing order of thermal stability: H₂O > H₂S > H₂Se > H₂Te

Oxygen (O) typically has an oxidation state of −2.

In TeO₃, there are three oxygen atoms, contributing a total of −6.

Therefore, to balance this, tellurium must have an oxidation state of +6.

Oxidation state of Te in TeO₃ is +6

Two gases deplete ozone layer :

(i) Nitrogen oxide (NO) released from exhaust systems of car or supersonic jet aeroplanes

(ii) Chlorofluorocarbons (Freons) used in aerosol sprays and refrigerators

• ClO₂ is used as a bleaching agent for paper pulp and textiles.
• It is also used in water treatment.

Bromine reacts instantly with magnesium metal to give magnesium bromide.

Mg(s) + Br₂(l) → MgBr₂(s)

Selenium has two allotropic forms as follows :

• Red (non-metallic) form
• Grey (metallic) form

(Oxidation number of H) + (Oxidation number of S) + (Oxidation number of O) = 0

2 x (+1) + (Oxidation number of S) + 4 x (-2) = 0

Oxidation number of S + 2 - 8 = 0

Hence, oxidation number of 'S' in H₂SO₄ = +6

For HCl, pK_a = − 7.0, hence its dissociation constant is, Ka = 1 x 10⁻⁷.

For HI pK_a = − 10.0, hence its dissociation constant is K_a = 1 x 10⁻¹⁰.

Hence HCl dissociates more than HI.

Therefore HCl is a stronger acid than HI.

Ozone as a reducing agent : Ozone reduces peroxides (O^1-) to oxides (O^2-).

Ozone reduces barium peroxide, BaO₂ to barium oxide, BaO.

BaO₂ + O₃ → BaO + 2O₂

Ozone reduces hydrogen peroxide, H₂O₂ to water, H₂O.

H₂O₂ + O₃ → H₂O + 2O₂

Concentrated sulphuric acid when added to sugar, it is dehydrated giving carbon.

C₁₂H₂₂O₁₁  \(\underrightarrow{conc.\,H_2SO_4}\) 12C + 11H₂O

The carbon that is left behind is called sugar charcoal and the process is called charring.

Chlorine is used :

• for bleaching wood pulp required for the manufacture of paper and rayon, cotton and textiles.
• in the manufacture of organic compounds like CHÇl₃, CCl₄, DDT, dyes and drugs.

(1) 2ICl₃ + 3H₂O → 5HCl + HlO₃ + ICl

(2) I₂ + KClO₃ → ICl + KIO₃

(3) BrCl + H₂O → HOBr + HCl

(4) Cl₂ + ClF₃ → 3ClF

(5) CH₂ = CH₂ + ICL → CH₂l-CH₂Cl

(6) 2XeF₆ + SiO₂ → 2XeOF₄ + SiF₄

(7) XeF₆ + 3H₂O → XeO₃ + 6HF

(8) XeOF₄ + H₂O → XeO₂F₂ + 2HF

XeOF₂ - Xenon monooxydifluoride

XeO₂F₂ - Xenon dioxydifluoride

XeO₃F₂ - Xenon trioxydifluoride

XeO₂F₄ - Xenon dioxytetrafluoride

• The atomic number increases as, ₁₆S < ₁₇Cl < ₁₈
• Due to decrease in atomic size and increase in effective nuclear charge, Cl binds valence electrons strongly.
• Hence ionisation enthalpy of Cl (1256 kJ mol⁻¹) is higher than that of S(1000 kJ mol⁻¹)
• Ar has electronic configuration 3s²3p⁶. Since all electrons are paired and octet is complete, it has the highest ionisation enthalpy, (1520 kJ mol⁻¹)

• The thermal stability of the hydrides of group 16 elements decreases from H₂O to H₂ This is because the bond dissociation enthalpy of the H-E bond decreases down the group.
• Thus, the acidic character increases from H₂O to H₂

 Laboratory methods : By heating chlorates, nitrates and permanganates :

Potassium chlorate in the presence of manganese dioxide on heating decomposes to form potassium chloride and oxygen.

2KClO₃   \(\underrightarrow{ΔMnO_2}\)  2KCl + 3O₂(g)

(i) It oxidises lead sulphide (PbS) to lead sulphate (PbSO₄) changing oxidation state of S from -2 to +6.

PbS(s) + 4O₃(g) → PbSO₄(s) + 4O₂(g)

(ii) Ozone oxidises nitrogen oxide to nitrogen dioxide.

NO(g) + O₃(g) → NO₂(g) + O₂(g)

(i) Hot and concentrated H₂SO₄ acts as an oxidizing agent, since it gives nascent oxygen on heating.

H₂SO₄ \(\underrightarrow{Δ}\) H₂O + SO₂ + [O]

(conc.)

(ii) It oxidises metals and non-metals. For example,

Cu(s) + 2H₂SO₄ → CuSO₄ + 2H₂O + SO₂

(metal)    (conc.)

C  +  2H₂SO₄ → CO₂ +2SO₂ +2H₂O

(non-metal)

• SO₂ molecule has a bent V shaped structure with S-O-S bond angle 119.5° and bond dissociation enthalpy is 297 kJ mol^-1.
• Sulphur in SO₂ is sp² hybridised forming three hybrid orbitals. Due to lone pair electrons, bond angle is reduced from 120° to 119.5°.
• In SO₂, each oxygen atom is bonded to sulphur by a σ and a π
• σ bonds between S and O are formed by sp²-p overlapping.
• One of π bonds is formed by pπ -pπ overlapping while other n bond is formed by pπ -dπ
• Due to resonance both the bonds are identical having observed bond length 143 pm due to resonance,

Partially filled 3p¹_z and 3d¹ orbitals overlap to form π bonds with oxygen atoms.

• The fluorine atom has no d-orbitals in its valence shell and therefore, cannot expand its octet. Thus, fluorine is the most electronegative exhibit -1 oxidation state only.
• Cl, Br, and I exhibit -1, +1, +3, +5, and +7 oxidation states. This is because they are less electronegative than F and possess empty d-orbitals in the valence shell and therefore, can expand the octet.

(a) Chlorine reacts with Fe to give ferric chloride.

2Fe + 3Cl₂ → 2FeCl₃

(b) Chlorine reacts with the excess of ammonia to form ammonium chloride, NH₄Cl and nitrogen.

8NH₃    +   3Cl₂ → 6NH₄Cl  +  N₂

(excess)

• In the laboratory, hydrogen chloride is prepared by heating sodium chloride (common salt) with concentrated sulfuric acid.

NaCl + H₂SO₄ \(\underrightarrow{420K}\) NaHSO₄ + HCl

NaHSO₄ + NaCl \(\underrightarrow{420K}\) Na₂SO₄ + HCl

• HCl gas can be dried by passing it through concentrated sulfuric acid and then collected by upward displacement of air.

Interhalogen compounds : Compounds formed by the combination of atoms of two different halogens are called interhalogen compounds.

• In an interhalogen compound, of the two halogen atoms, one atom is more electropositive than the other.
• The interhalogen compound is regarded as the halide of the more electropositive halogen. For example, ClF, BrF₃, ICl

Hydrochloric acid reacts with ammonia to give white fumes of ammonium chloride.

NH₃ + HCl → NH₄Cl

Hydrochloric acid reacts with sodium carbonate to give sodium chloride, water with the liberation of carbon dioxide gas.

Na₂CO₃ + 2HCl → 2NaCl +H₂O + CO₂

Uses of hydrogen chloride (hydrochloric acid) :

• in the manufacture of chlorine and ammonium chloride,
• to manufacture glucose from corn, starch
• to manufacture dye
• in medicine and galvanising

(i) Hypochlorous acid, HClO :

(ii) Chlorous acid, HClO₂ :

(iii) Chloric acid, HClO₃ :

(iv) Perchloric acid, HClO₄ :

(a) Cl₂ reacts with F₂ in equal volumes at 437 K to give chlorine monofluoride ClF.

Cl2 + F2 \(\underrightarrow{437K}\) 2ClF

(b) Br₂ reacts with excess of F₂ to give bromine trifluoride BrF₃.

Br₂ + 3F₂ \(\underrightarrow{Δ}\) 2BrF₃.

(excess)

Preparation of Xenon Fluorides : Xenon fluorides are generally prepared by direct reaction of xenon and fluorine in different ratios and conditions, such as temperature, electric discharge and photochemical reaction.

(i) Preparation of XeF₂ :

Xe + F₂  \(\frac{sealed\,\,Ni\,\,tube}{400^0C}\)->  XeF₂

(ii) Preparation of XeF₄ :

Xe + F₂ \(\frac{Ni\,\,tube\,\,400^0C}{5-6\,atoms}\)-> XeF₄

1 :  5

Xe + 2F₂ \(\frac{electric\,\,discharge}{-80^0C}\)-> XeF₄

(ii) Preparation of XeF₆ :

Xe + 3F₂ \(\frac{electric\,\,discharge}{low\,\,temp.}-->\) XeF₆

Preparation of XeO₃ : Xenon trioxide ( XeO₃) is prepared by the hydrolysis of XeF₄ or XeF₆

(i) By hydrolysis of XeF₄ :

3XeF₄ + 6H₂O → 2Xe + XeO₃ + 12HF + O₂

(ii) By hydrolysis of XeF6 :

XeF₆ + 3H₂O → XeO₃ + 6HF

Preparation of XeOF₄ : Xenon oxytetrafluoride (XeOF₄) is prepared by the partial hydrolysis of XeF₆.

XeF + H₂O → XeOF₄ + 2HF

Uses of neon (Ne) :

• Neon is used in the production of neon discharge lamps and signs by filling Ne in glass discharge tubes.
• Neon signs are visible from a long distance and also have high penetrating power in mist or fog.
• A mixture of neon and helium is used in voltage stabilizers and current rectifiers.
• Neon is also used in the production of lasers and fluorescent tubes.

Uses of argon (Ar) :

• Argon is used to fill fluorescent tubes and radio valves.
• It is used to provide inert atmosphere for welding and production of steel.
• It is used along with neon in neon sign lamps to obtain different colours.
• A mixture of 85% Ar and 15% N₂ is used in electric bulbs to enhance the life of the filament.

Ozone has molecular formula O₃. The lewis dot and dash structures for O₃ are :

• Infrared and electron diffraction spectra show that O₃ molecule is angular with O-O-O bond angle 117°.
• Both O-O bonds are identical having bond length 128 pm which is intermediate between single and double bonds.
• This is explained by considering resonating structures and resonance hybrid.

Uses of Ozone :

• Sterilizes drinking water by oxidizing germs and bacteria.
• Used as bleaching agent for ivory, oils, starch, wax, and delicate fabrics.
• Purifies air in crowded places like cinema halls, railways, tunnels.
• In industry, used in manufacturing synthetic camphor, potassium permanganate.

• Group 16 elements have smaller atomic and ionic radii due to higher nuclear charge compared to group 15 elements.
• As the group moves down from oxygen to polonium, the atomic and ionic radii increase due to the addition of a new shell.
• The atomic radii increases in the order 0 < S < Se < Te < Po

• The ionisation enthalpy of group 16 elements has quite high values.
• Ionisation enthalpy decreases down the group from oxygen to polonium. This is due to the increase in atomic volume down the group.
• The first ionisation enthalpy of the lighter elements of group 16 (O, S, Se) have lower values than those of group 15 elements in the corresponding periods. This is due to difference in their electronic configurations.

Group 15 : (valence shell) ns² np_x¹ np_y¹ np_z¹

Group 16 : (valence shell) ns² np_x² np_y¹ np_z¹

• Group 15 elements have extra stability of half filled and more symmetrical orbitals, while group 16 elements acquire extra stability by losing one of paired electrons from np_x −orbital forming half filled p−orbitals.

Hence group 16 elements have lower first ionisation enthalpy than group 15 elements.

• The electronegativity values of group 16 elements have higher values than corresponding group 15 elements in the same periods.
• Oxygen is the second most electronegative elements after fluorine.(O = 3.5, F = 4)
• On moving down the group electronegativity decreases from oxygen to polonium.
• On moving down the group atomic size increases, hence nuclear attraction decreases, therefore electro−negativity decreases.

(i) Hot and concentrated H₂SO₄ acts as an oxidizing agent, since it gives nascent oxygen on heating.

H₂SO₄ \(\underrightarrow{Δ}\) H₂O + SO₂ + [O]

(conc.)

(ii) It oxidises metals and non-metals. For example,

Cu(s) + 2H₂SO₄ → CuSO₄ + 2H₂O + SO₂

(metal)    (conc.)

C  +  2H₂SO₄ → CO₂ +2SO₂ +2H₂O

(non-metal)

(i) Sodium sulphite on treating with dilute H₂SO₄ forms SO₂

Na₂SO₃ + H₂SO₄(aq) → Na₂SO₄ + H₂O(l) + SO₂(g)

(ii) Sodium sulphite, Na₂SO₃ on reaction with dilute hydrochloric acid solution forms SO₂.

Na₂SO₃(aq) + 2HCl(aq) → 2NaCl(ag) + H₂O(l) + SO₂(g)

Physical properties of SO₂ :

• It is a colourless gas with a pungent smell.
• It is highly soluble in water and forms sulphurous acid, H₂SO₃

SO₂(g) + H₂O(l) → H2SO₃(aq)

• It is poisonous in nature.
• At room temperature, it liquefies at 2 atmospheres. It has boiling point 263 K.

.

Sulfuric acid is manufactured by Contact process, which involves the following three steps.

(i) Preparation of SO₂ : Sulfur or sulfide ore (iron pyrites) on burning or roasting in air produces sulfur dioxide.

S(s) + O₂(g) \(\underrightarrow{Δ}\) SO₂(g)

4FeS₂(s) + 11O₂ (g) \(\underrightarrow{Δ}\) 2Fe₂O₃(s) + 8SO₂ (g)

(ii) Oxidation of SO₂ to SO₃ : Sulfur dioxide is oxidised catalytically with oxygen to sulfur trioxide, in the presence of V₂O₅ catalyst.

2SO₂(g) + O₂ \(\underrightarrow{V_2O_5}\) 2SO₃(g)

The reaction is exothermic and reversible and the forward reaction leads to decrease in volume. Therefore low temperature (720K) and high pressure (2 bar) are favourable conditions for maximum yield of SO₃.

(iii) Dissolution of SO₃ : Sulfur trioxide gas (from the catalytic converter) is absorbed in concentrated H₂SO₄ to produce oleum.

Dilution of oleum with water gives sulfuric acid of desired concentration.

SO₃(g) + H₂SO₄ → H₂S₂O₇ (oleum)

H₂S₂O₇ + H₂O → 2H₂SO₄

The sulfuric acid obtained by contact process is 96 - 98 % pure.

Flow diagram for manufacture of Sulfuric acid