Notes-Part-2-Class-12-Chemistry-Chapter-7-Elements of Groups 16, 17, and 18-Maharashtra Board

Topics to be Learn : Part-1 Introduction Occurrence Electronic configuration of elements of groups 16, 17 and 18 Atomic and physical properties of elements of groups 16, 17 and 18 Anomalous behaviour of elements Chemical properties of elements of groups 16, 17 and 18 Topics to be Learn : Part-2 Allotropy Oxoacids Oxygen and compounds of oxygen Compounds of sulphur Topics to be Learn : Part-3 Chlorine and compounds of chlorine Interhalogen compounds Compounds of Xenon

Allotropy :

The property of some elements to exist in two or more different forms in the same physical state is called allotropy.

Property of allotropy in Group 16 elements :

• All the elements of group 16 exhibit allotropy.
• These elements exist in different allotropic modifications.
• Oxygen exists as O₂ and ozone as O₃.
• Sulphur exists as α-sulphur, β-sulphur, g-sulphur, homocyclic sulphur, plastic sulphur, etc. Rhombic sulphur (α sulphur) and mono-clinic sulphur (β sulphur) are the important allotropes. Both are non-metallic in nature.
• Selenium exists in two allotropic forms red (non-qxmetallic) and grey (metallic).
• Tellurium exists in two allotropic forms crystalline form and the amorphous form.
• Po has two fomqs namely α-form and β-form both being metallic.

Know This : Grey selenium allotrope of is a photoconductor used in photocells. The photocopying process : A selenium-coated rotating drum is given a uniform positive charge (step 1) and is then exposed to an image (step 2). Negatively charged toner particles are attracted to the charged area of the drum (step 3) and the image is transferred from the drum to a sheet of paper (step 4). Heating then fixes the image and the drum is flooded with light and cleaned to ready the machine for another cycle (step 5). Figure of photocopying process using Se is as shown below :

Allotropes of sulfur :

Sulfur exhibits numerous allotropic forms. However rhombic sulfur (α-sulfur) and monoclinic sulfur (β-sulfur) are the most important allotropes of sulfur.

Rhombic sulfur :

• Rhombic sulphur has orthorhombic crystals.
• This is the most stable form and common form of sulphur.
• It is pale yellow, having density 2.069 g/cm³ and melting point 385.8 K
• It is insoluble in water, but soluble in CS₂
• It is stable below 369 K and transforms to β-sulphur above this temperature.
• It exists as S8 molecules with a structure of a puckered ring.
• It is obtained by the evaporation of roll sulphur is CS₂.

Monoclinic sulfur :

• Monoclinic sulphur (β-sulphur or prismatic sulphur) has needle-shaped monoclinic crystals.
• It is bright yellow, having a density 1.989 g/cm³ and melting point 393 K.
• It is soluble in CS₂.
• It is stable above 369 K and transforms into α-sulphur below this temperature.
• It exists as, S8 molecules with a structure of a puckered ring.
• It is prepared by melting rhombic sulphur and cooling it till a crust is formed. Two holes are pierced in the crust and the remaining liquid is poured to obtain needle-shaped crystals of monoclinic sulphur.

Remember :

• Several modifications of sulfur containing 6-20 sulfur atoms per ring, have been synsthesised. In the S₈ molecule the ring is puckered and has a crown shape.
• In cyclo - S₆, the ring adopts the chair form. At elevated temperature (~ 1000 K), S₂ is the dominant species which like O₂ is paramagnetic.

Oxoacids

Oxoacids of sulfur : Sulfur forms a number of oxoacids. Some of them are unstable and cannot be isolated. They are known to exist in aqueous solutions or in the form of their salts.

Oxoacids of halogens :

• Fluorine forms only one oxyacid namely, hypofluorous acid or fluoric acid HOP while other halogens form several oxyacids.
• Only four oxoacids have been isolated in pure form: hypofluorous acid (HOF), perchloric acid (HClO4), iodic acid (HIO3), metaperiodic acid (H2IO6).
• The others are stable only in aqueous solutions or in the form of their salts.
• Oxidising power of oxyacids of halogens decreases as the oxidation number of halogens increases.

HClO (1+) > HClO₂ (3+) > HClO₃ (5+) > HClO₄ (7+)

(or HOCl > HOClO > HOClO₂ > HOClO₃ )

• The thermal stability of oxyacids of halogens increases with the increase in oxidation state of halogen. Hence the increasing order of thermal stability is,

HClO (1+) < HClO₂ (3+) < HClO₃ (5+) < HClO₄ (7+)

• The acid strength of the halogen oxoacids increases with the increasing oxidation state of halogen. For example, acid strength increases from HClO, a weak acid (Ka = 3.5 × 10^-8), to HClO₄, a very strong acid (Ka>>1).

Oxygen and Compounds of oxygen :

Dioxygen :

Preparation of Dioxygen by Laboratory methods :

(i) By heating oxygen containing salts such as chlorates, nitrates and permanganates.

2KClO₃(s)    \(\frac{Heat}{MnO_2}\)->   2KCl(s) + 3O₂(g)

(ii) By thermal decomposition of oxides of metals.

2Ag₂O(s)  \(\underrightarrow{heat}\)    4Ag(s) + O₂(g)

2HgO(s)   \(\underrightarrow{heat}\)   2Hg(l) + O₂(g)

2PbO₂(s)   \(\underrightarrow{heat}\)   2PbO(s) + O₂(g)

(iii) By decomposition of hydrogen peroxide in presence of catalyst such as finely divided metals and manganese dioxide.

2H₂O₂(aq)    (\frac{Heat}{MnO_2}\)->  2H₂O(l) + O₂(g)

Dioxygen on a large scale or commercial scale is obtained by two following methods :

Preparation of Dioxygen from water (By electrolysis) : By electrolysis of acidified water, H₂ gas is obtained at cathode and O₂ is obtained at anode.

H₂O(acidified) \(\underleftrightarrow{dilute\,\,H_2SO_4}\)  H⁺ + OH⁻

At cathode : 2H⁺ + 2e⁻ → H₂(g)

At anode ; 2OH⁻ → ½ O₂(g) + H₂ + 2e⁻

Preparation of Dioxygen from air (Industrial method) :

• Carbon dioxide and water vapour is removed fromair and the remaining gases are liquefied.
• O₂ on large scale is obtained by fractional distillation of liquid air.
• On distillation, liquid dinitrogen having low boiling point distils out first leaving behind liquid dioxygen. Then liquid O₂ is distilled out and separated.

Physical properties of Dioxygen  :

• Dioxygen is colourless and odourless gas.
• Dioxygen is sparingly soluble in water, 30.8 cm³ of O₂ dissolves in 1000 cm₃ of water at 293 K. A small amount of dissolved dioxygen is sufficient to sustain marine and aquatic life.
• It liquifies at 90 K and freezes at 55 K.
• Oxygen has three stable isotopes ¹⁶O, ¹⁷O and ¹⁸
• Molecular oxygen, O₂ exhibits paramagnetism.

Chemical Properties of Dioxygen :

(i) Reaction with metals : Dioxygen directly reacts with almost all metals except Au, Pt to form their oxides.

2Ca + O₂ → 2CaO

4Al + 3O₂ → 2Al₂O₃

(ii) Reaction with nonmetals : Dioxygen reacts with nonmetals (except noble gases) to form their oxides.

C + O₂ → CO₂

P₄ + 5O₂ → P₄O₁₀

(iii) Reaction with some compounds :

2ZnS + 3O₂ \(\underrightarrow{Δ}\) 2ZnO + 2SO₂

CH₄ + 2O₂ → CO₂ + 2H₂O

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

4HCl + O₂   \(\underrightarrow{CuCl_2}\)   2Cl₂ + 2H₂O

Uses of Dioxygen :

• Dioxygen is important for respiration to sustain animal and aquatic life.
• It is used in the manufacture of steel.
• It is used in oxyacetylene flame for welding and cutting of metals.
• Oxygen cylinders are widely used in hospitals, high altitude flying and mountaineering.
• It is used in combustion of fuels; for example, hydrazine in liquid oxygen provides tremendous thrust (energy) in rockets.

Simple Oxides :

A binary compound of oxygen with another element is called an oxide.

Oxides can be classified into

• Acidic oxides, CO₂, SO₂, etc.
• Basic oxides, CaO, BaO, etc.
• Amphoteric oxides, Al₂O₃, ZnO, etc
• Neutral oxides, NO, N₂O, CO, etc.

Ozone :

• The stratospheric pool of ozone which is a layer above eax1h’s surface and protects from harmful high energetic ultraviolet (UV) rays is called ozone umbrella or ozonosphere. .
• Ozone (O₃) is an allotrope of oxygen.

Ozone formation :

• In the atmosphere, ozone is naturally formed through photochemical reactions.
• Oxygen present in the lower mesosphere on absorption of solar radiations, is dissociated into two oxygen atoms which oxidise oxygen to ozone.
• One atomic oxygen combines with molecular oxygen to form O₃.

O₂ \(\underrightarrow{UV\,\,light}\)  O + O

O₂ + O → O₃

Laboratory preparation of ozone :

• When a slow dry stream of oxygen is passed through a silent electric discharge, oxygen is converted into ozone (about 10%). The mixture is called ozonized oxygen.

3O₂ \(\underrightarrow{Electric\,\,discharge}\) 2O₃           …ΔH = +142kJ

• It is an endothermic reaction.
• Silent electric discharge prevents the decomposition of ozone.

Physical properties of ozone :

• Gaseous ozone is blue, liquid ozone is dark blue and solid ozone is violet black.
• It has pungent odour hence the name is ozone.
• At higher concentration (about 100 ppm), it is harmful and results into nausea and headache while in small concentration it is harmless.
• Ozone is thermodynamically less stable than oxygen.
• The decomposition of ozone (2O₃ → 3O₂) is exothermic (ΔH < O) and has ΔS > 0. Therefore ΔG < O and hence decomposition of O₃ is spontaneous.
• Ozone is diamagnetic in nature.

Chemical Properties of Ozone :

(i) Oxidising property :

Ozone is a powerful oxidising agents as it easily decomposes to liberate nascent oxygen. (O₃ → O₂ + O).

Ozone oxidises lead sulfide to lead sulfate and iodide ions to iodine.

PbS(s) + 4O₃(g) → PbSO₄(s) + 4O₂(g)

2KI(aq) + H₂O(l) + O₃(g) → 2KOH(aq) + I₂(g) + O₂(g)

Ozone oxidises nitrogen oxide and gives nitrogen dioxide.

NO(g) + O₃(g) → NO₂(g) + O₂(g)

Hence the nitrogen oxide emitted from the exhaust systems of supersonic jet aeroplanes can bring forth depletion of ozone layer in the upper atmosphere.

(ii) Bleaching property : Ozone acts as a good bleaching agent due to its oxidising nature.

O3 → O + O2

Coloured matter + O → colourless matter

Ozone bleaches in absence of moisture so it is also known as dry bleach.

(iii) Reducing property : Ozone reduces peroxides to oxides. e.g.

H₂O₂ + O₃ → H₂O + 2O₂

BaO₂ + O₃ → BaO + 2O₂

(iv) Ozone depletion : Thinning of ozone layer in upper atmosphere is called ozone depletion.

• The ozone (O₃) layer in the upper atmosphere, absorbs harmful UV radiations from the sun, thus protecting people on the earth.
• Depletion of ozone layer in the upper atmosphere is caused by nitrogen oxide released from exhausts system of car or supersonic jet aeroplanes.

NO (g) + O₃ (g) → NO₂ (g) + O₂ (g)

• Depletion (thining) of ozone layer can also be caused by chlorofluoro carbons (freons) used in aerosol and refrigerators and their subsequent escape into the atmosphere.
• The depletion of ozone layer has been most pronounced in polar regions, especially over Antarctica.
• Ozone depletion is a major environmental problem because it increases the amount of ultraviolet (UV) radiation that reaches earth’s surface, thus causing an increase in rate of skin cancer, eye cataracts and genetic as well as immune system damage among people.

Q. How do the coolents in refrigerants deplete the concentration of ozone?

Know This : Ozone reacts with unsaturated compounds containing double bonds to form addition products called ozonides. Ozonides are decomposed by water or dilute acids to give aldehydes or ketones. This reaction is termed as ozonolysis

Compounds of sulfur :

Sulfur dioxide :

Sulphur dioxide, SO₂ is prepared by following methods :

(i) From sulfur : Sulfur dioxide gas can be prepared by burning of sulfur in air.

S(s) + O₂(g) → SO₂(g)

(ii) From sulfite : In the laboratory sulfur dioxide is prepared by treating sodium sulfite with dilute sulfuric acid.

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

(iii) From sulfides : (Industrial method) :

Sulfur dioxide can be prepared by roasting zinc sulfide and iron pyrites.

2ZnS(s) + 3O₂(g)    \(\underrightarrow{Δ}\)   2ZnO(s) + 2SO₂(s)

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

Physical properties of SO₂ :

• Sulfur dioxide is a colourless gas with a pungent smell.
• It is poisonous in nature.
• SO₂ is highly soluble in water and its solution in water is called sulfurous acid.
• It liquifies at room temperature under a pressure of 2 atm and boils at 263 K.

Chemical Properties :

(i) Reaction with Cl₂ : Sulfur dioxide reacts with chlorine in the presence of charcoal (catalyst) to form sulfuryl chloride.

SO₂ (g) + Cl₂ (g) \(\underrightarrow{Charcaol}\) SO₂Cl₂ (l)

(ii) Reaction with O₂ : Sulfur dioxide is oxidised by dioxygen in presence of vanadium (V) oxide to sulfur trioxide.

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

(iii) Reaction with NaOH : Sulfur dioxide readily reacts with sodium hydroxide solution to form sodium sulfite.

2NaOH + SO₂ → Na₂SO₃ + H₂O

(iv) Reaction with Na₂SO₃ : When SO₂ gas is passed through sodium hydroxide solution (NaOH), it forms sodium sulphite, Na₂SO₃ which further with excess of SO₂ forms sodium hydrogen sulphite, NaHSO₃.

Na₂SO₃ + H₂O(l) + SO₂ →  2NaHSO₃

(v) Reducing property :

• Sulfur dioxide acts as a reducing agent in the presence of moisture.
• Moist sulfur dioxide reduces ferric salts into ferrous salts.

2Fe³⁺ + SO₂ + 2H₂O → 2Fe²⁺ + SO₄²⁻ + 4H⁺

• Moist sulfur dioxide decolourises acidified potassium permangnate (VII) solution.

2KMnO₄ + 5SO₂ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄

• Moist sulfur dioxide reduces halogens to halogen acids.

I₂ + SO₂ + 2H₂O → H₂SO4 + 2HI

Uses of sulphur dioxide :

• SO₂ is used in the manufacture of H₂SO₄.
• In refining petroleum and also in sugar industry.
• In the manufacture of NaHSO₃.
• As an antichlor, as a disinfectant and preservative.
• Liquid SO₂ is used as a selective solvent to dissolve many inorganic and organic compounds.
• Sulphur dioxide is used for bleaching wool and silk.

As a bleaching agent in moist condition due to reduction reaction (SO₂ + 2H₂O → H₂SO₄+2[H])

colouring matter + [H] → colourless matter.

On exposing the bleached matter, the colour is restored due to oxidation.

Colourless matter + [O] → coloured matter.

Hence the bleaching action is temporary.

Sulfuric acid, H₂SO₄ :

Preparation : 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.

Q. How is a solution of H₂SO₄ prepared from concentrated sulphuric acid solution?

Physical properties of H₂SO₄ :

• It is a colourless, dense, oily liquid having specific gravity 1.84 at 298 K.
• It has freezing point 283 K and boiling point 611 K.
• It is a strong dehydrating agent and dissolves in water with the evolution of a large amount of heat.
• It is a strong dibasic acid and acts as an oxidizing agent.
• Due to hydrogen bonding, it is viscous.
• It is highly corrosive and produces severe burns on skin.

Chemical Properties of H₂SO₄ :

(i) Acidic Property : Sulfuric acid ionises in aqueous solution in two steps.

H₂SO₄ (aq) + H₂O(l) → H₃O⁺(aq)+ HSO₄(aq)  ..(Ka > 10)

HSO₄(aq)+ H₂O(l) → H₃O⁺ (aq) + SO₄²⁻(aq)  …(Ka = 1.2 × 10⁻²)

The greater value of Ka (Ka>10) means that H₂SO₄ is largely dissociated into H⁺ and HSO₄ ions. Thus H₂SO₄ is a strong acid.

(ii) Reaction with metals and nonmetals (oxidising property) : Metals and nonmetals both are oxidised by hot, concentrated sulfuric acid which itself gets reduced to SO₂.

Cu + 2H₂SO₄ (Conc.) → CuSO₄ + SO₂ + 2H₂O

S + 2H₂SO₄(Conc.) → 3SO₂ + 2H₂O

C + 2H₂SO₄(Conc.)  → CO₂ + 2SO₂ + 2H₂O

(iii) Dehydrating property : Concentrated sulfuric acid is a strong dehydrating agent.

Sulfuric acid removes water from sugar and carbohydrates. Carbon left behind is called sugar charcoal and the process is called charring.

C₁₂H₂₂O₁₁ conc. → H₂SO₄ 12C + 11H₂O

(iv) Reaction with salts : Concentrated sulfuric acid decomposes the salts of more volatile acids to the corresponding acid

e.g.

NaCl + H₂SO₄ → NaHSO₄ + HCl

KNO₃ + H₂SO₄ → KHSO₄ + HNO₃

CaF₂ + H₂SO₄ → CaSO₄ + 2HF

Remember : Oxidizing properties of sulfuric acid depend on its concentration and temperature. In dilute solutions, at room temperature, H₂SO₄ behaves like HCl, oxidizing metals that stand above hydrogen in the e.m.f. series.

Fe(s) + 2H⁺(aq) → Fe²⁺(aq) + H₂(g)

Hot, concentrated H₂SO₄ is a better oxidizing agent than the dilute, cold acid. It oxidises metals like copper.