NCERT Solution Class 12th Chemistry Chapter – 7 The p-Block Elements Notes

NCERT Solution Class 12th Chemistry Chapter – 7 The p-Block Elements

NCERT Solution Class 12th Chemistry Chapter – 7 The p-Block Elements

?Chapter – 7?

✍The p-Block Elements✍

?Notes?

The p-block elements are placed in groups 13 to 18 of the periodic table. Their valence shell electronic configuration is ns²np⁶ to ns²np⁶ (except He which has 1s² electronic configuration.)

Group 15 Elements: Group 15 elements include Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi).

Table 7.1: Atomic and Physical Properties of Group 15 Elements

^αE₁₁₁ single bond (E = element); ^bE^3-; ^cE³⁺; ^dWhite phosphorus; ^eGrey α-format 38.6 atm of Sublimation temperature; g At 63 K h Grey α-form; Molecular N₂

As we go down the group, there is a shift from non-metallic to the metallic character through metalloids character. N and P are non-metals, AS and Sb are metalloids and Bi is a typical metal.

Occurrence: N₂ comprises 78% by volume of the atmosphere. In the earth’s crust, it occurs as sodium nitrate (NaNO₃) called Chile saltpetre and potassium nitrate (KNO₃) called Indian saltpetre. It is found in the form of proteins in plants and animals. Phosphorus (P) occurs in minerals of the apatite family Ca₉ (PO₄)₆. CaX₂ [X = F, Cl, or OH] e.g. fluorapatite Ca₉ (PO₄)₆. CaF₂ which are the main components of the phosphate rocks. P is an important constituent of animal and plant matter and is present in bones and living cells.

Electronic Configuration: The valence shell electronic configuration of these elements is ns2np3. The s orbital is completely filled and p orbitals are exactly half-filled (p_x¹ p_y¹ P_z¹). Thus their electronic configuration is extra stable.

Atomic and Ionic Radii: Covalent and ionic radii increase in size down the group due to the addition of the orbit each time. There is a considerable increase from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and /or f orbitals in heavier members.

Ionization Enthalpy: I.E. values of Group 15 elements is much greater than that of Group 14 elements due to the extra-stability of half-filled p-orbitals and smaller size. It decreases from top to bottom within the group due to a gradual increase in atomic size. The order of successive ionisation enthalpies is Δ_iH₁ < Δ_iH₂ < Δ_iH₃.

Electronegativity: Electronegativity decrease down the group. The difference is not much pronounced amongst the heavier elements.

Physical Properties: All the elements are polyatomic. Dinitrogen (N₂) is a diatomic gas while others like P₄ etc. are solids. The b. pts; in general, increase from top to bottom in the group, but them. pts increase up to As and then decrease up to Bi.

Except for Nitrogen, all the elements show allotropy.

Chemical Properties:
Oxidation states and trends in chemical reactivity: Tire common oxidation states of these elements are – 3, + 3 and + 5. The tendency to exhibit – 3 oxidation state decreases down the group due to an increase in size and metallic character.

N does not show + 5 oxidation state due to the absence of d-orbitals and Bi does not show + 5 oxidation state due to the Inert pair Effect. The only well characterised Bi (V) compound is BiF₅. N exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen. P also exhibits + 1 and + 4 oxidation states in some oxoacids.

In the case of N, all oxidation states from + 1 to + 4 tend to disproportionate into acid solution1. For example
3 HNO₂ → HNO₃ + 2 NO + H₂O

Similarly, in the case of P nearly all intermediate oxidation states disproportionate into + 5 and – 3 both in alkali and acid,

Anomalous properties of Nitrogen: N differs from its congeners due to

1. its smaller size.
2. high electronegativity and high ionisation enthalpy.
3. non-availability of d-orbitals.

Nitrogen has a unique ability to form an-pit multiple bonds with itself and with other elements having small size and having high electronegativity (e.g., C, O)

N₂ has a triple bond: N ≡ N [one a and two K bonds]
∴ Its bond enthalpy (941.4 kJ mol^-1) is very high.
P and As can form dπ-pπ bonds with transition metals.

1. Reactivity towards hydrogen: All the elements of Group 15 form hydrides of the type EH₃ where E = N, P, As, Sb, Bi. The stability of hydrides decreases from NH₃ to BiH₃ and hence reducing character increases from NH₃ to BiH₃. Basicity of hydrides also decreases in the order NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃.

Table 7.2: Properties of Hydrides of Group 15 Elements

2. Reactivity towards oxygen: All these elements form oxides: E₂O₃ and E₂O₅. The oxide in the higher oxidation state is more acidic.
Their acidic character decreases down in a group. The oxides of the type EO₃ of N and Pare purely acidic, that of As, Sb amphoteric and those of Bi predominantly basic.

3. Reactivity towards halogens:
2E + 3X₂ → EX₃
2E + 5X₂ → 2EX₅
N does not form pentahalide due to absence of d-orbitais. All trihalides are stable except that of N.

Pentahalides are more covalent than trihalides.

4. Reactivity towards metals: These elements combine with metals showing – 3 oxidation state in compounds like Ca₃N₂, Ca₃P₂, Mg₃Bi₂.

Dinitrogen, N₂:
Preparation:
1. Commercially dinitrogen is produced by the liquefaction and fractional distillation of air when N₂ distils the first.

2. Lab. Method:
NH₄Cl (aq) + NaNO₂ (aq) → N₂ (g) + 2H₂O (l) + NaCl (aq)

3. Pure N₂ can be obtained from sodium or barium azide.

→ Properties: Dinitrogen is a colourless, odourless, tasteless and non-toxic gas. It has two stable isotopes: ¹⁴N, ¹⁵N. It has very low solubility in water.

Because of the high bond enthalpy of N ≡ N, dinitrogen is rather inert at room temperature.

At high temperature, it directly combines with some metals to form largely ionic nitrides and with non-metals, covalent nitrides.

Haber’s process:

Uses:

1. Manufacture of NH₃ and Calcium Cyanamide.
2. To create an inert atmosphere.
3. Liquid N₂ is used as a refrigerant,

Ammonia:
Preparation:
1. From Urea:
NH₂CONH₂ + 2H₂O → (NH₄)₂ CO₃ → 2NH₃ + H₂O + CO₂

2. From ammonium salts and caustic soda/lime.
2NH₄Cl + Ca(OH)₂ → 2NH₃ + 2H₂O + CaCl₂
(NH₄)₂ SO₄ + 2 NaOH → 2NH₃ + 2H₂O + Na₂SO₄

3. On a large scale, it is prepared by Haber’s process.

The pressure is 200 × 10⁵ Pa and iron oxide with small amounts of K₂O and Al₂O₃ is used as a catalyst,

Properties of Ammonia:

1. It is a colourless gas with a pungent smell.
2. It shows hydrogen bonding in solid and liquid states.

Structure of NH3:

3. It is highly soluble in water. Its aqueous solution is weakly basic.
NH₃ (g) + H₂O (l) ⇌ NH₄⁺ (aq) + OH^– (aq)

4. It forms ammonium salts with acids.
NH₃ + HCl → NH₄Cl

5. As a weak base it precipitates the hydroxides of many metals
from their salt solutions, e.g.,

6. Due to its tendency to donate a pair of electrons on N, it acts as a Lewis-Base, forms complexes with Cu²⁺, Ag⁺.

Uses:

1. Production of nitrogenous fertilizers.
2. Production of nitric acid.
3. Liquid ammonia is a non-aqueous solvent.
4. Liquid NH₃ is a referigerant.

Oxides of Nitrogen: Nitrogen forms a number of oxides in different oxidation state. The details of names, formulae, ox. state, preparation, colour etc. are given:

Lewis dot main resonance structures and bond parameters of oxides are given on the next page.

Table 7.4 Structures of Oxides of Nitrogen

Nitric Acid, HNO₃
Preparation: It is prepared in the laboratory by heating NaNO₃ or KNO₃ and cone. H₂SO₄.

Method of manufacture of Nitric acid by Ostwald’s process.

Step:
1. Catalytical oxidation of NH₃ by O, or air.

2. NO thus formed combines with 02 giving NO₂
2NO (g) + O₂ (g) ⇌ 2NO₂(g).

3. N02 SO formed dissolves in water to give HNO₃.
3NO₂ + H₂O(l) → 2 HNO₃ + NO (g)

Structure of HNO₃: In the gaseous state, HNO₃ exists as a planar molecule.

In an aqueous solution, it behaves as a strong acid.
HNO₃ (aq) + H₂O (Z) → (aq) + NO₃ (aq).
Concentrated HNO₃ is a strong oxidizing agent.

Reaction with Cu:
3Cu + 8 HNO₃ (dil) → 3 Cu (NO₃)₂ + 2 NO + 4 H₂O
Cu + 4 HNO₃ (cone.) → Cu (NO₃)₂ + 2NO₂ + 2H₂O

Reaction With Zinc
4 Zn + 10 HNO₃ (dil.) → 4 Zn (NO₃)₂ + 5 H₂O + N₂O
Zn + 4HNO₃ (cone.) → Zn (NO₃)₂ + 2H₂O + 2NO₂

Reaction with I₂

Reaction with Carbon
C + 4 HNO₃ → CO₂ + 2H₂O + 4NO₂

Reaction with Sulphur
S₈ + 48 HNO₃ (Cone.) → 8H₂SO₄ + 48 NO₂ + 12 H₂O

Reaction with Phosphorus
P₄ + 20 HNO₃ → 4H₃PO₄ + 20 NO₂ + 4H₂O

Brown Ring Test for nitrates (NO₃^–)
NO₃– + 3 Fe²⁺ + 4 H⁺ → NO + 3 Fe³⁺ + 2H₂O.

Uses of Nitric acid:

1. Oxidant in the laboratory.
2. manufacture of ammonium nitrate for fertilizers.
3. manufacture of explosive like TNT, nitroglycerine. Phosphorus: Phosphorus exists in three important allotropic forms: white, red and black.

White Phosphorus: White Phosphorus consists of discrete tetrahedral molecules as shown:

White phosphorus P₄

Red Phosphorus: Red Phosphorus is obtained by heating white phosphorus at 573 K in an inert atmosphere for several days. It is much less reactive than white phosphorus. It is polymeric, consisting of P4 tetrahedral linked together as shown in Fig. below.

Red Phosphorus

Black Phosphorus has two forms α-black phosphorus and β-black phosphorus. α-Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803 K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals. It does not oxidise in the air. β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not bum in the air up to 673 K.

Phosphine (PH₃) Preparation:
1. From calcium phosphide and H₂O/dil. HCl
Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃.
Ca₃P₂ + 6HCl → 3CaCl₂ + 2PH₃.

2. Lab. Method: By heating white phosphorus with cone. NaOH in an inert atom of CO₂.
P₄ + 3NaOH +- 3H₂O → PH₃ + NaH₂PO₂ (Sod. hypophosphite)

Properties:

1. Colourless gas with rotten fish smell, highly poisonous.
2. Explodes in contact with traces of oxidizing agents like HNO₃ Cl₂ etc.,
3. Slightly soluble in water.
4. When absorbed in CuSO₄ or HgCl₂ solutions, phosphides are obtained.
   3 CuSO4 + 2 PH₃ → Cu₃P₂ + 3H₂SO₄
   3HgCl₂ + 2PH₂  → Hg₂ P₂ + 6HCl.
5. PH₃ is weakly basic like NH₃.
   PH₃ + HBr → PH₄Br.

Phosphorus halides: P forms two types of halides, PX₃ [X = F, Cl, Br, I]and PX₅[X = F,Cl,Br]
Phosphorus trichloride (PCl₃)

Preparation:
P₄ + 6Cl₂ → 4PCl₃.

Properties:

1. Colourless oily liquid.
2. Hydrolysis in the presence of mixture.
   PCl₃ + 3H₂O → H₃PO₃ + 3HCl

Shape: It has a pyramidal shape as shown in which phosphorus is sp³ hybridized.

Shape of PCl₃

Phosphorus pentachloride (PCl5)
Preparation:

Properties:

1. It is a yellowish-white powder.
2. In moist air, it hydrolyses first to PoCl₃, and then H₃PO₄.
   PCl₅ + H₂O → POCl₃ + 2 HCl
   POCl₃ + 3H₂O → 4 H₃PO₄ + 3HCl
3. On heating, it sublimes but decomposes on stronger heating.

In gaseous and liquid phases, it has a trigonal bipyramidal structure shown below. The three equatorial P-Cl bonds are equivalent, while the two axial bonds are longer than equatorial bonds. This is due to the fact that the axial bond pairs suffer more repulsion as compared to equatorial bond pairs.

In the solid-state, it exits as an ionic solid, [PCl₄]⁺[PCl₆]^– in which the cation, [PCl₄]⁺ is tetrahedral and the anion, [PCl₆]⁺ octahedral.

Oxo-Acids of Phosphorus: P forms a number of oxo-acids like

1. Hypophosphorous acid (H₃PO₂) (ox. state + 1)
2. Orthophosphoric acid (H₃PO₃) (ox. state + 3)
3. Pyrophosphrus acid (H₄P₂O₅) (ox. state + 3)
4. Hypophosphoric acid (H₄P₂O₆) (ox. state +4)
5. Orthophosph’oric acid (H₃PO₄) (ox. state + 5)
6. Pyrophosphoric acid (H₄P₂O₇) (ox. state +5)
7. Metaphosphoric acid (HPO₃) (ox. state + 5)

The compositions of the oxo-acids are interrelated in terms of loss or gain of H₂O molecule or oxygen atom.
These acids in the + 3 oxidation state of P tend to disproportionate as follows:

The acids which contain the P-H bond are strong reducing agents.

Group 16 Elements: Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po) constitute Group 16 of the periodic table. They are also called Chalcogens (ore-forming).

Occurrence: Oxygen is the most abundant of all the elements of the earth. Dry air contains 20.95% oxygen by volume. However Sulphur is present in the earth’s crust to 0.03 – 0.1% only. Combined Sulphur is present in gypsum CaSO₄. 2H₂O, Epsom salt MgSO₄.7H₂O. Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores.

Electronic Configuration: The elements of Group 16 have six electrons in the outermost shell and have ns2np4 general electronic configuration.

Atomic and Ionic Radii: Due to the increase in the no. of shells, atomic and ionic radii increase from top to bottom, The size of the oxygen atom is, however, exceptionally small.

Ionisation Enthalpy: Ionisation enthalpy decreases down the group.

Table: Some Physical Properties of Group 16 Elements

^aSingle bond; ^bApproximate value; ^cAt the melting point; ^dRhombic sulphur; ^eHexagonal grey; ^fMonoclinic form, 673 K.
Oxygen shows oxidation states of + 2 and +1 in oxygen fluorides OF₂ and O₂F₂ respectively.

Electron Gain Enthalpy: Because of the compact nature of the oxygen atom, it has less negative electron gain enthalpy than sulphur.

Electronegativity: Next to F, oxygen has the maximum value of electronegativity value amongst the elements. Within the group, its value decreases from top to bottom implying that metallic character increases from O to Po.

Physical Properties: Oxygen and Sulphur are non-metals, Selenium and Tellurium metalloids, whereas Polonium is a metal. Pollonium is radioactive [t_1/2 = 13.8 days]

All these elements exhibit allotropy.
M.Pts and B.Pts. increase with the increase in atomic no. down the group. The large difference between the M.Pts and B.Pts. of oxygen and sulphur may be explained on the basis of their atomicity: Oxygen exists as diatomic molecules (O₂) whereas Sulphur exists as polyatomic (S₈).

Chemical Properties:
Oxidation states and trends in chemical reactivity:
O = – 2, -1, +1, (in O₂F₂), + 2 (in OF₂)
S = – 2, + 2, +4, + 6 (in SF₆)
Se = – 2, (+ 2), +4, + 6 (SeF₆)
Te=-2,(+2), +4, + 6(TeF₆)
Po = + 2, + 4
[Within brackets less stable.]

The stability of + 4 oxidation state increases and that of + 6 decreases due to the Inert-pair Effect. Bonding in + 4, + 6 oxidation states are primarily covalent.

The anomalous behaviour of oxygen is due to its

1. small size,
2. high electronegativity,
3. absence of d-orbitals.

1. Reactivity with Hydrogen.
(a) There is strong hydrogen bonding in H₂O which is not found in H₂S.
(b) All the elements of Group 16 form hydrides of the type H₂E[E = S, Se, Te, Po]
(c) H₂O is neutral. Other hydrides are acidic and acidic character increases.
H₂O < H₂S < H₂Se < H₂Te < H₂Po
(d) Thermal stability decreases:
H₂O > H₂S > H₂Se > H₂Te > H₂Po
(e) H₂O does do not have reducing character. Reducing character increases from H₂S < H₂Se < H₂Te.

2. Reactivity with Oxygen: All these elements form oxides of EO₂ and EO₃ types [E = S, Se, Te, or Po]. Both types are acidic in nature.

3. Reactivity towards halogens: Elements of Group 16 form a large no. of halides of the type, EX₆, EX₄, EX₂ where X = halogen. The stability of the halides decreases in the order F^– > Cl > Br > I^–

Dioxygen, O₂:

Industrially, O₂ is obtained from air by first removing CO₂ and water vapour and then, the remaining gases are liquified and fractionally distilled to give N₂ and O₂.

Properties of O₂:

1. Colourless, odourless gas.
2. 30.8 cm³3 of O₂ soluble per litre of water at 293 K.
3. It has three isotopes ¹⁶O, ¹⁷O, ¹⁸O.
4. Molecular oxygen (O₂) is paramagnetic.
5. Reactions with metals, non-metals and other compounds are as follows:

Uses:

• It is used in oxyacetylene welding.
• Manufacture of many metals, particularly steel.
• Oxygen cylinders are widely used in hospitals, high altitude flying and mountaineering.
• The contraction of fuels, e.g., hydrazines in liquid oxygen, provides tremendous thrust in rockets.

Simple Oxides: A binary compound of oxygen with another element is called oxide.

Oxides can be simple like MgO, Al₂O₃ or mixed like Pb₃O₄, Pe₃O₄.
In general metallic oxides like CaO, BaO, Na₂O are basic. Non-metal oxides like SO₂, CO₂ are acidic.

Some oxides like Al₂O₃ are amphoteric. They react with acids as well as bases.
Al₂O₃ + 6 HCl → 2AlCl₃ + 3H₂O.
Al₂O₃ + 2NaOH → 2Na AlO₂ + H₂O

Ozone, O₃: Ozone is an allotropic form of oxygen. At a height of 20 km from sear level, it is formed from atmospheric oxygen in the presence of sunlight.

The ozone layer protects the earth’s surface from UV rays.

Preparation:

ΔH° = + 142 kJ mol^-1

Properties:

1. It is a pale blue gas, dark blue liquid and violet-black solid.
2. O₃ is thermodynamically unstable w.r.t. O₂.
   ∴ It acts as a powerful oxidizing agent.

O₃ → O₂ + O
1.  It oxidizes lead sulphide to lead sulphate.
PbS + 4O₃(g) → PbSO₄ (s) + 4O₂

2. It oxidizes I^– ions to iodine
2I^– (aq) + H₂O (l) + O₃ (g) → 2 OH^– (aq) + I₂ (s) + O₂ (g)

3. It oxidizes NO to NO₂
NO (g) + O₃ (g) → NO₂ (g) + O₂ (g)

Structure: The two oxygen-oxygen bond lengths in the ozone molecule are identical (128 pm) and the molecule is angular as expected with a bond angle of about 117°. It is a resonance hybrid of two main forms:

Uses:

1. It is used as a germicide, disinfectant and for sterilising water.
2. It is also used for bleaching oils, ivory, flour, starch.
3. It is an oxidizing agent.

Sulphur-Allotropic Forms: Out of numerous allotropes, the two most important forms are:

• Yellow rhombic (α-sulphur),
• Monoclinic (β-Sulphur)

The temperature 369 K is called Transition temperature.

The structure of (a) S₈ ring in rhombic sulphur and (b) S₆ form:

Both rhombic and monoClinic Sulphur have S₈ molecules. These S₈ molecules are packed to give different crystal structures. The S₈ ring in both forms is puckered and has a crown shape.

In cyclo-S₆ the ring adopts the chair form. At elevated temperatures (~ 1000 K) S₂ is the dominant species and is paramagnetic like O₂

Sulphur Dioxide (SO₂):
Preparation:
1. By burning S in air or O₂.
S(s) + O₂(g) → SO₂(g)

2. Lab. Method: By treating a sulphite with dilute H₂SO₄4.
SO₃^2- (aq) + 2H⁺ (aq) → H₂O (l) + SO₂(g)

3. Industrially, it is produced by roasting of sulphide ores:

Properties:
1. It is a colourless gas with a pungent smell.

2. It is highly soluble in water.

3. It liquefies at room temp, under a pressure of 2 atmospheres.

4. Liquified SO₂ boils at 263 K.

5, SO₂ when dissolved in water, forms sulphurous acid.
SO₂(g) + H₂O (l) → H₂SO₃ (aq)

6. Reaction with NaOH:

7. Reaction with Cl₂:

8. Reaction with O₂:

9. Moist SO₂ behaves as a reducing agent. It reduces iron (III) ions to iron (II) ions and decolourised acidified potassium permanganate(VII) solution. It is a test for SO₂ gas.
2Fe³⁺ (aq) +SO₂ + 2H₂O → 2 Fe²⁺ (aq) + SO₄^2- + 4H⁺
5 SO₂ + 2 MnO₄^– + 2H₂O → 5 SO₄ ^2- + 4H⁺ + 2 Mn²⁺

Structure: The molecule of SO₂ is angular. It is a resonance hybrid of the two canonical structures

Uses:

• Refining of petroleum, sugar
• Bleaching wool and silk.
• Antichlor, disinfectant and preservative.
• Manufacture of H2S04 and other industrial chemicals.
• Liquid S02 is used as a solvent.

Oxo-Acids of Sulphur: Sulphur forms a number of oxo-acids such as H₂SO₂, H₂S₂O₃, H₂S₂O₄, H₂S₂O₅, H₂S_xO₆ (X = 2 to 5), H₂SO₄, H₂S₂O₇, H₂SO₅, H₂S₂O₈.

Structures of Sulpiwrous and Sulphuric acids

Sulphuric Acid (H₂SO₄): It is called the king of chemicals and is one of the most widely used industrial chemicals.

→ Manufacture: Sulphuric acid is manufactured by Contact Process which involves 3 steps:

1. Burning of Sulphur or Sulphide ores in the air to produce SO₂.
2. Conversion of SO₂ into SO₃ by reaction with air in the presence of catalyst V₂O₅.
3. Absorption of SO₃, in H₂SO_4, to give Oleum (H₂S₂O₇).
   The key step is the catalytical oxidation of SO₂ with O₂ to give SO₃ in the presence of V₂O₅.
   
   The reaction is exothermic, reversible and proceeds with a decrease in volume.
   ∴ Low temperature (720 K) and high pressure (2 bar) are favourable conditions for maximum yield.

The tempi should not be very low otherwise the rate of reaction will below.

The SO₃ gas from the catalytical converter is absorbed in concentrated H₂SO₄ to produced Oleum (H₂S₂O₇) which on dilution with water gives H₂SO₄ of the desired concentration.
SO₃ + H₂SO₄ → H₂S₂O₇ (oleum)
H₂S₂O₇ + H₂O → 2 H₂SO₄

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

Flow diagram the iiai1ufat true of sulphuric acid

Properties:

1. H₂SO₄ is a colourless, dense, oily liquid.
2. It dissolves in water producing a large amount of heat. Hence it should be slowly added to water (and never water in acid)

The chemical properties of H₂SO₄ are due to its
(a) low volatility,
(b) strong acidic character,
(c) strong affinity for water,
(d) ability to act as an oxidizing agent.

In water, it ionises in 2 steps:
1. H₂SO₄ (aq) + H₂O (l) → H₃O+ (aq) + HSO₄^– (aq)
K_a1 = very large (> 10)

2. HSO₄^–(aq) + H₂O (l) → H₃O⁺ (aq) + SO₄ ^2- (aq);
K_a2 = 1.2 × 10^-2.

(a) It is used to prepare more volatile acids:
NaNO₃ + H₂SO4 → NaHSO₄ + HNO₃.
2 NaCl + H₂SO₄ → Na₂SO₄ + 2 HCl

(b) Cone. H₂SO₄ is a strong dehydrating agent

(c) Hot cone. H₂SO₄ is a strong oxidising agent
H₂SO₄ → H₂O + SO₂ + O

(d) Reaction with metals and non-metals:
Cu + 2 H₂SO₄ → CuSO₄ + SO₂ + 2H₂O
C + 2H₂SO₄ (cone.) → CO₂ + 2SO₂ + 2H₂O

Uses of H₂SO₂ :

• It is used in the manufacture of hundreds of other compounds and also in many industrial processes.
• Manufacture of fertilizers.
• Petroleum refining and manufacture of pigments, paints and dyestuff intermediates.
• Detergent industry.
• Metallurgical applications.
• Storage batteries and so on.

Group 17 Elements: Fluorine, chlorine, .bromine, iodine and astatine are the members of Group 17. They are collectively called halogens (sea-salt producers)

Occurrence: Fluorine is present as CaF₂ (fluorspar) and Na₃ AlF₆ (cryolite). Seawater contains chlorides, bromides and iodides of sodium, potassium, magnesium and calcium. The deposits of dried up seas contain NaCl and carnallite, KCl. MgCl₂,6Fl₂O. Seaweeds contain 0.5% of iodine and Chile saltpetre contains up to 0.2% of sodium iodate.

Electronic Configuration: All these elements have seven electrons in their valence shell [ns²np³] which is 1 electron less than the next noble gas.

Atomic and ionic radii: The halogens have the smallest atomic radii in their respective periods due to the maximum effective nuclear charge. The size of the anions is greater than the corresponding atoms. The atomic and ionic radii increase from top to bottom.

Ionisation Enthalpy: They have little tendency to lose electrons. Thus they have very high ionisation enthalpy. I.E. decreases from top to bottom.

Electronegativity: They have high electronegativities. It decreases from top to bottom. Fluorine is the most electronegative element in the periodic table. [F = 4.0]

Electron Gain Enthalpy: Halogens have maximum negative- electron gain enthalpy in their corresponding periods. It becomes less negative down the group. Electron gain enthalpy of F is less than that of Cl due to the compact size of F.

Physical Properties: F₂ and Cl₂ are gases. Br₂ is a liquid and I₂ is solid. They are all diatomic. M.pts and B.pts increase from top to bottom. F₂ is yellow, Cl₂ is greenish-yellow, Br₂ is red and I₂ violet in colour.

F₂ has a smaller enthalpy of dissociation as compared to Cl₂.

Chemical Properties:
Oxidation States and Trends in Chemical Reactivity: All the halogens exhibit a -1 oxidation state. However, chlorine, bromine and iodine exhibit +1, + 3, +5 and +7 oxidation states also as explained below:

Table: Atomic and Physical Properties of Halogens (Group 17)

^aRadioactive; ^bPauling scale; ^cFor the liquid at temperatures (K) given in the parentheses; ^dsolid; ^eThe half-cell reaction is X₂(g) + 2e^– → 2X^–(aq).

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 a -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. F2 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.
F₂ + 2X^– → 2F + X₂ (X = Cl, Br or I)
Cl^– + 2X^– → 2CT + X₂ (X = Br or I)
Br₂ + 2I^– → 2Br + I₂

The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials which are dependent on the parameters indicated below:

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 non-spontaneous. In fact, I^– can be oxidised by oxygen in an acidic medium; just the reverse of the reaction observed with fluorine.
2F₂ (g) + 2 H₂O (l) → 4H⁺ (aq) + 4 F (aq) + O₂ (g)
X₂ (g) + H₂O (l) → HX (aq) + HOX (aq) (where X = Cl or Br)
4I^– (aq) + 4H⁺ (aq) + O₂ (g) → 2I₂ (s) + 2 H₂O (l)

Anomalous behaviour of Fluorine: It is due to

1. Highest electronegativity and Ionisation enthalpy in Group 17.
2. Small size.
3. Non-availability of d-orbitals.
4. The low bond dissociation energy of the F-F bond.

1. Reactivity towards hydrogen: They all react with hydrogen to give hydrogen halides, but an affinity for hydrogen decreases from F to I.

The acidic strength is in order HF < HCl < HBr < HI and stability is under HF > HCl > HBr > HI.

2. Reactivity towards oxygen: Halogens form many oxides with oxygen but most of them are unstable. F forms OF₂ and O₂F₂, but only OF₂ is stable. Oxides of Cl, Br, I am powerful oxidizing agents.

3. Reactivity towards metals
Mg + Br₂ (l) → MgBr₂(s)
The ionic character of halids decreases in the order MF > MCI > MBr > MI.

4. Formation of interhalogen compounds: They form interhalogen compounds among themselves of the type AB, AB₃, AB₅, AB₇ where A = larger size halogen and B = smaller size halogen.

Chlorine, Cl₂:
Preparation:
1. By heating MnO₂ with cone. HCl

However a mixture of common salt and cone. H₂SO₄ is used in place of HCl.
2 NaCl + MnO₂ + 3 H₂SO₄ → 2 NaHSO₄ + MnSO₄ + 2H₂O + Cl₂

2. By the action of HC1 on KMn04
2 KMnO₄ + 16HCl → 2 KCl + 2 MnCl₂ + 5 Cl₂ + 8 H₂O

Manufacture of Cl₂:
1. Deacon’s process:

2. Cl₂ is obtained by the electrolysis of blue solution.

Properties:
1. Greenish-yellow gas with a pungent smell.

2. 2.5 times heavier than air.

3. Can be liquified easily into greenish-yellow liquid.

4. Soluble in water.

5. It reacts with a number of metals and non-metals and other compounds.
2Al + 3Cl₂ → 2Al Cl₃
2 Na + Cl₂ → 2NaCl
2 Fe +3 Cl₂ → 2FeCl₃
P₄ +6 Cl₂ → 4PCl
S₃ +4 Cl₂ → 4 S₂Cl₂
H₂ + Cl₂ → 2HCl
H₂S + Cl₂ → 2HCl + S
C₁₀H₁₆ + 8 Cl₂ → 16HCl + 10 C.
8NH₃ (excess) + Cl₂ → 6NH₄Cl + N₂
NH₂  + 3Cl₂ (excess) → NCl₃ + 3HCl

6. Chlorine dissolved in water on standing loses its yellow colour due to the formation of HCl and HOCl. Hypochlorous acid (HOCl) so formed, gives nascent oxygen which is responsible for oxidising and bleaching properties: Its oxidising properties are:

(a) 2 FeSO4 + H₂SO₄ + Cl₂ → Fe₂(SO₄)₃ + 2 HCl
Na₂SO₃ + Cl₂ + H₂O → 4 Na₂SO₄ + 2 HCl
SO₂ + 2H₂O + Cl₂ → 4 H₂SO₄ + 2HCl
I₂ + 6 H₂O + 5 Cl₂ → 2 HIO₃ + 10 HCl.

(b) It is a powerful bleaching agent due to oxidation. It is permanent.

Uses:

1. For bleaching wood pulp, cotton and textiles.
2. In the extraction of gold and platinum.
3. Manufacture of dyes, drugs, DDT, refrigerants etc.
4. In sterilising water.

Hydrogen Chloride, HCl:
Preparation-Lab.
Method:
2 NaCl + H₂SO₂  → Na₂SO₄ + 2 HCl cone.

Properties:

1. It is a colourless gas with a pungent smell.
2. Highly soluble in water.
3. Can be liquified easily to a colourless liquid and solidified to a white crystalline solid.
4. It ionises in an aqueous solution.
   HCl + H₂O (l) → 4 H₃O+ (aq) + Cl^– (aq); Ka = 107
   The high value of Ka shows it is a strong acid in water.
5. Reaction with NH₃
   NH₃ + HCl → NH₄Cl (White dense fumes)
6. 3 parts of cone. HCl and one part of cone. HNO₃ on mixing give aqua regia which is used to dissolve, Au, Pt.
   Au + 4H⁺ + NO₃ + 4 Cl^– → AUCl₄^– + NO + 2H₂O
   3Pt + 16H⁺ + 4NO₃^– + 18Cl^– → 3PtCl₆^2- + 4NO + 8H₂O .
7. Reaction with salts:
   Na₂CO₃ + 2HCl → 2 NaCl + H₂O + CO₂ ↑
   NaHCO₃ + HCl → NaCl + H₂O + CO₂ ↑
   Na₂SO₃ + 2HCl → 2 NaCl + H₂O + SO₂ ↑

Uses of HCl:

• Manufacture of Cl₂, NH₄Cl and glucose from com starch.
• Extracting glue from bones and purifying bone black.
• In medicine and as a lab. reagent

Table: Oxo-Acids of Halogens

Structures of Oxoacids:

Inter Halogen Compounds:

Some Properties of Interhalogen Compounds:

Uses:

• Interhafogen compounds are very useful fluorinating agents.
• They can be used as non-aqueous solvents.
• ClF₃ and BrF₃ are used for the production of UF₆ in the enrichment of ²³⁶U.
  U (s) + 3 ClF₃ (l) → UF₆(g) + 3 Cl F (g) argon (Ar)

Group 18 Elements: Group 18 Elements are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). Because of ns2nph structure [1s² for He], They are chemically unreactive. They are called noble gases.

Occurrence: Except radon (Rn) which is radioactive, all noble gases occur in the atmosphere in traces. Noble gases constitute 1 % (by volume) of dry air. He is present in pitch blende, monazite, cleveite. Radon is obtained as a decay product of ²²⁶Ra.
²²⁶₈₈Ra → ²²²₈₆Ra + ⁴₂He

Electronic Configuration: All noble gases have ns²np⁶ electronic configuration except He which is 1s².

Ionisation Enthalpy: They have very high I.E. which decreases down the group.

Atomic radii: They increase down the group with an increase in at. number.

Electron Gain Enthalpy: They have large positive values because of their stable electronic configuration. They have no tendency to accept electrons.

Physical Properties:

1. These gases are monoatomic.
2. They are colourless, odourless and tasteless.
3. They are sparingly soluble in water.
4. Due to weak dispersion forces, they have very low m.pts and b. pts.
5. He has the lowest B.pt (4.2 K) of any known substance.

Chemical Properties: In general, noble gases are least reactive. It is due to

1. Stable electronic configuration of ns2np6 (or 1s² for He).
2. They have high Ionisation enthalpy and more positive electron gain enthalpy,

Neil Bartlett prepared the first compound of Xe: red coloured Xe⁺ Pt F₆^– by mixing PtF₆^– and Xenon. The compounds of Kr are fewer. He, Ne and Ar are not found to form any true compounds. Only KrF₂ has been studied. Compound involving Rn viz. RnF₂ has been identified but not isolated.

Xenon-fluorine Compounds: Xenon forms three binary compounds with fluorine: XeF₂, XeF₄, XeF₆ by the direct combination of elements under appropriate experimental conditions.

Xenon-oxygen compounds: Xenon and oxygen form XeO₃. Partial hydrolysis of XeF₄ and XeF₆ gives oxyfluorides XeOF₄ and XeO₂F₂.

structures of. (a) XeF₂, (b) XeF₄, (c) XeF₆, (d) XeOF₄ and (e) XeO₃
XeO₃ is a colourless explosive solid and has a pyramidal molecular structure. XeOF₄ is a colourless volatile liquid and has a square pyramidal molecular structure.

Uses: He is used in filling balloons because He is a light and non-inflammable gas. It is used in gas-cooled nuclear reactors. Liquid He is used in low temp, physics. It is used in the driving apparatus.

Neon is used in discharge tubes and fluorescent bulbs for display purposes. Argon is mainly used to provide an inert atmosphere in high-temperature metallurgical processes.

There are no significant uses of Xenon and Krypton. They are used in light bulbs designed for special purposes.