NCERT Revision Notes for The p-Block Elements Class 12 Chemistry

Answer

P – Block of the periodic table comprises six groups (group 13 to 18). The general electronic configuration of the elements present in this block is ns²np^1–6 .
All these groups have specific names mostly based on the first member of each group.

• Group 13 : Boron Family
• Group 14 : Carbon Family
• Group 15 : Nitrogen Family
• Group 16 : Oxygen Family
• Group 17 : Halogens
• Group 18 : Noble Gases

With the exception of the noble gases which are very little reactive chemically, the elements belonging to all these groups are reactive. This is the only block in the periodic table which includes all the three types of elements i.e., metals, non – metals and semi – metals. The p – block elements are also known as the Representative Elements.

Answer

In group 15 of the periodic table, the elements, nitrogen (₇N), phosphorus (₁₅P), arsenic (₃₃As), antimony (₅₁Sb) and bismuth (₈₃Bi) are present .

Atomic and Physical Properties

(a) Electronic configuration
The general outer electronic configuration of elements of group 15 is ns²np³. Hence these elements belong to p – block.

(b) Atomic and ionic radii: The atomic (covalent) and ionic radii (in a particular oxidation state) of the elements of nitrogen family (group 15) are smaller than the corresponding elements of carbon family (group 14). On moving down the group, the covalent and ionic radii (in a particular oxidation state) increase with increase in atomic number. There is a considerable increase in covalent radius from N to P. However, from As to Bi, only a small increase is observed. This is due to presence of completely filled d – and/or f – orbitals in heavier members.

(c) Ionisation enthalpy : As the atomic size increase down the group, the ionization enthalpy increases. The ionization enthalpy of nitrogen group element is more than the corresponding elements of oxygen group. This is because more stable half filled outermost p – subshell of nitrogen group elements. Ionisation enthalpy decreases from N to Bi down the group due to gradual increase in atomic size.

(d) Electronegativity : Group 15 elements are more electronegative than group 14 elements. It decreases on moving down the group from N to Bi due to gradual increases in atomic size. N is the 3^rd highest in periodic table.

(e) Physical state : Nitrogen is a diatomic gas while other members are solids.

(f) Metallic character : These is increase in metallic character down the group due to increase in atomic size and decrease in ionization enthalpy.

(g) Melting and boiling points : Melting point increases from N to As and then decreases to sb and Bi. Boiling point increase down the group.

(h) Allotropy : All elements except nitrogen show allotropy :

(i) Density : Density increases down the group.

(j) Catenation
They exhibit the property of catenation but due to weak M – M bond to less extent than 14 group elements.

Answer

(a) Oxidation state : The elements of this group can exhibit various oxidation states ranging between −3 to 5. Negative oxidation state will be exhibited when they combine with less electronegative element & positive oxidation state will be exhibited with more electronegative element. Positive oxidation state becomes more favourable as we move down the group due to increasing metallic character & electropositivity. Although due to inert pair effect the stability of +5 state will also decreases. The only stable compound of Bi (v) is BiF₅.

(b) Covalency : This is the number of covalent bonds formed by an atom. Maximum covalency of nitrogen is four because it does not have d– orbitals to expand its covalency. Whereas other elements can exhibit covalency of 5 and 6 as well by the use of their d– orbitals e.g. PCl₅, [PCl₆]⁻

(c) Reaction with metals : Elements of group 15 react with metals and form binary compounds in –3 oxidation states e.g. Ca3N₂
6Li + N₂ → 2Li₃N ; 6Mg + P₄ → 2Mg₃P₂

(d) Reaction with hydrogen : All elements of group 15 form gaseous hydrides of the type MH₃ . In all the hydrides the central atom is sp³ hybridized & their shape is pyramidal duo to presence of lone pair of electrons.

• The basic strength of the hydrides decreases as we move down the group. Thus, NH₃ is the strongest base.
  NH₃ > PH₃ > AsH₃ > SbH₃
• The thermal stability of the hydrides decreases as the atomic size increases, i.e., the M – H bond strength decreases which means reducing character increases.
• In the liquid state, the molecules of NH₃ are associated due to hydrogen bonding. The molecules of other hydrides are not associated.
• NH₃ is soluble in water whereas other hydrides are insoluble.
• All the hydrides, except NH₃, are strong reducing agents and react with metal ions (Ag⁺ , Cu²⁺ , etc.) to form phosphides, arsenides or antimonides.

Note :

(e) Halides : The elements of group 15 form two series of halides MX₃ and MX₅ .

• All the elements of the group form trihalides. The ionic character of trihalides increases as we move down the group. Except NCl₃ , all the trihalides are hydrolysed by water. This is due to the absence of d – orbitals in nitrogen.
• PF₃ is not hydrolysed because fluorine being more electronegative than oxygen forms more stable bonds with phosphorus than P – O bonds.
• N cannot form NX₅ because of non – availability of d – orbitals. Bi cannot form a BiX₅ because of reluctance of 6s electrons of Bi to participate in bond formation.
• The Hybridisation of M in MX₃ is sp³ and shape is pyramidal. In MX₅ M is sp³d hybridized and shape is trigonal bipyramid. The axial bonds in MX₅ are weaker and longer, So MX₅ are less stable and decompose on heating
  eg: 

(f) Oxides

• Nitrogen forms a number of oxides. The rest of the members (P, As, Sb and Bi) of the group form two types of oxides; E₂O₃ and E₂O₅.
• The reluctance of P, As, Sb and Bi to enter into pπ – pπ multiple bonding leads to cage structures of their oxides and they exist and dimers, E₄O₆ and E₅O₁₀ .
• The basic nature of the oxides increases with increase in atomic number of the element. Thus, the oxides of nitrogen (except N₂O and NO), P(III) and As (III) are acidic, Sb (III) oxide is amphoteric and Bi (III) oxide is basic.

Answer

The anomalous behaviour of nitrogen is due to its:
(a) Small size
(b) High E.N. and high I.E.
(c) Non availability of vacant d – orbital
(d) Tendency to form multiple bonds.

Nitrogen differs from other elements of its own group in following properties
(a) Nitrogen is a gas while other elements are solids
(b) Nitrogen is diatomic, while other elements are tetratomic [P₄, As₄, Sb₄]
(c) Nitrogen can form N₃⁻ion (due to small size and high E.N.)
(d) Nitrogen is chemically inert under ordinary conditions due to high dissociation energy of N = N bond.
(e)Nitrogen shows oxidation state from –3 to +5
(f) Hydride of nitrogen i.e., ammonia is stable and undergoes H – bonding.

Answer

Nitrogen constitutes about 78% by volume of the atmosphere. It occurs as chile saltpetre and Indian saltpetre. It is also an essential constituent of fertilizers, explosives and proteins.

Preparation
(a) Commercially by liquefaction and fractional distillation of air.
(b) In laboratory, by treating and aqueous solution of ammonium chloride with sodium nitrite.
NH₄Cl(aq) + NaNo₂(aq) → N₂(g) + 2H₂O(l) + NaCl(aq)
(c) By thermal decomposition of ammonium dichromate and of sodium or barium azide.

Physical properties
(a) It is a colourless, odourless, tasteless and non – toxic gas.
(b) it has two stable isotopes : ¹⁴N and ¹⁵N.
(c) It is inert at room temperature because of high bond enthalpy of N ≡ N bond but its reactivity increases rapidly with rise in temperature.

Uses
(a) Used in manufacture of NH₃ and other industrial chemicals containing nitrogen, e.g., calcium cyanamide.
(b) As inert diluent for reactive chemicals in iron and steel industry.
(c) As a refrigerant to preserve biological materials, food items and in cryosurgery.

Answer

Preparation
(a) Decay of nitrogenous organic matter, e.g., urea.

Iron oxide with small amount of K₂O and Al₂O₃ are used as catalysts, at a temperature of 700K and pressure of 200 atm.

Physical Properties
(a) Lighter than air
(b) Easily liquefied by cooling or compression
(c) Highly soluble in water. The solution is alkaline
(d) Forms H – bonding with water.
(e) Colourless gas with pungent odour

Chemical Properties
(a) It is basic and forms ammonium salts with acids.
NH₃ + HCl → NH₄Cl; 2NH₃ + H₂SO₄ → (NH₄)₂SO₄

(b) Reaction with metal salt

(c) Reaction with metals
Active metals liberate H₂ with NH₃

(d) Reaction with air
Ammonia burns in presence of catalyst to give NO

(e) Complex formation

(f) Oxidation (reducing property)

Uses :
(a) To produce various nitrogenous fertilisers.
(b) Manufacture of some inorganic nitrogen compounds like HNO₃ , (NH₄)CO₃, Na₂CO₃ etc.
(c) Liquid ammonia is used as a refrigerant.
(d) In making artificial silk and as a laboratory reagent.
Oxides or nitrogen: It forms a number of oxides in different oxidation states – N₂O, NO, N₂O₃, NO₂,N₂O₄, N₂O₅,

Answer

It is colourless neutral gas

Preparation :

Properties :
(a) Nitrous oxide is a colourless neutral gas with faint pleasant smell. It is heavier than air, fairly soluble in cold water but not in hot water. It is poisonous in nature; when inhaled in moderate quantities, it produces hysterical laughter, hence it is also known as laughing gas.

(b) Action of heat :

Supporter of Combustion : Although nitrous oxide itself is non – combustible, it supports combustion of glowing splinter, charcoal, burning phosphorus and magnesium ribbon.
Mg + N₂O → MgO + N₂

Uses :
(a) It is used as the propellant gas in ‘whipped’ cream bombs.
(b) Mixed with oxygen, it is used as an anaesthetic for small scale operations in dentistry and surgery.
Structure:

Note that N₂O is isoelectronic with CO₂.

Answer

Answer

O.S. (+4), brown gas, acidic in nature.
Preparation

Properties : Highly toxic, paramagnetic, reddish brown gas with choking odour, acidic
Uses :

• For manufacturing of HNO₃
• As a catalyst in lead chamber process for sulphuric acid

Structure :

Answer

Preparation
(a) 2HNO₃ + As₂O₃ + 2H₂O → NO + NO₂ + 2H₃AsO₄
(b) 2Cu + 6HNO₃ → 2Cu(NO₃)₂ + NO + NO₂ + 3H₂O

Properties :
The brown coloured mixture of NO and NO₂ on cooling condenses to a blue liquid which on freezing gives a blue solid of the formula N₂O₃.

It is an acidic oxide and hence dissolves in alkalies producing nitrites.
N₂O₃ + 2KOH → 2KNO₂ + H₂O
Thus dinitrogen trioxide is referred to as the anhydride of nitrous acid.
2HNO₂ → N₂O₃ + H₂O
Structure of N₂O₃ :

Answer

Answer

Answer

Preparation
(a) Lab method: by heating KNO₃ or NaNO₃ and conc. H₂SO₄.
NaNO₃ + H₂SO₄ → NaHSO₄ +HNO₃

(b) Commercially, by Ostwald’s process.

3NO₂(g) + H₂O(l) → 2HNO₃ (aq) + NO(g)

(c) From air (Birkeland – Eyde electric arc process).

Physical Properties
(a) It is colourless liquid but looks yellow due to dissolved NO₂.
(b) It freezes at 231.4 K and boils at 355.6 K.
(c) Fuming HNO₃ is pure HNO₃ with oxides of nitrogen dissolved in it.

Chemical Properties :

(b) It is a strong monobasic acid :

(c) It is a very strong oxidising agent, in concentrated as well as in dilute form :
2HNO₃(conc.) → H₂O + 2NO₂ + [O], 2HNO₃(dil.) → H₂O + 2NO + 3[O]
(d) Except gold and platinum, HNO₃ attacks all metals forming a variety of products.
(e) It oxidises many organic compounds and forms various products on reaction.
(f) Sugar on oxidation with HNO₃ gives oxalic acid.
(g) It is strong oxidising agent and attacks most metals except noble metals such as gold and platinum. However, these dissolve in aqua regia (a mixture of 3 parts of conc. HCl and 1 part of conc. HNO₃ ).
Uses
(a) Used in manufacture of ammonium nitrate for fertilisers and other nitrates for use in explosives and pyrotechniques.
(b) Preparation of nitroglycerin, trinitrotoluene and other organic nitro compounds.
(c) Pickling of stainless steel, etching of metals and as a oxidizer in rocket fuels.
Structure
Phosphorus
Phosphorus is obtained as phosphates in nature. For e.g. phosphorite (Ca₃(PO₄)₂), chlorapatite hydroxyapatite, Fluoraptite etc.

• Phosphorus occurs in bones, phosphate rocks. It is used in fertilizers.
• As, Sb, Bi, occur mostly as sulphides.

It is isolated by heating Ca₃(PO₄)₂ with coke and silica in an electric furnace.

Phosphorus exhibits allotropy. The various allotropes of phosphorus are white phosphorus, red phosphorus and black phosphorus. White phosphorus is the most common form whereas black phosphorus is the most stable form of phosphorus.

Answer

Answer

Answer

Phosphorus forms two types of halides, PX₃ and PX₅ (X = F, Cl, Br, l)

PX₃ (PCl₃)
Preparation
(a) P₄ + 6Cl₂ → 4PCl₃
(b) P₄ + 8SOCl₂ → 4PCl₂ + 4SO₂ + 2S₂Cl₂
Properties
(a) It is colourless oily liquid and undergo hydrolysis in presence of moisture.
PCl₃ + 3H₂O → H₃PO₃ + 3HCl
(b) It reacts with organic compounds containing –OH group
3CH₃COOH + PCl₃ → 3CH₃COCl + H₃PO₃
3C₂H₅OH + PCl₃ → 3C₂H₅Cl + H₃PO₃
Structure

It has a pyramidal shape in which P is sp³ –hybridized.

Answer

Answer

Answer

Answer

In group 16 of the periodic table, elements, oxygen (₈O), sulphur (₁₆S), selenium (₃₄Se), tellurium (₅₂Te) and polonium (₈₄Po) are present.

Atomic and Physical Properties

(a) Electronic configuration : The elements have the structure ns²np⁴ for their valence shells. The first element of the group 16 differs in its chemical behaviour from that of other members of the group due to its small size and high electronegativity.

(b) Metallic character : The metallic character increase with increase in atomic number. The first four elements are non – metallic in character. Non – metallic character is strongest in O and S, weaker in Se and Te while Po is metallic.

(c) Atomic and ionic radii : Atomic and ionic radii increases from top to bottom, due to increase in the number of shells.

(d) Ionisation enthalpy : Ionisation enthalpy decreases down the group, due to increase in size. Elements of group 16 have lower ionisation enthalpy values as compared to group 15 in the corresponding periods. This is due to the fact that group 15 elements have extra stable half – filled P – orbital electronic configurations.

(e) Electron gain enthalpy : Oxygen has less negative electron gain enthalpy than sulphur due to compact nature of oxygen atom. From sulphur onwards, electron gain enthalpy becomes less negative upto Po.
(f) Electronegativity : Next to F, O has highest electronegativity value among the elements. Within the group, electronegativity decreases with increase in atomic number.

(g) Catenation : The tendency for catenation decreases markedly as we go down group. Sulphur has a strong tendency for catenation. This is evident from the formation of polysulphided S_n²⁻ , polysulphonic acids, HO₃S–S_n–SO₃ Hand sulphanes, H–S_n–H .

(h) Density : Increases down the group regularly.

(i) Melting point and boiling point : Both show a regular increase down the group due to increase in molecular weight and Van der Wall’s forces of attraction .

(j) Atomicity : Oxygen is diatomic, sulphur and selenium octa atomic with puckered ring structure

(k) Allotropy : All the elements exhibit allotropy.

Chemical Properties

(a) Oxidation state : Oxygen shows oxidation state of −2 but in case of OF₂ and O₂F₂ oxidation state is +2 and +1 respectively. Sulphur, selenium, tellurium and polonium shows +2, +4 and +6 oxidation state. Stability of +6 oxidation state decreases down the group and +4 oxidation state increases down the group due to inert pair effect.

In OF₂ the oxidation state of oxygen is +2

(b) Reaction with hydrogen : All the elements of the group form volatile hydrides.

• The volatility increases markedly form water to hydrogen sulphide and then declines.
  This is evident in their boiling point. Increasing order of boiling points of hydrides is H₂S < H₂Se < H₂Te < H₂ Down the group boiling point increases because of increasing molecular weight increases the van der Waal’s interaction. H₂O has abnormally high b.p. due to hydrogen bonding.
• The thermal stability of the hydrides decreases in the order : H₂O > H₂S > H₂Se > H₂Te> H₂
• The strength of the hydrides as acids increases in the order: H₂O < H₂S < H₂Se < H₂
• The reducing power of the hydrides increases from H₂O to H₂

(c) Reaction with halogens : Group 16 elements form number of halides of type EX₄ and EX₆. Where E is an group 16 element and X is a halogen. Stability of halides decrease in the order F > Cl⁻ > Br⁻ > I⁻ .

• Hexahalides – These are formed by fluorine only (not by Cl, Br, I) where elements exhibit maximum valency of +6. SF₆, SeF₆, TeF₆ are colourless gases with sp³d² hybridisation and octahedral structure. These are covalent in nature. Due to bigger size of Cl, Br and I the coordination number of 6 is not achieved.
• Tetrahalides – With the exception of SBr₄, SI₄ and SeI₄ all tetrahalides are known.
  They have trigonal bipyramidal shape with sp³d hybridiation.
• Dihalides – The dihalides eg SCl₂, OF₂, TeBr₂ are sp³, hybridised and have distorted bond angles due to electron pair repulsions .
• Dimeric monohalides – The dimeric monohalides are given by sulphur and selenium eg. S₂F₂, S₂Cl₂, Se₂Cl₂, S_2­Br₂, Se₂Br₂ .

Answer

Preparation
(a) By heating oxygen containing salts, such asa chlorates, nitrates and permanganates.

(b) By thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some metals.
2Ag₂O(s) → 4Ag(s) + O₂ (g)
(c) By catalytic decomposition of H₂O₂.
2H₂O₂(aq) → 2H₂O(l)+ O₂(g)
(d) From water by electrolysis. Hydrogen is released at cathode and oxygen at anode.
(e) Industrial preparation : By liquefaction of air followed by its fractional distination. Carbon and water vapour are removed before liquefaction of the air.

Properties :
(a) Colourless and odourless gas soluble in water.
(b) It has three stable isotopes: ¹⁶O, ¹⁷O and ¹⁸O.
(c) It reacts with nearly all metals and non – metals except Au, Pt and some noble gases. Some reactions with metals, non – metals and other compounds are given as:
2Ca + O₂ → 2CaO; 2Mg+O₂ → 2MgO; 3Fe + 2O₂ → Fe₃O₄ : C + O₂ → CO₂; P₄ + 5O₂ → 2P₂O₅:

Answer

Binary compounds of oxygen with another element are called oxides. Oxides are of two types.
(a) Simple oxides – MgO, Al₂O₃.
(b) Mixed oxides – Fe₃O₄, Al₂O₃.

Simple oxides are of four types.
(i) Acidic (Non – metal oxides)
(ii) Basic (Metal oxides)
(iii) Amphoteric
(iv) Neutral

• Oxides which give an acid with water are called acidic oxides. General non – metal are acidic, e.g. SO₂, PO₅, SO₃, I₂O₅.
  Oxides of some metals in high oxidation state are also acidic, e.g., Mn₂O₇, Cr₂O₃, V₂O₅, etc.
• Oxides which give a base with water are basic oxides. Generally metal oxides are basic e.g. Na₂O, CaO, BaO, etc.
• Oxides which exhibit dual behaviour i.e., show properties of both acidic and basic oxides are called amphoteric oxides, e.g., ZnO, Al₂O₃ , SnO
• Oxides which are neither acidic nor basic are called neutral oxides, e.g., CO, NO and N₂

Uses :
(a) In respiration and combustion processes.
(b) In oxyacetylene welding and manufacture of steel.
(c) As oxygen cylinders in hospitals, high altitude flying and in mountaineering.

Answer

Ozone is an allotrope of ozone. It is formed from atmospheric oxygen in presence of sunlight.
Preparation:
When a slow dry stream of oxygen is passed through a silent electrical discharge in a special type of apparatus called ozoniser (Siemen’s and Brodie’s ozonisers) conversion of oxygen to ozone occurs. The product is known as ozonised oxygen :
3O₂ → 2O₃, ∆H°(298K) = +68kcal: 3O₂ →2O₃; ∆H° = +142kJ/mol

Properties
(a) It is a pale – blue gas with a characteristic smell. It changes to dark blue liquid and violet – black solid.
(b) It is harmless in small concentrations. In higher concentrations it can be explosive and can damage body tissues.
(c) It is heavier than air.
(d) O₃ is neutral to litmus.
(e) It is thermodynamically unstable, it is endothermic i.e. its standard heat of formation is positive.
(f) It acts as a powerful oxidising agent due to the ease with which it liberates atoms nascent oxygen (O₃ → O₂ + O).
Some of the examples of its oxidizing action are :

• Pbs + 4O₃ → PbSO₄ + 4O₂
• 2HCl + O₃ → Cl₂ + H₂O + O₂
• 2kl + H₂O + O₃ → 2KOH + I₂ + O₂

(g) Nitrogen oxides (emitted from jet aeroplanes) combine rapidly with ozone resulting in its depletion in upper atmosphere.
NO(g) + O₃(g) → NO₂(g) + O₂(g)

Note :
(a) This reaction forms the basis of quantitative analysis of O₃, O₃oxidize I in a solution buffered with borate buffer having pH 9.2. The I₂ formed is further reduced by a standard solution of Na₂S₂O₃. The amount of Na₂S₂O₃ gives the amount of I₂ produced in oxidation which is then related to amount of O₃ used in oxidizing I.

(b) Ozone layer in stratosphere protects earth from UV radiations but now it is under threat from nitrogenous oxides emitted from exhaust system of supersonic jets which are slowly depleting the concentration of ozone layer via reaction NO + O₃ → NO₂ + O₂

Structure

Uses :
(a) As germicide, disinfectant and for sterilising water.
(b) Bleaching oils, ivory, flour, starch, etc.
(c) Acts as an oxidising agent in the manufacture of KMnO₄.

Answer

Important allotropes of sulphur are yellow rhombic (α – sulphur) and monoclinic (β – sulphur).

Yellow or rhombic sulphur (α – Sulphur)

• It is stable form of sulphur at room temperature.
• Its m.p. is 385.8K
• Its specific gravity is 2.06
• It is insoluble in water, somewhat soluble in benzene, alcohol, ether & readily soluble in CS₂ .

Monoclinic sulphur (β – Sulphur)

• It is prepared by melting rhombic sulphur in a dish and cooling it.
• Its m.p. is 393 K
• Its specific gravity is 1.98
• It is soluble in CS₂
• It is stable only above 369K
• Both α and β sulphur have S₈ α – sulphur is stable below 369K and β – sulphur is stable above 369K and at 369K both the forms are stable.
• Other modifications of sulphur has 6 – 20 sulphur atoms per ring. In cyclo – S₆ the ring adopts chair form.
• In vapour state at higher termperatures sulphur exists as S₂ molecule which has two unpaired electrons in the antibonding π orbitals like O₂ and hence, exhibits paramagnetism.

Answer

Preparation
(a) S(s) + O₂(g) → SO₂(g)
(b) In laboratory, by treating a sulphite with dil. H₂SO₄ .
SO₃²⁻(aq) + 2H⁺ (aq) → H₂O (I) + SO₂ (g)
(c) Industrially, as a by product of roasting of sulphide ores.
4FeS₂(s) + 11O₂(g) → 2Fe₂O₃(s) + 8SO₂(g)

Properties
(a) Colourless gas with pungent smell and highly soluble in water.
(b) SO₂ with water forms sulphurous acid. Thus it is acidic in nature

Sodium sulphite reacts with more SO₂ to form sodium hydrogen sulphite.
Na₂SO₃ + H₂O + SO₂ → 2NaHSO₃
(d) Reaction with chlorine in presence of charcoal as catalyst:
SO₂(g) + Cl₂(g) → SO₂Cl₂ (l)
(e) Sulphur dioxide reacts with oxygen in the presence of vanadium (V) oxide.

(f) In moist condition, acts as reducing agent.
For example :
2Fe³⁺ + 2H₂O + SO₂ → 2Fe²⁺ + SO₄²⁻ + 4H⁺
It is an oxidising agent also,
Example of its oxidizing action are:
2H₂S + SO₂ → 2H₂O + 3S
5SO₂ + 2MnO₄⁻ + 2H₂O → 5SO₄²⁻ + 4H⁺ + 2Mn²⁺

Structure
It is angular and is a resonance hybrid of two canonical forms.

Uses:
(a) In refining petroleum and sugar.
(b) In bleaching wool and silk.
(c) As an anti – chlor, disinfectant and preservatives.
(d) In manufacture of NaHSO₃, calcium hydrogen sulphite.
(e) As a solvent.

Answer

Answer

Preparation : Sulphuric acid is manufactured by contact process :
(a) Sulphur or sulphide ores burns in O₂ to give SO₂
S + O₂ → SO₂
(b) Catalytic oxidation of SO₂ with O₂ to give SO₃ in the presence of V₂O₅.

(C) Then SO₃ made to react with sulphuric acid of suitable normality to obtain a thick oily liquid called oleum.
SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l)
Then oleum is diluted to obtain sulphuric acid of desired concentration.
H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(l)

Physical characteristic
(a) H₂SO₄ is colourless, dense, oily liquid.
(b) Freezing point and boiling point are 10.5°C and 340°C respectively.
(c) Density iss 1.84 gm cm⁻³.
(d) Forms hydrate with the evolution of heat :
H₂SO₄ . H₂O [ monohydrate], H₂SO₄ . 2H₂O [dehydrate],
H₂SO₄ .3H₂O [trihydrate]
(e) It has strong affinity for water and acts as oxidizing agent.
(f) High b.p. and viscosity of H₂SO₄ is due to H–bondig .

Chemical characteristics

(a) Acid character (dibasic acid)
H₂SO₄ → H⁺ + HSO₄⁻, HSO₄⁻ → H⁺ + SO₄²⁻

(b) Forms two type of salts (normal and acidic):
NaOH + H₂SO₄ →NaHSO₄ + H₂O, NaHSO₄ + NaOH → Na₂SO₄ + H₂O

(c) Dehydrating agent [due to high affinity for water]

(d) Acts as an oxidizing agent
P₄ + 10H₂SO₄ → P₄O₁₀ + 10SO₂ + 10H₂O, C + 2H₂SO₄ →CO₂↑ + 2SO₂↑ + 2H₂O,
2HBr + H₂SO₄ → 2H₂O + SO₂↑ + Br₂, H₂S + H₂SO₄ → 2H₂O + SO₂↑ + S

(e) Displaces more volatile acids :
NaCl + H₂SO₄ → NaHSO₄ + HCl, NaNO₃ + H₂SO₄ → NaHSO₄ + 2HNO₃
FeS + H₂SO₄ → FeSO₄ + H₂S, Ca₃(PO₄)₂ + 3H₂SO₄ → 3CaSO₄ + 2H₃PO₄

(f) Reaction with metals : Zn, Mg, Fe gives hydrogen
Zn + H₂SO₄ (dil) → ZnSO₄ + H₂
Cu gives SO₂.
Cu + 2H₂SO₄(conc.) → CuSO₄ + 2H₂O + SO₂
Note : Metals like Cu, Pb, Hg, Bi and noble metals are not attacked by dil. H₂SO₄.

(g) Fromation of insoluble sulphates :
BaCl₂ + H₂SO₄ → BaSO₄↓ + 2HCl, Pb(NO₃)₂ + H₂SO₄ → PbSO₄ ↓ + 2HNO₃

(h) Reaction with PCl₅ and KClO₃ :
PCl₅ + H₂SO₄ → ClSO₂OH + POCl₃ + HCl, 3KClO₃ + 3H₂SO₄ → 3KHSO₄ + HClO₄ + 2ClO₂ + H₂O

Uses :
(a) Manufacture of fertilisers.
(b) Used in petroleum refining.
(c) Manufacture of pigments, paints and dyestuff intermediates.
(d) Used in detergent industry.
(e) Metallurgical aapplication.
(f) Used in storage batteries.
(g) Manufacture of nitrocellulose products.
(h) Laboratory reagent.

Answer

Answer

(a) Electronic configuration : Electronic configuration is ns² np⁵ for valence shells.

(b) Atomic and ionic radii : Halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge. Down the group atomic and ionic radii increases from F to I due to increase in the number of quantum shells.

(c) Ionisation enthalpy : The first ionisation energies are relatively high out decrease down the group. Iodine can lose an electron and from I⁺ ion.

(d) Electron gain enthalpy : The electron affinities are high, increases from F to Cl and then decreases down the group. The electron affinity of F is less than that of Cl due to its small size with high electron density. Electron affinity varies as Cl > F > Br > I

(e) Electronegativity : Halogens have high electronegativities F is most electronegative element known. Electronegativity decreases on moving down the group.

(f) Physical state : Fluorine and chlorine are gases, bromine is a liquid while iodine is a solid.

(g)Melting and soiling points: Their melting and boiling points steadily increase with atomic number.

(h) Bond energy : Actually, the bond energy should have decreased from F – F to I – I but in practice it starts decreasing from Cl–Cl to I–I.

The lower value of bond dissociation energy of F – F bond Cl – Cl bond is due to larger inter – electron (electron – electron) repulsion between the non – bonding electrons in the 2p – orbitals of fluorine atom than those in the 3p – orbitals of chlorine atoms.
Moreover, the X – X bond in Cl₂, Br₂ and I₂ is much stronger due to hybridisation of p – and d – orbitals.

(i) Non – metallic character : There is a gradual decline in the non – metallic character as we move down the group with decrease in electronegativity hence iodine shows some metallic character, viz. presence of metallic lustre and formation of positive ion, I⁺ .

(j) Density: It increases down the group in a regular fashion and follows the order F > Cl > Br > I.

Answer

(a) Oxidation state : All halogens exhibit – l oxidation state : Other halogens except fluorine also exhibit +1, +3, + 5 and +7 oxidation stae.

(b) Oxidising nature : Fluorine does not has d – orbital in its valence shell and hence cannot expand its octet. Therefore shows only –1 oxidation state. Halogens are good oxidising agents. The tendency of oxidising power decreases down the group . F₂ is the strongest oxidizing agent and oxidizes other halide ions. A halogen oxidizes halide ions of higher atomic number.

(c) Reaction with metals and non – metals : X₂ reacts with metals to give halides. The ionic character of the halides decreases in the order MF > MCl > MBr > MI
Halides of metals having low IE are more ionic in comparison to halides of metals having high IE. For metals exhibiting more than one oxidation states, the halide with higher oxidation state of metal is more covalent in comparison to halide having metal with lower oxidation state. X₂ react with non – metals like S, P, As etc.

(d) Reaction with hydrogen : All halogens react with hydrogen to form volatile H – X, which dissolve in water to form hydrohalic acids. Acidic strength of hydrogen halides varies as :
HF < HCl < HBr < HI
The stability of these halides decreases down the group due to decrease in bond dissociation enthalpy in the order.
H–F > H–Cl > H–Br > H–I
Also volatility decreases from HCl to HI.
The order of b.p is HCl < Hbr < HI < HF

(e) Reaction with oxygen : Chlorine forms oxides Cl₂O, ClO₂ , Cl₂O₆ & Cl₂O₇.. These oxides are highly reactive and strong oxidising agents and tend to explode. ClO₂ is used as bleaching agent for paper and pulp industries. It is also used in textiles and in water treatment.
Among the halogen oxides, bromine oxides Br₂O₂, BrO₃ are least stable due to middle row anomaly. These are stable only at low temperature. These are powerful oxidants. Iodine forms I₂O₄, I₂O₅ and I₂O₇. These are insoluble solids. They decomposes in heating I₂O₅ is a good reducing agent. I₂O₅ is used in the estimation of CO.

(f) Reaction with metal : Halogens forms metal halides with metals. Ionic character of M – X bond decreases in the order M – F > M – Cl > M – Br > M – I
Halides in higher oxidation state are more covalent whereas halides in lower oxidation state are ionic

(g) Reaction with halogen : Halogens combine amongst themselves to form a number of compounds known as interhalogen of type XX’, XX₃’, XX₅’ and XX₇

(h) Reaction with alkalies : Cl₂, Br₂ and I₂ behave similarly when treated with alkali (they undergo disproportionation reaction) 

• Cold and dilute alkali :
  X₂ + 2NaOH → NaX + NaOX + H₂O
• Hot and concentrated alkali :
  3X₂ + 6NaOH → 5NaX + NaXO₃ + 3H₂O
  F₂ behaves differently with alkalies :
  F₂ + 2NaOH(dil) → 2NaF + OF₂ + H₂O, 2F₂ + 4NaOH(conc) → 4NaF + O₂ + 2H₂O

Answer

Answer

Preparation :
(a) By heating MnO₂ with conc. HCl.
MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O

(b) By action of HCl or KMnO₄.
2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 8H₂O + 5Cl₂

Commercial manufacture
(a) In the Deacon process HCl is oxidised by air in presence of CuCl₂ as catalyst at 400°C

(b) In the electrolytic process, it is obtained at anode in the electrolysis of concentrated solution of brine (NaCl).

2Na⁺ + 2e⁻ → 2Na + 2H₂O → 2NaOH + H₂ (at cathode); 2Cl⁻ −2e⁻ → 2Cl → Cl₂ (at anode)

Properties
(a) Greenish – yellow gas with pungent and suffocating odour.
(b) Cl₂ is 2 – 5 times heavier than air.
(c) Soluble in water its aqueous solution is known as chlorine water which on careful cooling gives chlorine hydrate Cl₂.8H₂O.
(d) Reacts with metals and non – metals to form chlorides.
For example : 2Fe + 3Cl₂ → 2FeCl₃ ; P₄ + 6Cl₂ → 4PCl₃
(e) Chlorine has great affinity for hydrogen and react with hydrogen containing compounds to form HCl.
For e.g: H₂S + Cl₂ → 2HCl + S; H₂ + Cl₂ → 2HCl
(f) Reaction with NH₃ :

(g) Reaction with NaOH

(h) Acts as oxidising and bleaching agent.
Oxidising action of Cl₂
2FeSO₄ + H₂SO₄ + Cl₂ → Fe₂ (SO₄)₃ + 2HCl
Bleaching action :
Bleaching action is due to oxidation
Cl₂ + H₂O → 2HCl + O, Coloured substance + O → Colourless substance

Uses :
(a) For bleaching pulp, cotton and textiles.
(b) Extraction of Au and Pt.
(c) Manufacture of dyes, drugs and organic compounds like CCl₄, CHCl₃, DDT, refrigerants, etc.
(d) Sterilising drinking water.
(e) Preparation of poisonous gases such as phosgene (COCl₂), tear gas (CCl₃NO₂), mustard gas (ClCH₂CH₂SCH₂CH₂Cl).