The p-Block Elements Class 11 Notes Chemistry

 Boron Family (Group 13 Elements )

• Members: B, Al, Ga, In & Tl
• Melting Point: Decreases from B to Ga and then increases up to Tl.
• Ionization Energies: 1^st <<< 2^nd < 3^rd
• Metallic Character:  Increases from B to Tl. B is non-metal

Boron
Preparation of Boron:

• From Boric Acid: B₂O₃(s) + 3Mg(s) →  2B(s) +3 MgO(s)
• From Boron Trichloride
  • (at 1270 k): 2BCl₃+ 3H₂ (g) →  2B(s) + 6HCl (g)
  • (at 900 ⁰C): 2BCl₃(g) + 3Zn (s) →  2B(s) + 3 ZnCl₂ (s)
• By electrolysis of fused mixture of boric anhydride (B₂O₃) and magnesium oxide (MgO) & Magnesium fluoride at 1100 ⁰C
  • 2 MgO- →  2Mg + O₂(g)
  • B₂O₃ + 3Mg →  2B + 3MgO
• By thermal decomposition of Boron hydrides & halides:
  B₂H₆ (g) + Δ →  2B(s) + 3H₂ (g)

Compounds of Boron:
Orthoboric acid (H₃BO₃)
Preparation of Orthoboric acid

• From borax : Na₂B₄O₇ + H2SO4 + 5H2O →  Na₂SO₄ + 4H₃BO₃
• From colemanite : Ca₂B₆O₁₁ + 2SO₂ + 11H₂O →  2Ca(HSO₃)₂ + 6H₃BO₃

Properties of Orthoboric acid

• Action of Heat:
  
• Weak monobasic acidic behavior:
  B(OH)3 ↔ H3BO3 ↔ H⁺ + H₂O +

Thus on titration with NaOH, it gives sodium metaborate salt

H₃BO₃ + NaOH ↔ NaBO₂  + 2H₂O 

• Reaction with Metaloxide:
  
• Reaction with Ammonium boro fluoride:
  

Borax (sodium tetraborate) Na₂B₄O₇. 10H₂O 

Preparation from Boric Acid
4H₃BO₃ + Na₂CO₃ --> Na₂B₄O₇ + 6H₂O + CO₂
Properties of Borax

• Basic Nature:-

Aqueous solution of borax  is alkaline in nature due to its hydrolysis
Na₂B₄O₇ + 3H₂O →  NaBO₂ + 3H₃BO₃
NaBO₂ + 2H₂O → NaOH + H₃BO₃

• Action of heat:
  

Diborabe( B₂H₆)
Preparation of Diborane:
Reduction of Boron Trifluoride:
BF₃ + 3LiAlH₄ →  2B₂H₆ + 3 LiAl F₄
From NaBH₄:
2NaBH₄ + H₂SO₄ →  B₂H₆ + 2H₂ + Na₂SO₄
2NaBH₄ + H₃PO₄ →  B₂H₆ + 2H₂ + NaH₂PO₄

Properties of Diborane:

• Reaction with water: B₂H₆ + H₂O -->2H3BO3 + 6H2
• Combustion: B₂H₆ +2O₂ --? B₂O₃ + 3H₂O  ΔH =  -2615 kJ/mol

Compounds of Aluminium:
Aluminium Oxide or Alumina (Al₂O₃)
2Al(OH)₃ +Heat → Al2O₃ + 2H₂O
2Al(SO₄)₃ +Heat → Al₂O₃ + 2SO₃
(NH₄)₂Al₂(SO₄)₃·24H₂O --> 2NH₃ +Al₂O₃ + 4SO₃ + 25 H₂O

Aluminum Chloride AlCl₃:
Structure of Aluminium Chloride:

Properties of Aluminium Chloride

• White, hygroscopic solid
• Sublimes at 183 ⁰C
• Forms addition compounds with NH₃, PH₃, COCl₂ etc.
• Hydrolysis: AlCl₃ + 3H₂O --> Al(OH)₃ + 3HCl + 3H₂O 
• Action of Heat: 2AlCl₃ .6H₂O --> 2Al(OH)₃ à Al₂O₃+ 6HCl + 3H₂O

Carbon Family (Group 14 Elements):

• Members: C, Si, Ge, Sn, & Pb
• Ionization Energies: Decreases from C to Sn and then increases up to Pb.
• Metallic Character:  C and Si are non metals, Ge is metalloid and Sn and Pb are metals
• Catenation:  C and Si show a tendency to combine with its own atoms to form long chain polymers

Compounds of Carbon:
Carbon Monoxide
Preparation  of Carbon Monoxide

• By heating carbon in limited supply of oxygen: C + 1/2O₂ --> CO.
• By heating oxides of heavy metals e.g. iron, zinc etc with carbon.
  • Fe₂O₃ + 3C →  2Fe + 3CO
  • ZnO + C →  Zn + CO
• By passing steam over hot coke: C + H₂O →  CO + H₂ (water gas)
• By passing air over hot coke: 2C + O₂ + 4N₂ →  2CO + 4N₂ (Producer gas)

 Properties of Carbon Monoxide:

• A powerful reducing agent : Fe₂O₃ + 3CO →  2Fe + 3CO₂
  CuO + CO →  Cu + CO₂
• Burns in air to give heat and carbon dioxide: CO + 1/2O₂ →  CO₂ + heat.

Tests For Carbon Monoxide:

• Burns with blue flame
• Turns the filter paper soaked in platinum or palladium chloride to pink or green.

Carbon di-oxide
Preparation of Carbon di-oxide

• By action of acids on carbonates: CaCO₃ + 2HCl →  CaCl₂ + H₂O + CO₂
• By combustion of carbon: C + O₂ →  CO₂

Properties of Carbon di-oxide

• It turns lime water milky Ca(OH)₂ + CO₂ →  CaCO₃ ¯ + H₂O,
• Milkiness disappears when CO₂ is passed in excess
  CaCO₃ + H₂O + CO₂ →  Ca(HCO₃)₂
• Solid carbon dioxide or dry ice is obtained by cooling CO₂ under pressure. It passes from the soild state straight to gaseous state without liquefying (hence dry ice).

Carbides:

• Salt like Carbides : These are the ionic salts containing either C₂^2- (acetylide ion) or C₄- (methanide ion)e.g. CaC₂, Al₄C₃, Be₂C.
• Covalent Carbides : These are the carbides of non-metals such as silicon and boron. In such carbides, the atoms of two elements are bonded to each other through covalent bonds. SiC also known as Carborundum.
• Interstitial Carbides : They are formed by transition elements and consist of metallic lattices with carbon atoms in the interstices. e.g. tungsten carbide WC, vanadium  carbide VC.

Compounds of Silicon:
Sodium Silicate (Na₂SiO₃):
Prepared by fusing soda ash with pure sand at high temperature:
Na₂CO₃+ SiO₃ →  Na₂SiO₃ +CO₂

Silicones:
Silicon polymers containing Si – O – Si linkages formed by the hydrolysis of alkyl or aryl substituted chlorosilanes and their subsequent polymerisation.

Silicates:
Salts of silicic acid, H₄SiO₄ comprised of SiO₄^4- units having tetrahedral structure formed as result of sp³ hybridization.

Nitrogen Family (Group 15 Elements)

• Members: N, P, As, Sb & Bi
• Atomic Radii: Increases down the group. Only a small increases from As to Bi.
• Oxidation state: +3, +4 & +5.Stability of +3 oxidation state increases down the group.
• Ionization energy: Increases from N to Bi.

Nitrogen
Preparation of Nitrogen:

• 3CuO + 2NH₃ + Heat --> N₂ + Cu + 3H₂O
• CaOCl₂ + 2NH₃ + Heat --> CaCl₂+ 3H₂O + N₂
• NH₄NO₂ +Heat --> 3H₂O + N₂ +Cr₂O₃

 Properties of Dinitrogen:

• Formation of Nitrides (with Li, Mg, Ca & Al):  Ca + N₂ +Heat →  Ca₃N₂
• Oxidation: N₂ + O₂ →  2NO
• Reaction with carbide (at 1273 K): CaC₂ + N₂ →  CaCN₂ + C

Oxides of Nitrogen

Oxy -Acids of Nitrogen :

Ammonia (NH₃):
Preparation of Ammonia:

• By heating an ammonium salt with a strong alkali ;NH₄Cl + NaOH --> NH_3­ + NaCl + H₂O
• By the hydrolysis of magnesium nitride: Mg₃N₂ + 6H₂O --> 3Mg(OH)₂ + 2NH₃.
• Haber's process :  N₂(g) + 3H₂(g) --> 2NH₃(g).

Properties of Ammonia:

• Basic nature : Its aq. solution is basic in nature and turns red litmus blue.
  NH₃ + H₂O  ↔  NH⁴⁺ + OH^-
• Reaction with halogens :
  • 8NH₃ + 3Cl₂ --> 6NH₄Cl + N₂
  • NH₃ + 3Cl₂ (in excess) → NCl₃ + 3HCl
  • 8NH₃ + 3Br₂ →  6NH₄Br + N₂
  •  NH₃ + 3Br₂ (in excess) →  NBr₃ + 3HBr
  • 2NH₃ + 3I₂ →  NH₃.NI₃ + 3HI
  • 8NH₃.NI₃ →  6NH₄I + 9I₂ + 6N₂
• Complex formation :
  • Ag⁺ + NH³ →  [Ag(NH₃)₂]⁺
  • Cu²⁺ + 4NH³ → [Cu(NH₃)₄]²⁺
  • Cd²⁺ + 4NH³ → [Cd(NH₃)₄]²⁺

Precipitation of heavy metal ions from the aq. solution of their salts :

•  FeCl₃ + 3NH₄OH →  Fe(OH)₃ + 3NH₄Cl
                                Brown ppt.
• AlCl₃ + 3NH₄OH →  Al(OH)₃ + 3NH₄Cl
                       White ppt.
• CrCl₃ + 3NH₄OH →  Cr(OH)₃ + 3NH₄Cl
                                Green ppt.

Phosphorus:
Allotropy of Phosphorus:
a) White phosphorus:

• Translucent white waxy solid
• Extremely reactive
• Poisonous and insoluble in water

b) Red Phosphorus:

• Formed by heating white phosphorus in absence of air.
• Does not burn spontaneously at room temperature.
  

c) Black Phosphorus:  Formed by further heating of red phosphorus.
Compounds of Phosphorus:

a) Phosphine, PH₃:
Preparation of Phosphine

• Ca₃P₂ + 6H₂O →  2 PH₃ + 3 Ca(OH)₂
• 4H₃PO₃ +Heat →  PH₃+ 3 H₃PO₄
• PH₄I +KOH →  PH₃+KI + H₂O
• P₄ + 3KOH + 3H₂O →  PH₃ +3KH₃PO₂

Properties of Phosphine:

• Formation of Phosphonic Iodide: PH₃ + HI  à PH₄I
• Combustion: PH₃ + 2O₂  à H₃PO₄

b) Phosphorous Halides:

Preparation:

• P₄+ 6Cl₂ →  4PCl₃
• P₄+ 10Cl₂ →  4PCl₅
• P₄+ 8SOCl₂ →  4PCl₃ + 4SO₂+ 2S₂Cl₂
• P₄+ 10SOCl₂ →  4PCl₅ + 10SO₂

Properties:

• PCl₃ + 3H₂O → H₃PO₃ + 3HCl
• PCl₅ + 4H₂O →  POCl₃ à H₃PO₄ +5HCl
• PCl₃ + 3CH₃COOH →  3 CH₃COCl +H₃PO₃
• PCl₅ + CH₃COOH →   CH₃COCl + POCl₃+ HCl
• 2Ag + PCl₅  →  2AgCl + PCl₃
• 2Sn + PCl₅ → SnCl₄ + 2PCl₃
• PCl₅ + Heat →  PCl₃ + Cl₂

 C) Oxides of Phosphorus:

d) Oxy – Acids of Phosphorus:

Oxygen Family (Group 16 Elements) :

Chemical Properties of Group 16:
Formation of volatile Hydrides:

Formation of Halides:

Formation of Oxide:
a) All elements (except Se) forms monoxide.
b) All elements form dioxide with formula MO₂, SO₂ is a gas, SeO₂ is volatile solid. While TeO₂ and PoO₂ are non – volatile crystalline solids.
c) Ozone: It is unstable and easily decomposes into oxygen. It acts as a strong oxidising agent due to the case with which it can liberate nascent oxygen.

Oxyacids:

Allotropes of Sulphur :

Rhombic sulphur:

• It has bright yellow colour.
• It is insoluble in water and carbon disulphide. Its density is 2.07 gm cm^-3 and exists as S8 molecules. The 8 sulphur atoms in S8­ molecule forms a puckered ring.

Monoclinic Sulphur :

• Stable only above 369 K. It is dull yellow coloured solid, also called b - sulphur. It is soluble in CS₂ but insoluble in H₂O.
• It slowly changes into rhombic sulphur. It also exist as S8 molecules which have puckered ring structure. It however, differs from the rhombic sulphur in the symmetry of the crystals

Plastic Sulphur: 

• It is obtained by pouring molten sulphur to cold water.
• It is amorphous form of sulphur.
• It is insoluble in water as well as CS₂.

Sulphuric Acid:

•  Due to strong affinity for water, H₂SO₄ acts as a powerful dehydrating agent. 
• Concentrated H₂SO₄ reacts with sugar, wood, paper etc to form black mass of carbon. This phenomenon is called charring.
• It is moderately strong oxidizing agent.
  
• Decomposes carbonates, bicarbonates, sulphides, sulphites, thiosulphates and nitrites at room temperatures.
  
• Salts like chlorides, fluorides, nitrates, acetates, oxalates are decomposed by hot conc. H₂SO₄ liberating their corresponding acids.
  

Halogen Family ( Group 17 Elements)
Inter halogen compounds:

Hydrogen Halides:
Properties of Hydrogen Halides:

• All the three acids are reducing agents  HCl is not attacked by H₂SO₄.
  • 2HBr + H2SO₄ →  2H₂O + SO₂ + Br₂ ­
  • 2HI + H₂SO₄  → 2H₂O + SO₂ + I₂
• All the three react with KMnO₄ and K₂Cr₂O₇
  • 2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 8H₂O + 5Cl₂
  •  K₂Cr₂O₇ + 14HBr → 2KBr + 2CrBr₃ + 7H₂O + 3Br₂­
• Other reactions are similar.
  • Dipole moment : HI < HBr < HCl < HF
  • Bond length: HF < HCl < HBr < HI
  • Bond strength: HI < HBr < HCl < HF
  • Thermal stability: HI < HBr < HCl < HF
  • Acid strength: HF < HCl < HBr < BI
  • Reducing power: HF < HCl < HBr < HI

Pseudohalide ions and pseudohalogens:
Ions which consist of two or more atoms of which at least one is nitrogen and have properties similar to those of halide ions are called pseudohalide ions. Some of these pseudohalide ions can be oxidised to form covalent dimers comparable to halogens (X­₂). Such covalent dimers of pseudohalide ions are called pseudohalogens.
The best known psuedohalide ion is CN^–

Pseudohalogen

• (CN)₂ cyanogen
• (SCM)₂ thiocyanogen

Some important stable compound of Xenon

• XeO₃            Pyramidal
• XeO₄            Tetrahedral
• XeOF₄          Square pyramidal
• XeO₂F₂           Distorted octahedral

First rare gas compound discovered was Xe+ (PtF₆]– by Bartlett.

Oxyacids of Chlorine

Acidic Character: Acidic character of the same halogen increases with the increase  in oxidation number of the halogen: HClO₄> HClO₃ > HClO₂ > HOCl

Preparation

HOCl :     

• Ca(OCl)₂ + 2HNO₃ →  Ca(NO₃)₂ + 2HOCl

HClO₂ :    

• BaO₂ + 2ClO₂ →  Ba(ClO₂)₂ (liquid) + O₂
• Ba(ClO₂)₂ + H₂SO₄(dil.) →  BaSO₄ ¯ + 2HClO₂

HClO₃ :    

• 6Ba(OH)₂ + 6Cl₂ →  5BaCl₂ + Ba(ClO₃)₂ + 6H₂O
• Ba(ClO₃)₂ + H₂SO₄(dil.) →  BaSO₄ ¯ + 2HClO₃

HClO₄ :    

• KClO₄ + H₂SO₄ →  KHSO₄ + HClO₄
• 3HClO₃ →  HClO₄ + 2ClO₂ + H₂O

The Noble Gases (Group 18 Elements):
The noble gases are inert in nature. They do not participate in the reactions easily because they have

• stable electronic configuration i.e. complete octet.
• high ionization energies.
• low electron affinity.

Compounds of Xenon

Uses of Nobles gas

The noble gases are used in following ways:

(A) Helium

• It is used to fill airships and observation balloons.
• In the oxygen mixture of deep sea divers.
• In treatment of asthma.
• Used in inflating aeroplane tyres.
• Used to provide inert atmosphere in melting and welding of easily oxidizable metals.

(B)  Neon

• It is used for filling discharge tubes, which have different characteristic colours and are used in advertising purposes.
• Also used in beacon lights for safety of air navigators as the light possesses fog and stram perpetrating power.

(C)  Argon
Along with nitrogen it is used in gas – filled electric lamps because argon is more inert than nitrogen.