Chemistry: p-Block Elements

Revision Notes on p-Block Elements:



 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: 1st <<< 2nd < 3rd


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)3à 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 C22- (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 SiO44- 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 :



Oxy Acids

Name of oxy – acid


1.   H2N2O2

Hyponitrous acid


2.   H2 NO2

Hydronitrous acid


3.   HNO2

Nitrous acid


4.   HNO3

Nitric acid


5.   HNO4

Per nitric acid




Ammonia (NH3):

Preparation of Ammonia:

By heating an ammonium salt with a strong alkali ;NH4Cl + NaOH --> NH3­ + 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:



Oxo acid

Name 


      H3PO2

Hypophosphorus acid


      H3PO3

Phosphorus acid


      H4P2O6

Hypophosphoric acid


      H3PO4

Orthophosphoric acid


      H4P2O7

Pyrophosphoric acid


      HPO3

Metaphosphoric acid







Oxygen Family (Group 16 Elements) :



Sr. No.

Property

Oxygen

Sulfur

Selenium

Tellurium

Polonium


1.

Configuration

[He]2s22p4

[Ne]3s23p4

[Ar]4s24p4

[Kr]5s25p4

[Xe]6s26p4


2.

Common oxidation state

-2

-2, +4, +6

+4, +6

+4, +6





3.

Atomic radius (pm)

66

104

116

143

167


4.

First ionization energy (KJ/mol)

1314

1000

941

869

812


5.

Electronegativity

3.5

2.5

2.4

2.1

2.0




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:




Sulphur

Selenium

Tellurium


Sulphurous acid H2SO3.

Sulphuric acid H2SO4

Peroxomonosulphuric acid H2SO5(Caro’s acid)

Peroxodisulphuric acid

H2S2O8 (Marshell’s acid)

Thio sulphuric acid H2S2O3

Dithiconic acid H2S2O6

Pyrosulphuric acid H2S2O7

Selenious acid H2SeO3

Selnenic acid H2SeO4

Tellurous acid H2TeO3.

Telluric acid H2TeO4.




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-3and 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:




Type XX’1 (n = 1) (with linear shape)

Type XX’3 (n = 3) (with T-shape)

XX’5 (n = 5)
(with square pyramidal shape)

XX’7 (n = 7) with pentagonal bipyramidal shape)


CIF

ClF3

ClF5





BrF BrCl

BrF3

BrF5





ICl, IBr, IF

ICl3, IF3

IF5

IF7




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–



Pseudohalide ions     

Name


CN–

Cyanide ion   


OCN–

Cyanate ion  


SCN–

Thiocyante ion


SeCN–

Selenocyanate ion


NCN2– 

Cyanamide ion


N3– 

Azide ion


OMC– 

Fulminate ion




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




Formula

Name

Corresponding Salt


HOCl

        Hypochlorous acid

Hypochlorites


HClO2

        Chlorous acid

Chlorites


HClO3

        Chloric acid

Chlorates


HClO4

        Perchloric acid

Perchlorates




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





Molecule

Total electron pairs (BP + LP)

Hybridisation

Shape


XeF2

5

Sp3d

Linear


XeF4

6

Sp3d2

Square planar


XeF6

7

sp3d3

Distorted octahedral




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.