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CBSE Class 12th Chemistry Notes: The p-Block Elements (Part - I)

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The p-Block Elements, is an important chapter of CBSE Class 12th Chemistry. This is chapter number 7th of NCERT Class 12th Chemistry textbook. Questions based on this chapter are frequently asked in CBSE Class 12th board examinations. Here you will get important notes on this chapter.

The main topics covered in these quick notes are:

•    Definition of p-block elements

•    Nitrogen family

•    General properties of nitrogen family

•    Anomalous properties of nitrogen

•    Reactivity of group 15 elements towards:

     o    Hydrogen

     o    Oxygen

     o    Halogens

     o    Metals

•    Preparation, properties and uses of dinitrogen (N2)

•    Preparation, properties and uses of ammonia (NH3)

•    Oxides of nitrogen

•    Preparation, properties and uses of nitric acid (HNO3)

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•    Brown ring test

•    Allotropes of phosphorus (P)

•    Uses of phosphorus

•    Preparation and properties phosphine (PH3)

•    Phosphorus halides

The notes of the chapter as follow:

p-block elements

The group number 13 to 18, in which the last electrons or the valence electrons enter in the p-orbital are called the p-block elements. The general electronic configuration of p-block elements is ns2 np1 ‒ 6

Nitrogen family

The elements of group 15: nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi) having general electronic configuration

ns2 np3, are known as the nitrogen family. The s-orbital in these elements is completely filled and p-orbitals are half-filled, making their electronic configuration extra stable.

General properties of nitrogen family:

Atomic and ionic radii: Covalent and ionic radii increase down the group. There is appreciable increase in covalent radii from N to P. However, from As to Bi only a small increase in covalent radius is observed.due to presence of completely filled d or f-orbitals in heavy elements.

Ionisation enthalpy: Ionisation enthalpy goes on decreasing down the group due to the increase in atomic size. Due to the stable electronic configuration with half filled p-orbital, group 15 elements have higher ionisation energy than group 16 elements. Also due to the smaller size of the elements, the group 15 elements have higher ionisation energy than group 14 elements.

Oxidation states: The common oxidation states are +3, +5, –3. The tendency to show –3 oxidation state decreases down the group due to increased size and hence decreased electronegativity. The stability of +5 oxidation state decreases whereas stability of +3 oxidation state increases due to inert pair effect. Nitrogen and phosphorus with oxidation states from +1 to +4 undergo oxidation as well as reduction in acidic medium. This process is called disproportionation.

For example: 3HNO2 → HNO3 + H2O + 2NO

Anomalous properties of nitrogen

Nitrogen has unique ability to form multiple bonds with itself and with other elements having small size and high electronegativity (e.g., C, O). Thus nitrogen exists as diatomc molecule, N2 with a triple bond between the two atoms. However, P cannot from  bond, therefore P exists as  P4 ‒ P4 is highly strained and is chemically reactive while N2 is chemically inert.

The behaviour of nitrogen differs from rest of the elements due to the following reasons:

(i) It has a small size

(ii) It does not have d – orbitals

(iii) It has high electronegativity

(iv) It has high ionization enthalpy

All the elements of Group 15 form trihydrides, MH3 having sp3 type hybridization.

Stability: The stability of hydrides decreases down the group due to decrease in bond dissociation energy down the group.

NH3 > PH3 > AsH3 > SbH3 > BiH3

Boiling point: Boiling point of hydrides (except NH3) increases down the group as with increase in size due to the Van der Waals forces also increase down the group. Boiling point of NH3 is highest due to the presence of hydrogen bonding.

PH3 < AsH3 < NH3 < SbH3 < BiH3

Bond angle: Electronegativity of N is highest. Therefore, the lone pairs will be towards nitrogen and hence more repulsion between bond pairs. Therefore bond angle is the highest. After nitrogen, the electronegativity decreases down the group.

NH3 (107.8°) > PH3 (93.6°) > AsH3 (91.8°) ≈ SbH3 (91.3°) > BiH3 (90°)

Basicity: Basicity decreases down the group.

NH3 > PH3 > AsH3 > SbH3 > BiH3.

Reactivity towards oxygen: All the group 15 elements form two types of oxides: trioxides, M2O3 and pentaoxides, M2O5. The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state. Their acidic character decreases down the group.

Reactivity towards halogens: Group 15 elements react to form two series of halides: trihalides, MX3 and pentahalides, MX5. Trihalides are sp3 hybridised with pyramidal shape whereas pentahalides are sp3d hybridized with trigonal bipyramidal shape. Nitrogen does not form pentahalide due to non-availability of the d- orbitals in its valence shell. All the trihalides of these elements except those of nitrogen are stable. In case of nitrogen, only NF3 is known to be stable.

Reactivity towards metals: All the group 15 elements react with metals to form binary compounds in –3 oxidation state.

Dinitrogen, N2

Preparation: N2 is produced commercially by the liquifaction of air followed by fractional distillation.

Liquid N2 distills out first leaving behind liquid oxygen. In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.

NH4Cl (aq) + NaNO2 (aq) → N2 (g) + 2H2O (l) + NaCl (aq)

Small amounts of NO and HNO3 formed as impurities can be removed by passing the gas through aqueous H2SO4 containing K2Cr2O7.

(NH4)2Cr2O7 → N2 + 4 H2O + Cr2O3

Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide

Ba (N3)2 → Ba + 3N2

Properties: Dinitrogen is a colourless, odourless gas, tasteless and non-toxic gas. It is chemically inert at room temperature due to the presence of triple bond with high bond dissociation energy.

Uses: Some important uses of dinitrogen are as follows:

(i) Liquid N2 is used as a refrigerant

(ii) It is used to provide inert atmosphere in iron and steel industry

(iii) It is used in the manufacture of HNO3 and NH3.

Ammonia, NH3

Ammonia molecule is trigonal pyramidal with nitrogen atom at the apex. It has 3 bond pairs and 1 lone pair. N is sp3 hybridised.

Preparation: Ammonia (NH3) is manufactured on the commercial scale by Haber's process.

Pressure ‒ 200 × 105 Pa

Temperature ‒ 773 K

Catalyst  ‒ Iron oxide with small amounts of K2O and Al2O3

In laboratory, ammonia is prepared by reacting NH4Cl with NaOH.

NH4Cl + NaOH → NaCl + H2O + NH3


Properties: Due to the presence of the lone pair of electrons on the nitrogen atoms, NH3 is a Lewis base. It can form coordinate covalent bond with the transition metal ion and form complexes, e.g.

Uses: Some important uses of ammonia are:
(i) It is used as a refrigerant
(ii) It is used in the manufacture of nitric acid,  
(iii) It is used in the production of nitrogenous fertilisers.

Oxides of nitrogen

Nitrogen forms a total of five oxides from +1 oxidation state to +5 oxidation state. The five oxides of nitrogen are: N2O, NO, N2O3, NO2 or N2O4, N2O5.

The following table gives the brief information of various oxides of nitrogen:

Image Source: NCERT Books

Image Source: NCERT Books

Nitric acid, HNO3

Nitric acid is the most important oxoacid formed by nitrogen. It is a colourless liquid. In the gaseous state, HNO3 exists as a planar molecule with the structure as shown below:

Preparation: Nitric acid is manufactured by the catalytic oxidation of ammonia in Ostwald process.

Properties: Concentrated HNO3 is a strong oxidising agent and attacks most metals except noble metals. Cr, Fe and AI do not dissolve in conc. HNO3  due to the formation of a passive film of oxide on the surface.

The oxidising action of HNO3 is depends on its concentration and the nature of the reducing agent. The principal product of reduction of HNO3  is NO when it is dilute but NO2 when it is concentrated.

For example:

3Cu + 8HNO3 (dil) → 3Cu(NO3)2 + 2NO + 4H2O

Cu + 4HNO3 (conc) → Cu(NO3)2 + 2NO2 + 2H2O

Brown ring test for the detection of nitrates:

Uses: Some important uses of nitric acid are:

(i) HNO3 is used in the manufacture of fertilisers.

(ii) It is used in the formation of explosives, dynamites, TNT, etc.

(iii) It is also used in the etching of metals.


Phosphorus is an essential constituent of elements and plants. It has many allotropic forms, the important ones are:

(i) White phosphorus

(ii) Red phosphorus

(iii) Black phosphorus

Properties of white phosphorus:

Highly toxic and can ignite when exposed to air.

Waxy, insoluble in H2O

More reactive than the other solid phases due to the angular strain in the P4 molecule

Glows in dark

Discrete tetrahedral P4 molecules.

Properties of red phosphorus:

Prepared by heating white phosphorus at 573K in an inert atmosphere

Odourless, nonpoisonous and insoluble in H2O

Less reactive than white phosphorus

Polymeric structure consisting of chains of P4 units linked together

Higher meltng point and density

Properties of black phosphorus:

Prepared by heating white phosphorus at 473 K under high pressure.

Exists in two forms - α black P and β black P

Thermodynamically most stable, i.e., least reactive

Has an opaque monoclinic or rhombohedral crystals

Uses: Some important uses of phosphorus are:

(i) Used in the manufacture of fertilisers and food grade phosphates.

(ii) Elemental P is used the manufacture of organo-phosphorus compounds used as pesticides.

Phosphine, PH3

Phosphine is a highly poisonous, colourless gas and has a smell of rotten fish.

Preparation: Phosphine is prepared by the reaction of calcium phosphide with water or dilute HCl.

Ca3P2 + 6H2O → 2PH3 + 3Ca(OH)2

Ca3P2 + 6HCl → 2PH3 + 3CaCl2

Properties: It is insoluble in water and is a weaker base than ammonia. Like ammonia, it gives phosphonium compounds with acids.

For example: PH3 + HBr → PH4Br 

PH3 is non-inflammable when pure but becomes inflammable owing to the presence of P2H4 or P4 vapours.

In water, PH3 decomposes in the presence of light to give red phosphorus and H2.

Phosphorus halides

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

Phosphorus trichloride, PCl3

It is a colourless oily liquid.

It is obtained by passing dry chlorine over heated white phosphorus or by the action of thionyl chloride with white phosphorus.

Properties: It has a pyramidal shape, in which phosphorus is sp3 hybridised.

It gets hydrolysed in the presence of moisture.

PCl3 + 3H2O → H3PO3 + 3HCl

Phosphorus pentachloride, PCl


Phosphorus pentachloride is prepared by the reaction of white phosphorus with excess of dry chlorine or can be prepared by the action of SO2Cl2 on phosphorus.

P4 + 10 Cl2 → 4 PCl5

P4 + 10SO2Cl2 → 4 PCl5 + 10SO2


Oxides of nitrogen

Nitrogen forms a total of five oxides from +1 oxidation state to +5 oxidation state. The five oxides of nitrogen are: N2O, NO, N2O3, NO2 or N2O4, N2O5.

The following table gives the brief information of various oxides of nitrogen:

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