p-Block Elements Chemistry Notes
→ Those elements, in which the last electron enter in subshell, are known as p-block elements. These p-block elements belongs to six main groups from 13 to 18 of the extended or long form of the periodic table. These elements are located at the extreme right of the periodic table. The general electronic configuration of these elements is ns2 np1-6 but helium has 1s2.
→ In other words, we can say that. in these elements the s-subshell is completely filled while the electron enters in p-subshell of valence shell. So the valence shell configuration of these elements changes from ns2 np1 to ns2 np6. The properties of these elements are depend upon the electrons present in p.orbitals. The groups of p.block are regarded by some typical names.
They are as follows:
- Group 13 ( B to TI) – Boron family
- Group 14 (C to Pb) – Carbon family
- Group 15 (N to Bi) – Nitrogen family or Pnicogens
- Group 16 (0 to Po) – Oxygen family or chalcogens
- Group 17 (F to At) – Halogens
- Group 18 ( He to Rn) – Inert gases or Noble gases
→ Except intert gases (which are chemically very less reactive) all other elements of this block are very 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 or metalloids. s-block elements and pb1ock elements are also known as representative elements. We have already discussed the chemistry of the elements of group 13 and 14 in class XI. Therefore, this unit includes the chemistry of the elements of group 15-18.
Chapter in Brief and Glossary :
→ Phosphorus, arsenic and antimony all exist as discrete tetratomic tetrahedral molecules i.e., P4, As4 and Sb4 because these elements cannot from multipie bonds. In these molecules each at’om is bonded to three atoms by single covalent bond and the angle between them is 60° due to the presence of lone pair of electron.
→ Phosphours exist in various allotropic forms i,e. white, red and black phosphorus, the yellow or white allotropic form of phosphorus is very reactive and poisonous in nature. It readily catches lire when exposed in air. It cause phossy jaw disease White phosphorus is stored into water.
Some properties of hydrides of group – 15 are as follows
- NH3 > PH3 > AsH3 > SbH3 > BiH3 (Bask character)
- NH3 > PH3 > AsH3 > SbH3 > BiH3 (Bond Angle)
- NH3 > PH3 > AsH3 > SbH3 > BiH3 (Thermal Stability)
- NH3 < PH3 < ASH3 < SbH3 < BiH3 (Reducing Nature)
→ The important oxo acids of phosphorus are H3PO4 (orthophosphoric acid), H3 PO2 (hypophoaphorus acid). H3PO3(phosphorus acid). HPO3 (mcta phosphoric acid. H4P2O6 (hypophorphoric acid) and H4P2O7 (pyrophosphonc acid). The basicity of these oxo acids is decided by the presence of – OH group.
→ Phosphorus can form two types of chloride i.e., PCl3 and PCl5, PCl3 has a pyramidal shape and sp hybridisation while PCl5 has triangular bipyramidal shape and si’d hybridization. PCl5 has two types of P—Cl bonds Le., axial and equatorial. Axial bonds become longer than equatorial bonds to minimise the lone pair bond pair repulsion present in molecule. PCl5 in solid state. exists as ionic compound. [PCl5]+ [PCl6]–
→ Group — 16 has five elements i.e., O, S, Se, Te and Po. The general electronic configuration of these elements is na2 np4. They show maximum oxidation state + 6. They are collectively called as chalcogens (ore forming).
→ In laboratory, oxygen can be prepared by heating KClO3 in presence of MnO2. Oxygen is a gas while other elements are oiId st room tempreture. Oxygen molecule is diatomic. It has one and one it bond while sulphur exist as staggered 8.atom rings.
→ All the elements of group – 16 rrirm hydrides of the type H2M where M O, S, SO, To, PO. H2O is liquid due to the tendency of t’romation of H-bonding while other hydrides of group — l6are poisonous gases with pungent smell, Some important properties of these hydrides are as follows
- H2O > H2O> S > H2Se > HeTe (Thermal st.ability)
- H2O < H2S < H2Se < HeTe (Reducing character)
- H2O < H2S < H2Se > H2Te (Volatility)
- H2O < H2S < H2Se < H2Te (Acidic behaviour)
- H2S < H2Se < H2Te < H2O (Bond angle)
- H2O < H2S < H2Se < H2Te (Covalent character)
- H2S < H2Se < HeTe < H2O (m.p. and b.p.)
→ Sulphur forms various oxyacids like H2SO4 (Sulphuric avid). H2SO3 (Suiphurous acid), H2S2O8 (Peroxodisulphuric acid), H2S2O7 (Perusulphuric acid). Among these oxyacida H2SO4 is very important. It is a strong dibasic acid. It is known as ‘oil of vitriol’. It can be prepared by Contact process.
→ Halogens are very reactive elements. Their reactivity decreases down the group Halogen mok’culcs have low dissociation energies. Fluorine due to greater lone pair-lone pair repulsion have lower bond dissociation energy than chlorine.
F2 < Cl2 < Br2 < I2 (Bond dissociation energy)
→ All the halogen form hydrides with hydrogen. They form volatile covalent hydride of fornnila HX. HF is liquid due to formation of hydrogen bonding while other are gases. Their properties are as follows
- HF > HCl > HBr > HI (Thermal stability)
- HF < HCl < HBr < Hl (Acidic strength)
- HCl < HBr < HI < HF (Boiling Point)
→ Fluorine due to very high electronegativity and small size form only one oxuacid (HOF) while other halogens forms ozo acids of th type HXO, HXO2, HXO3 and HXO4.
→ In the oxo acids ot’ same halogen, the acidic nature and thermal stability increases while oxidising nature decreases with increase in oxidation number.
- HClO < HClO2 < HClO3 <HClO4 (Acidic behaviour)
- ClO– < ClO2– < ClO3– < ClO4– (Thermal stability)
- HClO < HClO2 < HClO3 < HClO4 (Oxidising nature)
For different halogens,
- HClO > HBrO > HIO (Acidic strength)
- HClO > HBrO > HIO (Thcrmal stability)