CBSE Class 12 Chemistry Chapter 7 Revision Notes

Chapter 7: The p-Block Elements Revision Notes

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Group 15 Elements

• Nitrogen (₇N), phosphorus (₁₅P), arsenic (₃₃As), antimony (₅₁Sb), and bismuth (₈₃Bi) are all found in group 15 of the Periodic Table.
• This group’s elements can have various oxidation states ranging from -3 to + 5.
• Nitrogen has a maximum covalency of four because it lacks d- orbitals to expand its covalency.
• The atomic (covalent) and ionic radii (in a given oxidation state) of nitrogen family (group 15) elements are smaller than those of carbon family elements (group 14).
• From N to P, the covalent radius increases dramatically. However, there is only a small increase from As to Bi.
• Nitrogen has a strong proclivity for forming pπ-pπ multiple bonds with itself, carbon, and oxygen.
• As we progress through the group, the likelihood of pπ-pπ multiple bonding decreases.
• Elements in group 15 are more electronegative than those in group 14. As you move down the group from N to Bi, the electronegativity decreases.
• Gaseous hydrides of the type MH₃ are formed by all elements in group 15.
• As we progress through the group, the hydrides’ basic strength decreases. As a result, NH₃ is the most powerful base. NH₃ > PH₃ > ASH₃ > SbH₃
• As the atomic size of the hydrides increases, their thermal stability decreases.
• Due to non-availability of d-orbitals, N cannot form NX₅. Bi cannot form a BiX₅ due to the unwillingness of its 6s electrons to participate in bond formation.
• Nitrogen breaks down into a variety of oxides. The remaining members of the group (P, As, Sb, and Bi) form two types of oxides: E₂0₃ and E₂0₅.

Group 16 Elements

• The elements oxygen (₈O), sulphur (₁₆S), selenium (₃₄Se), tellurium (₅₂Te), and polonium (₈₄Po) are found in group 16 of the Periodic Table.
• The valence shells of the elements have the electronic configuration ns2np4.
• Due to its small size and high electronegativity, the first element of group 16 has a different chemical behavior than the other members of the group.
• As the atomic number rises, the metallic character rises with it. The first four elements have a nonmetallic quality to them. Non-metallic characters are strongest in O and S, weakest in Se and Te, and metallic in Po.
• As the number of shells increases,** the atomic and ionic radii increase** from top to bottom.
• As the group grows larger, the ionization enthalpy decreases.
• Elements in group 16 have lower ionization enthalpy values than elements in group 15. This is due to the extra stable half-filled p-orbital electronic configurations of group 15 elements.
• Due to the compact nature of the oxygen atom, it has a lower negative electron gain enthalpy than sulfur.
• O has the highest electronegativity value among the elements, second only to F. With increasing atomic number, electronegativity decreases within the group.
• As we progress through the group, the likelihood of catenation decreases.
• All the group’s elements form volatile hydrides.
• From water to hydrogen sulfide, the volatility rises and then falls.
• Their boiling point reflects this. H₂S < H₂Se < H₂Te < H₂0 are the hydrides with the highest order of boiling points.
• Due to the increase in molecular weight, the group boiling point increases, which increases the van der Waal’s forces of interaction.
• Because of hydrogen bonding, H20 has an abnormally high boiling point.
• Hydride thermal stability decreases in the following order:H₂0>H₂S>H₂Se>H₂Te>H₂Po.
• The strength of the hydrides as acids increases in the following order: H₂0 < H₂S< H₂Se < H₂Te
• Binary halides are formed by all of the elements in group 16.
• S, Se, and Te combine to form a variety of oxo-acids. Sulphuric acid is the most important of S’s oxo-acids.
• Sulphuric acid (H₂S0₃) and thiosulfuric acid (H₂S₂0₃) are unstable and therefore cannot be isolated. They can only be found in aqueous solutions or as their salts.

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Group 17 Elements

• Fluorine (₉F), chlorine (₁₇Cl), bromine (₃₅Br), iodine (₅₃I), and astatine (₅₅At) are all found in group 17 of the Periodic Table.
• For valence shells, the electronic configuration is ns2 np5.
• Due to maximum effective nuclear change, halogens have the smallest atomic radii in their respective periods.
• The initial ionization energies are relatively high, but they gradually decrease as the group progresses.
• Iodine can lose an electron, resulting in the formation of the I⁺ ion.
• Electron affinity varies as Cl > F > Br >I.
• F is the most electronegative element that has ever been discovered.
• Electronegativity decreases as you progress through the group.
• Halogens are excellent oxidizers. As you move down the oxidizing power of the group decreases.
• The order of reactivity is F₂ > Cl₂ > Br₂ > I₂.
• Hypohalous acids are all weak acids that can only be found in solution. The acid strength decreases as you progress through the group.
• For a given halogen atom, HOCl > HOBr > HOI
• Acid strength increases as the number of O-atoms increases. HOCl < HClO₂ < HClO₃ < HClO₄

Group 18 Elements

• Helium (₂He), neon (₁₀Ne), argon (₁₈Ar), krypton (₃₆Kr), xenon (₅₄Xe), and radon (₈₆Rn) are all found in group 18 of the Periodic table.
• They are referred to as noble gasses as a group.
• The orbitals of their valence shell are completely filled.
• They are monoatomic and water soluble only in small amounts.
• The fluorides XeF₂, XeF₄, and XeF₆ are formed by Xe.
• XeO₃ has a square pyramidal shape, whereas XeOF₄ has a trigonal pyramidal shape.
• Because he is a non-flammable gas that is lighter than air, He is used to fill meteorological balloons.
• Ar is used to create an inert environment.

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