P Block Elements Class 12 Notes

Key Definitions

Inert Pair Effect

The reluctance of the outermost s-electrons to participate in bond formation, leading to lower oxidation states becoming more stable for heavier elements in a group (e.g., +1 for Tl in Group 13, +2 for Pb in Group 14).

Allotropy

The property of an element to exist in two or more different forms in the same physical state, possessing different physical properties but similar chemical properties (e.g., carbon: diamond, graphite; phosphorus: white, red, black).

Catenation

The ability of an atom to form bonds with other atoms of the same element, forming long chains or rings. This property is most prominent in carbon, followed by silicon and phosphorus.

Electron Gain Enthalpy

The enthalpy change when an electron is added to a neutral gaseous atom to form an anion. For halogens, it is highly negative due to their strong tendency to gain an electron.

Interhalogen Compounds

Compounds formed between two different halogens (e.g., ClF, BrF3, IF5, IF7). They are generally more reactive than halogens, except F2.

Oxoacids

Acids containing oxygen, hydrogen, and at least one other element, with the hydrogen bonded to oxygen and the central atom bonded to at least one oxygen atom (e.g., HNO3, H2SO4, HClO4).

Group 15 Elements (Nitrogen Family)

Group 15 consists of Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), and Bismuth (Bi). Their electronic configuration is ns²np³.

General Trends:

• Atomic and Ionic Radii: Increase down the group due to the addition of new shells.
• Ionization Enthalpy: Decreases down the group due to increasing atomic size and shielding effect. Group 15 elements have higher ionization enthalpy than Group 14 elements due to the stable half-filled p-orbitals.
• Electronegativity: Decreases down the group. Nitrogen is highly electronegative.
• Metallic Character: Increases down the group. Nitrogen and Phosphorus are non-metals, Arsenic and Antimony are metalloids, and Bismuth is a metal.
• Oxidation States: Common oxidation states are –3, +3, and +5. The stability of the +5 oxidation state decreases down the group, while the stability of the +3 oxidation state increases due to the inert pair effect (especially for Bi).

Anomalous Behaviour of Nitrogen:
Nitrogen differs from other elements in its group due to its small size, high electronegativity, high ionization enthalpy, and absence of d-orbitals. It forms pπ-pπ multiple bonds (e.g., N≡N), while heavier elements do not. It does not show +5 oxidation state readily in simple compounds due to the absence of d-orbitals to expand its octet.

Important Compounds:

• Ammonia (NH₃): A colourless gas with a pungent smell, highly soluble in water. Manufactured by Haber's Process. It acts as a Lewis base due to the lone pair on nitrogen.
• Nitric Acid (HNO₃): A strong oxidizing agent. Manufactured by Ostwald's Process.
• Oxides of Nitrogen: N₂O (nitrous oxide), NO (nitric oxide), N₂O₃ (dinitrogen trioxide), NO₂ (nitrogen dioxide), N₂O₄ (dinitrogen tetroxide), N₂O₅ (dinitrogen pentoxide). They show varying oxidation states of nitrogen.
• Phosphorus Allotropes: White phosphorus (P₄ tetrahedal, highly reactive, ignites in air, soluble in CS₂), Red phosphorus (polymeric, less reactive, stable), Black phosphorus (most stable).
• Phosphine (PH₃): Highly poisonous gas. Prepared by the reaction of calcium phosphide with water or dilute HCl.

Haber's Process for Ammonia (NH₃) Synthesis

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Group 17 Elements (Halogens)

Group 17 consists of Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). Their electronic configuration is ns²np⁵. They are highly reactive non-metals.

General Trends:

• Atomic and Ionic Radii: Increase down the group.
• Ionization Enthalpy: Decreases down the group.
• Electronegativity: Decreases down the group. Fluorine is the most electronegative element.
• Electron Gain Enthalpy: Highly negative due to their strong tendency to gain an electron. Chlorine has the most negative electron gain enthalpy, followed by Fluorine.
• Bond Dissociation Enthalpy: Generally decreases down the group. However, F₂ has a lower bond dissociation enthalpy than Cl₂ due to repulsion between lone pairs on small F atoms.
• Reactivity: Decreases down the group. Fluorine is the most reactive halogen.

Important Reactions & Compounds:

• Oxidizing Power: Decreases down the group (F₂ > Cl₂ > Br₂ > I₂). A halogen higher in the group can oxidize halide ions of lower halogens (e.g., Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂).
• Interhalogen Compounds: Form compounds of type AX, AX₃, AX₅, AX₇ (where A is a larger halogen and X is a smaller halogen). Examples: ClF, BrF₃, IF₅, IF₇. They are more reactive than halogens (except F₂).

Worked Example: Oxidizing Power of Halogens

• Predict the reaction: A student adds an aqueous solution of Chlorine (Cl₂) to a solution containing Potassium Bromide (KBr). Solution: Chlorine is more electronegative and has a higher oxidizing power than Bromine. Therefore, Chlorine will oxidize Bromide ions (Br⁻) to Bromine (Br₂). Reaction: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
• Oxidation states of Phosphorus in Oxoacids Determine the oxidation state of Phosphorus in H₃PO₂ (Hypophosphorous acid). Solution: Let the oxidation state of P be x. For H₃PO₂: 3(+1) + x + 2(-2) = 0. So, 3 + x - 4 = 0, which gives x = +1.

Section 6

When studying P-block elements, pay close attention to anomalous behaviour of the first element in each group (N, O, F). These are frequent sources of 'reasoning type' questions. Understand the reasons (small size, high electronegativity, absence of d-orbitals, high bond dissociation enthalpy). Also, practice balancing redox reactions involving P-block compounds and drawing structures of oxoacids of phosphorus, sulphur, and halogens (e.g., H₂SO₄, H₃PO₄, HClO₄). Remember that the inert pair effect becomes significant for heavier elements, influencing their stable oxidation states.

Key Points to Remember

• Group 15 elements exhibit stable +3 and +5 oxidation states. The +3 state is more stable for heavier elements due to the inert pair effect.
• Nitrogen shows anomalous behaviour due to its small size, high electronegativity, and absence of d-orbitals, enabling pπ-pπ bonding.
• Phosphorus exists in various allotropic forms; white phosphorus is highly reactive and poisonous.
• Sulphur (Group 16) also shows allotropy (rhombic, monoclinic). H₂SO₄ is manufactured by the Contact Process.
• Fluorine (Group 17) is the most reactive halogen and the strongest oxidizing agent.
• Chlorine has the highest negative electron gain enthalpy among all elements.
• Noble gases (Group 18) are generally inert due to stable octet configuration, but Xe can form compounds with F and O (e.g., XeF₂, XeF₄, XeOF₄).
• Interhalogen compounds are generally more reactive than halogens (except F₂) because A-X bonds are weaker than X-X bonds.
• Oxoacids of halogens show increasing acidic strength from HClO to HClO₄ due to increasing stability of conjugate base and higher oxidation state of halogen.

Practice Questions with Solutions

• Q: Why is nitrogen less reactive at room temperature? A: Due to the high bond dissociation enthalpy of the triple bond (N≡N) in N₂ molecule.
• Q: Explain why bismuth is predominantly metallic in nature while nitrogen is non-metallic. A: Metallic character increases down Group 15 due to decreasing ionization enthalpy and increasing atomic size. Bismuth is a heavy element and exhibits metallic properties.
• Q: Arrange F₂, Cl₂, Br₂, I₂ in increasing order of bond dissociation enthalpy. A: I₂ < F₂ < Br₂ < Cl₂. (Note: F₂ is anomalously low due to inter-electronic repulsion in its small size).
• Q: Give the formula of an interhalogen compound where the central atom exhibits an oxidation state of +5. A: IF₅ (Iodine pentafluoride).

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