p-Block Elements Study Notes

Introduction to p-Block Elements

• In Class XI, students have already learned about the p-block elements, which are situated in groups 13 to 18 of the periodic table.

• The general valence shell electronic configuration for these elements is given as $ns^2 np^{1-6}$, with the exception of Helium which has a configuration of $1s^2$.

• Properties of p-block elements are influenced significantly by:

  • Atomic sizes

  • Ionisation enthalpy

  • Electron gain enthalpy

  • Electronegativity

Characteristics of p-Block Elements

• The absence of d-orbitals in the second period and the presence of d or f orbitals in heavier elements (from the third period onwards) impact their properties.

• These elements comprise metals, metalloids, and non-metals, leading to diverse chemical behaviors.

Objectives

After studying this unit:

• Students will appreciate trends in groups 15, 16, 17, and 18.

• Learn preparation, properties, and uses of compounds like dinitrogen (N₂) and phosphorus (P) and their important compounds.

• Describe preparation, properties, and uses of dioxygen (O₂) and ozone (O₃), including the chemistry of simple oxides.

• Understand allotropy of sulphur and important compounds.

• Prepare, properties and uses of chlorine and hydrochloric acid (HCl).

• Know the chemistry of interhalogens and structures of oxoacids of halogens.

• Enumerate uses of noble gases.

• Appreciate the importance of these elements in daily life.

Group 15 Elements

7.1 Occurrence

• Group 15 consists of nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi).

• The shift from non-metallic to metallic character is noted as one moves down the group:

  • Nitrogen and phosphorus are non-metals.

  • Arsenic and antimony are metalloids.

  • Bismuth is a typical metal.

Nitrogen

• Comprises 78% of the atmosphere by volume.

• Occurs in earth's crust as sodium nitrate (NaNO₃) and potassium nitrate (KNO₃).

• A vital component of proteins in plants and animals.

Phosphorus

• Found in minerals of the apatite family: $Ca_9(PO_4)_6.CaX_2$ (X = F, Cl, OH).

• Essential for animal and plant tissues, present in living cells and bones.

7.1.2 Electronic Configuration

Atomic and Physical Properties (Table 7.1)

Notes:

• a EIII refers to bonds.

• - b E³⁻ refers to ionic state.

• - c E³⁺ refers to ionic state.

• - d Refers to white phosphorus.

• - e Refers to grey α-form at 38.6 atm.

• - f Sublimation temperature.

• - g Indicates at 63 K.

• - h Refers to grey α-form.

• - * For molecular N₂.

Trends in Atomic and Chemical Properties

• The valence shell electronic configuration is $ns² np³$; this configuration contributes to their stability.

• As one moves down the group:

  • Covalent and ionic radii increase.

  • A notable increase in covalent radius is seen from N to P, with minor increases from As to Bi.

  • Ionisation enthalpy decreases due to increased atomic size.

  • Electronegativity generally decreases down the group.

  • All elements are polyatomic; nitrogen is diatomic while others are solids.

  • The metallic character increases from N to Bi, with comparison to ionisation enthalpy and atomic size.

  • All elements (save nitrogen) exhibit allotropy.

Oxidation States and Reactivity

• Group 15 elements exhibit common oxidation states: -3, +3, +5.

• The lower oxidation state (-3) is less stable down the group, especially with Bismuth.

• The +5 oxidation state stability decreases and conversely, +3 increases due to the inert pair effect.

• Nitrogen additionally forms oxidation states of +1, +2, +4.

• Reactive tendencies involving nitrogen and phosphorus result in disproportionation reactions.

  • Example of nitrogen:

  • $3HNO₂ <br>ightarrow HNO₃ + H₂O + 2NO$

  • Phosphorus disproportionates into +5 and -3 oxidation states.

7.1.5 Electronegativity

• Electronegativity values decrease slightly down the group; nitrogen being notably more electronegative than its heavier counterparts.

7.1.6 Physical Properties

• Physical properties, including boiling and melting points, exhibit general trends.

• For melting: increases up to arsenic, then decreases up to bismuth.

• For boiling: generally increases from top to bottom.

• Notably, nitrogen exists as a gas, while the others are solids under standard conditions.

Dinitrogen (N₂)

Preparation

• Commercially produced via:

  • Liquefaction and fractional distillation of air.

  • On a smaller scale from reactions such as $NH₄Cl(aq) + NaNO₂(aq) <br>ightarrow N_2(g) + 2H₂O(l) + NaCl(aq)$

Physical and Chemical Properties

• Dinitrogen is colorless, tasteless, and non-toxic with low solubility in water affects marine life.

• High bond enthalpy of the N≡N bond stabilizes N₂, leading to inertness at ambient temperature but increasing reactivity at elevated temperatures.

Examples of Reactions

• Nitrogen reacts with metals:

  • $6Li + N₂ <br>ightarrow 2Li₃N$

  • $3Mg + N₂ <br>ightarrow Mg₃N₂$

Ammonia (NH₃)

Preparation

• Primarily through the Haber Process:

  • $N₂(g) + 3H₂(g) ⇌ 2NH₃(g)$; ΔfH = -46.1 kJ mol⁻¹.

  • Industrial conditions: 200 atm, ~700 K, with catalysts.

Properties

• Ammonia is a colorless gas with a distinct odor and greater boiling point due to hydrogen bonding.

• It forms ammonium salts with acids and acts as a weak base in water.

Typical Reactions

• Ammonia reacts with Cu²⁺ to form complexes:

  • $Cu^{2+} (aq) + 4NH₃ (aq) <br>ightarrow [Cu(NH₃)_4]^{2+} (aq) (blue)$

Oxides of Nitrogen

• Nitrogen forms several oxides across different oxidation states, such as:

  • N₂O, NO, N₂O₃, NO₂, N₂O₄, N₂O₅.

• These oxides exhibit resonance structures and unique chemical properties.

Nitric Acid (HNO₃)

• Prepared by heating KNO₃ or NaNO₃ with concentrated H₂SO₄.

• The key method is the Ostwald process, oxidizing NH₃ using Pt/Rh catalyst:

  • $4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)$

• Nitric acid is a strong acid and powerful oxidizing agent in various reactions.

Group 16 Elements

• Contains O, S, Se, Te, and Po. Known as chalcogens, these elements are discussed in terms of their chemical, physical properties, and group trends.

• Unique properties relate to their layer of governance moderated via their electronic configuration of $ns^2 np^4$.

Oxidation States, Reactivity

• Predominantly exhibit oxidation states from -2 to +6 with oxygen primarily showing -2.

• Reactivity associated with hydrogen includes forming hydrides such as $H₂E (E = O, S, Se, Te).

• For example, the acidic character of hydrides varies with bond dissociation enthalpy down the group.

Sulfur Dioxide (SO₂) and Sulfuric Acid (H₂SO₄)

• Sulfur is produced directly from the combustion of sulfur in O₂, and undergoes various reactions leading to multiple oxoacid formations.

• The Contact Process is noted as a primary industrial method for producing sulfuric acid.

Group 17 Elements (Halogens)

• Elements include F, Cl, Br, I, At, characterized by their tendency to form salts and interactions.

• Each element typically shows oxidation states of -1 and positive values due to their ability to accept electrons.

  • Reactivity decreases down the group, and the elements exhibit unique chemical properties with strong oxidizing activity.

Preparation of Chlorine

• Commonly prepared through the reaction of HCl with KMnO₄, along with Deacon’s process and electrolytic methods.

Noble Gases (Group 18)

• Comprise a series of elements known for their very high ionization enthalpy and lack of reactivity.

• Well-known for their complete outer electron shell, leading them to react infrequently with other elements.

Summary

• The p-block consists of diverse elements with unique chemical and physical properties ranging from multiple oxidation states to significant reactivity patterns. Both practical applications and theoretical principles underpin the chemistry of these groups, forming essential knowledge for students.