2022 JC1 Periodic Table Lecture Notes

1. Introduction

  • Elements are arranged by increasing atomic number and electronic configuration of outer shell electrons, affecting chemical properties.

  • Periodic Law: Properties of elements recur periodically when arranged by atomic number.

  • Main types of substances:

    • Metals: Good conductors of electricity.

    • Non-metals (except graphite): Insulators.

    • Metalloids: Semi-conductors; conductivity increases with temperature.

  • Metallic Character: Tends to decrease across a Period and increase down a Group.

2. Periodicity of Physical Properties in the Third Period

2.1 Factors Affecting Radius and First Ionisation Energy

  • Nuclear Charge: Attractive force from nucleus, increases with protons.

  • Shielding Effect: Repulsion among electrons; valence electrons shielded by inner shell electrons.

  • Effective Nuclear Charge: Net attraction felt by valence electrons after accounting for shielding.

  • Strong effective nuclear charge leads to smaller atomic radii and higher ionisation energies.

2.2 Atomic Radius

  • Across Period: Radius decreases due to increased nuclear charge (+protons) with constant shielding (same inner electrons).

    • Example: Na (0.186 nm) to Cl (0.099 nm).

  • Down Group: Radius increases due to increased electron shells (more shielding).

    • Example: Be (0.112 nm) to Ra (0.220 nm).

2.3 Ionic Radius

  • Cations (Na+ to Si4+): Radius decreases across period as effective nuclear charge increases.

  • Anions (P3– to Cl–): Radius also decreases due to similar electron configuration but increases as charge increases due to electron-electron repulsion.

2.4 First Ionisation Energy (1st I.E.)

  • Energy needed to remove one mole of electrons from gaseous atoms.

  • Across Period: 1st I.E. increases; however, deviations: Al < Mg and S < P due to subshell energy levels and electron pairing repulsions.

2.5 Electronegativity

  • Ability to attract shared electrons increases across Period due to higher nuclear charge and constant shielding.

  • Example: Na (0.93) to Cl (3.16).

  • Down Group: Decreases due to increased atomic size and shielding.

2.6 Melting Point

  • Structure and bonding affect melting points.

  • Metals: High melting points due to strong metallic bonds.

  • Si: Very high melting point due to extensive covalent bonding.

  • P4, S8: Low melting points due to weak van der Waals forces.

  • Order of melting point: S8 > P4 > Cl2 > Ar.

2.7 Electrical Conductivity

  • Metals (Na, Mg, Al) high conductivity due to delocalised electrons; they act as mobile charge carriers.

  • Silicon: Moderate conductivity (semiconductor).

  • Non-metals are insulators due to lack of mobile charge carriers.

3. Periodicity of Chemical Properties in the Third Period

3.1 Variation in Oxidation Number of Oxides and Chlorides

  • Higher oxidation states as you move from Na to Cl.

    • Example Oxides: Na2O (+1), MgO (+2), Al2O3 (+3), SiO2 (+4), P4O10 (+5), SO3 (+6).

    • Example Chlorides: NaCl, MgCl2, AlCl3, SiCl4, PCl5.

3.2 Properties of Oxides

3.2.1 Structure and Bonding

  • Increase in covalent character across Period due to increasing electronegativity.

  • Ionic bonding in Na2O and MgO, partially covalent in Al2O3, covalent in SiO2, and acidic oxides in P4O10 and SO3.

3.2.2 Melting Point

  • High: Ionic oxides Na2O, MgO, Al2O3; breakdown strong ionic bonds during melting.

  • Lower: SiO2 due to covalent bonds.

  • Lowest: P4O10, SO3 due to weak intermolecular forces.

3.2.3 Reaction of Oxides with Water

  • Behavior depends on oxide type:

    • Basic (Na2O): Forms NaOH and alkaline solution.

    • Amphoteric (Al2O3): Insoluble in water.

    • Acidic (SiO2, SO3): Produce acidic solutions upon reacting with water.

3.3 Properties of Chlorides

3.3.1 Structure and Bonding

  • Change from ionic to covalent bonding across Period.

    • Example Chlorides: NaCl (ionic), MgCl2 (ionic + covalent), SiCl4, PCl5 (covalent).

3.3.2 Reaction of Chlorides with Water

  • Generally lead to neutral, slightly acidic, or strongly acidic solutions.

  • NaCl: Simple dissolution to neutral solution.

  • AlCl3: Hydrolysis leading to acidic solutions.

  • SiCl4, PCl5: React with water producing strong acids

4. Diagonal Relationships

  • Elements in the second and third periods show similar properties due to opposing trends in their atomic radii and electronegativity.

  • Examples include similarities between Li and Mg, Be and Al.

Summary of Periodic Trends for Third Period Elements

  • Atomic radius decreases across Period; ionic radius also decreases across Period; first ionisation energy and electronegativity increase across Period; melting point varies significantly due to bonding nature.