Periodic Trends: Atomic Radius, Ionisation Energy and Electropositivity

Periodic trends describe the systematic variations in elemental properties across periods and down groups of the periodic table. These patterns arise directly from changes in atomic number, effective nuclear charge, shielding effect, and principal energy levels. Understanding atomic radius, ionisation energy, and electropositivity explains chemical reactivity, bonding behaviour, and elemental classification.

Core Factors Influencing Periodic Trends

Four fundamental atomic parameters govern the magnitude and direction of periodic variations.

• Nuclear Charge (Z): The total positive charge in the nucleus determined by the number of protons.
• Effective Nuclear Charge (Z_eff): The actual positive charge experienced by valence electrons, calculated as Z_eff = Z – sigma, where sigma is the shielding or screening constant.
• Shielding (Screening) Effect: Inner shell electrons repel outer electrons, reducing the net electrostatic attraction from the nucleus. The shielding ability decreases in the order: s > p > d > f.
• Principal Quantum Number (n): Represents the main energy shell. As n increases, the average distance between the nucleus and the outermost electrons increases.

Atomic and Ionic Radii

Atomic radius represents the distance from the centre of the nucleus to the outermost electron shell.

Types of Atomic Radii

• Covalent Radius: Half the distance between nuclei of two identical non-metallic atoms joined by a single covalent bond (e.g., Cl₂).
• Van der Waals Radius: Half the internuclear distance between two non-bonded adjacent atoms belonging to neighbouring molecules in the solid state; it is always larger than the covalent radius.
• Metallic Radius: Half the distance between two adjacent metal ions in a metallic lattice.
• Ionic Radius: The effective radius of an ion in a crystal lattice. Cations are always smaller than their parent neutral atoms due to increased Z_eff. Anions are always larger than their parent neutral atoms due to increased electron-electron repulsion.

Periodic Variation of Radii

Isoelectronic Species

Isoelectronic ions contain the same total number of electrons but different nuclear charges (e.g., N^3-, O^2-, F^-, Na⁺, Mg²⁺, Al³⁺, all possessing 10 electrons). In an isoelectronic series, ionic radius decreases as nuclear charge increases: N^{3-} > O^{2-} > F^- > Na^+ > Mg^{2+} > Al^{3+}

Ionisation Energy (Ionisation Potential)

Ionisation energy (IE) is the minimum energy required to remove the most loosely bound valence electron from an isolated gaseous atom in its ground state: Xtext{(g)} + text{IE}_1 rightarrow X^+text{(g)} + e^-

Successive Ionisation Energies

Removing subsequent electrons requires progressively higher energy because each removal increases the effective nuclear charge on the remaining electrons: text{IE}_1 < text{IE}_2 < text{IE}_3 A large jump between successive ionisation energies indicates the removal of an electron from a stable, fully filled noble gas core (e.g., the jump from IE₁ to IE₂ in Sodium).

Periodic Variation and Major Anomalies

• General Trend: Ionisation energy increases across a period from left to right and decreases down a group from top to bottom.
• Beryllium vs. Boron: Beryllium (1s^{2 2s^2}) has a higher IE₁ (899 kJ/mol) than Boron (1s^{2 2s^2 2p^1}, 801 kJ/mol) because removing an electron from a fully filled, penetrating 2s subshell requires more energy than removing an electron from a single 2p subshell.
• Nitrogen vs. Oxygen: Nitrogen (1s^{2 2s^2 2p^3}) has a higher IE₁ (1402 kJ/mol) than Oxygen (1s^{2 2s^2 2p^4}, 1314 kJ/mol) because half-filled p subshells provide extra exchange energy and stability, while Oxygen experiences inter-electronic repulsion in its paired 2p orbital.

Electropositivity and Metallic Character

Electropositivity is the tendency of an atom to lose valence electrons and form positive cations. It directly measures metallic character.

Periodic Variation of Electropositivity

• Across a Period: Decreases from left to right as ionisation energy and Z_eff increase, making electron loss difficult.
• Down a Group: Increases from top to bottom as atomic size increases and valence electrons experience weaker nuclear pull.
• Group 1 Alkali Metals: Exhibit the highest electropositivity in their respective periods due to their low first ionisation energies and large atomic radii.

Summary Comparison of Key Periodic Trends

Important Facts

• Helium possesses the highest first ionisation energy of all elements (2372 kJ/mol), combined with the smallest covalent atomic radius.
• Caesium is the most electropositive stable element found in nature, exhibiting the lowest non-radioactive ionisation energy (376 kJ/mol).
• Francium is technically the most electropositive element, but its extreme radioactivity and short half-life (22 minutes for ²²³Fr) make Caesium the standard reference.
• Noble gases have the largest atomic radii in their respective periods when measured as Van der Waals radii, because they do not form standard covalent bonds in ground states.
• Lanthanoid Contraction causes the atomic radii of 4d and 5d transition elements of the same vertical group to be nearly identical (e.g., Zirconium radius ≈ Hafnium radius, Niobium radius ≈ Tantalum radius).
• Poor shielding by 3d electrons causes Gallium (135 pm) to have a slightly smaller atomic radius than Aluminium (143 pm), an anomaly in Group 13.
• Second ionisation energy (IE₂) of alkali metals is exceptionally high because removing the second electron requires breaking a stable octet configuration.
• Shielding efficiency follows the order s > p > d > f; poor shielding by inner d and f orbitals increases Z_eff on outer electrons.
• Metallic character directly correlates with standard oxidation potentials and reducing strength in gaseous states.
• Diagonal relationship pairs (Li–Mg, Be–Al, B–Si) display similar atomic radii and electropositive properties due to their nearly equal charge-to-radius ratios.