Electronegativity

Moving left to right across a period, the nuclear charge (number of protons) increases while electrons are added to the same principal quantum shell. The shielding effect remains approximately constant because the new electrons are in the same shell and shield each other poorly. As a result, the effective nuclear charge (Z_eff) experienced by bonding electrons increases across the period. A higher Z_eff means a stronger pull on shared electrons, so electronegativity increases from left to right. For example, in Period 2: Li (1.0) < Be (1.5) < B (2.0) < C (2.5) < N (3.0) < O (3.5) < F (4.0). Noble gases are typically excluded because they do not ordinarily form covalent bonds.

Moving down a group, each successive element has an additional electron shell. Although the nuclear charge increases, the inner (core) electrons shield the bonding electrons from the nucleus very effectively. This increased shielding, combined with the larger atomic radius, reduces Z_eff experienced by bonding pairs. Consequently, the atom's ability to attract shared electrons decreases, and electronegativity decreases down a group. For Group 17 (halogens): F (4.0) > Cl (3.2) > Br (3.0) > I (2.7). Fluorine is therefore the most electronegative element in the entire periodic table.

The difference in electronegativity (ΔEN) between two bonded atoms determines bond character. If ΔEN < 0.4, the bond is considered non-polar covalent (e.g., H–H, C–H). If 0.4 ≤ ΔEN ≤ 1.7, the bond is polar covalent, with a permanent dipole (e.g., H–Cl, ΔEN = 1.0; H–F, ΔEN = 1.9 — borderline ionic). If ΔEN > 1.7, the bond is predominantly ionic (e.g., Na–Cl, ΔEN = 2.1). These thresholds are approximate; bond character is a continuum rather than a sharp division. Polar covalent bonds are represented with δ+ and δ− symbols or with an arrow pointing toward the more electronegative atom.

Metallic character refers to the tendency of an element to behave as a metal — losing electrons to form cations. Low electronegativity is synonymous with high metallic character: an atom that attracts shared electrons weakly will readily give them up entirely. Moving left to right across a period, metallic character decreases and electronegativity increases. Moving down a group, metallic character increases and electronegativity decreases. This is why the most metallic elements (caesium, francium, barium) appear at the bottom-left of the periodic table, and the most electronegative elements (F, O, N, Cl) appear at the top-right. The metalloid staircase separates metals from non-metals and marks the transition zone of intermediate electronegativity (roughly 1.8–2.2 on the Pauling scale).

A molecule is polar overall if (1) it contains polar bonds AND (2) those bond dipoles do not cancel due to molecular symmetry. For water (H₂O), ΔEN(O–H) = 3.5 − 2.1 = 1.4, making each O–H bond polar. The bent geometry means the two dipoles add vectorially, producing a net molecular dipole moment of 1.85 D. In contrast, carbon dioxide (CO₂) has two polar C=O bonds (ΔEN = 1.0) but is linear, so the dipoles cancel and the molecule is non-polar overall. Predicting molecular polarity requires both electronegativity data and knowledge of molecular geometry (VSEPR theory).

• E_{A-B} — Bond dissociation energy of A–B (kJ mol⁻¹): Energy required to break one mole of A–B bonds homolytically in the gas phase.
• \sqrt{E_{A-A} \cdot E_{B-B}} — Geometric mean of homonuclear bond energies (kJ mol⁻¹): Predicted bond energy if A–B were a purely non-polar covalent bond.
• \Delta EN — Electronegativity difference (dimensionless (Pauling units)): Absolute difference between the Pauling electronegativity values of atoms A and B.
• 96.48 — Conversion constant (kJ mol⁻¹): Converts the square of ΔEN (in eV²) to kJ mol⁻¹; reflects the ionic resonance stabilisation energy.

Problem. Arrange the following bonds in order of increasing polarity and identify which end carries δ−: (a) C–H, (b) N–H, (c) O–H, (d) F–H. Use Pauling electronegativities: H = 2.1, C = 2.5, N = 3.0, O = 3.5, F = 4.0.

Answer. Increasing polarity: C–H < N–H < O–H < F–H. The δ− charge resides on C (negligible), N, O, and F respectively.

• Ionisation Energy — Ionisation energy and electronegativity both increase across a period and decrease down a group for the same reason — increasing effective nuclear charge. Ionisation energy measures the ability to remove an electron from an isolated atom, while electronegativity measures the pull on shared electrons in a bond.
• Atomic Radius — Atomic radius decreases as electronegativity increases. A smaller atom has its bonding electrons closer to the nucleus, experiencing stronger attraction, which is why top-right elements are most electronegative.
• Polar Covalent Bonds — Polar covalent bonds arise directly from electronegativity differences between bonded atoms. The magnitude of bond polarity (and the δ+/δ− charges) is determined by ΔEN on the Pauling scale.
• Hydrogen Bonding — Hydrogen bonds form when H is bonded to a highly electronegative atom (N, O, or F). The high ΔEN creates a strongly polarised X–H bond, making the δ+ hydrogen interact strongly with a lone pair on a neighbouring electronegative atom.
• Metallic Character — Metallic character is inversely related to electronegativity. Low-electronegativity elements readily lose electrons (metallic behaviour), while high-electronegativity elements gain electrons (non-metallic behaviour). Understanding this relationship helps explain reactivity trends across the periodic table.