Revise Electronegativity and bonding for Cambridge 9701 (3.1)

Electronegativity is formally defined as the ability of an atom to attract the shared pair of electrons in a covalent bond. It is a measure of an atom's 'pulling power' on electrons within a bond. We quantify this using a numerical scale, most commonly the Pauling scale, which runs from approximately 0.7 (for Francium) to 4.0 (for Fluorine). A higher value indicates a greater attraction for electrons.

The type of chemical bond that forms between two atoms is directly related to the difference in their electronegativity values (ΔEN). Rather than being distinct categories, bonding types exist on a continuum. When two atoms with identical electronegativity bond (e.g., Cl-Cl), the electrons are shared perfectly equally, forming a pure (non-polar) covalent bond. When there is a small to moderate difference, the more electronegative atom pulls the electron density towards itself. This creates a polar covalent bond, with partial positive and negative charges.

$H^{\delta+} - Cl^{\delta-}$

In the example of hydrogen chloride above, chlorine is more electronegative than hydrogen. The bonding electrons spend more time, on average, closer to the chlorine atom. This gives chlorine a partial negative charge ($\delta^{-}$) and hydrogen a partial positive charge ($\delta^{+}$). If the electronegativity difference is very large (typically > 1.7), the more electronegative atom effectively 'wins' the tug-of-war completely, and an electron is transferred from one atom to the other. This results in the formation of ions and an ionic bond.

It is crucial to distinguish between bond polarity and molecular polarity. A molecule's overall polarity depends on two factors: the polarity of its individual bonds AND its three-dimensional shape. A molecule can contain polar bonds but be non-polar overall if its shape is symmetrical, causing the individual bond dipoles to cancel each other out.