3.1 Electronegativity and bonding

Electronegativity is the power of an atom to attract the bonding pair of electrons towards itself in a covalent bond. It is a comparative property of bonded atoms, not the charge on an isolated ion and not the energy needed to remove an electron.

When two bonded atoms have different electronegativities, the bonding electrons are attracted more strongly towards the atom with the higher value. This unequal attraction can make the bond polar; similar values indicate more equal sharing.

Use electronegativity to describe the direction and relative strength of attraction within a bond. Periodic trends and the factors that cause them belong to the neighbouring objectives; numerical differences used to predict ionic or covalent bonding also belong to the later application card.

Do not treat electronegativity as a fixed whole-atom charge, confuse it with first ionisation energy, or infer the polarity of an entire molecule from one bond without considering the other bonds and molecular shape. This card establishes the definition only.

Greater nuclear charge tends to strengthen the attraction between the nucleus and a bonding electron pair. More protons increase the positive pull, but this effect must be considered with the electron's distance from the nucleus and the shielding provided by inner electrons.

A larger atomic radius places the bonding electron pair farther from the nucleus, so attraction is weaker and electronegativity is lower. A smaller radius brings the bonding electrons closer and strengthens the attraction.

Inner-shell electrons shield the outer region from some nuclear attraction. Greater shielding reduces the effective pull on bonding electrons; therefore proton number alone cannot determine electronegativity. Use nuclear charge, radius and shielding together.

Do not confuse shielding with the number of bonding electrons, or state that every increase in proton number automatically gives a stronger attraction. The combined factor explanation supports the periodic trends on the neighbouring card; numerical bond classification belongs to the later application card.

Across a period, electronegativity generally increases from left to right. Nuclear charge increases while the added electrons enter the same principal shell, so shielding changes relatively little; the stronger nuclear attraction and generally smaller atomic radius pull bonding electrons more strongly.

Down a group, electronegativity generally decreases. Although nuclear charge increases, new occupied shells increase atomic radius and inner-shell shielding, so the bonding electron pair is farther from the nucleus and feels a weaker effective attraction.

Explain a trend by linking direction to the three factors: nuclear charge, distance/atomic radius and shielding. State the overall trend first, then identify which factor strengthens or weakens attraction and why the dominant effect gives the observed direction.

Use ‘generally’ rather than claiming a perfectly smooth rule for every comparison. Do not reverse the down-group direction, explain the trend with proton number alone, or use bond-polarity classification before the neighbouring application objective.

Use the supplied Pauling electronegativity values in a fixed sequence: compare the two bonded atoms, determine the size of the difference, and then classify the bond using the stated syllabus rule. Similar values support equal or nearly equal sharing; a larger difference means the bonding pair is attracted more strongly towards one atom.

A small difference is consistent with a non-polar covalent bond, an intermediate difference with a polar covalent bond, and a sufficiently large difference with ionic bonding under the assessed classification. In a polar covalent bond, show the partial-charge direction towards the more electronegative atom rather than describing full electron transfer.

The electronegativity difference predicts how unevenly the bonding electrons are shared: the greater the difference, the greater the unequal attraction and bond polarity. Keep the bond-level conclusion tied to the supplied values and the stated classification boundary.

Do not invent a universal numerical cutoff when values are not supplied, confuse bond polarity with the overall polarity of a molecule, or treat a polar covalent bond as complete ionic electron transfer. Covalent character in ionic compounds is outside the assessed scope here.