Attracting a bonding pair
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself.
Electronegativity values can be expressed using the Pauling scale. Fluorine is the most electronegative element and has a Pauling value of .
The strength with which an atom attracts bonding electrons depends on the attraction between its positive nucleus and the electrons in the covalent bond.
Factors affecting electronegativity
| Factor | Change | Effect on electronegativity |
|---|---|---|
| Nuclear charge | More protons in the nucleus | The attraction for bonding electrons becomes stronger, increasing electronegativity. |
| Atomic radius | Greater distance between the nucleus and outer electrons | The attraction for the bonding pair becomes weaker, decreasing electronegativity. |
| Shielding | More filled inner electron shells | The bonding electrons experience less attraction from the nucleus, decreasing electronegativity. |
Periodic trends
Across a period, electronegativity generally increases.
- Nuclear charge increases as proton number increases.
- No new main electron shell is added, so shielding changes relatively little.
- Atomic radius decreases.
- The nucleus therefore attracts a shared bonding pair more strongly.
Down a group, electronegativity generally decreases.
- Additional occupied electron shells increase atomic radius.
- Greater shielding reduces the attraction experienced by bonding electrons.
- These effects outweigh the increase in nuclear charge.
For example, sodium is more electronegative than caesium because sodium has fewer occupied shells, less shielding and a smaller atomic radius.
Exam Tip: When explaining an electronegativity trend, link the trend to nuclear charge, atomic radius and shielding. Do not state only that the element is further across or further up the Periodic Table.