Electronegativity: definition

· Electronegativity = the power of an atom to attract electrons to itself.
· It applies to atoms within a bond, especially when comparing how strongly two bonded atoms attract the bonding pair of electrons.
· A more electronegative atom attracts the bonding electrons more strongly.
· A less electronegative atom attracts the bonding electrons less strongly.
· Electronegativity values are usually given on the Pauling scale in exam questions when needed.

Factors affecting electronegativity

· Nuclear charge: higher proton number gives stronger attraction for bonding electrons, so electronegativity increases.
· Atomic radius: smaller atoms attract bonding electrons more strongly because the bonding pair is closer to the nucleus.
· Shielding: more inner shells and sub-shells reduce the attraction between the nucleus and bonding electrons, so electronegativity decreases.
· Strong electronegativity usually requires: high nuclear charge, small atomic radius, and low shielding.
· Weak electronegativity usually results from: large atomic radius and high shielding.

Trend across a period

· Across a period, electronegativity increases.
· This is because nuclear charge increases as proton number increases.
· Electrons are added to the same principal shell, so shielding does not increase significantly.
· Atomic radius generally decreases, so the nucleus attracts bonding electrons more strongly.
· Example trend: Na < Mg < Al < Si < P < S < Cl.

This diagram helps visualise the repeated pattern in Pauling electronegativity values. It supports the idea that electronegativity generally increases across a period. It is useful for linking periodic trends to bonding predictions. Source

Trend down a group

· Down a group, electronegativity decreases.
· This is because atoms have more electron shells, so the bonding pair is further from the nucleus.
· Shielding increases, reducing the attraction from the nucleus to the bonding electrons.
· The increase in atomic radius and shielding outweighs the increase in nuclear charge.
· Example trend: F > Cl > Br > I.

Using electronegativity to predict bonding

· Compare the Pauling electronegativity values of the two bonded atoms.
· Small electronegativity difference → electrons are shared more equally → covalent bond.
· Large electronegativity difference → electron transfer is more likely → ionic bond.
· Metal + non-metal with a large electronegativity difference usually forms an ionic compound.
· Two non-metals usually form a covalent bond because both atoms tend to attract electrons.

Exam method: predicting ionic or covalent bonding

· Step 1: identify the two atoms involved in bonding.
· Step 2: use the given Pauling electronegativity values.
· Step 3: calculate the electronegativity difference: larger value – smaller value.
· Step 4: decide whether the difference is small or large.
· Step 5: predict covalent bonding for small differences and ionic bonding for large differences.
· Always support your answer using the terms nuclear charge, atomic radius, shielding, electron attraction, and electronegativity difference where relevant.

These diagrams show the difference between equal sharing, unequal sharing and electron transfer. They are useful for seeing why a larger electronegativity difference gives bonding with more ionic character. This supports exam explanations of how electronegativity predicts bonding type. Source

Common exam wording

· Define electronegativity: “the power of an atom to attract electrons to itself.”
· Explain increase across a period: nuclear charge increases, shielding similar, atomic radius decreases, so attraction for bonding electrons increases.
· Explain decrease down a group: atomic radius increases and shielding increases, so attraction for bonding electrons decreases.
· Predict bonding: use the difference in Pauling electronegativity values to decide whether bonding is mainly ionic or covalent.
· Avoid vague answers such as “it wants electrons more”; use attracts bonding electrons more strongly.