Electronegativity and Bond Character Notes

Electronegativity is the power of a covalently bonded atom to attract the shared pair of electrons toward itself in a molecule.
It is a dimensionless property (no unit) because it represents a tendency, not a measured quantity with units.
It essentially indicates the net result of the tendencies of atoms to attract bond-forming electron pairs.
Electronegativity is measured on several scales; the most commonly used scale was designed by Linus Pauling.

Size of the atom (atomic radius)
Nuclear charge (number of protons)
Screening (shielding) by inner electrons
How these factors interplay:
- As the number of protons in the nucleus increases, the nuclear charge increases, which tends to increase electronegativity.
- If the atomic size decreases (nuclear charge dominates and shielding is limited), bonding electrons feel a stronger pull toward the nucleus, increasing electronegativity.
- Conversely, when the number of inner electrons increases, shielding increases and the distance between the nucleus and valence electrons effectively grows, reducing the attraction of bonding electrons to the nucleus and decreasing electronegativity.

Down a group: electronegativity generally decreases from top to bottom.
- Although nuclear charge increases with more protons, the addition of electron shells increases shielding, which reduces the attraction of bonding electrons to the nucleus.
- Result: the effective attraction to bonding electrons weakens as you move down a group.
Across a period: electronegativity increases from left to right.
- Effective nuclear charge increases while shielding remains relatively constant, causing atomic size to shrink and the nucleus to attract bonding electrons more strongly.
Example in the halogens (F to At): electronegativity decreases as you move down the group (fluorine > chlorine > bromine > iodine > astatine).
Example in the third period: electronegativity increases from sodium to chlorine.

Ionic bonds: ΔEN is large, typically greater than ext{Δ}EN > 1.8.
Polar covalent bonds: ΔEN is intermediate, typically between 0.4 < ext{Δ}EN \le 1.8.
Nonpolar covalent bonds: ΔEN is small, typically ext{Δ}EN \le 0.4.
If the electronegativities of the two atoms are equal, the bond is nonpolar with the electron pair shared equally.

Covalent character in a compound depends on the polarizing power of the cation, which in turn depends on the oxidation state of the cation.
Higher polarizing power of the cation leads to greater covalent character in the compound.
This tendency helps explain why covalent character increases across a period (greater effective nuclear charge and polarizing ability).
Example: Aluminum chloride vs magnesium chloride.
- Aluminum chloride (AlCl3) is more covalent in nature than magnesium chloride (MgCl2).
- Reason: Aluminum is in the +3 oxidation state, magnesium in +2; higher oxidation state gives higher polarizing power, enabling greater polarization of the anion and sharing of electrons.
Therefore, higher cation polarizing power promotes covalent character in the compound.

Predicting bond type (ionic vs polar covalent vs nonpolar covalent) based on ΔEN guides expectations for salt formation, solubility, melting points, and reactivity.
Understanding covalent character helps explain deviations from purely ionic behavior in many compounds, especially those involving highly charged small cations (e.g., Al3+).
The trends tie back to foundational principles: electron configuration, nuclear charge, shielding, atomic size, and their impact on electron distribution in bonds.