Electronegativity
Electronegativity describes how strongly an atom attracts electron density while it is bonded to another atom. The idea is powerful because it links periodic structure to bond polarity, but it is a relative model—not a literal charge sitting on an isolated atom.
Electronegativity and bond polarity
Meaning and periodic behavior
On the Pauling scale, fluorine is the most electronegative element. Across a period, increasing effective nuclear charge generally increases attraction for shared electrons. Down a group, larger size and shielding usually weaken that attraction.
The environment matters. Oxidation state, coordination and bonding partners can shift the effective attraction, so electronegativity values are best used as chemically informed comparisons rather than immutable constants.
More than one electronegativity scale
Pauling
Derived from bond-energy differences; the most common scale for bond polarity.
Mulliken
Connects electronegativity to ionization energy and electron affinity.
Allred–Rochow
Connects attraction to effective nuclear charge and covalent radius.
The numerical scales differ, but their ordering is similar for most main-group chemistry.
Bond polarity is a continuum
If bonded atoms have different electronegativities, shared electron density shifts toward the more electronegative atom, creating partial charges δ⁺ and δ⁻.
No universal Δχ value changes a bond abruptly from covalent to ionic. Labels such as nonpolar covalent, polar covalent and ionic describe useful regions on a continuum and should be checked against structure and measured properties.
Bond-polarity comparer
Δχ = 1.0
From bonds to molecules
Molecular polarity requires geometry
A bond dipole is a vector. Molecular polarity is the vector sum of all bond dipoles, so shape matters as much as the polarity of each bond.
Polarization and covalent character
A small, highly charged cation can distort the electron cloud of a large, polarizable anion. This polarizing effect increases covalent character even in compounds conveniently written as ions.
AlCl₃ therefore behaves much more covalently than NaCl despite both containing a metal and chlorine. Bonding models are limits on a continuum, not rigid boxes.
Oxidation state is bookkeeping, not measured charge
Electronegativity helps assign bonding electrons when calculating oxidation states. The resulting oxidation number is a formal accounting device, not the physical charge on the atom.
Carbon is assigned +4 in CO₂ because oxygen is more electronegative, but carbon does not carry a literal +4e point charge in the molecule.
CO₂ has polar C=O bonds but is linear, so equal opposing dipoles cancel. SO₂ is bent, so its S–O dipoles do not cancel and the molecule has a net dipole.
Chemical consequences
Concrete values and chemical reading
On the Pauling scale, typical values are approximately H 2.20, C 2.55, N 3.04, O 3.44 and F 3.98. The numbers are most useful comparatively. In an O–H bond, oxygen draws electron density toward itself; in a C–H bond, the difference is much smaller.
The direction of a bond dipole follows the electronegativity difference, but its magnitude also depends on bond length and the actual electronic structure. Two bonds with similar Δχ need not have identical dipole moments.
Likewise, bond polarity is not bond strength. H–F is very polar and strong, but a larger Δχ does not universally mean a stronger bond. Bond strength depends on orbital overlap, bond order, bond length and the surrounding molecular environment.
Inductive effects in molecules
Electronegativity can influence atoms beyond the directly bonded pair. In organic molecules, an electronegative substituent such as F or Cl withdraws electron density through σ bonds. This inductive effect weakens with distance but can change acidity, stability of charged intermediates and reaction rates.
For example, replacing H atoms near a carboxylic acid with F atoms tends to stabilize the conjugate-base negative charge through induction, increasing acidity. The explanation is not “fluorine grabs the proton”; it is redistribution of electron density through the bonding framework.
This is why electronegativity is more powerful when connected to a mechanism than when used as a memorized arrow on the periodic table.
Comparing real bonds
| Bond | Approx. Δχ (Pauling) | Useful first reading |
|---|---|---|
| C–H | 0.35 | Weakly polar; often treated as nearly nonpolar. |
| N–H | 0.84 | Polar toward N. |
| O–H | 1.24 | Strongly polar toward O. |
| Na–Cl | 2.23 | Strong ionic character in the solid lattice. |
The table is a starting point, not a classification law. Bonding is affected by charge, geometry, polarization and the surrounding structure. A large Δχ suggests strong charge separation, but it does not by itself tell you whether the substance forms molecules, an extended lattice or something intermediate.
Electronegativity can change with chemical environment
An atom is not equally electron-attracting in every oxidation state or coordination environment. Removing electron density tends to make the remaining electron cloud contract and can increase the atom’s effective attraction for bonding electrons. Adding electron density can have the opposite effect.
This is one reason advanced electronegativity scales sometimes assign different values to the same element in different valence states. The familiar periodic-table number is therefore a convenient reference, not a complete description of every bond that element can form.
For transition metals this matters especially because d-electron occupancy, oxidation state and ligand environment can all change the distribution of electron density. A single Pauling value remains useful for orientation, but not for a full electronic-structure calculation.
Acidity and bond polarization
Electronegativity helps explain why O–H bonds in acids are polarized and why nearby electronegative substituents can strengthen acidity through induction. When the conjugate base forms, electron-withdrawing groups can stabilize its negative charge by shifting electron density through σ bonds.
This does not mean “the most electronegative atom always makes the strongest acid.” Bond strength, resonance, solvation and molecular structure also matter. Electronegativity supplies one part of the mechanism, and good chemical reasoning identifies when that part dominates.
Transition metals and variable oxidation state
For transition metals, a single textbook electronegativity value hides more variation than it does for many main-group elements. The metal can appear in several oxidation states and coordination environments, changing orbital energies, bond lengths and the distribution of electron density.
That is one reason electronegativity should not be used as a stand-alone predictor of whether a transition-metal bond is “ionic” or “covalent”. Metal–ligand bonding can combine electrostatic attraction, σ donation and π interactions, and the balance changes from one complex to another.
The periodic number remains useful as a first comparison, but the actual molecule or solid decides how the electrons are distributed.
Exercises
H–Cl
Which end of an H–Cl bond is δ⁻?
Solution
Cl, because chlorine is more electronegative.
CO₂ and SO₂
Both contain polar bonds. Why is only SO₂ polar overall?
Solution
CO₂ is linear and its bond dipoles cancel. SO₂ is bent, so the dipoles add to a nonzero vector.
Electron affinity comparison
Why is electronegativity not the same as electron affinity?
Solution
Electron affinity is an energy change for a gas-phase atom gaining an electron. Electronegativity is a relative attraction for electron density in a bond.