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General Chemistry I · Periodic Properties

Electronegativity

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On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Study tools
  7. Sources & references

In 30 seconds

Electronegativity (EN) is the ability of an atom in a chemical bond to attract the shared electrons toward itself. It is a relative, dimensionless property — most often given on the Pauling scale, where fluorine (4.0) is the most electronegative element. EN increases from left to right across a period and decreases down a group. The difference in electronegativity between two bonded atoms (ΔEN) tells you whether the bond is nonpolar covalent, polar covalent, or ionic.

Why this matters

Electronegativity differences explain almost everything about bond character: whether a bond is ionic or covalent, how polar a molecule is, and where electrons concentrate in a reaction. It predicts that water (O–H, ΔEN = 1.4) is highly polar while methane (C–H) is not, which in turn explains why oil and water don't mix. Electronegativity also rationalizes oxidation states — the more electronegative atom is assigned the electrons.

The college version

Key Ideas

  • Definition: the tendency of a bonded atom to pull shared electrons toward itself.
  • Pauling scale: fluorine = 4.0 (highest); cesium/francium ≈ 0.7 (lowest). It is a relative scale with no units.
  • Across a period (→): EN increases (higher Zeff pulls electrons more strongly).
  • Down a group (↓): EN decreases (larger, more shielded atoms attract shared electrons less).
  • Noble gases: usually excluded (they rarely form bonds).
  • Bond classification by ΔEN (approximate cutoffs):
    • ΔEN ≈ 0 → nonpolar covalent (electrons shared equally).
    • 0.4 < ΔEN < ~1.7 → polar covalent (unequal sharing).
    • ΔEN > ~1.7–2.0 → ionic (electrons effectively transferred).
  • Bond dipole: points toward the more electronegative atom (δ⁻).

Equations and Variables

  • ΔEN = |EN(A) − EN(B)| for a bond A–B.
    • ΔEN = electronegativity difference (dimensionless).
    • EN(A), EN(B) = Pauling electronegativities of the two atoms.
  • Useful Pauling values: F 4.0 · O 3.5 · N 3.0 · Cl 3.0 · Br 2.8 · C 2.5 · S 2.5 · H 2.1 · P 2.1 · Si 1.8 · Al 1.5 · Mg 1.2 · Na 0.9 · K 0.8 · Cs 0.7.

How It Works

  1. In a covalent bond, two atoms share a pair of electrons.
  2. If the atoms are identical (ΔEN = 0), the electrons are shared equally and the bond is nonpolar.
  3. If one atom has a higher EN, it pulls the shared pair closer, giving itself a partial negative charge (δ⁻) and its partner a partial positive charge (δ⁺) — a polar covalent bond.
  4. If ΔEN is very large, the more electronegative atom essentially takes the electron completely, forming ions — an ionic bond.
  5. Because Zeff rises across a period, atoms on the right attract electrons more strongly; because atoms get larger down a group, atoms on the bottom attract electrons less.

Worked Example

Classify the bond in each compound and state the direction of the dipole (if any).

  1. NaCl: ΔEN = |3.0 − 0.9| = 2.1 → ionic. Chlorine takes the electron from sodium (Na⁺ Cl⁻).
  2. HCl: ΔEN = |3.0 − 2.1| = 0.9 → polar covalent. Dipole points from H (δ⁺) toward Cl (δ⁻).
  3. H₂: ΔEN = |2.1 − 2.1| = 0 → nonpolar covalent, no dipole.
  4. C–H in CH₄: ΔEN = |2.5 − 2.1| = 0.4 → essentially nonpolar (the C–H bond is only weakly polar).

Ranking check: order these by electronegativity — F, O, N, C. From the Pauling scale: F (4.0) > O (3.5) > N (3.0) > C (2.5).

Common Confusions

  • "Electronegativity is an energy." — Wrong: it is a relative, unitless scale. Electron affinity and ionization energy are energies; electronegativity is a ranking.
  • "Fluorine has the highest electron affinity." — Wrong: fluorine has the highest electronegativity; chlorine has the highest electron affinity (see the electron-affinity note).
  • "A polar bond always makes a polar molecule." — Wrong: symmetric molecules like CO₂ and CCl₄ have polar bonds but their dipoles cancel, making them nonpolar.
  • "ΔEN > 1.7 is a hard, exact boundary for ionic bonds." — It is only approximate; bond type is a continuum, and context (metals vs. nonmetals) matters.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Think of a game of tug-of-war over a shared rope (the bonding electrons). Electronegativity is how strong each player is. Two equally strong players (same element) keep the rope centered — a nonpolar bond. A much stronger player pulls the rope toward herself — a polar bond, and the rope (electrons) spends more time on her side. If one player is overwhelmingly strong, she simply takes the whole rope — that's an ionic bond, where the electron is transferred, not shared. Fluorine is the strongest player on the table. (The analogy captures unequal sharing; real electronegativity comes from Zeff and atomic size, not muscles.)

Key takeaways

  • Electronegativity is a relative, unitless scale (Pauling); F = 4.0 is the maximum.
  • EN increases left → right; decreases top → bottom (noble gases excluded).
  • Most electronegative: F (4.0); then O (3.5), N/Cl (3.0). Least: Cs/Fr (0.7).
  • ΔEN ≈ 0 → nonpolar covalent; 0.4–1.7 → polar covalent; >1.7–2.0 → ionic.
  • Bond dipole points toward the more electronegative atom (δ⁻).
  • Electronegativity is not the same as electron affinity (an energy) or ionization energy (an energy); EN is a relative scale.
  • Polar bonds may still form a nonpolar molecule if their dipoles cancel (see molecular-polarity note).
  • EN = ability of a bonded atom to attract shared electrons; Pauling scale, F = 4.0.
  • Increases left → right; decreases top → bottom.
  • ΔEN classifies bonds: ≈0 nonpolar covalent; 0.4–1.7 polar covalent; >1.7–2.0 ionic.
  • Dipole points toward the more electronegative atom (δ⁻).
  • Most electronegative: F, O, N, Cl. Least: Cs, Fr.
  • Polar bonds ≠ automatically polar molecules (dipoles can cancel).

Keep learning

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Practice General Chemistry I

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Study tools & related lessonsYou’ll learn to · Related

You’ll learn to

  • Define electronegativity and distinguish it from electron affinity and ionization energy.
  • Describe the periodic trends in electronegativity (Pauling scale).
  • Use electronegativity differences (ΔEN) to classify bonds as nonpolar covalent, polar covalent, or ionic.
  • Predict the direction of a bond dipole from electronegativity values.

Sources & references

  1. OpenStax, *Chemistry 2e*, "6.5 Periodic Variations in Element Properties."
  2. LibreTexts Chemistry, "Electronegativity."
  3. OpenStax, *Chemistry 2e*, "7.2 Covalent Bonding."
  4. PubChem, "Fluorine (Element)."

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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