Electronegativity Difference Calculator

Compute the electronegativity difference between two atoms, classify the bond as nonpolar covalent, polar covalent, or ionic, and explore bond polarity with a built‑in Pauling electronegativity table.

Electronegativity Difference & Bond Type

Element A

Element B

How the electronegativity difference is calculated

Electronegativity is a measure of how strongly an atom attracts shared electrons in a chemical bond. To estimate bond polarity, we compare the electronegativities of the two bonded atoms on the same scale (here we use the Pauling scale).

Formula

For atoms A and B with electronegativities \( \chi_A \) and \( \chi_B \):

\[ \Delta EN = |\chi_A - \chi_B| \]

Typical bond type cutoffs (Pauling scale)

Textbooks often use approximate ranges like these:

  • Nonpolar covalent: \( \Delta EN \lesssim 0.4 \)
  • Polar covalent: \( 0.4 \lesssim \Delta EN \lesssim 1.7 \)
  • Ionic: \( \Delta EN \gtrsim 1.7 \)

These are guidelines, not strict rules. Real bonds often have mixed ionic and covalent character, and the bonding environment (solid vs. gas, crystal structure, oxidation state) also matters.

Estimating percent ionic character

One empirical relationship (Pauling) for the percent ionic character of a bond is:

\[ \% \text{ ionic character} \approx \left(1 - e^{-0.25 (\Delta EN)^2}\right) \times 100\% \]

The calculator uses this expression to give a rough idea of how “ionic” a bond is, based on the electronegativity difference.

Understanding electronegativity and bond polarity

When two atoms form a bond, they share electrons. If both atoms attract the electrons equally, the bond is nonpolar covalent. If one atom attracts the electrons more strongly, the bond is polar covalent. If the attraction is so unequal that electrons are effectively transferred, the bond is often described as ionic.

Electronegativity trends in the periodic table

  • Electronegativity generally increases across a period (left → right).
  • It generally decreases down a group (top → bottom).
  • Fluorine is the most electronegative element (≈ 3.98 on the Pauling scale).

These trends reflect effective nuclear charge and atomic radius: smaller atoms with higher effective nuclear charge attract bonding electrons more strongly.

Examples of electronegativity differences

  • H–H (H2): ΔEN = 0 → nonpolar covalent.
  • C–H: ΔEN ≈ |2.55 − 2.20| = 0.35 → often treated as essentially nonpolar.
  • H–Cl: ΔEN ≈ |3.16 − 2.20| = 0.96 → polar covalent bond.
  • Na–Cl: ΔEN ≈ |3.16 − 0.93| = 2.23 → strong ionic character.

Limitations of using electronegativity difference alone

Electronegativity difference is a useful first approximation, but it has limitations:

  • It does not account for molecular geometry (e.g., CO2 has polar bonds but is nonpolar overall).
  • It ignores resonance and delocalization of electrons.
  • Different electronegativity scales (Pauling, Mulliken, Allred–Rochow) give slightly different values.
  • Bond character can change with oxidation state and coordination environment.

Use ΔEN as a qualitative guide, and combine it with structural information and experimental data when possible.

Step‑by‑step: determining bond type from electronegativity

  1. Look up or input the electronegativities of the two atoms on the same scale.
  2. Compute the absolute difference \( \Delta EN = |\chi_A - \chi_B| \).
  3. Compare ΔEN to the guideline ranges for nonpolar, polar covalent, and ionic bonds.
  4. Identify the more electronegative atom; it will carry a partial negative charge (δ−).
  5. Consider the molecular geometry to decide whether the molecule is polar overall.

Electronegativity difference – FAQ

How do you calculate electronegativity difference?

Use the absolute difference between the electronegativities of the two atoms on the same scale: \( \Delta EN = |\chi_A - \chi_B| \). For example, for HCl on the Pauling scale: EN(H) = 2.20, EN(Cl) = 3.16, so ΔEN = |3.16 − 2.20| = 0.96.

What electronegativity difference is considered ionic, polar covalent, or nonpolar?

On the Pauling scale, common textbook cutoffs are:

  • Nonpolar covalent: ΔEN ≲ 0.4
  • Polar covalent: 0.4 ≲ ΔEN ≲ 1.7
  • Ionic: ΔEN ≳ 1.7

These are approximate; real bonds often have mixed character.

Does a higher electronegativity difference always mean an ionic bond?

Not always. A large ΔEN indicates strong polarity and significant ionic character, but many bonds are best described as partially ionic and partially covalent. Lattice energy, coordination, and the solid‑state structure also influence whether a compound behaves more “ionic” or “covalent”.

Which electronegativity scale does this calculator use?

The built‑in element list uses Pauling electronegativities, which are standard in most general chemistry courses. If you prefer another scale (e.g., Mulliken or Allred–Rochow), you can overwrite the numeric values in the input fields and still use the same ΔEN formula.

Can a molecule with polar bonds be nonpolar overall?

Yes. If the molecular geometry is symmetric, the individual bond dipoles can cancel out. For example, CO2 has polar C–O bonds, but the linear geometry makes the molecule nonpolar overall. Electronegativity difference tells you about bond polarity, not necessarily molecular polarity.