Types of elementsHigh School / HonorsAP Chemistry

Bond Type and Electronegativity Difference

What you'll be able to do: Use electronegativity difference to classify a bond as nonpolar covalent, polar covalent or ionic.

Best after: Metals, Nonmetals and Metalloids, Atomic Radius, Ionization Energy and Electronegativity

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Introduction

Bonding is a spectrum, not three separate boxes. Electronegativity difference tells you where on that spectrum a given bond sits.

These are recommended, not required. You can start this lesson at any time.

Learning objectives

Calculate the electronegativity difference for a bond
Classify bonds as nonpolar covalent, polar covalent or ionic
Assign partial charges and a bond dipole direction
Explain why bond type is a continuum

Electronegativity difference sets the bond

Take the absolute difference between the two electronegativity values. A common classroom scheme: below about 0.4 is nonpolar covalent, 0.4 to about 1.7 is polar covalent, and above about 1.7 behaves as ionic. The cutoffs are conventions, not laws of nature.

Use both values from the same electronegativity scale and take the absolute difference, so the order of subtraction cannot give a negative bond polarity. For O–H, |3.44 − 2.20| = 1.24. For C–H, |2.55 − 2.20| = 0.35. This ranks O–H as more polar without requiring a universal boundary between categories. Electronegativity difference is dimensionless; a value of 1.24 is not 1.24 electrons transferred.

delta_EN = |EN_A - EN_B|

Partial charges and the bond dipole

In a polar covalent bond the more electronegative atom carries a partial negative charge and the other a partial positive charge. The bond dipole arrow points from the partial positive atom toward the partial negative atom.

Write δ⁺ on H and δ⁻ on O for an O–H bond. The δ symbol means partial charge: the bond electrons are shared unequally rather than assigning a whole electron to one atom. The crossed chemistry dipole arrow points toward the electron-rich end, with its positive marker at the electron-poor end. A monatomic ion such as Na⁺ instead has a net integer charge. Formal charge is another bookkeeping convention and should not be substituted for a bond's partial charges.

H–H is shown without partial charges; H–Cl has H delta positive and Cl delta negative with a crossed-tail dipole arrow toward chlorine. A sodium cation and chloride anion are shown separately as constituents of an extended ionic lattice. Partial charges and integer ion charges are distinguished; the categories are models rather than exact electronegativity thresholds. — Partial charges on a covalent bond are different from integer ionic charges. The categories describe models along a continuum.

δ charges describe unequal sharing of electron density; they are not the ion’s integer charge or the formal charge.

Why the cutoffs are soft

HF has a difference of 1.9 yet is a molecular gas, while some metal halides with smaller differences behave ionically. Use the difference as a ranking tool and combine it with element type before declaring a compound ionic.

Ionic and covalent character form a continuum because electron density can be shared to different degrees. An ionic solid such as NaCl is an extended lattice of ions, not a collection of independent NaCl molecules. HF consists of covalently bonded molecules despite its large electronegativity difference. For a borderline case, combine element types with actual structure and properties; an arbitrary cutoff alone cannot establish the existence of a lattice.

Rank bonds by polarity confidently; classify borderline cases with element type as well.

Polar bond is not the same as polar molecule

CO₂ has two strongly polar bonds, but they point in opposite directions and cancel, so the molecule is nonpolar overall. Molecular polarity needs geometry, which is covered in the bonding unit.

Linear CO2 has equal C-to-O bond-dipole arrows in opposite directions, summing to zero. Bent water has both H-to-O bond-dipole arrows angled toward oxygen; their sideways components cancel while their components along the bisector add to a nonzero net dipole. Arrows represent direction and qualitative addition, not measured dipole magnitudes. — Add bond-dipole vectors using molecular geometry: linear CO₂ cancels while bent H₂O does not.

Add bond dipoles using the molecular geometry

Molecule	Polar bonds?	Arrangement	Net result
CO₂	C=O bonds are polar	Two equal opposing dipoles, linear	Cancel; nonpolar molecule
H₂O	O–H bonds are polar	Bent	Do not cancel; polar molecule
CCl₄	C–Cl bonds are polar	Four identical substituents, tetrahedral	Cancel; nonpolar molecule
CH₃Cl	C–Cl and C–H differ	Tetrahedral with unlike substituents	Do not cancel; polar molecule

The workflow is bond polarity, Lewis structure, molecular shape, then vector sum. “Tetrahedral” alone does not guarantee a nonpolar molecule: the equal-bond symmetry in CCl₄ matters. You can identify individual bond dipoles now and return to detailed geometry in Unit 5. A nonpolar molecule can still attract another nonpolar molecule through dispersion forces.

First ask whether each bond is polar; then ask whether the dipoles cancel.

Key ideas

Definition

Electronegativity

A measure of how strongly an atom attracts the shared electrons in a bond.

Rule

Classification bands

Approximate classroom bands such as <0.4, 0.4–1.7 and >1.7 are conventions for nonpolar, polar and ionic character, with exceptions; combine the estimate with structure and element type.

Rule

Dipole direction

The arrow points toward the more electronegative atom.

Key concept

Geometry matters

Polar bonds can cancel, so a molecule with polar bonds may still be nonpolar.

Equation

Electronegativity difference

delta_EN = |EN_A - EN_B|

delta_EN = electronegativity difference for the bond
EN_A = electronegativity of the first atom
EN_B = electronegativity of the second atom

Worked examples

Worked example 1

Classify the bonds in Cl₂, HCl and NaCl using EN values Cl 3.16, H 2.20, Na 0.93.

Try it first: Compute each difference before classifying anything.

Worked example 2

Rank the bonds C-H, O-H and N-H by polarity. EN values: H 2.20, C 2.55, N 3.04, O 3.44.

Try it first: All three share hydrogen, so only the partner matters.

Worked example 3

Why is CO₂ nonpolar overall while H₂O is polar, although each contains polar bonds?

Try it first: Sketch the orientation of the bonds in each molecule.

Common mistakes

Common mistakes and better checks
Mistake	Why it's wrong	Check instead
Treating 1.7 as an exact boundary.	It is a convention with well known exceptions such as HF.	Use the difference to rank polarity and check element types for borderline cases.
Saying a molecule with polar bonds must be polar.	Symmetric geometries cancel bond dipoles, as in CO₂ and CCl₄.	Check the shape before assigning molecular polarity.
Drawing the dipole arrow toward the less electronegative atom.	Electron density shifts toward the stronger attractor.	Point the arrow at the more electronegative atom.

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What you should now know

The difference in electronegativity between two bonded atoms sets how unevenly the shared electrons sit. Small differences give nonpolar covalent bonds, intermediate differences give polar covalent bonds with partial charges, and large differences give essentially ionic bonding.

Electronegativity difference places a bond on the ionic to covalent spectrum
Partial charges appear on polar covalent bonds
Cutoff values are conventions, not hard rules
Polar bonds can cancel in symmetric molecules

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Sources and further reading

Chemistry 2e, Section 6.5: Periodic Variations in Element Properties
Flowers, Theopold, Langley, Robinson · OpenStax, Rice University · Chapter 6.5
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The current online Chemistry 2e edition lists CC BY-NC-SA 4.0. The source version used for this lesson and its applicable rights remain unverified.
Recorded use: further reading. This describes the stored citation, not verified authorship or reuse permission.
Recorded attribution / license claim (applicability unverified): Access for free at openstax.org
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Citation coverage needs review: This broad periodic-trends reference does not establish the exact source used for bond-polarity explanations.