What electronegativity difference makes a bond ionic?

A common rule is ΔEN < 0.4 nonpolar covalent, 0.4–1.7 polar covalent and above 1.7 ionic; some textbooks use 1.8 or 2.0 instead. At ΔEN = 1.7 Pauling's formula gives about 50 % ionic character, which is where 1.7 comes from. These are only guidelines: HF has ΔEN = 1.78 but is a covalent molecule, while LiI (ΔEN 1.68) is an ionic solid. “Metal + nonmetal → ionic, nonmetal + nonmetal → covalent” is often more reliable.

How do you calculate percent ionic character?

Pauling's formula is 1 − exp(−ΔEN²/4): about 20.6 % for HCl (ΔEN 0.96) and 71 % for NaCl (ΔEN 2.23). Hannay and Smyth proposed 16ΔEN + 3.5ΔEN² (%), giving 18.6 % for HCl. A more direct method divides the measured dipole moment by the dipole of a fully ionic bond, e·d: for HCl, μ = 1.08 D and d = 127.5 pm, so e·d = 6.12 D and the ionic character is 1.08 / 6.12 ≈ 17.6 %.

How do I tell if a molecule is polar or nonpolar?

First check whether the bonds are polar, then whether the shape lets the bond dipoles cancel. Treat each bond dipole as a vector pointing from δ+ to δ−; if the vectors add to zero the molecule is nonpolar. Linear CO₂ has two equal, opposite C=O dipoles, so it is nonpolar; bent H₂O cannot cancel its O–H dipoles, so it is polar (1.85 D). Likewise BF₃ and CCl₄ are nonpolar and NH₃ and CHCl₃ are polar. O₃ has only O–O bonds yet a 0.53 D dipole, because of its formal charges and bent shape.

What intermolecular forces does a molecule have, and how do they affect boiling point?

Every molecule has London dispersion forces, which grow with size and number of electrons. Polar molecules add dipole–dipole attraction, and molecules with H bonded directly to N, O or F form hydrogen bonds, the strongest of the three. For similar molar masses the boiling point order is usually hydrogen bonding > dipole–dipole > dispersion only: H₂O boils at 100 °C versus −60 °C for H₂S, and HF at 20 °C versus −85 °C for HCl. When molar masses differ a lot, dispersion can dominate — CCl₄ (77 °C) boils far above CH₄ (−162 °C).

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🧷Bond Polarity Calculator - Electronegativity Difference, Percent Ionic Character & Molecular Polarity

Q: How do you calculate percent ionic character?

Pauling's formula is 1 − exp(−ΔEN²/4): about 20.6 % for HCl (ΔEN 0.96) and 71 % for NaCl (ΔEN 2.23). Hannay and Smyth proposed 16ΔEN + 3.5ΔEN² (%), giving 18.6 % for HCl. A more direct method divides the measured dipole moment by the dipole of a fully ionic bond, e·d: for HCl, μ = 1.08 D and d = 127.5 pm, so e·d = 6.12 D and the ionic character is 1.08 / 6.12 ≈ 17.6 %.

Q: How do I tell if a molecule is polar or nonpolar?

First check whether the bonds are polar, then whether the shape lets the bond dipoles cancel. Treat each bond dipole as a vector pointing from δ+ to δ−; if the vectors add to zero the molecule is nonpolar. Linear CO₂ has two equal, opposite C=O dipoles, so it is nonpolar; bent H₂O cannot cancel its O–H dipoles, so it is polar (1.85 D). Likewise BF₃ and CCl₄ are nonpolar and NH₃ and CHCl₃ are polar. O₃ has only O–O bonds yet a 0.53 D dipole, because of its formal charges and bent shape.

Q: What intermolecular forces does a molecule have, and how do they affect boiling point?

Every molecule has London dispersion forces, which grow with size and number of electrons. Polar molecules add dipole–dipole attraction, and molecules with H bonded directly to N, O or F form hydrogen bonds, the strongest of the three. For similar molar masses the boiling point order is usually hydrogen bonding > dipole–dipole > dispersion only: H₂O boils at 100 °C versus −60 °C for H₂S, and HF at 20 °C versus −85 °C for HCl. When molar masses differ a lot, dispersion can dominate — CCl₄ (77 °C) boils far above CH₄ (−162 °C).

Bond polarity comes from the electronegativity difference ΔEN between two atoms: the more electronegative atom pulls the bonding electrons towards itself and carries a partial negative charge δ−. For H–Cl, the Pauling electronegativities are 2.20 and 3.16, so ΔEN = 0.96 — a polar covalent bond. Pauling's formula 1 − exp(−ΔEN²/4) gives about 20.6 % ionic character and the Hannay–Smyth formula 16ΔEN + 3.5ΔEN² gives 18.6 %; NaCl, with ΔEN = 2.23, is about 71 % ionic. This tool shows ΔEN, the bond type under three common textbook cut-offs (0.4/1.7, 0.4/1.8 and 0.5/2.0 — rules of thumb, not hard limits), both ionic-character estimates, a δ+/δ− bond-dipole diagram, and places the compound on the van Arkel–Ketelaar triangle (average EN vs ΔEN) with ionic, covalent and metallic regions. In molecule mode it takes the VSEPR geometry and adds the bond dipoles as 3D vectors: CO₂, BF₃ and CCl₄ are nonpolar because their bond dipoles cancel, while H₂O, NH₃ and CHCl₃ are polar. It lists measured gas-phase dipole moments for about 30 common molecules (H₂O 1.85 D, NH₃ 1.47 D), identifies dispersion, dipole–dipole and hydrogen-bonding forces, and compares boiling points. The dipole calculator applies μ = q·d, converts between debye and C·m (1 D = 3.33564×10⁻³⁰ C·m), and finds percent ionic character from a measured dipole and bond length — HCl, 1.08 D and 127.5 pm, is 17.6 % ionic.

Free · No sign-up · Runs in your browserLast updated: 2026-10-29

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How to use Bond Polarity Calculator

1In Bond polarity, type two elements (H–Cl, C=O, NaCl) or pick them from the dropdowns
2Switch to Molecular polarity and enter a single-centre molecule such as CO2, H2O or CHCl3 to see dipole arrows, intermolecular forces and a boiling-point comparison
3Use the Dipole moment tab to compute μ = q·d, convert D ↔ C·m, or get percent ionic character from a measured dipole and bond length

FAQ

What electronegativity difference makes a bond ionic?: A common rule is ΔEN < 0.4 nonpolar covalent, 0.4–1.7 polar covalent and above 1.7 ionic; some textbooks use 1.8 or 2.0 instead. At ΔEN = 1.7 Pauling's formula gives about 50 % ionic character, which is where 1.7 comes from. These are only guidelines: HF has ΔEN = 1.78 but is a covalent molecule, while LiI (ΔEN 1.68) is an ionic solid. “Metal + nonmetal → ionic, nonmetal + nonmetal → covalent” is often more reliable.
How do you calculate percent ionic character?: Pauling's formula is 1 − exp(−ΔEN²/4): about 20.6 % for HCl (ΔEN 0.96) and 71 % for NaCl (ΔEN 2.23). Hannay and Smyth proposed 16ΔEN + 3.5ΔEN² (%), giving 18.6 % for HCl. A more direct method divides the measured dipole moment by the dipole of a fully ionic bond, e·d: for HCl, μ = 1.08 D and d = 127.5 pm, so e·d = 6.12 D and the ionic character is 1.08 / 6.12 ≈ 17.6 %.
How do I tell if a molecule is polar or nonpolar?: First check whether the bonds are polar, then whether the shape lets the bond dipoles cancel. Treat each bond dipole as a vector pointing from δ+ to δ−; if the vectors add to zero the molecule is nonpolar. Linear CO₂ has two equal, opposite C=O dipoles, so it is nonpolar; bent H₂O cannot cancel its O–H dipoles, so it is polar (1.85 D). Likewise BF₃ and CCl₄ are nonpolar and NH₃ and CHCl₃ are polar. O₃ has only O–O bonds yet a 0.53 D dipole, because of its formal charges and bent shape.
What intermolecular forces does a molecule have, and how do they affect boiling point?: Every molecule has London dispersion forces, which grow with size and number of electrons. Polar molecules add dipole–dipole attraction, and molecules with H bonded directly to N, O or F form hydrogen bonds, the strongest of the three. For similar molar masses the boiling point order is usually hydrogen bonding > dipole–dipole > dispersion only: H₂O boils at 100 °C versus −60 °C for H₂S, and HF at 20 °C versus −85 °C for HCl. When molar masses differ a lot, dispersion can dominate — CCl₄ (77 °C) boils far above CH₄ (−162 °C).

Features

Pauling electronegativity difference and average EN with a δ+/δ− bond-dipole diagram
Nonpolar / polar covalent / ionic classification under three textbook cut-offs
Percent ionic character by Pauling and Hannay–Smyth formulas
van Arkel–Ketelaar bond-type triangle with ionic, covalent and metallic regions
Molecular polarity from the 3D vector sum of bond dipoles, with a rotatable model
Measured gas-phase dipole moments for about 30 common molecules
Dispersion, dipole–dipole and hydrogen-bonding check plus a boiling-point comparison
μ = q·d, debye ↔ C·m conversion and percent ionic character from a measured dipole

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