You’re trying to figure out whether a bond is polar covalent or ionic, find the ΔEN between two elements, or look up a Pauling electronegativity value for an exam or lab. This calculator covers all 94 elements, classifies the bond type, calculates percentage ionic character, and tells you which direction the dipole points.
Select any element to get its Pauling electronegativity value, periodic table rank, and comparison to key reference elements.
All 94 elements with assigned Pauling valuesPlease select an element.
Select two elements to calculate ΔEN, bond type, percentage ionic character, and dipole direction.
First bonded atomPlease select Element A.
Second bonded atomPlease select Element B.
Electronegativity (Pauling)
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Nonpolar Covalent
Bond Polarity (ΔEN scale 0–3.3)
0 Nonpolar0.41.7 Ionic boundary3.3
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⚠️ Note: Pauling electronegativity values are empirical estimates that may vary slightly between sources by ±0.05. The ionic/covalent boundary at ΔEN = 1.7 is a guideline; some textbooks use 1.6 or 2.0. Percentage ionic character from the Hanney-Smyth formula is an approximation for qualitative comparison, not an exact measurable quantity.
Sources & Methodology
All Pauling electronegativity values and formulas used in this calculator are from verified authoritative chemistry sources:
Pauling, L. (1932) — Original derivation of the electronegativity concept and scale from bond dissociation energies. J. Am. Chem. Soc. 54, 3570–3582. Updated 1961 Allred revision included.
pubs.acs.org
IUPAC Recommendations — Current recommended Pauling electronegativity values for all standard elements, cross-referenced with NIST WebBook data.
iupac.org
Electronegativity Explained: Worked Examples, Trends & Bond Classification
Before defining the formula, here’s a result that makes it concrete: the H–Cl bond in hydrochloric acid. Hydrogen has EN = 2.20, chlorine has EN = 3.16. The difference is 0.96 — squarely in the polar covalent range. That 0.96 difference means chlorine is pulling the bonding electrons toward itself, creating a partial negative charge (δ−) on Cl and a partial positive charge (δ+) on H. The dipole arrow points from H toward Cl. That’s the entire concept of electronegativity applied in three numbers.
Linus Pauling defined electronegativity in 1932 as the tendency of an atom to attract shared electrons in a chemical bond. He noticed that when two different atoms bond, the bond releases more energy than you’d predict from a purely covalent model. That “extra energy” comes from the electron density shifting toward the more electronegative atom, creating partial ionic character. The larger the shift, the more ionic the bond.
The Pauling Electronegativity Formula Explained
Pauling didn’t measure EN directly — he derived it from bond energies. The electronegativity difference between two elements A and B is: ΔEN² = E(A–B) − ½[E(A–A) + E(B–B)], where E() represents bond dissociation energies in eV. He then assigned F = 3.98 as the reference (highest possible EN) and calculated all other values relative to that. The Allred (1961) revision updated several values using better experimental data.
Electronegativity of the Most Common Bonds — Reference Table
Most students searching this topic need to classify specific bonds. The table below covers the 12 most frequently asked-about element pairs in general chemistry — with exact ΔEN values, bond type, and percentage ionic character so you don’t have to look them up separately.
Bond
EN (A)
EN (B)
ΔEN
Bond Type
% Ionic
Dipole Direction
H–H
2.20
2.20
0.00
Nonpolar cov.
0%
None
C–H
2.55
2.20
0.35
Nonpolar cov.
4.3%
C ← H (very slight)
H–N
2.20
3.04
0.84
Polar cov.
16.8%
H → N
H–Cl
2.20
3.16
0.96
Polar cov.
18.6%
H → Cl
H–O
2.20
3.44
1.24
Polar cov.
26.3%
H → O (water)
H–F
2.20
3.98
1.78
Polar cov.*
44.7%
H → F (very polar)
C–O
2.55
3.44
0.89
Polar cov.
17.5%
C → O (carbonyl)
C–Cl
2.55
3.16
0.61
Polar cov.
10.9%
C → Cl
Li–F
0.98
3.98
3.00
Ionic
85.5%
Li → F
Na–Cl
0.93
3.16
2.23
Ionic
53.1%
Na → Cl (table salt)
K–Br
0.82
2.96
2.14
Ionic
50.0%
K → Br
Cl–Cl
3.16
3.16
0.00
Nonpolar cov.
0%
None
*H–F has ΔEN = 1.78, which is just over the 1.7 ionic threshold, yet it’s universally classified as polar covalent in general chemistry because HF is a gas at room temperature, not a salt. This is the most common student confusion with EN classification — the 1.7 boundary is a guideline, and bond type is better confirmed by physical properties (melting point, conductivity, crystal structure).
How Electronegativity Changes Across the Periodic Table
The trend is consistent and explainable by atomic structure. Moving left to right across a period, nuclear charge increases (more protons) while atomic radius decreases — electrons in the bonding shell are closer to a stronger nucleus. That combination pulls bonding electrons more strongly, so EN increases. Moving down a group, atomic radius increases as new electron shells are added, putting bonding electrons farther from the nucleus and reducing the attractive force. The top-right corner of the periodic table (excluding noble gases) has the highest EN values. The bottom-left has the lowest.
Which Bond Classification System Should You Use?
If you’ve taken general chemistry at two different institutions, you may have been taught different ΔEN thresholds. Some textbooks say ionic starts at 1.7. Others say 2.0. A few say 1.6 for polar covalent. None of them is wrong — they’re different conventions applied to a continuum that doesn’t have a sharp boundary. Here’s what actually matters for different use cases:
For General Chemistry (AP, Introductory University)
ΔEN 0.1–0.4: very slightly polar (often treated as nonpolar in practice)
ΔEN 0.5–1.7: polar covalent (this is where most real chemistry happens)
ΔEN > 1.7: ionic character dominates (salts, metal halides)
For Organic Chemistry
In organic chemistry, the key EN values to memorize are just four: C (2.55), H (2.20), O (3.44), N (3.04). Nearly every organic reaction depends on understanding where electrons are more dense. The C–O bond (ΔEN = 0.89) makes carbonyl carbon electrophilic. The N–H bond (ΔEN = 0.84) explains why amines are nucleophiles at nitrogen. The O–H bond (ΔEN = 1.24) explains the acidity of alcohols and carboxylic acids. EN differences drive organic mechanism arrows.
What Electronegativity Does Not Tell You
EN alone doesn’t determine molecular polarity. A molecule can have polar bonds and still be nonpolar overall if the bond dipoles cancel symmetrically. CO&sub2; has two very polar C=O bonds (ΔEN = 0.89 per bond) that point in exactly opposite directions, cancelling to give a nonpolar molecule. Methane has four slightly polar C–H bonds arranged tetrahedrally — they cancel perfectly. EN tells you about individual bonds; molecular geometry determines whether those bond dipoles add up or cancel out.
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What most students get wrong: They use EN to predict molecular polarity without considering geometry. BF&sub3; has three very polar B–F bonds (ΔEN = 2.0 each) but is nonpolar because of its trigonal planar shape — all three bond dipoles point outward at 120° and cancel exactly. Always combine EN analysis with VSEPR geometry before calling a molecule polar or nonpolar.
❓ Frequently Asked Questions
Electronegativity is the tendency of an atom to attract shared bonding electrons toward itself. Pauling scale runs 0.7 (francium, lowest) to 3.98 (fluorine, highest). It increases across a period and decreases down a group. The concept explains why electrons in H-Cl are closer to Cl (EN=3.16) than to H (EN=2.20) - the 0.96 difference creates partial charges.
delta-EN = |EN1 - EN2| - absolute value of the difference. For H-O in water: |2.20 - 3.44| = 1.24 (polar covalent). For Na-Cl: |0.93 - 3.16| = 2.23 (ionic). The absolute value is used because polarity depends only on the magnitude of the difference, not which atom is larger. Use the Two-Element Bond Analysis tab above for any pair.
delta-EN less than 0.4: nonpolar covalent (equal sharing - H2, Cl2, C-H). delta-EN 0.4 to 1.7: polar covalent (unequal sharing - H-Cl, H-O, H-N, C-O). delta-EN above 1.7: predominantly ionic (electron transfer - NaCl, KBr, NaF). The 1.7 threshold is a general guideline; some textbooks use 1.6 or 2.0. Bond type is ultimately confirmed by physical properties like melting point and conductivity.
% ionic = 16 x |delta-EN| + 3.5 x (delta-EN)^2 (Hanney-Smyth formula). HCl with delta-EN = 0.96: % ionic = 15.36 + 3.23 = 18.6%. NaCl with delta-EN = 2.23: % ionic = 35.7 + 17.4 = 53.1%. This is an empirical estimate - pure ionic bonds don't exist, and even "ionic" NaCl has nearly 47% covalent character. Use the Bond Analysis tab for automatic calculation.
Fluorine (F) = 3.98 on the Pauling scale. Small radius + 9 protons + one electron short of a full shell = extreme electron attraction. Ranking: F (3.98) > O (3.44) > Cl (3.16) > N (3.04) > Br (2.96) > I (2.66) > S (2.58) > C (2.55) > H (2.20). Noble gases have no assigned value since they form no typical chemical bonds.
H = 2.20 (Pauling). Close to carbon (2.55), making C-H bonds nearly nonpolar (delta-EN = 0.35). Hydrogen's intermediate EN explains its dual chemistry: it acts as H+ (acid) in bonds with electronegative elements like O (delta-EN = 1.24) and as H- (hydride) in bonds with very low EN metals. In water, the O-H bond polarity (delta-EN = 1.24) is responsible for hydrogen bonding and water's unusual properties.
Pauling (1932): from bond dissociation energies, dimensionless 0.7-3.98. Most widely taught. Mulliken (1934): EN = (IE + EA) / 2 in eV, more physically grounded but requires measured ionization energy and electron affinity. Allred-Rochow (1958): based on effective nuclear charge and covalent radius. All three correlate well and show the same periodic trends. The Pauling scale is used in virtually all general chemistry courses.
Increases left-to-right across a period (nuclear charge rises, atomic radius falls, bonding electrons pulled more strongly). Decreases top-to-bottom in a group (atomic radius grows, electrons farther from nucleus, more shielding). Top-right = highest EN: F (3.98), O (3.44), N (3.04), Cl (3.16). Bottom-left = lowest EN: Fr (0.70), Cs (0.79), Rb (0.82), K (0.82).
A dipole forms when electrons are distributed unequally in a bond. The more electronegative atom gets delta-minus (partial negative), the less electronegative gets delta-plus (partial positive). The dipole arrow points from positive toward negative - toward the more electronegative atom. H-Cl: arrow points H to Cl. Water's two O-H dipoles (pointing toward O) add to give a net dipole of 1.85 Debye - making water polar and able to dissolve ionic compounds.
Because ionic vs. covalent is a continuum, not a sharp switch. The 1.7 Pauling threshold corresponds to roughly 50% ionic character. Some textbooks set the threshold at 1.6 or 2.0 instead. HF has delta-EN = 1.78 yet is a gas (polar covalent behavior), while NaBr at delta-EN = 1.9 forms a salt crystal (ionic behavior). The rule: if a metal and nonmetal are bonded and delta-EN is above about 1.6-1.7, call it ionic. For nonmetal-nonmetal bonds, use 2.0 as the threshold for ionic. When in doubt, check physical properties.
Oxidation state rules assume all bonding electrons belong to the more electronegative atom. In H2O: O (EN=3.44) takes both bonding electrons from each H (EN=2.20), giving O the oxidation state -2 and each H +1. In CO2: O (3.44) is assigned both electrons from each C-O bond, giving C oxidation state +4. This is why "the more electronegative element gets the negative oxidation state" is a universal rule - it's just electronegativity applied formally.
Yes, absolutely - and this is one of the most important concepts in molecular polarity. CO2 has two very polar C=O bonds (delta-EN = 0.89 each), but the linear geometry means both bond dipoles point in exactly opposite directions and cancel perfectly. The result: CO2 is nonpolar. Similarly, methane (CH4) has four slightly polar C-H bonds arranged tetrahedrally - they cancel. BF3 is trigonal planar - three polar B-F bonds cancel. Always combine EN analysis with VSEPR geometry.