Electronegativity Calculator

Background.

Two atoms sharing a pair of electrons rarely share it equally. The atom with the stronger pull, the higher electronegativity, draws the bonding electrons toward itself and takes on a slight negative charge, leaving its partner slightly positive. That lopsidedness is why water molecules cling to one another, why table salt dissolves in water and why oil does not. The difference in electronegativity between the two atoms, Δχ = |χA − χB|, is the first number chemists reach for to judge how lopsided a bond is.

Electronegativity cannot be measured directly. Linus Pauling built the first scale in 1932 from bond energies: a bond between two different atoms is usually stronger than the average of the matching like-atom bonds, and he traced the extra energy to the uneven sharing. On the revised Pauling scale used in most textbooks, fluorine sits at the top with 3.98 and caesium near the bottom at 0.79, with oxygen at 3.44, carbon at 2.55 and hydrogen at 2.20. Mulliken's scale and Allen's scale rest on other atomic energies and give different numbers, so both values entered here must come from the same table.

A bigger difference means a more polar bond. Chlorine bonded to chlorine shares evenly, and the 0.35 between carbon and hydrogen is small enough that hydrocarbons are barely polar. Oxygen and hydrogen, 1.24 apart, make the strongly polar O–H bond; sodium and chlorine, 2.23 apart, make a bond usually called ionic. Where polar covalent ends and ionic begins is a convention that textbooks set in different places, so the page reports the difference and leaves the label to you.

Polarity is not the same thing as reactivity, and the carbon–halogen bonds of organic chemistry show the difference. Carbon–fluorine, 1.43 apart, is the most polar of them and also the hardest to break; carbon–iodine, only 0.11 apart, is almost nonpolar, yet alkyl iodides are the ones that react fastest when the halogen is displaced, because the long, weak C–I bond and the stable iodide ion matter more than the charge separation. The difference tells you where electron density sits in a bond; what happens next depends on bond strength, molecular shape and the reaction itself.

What is electronegativity calculator?

Electronegativity is the power of an atom in a molecule to attract the shared electrons of a bond toward itself, a concept introduced by Linus Pauling. It is a relative scale rather than a measured physical quantity, and on the Pauling scale it is treated as a pure number. The electronegativity difference is the absolute gap between the values of two bonded atoms on the same scale. It describes the polarity of one bond, not of a whole molecule. On the revised Pauling scale values run from about 0.79 for caesium to 3.98 for fluorine. The reactive metals sit at the low end, the halogens, oxygen and nitrogen at the high end, and carbon, hydrogen, phosphorus and sulfur cluster in the middle, between 2.19 and 2.58, which is why bonds among those four are only weakly polar. A tabulated value is an average over many compounds, and an atom's real pull shifts with its bonding: carbon with sp hybridisation, as in acetylene, holds electrons more tightly than the sp³ carbon of ethane, and a metal in a high oxidation state pulls harder than the same metal in a low one. Gold is a reminder that the trends have exceptions; at 2.54 it is almost as electronegative as carbon, and with caesium, 1.75 apart, it forms caesium auride, a salt containing the Au⁻ ion.

How to use this calculator.

1. 01Look up both atoms on one electronegativity scale; the Pauling scale is the one most textbooks and periodic tables print, usually to two decimals (nitrogen 3.04, sulfur 2.58, chlorine 3.16). Some older books round to one decimal, which is fine as long as both values come from the same book.
2. 02Enter the two values in either order. The difference is taken as a positive number, so swapping them changes nothing.
3. 03Do not mix scales. Mulliken values quoted in electronvolts are on a different footing from Pauling values and must be converted before they can be compared.
4. 04Note which atom had the higher value before reading the result. That atom carries the partial negative charge, and the difference alone does not record it.
5. 05For a molecule, take each kind of bond in turn and then look at the shape. In chloromethane, C–H gives 0.35 and C–Cl 0.61, and because the bonds point in different directions around a tetrahedral carbon they do not cancel, so the molecule is polar. In tetrachloromethane four identical C–Cl bonds do cancel, and the molecule is not.
6. 06Two atoms of the same element, as in Cl–Cl or a carbon–carbon bond, give a difference of 0: the electrons are shared evenly, and that is the reference point every other result is measured from.

The formula.

Δχ = |χA − χB|, the absolute value of the difference between the two electronegativities. There are no constants, because all the physics sits inside the tabulated values. In Pauling's construction the difference is actually the primary quantity: the square of Δχ, in electronvolts, matches the extra energy of an A–B bond over the average of the A–A and B–B bond energies, and individual atomic values follow once the scale is anchored to a reference element. That is why a difference read from one scale should not be compared with thresholds taken from another. The result depends only on the gap, so carbon at 2.55 paired with hydrogen at 2.20 gives the same 0.35 as any other pair 0.35 apart. Hydrogen chloride shows the construction working. Approximate bond energies are 436 kJ/mol for H–H, 242 for Cl–Cl and 431 for H–Cl. The average of the first two is 339, so H–Cl is about 92 kJ/mol stronger than the average; that is about 0.954 eV, whose square root is 0.98, close to the tabulated gap of 3.16 − 2.20 = 0.96. The match is close rather than exact because the scale is a best fit across many bonds, not a value read from any one of them, and that is also why Pauling values carry the formal unit eV^½ but are used as plain numbers. The model's weakness is that it assigns one fixed value per element, so it cannot register the hybridisation and oxidation-state effects that change an atom's pull from one compound to the next.

A worked example.

Frequently asked questions.

References& sources.

1. [1]Pauling L. The Nature of the Chemical Bond. 3rd ed. (PRINT).
2. [2]IUPAC Gold Book, electronegativity.
3. [3]BIPM, The International System of Units (SI Brochure), 9th ed., version 3.01, coherent derived units and quantity equations.
4. [4]NIST Special Publication 811, Guide for the Use of the International System of Units, 2008 edition.
5. [5]NIST Chemistry WebBook.