Periodic table with electronegativity values

What electronegativity means

When two atoms share a pair of electrons, the pair is not always shared equally. The more electronegative atom pulls the electrons closer and gains a slight negative charge, leaving the other atom slightly positive. Electronegativity cannot be measured directly on a single atom; it is a relative scale worked out from the energies of chemical bonds.

The values here use the Pauling scale, introduced by Linus Pauling in 1932, which has no units. Most metals have values below 2, the metalloids sit close to 2 and the nonmetals are above 2, with hydrogen at 2.2. Each cell shows the value under the element's name, and the legend under the table shows the color scale from low to high.

The electronegativity trend

Electronegativity increases from left to right across a period. Each step adds a proton to the nucleus, while the new electrons join the same outer shell and shield one another poorly, so the nucleus pulls harder on bonding electrons. In period 2 the values climb steadily from lithium to fluorine.

Electronegativity decreases down a group. Each new period adds an electron shell, so bonding electrons sit farther from the nucleus and are better shielded from it. Among the halogens, the values fall from fluorine through chlorine and bromine to iodine.

Exceptions and missing values

Helium, neon and argon have no Pauling value because they form no stable compounds, so their cells show a dash. Krypton, xenon and radon do have values, because they can form compounds, mainly with fluorine. Values for the heavier actinides are rough estimates, and none are known for the superheavy elements.

A few heavier elements break the smooth trend. Gallium is slightly more electronegative than aluminum above it, because the ten d electrons added in between shield the nucleus poorly. Gold is unusually electronegative for a metal, enough to form the auride ion, Au⁻, in compounds such as cesium auride (CsAu).

Using electronegativity to predict bonds

The difference in electronegativity between two atoms tells you what kind of bond they form. As a rule of thumb, a difference below about 0.4 gives a nonpolar covalent bond, a difference between about 0.4 and 1.7 gives a polar covalent bond, and a larger difference gives an ionic bond. Textbooks draw these cutoffs slightly differently, so treat them as a guide.

Example: in water, oxygen has an electronegativity of 3.44 and hydrogen 2.2. The difference is 3.44 − 2.2 = 1.24, so each O–H bond is polar covalent, with the oxygen end slightly negative. In sodium chloride the difference is 3.16 − 0.93 = 2.23, large enough for chlorine to take sodium's outer electron completely and form an ionic bond.