I recently learnt that electronegativity generally decreases as I move down a group and from right to left within a period. However, according to the table below, Pb has an electronegativity of $2.33$, which is higher than $1.96$ of Sn. There are also a lot other elements, such as Au, Hg, and W, that does not follow the general rule of electronegativity.

Can someone offer an explanation of why some elements have a higher electronegativity than the elements directly above them?


The explanation is actually well-known.

  • Elements in Period 6 get f orbitals for the first time. 4f subshell has relatively poor shielding effects, thus it's harder for the atoms to lose the outer electrons, increasing electronegativity. (similar reasons also account for some chemical properties of Period 4 elements, like $\ce{H2SeO4}$'s oxidizability)

  • Relativistic effects also contribute. For a very informal (and actually incorrect) understanding, just imagine that the electrons in these heavy atoms move fast and their masses increase and their energy decreases.


Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom to attract a shared pair of electrons towards itself.

Lead has an electronic configuration of [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p² Inorder to achieve half filled electronic configuration which stabilises the lead atom, it attracts on the shared pair of electrons with greater force making it more electronegative than Tin.

  • $\begingroup$ But Sn is also one electron away from half filled 5p. How would you explain that? $\endgroup$ – Larry Jan 14 '20 at 12:08

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