# Periodic Trends of Zeff and Electronegativity

The effective nuclear charge $Z_\mathrm{eff}$ increases from left to right and from top to bottom. Can you explain why it increases from top to bottom?

Also can we explain the periodic trend of electronegativity using the shielding effect?

EDIT: The calculation of $Z_\mathrm{eff}$ for Li would be 1 (3 protons − 2 core electrons). For Na is 1 (11 protons − 10 core electrons). So far it doesn't seem to me that $Z_\mathrm{eff}$ should increase from top to bottom. Considering the fact though that for Na there are more electrons shielding the only valence electron than Li as well as Na's electron is relatively more distant than that of Li, the first should be experiencing less $Z_\mathrm{eff}$. This is not true however. As far electronegativity is concerned I can't come up with a neat explanation.

• Looks like homework to me. Can you share what work you've done while working towards an answer? – Todd Minehardt Aug 7 '15 at 21:01
• Your misconception: Shielding by core electrons is perfect. Truth: shielding for Li is less than 2, shielding for Na is less than 10. – orthocresol Aug 8 '15 at 18:05

The Slater shielding rules give more quantitative information about the amount of shielding.

Put simply, valence electrons (as you go across a row) do not shield very well, roughly 35% of a full electron. So when you go from B to C to N, you keep increasing the nuclear charge by one proton, but the electrons don't fully shield the nucleus.

Core electrons, as suggested in comments and other answers do shield fully, since they're closer to the nucleus.

So yes, the shielding for Li and Na are similar. As you mention, the valence electron in Na is further from the nucleus, and that does effect the ionization energy and electronegativity as you suggest.

The electronegativity is the greatest at the top of the periodic table because fewer electrons are shielding the outermost electrons from the attraction of the nucleus. As more electrons are added the electrons closer to the nucleus repel some of the outermost electrons and block the nucleus's attraction. Within one left to right row the electronegativity increases as you move right because more protons exert an attractive pull on the electrons without adding more electrons that shield the charge of the nucleus.