Well Wikipedia says that the atomic radius is lowered for elements in a group closer to noble gases (same row). It also says that those elements want to keep their valence electrons more than metals, for example. But the atomic shielding of those elements is bigger than the metals one because there are more electrons to repel. How is that possible?
-1
$\begingroup$
$\endgroup$
-
$\begingroup$ Some insights: crystalmaker.com/support/tutorials/atomic-radii $\endgroup$ – Mathew Mahindaratne Mar 7 at 16:54
-
$\begingroup$ More tutorials: chem.libretexts.org/Bookshelves/Inorganic_Chemistry/… $\endgroup$ – Mathew Mahindaratne Mar 7 at 16:56
2
$\begingroup$
$\endgroup$
Unfortunately, that's not how shielding works.
Shielding is the ability of electrons to shield each other from the nuclear charge.
As you move right across the periodic table, the number of electrons increases but the new electrons are all in the same energy level and thus, roughly at the same distance from the nucleus. Therefore, they cannot shield each other from nucleus as well as, say, an inner (core) electron.
You've now increased the nuclear charge by a full unit of charge, but the increase in shielding is not commensurate. That means that the effective nuclear charge ($Z_{\mathrm{eff}}$) has increased, pulling all of the valence electrons closer to the nucleus.
