Final answer:
Core electrons efficiently shield outer electrons from the nucleus, reducing the nucleus's effective charge on these electrons.
Step-by-step explanation:
In the context of atomic structure, shielding is where inner-shell electrons reduce the attractive force exerted by the nucleus on the outer-shell or valence electrons. The core electrons are much more effective at this shielding process because they lie closer to the nucleus and therefore can intercept the nuclear charge more effectively. As we move across a period on the periodic table, the number of protons in the nucleus (Z) increases, enhancing the nuclear attraction.
However, because additional valence electrons do not efficiently shield each other from this attraction, the effective nuclear charge (Zeff) experienced by electrons increases, pulling them closer to the nucleus and reducing atomic radius.
For example, in a lithium atom, the two 1s electrons are very good at shielding the valence electron in the 2s orbit from the nuclear charge, due to their proximity to the nucleus.
Conversely, when we consider a sodium atom, where the 3s valence electron is being shielded by the 1s and 2s electrons, the increased positive charge in the nucleus attracts the 3s valence electron more strongly because there is only minor additional shielding from the 2p electrons, leading to an increased Zeff.