effective nuclear charge
/ Z-eff /
Imagine standing in a crowd, trying to feel the warmth of a bonfire on the far side. The people between you and the fire block part of the heat, so what reaches you is weaker than the fire's full blaze. An outer electron in an atom is in much the same situation: the inner electrons stand between it and the nucleus, soaking up part of the nuclear pull, so it feels less than the full charge.
The effective nuclear charge, written Z_eff, is the net positive pull that a given electron actually experiences from the nucleus, after the repulsion from other electrons (especially inner ones) is subtracted. Roughly, Z_eff equals the true nuclear charge Z minus a shielding constant: the more inner electrons stand in the way, the smaller the leftover pull on the outer electron.
This idea quietly explains a great deal: why outer electrons are held loosely, why atoms shrink across a period, and why ionisation gets harder. The caveat is that Z_eff is an approximation that compresses the messy many-electron problem into one number; estimates from simple shielding rules (like Slater's) are handy guides rather than exact values.
Sodium's nucleus carries +11, but its lone outer 3s electron sits behind ten inner electrons that screen most of that pull, so it feels an effective charge of only about +2.5. That weak grip is exactly why sodium gives up that electron so easily and is such a reactive metal.
Inner electrons screen the nucleus, so the outer ones feel a reduced pull.
Across a period Z_eff rises (protons added, shielding barely changes), which is the engine behind shrinking radii and rising ionisation energy; down a group the extra full shells keep outer Z_eff modest.