Q.What is effective nuclear charge?

Inner-shell electrons partly cancel the nucleus's pull on a valence electron by repelling it -- this is the shielding (screening) effect -- so the net attractive charge a valence electron actually feels, the effective nuclear charge, is less than the full nuclear charge:

where $S$, the screening constant, is estimated by Slater's rules:

1. Group the electron configuration as $(1s)(2s,2p)(3s,3p)(3d)(4s,4p)(4d)(4f)(5s,5p)\ldots$
2. Electrons in a group to the right of the electron of interest contribute nothing; electrons in the same group contribute 0.35 each (0.30 if the electron of interest is 1s).
3. For an s or p electron of interest: electrons one group inward $(n-1)$ contribute 0.85 each; electrons two or more groups inward contribute the full 1.00 each. For a d or f electron of interest: every electron to its left contributes the full 1.00, with no intermediate 0.85 tier.
4. Sum all contributions to get $S$.

Worked on scandium ($Z=21$, config $(1s)^2(2s,2p)^8(3s,3p)^8(3d)^1(4s)^2$): the 4s electron feels $S = (1)(0.35) + (9)(0.85) + (10)(1.00) = 18.00$, giving $Z_{eff} = 3$; the 3d electron feels $S = (0)(0.35) + (18)(1.00) = 18$ (the two 4s electrons sit to the right of 3d in the Slater ordering and so do not shield it at all), also giving $Z_{eff} = 3$.

Slater's rules are the mechanism behind both major radius/energy trends: effective nuclear charge rises across a period (driving atomic radius down and ionisation energy up) and the growing number of inner shells down a group increases total shielding (driving atomic radius up and ionisation energy down). The same Slater $Z_{eff}$ calculation, applied to an ion rather than a neutral atom, is also the basis of Pauling's method for finding ionic radii.