Effective nuclear charge, Zeff, measures how strongly the nucleus pulls on a specific electron, accounting for any electron–electron repulsions. For most atoms, the inner electrons partially shield/block the outer electrons from the pull of the nucleus, and thus:
Zeff = Z−shielding (blocking positive charge by other electrons)
Shielding is determined by the probability of another electron being between the electron of interest and the nucleus, blocking or shielding it from the full nuclear pull. Core electrons are good at shielding, while electrons in the same valence shell do not block (shield electrons from) the nuclear attraction experienced by each other as efficiently (have very minimal effect).
For hydrogen, there is only one electron and so the nuclear charge (Z) and the effective nuclear charge (Zeff) are equal.
However, for all other elements, each time we move from one element to the next across a period, Z increases by one, but the shielding increases only slightly.
Thus, as expect, the outermost or valence electrons are easiest to remove. They possess the highest energies, experience greater shielding, and are farthest from the nucleus.
Zeff then increases as we move from left to right across a period, however, very minimal if any change in Zeff occurs as we go down a group.