If there are more electrons closer to the nucleus, the outermost electron is protected from the strong nuclear attraction and can freely orbit the atom. The term “shielding” or “screening” describes this phenomenon in simple words.

Shielding effect

The shielding effect is similar to that of a protective coating that leads to a decrease in the nuclear charge felt by valence electrons.
Atoms with less shield have their electrons near the nucleus and therefore require more energy to remove the outer electron e.g. helium whereas atoms Cesium and Francium have less attraction between the valence shell electron and the nucleus, so they require less ionization energy to remove the outer electron.
The more electron shells are present, the greater the shielding effect experienced by the outermost electrons.
The shielding effect explains why the electrons of valence are so far away from the atom. The nucleus can pull the shell of valence firmly when the attraction is strong and small when firmness is weakened. In the event of further shielding, the valence shell may continue to open. As a result, atoms will be larger.

Example:

Why is cesium atomic size greater than elemental sodium?

Reason:
The sodium element has an electron 1s² 2s² 2p⁶ 3s¹ configuration. The outer shell is n = 3 and there is one electron. The attraction between the single electron and the nucleus having 11 protons is shielded by the other inner shell of 10 core electrons.
The electron configuration for cesium is 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶6s¹. While there are more protons in a cesium atom, there are also many more electrons that shield the outer electron from the nucleus. The outer electron, 6s1, therefore, is held very freely, resulting in a greater atomic size of cesium.