In hydrogen-like atoms (those with only one electron), the net force on the electron is just as large as the electric attraction from the nucleus. However, when more electrons are involved, each electron (in the n-shell) feels not only the electromagnetic attraction from the positive nucleus, but also repulsion forces from other electrons in shells from 1 to n. This causes the net force on electrons in outer shells to be significantly smaller in magnitude; therefore, these electrons are not as strongly bonded to the nucleus as electrons closer to the nucleus. This phenomenon is often referred to as the Orbital Penetration Effect. The shielding theory also explains why valence-shell electrons are more easily removed from the atom.
The size of the shielding effect is difficult to calculate precisely due to effects from quantum mechanics. As an approximation, we can estimate the effective nuclear charge on each electron by the following:
Where Z is the number of protons in the nucleus and is the average number of electrons between the nucleus and the electron in question. can be found by using quantum chemistry and the Schrödinger equation, or by using Slater's empirical formulas.
In Rutherford backscattering spectroscopy the correction due to electron screening modifies the Coulomb repulsion between the incident ion and the target nucleus at large distances.
- L. Brown, Theodore; H. Eugene LeMay, Jr., Bruce E. Bursten, Julia R. Burdge (2003). Chemistry: The Central Science (8th Edition ed.). US: Pearson Education. ISBN 0-13-061142-5.
- Dan Thomas, Shielding in Atoms, 
- Peter Atkins & Loretta Jones, Chemical principles: the quest for insight [Variation in shielding effect]
|This physical chemistry-related article is a stub. You can help Wikipedia by expanding it.|
|This atomic, molecular, and optical physics-related article is a stub. You can help Wikipedia by expanding it.|