Shielding effect

Description

In the hydrogen atom (which has only a single electron and therefore no inner-core electrons), the electron experiences the full electromagnetic attraction of the nucleus. However, in multi-electron atoms, each outer-shell electron experiences both the attraction from the positive nucleus and the electrostatic repulsion from electrons in lower shells.

This causes the net force on valence electrons to be significantly smaller in magnitude than the actual nuclear charge. Consequently, these outer electrons are less tightly bound to the nucleus, which explains why they are more easily removed during ionization or chemical bonding.

The magnitude of the shielding effect is difficult to calculate precisely due to quantum mechanical effects. As an approximation, we can estimate the effective nuclear charge (${\displaystyle Z_{\mathrm{e}\mathrm{f}\mathrm{f}}}$) on each electron using the following formula:

${\displaystyle Z_{\mathrm{e}\mathrm{f}\mathrm{f}}=Z-\sigma }$

Where:

• Z is the atomic number (number of protons in the nucleus);
• ${\displaystyle \sigma }$ is the screening constant, which can be determined using quantum chemistry methods (solving the Schrödinger equation) or estimated empirically via Slater's rules.

In Rutherford backscattering spectroscopy, the correction due to electron screening also modifies the Coulomb repulsion between the incident ion and the target nucleus at large distances, accounting for the screening effect caused by inner electrons.