The reducing effect of the nuclear charge by the inner electrons for an outer electron is    termed as shielding (or screening). As a result of shielding, the outer electrons in an atom always experience less nuclear charge than the actual nuclear charge Z. the effective nuclear charge $\left(Z^{*}\right)$as  experienced by an electron is then obtained by subtracting the total shielding contributions from all the other electrons (i.e., except the  one under consideration) from the actual nuclear charge. $Z^{*}=Z-s$

Where s = sum of the shielding contributions. The rules for estimating contributions to s are as follows (Slater’s rule) Contribution to shielding by each electron is  :

Electron group	All higher group	Same group	Group n-1	Group  $\le n-2$
Is	0	0.30	-	-
(ns, np)	0	0.35	0.85	1.00
(nd)  or  (nf)	0	0.35	1.00	1.00

 

According to Slater’s treatment, the energy of an electron in $n^{th}$shell of an atom having    atomic number Z is given by the empirical equation.$E=-13.6{\left(\frac{Z^{*}}{n}\right)}^{2}eV;Z^{*}=$ effective nuclear charge.