Shielding Effect
Shielding Effect
In a multielectron atom, the electrons of the valency shell (outermost shell) are attracted towards the nucleus and also these electrons are repelled by the electrons present in the inner shells. On account of this, the actual force of attraction between the nucleus and the valency electrons is somewhat decreased by the repulsive forces acting in opposite direction. This decrease in the force of attraction exerted by the nucleus on the valency electrons due to the presence of electrons in the inner shells, is called screening effect or shielding effect. The magnitude of the screening effect depends upon the number of inner electrons i.e. higher the number of inner electrons, greater shall be the value of screening effect. The screening effect constant is represented by the symbol s. The magnitude of ‘s’ is determined by the Slater’s rules. The contribution of inner electrons to the magnitude of ‘s’ is calculated in the following ways.
For ns or np orbital Electrons
(i) Write the electronic configuration of the element in the following order and group them as,
(1s), (2s 2p), (3s 3p), (3d), (4s 4p), (4d 4f), (5s 5p), (5d 5f), (6s 6p), etc.
(ii) Electrons to the right of the (ns, np) group are not effective in shielding the ns or np
electrons and contribute nothing to s.
(iii) All other electrons in the (ns, np) group contribute to the extent of 0.35 each to the
Screening constant (except for 1s for which the value is 0.30)
(iv) All the electrons in the (n-1)th shell contribute 0.85 each to the screening constant.
(v) All the electrons in the (n-2)thshell or lower contribute 1.0 each to the screening constant.
For d- or f- electron,
rules (i) to (iii) remain the same but rules (iv) and (v) get replaced by the rule (vi).(vi)
All the electrons in the groups lying left to (nd, nf) group contribute 1.0 each to the screening effect.
|
per electron of the orbit |
|||
|
Electron in orbitals |
n |
(n – 1) |
(n – 2) or (n – 3), etc |
|
(Shell) |
|
|
|
|
S or P orbital |
0.35 |
0.85 |
1.00 |
|
d or f orbital |
0.35 |
1.00 |
1.00 |
For is electron for a He like atom which has 2 electrons
Zeff = Z – 0.3 = 1.7
For hydrogen atom, Zeff = z
As we move left to right in a period table the value of Zeff increases by 0.65.
Effective Nuclear Charge
Due to screening effect the valency electron experiences less attraction towards nucleus. This brings decrease in the nuclear charge (Z) actually present on the nucleus. The reduced nuclear charge is termed effective nuclear charge and is represented by Z*. It is related to actual nuclear charge (Z) by the following formula:
Z* = (Z - s), where s is screening constant
It is observed that magnitude of effective nuclear charge increases in a period when we move from left to right.
2nd Period Li Be B C N O F Ne
Z 3 4 5 6 7 8 9 10
s 1.7 2.05 2.42 2.75 3.1 3.45 3.8 4.15
Z* = Z - s 1.3 1.95 2.6 3.25 3.9 4.55 5.2 5.85
In a subgroup of normal elements the magnitude of effective nuclear charge remains almost the same.
Alkali group Li Na K Rb Cs
Z 3 11 19 37 55
s 1.7 8.8 16.8 34.8 52.8
Z* = Z - s 1.3 2.2 2.2 2.2 2.2
