What is Shielding Effect
In a multi-electron atom, the electrons of the valence shell (outermost shell) are attracted toward 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 the opposite direction. This decrease in the force of attraction exerted by the nucleus on the valence electrons due to the presence of electrons in the inner shells is called the screening effect or shielding effect.
What are Slater’s Rules?
The magnitude of the screening effect depends upon the number of inner electrons i.e., the higher the number of inner electrons, the greater shall be the value of the screening effect. The screening effect constant is represented by the symbol σ. The magnitude of ‘σ’ is determined by Slater’s rules. The contribution of inner electrons to the magnitude of ‘σ’ is calculated in the following ways.
Rules to find Shielding Effect by Slater’s rules
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 σ.
(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) th shell 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.
| s per electron of the orbit | |||
|---|---|---|---|
| Electron in orbitals ø | n | (n – 1) | (n – 2) or (n – 3), etc |
| (Shell)ΔE | |||
| 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 a 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 the screening effect, the valence electron experiences less attraction towards the nucleus. This brings a decrease in the nuclear charge (Z) actually present on the nucleus. The reduced nuclear charge is termed the effective nuclear charge and is represented by Z*. It is related to the actual nuclear charge (Z) by the following formula:
Z* = (Z - σ), where s is the screening constant
It is observed that the magnitude of the 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 |
| σ | 1.7 | 2.05 | 2.42 | 2.75 | 3.1 | 3.45 | 3.8 | 4.15 |
| Z* = Z - σ | 1.3 | 1.95 | 3.25 | 3.9 | 4.55 | 5.2 | 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 |
| σ | 1.7 | 8.8 | 16.8 | 34.8 | 52.8 |
| Z* = Z - σ | 1.3 | 2.2 | 2.2 | 2.2 | 2.2 |
What are Slater’s Rules?
Slater’s Rules are a set of guidelines used to calculate the shielding constant and effective nuclear charge in multi-electron atoms. These rules help estimate how strongly the nucleus attracts electrons.
According to Slater’s rules:
- Electrons in the same shell partially shield each other
- Inner shell electrons provide stronger shielding
- Outer electrons provide almost no shielding to inner electrons
These rules are commonly used in atomic structure and quantum chemistry.
Importance of Shielding Effect
The shielding effect influences many atomic properties, such as:
- Atomic size
- Ionization energy
- Electron affinity
- Periodic table trends
As shielding increases, the attraction between the nucleus and outer electrons decreases.
Applications of Slater’s Rules
Slater’s rules are used to:
- Calculate effective nuclear charge
- Understand electron configuration
- Explain periodic trends
- Study chemical bonding and atomic behavior
Conclusion
The shielding effect explains how inner electrons reduce nuclear attraction on outer electrons, while Slater’s Rules provide a method to calculate this effect. These concepts are essential in chemistry for understanding atomic structure, periodic trends, and electron behavior.
FAQs on Shielding Effect and Slater’s Rules
The shielding effect is the reduction in the attraction between the nucleus and outer electrons due to the presence of inner electrons in an atom. Inner shell electrons act like a shield and block some of the positive nuclear charge from reaching the outermost electrons.
The effective nuclear charge is represented by:
Z_{eff}=Z-S
Where:
- (Z_{eff}) = effective nuclear charge
- (Z) = actual nuclear charge
- (S) = shielding constant
Because of the shielding effect, outer electrons are held less tightly by the nucleus. This concept is important in understanding atomic size, ionisation energy, and periodic table trends in chemistry.
The shielding effect increases down a group because the number of electron shells increases as we move from top to bottom in the periodic table. Each new shell adds more inner electrons between the nucleus and the outermost electrons.
These inner electrons block part of the nuclear attraction, reducing the force experienced by valence electrons. As a result, the outer electrons are held less tightly by the nucleus.
The effective nuclear charge is given by:
Z_{eff}=Z-S
As the shielding constant (S) increases, the effective nuclear charge on outer electrons decreases. This is why atomic size generally increases down the group in the periodic table.
The s-orbital has the highest shielding effect because it is closest to the nucleus.
The order of shielding effect is s > p > d > f, with s having the strongest effect.
The s orbital is closer to the nucleus and can block the nuclear pull more effectively.