Screening effect, also known as shielding effect, explains how inner shell electrons reduce the effective nuclear charge experienced by outer electrons in an atom. Due to this effect, the attraction between the nucleus and valence electrons decreases, influencing atomic size, ionization energy, and other periodic properties. The concept of effective or reduced nuclear charge is quantitatively calculated using Slater’s rules, which provide a systematic method to estimate the shielding contribution of electrons in different orbitals.
- What is The Screening Effect (Shielding Effect) ?
- Why Does the Screening Effect Occur ?
- What are the Factors Affecting the Screening Effect ?
- What is Effective Nuclear Charge ?
- What is Relation Between Actual and Effective Nuclear Charge (Screening Constant (σ))
- Slater’s Rules – A Clear Explanation
- Screening Effect Trends in the Periodic Table
- FAQs — Screening (Shielding) Effect
What is The Screening Effect (Shielding Effect) ?
In multi-electron atoms, electrons in the outermost shell (valence electrons) are not only attracted to the positively charged nucleus, but also repelled by the electrons present in the inner shells.
The net effect of Attractive force from the nucleus and Repulsive force from inner electrons causes the valence electrons to experience less nuclear attraction than they would in the absence of inner electrons. This reduction in the effective pull of the nucleus on the valence electrons is known as the Screening Effect or Shielding Effect.
Definition of shielding or screening effect
In multielectron atoms, the outermost electrons are shielded or screened from the nucleus by the inner electrons. This is known as shielding or screening effect. As a result of this, the outermost electron does not feel the full charge of the nucleus. The actual charge felt by an electron is termed as effective nuclear charge.
In Simple Words
The screening effect is like a crowd standing between you and a speaker’s voice — the more people between you and the speaker, the less clearly you can hear them.
In atoms, the inner electrons are the crowd, the nucleus is the speaker, and the valence electrons are you.
Why Does the Screening Effect Occur ?
In a polyelectronic atom (an atom with more than one electron), electrons are arranged in different shells. The inner-shell electrons lie between the nucleus and the valence electrons. Since electrons carry a negative charge, they repel each other. This repulsion pushes the outer electrons slightly away from the nucleus and reduces the effective nuclear pull acting on them.
As a result the outermost electrons do not feel the full positive charge of the nucleus and only a part of the nuclear charge is effective in attracting them.
What are the Factors Affecting the Screening Effect ?
Number of Inner Electrons : The greater the number of inner electrons, the stronger the repulsion they cause, leading to a greater screening effect.
Position of the Electron (Shell/Orbital Type) : Electrons in s-orbitals are closer to the nucleus and have lower shielding ability. Electrons in p, d, and f orbitals are farther away and have higher shielding ability.
Nature of the Atom : In heavier atoms (with more electrons), the screening effect is generally stronger due to the presence of multiple inner shells.
What is Effective Nuclear Charge ?
Due to the screening effect, the valence electron in a multi-electron atom experiences less attraction from the nucleus. This happens because the inner electrons shield or block part of the nuclear charge from reaching the valence electrons. As a result, the nuclear charge (Z) actually present on the nucleus is reduced for the valence electrons.
The reduced nuclear charge experienced by a valence electrons is called the Effective Nuclear Charge (denoted by Zeff).
What is Relation Between Actual and Effective Nuclear Charge (Screening Constant (σ))
The magnitude of the screening effect can be expressed by a screening constant (σ). Using Slater’s Rules, the effective nuclear charge (Zeff) or reduced nuclear charge felt by a valence electron can be calculated as :
Zeff = Z − σ
where :
Z = Atomic number (total or actual nuclear charge)
σ = Screening constant (total repulsion effect from other electrons)
| Key Point |
|---|
| The greater the screening constant (σ), the smaller the effective nuclear charge felt by the valence electron. |
Slater’s Rules – A Clear Explanation
A scientist named John C. Slater developed a set of rules to calculate the shielding constant (σ), which helps us find the Effective Nuclear Charge (Zeff) experienced by a valence electron.
1. First, why do we need shielding constant (σ) ?
In an atom, electrons repel each other due to their negative charges.
This repulsion reduces the full attractive pull of the nucleus on outer electrons.
The actual pull that an electron feels is called Effective Nuclear Charge (Zeff).
Formula: : Zeff = Z − σ
Where :
Z = actual nuclear charge (atomic number)
σ = shielding constant (total repulsion effect from other electrons)
2. Slater’s Golden Rules for σ
Slater divided electrons into groups and gave specific shielding values :
- Electrons in the same orbital (outermost shell)
Contribution per electron = 0.35 (except 1s, where it’s 0.30)
Do NOT count the electron for which you are calculating Zeff. i.e [0.35 × No. of nth electrons) – 1] - Electrons in (n – 1) shell (penultimate shell)
Contribution per electron = 0.85 - Electrons in (n – 2) or lower shells (inner core electrons)
Contribution per electron = 1.00
3. Example : Second Period Elements (n = 2)
For Li, Be, B, C, N, O, F, Ne → valence electrons are in n = 2 shell.
Let’s calculate step-by-step for a few cases:
Example 1 – Lithium (Li)
- Atomic number : Z = 3
- Electronic configuration : 1s2 2s1
- Electron of interest : the 2s electron.
Applying Slater’s Rules :
- Same group (2s, 2p) : No other electrons here ⇒ (1 – 1) × 0.35 = 0 × 0.35 = 0
- (n – 1) shell = 1s electrons: 2 × 0.85 = 1.70
Total shielding constant : σ = 0 + 1.70 = 1.70
Effective nuclear charge : Zeff = Z − σ = 3 − 1.70 = 1.30
The 2s electron in Li feels an effective charge of +1.30.
Example 2 – Boron (B)
- Atomic number : Z = 5
- Electronic configuration : 1s2 2s2 2p1
- Electron of interest : one 2p electron.
Applying Slater’s Rules :
- Same group (2s, 2p) : 2 other electrons in 2s +0 other electron in 2p = 2 electrons = 2 × 0.35 = 0.70
- (n – 1) shell = 1s electrons: 2 × 0.85 = 1.70
Total shielding constant : σ=0.70 + 1.70 = 2.40
Effective nuclear charge : Zeff = 5 − 2.40 = 2.60
A 2p electron in Boron feels an effective charge of +2.60.
Example 3 – Oxygen (O)
- Atomic number : Z = 8
- Electronic configuration : 1s2 2s2 2p4
- Electron of interest : one 2p electron.
Applying Slater’s Rules :
- Same group (2s, 2p) :
2 other electrons in 2s + 3 other electrons in 2p = 5 electrons
5 × 0.35 = 1.75 - (n – 1) shell = 1s electrons:
2 × 0.85 = 1.70
Total shielding constant: σ = 1.75 + 1.70 = 3.45
Effective nuclear charge: Zeff = 8 − 3.45 = 4.55
A 2p electron in Oxygen feels an effective charge of +4.55.
4. Summary Table – Second Period Elements
| Element | Z | [0.35 ×No. of nth electrons)-1] | [0.85×No. of (n–1)th electrons] | σ | Zeff |
|---|---|---|---|---|---|
| Li | 3 | 0 | 1.70 | 1.70 | 1.30 |
| Be | 4 | 0.35 | 1.70 | 2.05 | 1.95 |
| B | 5 | 0.70 | 1.70 | 2.40 | 2.60 |
| C | 6 | 1.05 | 1.70 | 2.75 | 3.25 |
| N | 7 | 1.40 | 1.70 | 3.10 | 3.90 |
| O | 8 | 1.75 | 1.70 | 3.45 | 4.55 |
| F | 9 | 2.10 | 1.70 | 3.80 | 5.20 |
| Ne | 10 | 2.45 | 1.70 | 4.15 | 5.85 |
Screening Effect Trends in the Periodic Table
For s- and p-block elements:
- Across a Period :
The screening effect increases slightly because the number of inner electrons increases as atomic number rises. - Down a Group :
The screening effect increases significantly because more inner shells are added, increasing electron repulsion.
FAQs — Screening (Shielding) Effect
What is the screening (shielding) effect?
The screening effect, also known as the shielding effect, refers to the phenomenon where inner-shell electrons reduce the effective nuclear charge experienced by valence electrons. This happens because these inner electrons repel the outer electrons, thereby diminishing the nucleus’s pull on them.
Is there a difference between the screening effect and the shielding effect?
No — both terms mean the same thing! They describe how inner electrons buffer or screen the attraction from the nucleus, weakening the force felt by the outermost electrons.
Why does the screening effect occur?
In atoms with multiple electrons, valence electrons are simultaneously attracted by the nucleus and repelled by inner electrons. This repulsion reduces the net attractive force, making the outer electrons less tightly bound.
What factors affect the magnitude of the shielding effect?
Number of inner shells : More shells between the nucleus and outer electrons lead to greater shielding.
Orbital type : Electrons in s-orbitals shield more effectively than those in p, d, or f orbitals, due to their higher penetration toward the nucleus.
How is the screening constant (σ) calculated?
The shielding constant (σ) can be estimated using Slater’s Rules, which assign specific shield values per electron based on its orbital and proximity to the valence electron. Examples include :
0.35 for electrons in the same shell (except 1s, which is 0.30)
0.85 for electrons in the penultimate shell
1.00 for electrons in deeper shells
How is the effective nuclear charge (Zeff) determined?
Once σ is known, effective nuclear charge is calculated as: Zeff = Z − σ
Where Z is the atomic number (actual nuclear charge). This describes the net attraction an outer electron experiences.
What is Screening Constant and Effective nuclear charge of Scandium ? [JEE]
Given : Sc (Z = 21), electronic configuration : $$1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^1 \, 4s^2$$
We find σ for the last 4s electron:
Same shell (n = 4): 1 other electron → 1 × 0.35 = 0.35
(n – 1) shell (n = 3): 9 electrons → 9 × 0.85 =7.65
(n – 2) and lower shells: 10 electrons → 10 × 1.00 = 10
σ = 0.35 + 7.65 + 10 = 18
Zeff = 21 − 18 = 3
Why is the concept of shielding important?
It helps explain periodic properties such as :
Atomic radius (decreases across a period due to rising Zeff)
Ionization energy (higher Zeff → stronger nuclear pull → higher ionization energy)
Chemical reactivity (metals tend to have less shielding, making valence electrons easier to remove)
Can shielding explain irregularities in periodic trends?
Yes, shielding helps clarify anomalies in elements like the Group 13 elements (B, Al, Ga…), transition series, and lanthanides & actinides where expected trends don’t always match the idealized pattern due to variable shielding.
This section provides a complete and interconnected study of Classification of Elements and Periodicity in Properties, starting with detailed theory and notes for Class 11 Chemistry to build a strong conceptual foundation. You can explore atomic radius and its types including covalent, van der Waals, metallic, and ionic radii to understand periodic trends in atomic size. The causes of periodicity explain why elements show repeating properties based on electronic configuration, which is further supported by the modern periodic law and structure of the modern periodic table including groups, periods, and blocks for elements even beyond atomic number 100. The historical development is covered through Mendeleev’s periodic law and table, leading to the modern classification of elements into s, p, d, and f blocks with prediction of period, group, and block. To strengthen exam preparation, you can practice JEE Main PYQs, IMU CET PYQs and Merchant Navy sponsorship exam MCQs, and other previous year questions with solutions, along with solved examples, conceptual questions, and practice problems on the modern periodic table. Additionally, complete study material, mock tests, and guidance are provided under Anand Classes Chemistry notes, along with expert support from Er Neeraj Anand, making this section a comprehensive resource for competitive exam preparation.