4.3.1 Shielding Effect

This topic explains the shielding effect — how inner electrons reduce the attraction between the nucleus and the valence electrons.

Definition

Shielding effect is the reduction in force of attraction between the nucleus and the valence electrons by the electrons present in the inner sub-shells.

Simple explanation: Inner electrons act like a shield, blocking some of the nucleus's positive charge from reaching the outer electrons.

Main Idea

Electrons present in the inner shells cut off attractive force between the nucleus and the valence electrons.

• The more inner shells there are, the greater the shielding effect.
• As you move down a group, the number of electronic shells increases, so the shielding effect increases.
• As you move across a period, the number of inner-shell electrons stays constant, so the shielding effect remains constant.

Simple explanation: Inner electrons "shield" outer electrons from the full pull of the nucleus.

Core Concept

How the Shielding Effect Works

Trend in the Periodic Table

Simple explanation: Going down a group adds more shells, so shielding increases. Going across a period adds electrons to the same shell, so shielding stays the same.

Detailed Explanation

Why the Shielding Effect Happens

The nucleus has a positive charge and attracts the negatively charged electrons. However:

• The inner electrons are between the nucleus and the outer (valence) electrons.
• These inner electrons also attract the nucleus and partly cancel out its pull on the outer electrons.
• So the valence electrons feel a weaker attraction than they would if the inner electrons weren't there.

This reduction in attraction is the shielding effect.

Simple explanation: The inner electrons act as a "buffer" — the nucleus's pull is partly used up on them, so less reaches the outer electrons.

Shielding Effect Down a Group

As you move from top to bottom in a group:

• The number of electronic shells increases.
• So the number of inner-shell electrons also increases.
• As a result, the shielding effect increases.

• Be (beryllium): 1s² 2s² → only 2 inner electrons (in 1s).
• Mg (magnesium): 1s² 2s² 2p⁶ 3s² → 10 inner electrons (in 1s, 2s, 2p).

Mg has more inner shells than Be, so Mg has a greater shielding effect than Be.

Simple explanation: Mg's valence electrons are shielded by 10 electrons; Be's are shielded by only 2. So Mg's valence electrons feel a weaker pull from the nucleus.

Shielding Effect Across a Period

As you move from left to right in a period:

• The number of electrons in the inner shells remains constant.
• Therefore, the shielding effect remains constant.

• All atoms have only the 1s² inner shell (2 inner electrons).
• The shielding effect stays the same across the whole period.

Simple explanation: Across a period, you're adding electrons to the same outer shell, so the inner shielding stays the same.

Book Example: Identifying the Element with Greater Shielding Effect

Question: Choose the elements whose atoms you expect to have greater shielding effect. (a) Be or Mg (b) C or Si

Problem Solving Strategy: Look at the periodic table and find the relative position of given elements. Apply the trend of increasing shielding effect in a group.

Solution: (a) Mg atoms will have greater shielding effect (Mg is below Be in Group 2). (b) Si atoms will have greater shielding effect (Si is below C in Group 14).

Concept Assessment Exercise 4.5

Choose the element whose atoms you expect to have smaller shielding effect. (a) F or Cl (b) Li or Na (c) B or Al

Answers: (a) F (F is above Cl in Group 17) (b) Li (Li is above Na in Group 1) (c) B (B is above Al in Group 13)

Intermediate Level

Why Shielding Increases Down a Group

• New shells are added.
• Each new shell adds more inner electrons.
• These inner electrons shield the valence electrons more effectively.
• So the valence electrons feel a weaker attraction to the nucleus.

Consequence: It becomes easier to remove the valence electron → ionization energy decreases down a group.

Why Shielding Stays Constant Across a Period

• Electrons are added to the same outer shell.
• The inner shells do not change.
• So the number of inner electrons stays the same.
• Therefore, the shielding effect remains constant.

• Atomic size to decrease
• Ionization energy to increase
• Electronegativity to increase

Effective Nuclear Charge

The effective nuclear charge (Z_eff) is the net positive charge experienced by a valence electron:

Z_eff = Z - S

• Z = atomic number (number of protons)
• S = shielding constant (shielding effect)

Across a period: Z increases, S stays constant → Z_eff increases.

Down a group: Z increases, S also increases → Z_eff increases only slightly.

Important Points

• - Shielding effect: reduction in force of attraction between nucleus and valence electrons by inner-shell electrons.
• - Down a group: shielding effect increases (more inner shells and electrons).
• - Across a period: shielding effect remains constant (inner shells stay the same).
• - Examples:
• - Mg has greater shielding effect than Be.
• - Si has greater shielding effect than C.
• - Na has greater shielding effect than Li.
• - Cl has greater shielding effect than F.
• - Shielding effect explains why ionization energy decreases down a group.

Remember

• - Shielding effect = inner electrons reduce nucleus's pull on valence electrons.
• - Down a group → shielding increases.
• - Across a period → shielding remains constant.
• - More inner shells → greater shielding.
• - Mg > Be in shielding; Si > C; Na > Li; Cl > F.
• - Shielding affects ionization energy, atomic size, electronegativity.

Common Mistakes

• 1. Thinking shielding increases across a period. It stays constant across a period.
• 2. Confusing shielding effect with effective nuclear charge. Shielding reduces the nuclear pull; effective nuclear charge is what's left.
• 3. Thinking shielding decreases down a group. It increases down a group.
• 4. Forgetting that shielding depends on inner-shell electrons, not valence electrons. Only inner electrons shield.
• 5. Thinking shielding and atomic size are the same thing. They are related but different — shielding is about electron–nucleus attraction; atomic size is about the actual size of the atom.

Quick Revision

• Shielding effect: reduction in force of attraction between nucleus and valence electrons by electrons in inner sub-shells.
• Down a group: number of electronic shells increases → shielding effect increases.
• Across a period: number of inner-shell electrons remains constant → shielding effect remains constant.
• Examples:
• Mg has greater shielding effect than Be.
• Si has greater shielding effect than C.
• Na has greater shielding effect than Li.
• Cl has greater shielding effect than F.
• Shielding affects other periodic trends: ionization energy decreases down a group because shielding increases.

60-Second Revision

• Shielding effect = inner electrons reduce pull on valence electrons.
• Down a group → shielding increases.
• Across a period → shielding constant.
• Mg > Be; Si > C; Na > Li; Cl > F.
• Warning: shielding does not increase across a period.