Short Answer

Shielding effect is the repulsion experienced by outer electrons due to inner electrons, which reduces the effective nuclear charge felt by valence electrons.

• The more inner electrons an atom has, the weaker the attraction between the nucleus and outermost electrons.
• This effect explains why atomic size increases down a group, and why ionization energy and electronegativity decrease as you go down the periodic table.

Detailed Explanation :

Definition of Shielding Effect

Shielding effect, also called electron shielding, refers to the reduction in the effective nuclear charge experienced by valence (outermost) electrons due to the repulsion from inner (core) electrons.

• The nucleus has a positive charge due to protons, which attracts electrons.
• However, inner electrons repel outer electrons, partially cancelling the nuclear attraction.
• This phenomenon is important for explaining atomic properties and periodic trends.

How Shielding Occurs

1. Inner Electron Repulsion:
   • Electrons in inner shells push outer electrons away due to their negative charge.
   • This reduces the net positive pull felt by outer electrons.
2. Core vs Valence Electrons:
   • Core electrons strongly shield valence electrons.
   • Valence electrons feel less attraction from the nucleus than inner electrons.
3. Effect on Atomic Properties:
   • Atomic size: Increases down a group because valence electrons are shielded from nucleus.
   • Ionization energy: Decreases as electrons are easier to remove due to shielding.
   • Electronegativity: Decreases down a group because outer electrons are less attracted to the nucleus.

Factors Affecting Shielding Effect

1. Number of Inner Electrons:
   • More core electrons → stronger shielding → weaker effective nuclear charge.
2. Number of Electron Shells:
   • More shells → valence electrons farther from nucleus → more shielding.
3. Nuclear Charge:
   • More protons → increases attraction, but shielding reduces net effect on outer electrons.
4. Position in Periodic Table:
   • Across a period: shielding effect is nearly constant because electrons are added to the same shell → effective nuclear charge increases.
   • Down a group: shielding increases as new shells are added → reduces pull on valence electrons.

Consequences of Shielding Effect

1. Atomic Radius:
   • Shielding reduces attraction → outer electrons are farther → larger atomic size down a group.
2. Ionization Energy:
   • Electrons easier to remove → ionization energy decreases down a group.
3. Electronegativity:
   • Weaker nuclear pull → lower tendency to attract bonding electrons → decrease in electronegativity down a group.
4. Chemical Reactivity:
   • Metals (low ionization energy) become more reactive down a group due to increased shielding.
   • Non-metals (high electronegativity) become less reactive down a group for the same reason.
5. Periodic Trends Explanation:
   • Explains why atomic size increases, ionization energy decreases, and electronegativity decreases down a group.
   • Explains why atomic properties change gradually across periods and groups.

Examples

• Alkali Metals (Group 1):
  • Lithium → small shielding → relatively higher ionization energy.
  • Cesium → more shells → strong shielding → much lower ionization energy.
• Halogens (Group 17):
  • Fluorine → less shielding → high electronegativity.
  • Iodine → more shielding → lower electronegativity.

Conclusion

Shielding effect is the repulsion of outer electrons by inner electrons, reducing the effective nuclear charge on valence electrons. It plays a crucial role in determining atomic size, ionization energy, electronegativity, and chemical reactivity. Understanding shielding effect is essential to explain periodic trends and chemical behavior of elements across the periodic table.