Screening Effect & Effective Nuclear Charge
In multi-electron atoms, the valence electrons (outermost electrons) are attracted by the nucleus but simultaneously repelled by the electrons present in the inner shells. This repulsion counteracts the attractive force of the nucleus. This phenomenon is called the Screening Effect or Shielding Effect.
1. Effective Nuclear Charge ($Z_{eff}$)
Due to shielding, the outer electron does not experience the full positive charge of the nucleus. The actual net positive charge experienced is called the Effective Nuclear Charge ($Z_{eff}$).
Where:
- $Z$ = Actual Nuclear Charge (Atomic Number).
- $\sigma$ (Sigma) = Screening Constant or Shielding Constant.
2. Screening Power of Orbitals
The ability of inner electrons to shield the outer electrons depends on the shape and penetration power of the orbitals they occupy.
Order of Shielding Power
$$ s > p > d > f $$
- s-orbitals: Spherical and penetrate closer to the nucleus. Strongest Shielding.
- f-orbitals: Diffused shape and far from the nucleus. Poorest Shielding.
3. Periodic Trends
A. Across a Period (Left to Right)
Electrons are added to the same shell. Electrons in the same shell do not shield each other very effectively. Meanwhile, the nuclear charge ($Z$) increases by +1 unit for each step.
- Result: $Z_{eff}$ Increases.
- Consequence: Atomic size decreases, Ionization Energy increases.
B. Down a Group (Top to Bottom)
New shells are added. The number of inner shell electrons increases significantly, increasing the shielding effect ($\sigma$).
- Result: Although $Z$ increases, $\sigma$ also increases. The $Z_{eff}$ remains roughly constant (or increases slightly), but the addition of new shells dominates.
- Consequence: Atomic size increases, Ionization Energy decreases.
4. Important Anomalies due to Poor Shielding
1. Gallium vs. Aluminum (Size Anomaly)
Normally, size increases down a group ($B < Al < Ga$). However, the atomic radius of Gallium (135 pm) is slightly smaller than Aluminum (143 pm).
Reason: Gallium follows the d-block elements. It has 10 inner 3d-electrons. Since d-electrons offer poor shielding, the outer electrons experience a higher effective nuclear charge ($Z_{eff}$), contracting the size.
2. Lanthanoid Contraction
In the 4f-series (Lanthanoids), electrons are filled in the 4f subshell. f-orbitals have very poor shielding.
Result: The nuclear charge ($Z$) increases by +14 across the series, but shielding does not compensate effectively. This causes a steady decrease in atomic size known as Lanthanoid Contraction.
Effect: Elements in the 5d series (e.g., Hf) have almost the same size as elements in the 4d series (e.g., Zr), making separation difficult.
3. Inert Pair Effect
In heavy p-block elements (Groups 13-16, e.g., Tl, Pb, Bi), the inner d- and f-electrons shield poorly. The nucleus holds the outer s-electrons ($ns^2$) very tightly, making them reluctant to participate in bonding.
Result: Lower oxidation states become more stable down the group (e.g., $Tl^{+1}$ is more stable than $Tl^{+3}$).
5. Slater's Rules (Brief Overview)
Slater's rules provide a method to calculate the screening constant ($\sigma$) quantitatively.
- Electrons in higher groups do not shield lower ones.
- Electrons in the same group shield by 0.35.
- Electrons in the ($n-1$) group shield by 0.85 (for s/p).
- Electrons in ($n-2$) and lower groups shield by 1.00 (full shielding).
Screening Effect Quiz
Test your concepts on Periodic Trends. 10 MCQs with explanations.
No comments:
Post a Comment