Summary of "Slater's Rule | Effective Nuclear Charge | Atomic structure - BSc 1st Year Inorganic Chemistry |"

Summary of the Video: "Slater's Rule | Effective Nuclear Charge | Atomic structure - BSc 1st Year Inorganic Chemistry"

Definition: Effective Nuclear Charge is the net positive charge experienced by an electron in an atom after accounting for the shielding effect of other electrons.
Key Concepts:
- Inner (core) electrons shield outer electrons from the full nuclear charge.
- Electrons in inner shells are called intervening electrons and act as a shield between the nucleus and outer shell electrons.
- This shielding reduces the actual nuclear charge (Z) to an Effective Nuclear Charge (Zeff) experienced by outer electrons.
- The Shielding Constant (σ) quantifies this reduction.
- Formula: Zeff = Z - σ where - Z = actual nuclear charge (atomic number) - σ = Shielding Constant (screening constant)
Shielding/Screening Effect: The phenomenon where inner electrons reduce the attraction between the nucleus and outer electrons.
Factors Affecting Zeff:
- Atomic size: - Increases down a group → Zeff decreases (due to more shielding and larger distance). - Decreases across a period → Zeff increases (due to increasing nuclear charge and similar shielding).
- Number of intervening electrons: More inner electrons → higher σ → lower Zeff.
Shielding varies by subshell: Shielding effectiveness order: s > p > d > f because s orbitals are closest to the nucleus and shield more effectively.

Purpose: To estimate the Shielding Constant σ and thus calculate Zeff for any electron in an atom.
General Guidelines:
- Write the electronic configuration in increasing order of principal quantum number (n), not the filling order.
- Divide electrons into groups based on their shells and subshells relative to the electron of interest.
Three Main Rules:
- Rule 1: For ns or np electrons - Electrons in the same shell (n) contribute 0.35 each (except for 1s where it is 0.30). - Electrons in the (n-1) shell contribute 0.85 each. - Electrons in shells with principal quantum number less than (n-1) contribute 1.00 each. - Electrons in shells with principal quantum number greater than n contribute 0.
- Rule 2: For nd or nf electrons - Electrons in the same (n) shell contribute 0. - Electrons in the (n-1) d or f subshell contribute 0.35 each. - Electrons in the (n-1) s or p subshells and all inner shells contribute 1.00 each.
- Rule 3: For calculating Zeff at the periphery (outermost electron) - Include the electron itself when calculating σ (unlike rules 1 and 2 where the electron itself is excluded). - Apply the same weighting factors as above.
Examples: Calculation of σ and Zeff for electrons in Copper (Z=29), Manganese (Z=25), Chromium (Z=24), Nickel (Z=28), Sodium (Z=11), Chlorine (Z=17), and Titanium (Z=22) were demonstrated step-by-step.

Why 4s orbital is filled before 3d in potassium: Zeff for 4s electron is higher than that for 3d electron → 4s electron experiences stronger nuclear attraction → lower energy → more stable → filled first.
Why 4s electrons are lost before 3d electrons in transition metals during cation formation: 3d electrons experience higher Zeff and are held more tightly by the nucleus than 4s electrons, so 4s electrons are removed first.
Why cations are smaller than parent atoms: Removal of electrons decreases shielding (σ), increasing Zeff on remaining electrons → stronger pull by nucleus → smaller atomic radius.
Why anions are larger than parent atoms: Addition of electrons increases shielding and decreases Zeff, leading to a larger atomic radius.