Periodic Table Electronic Configuration of Element: Overview, Questions, Preparation

Electrons occupy orbitals according to three fundamental principles (NCERT, p. 13.7):

• Aufbau Principle: Electrons fill orbitals in order of increasing energy, typically following the sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. The energy order is approximated by the  ( $n+l$ ) rule, where $n$ is the principal quantum number and  is the $l$ azimuthal quantum number.
• Paulis Exclusion Principle: Each orbital holds a maximum of two electrons with opposite spins.
• Hunds Rule: Electrons occupy degenerate orbitals singly with parallel spins before pairing, maximizing spin multiplicity for stability.

Orbitals are designated by:

• Principal Quantum Number ( $n$ ): Indicates the energy level or shell ( $\mathrm{n}=\mathrm{1,2},3,\dots$ ).
• Azimuthal Quantum Number ( $l$ ): Defines the subshell ( $\mathrm{s}:l=0,\mathrm{p}:l=1,\mathrm{d}:l=2,\mathrm{f}$ : $l=3$ ).
• Maximum Electrons: Each subshell holds $2(2l+1)$ electrons (s: $2,\mathrm{p}:6,\mathrm{d}:10,\mathrm{f}:14$ ).

The electronic configuration is written as, e.g., $1s^{2}2s^{2}2p^6$ for neon ( $Z=10$ ), indicating 2 electrons in 1s, 2 in 2s, and 6 in 2p orbitals (NCERT, p. 13.7).

The periodic tables structure is directly tied to electronic configurations, with elements grouped into blocks based on the subshell of the last electron. The table has 7 periods (rows) and 18 groups (columns), reflecting the filling of electron shells and subshells.

• Location: Groups 1 (alkali metals, except H) and 2 (alkaline earth metals, including He).
• Configuration: $n{s}^1$ (group 1) or $n{s}^2$ (group 2), where $n$ is the period number.

Examples:

• Lithium $(Z=3):1s^{2}2s^1$ (group 1, period 2).
• Magnesium (Z=12): $1s^{2}2s^{2}2p^{6}3s^2$ (group 2, period 3).

Total Elements: 14 (including H, He) (NCERT, p. 13.8).

• Location: Groups 1318 (boron family, carbon family, nitrogen family, oxygen family, halogens, noble gases).
• Configuration: $n{s}^{2}n{p}^{1-6}$ , where the last electron enters the p-orbital.

Examples:

• Boron $(Z=5):1s^{2}2s^{2}2p^1$ (group 13, period 2).
• Chlorine $(Z=17):1s^{2}2s^{2}2p^{6}3s^{2}3p^5$ (group 17, period 3).

Total Elements: 36 (NCERT, p. 13.9).

• Location: Groups 312 (transition elements).
• Configuration: $(n-1)d^{1-10}n{s}^{0-2}$ , where the last electron typically enters the $(n-1)d$ orbital.

Examples:

• Scandium (Z=21): [Ar] $3d^{1}4s^2$ (group 3, period 4).
• Copper (Z=29): [Ar] $3{\mathrm{d}}^{10}4{\mathrm{s}}^1$ (group 11, period 4, exception due to full dorbital stability).
• Exceptions: Some elements deviate from the Aufbau order for stability:
• Chromium (Z=24): $3d^{5}4s^1$ [Ar] (half-filled 3d for stability).
• Copper (Z=29): $\mathrm{}3d^{10}\mathrm{}4s^1$ [Ar] (fully filled 3d).

Total Elements: 39 (NCERT, p. 13.9).

• Location: Lanthanides (CeLu, Z=5871) and actinides (ThLr, Z=90103), placed below the main table in group 3.
• Configuration: $(n-2)f^{1-14}(n-1)d^{0-2}n{s}^2$ , where the last electron enters the (n2)f orbital.

Examples:

• Cerium (Z=58): [Xe] $4f^1\mathrm{}5d^1\mathrm{}6s^2$ (period 6).
• Uranium (Z=92): [Rn] $5f^3\mathrm{}6d^1\mathrm{}7s^2($ period 7).

Total Elements: 28 (14 lanthanides, 14 actinides) (NCERT, p. 13.10).