Simulation · Chemistry · Class 11
Fill the orbitals, one electron at a time
From the lesson Filling of orbitals and electronic configuration in Structure of Atom. Change the values and watch what happens.
Fill the orbitals, one electron at a timeChemistry · Class 11
The idea behind it
NCERT §2.6.4; §2.6.5; §2.6.6
- Aufbau principle: in the ground state, electrons enter the lowest-energy orbital available. The working order is 1s, 2s, 2p, 3s, 3p, then 4s, 3d, 4p, then 5s, 4d, 5p, then 6s, 4f, 5d, 6p, then 7s. It is a guide, and exceptions occur.
- Pauli exclusion principle: no two electrons in an atom share all four quantum numbers, so an orbital holds two electrons of opposite spin. Subshells hold 2 (s), 6 (p), 10 (d), 14 (f); a shell holds 2n².
- Hund's rule of maximum multiplicity: in a set of degenerate orbitals, electrons pair only after each orbital has one, so pairing starts with the 4th electron in p, the 6th in d and the 8th in f.
- Configurations are written as spdf notation (1s² 2s² 2p⁶ …) or as box diagrams with arrows; the box form shows all four quantum numbers.
- Examples: H 1s¹, He 1s², Li 1s² 2s¹, C 1s² 2s² 2p², N 1s² 2s² 2p³, Ne 1s² 2s² 2p⁶, Na [Ne] 3s¹, Ar [Ne] 3s² 3p⁶. Electrons of the filled inner shells are core electrons; those in the outermost shell are valence electrons.
- K and Ca put electrons into 4s before 3d; from Sc to Zn the 3d orbitals fill, then 4p from Ga to Kr.
- Exceptions: Cr is 3d⁵ 4s¹ and Cu is 3d¹⁰ 4s¹, not 3d⁴ 4s² and 3d⁹ 4s², because half-filled and completely filled subshells (p³, p⁶, d⁵, d¹⁰, f⁷, f¹⁴) are extra stable.
- That extra stability comes from the symmetrical electron distribution (small mutual shielding, stronger pull of the nucleus), smaller repulsion and the largest exchange energy, which is greatest when a subshell is half or completely filled and is also the basis of Hund's rule.