Electronic Configuration

Orbitals and 1st ionisation

Orbital Definition

• A region of space where there is a high probability of finding a electron

• An orbital can hold up to two electrons

• Electrons can only be in the same orbital if they have opposite spin

Equation for max number of electrons in a shell

2n²

How to fill orbitals

• Electrons enter lowest energy orbital available (Aufbau principle)

• Electrons prefer to occupy orbitals on their own (Hund’s rule)

S-block group numbers

Groups 1-2

D-block group numbers

Transition metals

P-block group numbers

Groups 3-8

Chromium electronic configuration

1s² 2s² 2p⁶ 3s² 3p⁶ 4s1 3d5

Copper electronic configuration

1s² 2s² 2p⁶ 3s² 3p⁶ 4s1 3d10

Transition metal exceptions explantation

• Copper and Chromium

• Due to 4s and 3d levels being close in energy

• Fills or half fills the 3d subshell with no gaps

• More stable when there is a completely full/half full subshell to maximise exchange of energy

• When they lose electrons, they’re lost from 4s before 3d

Be vs B (electron removed from p vs s subshell) - group 2 vs 3

Why is the ionisation energy for B lower than Be

• P-subshell is more shielded from the nucleus than an s-subshell

• Less attraction between nucleus and outer electrons

P vs S - groups 5 vs 6

Why is the ionisation energy for S lower than P

• Outer electron removed from orbital with has a pair of electrons

• Increased repulsion between electrons

• Electron lost more easily - already at higher energy level