Electrons, energy levels and atomic orbitals — common mistakes

A very common mistake is to assume the third shell (3s, 3p, 3d) fills completely before the fourth shell (4s) begins. This is incorrect! The 4s sub-shell is at a lower energy level than the 3d sub-shell, so it fills first. Remember the sequence: ...3p, then 4s, then 3d.

The anomalous configurations of Cr and Cu are very common exam questions, both in multiple choice and structured questions. Make sure you have memorised them. Remember the reason: the enhanced stability of a half-filled or fully-filled d sub-shell outweighs the energy cost of promoting a 4s electron.

Due to shielding and penetration effects, the 4s orbital is actually at a slightly lower energy level than the 3d orbitals. The Aufbau principle dictates that we fill the lowest energy orbitals first, so electrons go into 4s before starting on 3d. However, once electrons are in the 3d orbitals, the relative energy levels shift, which is why when forming ions, 4s electrons are removed first.

Degenerate orbitals are orbitals that have the exact same energy. For example, the three p orbitals in a given sub-shell ($p_x, p_y, p_z$) are degenerate. Hund's rule applies to the filling of degenerate orbitals.

No. For Cambridge A-Level Chemistry (9701), you are required to know and be able to sketch the shapes of s and p orbitals, but you only need to know that d-orbitals have more complex shapes. You are not expected to draw them.

It's a shorthand way of writing electron configurations. Instead of writing out the full configuration from 1s, you can use the symbol of the preceding noble gas in square brackets to represent the core electrons. For example, for Sodium ($1s^2 2s^2 2p^6 3s^1$), the configuration of Neon is $1s^2 2s^2 2p^6$. So, you can write Sodium's configuration as $[Ne] 3s^1$.