Electron configurations of atoms and ions up to Z = 36 in terms of shells…

The idea

Electrons fill sub-shells in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. An s sub-shell holds 2 electrons, p holds 6 and d holds 10, because each orbital takes two electrons of opposite spin. The 4s sub-shell lies just below 3d in energy so it fills first, but once the 3d sub-shell is occupied it drops below 4s, which is why transition metal ions lose their 4s electrons first: iron is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² but iron(II) is 3d⁶ with no 4s. Chromium and copper break the pattern, adopting 3d⁵ 4s¹ and 3d¹⁰ 4s¹ because a half-filled or full d sub-shell is more stable.

What earns the marks

• The sub-shells written in the correct filling order including 4s before 3d.
• The correct number of electrons in each sub-shell totalling the atomic number.
• Removing 4s electrons before 3d electrons when writing a transition metal ion.
• The exceptional configurations of chromium and copper.

Common mistakes and how to put them right

• Removing 3d electrons first when forming an ion such as iron(II).

• Writing chromium as 3d⁴ 4s² instead of 3d⁵ 4s¹.

• Putting more than six electrons in a p sub-shell or more than ten in a d sub-shell.

• Inventing a 2d sub-shell, which does not exist.

• Adding electrons to a positive ion rather than removing them.

Exam tips

• Count the electrons in your finished configuration and check the total against the atomic number or adjusted ion charge.
• Write the atom first, then edit it to make the ion; do not build an ion configuration from scratch.
• In electrons-in-boxes diagrams, fill each orbital singly before pairing, and show paired arrows pointing opposite ways.