Orbitals and Electron Configuration

Electron configuration questions are workflow questions: count electrons, fill orbitals in the correct energy order, then check the total. This lesson focuses on the exam-safe order and the common traps (ions, especially 4s vs 3d).

Definitions (Must Know)

A. Orbital

An orbital is a region of space around the nucleus in which there is a high probability of finding an electron. It can hold up to two electrons with opposite spins.

B. Subshells (s, p, d)

A subshell is a set of orbitals of the same type within a shell.

• s subshell: 1 orbital → holds 2 electrons
• p subshell: 3 orbitals → holds 6 electrons
• d subshell: 5 orbitals → holds 10 electrons

C. Shell (principal quantum number, n)

A shell is an energy level labelled by n = 1, 2, 3, ….

Key Ideas (What Earns Marks)

• Fill lowest-energy orbitals first.
• Maximum 2 electrons per orbital (opposite spins).
• In orbitals of the same energy (e.g. 2p/3p/3d), electrons occupy separate orbitals before pairing (Hund’s rule).
• Use the energy order (not just “highest shell number”), especially around 4s/3d.

Common filling order (enough for H2 9476): 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p

Detailed Explanations

A. Writing a configuration (workflow)

1. Count electrons (atomic number; adjust for charge if it is an ion).
2. Fill orbitals using the energy order.
3. Check the total electrons match what you started with.

If it is an ion, do the electron count first (e.g. Al³⁺ has 13-3 = 10 electrons).

Quick examples:

• O: 1s² 2s² 2p⁴
• Cl: 1s² 2s² 2p⁶ 3s² 3p⁵

B. Ions (what changes)

Use the charge to adjust the electron count:

• Xⁿ⁺: the atom has lost n electrons → electrons = Z - n.

• Xⁿ⁻: the atom has gained n electrons → electrons = Z + n.

• Anions: add electrons (more negative).

• Main-group cations: remove electrons from the outer shell (highest n).

• Transition-metal cations: remove 4s electrons before 3d (the most common exam trap).

Mini example: S²⁻ has 16 + 2 = 18 electrons, so its configuration ends at 3p⁶.

C. 4s vs 3d removal (the exam rule)

In neutral atoms, 4s fills before 3d. But when transition metals form cations, the 4s electrons are removed first because (once 3d is occupied) the 4s electrons are higher in energy and more exposed.

So if you see a transition-metal ion, remove 4s electrons before 3d electrons.

D. Chromium and copper exceptions

The simple filling order predicts [Ar]3d⁴ 4s² for chromium and [Ar]3d⁹ 4s² for copper, but their observed ground-state configurations are:

• Cr: [Ar]3d⁵ 4s¹
• Cu: [Ar]3d¹⁰ 4s¹

Apply the ion rule after writing the correct neutral-atom configuration:

• Cr³⁺: remove the 4s electron, then two 3d electrons → [Ar]3d³
• Cu²⁺: remove the 4s electron, then one 3d electron → [Ar]3d⁹

Step through the elements, add or remove electrons, and tap the orbital boxes to see which arrangements are the ground state.

Worked Examples

Common Mistakes

• Writing more than 2 electrons in one orbital.
• Forgetting that p has three orbitals (so p can hold 6).
• For transition-metal ions, removing 3d electrons before 4s.
• Using the simple filling pattern for Cr and Cu instead of their ground-state exceptions.

Exam Tips

• Group number (main-group) matches valence electrons for neutral atoms (e.g. Group 17 → 7 valence electrons).
• Use noble gas shortcuts if allowed (e.g. [Ne]3s² 3p⁵), but full notation is always safe.
• For transition-metal ions, electrons are removed from 4s before 3d (for further H2 study, see Definition and Electron Configurations).
• Memorise the neutral-atom exceptions Cr: [Ar]3d⁵ 4s¹ and Cu: [Ar]3d¹⁰ 4s¹ before forming their ions.

Mind Stretchers