What you will learn

• Write full and noble-gas-core configurations in s, p and d notation for elements with atomic numbers 1–36.
• Count electrons in positive and negative ions.
• Translate configurations into electron-in-boxes diagrams.
• Remove 4s electrons before 3d when forming transition-metal cations.
• Check chromium and copper ground-state exceptions.

A1.2.3 requires both notation forms for atoms and ions of elements with atomic numbers 1–36. The rules from A1.2.2 are applied here rather than retaught in full.

1. Read s, p and d notation

An electronic configuration states how electrons occupy the orbitals of an atom or ion. Unless told otherwise, use its ground state, the arrangement with the lowest total energy.

In 3p⁴: 3 = principal shell; p = subshell type; 4 = number of electrons in that subshell.

The superscript is an electron count, not the charge of the atom. Thus oxygen, 1s² 2s² 2p⁴, has 2 + 2 + 4 = 8 electrons. It does not have a 4+ charge. Keep orbital labels and ion charges distinct: O²⁻ is an ion symbol; 2p⁶ describes a subshell.

Use the usual filling sequence 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p, with the Cr and Cu exceptions below. A configuration may list 3d before 4s: [Ar] 3d⁶ 4s² and [Ar] 4s² 3d⁶ describe the same occupancy. Written order alone does not determine which electrons leave an ion.

Noble-gas shorthand: [He] = 1s² (2 electrons); [Ne] = 1s² 2s² 2p⁶ (10); [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶ (18). Write the remaining occupied subshells after the core. Expand the core if a question asks for a full configuration.

2. Count electrons before writing a configuration

Atomic number Z fixes the number of protons. A neutral atom has Z electrons. Forming an ion changes its electron number, not its atomic number.

Positive ion: electrons = Z − positive charge.
Negative ion: electrons = Z + magnitude of negative charge.

For example, Mg²⁺ has 12 − 2 = 10 electrons; Cl⁻ has 17 + 1 = 18. The mass number is not used to count electrons. A negative charge means extra electrons, because electrons themselves have negative charge.

Species with equal electron numbers are isoelectronic. Ne, Na⁺, Mg²⁺, F⁻ and O²⁻ each have ten electrons and the configuration 1s² 2s² 2p⁶. They remain different elements because their proton numbers differ.

Important limit: equal electron counts do not guarantee identical ground-state configurations for every transition-metal species. Ca and Ti²⁺ both have 20 electrons, but Ca is [Ar] 4s² whereas Ti²⁺ is [Ar] 3d². For transition-metal ions, start from the correct neutral atom and remove its outer electrons.

3. Translate into electron-in-boxes diagrams

Each box is one orbital; each arrow is one electron. Label the subshells. Draw one box for s, three for p and five for d, even when some are empty. A pair within one box must have opposite spins. Across equal-energy orbitals in a subshell, place electrons singly with parallel spins before pairing.

Any of the three equivalent p orbitals can contain the pair. Reversing every spin arrow also gives an equivalent diagram. Arrows represent spin states, not electron paths. A diagram that pairs all four 2p electrons into two boxes while leaving the third empty is not the ground-state arrangement.

When using a core such as [Ar] in a diagram, label it clearly as an omitted filled core. If asked for the complete diagram, draw its occupied orbitals too. The diagrams below state when a core is omitted.

4. Worked examples: main-group atoms and ions

5. Transition-metal atoms and ions

For elements in the first d series, include both the occupied 3d and 4s subshells. Although the usual building sequence enters 4s before 3d, 4s electrons are removed before 3d electrons when forming transition-metal cations. Their relative energies change with electron occupancy and ionisation; “4s always has lower energy” is incorrect.

Write the neutral atom correctly → remove 4s electrons first → remove further electrons from 3d → check charge and total.

Chromium: Cr (Z = 24) is [Ar] 3d⁵ 4s¹. Cr³⁺ loses the 4s electron and two 3d electrons, giving [Ar] 3d³ (21 electrons). Draw three singly occupied d boxes with parallel spins and two empty d boxes. Half-filled and filled d subshells are useful memory aids for Cr and Cu; the underlying criterion is the lowest total energy, not an instruction to rearrange every atom into a half-filled subshell.

6. Check configurations across atomic numbers 1–36

Use these checkpoints with the filling sequence and electron total. For each neutral atom between checkpoints, add the appropriate electrons; remember Cr and Cu. Ionic charge changes the count and removal route.

7. Common misconceptions and exam tips

• “A positive ion gains electrons.” Positive ions have lost negative electrons; subtract the charge from Z.
• “Fe²⁺ is [Ar] 3d⁴ 4s².” Remove the 4s electrons first: [Ar] 3d⁶.
• “The superscript is an orbital count.” It counts electrons; d⁶ requires five orbital boxes, not six.
• “Every electron must be paired.” Apply Hund to equal-energy orbitals; p⁴ has two unpaired electrons.
• “Cu follows [Ar] 3d⁹ 4s².” Use its observed ground state, [Ar] 3d¹⁰ 4s¹.
• “Same configuration means same element.” Proton number identifies an element; ions of different elements can have the same configuration.

Final check: correct species and charge; correct electron total; no subshell over capacity; all orbital boxes shown and labelled; opposite-spin pairs; parallel single occupants; 4s removed first for transition-metal cations.

Your revision summary

s, p and d notation records subshell occupancy. Electron-in-boxes diagrams also show the distribution and spin pairing within each subshell. Count from Z and charge before starting, use the correct neutral ground state and ion-removal route, then total every superscript or arrow.

Retrieval practice: write and draw N, N³⁻, Mg²⁺, Cr³⁺ and Cu²⁺. Explain each electron total and identify any omitted noble-gas core.

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