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⚗️ Chemistry Course
⚛️ Atomic Structure & the Periodic Table · Lesson 4 of 5

Electron Configuration

The address system for every electron — and the reason the periodic table is shaped the way it is.

High schoolIntro Gen ChemUni Year 1
⏱️ About 22 min
Electron Configuration — illustration
Illustrative image (AI-generated) — decorative only.

Electrons don't pile into an atom at random. They fill a strict set of energy levels in a predictable order — and once you know that order, you can write the electron arrangement of any element and predict how it will react.

💡
The big idea: Electrons occupy orbitals grouped into subshells (s, p, d, f) and shells. They fill from lowest energy upward, following three rules — Aufbau, Pauli, and Hund — and that filling pattern is exactly what gives the periodic table its blocks.
🎯 By the end, you'll be able to
  • Name the subshells (s, p, d, f) and their electron capacities
  • Apply the Aufbau order, the Pauli exclusion principle, and Hund's rule
  • Write the full and noble-gas electron configuration of an element
  • Explain why 4s fills before 3d — and know the chromium/copper exceptions
📎 Helpful to know first
  • Protons, neutrons & electrons

Orbitals, subshells and shells

Electrons live in orbitals — regions of space where an electron is likely to be. Orbitals come in types called subshells: s (holds 2 electrons), p (6), d (10) and f (14). Subshells group into shells labelled by the principal quantum number n (n = 1, 2, 3…).

Think of it as an address: an electron in 3p is in shell 3, subshell p. Chemistry mostly cares about the outermost (valence) shell — those are the electrons that bond.

🔑 Three rules for filling
Aufbau: fill the lowest-energy subshell first. Pauli: each orbital holds at most 2 electrons, with opposite spins. Hund: within a subshell, spread electrons out singly before pairing up.
✨ Why 4s fills before 3d
The energy ladder isn't a simple 1, 2, 3… by shell. The 4s subshell is actually slightly lower in energy than 3d, so it fills first. That's why potassium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ — its outer electron goes to shell 4, not into 3d. This single fact is why the periodic table has a d-block that starts one row 'late'.
e⁻ e⁻ n=1 (2 e⁻) e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ n=2 (8 e⁻) e⁻ n=3 (1 e⁻) 11 p⁺ 12 n Sodium · Na · Z=11 · mass number 23
Sodium's shells fill 2, 8, 1 — matching its configuration 1s² 2s² 2p⁶ 3s¹. That lone outer electron is why sodium is so reactive.
\[ \text{Na}:\ 1s^2\,2s^2\,2p^6\,3s^1 \quad=\quad [\text{Ne}]\,3s^1 \]
Full configuration and the shorthand noble-gas form (replace the inner core with the previous noble gas).
📝 Worked example: Write the electron configuration of sulfur (Z = 16).
  1. 16 electrons to place, filling in Aufbau order: 1s, 2s, 2p, 3s, 3p…
  2. 1s² (2) → 2s² (4) → 2p⁶ (10) → 3s² (12) → 3p⁴ (16). Stop at 16.
  3. So sulfur is 1s² 2s² 2p⁶ 3s² 3p⁴, or [Ne] 3s² 3p⁴.
✓ 1s² 2s² 2p⁶ 3s² 3p⁴ — six valence electrons (3s² 3p⁴).
⚠️ Two famous exceptions
A few elements break the naive pattern for extra stability. Chromium is [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²) and copper is [Ar] 3d¹⁰ 4s¹ (not 3d⁹ 4s²) — a half-full and a full d-subshell are especially stable, so one 4s electron drops into 3d. Worth memorising these two.
✏️ Practice: How many valence (outermost-shell) electrons does phosphorus (Z = 15) have? Hint: its configuration ends in 3s² 3p³.
electrons
Solution
  1. Phosphorus: 1s² 2s² 2p⁶ 3s² 3p³.
  2. The outermost shell is n = 3: it holds 3s² and 3p³.
  3. 2 + 3 = 5 valence electrons (which is why phosphorus sits in group 15).

Check your understanding

1. What is the maximum number of electrons in a single p subshell?
A p subshell has 3 orbitals × 2 electrons each = 6. (s holds 2, d holds 10, f holds 14.)
2. Which correctly shows why potassium's last electron enters the 4s subshell?
4s is slightly lower in energy than 3d, so it fills first — giving K the configuration [Ar] 4s¹. This is the normal rule, not an exception.
3. Chromium's ground-state configuration is:
Chromium is [Ar] 3d⁵ 4s¹ — a half-filled d-subshell is extra stable, so one 4s electron shifts into 3d. Copper does the same to reach a full 3d¹⁰.
🔎 Related lessons
Previous: Isotopes & Atomic Mass The Periodic Table & Its Trends
✅ Key takeaways
  • Electrons fill orbitals in subshells s (2), p (6), d (10), f (14), grouped into shells.
  • Fill by Aufbau (lowest energy first), Pauli (2 per orbital) and Hund (singly first).
  • 4s fills before 3d — the reason the d-block appears where it does.
  • Chromium ([Ar] 3d⁵ 4s¹) and copper ([Ar] 3d¹⁰ 4s¹) are the key exceptions to know.
➡️ Electron configuration explains the shape of the periodic table. Next, we read the table itself: how configuration drives the trends in size, ionization energy and electronegativity that predict how elements behave.
Want to test yourself on this? Try the Chemistry practice test →
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