Chemistry 2e · Electronic Structure and Periodic Properties of Elements
Electronic Structure of Atoms (Electron Configurations)
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In 30 seconds
The quantum numbers from the previous topic give every electron a unique address — but which addresses do an atom's electrons actually occupy? That is Electron configuration The distribution of an atom's electrons among orbitals, e.g., 1s2 2s2 2p6... Full entry →: the organized listing of how an atom's electrons are distributed among shells and subshells. Three rules govern the assignment: the Aufbau principle Orbitals fill in order of increasing energy Full entry → (fill the lowest-energy orbitals first), the Pauli exclusion principle No two electrons share all four quantum numbers; max two per orbital, opposite spins Full entry → (at most two electrons per orbital, with opposite spins), and Hund's rule Degenerate orbitals fill singly with parallel spins before pairing Full entry → (fill orbitals of the same energy singly before pairing up). Following these rules lets you write any element's configuration — and the pattern is written into the periodic table, with s-, p-, d-, and f-block elements matching the subshell being filled.
Why this matters
An atom's electron configuration is the single best predictor of its chemistry. The Valence electrons Electrons in the outermost shell Full entry → — those in the outermost shell — determine how an atom bonds, what ions it forms, and how reactive it is; the periodic trends in the next topic all trace back to configurations. Configuration explains why sodium is a soft metal forming +1 ions, why noble gases are unreactive, why transition metals show multiple oxidation states and colors, and why chromium and copper break the rules. In practice, configurations let you predict ion charges and identify paramagnetic vs. diamagnetic substances — skills tested constantly in exams.
The college version
Core Concepts
Rule 1: The aufbau principle — fill from the bottom up
Electrons occupy orbitals in order of increasing energy. The ordering for neutral atoms is the familiar sequence:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s …
The famous surprise is that 4s fills before 3d (and 5s before 4d, 6s before 4f/5d): 4s is lower in energy than 3d for neutral atoms, even though 4 > 3 by shell number. The periodic table encodes this — each period ends where a new s orbital begins filling, which is why the fourth period starts with potassium and calcium (4s¹, 4s²) before the 3d transition metals.
Rule 2: The Pauli exclusion principle — two per orbital, no more
No two electrons in an atom may share all four quantum numbers (n, ℓ, mℓ, ms). Since the first three identify a specific orbital, the fourth — spin — limits each orbital to two electrons, ms = +12 and -12. This is why a shell holds 2n2 electrons maximum and the periodic table has its widths: s block 2 wide, p 6, d 10, f 14.
Rule 3: Hund's rule — spread out before pairing up
When several orbitals have the same energy (degenerate orbitals, like the three p orbitals), electrons enter them singly with parallel spins before any pairing. Pairing costs energy — two electrons in one orbital repel — so the lowest-energy arrangement maximizes unpaired electrons. That is why nitrogen is 2p3 with three unpaired electrons, not one doubly occupied orbital. Unpaired electrons make atoms and ions paramagnetic (attracted to magnetic fields); all-paired species are diamagnetic.
Writing configurations: full, shorthand, and box diagrams
A full configuration lists every occupied subshell with its electron count as a superscript. For phosphorus (Z = 15):
1s2 2s2 2p6 3s2 3p3
Check: 2 + 2 + 6 + 2 + 3 = 15 electrons. Noble-gas shorthand replaces filled inner shells with the preceding noble gas: [Ne] 3s2 3p3. Orbital box diagrams show each orbital as a box and each electron as an arrow (↑ or ↓), making spin and Hund's rule visible. The valence electrons — for phosphorus, the five in 3s2 3p3 — are the ones chemistry cares about.
Exceptions to the rules: chromium and copper
Aufbau predicts [Ar] 4s2 3d4 for chromium (Z = 24) and [Ar] 4s2 3d9 for copper (Z = 29), but the actual ground states are:
Cr: [Ar] 4s1 3d5 Cu: [Ar] 4s1 3d10
Both promote one 4s electron into 3d to achieve a half-filled (d⁵) or filled (d¹⁰) d subshell — extra stability from reduced repulsion and favorable exchange energy among parallel-spin electrons. Similar effects appear for Mo, Ag, Au, and others; always check the actual configuration.
Ions: electrons leave the outer shell first
Cations lose electrons from the highest-n occupied subshell first, not the subshell filled last. Iron (Z = 26), [Ar] 4s2 3d6, loses its two 4s electrons to give Fe²⁺ ([Ar] 3d6); Fe³⁺ ([Ar] 3d5) then loses one 3d. Anions add electrons to the partially filled outer subshell — oxygen gains two to reach 2p6, giving O²⁻. Reaching a noble-gas configuration (octet) drives ion charges.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Filling order (aufbau) | Removal order for cations | Neutral atoms fill 4s before 3d; cations remove 4s (highest n) BEFORE 3d |
| Aufbau prediction | Actual configuration | Cr, Cu (and Mo, Ag, Au...) deviate to reach half-filled/filled d subshells |
| Fe²⁺ configuration | Fe configuration minus "last filled" | Fe²⁺ = [Ar]3d⁶ — 4s² leaves first, not 3d |
| [Ar] 3d6 (Fe²⁺) | [Ar] 4s2 3d4 (predicted Cr) | Same electron count (24) but different atoms — always state which element/ion |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Electrons are like kids boarding a school bus: front seats (lowest energy) first, two kids per seat, and when several empty seats are in a row, each kid takes their own seat before anyone doubles up. Chromium and copper are the kids who break the seating chart because they end up calmer (more stable) with a half-full or full back row. Write the seating chart of an atom — which "seat" each electron takes — and you've written its electron configuration, the secret code predicting how the atom behaves.
Worked example
Worked Example 1: Iron and its ions — where do electrons leave from?
Problem. Write the configuration of iron (Z = 26), then those of Fe²⁺ and Fe³⁺.
Strategy. Fill in order: 1s2 2s2 2p6 3s2 3p6 (18 electrons, argon core), then 4s before 3d:
Fe: [Ar] 4s2 3d6
Count check: 18 + 2 + 6 = 26 ✓. For Fe²⁺, remove the two highest-n electrons — the 4s electrons, NOT the 3d:
Fe2+: [Ar] 3d6 Fe3+: [Ar] 3d5
Count check: Fe²⁺ has 26 - 2 = 24 electrons; [Ar] = 18 + 6 = 24 ✓. Interpretation. Fe²⁺ (3d6) has four unpaired electrons; Fe³⁺ (3d5) five — both strongly paramagnetic. Exam trap: an ion's configuration is NOT "remove what was filled last" — strip the outermost shell (highest n) first.
Worked Example 2: The chromium and copper exceptions
Problem. Predict chromium's (Z = 24) and copper's (Z = 29) configurations by aufbau, then give the actual ground states and explain the difference.
Strategy. Aufbau prediction: Cr = [Ar] 4s2 3d4, Cu = [Ar] 4s2 3d9. Actual ground states:
Cr: [Ar] 4s1 3d5 Cu: [Ar] 4s1 3d10
Count check: Cr: 18 + 1 + 5 = 24 ✓; Cu: 18 + 1 + 10 = 29 ✓. Interpretation. One 4s electron is promoted into 3d, leaving 4s half-filled and 3d half-filled (d⁵) or filled (d¹⁰) — arrangements with reduced repulsion and favorable exchange energy. The same pattern appears in Mo ([Kr] 5s1 4d5), Ag ([Kr] 5s1 4d10), and Au — always verify against a periodic table.
Key takeaways
- Three rules: aufbau (lowest energy first), Pauli (max 2 per orbital, opposite spins), Hund (fill degenerate orbitals singly, parallel spins, before pairing).
- Filling order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s — 4s fills before 3d.
- Valence electrons = outermost-shell electrons; for main-group elements their count equals the group number.
- Exceptions: Cr = [Ar] 4s1 3d5, Cu = [Ar] 4s1 3d10 (half-filled/filled d).
- Cations: remove electrons from the highest-n subshell first — Fe²⁺ = [Ar] 3d6, not 4s2 3d4.
- Always verify: superscripts must sum to Z (and to Z ± charge for ions).
Check yourself
4 review questions from the chapter. Try each one, then open the answer.
State the aufbau principle, the Pauli exclusion principle, and Hund's rule in your own words.
Show answer
Aufbau: electrons occupy orbitals in order of increasing energy. Pauli: no two electrons share all four quantum numbers — max two per orbital, opposite spins. Hund: degenerate orbitals fill singly with parallel spins before pairing.
Write the full and shorthand configurations for sulfur (Z = 16). How many valence electrons does it have?
Show answer
1s2 2s2 2p6 3s2 3p4 = [Ne] 3s2 3p4. Valence electrons: 6 (group 16 — needs 2 more for an octet, so it forms S²⁻).
Why does 4s fill before 3d, yet Fe²⁺ loses 4s electrons first?
Show answer
For neutral atoms, 4s is lower in energy than 3d, so it fills first. When forming cations, the highest-n electrons (n = 4) are farthest out and least tightly bound, so 4s leaves first.
How many unpaired electrons does sulfur ([Ne] 3s2 3p4) have, and is it paramagnetic or diamagnetic?
Show answer
Two unpaired electrons: the three 3p orbitals fill singly first (three unpaired), then the fourth electron pairs with one, leaving two unpaired. Sulfur is paramagnetic.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Electron configuration
- The distribution of an atom's electrons among orbitals, e.g., 1s2 2s2 2p6...
- Aufbau principle
- Orbitals fill in order of increasing energy
- Pauli exclusion principle
- No two electrons share all four quantum numbers; max two per orbital, opposite spins
- Hund's rule
- Degenerate orbitals fill singly with parallel spins before pairing
- Valence electrons
- Electrons in the outermost shell
Sources & references
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