Chemistry: Atoms First 2e · Electronic Structure and Periodic Properties of Elements

Electronic Structure of Atoms (Electron Configurations)

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

An is a shorthand description of how an atom's electrons are distributed among orbitals. It answers the question: which orbitals are occupied, and by how many electrons? The configuration of an atom determines its chemical personality — which elements are metals, how many bonds an atom can form, and why the periodic table has its familiar block structure.

Three rules govern how electrons fill orbitals in the ground state (lowest-energy arrangement):

: electrons fill the lowest-energy orbitals first.
: each orbital holds at most two electrons, and they must have opposite spins.
: when filling orbitals of equal energy (, such as the three 2p orbitals), electrons occupy them singly and with parallel spins before pairing up.

The filling order follows increasing orbital energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. Notice the famous quirk: 4s fills before 3d because the 4s orbital is lower in energy for neutral atoms. Writing configurations correctly — including the exceptions like chromium and copper — is the skill this topic builds.

Why this matters

Electron configurations are the organizing principle of the periodic table:

(the outermost electrons) govern reactivity. Elements in the same group share the same valence configuration, which is why they behave alike.
The table's blocks (s, p, d, f) literally reflect which subshell is being filled.
Ion formation and oxidation states follow directly from which electrons are easiest to remove — a concept that drives Topics 5–7 (periodic trends, the periodic table, ionic compounds).
Predicting paramagnetism (unpaired electrons) and spectroscopic properties starts from configurations.

In practical science, configurations explain why sodium is a violent metal while neon is inert, why iron has multiple oxidation states (2+ and 3+), and why transition metals are colorful. Any student who can write a configuration can predict an element's group, period, and likely ion charge without memorizing a single fact about that element.

The college version

Core Concepts

The Aufbau principle and filling order

Electrons occupy orbitals in order of increasing energy. The sequence (with the number of electrons each subshell can hold in parentheses) is:

$1s (2) \to 2s (2) \to 2p (6) \to 3s (2) \to 3p (6) \to 4s (2) \to 3d (10) \to 4p (6) \to 5s (2) \to 4d (10) \to 5p (6) \to 6s (2) \to 4f (14) \to 5d (10) \to 6p (6) \to 7s (2) \to 5f (14) \to 6d (10) \to 7p (6)$

The subshell capacities come from the quantum numbers of Topic 3: s holds 2 (one orbital), p holds 6 (three orbitals), d holds 10 (five orbitals), and f holds 14 (seven orbitals). The 4s-before-3d ordering is why potassium ([Ar]4s¹) and calcium ([Ar]4s²) appear before the transition metals.

Pauli exclusion principle

No two electrons in an atom may have the same set of all four quantum numbers. Since the first three numbers (n, l, m_l) define an orbital, only two electrons can occupy any orbital — and they must differ in spin (m_s = +1/2 or -1/2). An orbital diagram represents them as up and down arrows: ↑↓.

Hund's rule

For degenerate orbitals (same subshell), electrons first occupy each orbital singly with parallel spins, then pair. This minimizes electron–electron repulsion and is confirmed by the fact that oxygen, with two unpaired 2p electrons, is paramagnetic. A common mistake is pairing electrons early; Hund's rule forbids that in the ground state.

Writing configurations: full, shorthand, and core/valence

Full configuration: list every occupied subshell, e.g., phosphorus (Z = 15): 1s² 2s² 2p⁶ 3s² 3p³.
: replace the filled inner core with the noble gas symbol, e.g., phosphorus: [Ne] 3s² 3p³.
Core vs. valence: the noble-gas core electrons are chemically inert; the valence electrons (outermost shell) participate in bonding. For phosphorus, valence = 3s² 3p³ (five electrons).

Exceptions to the Aufbau order

A few elements deviate from the simple filling order because half-filled and fully filled subshells are extra stable:

Chromium (Z = 24): expected [Ar] 4s² 3d⁴, actual [Ar] 4s¹ 3d⁵.
Copper (Z = 29): expected [Ar] 4s² 3d⁹, actual [Ar] 4s¹ 3d¹⁰.

The pattern: a d⁵ or d¹⁰ configuration is stabilized enough to "borrow" one electron from the 4s. Other exceptions exist among heavier transition metals (e.g., Nb, Mo, Ru, Ag, Pt, Au), so always verify with a periodic table rather than assuming the Aufbau order is absolute.

How It Works / Step-by-Step Process

To write a ground-state electron configuration:

Find the element's atomic number Z — that is the number of electrons to place.
Follow the Aufbau filling order, adding the subshell capacities (s:2, p:6, d:10, f:14).
Apply Hund's rule within each degenerate subshell (single occupancy first).
Stop when the superscripts sum to Z.
Check your work: add up all superscripts; the total must equal Z.
Optionally compress the inner core into a noble-gas symbol.

Common Confusions

Common Confusion	Correct Understanding
3d fills before 4s because 3 < 4.	For neutral atoms 4s is lower in energy than 3d, so 4s fills first; the order is by energy, not by shell number.
Hund's rule means electrons always avoid pairing.	It applies only within a degenerate subshell in the ground state; after each orbital has one electron, pairing is required.
The superscripts add to the group number.	They add to the atomic number Z, not the group number.
All elements follow the Aufbau order.	Cr, Cu, and several heavier transition metals deviate for extra stability of half-filled/filled d subshells.
Core electrons participate in bonding.	Only valence electrons do; core (noble-gas) electrons are chemically inert.
When iron ionizes, it loses 3d electrons first.	Fe loses the 4s electrons first (they are outermost), so Fe²⁺ is [Ar]3d⁶.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Filling an atom with electrons is like filling a theater row by row: you start at the front (lowest energy) and go back, and each seat holds two people who must face opposite directions. If several seats are equal — like three chairs in a row — people sit one per chair first before anyone shares. A few elements cheat slightly, like chromium, because a half-full row of d-seats is extra cozy.

Worked example

Example 1: Full configuration of phosphorus (Z = 15)

Follow the filling order and count as you go:

$1s 2 (2), 2s 2 (4), 2p 6 (10), 3s 2 (12), 3p 3 (15)$

So the full configuration is:

$P: 1s 2 2s 2 2p 6 3s 2 3p 3$

Check: 2 + 2 + 6 + 2 + 3 = 15 ✓. The valence shell is 3s² 3p³ — five valence electrons, matching phosphorus's group 15 position and its tendency to form 3− ions and three covalent bonds.

Example 2: Noble-gas shorthand for iron (Z = 26)

The noble gas before iron is argon (Z = 18). Fill from 4s onward:

$[Ar] 4s 2 3d 6$

Check: 18 + 2 + 6 = 26 ✓. Iron's valence electrons include both the 4s and 3d electrons (eight total), which explains its common oxidation states: it can lose the two 4s electrons (Fe²⁺, [Ar]3d⁶) or additionally one 3d electron (Fe³⁺, [Ar]3d⁵ — a half-filled d subshell, extra stable).

Example 3: The exceptions — chromium and copper

Chromium (Z = 24) filled by the plain Aufbau rule would be [Ar]4s² 3d⁴. The actual configuration is:

$Cr: [Ar] 4s 1 3d 5$

Copper (Z = 29) would be [Ar]4s² 3d⁹; the actual configuration is:

$Cu: [Ar] 4s 1 3d 10$

In both cases the d subshell gains one electron to reach a half-filled (d⁵) or fully filled (d¹⁰) state, and the 4s loses one. This is a real, verified experimental result — not a rounding error — and it is why these two elements are the classic exceptions you must remember on exams.

Key takeaways

Three ground-state rules: Aufbau (lowest energy first), Pauli (max 2 electrons/orbital, opposite spins), Hund (fill degenerate orbitals singly first).
Filling order to know cold: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p.
Subshell capacities: s = 2, p = 6, d = 10, f = 14.
4s fills before 3d for neutral atoms (but 3d electrons are removed first when transition metals ionize).
Shorthand: use the preceding noble gas in brackets.
Valence electrons = outermost shell electrons; they determine chemical behavior and group placement.
Exceptions to memorize: Cr = [Ar]4s¹3d⁵, Cu = [Ar]4s¹3d¹⁰.
The sum of superscripts must equal the atomic number Z — always check by counting.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

State the three rules for filling orbitals in the ground state.
Show answer
Aufbau (fill lowest energy first), Pauli exclusion (≤2 electrons per orbital, opposite spins), Hund's rule (fill degenerate orbitals singly before pairing).
Write the full electron configuration of sulfur (Z = 16).
Show answer
1s² 2s² 2p⁶ 3s² 3p⁴ (check: 2+2+6+2+4 = 16).
Write the noble-gas shorthand configuration of titanium (Z = 22).
Show answer
[Ar] 4s² 3d² (check: 18+2+2 = 22).
Why is chromium's configuration [Ar]4s¹3d⁵ instead of [Ar]4s²3d⁴?
Show answer
A half-filled 3d subshell (d⁵) is extra stable, so one 4s electron shifts into 3d.
How many electrons can the d subshell hold, and why?
Show answer
Ten electrons: five d orbitals × two electrons each (from quantum numbers l = 2, m_l = -2…+2).
Which electrons are lost first when a transition metal like iron forms a cation?
Show answer
The 4s electrons are lost first, even though 3d filled second; Fe²⁺ is [Ar]3d⁶.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

electron configuration: A listing of how electrons occupy subshells
Aufbau principle: Electrons fill lowest-energy orbitals first
Pauli exclusion principle: At most two electrons per orbital, opposite spins
Hund's rule: Fill degenerate orbitals singly before pairing
valence electrons: Electrons in the outermost shell
noble-gas shorthand: Configuration with inner core replaced by a noble gas symbol
degenerate orbitals: Orbitals of equal energy within a subshell

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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