Working out electron configuration by hand means memorising orbital fill order, applying three separate rules, and checking for exceptions. Select any of the 118 elements below and get the full notation, noble gas shorthand, and valence electron count instantly. The Ion tab handles cations and anions too.
✓Aufbau Principle • Hund's Rule • Pauli Exclusion • IUPAC • NIST • April 2026
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Full Electron Configuration
—
select an element above
Noble Gas Notation
—
condensed shorthand
Valence Electrons
—
outer shell
Atomic Number
—
protons / electrons
Total Electrons
—
in this species
⚠️ Educational reference. Configurations follow standard undergraduate rules (Aufbau, Hund's rule, Pauli exclusion) with common ground-state exceptions applied. For advanced spectroscopy, publication work, or excited states, verify against the NIST Atomic Spectra Database.
Sources & Methodology
✅ Configurations verified from IUPAC periodic table and NIST Atomic Spectra Database. Exception elements cross-checked against primary literature. Ion electron removal order follows standard undergraduate convention: ns before (n−1)d for transition metals.
Primary reference for element data, standard ground-state configurations, and group conventions for valence electron counting. Used for all 118 elements including synthetic and transactinide elements.
Used to verify ground-state electron configurations for all Aufbau exception elements including Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Ce, Gd, Pt, Au, and actinide series exceptions.
Reference for Aufbau principle, Hund's rule, Pauli exclusion principle, noble gas notation methodology, and ion electron configuration removal order rules used in content sections.
What Is Electron Configuration and How Do You Calculate It?
If you're about to sit a chemistry exam and someone asks you for the electron configuration of iron, the last thing you want to do is guess. Iron has 26 electrons. You need to know exactly where each one goes — and one wrong subshell label fails the whole answer. That's what this section is for.
The Worked Example First — Oxygen
Oxygen has atomic number 8, meaning 8 electrons to place. Fill from the lowest energy orbital upwards:
1s orbital holds 2 electrons maximum → fills completely: 1s²
2s orbital holds 2 electrons maximum → fills completely: 2s²
2p orbital holds 6 electrons maximum → 4 electrons remain: 2p⁴
Oxygen's full configuration: 1s² 2s² 2p⁴. Noble gas notation: [He] 2s² 2p⁴. Six valence electrons. Two short of a full outer shell — which is exactly why oxygen forms double bonds and O²− ions so readily. The configuration explains the chemistry directly.
Subshell capacity: s = 2 electrons (1 orbital) • p = 6 electrons (3 orbitals) • d = 10 electrons (5 orbitals) • f = 14 electrons (7 orbitals) Note: 4s fills before 3d even though 3 is the lower principal quantum number. This is because 4s sits at lower energy than 3d in neutral ground-state atoms.
The Three Rules You Must Apply
Aufbau principle: Fill from lowest energy upwards. The ordering above isn't alphabetical — 4s fills before 3d because 4s has lower energy in neutral atoms. Students who write 3d before 4s for potassium and calcium are wrong. K is [Ar] 4s¹. Ca is [Ar] 4s². The 3d subshell doesn't appear until scandium.
Hund's rule: Within any subshell, put one electron in each orbital before pairing any. Nitrogen has three 2p electrons — they go one per orbital (↑ ↑ ↑), not two in one (↑↓ ↑ ·). This minimises electron-electron repulsion. All unpaired electrons in a subshell have the same spin direction.
Pauli exclusion principle: No two electrons in the same atom can have identical quantum numbers. In practice: each orbital holds exactly 2 electrons, and they must have opposite spins. This is why subshell capacities are fixed at 2, 6, 10, and 14.
How Noble Gas Notation Works
Iron's full notation — 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ — contains 26 electrons worth of information. But the first 18 (the argon core) don't participate in chemistry. Noble gas notation writes those 18 as [Ar] and highlights only the reactive outer electrons: [Ar] 4s² 3d⁶. Everything interesting about iron's chemistry lives in those 8 electrons.
To convert: find the nearest preceding noble gas with a lower atomic number. For iron (Z=26), that's argon (Z=18). Write [Ar] then continue the fill sequence from there.
💡 What the configuration tells you immediately: Which block the element sits in (last subshell type filled = s, p, d, or f block). How many valence electrons it has. Whether it'll form positive or negative ions and what charge. Whether it's paramagnetic (has unpaired electrons) or diamagnetic (all paired). A single configuration string carries all of this.
Electron Configuration Reference Table — First 30 Elements
The table below covers the first 30 elements — the range tested in most general chemistry courses. Elements with ⚠ exception break the Aufbau prediction and have actual configurations that differ from what the filling order would give you.
Z
Element
Full Configuration
Noble Gas Notation
Valence e−
1
H
1s¹
1s¹
1
2
He
1s²
1s²
2
3
Li
1s² 2s¹
[He] 2s¹
1
4
Be
1s² 2s²
[He] 2s²
2
5
B
1s² 2s² 2p¹
[He] 2s² 2p¹
3
6
C
1s² 2s² 2p²
[He] 2s² 2p²
4
7
N
1s² 2s² 2p³
[He] 2s² 2p³
5
8
O
1s² 2s² 2p⁴
[He] 2s² 2p⁴
6
9
F
1s² 2s² 2p⁵
[He] 2s² 2p⁵
7
10
Ne
1s² 2s² 2p⁶
[He] 2s² 2p⁶
8
11
Na
1s² 2s² 2p⁶ 3s¹
[Ne] 3s¹
1
12
Mg
1s² 2s² 2p⁶ 3s²
[Ne] 3s²
2
13
Al
1s² 2s² 2p⁶ 3s² 3p¹
[Ne] 3s² 3p¹
3
14
Si
1s² 2s² 2p⁶ 3s² 3p²
[Ne] 3s² 3p²
4
15
P
1s² 2s² 2p⁶ 3s² 3p³
[Ne] 3s² 3p³
5
16
S
1s² 2s² 2p⁶ 3s² 3p⁴
[Ne] 3s² 3p⁴
6
17
Cl
1s² 2s² 2p⁶ 3s² 3p⁵
[Ne] 3s² 3p⁵
7
18
Ar
1s² 2s² 2p⁶ 3s² 3p⁶
[Ne] 3s² 3p⁶
8
19
K
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
[Ar] 4s¹
1
20
Ca
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
[Ar] 4s²
2
21
Sc
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
[Ar] 4s² 3d¹
3
22
Ti
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²
[Ar] 4s² 3d²
4
23
V
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
[Ar] 4s² 3d³
5
24
Cr ⚠ exception
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵
[Ar] 4s¹ 3d⁵
6
25
Mn
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵
[Ar] 4s² 3d⁵
7
26
Fe
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
[Ar] 4s² 3d⁶
8
27
Co
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷
[Ar] 4s² 3d⁷
9
28
Ni
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸
[Ar] 4s² 3d⁸
10
29
Cu ⚠ exception
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
[Ar] 4s¹ 3d¹⁰
11
30
Zn
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰
[Ar] 4s² 3d¹⁰
12
Elements 24 (Cr) and 29 (Cu) are the most commonly tested Aufbau exceptions in general chemistry. Both have 4s¹ instead of the predicted 4s² because half-filled or fully filled d subshells provide extra stability through exchange energy.
Ion Configuration — The Rule Most Students Get Wrong
Most students learn that 4s fills before 3d during neutral atom configuration. They then assume that when forming a cation you remove 3d electrons first. That's wrong — and it costs exam marks.
The correct rule: remove electrons from the highest principal quantum number (n) first. For transition metals, this means ns before (n−1)d electrons leave — even though 4s filled before 3d on the way in.
Species
Configuration
Electrons Removed
From Which Subshell
Fe (neutral)
[Ar] 4s² 3d⁶
—
—
Fe²+
[Ar] 3d⁶
2 electrons
4s first (not 3d)
Fe³+
[Ar] 3d⁵
3 electrons
4s² + one 3d
Cu (neutral)
[Ar] 4s¹ 3d¹⁰
—
—
Cu+
[Ar] 3d¹⁰
1 electron
4s¹
O (neutral)
[He] 2s² 2p⁴
—
—
O²−
[He] 2s² 2p⁶ = [Ne]
+2 electrons added
2p gains 2
Cl (neutral)
[Ne] 3s² 3p⁵
—
—
Cl−
[Ne] 3s² 3p⁶ = [Ar]
+1 electron added
3p gains 1
Aufbau Exceptions, Exam Mistakes, and Magnetic Properties
Three Mistakes Students Almost Always Make
Mistake 1 — Treating Aufbau as universal for d-block elements. It isn't. Chromium and copper are the two exceptions every student must know. Cr predicted: [Ar] 4s² 3d⁴. Cr actual: [Ar] 4s¹ 3d⁵. Cu predicted: [Ar] 4s² 3d⁹. Cu actual: [Ar] 4s¹ 3d¹⁰. The reason both break the rule: half-filled (d&sup5;) and fully filled (d¹⁰) subshells are energetically more stable than a partly filled subshell plus a fully filled s orbital. Nature picks the lower-energy arrangement.
Mistake 2 — Wrong d-block notation for period 4. In period 4, after filling 4s, d electrons go into 3d — not 4d. Scandium's configuration is [Ar] 4s² 3d¹, never [Ar] 4s² 4d¹. The d subshell principal quantum number is always one less than the period number: period 4 → 3d, period 5 → 4d, period 6 → 5d. Getting this wrong makes half of transition metal configurations incorrect.
Mistake 3 — Removing 3d electrons when forming transition metal cations. The correct order removes 4s first, then 3d. Fe²+ is [Ar] 3d⁶, not [Ar] 4s² 3d⁴. Writing [Ar] 4s² 3d⁴ for Fe²+ is a guaranteed wrong answer on any chemistry exam.
⚠️ The single most common exam error: Writing chromium as [Ar] 4s² 3d⁴. Every mark scheme flags this as incorrect. The verified answer is [Ar] 4s¹ 3d⁵. Molybdenum (Mo, Z=42) makes the identical exception in period 5: [Kr] 5s¹ 4d⁵, not [Kr] 5s² 4d⁴.
Complete Aufbau Exception Reference
Element (Z)
Predicted (Aufbau)
Actual Configuration
Reason
Cr (24)
[Ar] 4s² 3d⁴
[Ar] 4s¹ 3d⁵
Half-filled 3d⁵ stability
Cu (29)
[Ar] 4s² 3d⁹
[Ar] 4s¹ 3d¹⁰
Fully filled 3d¹⁰ stability
Nb (41)
[Kr] 5s² 4d³
[Kr] 5s¹ 4d⁴
d&sup4; near-half-fill stability
Mo (42)
[Kr] 5s² 4d⁴
[Kr] 5s¹ 4d⁵
Half-filled 4d⁵ stability
Pd (46)
[Kr] 5s² 4d⁸
[Kr] 4d¹⁰
Fully filled 4d, empty 5s
Ag (47)
[Kr] 5s² 4d⁹
[Kr] 5s¹ 4d¹⁰
Fully filled 4d¹⁰ stability
Pt (78)
[Xe] 6s² 4f¹⁴ 5d⁸
[Xe] 6s¹ 4f¹⁴ 5d⁹
Relativistic effects + 5d stability
Au (79)
[Xe] 6s² 4f¹⁴ 5d⁹
[Xe] 6s¹ 4f¹⁴ 5d¹⁰
Fully filled 5d¹⁰ stability
Reading Magnetic Properties from Configuration
Count unpaired electrons. That's it. One unpaired electron = paramagnetic. Zero unpaired electrons = diamagnetic. Iron ([Ar] 4s² 3d⁶) has 4 unpaired 3d electrons — strongly paramagnetic, which is why iron responds to magnets. Zinc ([Ar] 4s² 3d¹⁰) has every electron paired — diamagnetic, despite sitting directly below iron on the periodic table and having a similar-looking configuration. Manganese ([Ar] 4s² 3d⁵) has 5 unpaired electrons — the highest of any first-row transition metal.
💡 5-point check before submitting a configuration: (1) Is this Cr, Cu, Mo, Ag, Pd, Au, or Pt? If yes, override Aufbau. (2) For period 4 d-block, is it 3d not 4d? (3) For an ion, did you remove ns before (n−1)d? (4) Does the electron count match Z minus charge? (5) Does the noble gas symbol represent the last noble gas with Z less than this element?
Frequently Asked Questions
Electron configuration is the distribution of electrons across an atom's atomic orbitals. It uses subshell notation (1s, 2s, 2p, 3d, 4f etc.) with superscripts showing how many electrons are in each. Oxygen's configuration 1s² 2s² 2p⁴ tells you it has 8 electrons across three subshells — 2 in 1s, 2 in 2s, and 4 in 2p — and that it's 2 electrons short of a full outer shell. That single fact explains most of oxygen's chemistry.
Noble gas notation replaces the inner core electrons with the symbol of the nearest preceding noble gas in square brackets. Sodium's full configuration 1s² 2s² 2p⁶ 3s¹ becomes [Ne] 3s¹ because neon's configuration is 1s² 2s² 2p⁶. Shorter, cleaner, and puts the chemically active outer electrons front and centre — the inner core doesn't bond or react anyway.
Valence electrons sit in the outermost principal energy level (highest n value). For main group (s and p block) elements, count all electrons in the highest n shell. Oxygen (1s² 2s² 2p⁴) has 6 valence electrons at n=2. For transition metals the situation's more complex — both outermost s and d electrons participate in bonding, which is why iron forms both Fe²+ and Fe³+ ions without the atom falling apart.
[Ar] 4s¹ 3d⁵ is lower in energy than [Ar] 4s² 3d⁴. A half-filled 3d subshell — five electrons, one per d orbital, all same spin — gains extra stability from exchange energy, the quantum mechanical energy released when electrons with parallel spins can exchange positions. With both 4s and 3d half-filled, chromium achieves double stability. The Aufbau prediction of 4s² 3d⁴ is simply less stable, so nature doesn't use it.
A completely filled 3d¹⁰ subshell has exceptional stability — maximum symmetry, maximum exchange energy, all orbitals paired. The energy gained by completely filling 3d outweighs the cost of leaving 4s with only one electron. So one electron promotes from 4s to 3d. Silver ([Kr] 5s¹ 4d¹⁰) and gold ([Xe] 6s¹ 4f¹⁴ 5d¹⁰) both follow the same pattern — all three are in group 11 and all three are Aufbau exceptions.
Start with the neutral atom. For cations, remove electrons from the highest principal quantum number first. For transition metals remove ns before (n−1)d. Iron (Fe): [Ar] 4s² 3d⁶. Fe²+: remove 4s² → [Ar] 3d⁶. Fe³+: remove 4s² plus one 3d → [Ar] 3d⁵. For anions add electrons to the next available orbital: Cl− adds one to 3p giving [Ar]. O²− adds two to 2p giving [Ne].
The Aufbau principle (from German "aufbauen" — to build up) states electrons fill atomic orbitals in order of increasing energy. The fill order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. The order looks strange because 4s fills before 3d — that's because 4s has lower energy than 3d in neutral ground-state atoms. The principle works for most elements but has exceptions in transition metals and heavier elements where subshell stability effects override the standard sequence.
Hund's rule states electrons occupy each orbital of a subshell singly before pairing begins, and all singly-occupied orbitals in a subshell have the same spin. Nitrogen's three 2p electrons each take a separate orbital rather than two in one and one in another. Spreading electrons across separate orbitals reduces repulsion and lowers energy. Only after every orbital in a subshell has one electron do pairs start forming.
The Pauli exclusion principle states no two electrons in the same atom can have an identical set of four quantum numbers (n, l, mₗ, mₛ). Since two electrons in the same orbital share the same n, l, and mₗ values, they must have opposite spins. This limits every orbital to a maximum of two electrons — one spin-up and one spin-down. It's why subshell capacities are exactly 2 (s), 6 (p), 10 (d), and 14 (f).
1s² 2s² 2p⁴. Noble gas notation: [He] 2s² 2p⁴. Six valence electrons — two short of a full outer shell. That's why oxygen is so reactive: it desperately wants two more electrons, which it gets either by forming two covalent bonds (water, CO&sub2;) or by taking them as O²− in ionic compounds like rust and metal oxides.
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. Noble gas notation: [Ar] 4s² 3d⁶. The 3d⁶ subshell has 4 unpaired electrons, making iron strongly paramagnetic. Fe²+ (common in iron(II) compounds) removes 4s² first: [Ar] 3d⁶. Fe³+ (most stable iron oxidation state) removes 4s² plus one 3d: [Ar] 3d⁵ — a half-filled d subshell, which is notably stable.
Count unpaired electrons. Iron ([Ar] 4s² 3d⁶) has 4 unpaired d electrons — strongly paramagnetic. Zinc ([Ar] 4s² 3d¹⁰) has every electron paired — diamagnetic, despite being in the same row as iron. Manganese ([Ar] 4s² 3d⁵) has 5 unpaired electrons — the maximum for first-row transition metals, making it the most paramagnetic of the group. One unpaired electron gives weak paramagnetism; 5 gives strong.
For general chemistry: Cr (24): [Ar] 4s¹ 3d⁵. Cu (29): [Ar] 4s¹ 3d¹⁰. Mo (42): [Kr] 5s¹ 4d⁵. Ag (47): [Kr] 5s¹ 4d¹⁰. Pd (46): [Kr] 4d¹⁰ (no 5s at all). Au (79): [Xe] 6s¹ 4f¹⁴ 5d¹⁰. Pt (78): [Xe] 6s¹ 4f¹⁴ 5d⁹. The pattern: exceptions cluster in groups 6 and 11 (half-filled or fully filled d subshells) and become more common in heavier elements where relativistic effects reduce energy differences between subshells.
4s has lower energy than 3d in neutral ground-state atoms. Despite having a higher principal quantum number (4 vs 3), the 4s electrons experience less electron-electron repulsion than 3d electrons, which are more compact and more shielded. So 4s fills first. Confusingly, once a transition metal loses electrons to form a cation, the energy ordering reverses — 4s becomes higher energy than 3d in ionic form. That's why 4s electrons leave first when forming cations.
Full configuration lists every subshell from 1s onwards. Iron full: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. Condensed (noble gas) notation: [Ar] 4s² 3d⁶. Both correct. Condensed is preferred in most chemistry because it highlights only the electrons that actually matter for bonding and reactivity. The [Ar] core never participates in iron's chemistry — only the 4s² 3d⁶ electrons do.