What an Electronic Configuration Is
The distribution of electrons into the orbitals of an atom is called its electronic configuration. The ground-state configuration of any element always corresponds to the arrangement of lowest total electronic energy. If the basic rules that govern the filling of orbitals are kept in mind, the configuration of any atom can be written very easily.
Three rules, treated in detail in the sibling note on Aufbau, Pauli and Hund rules, jointly fix the ground-state arrangement: the aufbau principle (orbitals are filled in order of increasing energy), the Pauli exclusion principle (no two electrons in an atom share all four quantum numbers, so an orbital holds at most two electrons of opposite spin), and Hund's rule of maximum multiplicity (degenerate orbitals are singly occupied with parallel spins before any pairing begins). This page assumes those rules and applies them element by element.
The Two Notations: spdf and Orbital Diagram
A configuration is represented in two complementary ways. In the $s^a\,p^b\,d^c$ notation, each subshell is written as its letter symbol with the number of electrons it contains as a superscript, and the principal quantum number written in front of the subshell to distinguish the same subshell in different shells. In the orbital diagram, each orbital of a subshell is drawn as a box and each electron as an arrow — ↑ for one spin, ↓ for the opposite. The orbital diagram carries more information: it displays all four quantum numbers, including the spin that the bare $s^a p^b$ string omits.
Ground-state nitrogen, $\ce{N}$ ($Z = 7$): $1s^2\,2s^2\,2p^3$
1s ⇅ 2s ⇅ 2p ↑↑↑
By Hund's rule the three 2p electrons occupy the three degenerate 2p orbitals singly with parallel spins — they do not pair up. This is exactly why NEET 2018 marked the configuration written as $2p^1_x\,2p^1_y\,2p^1_z$ (paired notation issues aside) as a point of caution: in degenerate orbitals all unpaired electrons must show the same spin.
The Order of Filling and the 2n² Rule
Orbital energies do not increase in a single tidy sequence of principal quantum number; effective nuclear charge affects different orbital types unequally. There is therefore no ordering universally correct for every atom, but the following sequence is extremely useful and accurate for the placement of outer valence electrons:
$1s,\ 2s,\ 2p,\ 3s,\ 3p,\ 4s,\ 3d,\ 4p,\ 5s,\ 4d,\ 5p,\ 6s,\ 4f,\ 5d,\ 6p,\ 7s\ldots$
The valence electron of potassium, for instance, must choose between 3d and 4s, and as this sequence predicts, it is found in 4s because 4s lies lower in energy than 3d at that point. The sequence should be treated as a reliable guide rather than an absolute law — orbitals close in energy can swap order, and exceptions do occur.
The Pauli principle fixes the capacity of each level. A subshell holds twice as many electrons as it has orbitals: $s$ holds 2, $p$ holds 6, $d$ holds 10, $f$ holds 14. Summed over a shell, the maximum number of electrons in the shell of principal quantum number $n$ is:
$\text{Maximum electrons in shell } n = 2n^2$
| Shell (n) | Subshells present | Orbitals | Max electrons = $2n^2$ |
|---|---|---|---|
| 1 (K) | 1s | 1 | $2(1)^2 = 2$ |
| 2 (L) | 2s, 2p | 4 | $2(2)^2 = 8$ |
| 3 (M) | 3s, 3p, 3d | 9 | $2(3)^2 = 18$ |
| 4 (N) | 4s, 4p, 4d, 4f | 16 | $2(4)^2 = 32$ |
Building Up: Hydrogen to Calcium
Hydrogen has one electron, which goes into the lowest-energy orbital, 1s, giving $1s^1$. The second electron of helium also enters 1s with opposite spin, completing it as $1s^2$. Lithium's third electron is barred from 1s by the Pauli principle and takes the next orbital, 2s, giving $1s^2\,2s^1$; beryllium completes 2s as $1s^2\,2s^2$. From boron to neon the three 2p orbitals fill progressively, ending at neon, $1s^2\,2s^2\,2p^6$.
Sodium to argon repeat the lithium-to-neon pattern one shell out, filling 3s and 3p. At potassium and calcium the 4s orbital — lower in energy than 3d — takes one and two electrons respectively. A new pattern then begins at scandium: 3d, now lower than 4p, fills across the ten elements scandium to zinc.
| Z | Element | Full configuration | Noble-gas core notation |
|---|---|---|---|
| 1 | H | 1s¹ | 1s¹ |
| 2 | He | 1s² | 1s² |
| 3 | Li | 1s² 2s¹ | [He] 2s¹ |
| 4 | Be | 1s² 2s² | [He] 2s² |
| 5 | B | 1s² 2s² 2p¹ | [He] 2s² 2p¹ |
| 6 | C | 1s² 2s² 2p² | [He] 2s² 2p² |
| 7 | N | 1s² 2s² 2p³ | [He] 2s² 2p³ |
| 8 | O | 1s² 2s² 2p⁴ | [He] 2s² 2p⁴ |
| 9 | F | 1s² 2s² 2p⁵ | [He] 2s² 2p⁵ |
| 10 | Ne | 1s² 2s² 2p⁶ | [He] 2s² 2p⁶ |
| 11 | Na | 1s² 2s² 2p⁶ 3s¹ | [Ne] 3s¹ |
| 12 | Mg | 1s² 2s² 2p⁶ 3s² | [Ne] 3s² |
| 13 | Al | … 3s² 3p¹ | [Ne] 3s² 3p¹ |
| 14 | Si | … 3s² 3p² | [Ne] 3s² 3p² |
| 15 | P | … 3s² 3p³ | [Ne] 3s² 3p³ |
| 16 | S | … 3s² 3p⁴ | [Ne] 3s² 3p⁴ |
| 17 | Cl | … 3s² 3p⁵ | [Ne] 3s² 3p⁵ |
| 18 | Ar | … 3s² 3p⁶ | [Ne] 3s² 3p⁶ |
| 19 | K | … 3p⁶ 4s¹ | [Ar] 4s¹ |
| 20 | Ca | … 3p⁶ 4s² | [Ar] 4s² |
| 21 | Sc | [Ar] 3d¹ 4s² | [Ar] 3d¹ 4s² |
| 22 | Ti | [Ar] 3d² 4s² | [Ar] 3d² 4s² |
| 23 | V | [Ar] 3d³ 4s² | [Ar] 3d³ 4s² |
| 24 | Cr * | [Ar] 3d⁵ 4s¹ | [Ar] 3d⁵ 4s¹ |
| 25 | Mn | [Ar] 3d⁵ 4s² | [Ar] 3d⁵ 4s² |
| 26 | Fe | [Ar] 3d⁶ 4s² | [Ar] 3d⁶ 4s² |
| 27 | Co | [Ar] 3d⁷ 4s² | [Ar] 3d⁷ 4s² |
| 28 | Ni | [Ar] 3d⁸ 4s² | [Ar] 3d⁸ 4s² |
| 29 | Cu * | [Ar] 3d¹⁰ 4s¹ | [Ar] 3d¹⁰ 4s¹ |
| 30 | Zn | [Ar] 3d¹⁰ 4s² | [Ar] 3d¹⁰ 4s² |
The asterisked entries, chromium and copper, are the two exceptional configurations among the first thirty elements and are addressed in their own section below. Beyond zinc, the 4p orbital fills from gallium to krypton, after which the 5s, 4d and 5p pattern repeats the 4s, 3d and 4p story, followed by 6s, then the 4f and 5d filling from lanthanum to mercury.
Noble-Gas Core (Condensed) Notation
Writing $1s^2\,2s^2\,2p^6\,3s^1$ for sodium each time is wasteful, because the first ten electrons are simply the neon configuration. The noble-gas core notation replaces the inner completely filled shells by the symbol of the preceding noble gas in square brackets, leaving only the outer electrons explicit: sodium becomes $\ce{[Ne]}\,3s^1$, and the elements from sodium to argon run from $\ce{[Ne]}\,3s^1$ to $\ce{[Ne]}\,3s^2\,3p^6$.
This notation also names two categories of electron precisely. Electrons in the completely filled inner shells — those folded into the bracketed core — are the core electrons. Electrons added to the shell with the highest principal quantum number are the valence electrons, and it is these that govern chemical behaviour. In neon-cored sodium, the $\ce{[Ne]}$ ten are core electrons and the single $3s$ electron is the valence electron.
Every step here rests on three filling rules. If the order or Hund's pairing feels shaky, revisit Aufbau, Pauli and Hund's rules first.
The Chromium and Copper Exceptions
Read straight off the aufbau order, chromium ($Z = 24$) and copper ($Z = 29$) would be expected as $\ce{[Ar]}\,3d^4\,4s^2$ and $\ce{[Ar]}\,3d^9\,4s^2$. Spectroscopy shows otherwise: chromium is $\ce{[Ar]}\,3d^5\,4s^1$ and copper is $\ce{[Ar]}\,3d^{10}\,4s^1$. In each case one electron has shifted out of 4s into 3d.
The reason is the very small energy gap between the 4s and 3d subshells. When the gap is small, an electron will move from the lower subshell (4s) to the slightly higher one (3d) if doing so produces an exactly half-filled or fully filled 3d set. For chromium that shift produces a half-filled $3d^5$; for copper it produces a fully filled $3d^{10}$. Configurations of the form $p^3,\,p^6,\,d^5,\,d^{10},\,f^7,\,f^{14}$ — half-filled or fully filled — carry extra stability, so the shifted arrangement is the true ground state.
Expected vs actual valence configuration of chromium
Expected ($3d^4\,4s^2$):
3d
↑↑↑↑
4s
⇅
Actual ($3d^5\,4s^1$):
3d
↑↑↑↑↑
4s
↑
Moving one 4s electron into 3d makes both subshells exactly half-filled — five parallel-spin 3d electrons plus one 4s — a symmetric, low-energy arrangement. The same logic gives copper $3d^{10}\,4s^1$.
Why Half-Filled and Fully Filled Subshells Are Stable
The extra stability of completely filled and exactly half-filled subshells, which drives the chromium and copper exceptions, rests on two physical effects set out in NCERT Section 2.6.7.
| Cause | What it means | Effect on energy |
|---|---|---|
| Symmetrical distribution | Electrons of equal energy but different spatial distribution are spread symmetrically, so they shield one another only weakly and feel a stronger nuclear pull. | Lower energy → more stable |
| Exchange energy | Two or more parallel-spin electrons in degenerate orbitals can exchange positions; the energy released is the exchange energy. The number of possible exchanges is maximum when a subshell is half-filled or fully filled. | Maximum exchange energy → maximum stability |
Taken together, the extra stability of half-filled and fully filled subshells comes from three linked factors: relatively small shielding, smaller coulombic repulsion, and larger exchange energy. Exchange energy is also the deeper reason behind Hund's rule — electrons entering degenerate orbitals adopt parallel spins precisely to maximise the number of stabilising exchanges.
Configurations of Ions
For a cation, write the neutral-atom configuration first, then remove electrons from the orbital with the highest principal quantum number $n$. This matters most for the transition metals, where 4s fills before 3d but is also removed before 3d. For an anion, simply add the extra electrons to the next available orbital in the aufbau order.
Write the configurations of $\ce{Fe^2+}$ and $\ce{Fe^3+}$ from neutral $\ce{Fe}$ ($Z = 26$).
Neutral iron: $\ce{[Ar]}\,3d^6\,4s^2$.
$\ce{Fe^2+}$: remove the two 4s electrons (highest $n$, not 3d) → $\ce{[Ar]}\,3d^6$. It is not $\ce{[Ar]}\,3d^4\,4s^2$.
$\ce{Fe^3+}$: remove one more, now from 3d → $\ce{[Ar]}\,3d^5$, a stable half-filled set. This is exactly why $\ce{Fe^3+}$ (23 electrons) is isoelectronic with $\ce{Mn^2+}$ (also 23 electrons), a pairing NEET has tested directly.
Although 4s fills before 3d on the way up, it empties first on the way down: the two 4s electrons (highest $n$) are removed before any 3d electron, giving $\ce{Fe^2+}=\ce{[Ar]}\,3d^6$. Stripping a third electron then comes from 3d, leaving the half-filled $\ce{[Ar]}\,3d^5$ of $\ce{Fe^3+}$.
Filling order is not removal order
4s fills before 3d during aufbau, which tempts students to also remove 3d before 4s when forming cations. That is wrong. Electrons are stripped from the highest principal quantum number first, so the 4s electrons leave before any 3d electron does.
Cation rule: take electrons out of the highest-$n$ orbital first. $\ce{Fe^2+}$ is $\ce{[Ar]}\,3d^6$, never $\ce{[Ar]}\,3d^4\,4s^2$.
Configuration and the Blocks of the Periodic Table
The subshell that receives the last (differentiating) electron defines the block to which an element belongs, which is why the periodic table can be read directly as a chart of electronic configurations. This block relationship is developed further in Classification of Elements and Periodicity.
Knowing the configuration is not an end in itself. The modern account of chemistry rests almost entirely on electronic distribution: it explains why atoms combine into molecules, why some elements are metals and others non-metals, and why helium and argon are inert while the halogens are highly reactive. The closed, fully filled shells of the noble gases place them at the foot of the p-block and account for their unwillingness to react, whereas a single electron short of, or beyond, such a closed shell makes the halogens and alkali metals chemically eager. None of these questions could be answered in the Daltonian picture of the atom; all of them fall out naturally once the configuration is known.
The last electron in an s-block element enters an $ns$ orbital; in a p-block element, an $np$ orbital; in a d-block (transition) element, an $(n-1)d$ orbital; in an f-block element, an $(n-2)f$ orbital. Scandium to zinc therefore form the 3d series of the d-block.
Electronic Configurations in one screen
- Configuration = distribution of electrons in orbitals; ground state = lowest total energy. Written as $s^a p^b d^c$ or as an orbital (box) diagram, the latter showing spin.
- Fill in order $1s,2s,2p,3s,3p,4s,3d,4p,5s\ldots$; a shell of quantum number $n$ holds at most $2n^2$ electrons.
- Use noble-gas cores to condense: $\ce{Na} = \ce{[Ne]}\,3s^1$. Core electrons are inner; valence electrons sit in the highest-$n$ shell.
- Cr is $\ce{[Ar]}\,3d^5\,4s^1$ and Cu is $\ce{[Ar]}\,3d^{10}\,4s^1$ — half-filled and fully filled subshells gain stability from symmetry and exchange energy.
- For cations, remove electrons from the highest-$n$ orbital first: $\ce{Fe^2+}=\ce{[Ar]}\,3d^6$, $\ce{Fe^3+}=\ce{[Ar]}\,3d^5$.
- The last-filled subshell fixes the block: s, p, d or f.