Electron Configuration Explained: Orbitals, Shells, and the Aufbau Principle
Learn how electrons fill atomic orbitals, how to write electron configurations, and why exceptions like copper and chromium break the rules.
· 8 min read
Why Electron Configuration Matters
An atom's chemical behavior is almost entirely determined by how its electrons are arranged. Electron configuration tells you which electrons are available for bonding, predicts whether an element will be reactive or stable, and explains the periodic trends you see on the table — ionization energy, atomic radius, electronegativity. Before you can understand why sodium reacts violently with water or why noble gases barely react at all, you need to understand where the electrons actually live.
Shells, Subshells, and Orbitals
Electrons don't orbit the nucleus in neat circular paths. Instead, they occupy regions of space called orbitals, each described by four quantum numbers.
Principal quantum number (n) defines the shell. n = 1 is the first shell, closest to the nucleus and lowest in energy. n = 2 is the second shell, and so on. As n increases, electrons are farther from the nucleus and higher in energy.
Angular momentum quantum number (l) defines the subshell within a shell. l can be any integer from 0 to n − 1. Each value of l corresponds to a subshell type:
| l value | Subshell | Shape | |---------|----------|-------| | 0 | s | Spherical | | 1 | p | Dumbbell | | 2 | d | Cloverleaf | | 3 | f | Complex |
Magnetic quantum number (m_l) specifies the orientation of an orbital in space. It ranges from −l to +l. This tells you how many orbitals exist in each subshell:
- s subshell: 1 orbital (holds 2 electrons)
- p subshell: 3 orbitals (holds 6 electrons)
- d subshell: 5 orbitals (holds 10 electrons)
- f subshell: 7 orbitals (holds 14 electrons)
Spin quantum number (m_s) describes the spin of an electron. It can only be +½ (spin-up) or −½ (spin-down).
The Aufbau Principle
Aufbau is German for "building up." The principle states that electrons fill orbitals starting from the lowest available energy level and work upward. The tricky part is that energy levels don't always follow the order you'd expect from the shell number alone.
The Madelung rule (also called the diagonal rule) gives the filling order. A simple way to remember it: add the n and l values for each subshell. Orbitals fill in order of increasing (n + l). When two subshells have the same (n + l) sum, the one with the lower n fills first.
The correct filling order is:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
Notice that 4s fills before 3d. For the 4s subshell, n + l = 4 + 0 = 4. For 3d, n + l = 3 + 2 = 5. So 4s has a lower sum and fills first.
The Pauli Exclusion Principle
No two electrons in the same atom can have identical sets of all four quantum numbers. In practice, this means each orbital can hold a maximum of two electrons, and those two electrons must have opposite spins — one spin-up (+½) and one spin-down (−½). You cannot put three electrons into a single orbital.
Hund's Rule
When electrons fill orbitals of equal energy — called degenerate orbitals — they spread out one per orbital before any orbital gets a second electron. All these singly-occupied orbitals have parallel spins. This arrangement minimizes electron-electron repulsion and produces a lower-energy, more stable atom.
Think of it like seating on a bus: people prefer to sit alone before they sit next to someone.
Writing Electron Configurations
Write the subshells in filling order, with a superscript showing how many electrons occupy each subshell.
Hydrogen (Z = 1): 1 electron
1s¹
Helium (Z = 2): 2 electrons
1s²
Carbon (Z = 6): 6 electrons
1s² 2s² 2p²
The two 2p electrons each occupy separate 2p orbitals with parallel spins (Hund's rule).
Nitrogen (Z = 7): 7 electrons
1s² 2s² 2p³
Each of the three 2p orbitals holds exactly one electron.
Oxygen (Z = 8): 8 electrons
1s² 2s² 2p⁴
Now one 2p orbital must hold two paired electrons while the other two hold one each.
Sodium (Z = 11): 11 electrons
1s² 2s² 2p⁶ 3s¹
The first 10 electrons fill through the second shell, and the 11th electron starts the third shell.
Chlorine (Z = 17): 17 electrons
1s² 2s² 2p⁶ 3s² 3p⁵
One electron short of a full 3p subshell — this is why chlorine is so reactive and eager to gain one electron.
Iron (Z = 26): 26 electrons
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
After filling through 4s, the remaining 6 electrons enter the 3d subshell. By Hund's rule, the first five 3d electrons each occupy a separate orbital; the sixth pairs up in one of them.
Noble Gas Shorthand Notation
Writing out the full configuration for every element gets tedious. The noble gas shorthand replaces the filled inner-shell configuration with the symbol of the preceding noble gas in brackets, then continues from there.
Examples:
- Nitrogen (Z = 7): [He] 2s² 2p³
- Sodium (Z = 11): [Ne] 3s¹
- Chlorine (Z = 17): [Ne] 3s² 3p⁵
- Iron (Z = 26): [Ar] 3d⁶ 4s²
Note that in the shorthand for transition metals like iron, the 3d subshell is written before 4s even though 4s filled first. This reflects the actual energy ordering once the atom is fully assembled — 3d sits lower in energy than 4s in a filled atom.
Exceptions: Copper and Chromium
Two of the most common exceptions to the Aufbau principle are chromium and copper. Based on the filling rules, you would predict:
- Chromium (Z = 24): [Ar] 3d⁴ 4s²
- Copper (Z = 29): [Ar] 3d⁹ 4s²
The actual configurations are:
- Chromium (Z = 24): [Ar] 3d⁵ 4s¹
- Copper (Z = 29): [Ar] 3d¹⁰ 4s¹
One electron from the 4s subshell shifts into the 3d subshell in each case. The reason is extra stability associated with half-filled (d⁵) and completely filled (d¹⁰) d subshells. In chromium, having five singly occupied 3d orbitals (all parallel spins) and a singly occupied 4s creates a particularly stable, symmetrical arrangement. In copper, a completely filled 3d subshell is more stable than a nearly full one. The small energy cost of moving one electron out of 4s is more than recovered by achieving these stable configurations.
These exceptions matter in practice — copper's full d subshell contributes to its excellent electrical conductivity, and chromium's configuration affects its oxidation states and the colors of its compounds.
Connection to the Periodic Table Blocks
The periodic table's structure directly maps onto electron configurations.
s-block (Groups 1–2 and He): The last electron fills an s subshell. Hydrogen and helium are in the s-block, as are all alkali metals and alkaline earth metals.
p-block (Groups 13–18): The last electron fills a p subshell. This block contains the nonmetals, metalloids, and most of the elements with predictable reactivity patterns like the halogens (one electron short of a full p subshell) and noble gases (full p subshell).
d-block (Groups 3–12): The last electron fills a d subshell. These are the transition metals. Their partially filled d orbitals are responsible for variable oxidation states, colored ions, and catalytic activity.
f-block (Lanthanides and Actinides): The last electron fills an f subshell. These elements are pulled out into the two rows below the main table to keep the layout manageable.
Once you know an element's position on the periodic table, you can often write its electron configuration directly without memorizing the diagonal rule — the block tells you which subshell is being filled, and the element's position within that block tells you how many electrons are in it.
Practice while it's fresh
Test what you just learned on the interactive periodic table.