Electron Configuration — Rules, Notation & Examples

Think of an atom's electrons as guests filling a parking structure: lower levels fill first, each spot holds at most two, and identical-level spots fill one car at a time before doubling up. Electron configuration is that parking map — it shows how an atom's electrons are arranged in orbitals, and it explains valence electrons, bonding patterns, magnetism, and periodic trends.

The Three Rules and Their Notation

A configuration like $1s^2 2s^2 2p^6$ has three parts: the number gives the main energy level, the letter gives the subshell ($s$, $p$, $d$, $f$), and the superscript gives the electron count. So $2p^6$ means "six electrons in the $2p$ subshell." You will also meet noble-gas shorthand — sulfur as $[Ne]3s^2 3p^4$, where $[Ne]$ stands for neon's filled $1s^2 2s^2 2p^6$.

Three rules do most of the work:

• Aufbau principle: electrons fill lower-energy orbitals first, giving the common order $1s,\ 2s,\ 2p,\ 3s,\ 3p,\ 4s,\ 3d,\ 4p,\dots$ — practical for classroom problems, not a promise that every atom behaves identically.
• Pauli exclusion principle: one orbital holds at most two electrons, and a shared pair must have opposite spins — so an $s$ subshell holds at most two electrons and a $p$ subshell at most six.
• Hund's rule: among equal-energy orbitals, electrons fill singly before pairing — in a $p$ subshell, all three orbitals get one electron before any gets a second.

Why the Rules Are Built This Way

The three rules are not arbitrary conventions; each reflects a physical tendency. Aufbau filling reflects that systems settle into the lowest available energy. The Pauli limit of two-per-orbital with opposite spins reflects that no two electrons can share the same quantum state. Hund's rule reflects that electrons, being mutually repulsive, spread out to avoid sharing an orbital until forced — pairing costs energy. Read together, they say: go low, cap at two, and spread before pairing.

Worked Example: Sulfur

A neutral sulfur atom (atomic number 16) has 16 electrons. Fill in order:

1. $1s^2$ — 2 electrons.
2. $2s^2$ — total 4.
3. $2p^6$ — total 10.
4. $3s^2$ — total 12.
5. The remaining 4 go into $3p$, giving $3p^4$.

Full configuration:

Shorthand:

The part people trip on is $3p^4$. By Hund's rule the first three $3p$ electrons each take a separate $p$ orbital, and only the fourth pairs with one of them — so sulfur does not start by making two pairs in the $3p$ set.

Practice It Yourself

1. Write the configuration for phosphorus (atomic number 15) and compare with sulfur. Answer check: $1s^2 2s^2 2p^6 3s^2 3p^3$, or $[Ne]3s^2 3p^3$ — phosphorus ends at $3p^3$ (three singly occupied orbitals) while sulfur reaches $3p^4$, so you can watch Hund's rule operate across one step.
2. Write the configuration for the chloride ion $Cl^-$. Answer check: neutral $Cl$ has 17 electrons, so $Cl^-$ has 18, giving $1s^2 2s^2 2p^6 3s^2 3p^6$, or $[Ar]$ — a filled $3p$ subshell.

Common Mistakes

Forgetting to Change the Electron Count for Ions

A neutral atom and its ion differ in electron count. For example, $Cl$ has 17 electrons but $Cl^-$ has 18.

Pairing Too Early in a $p$ Subshell

For $p^2$ or $p^3$, electrons should spread out before pairing. Pairing too early breaks Hund's rule.

Treating the Filling Order as Untouchable

The standard order works for many beginner problems, but some transition-metal cases are exceptions, and ions need extra care — especially when electrons are removed from transition metals.

Not Checking the Total Electron Count

A configuration can look neat and still be wrong if the superscripts do not add up to the correct number of electrons.

A Fast Way to Check Your Answer

Before moving on, ask: (1) Do the superscripts add to the correct number of electrons? (2) Did any orbital get more than two? (3) Did equal-energy orbitals fill singly before pairing? Those three checks catch most beginner errors.

Why It Is Useful

Electron configuration predicts how many valence electrons an atom has, whether a species tends to gain or lose electrons, and whether unpaired electrons are present — which is why it underlies atomic structure, periodic trends, bonding, and magnetism. In more advanced chemistry it also supports spectroscopy and transition-metal chemistry. The notation is simple, but the consequences are broad, and if the configuration is wrong, later reasoning usually is too.