Electron Configuration — Rules, Notation & Examples

Electron configuration shows how an atom's electrons are arranged in orbitals. If you want the fast version, read it as a map of where the electrons go and in what order they fill. That map helps explain valence electrons, bonding patterns, magnetism, and periodic trends.

For most introductory chemistry problems, three rules do most of the work: fill lower-energy orbitals first, put at most two electrons in one orbital, and spread electrons across equal-energy orbitals before pairing them. Those rules are often introduced as the Aufbau principle, the Pauli exclusion principle, and Hund's rule.

What The Notation Means

A configuration such as $1s^2 2s^2 2p^6$ has three parts:

• the number tells you the main energy level
• the letter tells you the subshell, such as $s$, $p$, $d$, or $f$
• the superscript tells you how many electrons are in that subshell

So $2p^6$ means "six electrons in the $2p$ subshell."

You will also see noble-gas shorthand. For example, sulfur can be written as $[Ne]3s^2 3p^4$. The $[Ne]$ part stands for the filled inner configuration of neon, which is $1s^2 2s^2 2p^6$.

The Three Rules That Matter Most

Aufbau Principle

In the usual introductory filling scheme, electrons go into lower-energy orbitals before higher-energy ones. That gives the common order

$$1s,\ 2s,\ 2p,\ 3s,\ 3p,\ 4s,\ 3d,\ 4p,\dots$$

This is a practical rule for many classroom problems, not a promise that every atom behaves the same way in every context.

Pauli Exclusion Principle

One orbital can hold at most two electrons, and if two electrons occupy the same orbital, they must have opposite spins. That is why an $s$ subshell holds at most two electrons and a $p$ subshell holds at most six.

Hund's Rule

If several orbitals have the same energy, electrons fill them one at a time before pairing. In a $p$ subshell, that means the three $p$ orbitals each get one electron before any orbital gets a second one.

Worked Example: Sulfur

A neutral sulfur atom has atomic number 16, so it has 16 electrons.

Fill them in order:

1. $1s^2$ uses 2 electrons.
2. $2s^2$ brings the total to 4.
3. $2p^6$ brings the total to 10.
4. $3s^2$ brings the total to 12.
5. The remaining 4 electrons go into $3p$, so the last part is $3p^4$.

The full electron configuration is

$$1s^2 2s^2 2p^6 3s^2 3p^4$$

The shorthand version is

$$[Ne]3s^2 3p^4$$

The part that often trips people up is $3p^4$. Because of Hund's rule, the first three $3p$ electrons go into separate $p$ orbitals. The fourth electron pairs with one of them. So sulfur does not start by making two pairs in the $3p$ set.

Why Electron Configuration Is Useful

Electron configuration is not just notation to memorize. It helps you predict how many valence electrons an atom has, whether a species is likely to gain or lose electrons, and whether unpaired electrons are present.

That is why the idea appears in atomic structure, periodic trends, bonding, and magnetism. If the configuration is wrong, later reasoning is usually wrong too.

Common Mistakes

Forgetting To Change The Electron Count For Ions

A neutral atom and its ion do not have the same number of electrons. 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. If you pair too early, you break Hund's rule.

Treating The Filling Order As An Untouchable Rule

The standard order works well for many beginner problems, but some transition-metal cases are exceptions. Ions can also require 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.

When This Concept Is Used

Use electron configuration when you need to connect an element's position in the periodic table to its behavior. It is especially useful for valence electrons, common ion formation, magnetic behavior, and introductory bonding questions.

In more advanced chemistry, the same idea also supports topics such as spectroscopy and transition-metal chemistry. The notation is simple, but the consequences are broad.

A Fast Way To Check Your Answer

Before you move on, ask three questions:

1. Do the superscripts add to the correct number of electrons?
2. Did any orbital get more than two electrons?
3. Did equal-energy orbitals fill singly before pairing?

Those three checks catch most beginner errors quickly.

Try A Similar Case

Try writing the configuration for phosphorus, then compare it with sulfur. That one-step comparison is useful because phosphorus ends at $3p^3$ while sulfur ends at $3p^4$, so you can see Hund's rule in action instead of just reading about it.