Electron Configuration Practice: Rules, Examples & Quiz

There are two ways students approach electron configuration practice. Some memorize the filling order by rote and hope for the best on test day. Others understand the three underlying rules so well that they can reconstruct any configuration from scratch, even for transition metals and ions, without a cheat sheet. One approach takes minutes to learn but fails under pressure. The other takes an hour to master and holds up for life.

In my experience teaching chemistry to thousands of students, the second approach wins every time. This guide gives you both the rules and the practice to make them automatic.

Electron configuration practice involves placing an element’s electrons into subshells (s, p, d, f) in the correct order using three rules: the Aufbau principle (fill lowest energy first), the Pauli Exclusion Principle (max 2 electrons per orbital, opposite spins), and Hund’s rule (half-fill before pairing). Mastering these three rules lets you write the configuration for any element on the periodic table.

After reading this guide, you will be able to:

• State and apply all three electron configuration rules correctly
• Write full and noble-gas shorthand configurations for any main-group or transition-metal element
• Handle common exceptions like Cr and Cu without confusion
• Write configurations for cations and anions
• Spot and fix the five most common student mistakes

Two Approaches to Electron Configuration

The rote-memorization approach gives you a list of configurations to copy. It breaks the moment an exam asks about an element you did not memorize, or asks you to write the configuration of an ion or an excited state. The rule-based approach takes slightly longer to learn, but it generates the correct answer for every element, every ion, and every exception.

Here is a direct comparison of the two approaches so you can see exactly what you are choosing between:

The rest of this guide is built entirely around the rule-based approach. Let’s start with the rules themselves.

The Three Rules: A Comparison Table

All electron configuration practice reduces to three rules. Each rule answers a different question about where electrons go. Understanding what each rule does — and what it does NOT do — prevents most student errors.

What Is Electron Configuration?

An electron configuration is a notation that shows how electrons are distributed among the energy subshells of an atom. It tells you exactly how many electrons occupy each subshell, written in the form: subshell notation + superscript count.

For example, carbon (atomic number 6) has the configuration: 1s² 2s² 2p². This means 2 electrons in the 1s subshell, 2 in the 2s subshell, and 2 in the 2p subshell. The total is 6, matching the atomic number.

Electron configurations matter because they predict chemical behavior, bonding patterns, magnetic properties, and an element’s position on the periodic table. Every trend you study in chemistry — ionization energy, atomic radius, electronegativity — traces back to electron configuration.

The Aufbau Principle and Filling Order

The Aufbau principle (from the German word for “building up”) states that electrons occupy the lowest available energy subshell first. The energy order does not follow a simple numerical sequence, which is why students need a memory tool.

The correct filling order is:

Notice that 4s fills before 3d, and 6s fills before 4f. This surprises many students. The reason is that 4s has slightly lower energy than 3d in a neutral atom, so it fills first.

  1s
  2s  2p
  3s  3p  3d
  4s  4p  4d  4f
  5s  5p  5d  5f
  6s  6p  6d
  7s  7p

  Read diagonally from top-right to bottom-left:
  1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s...
  

Each subshell holds a fixed maximum number of electrons: s holds 2, p holds 6, d holds 10, and f holds 14. These numbers come from the number of orbitals in each subshell (1, 3, 5, and 7 orbitals respectively) multiplied by 2 (Pauli’s limit per orbital).

Pauli Exclusion Principle and Hund’s Rule

Once you know the filling order, two more rules govern exactly how electrons occupy individual orbitals within each subshell.

Pauli Exclusion Principle

No two electrons in the same atom can share all four quantum numbers. In practical terms: each orbital holds a maximum of two electrons, and those two electrons must have opposite spins. You represent this in orbital box notation as one arrow pointing up and one pointing down inside the same box.

Hund’s Rule

When filling orbitals within the same subshell (for example, the three 2p orbitals), place one electron in each orbital before placing a second electron in any orbital. All singly occupied orbitals must have the same spin direction (parallel spins). This minimizes electron-electron repulsion and produces the most stable arrangement.

Step-by-Step: How to Write Any Configuration

Follow these seven steps for any element and you will get the correct configuration every time.

1. Find the atomic number. This equals the total number of electrons in a neutral atom.
2. Write out the Aufbau filling order. Use the diagonal mnemonic if needed.
3. Fill each subshell in order until you have placed all electrons. Respect the capacity limits (s=2, p=6, d=10, f=14).
4. Apply Hund’s rule within each subshell: one electron per orbital before pairing.
5. Apply Pauli’s principle when pairing: opposite spins only.
6. Check for exceptions (Cr, Cu, Mo, Ag, and a few others). If the element is one of these, adjust accordingly.
7. Optionally condense using noble gas shorthand for speed.

Worked Examples (H to Fe)

Let’s apply the steps to real elements. These cover the range of difficulty you will encounter in most courses.

For iron, notice that 4s² is written after 3d⁶ in the noble gas shorthand. This is the conventional way to write transition metal configurations — list the d subshell before the s subshell in the shorthand, even though 4s filled first.

  3d subshell (5 orbitals, 6 electrons — Hund's rule applied):
  [ ↑↓ ] [ ↑ ] [ ↑ ] [ ↑ ] [ ↑ ]
   3d      3d    3d    3d    3d

  4s subshell (1 orbital, 2 electrons):
  [ ↑↓ ]
    4s

  Result: [Ar] 3d6 4s2
  (One 3d orbital is paired; the other four hold single electrons)
  

Noble Gas Shorthand

Noble gas shorthand condenses the core electron configuration into the symbol of the preceding noble gas in square brackets. This saves time, reduces errors in long configurations, and is the standard format used in most textbooks and exams.

To use it: find the noble gas that comes just before your element in the periodic table, write its symbol in brackets, then write only the remaining electrons.

The noble gas shorthand also highlights valence electrons immediately. Everything outside the brackets is in the valence shell region, making it easy to predict bonding behavior.

Exceptions: Cr, Cu, and Why They Break the Rules

Chromium (Cr, Z=24) and Copper (Cu, Z=29) are the two most commonly tested exceptions to the Aufbau principle. Both involve the 3d and 4s subshells, and both occur because a half-filled or fully-filled d subshell is energetically more stable than the standard prediction.

Other elements with similar exceptions include Molybdenum (Mo), Silver (Ag), Gold (Au), and several lanthanides and actinides. For most high school and introductory university courses, Cr and Cu are the only two you need to know by heart.

Common Mistakes and How to Fix Them

After reviewing thousands of student answers, I have identified five mistakes that account for the vast majority of lost marks in electron configuration problems.

Electron Configurations of Ions

Writing configurations for ions follows the same three rules, with one extra step: adjust the electron count for the charge, then remove or add electrons from the correct subshell.

For cations (positive charge, electrons removed): always remove electrons from the subshell with the highest principal quantum number first. For transition metals, this means removing from 4s before 3d, even though 4s filled first.

For anions (negative charge, electrons added): add electrons to the next available orbital following the Aufbau order.

Notice that Fe³⁺ with [Ar] 3d⁵ has a half-filled d subshell — this is one reason Fe³⁺ is particularly stable compared to Fe²⁺.

Sources & References

• Wikipedia — Electron configuration

Reviewed by Dr. Irfan Mansuri. External links open in a new tab and are provided for further reading and verification.