D Orbitals Electron Config (Chemistry Guide)

A d subshell contains five orbitals and can hold up to 10 electrons. To write its electron configuration, count the atom’s electrons, follow the usual filling order, and check for known exceptions. For transition-metal ions, remove the outer (ns) electrons before the ((n-1)d) electrons. These steps help you avoid common errors.

Traditionally, chemistry courses teach electron configurations by filling orbitals in a set order. That gives you a useful starting point, but transition metals bring two details that often cause confusion: a few neutral atoms do not match the simplest prediction, and ions lose electrons in an order that may seem to reverse the filling order.

I’ll build the method from the basics, then work through neutral atoms and ions. The goal is not just to copy a configuration. It is to know how to check it.

D-Orbital Electron Configurations: Rules, Notation & Transition-Metal Overview

A d-orbital configuration shows how many electrons occupy d subshells in an atom or ion. Each d subshell has five orbitals and holds up to 10 electrons. The first d subshell is 3d, and d subshells are important when writing configurations for transition metals.

The number before a subshell letter gives its principal energy level, or shell. The letter names the subshell: s, p, d, or f. For example, (3d^6) means six electrons occupy the 3d subshell.

Each d subshell has five orbitals, commonly labeled (d_{xy}), (d_{xz}), (d_{yz}), (d_{x^2-y^2}), and (d_{z^2}). An orbital is a region where an electron is likely to be found. Each orbital can hold two electrons, so a d subshell can hold 10 in total.

The first d subshell is 3d because d subshells begin at principal energy level (n=3). In transition-metal configurations, you will also see 4d, 5d, and 6d. The element’s position in the periodic table can help you predict which d subshell is being filled.

A compact configuration uses a noble-gas core to stand for inner electrons. For example, ([\mathrm{Ar}]) represents the electron configuration of argon. So ([\mathrm{Ar}]\,3d^6 4s^2) shows the electrons beyond argon without writing every subshell again.

Aufbau Order, Orbital Diagrams & Configuration Requirements

The Aufbau principle guides the usual filling sequence: electrons enter lower-energy subshells before higher-energy ones. Pauli’s principle limits each orbital to two electrons with opposite spins, while Hund’s rule says equal-energy orbitals fill one at a time before electrons pair.

For the common sequence used to write neutral-atom configurations, fill subshells in this order:

(1s,\ 2s,\ 2p,\ 3s,\ 3p,\ 4s,\ 3d,\ 4p,\ 5s,\ 4d,\ 5p,\ 6s,\ 4f,\ 5d)

This order is a practical guide, not an exception-free rule. In particular, the sequence places 4s before 3d for many neutral atoms. It does not mean that 4s must always be lower in energy than 3d in every atom or ion.

Subshell capacities help you check your work:

A d-orbital diagram can make the rules visible. Represent the five d orbitals as five boxes. Place one electron in each box before pairing any of them. Then add a second electron to each box as needed, with opposite spin. This follows Hund’s rule and Pauli’s principle.

For example, a (3d^6) subshell has one paired box and four boxes with one electron each. The exact box chosen for the pair does not change the number of unpaired electrons. In this simple diagram, (3d^6) has four unpaired electrons.

To write a neutral atom’s configuration:

• Find its atomic number. A neutral atom has the same number of electrons as its atomic number.
• Write the nearest earlier noble-gas core.
• Add electrons in the usual filling order.
• Check that the electron count matches the atomic number and that no subshell exceeds its capacity.

Here are two examples. Scandium has 21 electrons, so its configuration is ([\mathrm{Ar}]\,4s^2 3d^1). Iron has 26 electrons, giving ([\mathrm{Ar}]\,4s^2 3d^6). Both examples show the common neutral-atom pattern in which 4s fills before 3d.

Exceptions, Ions & Strategies for Accurate Problem Solving

Some neutral transition-metal atoms have ground-state configurations that differ from the simplest Aufbau prediction. Chromium and copper are common examples. When forming transition-metal cations, use a separate rule: remove the outer (ns) electrons before removing ((n-1)d) electrons.

The standard introductory configurations for chromium and copper are:

These are exceptions to the simplified filling prediction, not violations of the rules for electrons within orbitals. The actual ground-state arrangement reflects the relative energies of the subshells in each atom. For introductory work, learn these two standard exceptions and check a reliable periodic-table reference if a less familiar element comes up.

For ions, first adjust the electron count by the charge. A positive charge means electrons were lost; a negative charge means electrons were gained. Then write the configuration, removing transition-metal electrons from the outer (ns) subshell before the ((n-1)d) subshell.

Consider iron:

• Neutral Fe has 26 electrons: ([\mathrm{Ar}]\,4s^2 3d^6).
• Fe²⁺ has lost two electrons. Remove both 4s electrons: ([\mathrm{Ar}]\,3d^6).
• Fe³⁺ has lost one more electron, now from 3d: ([\mathrm{Ar}]\,3d^5).

This is a key point: neutral-atom filling order does not tell you to remove 3d electrons first. If you do that, you get the wrong transition-metal cation configuration. For copper, neutral Cu is ([\mathrm{Ar}]\,3d^{10}4s^1); Cu⁺ is ([\mathrm{Ar}]\,3d^{10}), after losing its 4s electron.

Use this short check when a problem asks for an ion:

• Count electrons from the atomic number and charge.
• Write or identify the neutral configuration, including known exceptions.
• Remove electrons from the highest outer (ns) subshell first for a transition-metal cation.
• Recount the remaining electrons and check each subshell’s capacity.

Applications to Oxidation States, Magnetism & Chemical Behavior

A correct d-electron count helps you interpret oxidation states and unpaired electrons. It does not, by itself, predict every feature of a compound. In particular, the magnetic behavior of a transition-metal ion can depend on its chemical setting, so distinguish a basic atomic-orbital count from a full analysis of a compound.

An oxidation state is a bookkeeping charge assigned to an atom in a compound. For a simple ion such as Fe²⁺, the charge means two electrons have been removed from neutral iron. The configuration ([\mathrm{Ar}]\,3d^6) then gives the ion’s d-electron count for introductory problems.

Unpaired electrons matter because they can affect magnetic behavior. In a simple orbital diagram, an electron that occupies an orbital alone is unpaired. For example, the diagram for a free Fe³⁺ ion’s (3d^5) configuration has five singly occupied orbitals under Hund’s rule.

However, do not assume that a free-ion diagram fully describes every transition-metal compound. The surrounding atoms or ions can affect the relative energies of d orbitals and how electrons occupy them. For a basic configuration question, use the stated atom or ion and the rules above; for a compound’s magnetism, follow the model and conditions given in your course.

A useful final check is to ask three questions: Does the total electron count match the species? Did I use the correct neutral-atom exception, if needed? For a cation, did I remove the (ns) electrons first? Those checks catch many common configuration errors.

Frequently Asked Questions

These answers focus on the most common questions about d subshells, transition-metal configurations, and ions. Use them to review the rules, but check your course’s instructions when a problem specifies a different notation style or asks for a more advanced model.

How many electrons can a d subshell hold?
A d subshell has five orbitals, and each orbital holds up to two electrons. Its maximum capacity is 10 electrons.

What is the first d subshell?
The first d subshell is 3d. A d subshell first occurs at principal energy level (n=3).

Why does 4s fill before 3d in many neutral atoms?
The usual filling sequence places 4s before 3d for many neutral atoms. This is a useful rule for writing configurations, but it does not mean 4s is always lower in energy in every atom or ion.

Which elements are common electron-configuration exceptions?
Chromium and copper are standard introductory examples. Their common ground-state configurations are ([\mathrm{Ar}]\,3d^5 4s^1) and ([\mathrm{Ar}]\,3d^{10}4s^1), respectively.

When forming a transition-metal cation, which electrons are removed first?
Remove electrons from the outer (ns) subshell before removing electrons from the ((n-1)d) subshell. For example, Fe²⁺ is ([\mathrm{Ar}]\,3d^6).

How do I find the number of electrons in an ion?
Start with the atomic number. Subtract the positive charge for a cation, or add the size of the negative charge for an anion.

What does (3d^6) mean?
It means six electrons occupy the 3d subshell. The superscript gives the number of electrons in that subshell.

How many unpaired electrons are in a (3d^5) subshell?
In the simple orbital diagram, there are five unpaired electrons. Hund’s rule places one electron in each of the five orbitals before pairing begins.

(This article was written by one of our staff writers, Alan Westbrook. Visit our Meet the Team page.)