Lesson Explainer: Electronic Configurations of Transition Metals Chemistry

In this explainer, we will learn how to describe the electronic configurations of transition metals and the formation of their ions.

The transition elements are metals that can be found in groups 3–11 of the periodic table. They have many fascinating properties; these include their capacity to form brightly colored solutions and their ability to form ions that have variable oxidation states. Most of the unique chemical properties of the transition metal elements can be understood by considering their electron configuration.

Definition: Transition Element

A transition element is an element whose atoms have an incomplete d subshell or which can give rise to cations with an incomplete d subshell.

The aufbau principle states that the electrons of any one atom must fill the lowest-energy atomic orbitals before they can fill the other higher-energy atomic orbitals. Electrons will always fill the lowest-energy, 1s, atomic orbital and then they fill the higher-energy, 2s and 2p, atomic orbitals. The following figure shows the relative energy values of the first eighteen electron subshells.

Definition: Aufbau Principle

The Aufbau Principle states that electrons fill the lowest energy atomic orbitals before they fill higher energy atomic orbitals.

The energy-level diagram can be used to understand that electrons will always fill the 1s subshell before they start to fill the 2s subshell and that the 2p subshell will always be filled after the 2s subshell. The energy-level diagram can also be used to infer that the filling order is not straightforward. Some atomic orbitals with relatively high principal quantum numbers have to be filled before other atomic orbitals that have lower principal quantum numbers. For example, the 4s atomic orbital tends to be filled before the 3d atomic orbital because the 4s atomic orbital has the lower energy value.

The electron configuration of an element is a series of baseline and superscript symbols that describes the distribution of electrons in atomic orbitals. The period four transition metal elements have the simplest electron configurations of all the different types of transition metal elements. They have electron configurations that are made up of just seven subshell labels.

There are two common conventions for writing the electron configurations of the period four transition elements. We can appreciate the different conventions by considering the electron configuration of scandium. Some scientists choose to write the electron configuration of scandium as 1223343sspspsd and other scientists choose to write the electron configuration of scandium as 1223334.sspspds

The first sequence is ordered in terms of subshell position in the periodic table and the second sequence is ordered in terms of increasing quantum number. We will use the convention in the first of these two sequences throughout this explainer.

Scandium has the fewest electrons of any period four transition metal element because it contains just one 3d subshell electron. Other period four transition metal elements have similar electron configurations, but they all have higher 3d subshell occupation numbers. The electron configuration of titanium is 1223343.sspspsd

The electron configuration of scandium and that of titanium are almost identical, but titanium has one more 3d subshell electron than scandium.

Example 1: Determining the Electronic Configuration of Titanium

Which of the following is the electronic configuration of Ti?

  1. [Ar]4s3d13
  2. [Ar]3s4d22
  3. [Kr]4s3d22
  4. [Kr]5s4d22
  5. [Ar]4s3d22


We can use the periodic table to determine that titanium has the following electron configuration: 1223343sspspsd. Krypton has the electron configuration 12233434sspspsdp and argon has the electron configuration 12233sspsp. This means that we will have to use the [Ar] prefix to represent the electron configuration of titanium. Titanium has two more 4s and 3d subshell electrons than argon and we can therefore surmise that the condensed notation of titanium is [Ar]4s3d22, or answer E.

The following table shows the electron configurations of the other period four transition metal elements. It uses the bracketed [Ar] noble gas term to represent the following five subshell terms: 12233.sspsp

ElementElectronic Configuration4s3d3d3d3d3d

The electron configurations are depicted on the right-hand side as a series of upward- and downward-facing arrows. Quantum chemists regularly use upward- and downward-facing arrows to show how electrons progressively fill atomic orbitals according to Hund’s rule. Hund’s rule states that each orbital of a given subshell is filled with one spin-up state electron () before any of the other subshell orbitals can be filled with secondary spin-down state electrons (). The 3d subshell orbitals should always be filled with upward-facing arrows before they are filled with any secondary downward-facing arrows.

Definition: Hund’s Rule

Every orbital in a subshell is singly occupied before any orbital is doubly occupied.

It is important to note here that the electron configurations of the period four transition metal elements can be described with a single mathematical expression. The period four transition metal electron configurations can be described with the expression [Ar] 𝑛(𝑛1)sd, where 𝑛=4. Similar expressions can be used to describe the electron configurations of the period five and six d-block transition metal elements.

Determining the electron configuration of the first-row transition elements is not always straightforward because some of them have a seemingly anomalous arrangement of valence electrons. Most of the transition metals have two electrons in the 4s subshell but chromium and copper have a single electron in their 4s subshell. The electron configurations for chromium and copper are shown below: Chromium:[Ar]4s3dCopper:[Ar]4s3d15110,.

It is quite difficult to fully rationalize the seemingly anomalous arrangement of electrons in copper and chromium atoms without introducing ideas that are beyond the scope of this explainer. It is customary for high school students to simply accept that chromium and copper atoms are most stable when they have one electron in the 4s subshell and either five or ten electrons in the 3d subshell.

Example 2: Identifying the Electron Configuration of Chromium Atoms from Simple Schematic Illustrations of Electrons in the 3d and 4s Subshells

Which of the following diagrams represents the electronic configuration of a chromium atom?


The electron configuration of an element can be presented as a series of baseline and superscript symbols or it can be depicted schematically with a combination of upward- and downward-facing arrows. The inner shells of electrons can all be presented explicitly, or they can be represented with bracketed noble gas terms such as the [Ar] prefix. The [Ar] prefix is used to represent the 12233sspsp electrons.

Chromium has the electronic configuration 1223343sspspsd. The electron configuration could alternatively be written as [Ar]4s3d15, because the [Ar] prefix can be used in place of the 12233sspsp electron configuration. This electron configuration could also be represented schematically with a series of arrows. All of the arrows would have to be upward facing because Hund’s rule states that atomic orbitals are singly occupied with spin-up () electrons before they are doubly occupied with spin-down () state electrons. There are only five 3d subshell electrons and one 4s subshell electron and this means that all of the 3d and 4s electrons must be in a spin-up state. We can use this line of reasoning to determine that option A must be the correct answer for this question.

Most transition elements can form different types of ions during chemical reactions because they can lose electrons from both d- and s-type subshells. Almost all of the period four transition elements can form one type of ion as they lose electrons from the 4s subshell and at least one other type of ion as they lose more electrons from the 3d subshell. Iron will initially form the Fe2+ ion if it loses its pair of 4s electrons and it can then go on to form the Fe3+ ion if it loses another single 3d subshell electron. The Fe atom electron configuration is shown below with the Fe2+ and Fe3+ ion electron configurations: Featom:[Ar]4s3dFeion:[Ar]4s3dFeion:[Ar]4s3d262+063+05,,.

The following table shows some common and less common oxidation states of the period four transition metal elements. It is clear from this table that most of the transition metals can form at least two different ions during chemical reactions. Some of the transition metals form as many as six different types of ions during standard chemical reactions. It is interesting to note from this table that some maximum oxidation states are simply determined as the total number of electrons in the 3d and 4s subshells.

ElementElectronic ConfigurationCommon and Less Common Oxidation StatesExample Compounds
Ti[Ar]4s3d22+2+3+4TiO, TiO23, TiO2
V[Ar]4s3d23+2+3+4+5VO, VO23, VO2, VO25
Cr[Ar]4s3d15+2+3+4+5+6CrO, CrO23, CrO3
Mn[Ar]4s3d25+2+3+4+5+6+7MnO, MnO23, MnO2, KMnO24, KMnO4
Fe[Ar]4s3d26+2+3+4+5+6FeO, FeO23
Co[Ar]4s3d27+2+3+4+5CoCl2, CoCl3
Ni[Ar]4s3d28+2+3+4NiO, NiO23, NiO2
Cu[Ar]4s3d110+1+2+3CuO2, CuO

Zinc atoms are very interesting because they are d-block elements which tend to only form one 2+ ion rather than a combination of multiple different types of positively charged ions. The electron configurations of the neutrally charged zinc atom and the positively charged Zn2+ ion are shown below: Znatom:[Ar]4s3dZnion:[Ar]4s3d2102+010,.

Zinc tends to only form one ion because it has a completely full 3d subshell and it is not energetically favorable to remove one or just a few electrons from a completely filled subshell. Zinc is not considered to be a transition metal by the IUPAC because it has a completely filled 3d subshell and it does not generally form ions that have incomplete 3d subshells.

Example 3: Identifying Which Element Is Not a Transition Metal

Which element in the 4th period of the periodic table and in the d block is not a transition metal when considering its electronic configuration?

  1. Chromium
  2. Scandium
  3. Copper
  4. Zinc
  5. Iron


Transition metals have atoms with incomplete d subshells or atoms that can produce cations with incomplete d subshells. Chromium and scandium should be defined as transition metals because they have the [Ar]4s3d15 and [Ar]4s3d21 electron configurations and they can both form ions that have incomplete d subshells. Copper and iron should be defined as transition metals because they have the [Ar]4s3d110 and [Ar]4s3d26 electron configurations and they can both form ions that have incomplete d subshells. Zinc should not be defined as a transition metal because it has the [Ar]4s3d210 electron configuration and it forms the Zn2+ ion which has the [Ar]4s3d010 electron configuration. Neutrally charged zinc atoms and positively charged zinc (Zn2+) ions both have a complete or full d electron subshell, and this demonstrates that they do not agree with the standard definition for a transition metal. We can use these statements to determine that option D must be the correct answer for this question.

The oxidation state of a transition metal ion can be used to determine its electron configuration because oxidation numbers show how many electrons have been lost from an atom. The +2, +3, +4, and +5 oxidation states of vanadium are shown below: Vatom:[Ar]4s3dVion:[Ar]4s3dVion:[Ar]4s3dVion:[Ar]4s3dVion:[Ar]4s3d232+033+024+015+00,,,,.

Vanadium has 5 valence electrons and it forms different types of positively charged ions when it loses its 4s and then 3d subshell electrons. The vanadium 2+ ion is formed when a vanadium atom loses two 4s subshell electrons. The vanadium 3+ ion is formed when another 3d subshell electron is removed from the V2+ ion. The vanadium 5+ ion is formed when a vanadium atom loses all of its 3d and 4s subshell electrons.

Oxidation numbers can always be used to understand how many 4s and 3d electrons have been removed from a transition element atom. The maximum oxidation state of any transition element can generally be inferred from its total number of valence electrons.

Example 4: Determining the Maximum Oxidation State of Manganese from Its Electron Configuration

Given that the electronic configuration of manganese is [Ar]4s3d25, what is the maximum oxidation state for a transition metal?


The oxidation number indicates the number of electrons that are lost or gained from an atom. The period four transition metal elements tend to lose electrons from their 3d and 4s electron subshells. Manganese atoms have seven electrons that can be lost from the 4s and 3d subshells. We can therefore determine that manganese can form an ion with the +7 oxidation state.

Key Points

  • The first-row transition metal elements have similar electron configurations with full 4s subshells.
  • The 3d subshell contains more electrons as we move from scandium through to copper.
  • Chromium and copper have unusual electron configurations with half-filled 4s orbitals: Chromium:[Ar]4s3dCopper:[Ar]4s3d15110,.
  • Oxidation numbers show how many electrons have been lost or gained from a transition element.
  • Elements titanium through copper have multiple oxidation states.
  • All subshell orbitals are singly occupied before any orbitals are doubly occupied.
  • Transition metal electronic configurations can be expressed schematically using combinations of upward-facing () and downward-facing () arrows.

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