Lesson Video: Electronic Configurations of Transition Metals Chemistry • Third Year of Secondary School

Video Transcript

In this video, we will learn how to describe the electronic configurations of transition metals and the formation of their ions. We will state the electronic configurations and oxidation states of the first row of d-block elements and their ions, show how they relate to each other, and use the electronic configurations of transition metals to help us define a transition metal.

The electronic configuration of an atom describes how many electrons an atom has and how these electrons are arranged in different electron shells and subshells. On the periodic table, a period or horizontal row contains elements which have valence electrons in the same highest occupied electron shell. For example, sodium and all of the elements in period three will have valence electrons in the third electron shell. However, the arrangement of electrons within electron shells is more complicated than what we’ve shown here. Inside the electron shells are subshells which have a letter code based on the type of orbitals they contain. There are s, p, d, and f subshells.

The periodic table can be divided into blocks that represent these subshells. Elements found in the same block have valence electrons in the same type of subshell. Elements found in the d block, which are the elements in groups three through 12, have one or more valence electrons in a d subshell. The order in which electrons fill up subshells of an atom is based on increasing energy. According to the Aufbau principle, electrons fill the lowest-energy subshells before they fill higher-energy subshells. The electron configuration of an atom can be written by moving period by period across the periodic table, starting with hydrogen until we reach the element that we want to write the electron configuration for.

Let’s write an electron configuration for the first d-block element found in period four, which is scandium. Scandium has an atomic number of 21, which means that a scandium atom has a total of 21 electrons. To write the electron configuration of scandium, let’s begin at hydrogen and move across period one of the periodic table, which represents the 1s subshell. Every time we move to a new element, we add another electron to the subshell. So we will need to fill the 1s subshell with two electrons. When writing the electron configuration, we start each part with the subshell label and use a superscript to indicate the number of electrons in the subshell.

Moving into period two, we fill the 2s subshell with two electrons and the 2p subshell with six electrons. Moving across period three, we fill the 3s subshell with two electrons and the 3p subshell with six electrons. Finally, moving across period four, we fill the 4s subshell with two electrons and the 3d subshell with one electron. When we write the electronic configuration using the periodic table, the value used for the subshell label generally matches the period number. However, when filling d subshells, the value used is always one less than the period number.

Now, you’ll often see the electron configurations of transition metals written so that the 3d subshell comes before the 4s subshell. It is acceptable to write it either way. The electron configurations for elements beyond period four can be quite long. So chemists simplify the electron configuration by using condensed notation. Scandium is in period four. The part of the electron configuration that we wrote before reaching period four corresponds to the electron configuration of argon. The remaining subshells contain the scandium atom’s valence electrons.

Condensed notation consists of a noble gas in brackets followed by the electronic configuration for the valence electrons. After combining these two parts, the condensed notation of scandium is argon in brackets 4s2 3d1.

Now that we’ve written an electron configuration for the first element in the d block, we’re going to take a deeper look at the electron configurations of the transition metals. The transition elements are metals that are found in groups three through 11 on the periodic table. The elements which are found in group 12, which are zinc, cadmium, mercury, and copernicium, are generally not considered transition metals.

A transition element is defined as an element whose atoms have an incomplete d subshell, or which can give rise to cations with incomplete d subshells. We’re going to focus on the transition metals in period four because they have the simplest electron configurations of all of the transition metals. Let’s write electron configurations using the condensed notation for the next two transition metals after scandium, which are titanium and vanadium.

Orbital diagrams can be used to show how the electrons are organized in the valence shells of these three transition-metal atoms. According to Hund’s rule, electron orbitals in the same subshell are filled with one spin-up state electron before they can be paired with a spin-down state electron. When filling the valence shell of a scandium atom, we’ll start by filling the 4s subshell with two electrons with opposite spins. The 4s subshell only has one orbital, so we can fill it with two electrons. Then, we’ll need to fill the first orbital in the 3d subshell with one electron spin up.

To complete the orbital diagram for the valence shell of titanium, we’ll need to fill the 4s subshell with two electrons with opposite spins. And then in the 3d subshell, we’ll put one spin-up electron in the first orbital and the second orbital. We do not put two electrons in the first orbital in the 3d subshell because that would violate Hund’s rule. A similar process is followed to complete the diagram for vanadium. In the 3d subshell of vanadium, the first three orbitals will contain one spin-up electron each.

Now we’re ready to take a look at the electron configurations for the remaining transition metals in period four. In each successive electron configuration, we would expect to see one more electron present in the d subshell. However, in the electron configurations of chromium and copper, the d subshell contains one more electron than we would expect.

By closely inspecting the electron configurations and orbital diagrams of chromium and copper, we see that the 4s subshell only contains one electron instead of two. The result is that in chromium the 4s and 3d subshells are exactly half full. And in copper, the 4s subshell is half full and the 3d subshell is completely full. The reason for this is beyond the scope of this video. But it is important to identify chromium and copper as transition metals in period four that are exceptions to the rules of writing electron configurations.

When looking at the electron configurations of all of the transition elements in period four, we see that, with the exception of copper, all have an incomplete d subshell. This is a defining characteristic of the transition elements.

Most transition metals can form more than one type of ion. In fact, all of the period four transition metals can form three or more different ions, except for scandium, which can only form one type of ion. Period four transition metals form cations or positively charged ions by losing valence electrons from their s and d subshells. The oxidation state of a transition metal represents the number of electrons that have been lost by the metal atom to form an ion.

Nickel has three different oxidation states: plus two, plus three, and plus four. Because of this variety of oxidation states, nickel can form more than one type of compound with some elements. In nickel(II) oxide, it’s the nickel two plus ion that’s combined with the oxide ion, and in nickel(III) oxide it’s the nickel three plus ion. So, how exactly do these different ions form?

Let’s begin by looking at the electron configuration and orbital diagram for a nickel atom. Transition metals in period four tend to lose electrons from the 4s subshell before they lose electrons from the 3d subshell. Let’s say a nickel atom loses two electrons to form a nickel two plus ion. Because the two electrons are lost from the 4s subshell, the electron configuration for the ion will not include the 4s subshell. If a nickel atom lost three electrons to form the nickel three plus ion, then the third electron could be taken from the 3d subshell.

In general, when writing the electron configurations of transition metals in period four, the 4s subshell is filled with electrons before the 3d subshell. In comparison, when writing the electron configurations for transition-metal ions, electrons are removed from the 4s subshell before the 3d subshell. Before we summarize what we’ve learned about transition metals and their electron configurations in this video, let’s take a look at a question.

Which of the following is the electronic configuration of Ti? (A) Ar 4s1 3d3, (B) Ar 3s2 4d2, (C) Kr 4s2 3d2, (D) Kr 5s2 4d2, (E) Ar 4s2 3d2.

To solve this problem, we need to select the answer choice that shows the correct electron configuration of the element titanium. The electronic configuration of an atom describes how many electrons the atom has and how these electrons are arranged into different electron shells and subshells. The atomic number of titanium is 22, which means that a titanium atom has a total of 22 electrons.

When looking at the answer choices, we see that the electron configurations provided are given in condensed notation. In general, condensed notation has the form of the chemical symbol of a noble gas followed by the electron configuration of the subshells that hold the valence electrons. Titanium is the second transition-metal element in period four on the periodic table. Let’s use the periodic table to fill the subshells in a titanium atom with electrons.

Starting with hydrogen and moving across period one, we fill the 1s subshell with two electrons. So we write 1s2. Moving across period two, we fill the 2s subshell with two electrons and the 2p subshell with six electrons. So we can write 2s2 2p6. Moving across period three, the pattern is similar to period two. We will fill the 3s subshell with two electrons and the 3p subshell with six electrons. So we can write 3s2 3p6. Finally, in period four, we fill the 4s subshell with two electrons and the 3d subshell with two electrons. So we can write 4s2 3d2.

To write the full electron configuration for titanium, let’s put the subshells in order from left to right. To determine the identity of the noble gas which should be used in the condensed notation, we simply locate the noble gas at the end of the period prior to the period our element is in. Because titanium is in period four, we look to find the noble gas at the end of period three, which is argon. We can condense the electron configuration that we wrote for periods one through three as argon in brackets. In other words, the condensed form of the electron configuration for titanium can be written as argon inside brackets followed by 4s2 3d2. Therefore, the correct answer is answer choice (E) argon in brackets 4s2 3d2.

Now, let’s summarize what we’ve learned. Transition elements are elements which have incomplete d subshells or form cations with incomplete d subshells. All period four transition metals have electron configurations with full 4s subshells and incomplete 3d subshells, except for chromium and copper. In chromium, the 4s and 3d subshells are each half full. And in copper, the 4s subshell is half full and the 3d subshell is completely full.

When writing electron configurations for atoms of period four transition metals, electrons fill the 4s subshell before the 3d subshell. However, when writing the electron configuration for period four transition-metal ions, electrons are removed from the 4s subshell before the 3d subshell. Because most transition metals have multiple oxidation states, they can therefore form more than one type of cation.