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Question Video: Determining the Electronic Configuration of Iron Chemistry

What is the electronic configuration of iron?

05:01

Video Transcript

What is the electronic configuration of iron? (A) 1s2 2s2 2p6 3s2 3p6 3d6. (B) 1s2 2s2 2p6 3s2 3p6 3d5 4s3. (C) 1s2 2s2 2p6 3s2 3p6 3d6 4s2. (D) 1s2 2s2 2p6 3s2 3p6 3d6 4p2. (E) 1s2 2s2 2p6 3s2 3p6 3d6 4d2.

The electronic configuration describes how many electrons an atom has and how these electrons are arranged in electron shells and subshells. There are s, p, d, and f subshells, and the letter tells us which type of atomic orbitals the subshell contains. The periodic table can be divided into blocks which represent these subshells. For the purpose of this video, we will ignore the elements found in the f block.

Iron is a transition metal located in period four and group eight of the periodic table, and we can see that it is located in the d block. Iron has an atomic number of 26. This means that an iron atom has 26 protons. Atoms are electrically neutral, so the number of protons must equal the number of electrons. Therefore, there are 26 electrons in an iron atom.

To help us write the electronic configuration of iron, we need to use the aufbau principle. The aufbau principle states that electrons fill the lowest energy subshells before they fill higher energy ones. On the periodic table, the subshells are in order of increasing energy. When writing the electronic configuration, each subshell label has a coefficient, which is the principal quantum number 𝑛 and represents the energy level in which electrons are found. The letter represents the specific type of subshell in which the electrons are found. Finally, the superscript represents the number of electrons in the subshell. The following table shows the maximum number of electrons that each type of subshell can hold.

Now we’re ready to begin writing the electronic configuration of iron. Starting with hydrogen, we move across period one and write 1s2. Until we reach the final subshell, we should always fill the subshell with the maximum amount of electrons it can hold. This is why we placed two electrons in the 1s subshell. Moving down to period two, we fill the 2s subshell with two electrons and write 2s2. Continuing to move across period two, we fill the 2p subshell with six electrons and write 2p6. Then, we move across period three to fill both the 3s and 3p subshells and write 3s2 3p6. Finally, we have reached period four, which represents the valence electron shell of the iron atom. First, we fill the 4s subshell with two electrons and write 4s2. Next, we move across the 3d subshell until we reach iron. Iron is the sixth element in the 3d subshell, so we write 3d6.

It’s important to note that the coefficient used for d subshells is one less than that used for s and p subshells in the same period.

There are two common ways to write the electronic configuration. The electronic configuration we just wrote puts the subshells in order of increasing energy. The second way to write the electronic configuration is to put the subshells in order of increasing principal quantum number 𝑛. This means that 3d6 is written before 4s2. To ensure the electronic configuration has the correct number of electrons, we can simply add together the superscripts. The sum of the superscripts is 26, which is the correct number of electrons for an iron atom.

When looking at the answer choices, we see that answer choice (C) matches the second electronic configuration we wrote. In conclusion, the electronic configuration of iron is 1s2 2s2 2p6 3s2 3p6 3d6 4s2, or answer choice (C).

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