Lesson 7: Electron Configurations

“Nameless” Rule

Transition metals and F block elements’ electron configurations occasionally do not follow the three principles we have previously learned. This is an observation (or a fourth “rule”) that states that atoms are generally more stable when they have more completely full, completely empty, or exactly half-full subshells.

For example:

The electron configuration of Cr is [Ar]3d⁵4s¹ instead of [Ar]3d⁴4s² because half of the d sublevel’s electron capacity is 5 and half of the s sublevel’s electron capacity is 1. This ends up with five full subshells shown by the [Ar] part (1s²2s²2p⁶3s²3p⁶) and two half-full subshells, instead of six full subshells and one almost-full subshell, overall increasing the stability of the atom. The reason it is not [Ar]3d⁵4s² is that it must retain the correct number of electrons, so the d subshell gains one electron from the s subshell. This differs from Mn, which is actually [Ar]3d⁵4s² because that’s how many electrons it has in total.

Ion Formation

To find the electron configurations of ions, we need to know which shells and subshells to add or take away electrons.

For cations, electrons need to be removed from the highest shells (e.g. n = 4) first and then from the highest subshells in those shells (e.g. 4f).

For example:

The electron configuration of Na is [Ne]3s¹. To find the electron configuration of Na⁺, we remove one electron from the highest shell and subshell. Writing out the whole electron configuration of Na gives us 1s²2s²2p⁶3s¹, showing that 3 is the highest energy shell and 3s is the highest energy subshell. Since Na⁺ has one less electron than Na, one electron is removed from the 3s sublevel, which gives us an electron configuration of 1s²2s²2p⁶ (or [Ne]) for Na+.

The electron configuration of Fe is [Ar]4s²3d⁶. To find the electron configuration of Fe3⁺, we remove three electrons from the highest shells and subshells. Writing out the whole electron configuration of Fe gives us 1s²2s²2p⁶3s²3p⁶4s²3d⁶, showing that 4 is the highest energy shell and 4s is the highest energy subshell. Since Fe3⁺ has three less electrons than Fe and the 4s sublevel only has two electrons, two electrons are removed from the 4s sublevel. Since there is still one electron left to remove, it has to be taken from the next highest shell and subshell: 3d. Only one more electron needs to be removed, so one is taken from the 3d shell and the electron configuration that is left is 1s²2s²2p⁶3s²3p⁶3d⁵ or [Ar]3d⁵ for Fe3⁺.

To find the electron configuration of an anion, you take the original element’s atomic number and add the absolute value of the charge of the anion to it. The new atomic number you find will give you a different element, and that element’s electron configuration is the same as this anion.

For example:

To find the electron configuration of Cl⁻, we know that the charge is -1 and the atomic number of Cl is 17. We take the absolute value of -1 (which is 1) and add it to 17, giving us 18, which is the atomic number of Ar. This means that the electron configuration of Cl⁻ is the same as the electron configuration of Ar, which is 1s²2s²2p⁶3s²3p⁶ or [Ne]3s²3p⁶.

Combining Both Concepts

As previously established, the electron configuration of Fe is [Ar]4s²3d⁶. To find the electron configuration of Fe²⁺, two electrons need to be removed. In most cases, two would be taken from the highest shell and subshell, 4s, leaving [Ar]3d⁶. However, removing one electron from the 4s subshell and one electron from the 3d subshell would leave [Ar] 4s¹3d⁵: five full subshells and two half-full subshells. This, as stated earlier, creates a more stable atom. So, the resulting electron configuration is 1s²2s²2p⁶3s²3p⁶4s¹3d⁵ or [Ar]4s¹3d⁵ for Fe²⁺.

Magnetic Behavior

In atoms, electrons spin. When a charged particle—such as a (negatively charged) electron—spins, the direction it is spinning in determines which pole it will be attracted to on a magnet.

In atoms specifically, unpaired electrons cause magnetic behavior in elements or compounds. “Unpaired electrons” refers to electrons that are not in an orbital with another electron, they are instead alone in their own orbital. Magnetic behavior in elements or compounds due to unpaired electrons is called unopposed magnetic attraction. When an atom or compound has more unpaired electrons, it is more magnetic.

Diamagnetic substances have no unpaired electrons, and are slightly repelled by strong magnets.

Paramagnetic substances have one or more (usually just a few) unpaired electrons, and are slightly attracted to strong magnets.

Ferromagnetic substances are substances in which most atoms have unpaired electrons, causing them to be strongly attracted to strong magnets.