# Why does oxygen form an O^(2-) ion and not an O^(3-) ion?

Jun 8, 2018

The electron configuration of ${\text{O}}^{2 -}$ ensure the atom an octet and therefore high chemical stability whereas ${\text{O}}^{3 -}$ doesn't.

#### Explanation:

Ground-state oxygen atoms have electron shell arrangement (starting from the innermost principle energy level)

$2 , 6 , \textcolor{g r e y}{0 , 0 , \cdot \cdot \cdot}$

Each of the electron shells holds a different maximum number of electrons:

$2 , 8 , 18 , 32 , \cdot \cdot \cdot$

An atom is most stable when all of its electron shells are either filled with electrons or left unoccupied. It will form ions or share electrons with another atom seeking to achieve such configurations.

An atom either loses or gains one or more electrons to form an ion. Each electron carries a charge of $- 1$ whereas each proton within the atomic nucleus carries a charge of $+ 1$. With the number of electrons different from that of protons, positive and negative charges no longer cancels such that unlike atoms- which are neutral for containing equal numbers of electrons and protons- ions carry charges. The charge on an ion is equal to the difference in the number of electrons and that of protons it contains- in other words, the number of electrons its parent atom has gained or lost.

${\text{O" + 2color(white)(l)e^(-) to "O}}^{2 -}$

An electrically-neutral oxygen atom gains two electrons to form an oxygen ion with two negative charges. Notice how the charge conserves in this process.

The two electrons will end up in the main energy level of the lowest potential energy possible- that is, closest to the atom- to minimize the potential energy of this atom. The atom would thus have the electron arrangement:

$2 , \textcolor{n a v y}{8} , \textcolor{g r e y}{0 , 0 , \cdot \cdot \cdot}$

This particular arrangement ensures two filled main energy levels while leaving the rest empty. As a result, the ${\text{O}}^{2 -}$ ion should be relatively chemically stable.

On the other hand, forming an ${\text{O}}^{3 -}$ ion requires adding an additional electron to an ${\text{O}}^{2 -}$ ion (or equivalently three electrons to a neutral oxygen atom.) However, the second main energy level is already full meaning that the last electron would be added to the third energy level. An ${\text{O}}^{3 -}$ ion would thus have the electron arrangement:

$2 , \textcolor{n a v y}{8} , \textcolor{red}{1} , \textcolor{g r e y}{0 , \cdot \cdot \cdot}$

The extra electron on the third main energy shell- the only electron on that energy level- would be subjected to significant electrostatic repulsion known as the shielding effect from the rest ten electrons. The ion would thus be highly unstable and potentially readily lose that outermost electron and return to ${\text{O}}^{3 -}$ as the ten inner shell electrons seek to get rid of the additional electron.