Electron configuration - Negative Ions

Now that you have a sense of electron configurations and how electrons are arranged in the orbitals of an atom, it is time to look at ions, specifically negative ions.

Remember that a set of orbitals with the same energy is most stable when it is empty, half-filled or filled. To understand negative ions, we will focus on the last of those three possibilities - when the orbitals are filled and in particular how it applies to the p block of the periodic table.

The p block can hold up to 6 electrons (3 orbitals, 2 electrons in each). The result of that is a column of remarkably stable elements - the noble gases.

Elements that are close to those elements often take electrons from other atoms to achieve that same stable electron configuration. For instance fluorine has an electron configuration of 1s2 2s2 2p5, which is almost the same as neon (1s2 2s2 2p6). If fluorine steals an electron from another atom, it achieves the same, stable structure that neon has. As a result the F-1 ion is very stable. 

This can be confusing, so let’s make sure that what we’re saying here is clear.

• Fluorine is unstable and dangerous (1s2 2s2 2p5)

• Fluorine often steals an electron from another atom to become stable (1s2 2s2 2p6)

• The F-1ion (called fluoride) is so stable and safe that we rub it on our teeth.

The same is true of other elements. For instance, chlorine is a dangerous corrosive gas. In fact, it was the first gas used as a chemical warfare agent by the German’s in WWI. Chloride (the Cl-1 ion) created when chlorine steals an electron from another atom is so safe we put it in almost everything we cook (It’s half of table salt).

Other elements will steal more electrons as needed to achieve a "noble gas structure." For instance, the most common charge that oxygen (normally 1s2 2s2 2p4 ) has in nature is -2 because the O-2 ion has an electron configuration of 1s2 2s2 2p6.