Ion Configurations

Neutral atoms gain or lose electrons to form ions. Writing the electron configuration of an ion follows different rules for anions and cations, and there is one critical trap that the MCAT loves to test.

Anions (Negative Ions)

Anions are formed when atoms gain electrons. The extra electrons fill into the next available subshell following the same Aufbau rules as neutral atoms.

Example: Fluorine (Z = 9) has the configuration 1s²2s²2p⁵. The fluoride ion (F⁻) gains one electron: 1s²2s²2p⁶ - the same configuration as neon.

Example: Oxygen (Z = 8) is 1s²2s²2p⁴. The oxide ion (O²⁻) gains two electrons: 1s²2s²2p⁶ - also isoelectronic with neon.

Cations (Positive Ions)

Cations are formed when atoms lose electrons. Here is where the critical rule comes in:

For transition metals, electrons are removed from the highest n value first - even if that is not the subshell that filled last.

This means for transition metals: s electrons are removed before d electrons, even though 4s fills before 3d.

Why the s Electrons Leave First

Although 4s fills before 3d (because 4s has a lower n + l value), once the d subshell begins to fill, the 3d electrons actually become lower in energy than the 4s electrons. This happens because d electrons penetrate more closely to the nucleus and become more tightly held. The 4s electrons, being in a higher shell, are farther from the nucleus and easier to remove.

Common Transition Metal Ions

Isoelectronic Species

Isoelectronic species are atoms or ions that have the same number of electrons (and therefore the same electron configuration). For example:

N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺ - all have 10 electrons and the configuration 1s²2s²2p⁶.

Despite having the same electron configuration, these species have different sizes because they have different numbers of protons pulling on those 10 electrons. More protons = smaller radius. This concept is essential for understanding ionic radius trends (covered in Chapter 2).