CHM 1150: Formation of Ions, Size, and Ionization Energy

Main group elements form ions by achieving the electron configuration of the nearest noble gas, becoming isoelectronic.
Cations: formed by losing highest energy electrons.
Anions: formed by adding electrons into the lowest energy orbital.
Transition metals: form cations of various oxidation states, not necessarily isoelectronic with noble gases.

Anion is always larger than its neutral element ( $O^{2-}$ > O) due to increased electron-electron repulsions.
Cation is always smaller than its neutral element ( $Al^{3+}$ < Al) due to:
- Valence electrons in smaller $(n-1)$ shell orbitals for cations vs. $n$ shell for neutral.
- Larger effective nuclear charge ( $Z_{eff}$ ) in cations, increasing electron attraction.
For transition metals, a higher cation charge results in a smaller cation ( $Fe^{3+}$ < $Fe^{2+}$ ).

Definition: Energy required to remove the highest energy electron from an isolated gaseous atom or ion.
First Ionization Energy ( $E_{i1}$ ): $M(g) \rightarrow M^+(g) + e^-$ .
Successive Ionization Energies: Ei1<Ei2<Ei3… (removing electrons from increasingly positive ions).
**Trends in $E_{i1}$ :
- Across a period (left to right): Generally increases due to increasing $Z_{eff}$ . Lowest for alkali metals, highest for noble gases.
- Down a group: Generally decreases as the highest energy electron is in a higher $n$ shell, further from the nucleus, and easier to remove.
"Large jump" in $E_i$ : Occurs when an electron is removed from a stable noble gas electron configuration, providing evidence for electron shell structure.