Electron Configurations of Ions and Periodic Trends

Electron Configurations of Ions

Noble-Gas Electron Configurations

Noble gases possess completely filled s and p sublevels in their outermost electron shell, making them particularly stable and chemically unreactive. This stability is often used as a reference when writing abbreviated electron configurations for other elements and ions.

• Noble-gas configuration: The electron configuration of an element or ion written by starting with the symbol of the nearest preceding noble gas in brackets, followed by the remaining electrons.

• Example: Sodium atom: $1s^2 2s^2 2p^6 3s^1$ can be abbreviated as $[Ne] 3s^1$.

• Application: Ions often achieve noble-gas configurations by losing or gaining electrons.

Abbreviated Electron Configurations of Cations

Alkali-Metal Cations

Alkali metals (Group 1) lose one electron to form cations with a noble-gas configuration.

• Li+: $1s^2$ or $[He]$

• Na+: $1s^2 2s^2 2p^6$ or $[Ne]$

• K+: $1s^2 2s^2 2p^6 3s^2 3p^6$ or $[Ar]$

• Rb+: $[Kr]$

• Cs+: $[Xe]$

Alkaline-Earth-Metal Cations

Alkaline earth metals (Group 2) lose two electrons to form cations with a noble-gas configuration.

• Be2+: $1s^2$ or $[He]$

• Mg2+: $[Ne]$

• Ca2+: $[Ar]$

• Sr2+: $[Kr]$

• Ba2+: $[Xe]$

Other Metal Cations

Transition metals and other metals can form cations by losing electrons, often from the outermost s orbital first, then from d orbitals.

• Example: Fe2+ loses two electrons: $[Ar] 3d^6$

• Example: Fe3+ loses three electrons: $[Ar] 3d^5$

Aluminum Ion Formation

Aluminum (Al) typically loses three electrons to form the Al3+ ion, achieving a noble-gas configuration.

• Al3+: $[Ne]$

• Expected charge: 3+

Unpaired Electrons in Transition Metal Ions

Transition metal ions may have unpaired electrons depending on their electron configuration.

• Rh3+: Remove three electrons from Rh ($[Kr] 4d^8$), resulting in $[Kr] 4d^6$.

• Unpaired electrons: For $4d^6$, typically 4 unpaired electrons (depends on crystal field splitting).

• Example: Sn2+ ($[Kr] 4d^{10} 5s^2$) has 0 unpaired electrons.

Abbreviated Electron Configurations of Anions

Halide Anions

Halogens (Group 17) gain one electron to form anions with a noble-gas configuration.

• F-: $[Ne]$

• Cl-: $[Ar]$

• Br-: $[Kr]$

• I-: $[Xe]$

Chalcogen Anions

Chalcogens (Group 16) gain two electrons to form anions with a noble-gas configuration.

• O2-: $[He]$

• S2-: $[Ne]$

• Se2-: $[Ar]$

• Te2-: $[Kr]$

Phosphorus Anion Formation

Phosphorus (P) can gain three electrons to form the P3- ion, achieving a noble-gas configuration.

• P3-: $[Ne]$

• Expected charge: 3-

Periodic Trends: Atomic and Ionic Radii, Ionization Energy

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron shell. It generally increases down a group and decreases across a period.

• Smallest atom: Helium (He) is the smallest; among alkali metals, Li is smallest.

• Largest atom: Cesium (Cs) is the largest alkali metal.

Ionic Radius

Ionic radius refers to the size of an ion. Cations are smaller than their parent atoms, while anions are larger.

Ion	Group 1	Group 2	Group 13	Group 16	Group 17
Na+	102 pm
Mg2+		72 pm
Al3+			53 pm
O2-				140 pm
F-					133 pm

• Order of increasing radius: Na+ < F- < O2-

• Largest ion: O2- among those listed

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom or ion in the gas phase. It increases across a period and decreases down a group.

• Smallest ionization energy: Alkali metals (e.g., Na) have the lowest first ionization energies.

• Second ionization energy: Removing a second electron is much harder for alkali metals, as it requires removing an electron from a noble-gas configuration.

Summary Table: Electron Configurations and Charges

Element/Ion	Abbreviated Configuration	Common Charge
Na+	[Ne]	+1
Mg2+	[Ne]	+2
Al3+	[Ne]	+3
Cl-	[Ar]	-1
O2-	[Ne]	-2

Additional info: The number of unpaired electrons in transition metal ions depends on the specific electron configuration and crystal field effects, which may be covered in more advanced courses.