Electron Configurations of Ions and Periodic Trends

Electron Configurations of Ions and Periodic Trends

Periodicity and Electron Configuration

The periodic table is organized so that elements with similar valence electron configurations are grouped together, resulting in periodic trends in their chemical properties.

• Valence electrons are the outermost electrons, responsible for an element's chemical behavior.

• Members of the same group have the same number of valence electrons in the same types of orbitals.

• Elemental properties are periodic because valence electron configurations repeat periodically.

• The strong reducing ability of alkali metals is due to their ns1 valence-shell electron configurations.

• The strong oxidizing ability of halogens is due to their ns2np5 valence-shell electron configurations.

• The unreactivity of noble gases is due to their ns2np6 valence-shell electron configurations.

• The electronic structure of an element determines its chemical properties.

Cation Formation

Main-group metals lose electrons to form cations. The number of electrons lost is equal to the number of spaces moved backward on the periodic table to reach the nearest noble gas.

• Electrons are lost from the highest-energy shell first.

• Example: Sodium and calcium cation formation:

• $\mathrm{Na}(1s^2 2s^2 2p^6 3s^1) \rightarrow \mathrm{Na}^+(1s^2 2s^2 2p^6) + e^-$

• $\mathrm{Ca}(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2) \rightarrow \mathrm{Ca}^{2+}(1s^2 2s^2 2p^6 3s^2 3p^6) + 2e^-$

• Na+ has the same electron configuration as Ne.

• Ca2+ has the same electron configuration as Ar.

• Counting the number of spaces moved backward to the nearest noble gas gives the charge of the cation.

• Both Na+ and Ca2+ have filled valence shells (eight valence electrons).

Anion Formation

Main-group nonmetals gain electrons to form anions. The number of electrons gained is equal to the number of spaces moved forward on the periodic table to reach the nearest noble gas.

• Electrons are added to the highest-energy shell until it is filled.

• Example: Fluorine and sulfur anion formation:

• $\mathrm{F}(1s^2 2s^2 2p^5) + e^- \rightarrow \mathrm{F}^-(1s^2 2s^2 2p^6)$

• $\mathrm{S}(1s^2 2s^2 2p^6 3s^2 3p^4) + 2e^- \rightarrow \mathrm{S}^{2-}(1s^2 2s^2 2p^6 3s^2 3p^6)$

• F- has the same electron configuration as Ne.

• S2- has the same electron configuration as Ar.

• Counting the number of spaces moved forward to the nearest noble gas gives the charge of the anion.

• Both F- and S2- have filled valence shells (eight valence electrons).

Metals with Variable Charges

Some main-group metals can form more than one cation, often with different charges. These are called variable charge metals.

• Tin forms Sn2+ and Sn4+.

• Thallium forms Tl+ and Tl3+.

• Lead forms Pb2+ and Pb4+.

• The cation with the lower charge is formed by losing electrons from the highest-energy p-subshell.

• The cation with the higher charge is formed by losing electrons from both the highest-energy p- and s-subshells.

• Example (Lead):

• $[\mathrm{Xe}]6s^2 4f^{14} 5d^{10} 6p^2 \rightarrow \mathrm{Pb}^{2+}([\mathrm{Xe}]6s^2 4f^{14} 5d^{10}) + 2e^-$

• $[\mathrm{Xe}]6s^2 4f^{14} 5d^{10} 6p^2 \rightarrow \mathrm{Pb}^{4+}([\mathrm{Xe}]4f^{14} 5d^{10}) + 4e^-$

Additional info: The order of electron removal can differ from the order of filling due to energy changes upon ionization.

Cations of Transition Metals

Transition metals form cations by first losing electrons from their highest-energy s-subshell, then from the d-subshell if necessary.

• Example (Iron):

• $[\mathrm{Ar}]4s^2 3d^6 \rightarrow \mathrm{Fe}^{2+}([\mathrm{Ar}]3d^6) + 2e^-$

• $[\mathrm{Ar}]4s^2 3d^6 \rightarrow \mathrm{Fe}^{3+}([\mathrm{Ar}]3d^5) + 3e^-$

• When building up the periodic table, the d-subshell is higher in energy than the s-subshell, but when forming cations, the s-electrons are lost first.

Periodic Trends in Atomic Radii

Atomic radii show predictable trends across the periodic table.

• Size increases down a group due to increasing principal quantum number (n), which means valence electrons are farther from the nucleus.

• Size decreases across a period due to increasing effective nuclear charge ($Z_{\text{eff}}$), pulling electrons closer to the nucleus.

• Across a period, electrons are added to the same shell, so electron shielding increases only slightly, but nuclear charge increases significantly.

• $Z_{\text{eff}} = Z - \text{electron shielding}$

• As $Z_{\text{eff}}$ increases, the atom contracts.

Periodic Trends in Ionic Radii

Ionic radii also display periodic trends, which differ for cations and anions.

• Cations are smaller than their parent atoms because electrons are removed, reducing electron-electron repulsion and increasing $Z_{\text{eff}}$.

• Anions are larger than their parent atoms because adding electrons increases electron-electron repulsion and decreases $Z_{\text{eff}}$.

• Ion size increases down a group for both cations and anions.

Comparing Cations and Anions to Parent Atoms

• For cations: Electrons are removed from the outermost shell, resulting in a smaller radius.

• For anions: Electrons are added to the outermost shell, resulting in a larger radius.

• Reducing the number of electrons reduces electron shielding, increasing $Z_{\text{eff}}$ and shrinking the ion.

• Increasing the number of electrons increases electron shielding, decreasing $Z_{\text{eff}}$ and expanding the ion.

Sizes of Isoelectronic Species

Isoelectronic species have the same electron configuration but different nuclear charges. Their sizes depend on $Z_{\text{eff}}$.

Ion	Approximate Radius (pm)	Electron Configuration
F-	130	1s22s22p6
Na+	100	1s22s22p6
Mg2+	60	1s22s22p6

• All three ions are isoelectronic, but Mg2+ has the highest $Z_{\text{eff}}$ and is the smallest, while F- has the lowest $Z_{\text{eff}}$ and is the largest.

• For isoelectronic species, size decreases as nuclear charge increases.

Summary Table: Periodic Trends in Atomic and Ionic Radii

Trend	Atomic Radius	Cation Radius	Anion Radius
Down a group	Increases	Increases	Increases
Across a period	Decreases	Decreases	Decreases
Compared to parent atom	—	Smaller	Larger

Key Points

• Main-group metals form cations by losing electrons from their highest-energy shell.

• Main-group nonmetals form anions by gaining electrons to fill their highest-energy shell.

• Metals with variable charges can lose different numbers of electrons, forming cations with different charges.

• Transition metals lose s-electrons before d-electrons when forming cations.

• Atomic and ionic radii show predictable periodic trends based on electron configuration and effective nuclear charge.