Ionic Compounds: Periodic Trends and Bonding Theory

Chapter 6: Ionic Compounds – Periodic Trends and Bonding Theory

Electron Configurations of Ions

Understanding the electron configurations of ions is essential for predicting their chemical behavior and stability. Main-group elements tend to form ions that achieve noble-gas configurations, while transition metals may have additional filled d subshells.

• Electron Configuration: The arrangement of electrons in an atom or ion, described by principal energy levels and sublevels (s, p, d, f).

• Formation of Cations: Atoms lose electrons, typically from the highest principal quantum number (n), to form positively charged ions.

• Formation of Anions: Atoms gain electrons to fill their valence shell, forming negatively charged ions.

• Noble-Gas Configuration: Ions often achieve the same electron configuration as the nearest noble gas, with filled s and p sublevels in the outermost shell.

Example: Sodium (Na) loses one electron to form Na+ with the same configuration as neon (Ne): $\mathrm{Na}: 1s^2\ 2s^2\ 2p^6\ 3s^1 \rightarrow Na^+: 1s^2\ 2s^2\ 2p^6$

Table 6.1: Some Common Main-Group Ions and Their Noble-Gas Electron Configurations

Ion	Electron Configuration	Noble Gas
Na+	1s2 2s2 2p6	[Ne]
Mg2+	1s2 2s2 2p6	[Ne]
Cl-	1s2 2s2 2p6 3s2 3p6	[Ar]
Ca2+	1s2 2s2 2p6 3s2 3p6	[Ar]
O2-	1s2 2s2 2p6	[Ne]
Al3+	1s2 2s2 2p6	[Ne]
Fe2+	[Ar] 3d6	*
Cu+	[Ar] 3d10	*

*These ions do not have a true noble-gas configuration due to filled d subshells.

Ionic Radii

The size of an ion (ionic radius) depends on its charge and the arrangement of electrons. Ionic radii influence the structure and properties of ionic compounds.

• Cations: Smaller than their parent atoms because they lose electrons, resulting in a smaller principal quantum number for the valence shell and increased effective nuclear charge.

• Anions: Larger than their parent atoms due to the addition of electrons, which increases electron–electron repulsions and decreases effective nuclear charge per electron.

Example: The radius of Na+ is much smaller than Na, while Cl- is larger than Cl.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom. It reflects how strongly an atom holds its electrons and varies across the periodic table.

• Definition: The amount of energy necessary to remove the highest-energy electron from an isolated neutral atom in the gaseous state.

• Trends:

  • Increases across a period (left to right) due to increasing nuclear charge.

  • Decreases down a group due to increasing atomic radius and electron shielding.

• Exceptions:

  • Group 2A elements (Be, Mg, Ca) have slightly higher ionization energies than expected.

  • Group 6A elements (O, S) have slightly lower ionization energies due to electron repulsion in filled orbitals.

  • Boron has a lower ionization energy due to increased shielding by 2s electrons.

Equation: $\mathrm{X}(g) \rightarrow \mathrm{X}^+(g) + e^- \qquad \Delta E = \text{Ionization Energy}$

Higher Ionization Energies

Successive ionization energies refer to the energy required to remove additional electrons after the first. Each successive ionization energy is higher than the previous one, with large jumps when removing electrons from a new shell.

• Trend: Large increases in ionization energy occur when an electron is removed from a stable, noble-gas configuration.

Table 6.2: Higher Ionization Energies for Main-Group Third Row Elements

Element	1st IE (kJ/mol)	2nd IE (kJ/mol)	3rd IE (kJ/mol)
Na	496	4562	6910
Mg	738	1451	7733
Al	578	1817	2745
Si	786	1577	3232
P	1012	1907	2914
S	1000	2252	3357
Cl	1251	2298	3822

Additional info: The zigzag line in the table marks large jumps in ionization energies, indicating removal from a new shell.

Electron Affinity

Electron affinity is the energy change when an electron is added to a neutral atom in the gaseous state. It measures the tendency of an atom to accept an electron.

• Negative Value: Energy is released (exothermic) when an electron is added, as seen in halogens (Group 7A).

• Zero Value: Energy is absorbed (endothermic) or not measurable, as in Group 2A (alkaline earths) and Group 8A (noble gases).

Equation: $\mathrm{X}(g) + e^- \rightarrow \mathrm{X}^-(g) \qquad \Delta E = \text{Electron Affinity}$

Example: Chlorine has a high (negative) electron affinity, favoring the formation of Cl-.

The Octet Rule

The octet rule states that main-group elements tend to react to achieve eight electrons in their outer shell, attaining a noble-gas configuration with filled s and p sublevels.

• Metals: Low ionization energies and electron affinities; tend to lose electrons and form cations.

• Nonmetals: High ionization energies and electron affinities; tend to gain electrons and form anions.

Example: Sodium (Na) loses one electron to form Na+, while chlorine (Cl) gains one electron to form Cl-, both achieving octets.

Ionic Bonds and the Formation of Ionic Solids

Ionic bonds form when electrons are transferred from metals to nonmetals, resulting in oppositely charged ions that attract each other and form ionic solids.

• Ionic Bond: The electrostatic attraction between cations and anions in an ionic compound.

• Formation: Typically involves a metal losing electrons and a nonmetal gaining electrons.

Born-Haber Cycle: A thermochemical cycle used to analyze the steps in the formation of an ionic solid from its elements.

• Step 1: Sublimation of the metal (solid to gas)

• Step 2: Ionization of the metal atom

• Step 3: Dissociation of the nonmetal molecule

• Step 4: Addition of an electron to the nonmetal atom (electron affinity)

• Step 5: Formation of the ionic solid (lattice energy)

Equation (Lattice Energy): $\mathrm{MX}(s) \rightarrow \mathrm{M}^+(g) + \mathrm{X}^-(g) \qquad U = \text{Lattice Energy}$

Lattice Energies in Ionic Solids

Lattice energy is the energy required to separate one mole of an ionic solid into its gaseous ions. It is a measure of the strength of the ionic bonds in a solid.

• Trend: Lattice energy increases with higher ionic charges and smaller ionic radii.

Table 6.3: Lattice Energies of Some Ionic Solids (kJ/mol)

Compound	Lattice Energy (kJ/mol)
NaCl	786
MgO	3795
CaO	3414
LiF	1036
KBr	671

Additional info: Higher lattice energies correspond to stronger ionic bonds and higher melting points.