Ionic Bonding and Lattice Structures

Ionic bonds typically form between metals (Groups 1, 2, and some 13) and non-metals (Groups 16, 17). Metals tend to lose valence electrons to achieve a stable noble gas electron configuration, forming positively charged cations. For example, sodium (Na) loses one electron to become Na$^+$. Non-metals tend to gain valence electrons to achieve a stable noble gas electron configuration, forming negatively charged anions. For instance, chlorine (Cl) gains one electron to become Cl$^-$. The driving force behind this electron transfer is the octet rule, which states that atoms tend to achieve a full outer shell of eight electrons (or two for elements like hydrogen and lithium). The electrostatic attraction between these oppositely charged ions constitutes the ionic bond. This process is highly exothermic, releasing a significant amount of energy, which contributes to the stability of the ionic compound. The difference in electronegativity between the two atoms involved is usually large (>1.7 on the Pauling scale).

To illustrate the formation of ionic bonds, we use dot-and-cross diagrams. These diagrams show the valence electrons of each atom before and after electron transfer. The original valence electrons of one atom are represented by dots, and those of the other by crosses.

Example: Sodium Chloride (NaCl)

• Sodium atom (Na): 2, 8, 1 (1 valence electron, represented by a dot)
• Chlorine atom (Cl): 2, 8, 7 (7 valence electrons, represented by crosses)

Sodium transfers its one valence electron to chlorine.

• Sodium ion (Na$^+$): [Na]$^+$ (2, 8) - now has a full outer shell.
• Chloride ion (Cl$^-$): [Cl]$^-$ (2, 8, 8) - now has a full outer shell, with the gained electron shown as a dot.

The electrostatic attraction between Na$^+$ and Cl$^-$ forms the ionic bond. Remember to include the charges on the ions and square brackets around the electron shells of the ions in your diagrams.

Ionic compounds exhibit a characteristic set of properties due to the strong electrostatic forces within their crystal lattices:

• High Melting and Boiling Points: A large amount of energy is required to overcome the strong electrostatic forces of attraction between ions in the lattice, hence high melting and boiling points.
• Solubility in Polar Solvents: Many ionic compounds are soluble in polar solvents like water. Water molecules, being polar, can surround and separate the individual ions from the lattice through a process called hydration, overcoming the lattice forces.
• Electrical Conductivity:
  • Solid State: Ionic compounds do not conduct electricity in the solid state because the ions are fixed in the lattice and cannot move to carry charge.
  • Molten (Liquid) State: They conduct electricity when molten because the ions are free to move and carry charge.
  • Aqueous Solution: They conduct electricity when dissolved in water because the dissociated ions are free to move.
• Brittleness: When a force is applied to an ionic crystal, layers of ions can shift. This brings ions of the same charge into alignment, leading to strong repulsive forces that cause the crystal to cleave or shatter.