Ionic or Electrovalent Bond

Full Notes

Introduction to Ionic Bonding

An ionic bond (or electrovalent bond) is the electrostatic attraction between oppositely charged ions.

The ions are formed when one atom donates an electron and another accepts it, resulting in positively and negatively charged ions.

Ionic bonding typically occurs between metals and non-metals.

• Metals (like Na, Mg, Ca) lose electrons and become cations.
• Non-metals (like Cl, O, F) gain electrons and become anions.

Example Formation of Sodium Chloride (NaCl)

Below shows the electron transfer and ion formation for Na and Cl.

Sodium (Na): 1 valence electron → loses 1e⁻ → Na⁺

Chlorine (Cl): 7 valence electrons → gains 1e⁻ → Cl⁻

These ions arrange to maximise attraction and form NaCl, forming a giant lattice that is held together by electrostatic attraction.

Factors Favouring Formation of Ionic Bonds

Ionic bonding is influenced by energy considerations:

1. Low ionisation enthalpy of the metal.
2. High (more negative) electron gain enthalpy of the non-metal.
3. High lattice enthalpy (explained next).

Ionic bonds are formed more easily between elements with comparatively low ionisation enthalpies and elements with comparatively high negative value of electron gain enthalpy. In other words, ionic bonds form easily between an element that easily loses electrons and an element that easily gains an electron.

4.2.1 Lattice Enthalpy

Lattice enthalpy is defined as the energy released when one mole of an ionic crystalline solid is formed from its gaseous ions.

• It is also described as the energy required to completely separate one mole of an ionic compound into gaseous ions.
• It reflects the strength of the ionic bond in a crystal lattice.
• As energy is required to overcome attraction between ions, lattice enthalpies are endothermic (+ΔH). They are sometimes referred to as lattice energies.

For exampleThe lattice enthalpy of NaCl(s) is the enthalpy change that occurs when 1 mole of NaCl(s) is separated completely into 1 mole of Na⁺(g) and 1 mole of Cl⁻(g) ions.

Key Point

The greater the lattice enthalpy the stronger the ionic bond and the more stable the ionic compound.

Factors Affecting Lattice Enthalpy

• Charge: Higher ion charge = stronger electrostatic attraction between ions = larger magnitude lattice enthalpy.
  Example: Mg²⁺ forms greater lattice enthalpies than Na⁺.
• Ionic radius: Smaller radius → stronger attraction → larger magnitude lattice enthalpy.
  Down a group, ionic radius increases and lattice enthalpy decreases in magnitude.

Summary

• Ionic bonds form by electron transfer and the electrostatic attraction between cations and anions.
• Lattice enthalpy measures the strength of the ionic lattice and can be viewed as energy released on formation or energy required for complete separation.
• Lattice enthalpy increases in magnitude with higher ionic charge and smaller ionic radius.
• Low ionisation enthalpy (metals) and highly negative electron gain enthalpy (non-metals) favour ionic bond formation.