Ionic or Electrovalent Bond

4.2 Ionic or Electrovalent Bond

The ionic bond is the electrostatic force of attraction that holds together oppositely charged ions. To understand when and why such a bond forms, we need to look at the energy changes involved in creating the ions and then assembling them into a solid.

The Kössel–Lewis Picture

The formation of an ionic compound depends on two main factors:

1. How easily the positive and negative ions can be made from their neutral atoms — this is about the energy cost or gain of removing or adding electrons.
2. How the ions arrange themselves in the solid — the crystal lattice that forms when the ions come together.

Making a positive ion means pulling an electron off a neutral atom. That always costs energy — it is an endothermic process called ionization. Making a negative ion means adding an electron to a neutral atom. That can release energy (exothermic) or require energy (endothermic), depending on the atom.

The three steps can be written as:

The electron gain enthalpy, $\Delta_{\text{eg}}H$, is the enthalpy change when a gas-phase atom in its ground state gains an electron. It can be exothermic (negative $\Delta H$) or endothermic (positive $\Delta H$). The term electron affinity is the negative of the energy change accompanying electron gain — so a high negative electron gain enthalpy means a strong tendency to accept an electron.

Ionization, by contrast, is always endothermic — you must put energy in to remove an electron.

Which Elements Form Ionic Bonds?

Ionic bonds form most easily between elements that have low ionization enthalpies (so they give up electrons readily) and elements that have highly negative electron gain enthalpies (so they accept electrons eagerly).

Most cations come from metals; most anions come from non-metals. There is one important exception: the ammonium ion, $\text{NH}_4^+$, is made entirely of non-metals but acts as a cation in many ionic compounds like ammonium chloride.

The Crystal Lattice

Ionic compounds in the solid state are not made of isolated ion pairs. Instead, they form orderly three-dimensional arrangements — crystal lattices — where each cation is surrounded by anions and each anion by cations, held together by coulombic (electrostatic) forces.

The specific crystal structure depends on:

• The sizes of the ions
• How the ions pack together
• Other factors like charge

The classic example is sodium chloride, NaCl, which crystallizes in the rock salt structure. In this structure, each Na$^+$ ion is surrounded by six Cl$^-$ ions, and each Cl$^-$ by six Na$^+$ ions, in a cubic arrangement.

Why Ionic Compounds Are Stable: The Role of Lattice Enthalpy

Here is the key insight that the Kössel–Lewis picture explains. Even if the sum of the ionization enthalpy and the electron gain enthalpy is positive — meaning the gas-phase ion formation costs energy overall — the compound can still be stable because of the energy released when the ions come together to form the crystal lattice.

Consider sodium chloride:

• Ionization enthalpy of Na: $\text{Na(g)} \rightarrow \text{Na}^+\text{(g)} + e^-$ requires $+495.8\ \text{kJ mol}^{-1}$
• Electron gain enthalpy of Cl: $\text{Cl(g)} + e^- \rightarrow \text{Cl}^-\text{(g)}$ releases $-348.7\ \text{kJ mol}^{-1}$

The sum of these two steps is:

So forming the gaseous ions costs 147.1 kJ per mole — it is energetically unfavourable. But that is not the whole story. When the gaseous Na$^+$ and Cl$^-$ ions come together to form solid NaCl, a huge amount of energy is released: …