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Ionic bonding is the electrostatic force of attraction between positively (+ve) and negatively (−ve) charged ions (primarily between non-metals such as chloride or fluoride ions and metals such as sodium or potassium ions). Most ionic compounds are crystalline solids at room temperature.

The crucial feature of an ionic bond is that each atom either gains or loses an electron so that the resulting ion has its lowest energy (noble gas-like) configuration. Table salt, NaCl, is a typical example of ionic bonding, and you can see its structure in Figure 3.4. In this salt, sodium gives up an electron, and chlorine gains this electron so that both ions gain a noble gas configuration, as we saw in Section 3.4. In other words, the Na atom has transferred its electron to the Cl atom, and the result is two ions, Na⁺ and Cl⁻.

Figure 3.4 The structure of NaCl showing the alternating sodium and chloride ions.

Source: Reproduced with permission of B. Blaus, https://commons.wikimedia.org/wiki/Category:Crystal_structure_of_sodium_chloride#/media/File:Blausen_0660_NaCl.png.

After transferring an electron, we now have two ions, Na⁺ and Cl⁻, with opposite charges. They are attracted to one another. However, if we have many of these ions, then things get more complicated. Clearly, negatively charged chloride ions will be repelled from each other, and positively charged sodium ions will be repelled from each other. If we place many Na⁺ and Cl⁻ ions together, the natural configuration they take up to maximize attraction and minimize repulsion is an alternating packed cubic structure (Figure 3.4). Other similar examples are cesium chloride (CsCl) and sodium fluoride (NaF).

Ionic bonds are typically very strong. We can consider the stability of these bonds from an energetic (thermodynamic) point of view. Let's consider sodium fluoride, NaF. The energy required to break an ionic chemical bond in this structure is about 3 × 10⁻¹⁹ J. We can also calculate the typical thermal energy in a bond at a specified temperature. This is the thermal energy that would be in the bond when it is in equilibrium with a given environment at a particular temperature. This can be approximated by ∼k_BT. k_B is the Boltzmann constant, which has a value of 1.381 × 10⁻²³ JK⁻¹, and T is the temperature, measured in Kelvin. At room temperature (about 300 K), the value of k_BT is 4.1 × 10⁻²¹ J. The ionic bond energy is therefore about 75 times the thermal energy in the bond. In other words, the thermal energy in NaF at room temperature will not cause the ions to break apart. You need to put in a lot more energy to break apart the ionic bonds, explaining why this compound is stable at room temperature. This is the case for many ionic bonds.