7.5 Strengths of Ionic and Covalent Bonds

Bond Energies and Enthalpies

Chemical bonds store energy. The amount of energy it takes to break a bond (or the amount released when one forms) tells you how strong that bond is. By tracking these energies across a whole reaction, you can estimate whether the reaction releases or absorbs energy overall.

Energy Changes in Bond Types

Covalent bonds form when atoms share electrons. Forming a covalent bond releases energy (exothermic), and the amount released is called the bond dissociation energy. That same amount of energy must be put back in to break the bond (endothermic).

• The H–H bond has a dissociation energy of 436 kJ/mol, meaning you need 436 kJ to break one mole of H–H bonds.
• The Cl–Cl bond is weaker at 243 kJ/mol, so it takes less energy to pull apart.
• Larger electronegativity differences between the two atoms generally make covalent bonds more polar and stronger.

Ionic bonds involve a full transfer of electrons from a metal to a nonmetal, producing oppositely charged ions that attract each other. Forming an ionic compound from its elements isn't a single step. It's a sequence of energy changes:

1. Sublimation of the metal: solid → gas (endothermic). For Na: +107 kJ/mol.
2. Ionization of the metal atom: removes an electron to form a cation (endothermic). For Na: +496 kJ/mol.
3. Electron affinity of the nonmetal: adds an electron to form an anion (exothermic). For Cl: −349 kJ/mol.
4. Lattice energy released when gaseous ions come together to form a solid crystal (exothermic). For NaCl: −787 kJ/mol.

The lattice energy is so large that it more than compensates for the endothermic steps. That's why the overall formation of NaCl is exothermic (−411 kJ/mol).

Breaking an ionic compound apart requires energy input equal to the lattice energy, because you're overcoming the strong electrostatic attractions between oppositely charged ions packed in a crystal. LiF, for example, has a lattice energy of 1030 kJ/mol, reflecting the very strong attraction between small, highly charged ions.

Lattice Energy Calculations

The Born-Haber cycle is a thermochemical cycle that lets you calculate lattice energy indirectly, using Hess's law. You arrange all the individual steps of ionic compound formation into a cycle and solve for the one value you can't measure directly.

Here are the steps, using NaCl as an example:

1. Sublimation of the metal: $\Delta H_{sub}$ = +107 kJ/mol (Na solid → Na gas)
2. Ionization of the metal: $\Delta H_{ion}$ = +496 kJ/mol (Na gas → Na⁺ gas)
3. Dissociation of the nonmetal molecule: $\frac{1}{2}\Delta H_{diss}$ = +121 kJ/mol (½ Cl₂ gas → Cl gas)
4. Electron affinity of the nonmetal: $\Delta H_{ea}$ = −349 kJ/mol (Cl gas → Cl⁻ gas)
5. Formation of the ionic solid from gaseous ions: $\Delta H_{lattice}$ (this is what you're solving for)

Since the overall formation enthalpy $\Delta H_f$ of NaCl is −411 kJ/mol, you can rearrange:

$\Delta H_{lattice} = \Delta H_f - \Delta H_{sub} - \Delta H_{ion} - \frac{1}{2}\Delta H_{diss} - \Delta H_{ea}$

Note that step 3 uses half the dissociation energy of $Cl_2$ (242 kJ/mol total) because you only need one Cl atom per formula unit of NaCl.

Lattice energy varies a lot between compounds. NaCl has a lattice energy of −787 kJ/mol, while MgO has −3791 kJ/mol. The difference comes down to charge and size: Mg²⁺ and O²⁻ have higher charges and smaller radii than Na⁺ and Cl⁻, so the electrostatic attraction is much stronger.

Reaction Enthalpy Estimations

You can estimate the enthalpy change of a reaction using average bond energies. The idea is straightforward: energy goes in to break bonds in the reactants, and energy comes out when new bonds form in the products.

1. Draw out the structural formulas and identify every bond broken in the reactants and every bond formed in the products.

2. Add up the energy required to break all reactant bonds. Multiply the count of each bond type by its average bond energy.

   • Common values: C–H = 413 kJ/mol, O=O = 498 kJ/mol

3. Add up the energy released when forming all product bonds.

   • Common values: C=O = 745 kJ/mol, O–H = 463 kJ/mol

4. Apply the formula:

$\Delta H_{rxn} = \Sigma \Delta H_{bonds\ broken} - \Sigma \Delta H_{bonds\ formed}$

If the result is negative, the reaction is exothermic (more energy released than absorbed). If positive, it's endothermic. For the combustion of methane, this method gives approximately −802 kJ/mol, confirming it releases a lot of energy.

This approach gives an estimate, not an exact value. Average bond energies are averages across many different molecules, so they don't account for the specific molecular environment of each bond or intermolecular forces. Still, it's a reliable way to predict whether a reaction is exothermic or endothermic.

Factors Affecting Bond Strength

Several factors determine how strong a particular bond will be:

• Bond length: Shorter bonds are generally stronger. When atoms are closer together, their orbitals overlap more, creating a stronger attraction.
• Atomic radius: Smaller atoms form shorter, stronger bonds because their nuclei are closer to the shared electrons.
• Bond polarity: A greater electronegativity difference pulls electron density unevenly, which can strengthen the bond. Compare H–F (568 kJ/mol) to H–H (436 kJ/mol).
• Valence electron arrangement: The number of available valence electrons affects how many bonds an atom can form and how strong those bonds are. Triple bonds (like N≡N at 946 kJ/mol) are stronger than double bonds, which are stronger than single bonds between the same atoms.