Chemical bonds are the forces that hold atoms together, forming molecules and compounds. The strength of a chemical bond is measured by the energy required to break that bond. For covalent bonds, this measure is typically quantified as the bond dissociation energy (BDE), which is the energy needed to cleave one mole of a specific bond in a gaseous molecule. Ionic compounds form extended crystal lattices, and their strength is measured by lattice energy, representing the energy required to separate one mole of the solid compound into its gaseous ions. Understanding the factors that influence these energy values dictates a substance’s chemical reactivity, physical properties, and overall stability.
The Role of Bond Multiplicity
The number of electron pairs shared between two atoms, known as bond multiplicity or bond order, determines covalent bond strength. As the number of shared electrons increases, the bond becomes both shorter and stronger. A single bond involves one electron pair, a double bond involves two, and a triple bond involves three.
An increase in shared electron pairs results in a greater concentration of electron density between the two atomic nuclei. This enhanced density increases the attractive electrostatic force holding the nuclei together. For example, a C-C single bond is longer and weaker than a C=C double bond, which is in turn longer and weaker than a C≡C triple bond.
The formation of multiple bonds involves the overlap of atomic orbitals. A single bond consists of a sigma (s) bond formed by the head-on overlap of orbitals. Double and triple bonds add one or two pi (p) bonds, respectively, formed by the side-by-side overlap of p-orbitals. Each additional pi bond contributes significantly to the total bond energy, pulling the nuclei closer and requiring greater energy to break the connection.
Impact of Atomic Size and Bond Length
The distance between the nuclei of two bonded atoms, known as bond length, has an inverse relationship with bond strength. Shorter bonds are generally stronger because the attractive forces between the shared electrons and the two nuclei are more effective over shorter distances. This internuclear distance is influenced by the size of the atoms involved.
As one moves down a group on the periodic table, the atomic radius increases because valence electrons occupy shells farther from the nucleus. When these larger atoms form bonds, the resulting bond length is longer. For instance, the bond between hydrogen and a halogen atom weakens as the halogen size increases, meaning the H-F bond is significantly stronger than the H-I bond.
The larger size of the orbitals in heavier atoms leads to less efficient overlap when forming a bond. This reduced orbital overlap results in a lower electron density concentrated between the nuclei, which weakens the overall electrostatic attraction. Comparing carbon-halogen bonds, the C-I bond is noticeably longer and weaker than the C-Br bond because iodine is a larger atom than bromine.
Influence of Electronegativity and Charge Distribution
The distribution of electron density within a covalent bond is governed by the electronegativity of the bonded atoms, which significantly influences the bond’s strength. When two atoms of different electronegativities form a bond, the electrons are shared unequally, creating a polar covalent bond.
This unequal sharing generates a partial negative charge (delta-) on the more electronegative atom and a partial positive charge (delta+) on the less electronegative atom, forming a molecular dipole. The creation of these partial charges introduces an ionic character to the bond, adding an electrostatic attraction component. The greater the difference in electronegativity, the more polarized the bond becomes.
The resulting electrostatic attraction between the partial poles adds significantly to the strength derived from the covalent sharing of electrons. For example, the bond in hydrogen fluoride (H-F) is stronger than the bond in methane (C-H) because fluorine is much more electronegative. This greater polarity in the H-F bond leads to a higher bond dissociation energy.
The total strength of the polar covalent bond is greater than what would be predicted from simple covalent attraction alone. This trend continues until the electronegativity difference becomes so large that a complete electron transfer occurs, resulting in a fully ionic bond.
How Ion Charge Affects Ionic Bond Strength
For compounds that form purely ionic bonds, the primary determinant of strength is the electrostatic attraction between the oppositely charged ions, quantified by the lattice energy. This attraction is governed by Coulomb’s Law, meaning the energy required to break the ionic bonds increases dramatically with the magnitude of the ion charges.
When comparing two different ionic compounds, the product of the charges on the cation and the anion is the most important factor. For instance, sodium chloride (NaCl) consists of ions with charges of +1 and -1, resulting in a charge product of one. In contrast, magnesium oxide (MgO) consists of ions with charges of +2 and -2, resulting in a charge product of four.
The four-fold increase in the charge product for MgO translates into a significantly greater electrostatic force and a much higher lattice energy compared to NaCl. Even though the internuclear distances are similar, the bond in magnesium oxide is substantially stronger. This difference is reflected in properties like the melting point, which is higher for MgO than for NaCl.