Lewis structures show the arrangement of valence electrons in covalently bonded molecules and ions. We can use Lewis structures to predict the properties of the bonds in molecules and ions, such as the properties of bond length and bond energy, and use these bond properties to predict physical and chemical properties of molecules and ions.
One factor that has a great influence on bond properties is bond order, the number of bonding electron pairs between two bonded atoms. The bond order of the carbon–carbon single bond in acetone is 1, for example, whereas the carbon–oxygen double bond in acetone has a bond order of 2. The carbon–nitrogen triple bond in the cyanide ion has a bond order of 3.
As we saw earlier, chemical bonds form when attractive forces between atoms are stronger than repulsive forces. The distance between the atomic nuclei when energy is minimized is the bond length.
Accurate bond distances are determined from careful measurements using techniques such as x-ray crystallography. Bond lengths between two different elements vary slightly between compounds. For that reason, average bond lengths are reported in tables, such as Interactive Table 8.3.1.
|Multiple Bonds C≡O|
|O=O (in O2)||112||N≡O||108|
The bond lengths in Interactive Table 8.3.1 have units of picometers (1 pm = 10-12 m). Other common units for bond lengths are nanometers (1 nm = 10-9 nm) and Ångstroms (1 Å = 10-10 m). The data in Interactive Table 8.3.1 demonstrate two general trends in bond lengths. First, bond lengths increase with increasing atom size. Consider the trend in H—X bond lengths (where X is a halogen). Atomic radii increase as you move down the periodic table (F < Cl < Br < I), and therefore the H—X bond lengths increase as the halogen radius increases.
The second trend demonstrated in Interactive Table 8.3.1 is the relationship between bond length and bond order. As bond order increases, there is an increase in electron density between two nuclei. This results in stronger attractive forces between electrons and nuclei, decreasing the distance between the nuclei. A carbon–carbon single bond has a bond order of 1 and is longer than a carbon–carbon double bond with a bond order of 2. In general, for a series of bonds that differ only in bond order, an increase in bond order results in a decrease in bond length.
C—O 143 pm > C=O 122 pm > C≡O 113 pm
As shown previously in Interactive Figure 8.1.1, a chemical bond forms when the attractive and repulsive forces between atoms results in an energy minimum. Bond energy is the energy required to break a chemical bond in a gas-phase molecule. For example, the bond energy of an H—H bond is 436 kJ/mol (at 298 K). This means that 436 kJ of energy is required to break 1 mol of H—H bonds, forming 2 mol of H atoms.
H—H(g) → H(g) + H(g) ΔH° = +436 kJ/mol
Breaking bonds is an endothermic process; therefore, bond energies are always positive values. Conversely, bond formation is an exothermic process that always releases energy. Just as bond lengths vary between compounds, so do bond energies. Some average bond energies are shown in Interactive Table 8.3.2.
|Multiple Bonds C≡O|
|C=N||616||C=O (as in CO2, O=C=O)||803|
|C≡N||866||C≡O (as in H2C=O)||695|
|O=O (in O2)||498||C≡O||1073|
A clear trend that can be observed from the data in Interactive Table 8.3.2 is the relationship between bond energy and bond order. Bond energy increases with increasing bond order:
C—O 351 kJ/mol < C=O 745 kJ/mol < C≡O 1075 kJ/mol
As bond order increases, the stronger attractive forces between bonding electrons and nuclei mean that more and more energy is required to separate the bonded nuclei.
You are asked to identify the bond order in a molecule or ion and to use bond order to make predictions about relative bond length or bond order in a series of molecules or ions.
You are given the chemical formula of a molecule or ion, or a series of molecules or ions.