a) Write the Lewis structures for \(\mathrm{N}_{2}\) and \(\mathrm{O}_{2}\). b) Explain how these Lewis structures are consistent with the relative enthalpies of atom combination for \(\mathrm{N}_{2}(\mathrm{~g})\) and \(\mathrm{O}_{2}(\mathrm{~g})\).

The enthalpy of atom combination, also known as bond enthalpy, is a measure of the energy change when atoms combine to form molecules. When bonds are formed during the combination of atoms, energy is released, and this release is reflected as a negative change in enthalpy. The magnitude of enthalpy change can be directly linked to the strength and number of bonds formed.
In the case of diatomic molecules like nitrogen (N₂) and oxygen (O₂), they consist of triple and double bonds, respectively. Since more energy is required to break the triple bond in nitrogen compared to the double bond in oxygen, the enthalpy change for the formation of N₂ is more negative than that of O₂. This means that forming a molecule of nitrogen releases more energy compared to forming a molecule of oxygen, highlighting a stronger bond in nitrogen.
Furthermore, the concept of enthalpy of atom combination helps in predicting the relative stability of molecules. Molecules with more negative enthalpy values are generally more stable, as it would require more energy to break them apart.

Valence electrons are the outermost electrons of an atom and are crucial in determining an element's chemical properties. They are responsible for the atom's ability to form bonds with other atoms. For instance, in the Lewis structures exercise, nitrogen has five valence electrons, while oxygen has six.
By using the valence electrons to create Lewis structures, we can graphically represent possible bonding scenarios. Each element follows the octet rule by striving to complete an outer shell of eight electrons, either by sharing, losing, or gaining electrons.
Nitrogen atoms reach an octet by sharing three electrons with another nitrogen atom, resulting in a triple bond. On the other hand, oxygen atoms achieve an octet by sharing two electrons with another oxygen atom, forming a double bond. These bonding patterns are all dictated by the number of valence electrons and result in stable molecular structures with lower potential energy.

Bond formation is a fundamental concept in chemistry where two atoms share, donate, or accept electrons to create a stable arrangement. By forming bonds, atoms attain a lower energy state which is depicted as a negative change in enthalpy. There are different types of bonds that can be formed, such as covalent, ionic, and metallic bonds.
In the exercise concerning N₂ and O₂, we observe the formation of covalent bonds, where the atoms share valence electrons. The triple bond in nitrogen and the double bond in oxygen are examples of covalent bonding, with shared electron pairs holding the atoms together.

Energy and Stability

The sharing of electrons is not arbitrary but follows principles that minimize the potential energy of the molecule. For example, the triple bond in N₂ implies a stronger and more energy-storing bond compared to the double bond in O₂. As such, when these bonds form, the energy released in the formation of a triple bond is more significant, which in turn, reflects on the greater negative value in the enthalpy of atom combination for nitrogen.