Benzene is one of the compounds used as an octane enhancer in unleaded gasoline. It is manufactured by the catalytic conversion of acetylene to benzene: \(3 \mathrm{C}_{2} \mathrm{H}_{2} \longrightarrow \mathrm{C}_{6} \mathrm{H}_{6}\) Draw Lewis structures for these compounds, with resonance structures as appropriate, and determine the hybridization of the carbon atoms in each.

When drawing molecular structures, sometimes a single Lewis structure is not enough to represent the distribution of electrons accurately. This is where resonance structures come into play. They are alternative Lewis structures for a molecule where the arrangement of electrons can differ, but the arrangement of atoms remains the same.

Take benzene, for example. Its resonance structures show alternating single and double bonds between carbon atoms in a hexagonal ring. This concept is an average of these structures, depicting electrons being delocalized or spread out throughout the molecule, leading to increased stability. It's essential to recognize that resonance structures are not real, individual structures; they are a tool used to describe the behavior of electrons in certain molecules.

Understanding valence electrons is fundamental in chemistry, as they are the electrons that participate in bonding. The number of valence electrons an atom has dictates how it can bond with other atoms. In Lewis structures, these electrons are represented as dots surrounding the atomic symbols.

For carbon, with four valence electrons, these can form four covalent bonds as seen in hydrocarbons. In acetylene, each carbon uses three valence electrons to form a triple bond with the other carbon, and one to bond with hydrogen, thus adhering to the octet rule – which states that atoms are most stable when they have eight electrons in their valence shell.

In acetylene, the carbon atoms exhibit sp hybridization. Hybridization is the concept of atomic orbitals mixing to form new hybrid orbitals, which are used to make sigma and pi bonds. The sp hybridization involves one s orbital mixing with one p orbital, forming two sp hybrid orbitals. These two orbitals lay on a linear axis and form sigma bonds, leaving two unhybridized p orbitals per carbon atom free to form pi bonds. In acetylene, each carbon has one sigma bond with a hydrogen atom and another with the other carbon, and the remaining two p orbitals form two pi bonds, making up the triple bond.

In benzene, we encounter a different hybridization: sp² hybridization. Each carbon atom uses one s orbital and two p orbitals to create three sp² hybrid orbitals that are oriented in a trigonal planar fashion. This allows the formation of three sigma bonds—the two with adjacent carbon atoms and one with a hydrogen atom. The remaining unhybridized p orbital on each carbon overlaps with those on adjacent carbons to form pi bonds, which are the basis of the delocalized electrons across the benzene ring. This delocalization is indicative of the molecular stability and the unique properties of aromatic compounds like benzene.

There are two main types of covalent bonds we see in these structures: sigma (σ) and pi (π) bonds. Sigma bonds are the first bonds formed between atoms, characterized by head-on overlapping of orbitals. These bonds are single bonds and allow for the free rotation of bonded atoms.

Pi bonds are formed when parallel orbitals overlap and are found in double and triple bonds along with sigma bonds. They result from the sideways overlap of p orbitals. Pi bonds restrict the rotation of bonded atoms due to their electron cloud's alignment. Each double bond consists of one sigma and one pi bond, while a triple bond comprises one sigma and two pi bonds, as shown in acetylene's carbon-carbon bond.