Learning Objectives Describe multiple covalent bonding in terms of atomic orbital overlap Relate the concept of resonance to π-bonding and electron delocalization The hybrid orbital model accounts for the geometry of molecules with single covalent bonds. It also describes how double and triple bonds form. Molecules with dobule and triple bonds use both σ and π bonds. This section describes how the hybrid orbital model works to visualize the creation of multiple bonds using σ and π bonds. Double Bonds The Lewis structure of ethene, C 2 H 4 , shows each carbon atom is connected to the other carbon atom and to two hydrogen atoms. As a result each carbon atom has trigonal planar geometry - according to VSEPR - and is sp 2 hybridized. Figure \(\PageIndex{1}\) shows the energy diagram for the atomic orbitals and the hybrid orbitals. The sp 2 hybrid orbitals overlap to make σ bonds to the hydrogen atoms and between the carbon atoms. The π bond in the C=C double bond is formed by overlap of the unused 2 p orbital, only two of the three p orbitals are used for hybridization so one is left over. This unhybridized p orbital is shown with the red and blue lobes in Figure \(\PageIndex{2}\). This unhybridized 2 p orbital is perpendicular to the plane of the sp 2 hybrid orbitals. This lets the unhybridized 2 p orbitals on the carbon atoms overlap side-by-side, above and below the internuclear axis, and form a π bond . For the unhybridized 2 p orbitals in ethene to overlap and form a π bond, the four hydrogen atoms and the two carbon atoms all must be in the same plane so that the ethene molecule is flat. If the molecule is twisted around the C=C bond the p orbitals are not overlapping and the π bond can not form. The planar configuration for the ethene molecule occurs because it is the most stable bonding arrangement. This demonstrates a significant difference between σ bonds, which rotate or spin easily, and π bonds, which do not allow rotation. Rotation around the internuclear axis does not change the extent to which the σ bonding orbitals overlap because the bonding electron density is symmetric about the axis. Rotation about the internuclear axis is much more difficult for multiple bonds because it would change the overlap of the π bonding orbitals and require breaking the π bond. Triple Bonds In molecules with sp hybrid orbitals there are two unhybridized p orbitals, shown in Figure \(\PageIndex{3}\) so they are capable of forming two π bonds. Figure \(\PageIndex{4}\) shows the bonding arrangement for acetylene, H−C≡C−H, a linear molecule. The sp hybrid orbitals of the two carbon atoms overlap end to end to form a σ bond between the two carbon atoms. The remaining sp orbitals form σ bonds with hydrogen atoms. This leaves two unhybridized p orbitals on each carbon atom that can overlap to form two π bonds. The two carbon atoms of acetylene are bound by one σ bond and two π bonds, giving a triple bond. Resonance Structures Hybridization involves only σ bonds, lone pairs of electrons, and single unpaired electrons (radicals). Structures that account for these features describe the correct hybridization of the atoms. However, many structures also include resonance forms. Remember that resonance forms occur when various arrangements of π bonds are possible without changing the three dimensional structure of the molecule. The orientation of the atoms stays the same, so the hybridization is the same for each resonance structure. The difference between resonances structures is the arrangement of π bonds between the unhybridized p orbitals. Resonance does not influence the assignment of hybridization. For example, benzene has two resonance forms (Figure \(\PageIndex{5}\)). Either of these forms show each carbon atom is bonded to three other atoms with no lone pairs, so the hybridization is sp 2 . The electrons in the unhybridized p orbitals form π bonds but there are two possible configurations. The difference between these two structures is only dependent on which p orbitals are overlapping. Neither resonance structure completely describes the electrons in the π bonds. They are not located in one position or the other, but are in both configurations at the same time. Example \(\PageIndex{1}\): Assignment of Hybridization Involving Resonance One component of acid rain is produced by the reaction of sulfur dioxide with atmospheric water vapor to form sulfuric acid. Sulfur dioxide, \(\ce{SO2}\), is a major component of volcanic gases and a product of the combustion of sulfur-containing coal. What is the hybridization of the \(S\) atom in \(\ce{SO2}\)? Solution The resonance structures of \(\ce{SO2}\) are The sulfur atom is surrounded by two bonds and one lone pair of electrons in either resonance structure. Therefore, the electron-pair geometry is trigonal planar, and the hybridization of the sulfur atom is sp 2 . Exercise \(\PageIndex{1}\) Another acid in acid rain is nitric acid, HNO 3 , which is produced by the reaction of nitrogen dioxide, NO 2 , with atmospheric water vapor. What is the hybridization of the nitrogen atom in NO 2 ? (Note: the lone electron on nitrogen occupies a hybridized orbital just as a lone pair would.) Answer sp 2 Summary Multiple bonds consist of a σ bond located along the axis between two atoms and one or two π bonds. The σ bonds are usually formed by the overlap of hybridized atomic orbitals, while the π bonds are formed by the side-by-side overlap of unhybridized orbitals. Resonance occurs when there are multiple unhybridized orbitals with the appropriate alignment to overlap, so the placement of π bonds can vary.