16.5 Molecular Structure of Ethene

Ethene (${\mathrm{C}}_2{\mathrm{H}}_4$) is an alkene consisting of two carbon atoms joined by a double covalent bond, with two hydrogen atoms singly bonded to each carbon.

• Hybridization: Each carbon atom in ethene is $s{p}^2$ hybridized.
• Geometry: The $s{p}^2$ hybridization results in a trigonal planar arrangement around each carbon atom.
• Planarity: The entire ethene molecule is planar, meaning all six atoms (2 carbons and 4 hydrogens) lie in the same plane.
• Bond Angles: The $H-C-H$ bond angle is approximately $117.6^{\circ }$ and the $H-C=C$ bond angle is approximately $121.3^{\circ }$ (both close to the ideal $120^{\circ }$).
• Bond Length: The $C=C$ bond length is $1.34\text{A˚}$, shorter than the $C-C$ single bond ($1.54\text{A˚}$) in ethane.

Bonding in Ethene

The double bond between the two carbon atoms is a combination of one sigma ($\sigma$) bond and one pi ($\pi$) bond.

Sigma ($\sigma$) Bonds

Sigma bonds are formed by the direct, head-on overlap of atomic or hybrid orbitals. They are strong and hold electrons tightly.

• Carbon-Carbon ($\sigma$) Bond: Formed by the head-on overlap of two $s{p}^2$ hybridized orbitals, one from each carbon atom ($s{p}^2-s{p}^2$ overlap).
• Carbon-Hydrogen ($\sigma$) Bonds: Each of the four C-H bonds is formed by the overlap of an $s{p}^2$ orbital from a carbon atom and the $s$-orbital of a hydrogen atom ($s{p}^2-s$ overlap).

Pi ($\pi$) Bond

Pi bonds are formed by the parallel (sideways) overlap of unhybridized p-orbitals. For more on orbital shapes, see Shapes of Orbitals→.

• Carbon-Carbon ($\pi$) Bond: After $s{p}^2$ hybridization, each carbon atom has one unhybridized $p_z$-orbital perpendicular to the molecular plane. These two parallel $p_z$-orbitals overlap sideways to form the $\pi$ bond.
• Characteristics: The electrons in a $\pi$ bond are held less tightly and are located above and below the plane of the sigma bonds. This makes the $\pi$ bond weaker and more exposed than a sigma bond.

Why is the C=C Bond Shorter than C-C?

Two factors contribute to the shorter $C=C$ bond ($1.34\text{A˚}$) compared to $C-C$ ($1.54\text{A˚}$):

1. Greater $s$-character: $s{p}^2$ orbitals have 33% $s$-character vs. 25% in $s{p}^3$ orbitals, making them more compact and pulling the nuclei closer.
2. Additional $\pi$ bond: The extra electron density between the nuclei from the $\pi$ bond pulls them closer together.

Reactivity of Ethene

The presence of the weak and accessible $\pi$ bond makes ethene significantly more reactive than alkanes like ethane.

Ethene vs. Ethane

Ethene is more reactive than ethane because the $\pi$ bond can be easily broken to form new single bonds (addition reactions). Ethane only contains strong $\sigma$ bonds, which are difficult to break.

Ethene vs. Ethyne

Ethene is more reactive than ethyne. The $\pi$ electrons in ethene are more exposed and accessible for reaction. The shorter $C\equiv C$ triple bond in ethyne holds its $\pi$ electrons more tightly between the carbon nuclei (due to greater nuclear attraction), making them less available for electrophilic attack compared to the more diffuse $\pi$ cloud of the $C=C$ double bond in ethene.

Electrophilic Addition

Alkenes typically undergo electrophilic addition reactions:

• The exposed $\pi$ electrons act as a nucleophile (electron donor)
• An electrophile (electron acceptor) is attracted to the $\pi$ cloud
• The $\pi$ bond breaks and two new $\sigma$ bonds form

Example: Addition of $\mathrm{HBr}$ to ethene: $\mathrm{C}{\mathrm{H}}_2=\mathrm{C}{\mathrm{H}}_2+\mathrm{HBr}\to \mathrm{C}{\mathrm{H}}_3-\mathrm{C}{\mathrm{H}}_2\mathrm{Br}$