Ethene, commonly known as ethylene, is the simplest alkene, featuring the chemical formula C2H4 and a double bond between its two carbon atoms. The physical arrangement of atoms dictates its chemical behavior. The molecular shape of C2H4 is trigonal planar, a geometry resulting from sp2 hybridization, which involves the mixing of atomic orbitals on the carbon atoms.
The Observed Molecular Geometry
The trigonal planar shape means that all six atoms of the ethene molecule (two carbon and four hydrogen) lie in a single plane, making it a flat molecule. This two-dimensional arrangement is a consequence of the double bond connecting the two carbon atoms. The spatial arrangement around each carbon atom resembles a triangle with the carbon at the center.
Each carbon atom is bonded to two hydrogen atoms and one other carbon atom, totaling three surrounding atoms. Repulsion between the electron pairs forming these bonds pushes the atoms as far apart as possible in the plane. This results in bond angles close to the ideal 120 degrees. The actual measured H-C-H bond angle is slightly smaller, around 117.4 degrees, due to the differing electron density of the C-H single bonds compared to the C=C double bond.
Understanding sp2 Hybridization
The planar shape and specific bond angles of ethene are explained by sp2 hybridization occurring at each carbon atom. In this process, one s orbital and two p orbitals combine to form three equivalent sp2 hybrid orbitals. These sp2 orbitals arrange themselves in the trigonal planar geometry, pointing toward the corners of an equilateral triangle at 120-degree angles.
The three sp2 orbitals on each carbon atom form strong, head-on overlap bonds known as sigma (\(\sigma\)) bonds. One sp2 orbital from each carbon overlaps with the other carbon to form the central C-C sigma bond. The remaining two sp2 orbitals on each carbon overlap with the s orbitals of the two hydrogen atoms, forming the four C-H sigma bonds. This head-on overlap creates the rigid structural backbone of the molecule.
One p atomic orbital on each carbon atom does not participate in hybridization. These unhybridized p orbitals are positioned perpendicular to the plane formed by the sp2 orbitals and the sigma bonds. They are parallel to each other and overlap sideways, above and below the molecular plane. This sideways overlap forms a weaker pi (\(\pi\)) bond, which constitutes the second part of the carbon-carbon double bond.
How the Structure Determines Chemical Behavior
The combination of the sigma and pi bonds creates a region of high electron density between the two carbon atoms, making the molecule highly reactive. This concentration of electrons is readily accessible for chemical reactions. Ethene’s primary chemical property is its tendency to undergo addition reactions, where the weaker pi bond breaks easily to form two new, stronger sigma bonds with incoming groups.
The pi bond imposes a structural constraint on the molecule. Unlike single bonds, where free rotation is possible, the pi bond requires the unhybridized p orbitals to remain parallel to maintain their sideways overlap. Rotation around the C=C axis would break the pi bond, requiring substantial energy input. This restricted rotation locks the molecule into its rigid, planar geometry.
This structural rigidity, defined by the sp2 hybridization and the double bond, applies to all alkenes. For larger molecules, this lack of rotation can lead to the existence of cis and trans isomers. The planar shape and exposed pi bond make ethene a widely used starting material in the chemical industry, particularly for polymerization reactions that create plastics like polyethylene.