The molecule with the chemical formula \(\text{C}_2\text{H}_2\), commonly known as acetylene or ethyne, possesses a definitive and highly symmetrical shape. The molecular geometry of \(\text{C}_2\text{H}_2\) is linear. This straight-line arrangement means that all four atoms—the two carbon atoms and the two hydrogen atoms—are positioned along a single axis in three-dimensional space. This linear structure is a direct consequence of how the carbon and hydrogen atoms share their electrons, which dictates the maximum separation between the atoms.
Molecular Composition and Bonding
Acetylene is the simplest member of the alkyne family of hydrocarbons, composed only of carbon and hydrogen atoms, featuring a carbon-carbon triple bond (\(\text{H} – \text{C} \equiv \text{C} – \text{H}\)). The molecule uses ten valence electrons: four from each carbon and one from each hydrogen.
The triple bond between the two carbon atoms is composed of one sigma (\(\sigma\)) bond and two pi (\(\pi\)) bonds. The sigma bond forms from the direct, head-to-head overlap of atomic orbitals along the internuclear axis, providing the primary link between the two carbons.
The two pi bonds are formed by the sideways overlap of unhybridized atomic orbitals, creating two electron clouds positioned perpendicular to the central sigma bond. Each carbon atom also forms a sigma bond with a hydrogen atom. The complete molecule contains three sigma bonds (one \(\text{C}-\text{C}\) and two \(\text{C}-\text{H}\)) and two \(\text{C}-\text{C}\) pi bonds.
Determining the Geometry: Hybridization and Electron Domains
The linear shape is explained by two primary chemical theories: hybridization and the Valence Shell Electron Pair Repulsion (VSEPR) theory.
Hybridization
To form the triple bond, each carbon atom undergoes \(sp\) hybridization, mixing one \(s\) orbital and one \(p\) orbital from the carbon’s valence shell to create two identical \(sp\) hybrid orbitals. These two \(sp\) orbitals orient themselves \(180^\circ\) apart to minimize electron repulsion, resulting in a straight line.
One \(sp\) orbital on each carbon forms a \(\text{C}-\text{H}\) sigma bond with a hydrogen atom. The other \(sp\) orbital forms the \(\text{C}-\text{C}\) sigma bond. The two remaining \(p\) orbitals on each carbon do not hybridize. They remain perpendicular to the \(sp\) orbitals and form the two \(\pi\) bonds through sideways overlap. Since the \(sp\) hybrid orbitals are directed \(180^\circ\) from each other, the four atoms are forced into a linear arrangement.
VSEPR Theory
VSEPR theory predicts this geometry by counting electron domains around the central atoms. An electron domain is any region of electron density (single bonds, double bonds, triple bonds, or lone pairs).
In acetylene, each carbon atom has only two electron domains: one for the single bond to hydrogen and one for the triple bond to the other carbon. The VSEPR principle dictates that these two domains maximize their distance to minimize repulsion by orienting \(180^\circ\) apart, mandating the linear geometry.
Physical Implications of the Linear Shape
The linear geometry of acetylene has several measurable physical consequences, beginning with the precise \(180^\circ\) bond angle between the \(\text{H}-\text{C}-\text{C}\) atoms. This complete linearity is responsible for the molecule’s overall nonpolar nature.
Although there is a slight difference in electronegativity between carbon and hydrogen, creating a small, localized bond dipole in each \(\text{C}-\text{H}\) bond, these dipoles perfectly cancel each other out. The two \(\text{C}-\text{H}\) bond dipoles are equal in magnitude but point in exactly opposite directions along the straight molecular axis, resulting in a net dipole moment of zero. This perfect symmetry makes acetylene a nonpolar molecule, which affects its solubility and intermolecular forces.
The linear shape and the presence of the triple bond also contribute to acetylene’s high chemical reactivity. The two \(\pi\) bonds are located outside the main \(\sigma\)-bond axis, making them relatively exposed and energetically accessible. This makes the molecule highly susceptible to addition reactions, where other atoms can easily break the \(\pi\) bonds and add themselves across the carbon-carbon framework. This reactivity is the basis for its widespread use in applications like oxyacetylene welding.