Molecular Geometry

In accord with VSEPR theory, molecules are arranged so as to minimize repulsion between electrons. Because they are all negatively charged, electrons repel one another. As a result of this electrical repulsion, the atoms of a covalently bonded molecule assume a shape around the central atom that maximizes the distance between the outermost or valence electrons. This means that repulsion between electrons in a molecule is at a minimum when the angles between the bonds (bond angles) allow for the greatest separation of the valence electrons. Bond angles are calculated using the central atom as the vertex of the bond angle.

In attempting to predict molecular shapes, it is often useful to consider the oversimplified view of molecules with independent electron orbitals (e.g., s and p orbitals). In the case of methane (CH₄), the greatest distance in space that can separate the four carbon-hydrogen bonds around the central carbon atom occurs when the bonds are pointed at the corners of a tetrahedron. When the bonds are pointed toward the corners of a tetrahedron the bond angles are 109.5° and the molecule is said to be a tetrahedral molecule.

If three atoms are bonded to a central atom and all of the bonds lie in the same plane, then the three bond angles must be 120° apart. Such a planar molecule is termed a trigonal planar molecule.

If only two atoms are bonded together, as is the case with molecular oxygen (O₂), the resulting bond angles must be 180° and the molecule is described as a linear molecule.