Borane (\(\text{BH}_3\)) does not possess a net dipole moment, classifying it as a nonpolar molecule. Molecular polarity is determined by the polarity of individual chemical bonds and the precise three-dimensional shape. Although the bonds within borane are slightly polar, the molecule’s perfect symmetry causes all forces to cancel out completely.
Understanding Dipole Moments
A dipole moment (\(\mu\)) measures the overall polarity of a molecule, describing how electrical charge is distributed across its structure. This moment arises when electrons are shared unequally between two bonded atoms due to differences in electronegativity. When atoms with different electronegativities bond, the shared electron density shifts toward the more attractive atom, creating partial negative (\(\delta-\)) and positive (\(\delta+\)) charges. This charge separation generates an individual bond dipole. The molecular dipole moment is the vector sum of all these individual bond dipoles; a molecule is considered nonpolar if this sum is zero.
The Geometry of Borane (\(\text{BH}_3\))
The three-dimensional structure of \(\text{BH}_3\) is the primary reason for its lack of a net dipole moment. Boron, the central atom, forms single covalent bonds with the three surrounding hydrogen atoms using its three valence electrons. To determine the molecular shape, chemists use the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR dictates that the three B-H bonds around the central boron atom arrange themselves as far apart as possible to minimize electron-electron repulsion. The resulting geometry is perfectly flat, or trigonal planar, placing the three hydrogen atoms at \(120^\circ\) angles from each other.
Why the Dipoles Cancel Out
Although the \(\text{BH}_3\) molecule is nonpolar overall, the B-H bonds themselves are slightly polar. The small difference in electronegativity means the electrons in the B-H bond are slightly more attracted to the hydrogen atom, giving each hydrogen a partial negative charge and the central boron a partial positive charge. Each B-H bond possesses a small, real bond dipole moment pointing from the boron atom toward each hydrogen atom. However, because the molecule has a perfectly symmetrical trigonal planar geometry, these three equal bond dipoles are oriented precisely \(120^\circ\) apart from one another. This arrangement causes the vector components of the three dipoles to perfectly oppose and cancel each other out, resulting in a net molecular dipole moment of zero.