Borane (\(\text{BH}_3\)) is a classic example illustrating molecular polarity. Polarity in a molecule measures how electrons are shared and distributed across its structure. A molecule is nonpolar when the electrical charge is distributed equally, resulting in no net charge separation. Borane provides a clear case where the molecule’s physical arrangement overcomes the slight charge separation present in its individual chemical bonds.
The Shape of Borane
The overall polarity of any molecule is fundamentally determined by its three-dimensional structure. Chemists use the Valence Shell Electron Pair Repulsion (VSEPR) theory to predict the geometry of \(\text{BH}_3\). VSEPR minimizes the repulsion between the electron domains around the central atom.
In Borane, the central Boron atom is bonded to three Hydrogen atoms. Boron contributes three valence electrons, and each Hydrogen contributes one, resulting in three bonding pairs around the central atom. The Boron atom has no non-bonding or lone pairs of electrons.
This arrangement is designated as \(\text{AX}_3\) in VSEPR theory. The three electron domains spread out as far as possible to minimize repulsion. The resulting molecular geometry is the highly symmetrical “trigonal planar” shape. All four atoms lie in the same plane, with the three Hydrogen atoms situated at the corners of an equilateral triangle. The angle between any two B-H bonds is \(120^\circ\).
Polarity of Individual Bonds
Bond polarity is determined by electronegativity, which is an atom’s ability to attract shared electrons toward itself. When two atoms with different electronegativity values bond, the shared electrons are pulled closer to the more electronegative atom, creating a slight charge separation called a bond dipole.
Boron’s electronegativity is approximately 2.04, while Hydrogen’s is slightly higher at about 2.20. This small difference indicates that the B-H bond is technically a polar covalent bond.
The electrons in each B-H bond are slightly drawn toward the Hydrogen atoms, creating a small bond dipole moment. This makes the Hydrogen end marginally negative and the Boron end marginally positive.
Molecular Symmetry and Dipole Cancellation
The key to understanding the nonpolar nature of Borane lies in how these three individual bond dipoles interact in three-dimensional space. The overall polarity of a molecule is determined by its net dipole moment, which is the vector sum of all the individual bond dipoles.
In the trigonal planar structure of \(\text{BH}_3\), the three B-H bond dipoles are equal in magnitude because they are geometrically identical. These three dipoles are oriented symmetrically, lying in the same plane and pointing outward from the central Boron atom at \(120^\circ\) angles from each other.
When three vectors of equal strength are arranged symmetrically in a plane with \(120^\circ\) separation, they perfectly oppose one another. The vector sum of these three bond dipoles is zero. Since the net dipole moment is zero, the \(\text{BH}_3\) molecule is nonpolar. This perfect cancellation due to the molecule’s high symmetry is the definitive reason why Borane is nonpolar.