Phosphine (\(\text{PH}_3\)) consists of one phosphorus atom bonded to three hydrogen atoms. Determining if \(\text{PH}_3\) is polar or nonpolar relies on two concepts: the unequal sharing of electrons and the molecule’s three-dimensional shape. By examining phosphine’s electronic distribution and geometry, we can conclude that \(\text{PH}_3\) is a polar molecule.
Defining Chemical Polarity
Molecular polarity begins with electronegativity, which is an atom’s ability to attract shared electrons within a chemical bond. When two atoms with different electronegativities bond, electrons are pulled closer to the more electronegative atom. This creates a polar bond with a slight negative charge (\(\delta^-\)) on one end and a slight positive charge (\(\delta^+\)) on the other, known as a bond dipole.
The \(\text{P-H}\) bond polarity is determined by comparing the electronegativity values of phosphorus (\(\text{P}\)) and hydrogen (\(\text{H}\)). Hydrogen has an electronegativity of \(2.20\) and phosphorus has a value of \(2.19\). This extremely small difference of \(0.01\) makes the \(\text{P-H}\) bond essentially nonpolar, or very weakly polar covalent.
It is important to differentiate between a polar bond and a polar molecule. A molecule’s overall polarity depends on its entire three-dimensional structure, not just the sum of its bond polarities. Even if individual bonds possess a small dipole, the molecule can be nonpolar if those dipoles are arranged symmetrically and cancel each other out. Conversely, a molecule with weakly polar bonds can still be polar if its structure is asymmetrical, which is the case for \(\text{PH}_3\).
The Shape of Phosphine
The shape of the phosphine molecule ultimately dictates its polarity. Geometry is predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron groups around a central atom arrange themselves to minimize repulsion. The central phosphorus atom in \(\text{PH}_3\) is surrounded by four electron groups: three single bonds to hydrogen atoms and one non-bonding lone pair.
These four electron groups adopt a tetrahedral electron pair geometry to maximize separation. However, the molecular shape, determined only by the positions of the atoms, is not tetrahedral. The non-bonding lone pair significantly alters the final shape because it occupies more space than the bonding pairs, exerting a greater repulsive force.
This increased repulsion pushes the three hydrogen atoms closer together. As a result, the molecular geometry of phosphine is described as trigonal pyramidal. This asymmetrical structure is characterized by \(\text{H-P-H}\) bond angles of approximately \(93.5^\circ\), substantially smaller than the \(109.5^\circ\) found in a symmetrical tetrahedron.
Calculating the Molecular Dipole
The overall polarity of \(\text{PH}_3\) is determined by calculating its molecular dipole moment, which is the vector sum of all individual bond dipoles. Since the trigonal pyramidal shape is asymmetrical, the slight individual dipoles from the three \(\text{P-H}\) bonds cannot cancel each other out.
While the \(\text{P-H}\) bonds are only minimally polar, the asymmetrical geometry ensures that their small dipole moments combine to produce a net overall dipole moment. The significant contribution to the molecule’s polarity comes from the lone pair of electrons on the phosphorus atom.
The lone pair creates a region of high electron density concentrated at the apex of the pyramid, strongly contributing to the molecule’s overall charge separation. This concentration of electron density, combined with the structural asymmetry, results in a non-zero net dipole moment. Phosphine has a measured net dipole moment of \(0.58\) Debye.
Because the vector sum of the bond dipoles and the strong influence of the lone pair is not zero, the phosphine molecule possesses a permanent net dipole moment. This uneven distribution of charge confirms that \(\text{PH}_3\) is an overall polar molecule, despite the near-nonpolar nature of its individual \(\text{P-H}\) bonds.