Intermolecular forces (IMFs) are the attractive forces that exist between separate molecules. These forces determine a substance’s physical properties, such as its boiling point and melting point. Understanding the IMFs in phosphine (\(\text{PH}_3\)) requires a detailed look at its molecular structure and electrical charge distribution. This analysis clarifies the specific forces, including dipole-dipole interactions, that govern how \(\text{PH}_3\) molecules interact.
Molecular Geometry and Structure of Phosphine (\(\text{PH}_3\))
The phosphine molecule consists of a central phosphorus atom bonded to three hydrogen atoms. Phosphorus is in Group 15, possessing five valence electrons. Three electrons form single covalent bonds with the hydrogen atoms, leaving a lone pair on the central phosphorus atom.
The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the molecule’s three-dimensional shape by minimizing repulsion between the electron domains. With three bonding pairs and one lone pair, there are four regions of electron density around the phosphorus atom. These four electron domains arrange themselves in a tetrahedral electron geometry.
The lone pair occupies more space than the bonding pairs, exerting a stronger repulsive force on the \(\text{P-H}\) bonds. This repulsion pushes the hydrogen atoms closer together, distorting the perfect tetrahedral arrangement. The resulting molecular geometry is called trigonal pyramidal, with the phosphorus atom at the apex.
Assessing Bond Polarity vs. Molecular Polarity
Molecular polarity is determined by the combination of bond polarity and the molecule’s overall geometry. Bond polarity results from the difference in electronegativity between the bonded atoms. For the \(\text{P-H}\) bond, the electronegativity of phosphorus is approximately 2.19, and that of hydrogen is approximately 2.20.
The difference in electronegativity is only about 0.01, making the \(\text{P-H}\) bond essentially nonpolar or very weakly polar. Even with this minimal bond polarity, the molecule’s overall shape still dictates its polarity.
Molecular polarity depends on the net dipole moment, which is the vector sum of all individual bond dipoles. Since \(\text{PH}_3\) has a trigonal pyramidal shape, it is asymmetrical. The slight bond dipoles, combined with the significant electron density from the lone pair on the phosphorus atom, do not cancel out. This asymmetry results in a small but measurable net dipole moment, classifying \(\text{PH}_3\) as a polar molecule with a dipole moment of approximately 0.58 Debye.
Identifying the Intermolecular Forces in \(\text{PH}_3\)
Because \(\text{PH}_3\) is a polar molecule, it exhibits dipole-dipole forces. Dipole-dipole forces are attractive interactions that occur between the permanent positive end of one polar molecule and the permanent negative end of an adjacent polar molecule. These attractions are stronger than London Dispersion Forces but weaker than covalent or ionic bonds.
In addition to dipole-dipole forces, \(\text{PH}_3\) also experiences London Dispersion Forces (LDFs). LDFs are present in all molecules, whether they are polar or nonpolar. These forces result from temporary, instantaneous dipoles created by the random movement of electrons, which then induce a corresponding dipole in a neighboring molecule. Since \(\text{PH}_3\) is a relatively small molecule, LDFs still contribute to the overall intermolecular attraction.
Comparing Phosphine to Ammonia (\(\text{NH}_3\))
Phosphine (\(\text{PH}_3\)) is often compared to ammonia (\(\text{NH}_3\)) because both molecules share the trigonal pyramidal geometry and are polar. The central nitrogen atom in \(\text{NH}_3\) also has one lone pair and is bonded to three hydrogen atoms. However, the difference in the strength of their intermolecular forces is significant.
Nitrogen is substantially more electronegative than phosphorus, creating much more polar \(\text{N-H}\) bonds than the \(\text{P-H}\) bonds. This higher polarity gives ammonia a much larger net dipole moment (approximately 1.4 D) compared to phosphine’s 0.58 D, making ammonia a much more polar molecule. Furthermore, the strong polarity in \(\text{NH}_3\) allows it to form a distinct and powerful IMF called hydrogen bonding.
Hydrogen bonding is a particularly strong type of dipole-dipole force that only occurs when hydrogen is directly bonded to one of the three highly electronegative atoms: nitrogen (\(\text{N}\)), oxygen (\(\text{O}\)), or fluorine (\(\text{F}\)). Since \(\text{PH}_3\) contains phosphorus, it cannot form hydrogen bonds. The absence of this powerful force explains why phosphine has a much lower boiling point (approximately \(-87.7^\circ \text{C}\)) than ammonia (approximately \(-33.3^\circ \text{C}\)), even though both are polar and experience dipole-dipole forces.