Phosphine (\(\text{PH}_3\)) is a molecule composed of one phosphorus atom and three hydrogen atoms. The nature of the chemical bond determines a substance’s physical and chemical properties, making the classification of the phosphorus-hydrogen bond as ionic or covalent essential. This distinction is based on whether electrons are fully transferred between atoms or merely shared between them.
Defining the Extremes: Ionic and Covalent Bonds
Chemical bonds exist on a spectrum, but they are broadly categorized into two major types: ionic and covalent. An ionic bond forms when there is a complete transfer of one or more valence electrons from one atom to another. This transfer typically occurs between a metal and a non-metal, resulting in oppositely charged ions (cations and anions) held together by a strong electrostatic attraction.
Covalent bonds involve electrons shared between two atoms rather than transferred. This type of bonding is most common between two non-metal atoms, like the phosphorus and hydrogen found in phosphine. The sharing of electrons allows each atom to achieve a more stable electron configuration, generally satisfying the octet rule.
How Electronegativity Determines Bond Type
Chemists use the concept of electronegativity to quantify where a specific bond falls on this spectrum. Electronegativity is a measure of an atom’s ability to attract a shared pair of electrons in a chemical bond. Linus Pauling developed the most common scale, where values range from about 0.7 to 3.98.
The difference in electronegativity (\(\Delta \text{EN}\)) between the two bonded atoms is the metric used to predict the bond type. A large difference indicates that the more electronegative atom pulls the electron pair strongly, classifying the bond as ionic. Generally, a \(\Delta \text{EN}\) greater than 1.7 suggests an ionic bond.
Bonds with smaller differences are considered covalent, with the degree of electron sharing dictating the bond’s polarity. When the \(\Delta \text{EN}\) is between 0.4 and 1.7, the sharing is unequal, creating a polar covalent bond. If the difference is very small, typically less than 0.4, the electrons are shared nearly equally, resulting in a nonpolar covalent bond.
Calculating the Bond Type for Phosphine (\(\text{PH}_3\))
Applying the electronegativity concept to phosphine requires determining the difference between the phosphorus (P) and hydrogen (H) atoms. On the Pauling scale, the electronegativity of phosphorus is 2.19, while the value for hydrogen is 2.20. Both elements are non-metals, which is an initial strong indicator of covalent bonding.
The calculation for the electronegativity difference is \(\Delta \text{EN} = |2.20 – 2.19|\), which results in a value of 0.01. This minimal difference falls far below the 0.4 threshold for nonpolar covalent bonds and the 1.7 threshold for ionic bonds. The bond in phosphine is therefore classified as covalent, specifically a nonpolar covalent bond due to the almost equal sharing of the electron pair.
The Resulting Molecular Shape and Polarity
The covalent nature of the P-H bonds dictates the molecular structure of phosphine, which is predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory. The central phosphorus atom possesses three single bonds to hydrogen and one lone pair of non-bonding electrons. This arrangement results in a tetrahedral electron geometry.
The actual shape of the molecule, known as the molecular geometry, is determined by the positions of the atoms. The repulsion from the lone pair compresses the bond angles between the P-H bonds, distorting the shape into a trigonal pyramidal structure.
While the individual P-H bonds are nonpolar covalent, the presence of the lone pair on the phosphorus atom creates an uneven distribution of electron density across the entire molecule. This asymmetry gives phosphine a slight dipole moment, meaning the molecule as a whole is slightly polar.