Molecular polarity, which describes the distribution of electrical charge within a molecule, is a fundamental characteristic that governs how a substance behaves and interacts with others. This concept determines whether a molecule has distinct positive and negative ends. Understanding molecular polarity is important for comprehending the physical and chemical properties of any substance, including dinitrogen (\(\text{N}_2\)).
Understanding Molecular Polarity
Molecular polarity is ultimately determined by two main factors: the polarity of the individual bonds and the overall geometric shape of the molecule. Bond polarity arises from a concept called electronegativity, which is an atom’s inherent ability to attract a shared pair of electrons toward itself in a chemical bond. When two atoms with different electronegativity values bond, the shared electrons spend more time near the more attractive atom, creating a polar covalent bond with a slight negative charge (\(\text{\delta}^-\)) on one end and a slight positive charge (\(\text{\delta}^+\)) on the other.
The second factor is the molecule’s geometry, which determines if these individual bond polarities cancel each other out. Each polar bond generates a bond dipole moment, which is a vector quantity. If the molecule’s shape is symmetrical, the individual dipole moments may be oriented in opposite directions, causing them to add up to zero. Even a molecule with multiple polar bonds can be nonpolar if its perfect symmetry leads to this cancellation, resulting in a zero net dipole moment. A molecule is considered polar only if it contains polar bonds and possesses an asymmetrical shape where the dipole moments do not cancel, leading to a net separation of charge.
Analyzing the Nitrogen Molecule N2
Applying the rules of polarity to the dinitrogen molecule (\(\text{N}_2\)) provides a definitive answer regarding its charge distribution. The \(\text{N}_2\) molecule consists of two nitrogen atoms connected by a triple covalent bond. This bond is strong because each nitrogen atom shares three electrons with the other to satisfy the octet rule.
To determine the polarity of the bond, the electronegativity values of the bonded atoms must be compared. Since the molecule is diatomic and composed of two identical nitrogen atoms, the electronegativity values are exactly the same. This means the shared electrons in the triple bond are pulled equally by both nuclei. The result is a nonpolar covalent bond because there is no difference in electron attraction between the two identical atoms, leading to an even distribution of charge.
Furthermore, the \(\text{N}_2\) molecule exhibits a linear geometric shape. Because the bond itself is nonpolar, there is no separation of charge to create a dipole moment in the first place. Consequently, the molecule has a net dipole moment of zero, confirming that dinitrogen is a nonpolar molecule. This complete lack of charge separation makes \(\text{N}_2\) chemically stable and unreactive under ordinary conditions.
How Polarity Affects Chemical and Physical Behavior
The nonpolar nature of the \(\text{N}_2\) molecule directly influences its physical and chemical behavior. Nonpolar molecules experience very weak intermolecular forces, specifically London Dispersion Forces, which arise from temporary, instantaneous shifts in electron density. These transient forces are the only means by which one dinitrogen molecule attracts another.
Because these attractive forces are so weak, very little energy is required to separate the molecules from each other. This accounts for nitrogen’s extremely low boiling point of approximately \(\text{-196}^{\circ}\text{C}\) and its state as a gas at room temperature. The nonpolar quality also governs its solubility based on the principle that “like dissolves like.” Dinitrogen is therefore not readily soluble in highly polar solvents such as water, which is a trait typical of nonpolar substances.