Dinitrogen, commonly known as nitrogen gas (\(\text{N}_2\)), is the dominant component of Earth’s atmosphere, making up approximately 78% of the air we breathe. Understanding the forces that hold this molecule together is fundamental to chemistry, as the nature of a chemical bond dictates a substance’s physical and chemical behavior. The classification of chemical bonds depends on how the electrons are distributed between the participating atoms.
The Covalent Nature of N2
The bond in the dinitrogen molecule is classified as covalent, characterized by the sharing of valence electrons between two atoms. Since nitrogen is a nonmetal, bonds formed between two nitrogen atoms are typically covalent. In contrast, an ionic bond involves the complete transfer of electrons from one atom to another, resulting in charged ions.
Since the two atoms forming the molecule are identical, they share the electrons equally. This equal sharing results in a nonpolar covalent bond. Molecules composed of two atoms of the same element, such as \(\text{H}_2\) or \(\text{O}_2\), are always nonpolar covalent because there is no difference in the attractive pull on the shared electrons. For the \(\text{N}_2\) molecule, the electron cloud is symmetrically distributed around both nitrogen nuclei, preventing the formation of positive and negative poles.
Determining Bond Type Using Electronegativity
The most reliable scientific method for classifying a bond as ionic or covalent involves calculating the difference in electronegativity (\(\Delta\text{EN}\)) between the two bonded atoms. Electronegativity is an intrinsic property that measures an atom’s tendency to attract a shared pair of electrons toward itself. Nitrogen has a high electronegativity value of 3.04 on the Pauling scale.
To determine the bond type in \(\text{N}_2\), the calculation is \(3.04 – 3.04\), which results in a \(\Delta\text{EN}\) of zero. This zero difference places the bond squarely within the range for nonpolar covalent bonds, typically defined as a \(\Delta\text{EN}\) of less than 0.4.
A polar covalent bond occurs when the \(\Delta\text{EN}\) is between 0.4 and 1.7. Bonds with a \(\Delta\text{EN}\) greater than 1.7 are classified as ionic, reflecting a near-complete transfer of the electron. The zero difference for \(\text{N}_2\) confirms the perfect symmetry of the electron sharing, confirming the bond is purely covalent.
The Stability of the Nitrogen Triple Bond
The specific structure of the \(\text{N}_2\) bond, beyond its covalent nature, is what gives the molecule its remarkable properties. To satisfy the octet rule, which states that atoms tend to bond in such a way that they each have eight electrons in their valence shell, the two nitrogen atoms share three pairs of electrons. This sharing of six total electrons forms a triple bond, often represented as \(\text{N}\equiv\text{N}\).
This triple bond is exceptionally strong, possessing one of the highest bond energies found in any diatomic molecule. Breaking this bond requires a substantial energy input of approximately 945 kilojoules per mole (\(\text{kJ/mol}\)). This high bond energy is the direct reason why \(\text{N}_2\) is largely unreactive, or inert, under normal atmospheric conditions.
The strength of the triple bond must be overcome for nitrogen to participate in chemical reactions. This inertness is beneficial because it prevents the atmosphere from undergoing rapid chemical change. However, this stability poses a challenge for biological systems and industrial processes, which must employ specialized enzymes or high-energy methods, such as the Haber process, to “fix” nitrogen into usable compounds like ammonia.