Molecular hydrogen (\(\text{H}_2\)) is a simple two-atom molecule often used as a reference point for understanding how atoms interact. Can molecular hydrogen participate in the strong, specific interaction known as hydrogen bonding? Understanding the fundamental requirements for this type of bond is necessary to determine the answer for the \(\text{H}_2\) molecule.
Defining the Strict Requirements for Hydrogen Bonding
Hydrogen bonding is a particularly strong type of intermolecular force, which is an attraction that occurs between molecules. For this special attraction to form, a molecule must meet two specific criteria. The first requirement is the presence of a hydrogen atom covalently bonded to one of three highly electronegative elements: Fluorine (\(\text{F}\)), Oxygen (\(\text{O}\)), or Nitrogen (\(\text{N}\)).
The high electronegativity of these three atoms means they pull the shared electrons away from the hydrogen atom. This uneven sharing creates a highly polarized bond, resulting in the hydrogen atom acquiring a significant partial positive charge (\(\delta^+\)). This partially positive hydrogen atom is then strongly attracted to a lone pair of electrons on a neighboring \(\text{F}\), \(\text{O}\), or \(\text{N}\) atom in an adjacent molecule, forming the hydrogen bond. Water (\(\text{H}_2\text{O}\)) provides the classic example of this powerful attraction.
Why Molecular Hydrogen Cannot Form Hydrogen Bonds
Molecular hydrogen (\(\text{H}_2\)) fails to meet the fundamental structural requirement for forming a hydrogen bond because it lacks the necessary partial charge separation. The \(\text{H}_2\) molecule consists of two identical hydrogen atoms joined by a covalent bond. Since the two atoms possess the exact same electronegativity, they share the bonding electrons perfectly equally. This equal sharing means the electron density is uniformly distributed across the molecule, making \(\text{H}_2\) an entirely nonpolar molecule. As a result, neither hydrogen atom carries the significant partial positive charge (\(\delta^+\)) required to act as a hydrogen bond donor.
The Actual Intermolecular Forces Governing Molecular Hydrogen
Since the strong, specific attraction of hydrogen bonding is absent, the interactions between \(\text{H}_2\) molecules are governed by the weakest of all intermolecular forces: London Dispersion Forces (LDFs). These forces exist in all molecules, regardless of their polarity, but they become the dominant and only type of attraction in nonpolar molecules like molecular hydrogen.
LDFs arise from the continuous, random movement of electrons within the molecule. At any given instant, the electrons may momentarily shift to one side, creating a brief, temporary, and instantaneous dipole. This fleeting charge separation in one \(\text{H}_2\) molecule can then momentarily induce an opposing dipole in a neighboring molecule, leading to a very weak, short-lived attraction. Because molecular hydrogen is a very small molecule with only two electrons, its LDFs are exceptionally weak. This minimal attraction is why hydrogen must be cooled to extremely low temperatures, specifically below \(\text{-253}^{\circ}\text{C}\) (\(\text{20 K}\)), before it can transition into its liquid state.