The hydrogen molecule (\(H_2\)) exists as a gas under normal conditions, reflecting its extremely low boiling point of approximately \(-253^\circ\text{C}\), or \(20\text{ K}\). This temperature is one of the lowest found in nature, second only to that of helium. The physical state of any substance is determined by the forces acting upon its molecules, and for hydrogen, these forces are uniquely weak.
The Forces That Determine Boiling Point
Boiling is a phase change that requires energy to overcome the attractive forces that hold liquid molecules close together. These attractions are known as intermolecular forces (IMFs), which operate between separate molecules. IMFs are distinctly different from the much stronger intramolecular forces, such as the covalent bonds that hold the two hydrogen atoms together within the \(H_2\) molecule.
The covalent bond inside the hydrogen molecule is stable and requires a massive input of energy to break. However, the boiling process does not break this internal bond; it only supplies enough thermal energy to separate one intact \(H_2\) molecule from its neighbors. The boiling point is therefore defined by the amount of energy needed to disrupt the weak, external forces of attraction between molecules. In the case of \(H_2\), the weakness of these external forces is the direct cause of its low boiling temperature.
The Nature of London Dispersion Forces
The hydrogen molecule is nonpolar, meaning the two electrons are shared evenly, resulting in no permanent positive or negative ends. For such molecules, the only attraction available is the London Dispersion Force (LDF), a temporary and inherently weak type of intermolecular force. These forces are present in all molecules but become the dominant factor for nonpolar substances.
LDFs arise from the continuous, random movement of electrons around the molecule’s nucleus. At any given instant, electrons may be momentarily distributed unevenly, creating a fleeting, instantaneous dipole moment. This temporary charge imbalance in one molecule influences the electron cloud of a nearby molecule, inducing a corresponding temporary dipole in the neighbor. The resulting momentary electrostatic attraction between these short-lived dipoles constitutes the London Dispersion Force.
The Minimal Forces of the Hydrogen Molecule
While LDFs are the weakest class of intermolecular force, they are exceptionally minimal in molecular hydrogen. The strength of LDFs is directly related to a molecule’s polarizability, which is the ease with which its electron cloud can be distorted to form a temporary dipole. Molecules with more electrons and larger size are more polarizable because their outer electrons are held less tightly by the nucleus.
Molecular hydrogen is the smallest possible molecule, consisting of only two protons and two electrons. These two electrons are held tightly in a small, compact electron cloud. Because the cloud is small, it is difficult to distort the electron distribution and create a significant temporary dipole. The resulting attractions are extremely weak and fleeting, requiring very little kinetic energy to overcome them. This minimal energy requirement translates directly into the extremely low boiling point of \(20\text{ K}\).
How Hydrogen Compares to Other Substances
Comparing \(H_2\) to other common substances highlights that boiling point is a function of both the type and magnitude of intermolecular forces. Water (\(H_2O\)), for example, has a boiling point of \(100^\circ\text{C}\). Water is a polar molecule that uses much stronger attractions, specifically hydrogen bonds and dipole-dipole forces, which require significantly more energy to break than the LDFs in \(H_2\).
Even when comparing \(H_2\) to other nonpolar molecules that rely only on LDFs, the effect of size is clear. Methane (\(CH_4\)), the main component of natural gas, is nonpolar, but its boiling point is \(\approx -161^\circ\text{C}\) (\(112\text{ K}\)). Methane is significantly larger, with 10 electrons compared to \(H_2\)‘s two electrons. This greater number of electrons makes methane’s electron cloud more polarizable, resulting in substantially stronger LDFs and a much higher boiling point.