Solubility defines the maximum amount of a substance (solute) that can dissolve into a liquid (solvent) at a specific temperature and pressure. Hydrogen gas (H₂) is remarkably limited in its ability to dissolve into water (H₂O). Understanding this dissolution process requires examining the fundamental molecular interactions and the external physical conditions that dictate how much of this gas can be held in a liquid medium.
The Molecular Reason for Low Solubility
The primary factor limiting hydrogen’s solubility in water is the fundamental mismatch between the chemical properties of the two molecules. Water is a highly polar solvent, meaning the oxygen atom strongly pulls electrons away from the two hydrogen atoms, creating a partial negative charge on the oxygen side and partial positive charges on the hydrogen side. This polarity allows water molecules to form strong, interconnected networks through hydrogen bonds, which are powerful intermolecular forces.
Hydrogen gas, however, is a nonpolar molecule; the two identical hydrogen atoms share electrons equally, resulting in no net charge separation. The principle of “like dissolves like” dictates that polar solvents readily dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes. For hydrogen to dissolve, it must disrupt water’s highly structured hydrogen-bonded network, a process that requires a significant input of energy.
The only attractive forces that form between a nonpolar H₂ molecule and a polar H₂O molecule are weak, temporary electrical forces called London dispersion forces. These forces are minimal and cannot compensate for the energy lost by breaking the strong hydrogen bonds between water molecules to create space for the H₂. Consequently, water molecules prefer to remain bonded to each other rather than accommodate the nonpolar hydrogen molecules, resulting in the low solubility of the gas.
Quantifying Solubility: The Role of Pressure
For gases, unlike most solids, solubility is highly dependent on the pressure of the gas directly above the liquid surface. The physical law that governs this relationship is known as Henry’s Law, which states that the amount of gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid. In simpler terms, to force more hydrogen molecules into the water, scientists must increase the pressure exerted by the hydrogen gas.
Increasing the partial pressure above the liquid effectively increases the frequency of collisions between the gaseous hydrogen molecules and the water surface, driving more of them into the solution. At standard atmospheric pressure (1 atm), water can only hold a maximum concentration, or saturation point, of about 1.6 milligrams of hydrogen per liter. To achieve higher concentrations, such as those used in commercial applications, the hydrogen gas must be pressurized to multiple atmospheres above the water.
This technique is similar to how carbonated beverages are made, where high pressure forces carbon dioxide into the liquid, and the gas rapidly escapes when the pressure is released.
How Temperature Influences Dissolution
Temperature introduces an inverse relationship over hydrogen dissolution: as the temperature of the water increases, the solubility of hydrogen gas decreases. Dissolving a gas in a liquid is an exothermic process, meaning it releases a small amount of heat into the surroundings. According to thermodynamic principles, increasing the temperature of the system shifts the equilibrium away from the heat-releasing process.
The added thermal energy increases the kinetic energy of both the water molecules and the dissolved hydrogen molecules. This greater agitation makes it easier for the weakly bound hydrogen molecules to overcome the attractive forces of the water molecules and escape back into the atmosphere above the liquid.
Real-World Relevance of Dissolved Hydrogen
The factors determining hydrogen solubility have important implications across various practical fields, most notably in the emerging market of “hydrogen water.” Because of the low saturation concentration at normal atmospheric pressure, manufacturers must use high-pressure dissolution or other methods to achieve the higher concentrations marketed for therapeutic effects (sometimes exceeding 3.5 mg/L). The rapid dissipation of the gas, governed by temperature and agitation, presents a challenge for consumers, as the dissolved hydrogen can be lost quickly after the container is opened.
In industrial settings, managing hydrogen solubility is crucial for processes like hydrogenation, where hydrogen gas is added to organic compounds to create products like margarine or pharmaceuticals.
In the development of hydrogen fuel cells and energy storage, engineers must precisely manage the dissolution of hydrogen and its interaction with water, which is a byproduct of the energy conversion process. The precise control of temperature and pressure is necessary to ensure efficient gas-liquid transfer and prevent the formation of gas bubbles that could interfere with the performance and safety of these high-tech systems.