Hydrophobe

In chemistry, hydrophobicity is the chemical property of a molecule (called a hydrophobe) that is seemingly repelled from a mass of water. In contrast, hydrophiles are attracted to water.

Hydrophobic molecules tend to be nonpolar and, thus, prefer other neutral molecules and nonpolar solvents. Because water molecules are polar, hydrophobes do not dissolve well among them. Hydrophobic molecules in water often cluster together, forming micelles. Water on hydrophobic surfaces will exhibit a high contact angle.

Examples of hydrophobic molecules include the alkanes, oils, fats, and greasy substances in general. Hydrophobic materials are used for oil removal from water, the management of oil spills, and chemical separation processes to remove non-polar substances from polar compounds.

The term hydrophobic—which comes from the Ancient Greek ὑδρόφοβος (hydróphobos), "having a fear of water", constructed from Ancient Greek ὕδωρ (húdōr) 'water' and Ancient Greek φόβος (phóbos) 'fear'—is often used interchangeably with lipophilic, "fat-loving". However, the two terms are not synonymous. While hydrophobic substances are usually lipophilic, there are exceptions, such as the silicones and fluorocarbons.

For small solutes, the hydrophobic interaction is mostly an entropic effect originating from the disruption of the highly dynamic hydrogen bonds between molecules of liquid water by the nonpolar solute, causing the water to compensate by forming a clathrate-like cage structure around the non-polar molecules. This structure is more highly ordered than free water molecules due to the water molecules arranging themselves to interact as much as possible with themselves, and thus results in a lower entropic state at the interface. This causes non-polar molecules to clump together to reduce the surface area exposed to water and thereby increase the entropy of the system. Thus, the two immiscible phases (hydrophilic vs. hydrophobic) will change so that their corresponding interfacial area will be minimal. This effect can be visualized in the phenomenon called phase separation.^{[citation needed]}

For larger nonpolar solutes that cannot be adequately "clathrated" by the hydrogen-bond network of water, the disruption of these bonds becomes inevitable, leading to a high enthalpic cost. Under ambient conditions, this transition from an entropy-dominated regime to one governed by enthalpy occurs at around ~1 nm in size, reflecting a shift in hydration free energy behavior from scaling with the solute volume to depending on the exposed surface area.

In this context, a quantitative molecular definition of hydrophobicity has been proposed, based on the energetic cost for a system to induce hydrogen-bond defects in its hydration shell. According to this approach, a system is considered hydrophobic if it cannot compensate for the missing hydrogen bonds with an energy at least as favorable as the cost of generating such a defect in pure water, a value known as the Defect Interaction Threshold (DIT), estimated at approximately −6 kJ/mol (around 30% of the typical energy of a hydrogen bond). This criterion coincides with the classical 90° contact angle threshold, thus providing a molecular justification for the transition to hydrophobic behavior.

Additionally, the DIT helps determine the regimes of filling, partial filling, and drying in nanoconfined water, depending on how many of the water molecule's interaction sites (among its four tetrahedral sites) exceed this threshold. This analysis for quantifying hydrophobicity or wetting can be performed using a structural indicator, the V_4S index, which reveals the existence of two inherently preferential interaction states for water.