A topic in the Open Knowledge Graph — a free, open map of 15,290 topics and the order to learn them in.
Adsorption is the accumulation of adsorbate molecules on a surface, driven by enthalpy reduction but opposed by entropy loss (adsorbate loses translational freedom). The Gibbs free energy for adsorption depends on temperature: at low T, ΔH dominates and adsorption is favorable; at high T, the -TΔS term dominates and desorption is favored. Surface coverage and desorption temperature depend on temperature and adsorbate pressure through thermodynamic relations.
From your study of Gibbs free energy, you know that a process occurs spontaneously when ΔG = ΔH − TΔS is negative. Adsorption — the binding of gas or liquid molecules onto a solid surface — is a beautiful case study in the competition between enthalpy and entropy. When a gas molecule lands on a surface and forms a bond (whether a weak van der Waals interaction in physisorption or a strong chemical bond in chemisorption), the system releases energy: ΔH is negative. But that same molecule, which was freely translating and rotating in three dimensions, is now pinned to a two-dimensional surface with restricted motion. It has lost degrees of freedom, and its entropy has decreased: ΔS is negative for the adsorbate.
At low temperatures, the TΔS penalty is small, and the favorable (negative) ΔH drives ΔG negative — adsorption proceeds spontaneously and surface coverage builds up. As temperature increases, the TΔS term grows. At some crossover temperature, the entropy penalty overwhelms the enthalpy gain, ΔG turns positive, and the surface begins to clear as molecules desorb. This is why catalytic surfaces clean themselves at high temperatures and why adsorption experiments typically show decreasing coverage with increasing temperature at a fixed pressure. The desorption temperature — where coverage drops sharply — is a direct readout of the adsorption enthalpy: stronger surface bonds require higher temperatures to overcome.
The quantitative tool for analyzing these relationships is the Clausius-Clapeyron equation for adsorption, which relates the change in equilibrium pressure with temperature to the enthalpy of adsorption. By measuring adsorption isotherms (coverage versus pressure) at several temperatures, you can extract ΔH_ads from the slope of ln(P) versus 1/T at constant coverage. More negative ΔH_ads values correspond to stronger surface binding and higher desorption temperatures. The entropy of adsorption can also be extracted and provides insight into the mobility of the adsorbate: a molecule that retains some translational freedom along the surface (mobile adsorption) has a smaller entropy loss than one locked into a fixed site (localized adsorption).
Understanding these thermodynamic balances is essential for designing catalysts and adsorbents. An ideal catalyst binds reactants strongly enough to hold them on the surface and lower activation barriers, but not so strongly that the products cannot desorb — a principle known as Sabatier's principle. Similarly, industrial adsorbents for gas separation (like zeolites or activated carbon) are engineered so that the target molecule adsorbs preferentially at operating temperature but can be regenerated by heating. In every case, it is the interplay between ΔH and TΔS that determines where the thermodynamic sweet spot lies.