Animated Solution for Chemistry - Surface Chemistry: Which of the following statements is incorrect regarding physisorptions?

Understanding Physisorption

Imagine you are standing in a dusty room. Over time, you notice a thin layer of dust settling on every surface—your desk, your books, the floor. The dust isn't glued or welded to the desk; it's just resting there, held by incredibly weak forces. This everyday phenomenon is a perfect analogy for physisorption (physical adsorption) in chemistry.

In the microscopic world, a solid surface acts like a subtle magnet for gas molecules. However, instead of forming strong, permanent chemical bonds (which would be chemisorption), the gas molecules are held to the solid surface by weak, transient forces known as van der Waals' forces. Because these forces are universal, physisorption is not specific; any gas will adsorb on any solid to some extent, provided the conditions are right.

Not all gases are created equal when it comes to sticking to a surface. Think about water vapor versus helium. Water vapor condenses into a liquid very easily because its molecules have strong intermolecular attractions. Helium, on the other hand, is notoriously difficult to liquefy.

In physisorption, gases that are easily liquefiable (those with higher critical temperatures) are adsorbed much more readily. Why? Because the very same intermolecular forces that make a gas easy to liquefy also make it easier for the gas molecules to grab onto the solid surface. Therefore, a gas like ammonia or sulfur dioxide will show a much higher extent of physisorption compared to hydrogen or nitrogen under the same conditions.

Multimolecular Layers

What happens if we start pumping more gas into our system, increasing the pressure? In chemisorption, once the surface is covered with a single layer of molecules (a monolayer), the process stops because all the chemical bonding sites are occupied.

Physisorption plays by different rules. Because it relies on van der Waals' forces, a gas molecule doesn't need to touch the bare solid surface directly. It can happily stick to another gas molecule that is already adsorbed. As we increase the pressure, these molecules start piling up on top of each other, forming multimolecular layers. This is a hallmark characteristic of physical adsorption.

Why Heat is Released

Now, let's tackle the core of the problem: the thermodynamics of adsorption. When gas molecules are floating freely in a container, they have a high degree of kinetic energy and randomness (high entropy). When they adsorb onto a surface, their movement is severely restricted. This means the change in entropy, $\Delta S$, is negative.

For any process to occur spontaneously, the Gibbs free energy change, $\Delta G$, must be negative. The famous equation relates these thermodynamic quantities:

Since $\Delta S$ is negative, the term $-T\Delta S$ becomes positive. For $\Delta G$ to be negative overall, the enthalpy change, $\Delta H$, must be highly negative.

Physically, what does this mean? When molecules bind to the surface, the residual attractive forces of the solid are satisfied, leading to a decrease in the surface energy of the system. This lost energy is released into the surroundings as heat. Therefore, adsorption is inherently an exothermic process.

The statement claiming that the enthalpy of adsorption ($\Delta H_{\text{adsorption}}$) is positive directly contradicts the fundamental thermodynamic reality of the process. The enthalpy of physisorption is indeed low (typically between $-20$ to $-40\text{ kJ/mol}$), but it is strictly negative, never positive.