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Adsorption – Definition, Mechanism and Types

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  • Last Updated : 09 Nov, 2021

Adsorption is the adhesion of atoms, ions, or molecules to a surface from a gas, liquid, or dissolved solids. This process forms an adsorbate film on the adsorbent’s surface.

This differs from absorption, which occurs when a fluid dissolves or permeates a liquid or solid. Adsorption is a surface phenomenon, whereas absorption occurs throughout the entire volume of the material, though adsorption frequently precedes absorption. The term sorption refers to both processes, whereas desorption refers to the opposite.

Mechanism of Adsorption

Adsorption occurs because the adsorbent’s surface particles are not in the same environment as the particles within the bulk. Inside the adsorbent, all of the forces acting between the particles are mutually balanced, but on the surface, the particles are not surrounded on all sides by atoms or molecules of their kind, and thus they have unbalanced or residual attractive forces. These adsorbent forces are responsible for attracting adsorbate particles to their surface.

At a given temperature and pressure, the extent of adsorption increases as the surface area per unit mass of the adsorbent increases. The heat of adsorption is another important aspect of adsorption. During adsorption, there is always a decrease in surface residual forces, i.e., a decrease in surface energy, which appears as heat.

Adsorption is always an exothermic reaction. To put it another way, the ∆H of adsorption is always negative. When a gas is adsorbed, the molecules’ freedom of movement is restricted. This corresponds to a decrease in the entropy of the gas following adsorption, i.e., ∆S is negative. Adsorption is thus accompanied by a decrease in the system’s enthalpy as well as entropy. 

The thermodynamic requirement for a process to be spontaneous is that ∆G must be negative at constant temperature and pressure, implying a decrease in Gibbs energy. Based on the equation, 

∆G = ∆H – T∆S

If H has a sufficiently high negative value and – ∆TS is positive, ∆G can be negative. Thus, in a spontaneous adsorption process, the combination of these two factors causes ∆G to be negative. As adsorption progresses, H becomes increasingly negative. At some point, ∆H equals ∆TS, and ∆G equals zero. At this point, equilibrium has been achieved.

Types of Adsorption

There are primarily two types of gas adsorption on solids. Adsorption is defined as physical adsorption or physisorption when gas accumulates on the surface of a solid due to weak van der Waals forces.

Chemical adsorption or chemisorption occurs when gas molecules or atoms are held to a solid surface by chemical bonds. 

Chemical bonds can be either covalent or ionic in nature. Chemisorption has high activation energy and is thus often referred to as activated adsorption. Sometimes these two processes occur concurrently, making it difficult to determine the type of adsorption. Physical adsorption at low temperatures can progress to chemisorption as the temperature rises. Dihydrogen, for example, is first adsorbed on nickel by van der Waals forces. Hydrogen molecules then dissociate to form hydrogen atoms, which are held on the surface via chemisorption.

Characteristics of Physisorption 

  • Lack of specificity: Since van der Waals forces are universal, a given adsorbent surface has no preference for a specific gas.
  • Nature of adsorbate: The amount of gas absorbed by a solid is determined by the nature of the gas. In general, gases that are easily liquefiable (i.e., have higher critical temperatures) are readily adsorbed because van der Waals forces are stronger near the critical temperatures. As a result, 1g of activated charcoal adsorbs more sulphur dioxide (critical temperature 630K) than methane (critical temperature 190K), which is still more than 4.5 mL of dihydrogen (critical temperature 33K).
  • Reversible nature: Adsorption of a gas by a solid is usually reversible. When pressure is increased, more gas is adsorbed as the volume of the gas decreases (Le–Chateliers’s’ principle), and the gas can be removed by decreasing pressure. Due to the exothermic nature of the adsorption process, physical adsorption occurs readily at low temperatures and decreases with increasing temperature (Le-Chateliers’s’ principle).
  • Surface area of adsorbent: The extent of adsorption increases as the adsorbent’s surface area increases. As a result, finely divided metals and porous substances with large surface areas are effective adsorbents.
  • Enthalpy of adsorption: Physical adsorption is, without a doubt, an exothermic process, but its enthalpy of adsorption is quite low (20–40 kJ mol-1). This is due to the fact that the attraction between gas molecules and solid surfaces is caused solely by weak van der Waals forces.

Characteristics of Chemisorption

  • High specificity: Chemisorption is highly specific, and it will only occur if chemical bonding between adsorbent and adsorbate is possible. Oxygen, for example, is adsorbed on metals as a result of oxide formation, and hydrogen is adsorbed on transition metals as a result of hydride formation.
  • Irreversibility: Chemisorption is usually irreversible in nature because it involves compound formation. Chemisorption is also an exothermic process, but it is very slow at low temperatures due to the high activation energy. Adsorption, like most chemical changes, often increases with increasing temperature. At high temperatures, the physisorption of a gas adsorbed at low temperatures can be converted to chemisorption. High pressure is usually advantageous for chemisorption.
  • Surface area: Chemisorption, like physical adsorption, increases with increasing adsorbent surface area.
  • Adsorption enthalpy: Chemisorption has a high enthalpy (80-240 kJ mol-1) because it involves the formation of chemical bonds.

Sample Questions

Question 1: Why does adsorption happen?

Answer:

Adsorption occurs because the adsorbent’s surface particles are not in the same environment as the particles within the bulk. Inside the adsorbent, all of the forces acting between the particles are mutually balanced, but on the surface, the particles are not surrounded on all sides by atoms or molecules of their kind, and thus they have unbalanced or residual attractive forces. These adsorbent forces are responsible for attracting adsorbate particles to its surface.

Question 2: Is adsorption an exothermic process? How?

Answer:

Adsorption is always an exothermic reaction. To put it another way, the ∆H of adsorption is always negative. When a gas is adsorbed, the molecules’ freedom of movement is restricted. This corresponds to a decrease in the entropy of the gas following adsorption, i.e., ∆S is negative. Adsorption is thus accompanied by a decrease in the system’s enthalpy as well as entropy.

Question 3: What is the relation between Gibbs energy and the process to be spontaneous?

Answer

The thermodynamic requirement for a process to be spontaneous is that ∆G must be negative at constant temperature and pressure, implying a decrease in Gibbs energy. Based on the equation,

∆G = ∆H – T∆S

If H has a sufficiently high negative value and – T∆S is positive, ∆G can be negative. Thus, in a spontaneous adsorption process, the combination of these two factors causes G to be negative.

Question 4: What is Le-Chateliers’s principle?

Answer:

It states that when a system that has been in equilibrium for a long time is subjected to a change in concentration, temperature, volume, or pressure, the system shifts to a new equilibrium, which partially offsets the applied change.

Question 5: What are Van der Waals forces?

Answer:

A distance-dependent interaction between atoms or molecules is the Van der Waals force. These attractions, unlike ionic or covalent bonds, are not the result of a chemical electronic bond; they are therefore comparatively weak and more susceptible to disturbance. At greater distances between interacting molecules, the van der Waals force quickly vanishes.

Question  6: What is critical temperature? What is the relation between critical temperature and the gas adsorbed?

Answer:

The critical temperature is the temperature above which a gas cannot be liquefied even with high pressure. The higher the critical temperature, the more gas adsorbed on the surface.

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