Surface Chemistry (Adsorption Basics)
Adsorption
Surface chemistry deals with phenomena that occur at the surfaces or interfaces of substances. This includes processes like adsorption, catalysis, and colloid formation. The interface or surface represents the boundary between two bulk phases, which can be solid-liquid, solid-gas, liquid-gas, or liquid-liquid. Interfaces between solid-solid or gas-gas phases are generally not possible (gas-gas interfaces do not exist as gases are completely miscible).
At the interface, the particles (atoms, molecules, or ions) are not in the same environment as in the bulk. For example, a molecule in the bulk of a liquid is surrounded by similar molecules, experiencing balanced intermolecular forces from all sides. A molecule on the surface, however, is only surrounded by molecules from the sides and below, experiencing unbalanced forces, typically a net inward pull (leading to surface tension).
This imbalance of forces or residual attractive forces at the surface is responsible for the phenomenon of adsorption.
Adsorption is defined as the accumulation of molecular species at the surface rather than in the bulk of a solid or liquid. The substance that gets adsorbed is called the adsorbate, and the surface on which adsorption takes place is called the adsorbent.
For example, when a gas comes into contact with a solid surface, the gas molecules may accumulate on the surface of the solid. The gas is the adsorbate, and the solid is the adsorbent.
Examples of adsorption:
- When ammonia gas is passed over charcoal, the ammonia molecules are adsorbed on the surface of the charcoal.
- Silica gel adsorbs water vapour from the air.
- Activated charcoal is used in gas masks to adsorb poisonous gases.
It is important to note that adsorption is a surface phenomenon. The adsorbate particles are concentrated only on the surface of the adsorbent.
Distinction Between Adsorption And Absorption
While often used interchangeably in casual conversation, adsorption and absorption are distinct phenomena.
Adsorption: As defined above, it is the surface phenomenon where the adsorbate accumulates only on the surface of the adsorbent.
Absorption: It is a bulk phenomenon where the absorbed substance is uniformly distributed throughout the body of the absorbent.
Consider the example of a sponge placed in water. The water is soaked up and distributed throughout the sponge's volume. This is absorption. However, if you place activated charcoal in a solution of a coloured dye, the dye molecules will stick to the surface of the charcoal particles, while the water permeates the bulk. This is adsorption of the dye by charcoal.
Sometimes, both adsorption and absorption occur simultaneously. This process is called sorption. For example, when water vapour is absorbed by anhydrous calcium chloride, it is absorption. When water vapour is adsorbed by silica gel, it is adsorption. If a gas is dissolved and distributed in a liquid or solid, it is absorption.
Comparison between Adsorption and Absorption:
| Feature | Adsorption | Absorption |
|---|---|---|
| Location | Surface phenomenon (accumulation only on the surface) | Bulk phenomenon (uniform distribution throughout the bulk) |
| Rate | Initially rapid, then decreases until equilibrium is reached | Occurs at a uniform rate throughout the process |
| Concentration | Concentration of adsorbate is higher on the surface than in the bulk | Concentration of absorbed substance is uniform throughout the bulk |
| Nature | Surface process | Bulk process |
| Examples | Water vapour on silica gel, Gas on charcoal | Water absorbed by a sponge, Water absorbed by anhydrous CaCl$_2$ |
Mechanism Of Adsorption
Adsorption occurs due to the presence of unbalanced or residual attractive forces at the surface of the adsorbent.
Particles in the bulk of a solid or liquid are surrounded by neighbouring particles, and the forces between them are balanced. However, particles on the surface are not surrounded by similar particles on all sides. They have unused or residual attractive forces (like van der Waals forces or sometimes stronger chemical forces).
These residual attractive forces on the surface of the adsorbent attract and hold the adsorbate particles onto the surface. The extent of adsorption depends on the strength of these residual forces and the nature of the adsorbent and adsorbate.
The surface energy of the adsorbent is lowered during adsorption. Adsorption is thus generally an exothermic process ($\Delta H_{adsorption} < 0$). The enthalpy change during adsorption is called the enthalpy of adsorption.
When a gas is adsorbed on a solid surface, its freedom of movement is restricted, leading to a decrease in entropy ($\Delta S < 0$). According to the second law of thermodynamics, for a process to be spontaneous, the Gibbs free energy change ($\Delta G$) must be negative:
$ \Delta G = \Delta H - T\Delta S $
Since $\Delta H$ is negative (exothermic) and $\Delta S$ is negative (decrease in disorder), $\Delta G$ will be negative if the term $T\Delta S$ is less negative than $\Delta H$. This is true for spontaneous adsorption. As adsorption proceeds, $\Delta H_{adsorption}$ becomes less negative (as fewer active sites are available, or the attractive forces become weaker), and $\Delta S$ becomes more negative (as the surface becomes more ordered). Eventually, a state of equilibrium is reached where $\Delta G = 0$, and the rate of adsorption equals the rate of desorption (removal of adsorbate from the surface).
Types Of Adsorption
Adsorption can be classified into two main types based on the nature of the forces of attraction between the adsorbate and the adsorbent:
- Physisorption (Physical Adsorption):
- Forces: Adsorbate is held to the adsorbent surface by weak van der Waals forces (Dispersion, Dipole-dipole, etc.).
- Nature: Non-specific. Occurs between any gas and any solid, although the extent varies.
- Enthalpy of adsorption: Low, typically 20-40 kJ/mol. This energy is comparable to the latent heat of vaporisation of the adsorbate.
- Reversibility: Generally reversible. The gas can be easily removed from the surface by increasing temperature or decreasing pressure.
- Effect of temperature: Decreases with increasing temperature. Since it's an exothermic process, according to Le Chatelier's principle, increasing temperature shifts the equilibrium towards desorption. Occurs significantly at low temperatures.
- Effect of pressure: Increases with increasing pressure. Increasing pressure of the gas above the surface increases the number of molecules striking the surface and thus the rate of adsorption.
- Nature of adsorbate: More easily liquefiable gases (those with higher critical temperatures, indicating stronger intermolecular forces) are adsorbed to a greater extent.
- Surface area: Increases with increasing surface area of the adsorbent.
- Formation of layers: Can form multi-molecular layers on the adsorbent surface.
- Activation energy: Low or negligible activation energy.
- Chemisorption (Chemical Adsorption):
- Forces: Adsorbate is held to the adsorbent surface by strong chemical bonds (covalent or ionic).
- Nature: Specific. Occurs only when there is a possibility of chemical bonding between the adsorbate and adsorbent.
- Enthalpy of adsorption: High, typically 80-240 kJ/mol. This energy is comparable to the enthalpy of chemical bond formation.
- Reversibility: Generally irreversible. Formation of chemical bonds is usually a permanent change. Desorption requires higher temperatures.
- Effect of temperature: Initially increases with increasing temperature (as it often requires activation energy for bond formation), then decreases at very high temperatures (due to increased desorption). Favoured at higher temperatures.
- Effect of pressure: Increases with increasing pressure, as it increases the concentration of adsorbate particles available to react with the surface.
- Nature of adsorbate: Specific to the chemical nature of the gas and the solid.
- Surface area: Increases with increasing surface area of the adsorbent.
- Formation of layers: Usually forms a mono-molecular layer (a single layer of adsorbate molecules) on the adsorbent surface, as once the surface sites form chemical bonds, they are occupied.
- Activation energy: Often involves significant activation energy for the formation of chemical bonds.
Comparison between Physisorption and Chemisorption:
| Property | Physisorption | Chemisorption |
|---|---|---|
| Nature of forces | Van der Waals forces | Chemical bonds (covalent or ionic) |
| Enthalpy of adsorption | Low (20-40 kJ/mol) | High (80-240 kJ/mol) |
| Specificity | Non-specific | Highly specific |
| Reversibility | Reversible | Irreversible |
| Effect of temperature | Decreases with increasing temperature | Increases initially, then decreases |
| Effect of pressure | Increases with increasing pressure | Increases with increasing pressure |
| Surface area | Increases with increasing surface area | Increases with increasing surface area |
| Layers | Multimolecular layers | Unimolecular layer |
| Activation energy | Low/Negligible | High |
It's possible for physisorption to occur first, and then transition to chemisorption at higher temperatures as sufficient activation energy is available for chemical bond formation.
Adsorption Isotherms
An adsorption isotherm is a graph that shows the relationship between the amount of adsorbate adsorbed per unit mass of adsorbent and the pressure of the gas (or concentration of the solute) at a constant temperature.
Commonly, it is plotted as $\frac{x}{m}$ vs $P$, where $x$ is the mass of the adsorbate, $m$ is the mass of the adsorbent, and $P$ is the equilibrium pressure of the adsorbate gas.
Different models have been proposed to describe adsorption isotherms. Two well-known ones are Freundlich and Langmuir adsorption isotherms.
Freundlich Adsorption Isotherm
This is an empirical relationship proposed by Freundlich in 1909.
Statement: The amount of gas adsorbed per unit mass of solid adsorbent is related to the pressure at constant temperature by the following equation:
$ \frac{x}{m} = k P^{1/n} $
Where:
- $x$ is the mass of the adsorbate.
- $m$ is the mass of the adsorbent.
- $P$ is the equilibrium pressure of the adsorbate gas.
- $k$ and $n$ are constants that depend on the nature of the adsorbent and adsorbate, and the temperature ($n > 1$).
This equation is valid over a limited range of pressures.
- At low pressures, $1/n \approx 1$, so $\frac{x}{m} \propto P$. Adsorption is nearly proportional to pressure.
- At high pressures, $1/n \approx 0$, so $\frac{x}{m} \propto P^0 =$ constant. Adsorption approaches a limiting value (saturation) and becomes independent of pressure.
- In the intermediate range, $1/n$ is between 0 and 1.
Taking the logarithm of the Freundlich equation gives a linear form:
$ \log \left(\frac{x}{m}\right) = \log k + \frac{1}{n} \log P $
A plot of $\log (\frac{x}{m})$ versus $\log P$ is a straight line with slope $1/n$ and intercept $\log k$.
Langmuir Adsorption Isotherm
Langmuir proposed a theoretical adsorption isotherm based on specific assumptions (e.g., monolayer adsorption on homogeneous sites, dynamic equilibrium between adsorption and desorption). The Langmuir equation is:
$ \frac{x}{m} = \frac{aP}{1+bP} $
Where $a$ and $b$ are constants.
This model explains the saturation observed at high pressures where the surface becomes completely covered by a single layer of adsorbate molecules.
Adsorption From Solution Phase
Solids can also adsorb solutes from solutions. For example, activated charcoal is used to decolourise sugar solutions or vegetable oils by adsorbing the coloured impurities.
The Freundlich isotherm can also be applied to adsorption from solution, by replacing the pressure term ($P$) with the equilibrium concentration ($C$) of the solute in the solution:
$ \frac{x}{m} = k C^{1/n} $
Where $x$ is the mass of solute adsorbed, $m$ is the mass of adsorbent, and $C$ is the equilibrium concentration of the solute in the solution. $k$ and $n$ are constants.
Taking the logarithm gives:
$ \log \left(\frac{x}{m}\right) = \log k + \frac{1}{n} \log C $
A plot of $\log (\frac{x}{m})$ versus $\log C$ gives a straight line, confirming the applicability of the Freundlich isotherm.
Factors affecting adsorption from solution phase:
- Nature of Adsorbent and Adsorbate: Specific interactions between the solute and adsorbent are important. Activated charcoal, fullers earth, and silica gel are common adsorbents.
- Concentration of Solute in Solution: Adsorption increases with increasing concentration of the solute up to a certain limit (saturation).
- Temperature: Adsorption from solution is generally an exothermic process, so it decreases with increasing temperature.
- Surface Area of Adsorbent: Adsorption increases with increasing surface area of the adsorbent.
Applications of adsorption from solution include decolourisation, purification, and separation processes (like chromatography).