Surface ChemistryNEET Chemistry · Class 12 · NCERT Chapter 14

Adsorption: Basics and Types

What is Adsorption?

Adsorption is the process by which molecules of a gas or liquid (or dissolved substance) accumulate on the surface of a solid or liquid. The substance that adsorbs is called the adsorbate. The substance on whose surface adsorption happens is called the adsorbent.

NEET key: adsorption is a surface phenomenon. It happens only on the surface, not inside the bulk of the material.

Adsorption vs Absorption

These two words look similar but mean very different things. You must know the difference for NEET.

The term sorption is used when both adsorption and absorption occur simultaneously.

Why Does Adsorption Happen?

Molecules on the surface of a solid have unbalanced forces (their bonds are not fully satisfied on the free surface side). This gives the surface extra energy called surface energy. When an adsorbate molecule sticks to the surface, it reduces this surface energy and releases energy as heat. This is why adsorption is always exothermic.

Since adsorption is exothermic, the enthalpy change (delta H) is always negative for adsorption. The energy released per mole of adsorbate adsorbed is called the enthalpy of adsorption.

Physisorption

When the adsorbate is held on the surface by weak van der Waals forces, the process is called physical adsorption or physisorption.

• Force: weak van der Waals forces
• Enthalpy: low, 20-40 kJ/mol
• Reversibility: reversible (remove by reducing pressure or raising temperature slightly)
• Layers: multilayer adsorption possible
• Temperature effect: decreases with increasing temperature (exothermic process)
• Activation energy: not needed
• Specificity: not specific (any gas adsorbs on any surface)

Chemisorption

When the adsorbate forms a strong chemical bond (covalent or ionic) with the adsorbent surface, the process is called chemical adsorption or chemisorption.

• Force: strong covalent or ionic bonds
• Enthalpy: high, 40-400 kJ/mol
• Reversibility: irreversible (or very difficult to reverse)
• Layers: only monolayer adsorption
• Temperature effect: first increases (needs activation energy), then decreases
• Activation energy: needed (like a chemical reaction)
• Specificity: highly specific

Side-by-Side Comparison

Factors Affecting Adsorption

Several factors determine how much adsorption occurs:

1. Surface area of adsorbent: greater surface area = more adsorption. Finely divided solids and porous materials (activated charcoal, silica gel) adsorb more because they have more surface area per gram.
2. Nature of adsorbate: gases that are easily liquefied (high critical temperature, strong intermolecular forces) adsorb more readily. For example, SO2 adsorbs more than H2 on charcoal.
3. Nature of adsorbent: activated charcoal, silica gel, alumina, and zeolites are common adsorbents with large surface areas.
4. Temperature: for physisorption, adsorption decreases with temperature (exothermic process); for chemisorption, it increases up to a point then decreases.
5. Pressure: for gases, adsorption increases with pressure (more molecules available). At very high pressures, the surface becomes saturated.

Applications of Adsorption

• Activated charcoal in gas masks (adsorbs toxic gases)
• Silica gel as desiccant (adsorbs water vapour)
• Decolourisation of sugar solution by charcoal (adsorbs coloured impurities)
• Chromatography relies on differential adsorption
• Heterogeneous catalysis (reactants adsorb on catalyst surface)

Adsorption Isotherms

An adsorption isotherm is a graph that shows how the amount of gas adsorbed per gram of adsorbent (x/m) varies with the pressure of the gas at a constant temperature. "Isotherm" means "at constant temperature."

Freundlich Adsorption Isotherm

H.W. Freundlich proposed an empirical equation in 1909:

where:

• x = mass of adsorbate adsorbed
• m = mass of adsorbent
• k = Freundlich constant (depends on nature of adsorbate and adsorbent)
• P = pressure of adsorbate gas
• n = Freundlich constant (always greater than 1)

Taking logarithm of both sides gives the linearised form:

This is of the form y = c + mx (a straight line). When you plot log(x/m)on the y-axis and log P on the x-axis, you get a straight line where:

• Slope = 1/n
• y-intercept = log k

Freundlich Isotherm: Limiting Cases

• When 1/n = 1: x/m = kP (adsorption is directly proportional to pressure, linear)
• When 1/n = 0: x/m = k (adsorption is independent of pressure, flat line)
• In practice 1/n is between 0 and 1, so x/m increases with P but not linearly

Limitations of Freundlich Isotherm

• It is empirical (no theoretical derivation)
• It fails at very high pressures (predicts continuous increase, but surface saturates)
• It is not valid near saturation vapour pressure

Langmuir Adsorption Isotherm

Irving Langmuir (1916) proposed a theoretically derived isotherm based on four assumptions:

1. Adsorption forms only a monolayer (single molecule thick layer)
2. All adsorption sites on the surface are equivalent (same adsorption energy)
3. Once a site is occupied, no further adsorption occurs at that site
4. There is no interaction between adsorbed molecules

The Langmuir equation is:

where a and b are Langmuir constants. At low pressure (bP << 1): x/m ≈ abP (linear). At high pressure (bP >> 1): x/m ≈ a (constant, monolayer complete). This built-in saturation is the key advantage over Freundlich.

Catalysis

A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed. The phenomenon is called catalysis. The catalyst changes the mechanism of the reaction and provides a lower-energy pathway, reducing the activation energy.

Key point: a catalyst does NOT change the equilibrium position or the overall free energy (delta G) of the reaction. It only makes the reaction reach equilibrium faster.

Homogeneous Catalysis

In homogeneous catalysis, the catalyst and the reactants are in the same phase.

Heterogeneous Catalysis

In heterogeneous catalysis, the catalyst and reactants are in different phases. The catalyst is almost always a solid, and the reactants are gases or liquids.

Mechanism of Heterogeneous Catalysis (Adsorption Theory)

The mechanism involves five steps:

1. Diffusion: reactant molecules (gas or liquid) move toward the catalyst surface
2. Adsorption: reactant molecules adsorb on the catalyst surface (physisorption or chemisorption)
3. Bond weakening: adsorption weakens the bonds in the reactant molecules, making them more reactive
4. Reaction: the adsorbed molecules react on the surface to form products
5. Desorption: products desorb from the surface and diffuse away, freeing sites for more reactant molecules

Promoters and Poisons

A promoter (or activator) is a substance that increases the activity or selectivity of a catalyst without itself being a catalyst. Example: in the Haber process, Mo (molybdenum) and K2O are promoters for the Fe catalyst.

A catalyst poison is a substance that decreases or destroys the activity of a catalyst. Example: CO poisons the Fe catalyst in the Haber process. H2S poisons Pt catalyst in catalytic converters.

Enzyme Catalysis

Enzymes are biological catalysts (mostly proteins). They catalyse biochemical reactions with extraordinary specificity and efficiency at mild conditions (37°C, neutral pH).

Lock-and-key model: the active site of an enzyme has a specific shape that exactly fits one substrate (like a lock and key). Only the correct substrate binds.

Steps in enzyme catalysis:

1. Substrate (S) binds to the active site of enzyme (E): E + S → ES (enzyme-substrate complex)
2. Bonds in the substrate are strained or activated within the complex
3. Product (P) forms and leaves the active site: ES → E + P
4. Enzyme is regenerated and can catalyse the next reaction

Features of enzyme catalysis:

• High efficiency: one enzyme molecule can catalyse millions of reactions per second
• High specificity: one enzyme, one reaction (or one substrate)
• Optimal conditions: each enzyme has an optimum temperature and pH
• Denaturation: above optimum temperature, the enzyme unfolds and loses activity

Examples tested in NEET: zymase (glucose to ethanol), urease (urea to NH3 + CO2), amylase (starch to maltose), maltase (maltose to glucose), invertase (sucrose to glucose + fructose).

Colloids: Introduction and Types

Particle Size Classification

The particle size range of 1-1000 nm is also expressed as 1-1000 nm or 1 nm-1 micrometer. Colloidal particles are visible under an electron microscope but not under an ordinary optical microscope.

Types of Colloids by Phase

A colloidal system has two components: the dispersed phase (the substance in smaller amount, broken into colloidal-size particles) and the dispersion medium (the continuous phase in which the dispersed phase is distributed). Any of the three states of matter (solid, liquid, gas) can be the dispersed phase or the medium, giving 8 possible combinations (gas-gas gives a true mixture, not a colloid, so there are actually 8 types).

Lyophilic and Lyophobic Colloids

Lyophilic colloids ("solvent-loving"): the dispersed phase has a strong affinity for the dispersion medium.

• Examples: starch, gum, gelatin, proteins, rubber
• Stable: particles are heavily solvated (solvent molecules coat them)
• Reversible: evaporate the solvent, add it back, and the colloid re-forms
• Hard to coagulate: need large amounts of electrolyte
• Do not need a stabiliser

Lyophobic colloids ("solvent-fearing"): the dispersed phase has no affinity for the medium.

• Examples: gold sol, arsenic sulfide (As2S3) sol, silver sol, sulfur sol
• Unstable: easily coagulated by small amounts of electrolyte
• Irreversible: cannot be re-dispersed once coagulated without special treatment
• Need a stabiliser (emulsifier or protective colloid)

Association Colloids (Micelles)

Substances like soaps and detergents can exist either as true solutions (at low concentration) or as colloidal solutions (above a threshold concentration). When soap concentration exceeds the critical micelle concentration (CMC), the molecules aggregate into structures called micelles.

In a micelle:

• The hydrophobic (non-polar) tails of soap molecules point inward (away from water)
• The hydrophilic (polar/ionic) heads point outward toward water
• The interior of the micelle can dissolve non-polar (greasy) substances
• This is why soap cleans oily stains

Micelles are examples of association (lyophilic) colloids. They form spontaneously above the CMC and break up below it.

Properties of Colloids

1. Tyndall Effect

When a beam of light is passed through a colloidal solution in a dark room, the path of the beam becomes visible. This scattering of light by colloidal particles is called the Tyndall effect, named after John Tyndall (1869).

Why does it happen? Colloidal particles have sizes (1-1000 nm) comparable to the wavelength of visible light (400-700 nm). These particles scatter the light in all directions, making the beam visible.

True solutions do NOT show the Tyndall effect because dissolved particles (ions, molecules) are too small to scatter visible light.

Everyday examples: visible shaft of sunlight through a dusty room; car headlights creating a visible cone in fog; blue colour of the sky (scattering by atmospheric particles).

NEET application: "Which solution shows the Tyndall effect?" Always pick the colloid. Never true solutions.

2. Brownian Motion

Colloidal particles show a continuous, random, zigzag movement. This is called Brownian motion, named after Robert Brown who first observed it in pollen grains (1827).

Cause: the fast-moving molecules of the dispersion medium constantly bombard the colloidal particles from all sides. At any instant, the number of collisions from different directions is not equal, so the net force is never zero. This gives the colloidal particle a random kick.

Significance: Brownian motion provides kinetic energy to the colloidal particles, opposing their tendency to settle under gravity. This is why colloidal solutions are stable and particles do not settle.

3. Electrophoresis

Colloidal particles carry an electric charge (all particles in a sol carry the same sign of charge, which prevents them from coming together and coagulating).

When an electric field is applied across a colloidal solution, the charged colloidal particles migrate toward the electrode of opposite charge. This movement is called electrophoresis.

• Positively charged colloids (like Fe(OH)3 sol) move toward the cathode (negative electrode)
• Negatively charged colloids (like As2S3 sol, starch sol, gold sol, clay) move toward the anode (positive electrode)

Electrophoresis is used to determine the charge on colloidal particles and in practical applications like painting car bodies (electropainting), removal of smoke particles in Cottrell precipitators.

4. Charge on Colloidal Particles

The charge on colloidal particles arises by:

• Adsorption of ions from solution: the colloid preferentially adsorbs ions from solution. Example: As2S3 particles adsorb S2- ions from its surroundings, becoming negatively charged. AgI particles adsorb I- ions when excess KI is used, so they become negatively charged; they adsorb Ag+ ions when excess AgNO3 is used, so they become positively charged.
• Ionisation of surface groups: proteins have -COOH and -NH2 groups that ionise depending on pH. At high pH, proteins carry net negative charge; at low pH, they carry net positive charge.

5. Dialysis

Colloidal solutions prepared in the lab often contain dissolved impurities (electrolytes, small molecules). These can be removed by dialysis.

In dialysis, the colloidal solution is placed in a bag made of a semipermeable membrane (like cellophane or parchment paper) and kept in pure water or running water. Small ions and molecules (crystalloids) pass out through the membrane. Colloidal particles are too large to pass through the membrane and stay inside.

Electrodialysis is a faster version where an electric field is applied to accelerate the movement of ions through the membrane.

Practical application: kidney dialysis in patients with kidney failure uses the same principle. Blood (containing waste small molecules like urea) is passed along a semipermeable membrane, and waste molecules diffuse out.

Coagulation and Hardy-Schulze Rule

What is Coagulation?

Colloidal particles carry the same electric charge, which keeps them apart (like charges repel). If you neutralise this charge, the particles come together and aggregate into larger particles. These larger particles then settle under gravity. This process of settling is called coagulation or flocculation.

You can cause coagulation by:

• Adding an electrolyte (most common NEET method)
• Applying an electric field (electrophoresis causes particles to reach the electrode where they lose their charge and coagulate)
• Boiling (increases kinetic energy, breaks down the double layer)
• Mixing two oppositely charged colloids together

Hardy-Schulze Rule

The Hardy-Schulze rule states: the coagulating power of an ion is directly proportional to its valency (charge). A higher-valency coagulating ion is much more effective at coagulating a colloid. Quantitatively, coagulating power increases with the cube of the valency (Schulze-Hardy rule).

The coagulating ion is always the ion with the opposite charge to the colloidal particles.

For a Negatively Charged Colloid (e.g. As2S3 sol, starch sol, gold sol)

The coagulating ions are cations. Their coagulating power:

For a Positively Charged Colloid (e.g. Fe(OH)3 sol)

The coagulating ions are anions. Their coagulating power:

Coagulating Value

The coagulating value (or flocculation value) is the minimum concentration of electrolyte (in millimoles per litre) needed to just cause coagulation of a colloid. It is the inverse of coagulating power: a lower coagulating value means higher coagulating power.

Gold Number

Gold number is defined as the minimum amount of a protective colloid (in milligrams) that must be added to 10 mL of a gold sol to prevent coagulation when 1 mL of 10% NaCl solution is added.

Key rule: lower gold number = better (more efficient) protective colloid.

Lyophilic colloids (proteins, starch, gum) act as protective colloids by coating lyophobic particles and preventing electrolytes from reaching the particle surface.

Emulsions

An emulsion is a colloidal system in which both the dispersed phase and the dispersion medium are liquids. At least one of them must be water.

Types of Emulsions

Emulsifiers

Emulsions are inherently unstable (the two liquids tend to separate over time). To stabilise an emulsion, you need an emulsifying agent (or emulsifier).

An emulsifier is a substance that has both hydrophilic (water-loving) and hydrophobic (oil-loving) parts. It forms a layer at the interface between oil and water droplets, preventing them from coalescing.

• For O/W emulsions: proteins (casein in milk), soaps (sodium stearate), gum arabic
• For W/O emulsions: long-chain alcohols (cetyl alcohol), metallic soaps (calcium stearate)

Demulsification

The process of breaking an emulsion is called demulsification. Methods:

• Heating (breaks the emulsifier film)
• Adding electrolytes (coagulates the emulsifier)
• Centrifugation (cream separation from milk)
• Freeze-thaw (breaking butter back to cream)

Worked NEET Problems

Summary Cheat Sheet

Adsorption Essentials

Catalysis Essentials

Colloids Quick Reference

Coagulation Rules

• Hardy-Schulze rule: higher charge of coagulating ion = more coagulating power = less amount needed
• Negatively charged colloid(As2S3, gold, starch, clay): coagulated by cations. Power: Al3+ > Ca2+ > Na+
• Positively charged colloid(Fe(OH)3, Al(OH)3): coagulated by anions. Power: PO43- > SO42- > Cl-
• Gold number: lower = better protective colloid. Gelatin (0.005-0.01 mg) is the best; starch (25 mg) is the worst.

Emulsions Quick Reference

Frequently Confused Pairs in NEET