Adsorption

Surface chemistry studies phenomena occurring at surfaces or interfaces between different bulk phases (solid-liquid, liquid-gas, solid-gas, solid-solid, liquid-liquid interfaces). No interface exists between completely miscible gases. Interfaces are typically a few molecules thick. Many important processes like corrosion, catalysis, dissolution, and crystallisation take place at interfaces.

Adsorption is the phenomenon where molecular species accumulate at the surface of a solid or liquid rather than entering the bulk. The substance that gets adsorbed is called the adsorbate, and the surface on which adsorption occurs is the adsorbent.

Solids, especially porous or finely divided ones with large surface areas (e.g., charcoal, silica gel, alumina gel, clay), are effective adsorbents.

Examples of adsorption:

• Decrease in gas pressure in a vessel with powdered charcoal (gas molecules adsorb on charcoal surface).
• Decolourisation of dye solutions by animal charcoal (dye molecules adsorb on charcoal).
• Decolourisation of raw sugar solution by animal charcoal (colouring substances adsorb).
• Drying of air by silica gel (water molecules adsorb).

Desorption is the reverse process of removing an adsorbed substance from the surface.

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Distinction Between Adsorption And Absorption

In adsorption, accumulation is confined to the surface, not penetrating the bulk. In absorption, the substance is uniformly distributed throughout the bulk of the material.

Example: A chalk stick in ink. Colored molecules adsorb on the surface, giving it color, but the solvent is absorbed into the bulk, keeping the inside white. Water vapor is absorbed by anhydrous calcium chloride but adsorbed by silica gel.

When both adsorption and absorption occur simultaneously, the term sorption is used.

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Mechanism Of Adsorption

Adsorption is due to unbalanced or residual attractive forces on the surface particles of the adsorbent. Unlike bulk particles, surface particles are not uniformly surrounded by other particles of their kind, resulting in unsatisfied valencies or attractive forces on the surface. These forces attract and hold adsorbate particles.

The extent of adsorption increases with the surface area of the adsorbent at a given temperature and pressure.

Adsorption is always an exothermic process ($\Delta \textsf{H} < 0$). Surface energy decreases as residual forces are reduced. When a gas adsorbs, its freedom decreases, leading to a decrease in entropy ($\Delta \textsf{S} < 0$). For adsorption to be spontaneous ($\Delta \textsf{G} < 0$), $\Delta \textsf{G} = \Delta \textsf{H} - \textsf{T}\Delta \textsf{S}$ must be negative. This requires $\Delta \textsf{H}$ to be sufficiently negative to outweigh the positive value of $-\textsf{T}\Delta \textsf{S}$. As adsorption proceeds, $\Delta \textsf{H}$ becomes less negative. Equilibrium is reached when $\Delta \textsf{H} = \textsf{T}\Delta \textsf{S}$, and $\Delta \textsf{G} = 0$.

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Types Of Adsorption

Adsorption of gases on solids is classified into two types based on the nature of forces between adsorbate and adsorbent:

1. Physical adsorption (Physisorption): Adsorbate molecules are held to the solid surface by weak van der Waals’ forces.
2. Chemical adsorption (Chemisorption): Adsorbate molecules or atoms are held to the solid surface by strong chemical bonds (covalent or ionic). Often involves high activation energy (activated adsorption).

Sometimes, physisorption occurs at low temperatures and transitions to chemisorption at higher temperatures (e.g., $\textsf{H}_2$ on $\textsf{Ni}$).

Comparison of Physisorption and Chemisorption:

Property	Physisorption	Chemisorption
Forces	Van der Waals’ forces	Chemical bonds
Specificity	Not specific (van der Waals’ forces are universal)	Highly specific (requires chemical bonding possibility)
Reversibility	Reversible (easily reversible by changing pressure or temperature)	Usually irreversible (involves compound formation)
Nature of adsorbate	Adsorbs readily liquefiable gases (higher critical temperature)	Adsorbs gases that can form chemical bonds with adsorbent
Enthalpy of adsorption	Low (20-40 kJ mol$^{-1}$)	High (80-240 kJ mol$^{-1}$)
Temperature	Favourable at low temperature; decreases with increasing temperature	Favourable at high temperature (requires activation energy); increases with increasing temperature
Activation Energy	Not appreciable activation energy needed	High activation energy sometimes needed
Surface Area	Increases with increasing surface area	Increases with increasing surface area
Layers	Forms multimolecular layers at high pressure	Forms unimolecular layer

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Adsorption Isotherms

An adsorption isotherm is a curve that shows the variation in the amount of gas adsorbed by an adsorbent with pressure at a constant temperature.

Freundlich adsorption isotherm: An empirical relation proposed by Freundlich (1909) for the adsorption of a gas on a solid:

$\frac{\textsf{x}}{\textsf{m}} = \textsf{k} \cdot \textsf{p}^{1/\textsf{n}}$ (where n > 1)

Here, x = mass of gas adsorbed, m = mass of adsorbent, P = pressure, k and n are constants depending on the gas, adsorbent, and temperature. $1/\textsf{n}$ has values between 0 and 1 (typically 0.1 to 0.5).

The isotherm shows that at a fixed pressure, adsorption decreases with increasing temperature. At high pressures, the curve approaches saturation (adsorption becomes independent of pressure).

Taking logarithm of the equation: $\log \left(\frac{\textsf{x}}{\textsf{m}}\right) = \log \textsf{k} + \frac{1}{\textsf{n}} \log \textsf{p}$

Plotting $\log(\textsf{x}/\textsf{m})$ vs $\log \text{p}$ gives a straight line if the isotherm is valid. The slope gives $1/\textsf{n}$, and the intercept gives $\log \textsf{k}$.

Freundlich isotherm is approximate and fails at high pressures where saturation is reached (not predicted by the equation as $1/\textsf{n}$ would be 0).

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Adsorption From Solution Phase

Solids can also adsorb solutes from solutions (e.g., acetic acid from water by charcoal, litmus solution decolourised by charcoal, dye adsorption by precipitates). Observations:

• Adsorption decreases with increasing temperature.
• Adsorption increases with increasing surface area of adsorbent.
• Adsorption depends on the concentration of solute in solution.
• Adsorption depends on the nature of adsorbent and adsorbate.

Freundlich equation can be applied, replacing pressure (p) with equilibrium concentration (C) of the solute in solution:

$\frac{\textsf{x}}{\textsf{m}} = \textsf{k} \cdot \textsf{C}^{1/\textsf{n}}$

Taking logarithm: $\log \left(\frac{\textsf{x}}{\textsf{m}}\right) = \log \textsf{k} + \frac{1}{\textsf{n}} \log \textsf{C}$

A plot of $\log(\textsf{x}/\textsf{m})$ vs $\log \text{C}$ should be a straight line if the relationship holds. This can be verified experimentally.

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Applications Of Adsorption

Adsorption has numerous practical applications:

• Production of high vacuum (charcoal adsorbs remaining gases).
• Gas masks (adsorbent removes poisonous gases).
• Control of humidity (silica/alumina gels as adsorbents).
• Removal of colouring matter (animal charcoal decolourises solutions).
• Heterogeneous catalysis (adsorption of reactants on solid catalyst surface).
• Separation of inert gases (differential adsorption by charcoal).
• In curing diseases (drugs adsorb on germs).
• Froth floatation process (pine oil selectively adsorbs on sulphide ore particles).
• Adsorption indicators (dyes adsorbed by precipitates at endpoint, e.g., eosin on silver halide).
• Chromatographic analysis (separation based on differential adsorption).

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Catalysis

Catalysis is the phenomenon where the rate of a chemical reaction is altered by the presence of a substance called a catalyst, which itself remains chemically and quantitatively unchanged after the reaction. Catalysts generally accelerate reaction rates; those that reduce the rate are sometimes called inhibitors.

Catalysts function by providing an alternative reaction pathway with a lower activation energy, allowing a larger fraction of reactant molecules to overcome the energy barrier and react (discussed in Chemical Kinetics, Unit 4). They participate by forming temporary bonds with reactants to create intermediate complexes.

Substances that enhance the activity of a catalyst are promoters (e.g., Mo as promoter for Fe in Haber's process), while substances that decrease activity are poisons (e.g., CO poisoning of metal catalysts).

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Homogeneous And Heterogeneous Catalysis

Catalysis is classified based on the phase of the reactants, products, and catalyst:

1. Homogeneous catalysis: Reactants, products, and catalyst are all in the same phase (gas or liquid).
   • Example: Oxidation of $\textsf{SO}_2$ to $\textsf{SO}_3$ with $\textsf{O}_2$ using $\textsf{NO(g)}$ as catalyst (all gases).
   • Example: Hydrolysis of methyl acetate or sugar catalysed by $\textsf{H}^{+}$ ions in aqueous solution (all in liquid phase).
2. Heterogeneous catalysis: Reactants and catalyst are in different phases. Typically, solid catalysts are used with gaseous or liquid reactants.
   • Example: Oxidation of $\textsf{SO}_2$ to $\textsf{SO}_3$ using $\textsf{Pt(s)}$ or $\textsf{V}_2\text{O}_5\text{(s)}$ catalyst (gas reactants, solid catalyst).
   • Example: Haber's process for $\textsf{NH}_3$ synthesis using $\textsf{Fe(s)}$ catalyst (gas reactants, solid catalyst).
   • Example: Hydrogenation of vegetable oils using finely divided $\textsf{Ni(s)}$ catalyst (liquid/gas reactants, solid catalyst).

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Adsorption Theory Of Heterogeneous Catalysis

The modern adsorption theory explains heterogeneous catalysis, particularly involving solid catalysts and gaseous reactants. It combines the concepts of adsorption and intermediate compound formation. Mechanism involves five steps:

1. Diffusion of reactants from the bulk phase to the surface of the solid catalyst.
2. Adsorption of reactant molecules onto the catalyst surface. This is often chemisorption due to free valencies on the surface.
3. Occurrence of chemical reaction on the catalyst surface. Adsorbed reactant molecules interact to form an intermediate complex, which then reacts to form product molecules (often via surface reaction mechanisms).
4. Desorption of reaction products from the catalyst surface. Products detach from the surface once formed.
5. Diffusion of reaction products away from the catalyst surface into the bulk phase.

The catalytic activity is primarily located on the catalyst's surface. Free valencies on the surface provide sites for adsorption and reaction.

This theory explains why a catalyst is effective in small amounts (it's regenerated) and unchanged chemically. However, it doesn't fully explain promoter/poison action.

Important features of solid catalysts:

• Activity: Depends on the strength of chemisorption. Reactants must adsorb strongly enough to become active but not too strongly, which would block the surface. Activity varies greatly among different metals/substances (e.g., maximum activity for hydrogenation shown by Group 7-9 metals).
• Selectivity: Ability to direct a reaction towards a specific product when multiple products are possible from the same reactants. Selectivity is highly specific to the catalyst used. Example: $\textsf{H}_2$ and $\textsf{CO}$ can form different products ($\textsf{CH}_4$, $\textsf{CH}_3\text{OH}$) depending on the catalyst used (e.g., $\textsf{Ni}$ vs $\textsf{Cu/ZnO/Cr}_2\text{O}_3$).

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Shape-Selective Catalysis By Zeolites

Shape-selective catalysis is a type of catalysis where the reaction rate depends on the pore structure of the catalyst and the size/shape of reactant and product molecules. Zeolites are prominent examples of shape-selective catalysts.

• Zeolites are microporous aluminosilicates with a three-dimensional network structure containing pores and cavities of specific sizes and shapes.
• Their structure is like a honeycomb. The size and shape of reactants must fit into the pores to reach the active sites, and the products must be able to diffuse out.
• They are used extensively in petrochemical industries for reactions like cracking (breaking down large hydrocarbons) and isomerisation.
• Example: ZSM-5 is a zeolite catalyst used to convert alcohols directly into gasoline by dehydration followed by hydrocarbon formation.

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Enzyme Catalysis

Enzymes are complex protein molecules produced by living organisms. They are highly effective catalysts for numerous biochemical reactions occurring within living cells (biochemical catalysis). Enzymes are often colloidal in nature.

Examples of enzyme-catalysed reactions:

• Inversion of cane sugar by invertase.
• Conversion of glucose to ethanol by zymase.
• Conversion of starch to maltose by diastase.
• Conversion of maltose to glucose by maltase.
• Hydrolysis of urea by urease.
• Breakdown of proteins by pepsin and trypsin.

Enzyme	Source	Enzymatic reaction
Invertase	Yeast	Sucrose $\to$ Glucose + fructose
Zymase	Yeast	Glucose $\to$ Ethyl alcohol + carbon dioxide
Diastase	Malt	Starch $\to$ Maltose
Maltase	Yeast	Maltose $\to$ Glucose
Urease	Soyabean	Urea $\to$ Ammonia + carbon dioxide
Pepsin	Stomach	Proteins $\to$ Peptides
Trypsin	Pancreas	Proteins $\to$ Amino acids

Characteristics of enzyme catalysis:

• High efficiency: Highly effective, one enzyme molecule can catalyse millions of reactant molecules per minute.
• High specificity: Highly specific, each enzyme typically catalyses only one specific reaction or type of reaction.
• Optimum temperature: Most active within a narrow temperature range (optimum temperature), typically 298-310 K. Activity decreases outside this range.
• Optimum pH: Maximum activity at a specific pH value (optimum pH), typically between 5-7.
• Activators and Co-enzymes: Activity is enhanced by certain substances called activators (often metal ions like $\textsf{Na}^{+}, \textsf{Mn}^{2+}$) or co-enzymes (non-protein organic molecules, often vitamins).
• Inhibitors and poisons: Activity can be reduced or destroyed by inhibitors or poisons that interact with active sites.

Mechanism of enzyme catalysis: Enzymes have specific cavities on their surface called active sites with characteristic shapes and functional groups. The reactant molecules (substrate) with a complementary shape fit into these active sites (like a key in a lock) to form an enzyme-substrate complex (ES). This complex is activated ($\textsf{ES}^{\ne}$) and then decomposes to form products and release the enzyme.

Steps:

1. $\textsf{E} + \textsf{S} \to \textsf{ES}$ (Enzyme + Substrate form Enzyme-Substrate complex)

2. $\textsf{ES} \to \textsf{ES}^{\ne}$ (Formation of activated complex)

3. $\textsf{ES}^{\ne} \to \textsf{E} + \textsf{P}$ (Decomposition to Enzyme + Products)

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Catalysts In Industry

Catalysts are widely used in industrial chemical processes to increase reaction rates and improve efficiency. Some important examples:

Process	Catalyst	Conditions
Haber’s process for ammonia synthesis (N$_2$ + 3H$_2$ $\to$ 2NH$_3$)	Finely divided iron; Molybdenum as promoter	200 bar pressure, 723-773K temperature
Ostwald’s process for nitric acid manufacture (4NH$_3$ + 5O$_2$ $\to$ 4NO + 6H$_2$O, etc.)	Platinised asbestos	573K temperature
Contact process for sulphuric acid manufacture (2SO$_2$ + O$_2$ $\to$ 2SO$_3$, etc.)	Platinised asbestos or Vanadium pentoxide (V$_2$O$_5$)	673-723K temperature

Answer:

Colloids

Beyond homogeneous solutions and heterogeneous suspensions, there exists an intermediate category called colloidal dispersions or simply colloids. A colloid is a heterogeneous system where one substance (dispersed phase) is dispersed as very fine particles within another substance (dispersion medium).

The key difference between solutions, colloids, and suspensions is particle size:

• True solutions: Particles are ions or small molecules (< 1 nm diameter). Homogeneous.
• Colloidal dispersions: Dispersed particles are larger than small molecules but small enough to remain suspended, typically with diameters between 1 nm and 1000 nm. Heterogeneous (though may appear homogeneous).
• Suspensions: Particles are large (> 1000 nm), visible to the naked eye, and settle down over time (e.g., sand in water). Heterogeneous.

Colloidal particles have an enormous surface area per unit mass due to their small size. This large surface area contributes to the special properties of colloids.

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Classification Of Colloids

Colloids are classified based on various criteria:

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Classification Based On Physical State Of Dispersed Phase And Dispersion Medium

Based on whether the dispersed phase and dispersion medium are solid, liquid, or gas, eight types of colloidal systems are possible (gas in gas is a homogeneous mixture). Common types include sols (solid dispersed in liquid), gels (liquid dispersed in solid), and emulsions (liquid dispersed in liquid).

Dispersed phase	Dispersion medium	Type of colloid	Examples
Solid	Solid	Solid sol	Some coloured glasses, gem stones
Solid	Liquid	Sol	Paints, cell fluids
Solid	Gas	Aerosol	Smoke, dust
Liquid	Solid	Gel	Cheese, jellies
Liquid	Liquid	Emulsion	Milk, hair cream, butter
Liquid	Gas	Aerosol	Fog, mist, cloud, insecticide sprays
Gas	Solid	Solid sol (Solid foam)	Pumice stone, foam rubber
Gas	Liquid	Foam	Froth, whipped cream, soap lather

If the dispersion medium is water, the sol is called aquasol or hydrosol. If alcohol, it's alcosol, etc.

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Classification Based On Nature Of Interaction Between Dispersed Phase And Dispersion Medium

Based on the affinity between the dispersed phase and dispersion medium, colloidal sols are divided into:

1. Lyophilic colloids (Solvent-attracting): Formed directly by mixing substances like gum, gelatin, starch, etc., with a suitable liquid medium. These sols are quite stable. If the medium is removed (e.g., by evaporation), the sol can be reconstituted by simple remixing (reversible sols). If water is the medium, they are hydrophilic colloids.
2. Lyophobic colloids (Solvent-repelling): Substances like metals, sulfides, etc., do not form sols by simple mixing. Special methods are required for their preparation. These sols are less stable and readily precipitated by adding small amounts of electrolytes, heating, or shaking (irreversible sols). If water is the medium, they are hydrophobic colloids. Lyophobic sols need stabilising agents.

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Classification Based On Type Of Particles Of The Dispersed Phase, Multimolecular, Macromolecular And Associated Colloids

Based on how the particles of the dispersed phase are formed:

1. Multimolecular colloids: Formed when a large number of atoms or small molecules aggregate together to form particles within the colloidal size range. Example: Sulphur sol (aggregates of $\textsf{S}_8$ molecules), gold sol (aggregates of gold atoms).
2. Macromolecular colloids: Formed when individual molecules are large enough (macromolecules) to be within the colloidal size range. Solutions of macromolecules in suitable solvents form these colloids. These are quite stable and resemble true solutions. Examples: Naturally occurring - starch, cellulose, proteins, enzymes. Man-made - polythene, nylon, synthetic rubber.
3. Associated colloids (Micelles): Formed by substances that behave as normal electrolytes at low concentrations but form aggregates (micelles) at higher concentrations. Formation occurs above a specific Kraft temperature (Tk) and above a critical micelle concentration (CMC). These colloids have both lyophobic (hydrocarbon tail) and lyophilic (polar head) parts.
   • Example: Soaps ($\textsf{RCOO}^{-}\textsf{Na}^{+}$) and detergents. In water, $\textsf{RCOO}^{-}$ ions have a hydrophobic tail (R) and a hydrophilic head ($\textsf{COO}^{-}$). At low concentration, they stay on the surface. Above CMC, they aggregate to form micelles, typically spherical, with hydrophobic tails pointing inwards and hydrophilic heads outwards, interacting with water.
   • Cleansing action of soaps: Micelles encapsulate oil/grease droplets, with the hydrophobic tails dissolving in the grease and hydrophilic heads facing outwards into water. The negatively charged micelle surface allows the grease droplet to be suspended in water and washed away.

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Preparation Of Colloids

Colloidal solutions are prepared by various methods, broadly classified as dispersion methods (breaking down larger particles) and condensation methods (aggregating smaller particles).

• Chemical methods (Condensation): Forming molecules by chemical reactions (double decomposition, oxidation, reduction, hydrolysis), which then aggregate to colloidal size.
  • Double decomposition: $\textsf{As}_2\text{O}_3 + 3\textsf{H}_2\textsf{S} \to \textsf{As}_2\textsf{S}_3 \text{(sol)} + 3\textsf{H}_2\textsf{O}$
  • Oxidation: $\textsf{SO}_2 + 2\textsf{H}_2\textsf{S} \to 3\textsf{S} \text{(sol)} + 2\textsf{H}_2\textsf{O}$
  • Reduction: $2\textsf{AuCl}_3 + 3\textsf{HCHO} + 3\textsf{H}_2\textsf{O} \to 2\textsf{Au} \text{(sol)} + 3\textsf{HCOOH} + 6\textsf{HCl}$
  • Hydrolysis: $\textsf{FeCl}_3 + 3\textsf{H}_2\textsf{O} \to \textsf{Fe(OH)}_3 \text{(sol)} + 3\textsf{HCl}$
• Electrical disintegration or Bredig’s Arc method (Dispersion and Condensation): Used for preparing sols of metals (Au, Ag, Pt). An electric arc is struck between metal electrodes immersed in the dispersion medium. Intense heat vaporises the metal, which condenses into colloidal-sized particles.
• Peptization (Dispersion): Converting a precipitate into a colloidal sol by shaking it with the dispersion medium in the presence of a small amount of electrolyte (peptizing agent). The precipitate adsorbs ions from the electrolyte on its surface, developing a charge that causes it to break down into colloidal particles and disperse.

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Purification Of Colloidal Solutions

Colloidal solutions often contain excess electrolytes and other soluble impurities that can affect stability (large amounts coagulate, traces are needed). Purification methods reduce impurity concentration to a minimum:

• Dialysis: Removing dissolved substances (ions, small molecules) from a colloidal solution by diffusion through a semipermeable membrane (like parchment or cellophane). Small particles pass through, but colloidal particles are retained. The apparatus is a dialyser.
• Electro-dialysis: Speeding up dialysis by applying an electric field. Useful when impurities are electrolytes. Ions migrate towards oppositely charged electrodes, leaving the colloidal solution.
• Ultrafiltration: Separating colloidal particles from solvent and soluble solutes using special filters (ultrafilters) with pores small enough to retain colloidal particles but allow smaller particles/solvent to pass. Ordinary filter paper has large pores. Ultrafilters are prepared by impregnating filter paper with collodion solution (nitro-cellulose in alcohol-ether) and hardening. Pressure or suction is used to speed up the slow process. Colloidal particles are left on the filter and redispersed in fresh medium.

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Properties Of Colloidal Solutions

Colloidal solutions exhibit various properties due to their unique particle size and surface area:

• Colligative properties: Number of colloidal particles is much smaller than in true solutions at the same concentration (as colloidal particles are aggregates). Thus, colligative properties (osmotic pressure, vapour pressure lowering, boiling point elevation, freezing point depression) show smaller values compared to true solutions.
• Tyndall effect: Scattering of light by colloidal particles when a beam of light is passed through the colloidal solution. The path of the light beam becomes visible as a bright cone (Tyndall cone). True solutions do not show this effect as particles are too small to scatter light. Conditions for Tyndall effect: particle size not much smaller than wavelength, significant difference in refractive indices of dispersed phase and medium. Used to distinguish colloids from true solutions (ultramicroscope utilises this).
• Colour: Depends on the wavelength of light scattered by particles, which in turn depends on particle size and nature. Colour can also change depending on how light is observed (reflected vs transmitted).
• Brownian movement: Continuous, random, zig-zag motion of colloidal particles observed under an ultramicroscope. Caused by unbalanced bombardment of colloidal particles by molecules of the dispersion medium. Brownian movement prevents particles from settling, contributing to sol stability.
• Charge on colloidal particles: Colloidal particles always carry an electric charge, which is the same (either positive or negative) for all particles in a given sol. The charge arises from electron capture during electrodispersion, preferential adsorption of ions from solution, or formation of an electrical double layer. Preferential adsorption of common ions is a major cause (e.g., $\textsf{AgI}$ precipitating in excess $\textsf{KI}$ adsorbs $\textsf{I}^{-}$ and is negative; in excess $\textsf{AgNO}_3$ adsorbs $\textsf{Ag}^{+}$ and is positive).

The layer of specifically adsorbed ions attracts a second layer of counter-ions from the medium, forming the Helmholtz electrical double layer (fixed layer of adsorbed ions + diffused layer of counter-ions). The potential difference between these layers is the electrokinetic potential or zeta potential. Repulsion between particles due to zeta potential provides stability, preventing aggregation.

• Electrophoresis: Movement of charged colloidal particles under an applied electric potential. Positively charged particles move towards the cathode, negative particles towards the anode. Used to confirm charge presence and separate charged particles. If particle movement is prevented, the medium moves (electroosmosis).
• Coagulation or precipitation: Process of destabilising a colloidal sol, causing particles to aggregate and settle down. Stability of lyophobic sols depends on charge. Removing charge leads to coagulation. Methods:
  • Electrophoresis: Particles discharge at electrodes and coagulate.
  • Mixing oppositely charged sols: Charges are neutralised, causing mutual coagulation.
  • Boiling: Disturbs the adsorbed layer, reducing charge and stability.
  • Persistent dialysis: Removes essential traces of electrolyte, reducing stability.
  • Addition of electrolytes: Excess electrolytes provide oppositely charged ions that neutralise particle charge. The ion responsible for coagulation is the coagulating ion.

  Hardy-Schulze rule: States that the greater the valence of the coagulating ion, the greater its coagulating power. For a negative sol, coagulating power order is $\textsf{Al}^{3+}> \textsf{Ba}^{2+} > \textsf{Na}^{+}$. For a positive sol, it is $[\textsf{Fe(CN)}_6]^{4-}> \textsf{PO}_4^{3-} > \textsf{SO}_4^{2-} > \textsf{Cl}^{-}$.

  Coagulating value: Minimum concentration of electrolyte (millimoles/L) needed for precipitation in two hours. Lower coagulating value means higher coagulating power.

  Coagulation of lyophilic sols: More stable than lyophobic sols due to charge and solvation (particles surrounded by medium sheath). Coagulated by removing solvation (adding solvent like alcohol) and adding electrolyte.

  Protection of colloids: Lyophilic sols can protect lyophobic sols. Adding a lyophilic sol to a lyophobic sol forms a layer of lyophilic particles around the lyophobic ones, protecting them from electrolytes. Lyophilic colloids used for this are protective colloids.

Answer:

Emulsions

Emulsions are liquid-liquid colloidal systems where finely divided droplets of one liquid are dispersed in another liquid. Typically, one liquid is water. There are two types:

• Oil dispersed in water (O/W type): Water is the dispersion medium. Examples: Milk (fat in water), vanishing cream.
• Water dispersed in oil (W/O type): Oil is the dispersion medium. Examples: Butter (water in fat), cream.

Emulsions are often unstable and separate into layers. An emulsifying agent is added to stabilise the emulsion by forming an interfacial film between the droplets and the medium. Emulsifying agents for O/W include proteins, gums, soaps. For W/O, they include heavy metal salts of fatty acids, long-chain alcohols.

Emulsions can be diluted with the dispersion medium. They show Brownian movement and Tyndall effect. Emulsions can be broken into constituent liquids by heating, freezing, or centrifuging.

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Colloids Around Us

Many everyday substances and natural phenomena involve colloids:

• Blue colour of the sky (scattering of blue light by colloidal dust/water particles in air).
• Fog, mist, rain (condensation of water vapor on colloidal particles in air). Clouds are aerosols (liquid in gas).
• Food articles (milk, butter, ice creams, juices).
• Blood (colloidal solution of albuminoids; coagulation by alum/ferric chloride).
• Soils (colloidal nature with humus as protective colloid, helps adsorb moisture/nutrients).
• Formation of delta (coagulation of colloidal clay in river water by sea water electrolytes).

Applications of colloids in industry:

• Electrical precipitation of smoke (Cottrell precipitator). Smoke particles are charged and precipitated on oppositely charged plates before exiting chimneys.
• Purification of drinking water (alum addition coagulates suspended impurities).
• Medicines (many are colloidal, e.g., argyrol, colloidal antimony, gold, milk of magnesia - effective due to large surface area).
• Tanning (coagulation of positively charged animal hides by negatively charged tannin).
• Cleansing action of soaps and detergents (micelle formation and emulsification).
• Photographic plates/films (emulsion of silver bromide in gelatin).
• Rubber industry (coagulation of latex, a rubber sol).
• Industrial products (paints, inks, plastics, lubricants, cement).

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