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By going through these CBSE Class 12 Chemistry Notes Chapter 5 Surface Chemistry, students can recall all the concepts quickly.

Surface Chemistry Notes Class 12 Chemistry Chapter 5

The branch of chemistry which deals with the nature of surface and species present on it is called surface chemistry. Adsorption on solid or on solution surfaces and colloidal properties are important surface effects.

Adsorption: The phenomenon of higher concentration of molecular species (gases or liquids) on the surface of solids than in the bulk is called adsorption.

The solid on the surface of which adsorption occurs is called adsorbent. The substances that get adsorbed on the solid surface is called adsorbate. The adsorbent may be a solid or a liquid and the adsorbate may be a gas or a liquid.

The phenomenon of absorption differs from adsorption as:

Absorption	Adsorption
1. It is the phenomenon in which the particles of gas or liquid get uniformly distributed throughout the body of the solid.	1. It is the phenomenon of higher concentration of particles of gas or liquid on the surface than if! the bulk of the solid.
2. Absorption occurs at uniform rate.	2. Adsorption is rapid in the beginning and its rate slowly decreases.

Types Of Adsorption:
Depending upon the nature of forces between the molecules of the adsorbate and the adsorbent, the adsorption maybe classified as: physical adsorption and chemical adsorption.
1. Physical adsorption: When the particles of the adsorbate are held to the surface of the adsorbent by the weak forces such as van der Waals forces, the adsorption is called physical adsorption or physisorption. The attractive forces are weak and therefore, these can be easily overcome either by increasing the temperature or by decreasing the pressure. In other words, physical adsorption can be easily reversed or decreased.

2. Chemical adsorption: When the molecules of the adsorbate are held to the surface of the adsorbent by the chemical forces, the adsorption is called chemical adsorption or chemisorption, hi this case, a chemical reaction occurs between the adsorbed molecules and the molecules or atoms of adsorbent.

Table: Comparison between Physisorption and Chemisorption:

Physisorption	Chemisorption
1. Enthalpy of adsorption, usually is of the order of 20 – 40 kJ mol-1	1. Enthalpy of adsorption, is of the order 40-400 kJ mol-1.
2. Molecules of adsorbate and adsorbent are held by van der Waals interactions.	2. Molecules of adsorbate and adsorbent are held by chemical bonds.
3. It usually takes place at low temperature and decreases with increasing temperature.	3. It takes place at relatively high temperatures.
4. It is not very specific i.e., all gases are absorbed on all solids to some extent.	4. It is highly specific and takes place when there is some possibility of compound formation between the adsorbate and the adsorbent molecules.
5. Multi-molecular layers may be formed on the adsorbent.	5. Usually mono-molecular layer is formed on the adsorbqnt,
6. It does not require any activation energy.	6. It requires activation energy.
7. The amount of gas adsorbed is related to the ease of liquefaction of the gas.	7. There is no such correlation.
8. it is reversible in nature.	8. It is irreversible in nature.

Freundlich Adsorption Isotherm:
The adsorption of a gas on the surface of the solid depends upon the pressure of the gas. The extent of adsorption is generally expressed as x/m where m is the mass of the adsorbent and x is the mass of adsorbate when equilibrium has been attained. On the basis of experimental studies, Freundlich gave the following relationship between the amount of gas adsorbed (x) per unit mass of the adsorbent (m) and the pressure (p).
\(\frac{x}{m}\) = kp^1/n
where n is a constant (whole number) which depends upon the nature of adsorba te and adsorbent.

A graph between the amount (x/m) adsorbed by an adsorbent and the equilibrium pressure (or concentration for solutions) of the adsorbate at constant temperature is called Adsorption Isotherm.

At low pressure \(\frac{x}{m}\) ∝ p¹ ………..(i)

At high pressure \(\frac{x}{m}\) ∝ p⁰ ………..(ii)

In the intermediate range of pressure, combining (i) and (ii)
\(\frac{x}{m}\) ∝ p^0-1
∝ p1/n where n is an integer (n > 1)
or
\(\frac{x}{m}\) = kp^1/n ………..(iii)
where k is a constant depending upon the nature of the adsorbate and adsorbent.

This relationship is called Freundlich Adsorption Isotherm
Taking logs on both sides of (iii)
log \(\frac{x}{m}\) = log k + \(\frac{1}{n}\) log p

A graph between log \(\frac{x}{m}\) against log p should, therefore, be a straight line with slope equal to \(\frac{1}{n}\) and ordinate intercept equal to log K.

Adsorption From Solution Phase:
Solids can adsorb solutes from solutions so when a solution of acetic acid in water is shaken with charcoal, a part of the acid is adsorbed by charcoal and the concentration of the acid decreases in the solution.

The following conclusions have been made regarding adsorption from solution phase:

1. The extent of adsorption decreases with increase in temperature.
2. The extent of adsorption increases with the increase in the surface area of the absorbent.
3. The extent of adsorption depends upon the concentration of the solute in the solution.
4. The extent of adsorption depends upon the nature of the absorbent and the adsorbate.

Freundlich equation as applied to solutions is modified as \(\frac{x}{m}\) = k C^1/n  i.e., where C is the equilibrium concentration,
or
log \(\frac{x}{m}\) = log k + \(\frac{1}{n}\) log C

Applications of Adsorption:
1. Production of high vacuum: The remaining traces of air can be adsorbed by charcoal from a vessel evacuated by a vacuum pump to give a very high vacuum.

2. Gas masks: Gas mask (a device which consists of activated charcoal or mixture of adsorbents) is usually used for breathing in coal mines to adsorb poisonous gases.

3. Control of humidity: Silica and aluminium gels are used as adsorbents for removing moisture and controlling humidity.

4. Removal of colouring matter from solutions: Animal charcoal removes colours of solutions by adsorbing coloured impurities.

5. Heterogeneous catalysis: Adsorption of reactants on the solid surface of the catalysts increases the rate of reaction. There are many gaseous reactions of industrial importance involving solid catalysts. Manufacture of ammonia using iron as a catalyst, manufacture of H₂SO₄ by contact process and use of finely divided nickel in the hydrogenation of oils are excellent examples of heterogeneous catalysis.

6. Separation of inert gases: Due to the difference in degree of adsorption of gases by charcoal, a mixture of noble gases can be separated by adsorption on coconut charcoal at different temperatures.

7. In curing diseases: A number of drugs are used to kill germs by getting adsorbed on them.

8. Froth floatation process: A low grade sulphide ore is concentrated by separating it from silica and other earthy matter by this method using pine oil and frothing agent (see Unit 6).

9. Adsorption indicators: Surfaces of certain precipitates such as silver halides have the property of adsorbing some dyes like eosfn, fluorescein, etc. and thereby producing a characteristic colour at the end point.

10. Chromatographic analysis: Chromatographic analysis based on the phenomenon of adsorption finds a number of applications in analytical and industrial fields.

Catalysis: Potassium chlorate when heated strongly decomposes slowly giving dioxygen. The decomposition between 653 – 873 K.
2 KClO₃ → 2 KCl + 3O₂

However with a little of MnO a decomposition occurs only at 473 – 633 K and also at a much accelerated rate. MnO2 is a catalyst for this reaction.

A substance which accelerates the rate of a chemical reaction, itself remaining chemically and quantitatively unchanged after the reaction is called a catalyst.

Promoter is a substance that enhances the activity of a catalyst, while Poison is a substance that decreases the activity of a catalyst. In the reaction for the manufacture of NH₃ by Haber’s process.

Fe is catalyst and Molybdenum (MO) is a promoter.

Homogeneous And Heterogeneous Catalysis:
When the reactants and the catalyst are in the same phase (liquid or gas), the process is said to be homogeneous catalysis.

Example:

When the reactants and the catalyst are in different phases, the process is called heterogeneous catalysis.

Important Features of solid catalysts are:
(a) Activity: It depends upon the strength of chemisorption to a large extent. .
(b) Selectivity: Selectivity of a catalyst is its ability to direct a reaction to yield a particular product. For example, starting with H₂ and CO and using different catalysts, we get different products.

Thus a catalyst is highly selective in nature, i.e., a given substance can act as a catalyst only in a particular reaction and not for all the reactions. A catalyst for a particular reaction may fail to catalyse any other reaction.

Shape-Selective Catalysis by Zeolites: The catalytical reaction that depends upon the pore structure of the catalyst and the size of the reactant and product molecules is called shape selective catalysis. Zeolites are good shape-selective catalysts because of their honey comb¬like structure.

An important Zeolite catalyst used in the petroleum industry is ZSM-5. It converts alcohols directly into gasoline (petrol) by dehydrating them to give a mixture of hydrocarbons.

Enzyme Catalysis: Enzymes are complex nitrogeneous organic compounds which are produced by living plants and animals. They are actually protein molecules of high molecular mass, and form colloidal solutions in water. The enzymes are also referred to as Bio-Chemical Catalysts as they also occur in the bodies of animals and plants and such a phenomenon is known as Bio-Chemical Catalysis.

The following are examples of enzyme-catalysed reactions:
1. Inversion of cane-sugar

2. Conversion of glucose into ethyl alcohol.

3. Decomposition of urea into ammonia and CO₂.

Enzyme	Source	Enzymatic Reaction
Invertase	Yeast	Sucrose → Glucose and fructose
Zymase	Yeast	Glucose → Ethyl alcohol and carbondioxide
Diastase	Malt	Starch → Maltose
Maltase	Yeast	Maltose → Glucose
Urease	Soyabean	Urea → Ammonia and carbon dioxide
Pepsin	Stomach	Proteins → Amino acids

Characteristics Of Enzyme Catalysis:
Enzyme catalysis is unique in its efficiency and high degree of specificity. The following characteristics are exhibited by enzyme catalysts:
1. Most highly efficient: One molecule of an enzyme may transform one million molecules of the reactant per minute.

2. Highly specific nature: Each enzyme is specific for a given reaction, i.e., one catalyst cannot catalyse more than one reaction. For example, the enzyme urease catalyses the hydrolysis of urea only. It does not catalyse hydrolysis of any other amide.

3. Highly active under optimum pH: The rate of an enzyme- catalysed reaction is maximum at a particular pH called optimum pH, which is between pH value 5-7.

4. Highly active under optimum temperature: The rate of an enzyme reaction becomes,maximum at a definite temperature, called the optimum temperature. On either side of the optimum temperature, the enzyme activity decreases. The optimum temperature range for enzymatic activity is 298-310 K. Human body temperature being 310 K is suited to enzyme-catalysed reaction.

5. Increasing activity in presence of activators and co¬enzymes: The enzymatic activity is increased in the presence of certain substances, known as co-enzymes. It has been observed that when a small non-protein (vitamin) is present along with an enzyme, the catalytic activity is enhanced considerably.

Activators are generally metal ions such as Na⁺ Mn²⁺, CO²⁺, Cu²⁺, etc. These metal ions, when weakly bonded to enzyme molecules, increase their catalytic activity. Amylase in presence of sodium chloride i.e., Na⁺ ions are catalytically very active.

6. Influence of inhibitors and poisons: Like ordinary catalysts enzymes are also inhibited or poisoned by the presence of certain substances. The inhibitors or poisons interact with the active functional groups on the enzyme surface and often reduce or completely destroy the catalytic activity of the enzymes. The use of many drugs is related to their action as enzyme inhibitors in the body.

Mechanism of Enzyme Catalysis: There are a number of cavities present on the surface of collodial particles of enzymes. These cavities are of characteristic shape and possess active groups such as – NH₂, – COOH, – SH, – OH etc. These are actually the active centres on the surface of enzyme particles. The molecules of the reactant (substrate), which have complementary shape, fit into these cavities just like a key fits into a lock.

An activated complex is formed which then decomposes to yield the products in two steps as outlined below:
Step I: E + S → ES*
Step II: ES* → E + P

Some Industrial Catalytical Processes:

Process	Catalyst
1. Haber’s process for the manu-facture of ammonia N2(g) + 3H2(g) → 2NH3(g)	1. Finely divided iron, molybdenum as promoter; conditions: 200 bar pressure and 723 – 773K temperature.
2. Ostwald’s process for the manu-facture of nitric acid. 4NH3(g) + 5O2(g) → 4N0(g) + 6H2O(g) 2NO(g) + O2(g) → 2NO2(g) 4NO2(g) + 2H2O(l) + O2(g) → 4HNO3(l)	2. Platinised asbestos; temperature: 573K.
3. Contact process for the manu-facture of sulphuric acid. 2SO2(g) + O2 (g) ⇌ 2SO3(g) SO3(g) + H2SO4 (l) → H2S2O7 (l)oleum H2S2O7 (l)+H2O (l) → 2H2SO4 (l)	3. Platinised asbestos or vanadium pentoxide (V205); temperature 673-723K.

Colloids: A colloid is a heterogeneous system in which one substance is dispersed (dispersed phase) as very fine particles in another substance called dispersion medium.

The essential difference between a solution and a colloid is that of particle size. Their size is in between that of true solution and suspension

Classification of Colloids: Colloids are classified on the basis of the following criteria:
(a) Physical state of dispersed phase and dispersion medium.
(b) Nature of interaction between dispersed phase and dispersion medium.
(c) Type of particles of the dispersed phase.
(a) Physical state of dispersed phase and dispersion medium: Depending upon whether the dispersed phase and the dispersion medium are solids, liquids or gases, eight types of colloidal systems are possible.

Types of Collodial Systems:

(b) Depending upon the nature of interactions between the dispersed phase and dispersion medium, the colloids can be classified as Lyophilic Colloids and Lyophobic Colloids.

1. Lyophilic collids: The colloidal solutions in which the particles of the dispersed phase have a great affinity (or love) for the dispersion medium are called lyophilic colloids. Such solutions are reversible in nature. In case water acts as the dispersion medium, the lyophilic colloid is called hydrophilic colloid. The common examples of lyophilic colloids are glue, gelatin, starch, proteins, rubber etc.

2. Lyophobic colloids: The colloidal solutions in which the particles of the dispersed phase have no affinity or love, rather have hatred for the dispersion medium, are called lyophobic collids. The solutions of metals like Ag and Au, hydroxides like Al (OH)₃ and Fe (OH)₃ and metal sulphides like As₂S₃ are examples of lyophobic colloids.
Such sols are formed with difficulty. They are irreversible in nature.

Multimolecular Macromolecular And Associated Colloids:
Depending upon the molecular size, the colloids can be classified as:
1. Multimolecular colloids: In this type, the particles consist of an aggregate of atoms or small molecules with molecular size less than 1 nm. For example, sols of gold atoms and sulphur (S₈) molecules. In these colloids, the particles are held together by Van der Waals forces.

2. Macromolecular colloids: In this type, the particles of the dispersed phase are sufficiently big in size (macro) to be of colloidal dimensions. In this case, a large number of small molecules are joined together through their primary valencies to form giant molecules.

These molecules are called macro molecules and each macromolecule may consist of hundreds or thousands of simple molecules. The solution of such moleucles are called macromolecular soluions. For example, colloidal solution of starch, cellulose, etc.

3. Associated colloids: These are the substances which behave as normal electrolytes at low concentration but behave, as colloidal particles at higher concentration. These associated particles are also called miscelles. For example, in aqueous solution, soap (sodium stearate) ionises as:

In concentrated solutions, these ions get associated to form an aggregate of colloidal size.

General Methods Of Preparation Of Sols:
Lyophilic sols are readily formed by simply mixing the dispersed phase and the dispersion medium under ordinary conditions.

Lyophobic sols can generally be prepared by two methods.
1. Condensation methods,
2. Dispersion methods.

→ Condensation methods: In these methods, the smaller particles are condensed suitably to be of colloidal size. This can be done by chemical reactions or by exchange of solvent.

→ Dispersion methods: In these methods, the large particles of a substance (suspension) are broken into smaller particles. This can be done by mechanical dispersion, by electrical dispersion or Bredig’s arc method and by peptisation.

Purification Of Colloidal Solutions:
The colloidal solutions prepared usually contain impurities especially electrolytes which can destabilize the sols. These impurities must be eliminated to make the colloidal solution stable.

The following methods are commonly used for the purification of colloidal solutions.
1. Dialysis: The method is based upon the fact that colloidal particles cannot pass through a parchment or cellophane membrane while the ions of the electrolyte can pass through it. The colloidal solution is taken in a bag made of cellophane or parchment.

The bag is suspended in fresh water. The impurities slowly diffuse out of the bag leaving behind pure colloidal solution. For example, dialysis can be used for removing HCl from the ferric hydroxide sol.

2. Electrodialysis: The ordinary process of dialysis is slow:

An apparatus for electrodialysis

To speed up the process of purification, the dialysis is carried out by applying electric field. This process is called electrodialysis.

3. Ultra-filtration: It is the process of removing the impurities from the colloidal solution by passing it through graded filter paper called ultrafilter papers. These filter papers are made from ordinary filter papers by impregnating them with colloidal solutions.

As a result, the size of the pores gets reduced. These filter papers allow the ions and molecules of the impurities to pass but retain colloidal particles. Ordinary filter papers cannot be used for this purpose since the colloidal particles also easily pass through the pores of these papers.

Properties Of Colloidal Solutions: The important properties of colloidal solutions are:
1. Heterogeneous nature: The colloidal solutions are heterogenous in nature consisting of dispersed phase and dispersion medium.

2. Visibility: The colloidals are not visible to naked eye and these can be seen with ultra microscopes.

3. Brownian movement: The colloidal particles have continuous zigzag motion called Brownian movement.

Brownian Movement

4. Tyndall effect: When the light is passed through the colloidal solution, the path of the light becomes visible when viewed from a direction at right angle to the incident beam. This phenomenon was studied by Tyndall and is known as Tyndall effect. The phenomenon of scattering oflight’by colloidal particles as a result of which the path of the beam becomes visible is called Tyndall effect.

5. Electrical properties: The particles of the colloidal solutions possess electrical charge, positive or negative. The presence of charge is responsible for the stability of these solutions. It may be noted that only the sol particles carry some charge while the dispersion medium has no charge. For example, the collodial solutions of gold, arsenious sulphide (AS₂S₃) are negatively charged while those of Fe (OH)₃ and Al (OH)₃ have positive charge. In the case of silver chloride sol, the particles may either be positively or negatively charged.

The presence of the charge on the sol particles and its nature whether positive of negative can be determined with the help of a phenomenon known as electrophoresis. In this experiment, the colloidal particles move towards positive or negative electrodes depending upon their charge under the influence of electrical field.

The phenomenon of movement of colloidal particles under an applied electric field is called electrophoresis.

If the particles accumulate near the negative electrode, the charge on the particles is positive. On the other hand, if the sol particles accumulate near the positive electrode the charge on the particles is negative.

A set up for electrophoresis

Origin of charge: The charge on the colloidal particles may be due to selective adsorption of ions. The particles contributing the dispersed phase adsorb only those ions preferentially which are common with their own lattice ions. For example, if silver nitrate solution is added to an aqueous solution of potassium iodide, the silver iodide will adsorb negative ions (I^–) from the dispersion medium to form a negatively charged sol.

However, if silver iodide is formed by adding potassium iodide to silver nitrate solution, the sol will be positively charged due to the adsorption of Ag⁺ ions present in the dispersion medium.

Coagulation Of Colloidal Solution: The phenomenon of precipitation of a colloidal solution by the addition of excess of an electrolyte is called coagulation or floculation.

→ Factors governing coagulation
1. Nature of the electrolytes: The coagulation capacity of different electrolytes depends upon the valency cf the active ion or called floculating ion. It is the ion carrying charge opposite to the charge on the colloidal particles. According to Hardy Schulz law, greater the valency of the active ion or floculating ion greater will be its coagulating power. Thus, to coagulate negative sol of As₂S₃, the coagulating power of different cations has been found to decrease in the order as:
Al³⁺ > Mg²⁺ > Na⁺

Similarly, to coagulate a positive sol such as Fe (OH)₃, the coagulating power of different anions has been found to decrease in the order:
[Fe(CN)₆]^4- > PO₄ ^3- > SO₄^2- > Cl^–

The minimum concentration of an electrolyte which is required to cause the coagulation or flocculation of a sol is known as flocculation value. It is usually expressed as milli moles per litre.

Protection of Colloids: The process of protecting the lyophobic colloidal solution from precipitation by the electrolytes due to the previous addition of some lyophilic colloid is called protection. The colloid which is added to achieve such a protection is called protecting colloid.

Gold Number: The different protecting colloids differ in their portecting powers. Zsigmondy introduced a term called gold number to describe the protective power of different colloids. This is defined as the minimum number of milligrams of the protective colloid required to just prevent the coagulation of a 10 ml of a given gold sol when 1 ml of a 10% solution of sodium chloride is added to it.

The coagulation of gold sol is indicated by change incolour from red to blue. The gold number of a few protective colloids are as follows:

It may be noted that smaller the value of gold number, greater will be protecting power of the protective colloid. Therefore, reciprocal of gold number is a measure of the protective power of a colloid. Thus, out of the list given above, gelatin is the best protective colloid.

Emulsions: Emulsions are the colloidal solutions of two immiscible liquids in which the liquid acts as the dispersed phase as well as the dispersion medium. Normally they are obtained by mixing an oil with water. Since the two do not mix well, the emulsion is generally unstable and is stabilised by adding a suitable outside reagent called emulsifier or emulsifying agent. The substances that are commonly employed for the purposes are gum, soap, glass powder, etc.

Types of Emulsions: These are of two types:
1. Oil-in-water emulsions: In this case, oils acts as the dispersed phase (small amount) and water as the dispersion medium (excess) e.g., milk is an emulsion of soluble fats in water and here casein acts as an emulsifier. Vanishing cream is another example of this class. Such emulsions are called aqueous emulsions.

2. Water-in-oil emulsions: In this case water acts as the dispersed phase while the oil behaves as the dispersion medium e.g., butter, cod liver oil, cold cream etc. Such types of emulsions are called oily emulsions.

Types of Emulsions

Demulsification: It is the process of decomposing an emulsion back into its constituent liquids. The demulsification can be done by centrifugation, filtration, boiling, freezing and some chemical methods.

Following are the interesting noteworthy examples of colloids which we come across in daily life.

1. Blue colour of the sky: Dust particles along with water suspended scatter blue light due to which sky looks blue to, us.
2. Fog, mist and rain: Clouds are aerosols. It is possible to cause artificial rain by throwing electrified sand.
3. Food articles like milk, butter, fruit juices are all colloids.
4. Blood-: Blood is a colloidal solution of an albuminoid substahce. Alum and FeCl₃ solution stop oozing blood due to coagulation. ,
5. Soils: Fertile soils are colloidal in nature in which humus . acts as a protective colloid.
6. Formation of delta: River water is a colloidal solution of clay.

Sea water contains several electrolytes. When river water meets sea water, the electrolytes present in sea water, coagulate the colloidal solution of clay resulting in its deposition with the formation of delta.

Applications of colloids: Colloids are widely used in the industry.
Some examples are:
1. Cottrell Smoke Precipitator: Smoke is a colloidal solution of solid particles such as carbon, arsenic compounds, dust etc. in air. The smoke is led through a chamber containing platgs having a charge opposite to that carried by smoke particles.” The particles on coming in contact with these plates lose their charge and get precipitated. The particles thus settle down on the floor of the chamber.

Cottrell Smoke Precipitator

2. Purification of drinking water: Alum is added to water to coagulate the suspended impurities present in water from natural resources to make it fit for drinking purposes.

3. Medicines: Most of the medicines are colloidal in nature. . Milk of magnesia, an emulsion, is used for stomach disorder. Argyrol is a silver sol used as an eye lotion. Colloidal antimony is used in curing Kalazar. Colloidal gold is used for intramuscular injection. Colloidal medicines are more effective because of large surface area and hence are easily assimilated.

4. Tanning: Animal hides due to colloidal nature bear positive charge when it is soaked in tanning or chromium salts which Contain negatively charged colloidal particles, mutual coagulation takes place leading to hardening of leather. The process is called tanning.

5. Cleansing action of soaps and detergents: Soaps and detergents act as emulsifiers and remove greasy impurities and dust of clothes which is washed away.

6. Photographic plates & films are prepared by coating an emulsion of light sensitive AgBr in gelatin over glass plates or celluloid films.

7. Rubber industry: Latex is a colloidal solution of rubber particles which are negatively charged. Rubber is ob tained by coagulation of latex.

8. Industrial products: Paints, inks, synthetic plastics and rubber, graphite lubricants, cement etc. are all colloidal solutions.

By going through these CBSE Class 12 Chemistry Notes Chapter 4 Chemical Kinetics, students can recall all the concepts quickly.

Chemical Kinetics Notes Class 12 Chemistry Chapter 4

Thermodynamics tells us about the feasibility of a reaction. A reaction is feasible if ΔG < 0 at constant temperature and pressure. Chemical equilibrium tells us the extent to which a reaction can proceed. But none of them tell us the speed of a reaction, i.e., time taken by a reaction to reach equilibrium.

The branch of chemistry which deals with the rates of reactions and their mechanisms is called Chemical Kinetics.
Rate Of A Chemical Reaction

From a kinetics point of view, reactions can be classified into 3 categories:
1. Very fast reactions: Which take place instantaneously e.g., ionic reactions like
AgNO₃ (aq) + NaCl (aq) → ↓AgCl (s) + NaNO₃ (aq)
BaCl₂ (aq) + H₂SO₄ (aq) → BaS04 (s) + 2 HCl (aq)
NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)

2. Very slow reactions: Which may take days or months for

Decomposition of NH₄NO₂

Decomposition of Hydrogen peroxide

Rate Of A Chemical Reaction: The rate of the reaction means the speed of the reaction.
It is defined as the change in the concentration of any one of the reactants or products per unit time.
Rate of the reaction = \(\frac{\text { Decrease in the concentration of a reactant }}{\text { Time taken }}\)
= \(\frac{\text { Increase in the concentration of a product }}{\text { Time taken }}\)

Consider a hypothetical reaction, R → P assuming the volume of the system is constant. One mole of reactant R produces one mole of the product P. If [R]₁ and [P]₁ are the concentrations of R and P at time t₁ and [R]₂ and [P]₂ are their concentrations at time t₂ then
ΔT = t₂ – t₁
Δ[R] = [R]₂ – [R]₁
Δ[P] = [P]₂ – [P]₁
The square brackets are used to express molar concentrations.
Rate of disappearance of R = – \(\frac{\Delta[\mathrm{R}]}{\Delta \mathrm{t}}\) …….(i)
Rate of appearance of P = + \(\frac{\Delta[\mathrm{P}]}{\Delta \mathrm{t}}\) ………(ii)

Equations (i) and (ii) given above represent the average rate of a. reaction, r_ave. r_ave depends upon (a) change in the cone, of reactants or products and (ii) time taken for that change to occur.

Units of Rate of a Reaction: Units of rate are concentration time-1, i.e., mol L^-1 S^-1. However for gaseous reactions, when the concentrations are expressed in terms of partial pressure, then the units of the rate will be atm s^-1.

For example,
PCl₅ → PCl₃ + Cl₂

Any general reaction
A + B → C + D

For this reaction, the rate of decomposition (disappearance) of

Rate of formation of O₂ = + \(\frac{\mathrm{d}\left[\mathrm{NO}_{2}\right]}{\mathrm{d} \mathrm{t}}\)

Rate law: The dependence of the concentration of reactants on the rate of reaction is given by the law. It gives the experimental dependence on concentration of reactants. For example, for the reaction
aA + bB → Products

From the kinetic study of the reaction, the dependence of the concentration of reactants on the rate of the reaction has been found to be

Rate = k[A]^m[B]ⁿ
where m and n are constant numbers or the powers of the concentration terms of the reactants A and B respectively on which the rate of reaction depends. It may be noted that the values of m and n are determined experimentally and may or may not be equal to a and b coefficients in the reaction. The above expression is the rate law. It may be defined as:

The mathematical expression denotes the observed or actual rate of a reaction in terms of the molar concentration of the reacting species which influences the rate of the reaction.

Order And Molecularity Of A Reaction:
Order of a Reaction: The dependence of reaction rate on concentration may be expressed in terms of the order of a reaction. The order of a reaction is defined as

the sum of the powers to which the concentration terms in the rate law are raised to express the observed rate of a reaction.
Thus, if the rate of a reaction,
xA + yB + zC → Products is given by the rate law,

Rate = –\(\frac{\mathrm{d} \mathrm{x}}{\mathrm{dt}}\) = k[A]x [B]y [C]z then, the order of the reaction, n is
n = x + y + z
where x, y and z are the order with respect to individual reactants and overall order of the reaction is sum of these exponents, i.e., x + y + z.

For example,
1. decomposition of ammonium nitrite is a first-order reaction.
NH₄NO₂ → N₂ + 2H₂O
Rate = k [NH₄NO₂] Order = 1

2. decomposition of hydrogen iodide is a second order reaction
2HI → H₂ + I₂
Rate = k[HI]₂ Order = 2

3. the reaction between nitric oxide and oxygen is a third-order reaction.
2NO (g) + O₂ (g) → 2NO₂ (g)
Rate = k[NO]²[O₂] Order = 3

The order of a reaction may be fractional or even zero.
For example,
CH₃CHO → CH₄ + CO
Rate = k[CH₃CHO]^3/2 Order = 3/2

Rate = k [NH₃]° Order = 0

Molecularity of Reaction: The molecularity of a reaction is defined as the number of reacting molecules that collide simultaneously to bring about a chemical reaction.

For example, the decomposition of H₂O₂ involves only one molecule, so it is a unimolecular reaction.
H₂O₂ → H₂O + \(\frac{1}{2}\)O₂

Dissociation of hydrogen iodide is a biomolecular reaction.
2HI (g) → H₂ (g) + I₂ (g)

The reaction involving three or more than three molecules is uncommon because the occurrence of such reactions would require simultaneous collisions of three or more than three molecules. The chances of the occurrence of such collisions are very small.

However, some of these reactions are found to be quite fast. This means that even though the balanced equation involves a large number of molecules, yet the reaction does not proceed by the simultaneous collision of all these reacting particles.

Such types of reactions take place through a sequence of two or more consecutive steps and are called complex reactions. The detailed description of various steps by which reactants change into the products is called the mechanism of the reaction. The steps which contribute to the overall reaction are called elementary processes.

Mechanism and Rate Law: In the case of multi-step reactions, some of the steps will be very fast while others will be slow. If one step takes place much more slowly than all other steps, it controls the overall reaction rate. This means that all the steps have to wait for the occurrence of this slowest step. But once this slowest step has occurred, the other steps will take place to form the products. In other words, the rate of the reaction is determined by the slowest step in the sequence.

Let us consider reaction between NO₂ and F₂ to form NO₂F;
2NO₂(g) + F₂(g) → 2NO₂F(g)

The experimental observations reveal that the rate of the reaction is proportional to the product of the concentrations of nitrogen peroxide and fluorine. This indicates that the rate-determining step in the mechanism of this reaction must be the reaction between NO₂ and F₂ only. Keeping this in mind, a mechanism of this reaction may be suggested as:

The rate of the overall reaction is determined by the first step which is the slower of the two steps. Accordingly, the experimentally observed rate of the reaction is given by the expression :
fix
Rate = – \(\frac{\mathrm{d} x}{\mathrm{dt}}\) = k[NO₂][F₂]

The above is the rate law for the reaction. It is, therefore, evident that complex reactions involving more than three molecules in the stoichiometric equation must take place in more than one step.
e.g.KClO₃ + 6 FeSO₄ + 3 H₂SO₄ → KCl + 3Fe₂ (SO₄)₃ + 3H₂O
should be, apparently, of the tenth order is actually a second-order reaction. This shows that the reaction occurs in several steps. The slowest step is called the rate-determining step.

Consider

The rate equation for this reaction is found to be
Rate = – \(\frac{\mathrm{d}\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]}{\mathrm{dt}}\) = k [H202][I-]

This reaction is of first order w.r.t. both H2Oz and I-. Evidences suggest that this reaction takes place in 2 steps :

1. H₂O₂ + I → H₂O + IO^– – Slow Step
2. H₂O₂ + IO^– → H₂O + I^– + O₂ – Fast Step

∴ The first step, being slow, is the rate determining step.

It can be concluded from above

1. The order of a reaction is an experimental quantity. It can be zero and even a fraction, but molecularity cannot be zero or a non-integer.
2. Order is applicable to elementary as well as complex reactions whereas molecularity is applicable only for elementary reactions. For complex reactions, molecularity has no meaning/significance.
3. For a complex reaction, the order is given by the slowest step and generally, the molecularity of the slowest step is the same as the order of the overall reaction.

Units of Reaction Rate Constants: The rate is the change in concentration with time. Therefore, the rate of reaction is expressed by concentration units divided by time. If the concentrations are expressed in mol litre^-1 and time in seconds, then units of rate constants for different orders are:

• Units for first-order reaction = sec^-1
• Units for second-order reaction = litre mol^-1 sec^-1
• Units for third-order reaction = litre² mol^-2 sec^-1

In the case of gaseous reactions, if concentrations are expressed in units of atm, then

• Units of rate constant for the first-order reaction = sec^-1
• Units of rate constant for second-order reaction = atm-1 sec^-1
• Units of rate constant for third-order reaction = atm^-2 sec^-1

Rate Constant For First Order Reaction:
For a first-order reaction,
A → Products
Let the initial cone, of [A] be an ML^-1.

Let x change into products so that equilibrium concentration after time t is


If the initial concentration is ‘a’ moles per liter, x moles of A change in time t and k is the rate constant, then the integrated rate equation is

where [A]₀ is the initial concentration and [A] is the concentration at time t. The value of k can be calculated by substituting the values of a, t, and x.

A plot between In (A) and t gives a straight line with a slope equal to – k.

Half-Life Period Of A Reaction: The half-life period of a reaction is defined as the time during which the concentration of a reactant is reduced to one-half of its initial concentration. It is generally denoted as t_1/2. The half-life period of a first-order reaction may be calculated as given below:

For the first-order reaction,
t = \(\frac{2.303}{\mathrm{k}}\) log \(\frac{[\mathrm{A}]_{0}}{[\mathrm{~A}]}\)

Now half-life period corresponds to the time during which the initial concentration, [A]₀ = a, is reduced to half i.e. [A] = a/2

The half-life period, t_1/2becomes

Thus, half life period of a first-order reaction is independent of the initial concentration of the reactant.

Similarly, the relation for the time required to reduce the concentration of the reactant to any fraction of the initial concentration can be calculated. For example,
∴ t_3/4 = \(\frac{2.303}{\mathrm{k}}\) log \(\frac{\mathrm{a}}{\mathrm{c} / 4}\) = \(\frac{2.303}{\mathrm{k}}\) log 4

Zero order Reactions are those reactions in which the rate of the reaction is proportional to zero power of the concentration of the reactants.
Rate = – \(\frac{\mathrm{d}[\mathrm{R}]}{\mathrm{dt}}\) = k[R]_o = k
or
R = \(\frac{[\mathrm{R}]_{0}-[\mathrm{R}]}{\mathrm{t}}\)

Collision Theory of Reaction Rate: The number of collisions that take place per second per unit volume of the reaction mixture is known as collision frequency (Z). All the collisions are not effective and do not give products. The collisions which actually produce the products and therefore, result in chemical reactions are called effective collisions.

There are two conditions for effective collisions.
1. Energy barrier: For the reacting species to make effective collisions, they should have sufficient energy to break the chemical bond in the reacting molecules. The minimum amount of energy that the colliding molecules must possess is known as threshold energy. Thus, only those collisions of reactants will give products that possess energies greater than threshold energy.

2. Orientation barrier: The colliding molecules should also have proper orientation so that the old bonds may break and new bonds are formed.

Thus, the collisions in which the colliding molecules do not possess the threshold energy of proper orientation do not form products. Therefore, only a small fraction of collision is effective.

Orientation of colliding NO₂ molecules is proper in (a) but ‘not’ in (b) for the reaction to take place.

Dependance Of Reaction Rate On Temperature: Temperature has a great influence on reaction rates. In general, the rate of a reaction becomes almost double for every 10° rise in temperature. The increase in reaction rate is not due to an increase in collision frequency but it is due to an increase in the fraction of effective collisions. It has been found that the fraction of effective collisions becomes almost double for a 10° rise in temperature.

Arrhenius Equation: The quantitative relationship between the rate constant and temperature was proposed by Arrhenius known as the Arrhenius equation.
k = A e^-E_a/RT

where k is a constant called frequency factor, E_a is the activation energy. If k₁ and k₂ are rate constants at two different temperatures T₁ and T₂ respectively then, integrated form of Arrhenius equation is
log\(\frac{\mathrm{k}_{2}}{\mathrm{k}_{1}}=\frac{\mathrm{E}_{\mathrm{a}}}{2.303 \mathrm{R}}\left[\frac{\mathrm{T}_{2}-\mathrm{T}_{1}}{\mathrm{~T}_{1} \mathrm{~T}_{2}}\right]\)

Activation Energy: The excess of energy (over and above the average energy of the reactants) required by the reactants to undergo chemical reactions is called activation energy. It is equal to the difference between threshold energy needed for the reaction and the average energy of reactant molecules.

Activation Energy = Threshold energy – Average energy of all reacting molecules.

When the molecules possess the energy equal to E , the atomic configuration of species formed at this stage is different from the reactants as well as the products. This stage is called the activated state or the transition state and specific configuration of the state is called Activated Complex. For example in the reaction between H₂ (g) and I₂ (g), activated complex has configuration in which H-H arid 1-1 bonds are breacking and H-I bonds are forming as shown below.

The change of reactants to products, i.e., progress of a reaction is shown below.

E_a = Activation energy for forward reaction
E_a‘ = Activation energy for backward reaction

Effect Of Catalyst On Reaction Rate:
The substances which increase the rate of a reaction and can be recovered chemically unchanged in mass and composition after the reaction are called catalysts. The phenomenon of increasing the rate of reactions by the use of catalyst is known as catalysis.

The function of a catalyst is that it provides a new path for the reaction, in which the reactants are converted into products quickly. It is believed that the catalyst forms a new activated complex of lower potential energy. This means that the activation energy becomes lower for the catalysed reaction than that for uncatalysed reaction.

Consequently, the fraction of the total number of collisions possessing lower activation energy is increased and hence, the rate of reaction also increases. This is shown below. The solid lines shows the path for uncatalysed reaction and dotted line shows the path adopted by catalysed reaction.

E_a = Activation energy without catalyst
E’_a (Cat) = Activation energy with catalyst

Effect Of Radiation On Reaction Rate:
Rate of reaction is increased by use of certain radiations. Such type of reactions which are initiated by the absorption of radiation are called photochemical reactions. For example, reaction of hydrogen and chlorine takes pla,ce very slowly in the absence of light. However, in the presence of light, the reaction occurs rapidly.

Reversible Reactions: At any stage, the net rate of the reaction is determined by the difference in the rates of forward and backward reactions. At equilibrium the overall rate of reaction in either direction becomes zero. Therefore,
Net rate of reaction = Rate of forward reaction – Rate of backward reaction At equilibrium, net rate is zero, so that

Rate of forward reaction = Rate of backward reaction
For the simple reaction A + B ⇌ C + D

If the rate constant for the forward reaction is Rf and that for the backward reaction is kf then.
Rate of forward reaction a [A] [B] = k [A] [B]
Similarly, rate of backward reaction a [C] [D] = kb [C] [D]
∴ At equilibrium rate of forward reaction = rate of backward reaction

where K = Equilibrium constant
Thus, K = \(\frac{[C] \times[D]}{[A] \times[B]}\)

By going through these CBSE Class 12 Chemistry Notes Chapter 3 Electrochemistry, students can recall all the concepts quickly.

Electrochemistry Notes Class 12 Chemistry Chapter 3

Electrochemical Cell: It is a device to convert the chemical energy of a spontaneous redox reaction into electrical energy. It is also called Galvanic Or Voltaic Cell. One such example of an electrochemical cell is Daniel Cell.

Daniel’s cell converts the chemical energy liberated during the redox reactions.
Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + Cu (s) into electrical energy and has an electrical potential [E.M.F. ] equal to 1.1 volt when concentration of both Zn²⁺ and Cu²⁺ ions is unity (1 mol dm^-3)

In the left-hand beaker, oxidation occurs. Zn plate dissolves to form Zn²⁺. Zn plate loses weight
Oxidation Half-reaction:
At anode: Zn (s) → Zn²⁺ (aq) + 2e^–

Reduction Half-Reaction: In the right-hand beaker reduction occurs. Cu²⁺ ions from the solution deposit on the copper plate. It gains weight.
At Cathode: Cu²⁺ (aq) + 2e^– → Cu (s)

Electrons flow from zinc plate to copper plate in the external circuit. Conventional current flows from copper to zinc plate (as. shown above) Combining the two half-reactions, we get the complete cell reaction or redox reaction.
Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + 2e^–

Electrolytic Cell: It is a device to convert electrical energy from an external source to produce a chemical change. In the above cell if E_ext is > 1.1. the reaction will again start but in the opposite reaction. It now will function as an electrolytic cell.

(a) when Eext < 1.1 V

1. Electrons flow from Zn rod to Cu rod hence current flows from Cu to Zn.
2. Zn dissolves at anode and copper deposits at the cathode.

(b) when E_ext = 1.1 V

1. No flow of electrons or current.
2. No chemical reaction.

Functioning of Daniell cell when external voltage E_ext opposing the cell potential is applied.
(c) When E_ext > 1.1 V

1. Electrons flow from Cu to Zn arid current flows from Zn to Cu.
2. Zinc is deposited at the zinc electrode and copper dissolves at the copper electrode

In the electrochemical cell at each electrode-electrolyte interface, there is a tendency of metal ions from the solution to deposit on the metal electrode trying to make it positively charged. At the same time, metal atoms of the electrode have a tendency to go into the solution as ions and leave behind the electrons at the electrode trying to make it negatively charged.

At equilibrium, there is a separation of charges and depending on the tendencies of the two opposing reactions, the electrode may be positively or negatively charged with respect to the solution. A potential difference develops between the electrode and the electrolyte which is called Electrode Potential.

When the concentration of all species involved in half-cells is unity, the electrode potential is called Standard Electrode Potential.

According to the IUPAC convention, standard reduction potentials are now called standard electrode potentials. In a galvanic cell, the half¬cell in which oxidation takes place is called Anode and it has a negative potential with respect to the solution. The other half-cell in which reduction takes place is called Cathode and it has a positive potential with respect to the solution. The direction of the flow of current is opposite to that of the flow of electrons.

The cell potential is the difference between the electrode potentials (reduction potentials) of the cathode and anode. It is called electromotive force (emf) of the cell when no current flows through the circuit. Internally the two half-cells/beakers are connected through the salt bridge.
E_cell = E_right – E_left

In the Cu^– AgNO₃ cell cell reaction is
Cu (s) + 2 Ag⁺ (aq) → Cu²⁺ (aq) + 2 Ag (s)
At anode: Cu (s) → Cu²⁺ (aq) + 2 e^–] oxidation
At cathode: 2 Ag⁺ (aq) + 2 e^– → 2Ag (s)] reduction

Here Ag electrode acts as cathode and copper electrode, as anode.
The cell can be represented by
Cu (s) | Cu²⁺ (aq) || Ag⁺ (aq) | Ag (s)
E_cell = E_right – E_left
= E_Ag⁺/Ag – E_Cu²⁺/Cu

Similarly Daxtiel cell can be represented by
Zn (s) | Zn²⁺ (aq) || Cu²⁺ (aq)| Cu (s)
E_cell = E_right – Eleft
= E_Cu²⁺/Cu – E_Zn²⁺/Zn
Measurement of Electrode Potential: The measurement of the potential of half-cells is not possible.

Only the difference between the two half-cell potentials can be determined that gives the e.m.f. of the cell. By arbitrarily fixing the electrode potential of one half-cell, that of the other can be determined.

According to the convention, Standard Hydrogen Electrode represented by Pt (s), H₂(g), H₂(g)/H⁺ (aq) is assigned a zero potential at all temperatures according to the reaction
H⁺(aq) + e^– → H₂(g)

Standard Hydrogen Electrode: It consists of platinum wire sealed in a glass tube arid has a platinum foil attached to it. The foil is coated with finely divided platinum and acts as a platinum electrode. It is dipped into an acid solution containing H⁺ ions in 1 M concentration (1M HCl). Pure H₂ gas at 1 bar pressure is constantly bubbled into the solution at a constant temperature of 298 K. The following reaction occurs in the half-cell depending upon whether it acts as an anode or as a cathode.

Normal Hydrogen Electrode (NHE)

If S.H.E. (orN.H.E.) acts as Anode
H₂(g) → 2H⁺ (aq) + 2e^– oxidation half reaction

If S.H.E. (or N..H.E.) acts as Cathode
2H⁺ (aq) + 2e^– → H₂ (g)] reduction half reaction

The standard hydrogen electrode is also regarded as a reversible electrode with respect to H+ ions:
H₂(g) ⇌ 2H⁺ (aq) + 2e^–

Arbitrarily, the standard electrode potential of this electrode is fixed to be 0.000 V.
The electrode potential of an electrode can be determined by connecting this half-cell with a standard hydrogen electrode.

The electrode potential of a metal electrode as determined with respect to S.H.E. (or N.H.E.) is called Standard Electrode Potential E°.

Conventionally: The reduction (standard) potential of an electrode that acts as a Cathode when attached to S.H.E. is given a positive sign, e.g., Cu/Cu⁺⁺ electrode when attached to S.H.E. acts as a cathode.

Let us calculate the electrode (reduction) potential of Zn/Zn²⁺. The cell will be
Zn (s) | Zn²⁺_1.0 (aq)| | H⁺ _1.0 (aq), H₂ (g), Pt(s).

Let us measure the electrode potential of Cu/Cu²⁺ electrode concentration of Cu²⁺ is 1.0 M and the pressure of H₂ gas is one bar. At 298 K, the emf of the cell

Standard hydrogen electrode 11 second half cell.
E°_cell = E°_R – E°_L
0.34 = E°_Cu⁺⁺ – 0 [emf of the cell = 0.34 V]
∴ E_Red° of Cu electrode dipping in Cu²⁺ ions = 0.34 V

Electrochemical Series or The Standard Electrode Potentials at 298 K: Ions are present as aqueous species and H₂O as liquid, gases and solids are shown by aq, g and s.

1. A negative Ev means that the redox couple is a stronger reducing agent than the H⁺/ H₂ couple.
2. A positive Ev means that the redox couple is a weaker reducing agent than the H⁺/H₂ couple.

Applications Of Electrochemical Series:
1. Relative strength of oxidizing and reducing agents: In this series, metals are arranged in the decreasing order of their standard electrode potentials or decreasing order of their oxidizing character. e.g. Li⁺ (aq) which has least reduction potential (- 3.05) is the weakest oxidizing agent and F₂ (g) which has maximum value of p_Red = + 2.85 is the strongest oxidizing agent.

On the other hand, Li metal which has the highest standard oxidation potential (= + 3.05) is the strongest reducing agent while F- ( E°_ox = – 2.85 V) is the least reducing agent.

2. Calculation of e.m.f. of the cell: The e.m.f. of the desired cell can be calculated knowing the standard reducing potentials of the two half cells constituting the cell from the electrochemical series.
E°_cell = E°_Red [RIGHT] – E°Red [LEFT]
e.g. Zn (s) | Zn²⁺ (1.0 M) || Cu²⁺ (1.0 M) | Cu (s).
E°_cell = E°_Cu²⁺,_Cu – E°_Zn²⁺,_Zn.
= 0.34 – (- 0.76) V = 1.10 V

3. Predicting feasibility of redox reaction: In general, a redox reaction is feasible only if the species which has higher standard reduction potential is reduced i.e., accepts the electrons and the species which has lower reduction potential is oxidized i.e., loses the electrons. Otherwise, a redox reaction is not feasible. In other words, the species to release the electrons must have lower reduction potentials as compared to the species which is to accept electrons. In a nutshell, if the e.m.f. of the hypothetical cell is +ve, a redox reaction takes place; if it is – ve, a redox reaction is not feasible.

4. To predict whether a metal can librate H₂(g) from acid or not. The metals which have only negative reduction potentials, i.e., are lying above N.H.E. in the electrochemical series, can only liberate H₂ (g) from dilute acid solutions.

Nernst Equation: The relationship between electrode potentials and the concentration of the electrolytic solutions is called Nernst Equation.
1. The Nernst Equation is
E = E° + \(\frac{\mathrm{RT}}{\mathrm{nF}}\) In [Mn+]/[M] Since[M] = 1
E = E° + \(\frac{2.303 \mathrm{RT}}{\mathrm{nF}}\) log [Mn+]
E = reduction electrode potential
E° = reduction electrode potential in standard state [1 M solution of metal ions at 298 K]
R = Gas constant = 8.314 JK-1.M0l-1
T = Temperature
n = no. of electrons accepted during the change
| F | = Faraday t = 96500 C.

Putting the values of R, T = 298 K; F
| E | = E° + \(\frac{0.059}{\mathrm{n}}\) log [Mn+]

For the complete cell reaction
Zn (s) + Cu²⁺(aq) ⇌ Zn²⁺ (aq) + Cu (s)

2. Calculation of Gibbs free energy/Maximum work that can be obtained from the Galvanic cell
ΛG° = – n F E°elI = – 2.303 RT log Kc
where AG° is the standard Gibbs energy change in Gibbs function
n = number of moles of electrons involved in the cell reaction
F = Faraday = 96,500 coulombs
R = Gas Constant
Kc = Equilibrium constant.

Conductance of Electrolytic Solutions:
Resistance R (ohm) = \(\frac{\text { Potential Difference in volts }}{\text { Current strength in amperes }}\)

Resistance can be measured by the principle of the Wheatstone bridge. It depends upon the length and inversely on the area of cross-section A of the wire.
R ∝ l
R ∝ \(\frac{l}{\mathrm{~A}}\)
or
R = ρ\(\frac{l}{\mathrm{~A}}\);
where ρ (rho) is called specific resistance/Resistivity.

The resistivity of a substance is its resistance to the passage of electricity when it is one metre long and its area of cross-section is one m2.

The inverse of resistance R is called conductance.
Conductance = \(\frac{1}{\mathrm{R}}=\frac{\mathrm{A}}{\rho l}=\frac{\mathrm{A} \kappa}{l}\) ( κ = Kappa)

The SI unit of conductance is Siemens.
1S = 1 ohm^-1 = mho = Ω^-1.

The inverse of resistivity is called Conductivity or Specific Conductance.
SI unit of conductivity is Sm-1, but often (Greek Kappa) is expressed in S cm^-1.

The magnitude of conductivity varies a great deal and depends upon the nature of the material. It also depends upon the temperature and pressure of measurement.

Materials are classified as

1. Conductors,
2. Insulators,
3. Semi-conductors depending upon the magnitude of their conductivity.

Metals and their alloys have very large conductivity and are known as conductors. Certain, non-metals like carbon black, graphite and some organic polymers are also electronically conductors. Substance like glass, ceramics etc. have very low conductivity are known as insulators. Substance like silicon, doped silicon, gallium arsenide is semi-conductors. Electrical conductance ( through metals called metallic or electronic conductance is due to the flow of mobile electrons.

The metallic conductance depends upon,

1. The nature and structure of the metal.
2. The number of valence electrons per atom.
3. Temperature. It decreases with an increase in temperature.
4. The composition of the metallic conductor remains unchanged.

Electrolytic or Ionic Conductance: When electrolytes are dissolved in water, they furnish their own ions in the solution. They conduct electricity through their ions in the solution and is called electrolytic or ionic conductance.

The conductivity of electrolytic or ionic solutions depends upon:

1. The nature of the electrolyte added
2. Size of the ions produced and their solvation.
3. The nature of the solvent and its viscosity.
4. The concentration of the electrolyte.
5. Temperature conductivity increases with an increase in temperature.
6. It leads to a change in its composition with the passage of direct current through the solution over a prolonged period.

Measurement of Conductivity of Ionic Solutions: It is based upon the measurement of resistance by a wheat stone bridge. The cell is called a conductivity cell.

Unknown resistance of the cell R2 = \(\frac{\mathrm{R}_{1} \mathrm{R}_{4}}{\mathrm{R}_{3}}\)
Conductance C of the cell = \(\frac{1}{\mathrm{R}_{2}}\)

Cell Constant = \(\frac{\text { Specific conductance }}{\text { Observated conductance }}\)

Molar Conductivity: It is the conducting power of all ions obtained by dissolving 1 mole of an electrolyte in a given volume of the solution,

Molar conductivity = Λm = \(\frac{x}{c}\)
= \(\frac{\mathrm{Sm}^{-1}}{1000 \mathrm{Lm}^{-3} \times \text { molarity }\left(\mathrm{mL}^{-1}\right)}\)
= S m² mol^-1

Variation of Conductivity and Molar Conductivity with concentration: Both conductivity and molar conductivity change with the concentration of the electrolyte. Conductivity always decreases with the decrease in concentration both for weak and strong electrolytes. This can be explained by the fact that the number of ions per unit volume that carry the current in a solution decreases on dilution. The conductivity of a solution at any given concentration is the conductance of one unit volume of solution kept between two platinum electrodes with a unit area of cross-section and at a distance of unit length. This is clear from the equation:

C = \(\frac{\mathrm{κA}}{l}\)= κ (both A and l are unity in their appropriate Units in m or cm)

Molar conductivity of a solution at a given concentration is the conductance of volume V of a solution containing one mole of electrolyte kept between two electrodes with an area of cross-section A and distance of unit length. Therefore,
Λ_m = \(\frac{\mathrm{κA}}{l}\)= κ
Since l = 1 and A = V (volume containing 1 mole of electrolyte)
Λ_m = κV

Molar conductivity increases with a decrease in concentration. This is because the total volume, V, of a solution containing one mole of electrolyte also increases. It has been found that a decrease in κ on dilution of a solution is more than compensated by an increase in its volume.

Physically, it means that at a given concentration, Λm can be defined as the conductance of the solution of an electrolyte kept between the electrodes of a conductivity cell at a unit distance but having an area of cross-section large enough to accommodate sufficient volume of solution that contains one mole of the electrolyte. When concentration approaches zero, the molar conductivity is known as limiting molar conductivity and is represented by the symbol Λ°_m The variation in Λm with concentration is different for strong and weak electrolytes.

Strong Electrolytes: For strong electrolytes, A increases slowly with dilution and can be represented by the equation :
Λ_m = Λ°_m – A c^1/2
It can be seen that if we plot Λm against c^1/2, we obtain a straight line with an intercept equal to and slope equal to ‘A’. The value of the constant ‘A’ for a given solvent and temperature depends on the type of electrolyte i.e., the charges on the cation and anion produced on the dissociation of the electrolyte in the solution. Thus, NaCl, CaCl₂, MgS04 are known as 1-1, 2-1 and 2-2 electrolytes respectively. All electrolytes Of a particular type have the same value for ‘A’.

Molar conductivity versus c^1/2 for acetic acid (weak electrolyte) and potassium chloride (strong electrolyte) in aqueous solutions.

Kohlrausch law of independent migration of ions. The law states that limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte. Thus, if λ°_Na and λ°_Cl^– are limiting molar conductivity of the sodium and chloride ions respectively, then the limiting molar conductivity for sodium chloride is given by the equation :
Λ°_m(NaCl) = λ°_Na + λ°_Cl^–

In general: Λ°_m = v₊ λ°₊ + n_– λ°_–
Here λ°₊ and λ°_– are the limiting molar conductivities of the cation and anion respectively.

Weak Electrolytes: Weak electrolytes like acetic acid have a lower degree of dissociation at higher concentration and hence for such electrolytes, the change in Λ_m with dilution is due to an increase in the degree of dissociation and consequently the number of ions in the total volume of solution that contains 1 mol of electrolyte. In such cases, Λ_m increases steeply on dilution, especially near lower concentrations.

Therefore, Λ°m cannot be obtained by extrapolation of Λm to zero concentration. At infinite dilution (i.e., concentration c → zero) electrolyte 5 dissociates completely (α = 1), but at such low concentration, the conductivity of the solution is so low that it cannot be measured accurately. Therefore, Λ°_m for weak electrolytes is obtained by using Kohlrausch law of independent migration of ions. At any concentration c, if a is the degree of dissociation then it can be approximated to the ratio of molar conductivity A m at the concentration c to limiting molar conductivity, Λ°_m.

Thus we have:
α = \(\frac{\Lambda_{\mathrm{m}}}{\Lambda_{\mathrm{m}}^{\circ}}\)
But we know that for a weak electrolyte like acetic acid,

Applications of Kohlrausch law: Using Kohlrausch law of independent migration of ions, it is possible to calculate Λ°_m for any electrolyte from the λ° of individual ions. Moreover, for weak electrolytes like acetic acid, it is possible to determine the value of its dissociation constant once we know the Λ°_m and Λ_m at a given concentration c

Electrolytic Cells and Electrolysis: In an electrolytic cell external source of voltage is used to bring about a chemical reaction. The electrochemical processes are of great importance in the laboratory and the chemical industry. One of the simplest electrolytic cells consists of two copper strips dipping in an aqueous solution of copper sulphate. If a DC voltage is applied to the two electrodes, then Cu²⁺ ions discharge at the cathode (negatively charged) and the following reaction takes place:
Cu²⁺ (aq) + 2e^– → Cu (s)
Copper metal is deposited on the cathode. At the anode copper is converted into Cu²⁺ ions by the reaction:
Cu (s) → Cu²⁺ (s) + 2e^–

Thus copper is dissolved (oxidised) at the anode and deposited (reduced) at the cathode. This is the basis for an industrial process in which impure copper is converted into the copper of high purity. The impure copper is made an anode that dissolves on passing current and pure copper is deposited at the cathode.

Many metals like Na, Mg, Al, etc. are produced on large scale by electrochemical reduction of their respective cations where no suitable chemical reducing agents are available for this purpose. Sodium and magnesium metals are produced by the electrolysis of their fused chlorides and aluminium is produced by electrolysis of aluminium oxide in presence of cryolite.

Faraday’s Laws of Electrolysis:
1. First Law: The amount of chemical reaction which occurs at any electrode during electrolysis by a current is proportional to the quantity of electricity passed through the electrolyte (solution or melt).

2. Second Law: The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their chemical equivalent weights (Atomic Mass of Metal. Number of electrons required to reduce the cation).

Some Commercial Cells: The electrochemical cells can be used to generate electricity and these are called batteries. The battery is generally used for two or more galvanic cells connected in series.

There are two types of commercial cells:

1. Primary cells, in which electrode reactions cannot be reversed by external energy source. Therefore, these are not chargeable. For example, dry cell, mercury cell.
2. Secondary cells are those which can be recharged. For example, lead storage cell, nickel-cadmium cell.

→ Fuel Cells: These are voltaic cells in which the reactants are continuously supplied to the electrodes. These are designed to convert the energy from the combustion of fuels such as H₂, CO, CH₄ etc. directly into electrical energy. A common example is a hydrogen-oxygen fuel cell.

→ Corrosion: The process of deterioration of a metal as a result of its reaction with air or water surrounding it is called corrosion. In the case of iron, corrosion is called rusting. Chemically rust is a hydrated form of ferric oxide, Fe₂O₃ x H₂O. It is caused by moisture, carbon dioxide and oxygen present in the air.

The chemistry of corrosion is essentially an electrochemical phenomenon. At a particular spot of an iron object, oxidation takes place and that place behaves as an anode:
Anode 2 Fe(s) → 2Fe²⁺ + 4e^– E°_Fe²⁺/Fe = – 0.44 V

Electrons released at anodic spot move through the metal and go to another spot on the metal and reduce oxygen in presence of H⁺ (which is believed to be .available from H2C03 formed due to dissolution of carbon dioxide from the air into water. Hydrogen ion in water may also be available due to the dissolution of other acidic oxides from the atmosphere).

This spot behaves as cathode with the reaction Cathode:
O₂(g) + 4H⁺ (aq) + 4e^– → 2 H₂O
E_H⁺|O2|H2O = 1.23V

The overall reaction being:
2Fe(s) + O₂(g) + 4H⁺ (aq) → 2Fe₂ (aq) + 2 H₂O (l)
E^V_cell = 1.67 V

The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as rust in the form of hydrated ferric oxide (Fe₂O₂ . x H₂0) and with further production of hydrogen ions.

Corrosion of iron in the atmosphere

Oxidation: Fe (s) → Fe²⁺ (aq) + 2e^–
Reduction: O₂(g) + 4H⁺ (aq) + 4e^– → 2H₂O (1)

Atmospheric
oxidation : 2 Fe²⁺ (aq) + 2H₂O (l) + 1/2 O₂ (g) → Fe₂O₃ (s) + 4H⁺ (aq)

Prevention of Corrosion: The rusting of iron can be prevented or decreased by the following methods:
1. Barrier protection: In this method a barrier is placed between the iron and the atmospheric air by coating the surface with paint, applying a thin film of oil or grease or electroplating.

2. Sacrificial protection: In this method, iron is protected by covering it with a layer of metal more active than iron such as zinc, tin, etc. The process of covering iron with zinc is called galvanization.

3. Electrical protection: In this method, the iron is connected with a more active metal like magnesium or zinc.

4. Using anti-rust solution: To retard the corrosion of iron certain anti-rust solutions such as alkaline phosphates and alkaline chromates are used.

Commercial Production of Chemicals: Many metals and their compounds are prepared by using basic principles of electrolysis:

Na is prepared by the electrolysis of molten Na⁺ Cl^– by Down’s Cell

Mg is prepared from fured Mg Cl₂
MgCl₂(Z) ⇌ Mg²⁺ (l) + 2 Cl^– (l)
At anode 2Cl^– → Cl₂ (g) + 2e^–
At cathode Mg²⁺ + 2e^– → Mg (s).

The Hydrogen Economy: Both the production of hydrogen by electrolysis of water
2 H₂O (l) → 2H₂ (g) + O₂ (g) and hydrogen combustion in a fuel cell
2 H₂ (g) + O₂ (g) 2H₂O (l). will be important in future. Both are based on electrochemical principles.