Class 12 Chemistry Study Material: Interactive Q&A Guide

12th Chemistry Study Material

Interactive Question & Answer Guide

1. Write the differences between minerals and ores?

Minerals	Ores
Naturally occurring substances obtained by mining which contain the metals.	Minerals that contain a high percentage of metal from which it can be extracted conveniently and economically.
Mineral of Al is Bauxite and Clay.	Ore of Al is Bauxite.
All the minerals are not ores.	All the ores are minerals.

2. Steps involved in the extraction of metals

1. Concentration of ores
2. Extraction of crude metal
3. Refining of crude metal

3. Write a note about the Gravity separation method.

• Used for heavy oxide ores.
• Example: Haematite (Fe₂O₃), Tinstone.
• Principle: High specific gravity ore is separated from low specific gravity impurities.
• In this method, powdered ore is washed by running water, and lighter impurities are washed away.

4. Which type of ores can be concentrated by froth flotation methods? Give two examples.

Sulphide ores are concentrated by froth flotation.

Examples: Galena (PbS), Zinc blende (ZnS).

5. Explain the following terms with suitable examples: (i) Gangue (ii) Slag (iii) Flux

(i) Gangue: The nonmetallic impurities and rocky materials associated with the ore. Example: SiO₂

(ii) Flux: Flux is a substance that forms fusible slag with Gangue.

(iii) Slag: Fusible product formed by the reaction between flux and gangue.
Flux + Gangue → Slag
CaO + SiO₂ → CaSiO₃

6. What is the role of quick lime in the extraction of Iron?

Quick lime (CaO) acts as a basic flux to remove acidic gangue (SiO₂) as slag.

CaO + SiO₂ → CaSiO₃ (Slag)

7. What is Calcination?

Heating the ore strongly in the absence of air to remove volatile impurities. It is generally used for carbonate ores.

CaCO₃ → CaO + CO₂↑

8. What is Auto reduction?

A process where the ore itself reduces to crude metal upon roasting without adding any separate reducing agent.

HgS + O₂ → Hg + SO₂

Mercury is obtained by roasting its ore cinnabar.

9. What is the role of silica in the extraction of copper?

Silica acts as an acidic flux to remove iron oxide gangue as iron silicate slag.

FeO + SiO₂ → FeSiO₃

10. Explain Aluminothermic process

Chromic oxide is mixed with aluminium powder and heated. Chromic oxide is reduced to chromium by aluminium.

Cr₂O₃ + 2Al → 2Cr + Al₂O₃

Another example: BaO₃ + Mg → BaO + MgO

11. Explain zone refining process with an example

• Principle: Fractional crystallisation (impurities are more soluble in the molten state than in the solid state).
• The impure metal is taken in the form of a rod.
• The metal rod is heated at one end with a mobile induction heater, melting that zone.
• As the heater moves, pure metal crystallises out while the impurities dissolve and shift into the adjacent molten zone.
• Example: Si, Ga, Ge (semiconductor materials).

12. Give the limitations of Ellingham diagram.

• It does not explain the rate of the reaction (kinetics).
• It does not give any idea about the possibility of other secondary reactions taking place.
• ΔG is assumed at equilibrium condition, which is not always true.

13. Give the uses of zinc.

• Galvanizing iron to protect it from rusting.
• Zinc sulphide is used to produce luminous paints, fluorescent lights, and X-ray screens.
• Zinc oxide is used to prepare paints, rubber, plastics, and ink.

14. Give the basic requirements for vapour phase refining. Give examples.

• The metal should form a volatile compound with a suitable reagent.
• The volatile compound should easily decompose to give back the pure metal.
• Examples: Mond process (for Nickel), Van Arkel method (for Titanium/Zirconium).

15. Describe a method for refining nickel.

Mond process: Impure nickel is heated with carbon monoxide to form volatile nickel tetracarbonyl, which is then decomposed at higher temperatures to yield pure nickel.

Ni + 4CO → [Ni(CO)₄] (at 350K)
[Ni(CO)₄] → Ni + 4CO (at 460K)

16. Van Arkel method (Titanium refining)

Impure titanium is heated with iodine to form volatile titanium tetraiodide, which is decomposed on a hot filament to get pure titanium.

Ti + 2I₂ → TiI₄ (at 550K)
TiI₄ → Ti + 2I₂ (at 1800K)

17. Describe the role of the following in the process mentioned.

(i) Silica in the extraction of copper: Acts as an acidic flux to remove basic FeO impurity as slag.

(ii) Cryolite in the extraction of aluminium: Lowers the melting point of the alumina mixture and increases electrical conductivity.

(iii) Sodium cyanide in froth flotation: Acts as a depressing agent to selectively separate ores (e.g., separating ZnS from PbS).

(iv) Iodine in refining of zirconium: Forms the volatile Zirconium tetraiodide compound to separate it from impurities.

18. Explain the Froth Floatation process.

• Used to concentrate sulphide ores (e.g., PbS, ZnS).
• Frothing agent: Pine oil; Collector: Sodium Ethyl Xanthate; Depressing agent: Sodium cyanide.
• The powdered ore is mixed with water and pine oil.
• When air is blasted through the mixture, it produces froth.
• The ore particles preferentialy wet by oil rise to the surface with the froth and are collected, while impurities settle at the bottom.

19. Explain the principle of electrolytic refining with an example (Refining of Silver).

• Anode: Impure metal (e.g., Impure Silver)
• Cathode: Thin strip of pure metal (e.g., Pure Silver)
• Electrolyte: Acidified aqueous solution of metal salt (e.g., Silver nitrate solution)
• During electrolysis, pure metal deposits at the cathode:

At anode: Ag → Ag⁺ + e^-
At cathode: Ag⁺ + e^- → Ag

20. Extraction of aluminium - Hall-Heroult method

• Anode (+): Carbon rods
• Cathode (-): Carbon lining inside the iron tank
• Electrolyte: Molten mixture of 20% Alumina (Al₂O₃), cryolite, and calcium fluoride maintained above 1270 K.

At Cathode (Reduction): 4Al⁺³ + 12e^- → 4Al
At Anode (Oxidation): 6O^-2 → 3O₂ + 12e^-
Overall Reaction: 4Al⁺³ + 6O^-2 + 3C → 4Al + 3CO₂

21. Explain the magnetic separation process.

• Used for separating ferromagnetic components from non-magnetic matter.
• Based on differences in the magnetic properties of the ore and impurities.
• Example: Tinstone (non-magnetic) can be separated from wolframite (magnetic impurity).
• The crushed ore is poured onto a moving belt over an electromagnetic roller; magnetic particles fall closer to the roller while non-magnetic particles fall farther away.

22. Explain Acid leaching with an example.

Sulphide ores (e.g., ZnS, PbS) are treated with hot aqueous sulphuric acid in the presence of oxygen, converting insoluble sulphide into soluble sulphate and elemental sulphur.

2ZnS + 2H₂SO₄ + O₂ → 2ZnSO₄ + 2S + 2H₂O

23. Explain how gold ore is leached by the cyanide process.

Crushed gold ore is leached with an aerated dilute solution of sodium cyanide to form a soluble complex:

4Au + 8CN^- + 2H₂O + O₂ → 4[Au(CN)₂]^- + 4OH^-

The gold is later recovered by displacement with zinc (cementation):

2[Au(CN)₂]^- + Zn → [Zn(CN)₄]^-2 + 2Au

24. Write the applications of the Ellingham diagram.

1. To predict the thermodynamic feasibility of thermal reduction of an ore.
2. To select a suitable reducing agent for a particular metal oxide.
3. To select an appropriate temperature range for the reduction process.

25. Write a short note on electrochemical principles of metallurgy.

The thermodynamic relationship is given by:

ΔG° = -nFE°

If E° is positive, then ΔG is negative and the reduction is spontaneous. A more reactive metal can displace a less reactive metal from its salt solution because the net EMF of the combined redox reaction remains positive.

Example: Zn + Cu⁺² → Cu + Zn⁺²

1. Give examples for each of the following: Icosogens, Tetragens, Pnictogens, Chalcogens.

• Icosogens: Boron (B) group
• Tetragens: Carbon (C) group
• Pnictogens: Nitrogen (N) group
• Chalcogens: Oxygen (O) group

2. Write a note on metallic nature of P block elements.

The tendency of an element to form a cation by losing electrons is its metallic character. Down a group, the metallic character increases due to a decrease in ionization energy.

3. Write a short note on anomalous properties of the first element of p-block.

The first element of each group in the p-block differs significantly from the other elements due to:

1. Small size of the atom/ion.
2. Absence of d-orbitals in their valence shell.
3. High ionization energy and high electronegativity.

4. What is Inert Pair Effect?

In heavier post-transition p-block elements (groups 13, 14, 15, and 16), the outer ns² electrons show a reluctance to participate in bonding due to poor shielding of d and f electrons. This preference of the inner s-pair to remain inert is called the inert pair effect.

5. Describe briefly allotropism in p-block elements with specific reference to carbon.

Allotropism is the phenomenon where an element exists in more than one crystalline or physical form in the same state.

Allotropes of carbon: Diamond, Graphite, Graphene, Fullerenes, and Carbon nanotubes.

6. Write uses of Boron.

• B-10 isotope absorbs neutrons and is used as a moderator/control rod in nuclear reactors.
• Amorphous boron is used as a rocket fuel igniter.
• Boron is an essential micronutrient for plant cell walls.
• Boric acid and borax are used in eye drops and antiseptics.

7. Write Preparation of Borax from Colemanite Ore.

Colemanite ore is boiled with an aqueous solution of sodium carbonate:

2Ca₂B₆O₁₁ + 3Na₂CO₃ + H₂O → 3Na₂B₄O₇ + 3CaCO₃↓ + Ca(OH)₂

8. Borax Bead Test or Action of heat on borax

On heating, borax loses its water of crystallization, swells up, and turns into a transparent glassy bead containing sodium metaborate and boric anhydride.

Na₂B₄O₇·10H₂O → Na₂B₄O₇ + 10H₂O
Na₂B₄O₇ → 2NaBO₂ + B₂O₃

Transition metal metaborates show characteristic colours. Example: Co(BO₂)₂ gives a blue colour bead.

9. Write Uses of Borax.

• Identification of coloured metal ions via the borax bead test.
• Manufacture of optical and borosilicate glasses.
• Used as a flux in metallurgy.
• Acts as a food/analytical preservative.

10. Write Action of heat on boric acid.

4H₃BO₃ → 4HBO₂ + 4H₂O (at 373 K / Metaboric acid)
4HBO₂ → H₂B₄O₇ + H₂O (at 413 K / Tetraboric acid)
H₂B₄O₇ → 2B₂O₃ + H₂O (at Red Hot / Boric oxide)

11. How will you convert boric acid to Boron nitride?

Boric acid heated with ammonia gives boron nitride:

B(OH)₃ + NH₃ → BN + 3H₂O

12. How will you identify borate radical? or Write Ethyl Borate Test

Boric acid or borate salt reacts with ethanol in the presence of concentrated H₂SO₄ to form triethyl borate gas which burns with a characteristic green-edged flame.

B(OH)₃ + 3C₂H₅OH → B(OC₂H₅)₃ + 3H₂O

13. Write Uses of Boric Acid

• Manufacture of pottery glass, enamels, and pigments.
• Used as an antiseptic and mild eye lotion.
• Used as a food preservative.

14. What is inorganic benzene? How is it prepared?

Borazine (B₃N₃H₆) is called inorganic benzene because its structure is isoelectronic and isostructural with benzene. It is prepared by reacting diborane with ammonia:

3B₂H₆ + 6NH₃ → 2B₃N₃H₆ + 12H₂

15. Write a note on hydroboration.

Diborane adds across alkenes in ether solvents at room temperature to form trialkylboranes.

6CH₃-CH=CH₂ + B₂H₆ → 2(CH₃-CH₂-CH₂)₃B

Further oxidation yields alcohols:

(CH₃-CH₂-CH₂)₃B + 3H₂O₂ → 3CH₃-CH₂-CH₂-OH + B(OH)₃

16. Describe the structure of Diborane

• Contains two BH₂ units linked by two bridging hydrogen atoms.
• It has 8 B-H bonds but only 12 valence electrons (electron-deficient compound).
• Four terminal B-H bonds are normal 2-centre-2-electron (2c-2e) bonds.
• The two bridging B-H-B bonds are special 3-centre-2-electron (3c-2e) banana bonds.
• Both boron atoms are sp³ hybridized.

17. Write Uses of Diborane

• Used as a high-energy propellant/fuel.
• Used as a selective reducing agent in organic chemistry.
• Used in welding torches.

18. McAfee Process

Aluminium chloride (AlCl₃) is manufactured commercially by heating alumina and coke in a current of dry chlorine gas:

2Al₂O₃ + 3C + 6Cl₂ → 4AlCl₃ + 3CO₂

19. Write Uses of Aluminium chloride

1. Used extensively as a Lewis acid catalyst in Friedel-Crafts reactions.
2. Used in the manufacture of petrol by cracking mineral oils.

20. Write Preparation of Potash Alum

Potash alum is prepared commercially from alum stone by treating it with sulphuric acid and adding a calculated quantity of potassium sulphate:

K₂SO₄·Al₂(SO₄)₃·4Al(OH)₃ + 6H₂SO₄ → K₂SO₄ + 3Al₂(SO₄)₃ + 12H₂O
K₂SO₄ + Al₂(SO₄)₃ + 24H₂O → K₂SO₄·Al₂(SO₄)₃·24H₂O (Potash Alum)

21. Action of Heat on Potash Alum

At 475 K, it loses all its water of crystallization and swells to form a porous mass known as burnt alum.

K₂SO₄·Al₂(SO₄)₃·24H₂O → K₂SO₄·Al₂(SO₄)₃ + 24H₂O

On further strong red-heating, it decomposes into potassium sulphate, alumina, and sulphur trioxide:

K₂SO₄·Al₂(SO₄)₃ → K₂SO₄ + Al₂O₃ + 3SO₃

22. Uses of Potash Alum

• Used for the purification of drinking water.
• Used in waterproofing textiles.
• Used as a mordant in dyeing and leather tanning.
• Used as a styptic agent to arrest minor bleeding.

23. What is Catenation? Write its necessary conditions.

Catenation: The inherent ability of an element to form a long chain or ring of atoms linked through covalent bonds.

Conditions:

• The valency of the element must be greater than or equal to 2.
• The element should have the ability to bond with itself.
• The self-bond strength must be high.
• The catenated compounds should be kinetically inert towards other molecules.

24. Write a Note on Carbon Allotropes.

(i) Graphite: Soft, conducts electricity, planar hexagonal lattice sheets, sp² hybridised, C-C bond length is 1.41Å. Used as a solid lubricant.

(ii) Diamond: Extremely hard insulator, 3D tetrahedral framework, sp³ hybridised, C-C bond length is 1.54Å. Used for cutting glass and sharpening hard tools.

(iii) Fullerenes: Discrete cage-like structures. C₆₀ (Buckminsterfullerene) resembles a soccer ball, consisting of 12 five-membered rings fused to 20 six-membered rings with sp² hybridisation.

(iv) Carbon Nanotubes: Graphite-like cylindrical tubes closed with fullerene caps. They are mechanically stronger than steel, conduct electricity, and are used in nanotechnology.

(v) Graphene: A single isolated atomic layer sheet of carbon atoms organized in a 2D honeycomb lattice structure with sp² hybridisation.

25. Phosgene Gas

Carbonyl chloride (COCl₂) is commonly called phosgene gas. It is prepared by passing carbon monoxide and chlorine through charcoal:

CO + Cl₂ → COCl₂

It is highly poisonous and is used industrially in the synthesis of organic isocyanates.

26. Fisher-Tropsch Synthesis

The reaction of carbon monoxide with hydrogen at 500-700K under less than 50 atm pressure in the presence of a metal catalyst to form liquid hydrocarbons:

nCO + (2n+1)H₂ → C_nH_2n+2 + nH₂O
nCO + 2nH₂ → C_nH_2n + nH₂O

27. What are silicones? Write their preparation, characteristics and uses.

Definition: Silicones (or polysiloxanes) are organosilicon polymers containing recurring (-R₂Si-O-) links with an empirical formula (R₂SiO)_n.

Preparation: Prepared by the hydrolysis of dialkyldichlorosilanes:

2RCl + Si → R₂SiCl₂ → (+H₂O) → R₂Si(OH)₂ → (Polymerisation) → (R₂SiO)_n

Characteristics	Uses
Highly water-repellent due to organic side groups.	Used as high-temperature lubricants.
Excellent thermal and electrical insulators.	Used in vacuum pumps and oil baths.
Chemically inert and resistant to oxidation.	Used for waterproofing clothes and fabrics.
Vary from oily liquids to waxy solids depending on chain length.	Mixed with paints and enamels for superior heat resistance.

28. What are silicates? Name their types.

Silicates are minerals containing metal cations bonded to tetrahedral orthosilicate structural units [SiO₄]^4- linked together in different patterns.

Types:

1. Ortho silicates (Island silicates) - e.g., Phenacite
2. Pyro silicates (Sorosilates) - e.g., Thortveitite
3. Cyclic silicates (Ring silicates) - e.g., Beryl
4. Inosilicates (Chain / Double chain) - e.g., Spodumene, Asbestos
5. Sheet or phyllo silicates - e.g., Talc, Mica
6. Three-dimensional silicates (Tektosilicates) - e.g., Quartz

29. Write a note about Zeolites.

1. Zeolites are microporous three-dimensional crystalline aluminosilicate solids.
2. They are hydrated sodium aluminium silicates with the general formula Na₂O·Al₂O₃·xSiO₂·yH₂O.
3. They possess a regular cage-like honeycomb network consisting of interconnected channels and tunnels.
4. The monovalent sodium ions and water molecules are loosely held inside these pores.
5. Zeolites are extensively used as molecular sieves and to remove permanent hardness of water.

1. General characteristics of solids

1. They possess definite volume, mass, and shape.
2. They are rigid and virtually incompressible.
3. They have short interatomic distances and strong cohesive forces.

2. Differences between crystalline and amorphous solids

Crystalline solids	Amorphous solids
Long-range orderly arrangement of constituent particles.	Short-range or random arrangement of constituents.
Definite, characteristic geometrical shape.	Irregular shape.
Considered true solids.	Considered pseudo-solids or supercooled liquids.
Anisotropic in nature.	Isotropic in nature.
Example: NaCl, Quartz.	Example: Rubber, Glass, Plastics.

3. Isotropy vs Anisotropy

Anisotropy	Isotropy
Showing different physical properties (refractive index, electrical conductivity) when measured along different directions.	Showing identical physical properties in all directions.
Characteristic of crystalline solids. Example: NaCl	Characteristic of amorphous solids. Example: Rubber, Glass

4. Classify types of solids based on binding forces: Ionic, Covalent, Molecular, Metallic

Ionic solids	Covalent solids	Molecular solids
Lattice points are occupied by alternating cations and anions.	Lattice points are occupied by atoms linked continuously.	Lattice points contain neutral discrete molecules.
Held together by strong electrostatic coulombic attractions.	Interconnected throughout by strong directional covalent bonds.	Held together by weak intermolecular van der Waals forces.
Hard, brittle, high melting points. Ex: NaCl	Very hard, network structures. Ex: Diamond, SiC	Soft, lower melting points. Ex: Solid Ice, Iodine

Metallic solids: The lattice points are occupied by positive metal ions immersed in a mobile cloud of delocalized valence electrons. Example: Fe, Ag, Copper.

5. Why are ionic crystals hard and brittle?

The ions are held fixed in their lattice position by extremely powerful, non-directional electrostatic attractive forces making them hard. However, when an external shearing force is applied, layers of like-sign ions shift to face each other, resulting in intense electrostatic repulsion that cleaves the crystal, making it brittle.

6. Classify the following into Covalent, Molecular, Ionic and Metallic solids: Diamond, Brass, SiC, Naphthalene, Glucose, NaCl, P4, Iodine.

• Ionic solid: NaCl
• Covalent solid: Diamond, Silicon carbide (SiC)
• Molecular solid: Naphthalene, Glucose, Phosphorus (P₄), Iodine (I₂)
• Metallic solid: Brass

7. Explain the types of molecular solids

Types	Binding Interactions	Characteristics	Examples
Non-polar molecular solids	Weak London dispersion forces	Very low melting points, volatile soft solids	Naphthalene, Anthracene, Solid Ar
Polar molecular solids	Dipole-dipole interactions	Moderate melting points, insulators	Solid CO2, Solid NH3, Solid HCl
Hydrogen-bonded molecular solids	Intermolecular Hydrogen bonds	Soft solids at RT, non-conductors	Glucose, Urea, Solid Ice

8. Define Crystal lattice & Unit cell

Crystal lattice: The regular, repeating three-dimensional spatial arrangement of constituent particles (atoms, ions, or molecules) throughout a crystalline solid.

Unit cell: The fundamental basic repeating structural unit of a crystalline solid which, when repeated over and over in all directions, generates the entire macroscopic crystal lattice.

9. What is meant by coordination number? What is the coordination number of atoms in a bcc structure?

The coordination number is the total number of nearest neighbouring particles directly surrounding or touching a central particle in a crystal lattice.

The coordination number of a Body-Centered Cubic (BCC) structure is 8.

10. Calculate the number of atoms per unit cell in cubic lattices (SC, BCC, FCC)

Simple Cubic (SC)	Body-Centered Cubic (BCC)	Face-Centered Cubic (FCC)
Nc / 8 = 8 / 8	(Nc / 8) + (Nb / 1) = (8 / 8) + (1 / 1)	(Nc / 8) + (Nf / 2) = (8 / 8) + (6 / 2)
1 atom	1 + 1 = 2 atoms	1 + 3 = 4 atoms
Example: Polonium	Example: CsCl (lattice type)	Example: NaCl (lattice type)

11. Write the formulas for Unit Cell Density and Bragg's Law.

Density of a unit cell:

ρ = (n × M) / (a³ × N_A)

Where: n = number of atoms per cell, M = molar mass, a³ = volume of unit cell, N_A = Avogadro's number.

Bragg's Equation:

nλ = 2d sinθ

Where: n = order of diffraction, λ = wavelength of X-rays, d = interplanar spacing, θ = angle of diffraction.

12. Tabulate Packing Efficiency in Cubic Crystals

Lattice Type	Atomic Radius (r) formula	Packing Efficiency (%)	Void / Vacant Space (%)
Simple Cubic (SC)	r = a / 2	52.33%	47.67%
Body-Centered Cubic (BCC)	r = √3a / 4	68.00%	32.00%
Face-Centered Cubic (FCC)	r = √2a / 4	74.00%	26.00%

13. Distinguish between Hexagonal close packing (HCP) and Cubic close packing (CCP)

Hexagonal close packing (hcp)	Cubic close packing (ccp)
ABAB... stacking pattern.	ABCABC... stacking pattern.
Tetrahedral voids of the second layer are directly covered by spheres of the third layer.	Octahedral voids of the second layer are covered by spheres of the third layer.
Based on a hexagonal unit cell (6 spheres/cell).	Based on a face-centered cubic (FCC) unit cell (4 spheres/cell).

14. What is the radius ratio rule? Tabulate the structure relationships.

The radius ratio is the ratio of the radius of the cation (r₊) to the radius of the anion (r_-). It helps predict the coordination number and structure of ionic solids.

Radius Ratio Range	Coordination Number	Structural Geometry	Example
0.155 - 0.225	3	Trigonal planar	B2O3
0.225 - 0.414	4	Tetrahedral	ZnS
0.414 - 0.732	6	Octahedral	NaCl
0.732 - 1.000	8	Cubic	CsCl

15. Distinguish between tetrahedral and octahedral voids

Tetrahedral voids	Octahedral voids
A void enclosed or surrounded by 4 spheres in a tetrahedral geometry.	A void enclosed or surrounded by 6 spheres in an octahedral geometry.
Coordination number = 4	Coordination number = 6
Number of tetrahedral voids in a close packing = 2n (where n = number of spheres).	Number of octahedral voids in a close packing = n.

16. Explain Point Defects in Solids: Schottky vs Frenkel, Excess vs Deficiency

Stoichiometric Point Defects:

Schottky Defect	Frenkel Defect
Arises due to missing equal numbers of cations and anions from the crystal lattice.	Arises when an ion leaves its normal lattice site and occupies an interstitial position.
Observed when cations and anions are of similar sizes.	Observed when there is a significant size difference between ions.
Density of the crystal decreases. Example: NaCl	Density remains unchanged. Example: AgBr, AgCl

Non-Stoichiometric Defects:

• Metal excess defect: Occurs due to anion vacancies occupied by unpaired electrons (called F-centers which impart colour, like yellow in NaCl) or extra interstitial cations.
  Example: ZnO turns yellow on heating because it loses oxygen, leaving Zn²⁺ and free electrons to occupy interstitial spaces.
• Metal deficiency defect: Occurs when a cation is missing from its lattice site, and a nearby metal ion acquires a higher oxidation state to maintain charge neutrality.
  Example: FeO, FeS.

1. Distinguish between Average Rate and Instantaneous Rate.

Average Rate	Instantaneous Rate
The change in concentration of reactants or products divided by a finite measurable time interval (Δt).	The rate of a chemical reaction at a specific single instant of time (dt).
Rate = -Δ[A] / Δt = +Δ[B] / Δt	Rate = -d[A] / dt = +d[B] / dt

2. Rate Law, Rate Constant and Order of Reaction

For a general reaction: xA + yB → Products

Rate Law Expression: Rate = k[A]^m[B]ⁿ

Where k is the rate constant (proportionality constant equal to the rate when concentrations are unity). The powers m and n represent orders with respect to A and B. The overall order = m + n (determined purely by experiments).

Units of k:

• Zero order: mol L^-1 s^-1
• First order: s^-1
• Second order: L mol^-1 s^-1

3. Differences between Rate and Rate Constant of a Reaction

Rate of a Reaction	Rate Constant of a Reaction
Represents the actual speed at which reactants are converted to products at any instant.	It is a fundamental proportionality factor in the rate law equation.
Measured as decrease in concentration of reactants per unit time.	Equal to the reaction rate when the concentration of all reacting species is 1 M.
Depends directly on the initial/instantaneous concentration of reactants.	Independent of the initial concentrations of the reactants.

4. Explain Rate Determining Step (RDS) with an example.

In a complex multi-step mechanism, the slowest individual elementary step controls the overall speed of the reaction and is called the Rate Determining Step.

Example: Decomposition of Hydrogen Peroxide (2H₂O₂ → 2H₂O + O₂)

• Step 1: H₂O₂ + I^- → H₂O + OI^- (Slow, RDS)
• Step 2: H₂O₂ + OI^- → H₂O + I^- + O₂ (Fast)

Because two species are involved in the slow RDS step, it follows second-order (bimolecular) kinetics overall.

5. Difference between Order and Molecularity

Order of a Reaction	Molecularity of a Reaction
Sum of the exponential concentration terms in the experimentally determined rate law.	The total number of reactant species taking part in an individual elementary chemical step.
Can be zero, fractional, negative, or an integer.	Must always be a whole positive integer (cannot be zero or fractional).
Applicable to the entire overall reaction.	Assigned separately for each individual single elementary step.

6. Derivations of Integrated Rate Laws & Half Life for First and Zero-Order Reactions

(A) First-Order Reaction (A → Products)

Differential Rate: -d[A]/dt = k[A] ⇒ -d[A]/[A] = k dt

Integrating between boundaries [A₀] at t=0 to [A] at t=t:

k = (2.303 / t) × log([A₀] / [A])

Half life (t_1/2): At t = t_1/2, [A] = [A₀]/2:

t_1/2 = 0.693 / k

Conclusion: The half-life of a first-order reaction is completely independent of the initial concentration of reactants.

(B) Zero-Order Reaction (A → Products)

Differential Rate: -d[A]/dt = k[A]⁰ = k ⇒ -d[A] = k dt

Integrating gives:

k = ([A₀] - [A]) / t

Half life (t_1/2): At t = t_1/2, [A] = [A₀]/2:

t_1/2 = [A₀] / 2k

Conclusion: The half-life of a zero-order reaction is directly proportional to the initial reactant concentration.

7. Give examples for First-Order and Zero-Order reactions.

First Order Examples:

1. Decomposition of N₂O₅ → 2NO₂ + ½O₂
2. Decomposition of sulphuryl chloride: SO₂Cl₂ → SO₂ + Cl₂
3. Decomposition of H₂O₂ in aqueous medium.
4. Isomerisation of cyclopropane to propene at 780 K.

Zero Order Examples:

1. Photochemical reaction between Hydrogen and Chlorine: H₂ + Cl₂ → 2HCl (in presence of hν)
2. Decomposition of nitrous oxide (N₂O) on a hot Platinum catalyst surface.
3. Iodination of acetone in an acidic medium (zero order with respect to Iodine).

8. What is a Pseudo First Order Reaction? Give an example.

A higher-order reaction can be altered kinetically to behave as a first-order reaction by choosing one of the reactants in a massive excess compared to the other, making its concentration virtually constant during the process.

Example: Acid-catalyzed hydrolysis of an Ester

CH₃COOCH₃ + H₂O (excess) → (+H⁺) → CH₃COOH + CH₃OH

Rate = k[CH₃COOCH₃][H₂O] ≈ k'[CH₃COOCH₃]

9. Explain briefly the collision theory of bimolecular reactions

• Chemical reactions occur primarily due to collisions between reacting molecules.
• Collision rate is directly proportional to concentration: Collision rate = Z[A₂][B₂]
• Not all collisions yield products; reacting molecules must possess a minimum threshold energy known as Activation Energy (E_a).
• Furthermore, molecules must collide with a specific geometric orientation (Steric orientation factor, p).
• The final rate constant is given by: k = p × Z × e^(-Ea/RT)

10. Write the Arrhenius equation and state the factors affecting reaction rates.

Arrhenius Equation:

k = A × e^(-Ea/RT)

Where k = rate constant, A = frequency factor, E_a = activation energy, R = gas constant, T = absolute temperature.

Factors affecting reaction rates:

• Nature and state of reactants: Gaseous reactions are faster than solids. (e.g., Iodine vapours react faster with Na than solid Iodine).
• Concentration of reactants: High concentrations increase collision frequency.
• Surface area: Powdered reactants react faster due to a greater surface area (e.g., powdered CaCO₃ reacts faster with HCl than marble chips).
• Temperature: Reaction rate generally doubles with every 10°C increase.
• Presence of a catalyst: Accelerates the rate by offering an alternate pathway with a lower activation energy barrier.

```