Chemistry class 12 Isc

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ISC Class 12 Chemistry Notes

ISC Class 12 Chemistry Notes

1. Solid State

  • Definition: Study of solids, their structures, and properties.
  • Key Concepts:
    • Unit Cell: Smallest repeating unit in a crystal lattice.
    • Types of Solids: Crystalline and Amorphous.
    • Defects in Solids: Point defects, line defects, etc.

2. Solutions

  • Definition: Homogeneous mixtures of two or more substances.
  • Key Concepts:
    • Concentration Terms: Molarity, Molality, Normality.
    • Colligative Properties: Boiling point elevation, Freezing point depression.

3. Electrochemistry

  • Definition: Study of chemical reactions that produce electricity.
  • Key Concepts:
    • Electrolysis: Decomposition of a compound by electric current.
    • Electrochemical Cells: Galvanic and electrolytic cells.
    • Nernst Equation: Determines the equilibrium potential of an ion.

4. Chemical Kinetics

  • Definition: Study of the rate of chemical reactions.
  • Key Concepts:
    • Rate of Reaction: Change in concentration of reactants/products per unit time.
    • Order of Reaction: Sum of the powers of concentration terms in the rate law.
    • Arrhenius Equation: Relates temperature to reaction rate.

5. Surface Chemistry

  • Definition: Study of phenomena at the surface of materials.
  • Key Concepts:
    • Adsorption: Accumulation of molecules at the surface.
    • Catalysis: Acceleration of a chemical reaction by a catalyst.

6. General Principles and Processes of Isolation of Elements

  • Definition: Methods of extracting metals from their ores.
  • Key Concepts:
    • Concentration of Ore: Removal of impurities from the ore.
    • Extraction Methods: Reduction, Electrolytic reduction.
    • Refining: Purification of metals.

7. p-Block Elements

  • Definition: Elements in groups 13-18 of the periodic table.
  • Key Concepts:
    • Group 15 Elements: Nitrogen, Phosphorus.
    • Group 16 Elements: Oxygen, Sulfur.
    • Group 17 Elements: Halogens.
    • Group 18 Elements: Noble gases.

8. d- and f-Block Elements

  • Definition: Transition elements and inner transition elements.
  • Key Concepts:
    • Properties: Variable oxidation states, Magnetic properties.
    • Lanthanides and Actinides: Inner transition series.

9. Coordination Compounds

  • Definition: Compounds containing coordinate bonds.
  • Key Concepts:
    • Ligands: Molecules or ions that donate electron pairs.
    • Coordination Number: Number of ligand donor atoms attached to the metal atom.
    • Solid State – ISC Class 12 Chemistry

      Chapter 1: Solid State

      Introduction

      The solid state is one of the fundamental states of matter. It is characterized by strong intermolecular forces and a definite shape and volume. The properties of solids depend on the arrangement of particles within them.

      Types of Solids

      • Crystalline Solids: Particles arranged in a repeating pattern (e.g., NaCl, Quartz).
      • Amorphous Solids: Irregular arrangement of particles (e.g., glass, rubber).
      Key Point: Crystalline solids have sharp melting points and long-range order; amorphous solids do not.

      Classification of Crystalline Solids

      • Ionic Solids: Ions as constituents; strong electrostatic forces (e.g., NaCl).
      • Molecular Solids: Molecules held by van der Waals forces or hydrogen bonds (e.g., ice).
      • Covalent Solids: Atoms bonded by covalent bonds throughout the crystal (e.g., diamond).
      • Metallic Solids: Positive ions in a sea of delocalized electrons (e.g., copper).

      Crystal Lattices and Unit Cells

      A crystal lattice is a 3D arrangement of constituent particles. The smallest repeating unit in a lattice is the unit cell.

      Types of Unit Cells:

      • Simple Cubic: Atoms at the corners only.
      • Body-Centered Cubic (BCC): Corners + one atom at the center.
      • Face-Centered Cubic (FCC): Corners + one atom on each face center.

      Number of Atoms in a Unit Cell

      • Simple Cubic: 1 atom/unit cell
      • BCC: 2 atoms/unit cell
      • FCC: 4 atoms/unit cell

      Packing Efficiency

      It is the percentage of total space filled by particles in a unit cell.

      • Simple Cubic: ~52.4%
      • BCC: ~68%
      • FCC: ~74%

      Void & Density of Unit Cell

      Density (ρ) = (Z × M) / (a³ × NA)
      Where:
      • Z = Number of atoms per unit cell
      • M = Molar mass (g/mol)
      • a = Edge length in cm
      • NA = Avogadro number (6.022 × 10²³)

      Imperfections in Solids

      Real solids may have defects that affect their physical and electrical properties.

      Types of Defects:

      • Point Defects: Missing/extra atoms (vacancy, interstitial).
      • Stoichiometric Defects: Maintain stoichiometry (Schottky, Frenkel).
      • Non-Stoichiometric Defects: Excess/deficiency of ions (metal excess, metal deficiency).

      Electrical & Magnetic Properties

      • Conductors: Allow flow of electricity (e.g., metals).
      • Semiconductors: Conductivity increases with temperature (e.g., silicon).
      • Insulators: Do not conduct (e.g., glass).
      • Magnetic Types: Diamagnetic, paramagnetic, ferromagnetic.

      Applications

      • Metals and alloys in construction and electronics
      • Semiconductors in computers and solar cells
      • Ceramics in industry due to hardness and heat resistance
      Solutions – ISC Class 12 Chemistry

      Chapter 2: Solutions

      Introduction

      A solution is a homogeneous mixture of two or more substances. The component in larger quantity is called the solvent, and the one in smaller quantity is called the solute.

      Types of Solutions

      • Gaseous Solutions: Gas in gas (e.g., air).
      • Liquid Solutions: Liquid in liquid, solid in liquid, gas in liquid (e.g., salt water).
      • Solid Solutions: Solid in solid (e.g., alloys).

      Concentration Terms

      • Mass Percent: (mass of solute / mass of solution) × 100
      • Volume Percent: (volume of solute / volume of solution) × 100
      • Molarity (M): moles of solute / volume of solution (L)
      • Molality (m): moles of solute / mass of solvent (kg)
      • Mole Fraction (χ): moles of component / total moles of all components

      Solubility

      Solubility is the maximum amount of solute that can be dissolved in a given amount of solvent at a specified temperature.

      Henry’s Law:
      p = k_H · x
      where p is the partial pressure, x is the mole fraction, and k_H is Henry’s constant.

      Raoult’s Law

      The partial vapor pressure of a component in a solution is directly proportional to its mole fraction.

      Raoult’s Law for Ideal Solutions:
      P = PAxA + PBxB

      Colligative Properties

      These properties depend only on the number of solute particles, not on their nature.

      • Relative Lowering of Vapor Pressure: (P1⁰ - P1) / P1⁰ = x2
      • Elevation in Boiling Point: ΔTb = Kb · m
      • Depression in Freezing Point: ΔTf = Kf · m
      • Osmotic Pressure: π = C R T or π = (n/V)RT

      Abnormal Molar Mass

      Due to association or dissociation of solute particles, the observed molar mass differs from the expected value.

      van’t Hoff Factor (i):
      i = observed colligative property / normal colligative property

      Important Points

      • Colligative properties are used to determine molar mass.
      • Raoult’s law is applicable to ideal solutions only.
      • Molality is preferred over molarity because it is temperature-independent.
      Electrochemistry – ISC Class 12 Chemistry

      Chapter 3: Electrochemistry

      Introduction

      Electrochemistry is the study of the relationship between electrical energy and chemical changes. It involves redox reactions where oxidation and reduction occur simultaneously in separate half-cells.

      Types of Electrochemical Cells

      • Galvanic (Voltaic) Cell: Converts chemical energy to electrical energy.
      • Electrolytic Cell: Uses electrical energy to bring about chemical change.

      Galvanic Cell Components

      • Anode: Site of oxidation (negative electrode)
      • Cathode: Site of reduction (positive electrode)
      • Salt Bridge: Maintains electrical neutrality
      • External Circuit: Electrons flow from anode to cathode
      Example: Daniell Cell
      Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

      Electrode Potential

      Potential difference between electrode and its solution.

      • Standard Electrode Potential (E°): Measured under 1 M, 1 atm, 25°C.
      • Standard Hydrogen Electrode (SHE): Reference electrode with E° = 0 V.

      Nernst Equation

      Used to calculate electrode potential under non-standard conditions.

      E = E° - (0.0591/n) log Q
      Where:
      • E = electrode potential
      • E° = standard electrode potential
      • n = number of electrons transferred
      • Q = reaction quotient

      Gibbs Free Energy and EMF

      ΔG = -nFE
      Where:
      • ΔG = Gibbs free energy change
      • n = number of electrons
      • F = Faraday constant (96,500 C/mol)
      • E = Cell EMF

      Conductance of Electrolytes

      • Specific Conductance (κ): Conductance of 1 cm³ of solution
      • Molar Conductance (Λm): Λm = κ × 1000 / C
      • Cell Constant (G*): G* = L/A

      Kohlrausch’s Law

      Statement: At infinite dilution, each ion contributes independently to the total molar conductivity.
      Λm° = λ+ + λ-

      Electrolysis and Faraday’s Laws

      • First Law: Mass of substance is proportional to quantity of charge passed.
      • Second Law: Different substances liberated are proportional to their equivalent weights.
      Formula: W = (E × Q) / 96500
      Where W = mass, E = equivalent weight, Q = charge in coulombs

      Batteries and Fuel Cells

      • Primary Cells: Cannot be recharged (e.g., dry cell)
      • Secondary Cells: Rechargeable (e.g., lead-acid battery)
      • Fuel Cell: Converts chemical energy directly into electricity (e.g., H₂-O₂ fuel cell)
      Chemical Kinetics – ISC Class 12 Chemistry

      Chapter 4: Chemical Kinetics

      Introduction

      Chemical kinetics is the study of the speed or rate of a chemical reaction and the factors that influence this rate.

      Rate of Reaction

      • Average Rate: Change in concentration per unit time over a time interval.
      • Instantaneous Rate: Rate at a specific instant, given by the slope of the concentration-time curve.
      • Units: mol L−1 s−1
      Rate expression:
      Rate = k [A]m[B]n
      Where k is the rate constant, and m, n are orders with respect to A and B.

      Order and Molecularity

      • Order of Reaction: Sum of the powers of concentration in the rate law.
      • Molecularity: Number of molecules taking part in a single step of the reaction mechanism.
      • Order can be zero, fractional, or whole number; molecularity is always a whole number.

      Integrated Rate Equations

      1. Zero Order Reaction

      [A] = [A]0 − kt
      Half-life: t1/2 = [A]0 / 2k

      2. First Order Reaction

      ln[A] = ln[A]0 − kt or [A] = [A]0e−kt
      Half-life: t1/2 = 0.693 / k

      3. Second Order Reaction

      1/[A] = 1/[A]0 + kt

      Arrhenius Equation

      k = A e−Ea/RT
      Where:
      • k = rate constant
      • A = frequency factor
      • Ea = activation energy
      • R = gas constant (8.314 J/mol·K)
      • T = temperature in Kelvin

      Logarithmic form:

      ln k = ln A − Ea/RT or log k = log A − (Ea / 2.303RT)

      Factors Affecting Reaction Rate

      • Nature and concentration of reactants
      • Temperature
      • Presence of a catalyst
      • Surface area (for solids)

      Activation Energy and Reaction Pathway

      Activation energy is the minimum energy required for reactants to form products. A catalyst lowers the activation energy, increasing the reaction rate.

      Graphical Representations

      • Zero order: [A] vs. time → straight line
      • First order: ln[A] vs. time → straight line with negative slope
      • Second order: 1/[A] vs. time → straight line
      Reaction Order Graphs – Chemical Kinetics

      Graphical Representation of Different Reaction Orders

      Understanding how concentration changes with time helps determine the order of a reaction. Each order has a characteristic plot.

      Zero Order Reaction

      Plot: [A] vs. Time → Straight line (negative slope)

      Equation: [A] = [A]0 − kt

      Zero order reaction graph

      First Order Reaction

      Plot: ln[A] vs. Time → Straight line (negative slope)

      Equation: ln[A] = ln[A]0 − kt

      First order reaction graph

      Second Order Reaction

      Plot: 1/[A] vs. Time → Straight line (positive slope)

      Equation: 1/[A] = 1/[A]0 + kt

      Second order reaction graph
      Surface Chemistry – ISC Class 12

      Chapter 5: Surface Chemistry

      1. Introduction

      Surface chemistry deals with phenomena that occur at the surfaces or interfaces of substances. Important topics include adsorption, catalysis, and colloids.

      2. Adsorption

      Adsorption is the accumulation of molecules at the surface rather than in the bulk.

      Types of Adsorption

      • Physisorption (Physical): Weak van der Waals forces, low heat of adsorption
      • Chemisorption (Chemical): Strong chemical bonding, high heat of adsorption

      Factors Affecting Adsorption

      • Nature of adsorbent and adsorbate
      • Surface area of adsorbent
      • Temperature and pressure

      Freundlich Adsorption Isotherm

      x/m = kP1/n
      where x = mass of adsorbate, m = mass of adsorbent, P = pressure

      3. Catalysis

      Catalysis is the process of increasing the rate of a reaction by adding a substance called a catalyst.

      Types of Catalysis

      • Homogeneous: Catalyst in same phase as reactants (e.g., NO in oxidation of SO2)
      • Heterogeneous: Catalyst in different phase (e.g., Fe in Haber process)

      Characteristics of Catalysts

      • Not consumed in the reaction
      • Effective in small amounts
      • Do not alter equilibrium position

      Enzyme Catalysis

      Biochemical reactions use enzymes as highly specific catalysts. They work best at optimum pH and temperature.

      4. Colloids

      A colloid is a heterogeneous system in which one substance is dispersed in another.

      Types of Colloidal Systems

      • Sol: Solid in liquid (e.g., paint)
      • Gel: Liquid in solid (e.g., jelly)
      • Emulsion: Liquid in liquid (e.g., milk)
      • Aerosol: Solid or liquid in gas (e.g., smoke, fog)
      • Foam: Gas in liquid (e.g., shaving cream)

      Preparation of Colloids

      • Dispersion methods: Breaking large particles (e.g., mechanical dispersion)
      • Condensation methods: Aggregating small particles (e.g., chemical reactions)

      Purification of Colloids

      • Dialysis
      • Electrodialysis
      • Ultrafiltration

      Properties of Colloids

      • Tyndall Effect: Scattering of light by colloidal particles
      • Brownian Movement: Zig-zag motion due to collisions
      • Electrophoresis: Movement of colloidal particles in electric field
      • Coagulation: Precipitation of colloids

      5. Emulsions

      An emulsion is a colloidal dispersion of one liquid in another immiscible liquid.

      • Oil-in-water (O/W): Milk
      • Water-in-oil (W/O): Butter

      Applications of Colloids

      • In medicines and drug delivery
      • Purification of water (coagulation)
      • Food products (emulsions)
      • Rubber and plastics industry
      Isolation of Elements – ISC Class 12 Chemistry

      Chapter 6: General Principles and Processes of Isolation of Elements

      Introduction

      This chapter deals with the extraction and purification of metals from their ores — a field known as metallurgy.

      1. Important Terms

      • Minerals: Naturally occurring chemical substances in the earth’s crust.
      • Ores: Minerals from which metals can be extracted profitably.
      • Gangue: Earthy impurities present in ores.

      2. Steps in Metallurgy

      1. Concentration of Ore

      Removal of gangue using physical or chemical methods:

      • Gravity separation: For heavier ore particles.
      • Froth flotation: For sulphide ores.
      • Magnetic separation: For ores with magnetic properties.

      2. Calcination and Roasting

      • Calcination: Heating in absence of air. Removes volatile impurities and water.
      • Roasting: Heating in presence of air. Converts sulphides to oxides.

      3. Reduction of Oxide to Metal

      • Smelting: Metal is extracted by heating with a reducing agent like carbon or CO.
      • Aluminothermic process: For metals like chromium and manganese.
        Cr2O3 + 2Al → 2Cr + Al2O3
      • Electrolytic reduction: For highly reactive metals like Na, K, Al.

      4. Refining of Metals

      • Distillation: For low boiling metals like Zn and Hg.
      • Liquation: Metal is melted and flows away from impurities.
      • Electrolytic refining: Impure metal is made the anode.
        Example: Cu refining using CuSO4 solution.
      • Zone refining: For ultra-pure semiconductors like Si and Ge.

      3. Thermodynamic Principles of Metallurgy

      Based on Gibbs Free Energy change: ΔG = ΔH − TΔS

      Ellingham Diagram:
      A graph of ΔG vs. T that helps predict the feasibility of oxide reduction reactions.

      4. Extraction of Important Metals

      • Aluminium: Extracted by Hall-Héroult process using electrolysis of molten alumina (Al2O3) in cryolite (Na3AlF6).
      • Iron: Extracted from hematite (Fe2O3) in blast furnace.
      • Copper: Roasting of copper pyrites followed by Bessemer converter.
      • Gold and Silver: Cyanide process (leaching with NaCN or KCN solution).

      5. Environmental Issues in Metallurgy

      • Air pollution due to SO2 during roasting.
      • Water contamination from leaching processes.
      • Recycling and bioleaching as eco-friendly methods.

      Applications

      • Used in metal extraction industries
      • Electronics (purification of semiconductors)
      • Jewelry and alloy production
      The p-Block Elements – ISC Class 12 Chemistry

      Chapter 7: The p-Block Elements

      This chapter deals with the chemistry of Group 15 and 16 elements (Nitrogen, Phosphorus, Oxygen, Sulfur, and their compounds).

      Group 15 Elements: Nitrogen Family

      • Members: N, P, As, Sb, Bi
      • General electronic configuration: ns2np3
      • Exhibit oxidation states: –3, +3, +5
      • Tendency to form covalent compounds
      • Nitrogen shows unique behavior due to small size, high electronegativity, and ability to form pπ–pπ multiple bonds

      Important Compounds of Nitrogen

      • Ammonia (NH₃): Forms H-bonds, acts as a weak base
      • Nitric acid (HNO₃): Strong oxidizing agent
      • Oxides of Nitrogen: NO, NO₂, N₂O, N₂O₃, N₂O₅ (varied oxidation states)

      Group 16 Elements: Oxygen Family (Chalcogens)

      • Members: O, S, Se, Te, Po
      • Electronic configuration: ns2np4
      • Show oxidation states: –2, +2, +4, +6
      • Oxygen forms pπ–pπ multiple bonds; sulfur shows catenation

      Important Compounds of Sulfur

      • Oxides: SO₂, SO₃
      • Oxyacids: H₂SO₃, H₂SO₄
      • Structure of SO₂: Angular, resonance structures
      • H₂SO₄: Strong acid, dehydrating agent, oxidizer

      Trends in Groups 15 and 16

      • Atomic size increases down the group
      • Ionization enthalpy decreases
      • Electronegativity decreases
      • Metallic character increases
      • Oxidizing power: O > S > Se

      Anomalous Behavior of Oxygen and Nitrogen

      • Due to small size, high electronegativity, absence of d-orbitals
      • Form strong hydrogen bonds
      • Oxygen exists as O₂ (diatomic gas), while others form solids (e.g., S₈)
      • Nitrogen has very strong triple bond (N≡N)

      Uses of Group 15 & 16 Compounds

      • NH₃: Fertilizers (urea), refrigerant
      • HNO₃: Explosives (TNT), fertilizers
      • H₂SO₄: Battery acid, manufacturing chemicals
      • SO₂: Preservative, bleaching agent
      Group 17 and 18 Elements – ISC Class 12 Chemistry

      Group 17 and 18 Elements

      This section covers the properties, trends, and uses of the Halogens (Group 17) and Noble Gases (Group 18).

      Group 17: Halogens

      Members: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At)

      General configuration: ns2np5

      Physical Properties

      • Exist as diatomic molecules (F₂, Cl₂, etc.)
      • Colored and volatile
      • Trend in state: F₂, Cl₂ – gases; Br₂ – liquid; I₂ – solid

      Chemical Properties

      • Highly reactive non-metals
      • Strong oxidizing agents (F₂ > Cl₂ > Br₂ > I₂)
      • React with metals to form halides
      • Form interhalogen compounds and oxoacids

      Important Compounds

      • Hydrogen halides (HX): HF, HCl, HBr, HI – acidic in nature
      • Interhalogen compounds: e.g. ClF₃, BrF₅
      • Oxoacids: HClO, HClO₂, HClO₃, HClO₄

      Trends in Group 17

      • Electronegativity decreases from F to I
      • Atomic size and melting/boiling point increase down the group
      • Oxidizing power decreases down the group

      Uses

      • Cl₂ – Disinfectant, bleaching powder production
      • HF – Etching of glass
      • Br₂ – Fire retardants, dyes
      • I₂ – Antiseptic (tincture iodine)

      Group 18: Noble Gases

      Members: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn)

      General configuration: ns2np6 (except He: 1s2)

      Properties

      • Monoatomic, colorless, odorless gases
      • Very low reactivity due to complete octet
      • Boiling points increase down the group

      Chemical Reactivity

      • He, Ne, Ar – Chemically inert
      • Kr, Xe – Can form compounds (e.g., XeF₂, XeF₄, XeOF₄)

      Important Xenon Compounds

      • Xenon fluorides: XeF₂, XeF₄, XeF₆
      • Xenon oxyfluorides: XeOF₄, XeO₃F₂

      Uses of Noble Gases

      • Helium: Balloons, cryogenics, arc welding
      • Neon: Advertising signs (neon lights)
      • Argon: Inert atmosphere in welding
      • Xenon: Flash lamps, ion propulsion systems
      d- and f-Block Elements – ISC Class 12 Chemistry

      Chapter 8: The d- and f-Block Elements

      Definition

      • Transition Elements (d-block): Elements in groups 3 to 12 of the periodic table with partially filled d-orbitals.
      • Inner Transition Elements (f-block): Elements in which the last electron enters the f-orbital. Includes lanthanides and actinides.

      Key Concepts & Properties

      • Variable Oxidation States: Due to similar energies of (n−1)d and ns orbitals.
      • Magnetic Properties: Unpaired d- or f-electrons contribute to paramagnetism.
      • Complex Formation: Transition metals form colored complexes due to d–d transitions.
      • Catalytic Properties: Many act as catalysts (e.g., Fe in Haber process).
      • Formation of Alloys: Transition metals mix easily to form alloys (e.g., steel, bronze).
      • Colored Ions: Due to d–d electronic transitions (e.g., Cr³⁺ – green, Cu²⁺ – blue).

      Lanthanides (4f series)

      • Atomic numbers 58–71 (Ce to Lu)
      • Common oxidation state: +3
      • Show lanthanide contraction (gradual decrease in ionic radii)
      • Form colored compounds due to f–f transitions
      • Used in lasers, alloys, and catalysts

      Actinides (5f series)

      • Atomic numbers 90–103 (Th to Lr)
      • Mostly radioactive elements
      • Show more variable oxidation states than lanthanides
      • Used in nuclear reactors and weapons (e.g., U, Pu)

      Differences: d-block vs f-block

      • d-block: Involve (n−1)d orbitals; form colored ions and complex compounds
      • f-block: Involve (n−2)f orbitals; less complex formation; mostly +3 oxidation state

      Uses of Transition Elements

      • Iron (Fe): Steel manufacturing
      • Copper (Cu): Electrical wiring
      • Titanium (Ti): Aerospace applications
      • Platinum (Pt): Catalytic converters
      Coordination Compounds – ISC Class 12 Chemistry

      Chapter 9: Coordination Compounds

      This chapter introduces coordination chemistry involving metal ions and ligands forming complex ions and compounds.

      1. Basic Terms

      • Coordination Compound: A compound consisting of a central metal atom/ion bonded to ligands.
      • Ligands: Ions or molecules that donate a lone pair of electrons to the metal (e.g., H₂O, NH₃, Cl⁻).
      • Coordination Number: Number of ligand donor atoms bonded to the central metal.
      • Coordination Sphere: The metal ion and ligands enclosed in brackets [ ].

      2. Types of Ligands

      • Monodentate: Donates one electron pair (e.g., NH₃, Cl⁻)
      • Bidentate: Donates two pairs from different atoms (e.g., ethylenediamine: en)
      • Polydentate: Ligands that can donate more than two electron pairs (e.g., EDTA⁴⁻)
      • Ambidentate: Ligands that can bind through two different atoms (e.g., NO₂⁻ via N or O)

      3. Nomenclature Rules

      • Name the ligands in alphabetical order before the metal.
      • Use prefixes like di-, tri-, tetra- for multiple ligands.
      • Neutral ligands: H₂O (aqua), NH₃ (ammine), CO (carbonyl)
      • If the complex is an anion, metal ends in -ate (e.g., ferrate for Fe).

      Example: [Co(NH₃)₆]Cl₃ → Hexaamminecobalt(III) chloride

      4. Bonding in Coordination Compounds

      • Werner’s Theory: Differentiates between primary (ionisable) and secondary (non-ionisable) valency.
      • Valence Bond Theory (VBT): Involves hybridization of metal orbitals (e.g., d²sp³ or sp³d²).
      • Crystal Field Theory (CFT): Describes d-orbital splitting in ligand fields (octahedral or tetrahedral).

      5. Isomerism

      • Structural Isomerism:
        • Ionisation isomerism
        • Linkage isomerism
        • Coordination isomerism
      • Stereoisomerism:
        • Geometrical (cis-trans) isomerism
        • Optical isomerism (non-superimposable mirror images)

      6. Important Examples

      • [Fe(CN)₆]³⁻ – Hexacyanoferrate(III)
      • [Cr(H₂O)₆]³⁺ – Hexaaquachromium(III)
      • [Ni(CO)₄] – Tetracarbonylnickel(0), tetrahedral
      • [Pt(NH₃)₂Cl₂] – Diamminedichloroplatinum(II), shows cis-trans isomerism

      7. Applications of Coordination Compounds

      • Biological: Hemoglobin (Fe²⁺) and Chlorophyll (Mg²⁺)
      • Medicinal: Cisplatin used in chemotherapy
      • Analytical: Complexometric titrations using EDTA
      • Industrial: Electroplating and dye manufacture

      Organic Chemistry Class 12 Isc

      Haloalkanes and Haloarenes – ISC Class 12 Chemistry

      Chapter 1: Haloalkanes and Haloarenes

      This chapter covers the chemistry of haloalkanes and haloarenes including their nomenclature, preparation, properties, and important reactions.

      Haloalkanes

      Nomenclature

      • Derived from alkanes by replacing hydrogen with halogen (F, Cl, Br, I).
      • Common name: Alkyl halide (e.g., methyl chloride).
      • IUPAC name: Halogen as a prefix (e.g., chloromethane).

      Preparation Methods

      • From alcohols: R–OH + HX → R–X + H₂O
      • By halogenation of alkanes: R–H + X₂ (UV light) → R–X
      • From alkenes: Addition of HX across double bonds

      Physical Properties

      • Boiling point increases with molecular mass and halogen size.
      • Insoluble in water; soluble in organic solvents.

      Chemical Reactions

      • Nucleophilic Substitution: R–X + OH⁻ → R–OH + X⁻
      • Elimination Reactions: Dehydrohalogenation in presence of alcoholic KOH → Alkene
      • Saytzeff’s Rule: More substituted alkene is the major product in elimination.
      • Wurtz Reaction: 2R–X + 2Na → R–R + 2NaX (in dry ether)
      • Grignard Reagent Formation: R–X + Mg → R–MgX (in dry ether)

      Freons

      • Chlorofluorocarbons (CFCs), e.g., CF₂Cl₂ (Freon-12).
      • Used as refrigerants but cause ozone depletion.

      Haloarenes

      Preparation

      • Sandmeyer Reaction: Ar–N₂⁺ + CuX → Ar–X + N₂
      • Gattermann Reaction: Ar–N₂⁺ + HX + Cu → Ar–X + N₂

      Chemical Properties

      • Electrophilic Substitution:
        • Nitration: Ar–X + HNO₃ → Ar–NO₂
        • Halogenation: Ar–X + X₂ → Ar–X₂
        • Sulphonation: Ar–X + H₂SO₄ → Ar–SO₃H
      • Less reactive toward nucleophilic substitution than haloalkanes due to resonance and partial double bond character of C–X bond.
      Alcohols, Phenols, and Ethers – ISC Class 12 Chemistry

      Chapter 2: Alcohols, Phenols, and Ethers

      This chapter explores the structure, preparation, properties, and uses of alcohols, phenols, and ethers.

      Alcohols

      Classification

      • Primary (1°): –OH attached to a carbon with one alkyl group
      • Secondary (2°): –OH attached to a carbon with two alkyl groups
      • Tertiary (3°): –OH attached to a carbon with three alkyl groups

      Nomenclature

      • Common: Methyl alcohol, ethyl alcohol
      • IUPAC: Methanol, ethanol, propan-2-ol

      Methods of Preparation

      • From alkenes: Hydration of ethene
      • From alkyl halides: R–X + OH⁻ → R–OH
      • By reduction of aldehydes and ketones

      Physical Properties

      • High boiling point due to hydrogen bonding
      • Soluble in water (lower members)

      Chemical Properties

      • Oxidation: Forms aldehydes/ketones or acids
      • Dehydration: Produces alkenes (acid-catalyzed)
      • Esterification: R–OH + R’COOH → Ester

      Identification of Alcohols

      • Lucas Test: Distinguishes between 1°, 2°, and 3° alcohols using ZnCl₂/HCl

      Mechanism of Dehydration

      • In acidic medium, forms carbocation → alkene

      Uses

      • Solvent, antiseptic, fuel (ethanol), manufacturing chemicals

      Phenols

      Classification & Nomenclature

      • Phenols: –OH group attached directly to an aromatic ring
      • Naming: Phenol, o-cresol, p-nitrophenol, etc.

      Preparation

      • From benzene sulfonic acid
      • From diazonium salts
      • By hydrolysis of chlorobenzene

      Physical & Chemical Properties

      • Higher boiling points than hydrocarbons
      • Undergo electrophilic substitution: nitration, halogenation

      Acidic Nature

      • More acidic than alcohols due to resonance stabilization of phenoxide ion

      Uses

      • Disinfectants, synthesis of plastics, dyes, and pharmaceuticals

      Ethers

      General Formula and Structure

      • General formula: R–O–R’
      • Straight-chain and cyclic ethers

      Nomenclature

      • Common: Diethyl ether, methyl propyl ether
      • IUPAC: Methoxyethane, ethoxyethane

      Preparation

      • Williamson synthesis: R–X + R’O⁻ → R–O–R’
      • Dehydration of alcohols

      Physical Properties

      • Low boiling point compared to alcohols
      • Slightly polar and used as solvents

      Uses

      • As solvents, anesthetics (diethyl ether), chemical synthesis
      Aldehydes, Ketones, and Carboxylic Acids – ISC Class 12 Chemistry

      Chapter 3: Aldehydes, Ketones, and Carboxylic Acids

      This chapter covers the structure, preparation, properties, and reactions of carbonyl and carboxyl compounds.

      Aldehydes and Ketones

      Nomenclature

      • Aldehydes: Suffix “-al” (e.g., ethanal)
      • Ketones: Suffix “-one” (e.g., propanone)

      Structure

      • Contain carbonyl group (>C=O)
      • Aldehyde: Carbonyl at the end of chain; Ketone: Carbonyl within the chain

      Methods of Preparation

      • Oxidation of alcohols:
        • 1° alcohol → Aldehyde
        • 2° alcohol → Ketone
      • Hydration of alkynes
      • Rosenmund and Stephen’s reactions (for aldehydes)

      Physical Properties

      • Pleasant smell; polar; soluble in water (lower members)
      • Boiling points lower than alcohols but higher than alkanes

      Chemical Properties

      • Nucleophilic Addition:
        • Addition of HCN, NaHSO₃, alcohols, Grignard reagents
        • Mechanism: Attack of nucleophile on electrophilic carbonyl carbon
      • Reduction: Forms alcohols
      • Oxidation:
        • Aldehyde → Carboxylic acid
        • Ketone → No further oxidation under mild conditions
      • Haloform Reaction: Methyl ketones react with halogens + base to give haloform

      Reactivity of Alpha Hydrogen

      • Due to resonance stabilization of enolate ion
      • Undergo aldol condensation: 2 aldehydes/ketones → β-hydroxy aldehyde/ketone

      Uses

      • Solvents, preservatives (formalin), intermediates in synthesis

      Carboxylic Acids

      Classification

      • Monocarboxylic: One –COOH group (e.g., formic acid)
      • Dicarboxylic: Two –COOH groups (e.g., oxalic acid)

      General Formula and Structure

      • General formula: R–COOH
      • Planar carboxyl group, stabilized by resonance

      Nomenclature

      • IUPAC: Add suffix “-oic acid” (e.g., ethanoic acid)
      • Common: Acetic acid, formic acid

      Acidic Nature

      • Stronger acids than alcohols due to resonance stabilization of carboxylate ion
      • Electron-withdrawing groups increase acidity

      Methods of Preparation

      • Oxidation of primary alcohols or aldehydes
      • Hydrolysis of nitriles: R–CN + 2H₂O → R–COOH + NH₃
      • Grignard reagent + CO₂ → Carboxylic acid

      Physical Properties

      • High boiling points due to hydrogen bonding
      • Soluble in water (lower members)

      Chemical Properties

      • Acidic reactions: With metals, carbonates, and bases
      • Decarboxylation: Removal of –COOH group to form hydrocarbons
      • Esterification: R–COOH + R’–OH → Ester + H₂O

      Uses

      • Manufacture of esters, preservatives, pharmaceuticals, vinegar (acetic acid)
      Organic Compounds Containing Nitrogen – ISC Class 12 Chemistry

      Chapter 4: Organic Compounds Containing Nitrogen

      This chapter covers amines, aniline, cyanides, isocyanides, and diazonium salts.

      Amines

      General Formula & Classification

      • General formula: R–NH₂ (primary), R₂NH (secondary), R₃N (tertiary)
      • Aliphatic or aromatic based on the R group

      Structure of Amino Group

      • Trigonal pyramidal geometry
      • Lone pair on nitrogen contributes to basic character

      Nomenclature

      • Common: Methylamine, dimethylamine
      • IUPAC: Methanamine, ethanamine

      Preparation

      • Reduction of nitro compounds
      • Gabriel phthalimide synthesis (for primary amines)
      • Ammonolysis of alkyl halides

      Physical Properties

      • Lower amines are gases/liquids; higher ones are solids
      • Soluble in water due to hydrogen bonding

      Chemical Properties

      • Basic nature: Forms salts with acids
      • Acylation: Forms amides
      • Carbylamine test (for primary amines only)

      Uses

      • Pharmaceuticals, dyes, polymers, surfactants

      Identification of Amines

      • Primary: Positive carbylamine test (offensive smell)
      • Secondary: Forms N-substituted amide with acyl chloride
      • Tertiary: No reaction with carbylamine or acylation

      Aniline

      Preparation

      • By reduction of nitrobenzene using Sn/HCl or Fe/HCl

      Physical Properties

      • Oily liquid, slightly soluble in water
      • Unpleasant odor

      Chemical Properties

      • Undergoes electrophilic substitution: bromination, nitration, sulfonation
      • Acylation and diazotization reactions

      Cyanides and Isocyanides

      Properties

      • Cyanides: Contain –CN group, toxic, polar, moderately soluble
      • Isocyanides: Contain –NC group, extremely foul-smelling

      Uses

      • Used in organic synthesis and coordination chemistry
      • Cyanides used in metallurgy (e.g., gold extraction)

      Diazonium Salts

      Preparation

      • By reaction of primary aromatic amine with nitrous acid (NaNO₂ + HCl) at 0–5°C
      • Example: Aniline + NaNO₂ + HCl → Benzene diazonium chloride

      Reactions

      • Sandmeyer Reaction: Replacement of –N₂⁺ with Cl, Br, CN using Cu(I) salts
      • Gattermann Reaction: Similar to Sandmeyer using Cu powder
      • Coupling Reaction: Forms azo dyes with phenols or aromatic amines
      Biomolecules – ISC Class 12 Chemistry

      Chapter 5: Biomolecules

      Important concepts on carbohydrates, proteins, vitamins, and nucleic acids.

      Carbohydrates

      • Definition: Polyhydroxy aldehydes or ketones and their derivatives.
      • Classification:
        • Aldoses (with aldehyde group)
        • Ketoses (with ketone group)
      • Monosaccharides: Glucose, Fructose
      • Oligosaccharides: Sucrose, Lactose, Maltose
      • Polysaccharides: Starch, Cellulose, Glycogen
      • Importance: Energy source, structural component
      • Tests for glucose and fructose: Benedict’s test, Fehling’s test, Seliwanoff’s test

      Proteins

      • Structural units: Amino acids
      • Basic idea: Amino acids linked by peptide bonds form polypeptides and proteins
      • Protein structure:
        • Primary structure: Sequence of amino acids
        • Secondary structure: α-helix and β-pleated sheets
        • Tertiary structure: 3D folding
        • Quaternary structure: Association of multiple polypeptide chains
      • Denaturation: Loss of structure and function due to heat, pH changes
      • Enzymes: Biological catalysts
      • Hormones: Elementary idea only (chemical messengers)

      Vitamins

      • Classification:
        • Fat-soluble: Vitamins A, D, E, K
        • Water-soluble: Vitamins B-complex, C
      • Deficiency diseases: Rickets (D), Scurvy (C), Beriberi (B1), Night blindness (A)

      Nucleic Acids

      • DNA: Deoxyribonucleic acid, stores genetic information
      • RNA: Ribonucleic acid, involved in protein synthesis
      Most Asked Points – ISC Class 12 Chemistry

      Most Asked Points in ISC Class 12 Chemistry

      1. Electrochemistry

      • Nernst equation: Derivation and applications.
      • Galvanic cells: Cell notation and calculation of emf.
      • Conductance: Molar conductance and its variation with concentration.
      • Electrolysis: Faraday’s laws and calculations.

      2. Solutions

      • Raoult’s law: Ideal solutions and colligative properties.
      • Vapor pressure lowering, boiling point elevation, freezing point depression, osmotic pressure.
      • Calculations: Molality, molarity, mole fraction in colligative property problems.

      3. Chemical Kinetics

      • Rate laws: Zero, first, second order reactions.
      • Half-life: Expressions for different orders.
      • Factors affecting rate: Temperature, concentration, catalyst.

      4. Surface Chemistry

      • Adsorption vs Absorption.
      • Catalysis: Homogeneous and heterogeneous.
      • Colloids: Types, preparation, properties (Tyndall effect, Brownian movement).

      5. The p-Block Elements

      • Group 17: Properties, uses of halogens.
      • Group 18: Inertness and uses of noble gases.
      • Oxides and halides: Acid-base behavior.

      6. The d- and f-Block Elements

      • Transition metals: Variable oxidation states, color, magnetism.
      • Lanthanides and actinides: Inner transition metals and their uses.

      7. Coordination Compounds

      • Ligands: Types (monodentate, bidentate).
      • Coordination number and geometry.
      • Naming rules (IUPAC) and isomerism.
      • Importance and applications.

      8. Organic Chemistry

      • Haloalkanes & Haloarenes: Nucleophilic substitution, elimination.
      • Alcohols, phenols, ethers: Preparation, reactions, acidic nature.
      • Aldehydes & ketones: Nucleophilic addition mechanisms.
      • Carboxylic acids: Acidic nature, preparation.
      • Organic nitrogen compounds: Amines, diazonium salts, cyanides.

      9. Biomolecules

      • Carbohydrates: Classification and tests.
      • Proteins: Structure, denaturation.
      • Vitamins: Types and deficiency diseases.
      • Nucleic acids: DNA & RNA basics.

      10. Practical & Experimental Chemistry

      • Titrations: Acid-base, redox titrations.
      • Qualitative analysis: Tests for cations and anions.
      • Preparation of salts and crystals.

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