Electromagnetic Spectrum – Arrangement of electromagnetic radiations in increasing order of wavelength or decreasing frequency.

Conductance and its measurement– Measure of the ability of a solution to conduct electric current.

Kohlrausch’s Law – Limiting molar conductance is the sum of individual ionic conductances.

Hittorf’s Rule and Transport Number – Fraction of total current carried by a particular ion.

Conductometric Titrations – Quantitative determination of concentration using a standard solution.

Debye–Hückel Theory – Explains deviation of strong electrolytes from ideal behavior.

Activity Coefficient – Factor relating activity to concentration of an ion.

Redox Reactions – Reactions involving simultaneous oxidation and reduction.

Spontaneous and Non spontaneous Reactions

Electrochemical and Electrolytic Cells – Devices converting chemical energy into electrical energy and vice versa.

Electrode Potentials – Measure of tendency of an electrode to lose or gain electrons.

Liquid Junction Potential

Salt Bridge

Electrochemial Series

Nernst Equation – Relates electrode potential to concentration of ions.

Fuel Cells – Electrochemical cells producing electricity directly from fuel oxidation.

Corrosion – Gradual destruction of metals due to electrochemical reactions.

Equation of States

• Van der Waals Equation – Modified gas equation correcting intermolecular forces and volume.

Gas Laws – Mathematical relations between pressure, volume, temperature, and amount of gas.

Ideal and Real Gases – Gases that deviate from ideal behavior at high pressure and low temperature.

Critical Constants and Critical Phenomena – Critical temperature, pressure, and volume defining gas–liquid equilibrium.

Laws of Thermodynamics – Fundamental laws governing energy and entropy changes.

• Zeroth Law of Thermodynamics –
• First Law of Thermodynamics
• Second Law of Thermodynamics
• Third Law of Thermodynamics

Thermochemistry – Study of heat changes accompanying chemical reactions.

Calorimetry – Experimental measurement of heat absorbed or released.

Heat Capacities and Specific Heat – Amount of heat required to raise temperature of a system.

Reversible and Irreversible Processes –

Hess’s Law and Born Haber Cycle – Enthalpy change depends only on initial and final states. Thermochemical cycle for calculating lattice energy of ionic solids.

Entropy (S) and Free Energy and Equilibrim Relation – Measure of randomness or disorder of a system.

Gibbs Free Energy (G) and Gibbs Equation – Criterion for spontaneity of reactions at constant T and P.

Equilibrium Constant – Ratio expressing extent of reaction at equilibrium.

Helmholtz Equation – Relates Helmholtz free energy with work at constant volume.

Fugacity and Activity – Effective pressure accounting for non-ideality of real gases and Effective concentration of a species in non-ideal systems.

Molecularity and Order of Reaction – Number of reacting species involved in an elementary reaction, Sum of powers of concentration terms in the rate law:

Reaction Rates – Change in concentration of reactants or products per unit time.

Order of Reaction –

• Zero Order – Rate independent of reactant concentration.
• First Order – Rate proportional to concentration of one reactant.
• Second Order – Rate proportional to square of concentration or product of two concentrations.
• Third Order – Rate proportional to cube of concentration or combination of three species.

Half-Life – Time required for concentration of a reactant to become half of its initial value.

Methods for Order Determination – Experimental methods to determine reaction order.

Collision Theory – Reaction occurs due to effective collisions with proper energy and orientation.

Transition State Theory – Reaction proceeds through formation of an activated complex.

Activation Energy and Catalysis –

Arrhenius Equation – Relates rate constant with temperature and activation energy.

Complex Reactions – Reactions occurring via more than one elementary step:

• Chain Reactions – Reactions involving initiation, propagation, and termination steps.

• Parallel Reactions – One reactant forms different products simultaneously.
• Series (Consecutive) Reactions – Product of one step becomes reactant of next step.
• Reversible Reactions – Forward and backward reactions occur simultaneously.

Properties of Liquids – Intermolecular forces, viscosity, surface tension, capillarity; physical and chemical properties of surfaces.

Surface Area Determination – Methods to measure surface area of solids, powders, and catalysts.

Adsorption Explained –

• Physical Adsorption (Physisorption) – Weak van der Waals forces, reversible.
• Chemical Adsorption (Chemisorption) – Strong chemical bonds, usually irreversible.

Adsorption Isotherms – Relationship between amount adsorbed and pressure/concentration at constant temperature:

• Langmuir Isotherm – Monolayer adsorption on homogeneous surfaces.
• Freundlich Isotherm – Multilayer adsorption on heterogeneous surfaces.

Colloids – Dispersed systems:

• Properties – Tyndall effect, Brownian motion, stability.
• Classification – Lyophilic, lyophobic, association colloids.
• Preparation – Condensation, dispersion methods.

Surfactants – Surface-active agents; reduce surface tension, form micelles; used in detergents and emulsions.

Gibbs Phase Rule – Gibbs phase rule: F = C – P + 2; applied to:

• One-Component Systems – e.g., water, CO₂.
• Two-Component Systems – e.g., salt-water, alloys.

Catalysis Explained Thoroughly – Process speeding up reactions:

• Enzyme Catalysis – Biological catalysts; substrate binding, active site, specificity.
• Homogeneous Catalysis – Catalyst in same phase as reactants.
• Heterogeneous Catalysis – Catalyst in different phase; adsorption important.
• Acid-Base Catalysis – Proton donor/acceptor accelerates reaction.

Types of Chemical Bonding – Forces that hold atoms together:

• Ionic Bonding – Electrostatic attraction between oppositely charged ions.
• Covalent Bonding – Sharing of electron pairs between atoms.
• Metallic Bonding – Delocalized electrons holding metal cations together.
• Coordinate (Dative) Bonding – Both bonding electrons provided by one atom.
• Hydrogen Bonding – Attraction between hydrogen and electronegative atoms.
• Van der Waals Forces – Weak intermolecular forces (dipole-dipole, London, hydrogen bonding).

Valence Bond Theory (VBT) – Explains bonding through overlap of atomic orbitals.

Hybridization – Mixing of atomic orbitals to form equivalent hybrid orbitals.

Resonance – Representation of a molecule by two or more contributing structures.

VSEPR Theory – Predicts molecular shape based on electron pair repulsions.

Molecular Orbital Theory (MOT) – Describes bonding using molecular orbitals:

• Diatomic Molecules – Application of MOT to homonuclear and heteronuclear diatomics.

Delocalized Bonding – Bonding involving electrons spread over several atoms.

Electron-Deficient Bonding – Bonding with fewer electrons than required by classical theory.

p-Block Elements – Elements of Groups 13–18 and their physical and chemical properties.

Chemical Equilibrium – State where forward and reverse reaction rates are equal.

Acid–Base Theories – Concepts explaining acid and base behavior:

• Arrhenius Theory – Acids produce H⁺, bases produce OH⁻ in water.
• Brønsted–Lowry Theory – Acids donate protons, bases accept protons.
• Lewis Theory – Acids accept electron pairs, bases donate electron pairs.
• SHAB Concept – Hard and soft acids and bases classification.

Acid/Base Strength – Relative tendency to donate or accept protons.

pH – Measure of hydrogen ion concentration in solution.

Significance of pKw, pKa & pKb – Negative logarithms of acid and base dissociation constants.

Buffers – Solutions resisting change in pH upon addition of acid or base.

Indicators – Substances showing color change at specific pH range.

Solubility – Maximum amount of solute that dissolves in a solvent.

Solubility Product (Ksp) – Equilibrium constant for dissolution of sparingly soluble salts.

Common Ion Effect – Decrease in solubility of a salt due to presence of a common ion.

Characteristics of d- and f-Block Elements – General properties, trends, and reactivity.

Coordination Chemistry – Study of complexes formed by central metal ions with ligands:

• Nomenclature – IUPAC rules for naming coordination compounds.
• Structures & Coordination Number (2–10) – Geometry of complexes based on CN.
• Chelates & Chelate Effect – Multidentate ligands forming stable ring complexes.

Theories of Coordination Compounds – Explaining bonding and structure:

• Werner’s Theory – Primary and secondary valency concept.
• Valence Bond Theory (VBT) – Hybridization and bonding in complexes.
• Crystal Field Theory (CFT) – Splitting of d-orbitals in ligand field.
• Molecular Orbital Theory (MOT) – Molecular orbital approach for complexes.

Jahn–Teller Effect – Distortion in certain octahedral complexes to lower energy.

Magnetic & Spectral Properties – Paramagnetism, diamagnetism, electronic spectra.

Isomerism – Structural and stereoisomerism in complexes.

Lanthanides – General characteristics, oxidation states, lanthanide contraction, separation.

Actinides – General properties, oxidation states, electronic configuration, half-life & decay laws.

Bonding – Types of chemical bonds and interactions between atoms.

Hybridization – Mixing of atomic orbitals to form equivalent hybrid orbitals.

Inductive Effect – Electron-withdrawing or donating effect transmitted through sigma bonds.

Resonance – Delocalization of electrons represented by multiple contributing structures.

Hyperconjugation – Delocalization of electrons through sigma bonds adjacent to π systems or carbocations.

Dipole Moment – Measure of polarity of a molecule due to separation of charges.

Alkanes – Saturated hydrocarbons:

• Nomenclature – IUPAC rules for naming alkanes.
• Physical Properties – Boiling/melting points, solubility, density trends.
• Preparation – Methods like Wurtz reaction, hydrogenation of alkenes, halide reduction.
• Reactions – Combustion, free radical halogenation, cracking, isomerization.

Alkenes – Unsaturated hydrocarbons with C=C:

• Nomenclature – IUPAC rules for naming alkenes.
• Physical Properties – Boiling/melting points, solubility, density trends.
• Preparation – Dehydration of alcohols, dehydrohalogenation, cracking of alkanes.
• Reactions – Electrophilic addition, hydrogenation, polymerization, oxidation.

Alkynes – Unsaturated hydrocarbons with C≡C:

Reactions – Electrophilic addition, hydration, halogenation, oxidation, polymerization.

Nomenclature – IUPAC rules for naming alkynes.

Physical Properties – Boiling/melting points, acidity of terminal H, solubility trends.

Preparation – Dehalogenation of dihalides, acetylene from CaC₂, alkyne synthesis via alkylation.

Benzene Structure – Planar cyclic molecule with delocalized π-electrons forming a stable ring.

Aromaticity – Concept of extra stability in cyclic, conjugated, planar molecules following Hückel’s rule (4n+2 π electrons).

Electrophilic Substitution – Reactions where a hydrogen atom is replaced by an electrophile (e.g., nitration, halogenation, sulfonation).

Substituent Effects – Influence of existing groups on reactivity and orientation:

Deactivating Groups – Decrease reactivity, meta-directing.

Activating Groups – Increase reactivity, ortho/para-directing.

Alcohols – Compounds with –OH group:

• Physical & Chemical Properties, Applications – Hydrogen bonding, high boiling points, oxidation, dehydration, esterification; used as solvents, fuels, intermediates.
• Preparation – Hydration of alkenes, reduction of carbonyls, fermentation of sugars.
• Reactions – Oxidation (to aldehydes/ketones/acids), dehydration (to alkenes), esterification; mechanism included.

2. Phenols – Aromatic –OH compounds:

• Physical & Chemical Properties, Applications – Acidic, hydrogen bonding, undergo electrophilic substitution; used in antiseptics, dyes.
• Preparation – From aromatic halides, cumene process, diazonium salts.
• Reactions – Electrophilic substitution (halogenation, nitration), oxidation, esterification; mechanisms.

3. Ethers – R–O–R’ compounds:

• Physical & Chemical Properties, Applications – Low polarity, stable under neutral conditions, cleavage with HX; used as solvents.
• Preparation – Williamson synthesis, dehydration of alcohols.
• Reactions – Cleavage with hydrogen halides, limited reactivity.

4. Amines – Nitrogen-containing compounds (R–NH₂, R₂NH, R₃N):

• Physical & Chemical Properties, Applications – Basic, hydrogen bonding, form amides, used in dyes and pharmaceuticals.
• Preparation – Reduction of nitro compounds, amides, alkyl halides.
• Reactions – Acylation, alkylation, diazotization; reaction mechanisms.

5. Alkyl Halides – R–X compounds:

• Physical & Chemical Properties, Applications – Polar, higher boiling points, intermediates in synthesis.
• Preparation – Halogenation of alkanes, alcohol conversion.
• Reactions – Nucleophilic substitution (SN1/SN2), elimination (E1/E2).

6. Grignard Reagent – RMgX compounds:

• Physical & Chemical Properties, Applications – Highly nucleophilic/basic; used for forming alcohols, acids, ketones.
• Preparation – Reaction of alkyl/aryl halides with Mg in dry ether.
• Reactions – Nucleophilic addition to carbonyl compounds; mechanism.

7. Carbonyl Compounds (Aldehydes & Ketones):

• Physical & Chemical Properties, Applications – Polar, higher boiling points, undergo nucleophilic addition; industrial and lab applications.
• Preparation – Oxidation of alcohols, ozonolysis, Friedel–Crafts acylation.
• Reactions – Nucleophilic addition, condensation, oxidation/reduction; mechanisms.

8. Carboxylic Acids & Derivatives (R–COOH, R–COX, R–COOR, R–CONH₂):

• Physical & Chemical Properties, Applications – Hydrogen bonding, acidic, high boiling points; used in ester, amide synthesis.
• Preparation – Oxidation of aldehydes/alcohols, hydrolysis of nitriles/esters.
• Reactions – Formation of acid halides, anhydrides, esters, amides via nucleophilic acyl substitution; mechanisms included.

Nucleophilic Substitution (S) – Replacement of leaving group by nucleophile (SN1 & SN2); used in synthesis of alcohols, amines.

Elimination (E) – Removal of HX to form alkenes (E1 & E2); depends on substrate, base, and conditions.

Zaitsev Rule – More substituted alkene is major product in elimination.

Hofmann Rule – Less substituted alkene is major when bulky base is used.

Competition – Substitution vs elimination depends on nucleophile/base strength, solvent, temperature.

UV/Visible Spectroscopy – Study of electronic transitions; used for conjugated systems, determination of λmax.

IR Spectroscopy – Study of vibrational transitions; identifies functional groups via characteristic absorption.

¹H NMR Spectroscopy – Study of hydrogen nuclei environments; provides information on structure, number of protons, and splitting patterns.

Mass Spectrometry – Study of molecular mass and fragmentation; used for molecular weight determination and structure elucidation.