NEET 2025 • Class XI Chemistry

NCERT 2025–26: Class 11 Chemistry Hub — Notes, Figures, Summaries, Quizzes & Downloads

This hub covers the nine NCERT units for Class 11 Chemistry — from Some Basic Concepts of Chemistry to Hydrocarbons — with chapters presented in a consistent, exam-ready format. Each chapter block follows the same layout: Notes → Figures → Quick Summary → 10-MCQ Quiz → Downloads. Use the sticky contents at left to jump between chapters, and the search box to filter chapters live on this page.

Syllabus verified • Updated: 04 Sep 2025

Unit Overview

This page is your master table of contents for NCERT Class XI Chemistry (2025–26). Units are ordered as in CBSE’s official syllabus. Each chapter block is self-contained for teaching and revision: topic-wise notes with definitions and logic, diagram callouts, a quick summary for last-minute revision, a 10-MCQ quiz (NEET pattern), and downloads (revision sheets, tables, mechanisms). Progress chips flag weightage (High/Medium), diagram-heavy topics, and PYQ density to help prioritise.

High-yield PYQs Diagram & mechanism focus 10-MCQ quizzes Quick revision sheets

01 Some Basic Concepts of Chemistry

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Why this chapter matters: It teaches the “language” of chemistry—mole, mass, formula, and simple laws—so you can solve most NEET numericals confidently.
  • Laws of Chemical Combination (simple ideas):
    • Conservation of mass: Total mass stays the same during a reaction.
    • Definite proportions: A pure compound always has the same fixed ratio of elements.
    • Multiple proportions: Same elements can form different compounds in small whole-number mass ratios.
    • Combining volumes: Gases react in simple volume ratios (same T, P).
    • Avogadro’s law: Equal volumes of gases (same T, P) have equal number of molecules.
  • Atomic & Molar Masses:
    • Relative atomic mass (Ar): Based on the 12C scale.
    • Average atomic mass: Use isotopic abundance × isotopic mass, then add up.
    • Molar mass (M): Mass of 1 mole, in g·mol−1. Formula/molecular mass = sum of atomic masses.
  • Mole Concept (quick links):
    n (mol) = mass ÷ M = particles ÷ NA = gas volume at STP ÷ 22.4 L.
    NA = 6.022 × 1023 (entities per mole).
  • Empirical vs Molecular Formula (step idea):
    • From % composition → convert to moles → divide by the smallest → get simplest ratio = empirical formula.
    • Molecular formula = (empirical formula) × n, where n = molar mass ÷ empirical formula mass.
  • Stoichiometry (how to plan a sum):
    • Write a balanced equation → convert given data to moles → use mole ratio → convert to required unit.
    • Limiting reagent decides how much product forms; the other is in excess.
    • % yield = (actual ÷ theoretical) × 100.
  • Concentration Units (know the use):
    • Mass % = (mass solute ÷ mass solution) × 100; ppm = parts per million.
    • Molarity (M) = moles solute ÷ L solution (changes with T).
    • Molality (m) = moles solute ÷ kg solvent (does not change with T).
    • Mole fraction Xi = ni/Σn (all X add to 1).
  • Gas volumes (for quick estimates): At STP (273 K, 1 atm), 1 mol ≈ 22.4 L; at 298 K, 1 atm, ≈ 24.5 L (near-ideal).
  • Significant figures (fast rules):
    • Leading zeros don’t count; trapped zeros do; trailing zeros count only if a decimal is shown.
    • ×/÷ → keep least sig figs; +/− → keep least decimal places.
  • Check with units (dimensional sense): Before final answer, check if your units match what the question asked.
  • NEET Focus: mole↔mass↔volume conversions, limiting reagent, % composition → empirical/molecular formula, choosing correct concentration unit, and sig-fig discipline.
Important Figures
Key laws: conservation, definite and multiple proportions, combining volumes, Avogadro
Key laws at a glance (what each law says and where it helps).
Flowchart linking mass, moles, particles, and gas volumes
Mole roadmap: mass ↔ moles ↔ particles/volume (STP).
Steps from percent composition to empirical and molecular formula
From % composition to final formula: the quick path.
Decision tree to find the limiting reagent
Limiting reagent: a simple decision tree.
Comparison of molarity, molality, mole fraction, mass percent
Which concentration unit to use and why.
Quick Summary

Use moles to move between mass, particles, and gas volume. Laws of combination explain how formulas work. Find empirical formula from percentages, then scale to molecular formula using molar mass. In calculations, identify the limiting reagent and report answers with correct units and significant figures.

Practice Quiz (10 MCQs)
Downloads
Stoichiometry Formula Sheet (PDF) Sig-figs & Rounding Guide (PDF) Empirical vs Molecular Flowchart (PNG)
Cross-links

Next topics to reinforce these skills: Structure of Atom · Periodicity · States of Matter

02 Structure of Atom

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Discovery of subatomic particles: Cathode ray tube → electron (Thomson); Gold foil experiment → nucleus (Rutherford); discovery of neutron (Chadwick).
  • Atomic models:
    • Thomson: Plum pudding model, electrons in a positively charged sphere.
    • Rutherford: Nucleus-centric model; electron cloud, mostly empty space.
    • Bohr model: Quantized energy levels, electrons revolve in fixed orbits; energy emitted/absorbed as photons (ΔE = hν).
  • Key constants: Planck’s constant (h), Rydberg constant (RH), speed of light (c).
  • Dual nature of matter and light: Wave–particle duality; De Broglie wavelength λ = h/mv; Photoelectric effect explained by Einstein.
  • Quantum mechanical model: Schrödinger equation → orbitals (ψ); Heisenberg uncertainty principle (Δx·Δp ≥ h/4π).
  • Quantum numbers: n (shell), l (subshell), ml (orientation), ms (spin); shapes: s (sphere), p (dumbbell), d (clover).
  • Electronic configuration rules: Aufbau principle, Pauli exclusion, Hund’s rule; order of orbital filling (n+l rule).
  • NEET Focus: Numerical on λ, frequency, energy; Bohr radius and energy levels; orbital diagrams; quantum numbers; exceptional configurations.
Important Figures
Diagram of Rutherford gold foil experiment
Rutherford’s gold foil experiment setup and findings.
Bohr model showing electron orbits
Bohr’s quantized orbit model of hydrogen atom.
Shapes of s, p, and d orbitals
Shapes of s, p, and d atomic orbitals.
Quick Summary

The structure of atom evolved from Thomson’s plum pudding to Bohr’s quantized orbit model to the modern quantum mechanical model. Light and matter show dual nature, and electrons are described by wavefunctions (ψ) and quantum numbers. Master numerical problems on energy levels, λ, frequency, and orbital filling patterns.

Practice Quiz (10 MCQs)
Downloads
Bohr Model & Energy Levels (PDF) Quantum Numbers Table (PNG) Orbital Shapes Diagram Sheet (PDF)
Cross-links

Next topics to explore: Classification & Periodicity · Chemical Bonding

03 Classification of Elements and Periodicity in Properties

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Why we classify: Grouping elements by repeating (periodic) properties makes prediction easy—valency, reactivity, and trends become systematic.
  • From early attempts to modern table:
    • Dobereiner triads → groups of three with middle atomic mass ≈ average of the other two.
    • Newlands octaves → every 8th element similar; failed after Ca.
    • Mendeleev → arranged by atomic mass, left gaps, predicted new elements.
    • Modern periodic law → properties are periodic functions of atomic number (Z).
  • Layout (long form): 7 periods (rows), 18 groups (columns); s, p, d, f blocks based on the subshell of the differentiating electron.
  • Effective nuclear charge (Zeff) & shielding: Across a period, Zeff ↑ (poor shielding by same-shell electrons) → radius ↓, IE ↑, EN ↑. Down a group, added shells ↑ shielding → radius ↑, IE ↓, EN ↓.
  • Key periodic trends (general):
    • Atomic/ionic radius: decreases across, increases down. For isoelectronic species: higher Z → smaller radius.
    • Ionization enthalpy (IE): increases across, decreases down; notable dips: B > Be? (actually IE: Be > B), N > O anomaly.
    • Electron gain enthalpy (EGE): generally becomes more negative across; exceptions: noble gases (+), Be/Mg/N somewhat less negative.
    • Electronegativity (EN): increases across, decreases down; F is highest.
    • Metallic → non-metallic character: decreases across; increases down.
    • Valency/oxidation state: related to ns and np electrons; typical valencies repeat periodically.
  • Special patterns:
    • Diagonal relationship: Li–Mg, Be–Al show similarities (size/charge density).
    • Inert pair effect (intro idea): Heavier p-block elements may show lower oxidation states (e.g., Tl(+1), Pb(+2)).
    • Transition elements: d-block shows smaller variation across a period and multiple oxidation states (detail in XII/other units).
  • NEET Focus: Order questions (radius, IE, EN), isoelectronic size, identifying anomalies (Be/B, N/O), predicting valency/character, locating elements from configuration.
Important Figures
Annotated long-form periodic table with blocks and group numbers
Long-form periodic table with blocks and group numbering.
Trend arrows for atomic radius, ionization enthalpy, electron gain enthalpy, electronegativity
Direction of major trends (radius, IE, EGE, EN).
Schematic of shielding and effective nuclear charge across periods
Shielding vs Zeff: why trends occur.
Quick Summary

Properties repeat with atomic number. Across a period, Zeff increases → radius decreases → IE and EN increase; down a group, added shells reverse these trends. Watch typical exceptions (Be/B, N/O; noble gases in EGE). Use isoelectronic logic and position in the table to predict size, reactivity, and valency quickly.

Practice Quiz (10 MCQs)
Downloads
Periodic Trends Summary Sheet (PDF) Common Trend Exceptions (PNG) Isoelectronic Series Practice (PDF)
Cross-links

Continue with: Chemical Bonding & Molecular Structure · Structure of Atom (revision)

04 Chemical Bonding and Molecular Structure

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Why bonding occurs: Atoms combine to achieve stable electronic configurations (octet rule, duet rule for H/He).
  • Main types of bonds:
    • Ionic bond: Complete transfer of electrons; strong electrostatic forces; lattice energy concept (Born–Haber cycle).
    • Covalent bond: Shared electron pairs; single, double, triple bonds.
    • Coordinate (dative) bond: Both electrons from one atom.
    • Hydrogen bond: Strong dipole–dipole interaction; intermolecular and intramolecular.
    • Van der Waals forces: London dispersion, dipole–dipole interactions.
  • Theories of bonding:
    • Valence Shell Electron Pair Repulsion (VSEPR): Predicts shape by minimizing electron-pair repulsion.
    • Valence Bond Theory (VBT): Bond forms by overlap of half-filled orbitals; hybridization explains geometry (sp, sp², sp³, etc.).
    • Molecular Orbital (MO) Theory: Atomic orbitals combine to form bonding and antibonding orbitals; bond order predicts stability.
  • Important shapes and examples:
    • Linear (BeCl₂), trigonal planar (BF₃), tetrahedral (CH₄), trigonal bipyramidal (PCl₅), octahedral (SF₆), bent (H₂O), trigonal pyramidal (NH₃).
  • Formal charge: FC = valence electrons − (lone pair e⁻ + ½ bonding e⁻).
  • Bond characteristics: Bond length, bond energy, bond angle; bond order = ½(bonding − antibonding e⁻).
  • Resonance: Delocalization; real structure is a resonance hybrid (e.g., O₃, CO₃²⁻).
  • NEET Focus: Shapes and hybridization, bond order, MO diagrams of H₂, He₂, B₂, O₂, N₂; identification of paramagnetic/diamagnetic; hydrogen bonding trends; exceptions to octet rule.
Important Figures
VSEPR model shapes for different electron pairs
VSEPR shapes: predicting geometry from electron pairs.
Hybridization scheme chart for common molecules
Hybridization chart with geometry examples.
MO energy level diagrams for diatomic molecules
MO diagrams for homonuclear diatomics.
Quick Summary

Bonding arises from electron sharing or transfer to achieve stability. Use VSEPR to predict shapes, VBT for overlap and hybridization, and MO theory for bond order and magnetism. Hydrogen bonding and resonance explain unusual properties. Master bond order and shape questions—they are NEET favorites.

Practice Quiz (10 MCQs)
Downloads
VSEPR Shapes & Hybridization Chart (PDF) MO Diagrams Practice Sheet (PNG) Bond Order & Paramagnetism Table (PDF)
Cross-links

Explore next: Chemical Thermodynamics · Periodicity (Revision)

05 Chemical Thermodynamics

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Basic terms: System (part studied) vs surroundings; types—open/closed/isolated. State (P, V, T, n), state function (H, U, S, G), path function (q, w). Extensive (depend on amount) vs intensive (don’t).
  • First Law (energy conservation): ΔU = q + w (sign convention: q to system positive; w on system positive). PV work at constant external pressure: w = −PextΔV.
  • Enthalpy (H): H = U + pV; at constant pressure, ΔH = qp. For ideal gas: ΔH = ΔU + ΔngRT. Heat capacities: C = dq/dT; Cp−Cv = nR (ideal gas).
  • Standard Enthalpies: formation (ΔH°f), combustion, atomization, solution, neutralization. Hess’s Law: total enthalpy change is path-independent (add/subtract equations to get ΔH).
  • Bond enthalpy (gas phase): ΔH ≈ ΣD(reactants bonds broken) − ΣD(products bonds formed).
  • Spontaneity & Second Law: Spontaneous processes overall increase entropy of universe (ΔSuniv > 0). Entropy (S) measures dispersal: ΔS > 0 for melting, vaporization, mixing (generally).
  • Gibbs Free Energy: G = H − TS; criterion at constant T, p: ΔG < 0 ⇒ spontaneous, ΔG = 0 at equilibrium, ΔG > 0 non-spontaneous. Relation: ΔG = ΔH − TΔS.
  • Temperature effect: Endothermic (ΔH > 0) may be spontaneous if TΔS term dominates; exothermic (ΔH < 0) usually favorable.
  • Isothermal reversible expansion (ideal gas): wrev = −nRT ln(V2/V1) = −nRT ln(p1/p2). For ideal gas isothermal: ΔU = 0, so q = −w.
  • Kirchhoff’s idea (qualitative): ΔH changes with temperature via heat capacities (detail beyond NEET numericals; remember trend only).
  • NEET Focus: ΔU/ΔH sign and units, w at constant p, Hess’s law problems, bond enthalpy estimates, ΔG = ΔH − TΔS reasoning, Cp−Cv = nR, phase-change enthalpies, isothermal work formulae.
Important Figures
Diagrams of open, closed, isolated systems with boundaries
Systems and boundaries: open, closed, isolated.
Hess's law enthalpy cycle box method
Hess’s law: adding reactions to compute ΔH.
Gibbs free energy versus reaction coordinate with ΔG negative
ΔG landscape for spontaneity and equilibrium.
Quick Summary

Thermodynamics links heat, work, and energy. Use ΔU = q + w and ΔH = qp for calculations; apply Hess’s law or bond enthalpies for ΔH. Spontaneity at constant T, p depends on ΔG = ΔH − TΔS. Remember Cp−Cv = nR (ideal gas) and w = −PextΔV (or −nRT ln V₂/V₁ for reversible isothermal ideal gas).

Practice Quiz (10 MCQs)
Downloads
Thermodynamics Formula Sheet (PDF) Hess’s Law Worked Examples (PNG) ΔG Decision Chart (PDF)
Cross-links

Continue with: Equilibrium · Chemical Bonding (revision)

06 Equilibrium

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Dynamic equilibrium: In reversible reactions, forward and reverse rates are equal; concentrations remain constant but reactions continue microscopically.
  • Equilibrium constant (K):
  • Le Chatelier’s principle: System shifts to counteract change in concentration, temperature, or pressure; catalyst affects rate but not equilibrium position.
  • Factors affecting K: Only temperature affects K; concentration and pressure changes cause shifts but do not change K.
  • Acids and bases:
    • Arrhenius: H⁺ donor, OH⁻ donor; Brønsted–Lowry: proton donor/acceptor; Lewis: electron pair acceptor/donor.
    • Water is amphoteric; conjugate acid–base pairs.
  • pH and pOH: pH = −log[H⁺]; pOH = −log[OH⁻]; pH + pOH = 14 (at 25°C).
  • Ka and Kb: Strength of acids and bases; pKa = −log Ka. Strong acids/bases: complete dissociation, very large K.
  • Kw: Ionic product of water at 25°C: 1.0×10⁻¹⁴.
  • Salt hydrolysis and buffer solutions: pH calculation using Henderson–Hasselbalch equation: pH = pKa + log([A⁻]/[HA]).
  • Solubility product (Ksp): Predict precipitation; for salt AB, Ksp = [A⁺][B⁻].
  • Common ion effect: Solubility decreases with addition of a common ion; used in qualitative analysis.
  • NEET Focus: pH of weak acids/bases, buffer calculation, predicting precipitation, Le Chatelier’s shifts, Kp–Kc conversions, qualitative acid–base theory questions.
Important Figures
Graph showing forward and reverse rate equality at equilibrium
Rates of forward and reverse reactions at equilibrium.
Effect of concentration, temperature, and pressure changes
Le Chatelier’s principle illustrated with shifts in equilibrium.
pH change in buffer solution on acid/base addition
Buffer action and pH stabilization curve.
Quick Summary

Equilibrium is dynamic, with no net concentration change at equal rates. The equilibrium constant predicts reaction extent, and Le Chatelier’s principle predicts direction of shifts. Strong acids/bases fully dissociate; weak ones use Ka/Kb. Buffers resist pH change; solubility product predicts precipitation. Focus on Kp–Kc relations, pH calculations, and common ion effect.

Practice Quiz (10 MCQs)
Downloads
Equilibrium Constants Quick Sheet (PDF) Buffer Calculations Guide (PNG) Solubility Product Practice Problems (PDF)
Cross-links

Next topics: Redox Reactions · Thermodynamics (revision)

07 Redox Reactions

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Idea of redox: Redox = reduction + oxidation. Oxidation: loss of e⁻ / increase in oxidation number (O.N.); Reduction: gain of e⁻ / decrease in O.N.
  • Oxidation number (O.N.) rules (quick):
    • Free element = 0; monoatomic ion = its charge.
    • F = −1; O usually = −2 (peroxides: −1; OF₂: +2); H = +1 (in hydrides: −1).
    • Sum of O.N. in neutral compound = 0; in ion = charge on ion.
  • Oxidizing vs reducing agents: Oxidizing agent (OA) gets reduced (O.N. ↓); Reducing agent (RA) gets oxidized (O.N. ↑). Identify species with O.N. change.
  • Types of redox:
    • Combination (A + B → AB), decomposition, displacement (metal/halogen), disproportionation (same element: both oxidized and reduced), comproportionation (opposite of disproportionation).
  • Balancing redox equations:
    • Oxidation number method (molecular medium): equate total increase = total decrease in O.N., then balance O, H.
    • Ion–electron (half-reaction) method:
      • Acidic medium: balance O with H₂O, H with H⁺, then add e⁻; equalize e⁻ and add half-reactions.
      • Basic medium: first as acidic, then add OH⁻ to neutralize H⁺ → form H₂O; simplify waters.
  • n-factor & equivalents (for titrations):
    • n-factor = number of electrons lost/gained per formula unit in the reaction context.
    • Eq wt = molar mass / n-factor; N₁V₁ = N₂V₂ for redox titrations.
    • Common n-factors (in given medium): KMnO₄: +7→+2 in acid (n = 5); +7→+4 in neutral/alkaline (n = 3). Fe²⁺→Fe³⁺ (n = 1). C₂O₄²⁻→CO₂ (n = 2 per oxalate).
  • Redox titration basics: KMnO₄ (self-indicator) and K₂Cr₂O₇ often used; indicators: diphenylamine/ferroin for dichromate; end-point color change cues are standard.
  • NEET Focus: Calculate O.N., identify OA/RA, spot disproportionation, balance in acidic/basic media, compute n-factor and use N₁V₁=N₂V₂, quick recognition of medium-dependent products of common oxidants.
Important Figures
Table of oxidation number rules with common exceptions
Oxidation number rules and exceptions at a glance.
Flowchart for half-reaction method in acidic and basic media
Half-reaction method: acidic vs basic medium flowchart.
KMnO4 titration setup and color change at end point
KMnO₄ redox titration: setup and end-point color cue.
Quick Summary

Redox reactions are electron-transfer processes recognized by changes in oxidation numbers. Use OA/RA logic, pick a balancing method (oxidation number or half-reaction), and remember medium effects. For titrations, compute n-factor correctly and use N₁V₁ = N₂V₂. Disproportionation and common oxidants (KMnO₄/K₂Cr₂O₇) are frequent NEET checks.

Practice Quiz (10 MCQs)
Downloads
Redox Balancing (Acidic/Basic) Flowchart (PDF) Oxidation Number Rules Cheat Sheet (PNG) Redox Titration n-factor Drill (PDF)
Cross-links

Continue with: Organic Chemistry — Basics & Techniques · Equilibrium (revision)

08 Organic Chemistry — Some Basic Principles and Techniques

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Introduction: Organic compounds mainly contain C–C and C–H bonds; carbon’s tetravalency and catenation enable huge structural diversity.
  • Classification:
    • Acyclic (open-chain): straight/branched chains (alkanes, alkenes, alkynes).
    • Cyclic: alicyclic, aromatic, heterocyclic.
    • Based on functional groups: alcohols, aldehydes, ketones, acids, amines, etc.
  • Nomenclature (IUPAC):
    • Select longest chain → locate substituents → assign numbers → name substituents alphabetically → suffix indicates main functional group.
    • Examples: CH₃CH₂OH: ethanol; CH₃COOH: ethanoic acid.
  • Isomerism:
    • Structural: chain, position, functional group, tautomerism.
    • Stereoisomerism: geometrical (cis–trans), optical (chiral centers, enantiomers).
  • Hybridization & bonding in organic molecules:
    • sp³: tetrahedral (109.5°); sp²: trigonal planar (120°); sp: linear (180°).
    • σ and π bonds; delocalization in conjugated systems.
  • Electronic effects:
    • Inductive effect (I): electron-withdrawing/donating through σ bonds.
    • Resonance (R): delocalization of π-electrons; represented by resonance structures.
    • Hyperconjugation: delocalization of σ-bond electrons of C–H adjacent to a double bond or carbocation.
    • Electromeric effect: temporary shift of π-electrons under attacking reagent (conceptual).
  • Reactive intermediates:
    • Carbocations (sp², +ve), carbanions (sp³, −ve), free radicals (sp², unpaired e⁻).
    • Stability order: resonance > hyperconjugation > inductive effects.
  • Organic reactions: Substitution (SN1, SN2), addition, elimination (E1, E2), rearrangements; electrophile vs nucleophile.
  • Purification techniques:
    • Distillation (simple, fractional, steam), sublimation, crystallization, chromatography (adsorption, TLC, column), differential extraction.
  • Qualitative analysis: Lassaigne’s test for N, S, halogens; Baeyer's test for unsaturation; Tollen’s/Fehling’s for aldehydes.
  • Quantitative analysis: Estimation of C, H by combustion; N by Kjeldahl; halogens by Carius method.
  • NEET Focus: IUPAC naming, electronic effects, resonance/hyperconjugation, stability of intermediates, test reagents, chromatography principles, functional group conversions.
Important Figures
Stepwise flowchart for IUPAC naming
IUPAC naming steps with functional group priority.
Illustration of inductive, resonance, and hyperconjugation effects
Major electronic effects and their symbols (+I, −I, +R, −R).
Diagram of chromatography separation
Column/TLC chromatography principle of separation.
Quick Summary

Organic chemistry builds on carbon’s tetravalency and catenation. Master IUPAC rules, structural/optical isomerism, and the impact of inductive, resonance, and hyperconjugation effects on reactivity. Know purification and analysis methods (chromatography, Kjeldahl, Carius) and basic mechanisms (SN1, SN2, E1, E2).

Practice Quiz (10 MCQs)
Downloads
IUPAC Naming Practice Sheet (PDF) Electronic Effects Cheatsheet (PNG) Purification Techniques Flowchart (PDF)
Cross-links

Continue with: Hydrocarbons · Redox Reactions (revision)

09 Hydrocarbons

Summative Weightage: High Updated: 04 Sep 2025
Chapter Notes
  • Definition: Compounds made up of only carbon and hydrogen; classified as alkanes, alkenes, alkynes, and aromatic hydrocarbons.
  • Alkanes (CnH2n+2):
    • Saturated hydrocarbons, sp³ hybridized C atoms, tetrahedral geometry.
    • Reactions: free radical halogenation, combustion, isomerization, cracking.
  • Alkenes (CnH2n):
    • Unsaturated with one double bond, sp² hybridization.
    • Addition reactions: hydrogenation, halogenation, hydrohalogenation (Markovnikov’s rule, peroxide effect for HBr), polymerization.
    • Oxidation: Baeyer's test (decolorizes KMnO₄); epoxidation; ozonolysis gives carbonyl compounds.
  • Alkynes (CnH2n-2):
    • Triple bond (sp hybridized), linear geometry; acidic hydrogen in terminal alkynes.
    • Addition of halogens, hydrogen halides, water (Hg²⁺ catalysis → ketones).
  • Aromatic hydrocarbons:
    • Benzene and homologues; cyclic, planar, conjugated, 4n+2 π-electrons (Hückel's rule).
    • Electrophilic substitution: nitration, halogenation, sulphonation, Friedel–Crafts alkylation/acylation; directive effects of substituents.
  • Source and importance: Petroleum refining; cracking, reforming; LPG, CNG as fuels; industrial feedstock for plastics, dyes, drugs.
  • Tests: Bromine water/KMnO₄ test for unsaturation; aromatic compounds often resistant to addition.
  • Concept focus:
    • Stability order: alkynes < alkenes < alkanes < aromatics (resonance energy in benzene).
    • Resonance and delocalization explain aromatic reactivity and substitution patterns.
  • NEET Focus: Markovnikov and anti-Markovnikov addition, ozonolysis interpretation, aromatic directive effects, stability trends, and tests for unsaturation.
Important Figures
Free radical substitution and halogenation of alkanes
Alkane halogenation mechanism via free radical substitution.
Markovnikov addition in alkenes and peroxide effect
Markovnikov’s rule and peroxide effect in alkene reactions.
Electrophilic aromatic substitution mechanism in benzene
Benzene electrophilic substitution and resonance stabilization.
Quick Summary

Hydrocarbons form the base of organic chemistry. Alkanes are saturated and show substitution; alkenes/alkynes undergo addition, and aromatics show electrophilic substitution due to delocalization. Learn key mechanisms, directing effects, Markovnikov's and peroxide rules, and methods like ozonolysis for structural analysis.

Practice Quiz (10 MCQs)
Downloads
Hydrocarbon Reactions Sheet (PDF) Directive Effects Chart (PNG) Ozonolysis and Markovnikov Rules Guide (PDF)
Cross-links

Next topics: Environmental Chemistry · Organic Basics & Techniques (revision)