Full Syllabus Cheat Sheet Chemistry Class 11 - Chemistry Class 11 - NEET

1. Basic Concepts of Chemistry

1.1 Fundamental Laws

1.2 Mole Concept

2. Structure of Atom

2.1 Discovery of Subatomic Particles

2.2 Atomic Models

2.3 Quantum Mechanical Model

2.4 Quantum Numbers

2.5 Electronic Configuration Rules

• Aufbau Principle: Electrons fill orbitals from lowest to highest energy (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...)
• Pauli Exclusion Principle: No two electrons can have same set of all four quantum numbers
• Hund's Rule: Electrons fill degenerate orbitals singly with parallel spins before pairing
• Orbital Capacity: s=2, p=6, d=10, f=14 electrons

3. Classification and Periodicity

3.1 Modern Periodic Law

• Properties of elements are periodic functions of their atomic numbers
• 18 groups (vertical), 7 periods (horizontal)
• s-block (Groups 1-2), p-block (Groups 13-18), d-block (Groups 3-12), f-block (lanthanides, actinides)

3.2 Periodic Trends

3.3 Key Definitions

4. Chemical Bonding

4.1 Types of Bonds

4.2 Lewis Structures and Octet Rule

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons
• Exceptions: H (2 electrons), Be (4), B (6), elements beyond 2nd period can expand octet
• Formal Charge: V - N - B/2 (V=valence, N=non-bonding, B=bonding electrons)
• Resonance: Multiple valid Lewis structures; actual structure is hybrid

4.3 Hybridization and Geometry

4.4 VSEPR Theory

• Valence Shell Electron Pair Repulsion: electron pairs arrange to minimize repulsion
• Repulsion order: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
• Lone pairs cause deviation from ideal bond angles

4.5 Molecular Orbital Theory

6. Equilibrium

6.1 Chemical Equilibrium

6.2 Le Chatelier's Principle

• System at equilibrium shifts to counteract applied stress
• Concentration: Add reactant/remove product → forward shift; add product/remove reactant → backward shift
• Pressure: Increase pressure → shift to side with fewer gas moles; decrease pressure → shift to side with more gas moles
• Temperature: Increase T → shift in endothermic direction; decrease T → shift in exothermic direction
• Catalyst: No effect on equilibrium position; only speeds up attainment

6.3 Ionic Equilibrium

6.4 Buffer Solutions

6.5 Solubility Equilibrium

7. Redox Reactions

7.1 Basic Concepts

7.2 Oxidation Number Rules

• Free element = 0
• Monatomic ion = charge on ion
• H = +1 (except in metal hydrides: -1)
• O = -2 (except in peroxides: -1, OF₂: +2)
• Alkali metals = +1; Alkaline earth metals = +2
• Halogens = -1 (except when combined with O or more electronegative halogen)
• Sum of oxidation numbers = 0 (neutral molecule) or charge (ion)

7.3 Balancing Redox Reactions

7.3.1 Half-Reaction Method (Acidic Medium)

1. Write separate oxidation and reduction half-reactions
2. Balance atoms other than O and H
3. Balance O by adding H₂O
4. Balance H by adding H⁺
5. Balance charge by adding electrons
6. Multiply to equalize electrons and add half-reactions

7.3.2 For Basic Medium

• Balance as in acidic medium, then add OH⁻ to neutralize H⁺
• Combine H⁺ and OH⁻ to form H₂O and simplify

8. Organic Chemistry: Basic Principles

8.1 Classification and Nomenclature

8.2 Functional Groups

8.3 Isomerism

8.4 Electronic Effects

8.5 Reactive Intermediates

8.6 Reaction Mechanisms

9. Hydrocarbons

9.1 Alkanes (CₙH₂ₙ₊₂)

9.2 Alkenes (CₙH₂ₙ)

9.3 Alkynes (CₙH₂ₙ₋₂)

9.4 Aromatic Hydrocarbons

9.5 Electrophilic Substitution in Benzene

9.6 Directive Influence