JEE Main Physical Chemistry Syllabus

Physical Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Physical Chemistry in JEE Main, not a summary of it.

1. 1

   Some Basic Concepts in Chemistry

   8 topics

   • Matter and its Nature
   • Dalton's Atomic Theory
     • Concept of atom, molecule, element, and compound
   • Laws of Chemical Combination
   • Atomic and Molecular Masses
   • Mole Concept
   • Percentage Composition
   • Empirical and Molecular Formulae
   • Chemical Equations and Stoichiometry
2. 2

   Atomic Structure

   19 topics

   • Nature of Electromagnetic Radiation
   • Photoelectric Effect
   • Spectrum of the Hydrogen Atom
   • Bohr Model of a Hydrogen Atom
     • Postulates
     • Derivation of the Relations for the Energy of the Electron and Radii of the Different Orbits
     • Limitations of Bohr's Model
   • Dual Nature of Matter
   • De Broglie's Relationship
   • Heisenberg Uncertainty Principle
   • Elementary Ideas of Quantum Mechanics
   • Quantum Mechanical Model of the Atom
   • Important Features of Quantum Mechanical Model of the Atom
   • Concept of Atomic Orbitals as One-Electron Wave Functions
   • Variation of Ψ and Ψ2 with r for 1s and 2s Orbitals
   • Various Quantum Numbers
     • Principal Quantum Number
     • Angular Momentum Quantum Number
     • Magnetic Quantum Number
   • Significance of Quantum Numbers
   • Shapes of s, p, and d - Orbitals
   • Electron Spin and Spin Quantum Number
   • Rules for Filling Electrons in Orbitals
     • Aufbau Principle
     • Pauli's Exclusion Principle
     • Hund's Rule
   • Electronic Configuration of Elements
   • Extra Stability of Half-filled and Completely Filled Orbitals
3. 3

   Chemical Bonding and Molecular Structure

   9 topics

   • Kossel-Lewis Approach to Chemical Bond Formation
   • Concept of Ionic and Covalent Bonds
   • Ionic Bonding
     • Formation of Ionic Bonds
     • Factors Affecting the Formation of Ionic Bonds
     • Calculation of Lattice Enthalpy
   • Covalent Bonding
     • Concept of Electronegativity
     • Fajan’s Rule
     • Dipole Moment
   • Valence Shell Electron Pair Repulsion (VSEPR) Theory and Shapes of Simple Molecules
   • Quantum Mechanical Approach to Covalent Bonding
     • Valence Bond Theory
     • Concept of Hybridization Involving s, p, and d Orbitals
     • Resonance
   • Molecular Orbital Theory
     • LCAOs
     • Types of Molecular Orbitals (Bonding, Antibonding)
     • Sigma and Pi-bonds
     • Molecular Orbital Electronic Configurations of Homonuclear Diatomic Molecules
     • Concept of Bond Order, Bond Length, and Bond Energy
   • Elementary Idea of Metallic Bonding
   • Hydrogen Bonding and Its Applications
4. 4

   Chemical Thermodynamics

   3 topics

   • Fundamentals of Thermodynamics
     • System and Surroundings
     • Extensive and Intensive Properties
     • State Functions
     • Entropy
     • Types of Processes
   • The First Law of Thermodynamics
     • Concept of Work, Heat Internal Energy and Enthalpy
     • Heat Capacity, Molar Heat Capacity
     • Hess’s Law of Constant Heat Summation
     • Enthalpies of Bond Dissociation, Combustion, Formation, Atomization, Sublimation, Phase Transition, Hydration, Ionization, and Solution
   • The Second Law of Thermodynamics
     • Spontaneity of Processes
     • S of the Universe and G of the System as Criteria for Spontaneity
     • ΔG° (Standard Gibbs Energy Change) and Equilibrium Constant
5. 5

   Solutions

   3 topics

   • Different methods for expressing the concentration of solution
     • Molality
     • Molarity
     • Mole fraction
     • Percentage (by volume and mass both)
   • Vapour pressure of solutions and Raoult's Law
     • Ideal and non-ideal solutions
     • Vapour pressure-composition plots for ideal and non-ideal solutions
   • Colligative properties of dilute solutions
     • Relative lowering of vapour pressure
     • Depression of freezing point
     • Elevation of boiling point
     • Osmotic pressure
     • Determination of molecular mass using colligative properties
     • Abnormal value of molar mass
     • Van’t Hoff factor and its significance
6. 6

   Equilibrium

   4 topics

   • Meaning of equilibrium and concept of dynamic equilibrium
   • Equilibria involving physical processes
     • Solid-liquid equilibrium
     • Liquid-gas equilibrium
     • Gas and solid-gas equilibria
     • Henry's law
     • General characteristics of equilibrium involving physical processes
   • Equilibrium involving chemical processes
     • Law of chemical equilibrium
     • Equilibrium constants (Kp and Kc) and their significance
     • Significance of G and G in chemical equilibrium
     • Factors affecting equilibrium concentration, pressure, temperature, effect of catalyst
     • Le Chatelier’s principle
   • Ionic equilibrium
     • Weak and strong electrolytes
     • Ionization of electrolytes
     • Various concepts of acids and bases (Arrhenius, Bronsted-Lowry, and Lewis) and their ionization
     • Acid-base equilibria including multistage ionization and ionization constants
     • Ionization of water
     • pH scale
     • Common ion effect
     • Hydrolysis of salts and pH of their solutions
     • Solubility of sparingly soluble salts and solubility products
     • Buffer solutions
7. 7

   Redox Reactions and Electrochemistry

   4 topics

   • Electronic concepts of oxidation and reduction
   • Redox reactions
     • Oxidation number
     • Rules for assigning oxidation number
     • Balancing of redox reactions
   • Electrolytic and metallic conduction
     • Conductance in electrolytic solutions
     • Molar conductivities and their variation with concentration
     • Kohlrausch’s law and its applications
   • Electrochemical cells
     • Electrolytic and Galvanic cells
     • Different types of electrodes
     • Electrode potentials including standard electrode potential, half-cell and cell reactions
     • Emf of a Galvanic cell and its measurement
     • Nernst equation and its applications
     • Relationship between cell potential and Gibbs' energy change
     • Dry cell and lead accumulator
     • Fuel cells
8. 8

   Chemical Kinetics

   3 topics

   • Rate of a chemical reaction
   • Factors affecting the rate of reactions
     • Concentration
     • Temperature
     • Pressure
     • Catalyst
   • Elementary and complex reactions
     • Order and molecularity of reactions
     • Rate law, rate constant and its units
     • Differential and integral forms of zero and first-order reactions, their characteristics and half-lives
     • Effect of temperature on the rate of reactions
     • Arrhenius theory
     • Activation energy and its calculation
     • Collision theory of bimolecular gaseous reactions (no derivation)

Physical Chemistry flashcards for JEE Main

19 of 52 cards from the Physical Chemistry deck — real questions with worked answers.

1. What is matter, and what are its three physical states?

   Matter is anything that has mass and occupies space (volume). Its three common physical states are solid, liquid, and gas, differing in the arrangement, energy, and freedom of motion of their constituent particles.

2. State the main postulates of Dalton's Atomic Theory.

   (1) Matter consists of indivisible atoms. (2) All atoms of a given element are identical in mass and properties. (3) Compounds form when atoms of different elements combine in fixed whole-number ratios. (4) Chemical reactions involve rearrangement of atoms; atoms are neither created nor destroyed.

3. Which observations/laws could Dalton's Atomic Theory NOT explain (its limitations)?

   It could not explain the existence of isotopes (atoms of same element with different masses), isobars, the divisibility of atoms into subatomic particles, the law of gaseous volumes (Gay-Lussac), or why elements combine in fixed ratios at a deeper level.

4. Distinguish between an element and a compound.

   An element is a pure substance made of only one kind of atom and cannot be broken down chemically (e.g. $\ce{O2}$, Fe). A compound is a pure substance formed by chemical combination of two or more elements in a fixed mass ratio, with properties different from its constituents (e.g. $\ce{H2O}$).

5. Define a molecule and distinguish between a molecule of an element and a molecule of a compound.

   A molecule is the smallest particle of a substance that can exist independently and retains its chemical properties. A molecule of an element contains atoms of the same kind (e.g. $\ce{O2}$, $\ce{P4}$); a molecule of a compound contains atoms of different elements (e.g. $\ce{CO2}$, $\ce{H2O}$).

6. State the Law of Conservation of Mass.

   In a chemical reaction, mass is neither created nor destroyed; the total mass of the reactants equals the total mass of the products. Proposed by Antoine Lavoisier.

7. State the Law of Definite (Constant) Proportions.

   A given chemical compound always contains its constituent elements in a fixed ratio by mass, regardless of its source or method of preparation. Proposed by Proust. Example: water always has H : O = $1 : 8$ by mass.

8. State the Law of Multiple Proportions with an example.

   When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in a simple whole-number ratio. Example: in $\ce{CO}$ and $\ce{CO2}$, the masses of oxygen per fixed mass of carbon are in the ratio $1 : 2$.

9. State Gay-Lussac's Law of Gaseous Volumes.

   When gases react, they do so in volumes that bear a simple whole-number ratio to one another and to the volumes of gaseous products, provided temperature and pressure are constant. Example: $\ce{H2 + Cl2 -> 2HCl}$ reacts in volume ratio $1:1:2$.

10. State Avogadro's Law.

    Equal volumes of all gases under the same conditions of temperature and pressure contain an equal number of molecules. Consequently, $1$ mole of any ideal gas at STP occupies $22.4\ \text{L}$ (or $22.7\ \text{L}$ at $1\ \text{bar}$).

11. Define atomic mass and atomic mass unit (amu/u).

    Atomic mass is the mass of an atom expressed in unified atomic mass units. One amu is defined as $\frac{1}{12}$ the mass of one atom of the carbon-12 isotope: $1\ \text{u} = 1.66 \times 10^{-24}\ \text{g}$.

12. How is average atomic mass calculated for an element with isotopes?

    It is the weighted average of the isotopic masses based on their relative (fractional) abundances: $$\bar{A} = \sum_i (f_i \times m_i)$$ where $f_i$ is the fractional abundance and $m_i$ the isotopic mass.

13. Define molecular mass and formula mass.

    Molecular mass is the sum of the atomic masses of all atoms in a molecule (used for covalent/molecular substances). Formula mass is the sum of atomic masses of atoms in a formula unit, used for ionic compounds (e.g. NaCl) that do not exist as discrete molecules.

14. Calculate the molecular mass of $\ce{H2SO4}$. (H=1, S=32, O=16)

    $$M = 2(1) + 32 + 4(16) = 2 + 32 + 64 = 98\ \text{u}$$

15. Define the mole and state the value of Avogadro's number.

    A mole is the amount of substance containing as many elementary entities as there are atoms in exactly $12\ \text{g}$ of carbon-12. This number is Avogadro's number: $$N_A = 6.022 \times 10^{23}\ \text{mol}^{-1}$$

16. Write the formula relating number of moles to mass and to number of particles.

    $$n = \frac{\text{given mass}}{\text{molar mass}} = \frac{m}{M}, \qquad n = \frac{N}{N_A}$$ where $N$ is the number of particles and $N_A = 6.022\times10^{23}$.

17. What volume does 1 mole of an ideal gas occupy at STP, and how is gas volume related to moles?

    At STP ($273.15\ \text{K}$, $1\ \text{bar}$) one mole occupies the molar volume $22.7\ \text{L}$ (classically $22.4\ \text{L}$ at $1\ \text{atm}$). $$n = \frac{V_{\text{gas at STP}}}{22.4\ \text{L}}$$

18. How many molecules and how many atoms are present in $2$ moles of $\ce{CO2}$?

    Molecules: $2 \times 6.022\times10^{23} = 1.204\times10^{24}$. Each $\ce{CO2}$ has 3 atoms, so atoms $= 3 \times 1.204\times10^{24} = 3.612\times10^{24}$.

19. How is percentage composition of an element in a compound calculated?

    $$\% \text{ of element} = \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100$$

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