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JEE Main Physical Chemistry Syllabus

The Physical Chemistry syllabus for JEE Main in full: 8 chapters, 53 topics and 94 sub-topics, every one listed below, plus 52 Physical Chemistry flashcards with answers.

8Chapters
53Topics
94Sub-topics
52Flashcards

Physical Chemistry syllabus: all 8 chapters

Expand any chapter to see its topics and sub-topics.

  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. 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. 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. 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. 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. 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. 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. 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)

JEE Main Physical Chemistry flashcards — 52 questions with 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?

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

  20. Calculate the percentage of nitrogen in $\ce{NH3}$. (N=14, H=1)

    Molar mass $= 14 + 3 = 17\ \text{g/mol}$. $$\%\text{N} = \frac{14}{17}\times 100 = 82.35\%$$

Show the other 32 Physical Chemistry flashcards

  1. Define empirical formula and molecular formula.

    The empirical formula gives the simplest whole-number ratio of atoms of each element in a compound. The molecular formula gives the actual number of atoms of each element in a molecule. They are related by $\text{Molecular formula} = n \times \text{Empirical formula}$.

  2. How do you determine the value of $n$ relating empirical and molecular formulae?

    $$n = \frac{\text{molecular mass}}{\text{empirical formula mass}}$$ The molecular formula is then obtained by multiplying each subscript of the empirical formula by $n$.

  3. A compound has empirical formula $\ce{CH2O}$ and molecular mass $180\ \text{u}$. Find its molecular formula.

    Empirical mass $= 12 + 2 + 16 = 30$. $n = \frac{180}{30} = 6$. Molecular formula $= \ce{C6H12O6}$ (glucose).

  4. Outline the steps to calculate an empirical formula from percentage composition.

    (1) Take $100\ \text{g}$, so percentages become grams. (2) Divide each mass by the element's atomic mass to get moles. (3) Divide all mole values by the smallest. (4) Convert to nearest whole numbers (multiply if needed) to get the simplest ratio.

  5. What is a balanced chemical equation, and what principle requires balancing?

    A balanced chemical equation has equal numbers of each type of atom on both sides, reflecting the Law of Conservation of Mass. Coefficients are adjusted (not subscripts) to achieve this.

  6. Define limiting reagent.

    The limiting reagent is the reactant that is completely consumed first in a reaction; it limits the amount of product formed. The other reactants are in excess. Yield calculations are based on the limiting reagent.

  7. For $\ce{N2 + 3H2 -> 2NH3}$, how many moles of $\ce{NH3}$ form from $2$ mol $\ce{N2}$ and $3$ mol $\ce{H2}$? Identify the limiting reagent.

    $\ce{H2}$ is limiting: $3$ mol $\ce{H2}$ requires only $1$ mol $\ce{N2}$. From $3$ mol $\ce{H2}$, $\ce{NH3} = \frac{2}{3}\times 3 = 2\ \text{mol}$. $\ce{N2}$ is in excess.

  8. Define molarity and write its formula.

    Molarity is the number of moles of solute dissolved per litre of solution. $$M = \frac{n_{\text{solute}}}{V_{\text{solution (L)}}} = \frac{\text{mass}/\text{molar mass}}{V(\text{L})}$$ Units: $\text{mol L}^{-1}$.

  9. What constitutes electromagnetic radiation, and how do the electric and magnetic fields oscillate?

    EM radiation consists of oscillating electric and magnetic fields that are perpendicular to each other and to the direction of propagation. It is a transverse wave that travels through vacuum at the speed of light without needing a medium.

  10. Write the relationship between speed, frequency, and wavelength of EM radiation, and give the speed of light.

    $$c = \nu \lambda$$ where $c = 3 \times 10^{8}\ \text{m s}^{-1}$ is the speed of light, $\nu$ is frequency, and $\lambda$ is wavelength.

  11. Define wavenumber and give its formula and unit.

    Wavenumber $\bar{\nu}$ is the number of waves per unit length, the reciprocal of wavelength: $$\bar{\nu} = \frac{1}{\lambda}$$ Common unit: $\text{cm}^{-1}$ or $\text{m}^{-1}$.

  12. What is Planck's quantum theory, and write the energy of one quantum?

    Energy is emitted or absorbed not continuously but in discrete packets called quanta (photons). The energy of one quantum is $$E = h\nu = \frac{hc}{\lambda}$$ where $h = 6.626 \times 10^{-34}\ \text{J s}$ is Planck's constant.

  13. What is the photoelectric effect?

    The photoelectric effect is the ejection of electrons from a metal surface when light of suitable frequency strikes it. The ejected electrons are called photoelectrons.

  14. Define threshold frequency ($\nu_0$) and work function ($W_0$) in the photoelectric effect.

    Threshold frequency $\nu_0$ is the minimum frequency of incident light required to eject electrons from a metal. The work function $W_0 = h\nu_0$ is the minimum energy needed to remove an electron from the metal surface.

  15. Write Einstein's photoelectric equation.

    $$h\nu = h\nu_0 + \frac{1}{2}m v_{\max}^{2}$$ i.e. the photon's energy equals the work function plus the maximum kinetic energy of the emitted electron: $KE_{\max} = h(\nu - \nu_0)$.

  16. List the key experimental observations of the photoelectric effect that classical wave theory failed to explain.

    (1) Electrons are ejected only if $\nu \geq \nu_0$, regardless of intensity. (2) Kinetic energy of electrons depends on frequency, not intensity. (3) Number of electrons ejected depends on intensity. (4) Emission is instantaneous (no time lag).

  17. What type of spectrum is the hydrogen atom spectrum, and how is it produced?

    It is a line emission spectrum (discontinuous, discrete bright lines). It is produced when an electric discharge is passed through gaseous hydrogen; excited H atoms emit radiation at specific wavelengths as electrons fall to lower energy levels.

  18. Write the Rydberg formula for the hydrogen spectrum and give the Rydberg constant.

    $$\bar{\nu} = \frac{1}{\lambda} = R_H \left( \frac{1}{n_1^{2}} - \frac{1}{n_2^{2}} \right)$$ with $n_2 > n_1$ and Rydberg constant $R_H = 1.097 \times 10^{7}\ \text{m}^{-1}$ (or $109677\ \text{cm}^{-1}$).

  19. Name the spectral series of hydrogen and the regions in which they appear.

    Lyman ($n_1=1$, ultraviolet), Balmer ($n_1=2$, visible), Paschen ($n_1=3$, infrared), Brackett ($n_1=4$, infrared), Pfund ($n_1=5$, far infrared).

  20. State the main postulates of Bohr's model of the hydrogen atom.

    (1) The electron revolves around the nucleus in fixed circular orbits (stationary states) without radiating energy. (2) Only orbits with angular momentum $mvr = \frac{nh}{2\pi}$ are allowed. (3) Energy is absorbed or emitted only when an electron jumps between orbits, with $\Delta E = h\nu$.

  21. Write Bohr's quantization condition for angular momentum.

    $$mvr = \frac{nh}{2\pi}, \quad n = 1, 2, 3, \dots$$ The angular momentum of the electron is quantized in integral multiples of $\frac{h}{2\pi}$.

  22. Give the expression for the radius of the $n$th Bohr orbit of hydrogen and the value for $n=1$.

    $$r_n = \frac{n^{2} h^{2} \varepsilon_0}{\pi m e^{2}} = 0.529 \times n^{2}\ \text{\AA}$$ For $n=1$ (Bohr radius), $r_1 = 0.529\ \text{\AA} = 52.9\ \text{pm}$. For hydrogen-like ions, $r_n = 0.529\,\frac{n^2}{Z}\ \text{\AA}$.

  23. Give the expression for the energy of the electron in the $n$th Bohr orbit of hydrogen.

    $$E_n = -\frac{2\pi^{2} m e^{4}}{n^{2} h^{2}} = -\frac{13.6}{n^{2}}\ \text{eV} = -\frac{2.18 \times 10^{-18}}{n^{2}}\ \text{J}$$ For hydrogen-like species, $E_n = -13.6\,\frac{Z^2}{n^2}\ \text{eV}$.

  24. Why is the energy of the electron in a Bohr orbit negative?

    The negative sign indicates that the electron is bound to (attracted by) the nucleus; its energy is lower than that of a free electron at rest at infinity, which is taken as the zero of energy. Energy must be supplied to remove the electron.

  25. How does the velocity of the electron in the $n$th Bohr orbit depend on $n$ and $Z$?

    $$v_n = \frac{2\pi e^{2}}{n h} \cdot Z \;\propto\; \frac{Z}{n}$$ For hydrogen ($Z=1$), $v_1 = 2.18 \times 10^{6}\ \text{m s}^{-1}$. Velocity decreases as $n$ increases.

  26. Define the ionization energy of hydrogen using the Bohr model.

    It is the energy required to remove the electron from $n=1$ to $n=\infty$: $$E_{\text{ion}} = E_\infty - E_1 = 0 - (-13.6\ \text{eV}) = 13.6\ \text{eV}$$ per atom ($= 1312\ \text{kJ mol}^{-1}$).

  27. List the main limitations of Bohr's model.

    (1) Fails for atoms/ions with more than one electron. (2) Cannot explain fine structure or splitting of spectral lines (Zeeman and Stark effects). (3) Violates Heisenberg's uncertainty principle by assuming fixed orbits. (4) Cannot explain chemical bonding or the wave nature of the electron.

  28. What is the dual nature of matter (wave-particle duality)?

    Matter exhibits both particle-like and wave-like properties. Just as light shows particle nature (photons) and wave nature, every moving material particle has an associated wave (matter wave). This was proposed by Louis de Broglie.

  29. Write de Broglie's relationship and define its terms.

    $$\lambda = \frac{h}{p} = \frac{h}{mv}$$ where $\lambda$ is the wavelength of the matter wave, $h$ is Planck's constant, $p = mv$ is momentum, $m$ is mass, and $v$ is velocity.

  30. Why are de Broglie wavelengths significant for electrons but negligible for macroscopic objects?

    Since $\lambda = \frac{h}{mv}$ and $h$ is extremely small, large masses give immeasurably small wavelengths. Electrons have very small mass, so their wavelengths are appreciable (comparable to atomic dimensions) and observable, e.g. in electron diffraction.

  31. How does de Broglie's relationship justify Bohr's quantization of angular momentum?

    For a stable orbit, the electron wave must form a standing wave: the orbit circumference is an integral number of wavelengths, $2\pi r = n\lambda$. Substituting $\lambda = \frac{h}{mv}$ gives $mvr = \frac{nh}{2\pi}$, exactly Bohr's quantization condition.

  32. Express the de Broglie wavelength of an electron accelerated through a potential difference $V$.

    $$\lambda = \frac{h}{\sqrt{2meV}} = \frac{12.27}{\sqrt{V}}\ \text{\AA}$$ where $V$ is in volts. This follows from equating kinetic energy $\frac{1}{2}mv^2 = eV$.

JEE Main Physical Chemistry FAQ

What is in the JEE Main Physical Chemistry syllabus?

JEE Main Physical Chemistry is split into 8 chapters: Some Basic Concepts in Chemistry, Atomic Structure, Chemical Bonding and Molecular Structure, Chemical Thermodynamics, Solutions, Equilibrium, Redox Reactions and Electrochemistry and Chemical Kinetics. Together they contain 53 topics and 94 sub-topics.

How many chapters are in JEE Main Physical Chemistry?

8 chapters, containing 53 topics and 94 sub-topics. The largest chapter is Atomic Structure, with 19 topics.

Are there flashcards for JEE Main Physical Chemistry?

Yes — a 52-card Physical Chemistry deck, printed in full in the flashcards section of this page and available in the Examius app with spaced repetition. The free plan covers one exam of your choice, including its flashcards; Examius Pro, a one-time purchase, unlocks every other exam and deck.