AIIMS Nursing (BSc) Entrance Chemistry Syllabus

Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry in AIIMS Nursing (BSc) Entrance, not a summary of it.

1. 1

   Physical Chemistry

   6 topics

   • Some Basic Concepts of Chemistry
     • Mole concept and Avogadro number
     • Stoichiometry and limiting reagent
     • Empirical and molecular formula
   • Atomic Structure
     • Bohr's model and quantum numbers
     • Electronic configuration and orbitals
   • States of Matter
     • Gas laws and ideal gas equation
     • Liquids and intermolecular forces
   • Chemical Thermodynamics
     • Enthalpy, entropy and Gibbs energy
     • Hess's law and spontaneity
   • Equilibrium
     • Chemical equilibrium and Le Chatelier's principle
     • Ionic equilibrium, pH and buffers
   • Electrochemistry and Kinetics
     • Electrolysis and electrochemical cells
     • Rate of reaction and order
2. 2

   Inorganic Chemistry

   5 topics

   • Periodic Classification
     • Modern periodic law and table
     • Periodic trends in properties
   • Chemical Bonding
     • Ionic and covalent bonding
     • VSEPR theory and hybridization
     • Hydrogen bonding
   • s-Block and p-Block Elements
     • Alkali and alkaline earth metals
     • Groups 13 to 18 characteristics
   • d- and f-Block Elements
     • Transition elements and properties
     • Lanthanoids and actinoids
   • Coordination Compounds
     • Nomenclature and isomerism
     • Werner's theory and bonding
3. 3

   Organic Chemistry

   5 topics

   • Basic Principles of Organic Chemistry
     • IUPAC nomenclature
     • Isomerism and electronic effects
     • Reaction mechanisms and intermediates
   • Hydrocarbons
     • Alkanes, alkenes and alkynes
     • Aromatic hydrocarbons and benzene
   • Haloalkanes and Haloarenes
     • Nucleophilic substitution reactions
   • Oxygen-Containing Functional Groups
     • Alcohols, phenols and ethers
     • Aldehydes, ketones and carboxylic acids
   • Nitrogen-Containing Compounds and Biomolecules
     • Amines and diazonium salts
     • Carbohydrates, proteins and nucleic acids
4. 4

   Environmental and Applied Chemistry

   3 topics

   • Environmental Chemistry
     • Air and water pollution
     • Green chemistry
   • Polymers
     • Classification and types of polymerization
     • Common polymers and uses
   • Chemistry in Everyday Life
     • Drugs and medicines
     • Soaps, detergents and food chemistry

Chemistry flashcards for AIIMS Nursing (BSc) Entrance

21 of 51 cards from the Chemistry deck — real questions with worked answers.

1. State Avogadro's number and what it represents.

   6.022 × 10²³ — the number of elementary particles (atoms, molecules, ions) present in one mole of a substance.

2. Define molarity (M) and give its formula.

   Molarity = number of moles of solute / volume of solution in litres. Units: mol L⁻¹. It changes with temperature.

3. What is the difference between empirical formula and molecular formula?

   Empirical formula gives the simplest whole-number ratio of atoms (e.g., CH); molecular formula gives the actual number of atoms in a molecule (e.g., C₆H₆). Molecular = n × empirical.

4. State the law of conservation of mass.

   In a chemical reaction, matter is neither created nor destroyed; the total mass of reactants equals the total mass of products.

5. What are the four quantum numbers and what does each describe?

   Principal (n) — shell/energy & size; Azimuthal (l) — subshell/shape; Magnetic (mₗ) — orbital orientation; Spin (mₛ) — electron spin (+½ or −½).

6. State the Heisenberg uncertainty principle and its mathematical form.

   It is impossible to determine simultaneously the exact position and momentum of an electron. Δx · Δp ≥ h/4π.

7. State Hund's rule of maximum multiplicity.

   Electrons fill degenerate orbitals singly with parallel spins first before pairing begins, giving maximum number of unpaired electrons.

8. Give the de Broglie equation relating wavelength to momentum.

   λ = h / mv = h / p, where h is Planck's constant, m is mass and v is velocity.

9. State the ideal gas equation and define each term.

   PV = nRT, where P = pressure, V = volume, n = moles, R = gas constant (0.0821 L atm K⁻¹ mol⁻¹), T = absolute temperature (K).

10. State Boyle's law and Charles's law.

    Boyle's law: at constant T, P ∝ 1/V (PV = constant). Charles's law: at constant P, V ∝ T (V/T = constant).

11. What is the value and significance of absolute zero?

    0 K = −273.15 °C; the theoretical temperature at which molecular motion (kinetic energy) of an ideal gas becomes minimum/ceases and gas volume would be zero.

12. State the first law of thermodynamics with its equation.

    Energy can neither be created nor destroyed. ΔU = q + w, where ΔU = change in internal energy, q = heat added, w = work done on the system.

13. Define enthalpy (H) and write the relation for ΔH.

    Enthalpy is the heat content at constant pressure: H = U + PV. At constant pressure ΔH = qₚ (heat exchanged); ΔH < 0 = exothermic, ΔH > 0 = endothermic.

14. State the second law of thermodynamics in terms of entropy.

    The entropy (S) of an isolated system always increases for a spontaneous process; total entropy of the universe tends to a maximum (ΔS_universe > 0).

15. Write the Gibbs free energy equation and the criterion for spontaneity.

    ΔG = ΔH − TΔS. A process is spontaneous when ΔG < 0, non-spontaneous when ΔG > 0, and at equilibrium when ΔG = 0.

16. State Le Chatelier's principle.

    If a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the equilibrium shifts in the direction that counteracts the change.

17. Define pH and pOH and state their relationship at 25 °C.

    pH = −log[H⁺], pOH = −log[OH⁻]. At 25 °C, pH + pOH = 14.

18. What is a buffer solution? Give an example.

    A solution that resists change in pH on addition of small amounts of acid or base. Example: acetic acid + sodium acetate (acidic buffer).

19. Define Kw, the ionic product of water, and its value at 25 °C.

    Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.

20. What is the difference between a strong and a weak electrolyte?

    A strong electrolyte ionizes completely in solution (e.g., NaCl, HCl); a weak electrolyte ionizes only partially (e.g., CH₃COOH, NH₄OH).

21. State Faraday's first law of electrolysis.

    The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity (charge) passed: m ∝ Q.

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