DPT Admission Test 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 DPT Admission Test, not a summary of it.

1. 1

   Atomic Structure

   4 topics

   • Subatomic Particles
   • Quantum Numbers and Orbitals
   • Electronic Configuration
   • Periodic Trends
2. 2

   Chemical Bonding

   3 topics

   • Ionic and Covalent Bonding
   • Hybridization and Molecular Geometry
   • Hydrogen Bonding and Intermolecular Forces
3. 3

   States of Matter

   3 topics

   • Gases and Gas Laws
   • Liquids
   • Solids and Crystal Lattices
4. 4

   Chemical Equilibrium

   3 topics

   • Law of Mass Action
   • Le Chatelier's Principle
   • Ionic Equilibria and pH
5. 5

   Thermochemistry and Energetics

   2 topics

   • Enthalpy and Hess's Law
   • Endothermic and Exothermic Reactions
6. 6

   Chemical Kinetics

   3 topics

   • Rate of Reaction
   • Factors Affecting Reaction Rate
   • Catalysis
7. 7

   Electrochemistry

   3 topics

   • Redox Reactions
   • Electrolytic and Galvanic Cells
   • Electrode Potential
8. 8

   Fundamentals of Organic Chemistry

   3 topics

   • Hydrocarbons
   • Functional Groups
   • Isomerism
9. 9

   Organic Compounds and Reactions

   4 topics

   • Alcohols and Phenols
   • Aldehydes and Ketones
   • Carboxylic Acids and Esters
   • Amines
10. 10

    Biochemistry

    3 topics

    • Carbohydrates and Proteins
    • Lipids and Nucleic Acids
    • Vitamins and Enzymes
11. 11

    Industrial Chemistry and Environment

    3 topics

    • Manufacturing Processes
    • Atmospheric Pollution
    • Water Treatment

Chemistry flashcards for DPT Admission Test

21 of 50 cards from the Chemistry deck — real questions with worked answers.

1. What are the three fundamental subatomic particles, and what is the charge of each?

   Proton (+1, positive), neutron (0, neutral), and electron (-1, negative).

2. What are the approximate relative masses of a proton, neutron, and electron in atomic mass units (amu)?

   Proton ≈ 1 amu, neutron ≈ 1 amu, and electron ≈ 1/1836 amu (negligible, about 0.00055 amu).

3. Who discovered the electron, proton, and neutron, respectively?

   Electron by J.J. Thomson (1897), proton by Ernest Rutherford (1919), and neutron by James Chadwick (1932).

4. What is the charge-to-mass (e/m) ratio of the electron, and who determined it?

   1.7588 × 10^11 C/kg, determined by J.J. Thomson.

5. What do the four quantum numbers (n, l, m, s) describe about an electron?

   n = principal (energy level/size), l = azimuthal (subshell/shape), m = magnetic (orbital orientation), s = spin (spin direction, +1/2 or -1/2).

6. What values of the azimuthal quantum number (l) correspond to the s, p, d, and f subshells?

   l = 0 (s), l = 1 (p), l = 2 (d), and l = 3 (f).

7. How many orbitals and electrons can s, p, d, and f subshells hold?

   s: 1 orbital/2 electrons; p: 3 orbitals/6 electrons; d: 5 orbitals/10 electrons; f: 7 orbitals/14 electrons.

8. For a given value of l, what are the allowed values of the magnetic quantum number (m)?

   m ranges from -l to +l, including 0 (a total of 2l + 1 values).

9. State the Pauli Exclusion Principle.

   No two electrons in an atom can have the same set of all four quantum numbers; an orbital holds at most two electrons with opposite spins.

10. State Hund's Rule of maximum multiplicity.

    Electrons fill degenerate (equal-energy) orbitals singly with parallel spins before pairing up.

11. State the Aufbau Principle for electronic configuration.

    Electrons fill orbitals in order of increasing energy, lowest-energy orbitals first.

12. What is the (n+l) rule for determining orbital filling order?

    Orbitals fill in order of increasing (n+l); if two orbitals have the same (n+l), the one with the lower n fills first.

13. Write the ground-state electronic configuration of chromium (Z=24) and explain the anomaly.

    [Ar] 3d^5 4s^1; one 4s electron shifts to 3d because a half-filled 3d subshell gives extra stability.

14. What is the electronic configuration of a copper atom (Z=29), and why is it anomalous?

    [Ar] 3d^10 4s^1; a fully filled 3d^10 subshell is more stable than 3d^9 4s^2.

15. How does atomic radius vary across a period and down a group in the periodic table?

    It decreases across a period (left to right) due to increasing nuclear charge, and increases down a group due to added shells.

16. Define ionization energy and describe its trend across a period and down a group.

    Energy required to remove the most loosely held electron from a gaseous atom; it increases across a period and decreases down a group.

17. How does electronegativity vary across a period and down a group?

    Increases across a period (left to right) and decreases down a group; fluorine is the most electronegative element.

18. Define electron affinity and state its general periodic trend.

    Energy released when an electron is added to a gaseous atom; it generally increases (more negative/exothermic) across a period and decreases down a group.

19. Why is the second ionization energy always greater than the first ionization energy for a given element?

    After losing one electron the cation has a higher effective nuclear charge per electron, so the remaining electrons are held more tightly.

20. What is shielding (screening) effect and how does it influence ionization energy?

    Inner electrons repel and reduce the nuclear attraction felt by outer electrons; greater shielding lowers ionization energy.

21. Define an ionic bond and give a typical example.

    An electrostatic attraction formed by complete transfer of electrons from a metal to a non-metal, e.g., NaCl (Na+ and Cl-).

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