PIEAS 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 PIEAS Admission Test, not a summary of it.

1. 1

   Atomic Structure

   3 topics

   • Sub-Atomic Particles and Isotopes
   • Quantum Numbers and Orbitals
   • Electronic Configuration
2. 2

   Periodicity and Periodic Table

   2 topics

   • Periodic Trends
   • Classification of Elements
3. 3

   Chemical Bonding

   4 topics

   • Ionic and Covalent Bonds
   • Molecular Geometry (VSEPR)
   • Hybridization
   • Intermolecular Forces
4. 4

   States of Matter and Gases

   3 topics

   • Gas Laws and Ideal Gas Equation
   • Kinetic Molecular Theory
   • Liquids and Solids
5. 5

   Solutions and Colligative Properties

   2 topics

   • Concentration Units
   • Colligative Properties
6. 6

   Thermochemistry

   3 topics

   • Enthalpy and Heat of Reaction
   • Hess's Law
   • First Law of Thermodynamics
7. 7

   Chemical Equilibrium and Acids-Bases

   3 topics

   • Law of Mass Action and Kc
   • Le Chatelier's Principle
   • Acid-Base Theories and pH
8. 8

   Electrochemistry

   3 topics

   • Oxidation-Reduction Reactions
   • Galvanic and Electrolytic Cells
   • Electrode Potential
9. 9

   Chemical Kinetics

   3 topics

   • Rate of Reaction and Order
   • Factors Affecting Rate
   • Catalysis
10. 10

    Organic Chemistry Fundamentals

    4 topics

    • Hydrocarbons
    • Alcohols and Phenols
    • Carbonyl Compounds (Aldehydes and Ketones)
    • Carboxylic Acids and Derivatives

Chemistry flashcards for PIEAS Admission Test

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

1. What are the three sub-atomic particles, and what is the charge and relative mass of each?

   Proton (charge +1, mass ~1 amu), neutron (charge 0, mass ~1 amu), and electron (charge -1, mass ~1/1836 amu, negligible).

2. What is the atomic number (Z) and the mass number (A) of an atom?

   Atomic number Z = number of protons (defines the element). Mass number A = number of protons + number of neutrons.

3. What are isotopes?

   Atoms of the same element having the same atomic number (same protons) but different mass numbers (different numbers of neutrons).

4. Why do isotopes of an element have nearly identical chemical properties?

   Because chemical behaviour depends on the number and arrangement of electrons, which is the same for all isotopes (same atomic number); only the neutron count, affecting mass, differs.

5. How is the relative atomic mass of an element calculated from its isotopes?

   It is the weighted average of the isotopic masses: Σ(isotopic mass × fractional abundance).

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

   Principal (n) – energy level/shell size; Azimuthal/angular (l) – subshell/orbital shape; Magnetic (m_l) – orbital orientation; Spin (m_s) – electron spin direction (+1/2 or -1/2).

7. What values can the azimuthal quantum number (l) take, and which subshell does each correspond to?

   l ranges from 0 to (n-1): l=0 is s, l=1 is p, l=2 is d, l=3 is f.

8. What values can the magnetic quantum number (m_l) take for a given l, and how many orbitals result?

   m_l ranges from -l to +l including 0, giving (2l+1) orbitals: s=1, p=3, d=5, f=7.

9. State the Pauli Exclusion Principle.

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

10. What is the maximum number of electrons in a shell with principal quantum number n?

    2n² electrons.

11. Describe the shapes of s and p orbitals.

    An s orbital is spherical; a p orbital is dumbbell-shaped (two lobes) oriented along the x, y, or z axis.

12. State the Aufbau Principle.

    Electrons fill orbitals in order of increasing energy, lowest energy orbitals first (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d...).

13. State Hund's Rule of Maximum Multiplicity.

    Electrons occupy degenerate orbitals singly with parallel spins before any orbital is doubly occupied, minimising electron repulsion.

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

    [Ar]3d⁵4s¹. A half-filled 3d subshell plus half-filled 4s gives extra stability, so one 4s electron moves to 3d instead of the expected [Ar]3d⁴4s².

15. Write the ground-state electronic configuration of copper (Z=29) and explain the anomaly.

    [Ar]3d¹⁰4s¹. A completely filled 3d subshell is more stable, so an electron shifts from 4s to 3d instead of [Ar]3d⁹4s².

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

    Orbitals fill in order of increasing (n+l); for equal (n+l) values, the orbital with lower n fills first (e.g., 4s before 3d).

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

    It decreases across a period (increasing nuclear charge pulls electrons in) and increases down a group (new shells added, more shielding).

18. How does ionization energy vary across a period and down a group?

    It increases across a period (greater nuclear attraction) and decreases down a group (electrons farther out, more shielding).

19. Define first ionization energy.

    The minimum energy required to remove the most loosely bound (outermost) electron from one mole of gaseous atoms to form one mole of gaseous +1 ions.

20. How does electronegativity vary across a period and down a group, and which element is most electronegative?

    Increases across a period, decreases down a group. Fluorine is the most electronegative element.

21. How does electron affinity generally vary across a period?

    Electron affinity generally becomes more negative (more energy released) across a period as atoms more readily gain electrons; halogens have the highest electron affinities.

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