AKU BScN 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 AKU BScN Admission Test, not a summary of it.

1. 1

   Atomic Structure and Bonding

   3 topics

   • Atomic Models and Electronic Configuration
   • Periodic Table and Periodicity
   • Chemical Bonding
2. 2

   States of Matter and Gas Laws

   3 topics

   • Gas Laws
   • Liquids and Solids
   • Kinetic Molecular Theory
3. 3

   Stoichiometry and the Mole Concept

   3 topics

   • Moles and Avogadro's Number
   • Empirical and Molecular Formulae
   • Balancing Equations and Yield
4. 4

   Chemical Energetics and Kinetics

   3 topics

   • Enthalpy and Thermochemistry
   • Rate of Reaction
   • Chemical Equilibrium
5. 5

   Acids, Bases and Salts

   3 topics

   • pH and pOH
   • Acid-Base Theories
   • Buffers and Salt Hydrolysis
6. 6

   Electrochemistry

   3 topics

   • Oxidation and Reduction
   • Electrochemical Cells
   • Electrolysis
7. 7

   Organic Chemistry

   4 topics

   • Hydrocarbons: Alkanes, Alkenes, Alkynes
   • Functional Groups
   • Isomerism
   • Reaction Mechanisms
8. 8

   Biochemistry and Environmental Chemistry

   2 topics

   • Carbohydrates, Proteins and Lipids
   • Atmospheric Chemistry and Pollution

Chemistry flashcards for AKU BScN Admission Test

21 of 60 cards from the Chemistry deck — real questions with worked answers.

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

   Principal (n) = energy level/size; Azimuthal/Angular (l) = subshell/shape; Magnetic (m) = orbital orientation; Spin (s) = electron spin (+1/2 or -1/2).

2. State the three rules governing electron filling of orbitals.

   Aufbau principle (fill lowest energy orbitals first), Pauli exclusion principle (max 2 electrons per orbital with opposite spins), and Hund's rule (orbitals of equal energy fill singly before pairing).

3. How did Rutherford's model of the atom differ from Bohr's model?

   Rutherford proposed a dense positive nucleus with electrons around it but couldn't explain stability/spectra. Bohr added that electrons occupy fixed quantized energy levels (orbits) and emit/absorb energy only when jumping between them.

4. Write the electronic configuration of chromium (Z = 24) and explain the anomaly.

   1s2 2s2 2p6 3s2 3p6 3d5 4s1. One 4s electron shifts to 3d because a half-filled d subshell (3d5) is extra stable.

5. What is the maximum number of electrons in s, p, d, and f subshells?

   s = 2, p = 6, d = 10, f = 14.

6. What is the shape and number of orbitals for the p subshell?

   Dumbbell-shaped; 3 orbitals (px, py, pz).

7. Define periodicity in the periodic table.

   The recurrence of similar physical and chemical properties at regular intervals when elements are arranged in order of increasing atomic number.

8. How does atomic radius change across a period and down a group, and why?

   Decreases across a period (increasing nuclear charge pulls electrons closer); increases down a group (additional electron shells are added).

9. Define ionization energy and state its trend across a period and down a group.

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

10. Define electronegativity and identify the most electronegative element.

    The tendency of an atom to attract a shared (bonding) pair of electrons. Fluorine is the most electronegative element.

11. How does electron affinity generally trend across a period and down a group?

    Generally becomes more negative (more energy released) across a period and less negative down a group.

12. What is the difference between a metallic, ionic, and covalent bond?

    Metallic: attraction between metal cations and a 'sea' of delocalized electrons. Ionic: electrostatic attraction between oppositely charged ions (metal + nonmetal). Covalent: sharing of electron pairs between nonmetals.

13. State the octet rule.

    Atoms tend to gain, lose, or share electrons to achieve a stable outer shell of eight electrons (like a noble gas).

14. What is a coordinate (dative) covalent bond? Give an example.

    A covalent bond in which both shared electrons come from the same atom. Example: the bond between NH3 and H+ to form NH4+.

15. Differentiate between sigma and pi bonds.

    Sigma (σ) bonds form by head-on (axial) overlap of orbitals and are stronger; pi (π) bonds form by sideways (lateral) overlap and are weaker. A single bond is one σ; a double bond is one σ + one π.

16. What is hydrogen bonding and which atoms must be involved?

    A strong dipole-dipole attraction between a hydrogen atom bonded to a highly electronegative atom (F, O, or N) and a lone pair on another F, O, or N.

17. According to VSEPR theory, what shape and bond angle does a molecule with 4 bonding pairs and no lone pairs have?

    Tetrahedral shape with bond angles of 109.5° (e.g., CH4).

18. State Boyle's law with its formula.

    At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure: PV = constant, or P1V1 = P2V2.

19. State Charles's law with its formula.

    At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature: V/T = constant, or V1/T1 = V2/T2.

20. Write the ideal gas equation and define each term.

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

21. State Avogadro's law.

    Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules (V ∝ n at constant T and P).

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