NED 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 NED Admission Test, not a summary of it.

1. 1

   Atomic Structure

   3 topics

   • Subatomic Particles and Isotopes
   • Quantum Numbers and Orbitals
   • Electronic Configuration
2. 2

   Periodic Table and Periodicity

   2 topics

   • Classification of Elements
   • Periodic Trends
3. 3

   Chemical Bonding

   4 topics

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

   States of Matter

   3 topics

   • Gas Laws and Ideal Gas Equation
   • Liquids and Vapour Pressure
   • Solids and Crystal Structures
5. 5

   Thermochemistry and Energetics

   3 topics

   • Enthalpy and Heat of Reaction
   • Hess's Law
   • Born-Haber Cycle
6. 6

   Chemical Equilibrium

   3 topics

   • Reversible Reactions and Equilibrium Constant
   • Le Chatelier's Principle
   • Industrial Applications of Equilibrium
7. 7

   Acids, Bases and Salts

   4 topics

   • Theories of Acids and Bases
   • pH and pOH
   • Buffer Solutions
   • Salt Hydrolysis
8. 8

   Electrochemistry

   3 topics

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

   Chemical Kinetics

   3 topics

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

    s and p Block Elements

    3 topics

    • Group I and II Elements
    • Group IV to VII Elements
    • Properties and Reactions of Representative Elements
11. 11

    Transition Elements

    3 topics

    • Characteristics of d-Block Elements
    • Coordination Compounds
    • Colour and Catalytic Properties
12. 12

    Organic Chemistry

    4 topics

    • Hydrocarbons
    • IUPAC Nomenclature
    • Functional Groups
    • Macromolecules and Polymers

Chemistry flashcards for NED Admission Test

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

1. What are the three subatomic particles, and what are their relative charges and approximate masses (in amu)?

   Proton: charge +1, mass approx 1 amu (in the nucleus). Neutron: charge 0, mass approx 1 amu (in the nucleus). Electron: charge -1, mass approx 1/1836 amu (about 0.00055 amu, orbiting the nucleus).

2. Define atomic number (Z) and mass number (A). How do you find the number of neutrons?

   Atomic number Z = number of protons (and equals number of electrons in a neutral atom). Mass number A = number of protons + number of neutrons. Number of neutrons = A - Z.

3. What are isotopes? Give an example.

   Isotopes are atoms of the same element (same number of protons / same Z) but with different numbers of neutrons, hence different mass numbers. Example: carbon-12 and carbon-14 (both have 6 protons, but 6 and 8 neutrons respectively).

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

   It is the weighted average of the masses of the naturally occurring isotopes: relative atomic mass = sum of (isotopic mass x fractional abundance) for all isotopes.

5. Distinguish between isotopes, isobars, and isotones.

   Isotopes: same Z (protons), different A (neutrons). Isobars: same mass number A but different Z (different elements). Isotones: same number of neutrons but different Z.

6. Name the four quantum numbers and state what each describes.

   Principal (n): main energy level/shell and size. Azimuthal/angular (l): subshell and orbital shape. Magnetic (m_l): orbital orientation in space. Spin (m_s): direction of electron spin (+1/2 or -1/2).

7. What values can the azimuthal quantum number (l) take, and which subshells (s, p, d, f) do they correspond to?

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

8. For a given value of l, what values can the magnetic quantum number (m_l) take, and how many orbitals result?

   m_l ranges from -l to +l including 0, giving (2l+1) values. So s has 1 orbital, p has 3, d has 5, and f has 7 orbitals.

9. State the Pauli Exclusion Principle.

   No two electrons in the same atom can have the same set of all four quantum numbers. Consequently, an orbital can hold a maximum of two electrons, and they must have opposite spins.

10. State Hund's Rule of maximum multiplicity.

    When filling orbitals of equal energy (degenerate orbitals), electrons occupy them singly with parallel spins first before any orbital is doubly occupied. This minimizes electron-electron repulsion.

11. State the Aufbau Principle and give the order in which the first several subshells fill.

    Electrons fill orbitals from lowest to highest energy. The filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, ... (following the (n+l) rule).

12. What are the shapes of s, p, and d orbitals?

    An s orbital is spherical. A p orbital is dumbbell-shaped (two lobes). Most d orbitals are double-dumbbell (four-lobed/cloverleaf) shaped, with the d(z^2) orbital being an exception.

13. What is the maximum number of electrons that a shell with principal quantum number n can hold?

    A maximum of 2n^2 electrons. So n=1 holds 2, n=2 holds 8, n=3 holds 18, and n=4 holds 32.

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

    Cr: [Ar] 3d^5 4s^1, not [Ar] 3d^4 4s^2. A half-filled 3d subshell is more stable, so one 4s electron shifts to 3d for extra stability.

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

    Cu: [Ar] 3d^10 4s^1, not [Ar] 3d^9 4s^2. A completely filled 3d subshell gives extra stability, so one 4s electron shifts to 3d.

16. How are valence electrons defined, and why are they important?

    Valence electrons are the electrons in the outermost (highest n) shell of an atom. They determine an element's chemical bonding behavior, valency, and group placement in the periodic table.

17. On what basis is the modern periodic table arranged, and what is the Modern Periodic Law?

    It is arranged in order of increasing atomic number (Z). The Modern Periodic Law states that the physical and chemical properties of elements are a periodic function of their atomic numbers.

18. Classify the periodic table into its four blocks (s, p, d, f) by the subshell being filled.

    s-block: groups 1-2 (filling s). p-block: groups 13-18 (filling p). d-block (transition metals): groups 3-12 (filling d). f-block (lanthanides and actinides, filling f).

19. Distinguish between a period and a group in the periodic table.

    A period is a horizontal row (7 periods); elements in a period have the same number of shells. A group is a vertical column (18 groups); elements in a group have the same number of valence electrons and similar chemical properties.

20. Give the common names and group numbers of the alkali metals, alkaline earth metals, halogens, and noble gases.

    Alkali metals: Group 1. Alkaline earth metals: Group 2. Halogens: Group 17. Noble gases: Group 18.

21. How does atomic radius vary across a period and down a group? Explain why.

    Across a period (left to right) it decreases because increasing nuclear charge pulls electrons in. Down a group it increases because each new period adds a shell, increasing distance and shielding.

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