FBISE Examinations Chemistry (HSSC) Flashcards

24 sample cards from the Chemistry (HSSC) deck

Sampled from the end of the deck, so these are different cards from the ones shown on the syllabus page.

1. Distinguish between crystalline and amorphous solids.

   Crystalline solids have a regular, ordered, repeating arrangement of particles and sharp melting points; amorphous solids have a disordered arrangement and soften over a temperature range.

2. Define the lattice energy of an ionic solid.

   Lattice energy is the energy released when one mole of an ionic crystal is formed from its gaseous ions, or the energy required to separate one mole of the solid into gaseous ions.

3. What is observed in a discharge tube experiment at low pressure when high voltage is applied?

   Rays called cathode rays travel from the cathode to the anode, producing fluorescence; these rays consist of negatively charged particles (electrons).

4. List two key properties of cathode rays.

   Cathode rays travel in straight lines, are deflected by electric and magnetic fields toward the positive plate (showing negative charge), and can produce mechanical and heating effects.

5. What are positive (canal) rays and how are they produced in a discharge tube?

   Positive rays are streams of positively charged ions produced when gas atoms lose electrons during collisions with cathode rays; they travel toward the cathode and depend on the gas used.

6. What discovery resulted from J.J. Thomson's discharge tube experiments?

   J.J. Thomson discovered the electron and determined the charge-to-mass ratio (e/m) of the electron from cathode rays.

7. State the main postulate of Bohr's atomic model regarding electron orbits.

   Electrons revolve around the nucleus only in certain fixed circular orbits of definite energy (stationary states) without radiating energy; angular momentum is quantized as mvr = nh/2pi.

8. According to Bohr's model, when is energy emitted or absorbed by an atom?

   Energy is absorbed when an electron jumps to a higher orbit and emitted (as a photon) when it falls to a lower orbit; the energy difference equals hf = E2 - E1.

9. What are the spectral series of the hydrogen atom and the regions in which they appear?

   Lyman series (ultraviolet, transitions to n=1), Balmer series (visible, to n=2), Paschen series (infrared, to n=3), Brackett and Pfund (infrared, to n=4 and n=5).

10. Distinguish between a continuous spectrum and a line spectrum.

    A continuous spectrum contains all wavelengths merging without gaps; a line spectrum contains only discrete, separate lines characteristic of the atom emitting them.

11. What does the principal quantum number (n) describe?

    The principal quantum number (n = 1, 2, 3, ...) describes the main energy level (shell) and the size and energy of the orbital.

12. What does the azimuthal (subsidiary) quantum number (l) describe and what values can it take?

    The azimuthal quantum number (l) describes the shape of the orbital (subshell) and takes values from 0 to (n-1); l = 0, 1, 2, 3 correspond to s, p, d, f orbitals.

13. What does the magnetic quantum number (m) describe and what values can it take?

    The magnetic quantum number (m) describes the orientation of the orbital in space and takes values from -l to +l, including zero.

14. What does the spin quantum number (s) describe and what values can it take?

    The spin quantum number describes the direction of electron spin and takes values of +1/2 or -1/2.

15. State the Pauli Exclusion Principle.

    No two electrons in an atom can have the same set of all four quantum numbers; thus an orbital can hold a maximum of two electrons with opposite spins.

16. State Hund's Rule of maximum multiplicity.

    When filling orbitals of equal energy (degenerate orbitals), electrons occupy them singly with parallel spins before pairing up.

17. State the Aufbau principle.

    Electrons fill orbitals in order of increasing energy, starting with the lowest available energy orbital before filling higher ones.

18. Define an ionic (electrovalent) bond.

    An ionic bond is the electrostatic force of attraction between oppositely charged ions formed by the complete transfer of one or more electrons from a metal to a non-metal.

19. Define a covalent bond.

    A covalent bond is formed by the mutual sharing of one or more pairs of electrons between two atoms, usually non-metals.

20. Distinguish between a polar and a non-polar covalent bond.

    A non-polar covalent bond forms between atoms of equal electronegativity (shared pair equally distributed); a polar covalent bond forms between atoms of different electronegativity, creating partial charges.

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

    A coordinate covalent bond is a covalent bond in which the shared pair of electrons is donated by only one of the bonded atoms. Example: the bond in the ammonium ion (NH4+) or in H3O+.

22. State the basic principle of VSEPR theory.

    Valence Shell Electron Pair Repulsion theory states that electron pairs around a central atom arrange themselves as far apart as possible to minimise repulsion, determining molecular geometry.

23. According to VSEPR, predict the shape and bond angle of methane (CH4) and water (H2O).

    CH4 is tetrahedral with a bond angle of 109.5 degrees; H2O is bent/angular with a bond angle of about 104.5 degrees due to two lone pairs on oxygen.

24. Define hybridization and state the geometry and bond angle of sp, sp2, and sp3 hybridization.

    Hybridization is the mixing of atomic orbitals to form new equivalent hybrid orbitals. sp: linear, 180 degrees; sp2: trigonal planar, 120 degrees; sp3: tetrahedral, 109.5 degrees.

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