PIEAS Admission Test Chemistry Flashcards

24 sample cards from the Chemistry deck

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

1. What is the difference between an ionic bond and a covalent bond?

   Ionic bond: transfer of electrons between a metal and non-metal forming oppositely charged ions held by electrostatic attraction. Covalent bond: sharing of electron pairs between non-metals.

2. What is a coordinate (dative) covalent bond?

   A covalent bond in which both shared electrons are donated by the same atom (e.g., the N→B bond in NH₃·BF₃ or the bond to H⁺ in NH₄⁺).

3. Contrast the general properties of ionic and covalent compounds.

   Ionic: high melting points, conduct when molten/dissolved, often water-soluble, hard/brittle. Covalent (molecular): low melting points, poor conductors, often soluble in non-polar solvents.

4. What is the difference between a polar and a non-polar covalent bond?

   A polar covalent bond has unequal sharing of electrons due to an electronegativity difference, creating partial charges; a non-polar covalent bond has equal sharing (identical or similar electronegativities).

5. State the octet rule.

   Atoms tend to gain, lose, or share electrons to achieve a stable noble-gas configuration of eight electrons in their outermost shell (two for H and He).

6. What is the basic postulate of VSEPR theory?

   Electron pairs around a central atom arrange themselves to be as far apart as possible to minimise repulsion, determining molecular geometry; lone pairs repel more strongly than bonding pairs.

7. What is the shape and bond angle of a molecule with 2, 3, and 4 bonding pairs (no lone pairs)?

   2 pairs: linear, 180°. 3 pairs: trigonal planar, 120°. 4 pairs: tetrahedral, 109.5°.

8. Why is the bond angle in water (104.5°) less than in methane (109.5°)?

   Water's oxygen has two lone pairs whose stronger lone-pair–bonding-pair repulsion compresses the H–O–H angle below the ideal tetrahedral 109.5°.

9. What is the molecular shape of ammonia (NH₃) and its bond angle?

   Trigonal pyramidal with a bond angle of about 107°, due to one lone pair on nitrogen repelling the three bonding pairs.

10. What are the geometries for 5 and 6 electron domains (no lone pairs)?

    5 domains: trigonal bipyramidal (90° and 120°). 6 domains: octahedral (90°).

11. What is hybridization?

    The mixing of atomic orbitals of similar energy to form an equal number of new, equivalent hybrid orbitals oriented for effective bonding.

12. Relate sp, sp², and sp³ hybridization to geometry and bond angle.

    sp: linear, 180° (e.g., BeCl₂, C in alkynes). sp²: trigonal planar, 120° (e.g., BF₃, C in alkenes). sp³: tetrahedral, 109.5° (e.g., CH₄).

13. What hybridization corresponds to trigonal bipyramidal and octahedral geometries?

    sp³d gives trigonal bipyramidal (e.g., PCl₅); sp³d² gives octahedral (e.g., SF₆).

14. How do you determine the hybridization of a central atom from electron domains?

    Count the regions of electron density (bonded atoms + lone pairs): 2=sp, 3=sp², 4=sp³, 5=sp³d, 6=sp³d².

15. What is the difference between a sigma (σ) and a pi (π) bond?

    A σ bond forms by head-on (axial) orbital overlap and is stronger; a π bond forms by sideways (lateral) overlap of p orbitals and is weaker. Single bonds are σ; double = 1σ+1π; triple = 1σ+2π.

16. List the main types of intermolecular forces from weakest to strongest.

    London dispersion forces < dipole-dipole forces < hydrogen bonding (with ion-dipole forces typically strongest among these).

17. What conditions are required for hydrogen bonding?

    Hydrogen must be bonded to a small, highly electronegative atom (N, O, or F) and there must be a lone pair on an electronegative atom of a neighbouring molecule.

18. Why does water have unusually high boiling point and density anomalies compared to other hydrides of group 16?

    Extensive hydrogen bonding between water molecules raises its boiling point and causes ice to form an open lattice, making solid water less dense than liquid water.

19. State Boyle's Law and its equation.

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

20. State Charles's Law and its equation.

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

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

    PV = nRT, where P = pressure, V = volume, n = number of moles, R = universal gas constant (8.314 J·mol⁻¹·K⁻¹ or 0.0821 L·atm·mol⁻¹·K⁻¹), and T = absolute temperature in Kelvin.

22. State Dalton's Law of Partial Pressures.

    The total pressure of a mixture of non-reacting gases equals the sum of the partial pressures of the individual gases: P_total = P₁ + P₂ + P₃ + ...

23. State the main postulates of the Kinetic Molecular Theory of gases.

    Gas particles are tiny, in constant random motion, have negligible volume, exert no intermolecular forces, undergo perfectly elastic collisions, and their average kinetic energy is directly proportional to absolute temperature.

24. Compare the properties of solids, liquids, and gases in terms of shape, volume, and compressibility.

    Solids: fixed shape and volume, nearly incompressible, particles tightly packed/vibrating. Liquids: definite volume but take container's shape, slightly compressible. Gases: no fixed shape or volume, highly compressible, particles far apart.

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