Air University Entry 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. Define vapor pressure of a liquid.

   The pressure exerted by a vapor in equilibrium with its liquid in a closed container at a given temperature. It increases as temperature increases.

2. Define the boiling point of a liquid.

   The temperature at which a liquid's vapor pressure equals the external (atmospheric) pressure, so vaporization occurs throughout the liquid.

3. What is surface tension, and why does it occur?

   Surface tension is the energy/force that minimizes a liquid's surface area, arising because surface molecules experience a net inward attraction (unbalanced intermolecular forces). Stronger intermolecular forces give higher surface tension.

4. How does temperature affect the viscosity of a liquid?

   Viscosity (resistance to flow) decreases as temperature increases, because molecules gain kinetic energy and overcome intermolecular attractions more easily.

5. What are the seven crystal systems based on?

   They are classified by the relative lengths of the unit-cell axes (a, b, c) and the angles between them: cubic, tetragonal, orthorhombic, monoclinic, triclinic, hexagonal, and rhombohedral (trigonal).

6. Compare crystalline and amorphous solids.

   Crystalline solids have a regular, long-range ordered arrangement of particles, sharp melting points, and definite geometry. Amorphous solids (e.g., glass) lack long-range order and soften over a range of temperatures.

7. How many atoms per unit cell are in simple cubic, body-centered cubic (BCC), and face-centered cubic (FCC) lattices?

   Simple cubic = 1 atom; BCC = 2 atoms; FCC = 4 atoms per unit cell.

8. Name the four types of crystalline solids and give a property of each.

   Ionic (high melting, brittle, e.g., NaCl); covalent network (very hard, very high melting, e.g., diamond); metallic (conduct electricity, malleable); molecular (soft, low melting, e.g., ice).

9. State the first law of thermodynamics.

   Energy can neither be created nor destroyed, only converted from one form to another. ΔU = q + w (change in internal energy = heat added + work done on the system).

10. Define enthalpy of reaction (ΔH) and distinguish exothermic from endothermic.

    ΔH is the heat change at constant pressure. Exothermic reactions release heat (ΔH negative); endothermic reactions absorb heat (ΔH positive).

11. State Hess's law.

    The total enthalpy change of a reaction is the same regardless of the route taken, depending only on the initial and final states. Enthalpy changes of steps can be added to give the overall ΔH.

12. What is the standard enthalpy of formation (ΔHf)?

    The enthalpy change when one mole of a compound is formed from its elements in their standard states. The ΔHf of any element in its standard state is zero.

13. State Le Chatelier's principle.

    If a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the equilibrium shifts in the direction that counteracts (opposes) the change.

14. Write the equilibrium constant expression Kc for the reaction aA + bB <=> cC + dD.

    Kc = ([C]^c [D]^d) / ([A]^a [B]^b), where brackets are equilibrium molar concentrations and pure solids/liquids are omitted.

15. What does the magnitude of the equilibrium constant indicate?

    A large K (>>1) means products are favored at equilibrium; a small K (<<1) means reactants are favored; K ~ 1 means comparable amounts of both.

16. For the exothermic reaction N2 + 3H2 <=> 2NH3, how do increased pressure and decreased temperature affect the yield of ammonia?

    Increased pressure shifts equilibrium toward fewer gas moles (toward NH3), increasing yield. Decreased temperature favors the exothermic forward reaction, also increasing NH3 yield.

17. Does a catalyst change the position of equilibrium or the value of K?

    No. A catalyst speeds up both forward and reverse reactions equally, so equilibrium is reached faster but its position and K are unchanged.

18. Define the rate of a chemical reaction.

    The change in concentration of a reactant or product per unit time, e.g., mol L^-1 s^-1 (decrease for reactants, increase for products).

19. List the main factors that affect reaction rate.

    Concentration (or pressure for gases), temperature, surface area of solids, presence of a catalyst, and the nature of the reactants.

20. How does a catalyst increase reaction rate?

    It provides an alternative reaction pathway with a lower activation energy, allowing more colliding molecules to react, without being consumed in the reaction.

21. State the collision theory of reaction rates.

    Reactions occur when reactant particles collide with sufficient energy (equal to or greater than the activation energy) and with the correct orientation. Increasing collision frequency or energy increases the rate.

22. What is activation energy?

    The minimum energy that colliding reactant particles must possess for a reaction to occur (to form the activated complex/transition state).

23. Why does increasing temperature increase reaction rate?

    It increases the average kinetic energy of particles, so collisions are more frequent and a larger fraction of molecules have energy equal to or above the activation energy.

24. What is the order of reaction?

    The sum of the powers of the concentration terms in the experimentally determined rate law (rate = k[A]^x[B]^y; order = x + y). It indicates how rate depends on reactant concentrations.

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