Final Exam Sample Problems Chemistry 300
àNOTE: questions “crossed out” can be ignored for 2014
Make sure that you can solve problems like the ones shown here. Please remember, this should NOT be your only way of studying. Make sure to study all of your notes, homework, warm-ups, quizzes, lab quizzes, and tests. It is important to make your mind active when you study… simply “reviewing notes” and “looking over problems” DOES NOT WORK. Quiz yourself. Quiz each other. Do practice problems. Redo old problems under test taking conditions (i.e., with no notes) and analyze how you did and what you need to work on. Also, do these problems on separate paper - THESE QUESTIONS ARE NOT THE ONLY TOPICS ON THE EXAM - REVIEW THE TOPIC LIST ON MY WEBPAGE.
1. Bonding Unit: What are valence electrons? What is the “octet rule”?
2. Why do elements bond? Is bonding endothermic or exothermic?
3. What is electronegativity? How is the concept of electronegativity related to bonding? How can you tell if two elements will bond ionically or covalently?
4. Name some properties of ionic, molecular (covalent), and metallic substances.
5. Which type of bond would you expect in the following substances: nitrogen, water, lithium fluoride, carbon dioxide, copper, calcium sulfate, boron trihydride, and ammonia? Why?
6. Of the substances that you listed in #5 as covalent, which ones would you expect to be polar and which ones would you expect to be nonpolar and why?
7. Draw Lewis structures for all the substances in #5 except for the copper and calcium sulfate. Instead, draw a Lewis structure for the sulfate polyatomic ion.
8. What does VSEPR theory tell us about the shapes of molecules?
9. Identify the three-dimensional shape of all the substances in #7 that have a 3D shape that you can identify. Name the shape, draw a 3D picture of the molecule, and label the bond angle.
10. What makes a molecule as a whole polar or nonpolar?
11. Identify each substance in #9 as either being a polar or nonpolar molecule.
12. What are the three intermolecular forces we studied? How are these different from bonds?
13. Differentiate between London dispersion forces, dipole-dipole forces, ion-dipole forces, and hydrogen bond forces (where does one find them, how strong are they, etc.).
14. Identify the most dominant type of intermolecular force for each substance in #11.
15. If you were comparing a substance with London dispersion forces only to a polar substance to a substance with hydrogen bonding, which one would have the highest boiling point? Why?
16. If you were comparing three substances with only London dispersion forces, how could you predict which one would have the highest boiling point? Why?
17. Compare and contrast solids, liquids, and gases in terms of their intermolecular forces, shape, volume, compressibility, energy, and motion of particles.
18. Gas Unit: What are the ideas of Kinetic Molecular Theory? What are the two assumptions about gas particles that we make to simplify our calculations? Under what conditions of temperature and pressure do real gases behave most “ideally”?
19. Convert 1.5 atm into mmHg, torr, and kPa.
20. Convert 500 mmHg into kPa, torr, and atm.
21. Convert 25 degrees Celsius into Kelvin.
22. Convert 300 K into degrees Celsius.
23. Convert 450 mL into liters.
24. Convert 0.085 L into mL.
25. What is STP? What are the values for standard pressure in different units? What are the values for standard temperature in different units?
26. What is the volume of 1 mole of any gas at STP?
27. What does Avogadro’s hypothesis tell us about equal volumes of gases at the same pressure and temperature?
28. A sample of gas at a pressure of 750 torr occupies a volume of 75 mL. If the pressure increases to 800 torr, while the temperature is held constant, what is the new volume?
29. A sample of a gas occupies a volume of 4.5 L at 20 degrees Celsius. What volume will it occupy at -15 degrees Celsius, if the pressure remains constant?
30. A sample of a gas at 75 degrees Celsius has a pressure of 105 kPa. What pressure will it have if the temperature is increased to 85 degrees Celsius?
31. A 2.5 L sample of a gas at 75 degrees Celsius has a pressure of 105 kPa. What volume will it occupy at STP?
32. How many moles of gas have a pressure of 1.5 atm and a volume of 20.0 L at 300 K?
33. What is the volume of 15.0 g of nitrogen gas at 25 degrees Celsius and 770 torr of pressure?
34. A mixture of 2.0 moles of nitrogen, 1.5 moles of hydrogen, and 1.5 moles of helium is in a container. The total pressure in the container is 110 torr. Find the mole fraction and partial pressure of each gas.
35. Write an equation for the reaction of magnesium and hydrochloric acid to produce hydrogen gas and magnesium chloride. Balance the equation. How many mL of hydrogen gas are produced if 0.100 g of magnesium reacts with an excess of hydrochloric acid at STP?
36. Nuclear chemistry: what particles are alpha particles made of? What are beta particles made of? What is gamma radiation made of?
37. Write the equation for the alpha decay of Pu-239.
38. Write the equation for the beta decay of N-16.
39. Zirconium-84 has a half life of 26 minutes. If you start with 175g sample, how much remains after 104 minutes?
40. Half life question types: determine amount remaining(#39), determine half life, determine starting amount, determine fraction remaining.
41. What is the difference between fission and fusion?
42. Questions regarding the nuclear power project.
43. Phase Change Unit: List and define all six phase changes. Which ones are endothermic? Which ones are exothermic?
44. Evaporation and boiling are both types of vaporization. Differentiate between them.
45. Substance A and Substance B are both liquids. Substance A has stronger intermolecular forces than substance B. Draw a graph that shows how the vapor pressure of two substances, A and B, varies with increasing temperature. On your graph, show how you would determine the normal boiling point of each substance. Which substance will have the higher normal boiling point?
46. Draw and label a heating/cooling curve for a substance that has a melting point of -10 degrees Celsius and a boiling point of 50 degrees Celsius. Label the segments of the curve according to what state(s) of matter are present. Identify segments where potential energy is changing and segments where kinetic energy is changing.
47. Differentiate between the two types of solids discussed in class (crystalline and amorphous).
48. Go back over your solutions packet and redo the questions on solubility curves. Make sure you still understand how to get each answer using the graph and possibly other calculations.
49. Solutions Unit: Differentiate between unsaturated, saturated, and supersaturated solutions. What would happen in each case if you added one more crystal of solute to them? How can you tell if a solution is each type using a solubility curve?
50. Differentiate between molarity and molality.
51. What is the molarity of a solution made from 0.15 moles of solute in 500 mL of solution?
52. What is the molarity of a solution made from dissolving 10.5 g of calcium chloride in enough water to make 2.0 L of solution?
53. How would you make 500 mL of 0.10 M silver nitrate solution?
54. What is the molality of a solution made from 0.15 moles of solute in 500 g of water?
55. What is the molality of a solution made from dissolving 2.5 g of calcium chloride in 2.0 L of water?
56. What volume of 2.0 M sodium chloride solution is needed to make 500 mL of 0.15 M solution?
57. What is the concentration of a solution made by taking 350 mL of a 5.0 M aqueous solution and adding 500 mL of water to it?
58. Write an equation for the reaction of aqueous silver nitrate with aqueous copper(II) chloride to form aqueous copper(II) nitrate and solid silver chloride. Which substance in the reaction will show up as a precipitate in the lab? Balance the equation. How many grams of precipitate can be formed when 50 mL of 0.02 M copper(II) chloride reacts with an excess of silver chloride?
59. Define and explain vapor pressure lowering, boiling point elevation, and freezing point depression (including the equations for BPE and FPD).
60. Which will have a higher boiling point… a 0.5 m solution of sodium chloride or a 0.5 m solution of calcium chloride? Explain.
61. Which will have a higher freezing point… a 0.5 m solution of sodium chloride or a 0.5 m solution of calcium chloride? Explain.
62. Which will have a higher boiling point… a 0.5 m solution of sodium chloride or a 1.5 m solution of sodium chloride? Explain.
63. Which will have a higher boiling point… a 0.5 m solution of sodium chloride or a 1.5 m solution of sodium chloride? Explain.
64. Which will have a higher boiling point… a 1.5 m solution of sodium chloride or a 0.5 m solution of calcium chloride? Explain.
65. Define and differentiate between enthalpy, entropy, and spontaneity of a reaction.
66. Kinetics and Equilibrium Unit: Draw and label a potential energy diagram for an endothermic reaction and one for an exothermic reaction. Make sure to show the energy of the reactants, energy of the products, heat of the reaction, activation energy for the forward reaction, activation energy for the reverse reaction, and energy of the activated complex. Show how a catalyst would affect both graphs and point out any differences between them.
67. Name and explain (using graphs when possible) the factors that can influence the rate of a reaction.
68. What is a mechanism for a reaction? What is an intermediate?
69. N/A
70. N/A
71. Go back over your kinetics quiz and test and make sure that you can get the correct answers for all the kinetics problems.
72. Define equilibrium based on the kinetics of a reaction. Does a reaction stop when it reaches equilibrium? What is constant at equilibrium?
73. Explain how to write an equilibrium constant expression. What do you leave out and why?
74. Write out the equation for the reaction of hydrogen gas and nitrogen gas to form ammonia gas. Balance it. Write the expression for Keq.
75. Calculate Keq for the reaction in #67 if 2.0 moles of hydrogen, 1.0 mole of nitrogen, and 4.0 moles of ammonia are at equilibrium in a 5.0 L container (remember to use MOLARITY in the Keq) .
76. Using the Keq you got in #68, what concentration of nitrogen would be at equilibrium with 2.0 M hydrogen and 1.5 M ammonia?
77. State Le Chatelier’s principle.
78. Using the equation you wrote in #67, what direction will the reaction shift if the concentration of nitrogen is increased? Why?
79. Using the equation you wrote in #67, what direction will the reaction shift if the concentration of ammonia is decreased? Why?
80. Using the equation you wrote in #67 and if the reaction is exothermic, what direction will the reaction shift if the temperature is decreased? Why?
81. Using the equation you wrote in #67 and if the pressure is increased, what direction will the reaction shift? Why?
82. Acids & Bases Unit: Name some characteristics of acids and some of bases.
83. Differentiate between the Arrhenius and Bronsted-Lowry definitions of acid and base.
84. Write an equation for the reaction of water with hydrochloric acid to produce the hydronium ion and the chloride ion. Label the acid, base, conjugate acid, conjugate base, and conjugate acid base pairs.
85. Write an equation for the reaction of water with ammonia to produce the ammonium ion and the hydroxide ion. Label the acid, base, conjugate acid, conjugate base, and conjugate acid base pairs.
86. What does it mean that something is amphoteric or polyprotic? Give a few examples of things that can be amphoteric.
87. Name the strong acids. How are strong and weak acids different? How can you recognize weak acids?
88. Name the strong bases. How are strong and weak bases different? How can you recognize weak bases?
89. Make sure that you can name and write formulas for acids and bases.
90. Calculate the pH, pOH, hydrogen ion concentration, and hydroxide ion concentration for 0.011 M solutions of hydrochloric acid, sulfuric acid, sodium hydroxide, and strontium hydroxide.
91. Calculate the hydrogen concentration if the hydroxide concentration is 0.0005 M.
92. What are indicators? How are they used in titrations? What are their limitations?
93. How can you determine if a salt is acidic, basic, or neutral?
94. For the following salts (potassium chloride, magnesium chlorate, calcium carbonate, barium acetate, and magnesium fluoride), what are their parent acids and bases? Is each parent strong or weak? Is each salt acidic, basic, neutral, or can you not tell without additional information?
95. What volume of 0.002 M hydrochloric acid is required to neutralize 50.0 mL of 0.01 M sodium hydroxide?
96. What volume of 0.002 M sulfuric acid is required to neutralize 50.0 mL of 0.01 M phosphoric acid?
97. If it takes 75.0 mL of 0.1 M potassium hydroxide to neutralize 15.0 mL of hydrobromic acid, what was the concentration of the hydrobromic acid?
98. If it takes 25.0 mL of 0.1 M calcium hydroxide to neutralize 20.0 mL of hydrofluoric acid, what was the concentration of the hydrofluoric acid?