Osterberg-Chemistry II
Final Exam Review
Name _______________________________
Date ______________ Hour ____________
Chemistry Final Exam Review
Chapter 6
1. What is the difference between ionic and covalent bonds? Give an example of each type.
2. What is the difference between polar and nonpolar covalent bonds? Give an example of each type.
3. What is electronegativity? What does the difference of two electronegativity values of two bonded elements
indicate? What is the trend?
4. How can one determine which region in a polar molecule is slightly positive or negative? Indicate which
regions of a water molecule are slightly positive or negative using two methods: arrows and symbols.
5. For the following compounds: (1) Draw the Lewis Dot structure; (2) Determine the shape, according to the
VSEPR theory; (3) Label as polar or nonpolar; and (4) Determine the predominant intermolecular force.
a) NH3
b) H2S
c) CCl4
d) HF
6. What are the three types of intermolecular covalent bonding? What are their relative strengths? Which type
would most likely be involved in a material with a high boiling point or melting point? Low boiling point or
melting point?
7. How do London Dispersion forces occur?
Chapter 7/9-Part I
8. What would be the same and what would be different between 1 mole of lead and 1 mole of sodium?
9. Determine the percent composition of:
a) calcium phosphate
b) KNO2
10. What is an empirical formula? What is a molecular formula? Can they ever look the same?
11. What is the molecular formula of a compound with an empirical formula CHOCl and a molecular weight of
129 g?
12. Determine the empirical formula for a compound using the following information: (HINT: Use the table.)
a. 58% Rb, 9.50% N, 32.5% O
b. 49.62% C, 10.82% H, 39.56% O
Chapter 7/9-Part II
13. How can the mole ratio of two compounds involved in a chemical reaction be determined?
14. What is the mole ratio of hydrogen to water in the following unbalanced reaction: H2 + O2→H2O ?
15. What is a limiting reactant and an excess reactant?
16. Given unbalanced: Fe + H2O → Fe2O3 + H2
a) If the reaction begins with 0.79 moles of Fe, then how many grams of H2 are produced?
b) If 21.0 grams of water react, how many grams of iron (III) oxide are produced?
c) When 41 grams of water react with 167 grams of Fe, which is the limiting reactant?
Chapter 10
17. What is the difference between enthalpy and entropy? Give the symbols and units for each.
18. If the change in enthalpy is negative, then the reaction is ______thermic. If the change in enthalpy is
positive, then the reaction is ______thermic.
19. Draw a rough sketch of an energy diagram for an exothermic reaction and an endothermic reaction.
20. List three situations that would cause entropy to increase.
21. Which state of matter has the highest entropy? Lowest entropy?
22. Reaction with ____________________ entropy and _____________ enthalpy will be spontaneous.
23. If ∆H is ______ and ∆S is ______, the reaction will not occur spontaneously. If ∆H is ______ and ∆S is
______, the reaction will occur spontaneously.
24. a) A reaction has ∆H= -355 kJ and ∆S= -36 kJ/K and occurs at 525 K. Calculate the ∆G. Is the reaction
spontaneous?
b) A reaction has ∆H= -220 kJ and ∆S= -38 kJ/K and occurs at 354 K. Calculate the ∆G.
spontaneous?
Is the reaction
25. Using Hess' Law, calculate the change of enthalpy for the following reactions.
a) CH4 (g) + NH3 (g) → HCN (g) + 3 H2 (g), given:
N2 (g) + 3 H2 (g) →2 NH3 (g)
∆H = -91.8 kJ
C (s) + 2 H2 (g) → CH4 (g)
∆H = -74.9 kJ
H2 (g) + 2 C (s) + N2 (g) → 2 HCN (g) ∆H = +270.3 kJ
26. What amount of energy, must be lost from 110.0 grams of water at 78.0°C to cool it to 10.0°C? (Specific
heat capacity of water: 4.18 J/g/°C)
27. If a gold ring with a mass of 5.5 g changes in temperature from 25.0°C to 28.0°C, how much energy has it
absorbed? (Specific heat capacity of Au: .129 J/g/°C)
Chapter 13
28. What is molarity? What is its symbol? What is its unit?
29. What is the molarity of a NaOH solution that has a volume of 320.0 mL and 24.2 g of NaOH?
30. How would you prepare a 0.140 L solution of 0.80 M NaOH from a stock solution of 5.0 M?
31. What is a saturated, unsaturated, and supersaturated solution?
32. What does the phrase “likes dissolve likes” mean?
33. What are the two requirements to be an electrical conductor?
34. What is an electrolyte? What determines the strength of an electrolyte?
35. What are colligative properties?
36. What is boiling point elevation?
37. What is freezing point depression?
Chapter 14/16
38. What is the difference between a complete and incomplete reaction?
39. At equilibrium, how does the rate of the forward reaction compare to the reverse reaction? How do the
concentrations of the reactants compare to the products?
40. What does it mean if the Keq is greater than 1? equal to 1? less than 1?
41. How is the equilibrium constant calculated? What types of compounds affect the equilibrium constant?
42. What is Le Chatelier's Principle?
43. Which direction is favored if the following characteristics are increased during a reaction:
Pressure?
Temperature?
Concentration of the reactants?
44. At a temperature of 25oC, the concentrations (M) of the reactants and products for the reaction involving
carbonic acid and water are present. 2SO2 (g) + O2 (g) → 2SO3 (g)
[SO2]= 3.5 x 10-2 M
[O2]=5.4 x 10-4 M
[SO3]=3.6x 10-4 M
What is the Keq value for the reaction at equilibrium in a dilute aqueous solution?
45. What is activation energy? Draw a rough sketch of an energy diagram for a reaction and label the activation
energy. Then, draw a dotted line to show what a catalyst does to the activation energy. Label the change in
enthalpy for the reaction.
46. List the five main factors that affect the rate of reaction and explain how.
Concentration:
Temperature:
Particle Size:
Pressure:
Catalyst:
Chapter 15
47. What are the Arrhenius definitions for an acid and a base? What are the Bronsted-Lowry definitions of an
acid and a base? What is the significance of the Bronsted-Lowry definitions?
48. What is a conjugate acid? What is a conjugate base?
49. What is the pH of a 0.00033 M solution of HCl?
50. What is the pH for a solution that has [OH-] = 6.0 x 10-6?
51. What is a titration?
52. What is the equivalence point? How to know when you have reached the equivalence point?
53. What are the possible pH ranges for the equivalence point of the following Titrations:
a) Strong acid-Strong Base
b) Weak Acid-Strong Base c) Strong Acid-Weak Base
54. Use the Brønstead-Lowry and Arrehenius definitions of acids to identify each reactant as an acid, a base,
conjugate acid, and conjugate base.
a) KOH + HBr → KBr + H2O
b) HCl + H2O → Cl- + H3O+
Chapter 17
55. What is reduction? What is oxidation?
56. Where does oxidation and reduction take place in an electrochemical cell?
57. The more positive the value for electrode potential (voltage) means the electrode is more likely to be a(n):
cathode or anode?
58. Assign oxidation numbers: a) MnO2
b) NH3
c) KClO3
59. What is the standard cell voltage for Cd(s) + 2 Ag+(aq) → Cd2+(aq) + 2 Ag(s) ?
Ag+(aq) + e- → Ag(s) Eo = +0.799 V
Cd2+(aq) + 2 e- → Cd(s)
Eo = -0.402 V
60. Balance using the half-reaction method:
Cr0 + Pb+2 → Cr+3 + Pb0
______________________________________________________ (reduction)
______________________________________________________ (oxidation)
______________________________________________________ (overall)
Chapter 18
61. What does the top number in a nuclear symbol represent? Bottom number?
62. Write the particle symbol for alpha and beta particles.
63. Write the nuclear equation for the following reactions.
a)
198
90
Th→ __4 He + ____
173
0
b) 82 Pb→ −1 e
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+ _____
64. What is a half-life? How many half-lives are required for half of a sample of an isotope to decay? Threequarters?
65. An artifact is found to have 25% of carbon-14 what an object is found to have today. If the half-life of
carbon-14 is 5,730 years, how old is the object? Give your answer with correct sig figs and units!