Chemistry
Final Exam Review
Chemical Reactions
1. Why is it useful and important to balance chemical equations?
2.
What is a precipitation reaction?
3.
What happens in an oxidation-reduction reaction? (What is oxidation? What is reduction?)
4.
What is a combustion reaction?
5.
What are synthesis and decomposition reactions?
6.
How can you recognize an acid-base reaction?
7.
Balance the following reactions and identify the type(s) for each.
a.
FeO(s) + HNO3(aq)  Fe(NO3)2(aq) + H2O(l)
b.
Mg(s) + CO2(g) + O2(g)  MgCO3(s)
c.
NaOH(s) + CuSO4(aq)  Cu(OH)2(s) + Na2SO4(aq)
d.
HI(aq) + KOH(aq)  KI(aq) + H2O(l)
e.
C3H8(g) + O2(g)  CO2(g) + H2O(g)
f.
Co(NH3)6Cl2(s)  CoCl2(s) + NH3(g)
g.
HCl(aq) + Pb(C2H3O2)2(aq)  HC2H3O2(aq) + PbCl2(s)
h.
C12H22O11(s)  C(s) + H2O(g)
i.
Al(s) + HNO3(aq)  Al(NO3)3(aq) + H2(g)
j.
B(s) + O2(g)  B2O3(s)
Chemical Composition
8. How do we know that 16.00g of oxygen contains the same number of atoms as 12.01 g of
carbon and 22.99 g of sodium?
9.
Define, compare and contrast empirical and molecular formulas for a substance.
10. What is meant by the term limiting reactant? What is being limited?
11. How is the theoretical yield for a reaction different from the actual yield? How is each
calculated?
12. Calculate the number of moles in 1.25 g samples of each of the following:
a. AgNO3
b. Na2CO3
13. For the compounds in the previous question, calculate the percent by mass of the element
whose symbol occurs first in the formula.
14. Calculate the empirical formula of the compound with the following composition:
36.38% manganese, 21.24% sulfur, 42.38% oxygen.
Chemical Quantities
15. Calculate the grams of each product that would form from 12.5 g of SiC and excess Cl 2:
SiC + 2Cl2  SiCl4 + C
Chemistry
Final Exam Review
16. For the reaction above, calculate the grams of each product that would form from 12.5 g of
SiC and 5.0 g Cl2.
17. Calculate the grams of product that would form from 12.5 g of Li2O and excess H2O:
Li2O + H2O  2LiOH
18. For the reaction above, calculate the grams of product that would form from 12.5 g of Li 2O
and 5.0 g H2O.
Energy
19. Define temperature and heat; explain how they are different (include units).
20. Define: endothermic and exothermic reactions.
21. Define: calorie, Calorie, kcal, joule, specific heat capacity.
22. Calculate the heat required to warm 12.5 g of each substance from 25.0 oC to 50.0oC:
a. H2O(l) s = 4.184 J/goC
b.
Au(s)
s = 0.13 J/goC
c.
Fe(s)
s = 0.45 J/goC
23. Combustion of ethanol releases 1409 kJ of heat per mole of ethanol burned. Calculate the
heat released when 1.25 g of ethanol is burned.
Modern Atomic Theory
24. What is the difference between the ground state for an atom and an excited state? What
happens when an excited atom returns to the ground state?
25. In the flame test lab, you observed different colors for different metal ions. What caused the
colors you saw? Why do metal ions show different colors in the flame?
26. Write electron configurations:
a. F
b.
Na
c.
Al
d.
P
e.
Sr
27. Identify (underline or circle) the valance electrons in the configurations above.
Chemical Bonding
28. Compare and contrast: ionic bonding, covalent bonding, polar covalent bonding.
29. Define electronegativity. Relate electronegativity to bond type.
30. What is the difference between a polar bond and a polar molecule?
31. Why does a Lewis structure show only the valence electrons?
32. What is the octet rule? Why do we use it when drawing Lewis structures?
33. How do you know when to use double or triple bonds in Lewis structures?
Chemistry
Final Exam Review
34. Explain the main ideas of the valence shell electron pair repulsion (VSEPR) theory.
35. Draw Lewis structures for the following:
a. H2S
b. NH3
c.
d.
CCl4
SO42-
Gases; Liquids and Solids
36. What are some of the general properties that distinguish gases from liquids and solids?
37. Compare/contrast Charles’s law (volume-temperature) with Boyle’s law (volume-pressure).
38. What is meant by the partial pressure of an individual gas in a mixture?
39. What happens to a gas sample when it is collected by displacement of water? How is this
taken into account when calculating the pressure of the gas collected?
40. What does “STP” stand for? What conditions correspond to STP? What is the volume
occupied by one mole of an ideal gas at STP?
41. What is a dipole-dipole attraction?
42. What is hydrogen bonding? What conditions are required for hydrogen bonding to occur?
43. Define London dispersion forces. For what type of molecule are they the only major
intermolecular force?
44. If the pressure on a 255 mL sample of gas is doubled from 550 mmHg to 1100 mmHg at
constant temperature, what will the new volume be?
45. If a sample of gas expands from a volume of 225 mL to 450 mL at constant temperature,
what will happen to the pressure of the sample?
46. If the temperature of 135 mL of an ideal gas sample is increased at constant pressure from
25oC to 35oC, what will be the new volume?
47. To what temperature must 152 mL of an ideal gas at 75oC be cooled at constant pressure to
decrease its volume to 141 mL?
48. Calculate the volume at STP of a gas sample that occupies 245 mL at 29 oC and 1.51 atm.
49. Calculate the volume occupied by 21.6 g of N2 gas at STP
Solutions
50. Define a solution. Describe how NaCl dissolves in water. How is this different from when a
molecular solute (such as sugar) dissolves in water?
51. Define a saturated solution. Explain how this is different from a solution being concentrated.
52. Suppose 5.00 g NaCl were dissolved in 15.0 g of water, resulting in 16.1 mL of solution after
mixing. Explain how you would calculate the molarity.
53. Calculate the molarity:
5.25 g NaOH in 500. mL solution
54. Calculate the molarity: 35.2 g FeCl3 in 225 mL solution
Chemistry
Final Exam Review
55. 10.0 mL water is added to 5.05 mL of 0.201 M NaCl; calculate the new molarity
56. 25.2 mL of 0.561 M KBr is diluted to 100. mL; calculate the new molarity.
57. Calculate the volume of 0.271 M HCl required to neutralize 36.2 mL of 0.259 M NaOH.
Acids and Bases
58. How are weak acids (like acetic acid) different from strong acids (like HCl)?
59. What does it mean to say that water is amphoteric?
60. For pure water at 25oC, what are the concentrations of H+ and OH-?
61. How does [H+] compare to [OH-] in acidic solution? In basic solution?
62. What pH range do acidic solutions have? What is the pH range for basic solutions?
63. Write the conjugate acid for each of the following:
a. NH3
b. F-
c.
d.
OHH2O
64. Write the conjugate base for each of the following:
e. HCl
f. H3PO4
g.
h.
H2O
H2PO4-
65. Identify the conjugate acid-base pairs and label each component.
i.
NH3 + H2O  NH4+ + OH-
j.
H2SO4 + H2O  HSO4- + H3O+
k.
H2PO4- + OH-  HPO42- + H2O
66. [H+] = 2.11 x 10-3 M;
[OH-] = ?
69. pH = 4.85;
[H+] = ?
67. [OH-] = 9.21 x 10-8 M;
pH = ?
70. pH = 8.91;
pOH = ?
68. pOH = 7.91;
[H+] = ?
71. [OH-] = 7.21 x 10-3 M;
[H+] = ?
Radioactivity
72. The number of protons in a given nucleus is called the ________.
73. A nucleus that spontaneously decomposes is said to be ________.
74. The atomic number and the ___________ number are conserved in a nuclear equation.
75. Production of a helium nucleus from a heavy atom is referred to as _____ decay.
76. Many radioactive nuclei produce particles and _______ rays when they decay.
77. In alpha decay, the ________ of the nucleus decreases by four units.
78. The time required for half of a radioactive sample to decay is its ________.
79. The five most naturally abundant isotopes of zinc have mass numbers of 64, 66, 67, 68 and
70. Write the nuclear symbol for each isotope, plus the number of p+ and the number of n0.
Chemistry
Final Exam Review