Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
CHAPTER 1 REVIEW QUESTIONS
1. Classify each of the following substances as a pure substance or a mixture. If it is a mixture,
classify it as heterogeneous or homogeneous. If it is a pure substance, classify it as an element
or a compound.
a. air
b. tomato juice
c. iodine crystals
d. a nickel
e. graphite (carbon)
f. sodium chloride
g. orange juice with pulp
2. A solid white substance A is heated strongly in the absence of air. It decomposes to form a new
white substance B and a gas C. The gas has exactly the same properties as the product obtained
when carbon is burned in excess of oxygen. What can we say about whether solids A and B and
the gas C are elements or compounds?
3. Classify the following substances as solid, liquid, or gas:
a. flexible volume, high kinetic energy, particles can disperse freely
b. fixed volume, very low kinetic energy, orderly particles
c. fixed volume, low kinetic energy, particles can move past each other
4. Label each of the following as a physical or chemical process:
a. corrosion of aluminum metal
b. melting of ice
c. pulverizing an aspirin
d. digesting a candy bar
e. explosion of nitroglycerin
5. Determine the type of quantity described by each unit:
a. gram (g)
b. gram per milliliter (g/mL)
c. cubic centimeter (cm3)
d. Celcius (°C)
6. Convert:
a. 2.52 x 103 kg to g
e. Joule (J)
f. kilometer (km)
g. mole (mol)
h. gram per mole (g/mol)
b. 0.0023 mm to nm
c. 6.25 x 10-4 s to ms
7. a. A sample of carbon tetrachloride, a liquid once used in dry cleaning, has a mass of 39.75 g and
a volume of 25.0 mL. What is its density?
b. The density of platinum is 23.4 g/cm3. Calculate the mass of 75.0 cm3 of platinum.
c. The density of magnesium is 1.74 g/cm3. What is the volume of 275 g of this metal?
8. Indicate the number of significant figures in:
a. 1.689 x 10-3 km
b. 0.0234 m2
e. 204.080 g
c. 7,194,300 cm
d. 435.983 K
9. Round each of the following numbers to three significant digits, and express the result in
standard exponential notation:
a. 143700
b. 0.09750
c. 890,000
4
d. 6.764 x 10
e. 33,987.22
f. -6.5559
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Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
10. Carry out the following operations, and express the answer with the appropriate number of
significant figures:
a. 320.55 – (6104.5/2.3)
b. [(285.3 x 105) – (1.200 x 103)] x 2.8954
c. (0.0045 x 20,000.0) + (2813 x 12)
d. 863 x [1255 – (3.45 x 108)]
11. Perform the following conversions:
a. 0.076 L to mL
b. 5.0 x 10-8 m to nm
c. 6.88 x 105 ns to s
d. 1.55 kg/m3 to g/L
CHAPTER 2 REVIEW QUESTIONS
12. Each of the following isotopes is used in medicine. Indicate the number of protons and neutrons
in each:
a. phosphorus-32
b. chromium-51
c. cobalt-60
d. technecium-99
e. iodine-131
f. thallium-201
13. Fill in the gaps in the following table assuming each column represents a neutral atom:
46
SYMBOL
Ti
PROTONS
45
NEUTRONS
58
ELECTRONS
18
16
50
52
ATOMIC #
38
MASS #
127
14. Indicate whether each element in the periodic table is a metal, metalloid, or nonmetal &
describe its location:
a. Li
b. Sc
c. Ge
d. Yb
e. Mn
f. Au
g. Te
15. The element magnesium consists of three naturally occurring isotopes with masses 23.98504,
24.98584, and 25.98259 amu. The relative abundances of these three isotopes are 78.70, 10.13,
and 11.17 percent, respectively. From this data, calculate the average atomic mass of
magnesium.
16. How many protons and neutrons are in each atom below:
a. 26
b. 109
c. copper-65
13𝐴𝑙
47𝐴𝑔
17. How many hydrogen atoms are in each of the following:
a. C2H5OH
b. Ca(CH3COO)2
c. (NH4)3PO4
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d. fluorine-19
Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
18. Write the empirical formula corresponding to each of the following molecular formulas;
a. S4N4
b. C7H15
c. C6H10O2 d. P4O6
e. C6H10F8
f. Si3O9
19. Each of the following elements is capable of forming an ion in chemical reactions. By referring
to the periodic table, predict the charge of the most stable ion of each:
a. Al
b. Ca
c. S
d. I
e. Cs
20. Using the periodic table, predict the formula and name of the compound formed by the
following elements:
a. Ga and F
b. Li and H *exception c. Al and I
d. K and S
21. Which of the following are ionic and which are molecular? Name each chemical.
a. Sc2O3
b. NaI
c. SCl2
d. Ca(NO3)2
e. FeCl3
f. LaP
g. CoCO3
h. (NH4)2SO4
22. Name the following ionic compounds: a. Ag2S b. Ba3(PO4)2
d. SrSO3 (**Sulfite is SO32-)
e. CoBr2
c. Mg(ClO3)2
23. Give the chemical formulas for each of the following ionic compounds
a. magnesium nitride
b. iron(II) sulfite
c. chromium(III) carbonate
24. Give the name or chemical formula, as appropriate, for each of the following acids:
a. HBrO3
b. HBr
c. H3PO4
d. iodic acid
e. sulfurous acid
CHAPTER 3 REVIEW QUESTIONS
25. a. What is the difference between P4 and 4P in a chemical equation?
b. Is the following chemical equation, as written, consistent with the law of conservation of
mass?
Why or why not? 2HNO3(aq) + Mg(OH)2(s)  H2O(l) + Ca(NO3)2(aq)
26. Balance the following equations. Indicate the type of reaction for each.
a. KNO3  KNO2 + O2
b. La2O3 + H2O  La(OH)3
c. Li + MgCl2  LiCl + Mg
d. Mg3N2 + HCl  MgCl2 + NH4Cl
e. AgNO3 + K2SO4  Ag2SO4 + KNO3
27. Write balanced equations to correspond to each of the following descriptions:
a. Solid calcium carbide, CaC2, reacts with water to form an aqueous solution of calcium
hydroxide and acetylene gas, C2H2.
b. When solid potassium chlorate is heated, it decomposes to form solid potassium
chloride and oxygen gas.
c. Solid zinc metal reacts with sulfuric acid to form hydrogen gas and an aqueous solution
of zinc sulfate.
28. Balance the following equations and indicate whether they are combustion, combination, or
decomposition reactions:
a. C3H6 + O2  CO2 + H2O
b. NH4NO3  N2O + H2O
c. N2 + H2  NH3
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Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
29. Determine the molar mass of each of the following compounds:
a. N2O5
b. FeCO3
c. Ca(C2H3O2)2
30. Calculate the percentage by mass of oxygen in the following compounds:
a. NO2
b. CH3COOCH3
c. Cr(NO3)3
31. The molecular formula of allicin, the compound responsible for the smell of garlic, is C6H10OS2.
a. What is the molar mass of allicin?
b. How many moles of allicin are present in 5.00 mg of this substance?
c. How many molecules of allicin are in 5.00 mg of this substance?
d. How many S atoms are present in 5.00 mg of allicin?
32. a. How many moles of ammonium ions are in 0.557 g of ammonium carbonate?
b. What is the mass, in grams, of 0.0438 mol of iron(III) phosphate?
c. What is the mass, in grams, of 2.69 x 1023 molecules of aspirin, C9H8O4?
d. What is the molar mass of a substance if 0.05770 mol has a mass of 15.86 g?
33. Determine the empirical formula of each of the following compounds if a sample contains:
a. 0.104 mol K, 0.052 mol C, and 0.156 mol O
b. 5.28 g Sn and 3.37 g F
c. 87.5% N and 12.5% H by mass
34. What is the molecular formula of a compound with an empirical formula of CH2 and a molar
mass of 84 g/mol?
35. Hydrofluoric acid, HF(aq), cannot be stored in glass bottles because compounds called silicates
in the glass are attacked by the HF(aq). For example, sodium silicate, Na2SiO3, reacts in the
following way:
Na2SiO3 + 8HF  H2SiF6 + 2NaF + 3H2O
a. How many moles of HF are required to dissolve 0.50 mol of Na2SiO3?
b. How many grams of NaF form when 0.300 mol of HF reacts in this way?
c. How many grams of Na2SiO3 can be dissolved by 3.00 g of HF?
36. Aluminum sulfide reacts with water to form aluminum hydroxide and hydrogen sulfide.
a. Write the balanced equation for this reaction.
b. How many grams of aluminum hydroxide are obtained from 10.5 g of aluminum sulfide?
37. What is the molarity of a solution with 5.85 g of KI dissolved in enough water to make 0.125 L of
solution?
38. What volume of 3.00M NaCl is needed for a reaction that requires 146.3 g of NaCl?
39. Laughing gas, N2O, is produced when ammonium nitrate is decomposed according to the
following reaction.
NH4NO3  N2O + 2H2O
a. How many grams of NH4NO3 are required to produce 33.0 g of N2O?
b. How many grams of water are produced in this reaction?
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Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
40. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. How many
milliters of 3.00M HCl are required to react with 12.35 g of zinc?
41. How many grams of iron are needed to react with 31.0 L of chlorine gas at STP to produce
iron(III) chloride?
42. How many grams of iron(III) chloride are produced when 15.3 g of iron react with excess
chlorine gas?
43. a. Define the terms theoretical yield, actual yield, and percent yield.
44. The fizz produced when an Alka-Seltzer tablet is dissolved in water is due to the reaction
between sodium bicarbonate and citric acid. In a certain experiment, 1.00 g of sodium
bicarbonate and 1.00 g of citric acid are allowed to react.
3NaHCO3 + H3C6H5O7  3CO2 + 3H2O + Na3C6H5O7
a. What is the limiting reactant?
b. How many grams of carbon dioxide form?
c. How much of the excess reactant remains after the limiting reactant is completely
consumed?
45. How many grams of carbon dioxide can form when a mixture of 4.95 g of C2H4 and 3.25 g of O2 is
ignited, assuming complete combustion to form carbon dioxide and water?
46. A student reacts benzene, C6H6, with bromine, Br2, to form bromobenzene, C6H5Br.
C6H6 + Br2  C6H5Br + HBr
a. What is the theoretical yield of bromobenzene in this reaction when 30.0 g of benzene
reacts with 65.0 g of bromine?
b. If the actual yield of bromobenzene was 56.7 g, what was the percentage yield?
CHAPTER 5 REVIEW QUESTIONS
47. What is the difference between kinetic and potential energy?
48. Determine if the following processes are endothermic or exothermic:
a. breaking bonds
b. freezing
c. condensing
d. combustion
49. In a thermodynamic study, a scientist focuses on the properties of a solution in a flask that is
sealed with a stopper.
a. What is the system?
b. What are the surroundings?
50. What is heat? Under what conditions is heat transferred from one object to another?
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Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
51. Calculate ΔE, and determine whether the process is endothermic or exothermic for the
following cases:
a. A system releases 113 kJ of heat to the surroundings and does 39 kJ of work on the
surroundings.
b. q= 1.62 kJ and w= -874 J
c. the system absorbs 77.5 kJ of heat while doing 63.5 kJ of work on the surroundings
52. What is specific heat capacity? Water has a high specific heat capacity of 4.884 J/g·C. What
does that mean about water?
53. Draw a potential energy diagram for the following reaction:
N2 + 3H2  2NH3
ΔH = -91.8 kJ
Include a labeled y-axis, reactants, products, activated complex, and change in enthalpy.
54. What is the difference between a state function and a path function? Provide an example of
each.
55. Consider the following reaction:
CH3OH(g)  CO(g) + 2H2(g)
ΔH= +90.7 kJ
a. Is the reaction endothermic or exothermic?
b. Calculate the amount of heat transferred when 45.0 g of CH3OH(g) are decomposed by
this reaction.
c. If the enthalpy change is 16.5 kJ, how many grams of hydrogen gas are produced?
56. What is the heat capacity of liquid water? How many kJ of heat are needed to raise the
temperature of 10.00 kg of liquid water from 24.6°C to 46.2°C?
57. The specific heat of copper metal is 0.385 J/g·K. How many joules of heat are necessary to raise
the temperature of a 1.42-kg block of copper from 25.0°C to 88.5°C?
58. From the following enthalpies of reaction:
H2(g) + F2(g)  2HF(g)
ΔH= -537 kJ
C(s) + 2F2(g)  CF4(g)
ΔH= -680 kJ
2C(s) + 2H2(g)  C2H4(g)
ΔH= +52.3 kJ
Calculate ΔH for the reaction of ethylene with fluorine:
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C2H4(g) + 6F2(g)  2CF4(g) + 4HF(g)
Chemistry2 Midterm Review 2013 – Friday, January 18h, 1st exam
CHAPTER 6 REVIEW QUESTIONS
59. What is meant by wave-particle duality?
60. The equation that relates frequency to wavelength is:
61. The equation that is used to calculate the energy of a photon is:
62. Are the following variables inversely proportional or directly proportional:
a) frequency and wavelength
b) energy and wavelength
c) energy and frequency
63. An AM radio station broadcasts at 1440 kHz, and its FM partner broadcasts at 94.5 MHz.
Calculate and compare the energy of photons emitted by these radio stations.
64. The features of the Bohr model of the atom are:
65. How is an orbital, or electron cloud, different than an orbit in Bohr’s model? How is it the
same?
66. The Heisenburg Uncertainty Principle states that:
67. What are the names, symbols, descriptions, and values of the four quantum numbers?
68. The Pauli Exclusion Principle states that:
69. a) For n=4, what are the possible values of l?
b) For l=2, what are the possible values of ml?
70. Which of the following are permissible sets of quantum numbers for an electron in a hydrogen
atom?
a. n=2, l=1, ml=1
b. n=1, l=0, ml= -1
c. n=4, l=2, ml= -2
d. n=3, l=3, ml=0
e. For those that are permissible, write the appropriate designation for the subshell to
which the orbital belongs (ex. 1s or 3p, or so on).
71. Give the values for n, l, and ml for:
a. Each orbital in the 2p subshell
b. Each orbital in the 5d subshell
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