SAMPLE FE20132
Page 1 of 10
CHM111 SAMPLE FINAL EXAM
This examination should be taken closed book and closed notes. You may use a calculator, the attached
formula sheet and the attached periodic chart. It is generally best to work the problem, and then select
the appropriate answer. On the real final scrap paper will not be allowed; so, get use to showing your
work on the test itself using the space beside and/or below the question. When you take the real final you
will have until the end of the normal lab period to complete the exam (approximately three hours and
fifty minutes) though you will likely be finished well before that time. Please be sure to consider all the
choices.
I suggest that you use this exam only after you have studied extensively for the real final. Do a couple of
pages and then check your answers. If you missed any item, go back to your text and/or notes and learn
why your answer is wrong and and the key is correct.
Select the best answer.
0.0 340 mol
_____ 1. Determine the number of significant figures in the result of the calculation, 0.200 L .
A. 5
B. 4
C. 3
D. 2
_____ 2. Evaluate (4.60 X 1019)(5.98 X 10-4).
A. 2.75 X 1016 B. 2.75 X 1018 C. 1.06 X 1016
_____ 3. Convert 0.075 g to kg.
A. 7.5 X 10-5 kg
B. 75 kg
C. 7.5 kg
D. 1.06 X 1015
D. 7.5 X 103 kg
_____ 4. The density of gasoline is 0.68 g/mL. What is the mass of 6.2 ml of gasoline?
A. 6.9 g
B. 9.1 g
C. 0.11 g
D. 4.2 g
_____ 5. Convert 14 oF to oC.
A) 57 oC
B) –24 oC
C) –10 oC
D) 66 oC
_____ 6. A(n) _____ is a pure substance made up of two or more elements in a fixed ratio by mass.
A. mixture
B. compound
C. isotope
D. atom
_____ 7. The number of neutrons in Sr-90 is _____.
A) 87.62
B) 38
C) 52
D) 90
SAMPLE FE20132
Page 2 of 10
_____ 8. Which of the following is the correct noble gas notation for the electronic configuration of silicon?
A) [Ne]2s22p2
B) [He]2s22p2
C) [Ar]3s23p2
D) [Ne]3s23p2
_____ 9. The formula for potassium sulfide is _____.
A) KS2
B) KS
C) K2S
D) K2S3
_____ 10. What is the chemical name for FePO4?
A) ferrous phosphate
B) ferric phosphate
C) iron (II) phosphate
D) iron phosphorus tetraoxide
_____ 11. Which of the following is an ionic compound?
A) CaCl2
B) SO2
C) N2O5
_____ 12. What charge would an bromide anion be expected to have?
A) 1+
B) 2–
C) 1–
D) PCl3
D) 2+
_____ 13. Which of the following is an improper Lewis structure? .
_____ 14. Which Lewis structure shown below would indicate a predicted bond angle of about 120o around the
central atom?
_____ 15. Which covalent bond shown below would be the most polar?
A) H – C
B) H – Si
C) H – As
D) H – N
_____ 16. What is the formula weight of Mg(OH)2 in g/mol?
A) 41.3 g/mol
B) 58.3 g/mol
C) 116.6 g/mol
D) 34.0 g/mol
_____ 17. At high temperatures, strontium metal reacts with phosphorus in the reaction 6 Sr + P4 → 2 Sr3P2.
If 8.9 g of Sr react completely with excess P4, how many grams of Sr3P2 are produced?
A) 11 g
B) 66 g
C) 33 g
D) 0.034 g
_____ 18. Consider the balanced equation 2 KCl + 3 O2
2 KClO3. If 16 g of KCl are combined with
13 g of O2, how many moles of KClO3 are produced? Assume a complete reaction.
A) 0.27 mol
B) 0.14 mol
C) 0.21 mol
D) 0.41 mol
SAMPLE FE20132
Page 3 of 10
_____ 19. The theoretical yield of a reaction was 0.57 g. The actual yield of the reaction was 0.025 g. What
was the percent yield?
A) 23 %
B) 4.4 %
C) 0.23 %
D) 55 %
_____ 20. What is the strongest type of intermolecular force that would be predicted to occur between CH4
and N2?
A) hydrogen bonding B) dipole-dipole C) ionic bonding
D) London dispersion force
_____ 21. A gas occupies 2.00 L at 1.20 atm pressure. Calculate its volume, when its pressure is decreased to
0.850 atm at the same temperature.
A) 2.82 L
B) 0.354 L
C) 1.42 L
D) 0.510 L
_____ 22. Hygroscopic substances tend to absorb _____.
A) oxygen
B) water
C) metals
D) electrolytes
_____ 23. In general as the temperature of a liquid increases, the average kinetic energy of its
molecules _____.
A) increases B) doesn’t change C) decreases
D) could increase or decrease
_____ 24. Predict which substance would be the most soluble in water.
A) CH3CH3
B) O2
C) CH3OH
D) CH3CH2CH2CH2CH2OH
_____ 25. Some sodium chloride is completely dissolved in water. The resulting mixture is transparent.
This mixture is best described as a _____.
A) solution
B) colloidal dispersion C) nonelectrolyte
D) heterogeneous mixture
_____ 26. If 3.2 g of glucose are placed in a container and enough water is added to make the total
volume equal to 100 mL, then the concentration of the mixture is _____. Be careful when selecting
the proper symbols that represent the type of concentration measurement.
A) 3.2 % w/w
B) 0.32 % w/w C) 3.2 % v/v
D) 3.2 % w/v
_____ 27. The concentration of a reactant changes from 0.60 M to 0.75 M over a time of 30.0 s. What is the
average rate of the reaction in terms of this reactant?
A) + 5.0 X 10-3 M/s
B – 2.00 X 102 M/s
C) – 5.00 X 10-3 M/s D) + 7.3 X 10-1 M/s
_____ 28. Which letter in the diagram represents the energy of the products?
SAMPLE FE20132
Page 4 of 10
_____ 29. In general if the temperature of a reaction is increased, the rate of the reaction will _____?
A) increase B) remain constant C) halve
D) decrease
_____ 30. What is the equilibrium expression for the balanced equation 2Cl2O5(g)
K eq
A)
ClO2 O2 

Cl O 
2
5
B) K eq  ClO2  O2 
K eq 
4
C)
Cl O 
2
4ClO2(g) + O2(g)
2
K eq
5
4
ClO2  O2 
4

ClO 2  O2 

D)
Cl O 
2
2
5
_____ 31. Consider the chemical equation in question #30, and assume that it is endothermic. According to
Le Chatelier’s Principle, which of the following will shift the equilibrium to the right?
A) increasing the pressure
C) adding Cl2O5
B) decreasing the temperature D) adding a catalyst
_____ 32. Which species is called the hydroxide ion?
A) H3O+
B) OHC) H+
_____ 33. What is the name of HNO3(aq)?
A) nitric acid
B) hydronitric acid
D) NH4+
C) nitrous acid
D) hydronitrous acid
_____ 34. HI is the stronger acid of the two acids in comparison to H3PO4 . As a result, in the
reaction HI + H2PO4 –
I − + H3PO4 , the equilibrium lies to the _____, and
_____ is the stronger base.
A) left, H2PO4 – B) right, I −
C) left, I −
D) right, H2PO4 –
_____ 35. In the reaction H2CO3 + H2O → HCO3− + H3O+, ______ is the conjugate ______ of H2CO3.
A) HCO3−, acid
B) H2O, base
C) HCO3−, base
D) H3O+, base
_____ 36. 23.7 mL of a 0.540 molar solution of HCl(aq) is required to reach the end point during a titration of
16.0 mL of a KOH(aq) solution. What is the molarity of the KOH(aq) solution?
A) 1.3 M
B) 2.7 M
C) 2.0 X 102 M
D) 0.80 M
_____ 37. Which aqueous solution is most basic?
SAMPLE FE20132
A) [H3O+] = 1.0 X 10-8
Page 5 of 10
B) pH = 3
C) pOH = 11
D) [OH–] = 1.0 X 10-5
_____ 38. A solution whose pH changes very little when H3O+ or OH− ions are added to it is appropriately
called a(n) _____.
A) isotonic
B) neutralization
C) buffer
D) suspension
_____ 39. Which carbonic acid / sodium bicarbonate buffer system would have the greatest buffering
capacity?
A) 0.85 L solution of 0.400 M carbonic acid and 0.200 M sodium bicarbonate
B) 0.50 L solution of 0.400 M carbonic acid and 0.400 M sodium bicarbonate
C) 0.85 L solution of 0.200 M carbonic acid and 0.200 M sodium bicarbonate
D) 0.85 L solution of 0.400 M carbonic acid and 0.400 M sodium bicarbonate
_____ 40. Which acid is the strongest in comparison to the others shown?
A) boric acid, pKa = 9.14
B) dihydrogen phosphate ion, Ka = 6.2 X 10−8
C) phenol, pKa = 9.89
D) lactic acid, Ka = 8.4 X 10−4
_____ 41. For most solids and liquids that dissolve in liquids, solubility _____ with _____ temperature.
A) decreases, increasing
B) increases, increasing
C) increases, decreasing
D) remains constant, increasing
_____ 42. Which particle has the smallest mass?
A) neutron B) electron
C) proton
D) colloidal particle
_____ 43. Which element would likely have the greatest electronegativity?
A) Ba
B) Cu
C) O
D) Te
_____ 44. When the chemical equation P + Cl2 → PCl3 is balanced, the molar coefficient of PCl3 is _____.
A) 1
B) 2
C) 3
D) 6
_____ 45. A fixed mass of gas is placed in a cylinder with a movable sealed piston. The pressure of the
gas is held constant. If the volume of the gas increases, the temperature of the gas _____.
A) increases
B) decreases
C) doesn’t change
D) doubles
_____ 46. Identify the oxidizing agent in the reaction Al(s) + Fe3+(aq) → Al3+(aq) + Fe(s).
A) Al
B) Fe3+
C) Al3+
D) Fe
_____ 47. Which aqueous solution would have the greatest osmolarity?
A) 3.1 M MgF2 B) 2.5 M KBr C) 4.6 M HCl
D) 2.4 Na3N
_____ 48. 37.1 g of KCl are placed in a volumetric flask. Distilled water is added until the solution has a
volume of 1.08 L. What is the molarity of the solution?
A) 34.4 M
B) 2.91 X 10-2 M
C) 0.461 M
D) 1.86 M
SAMPLE FE20132
Page 6 of 10
_____ 49. What is the chemical name for PCl3?
A) phosphorus trichloride
B) potassium (III) chloride
C) potassium chloride
D) potassium trichloride
_____ 50. Which solution would have the highest pH?
A) 1.0 M H2SO4(aq)
B) 1.0 M NaOH(aq)
C) 1.0 M NH3(aq)
D) 1.0 M CH3COOH(aq)
_____ 51. Which of the following is NOT a physical change?
a. Formation of steam from water
b. Formation of salt from chlorine and sodium
c. Formation of liquid propane from gaseous propane
d. Formation of clear water and mud by filtration of muddy water.
_____ 52. What happens to the electrons when a polar covalent bond is formed?
a. One element has a larger electronegativity than the other and removes an electron from the other
making one element a cation and the other an anion. The bond is polar because the charges are
concentrated in different areas of the molecule.
b. The elements have nearly the same electronegativity and so the electrons are shared equally
forming a positive pole on one side of the molecule and a negative pole on the other.
c. One element has a larger electronegativity than the other but not strong enough to remove the
electron from the other element. The electron is shared but spends more time around the atom
with stronger electronegativity giving it a partial negative charge and leaving the other with a
slightly positive charge.
d. One element has a larger electronegativity than the other but not strong enough to remove it from
the element. The electron spends more time around the atom with stronger electronegativity
giving it a partial positive charge and leaving the other with a slightly negative charge.
_____ 53. How many electrons must be lost (-) or gained (+) for lithium to acquire the closest noble electron
configuration? a. +7
b. +1
c. +5
d. -7 e. -1
_____ 54. Which of the following elements would have chemical properties similar to magnesium?
a. Na
b. Calcium
c. O
d. Ar
e. Fe
_____ 55. Which of the following names indicates the unit should be multiplied by 10-9?
a. micro
b. giga
c. kilo
d. nano
e. mega
_____ 56. There are _______________ mm in 5.00 kilometers?
a. 5 x 106
b. 5,000,000
c. 5.00 x 10-6 d. 5 x 10-6
e. 5.00 x 106
_____ 57. Which of the following conversion factors should be used to convert megawatts (Mw)
to watts (w)?
a. Mw/106 w b. w/106 Mw
*c. 106 w/Mw
d. 109 w/Mw e. Mw/10-6 w
_____ 58. Which element pair below is most likely to form an ionic compound?
A. Sr & Cl B. C & O
C. Na & Mg
D. N & H
SAMPLE FE20132
Page 7 of 10
_____ 59. Which of the following compounds is a nonpolar covalent compound?
A. BrCl
B. CsCl
C. H2O
D. HCl
_____ 60. Which valence electron : Lewis diagram pair is correct for the indicated compound ?
A
B
C
D
4
18
32
30
HClO
HNO3
H3PO4
_____ 61. What is the correct name of CuNO3?
A. Cupric nitrate
B. Copper(I) nitrite
H3PO4
C. Cuperic nitrogen trioxide
D. Copper(I) nitrate
_____ 62. What is the correct formula for stannic nitrate?
A. Sn(NO2)4
B. Sn(NO3)4
C. Tn(NO3)4
D. Sn3N4
_____ 63. What is the correct name for P2O5?
A. Phosphorous pentaoxide
C. Diphosphorous pentaoxide
B. Phosphorous(V) oxide
D. Potassium(V) oxide
_____ 64. What is the formula for magnesium phosphite?
A. Mg3(PO4)2
B. Mg2(PO3)3
C. Mg3(PO3)2
D. Mn3(PO3)2
_____ 65. If 6.93 g of nitrogen dioxide gas occupies 5.38 L at a pressure of 0.830 atm, what is the temperature
in degrees Celsius? Show all steps for credit.
A. 361
B. -37.5
C. 236
D. 88.0
_____ 66. How many atoms of hydrogen are there in 0.500 moles of water?
a. 3.01 x 1022
b. 3.01 x 1023
c. 6.02 x 1023
d. 3.01 x 10-23
e. 1.50 x 1023
_____ 67. If the formula weight of a compound is 25g, how many moles are there in 100 g?
a. 4
b. 2
c. 1
d. 0.5
e. 0.25
Items 68-71: Consider the following chemical equation:
SAMPLE FE20132
Page 8 of 10
Calcium hydroxide reacts with phosphoric acid to yield calcium phosphate and water.
_____ 68. What is the sum of the molar coefficients of the balanced equation?
a. 16
b. 12
c. 11
d. 7
e. 5
_____ 69. Which of the following is a correct balanced net ionic equation for this reaction?
a. OH- + H+
 HOH
c. 6 H+ + 6OH-  6 H2O
-3
b. OH + PO4  (H0)3PO4
d. 2 H+ + Ca+2  CaH2
_____ 70. How many moles of calcium hydroxide will it take to react completely with 1.00 mole of phosphoric
acid? a. 0.667
b. 1.00
c. 1.50
d. 3.00
_____ 71. This reaction can be classified as a _______ reaction.
a. double displacement
b.
single displacement
c. decomposition
d.
synthetic
Items 72-77: Consider the following chemical reaction:
______Cr +
___Pb(NO3)4  ___Cr(NO3)3
+
___Pb
_____ 72. What is the sum of the molar coefficients of the balanced equation?
a. 8
b. 10
c. 12
d. 14
e. 16
_____ 73. Which component was oxidized in this reaction?
a. Cr(NO3)3
b. Pb(NO3)4
c. Cr
d. Pb
e. none of these
_____ 74. Which component was reduced in this reaction?
a. Cr(NO3)3
b. Pb(NO3)4
c. Cr
d. Pb
e. none of these
_____ 75. How many grams of chromium will be required to produce 100 g of Cr(NO3)3 (FW=238.0108)?
a. 100 g
b. 52.2 g
c. 21.8 g
d. 133 g
e. 76.5 g
_____ 76. Suppose you start with 300 g lead (IV) nitrate and 100 g of chromium. Which reactant is the
limiting reagent?
a. Cr
b. Pb(NO3)4
c. Pb
d. Cr(NO3)3
_____ 77. This reaction can be classified as a _______ reaction.
a. synthetic
b. decomposition
c. single displacement
d.
double displacement
SAMPLE FE20132
Page 9 of 10
_____ 78. The specific heat of steel is 0.118 cal/g ∙ oC. If 20.0 g of steel at 25.0 oC is warmed to 86.0 oC, how
much heat was absorbed? Show all steps for credit.
A. 203 cal
B. 144 cal
C. 59 cal D. 262 cal
_____ 79. Which of the following is not the same concentration as the other three sulfuric acid solutions?
A. 0.05 M
B. 0.1 N
C. pH=13 D. 0.15 osmolar
Exam Formula and Additional Information Sheet
d
m
V

o
C
5 o
F  32
9
o
F
9o
C  32
5

K = oC + 273
Avogadro’s number = 6.02 X 1023
P1V1 P2V2

T1
T2
PV = nRT
L  atm
R = 0.0821 mol  K
PT = P1 + P2 + P3 + …
M1V1 = M2V2
Osmolarity = M  i
pH = - log [H3O+]
Amount of heat = SH × m × (T2 – T1)
n = m/FM
n = M*L
O = M * TOTAL # IONS PRODUCED
SAMPLE FE20132
N = M * TOTAL # SPECIAL IONS
Keq = Product [products]/Product [reactants]
1 kcal = 4.184 kJ
1 kcal = 1 Cal
Page 10 of 10