Chem 101—General Chemistry
Practice Final Exam
Name __________________________
h = 6.626 x 10-34 J s (Planck’s Constant)
c = 3.00 x 108 m/s (speed of light)
RH = 1.097 x 10-7 m-1 (Rydberg Constant)
Multiple Choice (5 points each)
Identify the letter of the choice that best completes the statement or answers the question and legibly write the letter
in the space preceding the problem number.
____ 1. Which of the following is not a solution?
a. red wine
b. alcohol in water
c. brass
d. air
e. ice in water
____ 2. Which of the following statements concerning the atom is false?
a. Chemical reactions involve only the electrons.
b. The atom contains protons, neutrons and electrons.
c. Protons are found in orbits about the nucleus.
d. The nucleus contains both protons and neutrons.
e. Most of the mass of the atom is found in the nucleus.
____ 3. Which of the following statements about two isotopes is false?
a. They will have the same charge on the nucleus.
b. They will have different numbers of neutrons.
c. They will have essentially the same chemical reactivity.
d. They will have the same atomic numbers.
e. They will have the same atomic weights.
____ 4. Which element can be classified as an alkaline earth metal?
a. Al
b. Mg
c. Ag
d. Au
e. Na
____ 5. Which combination below corresponds to a cation with a +1 charge?
a. 7 protons and 8 electrons
b. 11 protons and 10 electrons
c. 18 protons and 18 electrons
d. 17 protons and 18 electrons
e. 20 protons and 18 electrons
____ 6. When 2.34 g of pure iron is allowed to react with an excess of oxygen, 3.35 g of the final compound is
produced. Determine the formula of the compound.
a. Fe2O3
b. FeO
c. FeO2
d. FeO3
e. FeO4
____ 7. The complete combustion of a hydrocarbon produces 90.36 g of CO2 and 46.25 g of H2O. What is the
empirical formula of the hydrocarbon?
a. CH
b. CH2
c. C2H5
d. C3H8
e. C3H4
____ 8. Ammonia and sulfuric acid react according to the equation given below. How many milliliters of 0.110 M
sulfuric acid are required to exactly neutralize 25.0 mL of 0.0840 M NH3 solution?
2 NH3(aq) + H2SO4(aq) → (NH4)2SO4(aq)
a.
b.
c.
d.
e.
1.46 mL
1.82 mL
3.64 mL
5.85 mL
9.55 mL
____ 9. Determine if each of the four situations below describes kinetic or potential energy.
I
II
III
IV
a.
b.
c.
d.
e.
a moving bullet
picture hanging on a wall
the bonds in a mixture of H 2 and O2
the movement of molecules
I
potential
potential
kinetic
kinetic
kinetic
II
kinetic
potential
potential
kinetic
potential
III
kinetic
kinetic
kinetic
potential
potential
IV
potential
kinetic
potential
potential
kinetic
____ 10. What is the molar heat capacity of table salt, NaCl (specific heat = 0.88 J g-1 °C-1)?
a. 5.30 × 1022 J mol-1°C-1
b. 24.6 J mol-1°C-1
c. 51.4 J mol-1°C-1
d. 117 J mol-1°C-1
e. 245 J mol-1°C-1
____ 11. A sample of water containing 2.00 moles is initially at 30.0°C. If the sample absorbs 2.00 kJ of heat, what
is the final temperature of the water? (specific heat of water = 4.184 J g-1 °C-1)
a. 13.3°C
b. 30.2°C
c. 43.3°C
d. 46.7°C
e. 269°C
____ 12. What is the enthalpy change when 225 g of C2H2 are burned in excess O2?
C2H2(g) + 5/2O2(g) → 2CO2(g) + H2O(l) ΔH° = -1300 kJ
a.
b.
c.
d.
e.
-1.1 × 104 kJ
-3.39 × 104 kJ
-2.93 × 105 kJ
+1.1 × 104 kJ
+2.93 × 105 kJ
____ 13. Determine the heat of reaction for the process
TiO2(s) + 4HCl(g) → TiCl4(l) + 2H2(g) + O2(g)
using the information given below:
Ti(s) + O2(g) → TiO2(s) ΔH° = -939.7 kJ
2HCl(g) → H2(g) + Cl2(g) ΔH° = -184.6 kJ
Ti(s) + 2Cl2(g) → TiCl4(l) ΔH° = -804.2 kJ
a.
b.
c.
d.
e.
-320.1 kJ
-233.7 kJ
233.7 kJ
320.1 kJ
504.7 kJ
____ 14. Which of the following is not a characteristic of the Bohr model of the atom?
a. An electron is located in an orbit around the nucleus.
b. Each orbit has a discrete energy associated with it.
c. Orbits have defined radii.
d. There is a continuum of energy levels that an electron can have around a nucleus.
e. Orbits have a defined circumference.
____ 15. Which azimuthal quantum numbers can exist for n = 3?
a. l = 0
b. l = 0, 1
c. l = 0, 1, 2
d. l = 0, 1, 2, 3
e. l = 0, 1, 2, 3, 4
____ 16. If l = 1, what value can ml have?
a. ml = -1
b. ml = +1
c. ml = 0, +1
d. ml = 0
e. ml = -1, 0, +1
____ 17. How many electrons can the second principal quantum level hold?
a. 2
b. 8
c. 16
d. 18
e. 32
____ 18. Which of the following has the electron configuration 1s22s22p63s23p6?
a. Ca
b. Cl
c. Ar
d. K+
e. Both c and d
____ 19. Which element has the largest atomic radius?
a. F
b. He
c. O
d. H
e. Na
____ 20. Which element will have five electrons in its Lewis symbol?
a. beryllium
b. neon
c. oxygen
d. carbon
e. nitrogen
____ 21. Which molecule does not contain a double bond?
a. CO2
b. CH2O
c. O2
d. HCOOH
e. HCN
____ 22. Write the correct Lewis dot structure for CO. Which statement correctly describes the structure?
a. The structure contains 1 single bond and 6 lone pairs.
b. The structure contains 1 single bond and 7 lone pairs.
c. The structure contains 1 double bond and 4 lone pairs.
d. The structure contains 1 double bond and 5 lone pairs.
e. The structure contains 1 triple bond and 2 lone pairs.
____ 23. Write the correct Lewis dot structures for the compounds given below. Arrange them in order of shortest
to longest bond lengths. Consider only the carbon-carbon and carbon-oxygen bonds.
C 2H 2
a.
b.
c.
d.
e.
C 2H 6
C 2H 4
CO
C2H6 < C2H4 < C2H2 < CO
C2H2 < CO < C2H4 < C2H6
C2H2 < C2H4 < C2H6 < CO
CO < C2H2 < C2H4 < C2H6
CO < C2H6 < C2H4 < C2H2
____ 24. From the data given below, calculate the approximate enthalpy change of reaction for the reaction below.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
bond enthalpy kJ/mol
C-H
414
C-C
347
C=C
611
C-O
351
C=O
803
O-H
463
O=O
498
H-H
436
a.
b.
c.
d.
e.
-806 kJ
-98 kJ
98 kJ
120 kJ
806 kJ
____ 25. Which statement properly describes the formal charges on the atoms in
a.
b.
c.
d.
e.
?
+2 on sulfur, -2 on oxygen
+2 on sulfur, -1 on oxygen
+1 on sulfur, -1 on oxygen
-1 on sulfur, +2 on oxygen
-2 on sulfur, 0 on oxygen
+
26. (15 points) Which species has a greater ionization energy: Ne or Na ? Explain why.
27. (15 points) List the following atoms in the correct order of increasing (smallest to largest) atomic
radii:
Bromine (Br), Fluorine (F), Sulfur (S), Arsenic (As)
28. (15 points) Draw the Lewis structure of carbonic acid, H2CO3, and determine the formal charge of
each atom in the structure.
29. (15 points) Draw the molecular orbital energy diagram for nitrogen monoxide, NO, and determine
the bond order of the molecule. (Reminder: place the atomic orbitals at appropriate relative energies
based on the electronegativity of the nitrogen and oxygen.)
30. (15 points) Iodine (I2) is a solid at room temperature but sublimes easily—ΔHosub = 62.42 kJ/mol. If
1.000 g of I2(s) is irradiated with 510 nm light, how many photons are required to vaporize this
amount of iodine?