Chem 11 - Exam #3 Sample
Sample Exams are designed to give you an idea of the types of questions you can expect. Questions may not exactly agree with currently assigned material, so you may need to examine other sample exams. Answers for Sample Exam #3 will be available prior to the exam.
1. One mole of an ideal gas is expanded from a volume of 1.00 liter to a volume of 10.00 liters against a constant external pressure of 1.00 atm. How much work (in joules) is performed on the surroundings? (T = 300 K; 1 L atm = 101.3 J)
2. A 140.0-g sample of water at 25.0oC is mixed with 100.0 g of a certain metal at 100.0oC. After thermal equilibrium is established, the (final) temperature of the mixture is 29.6oC. What is the heat capacity of the metal, assuming it is constant over the temperature range concerned?
3. For the reaction C2H5OH (l) + 3O2 → 2CO2 (g) + 3H2O(l) ΔH = -1.37 x 103 kJ
When a 15.1-g sample of ethyl alcohol (molar mass = 46.1 g/mol) is burned, how much energy is released as heat?
4. At 25oC, the following heats of reaction are known:
| 2ClF + O2 → Cl2O + F2O | ΔH = 167.4 kJ/mol |
| 2ClF3 + 2O2 → Cl2O + 3F2O | ΔH = 341.4 kJ/mol |
| 2F2 + O2 → 2 F2O | ΔH = -43.4 kJ/mol |
At the same temperature, calculate ΔH for the reaction:
| ClF + F2 → ClF3 | ΔHrxn = ______________________ kJ |
5. The following questions relate to ionization energy.
_____ Which atom would have the largest second ionization energy? a. Mg b. Cl c. S d. Ca e. Na
_____ If the first ionization energy of Mg is 735 kJ/mol, the second ionization energy is
a. 735 kJ/mol
c. greater than 735 kJ/mol
b. less than 735 kJ/mol
d. More information is needed to answer this question.
_____ Consider the ionization energy (IE) of the magnesium atom. Which of the following is not true?
a. The IE of Mg is lower than that of sodium.
c. The IE of Mg is lower than that of beryllium.
b. The IE of Mg is lower than that of neon.
d. The IE of Mg is higher than that of calcium.
e. The IE of Mg is lower than that of Mg+
6. Using the information below, calculate ΔHfo for PbO(s)
PbO(s) + CO(g) → Pb(s) + CO2(g) ΔHo = -131.4 kJ
ΔHfo for CO2(g) = -393.5 kJ/mol
Δ Hfo for CO(g) = -110.5 kJ/mol
7. The heat of formation of Fe2O3(s) is -826 kJ/mol. Calculate the heat of the reaction for
4Fe(s) + 3O2 (g) → 2Fe2O3(s) when a 55.8-g sample of iron is reacted.
8. Which is more likely to cause a severe burn, 1 g of steam at 100oC or 1 g of water at 100oC? Explain. How many joules is required to raise the temperature of 45 g water 45oC?
9. Write the electron configuration for the following:
P __________________________________________________________________
Cu __________________________________________________________________
S2- __________________________________________________________________
Cr __________________________________________________________________
K+ __________________________________________________________________
10. The blue color in fireworks is often achieved by heating copper(I) chloride to about 1200oC. Then the compound emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum) that is emitted at 4.5 x 102 nm by CuCl? (ΔE = hν ; h=6.626 x 10-34 J s)
11. How many orbitals in an atom can have the designation:
5p ___________________
4d ___________________
n=5 ___________________
n=4 ___________________
3px ___________________
11. Answer the following True-False questions:
___ T ___F In exothermic reaction, potential energy stored in chemical bonds is being converted to thermal energy via heat.
___ T ___F Heat and temperature are interchangeable terms.
___ T ___F The SI unit for frequency is cycles per second.
___ T ___F Diffraction results when light is scattered from a regular array of points or lines.
___ T ___F All matter exhibits either particulate or wave properties exclusively, not both.
___ T ___F The magnetic quantum number tells the polarity of the atom.
___ T ___F The electron correlation problem relates to the fact that repulsions among multiple electrons cannot be calculated exactly.
___ T ___F Valence electrons are found on the outermost principal quantum level of an atom.
___ T ___F The last electrons in the transition elements are filling p orbitals.
___ T ___F Only three quantum numbers are needed to uniquely describe an electron.
___ T ___F A property that is independent of the pathway is called an intensive property.
___ T ___F The first law of thermodynamics is the same as the law of conservation of mass.
___ T ___F In an endothermic process, heat is absorbed from the surroundings.
___ T ___F Frequency is the distance between two consecutive peaks or troughs in a wave.
___ T ___F Electron excitation is the energy required to remove an electron from a gaseous atom or ion.
___ T ___F The first quantum number relates to the shape of the atomic orbital.
___ T ___F Orbitals with the same principal quantum number and same energy are degenerate
___ T ___F The Pauli Exclusion Principle states that the lowest energy configuration for an atom has the maximum number of unpaired electrons in degenerate orbital.
___ T ___F The heat associated with a chemical reaction is measured in a calorimeter.
___ T ___F Under conditions of constant pressure, the heat flow that occurs during a chemical change is equal to DE.
EXTRA CREDIT (4 POINTS): Consider the following standard heats of formation:
| P4O10 (s) | ΔHfo = -3110 kJ/mol |
| H2O(l) | ΔHfo = -286 kJ/mol |
| H3PO4(s) | ΔHfo = -1279 kJ/mol |
Calculate the change in enthalpy for the following process starting with 50.00 g P4O10(s) :
P4O10(s) + 6H2O(l) 4H3PO4(s)
The above exam is the identical exam to that given in fall 1998. You can also review the following multiple choice questions which covers the material on the exam. This test is set up as multiple choice for most of the questions, so you need to work each of the problems. The correct answer is given as a choice, and the answers for all the questions are listed at the end of this page.
Keys: ΔH = change in enthalpy
liter-atm = Liters x atmospheres (L atm)
1. Calculate the work for the expansion of CO2 from 1.0 to 2.5 liters against a pressure of 1.0 atm at constant temperature.
a. 1.5 L.atm
b. 2.5 L.atm
c. 0
d. -1.5 L.atm
e. -2.5 L.atm
2. One mole of an ideal gas is expanded from a volume of 1.00 liter to a volume of 10.00 liters against a constant external pressure of 1.00 atm. How much work (in joules) is performed on the surroundings? (T = 300 K; 1 L.atm = 101.3 J)
a. 456 J
b. 912 J
c. 2740 J
d. 2870 J
e. none of these
3. Which of the following statements is correct?
a. The internal energy of a system increases when more work is done by the
system than heat was flowing into the system.
b. The internal energy of a system decreases when work is done on the system
and heat is flowing into the system.
c. The system does work on the surroundings when an ideal gas expands against
a constant external pressure.
d. All statements are true.
e. All statements are false.
4. You have a 28.2-g sample of a metal heated to 95.2oC. You drop it in a calorimeter with 100.g of water at 25.1oC. The final temperature of the water is 31.0oC. Assuming no heat loss to the surroundings nor the calorimeter, calculate the heat capacity of the metal.
a. 0.325 J/goC
b. 0.981 J/goC
c. 1.12 J/goC
d. 1.36 J/goC
e. none of these
5. What is the heat capacity of mercury if it requires 167 J to change the temperature of 15.0 g mercury from 25.0oC to 33.0oC?
a. 6.92 x 10-3 J/goC
b. 1.12 x 10-2 J/goC
c. 0.445 J/goC
d. 1.39 J/goC
e. 313 J/goC
6. Consider the following processes:
2A→ 1/2B + C ΔH = 5 kJ/mol
3/2B + 4C → 2A + C + 3D ΔH = -15 kJ/mol
E + 4A → C ΔH = 10 kJ/mol
Calculate ΔH for: C → E + 3D
a. 0 kJ/mol
b. 10 kJ/mol
c. -10 kJ/mol
d. -20 kJ/mol
e. 20 kJ/mol
7. Which of the following does not have a standard enthalpy of formation (ΔHfo) equal to zero at 25oC and 1.0 atm?
a. F2(g)
b. Al(s)
c. H2O(l)
d. H2(g)
e. They all have a ΔHfo equal to zero since these
are their standard states at 25oC and 1.0 atm.
8. Given: Cu2O(s) + 1/2 O2(g) → 2CuO(s) ΔH = -144 kJ
Cu2O(s) → Cu(s) + CuO(s) ΔH = +11 kJ
Calculate the standard enthalpy of formation of CuO(s).
a. -166 kJ
b. -299 kJ
c. +299 kJ
d. +155 kJ
e. -155 kJ
9. Which of the following is not being considered as an energy source for the future?
a. ethanol
b. methanol
c. seed oil
d. shale oil
e. carbon dioxide
10. When ignited, a uranium compound burns with a green flame. The wavelength of the light given off by this flame is greater than that of
a. red light.
b. infrared light.
c. radio waves.
d. ultraviolet light.
e. none of these
11. Which form of electromagnetic radiation has the longest wavelengths?
a. gamma rays
b. microwaves
c. radio waves
d. infrared radiation
e. x-rays
12. Which of the following frequencies corresponds to light with the longest wavelength?
a. 3.00 x 1013 s-1
b. 4.12 x 105 s-1
c. 8.50 x 1020 s-1
d. 9.12 x 1012 s-1
e. 3.20 x 109 s-1
13. Green light has a wavelength of 5.50 x 102 nm. The energy of a photon of green light is
a. 3.64 x 10-38 J
b. 2.17 x 105 J
c. 3.61 x 10-19 J
d. 1.09 x 10-27 J
e. 5.45 x 1012 J
14. In an investigation of the electronic absorption spectrum of a particular element, it is found that a photon having Lambda = 500 nm provides just enough energy to promote an electron from the second quantum level to the third. From this information, we can deduce
a. the energy of the n = 2 level.
b. the energy of the n = 3 level.
c. the sum of the energies of n = 2 and n = 3.
d. the difference in energies between n = 2 and n = 3.
e. all of these
15. What is the minimum wavelength of a photon of light that can excite an electron in the hydrogen atom from the n = 1 to the n = 8 energy level?
a. 92.7 nm
b. 122 nm
c. 40.1 nm
d. 60.4 nm
e. 300. nm
16. When a strontium salt is ignited, it burns with a red flame. The frequency of the light given off by this flame is greater than
a. yellow light
b. infrared light
c. ultraviolet light
d. radio waves
e. X-rays
17. Which of the following statements is true?
a. The exact location of an electron can be determined if we know its
energy.
b. An electron in a 2s orbital can have the same n, l, and
ml quantum numbers as an electron in a 3s orbital.
c. Ni has 2 unpaired electrons in its 3d orbitals.
d. In the buildup of atoms, electrons occupy the 4f orbitals before the 6s
orbitals.
e. Only three quantum numbers are needed to uniquely describe an electron.
18. How many electrons in an atom can have the quantum numbers n = 4, l = 2?
a. 14
b. 12
c. 5
d. 10
e. 6
19. What is the electron configuration of calcium?
a. 1s2 2s2 2p6 3s2 3p6
4s2
b. 1s2 2s2 2p6 2d10
c. 1s2 2s2 2p6 3s2 3p6
3d2
d. 1s2 2s3 2p6 3s3 3p3
4s2
e. 1s2 2s3 2p5 3s4 3p5
4s1
20. The electron configuration of indium is
a. 1s2 2s2 2p6 3s2 3p6
3d10 4s2 4p6 4d10 5s2
5p1 5d10
b. 1s2 2s2 2p6 3s2 3p6
3d10 4s2 4d10 4p1
c. 1s2 3s2 2p6 3s2 3p6
4s2 4p6 4d10 5s2
5d10 5p1
d. 1s2 2s2 2p6 3s2 3p6
3d10 4s2 4p6 4d10 5s2
5p1
e. none of these
21. Which of the following atoms would have the largest second ionization energy?
a. Mg
b. Cl
c. S
d. Ca
e. Na
22. Ti has __________ in its d orbitals.
a. 1 electron
b. 2 electrons
c. 3 electrons
d. 4 electrons
e. none of these
23. Fe has __________ that is (are) unpaired in its d orbitals.
a. 1 electron
b. 2 electrons
c. 3 electrons
d. 4 electrons
e. none of these
24. How many unpaired electrons are there in an atom of sulfur in its ground state?
a. 0
b. 1
c. 2
d. 3
e. 4
25. Which of the following have 10 electrons in the d orbitals?
a. Mn
b. Fe
c. Cu
d. Zn
e. two of these
26. Consider the following standard enthalpies of formation:
P4O10 (s) ΔHfo = -3110
kJ/mol
H2O(l) ΔHfo = -286 kJ/mol
H3PO4(s) ΔHfo = -1279 kJ/mol
Calculate the change in enthalpy for the following process:
P4O10 (s) + 6H2O(l) → 4H3PO4(s)
27. Give the quantum numbers for the last electron in:
a. gold
b. magnesium
c. iodine
d. cadmium
ANSWER KEY FOR TEST
1. d
2. b
3. c
4. d
5. d
6. c
7. c
8. e
9. e
10. d
11. c
12. b
13. c
14. d
15. a
16. b
17. c
18. d
19. a
20. d
21. e
22. b
23. d
24. c
25. e
26. -290 kJ
27. a. gold-5, 2, 2, 1/2
b. magnesium-3, 0, 0, 1/2
c. iodine-5, 1, 1, 1/2
d. cadmium-4, 2, 2, 1/2