Download PX211-0406 The following questions were randomly selected for

PX211-0406
The following questions were randomly selected for topics in Chapters 4 – 6 of the Chem 1411 lecture
outline, Dr. J. Pelezo, LSC-Tomball. Note: The question answers for each are attached at the end of the
practice exam and include selected suggestions for solving the problems and/or question. If needed, please
direct question for assistance to [email protected].
____
1. What is the molecular mass of the hydrocarbon styrene (shown in the figure)?
A)
B)
C)
D)
E)
104.1 amu.
91.1 amu.
103.1 amu.
13.0 amu.
78.1 amu.
____
2. What is the molar mass of zinc sulfate heptahydrate, ZnSO4 • 7H2O?
A) 180. g/mol
B) 288 g/mol
C) 384 g/mol
D) 162 g/mol
E) 582 g/mol
____
3. Which of the following samples contains the largest number of molecules?
A) 10. g Pb
B) 10. g Cl2
C) 10. g Kr
D) 10. g O2
E) 10. g S8
____
4. Which of the following compounds has the highest percentage of nitrogen by mass?
A) (NH4)2SO3
B) NaNO3
C) N2Cl4
D) NH4NO2
E) HNO3
____
5. How many grams of hydrogen atoms are present in 18.4 g of water?
A) 37.1 g
B) 1.02 g
C) 2.06 g
D) 1.96 g
E) 12.3 g
____
6. A compound containing only carbon, hydrogen, and oxygen is subjected to elemental analysis.
Upon complete combustion, a 0.1804-g sample of the compound produced 0.3051 g of CO2 and
0.1249 g of H2O. What is the empirical formula of the compound?
A) C3H6O3
B) C3H3O
C) C4H8O3
D) C2H2O
E) CH2O3
____
7. A compound composed of only C and F contains 17.39 % C by mass. What is its empirical
formula?
A) CF3
B) CF
C) C2F
D) CF4
E) CF2
____
8. The analysis of an organic compound showed that it contained 1.386 mol of C, 0.0660 mol of H,
0.924 mol of O, and 0.462 mol of N. How many nitrogen atoms are there in the empirical formula
for this compound?
A) 9
B) 7
C) 2
D) 4
E) 3
____
9. A sample containing 0.700 mol of a compound is composed of 4.21  1023 atoms of sodium,
24.79 g of chlorine atoms, and 33.57 g of oxygen atoms. The formula of the compound is
A) NaClO3.
B) NaClO5.
C) NaClO.
D) NaClO4.
E) NaClO2.
____
10. Complete combustion of a 0.30-mol sample of a hydrocarbon, CxHy, gives 1.20 mol of CO2 and
1.50 mol of H2O. The molecular formula of the original hydrocarbon is
A) C3H8.
B) C4H10.
C) C8H20.
D) C3H5.
E) C5H7.
____
11. One step in the isolation of pure rhodium metal (Rh) is the precipitation of rhodium(III) hydroxide
from a solution containing rhodium(III) sulfate according to the following balanced chemical
equation:
Rh2(SO4)3(aq) + 6NaOH(aq)  2Rh(OH)3(s) + 3Na2SO4(aq)
If 2.40 g of rhodium(III) sulfate reacts with excess sodium hydroxide, what mass of rhodium(III)
hydroxide may be produced?
A) 1.50 g
B) 4.80 g
C) 2.40 g
D) 0.374 g
E) 2.99 g
____
12. The limiting reactant is the reactant
A) that has the lowest coefficient in the balanced equation.
B) that has the lowest molar mass.
C) that is left over after the reaction has gone to completion.
D) for which there is the lowest mass in grams.
E) none of the above
____
13. The commercial production of phosphoric acid, H3PO4, can be represented by the equation
1500 g
300 g
Ca3(PO4)2 + 3SiO2 +
310 g/mol 60.1
g/mol
307 g
5C +
12.0
g/mol
1180 g
5O2 +
32.0
g/mol
300 g
3H2O
18.0
g/mol
 3CaSiO3 + 5CO2 + 2H3PO4
The molar mass for each reactant is shown below the reactant, and the mass of each reactant for
this problem is given above. Which substance is the limiting reactant?
A) H2O
B) C
C) O2
D) Ca3(PO4)2
E) SiO2
____
14. When 20.0 g C2H6 and 60.0 g O2 react to form CO2 and H2O, how many grams of water are
formed?
A) 14.5 g
B) 58.0 g
C) 18.0 g
D) 20.0 g
E) none of these
____
15. A 5.95-g sample of AgNO3 is reacted with BaCl2 according to the equation
2AgNO3(aq) + BaCl2(aq)  2AgCl(s) + Ba(NO3)2(aq)
to give 3.36 g of AgCl. What is the percent yield of AgCl?
A) 44.6 %
B) 33.5 %
C) 66.9 %
D) 56.5 %
E) 100 %
____
16. A 2.00-L glass soda bottle filled only with air is tightly capped at 24°C and 744.0 mmHg. If the
bottle is placed in water at 75°C, what is the pressure in the bottle?
A) 238 mmHg
B) 872 mmHg
C) 2330 mmHg
D) 635 mmHg
E) 383 mmHg
____
17. When the valve between the 2.00-L bulb, in which the gas pressure is 3.00 atm, and the 3.00-L
bulb, in which the gas pressure is 2.50 atm, is opened, what will be the final pressure in the two
bulbs? Assume the temperature remains constant.
A)
B)
C)
D)
E)
____
2.70 atm
5.50 atm
2.83 atm
2.80 atm
2.67 atm
18. Which of the following is a correct statement of Charles’s law,
A)
B)
C)
D)
E)
?
The volume of a gas varies proportionally with the pressure.
The volume of a gas sample varies directly with the absolute temperature.
All gas samples of the same volume at STP contain the same number of atoms.
The pressure of a gas sample varies inversely with the volume.
All gas samples of the same volume at STP contain the same number of molecules.
____
19. The pressure of 4.2 L of nitrogen gas in a flexible container is decreased to one-half its original
pressure, and its absolute temperature is increased to double the original temperature. The volume
is now
A) 2.1 L.
B) 4.2 L.
C) 8 L.
D) 17 L.
E) 1.1 L.
____
20. A fixed amount of gas in a rigid container is heated from 100°C to 700°C. Which of the following
responses best describes what will happen to the pressure of the gas?
A) The pressure will increase by a factor greater than 7.
B) The pressure will increase by a factor of 7.
C) The pressure will increase by a factor less than 7.
D) The pressure will decrease by a factor of 7.
E) The pressure will remain the same.
____
21. The volume of 1 mol of nitrogen
A) is lower than that of 1mol ammonia at high pressures.
B) is decreased by decreasing the pressure of the gas.
C) has the value of 22.4 L at 0°C and 1.00 atm.
D) is decreased by increasing its kinetic energy.
E) is increased by decreasing the temperature.
____
22. Which of the following graphs does not correctly describe the ideal gas law?
A)
B) They all correctly represent the ideal gas law.
C)
D)
E)
____
23. What is the pressure of a 59.6-L gas sample containing 3.01 mol of gas at 44.9°C?
(R = 0.0821 L • atm/(K • mol), 1 atm = 760 torr)
A) 1.41  102 mmHg
B) 1.73  10–3 mmHg
C) 1.32 mmHg
D) 1.00  103 mmHg
E) 5.77  102 mmHg
____
24. The density of ethane, C2H6 (30.1 g/mol), at 32°C and 1.31 atm pressure is
A) 1.57 g/L.
B) 19.2 g/L.
C) 1.34 g/L.
D) 0.635 g/L.
E) 0.162 g/L.
____
25. At 28°C and 452 mmHg, an unknown pure gas has a density of 0.674 g/L. Which of the following
gases could be the unknown gas?
A) F2
B) C3H6
C) N2O
D) Ne
E) N2
____
26. Calcium nitrate will react with ammonium chloride at slightly elevated temperatures, as
represented in the equation below.
Ca(NO3)2(s) + 2NH4Cl(s)  2N2O(g) + CaCl2(s) + 4H2O(g)
What is the maximum volume of N2O at STP that could be produced using a 3.40-mol sample of
each reactant?
A) 9.28  102 L
B) 152 L
C) 1.31  10–2 L
D) 76.2 L
E) 22.4 L
____
27. A 24.1-g mixture of nitrogen and carbon dioxide is found to occupy a volume of 15.1 L when
measured at 870.2 mmHg and 31.2oC. What is the mole fraction of nitrogen in this mixture?
A) 0.539
B) 0.500
C) 0.461
D) 0.426
E) 0.574
____
28. A sample of oxygen is collected over water at a total pressure of 690.7 mmHg at 19°C. The vapor
pressure of water at 19°C is 16.5 mmHg. The partial pressure of the O2 is
A) 0.9305 atm.
B) 0.9349 atm.
C) 0.9088 atm.
D) 1.070 atm.
E) 0.8871 atm.
____
29. Which of the following statements is least likely to be true of a sample of nitrogen gas at STP?
A) Collisions between the gaseous molecules are elastic.
B) The intermolecular forces between nitrogen molecules are not negligible.
C) Molecules of gaseous nitrogen are in constant random motion.
D) The average kinetic energy of the gaseous nitrogen is proportional to the absolute
temperature of the gas.
E) The pressure exerted by gaseous nitrogen is due to collisions of the molecules with
the walls of the container.
____
30. At STP, as the molar mass of the molecules that make up a pure gas increases, the
A) root mean square speed of the molecules increases.
B) root mean square speed of the molecules decreases.
C) root mean square speed of the molecules remains constant.
D) root mean square speed increases to a maximum, then decreases.
E) none of the above.
____
31. Molecular speed distributions for a gas at two different temperatures are shown below. Which of
the following graphs correctly describes the distributions at the two temperatures, where T2 > T1?
NOTE: The small vertical lines indicate average speed.
A)
B)
C)
D)
E) none of the above
____
32. Which of the following relates the rate of effusion of a gas to the square root of its molar mass?
A) Boyle’s law
B) Graham’s law
C) Charles’s law
D) Dalton’s law
E) Avogadro’s hypothesis
____
33. If 250 mL of methane, CH4, effuses through a small hole in 28 s, the time required for the same
volume of helium to pass through the hole under the same conditions will be
A) 56 s.
B) 7 s.
C) 1.8 s.
D) 14 s.
E) 112 s.
____
34. Using the van der Waals equation, determine the pressure of 459.0 g of SO2(g) in a 4.40-L vessel
at 625 K. For SO2(g), a = 6.865 L2 • atm/mol2 and b = 0.05679 L/mol.
(R = 0.0821 L • atm/(K • mol))
A) 110 atm
B) 8.06 atm
C) 73.9 atm
D) 83.6 atm
E) 8.36 atm
____
35. The air whipped up by a tornado possesses what type(s) of energy?
A) potential energy only
B) internal energy only
C) kinetic energy only
D) kinetic energy, potential energy, and internal energy
E) kinetic energy and potential energy only
____
36. Calculate U of a gas for a process in which the gas absorbs 9 J of heat and does 25 J of work by
expanding.
A) 16 J
B) 34 J
C) –34 J
D) 0, because U is a state function
E) –16 J
____
37. The phrase “the heat absorbed or released by a system undergoing a physical or chemical change
at constant pressure” is
A) the change in enthalpy of the system.
B) the change in internal energy of the system.
C) the definition of a state function.
D) the temperature change of the system.
E) a statement of Hess’s law.
____
38. What is the quantity of heat evolved at constant pressure when 60.3 g H2O(l) is formed from the
combustion of H2(g) and O2(g)?
H2(g) + O2(g)  H2O(l); H° = –285.8 kJ
A)
B)
C)
D)
E)
____
1.17  10–2 kJ
285.8 kJ
1.72  104 kJ
85.4 kJ
9.57  102 kJ
39. What mass of hydrogen is consumed when 587.9 kJ of energy is evolved from the combustion of a
mixture of H2(g) and O2(g)?
H2(g) + O2(g)  H2O(l); H° = –285.8 kJ
A)
B)
C)
D)
E)
____
4.147 g
2.073 g
0.2412 g
6.162 g
2.131 g
40. How much heat is liberated at constant pressure when 58.5 g of calcium oxide reacts with 83.9 L
of carbon dioxide gas, measured at 1.00 atm pressure and 25.0°C? (R = 0.0821 L • atm/(K • mol))
CaO(s) + CO2(g)  CaCO3(s); H° = –178.3 kJ
A)
B)
C)
D)
E)
____
–6.11  102 kJ
–1.04  104 kJ
–7.97  102 kJ
–1.86  102 kJ
–1.50  104 kJ
41. A 5.09-g sample of solid silver reacted in excess chlorine gas to give a 6.76-g sample of pure solid
AgCl. The heat given off in this reaction was 6.00 kJ at constant pressure. Given this information,
what is the enthalpy of formation of AgCl(s)?
A) –127 kJ/mol
B) –63.6 kJ/mol
C) 127 kJ/mol
D) –6.00 kJ/mol
E) 6.00 kJ/mol
____
42. The heat required to raise the temperature of 52.00 g of chromium by 1°C is called its
A) heat of vaporization.
B) specific heat.
C) heat of fusion.
D) entropy.
E) molar heat capacity.
____
43. The molar heat capacity of gaseous heptane at 25.0°C is 165.2 J/(mol°C). What is its specific
heat?
A) 0.6065 J/(g°C)
B) 1.649 J/(g°C)
C) 6.041  10–5 J/(g°C)
D) 165.2 J/(g°C)
E) 1.655  104 J/(g°C)
____
44. A 170.0-g sample of metal at 79.00°C is added to 170.0 g of H2O(l) at 14.00°C in an insulated
container. The temperature rises to 16.19°C. Neglecting the heat capacity of the container, what
is the specific heat of the metal? The specific heat of H2O(l) is 4.18 J/(g°C).
A) 4.18 J/(g°C)
B) 120 J/(g°C)
C) 0.146 J/(g°C)
D) –0.146 J/(g°C)
E) 28.6 J/(g°C)
____
45. A 94.7-g sample of silver (s = 0.237 J/(g°C)), initially at 348.25°C, is added to an insulated
vessel containing 143.6 g of water (s = 4.18 J/(g°C)), initially at 13.97°C. At equilibrium, the
final temperature of the metal–water mixture is 22.63°C. How much heat was absorbed by the
water? The heat capacity of the vessel is 0.244 kJ/°C.
A) 5.20 kJ
B) 3.09 kJ
C) 7.31 kJ
D) 9.12 kJ
E) 129 kJ
____
46. In a bomb calorimeter, reactions are carried out
A) at 1 atm pressure and 0°C.
B) at a constant pressure.
C) at a constant volume.
D) at a constant pressure and 25°C.
E) at 1 atm pressure and 25°C.
____
47. Which of the following has a standard enthalpy of formation value of zero at 25°C?
A) Cl(g)
B) Cl2(l)
C) Cl2(g)
D) Cl(s)
E) Cl2(s)
____
48. What is the standard enthalpy of formation of liquid butyraldehyde, CH3CH2CH2CHO(l)?
CH3CH2CH2CHO(l) + O2(g)  4H2O(l) + 4CO2(g); H° = –2471.8 kJ
Substance
CO2(g)
H2O(l)
A)
B)
C)
D)
E)
____
H°f (kJ/mol)
–393.5
–285.8
–245.4 kJ/mol
+245.4 kJ/mol
–1792.5 kJ/mol
–3151.1 kJ/mol
+3151.1 kJ/mol
49. What is the standard enthalpy change for the following reaction?
N2H4(l) + 2NO2(g)  2N2O(g) + 2H2O(l)
Substance
N2H4(l)
NO2(g)
N2O(g)
H2O(l)
A)
B)
C)
D)
E)
____
H°f (kJ/mol)
+50.6
+33.1
+82.1
–285.8
–290.6 kJ
–524.2 kJ
–119.7 kJ
+290.6 kJ
+119.7 kJ
50. From the following information, determine the enthalpy of formation of C2H4(g).
C2H4(g)  C(s) + H2(g); H = –26.2 kJ
A) –26.2 kJ/mol
B) 26.2 kJ/mol
C) 104.8 kJ/mol
D) –52.4 kJ/mol
E) 52.4 kJ/mol
PX211-0406
Answer Section
1. ANS: A
PTS: 1
DIF: easy
REF: 3.1
OBJ: Calculate the formula mass from molecular models. (Example 3.2)
TOP: stoichiometry | mass and moles of substance
KEY: molecular mass
MSC: general chemistry
2. ANS: B
PTS: 1
DIF: easy
REF: 3.2
OBJ: Understand how the molar mass is related to the formula weight of a substance.
TOP: stoichiometry | mass and moles of substance
KEY: formula mass
MSC: general chemistry
3. ANS: D
PTS: 1
DIF: moderate
REF: 3.2
OBJ: Calculate the number of molecules in a given mass of substance. (Example 3.6)
TOP: stoichiometry | mass and moles of substance
KEY: mole | mole calculations
MSC: general chemistry
4. ANS: D
PTS: 1
DIF: easy
REF: 3.3
OBJ: Calculate the percentage composition of the elements in a compound. (Example 3.7)
TOP: stoichiometry | determining chemical formulas
KEY: mass percentage
MSC: general chemistry
5. ANS: C
PTS: 1
DIF: easy
REF: 3.3
OBJ: Calculate the mass of an element in a given mass of compound. (Example 3.8)
TOP: stoichiometry | mass and moles of substance
KEY: mole | mole calculations
MSC: general chemistry
6. ANS: C
PTS: 1
DIF: moderate
REF: 3.4
OBJ: Calculate the percentage of C, H, and O from combustion data. (Example 3.9)
TOP: stoichiometry | determining chemical formulas
KEY: elemental analysis
MSC: general chemistry
7. ANS: A
PTS: 1
DIF: easy
REF: 3.5
OBJ: Determine the empirical formula from the percentage composition. (Example 3.11)
TOP: stoichiometry | determining chemical formulas
KEY: empirical formula
MSC: general chemistry
8. ANS: B
PTS: 1
DIF: easy
REF: 3.5
OBJ: Determine the empirical formula from the percentage composition. (Example 3.11)
TOP: stoichiometry | determining chemical formulas
KEY: empirical formula
MSC: general chemistry
9. ANS: A
PTS: 1
DIF: moderate
REF: 3.5
OBJ: Determine the empirical formula from the percentage composition. (Example 3.11)
TOP: stoichiometry | determining chemical formulas
KEY: empirical formula
MSC: general chemistry
10. ANS: B
PTS: 1
DIF: difficult
REF: 3.5
OBJ: Determine the molecular formula from the percentage composition and molecular mass.
(Example 3.12)
TOP: stoichiometry | determining chemical formulas
KEY: molecular formula
MSC: general chemistry
11. ANS: A
PTS: 1
DIF: easy
REF: 3.7
OBJ: Relate the quantities of reactant to the quantity of product. (Example 3.13)
TOP: stoichiometry | stoichiometry calculation
KEY: amounts of substances
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
MSC: general chemistry
ANS: E
PTS: 1
DIF: easy
REF: 3.8
OBJ: Understand how a limiting reactant or limiting reagent determines the moles of product
formed during a chemical reaction and how much excess reactant remains.
TOP: stoichiometry | stoichiometry calculation
KEY: limiting reactant
MSC: general chemistry
ANS: E
PTS: 1
DIF: moderate
REF: 3.8
OBJ: Calculate with a limiting reactant involving masses. (Example 3.16)
TOP: stoichiometry | stoichiometry calculation
KEY: limiting reactant
MSC: general chemistry
ANS: E
PTS: 1
DIF: moderate
REF: 3.8
OBJ: Calculate with a limiting reactant involving masses. (Example 3.16)
TOP: stoichiometry | stoichiometry calculation
KEY: limiting reactant
MSC: general chemistry
ANS: C
PTS: 1
DIF: moderate
REF: 3.8
OBJ: Determine the percentage yield of a chemical reaction.
TOP: stoichiometry | stoichiometry calculation
MSC: general chemistry
ANS: B
PTS: 1
DIF: easy
REF: 5.2
OBJ: Use Boyle's law. (Example 5.2)
TOP: phases | gas
KEY: empirical gas laws | Boyle's law
MSC: general chemistry
ANS: A
PTS: 1
DIF: moderate
REF: 5.2
OBJ: Use Boyle's law. (Example 5.2)
TOP: phases | gas
KEY: empirical gas laws | Boyle's law
MSC: general chemistry
ANS: B
PTS: 1
DIF: easy
REF: 5.2
OBJ: Express Charles's law in words and as an equation.
TOP: phases | gas
KEY: empirical gas laws | Charles's law MSC: general chemistry
ANS: D
PTS: 1
DIF: easy
REF: 5.2
OBJ: Use the combined gas law. (Example 5.4)
TOP: phases | gas
KEY: empirical gas laws | combined gas law
MSC: general chemistry
ANS: C
PTS: 1
DIF: easy
REF: 5.2
OBJ: Use the combined gas law. (Example 5.4)
TOP: phases | gas
KEY: empirical gas laws | combined gas law
MSC: general chemistry
ANS: C
PTS: 1
DIF: easy
REF: 5.3
OBJ: Learn the ideal gas law equation.
TOP: phases | gas KEY: ideal gas law
MSC: general chemistry
ANS: B
PTS: 1
DIF: moderate
REF: 5.3
OBJ: Derive the empirical gas laws from the ideal gas law. (Example 5.5)
TOP: phases | gas KEY: ideal gas law
MSC: general chemistry
ANS: D
PTS: 1
DIF: easy
REF: 5.3
OBJ: Use the ideal gas law. (Example 5.6)
TOP: phases | gas
KEY: ideal gas law | calculations with the ideal gas law
MSC: general chemistry
ANS: A
PTS: 1
DIF: moderate
REF: 5.3
OBJ: Calculate gas density. (Example 5.7)
TOP: phases | gas
KEY: ideal gas law | gas density
MSC: general chemistry
ANS: E
PTS: 1
DIF: moderate
REF: 5.3
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
OBJ: Determine the molecular mass of a vapor. (Example 5.8) TOP: phases | gas
KEY: ideal gas law | gas density
MSC: general chemistry
ANS: D
PTS: 1
DIF: difficult
REF: 5.4
OBJ: Solving stoichiometry problems involving gas volumes. (Example 5.9)
TOP: phases | gas KEY: ideal gas law | stoichiometry and gas volumes
MSC: general chemistry
ANS: E
PTS: 1
DIF: difficult
REF: 5.5
OBJ: Calculate the partial pressure and mole fractions of a gas in a mixture. (Example 5.10)
TOP: phases | gas KEY: gas mixtures MSC: general chemistry
ANS: E
PTS: 1
DIF: easy
REF: 5.5
OBJ: Calculate the amount of gas collected over water. (Example 5.11)
TOP: phases | gas KEY: gas mixtures | collecting gases over water
MSC: general chemistry
ANS: B
PTS: 1
DIF: easy
REF: 5.6
OBJ: List the five postulates of the kinetic theory.
TOP: phases | gas
KEY: kinetic-molecular theory
MSC: general chemistry
ANS: B
PTS: 1
DIF: easy
REF: 5.7
OBJ: Describe how the root-mean-square (rms) molecular speed of gas molecules varies with
temperature.
TOP: phases | gas
ANS: A
PTS: 1
DIF: moderate
REF: 5.7
OBJ: Describe the molecular-speed distribution of gas molecules at different temperatures.
TOP: phases | gas KEY: molecular speed
MSC: general chemistry
ANS: B
PTS: 1
DIF: easy
REF: 5.7
OBJ: Define effusion and diffusion.
TOP: phases | gas
KEY: molecular speed | effusion
MSC: general chemistry
ANS: D
PTS: 1
DIF: moderate
REF: 5.7
OBJ: Calculate the ratio of effusion rates of gases. (Example 5.13)
TOP: phases | gas KEY: molecular speed | effusion
MSC: general chemistry
ANS: C
PTS: 1
DIF: moderate
REF: 5.8
OBJ: Use the van der Waals equation. (Example 5.14)
TOP: phases | gas
KEY: real gases
MSC: general chemistry
ANS: D
PTS: 1
DIF: easy
REF: 6.1
OBJ: Define energy, kinetic energy, potential energy, and internal energy.
TOP: thermochemistry | heats of reaction
ANS: E
PTS: 1
DIF: easy
REF: 6.2
OBJ: State the law of conservation of energy.
TOP: thermochemistry | heats of reaction KEY: energy | law of conservation of energy
MSC: general chemistry
ANS: A
PTS: 1
DIF: easy
REF: 6.3
OBJ: Explain how the terms enthalpy of reaction and heat of reaction are related.
TOP: thermochemistry | heats of reaction KEY: enthalpy | enthalpy change
MSC: general chemistry
ANS: E
PTS: 1
DIF: easy
REF: 6.5
OBJ: Calculate the heat absorbed or evolved from a reaction given its enthalpy of reaction and
the mass of a reactant or product. (Example 6.4)
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
TOP: thermochemistry | heats of reaction
KEY: thermochemical equation | stoichiometry and heats of reaction
MSC: general chemistry
ANS: A
PTS: 1
DIF: easy
REF: 6.5
OBJ: Calculate the heat absorbed or evolved from a reaction given its enthalpy of reaction and
the mass of a reactant or product. (Example 6.4)
TOP: thermochemistry | heats of reaction
ANS: D
PTS: 1
DIF: difficult
REF: 6.5
OBJ: Calculate the heat absorbed or evolved from a reaction given its enthalpy of reaction and
the mass of a reactant or product. (Example 6.4)
TOP: thermochemistry | heats of reaction
KEY: thermochemical equation | stoichiometry and heats of reaction
MSC: general chemistry
ANS: A
PTS: 1
DIF: difficult
REF: 6.5
OBJ: Calculate the heat absorbed or evolved from a reaction given its enthalpy of reaction and
the mass of a reactant or product. (Example 6.4)
TOP: thermochemistry | heats of reaction
KEY: thermochemical equation | stoichiometry and heats of reaction
MSC: general chemistry
ANS: E
PTS: 1
DIF: easy
REF: 6.6
OBJ: Define heat capacity and specific heat.
TOP: thermochemistry | heats of reaction KEY: calorimetry | heat capacity
MSC: general chemistry
ANS: B
PTS: 1
DIF: easy
REF: 6.6
OBJ: Define heat capacity and specific heat.
TOP: thermochemistry | heats of reaction KEY: calorimetry | heat capacity
MSC: general chemistry
ANS: C
PTS: 1
DIF: difficult
REF: 6.6
OBJ: Calculate using this relation between heat and specific heat. (Example 6. 5)
TOP: thermochemistry | heats of reaction KEY: calorimetry | specific heat
MSC: general chemistry
ANS: A
PTS: 1
DIF: difficult
REF: 6.6
OBJ: Calculate using this relation between heat and specific heat. (Example 6. 5)
TOP: thermochemistry | heats of reaction KEY: calorimetry | specific heat
MSC: general chemistry
ANS: C
PTS: 1
DIF: easy
REF: 6.6
OBJ: Define calorimeter.
TOP: thermochemistry | heats of reaction
KEY: calorimetry | measuring heats of reaction
MSC: general chemistry
ANS: C
PTS: 1
DIF: easy
REF: 6.8
OBJ: Define standard state and reference form.
TOP: thermochemistry | heats of reaction KEY: standard enthalpies of formation
MSC: general chemistry
ANS: A
PTS: 1
DIF: moderate
REF: 6.8
OBJ: Calculate the heat (enthalpy) of reaction from the standard enthalpies of formation of the
substances in the reaction. (Example 6.9) TOP: thermochemistry | heats of reaction
KEY: standard enthalpies of formation
MSC: general chemistry
49. ANS: B
PTS: 1
DIF: moderate
REF: 6.8
OBJ: Calculate the heat (enthalpy) of reaction from the standard enthalpies of formation of the
substances in the reaction. (Example 6.9) TOP: thermochemistry | heats of reaction
KEY: standard enthalpies of formation
MSC: general chemistry
50. ANS: E
PTS: 1
DIF: easy
REF: 6.8
OBJ: Calculate the heat (enthalpy) of reaction from the standard enthalpies of formation of the
substances in the reaction. (Example 6.9) TOP: thermochemistry | heats of reaction
KEY: standard enthalpies of formation
MSC: general chemistry