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| 1 | |
A certain gas is compressed in a cylinder by a constant pressure of 9.0 atm from a volume of 35.0 L to 4.00 L. Calculate the work (w) for this process. |
| | A) | 2.8 x 104 J |
| | B) | 1.0 x 104 J |
| | C) | -2.8 x 102 J |
| | D) | -2.8 x 104 J |
| | E) | 2.8 x 102 J |
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| 2 | |
In a certain process, 750 J of work is done on a system which gives off 195 J of heat. The value of ΔE for the process is |
| | A) | -555 J |
| | B) | 750 J |
| | C) | 555 J |
| | D) | -195 J |
| | E) | 195 J |
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| 3 | |
A gas is allowed to expand at constant temperature from a volume of 2.0 L to 11.2 L against an external pressure of 0.750 atm. If the gas absorbs 128 J of heat from the surroundings, what are the values of q, w, and ΔE respectively? |
| | A) | 128 J, 6.9 J, 135 J |
| | B) | 128 J, -6.9 J, 121 J |
| | C) | 128 J, 697 J, 825 J |
| | D) | 128 J, -697 J, -569 J |
| | E) | -128 J, -6.9 J, -135 J |
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| 4 | |
Two solutions (the system), each of 25.0 mL volume and at 25.0°C, are mixed in a beaker. A reaction occurs between them, causing the temperature to drop to 20.0°C. After the products have equilibrated with the surroundings, the temperature is again 25.0°C and the total volume is 50.0 mL. No gases are involved in the reaction. Which one of the following relationships concerning the change from initial to final states (both at 25.0°C) is correct? |
| | A) | ΔE = 0 |
| | B) | ΔH = 0 |
| | C) | ΔE < 0 |
| | D) | w = 0 |
| | E) | q = 0 |
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| 5 | |
Which one of the following processes is exothermic? |
| | A) | H2(l) → H2(g) |
| | B) | CO2(s) → CO2(g) |
| | C) | H2O(g) → H2O(l) |
| | D) | 16CO2(g) + 18H2O(l) → 2C8H18(l) + 25O2(g) |
| | E) | H2(g) → 2H(g) |
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| 6 | |
Combustion of a hydrocarbon will release more energy than combustion of an equal mass of carbohydrate because |
| | A) | the C-O and O-H single bonds in carbohydrates are weak. |
| | B) | carbohydrates contain C-O and O-H bonds and hydrocarbons do not. |
| | C) | the total energy of the C-C and C-H bonds in hydrocarbons is greater than the total energy of the C=O and O-H bonds in the combustion products (carbon dioxide and water). |
| | D) | the total energy of the bonds in hydrocarbons is greater than the energy of the bonds in carbohydrates. |
| | E) | hydrocarbons have higher molar masses than carbohydrates. |
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| 7 | |
A 24.8 g piece of copper metal is heated from 22.5°C to 178.6°C. Calculate the heat absorbed by the metal. The specific heat of Cu is 0.385 J/g°C |
| | A) | 208 J |
| | B) | 9.5 J |
| | C) | 60.1 J |
| | D) | 1.49 x 103 J |
| | E) | 1.70 x 103 J |
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| 8 | |
The reaction of cyanamide, NH2CN(s), with oxygen was run in a bomb calorimeter, and ΔE was found to be -742.7 kJ/mol of NH2CN(s) at 298 K. Calculate ΔH for the reaction at this temperature.
Reaction: NH2CN(s) + 3/2O2(g) → N2(g) + CO2(g) + H2O(l) |
| | A) | -727.8 kJ/mol |
| | B) | -737.7 kJ |
| | C) | -741.5 kJ |
| | D) | -742.7 kJ |
| | E) | none of the above |
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| 9 | |
A quantity of 1.535 g of liquid methanol (CH3OH) was burned in a constant-volume bomb calorimeter. Consequently the temperature of the water rose from 20.27°C to 26.87°C. If the mass of water surrounding the calorimeter was exactly 1000 g and the heat capacity of the bomb calorimeter was 1.75 kJ/°C, calculate the molar heat of combustion of liquid methanol. |
| | A) | -8.17 x 105 kJ/mol |
| | B) | -817 kJ/mol |
| | C) | 1.88 kJ/mol |
| | D) | -0.8 kJ/mol |
| | E) | none of the above |
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| 10 | |
Data: C(s) + O2(g) → CO2(g) ΔH°rxn = -393 kJ
How many grams of C(s) must burn in this way to release 275 kJ of heat? |
| | A) | 0.70 g |
| | B) | 1.43 g |
| | C) | 8.40 g |
| | D) | 17.1 g |
| | E) | 275 g |
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| 11 | |
The combustion of 5.00 grams of C2H6(g), at constant pressure releases 259 kJ of heat.
What is ΔH for the reaction: 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(l)? |
| | A) | -259 kJ |
| | B) | -518 kJ |
| | C) | -1295 kJ |
| | D) | -1554 kJ |
| | E) | -3108 kJ |
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| 12 | |
Data:
C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l) ΔH°f = -1411 kJ
C(s) + O2(g) → CO2(g) ΔH°f = -393.5 kJ
H2(g) + ½O2(g) → H2O(l) ΔH°f = -285.8 kJ
Use the data given above to find the standard enthalpy of formation of ethylene, C2H4(g). |
| | A) | 731 kJ/mol |
| | B) | 87 kJ/mol |
| | C) | 1.41 x 103 kJ/mol |
| | D) | 52 kJ/mol |
| | E) | 2.77 x 103 kJ/mol |
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| 13 | |
Data:
2Fe(s) + O2(g) → 2FeO(s) ΔH° = -544.0 kJ
4Fe(s) + 3O2(g) → 2Fe2O3(s) ΔH° = -1648.4 kJ
Fe3O4(s) → 3Fe(s) + 2O2(g) ΔH° = +1118.4 kJ
Given the data above, determine the heat of reaction, ΔH°, for the reaction below:
Reaction: Fe2O3(s) + FeO(s) → Fe3O4(s) |
| | A) | -1074.0 kJ |
| | B) | -22.2 kJ |
| | C) | +22.2 kJ |
| | D) | +249.8 kJ |
| | E) | +1074.0 kJ |
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| 14 | |
Which of the following does not have a ΔH°f of zero under standard state conditions? |
| | A) | H2(g) |
| | B) | Li(s) |
| | C) | C (graphite) |
| | D) | Br2(l) |
| | E) | O3(g) |
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| 15 | |
Data:
2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(l) ΔH° = -11,200 kJ
Standard enthalpies of formation, ΔH°f:
CO2(g), -393.5 kJ/mol; H2O(l), -285.8 kJ/mol
Use the above data to calculate the standard enthalpy of formation of octane, C8H18(l). |
| | A) | -120 kJ/mol |
| | B) | -240 kJ/mol |
| | C) | -5144 kJ/mol |
| | D) | -6296 kJ/mol |
| | E) | -11,440 kJ/mol |
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