Chapter 7
Q 150 go to graph
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Given the standard heats of formation for the following compounds, calculate for the following reaction. CH4(g) + H2O(g) → CH3OH + H2(g) (kJ/mol) −75 −242 −238 0 a. +79 kJ b. −79 kJ c. +594 kcal d. −594 kcal e. −405 kJ
a. +79 kJ
A 8.56-g sample of solid silver reacted in excess chlorine gas to give a 11.4-g sample of pure solid AgCl. The heat given off in this reaction was 10.1 kJ at constant pressure. Given this information, what is the enthalpy of formation of AgCl(s)? a. -127 kJ/mol rxn b. -63.6 kJ/mol rxn c. 127 kJ/mol rxn d. -10.1 kJ/mol rxn e. 10.1 kJ/mol rxn
a. -127 kJ/mol rxn
Using the following thermochemical equation: 2NH3(g) + 3N2O(g) → 4N2(g) + 3H2O(g) ΔH = -880 kJ How much energy is released when 6.22g of ammonia, NH3, reacts with excess dinitrogen monoxide, N2O? a. -161 kJ b. -321 kJ c. -623 kJ d. -2740 kJ e. -5474 kJ
a. -161 kJ
An automobile engine provides 637 Joules of work to push the pistons. In this process the internal energy changes by -2767 Joules. Calculate the amount of heat that must be carried away by the cooling system. a. -2130 J b. 2130 J c. -3404 J d. 3404 J e. -2767 J
a. -2130 J
Using the standard heats of formation given in the table below, calculate the standard enthalpy change for the following reaction: 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) Compound ΔfHo (kJ/mol) NH3(g) -46.1 NO(g) 90.3 H2O(g) -241.8 a. -905.2 kJ b. -105.4 kJ c. -1033.6 kJ d. -1169.2 kJ e. -769.6 kJ
a. -905.2 kJ
How much heat is released or absorbed in the reaction of 10.0 grams of SiO2 (quartz) with excess hydrofluoric acid? SiO2(s) + 4HF(aq) → SiF4(g) + 2H2O (kJ/mol) −910.9 −320.8 −1615 −285.8 a. 1.25 kJ absorbed b. 1.25 kJ released c. 11.3 kJ absorbed d. 11.3 kJ released e. 6.56 kJ released
a. 1.25 kJ absorbed
How much heat would be released or absorbed if 575 g of H2 are produced? CH4(g) + H2O(g) → 3H2(g) + CO(g) = 205.9 kJ a. 1.97 × 104 kJ b. 5.90 × 104 kJ c. 3.54 × 105 kJ d. 7.08 × 105 kJ e. −1.97 × 105 kJ
a. 1.97 × 104 kJ
Calculate the enthalpy change for the reaction: 2 NiCl2(s) + O2(g) → 2NiO(s) + 2Cl2(g) Given the enthalpy changes for the following two reactions: 2Ni(s) + O2(g) → 2NiO(s) ΔHo = -479.4 kJ Ni(s) + Cl2(g) → NiCl2(s) ΔHo = -305.3 kJ a. 131.2 kJ b. -174.1 kJ c. 610.6 kJ d. -784.7 kJ e. 1090.0 kJ
a. 131.2 kJ
How much energy is required to decompose 765 g of PCl3, according to the following reaction: 4PCl3(g) → P4(s) + 6Cl2(g) ΔH = 1207 kJ a. 1680 kJ b. 2310 kJ c. 4330 kJ d. 5950 kJ e. 6720 kJ
a. 1680 kJ
An exothermic reaction liberates 7.58 kJ of heat in a coffee cup calorimeter containing 157 grams of solution. The temperature of the solution in the calorimeter increases by 11.2°C. How much heat was absorbed by the calorimeter? Assume the specific heat of the solution is 4.184 J/g•°C. a. 223 J b. 7.36 kJ c. 657 J d. 5820 J e. −223 J
a. 223 J
Calculate the amount of heat released in the complete combustion of 8.17 grams of Al to form Al2O3(s) at 25°C and 1 atm. for Al2O3(s) = −1676 kJ/mol 4Al(s) + 3O2(g) → 2Al2O3(s) a. 254 kJ b. 203 kJ c. 127 kJ d. 237 kJ e. 101 kJ
a. 254 kJ
Use average bond enthalpies (given in the table below) to calculate the enthalpy change for the following gas-phase reaction: Br2(g) + Cl2(g) → 2BrCl(g) Bond kJ/mol Br-Br 193 Cl-Cl 242 Br-Cl 216 a. 3 kJ b. -3 kJ c. 219 kJ d. -219 kJ e. 652 kJ
a. 3 kJ
What is the change in internal energy of the system (ΔE) if 18 kJ of heat energy is evolved by the system and 21 kJ of work is done on the system for a certain process? a. 3 kJ b. -39 kJ c. -18 kJ d. -3 kJ e. 39 kJ
a. 3 kJ
Given the standard enthalpy changes for the following two reactions: 4 C(s) + 5 H2(g) → C4H10(g) ΔHo = -125.6 kJ C2H2(g) → 2 C(s) + H2(g) ΔHo = 226.7 kJ What is the standard enthalpy change for the reaction: 2 C2H2(g) + 3 H2(g) → C4H10(g) ΔHo = ? a. 327.8 kJ b. 101.1 kJ c. 579.0 kJ d. 526.9 kJ e. -352.3 kJ
a. 327.8 kJ
When 32.1 g of H2 reacts with excess silicon to form SiH4(g) at standard conditions, 270.1 kJ of heat is absorbed. What is the DeltaHf^0 for SiH4? a. 33.7 kJ/mol b. 67.3 kJ/mol c. −33.7 kJ/mol d. 8.41 kJ/mol e. −67.3 kJ/mol
a. 33.7 kJ/mol
Calculate the standard heat of vaporization, , for tin(IV) chloride, SnCl4, in kJ per mole. = −511.3 kJ/mol for SnCl4 and −471.5 kJ/mol for SnCl4(g). a. 39.8 b. 16.4 c. 26.4 d. 44.8 e. 53.2
a. 39.8
Given: H−H bond energy = 435 kJ, Cl−Cl bond energy = 243 kJ, and the standard heat of formation of HCl(g) is −92 kJ/mol, calculate the H−Cl bond energy. a. 431 kJ b. 247 kJ c. 180 kJ d. 4.6 kJ e. 326 kJ
a. 431 kJ
Use the data below to calculate for benzene, C6H6 , at 25°C and 1 atm. 2C6H6 + 15O2(g) → 12CO2(g) + 6H2O ΔH0 = −6535 kJ = −393.5 kJ/mol, = −285.8 kJ/mol a. 49.1 kJ/mol b. 3.51 × 104 kJ/mol c. 103 kJ/mol d. 1.76 × 103 kJ/mol e. 561 kJ/mol
a. 49.1 kJ/mol
In the laboratory, a student burns a 0.415g sample of phenanthrene (C14H10) in a bomb calorimeter. The temperature increased from 25.90oC to 28.90oC. Calculate the heat capacity of the calorimeter, if the heat of combustion of phenanthrene is 7054 kJ/mol. a. 5.47 kJ/oC b. 49.3 kJ/oC c. 656 kJ/oC d. 976 kJ/oC e. 2350 kJ/oC
a. 5.47 kJ/oC
The following information is given for chromium at 1 atm: boiling point = 2672oC ΔHvap = 305 kJ/mol melting point = 1857oC ΔHfus = 14.6 kJ/mol specific heat solid = 0.460 J/goC specific heat liquid = 0.937 J/goC What is ΔH in kJ for the process of freezing a 24.7 g sample of liquid chromium at its normal melting point of 1857oC? a. 6.94 kJ b. 9.26 kJ c. 18.9 kJ d. 145 kJ e. 361 kJ
a. 6.94 kJ
How much heat is released when 6.38 grams of Ag(s) reacts by the equation shown below at standard state conditions? 4Ag(s) + 2H2S(g) + O2(g) → 2Ag2S(s) + 2H2O Substance (kJ/mol) Ag(s) 0 H2S(g) −20.6 O2(g) 0 Ag2S(s) −32.6 H2O −285.8 a. 8.80 kJ b. 69.9 kJ c. 22.1 kJ d. 90.8 kJ e. 40.5 kJ
a. 8.80 kJ
How much heat energy is liberated when 11.0 grams of manganese is converted to Mn2O3 at standard state conditions? is −962.3 kJ/mol. a. 96.2 kJ b. 192 kJ c. 289 kJ d. 460 kJ e. 964 kJ
a. 96.2 kJ
Which of the following changes represent a decrease in entropy? a. Condensation of steam on glass b. Evaporation of gasoline c. Decomposition of fallen leaves d. Melting snow e. Diffusion of perfume throughout a room
a. Condensation of steam on glass
Which chemical change listed below represents a decrease in entropy? a. N2(g) + 3H2(g) → NH3(g) b. CaCO3(s) → CaO(s) + CO2(g) c. 2NO2(g) → N2(g) + 2O2(g) d. 2C6H6 + 15O2(g) → 12CO2(g) + 6H2O(g) e. 2NaCl → 2Na + Cl2(g)
a. N2(g) + 3H2(g) → NH3(g)
A positive change in entropy represents: a. an increase in dispersal of matter (molecular disorder) b. release of thermal energy c. a decrease in thermal energy d. a process that is always spontaneous e. a process that cannot occur spontaneously
a. an increase in dispersal of matter (molecular disorder)
A 100 g sample of each of the following metals is heated from 35°C to 45°C. Which metal absorbs the greatest amount of heat energy? Metal Specific Heat copper 0.385 J/(g · °C) magnesium 1.02 J/(g · °C) mercury 0.138 J/(g · °C) silver 0.237 J/(g · °C) lead 0.129 J/(g · °C) a. magnesium b. lead c. copper d. mercury e. silver
a. magnesium
For the reaction given below, ΔH0 = −1516 kJ at 25°C and ΔS0 = −432.8 J/K at 25°C. This reaction is spontaneous ____. SiH4(g) + 2O2(g) → SiO2(s) + 2H2O a. only below a certain temperature b. only above a certain temperature c. at all temperatures d. at no temperatures e. cannot tell from the information available
a. only below a certain temperature
Which one of the following thermodynamic quantities is not a state function? a. work b. enthalpy c. entropy d. internal energy e. free energy
a. work
Calculate at 25°C for CO(g), given that ΔH0 at 25°C for the reaction below is −809.9 kJ. 2CH4(g) + O2(g) + 4Cl2(g) → 8HCl(g) + 2CO(g) (kJ/mol) −74.81 0 0 −92.31 ? a. −110.5 kJ/mol b. −177.5 kJ/mol c. −160.0 kJ/mol d. −437.7 kJ/mol e. −486.6 kJ/mol
a. −110.5 kJ/mol
The burning of 80.3 g of SiH4 at constant pressure gives off 3790 kJ of heat. Calculate ΔH for this reaction. SiH4(g) + 2O2(g) → SiO2(s) + 2H2O a. −1520 kJ/mol rxn b. −47.2 kJ/mol rxn c. −4340 kJ/mol rxn d. −2430 kJ/mol rxn e. +4340 kJ/mol rxn
a. −1520 kJ/mol rxn
What is the entropy change of the reaction below at 298 K and 1 atm pressure? N2(g) + 3H2(g) → 2NH3(g) (J/mol•K) 191.5 130.6 192.3 a. −198.7 J/K b. 76.32 J/K c. −129.7 J/K d. 303.2 J/K e. 384.7 J/K
a. −198.7 J/K
Estimate the enthalpy change for the reaction below from the average bond energies given. There are two Cl−Cl and two C−H bonds in CH2Cl2. Remember that energy is absorbed when bonds are broken and released when they are formed. CH4(g) + 2Cl2(g) → CH2Cl2(g) + 2HCl(g) Average Bond Energies C−H 413 kJ/mol Cl−Cl 242 kJ/mol H−Cl 432 kJ/mol C−Cl 339 kJ/mol a. −232 kJ/mol b. +578 kJ/mol c. +232 kJ/mol d. −578 kJ/mol e. +541 kJ/mol
a. −232 kJ/mol
Calculate ΔG0 for the reaction below. The standard molar entropy change for the reaction at 298 K is −287.5 J/mol•K. 3NO2(g) + H2O → 2HNO3(aq) + NO(g) + 136.8 kJ a. −51.2 kJ/mol b. 85,500 kJ/mol c. −68.4 kJ/mol d. −236 kJ/mol e. −222 kJ/mol
a. −51.2 kJ/mol
Estimate the heat of reaction at 298 K for the reaction shown, given the average bond energies below. Br2(g) + 3F2(g) → 2BrF3(g) Bond Bond Energy Br−Br 193 kJ/mol F−F 155 kJ/mol Br−F 249 kJ/mol a. −836 kJ b. −150 kJ c. −89 kJ d. −665 kJ e. −1222 kJ
a. −836 kJ
Evaluate for the following reaction at 25°C. 2N2(g) + 3O2(g) → 2N2O3(g) S0 N2(g) 191.5 J/mol•K O2(g) 205.0 J/mol•K N2O3(g) 83.72 kJ/mol 312.2 J/mol•K a. +540.0 kJ b. +278.8 kJ c. −540.0 kJ d. −56.1 kJ e. +56.1 kJ
b. +278.8 kJ
Use average bond enthalpies (given in the table below) to calculate the enthalpy change for the following gas-phase reaction: CH3OH(g) + HI(g) → CH3I(g) + H2O(g) Bond kJ/mol Bond kJ/mol C-H 413 C-O 351 C-C 348 C=O 728 C=C 615 O-H 463 C-I 213 H-I 299 a. 26 kJ b. -26 kJ c. 200 kJ d. 601 kJ e. -601 kJ
b. -26 kJ
A scientist measures the standard enthalpy change for the following reaction to be -389.6 kJ: P4O10(s) + 6 H2O(l) → 4 H3PO4(aq) Based on this value and the standard enthalpies of formation for the other substances, the standard enthalpy of formation of H2O(l) would be: Compound ΔfHo (kJ/mol) P4O10(s) -2984.0 H3PO4(aq) -1288.0 a. -129.1 kJ/mol b. -296.4 kJ/mol c. -290.4 kJ/mol d. -285.8 kJ/mol e. -241.8 kJ/mol
b. -296.4 kJ/mol
In the laboratory a student finds that it takes 49.3 Joules to increase the temperature of 10.2 grams of solid nickel from 24.7 to 36.3 degrees Celsius. The specific heat of nickel she has measured is: a. 0.0178 J/g.oC b. 0.417 J/g.oC c. 0.240 J/g.oC d. 0.561 J/g.oC e. 0.603 J/g.oC
b. 0.417 J/g.oC
A 51.6-mL dilute solution of acid at 23.85°C is mixed with 48.5 mL of a dilute solution of base, also at 23.85°C, in a coffee-cup calorimeter. After the reaction occurs, the temperature of the resulting mixture is 27.25°C. The density of the final solution is 1.03 g/mL. Calculate the amount of heat evolved. Assume the specific heat of the solution is 4.184 J/g•°C. The heat capacity of the calorimeter is 23.9 J/°C. a. 3.05 kJ b. 1.55 kJ c. 5.49 kJ d. 0.837 kJ e. 14.6 kJ
b. 1.55 kJ
The following reaction has ΔG = 0 at 9900°C. What is the value of ΔS0 for the reaction? At 298 K: CH4(g) + N2(g) + 163.8 kJ → HCN(g) + NH3(g) a. 16.5 J/mol•K b. 16.1 J/mol•K c. 62.1 J/mol•K d. 60.4 J/mol•K e. −16.1 J/mol•K
b. 16.1 J/mol•K
How much heat is evolved in the formation of 35.0 grams of Fe2O3(s) at 25°C and 1.00 atm pressure by the following reaction? 4Fe(s) + 3O2(g) → 2Fe2O3(s) (kJ/mol) 0 0 −824.2 a. 90.4 kJ b. 180.7 kJ c. 151 kJ d. 360.1 kJ e. 243. 9 kJ
b. 180.7 kJ
Based on the following data, what is the Br-Br bond energy? H2(g) + Br2(g) → HBr(g); ΔH = -36.44 kJ/mol rxn Bond Bond Energy (kJ/mol) H-H 435 H-Br 362 a. 399 kJ/mol b. 216 kJ/mol c. -216 kJ/mol d. -289 kJ/mol e. 289 kJ/mol
b. 216 kJ/mol
Assuming the gases are ideal, calculate the amount of work done, in joules, for the conversion of 2.00 mole of NO2 to N2O4 at 125°C in the reaction below. The value of R is 8.314 J/mol•K. 2NO2(g) → N2O4(g) a. 1,040 J b. 3,300 J c. −3,300 J d. 6,600 J e. −1,040 J
b. 3,300 J
A sample of solid tin is heated with an electrical coil. If 39.6 J of energy are added to a 14.3g sample initially at 24.2oC, what is the final temperature of the tin? cSn = 0.21J/g.oC a. 35.2oC b. 37.4oC c. 43.1oC d. 67.4oC e. 94.7oC
b. 37.4 C
Calculate ΔE0 for the reaction below at 25°C. SiO2(s) + 4HF(aq) → SiF4(g) + 2H2O (kJ/mol) −910.9 −320.8 −1615 −285.8 a. 7.5 kJ/mol b. 5.02 kJ/mol c. −5.02 kJ/mol d. 5.23 kJ/mol e. 12.5 kJ/mol
b. 5.02 kJ/mol
If 4.168 kJ of heat is added to a calorimeter containing 75.40 g of water, the temperature of the water and the calorimeter increases from 24.58°C to 35.82°C. Calculate the heat capacity of the calorimeter (in J/°C). The specific heat of water is 4.184 J/g•°C. a. 622 J/°C b. 55.34 J/°C c. 315.5 J/°C d. 25.31 J/°C e. 17.36 J/°C
b. 55.34 J/°C
Assuming the gases are ideal, calculate the amount of work done, in joules, for the conversion of 1.00 mole of Ni to Ni(CO)4 at 75°C in the reaction below. The value of R is 8.314 J/mol•K. Ni(s) + 4CO(g) → Ni(CO)4(g) a. 1.80 × 103 J b. 8.68 × 103 J c. −1.80 × 103 J d. −8.68 × 103 J e. −494 J
b. 8.68 × 103 J
Which of the following statements is incorrect? a. The thermochemical standard state of a substance is its most stable state under one atmosphere pressure and at some specific temperature (298 K if not specified). b. A superscript zero, such as ΔH0, indicates a specified temperature of 0°C. c. For a pure substance in the liquid or solid phase, the standard state is the pure liquid or solid. d. For a pure gas, the standard state is the gas at a pressure of one atmosphere. e. For a substance in solution, the standard state refers to one-molar concentration.
b. A superscript zero, such as ΔH0, indicates a specified temperature of 0°C.
Which statement about the dispersal of matter and energy is false? a. Dispersal of energy in a system results in energy being spread over many particles. b. Dispersal of matter results in a more ordered system. c. The greater the number of molecules and the higher the total energy of the system are, the less likely the energy will be concentrated in a few molecules. d. Matter being dispersed is statistically a more probable outcome. e. There are two ways that the final state can be more probable than its initial state.
b. Dispersal of matter results in a more ordered system.
Which of the following statements about the first law of thermodynamics and energy is false? a. Kinetic energy can be converted to potential energy. b. Kinetic energy = 1/2 mv. c. All the energy in the universe is conserved. d. A system can never decrease its energy. e. Potential energy is the energy of position or composition.
b. Kinetic energy = 1/2 mv.
Which of the following reactions is spontaneous at relatively low temperatures? a. NH4Br(s) + 188 kJ → NH3(g) + Br2 b. NH3(g) + HCl(g) → NH4Cl(s) + 176 kJ c. 2H2O2 → 2H2O + O2(g) + 196 kJ d. both (a) and (b) e. both (a) and (c)
b. NH3(g) + HCl(g) → NH4Cl(s) + 176 kJ
Which of the following statements regarding the third law of thermodynamics is incorrect? a. The absolute S is zero at 0 Kelvin. b. The absolute S at 298 K can be positive or negative. c. Pure substances have positive absolute S at T > 0 Kelvin. d. Absolute zero gives a reference point for determining absolute S. e. The absolute S is greater at 300 K than 100 K for a given substance.
b. The absolute S at 298 K can be positive or negative.
Consider the following equation carefully, and determine the sign of ΔS0 for the reaction it describes. NH4Br(s) → NH3(g) + HBr(g) = +188.3 kJ Which response describes the thermodynamic spontaneity of the reaction? a. The reaction is spontaneous at all temperatures. b. The reaction is spontaneous only at relatively high temperatures. c. The reaction is spontaneous only at relatively low temperatures. d. The reaction is not spontaneous at any temperatures. e. We cannot tell from information given.
b. The reaction is spontaneous only at relatively high temperatures.
Which of the following techniques cannot be used to calculate ΔHrxn? a. Calorimetry b. Using melting points of reactants and products c. Hess's Law d. Using of Heats of Formation of reactants and products e. Using bond energies of reactants and products
b. Using melting points of reactants and products
If q = 84 kJ for a certain process, that process a. requires a catalyst. b. is endothermic. c. occurs slowly. d. is exothermic. e. cannot occur.
b. is endothermic.
Which one of the following statements is false? For a reaction carried out at constant temperature and constant pressure in an open container, ____. a. the work done by the system can be set equal to −PΔV b. the work done by the system can be set equal to VΔP c. the work done by the system can be set equal to −ΔnRT where Δn is the number of moles of gaseous products minus the number of moles of gaseous reactants d. the heat absorbed by the system can be called qp e. the heat absorbed by the system can be called ΔH
b. the work done by the system can be set equal to VΔP
The following reaction is spontaneous only below 3000 K. What conclusions can be drawn regarding this reaction? A + B → AB a. ΔH is negative and ΔS is positive. b. ΔH is negative and ΔS is negative. c. ΔH is positive and ΔS is positive. d. ΔH is positive and ΔS is negative. e. Cannot conclude anything.
b. ΔH is negative and ΔS is negative.
For a certain process at 27°C, ΔG = +210.6 kJ and ΔH = −168.2 kJ. What is the entropy change for this process at this temperature? Express your answer in the form, ΔS = ____ J/K. a. 1.26 × 103 J/K b. −1.26 × 103 J/K c. −141.3 J/K d. +141.3 J/K e. +628.3 J/K
b. −1.26 × 103 J/K
From the following data at 25°C, H2(g) + Cl2(g) → 2HCl(g) ΔH0 = −185 kJ 2H2(g) + O2(g) → 2H2O(g) ΔH0 = −483.7 kJ Calculate ΔH0 at 25°C for the reaction below. 4HCl(g) + O2(g) → 2Cl2(g) + 2H2O(g) a. +299 kJ b. −114 kJ c. −299 kJ d. +114 kJ e. −86.8 kJ
b. −114 kJ
Given that ΔH0 for the oxidation of sucrose, C12H22O11(s), is −5648 kJ per mole of sucrose at 25°C, evaluate for sucrose. C12H22O11(s) + 12O2(g) → 12CO2(g) + 11H2O (kJ/mol) ? 0 −393.5 −285.8 a. −1676 kJ/mol b. −2218 kJ/mol c. −1431 kJ/mol d. −1067 kJ/mol e. −2640 kJ/mol
b. −2218 kJ/mol
The reaction of 1.00 mole of H2(g) with 0.500 mole of O2(g) to produce 1.00 mole of steam, H2O(g), at 100°C and 1.00 atm pressure evolves 242 kJ of heat. Calculate ΔE per mole of H2O(g) produced. The universal gas constant is 8.314 J/mol•K. a. +240 kJ b. −240 kJ c. +242 kJ d. −242 kJ e. −238 kJ
b. −240 kJ
Calculate ΔS0 for the reaction below at 25°C. S0 for SiH4 = 204.5 J/mol•K, for O2(g) = 205.0 J/mol•K, for SiO2(s) = 41.84 J/mol•K, for H2O = 69.91 J/mol•K. SiH4(g) + 2O2(g) → SiO2(s) + 2H2O a. −353.5 J/K b. −432.8 J/K c. 595.0 J/K d. −677.0 J/K e. −880.3 J/K
b. −432.8 J/K
The reaction of 5.5 grams of HCl with excess Ba(OH)2 releases 8300 J of heat. What is the molar heat of neutralization, ΔH, for the reaction? a. 55 kJ/mol b. −55 kJ/mol c. −110 kJ/mol d. −27.5 kJ/mol e. 1500 J/mol
b. −55 kJ/mol
Calculate ΔH0 at 25°C for the reaction below. 2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g) (kJ/mol) −205.6 0 −348.3 −296.8 a. −257.1 kJ b. −879.0 kJ c. +257.1 kJ d. −582.2 kJ e. +879.0 kJ
b. −879.0 kJ
Evaluate ΔH0 for the following reaction from the given bond energies. 2HBr(g) → H2(g) + Br2(g) ΔHH−H = 436 kJ/mol, ΔHBr−Br = 193 kJ/mol, ΔHH−Br = 366 kJ/mol a. −103 kJ b. −143 kJ c. +103 kJ d. +142 kJ e. 259 kJ
c. +103 kJ
What is the enthalpy change for the following reaction when 3.30 moles of oxygen react with excess methane? CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔH = -802 kJ a. -486.1 kJ b. -1203.0 kJ c. -1323.3 kJ d. -1604.0 kJ e. -2646.6 kJ
c. -1323.3 kJ
What is the change in enthalpy when 9.00 mol of sulfur trioxide decomposes to sulfur dioxide and oxygen gas? 2SO2(g) + O2(g) → 2SO3(g); ΔH° = 198 kJ/mol rxn a. 891 kJ b. -198 kJ c. -891 kJ d. 198 kJ e. 1782 kJ
c. -891 kJ
A 120.0-g sample of metal at 89.0°C is added to 120.0 g of H2O(l) at 17.0°C in an insulated container. The temperature rises to 23.2°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. 44.7 J/(g · °C) c. 0.391 J/(g · °C) d. -0.391 J/(g · °C) e. 10.7 J/(g · °C)
c. 0.391 J/(g · °C)
The heat of reaction of one of the following reactions is the average bond energy for the N-H bond in NH3. Which one? a. 2NH3(g) → N2(g) + 3H2(g) b. NH3(g) → 1/2N2(g) + 3/2H2(g) c. 1/3NH3(g) → 1/3N(g) + H(g) d. 2/3NH3(g) → 1/3N2(g) + H2(g) e. 1/3N(g) + H(g) → 1/3NH3(g)
c. 1/3NH3(g) → 1/3N(g) + H(g)
Calculate ΔH0 for the following reaction at 25.0°C. Fe3O4(s) + CO(g) → 3FeO(s) + CO2(g) (kJ/mol) −1118 −110.5 −272 −393.5 a. −263 kJ b. 54 kJ c. 19 kJ d. −50 kJ e. 109 kJ
c. 19 kJ
If the entropy change for the reaction below at 298 K and 1 atm pressure is 137 J/K and S0 = 205 J/mol•K for O2(g), what is S0 for O3(g)? 2O3(g) → 3O2(g) a. 364 J/mol•K b. 478 J/mol•K c. 239 J/mol•K d. −117 J/mol•K e. −59 J/mol•K
c. 239 J/mol•K
For a particular reaction at 25°C, ΔH0 = −297 kJ/mol, and ΔS0 = −113.3 J/mol•K. At which of the following temperatures would the reaction become spontaneous? a. 2750 K b. 3250 K c. 2450 K d. 10500 K e. 3750 K
c. 2450 K
A system is compressed from 50.0 L to 5.0 L at a constant pressure of 10.0 atm. What is the amount of work done? a. 2.5 × 105 J b. 450 J c. 4.6 × 104 J d. −450 J e. −4.6 × 104 J
c. 4.6 × 104 J
At 25°C ΔH = 128.9 kJ and ΔG = 33.5 kJ for a reaction. Above what minimum temperature will this reaction become spontaneous? a. 298 K b. 332 K c. 403 K d. 530 K e. 1150 K
c. 403 K
Calculate the average bond energy in kJ per mol of bonds for the C−H bond from the following data: C(graphite) + 2H2(g) → CH4(g) = −74.81 kJ for H(g) = 218.0 kJ for C(g) = 716.7 kJ a. 590.4 kJ/mol b. 1011 kJ/mol c. 415.9 kJ/mol d. 1665 kJ/mol e. 1229 kJ/mol
c. 415.9 kJ/mol
A 50.0 mL solution of 1.2 M HCl at 24.1°C is mixed with 50.0 mL of 1.3 M NaOH, also at 24.1°C, in a coffee-cup calorimeter. After the reaction occurs, the temperature of the resulting mixture is 29.8°C. The density of the final solution is 1.05 g/mL. Calculate the molar heat of neutralization. Assume the specific heat of the solution is 4.184 J/g•°C. The heat capacity of the calorimeter is 32.5 J/°C. a. 41.7 kJ/mol b. 58.5 kJ/mol c. 44.8 kJ/mol d. 13.0 kJ/mol e. 33.9 kJ/mol
c. 44.8 kJ/mol
Estimate the temperature at which ΔG = 0 for the following reaction. NH3(g) + HCl(g) → NH4Cl(s) ΔH = −176 kJ; ΔS = −284.5 J/K a. 467 K b. 582 K c. 619 K d. 634 K e. 680 K
c. 619 K
How much heat would be released if 12.0 g of methane, CH4, was completely burned in oxygen to form carbon dioxide and water at standard state conditions? for CH4(g) = −74.81 kJ/mol, for CO2(g) = −393.5 kJ/mol and for H2O = −285.8 kJ/mol a. 77.5 kJ b. 453 kJ c. 668 kJ d. 190. kJ e. 890. kJ
c. 668 kJ
How much energy is required to raise the temperature of 10.9g of water from 22.9oC to 38.2oC? a. 38.5 J b. 298 J c. 698 J d. 1040 J e. 1740 J
c. 698 J
The following information is given for water, H2O, at 1 atm: boiling point = 100oC ΔHvap = 40.7 kJ/mol melting point = 0.000oC ΔHfus = 6.01 kJ/mol specific heat liquid = 4.18 J/goC At a pressure of 1 atm, how many kJ of heat are needed to vaporize a 39.6g sample of liquid water at its normal boiling point of 100oC? a. 13.2 kJ b. 16.6 kJ c. 89.4 kJ d. 238 kJ e. 1610 kJ
c. 89.4 kJ
Estimate the temperature above which this reaction is spontaneous. ΔS0 = 16.1 J/K. CH4(g) + N2(g) + 163.8 kJ → HCN(g) + NH3(g) a. 9.91°C b. 1045 K c. 9.90 × 103°C d. 10.7 K e. 10.1°C
c. 9.90 × 103°C
Which one of the following statements is false? a. The change in internal energy, ΔE, for a process is equal to the amount of heat absorbed at constant volume, qv. b. The change in enthalpy, ΔH, for a process is equal to the amount of heat absorbed at constant pressure, qp. c. A bomb calorimeter measures ΔH directly. d. If qp for a process is negative, the process must be exothermic. e. The work done in a process occurring at constant pressure is zero if Δngases is zero.
c. A bomb calorimeter measures ΔH directly.
Which of the following substances is not correctly matched with its molar heat of formation, ? a. C6H6 / = 49.03 kJ/mol b. SO2(g) / = −296.8 kJ/mol c. Br2(g) / = 0 d. H2O / = −285.8 kJ/mol e. Ca(s) / = 0
c. Br2(g) / = 0
Which of the following reactions corresponds to the thermochemical equation for the standard enthalpy of formation of solid calcium nitrate? a. Ca2+(aq) + 2NO3-(aq) → Ca(NO3)2(s) b. Ca(OH)2(s) + 2HNO3(aq) → Ca(NO3)2(s) + 2H2O(l) c. Ca(s) + N2(g) + 3O2(g) → Ca(NO3)2(s) d. Ca(s) + 2HNO3(aq) → Ca(NO3)2(s) + H2(g) e. Ca(s) + 2N(g) + 6O(g) → Ca(NO3)2(s)
c. Ca(s) + N2(g) + 3O2(g) → Ca(NO3)2(s)
Which one of the following statements is false? a. The amount of heat absorbed by a system at constant volume, qv, is ΔE for the process. b. The amount of heat absorbed by a system at constant pressure, qp, is ΔH for the process. c. In the relationship ΔE = q + w, as applied to a typical chemical reaction, w is usually much larger than q. d. At constant temperature and pressure, the work done by a system involving gases is −Δngases(RT) where Δngas = nproduct gases − nreactant gases for the process of interest. e. At constant pressure, the work done in a process by a system involving gases can be expressed as −PΔV.
c. In the relationship ΔE = q + w, as applied to a typical chemical reaction, w is usually much larger than q.
For which of the following substances does DeltaHf^0=0 ? a. CO2(g) b. H2O(g) c. Na(s) d. Br2(g) e. C(diamond)
c. Na(s)
Consider the following reaction at constant pressure. Which response is true? N2(g) + O2(g) → 2NO(g) a. Work is done on the system as it occurs. b. Work is done by the system as it occurs. c. No work is done as the reaction occurs. d. Work may be done on or by the system as the reaction occurs, depending upon the temperature. e. The amount of work depends on the pressure.
c. No work is done as the reaction occurs.
Which of the following substances is not in its standard state? a. C, graphite b. Br2, c. O3, (g) d. H2, (g) e. Hg,
c. O3, (g)
Consider the reaction below at 25°C for which ΔS0 = 16.1 J/K. CH4(g) + N2(g) + 163.8 kJ → HCN(g) + NH3(g) Which one of the following statements describes the reaction? a. Spontaneous at all temperatures b. Spontaneous at relatively low temperatures only c. Spontaneous at relatively high temperatures only d. Nonspontaneous at all temperatures e. Insufficient information to estimate temperature range of spontaneity
c. Spontaneous at relatively high temperatures only
The enthalpy change, ΔH, of a process is defined as: a. The maximum amount of useful work that can be done in a system. b. The increase or decrease in temperature in a system. c. The quantity of heat transferred in or out of a system as it undergoes a change at constant pressure. d. The change in molecular disorder in a system. e. None of these are correct.
c. The quantity of heat transferred in or out of a system as it undergoes a change at constant pressure.
Which situation does not represent a dispersal of matter? a. A gas expands into a vacuum. b. A sugar cube dissolves in a cup of coffee. c. Water freezes to form ice. d. Soap and water are mixed in a sink. e. Dye is dropped into a glass of water.
c. Water freezes to form ice.
For which set of values of ΔH and ΔS will a reaction be spontaneous (product-favored) at all temperatures? a. ΔH = +10 kJ, ΔS = −5 J/K b. ΔH = −10 kJ, ΔS = −5 J/K c. ΔH = −10 kJ, ΔS = +5 J/K d. ΔH = +10 kJ, ΔS = +5 J/K e. no such values exist
c. ΔH = −10 kJ, ΔS = +5 J/K
Joseph Priestley prepared oxygen by heating mercury(II) oxide. The compound HgO is stable at room temperature but decomposes into its elements (Hg and O2) at high temperatures. What conclusions can be drawn concerning ΔH and ΔS for this decomposition reaction? a. ΔH is negative and ΔS is positive. b. ΔH is negative and ΔS is negative. c. ΔH is positive and ΔS is positive. d. ΔH is positive and ΔS is negative. e. ΔH becomes negative at high temperatures.
c. ΔH is positive and ΔS is positive.
Calculate ΔS0 at 25°C for the reaction below. PbS(s) + 2HCl(g) → PbCl2(s) + H2S(g) (kJ/mol) −100.4 −92.31 −359.4 −20.6 (kJ/mol) −98.7 −95.30 −314.1 −33.6 a. 686 J/K b. −741 J/K c. −123 J/K d. 1.33 × 103 J/K e. 515 J/K
c. −123 J/K
Calculate ΔG0 at 298 K for the reaction below. N2O4(g) + 2N2H4 → 3N2(g) + 4H2O(g) (kJ/mol) 97.82 149.0 0 −228.6 a. −518.1 kJ/mol b. −475.6 kJ/mol c. −1311 kJ/mol d. 1311 kJ/mol e. −667.1 kJ/mol
c. −1311 kJ/mol
Evaluate ΔS0 for the reaction below at 25°C. CH4(g) + 2Cl2(g) → CCl4 + 2H2(g) (kJ/mol) −74.81 0 −135.4 0 (kJ/mol) −50.75 0 −65.27 0 a. −360 J/K b. −66.9 J/K c. −155 J/K d. −487 J/K e. −387 J/K
c. −155 J/K
Determine ΔG0 for the reaction: 2CO(g) + 2H2(g) → CO2(g) + CH4(g) ΔH0 = −247.3 kJ/mol ΔS0 = −256.5 J/mol K a. −182.8 kJ/mol b. −253.7 kJ/mol c. −170.9 kJ/mol d. 76,200 J/mol e. 323.9 kJ/mol
c. −170.9 kJ/mol
For a certain process at 127°C, ΔG = −16.20 kJ and ΔH = −17.0 kJ. What is the entropy change for this process at this temperature? Express your answer in the form, ΔS = ____ J/K. a. −6.3 J/K b. +6.3 J/K c. −2.0 J/K d. +2.0 J/K e. −8.1 J/K
c. −2.0 J/K
Evaluate ΔS0 for the reaction below at 25°C and 1 atm. 3NO2(g) + H2O → 2HNO3(aq) + NO(g) S0 (J/mol•K) 240 69.91 146 210.7 a. +1.37 × 103 J/K b. +287.2 J/K c. −287.2 J/K d. +1.37 × 103 J/K e. −531.4 J/K
c. −287.2 J/K
For the following reaction at 25°C, ΔH0 = −26.88 kJ and ΔS0 = 11.2 J/K. Calculate ΔG0 for the reaction at 25°C in kilojoules. I2(g) + Cl2(g) → 2ICl(g) a. −102 kJ b. +50.6 kJ c. −30.2 kJ d. −50.6 kJ e. +77.0 kJ
c. −30.2 kJ
Calculate the at 298 K for PbCl2(s) from the following information. ΔG0 for the reaction below is −58.4 kJ at 298 K. PbS(s) + 2HCl(g) → PbCl2(s) + H2S(g) (kJ/mol) −98.7 −95.30 ? −33.6 a. −16.0 kJ/mol b. −47.6 kJ/mol c. −314.1 kJ/mol d. −36.2 kJ/mol e. −52.3 kJ/mol
c. −314.1 kJ/mol
Calculate the average N−H bond energy in NH3(g). for NH3(g) = −46.11 kJ/mol, for N(g) = 472.7 kJ/mol, for H(g) = 218.0 kJ/mol, for N2(g) = 0 kJ/mol, for H2(g) = 0 kJ/mol. a. −46.11 kJ b. −15.4 kJ c. −390.9 kJ d. 264.5 kJ e. 1173 kJ
c. −390.9 kJ
Calculate the standard Gibbs free energy change for the following reaction at 25°C. CaCO3(s) + 2HCl(g) → CaCl2(s) + CO2(g) + H2O (kJ/mol) −1129 −95.3 −750.2 −394.4 −237.2 a. −41 kJ b. −158 kJ c. −62 kJ d. −87 kJ e. −104 kJ
c. −62 kJ
Which of following would have the highest value of absolute entropy per mole? a. water at 50°C b. water at 10°C c. ice at −10°C d. 1 M NaCl at 50°C e. 1 M NaCl at 10°C
d. 1 M NaCl at 50°C
The heat of vaporization of methanol, CH3OH, is 35.20 kJ/mol. Its boiling point is 64.6°C. What is the change in entropy for the vaporization of methanol? a. −17.0 J/mol•K b. 3.25 J/mol•K c. 17.0 J/mol•K d. 104 J/mol•K e. 543 J/mol•K
d. 104 J/mol•K
How much heat is gained by cobalt when 23.4 g of cobalt is warmed from 26.2°C to 63.2°C? The specific heat of cobalt is 0.421 J/(g · °C). a. 2.58 × 102 J b. 26.61 J c. 15.58 J d. 3.65 × 102 J e. 6.23 × 102 J
d. 3.65 × 102 J
Calculate the average S−F bond energy in SF6. for SF6(g) = −1209 kJ/mol, for S(g) = 278.8 kJ/mol, and for F(g) = 78.99 kJ/mol. a. 1962 kJ b. 1209 kJ c. 200.8 kJ d. 327.0 kJ e. 1565 kJ
d. 327.0 kJ
How much heat is released when 75 g of octane is burned completely if the enthalpy of combustion is −5,500 kJ/mol C8H18? C8H18 + 25/2 O2 → 8CO2 + 9H2O a. 7200 kJ b. 8360 kJ c. 4.1 × 105 kJ d. 3600 kJ e. 5500 kJ
d. 3600 kJ
Estimate the boiling point of water if for H2O , = −285.8 kJ/mol and S0 = 69.91 J/mol•K and for H2O(g), = −241.8 kJ/mol and S0 = 188.7 J/mol•K. a. 101 K b. 387 K c. 398 K d. 370 K e. 274 K
d. 370 K
Evaluate ΔH0 for the reaction below at 25°C. SiO2(s) + 4HF(aq) → SiF4(g) + 2H2O (kJ/mol) −910.9 −320.8 −1615 −285.8 a. +293.3 kJ b. −954.9 kJ c. −366.5 kJ d. 7.5 kJ e. −1781.1 kJ
d. 7.5 kJ
The specific heat capacity of graphite, C(s), is 0.71J/g.oC. Calculate the molar heat capacity of graphite. a. 0.059 J/mol.oC b. 0.12 J/moloC c. 1.4 J/moloC d. 8.5 J/moloC e. 17 J/moloC
d. 8.5 J/moloC
The for gaseous acetylene, H−C≡C−H, is 227 kJ/mol. What is the C≡C bond energy? The bond energies are 423 kJ/mol for C−H and 436 kJ/mol for H−H. The heat of sublimation for carbon is 717 kJ/mol. a. 98 kJ/mol b. 348 kJ/mol c. 986 kJ/mol d. 817 kJ/mol e. 1251 kJ/mol
d. 817 kJ/mol
Priestley prepared oxygen by heating mercury(II) oxide. From the data given below estimate the temperature above which this reaction will become spontaneous. HgO(s) → Hg + O2(g) ΔH0 = 90.83 kJ S0(Hg) = 76.02 J/mol•K S0(HgO) = 70.29 J/mol•K S0(O2) = 205.0 J/mol•K a. 108 K b. 566 K c. 430 K d. 840 K e. 739 K
d. 840 K
1. Which statement is incorrect? a. Energy is the capacity to do work or to transfer heat. b. Kinetic energy is the energy of motion. c. Potential energy is the energy that a system possesses by virtue of its position or composition. d. A process that absorbs energy from its surroundings is called exothermic. e. The Law of Conservation of Energy is another statement of the First Law of Thermodynamics.
d. A process that absorbs energy from its surroundings is called exothermic.
Which of the following statements is a correct interpretation of the First Law of Thermodynamics? a. The combined amount of matter and energy in the universe is a constant. b. Energy is neither created nor destroyed in chemical reactions. c. Energy is neither created nor destroyed in physical changes. d. All of these are correct. e. None of these are correct.
d. All of these are correct.
A process occurs spontaneously and ΔSsystem < 0. Which statement below must be true? a. ΔSsurroundings > 0 b. ΔSuniverse > 0 c. The pressure is constant. d. Both (a) and (b) are correct. e. All of these answers are correct.
d. Both (a) and (b) are correct.
Which statement concerning sign conventions for ΔE = q + w is false? a. For heat absorbed by the system, q is positive. b. For work done by the system, w is negative. c. When energy is released by the reacting system, ΔE is negative. d. If ΔE is positive, energy can be written as a product in the equation for the reaction. e. For an expansion, w is negative.
d. If ΔE is positive, energy can be written as a product in the equation for the reaction.
Which statement regarding enthalpy change is incorrect? a. ΔH = Hfinal − Hinitial b. Enthalpy change is a state function. c. ΔH = qp d. The absolute enthalpy of a system can be experimentally measured. e. ΔHreaction = Hproducts − Hreactants
d. The absolute enthalpy of a system can be experimentally measured.
Which statement below is false? a. For reactions that release heat to the surroundings, ΔH is negative. b. For reactions in which the reacting system becomes more disordered, ΔS is positive. c. If the free energy change of reaction is positive, the reaction cannot occur to give predominantly products under the given conditions. d. The entropy of a system increases when order increases. e. Endothermic reactions may be spontaneous.
d. The entropy of a system increases when order increases.
Consider the following reaction occurring at constant pressure and temperature, for which the value of ΔE is negative. Which response is false? CH4(g) + 2O2(g) → CO2(g) + 2H2O a. Work is done by the surroundings on the system. b. Work is positive. c. Heat is released by the system. d. The volume must increase at constant pressure. e. All of these statements are true.
d. The volume must increase at constant pressure.
Which one of the following statements is not correct? a. When ΔG for a reaction is negative, the reaction is spontaneous. b. When ΔG for a reaction is positive, the reaction is nonspontaneous. c. When ΔG for a reaction is zero, the system is at equilibrium. d. When ΔH for a reaction is negative, the reaction is never spontaneous. e. When ΔH for a reaction is very positive, the reaction is not expected to be spontaneous.
d. When ΔH for a reaction is negative, the reaction is never spontaneous.
When HCl(g) reacts with NH3(g) to form NH4Cl(s) according to the following equation, energy is released into the surroundings. HCl(g) + NH3(g) → NH4Cl(s) Is this reaction endothermic or exothermic, and what is the sign of ΔH for this reaction? a. endothermic, + b. endothermic, - c. exothermic, + d. exothermic, - e. energy change is balanced by the surroundings, neutral
d. exothermic, -
A process cannot be spontaneous (product-favored) if ____. a. it is exothermic, and there is an increase in disorder b. it is endothermic, and there is an increase in disorder c. it is exothermic, and there is a decrease in disorder d. it is endothermic, and there is a decrease in disorder e. the entropy of the universe increases
d. it is endothermic, and there is a decrease in disorder
Which of the following statements about internal energy, E, is false? a. It represents all the energy within a certain amount of matter. b. In some processes ΔE = q c. In some processes ΔE = q + w d. ΔE is positive in exothermic reactions. e. Its absolute value cannot be determined.
d. ΔE is positive in exothermic reactions.
Consider the conversion of a substance from solid to liquid. Solid Liquid At one atmosphere pressure and at the melting point of the substance, ____. a. ΔH = 0 for the process b. ΔS = 0 for the process c. ΔE = 0 for the process d. ΔG = 0 for the process e. both ΔH = 0 and ΔE = 0 for the process
d. ΔG = 0 for the process
Given the enthalpy changes for the following reactions, calculate for CO(g). C (graphite) + O2(g) → CO2(g) = −393.5 kJ CO(g) + 1/2O2 → CO2(g) ΔH0 = −283.0 kJ a. 6.78 × 102 kJ b. −6.78 × 102 kJ c. +110.5 kJ d. −110.5 kJ e. −173 kJ
d. −110.5 kJ
Evaluate ΔG0 for the reaction below at 25°C. 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O (kJ/mol) 209.2 0 −394.4 −237.2 a. −1409 kJ b. −2599 kJ c. −1643 kJ d. −2470 kJ e. −766 kJ
d. −2470 kJ
For the following reaction, ΔH0 = −104.9 kJ/mol. If this reaction is spontaneous at all temperatures below 361 K, what is the value of ΔS0? a. −165 J/mol•K b. −0.291 J/mol•K c. 3.44 J/mol•K d. −291 J/mol•K e. 3.44 kJ
d. −291 J/mol•K
Consider the following reaction and its and values at 25°C. Evaluate at 25°C. 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O = −2599 kJ, = −2470 kJ a. +340 J/K b. −340 J/K c. +386 J/K d. −433 J/K e. −386 J/K
d. −433 J/K
Calculate ΔG0 at 298 K for the reaction below. Fe2O3(s) + 13CO(g) → 2Fe(CO)5(g) + 3CO2(g) (kJ/mol) −824.2 −110.5 −733.8 −393.5 S0 (J/mol•K) 87.4 197.6 445.2 213.6 a. +63.6 kJ b. +26.8 kJ c. −243.1 kJ d. −52.2 kJ e. −193.3 kJ
d. −52.2 kJ
Evaluate for the following reaction at 25°C. 2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g) (kJ/mol) −205.6 0 −348.3 −296.8 S0 (J/mol•K) 57.7 205.0 43.64 248.1 a. −951.1 kJ b. −922.6 kJ c. −704.2 kJ d. −835.2 kJ e. −1902 kJ
d. −835.2 kJ
The "roasting" of 48.7 g of ZnS at constant pressure gives off 220. kJ of heat. Calculate the ΔH for this reaction. 2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g) a. −110 kJ/mol rxn b. −293 kJ/mol rxn c. −440. kJ/mol rxn d. −881 kJ/mol rxn e. +440. kJ/mol rxn
d. −881 kJ/mol rxn
Given the following at 25°C, calculate for HCN(g) at 25°C. 2NH3(g) + 3O2(g) + 2CH4(g) → 2HCN(g) + 6H2O(g) = −870.8 kJ = −80.3 kJ/mol for NH3(g), −74.6 kJ/mol for CH4, −241.8 kJ/mol for H2O(g). a. −135 kJ/mol b. −147 kJ/mol c. +270 kJ/mol d. −870.8 kJ/mol e. +135 kJ/mol
e. +135 kJ/mol
A list of the calorie content of foods indicates that a 10 oz chocolate shake contains 353 Calories. Express this value in Joules. (1 Calorie = 1000 calories; 1 calorie = 4.18 Joules) a. 84.4 J b. 84,400 J c. 148 J d. 1480 J e. 1480,000 J
e. 1480,000
An automobile engine provides 551J of work to push the pistons and generates 2250J of heat that must be carried away by the cooling system. Calculate the change in internal energy of the engine. a. 551J b. 1102J c. 1699J d. 2250J e. 2801J
e. 2801J
Which one of the following reactions has a positive entropy change? a. H2O(g) → H2O b. BF3(g) + NH3(g) → F3BNH3(s) c. 2SO2(g) + O2(g) → 2SO3(g) d. N2(g) + 3H2(g) → 2NH3(g) e. 2NH4NO3(s) → 2N2(g) + 4H2O(g) + O2(g)
e. 2NH4NO3(s) → 2N2(g) + 4H2O(g) + O2(g)
How much heat is absorbed in the complete reaction of 3.00 grams of SiO2 with excess carbon in the reaction below? ΔH0 for the reaction is +624.6 kJ. SiO2(g) + 3C(s) → SiC(s) + 2CO(g) a. 366 kJ b. 1.13 × 105 kJ c. 5.06 kJ d. 1.33 × 104 kJ e. 31.2 kJ
e. 31.2 kJ
The following thermochemical equation is for the reaction of ammonia(g) with oxygen(g) to form nitrogen monoxide(g) and water(g): 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g) ΔH = -905 kJ How many grams of NH3(g) would have to react with excess O2(g) to produce 58.6 kJ of energy? a. 15.4 g b. 3.86 g c. 61.6 g d. 0.259 g e. 4.41 g
e. 4.41 g
Estimate the temperature above which this reaction is nonspontaneous. PbS(s) + 2HCl(g) → PbCl2(s) + H2S(g) (kJ/mol) −100.4 −92.31 −359.4 −20.6 (kJ/mol) −98.7 −95.30 −314.1 −33.6 a. −144°C b. 88°C c. 16°C d. −42°C e. 499°C
e. 499°C
Which of the following statements regarding spontaneous changes is false? a. Spontaneity is favored when heat is released. b. Spontaneity is favored when the dispersal of matter is increased. c. Spontaneous changes occur at a given state without any outside influence. d. Ice melting at 25°C is spontaneous primarily due to the increase in molecular disorder (dispersal of matter). e. All exothermic reactions are spontaneous.
e. All exothermic reactions are spontaneous.
A flask containing helium gas is released into a closed room. Which of the following ideas regarding entropy is false? a. ΔSsystem > 0 b. Matter is dispersed. c. ΔSuniverse > 0 d. This process is spontaneous. e. All of these statements are true.
e. All of these statements are true.
Which of the following is not a formation reaction? a. 1/2H2(g) + 1/2Br2 → HBr(g) b. H2(g) + 1/2O2(g) → H2O c. Ca(s) + 1/2O2(g) → CaO(s) d. 4Al(s) + 3/2O2(g) → Al2O3(s) e. H2O + SO3 → H2SO4
e. H2O + SO3 → H2SO4
Which physical change listed below would have a negative value of ΔS? a. The sublimation (vaporization) of dry ice (solid CO2). b. Boiling water. c. Evaporation of water from a lake. d. Sugar dissolving in coffee. e. Rain drops forming in a cloud.
e. Rain drops forming in a cloud.
The second law of thermodynamics states: a. All exothermic processes also increase entropy. b. The enthalpy of the universe always increases in spontaneous processes. c. A spontaneous process always increases entropy. d. ΔH < 0 and ΔS > 0 for all spontaneous processes e. The entropy of the universe always increases in spontaneous processes.
e. The entropy of the universe always increases in spontaneous processes.
The energy associated with a motionless rock sitting atop a mountain is called _____. a. heat b. internal energy c. temperature d. kinetic energy e. potential energy
e. potential energy
Which term is not correctly matched? a. endothermic / energy is absorbed b. universe / system plus surroundings c. exothermic / energy is released d. thermodynamic state / conditions specifying the properties of a system e. state function / property dependent on how the process takes place
e. state function / property dependent on how the process takes place
Which statement is false? a. The thermodynamic quantity most easily measured in a "coffee cup" calorimeter is ΔH. b. No work is done in a reaction occurring in a bomb calorimeter. c. ΔH is sometimes exactly equal to ΔE. d. ΔH is often nearly equal to ΔE. e. ΔH is equal to ΔE for the process: 2H2(g) + O2(g) → 2H2O(g).
e. ΔH is equal to ΔE for the process: 2H2(g) + O2(g) → 2H2O(g).
Which of the following statements about free energy is false? a. If ΔS is negative then ΔH must be negative for a spontaneous process. b. ΔS is positive for many spontaneous processes. c. ΔG is always negative for spontaneous processes. d. ΔG is always positive for nonspontaneous processes. e. ΔS must be positive for a process to be spontaneous.
e. ΔS must be positive for a process to be spontaneous.
Calculate the standard enthalpy change for the reaction below. C(graphite) + 4HNO3 → CO2(g) + 4NO2(g) + 2H2O (kJ/mol) 0 −174.1 −393.5 33.2 −285.8 a. −123.9 kJ b. −472.1 kJ c. −201.9 kJ d. −404.8 kJ e. −135.9 kJ
e. −135.9 kJ
Calculate the standard energy change, ΔE0, for the reaction below. 12NH3(g) + 21O2(g) → 8HNO3 + 4NO(g) + 14H2O(g) (kJ/mol) −45.9 0 −133.9 91.3 −241.8 a. −3,540 kJ/mol b. −201.3 kJ/mol c. −2,259 kJ/mol d. −4270 kJ/mol e. −3,503 kJ/mol
e. −3,503 kJ/mol
Calculate the standard enthalpy change for the reaction below. 12NH3(g) + 21O2(g) → 8HNO3 + 4NO(g) + 14H2O(g) (kJ/mol) −45.9 0 −133.9 91.3 −241.8 a. +3,540 kJ b. −4,650 kJ c. −2,259 kJ d. −4270 kJ e. −3,540 kJ
e. −3,540 kJ
Evaluate ΔG0 for the reaction below at 25°C. P4O10(s) + 6H2O → 4H3PO4(s) (kJ/mol) −2984 −285.8 −1281 S0 (J/mol•K) 228.9 69.91 110.5 a. −50.33 kJ b. −172.0 kJ c. −282.5 kJ d. −304.8 kJ e. −363.7 kJ
e. −363.7 kJ
A 0.900-g sample of toluene, C7H8, was completely burned in a bomb calorimeter containing 4560. g of water which increased in temperature from 23.800°C to 25.718°C. What is ΔE for the reaction in kJ/mol C7H8? The heat capacity of the calorimeter was 780. J/°C. The specific heat of water is 4.184 J/g•°C. a. −4520 kJ/mol b. +3500 kJ/mol c. −38.1 kJ/mol d. −2220 kJ/mol e. −3900 kJ/mol
e. −3900 kJ/mol
A 1.00-g sample of hexane, C6H14, undergoes complete combustion with excess O2 in a bomb calorimeter. The temperature of the 1500. g of water surrounding the bomb rises from 22.64°C to 29.30°C. The heat capacity of the calorimeter is 4.04 kJ/°C. What is ΔE for the reaction in kJ/mol of C6H14. The specific heat of water is 4.184 J/g•°C. a. −9.96 × 103 kJ/mol b. −4.52 × 103 kJ/mol c. −1.15 × 104 kJ/mol d. −7.40 × 104 kJ/mol e. −5.91 × 103 kJ/mol
e. −5.91 × 103 kJ/mol
Given the following at 25°C and 1.00 atm: 1/2N2(g) + O2(g) → NO2(g) ΔH0 = 33.2 kJ N2(g) + 2O2(g) → N2O4(g) ΔH0 = 11.1 kJ Calculate the ΔH0 for the reaction below at 25°C. 2NO2(g) → N2O4(g) a. +11.0 kJ b. +44.3 kJ c. +55.3 kJ d. −22.1 kJ e. −55.3 kJ
e. −55.3 kJ
