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chemistry
physical chemistry

Chemical Principles 7th edition Steven S. Zumdahl, Donald J. DeCoste - Solutions

Given the following data: 2ClF(g) + O2(g) → Cl2O(g) + F2O(g) ΔH = 167.4 kJ 2ClF3(g) + 2O2(g) → Cl2O(g) + 3F2O(g) ΔH = 341.4 kJ 2F2(g) + O2(g) → 2F2O(g) ΔH = –43.4 kJ Calculate ΔH for the reaction ClF(g) + F2(g) → ClF3(g)
Given the following data: Ca(s) + 2C(graphite) → CaC2(s) ΔH = –62.8 kJ Ca(s) + ½ O2(g) → CaO(s) ΔH = –635.5 kJ CaO(s) + H2O(l) → Ca(OH)2(aq) ΔH = –653.1 kJ C2H2(g) + 5/2 O2(g) → 2CO2(g) 1 H2O(l) ΔH = –1300 kJ C(graphite) + O2(g) → CO2(g) ΔH = –393.5 kJ Calculate ΔH for
Given the following data: Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) ΔH = –23 kJ 3Fe2O3(s) + CO(g) → 2Fe3O4(s) + CO2(g) ΔH = –39 kJ Fe3O4(s) + CO(g) → 3FeO(s) + CO2(g) ΔH = 18 kJ Calculate ΔH for the reaction FeO(s) + CO(g) → Fe(s) + CO2(g)
Combustion reactions involve reacting a substance with oxygen. When compounds containing carbon and hydrogen are combusted, carbon dioxide and water are the products. Using the enthalpies of combustion for C4H4 (–2341 kJ/ mol), C4H8 (–2755 kJ/ mol), and H2 (–286 kJ/ mol), calculate ΔH for
Given the following data: 2O3(g) → 3O2(g) ΔH = –427 kJ O2(g) → 2O(g) ΔH = 495 kJ NO(g) + O3(g) → NO2(g) + O2(g) ΔH = –199 kJ Calculate ΔH for the reaction NO(g) + O(g) → NO2(g)
The bombardier beetle uses an explosive discharge as a ΔEfensive measure. The chemical reaction involved is the oxidation of hydroquinone by hydrogen peroxide to produce quinone and water: C6H4(OH)2(aq) + H2O2(aq) → C6H4O2(aq) + 2H2O(l) Calculate ΔH for this reaction from the following
Given the following data: P4(s) + 6Cl2(g) → 4PCl3(g) ΔH = –1225.6 kJ P4(s) + 5O2(g) → P4O10(s) ΔH = –2967.3 kJ PCl3(g) + Cl2(g) → PCl5(g) ΔH = –84.2 kJ PCl3(g) + ½ O2(g) → Cl3PO(g) ΔH = –285.7 kJ calculate ΔH for the reaction P4O10(s) + 6PCl5(g) → 10 Cl3PO(g)
Given the definition of the standard enthalpy of formation for a substance, write separate reactions for the formation of NaCl, H2O, C6H12O6, and PbSO4 that have ΔHo values equal to ΔHof for each compound.
A system undergoes a process consisting of the following two steps: Step 1: The system absorbs 72 J of heat while 35 J of work is done on it. Step 2: The system absorbs 35 J of heat while performing 72 J of work. Calculate ΔE for the overall process.
The combustion of methane can be represented as follows:a. Use the information given above to determine the value of Î”H for the combustion of methane to form CO2(g) and 2H2O(l). b. What is Î”Hof for an element in its standard state? Why is this? Use the figure above to
Use the values of Î”Hof in Appendix 4 to calculate Î”Ho for the following reactions.a.b. Ca3(PO4)2(s) + 3H2SO4(l) †’ 3CaSO4(s) + 2H3PO4(l) c. NH3(g) + HCl(g) †’ NH4Cl(s) d. e. SiCl4(l) + 2H2O(l) †’ SiO2(s) + 4HCl(aq) f. MgO(s) + H2O(l)
The Ostwald process for the commercial production of nitric acid from ammonia and oxygen involves the following steps: 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g) 2NO(g) + O2(g) → 2NO2(g) 3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g) a. Use the values of ΔH8f in Appendix 4 to calculate the value of ΔHo for
Calculate ΔHo for each of the following reactions using the data in Appendix 4: 4Na(s) + O2(g) → 2Na2O(s) 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) 2Na(s) + CO2(g) → Na2O(s) + CO(g) Explain why a water or carbon dioxide fire extinguisher might not be effective in putting out a sodium fire.
The reusable booster rockets of the space shuttle use a mixture of aluminum and ammonium perchlorate as fuel. A possible reaction is 3Al(s) + 3NH4ClO4(s) → Al2O3(s) + AlCl3(s) + 3NO(g) + 6H2O(g) Calculate ΔHo for this reaction.
The space shuttle orbiter uses the oxidation of methyl hydrazine by dinitrogentetroxide for propulsion. The balanced reaction is 5N2O4(l) + 4N2H3CH3(l) → 12H2O(g) + 9N2(g) + 4CO2(g) Calculate ΔHo for this reaction.
Does the reaction in Exercise 70 or that in Exercise 71 produce more energy per kilogram of reactant mixture (stoichiometric amounts)? Reaction of exercise 70 3Al(s) + 3NH4ClO4(s) → Al2O3(s) + AlCl3(s) + 3NO(g) + 6H2O(g) Reaction of exercise 71 5N2O4(l) + 4N2H3CH3(l) → 12H2O(g) + 9N2(g) +
At 298 K, the standard enthalpies of formation for C2H2(g) and C6H6(l) are 227 kJ/ mol and 49 kJ/ mol, respectively. a. Calculate ΔHo for C6H6(l) → 3C2H2(g) b. Both acetylene (C2H2) and benzene (C6H6) can be used as fuels. Which compound would liberate more energy per gram when combusted in air?
Calculate ΔHo for each of the following reactions, which occur in the atmosphere. a. C2H4(g) + O3(g) → CH3CHO(g) + O2(g) b. O3(g) + NO(g) → NO2(g) + O2(g) c. SO3(g) + H2O(l) → H2SO4(aq) d. 2NO(g) + O2(g) → 2NO2(g)
Use the reaction 2ClF3(g) + 2NH3(g) → N2(g) + 6HF(g) + Cl2(g) ΔHo = –1196 kJ to calculate ΔHof for ClF3(g).
Calculate the internal energy change for each of the following. a. One hundred (100.) joules of work are required to compress a gas. At the same time, the gas releases 23 J of heat. b. A piston is compressed from a volume of 8.30 L to 2.80 L against a constant pressure of 1.90 atm. In the process,
The standard enthalpy of combustion of ethene gas [C2H4(g)] is –1411.1 kJ/ mol at 298 K. Given the following enthalpies of formation, calculate ΔHof for C2H4(g). CO2(g) –393.5 kJ/ mol H2O(l) –285.8 kJ/ mol
The complete combustion of acetylene [C2H2(g)] produces 1300. kJ of energy per mole of acetylene consumed. How many grams of acetylene must be burned to produce enough heat to raise the temperature of 1.00 gal of water by 10.0oC if the process is 80.0% efficient? Assume the Density of water is 1.00
Assume that 4.19 × 106 kJ of energy is needed to heat a home. If this energy is derived from the combustion of methane (CH4), what volume of methane, measured at STP, must be burned? (ΔHoCombustion for CH4 = –891 kJ/ mol)
Syngas can be burned directly or converted to methanol. Calculate ΔHo for the reaction CO(g) + 2H2(g) → CH3OH(l)
Ethanol (C2H5OH) has been proposed as an alternative fuel. Calculate the standard enthalpy of combustion per gram of liquid ethanol.
Methanol (CH3OH) has also been proposed as an alter-native fuel. Calculate the standard enthalpy of combustion per gram of liquid methanol, and compare this answer to that for ethanol in Exercise 80.
Some automobiles and buses have been equipped to burn propane (C3H8) as a fuel. Compare the amount of energy that can be obtained per gram of C3H8(g) with that per gram of gasoline, assuming that gasoline is octane [C8H18(l)]. Look up the physical properties of propane. What disadvantages are there
Consider the following cyclic process carried out in two steps on a gas: Step 1: 45 J of heat is added to the gas, and 10. J of expansion work is performed. Step 2: 60. J of heat is removed from the gas as the gas is compressed back to the initial state. Calculate the work for the gas compression
Determine ΔE for the process H2O(l) → H2O(g) at 25oC and 1 atm.
The standard enthalpy of formation of H2O(l) at 298 K is –285.8 kJ/ mol. Calculate the change in internal energy for the following process at 298 K and 1 atm: H2O(l) → H2(g) + ½ O2(g) ΔEo = ?
A balloon filled with 39.1 moles of helium has a volume of 876 L at 0.08C and 1.00 atm pressure. At constant pressure, the temperature of the balloon is increased to 38.08C, causing the balloon to expand to a volume of 998 L. Calculate q, w, and ΔE for the helium in the balloon. (The molar heat
A piece of chocolate cake contains about 400 Calories. A nutritional Calorie is equal to 1000 calories (thermo-chemical calories). How many 8-in-high steps must a 180-lb man climb to expend the 400 Cal from the piece of cake? See Exercise 15 for the formula for potential energy.
In a bomb calorimeter, the bomb is surrounded by water that must be added for each experiment. Since the amount of water is not constant from experiment to experiment, mass must be measured in each case. The heat capacity of the calorimeter is broken down into two parts: the water and the
The bomb calorimeter in Exercise 87 is filled with 987 g of water. The initial temperature of the calorimeter contents is 23.32oC. A 1.056-g sample of benzoic acid (ΔEcomb = 226.42 kJ/ g) is combusted in the calorimeter. What is the final temperature of the calorimeter contents?
If a student performs an endothermic reaction in a calorimeter, how does the calculated value of ΔH differ from the actual value if the heat exchanged with the calorimeter is not taken into account?
The enthalpy of neutralization for the reaction of a strong acid with a strong base is 256 kJ/ mol of water produced. How much energy will be released when 200.0 mL of 0.400 M HNO3 is mixed with 150.0 mL of 0.500 M KOH?
Three gas- phase reactions were run in a constant-pressure piston apparatus as illustrated. For each reaction, give the balanced reaction and predict the sign of w (the work done) for the reaction.If just the balanced reactions were given, how could you predict the sign of w for a reaction?
Consider the following changes: a. N2(g) → N2(l) b. CO(g) + H2O(g) → H2(g) + CO2(g) c. Ca3P2(s) + 6H2O(l) → 3Ca(OH)2(s) + 2PH3(g) d. 2CH3OH(l) + 3O2(g) → 2CO2(g) + 4H2O(l) e. I2(s) → I2(g) At constant temperature and pressure, in which of these changes is work done by the system on the
a. Draw what the container will look like after the reaction has gone to completion. Assume a constant pressure of 1 atm.b. Is the sign of work positive or negative, or is the value of work equal to zero for the reaction? Explain your answer.
Using the following data, calculate the standard heat of formation of ICl(g) in kJ/ mol: Cl2(g) → 2Cl(g) ΔHo = 242.3 kJ I2(g) → 2I(g) ΔHo = 151.0 kJ ICl(g) → I(g) + Cl(g) ΔHo = 211.3 kJ I2(s) → I2(g) ΔH8 = 62.8 kJ
Consider a mixture of air and gasoline vapor in a cylinder with a piston. The original volume is 40 cm3. If the combustion of this mixture releases 950. J of energy, to what volume will the gases expand against a constant pressure of 650. torr if all the energy of combustion is converted into work
High- quality audio amplifiers generate large amounts of heat. To dissipate the heat and prevent damage to the electronic devices, manufacturers use heat- radiating metal fins. Would it be better to make these fins out of iron or aluminum? Why?Table
Write reactions that correspond to the following enthalpy changes: a. ΔHof for solid aluminum oxide b. the standard enthalpy of combustion of liquid ethanol [C2H5OH(l)]
The heat required to raise the temperature from 300.0 K to 400.0 K for 1 mole of a gas at constant volume is 2079 J. The internal energy required to heat the same gas at constant pressure from 550.0 K to 600.0 K is 1305 J. The gas does 150. J of work during this expansion at constant pressure. Is
When water is supercooled, it freezes at a temperature below 0.0oC. If 10.9 kJ of heat is released when 2.00 moles of supercooled water at 215.0oC freezes, calculate the molar enthalpy of fusion for ice at 0.0oC and 1 atm. Assume the molar heat capacities for H2O(s) and H2O(l) are 37.5 J K-1 mol-1
The sun supplies energy at a rate of about 1.0 kilowatt per square meter of surface area (1 watt = 1 J/s). The plants in an agricultural field produce the equivalent of 20 kg of sucrose (C12H22O11) per hour per hectare (1 ha = 10,000 m2). Assuming that sucrose is produced by the reaction 12CO2(g)
The heat of vaporization of water at the normal boiling point, 373.2 K, is 40.66 kJ/mol. The specific heat capacity of liquid water is 4.184 JK-1g-1 and of gaseous water is 2.02 J K-1g-1. Assume that these values are independent of temperature. What is the heat of vaporization of water at 298.2 K?
Consider the following reaction at 248oC and 1.00 atm: CH3Cl(g) + H2(g) → CH4(g) + HCl(g) For this reaction, the enthalpy change at 248oC is 283.3 kJ/ mol. At constant pressure the molar heat capacities (Cp) for the compounds are as follows: CH3Cl (48.5 JK-1mol-1), H2 (28.9 JK-1 mol-1), CH4 (41.3
The best solar panels currently available are about 15% efficient in converting sunlight to electricity. A typical home will use about 40 kWh of electricity per day (1 kWh = 1 kilowatt hour; 1 kW = 1000 J/ s). Assuming 8.0 hours of useful sunlight per day, calculate the mini-mum solar panel surface
You have 2.4 moles of a gas contained in a 4.0-L bulb at a temperature of 32oC. This bulb is connected to a 20.0-L sealed, initially evacuated bulb via a valve. Assume the temperature remains constant. a. What should happen to the gas when you open the valve? Calculate any changes of conditions.
An isothermal process is one in which the temperatures of the system and surroundings remain constant at all times. With this in mind, what is wrong with the following statement: “For an isothermal expansion of an ideal gas against a constant pressure, ΔT = 0, so q = 0”? What is q equal to in
One mole of H2O(g) at 1.00 atm and 100oC occupies a volume of 30.6 L. When 1 mole of H2O(g) is condensed to 1 mole of H2O(l) at 1.00 atm and 100oC, 40.66 kJ of heat is released. If the Density of H2O(l) at this temperature and pressure is 0.996 g/ cm3, calculate ΔE for the condensation of 1 mole
You have a 1.00-mole sample of water at 230oC, and you heat it until you have gaseous water at 140.oC. Calculate q for the entire process. Use the following data: Specific heat capacity of ice = 2.03 JoC-1g-1 Specific heat capacity of water = 4.18 JoC-1g-1 Specific heat capacity of steam = 2.02
A sample consisting of 22.7 g of a nongaseous, unstable compound X is placed inside a metal cylinder with a radius of 8.00 cm, and a piston is carefully placed on the surface of the compound so that, for all practical purposes, the distance between the bottom of the cylinder and the piston is zero.
A gaseous hydrocarbon reacts completely with oxygen gas to form carbon dioxide and water vapor. Given the following data, determine ΔHof for the hydrocarbon: ΔHrxn = –2044.5 kJ/ mol ΔHof (CO2) = –393.5 kJ/ mol ΔHof (H2O) = –242 kJ/ mol Density of CO2 and H2O mixture at 1 atm, 200oC =
Choose the substance with the larger positional probability in each case. a. 1 mole of H2 (at STP) or 1 mole of H2 (at 100oC, 0.5 atm) b. 1 mole of N2 (at STP) or 1 mole of N2 (at 100 K, 2.0 atm) c. 1 mole of H2O(s) (at 0oC) or 1 mole of H2O(l) (at 20oC)
In the roll of two dice, what total number is the most likely to occur? Is there an energy reason why this number is favored? Would energy have to be spent to increase the probability of getting a particular number (that is, to cheat)?
Entropy can be calculated by a relationship proposed by Ludwig Boltzmann: S = kB ln V where kB = 1.38 Ã— 10€“23 J/K and V is the number of ways a particular state can be obtained. (This equation is engraved on Boltzmann€™s tombstone.) Calculate S for the three
Calculate the energy required to change the temperature of 1.00 kg of ethane (C2H6) from 25.0oC to 73.4oC in a rigid vessel. (Cv for C2H6 is 44.60 J K-1 mol-1.) Calculate the energy required for this same temperature change at constant pressure. Calculate the change in internal energy of the gas in
For nitrogen gas the values of Cv and Cp at 25oC are 20.8 J K-1 mol-1 and 29.1 J K-1 mol-1, respectively. When a sample of nitrogen is heated at constant pres-sure, what fraction of the energy is used to increase the internal energy of the gas? How is the remainder of the energy used? How much
Consider a rigid, insulated box containing 0.400 mole of He(g) at 20.0oC and 1.00 atm in one compartment and 0.600 mole of N2(g) at 100.0oC and 2.00 atm in the other compartment. These compartments are connected by a partition that transmits heat. What is the final temperature in the box at thermal
One mole of an ideal gas is contained in a cylinder with a movable piston. The temperature is constant at 77oC. Weights are removed suddenly from the piston to give the following sequence of three pressures: a. P1 = 5.00 atm (initial state) b. P2 = 2.24 atm P3 = 1.00 atm (final state) What is the
One mole of an ideal gas with a volume of 1.0 L and a pressure of 5.0 atm is allowed to expand isothermally into an evacuated bulb to give a total volume of 2.0 L. Calculate w and q. Also calculate qrev for this change of state.
A cylinder with an initial volume of 10.0 L is fitted with a frictionless piston and is filled with 1.00 mole of an ideal gas at 25oC. Assume that the surroundings are large enough so that if heat is withdrawn from or added to it, the temperature does not change. a. The gas expands isothermally and
The molar heat capacities for carbon dioxide at 298.0 K are Cv = 28.95 J K-1 mol-1 Cp = 37.27 J K-1mol-1 The molar entropy of carbon dioxide gas at 298.0 K and 1.000 atm is 213.64 JK-1mol-1. a. Calculate the energy required to change the temperature of 1.000 mole of carbon dioxide gas from 298.0 K
The molar entropy of helium gas at 25oC and 1.00 atm is 126.1 J K-1mol-1. Assuming ideal behavior, calculate the entropy of the following. a. 0.100 mole of He(g) at 25oC and a volume of 5.00 L b. 3.00 moles of He(g) at 25oC and a volume of 3000.0 L
Consider the processwhich is carried out at constant pressure. The total DS for this process is known to be 75.0 J K-1mol-1. For A(l) and A(g), the Cp values are 75.0 JK-1mol-1 and 29.0 JK-1mol-1, respectively, and are not dependent on temperature. Calculate Î”H vaporization for A(l) at
A sample of ice weighing 18.02 g, initially at 230.0oC, is heated to 140.0oC at a constant pressure of 1.00 atm. Calculate q, w, ΔE, ΔH, and ΔS for this process. The molar heat capacities (Cp) for solid, liquid, and gaseous water— 37.5 JK-1mol-1, 75.3 JK-1mol-1, and 36.4 JK-1mol-1,
Calculate the entropy change for a process in which 3.00 moles of liquid water at 0oC is mixed with 1.00 mole of water at 100.oC in a perfectly insulated container. (Assume that the molar heat capacity of water is constant at 75.3 J K-1mol-1.)
Calculate the change in entropy that occurs when 18.02 g of ice at 210.0oC is placed in 54.05 g of water at 100.0oC in a perfectly insulated vessel. Assume that the molar heat capacities for H2O(s) and H2O(l) are 37.5 JK-1mol-1 and 75.3 JK-1mol-1, respectively, and the molar enthalpy of fusion for
The synthesis of glucose directly from CO2 and H2O and the synthesis of proteins directly from amino acids are both non-spontaneous processes under standard conditions. Yet these processes must occur for life to exist. In light of the second law of thermodynamics, how can life exist?
A green plant synthesizes glucose by photosynthesis as shown in the reaction 6CO2(g) + 6H2O(l) → C6H12O6(s) + 6O2(g) Animals use glucose as a source of energy: C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l) If we were to assume that both of these processes occur to the same extent in a cyclic process,
What determines ΔSsurr for a process? To calculate ΔSsurr at constant pressure and temperature, we use the following equation: ΔSsurr = 2ΔH/ T. Why does a minus sign appear in the equation, and why is ΔSsurr inversely proportional to temperature?
Predict the sign of ΔSsurr for the following processes. a. H2O(l) → H2O(g) b. I2(g) → I2(s)
Calculate Î”Ssurr for the following reactions at 25oC and 1 atm.
For each of the following pairs of substances, which sub-stance has the greater value of S8 at 25oC and 1 atm? a. Cgraphite(s) or Cdiamond(s) b. C2H5OH(l) or C2H5OH(g) c. CO2(s) or CO2(g) d. N2O(g) or He(g) e. HF(g) or HCl(g)
Predict the sign of Î”So for each of the following changes.a.b. AgCl(s) †’ Ag+(aq) + Cl-(aq) c. 2H2(g) + O2(g) †’ 2H2O(l) d. Na(s) + 1/2 Cl2(g) †’ NaCl(s) e. HCl(g) †’ H+(aq) + Cl2(aq) f. KBr(s) †’ K+(aq) + Br-(aq)
Predict the sign of ΔSo and then calculate ΔSo for each of the following reactions. a. 2H2S(g) + SO2(g) → 3Srhombic(s) + 2H2O(g) b. 2SO3(g) → 2SO2(g) + O2(g) c. Fe2O3(s) + 3H2(g) → 2Fe(s) + 3H2O(g)
For the reaction C2H2(g) + 4F2(g) → 2CF4(g) + H2(g) ΔSo is equal to -358 J/ K. Use this value and data from Appendix 4 to calculate the value of So for CF4(g).
Using Appendix 4 and the following data, determine So for Fe(CO)5(g). Fe(s) + 5CO(g) → Fe(CO)5(g) ΔSo = ? Fe(CO)5(l) → Fe(CO)5(g) ΔSo = 107 J/ K Fe(s) + 5CO(g) → Fe(CO)5(l) ΔSo = –677 J/ K
For the reaction 2Al(s) + 3Br2(l) → 2AlBr3(s) ΔSo is equal to –144 J/ K. Use this value and data from Appendix 4 to calculate the value of So for solid alumi-num bromide.
Ethanethiol (C2H5SH; also called ethyl mercaptan) is commonly added to natural gas to provide the “ rotten egg” smell of a gas leak. The boiling point of ethanethiol is 35ºC and its heat of vaporization is 27.5 kJ/ mol. What is the entropy of vaporization for this substance?
For mercury at 1 atm, the enthalpy of vaporization is 58.51 kJ/mol and the entropy of vaporization is 92.92 J K-1mol-1. What is the boiling point of mercury?
The enthalpy of vaporization of ethanol is 38.7 kJ/ mol at its boiling point (78oC). Determine ΔSsys, ΔSsurr, and ΔSuniv when 1.00 mole of ethanol is vaporized at 78oC and 1.00 atm.
For ammonia (NH3) the enthalpy of fusion is 5.65 kJ/ mol, and the entropy of fusion is 28.9 J K-1 mol-1. a. Will NH3(s) spontaneously melt at 200. K? b. What is the approximate melting point of ammonia?
It is quite common for a solid to change from one structure to another at a temperature below its melting point. For example, sulfur undergoes a phase change from the rhombic crystal structure to the monoclinic crystal form at temperatures above 95oC. a. Predict the signs of ΔH and ΔS for the
As O2(l) is cooled at 1 atm, it freezes at 54.5 K to form solid I. At a lower temperature, solid I rearranges to solid II, which has a different crystal structure. Thermal measurements show that ΔH for the I → II phase transition is –743.1 J/ mol, and DS for the same transition is 217.0 J K-1
From data in Appendix 4, calculate ΔHo, ΔSo, and ΔGo for each of the following reactions at 25oC. a. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) b. 6CO2(g) + 6H2O(l) → C6H12O6(s) + 6O2(g) Glucose c. P4O10(s) + 6H2O(l) → 4H3PO4(s) d. HCl(g) + NH3(g) → NH4Cl(s)
The value of ΔGo for the reaction 2C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(l) is –5490. kJ. Use this value and data from Appendix 4 to calculate the standard free energy of formation for C4H10(g).
Of the functions ΔHo, ΔSo, and ΔGo, which depends most strongly on temperature? When ΔGo is calculated at temperatures other than 25oC, what assumptions are generally made concerning ΔHo and ΔSo?
For the reaction at 298 K, 2NO2(g) ⇌ N2O4(g) the values of ΔHo and ΔSo are –58.03 kJ and –176.6 J/ K, respectively. What is the value of ΔGo at 298 K? Assuming that ΔHo and ΔSo do not depend on temperature, at what temperature is ΔGo = 0? Is ΔGo negative above or below this temperature?
Acrylonitrile is the starting material used in the manufacture of acrylic fibers (U. S. annual production capacity is more than 2 million pounds). Three industrial processes for the production of acrylonitrile are given below. Using data from Appendix 4, calculate Î”So, Î”Ho,
Consider the reaction 2POCl3(g) → 2PCl3(g) + O2(g) a. Calculate ΔGo for this reaction. The ΔGof values for POCl3(g) and PCl3(g) are –502 kJ/ mol and –270. kJ/ mol, respectively. b. Is this reaction spontaneous under standard conditions at 298 K? c. The value of ΔSo for this reaction is 179
Consider two reactions for the production of ethanol: C2H4(g) + H2O(g) → CH3CH2OH(l) C2H6(g) + H2O(g) → CH3CH2OH(l) + H2(g) Which would be more thermodynamically feasible? Why? Assume standard conditions and assume that ΔHo and ΔSo are temperature independent.
Using data from Appendix 4, calculate Î”Ho, Î”So, and DGo for the following reactions that produce acetic acid:Which reaction would you choose as a commercial method for producing acetic acid (CH3CO2H)? What temperature conditions would you choose for the reaction? Assume
Given the following data: 2C6H6(l) + 15O2(g) → 12CO2(g) + 6H2O(l) ΔGo = –6399 kJ C(s) + O2(g) → CO2(g) ΔGo = –394 kJ H2(g) + 1/2 O2(g) → H2O(l) ΔGo = –237 kJ calculate ΔGo for the reaction 6C(s) + 3H2(g) → C6H6(l)
For the reaction 2O(g) → O2(g) a. Predict the signs of ΔH and ΔS. b. Would the reaction be more spontaneous at high or low temperatures?
Hydrogen cyanide is produced industrially by the following exothermic reaction:Is the high temperature needed for thermodynamic or for kinetic reasons?
A reaction at constant T and P is spontaneous as long as ΔG is negative; that is, reactions always proceed as long as the products have a lower free energy than the reactants. What is so special about equilibrium? Why don’t reactions move away from equilibrium?
ΔG predicts spontaneity for a reaction at constant T and P, whereas ΔGo predicts the equilibrium position. Explain what this statement means. Under what conditions can you use ΔGo to determine the spontaneity of a reaction?

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