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## Thermo-Chemistry

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**Thermo-Chemistry**Chapters 9 & 10 E-mail: benzene4president@gmail.com Web-site: http://clas.sa.ucsb.edu/staff/terri/**Thermodynamics - ch 9**1. For a particular process q = -10 kJ and w = 25 kJ. Which of the following statements is true? a. Heat flows from the surroundings to the system. b. The system does work on the surroundings. c. ∆E = -15 kJ d. All of these are true. e. None of these is true.**Thermodynamics - ch 9**q > 0 => endothermic => heat is flowing from the surroundings into the system q < 0 => exothermic => heat is flowing out of the system to the surroundings w > 0 => work is being done on the system by the surroundings when a gas is getting compressed by an external pressure w < 0 => the system is doing work on the surroundings when a gas is expanding against an external pressure ∆E > 0 => internal energy of a gas increases when the temperature of the gas increases ∆E < 0 => internal energy of a gas decreases when the temperature of the gas decreases ∆E = 0 => internal energy of a gas remains constant if isothermal**Thermodynamics - ch 9**2. A gas absorbs 4.8 J of heat and then performs 13.0 J of work. What is the change in internal energy of the gas?**Thermodynamics - ch 9**3. Which of the following statements is correct? a. The internal energy of a system increases when more work is done by the system than heat is flowing into the system. b. The internal energy of a system decreases when work is done on the system and heat is flowing into the system. c. The system does work on the surroundings when an ideal gas expands against a constant external pressure. d. All the statements are true. e. All the statements are false.**Thermodynamics - ch 9**4. Which of the following are always endothermic? a. Melting b. Combustion c. Condensation d. All of the above**Thermodynamics - ch 9**Breaking bonds requires energy => q > 0 Making bonds releases energy => q < 0**Thermodynamics - ch 9**5. The following reaction occurs at constant temperature and pressure. 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g) a. The system does work on the surroundings and w < 0 b. The surroundings do work on the system and w < 0 c. The system does work on the surroundings and w > 0 d. The surroundings do work on the system and w > 0 e. No work is done, w = 0**Thermodynamics - ch 9**3 equations for work 1) w = -Pext∆V 2) w = -∆nRT 3) w = -nR∆T equations 2 & 3 are derived from PV = nRT choose the equation by which variables are given in the problem**Thermodynamics - ch 9**6. There are two containers, one has 1 mole of CO2(g) and the other has 1 mole of Ne (g). Both gases are heated from 25 oC to 30 oC at constant pressure. For each of the following indicate true or false. a. No work is done in heating either gas b. More work is done in heating CO2(g) than for Ne (g) c. The change in internal energy of CO2(g) is greater than it is for Ne (g)**Thermodynamics - ch 9**7. For nitrogen gas at 25 oC, Cv = 20.8 J K-1 mol-1 and Cp = 29.1 J K-1 mol-1. When a sample of nitrogen gas is heated at constant pressure, what fraction of the energy is converted to heat? a. 1.00 b. 0.417 c. 0.285 d. 0.715 e. 0.588**Thermodynamics - ch 9**At constant pressure the fraction of energy converted to heat => Cv /Cp At constant volume the fraction of energy converted to heat => Cv /Cvor 1**Thermodynamics - ch 9**8. Which of the following are state functions? a. work, heat b. work, heat, enthalpy, energy c. enthalpy, energy d. work, heat, enthalpy e. heat, enthalpy, energy**Thermodynamics - ch 9**State Functions => pathway independent => since ∆ is defined as (final state – initial state) then ∆ anything can not depend on the states in between the initial and final states a.k.a. the “path” => if dealing with a state function you can use any path to solve for it => the easiest paths for ∆E and ∆H are… ∆E = qv ∆H = qp**Thermodynamics - ch 9**9. Which one of the following statements is false? a. The change in internal energy, ∆E, for a process is equal to the amount of heat absorbed at constant volume, qv. b. The change in enthalpy, ∆H, for a process is equal to the amount of heat absorbed at constant pressure, qp. c. A bomb calorimeter measures ∆H directly. d. If qp for a process is negative, the process is exothermic. e. The freezing of water is an example of an exothermic reaction.**Thermodynamics - ch 9**10. When a gas is cooled at constant volume from an initial temperature of 50 oC to 25 oC it loses 60 J of heat. When the same gas is cooled from 50 oC to 25 oC at constant pressure, it loses 100 J of heat. What is the value of the enthalpy change, ∆H, when the gas is heated from 25 oC to 50 oC at constant volume? a. + 60 J b. – 60 J c. – 100 J d. + 100 J e. – 40 J**Thermodynamics - ch 9**11. Suppose you add 45 J of heat to a system, let it do 10. J of expansion work, and then return the system to its initial state by cooling and compression. Which statement is true for this process? a. ∆H < ∆E b. The work done in compressing the system must exactly equal the work done by the system in the expansion step. c. ∆H = 70. J d. The change in the internal energy for this process is zero.**Thermodynamics - ch 9**12. Consider a gas in a 1.0-L bulb at STP that is connected via a valve to another bulb that is initially evacuated. After the valve is opened… a. q > 0 q < 0 q = 0 b. w > 0 w < 0 w = 0 c. ∆H> 0 ∆H< 0 ∆H= 0 d. ∆E> 0 ∆E< 0 ∆E= 0**Thermodynamics - ch 9**13. For the following reactions at constant pressure, predict if ∆H > ∆E, ∆H < ∆E, or ∆H = ∆E. Why is there a difference? Is work being done by the system or the surroundings? a. H2(g) + Cl2(g) 2 HCl (g) b. N2(g) + 3 F2(g) 2 NF3(g)**Thermodynamics - ch 9**14. Consider a sample containing 2 moles of xenon that undergoes the following changes: Pa = 10.0 atm Va = 10.0 L Pb = 10.0 atm Vb = 5.0 L Pc = 20.0 atm Vc = 25.0 L Step 1 Step 2 For each step, assume that the external pressure is constant and equals the final pressure of the gas for that step. Calculate q, w, ΔE andΔH (in kJ) for each step and calculate the total values. Which values would change if you went directly from state a to state c?**Thermodynamics - ch 9**15. Calculate ∆E (in kJ) when 1 mol of liquid is vaporized at 1 atm and its boiling point (73°C). ∆Hvap = 28.6 kJ/mol at 73°C.**Thermodynamics - ch 9**16.Consider the reaction: C2H5OH (l) + 3O2 (g) → 2CO2 (g) + 3H2O (l) ∆H = –1.37×103 kJ When a 15.6-g sample of ethyl alcohol (molar mass = 46.1 g/mol) is burned, how much energy is released as heat? A) 87.8 kJ B) 2.14 × 104 kJ C) 4.64 × 102 kJ D) 4.05 × 103 kJ E) 4.75 kJ**Thermodynamics - ch 9**17. Consider the following reaction: CH4 + 4 Cl2 (g) → CCl4 (g) + 4 HCl (g), ∆H = –434 kJ How many grams of methane must be burned to raise the temperature of 1.2 L of water 50°C (CH2O = 4.184 J/g°C)**Thermodynamics - ch 9**18. A 25.0 g piece of unknown metal was transferred from an oven at 115 °C into a coffee cup calorimeter containing 150. mL of water at 24 °C and allowed to come to equilibrium where the temperature was measured to be 28 °C. Calculate the specific heat capacity of the metal. (Ccal = 25 J/K and Cwater = 4.18 J/g°C)**Thermodynamics - ch 9**19. A 5.00 g of aluminum pellets (CAl = 0.89 J°C-1g-1) and 10.0 g of iron pellets (CFe = 0.45 J°C-1g-1) are heated to 100.0°C. The mixture of hot aluminum and iron is then dropped into 97.3 mL of water at 22.0°C. Calculate the equilibrium temperature assuming no heat is lost to the surroundings. CH2O = 4.18 Jg-1°C-1**Thermodynamics - ch 9**20. In a coffee cup calorimeter 50.0 mL of 0.10 M NaOH and 20.0 mL of 0.30 M HCl are mixed. The temperatures before and after the reaction were measured to be 23.0°C and 31.0°C. Calculate the molar enthalpy change for the neutralization reaction. Assume no heat is absorbed by the calorimeter. The density of each solution is 1.0 g/mL and the specific heat capacity for each solution is 4.18 J/g°C.**Thermodynamics - ch 9**21. Calculate the final temperature when 38 mL of 0.15 M Pb(NO3)2 is mixed with 42 mL of 0.23 M KI both initially at 23 °C. Assume no heat is lost to the surroundings. The specific heat capacity for the solution is 4.18 J/g°C and the density is 1.0 g/mL. Pb2+ + 2 I- PbI2 ∆H = -297 kJ/mol**Thermodynamics - ch 9**22. Given the following data calculate the ∆H° for the reaction: NO + O NO2 Rxn 1: 2 O3 3 O2 ∆H° = -427 kJ Rxn 2: O2 2 O ∆H° = +495 kJ Rxn 3: NO + O3 NO2 + O2 ∆H° = -195 kJ**Thermodynamics - ch 9**Hess’ Laws 1. If a reaction is flipped => change the sign of ∆H 2. If a reaction is multiplied by a number => multiply ∆H by the same number 3. If two or more reactions are added => take the sum of the ∆Hs**Thermodynamics - ch 9 - Feldwinn**23. Calculate ΔH for the reaction N2H4(l) + O2(g) N2(g) + 2 H2O (l) given the following data: 2 NH3(g)+3 N2O(g)4 N2(g)+3 H2O(l)ΔH = -1010. kJ N2O(g)+3 H2(g)N2H4(l)+H2O(l) ΔH= -317 kJ 2 NH3(g)+ ½ O2(g) N2H4(l)+H2O(l) ΔH= -143 kJ H2(g) + ½ O2(g) H2O (l) ΔH= -286 kJ**Thermodynamics - ch 9**24. Choose the correct equation for the standard enthalpy of formation of CO (g), where ∆H°ffor CO = –110.5 kJ/mol. a. 2 Cgraphite (s) + O2 (g) → 2 CO (g) ∆H° = –110.5 kJ b. Cgraphite (s) + O (g) → CO (g) ∆H° = –110.5 kJ c. Cgraphite (s) + ½ O2 (g) → CO (g) ∆H° = –110.5 kJ d. Cgraphite (s) + CO2 (g) → 2 CO (g) ∆H° = –110.5 kJ e. CO (g) → Cgraphite (s) + O (g) ∆H° = –110.5 kJ**Thermodynamics - ch 9**25. The combustion of methanol takes place according to the reaction 2 CH3OH (l) + 3 O2(g) → 2 CO2(g) + 4 H2O (l) Calculate ∆H for the combustion of 1 mol of methanol under standard conditions. Use the following standard enthalpies of formation: ∆H°f for CH3OH (l) = -238.5 kJ/mol ∆H°f for CO2(g) = -393.5 kJ/mol ∆H°f for H2O (l) = -285.6 kJ/mol**Thermodynamics - ch 9**Hess’ Equation: ∆H°rxn = Σ ∆H°f(products) - Σ ∆H°f(reactants)**Thermodynamics - ch 9**26. The heat combustion of acetylene, C2H2 (g), at 25°C, is –1299 kJ/mol. At this temperature, ∆H°f values for CO2 (g) and H2O (l) are –393 and –286 kJ/mol, respectively. Calculate ∆H°f for acetylene. a. 2376 kJ/mol b. 625 kJ/mol c. 227 kJ/mol d. –625 kJ/mol e. none of these**Thermodynamics - ch 10**27. Which has the greatest entropy? a. 1 mol of He at STP or 1 mol of He at 25°C b. 1 mol of Ne at STP or 1 mol of CH4 at STP c. 1 mol of Cl2 at STP or 1 mol of F2 at STP**Thermodynamics - ch 10**Entropy Considerations: 1. S(s) < S(l) << S(g) 2. There is more entropy at higher temperatures and/or larger volumes (lower pressures) 3. The more bonds per molecule the greater the positional probability ex: CH4 > H2 4. If there are the same number of atoms in the molecules/elements; then the one with more electrons has the greater the positional probability ex: Ar > He 5. For the same atom but different structures (allotropes) the positional probability is greater in the more disordered structure ex: C (graphite) > C (diamond)**Thermodynamics - ch 10**28. Predict if ∆Ssys and ∆Ssurr is positive or negative for the following under standard conditions. a. melting ice b. photosynthesis => 6 CO2(g) + 6 H2O (l) C6H12O6(s) + 6 O2(g) c. precipitation of AgCl**Thermodynamics - ch 10**29. One mole of an ideal gas at 25°C is expanded isothermally and reversibly from 125.0 L to 250.0 L. Which statement is correct? a. ∆Sgas = 0 b. ∆Ssurr= 0 c. ∆Suniv = 0**Thermodynamics - ch 10**30. One mole of an ideal gas is compressed isothermally and reversibly at 607.4 K from 5.60 atm to 8.90 atm. Calculate ∆S for the gas. a. 2.34 J/K b. -2.34 J/K c. -3.85 J/K d. 3.85 J/K e. 0 J/K**Thermodynamics - ch 10**31. In a certain reversible expansion, a system at 300. K absorbs exactly 6.00 × 102 J of heat. In the irreversible recompression to the original state of the system, twice as much work is done on the system as is performed on the surroundings in the expansion. What is the entropy change of the system in the recompression step? a) -4.00 J/K b) -2.00 J/K c) 0.00 J/K d) 2.00 J/K e) 4.00 J/K**Thermodynamics - ch 10**32. A system composed of a ideal gas expands spontaneously in one step from an initial volume of 1.00 L to a final volume of 2.00 L at a constant temperature of 200 K. During the process the gas does 200 J of work. What conclusion can be reached about the value of the entropy change, ΔS, for this process? a) ΔS = + 1.00 J/K b) ΔS = – 1.00 J/K c) ΔS is less than +1.00 J/K d) ΔS is greater than +1.00 J/K e) None of the above**Thermodynamics - ch 10**33. Calculate the change in entropy for a process in which 3.00 moles of liquid water at 0°C is mixed with 1.00 mole of liquid water at 100°C in a perfectly insulated container. The molar heat capacity of liquid water is 75.3 Jmol-lK-1**Thermodynamics - ch 10**34. Calculate ∆S when 54 g of water is heated from -22°C to 156 °C at a constant pressure of 1 atm. The heat capacities for solid, liquid and gaseous water are 2.03 J/g°C, 4.18 J/g°C and 2.02 J/g°C respectively. The enthalpies of fusion and vaporization are 6.01 kJ/mol and 40.7 kJ/mol respectively.**Thermodynamics - ch 10 - Felwinn**35. Consider the process A (l) at 75°C A (g) at 155°C which is carried out at constant pressure. The total ΔS for this process is 75.0 Jmol-1K-1. For A (l) and A (g) the Cp values are 75.0 Jmol-1K-1 and 29.0 Jmol-1K-1 respectively. Calculate ΔHvap at 125°C (its boiling point).**Thermodynamics - ch 10**36. Indicate true or false for each of the following statements. a. Spontaneous reactions must have a positive ΔSº for the reaction. b. When the change in free energy is less than zero for a chemical reaction, the reaction must be exothermic. c. For a spontaneous reaction, if ΔSº < 0 then the reaction must be exothermic.**Thermodynamics - ch 10**Spontaneous => wants to go forward on its own K > Q ∆Suniverse > 0 or ∆Gsystem < 0 Non-spontaneous => wants to go backward on its own K < Q ∆Suniverse < 0 or ∆Gsystem > 0 Equilibrium => doesn’t prefer one direction over the other K = Q ∆Suniverse = 0 or ∆Gsystem = 0**Thermodynamics - ch 10**** Note => If ∆H and ∆S are the same sign Teq = ∆H /∆S**Thermodynamics - ch 10**37. A 100-mL sample of water is placed in a coffee cup calorimeter. When 1.0 g of an ionic solid is added, the temperature decreases from 21.5°C to 20.8°C as the solid dissolves. Which of the following is true for the dissolving of the solid? a. ∆H < 0 b. ∆Suniv > 0 c. ∆Ssys < 0 d. ∆Ssurr > 0 e. none of these**Thermodynamics - ch 10**38. At what temperatures (high or low) would the dissociation of hydrogen be spontaneous? H2 (g) 2 H (g)**Thermodynamics - ch 10**39. For the freezing of water at -10 °C and 1 atm, predict whether ∆H, ∆S, and ∆G should be positive, negative or zero.