Worksheet 8 (thermodynamics) With Answers 544h5o

CHM2046 Discussion Class Worksheet 8 (Fall 2017) Student Name___________________________________________________

Section #__________

1. To what do the "system", "surroundings", and "universe" refer? "System" is the defined reaction or process in which you are interested; "surroundings" are the defined things other than the system that are relevant to the reaction or process (e.g., an aqueous reaction taking place in an insulated dewar or beaker is the system, whereas the water and the dewar/beaker are the surroundings); and the "universe" is the defined system plus the defined surroundings. 2. What does it mean when it is said that a system is "spontaneous" under certain conditions? A system undergos a "spontaneous" change when that change occurs by itself under specified conditions, without continuous input of energy from outside the system (i.e., from the surroundings). For example, once you crank the engine on your car, the combustion of fuel begins and continues spontaneously without your having to continually and repeatedly crank the engine. Likewise, once your start to roll a ball down a hill, it spontaneously rolls down the hill on its own without your having to continue to push it down the hill. A "nonspontaneous" system will be a system for which continuous input of energy from the surroundings is required. Also, and this is very important, a reaction will proceed spontaneously to equilibrium, regardless of whether the reaction is reactant- or product-favored, and any time something happens so that Q ≠ K, the reaction will spontaneously shift toward the right or left to re-establish equilibrium so that Q = K. 3. Can the spontaneity (or lack thereof) of a system be ascertained by the ΔH of the system alone? No because some systems proceed spontaneously under specified conditions, regardless of ΔH. For example, both of the following chemical processes occur spontaneously at 25°C: H2O(g) → H2O(l) ΔH < 0 (exothermic) CO2(s) → CO2(g) ΔH > 0 (endothermic) Therefore, other thermodynamic variables must be known before spontaneity of a system can be ascertained. 4. What is "entropy" ? Is entropy a state function like enthalpy ? The degree of randomness or disorder in a system. Yes entropy is a state function like enthalpy. 5. What is the 2nd law of thermodynamics ? The entropy change ΔSuniverse must be > 0 for a spontaneous process, which means that (ΔSsystem + ΔSsurroundings) must be > 0 for a spontaneous process. So, even if ΔSsystem is negative, you can still have a spontaneous reaction if ΔSsurroundings is more positive, so that ΔSuniverse = (ΔSsystem + ΔSsurroundings) > 0. At equilibrium ΔSuniverse = (ΔSsystem + ΔSsurroundings) = 0. 6. What is the 3rd law of thermodynamics ? A perfect defect-free crystalline substance exhibits zero entropy at absolute zero K. 7. Which of the following reactions would be predicted to be spontaneous at relatively high temperatures but nonspontaneous at relatively low temperatures ? Be sure to explain why in the space below the answer options. (1) an exothermic reaction with ΔS°reaction < 0 (2) an exothermic reaction with ΔS°reaction > 0 (3) an endothermic reaction with ΔS°reaction < 0 (4) an endothermic reaction with ΔS°reaction > 0 (5) a reaction with both ΔS°system < 0 and ΔS°surroundings < 0 An endothermic reaction (ΔH°reaction > 0) requires that the TΔS°reaction component be sufficiently > 0 and large enough in order to render the reaction spontaneous (ΔG°reaction < 0). Therefore ΔS°reaction must be > 0 and T must be relatively high. 8. If someone left the valve to a hydrogen gas tank open in a lab so that the partial pressures of H2 and N2 (in the air) in the lab reached 0.50 atm and 0.50 atm, respectively, would the following reaction forming highly toxic liquid hydrazine (N2H4) become a concern (i.e., is the following reaction likely to occur at 25°C, given these partial pressures of H2 and N2) ? 2H2(g) + N2(g) ↔ N2H4(l) ΔG°reaction = 149 kJ/mol (1) no concern because Qp > Kp and so ΔG > 0 (2) yes concern because Qp > Kp and so ΔG < 0 (3) yes concern because Qp < Kp and so ΔG < 0 (4) no concern because Qp < Kp and so ΔG > 0 (5) no concern because a reaction is always nonspontaneous when ΔS°reaction < 0 (answer continued on next page…)

Students need to calculate K from ΔG°reaction and then compare their calculated Q to their calculated K. If Q > K (which it is), then ΔG > 0 (which it is) and the reaction will favor reactants rather than products. 9. Select the false statement below. (1) if ΔG < 0, ΔG represents the maximum free energy or useful work obtainable from the reaction system (2) if ΔG > 0, ΔG represents the minimum work that must be done to the system for it to produce product(s) (3) if ΔG° = 0 then ΔG can be either = 0 or > 0 or < 0 (Students will likely need help with this one.) (4) ΔG is based upon the difference between Keq and Q of a reaction system at some temperature T (5) if ΔG° > 0 for a reaction at some temperature, that reaction is product-favored and Keq > 1 at that temperature 10. Consider the following reactions with their respective thermochemical data and determine whether the thermodynamic coupling of the reactions would result in the spontaneous (product-favored) reduction of Fe2O3 to Fe metal at 298K and/or at 1000K. (Assume that ΔS°rxn and ΔH°rxn do not change appreciably with temperature.) 2Fe2O3(s) ↔ 4Fe(s) + 3O2(g)ΔH° = +1651 kJ/mol ΔS° = +549 J/K•mol + 6H2(g) + 3O2(g) ↔ 6H2O(g) ΔH° = −1451 kJ/mol ΔS° = −266 J/K•mol = 2Fe2O3(s) + 6H2(g) ↔ 4Fe(s) + 6H2O(g) (hint: what is the ΔG° of this target reaction?) (1) nonspontaneous at both temperatures (2) spontaneous at 1000K but nonspontaneous at 298K (3) spontaneous at 298K but nonspontaneous at 1000K (4) spontaneous at both temperatures (5) based on the thermodynamics of the target reaction, it is nonspontaneous at all temperatures 11. Calculate the free energy (ΔG) for the following reaction if pieces of pure cium wire and pure iron wire were placed in a solution containing 1.0 M Fe2+ ions and 0.010 M Cd2+ ions at 25°C. Cd(s) + Fe2+(aq) ↔ Cd2+(aq) + Fe(s) ΔG° = +7.2 kJ/mol at 25°C (1) −2.1 kJ/mol (2) +2.1 kJ/mol (3) −4.2 kJ/mol (4) +4.2 kJ/mol (5) −8.4 kJ/mol (This is meant to be a preview of what we'll be getting into with electrochemistry later.) ________________________________________________________________________________________ Consider the formation reaction for CCl4(g) below when answering questions 12−14. C(graphite) + 2Cl2(g) ↔ CCl4(g) Kp = 2.57 x 109 at 25°C. 12. Which of the following statements is false regarding this reaction at equilibrium at 25°C ? (1) ΔG°system < 0 (2) ΔS°system < 0 (3) the reaction is endothermic (4) ΔGsystem = 0 (5) ΔS°surroundings > 0 13. In a reaction flask at 553K, 1.00 atm of gaseous Cl2 was allowed to react with the C(graphite) until equilibrium was reached with the gaseous CCl4 product. At this temperature of 553K, the %-yield of gaseous CCl4 was determined to be 90%. Determine ΔG°system at 553K. (hint: what was the actual yield at 553K ?) (1) −48.3 kJ/mol (2) −17.5 kJ/mol (3) −53.7 kJ/mol (4) −96.0 kJ/mol (5) −5.50 kJ/mol 14. If the reaction reaches standard-state equilibrium at 403°C, use this information along with the Kp information given for the reaction at 25°C to determine the standard molar enthalpy of formation (ΔH°f) of gaseous CCl4. (ln(K1/K2) = (ΔH°rxn / R){(1/T2) – (1/T1)}) (1) −48.3 kJ/mol (2) −17.5 kJ/mol (3) −53.7 kJ/mol (4) −96.0 kJ/mol (5) −5.50 kJ/mol Students will have to know that standard-state equilibrium means that ΔG°reaction = 0 and therefore Kp = 1, then use this Kp and the Kp above with their respective T values in Kelvin to solve for the ΔH°rxn which is ΔH°f of gaseous CCl4. Of course, R = 8.314 J/mol•K and ΔH°rxn has to be in Joules/mol.