<p>AP Chemistry UNIT #4 HW Packet (Chapters 6, 17, 12 & 18) Chapter 6 - Thermochemistry Section 1: Multiple Choice 1) CH4(g) + 2 O2(g) ---> CO2(g) + 2 H2O(l); ΔH = - 889.1 kJ ΔHf° H2O(l) = -285.8 kJ / mole ΔHf° CO2(g) = - 393.3 kJ / mole What is the standard heat of formation of methane, ΔHf° CH4(g), as calculated from the data above? (A) -210.0 kJ/mole (B) -107.5 kJ/mole (C) -75.8 kJ/mole (D) 75.8 kJ/mole (E) 210.0 kJ/mole </p><p>2) Which of the following is a graph that describes the pathway of reaction that is endothermic and has high activation energy? (A) (B)</p><p>(C) (D) (E) </p><p>3)  H2(g) + 1/2 O2(g) ---> H2O(l) Δ H° = x  2 Na(s) + 1/2 O2(g) ---> Na2O(s) Δ H° = y  Na(s) + 1/2 O2(g) + 1/2 H2(g) ---> NaOH(s) Δ H° = z Based on the information above, what is the standard enthalpy change for the following reaction? Na2O (s) + H2O (l) ---> 2 NaOH (s) (A) x + y + z (B) x + y - z (C) x + y - 2z (D) 2z - x - y (E) z - x - y 4) The energy diagram for the reaction X + Y ---> Z is shown below. The addition of a catalyst to this reaction would cause a change in which of the indicated energy differences?</p><p>(A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 5) Which of the following best describes the role of the spark from the spark plug in an automobile engine? (A) The spark decreases the energy of activation for the slow step. (B) The spark increases the concentration of the volatile reactant. (C) The spark supplies some of the energy of activation for the combustion reaction. (D) The spark provides a more favorable activated complex for the combustion reaction. (E) The spark provides the heat of vaporization for the volatile hydrocarbon. </p><p>6) C2H4(g) + 3 O2(g) --> 2 CO2(g) + 2 H2O(g) For the reaction of ethylene represented above, ΔH is - 1,323 kJ. What is the value of ΔH if the combustion produced liquid water H2O (l), rather than water vapor H2O(g)? (ΔH for the phase change H2O(g) --> H2O(l) is -44 kJ mol¯1.) (A) -1,235 kJ (B) -1,279 kJ (C) -1,323 kJ (D) -1,367 kJ (E) -1,411 kJ 7)</p><p>I2 (g) + 3 Cl2 (g) ---> 2 ICl3 (g) According to the data in the table below, what is the value of ΔH° for the reaction represented above? </p><p>Bond Average Bond Energy (kj/mole) </p><p>I---I 150 </p><p>Cl---Cl 240</p><p>I---Cl 210 </p><p>(A) - 870 kJ (B) - 390 kJ (C) + 180 kJ (D) + 450 kJ (E) + 1,260 kJ Section II: Free Response FRQ #1 Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking. (a)Write a balanced equation for the complete combustion of propane gas, which yields CO2(g) and H2O(l).</p><p>(b) Calculate the volume of air at 30°C and 1.00 atmosphere that is needed to burn completely 10.0 grams of propane. Assume that air is 21.0 percent O2 by volume.</p><p>(c) The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of formation, ΔHf°, of propane given that ΔHf° of H2O(l) = -285.3 kJ/mol and ΔHf° of CO2(g) = -393.5 kJ/mol.</p><p>(d) Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred to 8.00 kilograms of water (specific heat = 4.18 J/g.K), calculate the increase in temperature of water FRQ#2</p><p>–1 N2(g) + 3 F2(g) ® 2 NF3(g) ΔH° = – 264 kJ mol ; The enthalpy change in a chemical reaction is the difference between energy absorbed in breaking bonds in the reactants and energy released by bond formation in the products. (a) How many bonds are formed when two molecules of NF3 are produced according to the equation in the box above? (b) Use both the information above and the table of average bond enthalpies below to calculate the average enthalpy of the F–F bond.</p><p>Average Bond Enthalpy Bond (kJ mol-1) </p><p>N 946 </p><p>N–F 272 </p><p>F–F ? </p><p>FRQ#3</p><p>CO(g) + ½ O2(g) ® CO2(g) The combustion of carbon monoxide is represented by the equation above. (a) Determine the value of the standard enthalpy change, ∆H˚rxn for the combustion of CO(g) at 298 K using the following information.</p><p> C(s) + ½ O2(g) ® CO(g) ∆H˚298 = –110.5 kJ mol-1</p><p> C(s) + O2(g) ® CO2(g) ∆H˚298 = –393.5 kJ mol-1 FRQ #4 Enthalpy of Combustion, ∆H° Substance (kiloJoules/mol) C(s) -393.5 H2(g) -285.8 C2H5OH(l) -1366.7 H2O(l) -- (a) Write a separate, balanced chemical equation for the combustion of each of the following: C(s), H2(g), and C2H5OH(l). Consider the only products to be CO2 (g) and/or H2O(l). (b) In principle, ethanol can be prepared by the following reaction:</p><p>2 C(s) + 2H2(g) + H2O(l) à C2H5OH(l) Calculate the standard enthalpy change, ∆H°, for the preparation of ethanol, as shown in the reaction above. FRQ#5 ∆Hf° Compound (kilocalories/mole) H2O(l) -68.3 CO2(g) -94.1 O2(g) 0.0 C3H8 ? When 1.000 gram of propane gas, C3H8, is burned at 25°C and 1.00 atmosphere, H2O(l) and CO2(g) are formed with the evolution of 12.03 kilocalories. (a) Write a balanced equation for the combustion reaction. (b) Calculate the molar enthalpy of combustion, ∆H°comb, of propane. (c) Calculate the standard molar enthalpy of formation, ∆Hf°, of propane gas. Equations: For the reactions below write a balanced NET ionic reaction and predict what type of question will be asked with the reaction: a. carbon dioxide is bubbled into water b. equimolar solutions of ammonium chloride and potassium hydroxide are mixed. c. hydrogen peroxide is catalytically decomposed. d. the electrolysis of water e. equimolar solutions of barium nitrate and potassium fluoride are mixed. Chapter 17 - Spontaneity, Entropy, & Free Energy Section 1: Multiple Choice (1) Which of the following reactions has the largest positive value of entropy per mole of Cl2 ? (A) H2(g) + Cl2(g) ---> 2 HCl(g) (B) Cl2(g) + 1/2 O2(g) ---> Cl2O(g) (C) Mg(s) + Cl2(g) ---> MgCl2(s) (D) 2 NH4Cl (s) ---> N2(g) + 4 H2(g) + Cl2(g) (E) Cl2(g) ---> 2 Cl (g) </p><p>(2) For which of the following processes would ΔS have a negative value? I. 2 Fe2O3(s) ---> 4 Fe(s) + 3 O2(g) II. Mg2+ + 2 OH¯ ---> Mg(OH)2(s) III. 3H2(g) + 3C2H4(g) ---> 3 C2H6(g) (A) I only (B) I and II only (C) I and III only (D) II and III only (E) I, II, and III (3) N2 (g) + 3 H2 (g) ---> 2 NH3 (g) The reaction indicated above is thermodynamically spontaneous at 298 K, but becomes nonspontaneous at higher temperatures. Which of the following is true at 298 K? (A) ΔG, ΔH, and ΔS are all positive. (B) ΔG, ΔH, and ΔS are all negative. (C) ΔG and ΔH are negative, but ΔS is positive. (D) ΔG and ΔS are negative, but ΔH is positive. (E) ΔG and ΔH are positive, but ΔS is negative</p><p>(4) H2O (s) ---> H2O (l) When ice melts at its normal melting point, 298.16 K and 1 atmosphere, which of the following is true for the process shown above? (A) ΔH < 0, ΔS > 0, ΔG > 0 (B) ΔH < 0, ΔS < 0, ΔG > 0 (C) ΔH > 0, ΔS < 0, ΔG < 0 (D) ΔH > 0, ΔS > 0, ΔG > 0 (E) ΔH > 0, ΔS > 0, ΔG < 0</p><p>(5) Of the following reaction, which involves the largest decrease in Entropy? (A) CaCO3 (s) ---> CaO (s) + CO2 (g) (B) 2 CO (g) + O2 (g) ---> 2 CO2 (g) (C) Pb(NO3)3 (aq) + 2 KI (aq) ---> PbI2 (s) + 2 KNO3 (aq) (D) C3H8 (g) + 5O2 (g) ---> 3 CO2 (g) + 4 H2O (g) (E) 4 La (s) + 3 O2 (g) ---> 2 La2O3 (s) Section 2: FRQ FRQ #1 Answer the following questions about nitrogen, hydrogen, and ammonia. ▸ (a) In the boxes below, draw the complete Lewis electron-dot diagrams for N2 and NH3.</p><p>▸ (b) Calculate the standard free-energy change, ΔG°, that occurs when 12.0 g of H2(g) reacts with excess N2(g) at 298 K according to the reaction represented below: N2(g) + 3 H2(g) ® 2 NH3(g) ∆G˚298 = -34 kJ mol-1 ▸ (c) Given that ΔH˚298 for the reaction is −92.2 kJ mol-1, which is larger, the total bond dissociation energy of the reactants or the total bond dissociation energy of the products? Explain. ▸ (d) The value of the standard entropy change, ΔS˚298 , for the reaction is −199 J mol-1K-1. Explain why the value of ΔS˚298 is negative. ▸ (e) Assume that ΔH° and ΔS° for the reaction are independent of temperature. ◦ (i) Explain why there is a temperature above 298 K at which the algebraic sign of the value of ΔG° changes. ◦ (ii) Theoretically, the best yields of ammonia should be achieved at low temperatures and high pressures. Explain. FRQ #2 Use principles of thermodynamics to answer the following questions. ▸ (a) The gas N2O4 decomposes to form the gas NO2. ◦ (i) Predict the sign of ∆H˚ for the reaction. Justify your answer. ◦ (ii) Predict the sign of ∆S˚ for the reaction. Justify your answer. ▸ (b) One of the diagrams below best represents the relationship between ∆G˚ and temperature for the reaction given in part (a). Assume that ∆H˚ and ∆S˚ are independent of temperature. ▸ Draw a circle around the correct graph. Explain why you chose that graph in terms of the relationship: ∆G˚ = ∆H˚ - T∆S˚. FRQ #3</p><p>2 Fe (s) + 3/2 O2 (g) à Fe2O3 (s) ∆Hf˚ = -824 kJ mol–1 Iron reacts with oxygen to produce iron(III) oxide as represented above. A 75.0 g sample of Fe(s) is mixed with 11.5 L of O2(g) at 2.66 atm and 298 K. (a) Calculate the number of moles of each of the following before the reaction occurs. (i) Fe(s) (ii) O2(g) (b) Identify the limiting reactant when the mixture is heated to produce Fe2O3. Support your answer with calculations. (c) Calculate the number of moles of Fe2O3 produced when the reaction proceeds to completion. (d) The standard free energy of formation, ∆Gf˚ of Fe2O3 is –740. kJ mol–1 at 298 K. (i) Calculate the standard entropy of formation ∆Sf˚ of Fe2O3 at 298 K. Include units with your answer. (ii) Which is more responsible for the spontaneity of the formation reaction at 298K, the standard enthalpy or the standard entropy? The reaction represented below also produces iron(III) oxide. The value of ∆H˚ for the reaction is –280 kJ per mol.</p><p>2 FeO (s) + ½ O2 (g) à Fe2O3 (s) (e) Calculate the standard enthalpy of formation, ∆Hf˚ of FeO(s). FRQ #4 For the gaseous equilibrium represented below, it is observed that greater amounts of PCl3 and Cl2 are produced as the temperature is increased: PCl5(g) à PCl3(g) + Cl2(g) ▸ (a) What is the sign of ∆S° for the reaction? Explain. ▸ (b) What change, if any, will occur in DG° for the reaction as the temperature is increased? Explain your reasoning in terms of thermodynamic principles. ▸ (c) If He gas is added to the original reaction mixture at constant volume and temperature, what will happen to the partial pressure of Cl2? Explain. ▸ (d) If the volume of the reaction mixture is decreased at constant temperature to half the original volume, what will happen to the number of moles of Cl2 in the reaction vessel? Explain. FRQ #5</p><p>2 C4H10 (g) + 13 O2 (g) à 8 CO2 (g) + 10 H2O (l) The reaction represented above is spontaneous at 25°C. Assume that all reactants and products are in their standard state. ▸ (a) Predict the sign of ∆S° for the reaction and justify your prediction. ▸ (b) What is the sign of ∆G° for the reaction? How would the sign and magnitude of ∆G° be affected by an increase in temperature to 50°C? Explain your answer. ▸ (c) What must be the sign of ∆H° for the reaction at 25°C? How does the total bond energy of the reactants compare to that of the products? ▸ (d) When the reactants are place together in a container, no change is observed even though the reaction is known to be spontaneous. Explain this observation. Chapter 12 - Chemical Kinetics Section 1: Multiple Choice (1) Relatively slow rates of chemical reaction are associated with which of the following? (A) The presence of a catalyst (B) High temperature (C) High concentration of reactants (D) Strong bonds in reactant molecules (E) Low activation energy (2) The proposed steps for a catalyzed reaction between Ce4+ and Tl+ are represented above. The products of the overall catalyzed reaction are Step 1: Ce4+ + Mn2+ ---> Ce3+ + Mn3+ Step 2: Ce4+ + Mn3+ ---> Ce3+ + Mn4+ Step 3: Mn4+ + Tl+ ---> Tl3+ + Mn2+ (A) Ce4+ and Tl+ (B) Ce3+ and Tl3+ (C) Ce3+ and Mn3+ (D) Ce3+ and Mn4+ (E) Tl3+ and Mn2+ (3) (CH3)3CCl(aq) + OH¯ ---> (CH3)3COH(aq) + Cl¯ For the reaction represented above, the experimental rate law is given as follows. Rate = k [(CH3)3CCl] If some solid sodium solid hydroxide is added to a solution that is 0.010-molar in (CH3)3CCl and 0.10-molar in NaOH, which of the following is true? (Assume the temperature and volume remain constant.) (A) Both the reaction rate and k increase. (B) Both the reaction rate and k decrease. (C) Both the reaction rate and k remain the same. (D) The reaction rate increases but k remains the same. (E) The reaction rate decreases but k remains the Questions 4 & 5: H3AsO4 + 3I¯ + 2 H3O+ ---> H3AsO3 + I3¯ + H2O The oxidation of iodide ions by arsenic acid in acidic aqueous solution occurs according to the stoichiometry shown above. The experimental rate law of the reaction is: Rate = k [H3AsO4] [I¯] [H3O+] 4. What is the order of the reaction with respect to I¯? (A) 1 (B) 2 (C) 3 (D) 5 (E) 6 5. According to the rate law for the reaction, an increase in the concentration of hydronium ion has what effect on this reaction? (A) The rate of reaction increases. (B) The rate of reaction decreases. (C) The value of the equilibrium constant increases. (D) The value of the equilibrium constant decreases. (E) Neither the rate nor the value of the equilibrium constant is changed. Section 2: FRQ A(g) + B(g) ® C(g) + D(g)  For the gas-phase reaction represented above, the following experimental data were obtained:</p><p>Experiment Initial [A] (mol L-1) Initial [B] (mol L-1) Initial Reaction Rate (mol L-1 s-l) </p><p>1 0.033 0.034 6.67´10-4 </p><p>2 0.034 0.137 1.08´10-2</p><p>3 0.136 0.136 1.07´10-2</p><p>4 0.202 0.233 ?</p><p>(a) Determine the order of the reaction with respect to reactant A. Justify your answer. (b) Determine the order of the reaction with respect to reactant B. Justify your answer. (c) Write the rate law for the overall reaction. (d) Determine the value of the rate constant, k, for the reaction. Include units with your answer. (e) Calculate the initial reaction rate for experiment 4. (f) The following mechanism has been proposed for the reaction. Step 1: B + B ® E + D slow Step 2: E + A ® B + C fast equilibrium • Provide two reasons why the mechanism is acceptable. (g) In the mechanism in part (f), is species E a catalyst, or is it an intermediate? Justify your answer. Chapter 18 - Electrochemistry FRQ #1 Answer the following questions related to chemical reactions involving nitrogen monoxide, NO(g). The reaction between solid copper and nitric acid to form copper(II) ion, nitrogen monoxide gas, and water is represented by the following equation.</p><p>- + 2+ 3 Cu(s) + 2 NO3 (aq) + 8 H (aq)  3 Cu (aq) + 2 NO(g) + 4 H2O(l) E˚ = +0.62 V (a) Using the information above and in the table below, calculate the standard reduction potential, E˚, for - the reduction of NO3 in acidic solution. Half-Reaction Standard Reduction Potential, E0 Cu2+(aq) + 2 e-  Cu(s) +0.34 V</p><p>- + - NO3 (aq) + 4 H (aq) + 3 e  NO(g) + 2 ?</p><p>H2O(l)</p><p>(b) Calculate the value of the standard free energy change, ∆G˚, for the overall reaction between solid copper and nitric acid. (c) Predict whether the value of the standard entropy change, ∆S˚, for the overall reaction is greater than 0, less than 0, or equal to 0. Justify your prediction. Nitrogen monoxide gas, a product of the reaction above, can react with oxygen to produce nitrogen dioxide gas, as represented below. </p><p>2 NO(g) + O2(g)  2 NO2(g)</p><p>Initial Concentration of Initial Concentration of O2 Initial Rate of Formation Experiment -1 -1 -1 -l NO (mol L ) (mol L ) of NO2 (mol L s )</p><p>1 0.0200 0.0300 8.52  10-2 2 0.0200 0.0900 2.56  10-1 3 0.0600 0.0300 7.67  10-1</p><p>(d) Determine the order of the reaction with respect to each of the following reactants. Give details of your reasoning, clearly explaining or showing how you arrived at your answers. </p><p>(i) NO (ii) O2 (e) Write the expression for the rate law for the reaction as determined from the experimental data. (f) Determine the value of the rate constant for the reaction, clearly indicating the units. </p><p>FRQ #2 An external direct-current power supply is connected to two platinum electrodes immersed in a beaker containing 1.0 M CuSO4(aq) at 25˚C, as shown in the diagram above. As the cell operates, copper metal is deposited onto one electrode and O2(g) is produced at the other electrode. The two reduction half-reactions for the overall reaction that occurs in the cell are shown in the table below. Half-Reaction E0(V)</p><p>+ - O2(g) + 4 H (aq) + 4 e  2 +1.23</p><p>H2O(l) Cu2+(aq) + 2 e-  Cu(s) +0.34</p><p>(a) On the diagram, indicate the direction of electron flow in the wire. (b) Write a balanced net ionic equation for the electrolysis reaction that occurs in the cell. (c) Predict the algebraic sign of ∆G˚ for the reaction. Justify your prediction. (d) Calculate the value of ∆G˚ for the reaction. An electric current of 1.50 amps passes through the cell for 40.0 minutes. (e) Calculate the mass, in grams, of the Cu(s) that is deposited on the electrode.</p><p>(f) Calculate the dry volume, in liters measured at 25˚C and 1.16 atm, of the O2(g) that is produced.</p>