Electrochemistry - Be able to predict products of an electrolysis reaction that occurs in water. As well as design a redox titration experiment.

Transcription

1 Big Idea 3: Chemical Reactions Problem Set: / 25 pts DUE: Mon. 3/9 Big Idea 3: AP Chemistry is a rigorous course that requires students to use their analytical and problem solving skills as they apply into the 6 Big Idea Divisions the College Board has designated. As such, according to the College Board, in order to receive a 5 on the AP Test a student should have the following understandings and abilities. Physical and Chemical Processes - Capable of identifying connections between symbolic representations of reactions and energies associated with change and equilibrium, e.g., recounting that not all reactions go to completion. Able to classify evidence (data) as suggesting a physical versus chemical change, for ambiguous cases, e.g., dissolution of a salt Electrochemistry - Be able to predict products of an electrolysis reaction that occurs in water. As well as design a redox titration experiment. A great video concept approach to each Big Idea may be found on BOZEMANSCIENCE at Big Idea Sheet: Recall, your Big Idea Sheet will be graded based on the rubric you were provided when your notebook was setup. If you have misplaced it you may of course find a PDF copy on the website at When writing your Big Idea, be sure to refer to the rubric and check it again before you submit is. Big Idea Sheets are due at the end of each Big Idea Unit and may be submitted electronically or on paper. NO EXCEPTIONS OR LATE WORK. Rough drafts are welcome and will be handed/sent back with comments upon if arrangements are made with your teacher before the due date for use of a rough draft. Big IDEA 3 SHEET DUE: Mon. 3/9 Learning Objectives that should be covered on Big Idea Sheet = LO 1 Students can translate among macroscopic observations of change, chemical reactions and particle views. Denoting Change - A change may be represented by a molecular, ionic or net ionic equation. LO 2 The student can translate an observed chemical change into a balanced chemical equation and justify the choice of equation type (molecular, ionic or net ionic) in terms of utility for the given circumstances. Stoichiometry - Quantitative information can be derived from stoichiometric calculations that utilize the mole ratios from the balanced chemical reactions. The role of stoichiometry in real-world applications is important to note, so that it does not seem to be simply an exercise done only by chemists. LO 3 The student is able to use stoichiometric calculations to predict the results of performing a reaction in the laboratory and/or to analyze deviations from the expected results LO 4 The student is able to relate quantities (measured mass of substances, volumes of solutions, or volumes and pressures of gases) to identify stoichiometric relationships for a reaction, including situations involving limiting reactants and situations in which the reaction has not gone to completion. Reaction Types - Chemical reactions can be classified by considering what the reactants are, what the products are, or how they change from one into the other. Classes of chemical reactions include synthesis, decomposition, acid-base and oxidation-reduction reactions. LO 5 The student is able to design a plan in order to collect data on the synthesis or decomposition of a compound to confirm the conservation of matter and the law of definite proportions. Acid-Base Reaction Type - In a neutralization reaction, protons are transferred from an acid to a base. LO 6 The student is able to use data from synthesis or decomposition of a compound to confirm the conservation of matter and the law of definite proportions. LO 7 The student is able to identify compounds as Bronsted-Lowry acids, bases, and/or conjugate acid-base pairs, using proton-transfer reactions to justify the identification. Exclusion Statement: Lewis acid-base concepts are beyond the scope of this course and the AP Exam. Redox Reaction Types - In oxidation-reduction (redox) reactions, there is a net transfer of electrons. The species that loses electrons is oxidized, and the species that gains electrons is reduced. LO 8 The student is able to identify redox reactions and justify the identification in terms of electron transfer. LO 9 The student is able to design and/or interpret the result of an experiment involving a redox titration Exclusion Statement: Language of reducing agent and oxidizing agent is beyond the scope of this course and the AP Exam. Chemical Change - Chemical and physical transformations may be observed in several ways and typically involve a change in energy. Production of heat or light, formation of a gas, and formation of a precipitate and/or a color change are possible evidences that a chemical change has occurred. LO 10 The student is able to evaluate the classification of a process as a physical change, chemical change based on both macroscopic observations and the distinction between rearrangement of covalent interactions and noncovalent interactions. Energy of Change - New changes in energy for a chemical reaction can be endothermic or exothermic. LO 11 The student is able to interpret observations regarding macroscopic energy changes associated with a reaction or process to generate a relevant symbolic and/or graphical representation of the energy changes. Electrochemistry - Electrochemistry shows the interconversion between chemical and electrical energy in galvanic and electrolytic cells. LO 12 The student can make qualitative or quantitative predictions about galvanic or electrolytic reactions based on half-cell reactions and potentials and/or Faraday s laws. LO 13 The student can analyze data regarding galvanic or electrolytic cells to identify properties of the underlying redox reactions. Exclusion Statement: Labeling an electrode as positive or negative is beyond the scope of this course and the AP Exam. Exclusion Statement: The Nernst equation is beyond the scope of this course and the AP Exam.

2 College Board References and Resources: I have noted right the formulas/sections of the table of equations and constants that are pertinent to the Big Idea we are currently studying. Meaning that any mathematical procedures, constants, formulas, etc. not noted here, and used for the Big Idea, cannot be referenced during your AP examination. Instructions: Write the chemical symbol (including state to the best of your ability) for the following formulas. If the substance exists mainly as separate ions in solution, write the separate ions. Remember, on the AP exam you may only use the periodic table. No solubility chart. 1. solid aluminum oxide 2. a solution of sodium hydroxide 3. solid calcium oxide 4. sulfur trioxide gas molar sulfuric acid molar potassium hydroxide 7. calcium metal 8. nitrogen gas 9. solid copper(ii) sulfide 10. oxygen gas 11. concentrated hydrochloric acid 12. powdered manganese dioxide 13. concentrated solution of ammonia 14. a solution of zinc iodide 15. a solution of copper(ii) sulfate 16. a solution of barium hydroxide 17. a solution of magnesium nitrate 18. solid lithium hydride 19. water 20. a solution of ammonia 21. a solution of hydrofluoric acid 22. a piece of aluminum metal 23. a solution of silver nitrate 24. a solution of potassium iodide 25. solid potassium oxide 26. an excess of nitric acid solution 27. copper(ii) sulfate 28. carbon dioxide gas 29. a suspension of calcium carbonate 30. a strip of copper 31. dilute nitric acid 32. potassium permanganate solution 33. an acidic solution of hydrogen peroxide 34. solid manganese (II) sulfide 35. chlorine gas 36. hot iron filings 37. solid magnesium nitride 38. sulfur dioxide gas 39. a suspension of silver chloride 40. a solution of tri-potassium phosphate 41. a solution of zinc nitrate 42. sodium cyanide solution 43. a solution of manganese(ii) sulfate 44. a solution of ammonium sulfide 45. phosphorus(v) oxide powder 46. solid ammonium carbonate 47. solid potassium permanganate 48. a small piece of sodium metal 49. a solution of potassium dichromate 50. an acidified solution of iron(ii) chloride 51. Methane 52. solid barium oxide 53. a solution of iron(ii) nitrate 54. solid calcium phosphate 55. hydrogen sulfide gas 56. a solution of mercury(ii) chloride 57. solid calcium hydride 58. a bar of zinc metal

3 Instructions: Write net-ionic equations for these double replacement reactions. I should not have to tell you that they need to be balanced and states indicated. Remember that on the AP exam you may only use the periodic table. No solubility chart. 59. Aqueous solutions of zinc sulfate and sodium phosphate are mixed. 60. Hydrofluoric acid is combined with a solution of lead(ii) nitrate. 61. Solid calcium sulfide is sprinkled into dilute hydrochloric acid. 62. An aqueous solution of lead(ii) acetate reacts with hydrochloric acid. 63. Solid sodium carbonate is stirred into hydrobromic acid. 64. Nitric acid is reacted with an aqueous solution of calcium acetate. 65. Hydrochloric acid is poured over powdered potassium carbonate. 66. An aqueous solution of cadmium chloride is reacted with an aqueous solution of potassium phosphate. 67. A solution of hydrofluoric acid is poured over barium carbonate crystals. 68. Hydroiodic acid is poured over potassium sulfite. 69. An aqueous solution of barium hydroxide is reacted with an aqueous solution of iron(iii) sulfate. 70. A solution of sodium hydroxide is poured into a solution of magnesium chloride. 71. Aqueous lead(ii) nitrate is combined with potassium iodide Instructions: Write net-ionic equations for these acid-base reactions. Remember 73. A 0.40 M solution of acetic acid is reacted with a 0.20 M lithium hydroxide solution 74. A solution of nitric acid is combined with a suspension of magnesium hydroxide. 75. A solution of sulfuric acid is poured over copper(i) hydroxide crystals 76. A solution of sulfuric acid is added to a solution of barium hydroxide until the same number of moles of each compound has been added. 77. Hydrogen sulfide gas is bubbled through a solution of potassium hydroxide. 78. A solution of sodium hydroxide is added to a solution of sodium dihydrogen phosphate until the same number of moles of each compound has been added. 79. Equal volumes of 0.1-molar sulfuric acid and 0.1-molar potassium hydroxide are mixed. 80. Excess potassium hydroxide solution is added to a solution of potassium hydrogen phosphate 81. Excess hydrochloric acid solution is added to a solution of sodium dihydrogen phosphate M ammonia reacts with 0.2 M hydrobromic acid. 83. A solution of ammonia is added to a dilute solution of acetic acid. 84. A solution of ammonia and hydrofluoric acid are combined. Instructions: Write net-ionic equations for these single replacement reactions. Remember that on the AP exam you may only use the periodic table. No solubility chart. 85. A strip of magnesium is added to a solution of silver nitrate 86. Aluminum metal is dropped into an solution of zinc chloride 87. Solid silver is dropped into an solution of gold(ii) nitrate 88. Aluminum foil is dropped into a solution of nitric acid. 89. Solid barium is added to chlorous acid 90. Potassium metal is dropped into water 91. Chromium(II) nitrate solution is combined with iron(iii) nitrate solution. 92. Iron(II) nitrate solution is mixed with cobalt(iii) chloride solution 93. Liquid bromine is added to an aqueous sodium iodide solution 94. Hydrogen gas is passed over hot copper(ii) oxide. 95. Small chunks of solid sodium is added to water. 96. Magnesium metal is added to a dilute solution of nitric acid. 97. Chlorine gas is bubbled into a solution of potassium iodide. Instructions: Write net-ionic equations for these decomposition reactions. Remember that on the AP exam you may only use the periodic table. No solubility chart. 98. A 27 % hydrogen peroxide solution is catalytically decomposed 99. Solid potassium chlorate is heated in the presence of a catalyst 100. Solid magnesium carbonate is heated Solid ammonium carbonate is heated Calcium sulfite is heated in a vacuum. Instructions: Write net-ionic equations for these synthesis reactions. Remember 104. Bromine vapor is passed over cadmium powder Chlorine gas is reacted with aluminum Strontium oxide is reacted with sulfur trioxide Hot bromine vapor is reacted with aluminum foil 108. The gases boron trifluoride and ammonia are mixed A mixture of solid calcium oxide and solid tetraphosphorus decoxide is heated Calcium metal is heated strongly in nitrogen gas Magnesium ribbon is burned in oxygen Powdered magnesium oxide is added to a container of carbon dioxide gas. Instructions: Write net-ionic equations for these combustion reactions. Remember 113. Pentane is burned in oxygen 114. Propene is burned in air 115. Butene is burned in air Ethanol is burned in air Pentanol is combusted in air 118. Lithium metal is burned in air Gaseous diborane, B2H6, is burned in excess oxygen Carbon disulfide vapor is burned in excess oxygen Solid copper(ii) sulfide is heated strongly in oxygen gas Gaseous silane (silicon tetrahydride) is burned in oxygen 123. Propanol is burning in oxygen Zinc sulfide is heated in an excess of oxygen. Instructions: Write net-ionic equations for these anhydride reactions. Remember 125. Sodium oxide powder is sprinkled in water Dinitrogen pentoxide is bubbled through water Dinitrogen trioxide gas is bubbled into water Diphosphorus pentoxide is bubbled through water Rubidium hydride is placed into distilled water 130. Solid copper(ii) hydride is dropped into water 131. Solid strontium hydride reacts with hydrochloric acid 132. Powdered barium oxide is stirred into aqueous hydrochloric acid Solid potassium oxide is added to water Sulfur dioxide is bubbled through a dilute solution of sodium hydroxide Gaseous sulfur dioxide bubbled through an aqueous solution of calcium hydroxide M hydrochloric acid is poured over calcium oxide powder Ammonia and carbon dioxide gases are bubbled into cold distilled water Solid phosphorus pentachloride is added to excess water.

4 Instructions: Write net-ionic equations for these complex-ion reactions. Remember 139. Concentrated ammonia is is added to a solution of zinc chloride 140. A solution of aluminum chloride is reacted with excess, concentrated sodium hydroxide 141. An excess of nitric acid solution is added to a solution of tetraaminecopper(ii) sulfate A solution of diamminesilver(i) chloride is treated with dilute hydrochloric Excess concentrated ammonia solution is added to a suspension of silver chloride A suspension of zinc hydroxide is treated with concentrated sodium hydroxide solution Excess sodium cyanide is added to a solution of silver nitrate A concentrated solution of ammonia is added to a suspension of zinc hydroxide An aqueous solution of hydrochloric acid is added to a solution containing the tetraaminecadmium(ii) ion. Instructions: Determine the oxidation number of each element in the following ions or compounds: 148. BrO CaH C2O H4SiO F SO SF N2O H2AsO PCl UO XeO4 2 Instructions: Using whichever method works best for you, balance the following oxidation/reduction reactions Balance the following net ionic Redox reactions in acidic solution: a. MnO4 + Cl Mn 2+ + Cl2 b. Cu + SO4 2 SO2(g) + Cu 2+ c. Pb + PbO2 + SO4 2 PbSO4 d. MnO2 + PbO2 MnO4 + Pb 2+ e. Br2 + SO2 Br + SO4 2 f. P4(s) + NO3 H3PO4(aq) + NO(g) g. Cr2O7 2 + CH3OH Cr 3+ + CH2O h. NO3 + Fe Fe 2+ + NH4 + i. CuS + NO3 Cu 2+ + S(s) + NO(g) j. H2S + Fe 3+ FeS + S(s) 161. Balance the following net ionic Redox reactions in basic solution: a. MnO4 + BrO3 MnO2 + BrO4 b. Bi(OH)3 + SnO2 2 Bi(s) + SnO3 2 c. MnO4 + C2O4 2 MnO2 + CO3 2 d. Pb 2+ + OCl PbO2 + Cl e. Bi(OH)2 + Sn(OH)3 Bi + Sn(OH)6 2 f. Zn + NO3 NH3(aq) + Zn(OH)4 2 g. Al + OH Al(OH)4 + H2 h. H2O2 + I IO3 + H2O i. Fe(CN)6 3 + Re Fe(CN)6 4 + ReO4 j. SO3 2 + Cl2 SO4 2 + Cl Instructions: Answer the following questions related to electrochemistry. Standard Reduction Potential E (volts) Cl2(g) + 2e - 2Cl - (aq) O2(g) + 4H + (aq) + 4e - 2H2O(l) Ag + (aq) + e - Ag(s) I2(s) + 2e - 2I - (aq) Cu 2+ (aq) + 2e - Cu(s) SO4 2- (aq) + 4 H + (aq) + 2e - SO2(g) + 2 H2O H + (aq) + 2 e - H2(g) (reference electrode) H2O(l) + 2e - H2(g) + 2OH - (aq) Na + (aq) + e - Na(s) K + (aq) + e - K(s) All of the equations in the chart right are written as (oxidations/reductions) The chemicals at the upper left (Cl2 and O2) are the most likely to be (oxidized/reduced) and therefore the best (oxidizing agents/reducing agents) The chemicals at the lower right (Na and K) are the most likely to be (oxidized/reduced) and therefore the best (oxidizing agents/reducing agents) In an electrolytic cell, the (-) electrode is negative because it has (too many/too few) electrons. Chemicals that come into contact with the (-) electrode will (gain/lose) electrons and be (oxidized/reduced). The (-) electrode in electrolysis is called the (cathode/anode) Write the change that water goes through at the (+) electrode: 124. In an electrochemical cell, the (+) electrode is positive because is has (too many/too few) electrons. Chemicals that come into contact with the (+) electrode will (gain/lose) electrons and be (oxidized/reduced). The (+) electrode in electrolysis is called the (cathode/anode) Write the change that water goes through at the (-) electrode: 126. Add these two reactions (5. and 7.) together (make certain the electrons cancel) and write the overall reaction for the electrolysis of water We will perform this electrolysis using an aqueous solution of sodium sulfate. Both the Na + and H2O will be near the (-) electrode. Which chemical is more likely to be reduced? 128. Both the SO4 and H2O will be near the (+) electrode. Which chemical will be oxidized? 129. In the electrolysis of KI(aq) Both the K + and H2O will be near the (-) electrode. Which chemical is more likely to be reduced? Both the I and H2O will be near the (+) electrode. Which chemical is more likely to be oxidized? Write the reactions at each electrode and the overall reaction: Cathode: Anode: Overall:

5 130. In the electrolysis of CuSO4(aq) Both the Cu 2+ and H2O will be near the (-) electrode. Which chemical will be reduced? Both the SO4 and H2O will be near the (+) electrode. Which chemical will be oxidized? Write the reactions at each electrode and the overall reaction: Recall that 1 amp sec = 1 Coulomb and 96,500 Coulombs = 1 mole e s (Faraday s constant). If a cell is run for 200. seconds with a current of amps, how many grams of Ag will be deposited? Cathode: Anode: Overall: 131. Silver plating occurs when electrolysis of a Ag2SO4 solution is used because silver metal is formed at the (cathode/anode). This is the ( ) ( + / - )electrode. The reaction at this electrode is: A current of 10.0 amperes flows for 2.00 hours through an electrolytic cell containing a molten salt of metal X. This results in the decomposition of mole of metal X at the cathode. The oxidation state of X in the molten salt is (X +, X 2+, X 3+, X 4+ ) 133. Solutions of Ag +, Cu 2+, Fe 3+ and Ti 4+ are electrolyzed with a constant current until 0.10 mol of metal is deposited. Which will require the greatest length of time? Instructions: Clearly circle the response that best satisfies the prompt giving only one answer to each question. DO NOT USE A CALCULATOR. The College Board does not allow the use of a calculator for Section I (the MC part of the exam, so it is time to start practicing without one. You may use ONLY a periodic table and the equations. While you should practice working as fast as possible, it is more important at this point in the course that you practice without a calculator, even if it slows you down. Look for the easy math common factors and rough estimation do not do long division to try to get exact values. Remember it is a MC test, use the answers to guide you Consider this reaction: 3Au(s) + 8H + + 2NO3 3Au 2+ (aq) + 2NO(g) + 4H2O(Ll Mark each of the following statements (about the reaction above) as true or false. T F Au(s) is reduced during the reaction. T F The oxidation state of nitrogen changes from +6 to +2. T F Hydrogen ions are oxidized when they form H2O(l). T F 8 electrons are lost and 8 electrons are gained during this reaction Which substance contains the element that is reduced in the following reaction: Cr2O S2O H + 2Cr S4O H2O a. Cr2O7 2 b. S2O3 2 c. H + d. Cr 3+ e. S4O What is the oxidation number of manganese in the KMnO4: a. +1 b. +2 c. +5 d. +4 e How many electrons appear in the following halfreaction when it is balanced. Cr2O H + 2Cr H2O a. 1 b. 2 c. 3 d. 4 e The more the value of E red, the greater the driving force for the substance s own oxidation. a. positive b. negative c. exothermic d. endothermic e. extensive 139. How many electrons appear in the following halfreaction when it is balanced. S4O6 2 2S2O3 2 a. 6 b. 2 c. 4 d. 1 e The half-reaction occurring at the anode in the balanced reaction shown below is 3MnO4 + 24H + + 5Fe 3Mn Fe H2O a. MnO4 + 8H + + 5e - Mn H2O b. MnO4 + 5e Mn 2+ c. Fe Fe + + 3e d. Fe + 3e Fe 3+ e. MnO4 + 8H + Mn H2O + 5e 141. The oxidation half reaction occurring in the standard hydrogen electrode is a. H2 2H + + 2e b. 2H + + 2OH H2O c. O2 + 4H + + 4e 2H2O d. 2H + + 2e H2 e. 2H + + Cl2 2HCl H + + 5H2O2 + 2MnO4 5O2 + 2Mn H2O According to the balanced equation above, how many moles of the permanganate ion are required to react completely with 25.0 ml of M hydrogen peroxide? a mol b mol c mol d mol e mol 143. A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl2. What is the minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all of the Cl as AgCl(s)? a mol b mol c mol d mol e mol 144. When 400. milliliter of 0.10-molar sodium chloride is added to 200. milliliters of 0.10-molar aluminum chloride the number of moles of Pb 2+ that must be added to precipitate out all of the Cl would be a moles b moles c moles d moles e moles 145. A 20.0-milliliter sample of molar K2CO3 solution is added to 30.0 milliliters of molar Ba(NO3)2 solution. Barium carbonate precipitates. The concentration of barium ion, Ba 2+, in solution after reaction is a M b M c M d M e M 146. How many moles of solid Ba(NO3)2 should be added to 300. milliliters of 0.20-molar Fe(NO3)3 to increase the concentration of the NO3 ion to 1.0- molar? (Assume that the volume of the solution remains constant.) a mole b mole c mole d mole e mole 147. Given the following half reactions: Sn e Sn 2+ Fe e Fe 2+ Eº = 0.15 V Eº = 0.77 V Determine the standard cell potential (E cell) for the voltaic cell based on the reaction Sn Fe 3+ 2Fe 2+ + Sn 4+ a V b V c V d V e V 148. When this reaction is balanced, the coefficient on the Sn 2+ is. Sn 4+ + Cr Cr 3+ + Sn 2+ a. 1 b. 2 c. 3 d. 4 e Identify the redox reactions below. I. K2CrO4 + BaCl2 BaCrO4 + 2KCl II. Pb Br PbBr2 III. Cu + S CuS a. III only b. II only c. I only d. I and III e. II and III

6 150. Which one of the following reactions is a redox reaction? a. NaOH + HCl NaCl + H2O b. Pb Cl PbCl2 c. AgNO3 + HCl HNO3 + AgCl d. 2HC2H3O2 + Ca(OH)2 2H2O+ Ca(C2H3O2)2 e. None of the above are redox reactions Which substance is the reducing agent (the substance that caused reduction i.e the substance that is oxidized) in the following reaction? Fe2S3 + 12HNO3 2Fe(NO3)3 + 3S + 6NO2 + 6H2O a. HNO3 b. S c. NO2 d. Fe2S3 e. H2O 152. What is the coefficient of the bromide ion when the following equation is balanced? MnO4 + Br Mn 2+ + Br2 (acidic solution) a. 1 b. 2 c. 4 d. 5 e Consider the half reactions shown below to answer the following question: Fe e Fe 2+ Eº = 0.77 V Sn e Sn 2+ Eº = 0.15 V Fe e Fe Eº = 0.44 V Cr e Cr Eº = 0.74 V Which of the following reactions will occur spontaneously as written? a. Sn 4+ + Fe 3+ Sn 2+ + Fe 2+ b. 3Fe + 2Cr 3+ 2Cr + 3Fe 2+ c. Sn 2+ + Fe Sn 4+ + Fe 2+ d. 3Sn Cr 2Cr Sn 2+ e. 3Fe 2+ Fe + 2Fe In the oxidation-reduction reaction Sn Fe 2+ 2 Fe 3+ + Sn 2+ a. Sn 4+ is reduced and Fe 2+ is oxidized. b. Sn 4+ is oxidized and Fe 2+ is reduced. c. Sn 4+ is oxidized and Fe 3+ is reduced. d. Fe 3+ is reduced and Sn 2+ is oxidized. e. Fe 2+ is oxidized and Fe 3+ is reduced Given the standard reduction potentials Cu e Cu E = Volt Al e Al E = 1.66 Volt Calculate the standard voltage for the reaction 2Al(s) + 3Cu 2+ 2Al Cu(s) a Volts b Volts c Volts d Volts e Volts 156. Given the standard electrode (reduction) potentials: Cd e - Cd E = 0.40 v Ag + + e- Ag E = v What would be the E for a cadmium-silver cell? a. 0.0 V b. 0.4 V c. 0.5 V d. 1.2 V e. 2.0 V 157. In the reaction SO2 + 2 H2S 3 S + 2 H2O a. sulfur is oxidized and hydrogen is reduced b. sulfur is reduced and there is no oxidation c. sulfur is oxidized and there is no reduction d. sulfur is reduced and hydrogen is oxidized e. sulfur is both reduced and oxidized 158. Which ion, in solution, can be oxidized by chemical means but also can be reduced by a different chemical reaction? a. Fe 2+ b. F c. CO3 2 d. NO3 e. Al Zinc reacts with dilute acid to produce H2 and Zn 2+ but silver does not liberate hydrogen gas from an acid. This information enables one to predict that which of the following reactions has a positive emf? a. H2(g) + Zn 2+ 2H + + Zn b. 2 Ag + Zn 2+ 2Ag + + Zn c. 2 Ag+ + Zn 2Ag + Zn 2+ d. 2 Ag + 2H + H2(g) + 2 Ag + e. 2 Ag+ + Zn 2+ 2Ag + Zn 160. If solid nickel metal were added to separate aqueous solutions each containing 1M concentrations of Ag +, Cd 2+, Sn 2+, and Ni 2+ ions, which metals would plate out, based on the given standard reaction potentials? Standard Reduction Potentials Ag + /Ag V Sn 2+ /Sn V Ni 2+ /Ni V Cd 2+ /Cd V I. Ag II. Sn III. Ni IV. Cd a. I only b. IV only c. I and II only d. III and IV only e. I, II, and III 161. What voltage will be produced by the electrochemical cell? a V b V c V d V e V 162. Which statements below are true about the voltaic cell in the previous problem. I. electrons move right to left through the external circuit II. the lead electrode will gain mass III. the porous barrier maintains the charge balance in both half-cells a. I only b. II only c. I and II only d. II and III only e. I, II, and III 163. Which of these ions is the best at being oxidized given the following standard reduction potentials? Fe e Fe 2+ Eº = V Cu e Cu + Eº = V a. Fe 3+ b. Fe 2+ c. Cu 2+ d. Cu + e. None of them can act as reducing agents 164. Use these reduction potentials to determine which one of the reactions below is spontaneous. Standard Reduction Potentials Ag + /Ag V Pb 2+ /Pb V V 2+ /V V I. V Ag V + 2 Ag + II. V + Pb 2+ V 2+ + Pb III. 2Ag + + Pb 2+ 2Ag + Pb a. II only b. III only c. I and II only d. II and III only e. I, II, and III 165. For this reaction, E cell = 0.79 V. 6I + Cr2O H + 3I2 + 2Cr H2O Given that the standard reduction potential for Cr2O7 2 2Cr 3+ is 1.33 V, what is E red for I2(aq)? a V b V c V d V e V 166. All of the following may get oxidized EXCEPT a. NO2 b. Na c. Fe 2+ d. MnO4 e. Br 167. For this reaction, KClO4 + H2O2 KClO3 + O2 + H2O Choose the true statement from the following list. a. The Cl oxidation state is reduced form +8 to +6 b. This is not an oxidation-reduction reaction. c. H2O2 caused reduction. d. Hydrogen is reduced from +2 to +1 e. Potassium is reduced during the reaction

7 Practice Free Response Questions: We will be doing these together as a class to prepare you for your take home assessment and severity of grading that will be associated with these. You may attempt on your own if you like but these are not a part the problem set. ALL WORK must be done on a blank white sheet of paper that has your name on it. SFR # H2(g) + O2(g) 2 H2O(l) In a hydrogen-oxygen fuel cell, energy is produced by the overall reaction represented above. (a) When the fuel cell operates at 25 C and 1.00 atm for 78.0 minutes, mol of O2(g) is consumed. Calculate the volume of H2(g) consumed during the same time period. Express your answer in liters measured at 25 C and 1.00 atm. (b) Given that the fuel cell reaction takes place in an acidic medium, (i) write the two half reactions that occur as the cell operates, (ii) identify the half reaction that takes place at the cathode, and (iii) determine the value of the standard potential, E, of the cell. (c) Calculate the charge, in coulombs, that passes through the cell during the 78.0 minutes of operation as described in part (a). A student is given a standard galvanic cell, represented above, that has a Cu electrode and a Sn electrode. As current flows through the cell, the student determines that the Cu electrode increases in mass and the Sn electrode decreases in mass. (a) Identify the electrode at which oxidation is occurring. Explain your reasoning based on the student s observations. (b) As the mass of the Sn electrode decreases, where does the mass go? (c) In the expanded view of the center portion of the salt bridge shown in the diagram below, draw and label a particle view of what occurs in the salt bridge as the cell begins to operate. Omit solvent molecules and use arrows to show the movement of particles. SFR # Answer the following questions that relate to electrochemical reactions. (a) Under standard conditions at 25 C, Zn(s) reacts with Co 2+ (aq) to produce Co(s). (i) Write the balanced equation for the oxidation half reaction. (ii) Write the balanced net-ionic equation for the overall reaction. (iii) Calculate the standard potential, E, for the overall reaction at 25 C. (b)at 25 C, H2O2 decomposes according to the following equation. 2 H2O2(aq) 2 H2O(l) + O2(g) E = 0.55 V (i) Determine the value of the standard free energy change, ΔG, for the reaction at 25 C. (iii) The standard reduction potential, E, for the half reaction: O2 (g) + 4 H + (aq) + 4 e 2 H2O (l) has a value of 1.23 V. Using this information in addition to the information given above, determine the value of the standard reduction potential, E, for the half reaction below. O2 (g) + 2 H + (aq) + 2 e H2O2(aq) (c) In an electrolytic cell, Cu(s) is produced by the electrolysis of CuSO4(aq). Calculate the maximum mass of Cu(s) that can be deposited by a direct current of 100. amperes passed through 5.00 L of 2.00 M CuSO4(aq) for a period of 1.00 hour. (d) A nonstandard cell is made by replacing the 1.0 M solutions of Cu(NO3)2 and Sn(NO3)2 in the standard cell with 0.50 M solutions of Cu(NO3)2 and Sn(NO3)2. The volumes of solutions in the nonstandard cell are identical to those in the standard cell. (i) Is the cell potential of the nonstandard cell greater than, less than, or equal to the cell potential of the standard cell? Justify your answer. (ii) Both the standard and nonstandard cells can be used to power an electronic device. Would the nonstandard cell power the device for the same time, a longer time, or a shorter time as compared with the standard cell? Justify your answer. (e) In another experiment, the student places a new Sn electrode into a fresh solution of 1.0 M Cu(NO3)2. LFR # (i) Using information from the table above, write a net-ionic equation for the reaction between the Sn electrode and the Cu(NO3)2 solution that would be thermodynamically favorable. Justify that the reaction is thermodynamically favorable. (ii) Calculate the value of ΔG 0 for the reaction. Include units with your answer.