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Chemistry 7: Electrochemistry A. B. C. Name __________________________ Balancing Redox Equations (20.2) assign oxidation to each atom determine oxidized atom (oxidation # increases) and reduced atom (oxidation # decreases) split into oxidation and reduction half-reactions eliminate spectator ions (ion that doesn't contain atom that changes oxidation number—often cation) balance each half reaction o balance atoms except O and H o balance O, by adding H2O o balance H, by adding H+ o balance charge, by adding e multiply half-reactions to equalize electrons add half-reactions together simplify by reducing H2O and H+ and/or coefficients this process assumes reaction takes place in acid (H+), if in base, add an OH- for each H+ in the final equation (combine H+ and OH- to make water) Standard Reduction Potentials Chart (20.4 to 20.6) 1. reduction half reactions a. listed from “greatest” electron affinity to “least” b. 2 H+ + 2 e- H2: Eored = Eoox = 0 V c. Eo measured in volts, 1 V = 1 J/C 1. "o": standard conditions (25oC, 1 atm, 1 M) 2. not proportional to amount of chemical d. oxidation is reverse (Eoox = -Eored) 2. Eo = Eored + Eoox a. Eo > 0 is a spontaneous reaction (reduction listed above oxidation on chart) b. voltage under nonstandard conditions 1. Nernst equation: E = Eo – (RTo/nF)lnQ R (8.31), To (298) and F (96,500) are constant (E = Eo – (0.0257/n)lnQ) 2. Q (quotient) = Product/Reactants a. partial pressure (atm) gases, concentration (M) of ions b. solids and liquids excluded Voltaic (Galvanic) Cell (20.3) spontaneous redox reaction generates voltage electrons flow through wires from oxidation cell to reduction cell anode (–) Voltage > 0 salt bridge cathode (+) porous membrane site of oxidation site of reduction 1. oxidation half cell (– anode) a. reducing agent gives up electrons to external circuit (wires) b. anions flow toward anode through salt bridge/ porous membrane to maintain electrical neutrality 2. reduction half cell (+ cathode) a. oxidizing agent attract electrons from wire b. cations flow toward cathode through salt bridge 3. predict how change affects voltage a. reactant: [ions] or Pgases E (voltage) b. over time: reactant & product E c. size of electrode and chamber: no change d. remove salt bridge: E = 0 D. Electrolytic Cell (20.9) battery forces non-spontaneous redox reaction by pulling electrons from reducing agent and sending to oxidizing agent anode (+) + Battery – cathode (–) site of oxidation site of reduction 1. Eo < 0 (battery makes up for deficit) 2. oxidation at + anode, reduction at – cathode 3. electrolysis in water solutions (inert electrodes) a. cathode reduction: H2O or cation (greater Eored) 1. columns 1, 2 or Al3+: 2 H2O + 2 e- H2 + 2 OH2. acid (H+): 2 H+ + 2 e- H2 3. otherwise: Mx+ + X e- M b. anode oxidation: anion or H2O 1. Cl-, Br-, I-: 2 X- X2 + 2 e2. base (OH-): 4 OH- O2 + 2 H2O + 4 e3. otherwise: 2 H2O O2 + 4 H+ + 4 e4. electroplating (transition metal cations coat cathode) a. current, I, measured in amperes (amps—A) 1 A = 1 C/s (coulomb/second) b. mass plated given current, I, and time, t (t) s x (I) C x mol e- x mol Mx+ x (MM) g = __ g 1 s 96,500 C X mol e- 1 mol Mx+ c. time for plating—calculate right to left (2) Given that pink indicates the presence of OH-, write the half-reaction that occurs at the cathode. Experiments 1. Voltaic Cell Lab—Measure the voltage of six voltaic cells, determine the reduction half-cell potential for each metal and compare these results to the standard values. Polish the metal electrodes (1-Ag, 2-Cu, 3-Pb, and 4-Zn). Half fill each quadrant of the Petri dish with the 0.1 M nitrate salt of the corresponding metal. Drape the KNO3-soaked ring so that it is submerged in each quadrant. Measure the voltage of each system (reverse the electrodes when V < 0). Record the metal that is connected to the black electrode (oxidation) and the voltage (Elab) for the overall reaction. a. Complete the chart for each voltaic cell. Ag and Cu Ag and Pb Ag and Zn Elab Elab Elab Black Black Black (1) What color change occurs at the anode? (2) Given that amber indicates the presence of I2, write the half-reaction that occurs at the anode. c. Are these reactions consistent with the predicted reactions based on KI as the electrolyte? Explain Ag + Cu Part 2: Measure the change in mass of the zinc electrode and the volume of hydrogen gas produced during electrolysis, calculate the molar mass of zinc and compare the value to the periodic table. Fill a 150 mL beaker ¾ full with conducting solution. Fill the 25 mL volumetric flask with the conducting solution from the beaker. Place rubber dam over the mouth of the flask, invert it and place it in the beaker (mouth down). Scrap off the rubber dam and slip the J-shaped electrode inside the flask without getting any air in. Polish the zinc and mass it (m1). Connect the wire lead attached to the black side of the power supply to the J-electrode (cathode). Connect the wire lead attached to the red side of the power supply to the zinc (anode). Turn on the power supply. Hydrogen gas should be bubbling from the J-electrode. Record the start time (t1). When the water level in the flask is near the 25 mL line, carefully raise the flask so that the level of the conducing fluid solution inside the flask is the same as the level in the beaker. When the water line in the flask is at 25 mL, turn the power supply off and disconnect the wires from the zinc and J-electrode. Record the stop time (t2). Return the conducting solution to the stock bottle. Scrub the zinc with a scrub pad under running water. Thoroughly dry the zinc and then mass it (m2). Record the temperature (T), lab pressure (Plab) and water vapor pressure (PH2O). d. Record the data. m1 t1 m2 t2 V T Plab PH2O 25.0 mL e. Complete and balance the redox reaction between zinc metal and water. Reduction Ag + Pb Oxidation Ag + Zn d. Compare the voltages of the remaining systems with the table values. (There is no need to correct for 0.1 M) Eolab Eotable = Eored + Eoox Cell % Overall f. Calculate the molar mass of Zn by completing the chart. Cu + Pb Volume in L Cu and Pb Elab Black b. Ag + Cu Cu and Zn Elab Black Pb and Zn Elab Black Determine the standard voltage (Eolab) for each cell. Half Reaction Reduction Oxidation Overall Eolab Ag + Pb Reduction Oxidation Overall Eolab Ag + Zn Reduction Oxidation Overall Eolab c. Determine the lab values for Eored for Cu, Pb and Zn given Eored for Ag is 0.80. Eored = 0.80 – Eolab Eotable Cell % Cu + Zn 2. b. Pb + Zn Electrolysis Lab—Part 1: Observe the electrolysis of KI(aq). Fill a weighing boat with 0.5 M KI. Add 3 drops indicator. With the power supply OFF, attach a piece of graphite to each wire lead and then plug in the power supply into the electric outlet. Dip the two pieces of graphite into opposite sides of the weighing boat for 10 seconds. Observe changes at the cathode (black electrode) and anode (red electrode), record the observations. Turn off the power supply and remove the graphite. a. (1) What color change occurs at the cathode? Mass of Zinc lost Pressure in atm Temperature in K Moles of H2 Molar Mass of Zinc % g. Why is it necessary to equalize the water level inside the volumetric flask with the water level in the beaker? h. Calculate the average current during the electrolysis by completing the chart. 8. Total charge Q in C Total time T in s The mass percent of H2O2 in a hydrogen peroxide solution is determined by titration with an acidified solution of KMnO4. a. Balance the redox reaction: MnO4- + H2O2 + H+ Mn2+ + O2 + H2O. oxidation reduction Current I (I = Q/t) i. overall b. How many moles of MnO4- are needed to react all the H2O2 if 2.647 mL of 0.0200 M KMnO4 is added? Calculate the approximate voltage needed to perform the electrolysis. (Assume standard conditions) Practice Problems 1. 2. A. Oxidation-Reduction Reactions Balance the redox equation in acid. MnO4- + NO2- + H+ Mn2+ + NO3- + H2O Balanced the redox equation in base Cr(OH)3(s) + ClO- Cl- + CrO42- 3. Balance the redox equation in acid: Zn(s) + NO3- Zn2+ + NH4+. 4. Balanced the redox equation in base: NO2- + Al(s) NH3(aq) + Al(OH)4- 5. Balance the disproportionation reaction (same element undergoes oxidation and reduction): Cl2(g) Cl- + ClO-. in acid How many moles of H2O2 reacted with the MnO4-? d. How many grams of H2O2 reacted with the MnO4-? e. What is the % H2O2 if 0.150 g of solution is used? B. Standard Reduction Potentials Chart Consider the Standard Reduction Potential Chart. a. What are the standard conditions? Highlight the correct option. b. When comparing two reactants (on the left side of the chart), the reactant that is listed (higher/lower) on the chart is the stronger oxidizing agent. c. When comparing two products (on the right side of the chart), the product that is listed (higher/lower) on the chart is the stronger reducing agent. d. The strongest reducing agents are found in column (1/17) on the periodic table. e. The strongest oxidizing agents are found in column (1/17) on the periodic table. f. Reactive (metals/nonmetals) tend to lose electrons and act as good (oxidizing/reducing) agents. g. When combining two half reactions, the left side agent listed higher on the chart will (give/receive) electrons from the right side agent listed lower on the chart. h. Write the equations for each half-reaction and overall reaction. Calculate Eo. Eo Balanced reaction Zn Zn2+ in base 6. Balance the reaction: P4(s) H2PO2- + PH3(g) Br2 BrOverall H2 H+ in acid Ag+ Ag in base 7. Write net ionic equations for the following redox reactions. a. Solid iron + iron(III) sulfate b. 9. c. Potassium dichromate + acidified hydrogen peroxide. Overall i. j. A spontaneous reaction has a (positive/negative) Eotot. Is the redox reaction, Cu(s) + Cl2(g) CuCl2(s), spontaneous? Support your answer with calculations. 10. Indicate where on the Standard Reduction Potentials Chart you would find: a. The chemical species that is the easiest to oxidize. b. The chemical species that is the easiest to reduce. 11. Using standard reduction potentials, calculate the standard voltage for each of the following reactions. Cl2(g) + 2 I - 2 Cl- + I 2(s) Ni(s) + 2 Fe3+ Ni2+ + Fe2+ Fe(s) + 2 Fe3+ 2 Fe2+ 2 Al3+ + 3 Ca(s) 2 Al(s) + 3 Ca2+ 12. Using standard reduction potentials, highlight the stronger reducing agents for each of the following pairs. Fe(s) or Mg(s) Ca(s) or Al(s) H2(g) or H2S(g) Sn2+ or Fe2+ 13. The unbalanced reduction half-reactions that operate in a car battery follow. PbSO4(s) Pb(s) + SO42Eored = -0.356 V 2PbO2(s) + SO4 PbSO4(s) Eored = 1.687 V a. Write balanced half reactions and the overall reaction. Label which reaction occurs at the anode and cathode. b. Calculate Eotot for the cell. c. Write the Nernst equation for the overall reaction. Calculate the voltage when [H+] = 2.0 x 10-4 M and [SO42-] = 0.50 M. e. e- C. Voltaic (Galvanic) Cell 14. Answer the following questions based on the diagram. e- If you made a voltaic cell out of this, what half-reaction would be occurring at the cathode, and what halfreaction would be occurring at the anode? c. What would happen to the voltage of the battery if the concentration of [H+] increased justify your answer)? Voltage = oxidation half reaction reduction half reaction overall reaction 16. For the generic reaction: A + B A- + B+, for which Eo is a positive number, answer the following questions: a. What is being oxidized, and what is being reduced? b. d. 15. Complete the voltaic cell drawing using the half-reactions. Ni2+ + 2 e- Ni(s) Eored = -0.25 V Fe2+ + 2 e- Fe(s) Eored = -0.44 V Label the anode metal, cathode metal, site of oxidation, site of reduction, cation flow, anion flow, voltage, oxidation half-reaction, reduction half-reaction, and overall reaction. (You will need to figure out which reaction occurs at the anode, which at the cathode, and what the voltmeter would read under standard conditions. Assume each metal electrode is immersed in 1 M nitrate salt of the metal.) Which half-reaction from (b) is higher on the Standard Reduction Potential Chart? 17. A voltaic cell consists of a strip of aluminum in a solution of Al(NO3)3 in one beaker, and in the other beaker a strip of nickel in a solution of Ni(NO3)2. The overall reaction is: 2 Al(s) + 3 Ni2+ 2 Al3+ + 3 Ni(s) a. What is being oxidized, and what is being reduced? b. Write the half-reactions that occur in the beakers. (Indicate which reaction takes place at the anode and which takes place at the cathode.) c. Indicate the signs of the two electrodes. a. which Zn Zn2+ Cu Cu2+ species is reduced? species is oxidized? species gives up its electron? species accepts electrons? ion passes through salt bridge? reducing agent? oxidizing agent? electrode is the cathode? electrode is the anode? oxidation half cell reduction half cell b. What memory device can be used to remember that the cathode is the site of reduction and + ions pass through the porous membrane to the cathode? d. Do electrons flow from the aluminum to the nickel or from the nickel to the aluminum? e. In which directions do the cations migrate and in which direction do the anions migrate through the solution? 18. A voltaic cell consists of a strip of lead metal in a solution of Pb(NO3)2 in one beaker, and in the other beaker a platinum electrode is immersed in a NaCl solution, with Cl 2 gas bubbled around the electrode. The two beakers are connected with a salt bridge. a. Write the equation for the overall cell reaction. b. What is the overall voltage generated by the cell under standard conditions? c. Which electrode serves as the anode, and which serves as the cathode? 22. An electrolytic cell contains a solution of Cr(NO3)3. a. Write the anode, cathode and overall equations. anode d. Does the Pb electrode gain or lose mass as the cell reaction proceeds? cathode overall b. How long will it take to deposit 15.0 g of chromium metal, using a current of 4.50 A? D. Electrolytic Cell 19. Answer the following questions based on the electrolysis of fused (melted) sodium chloride. c. A current of 4.50 A for 30.0 minutes passed through the cell. The initial electrolyte contained 250 mL of 1.00 M Cr(NO3)3. Determine the (1) Initial moles of Cr3+. (2) moles of Cr3+ reacted. (3) Concentration of Cr3+ after 30.0 min. of electrolysis. a. which Cl2 Cl- Na species is reduced? species is oxidized? species gives up its electron? species accepts electrons? species has a higher electron affinity? species is the reducing agent? species is the oxidizing agent? b. Write the equations for the electrolysis of NaCl. oxidation half reaction (4) Concentration of H+ after 30.0 min. of electrolysis. Na+ reduction half reaction overall reaction c. What mass of Na(l) is produced using a current of 3.00 A for one hour? 20. Which cell... Electrolytic Voltaic a. has a battery? b. has a salt bridge? c. is spontaneous? d. has a positive anode? 21. Write equations for the oxidation, reduction and overall reactions for the electrolysis of the salt solutions. (5) Mass of Cr(s) that plate out at the cathode. Practice Multiple Choice Briefly explain why the answer is correct in the space provided. 1. H2Se + 4 O2F2 SeF6 + 2 HF + 4 O2 Which is true regarding the reaction represented above? (A) Oxidation number of O does not change. (B) Oxidation number of H changes from -1 to +1. (C) Oxidation number of F changes from +1 to -1. (D) Oxidation number of Se changes from -2 to +6. 2. 6 I- + 2 MnO4- + 4 H2O 3 I2 + 2 MnO2 + OHWhich statement regarding the reaction is correct? (A) Iodide ion is oxidized by hydroxide ion. (B) MnO4- is oxidized by iodide ion. (C) Manganese oxidation number changes from +7 to +2. (D) Oxidation number of iodine changes from -1 to 0. 3. 2 H2O + 4 MnO4- + 3 CIO2- 4 MnO2 + 3 CIO4- + 4 OHWhich species acts as an oxidizing agent in the reaction? (A) MnO4- (B) CIO4(C) CIO2(D) MnO2 4. Which species CANNOT function as an oxidizing agent? (A) Cr2O72- (B) MnO4- (C) NO3(D) I- 5. When acidified solutions of K2Cr2O7 and Na2S are mixed, Cr3+ and S are formed. Which is the best reducing agent? (A) K2Cr2O7 (B) Na2S (C) Cr3+ (D) S 6. _Ag+ + _AsH3 + _OH- _Ag + _H3AsO3 + _H2O What is the coefficient for OH- in the balanced equation? (A) 2 (B) 4 (C) 5 (D) 6 NaCl CuSO4 Ba(OH)2 HNO3 Na2CO3 KF 7. _Cr2O72- + _e- + _H+ _Cr3+ + _H2O What is the coefficient for H+ in the balanced half-reaction? (A) 2 (B) 6 (C) 7 (D) 14 18. 10 HI + 2 KMnO4 + 3 H2SO4 5 I2 + 2 MnSO4 + K2SO4 + 8 H2O How many moles of HI are used to produce 2.5 mol of I2? (A) 5.0 (B) 8.0 (C) 10. (D) 12 8. _Mg + _NO3– + _H+ _Mg2+ + _NH4+ + _H2O What is the coefficient for H+ in the balanced equation? (A) 1 (B) 3 (C) 5 (D) 10 19. 2 H2O + 4 MnO4- + 3 CIO2- 4 MnO2 + 3 CIO4- + 4 OHHow many moles of ClO2- react with 0.20 L of 0.20 M MnO4-? (A) 0.030 (B) 0.053 (C) 0.075 (D) 0.13 9. _CrO2– + _OH– _CrO42– + _H2O + _e– What is the ratio of the coefficients OH–/CrO2– in the balanced half-reaction (A) 1:1 (B) 2:1 (C) 3:1 (D) 4:1 20. 10. _Fe(OH)2 + _O2 + _H2O _Fe(OH)3 If 1 mole of O2 oxidizes Fe(OH)2, how many moles of Fe(OH)3 can be formed? (A) 2 (B) 3 (C) 4 (D) 5 11. In which species does sulfur have the same oxidation number as it does in H2SO4? (A) H2SO3 (B) S2O32- (C) S2(D) SO2Cl2 12. 2 HClO + 3 O2 2 HClO4 As the reaction proceeds to the right, the oxidation number of chlorine changes from (A) -1 to +3 (B) -1 to +5 (C) +1 to +7 (D) +3 to +7 13. Which will generate H2(g) when added to 1 M HCl? (A) CuS (B) Zn (C) CaCO3 (D) Mg(OH) 14. _Cr2O72- + _H2S + _H+ _Cr3+ + _S + _H2O What is the coefficient for H+ in the balanced equation? (A) 2 (B) 4 (C) 6 (D) 8 15. 3 Cu + 8 H+ + 2 NO3- 3 Cu2+ + 2 NO + 4 H2O Which statements about the reaction are true? I. Cu acts as an oxidizing agent. II. Nitrogen's oxidation state changes from +5 to +2. III. Hydrogen ions are oxidized to form H2O. (A) I only (B) II only (C) III only (D) I and II 16. In which reaction does the same element undergo both oxidation and reduction? (A) S8(s) + 8 O2(g) 8 SO2(g) (B) 3 Br2(aq) + 6 OH- 5 Br- + BrO3- + 3 H2O (C) Ca2+ + SO42- CaSO4(s) (D) PtCI4(s) + 2 CI- PtCI62- 17. Which reaction is an oxidation-reduction reaction? (A) HC2H3O2(aq) + NH3(aq) C2H3O2- + NH4+ (B) Ba2+ + SO42- BaSO4(s) (C) Zn(OH)2(s) + 2 OH- [Zn(OH)4]2(D) 2 K(s) + Br2(l) 2 KBr(s) 5 Fe2+ + MnO4- + 8 H+ 5 Fe3+ + Mn2+ + 4 H2O 25.0 mL of an acidified Fe2+ solution requires 14.0 mL of 0.10-M MnO4- solution to reach the equivalence point. The concentration of Fe2+ in the original solution is (A) 0.10 M (B) 0.56 M (C) 0.28 M (D) 0.14 M 21. Use the reduction potentials to determine which one of the reactions below is spontaneous. -0.5 V 0.5 V Cd2+ + 2 e- Cd Cu+ + 1 e- Cu 2+ Mn + 2 e Mn -1.2 V Fe3+ + 1 e- Fe2+ 0.7 V (A) Cd2+ + 2 Cu Cd + 2 Cu (B) Mn2+ + 2 Cu Mn + 2 Cu+ (C) Cd2+ + Mn Cd + Mn2+ (D) Cu+ + Fe3+ Cu + Fe2+ 22. According to the information below, what is the standard reduction potential for the half-reaction: M3+ + 3 e- M? M + 3 Ag+ 3 Ag + M3+ Eo = + 2.5 V + Ag + e Ag Eo = + 0.8 V (A) -1.7 V (B) -0.1 V (C) 0.1 V (D) 1.7 V 23. Magnesium reacts with dilute hydrochloric acid to produce hydrogen gas. Silver does not react in dilute hydrochloric acid. Based on this information, which of the following reactions will occur spontaneously? (A) H2(g) + Mg2+ 2 H+ + Mg(s) (B) 2 Ag(s) + Mg2+ 2 Ag+ + Mg(s) (C) 2 Ag+ + Mg(s) 2 Ag(s) + Mg2+ (D) 2 Ag + 2 H+ H2(g) + 2 Ag+ 24. Zn + Cu2+ Zn2+ + Cu Which could account for the observed voltage of 1.00 V instead of the standard cell potential, Eo, of 1.10 V? (A) The copper electrode was larger than the zinc electrode. (B) The Zn2+ electrolyte was Zn(NO3)2, while the Cu2+ electrolyte was CuSO4. (C) The concentration of the Zn2+ solution was greater than the Cu2+ solution. (D) The solutions in the half-cells had different volumes. Questions 25-26 refer to an electrolytic cell that involves the following half-reaction. AIF63- + 3 e- Al + 6 F25. Which of the following occurs in the reaction? (A) AIF63- is reduced at the cathode. (B) Al is oxidized at the anode. (C) Aluminum is converted from the -3 oxidation state to the 0 oxidation state. (D) F- acts as a reducing agent. 26. A steady current of 10 A is passed through an aluminumproduction cell for 15 minutes. Which of the following is the correct expression for calculating the number of grams of aluminum produced? (1 faraday = 96,500 coulombs) (A) (10)(15)(96,500)/(27)(60) (B) (10)(15)(27)/(60)(96,500) (C) (10)(15)(60)(27)/(96,500)(3) (D) (96,500)(27)/(10)(15)(60)(3) 27. Which of the following expressions is correct for the maximum mass of copper, in grams, that could be plated out by electrolyzing aqueous CuCl2 for 16 hours at a constant current of 3.0 A? (1 faraday = 96,500 coulombs) (A) (16)(3,600)(3.0)(63.55)(2)/(96,500) (B) (16)(3,600)(3.0)(63.55)/(96,500)(2) (C) (16)(3,600)(3.0)(63.55)/(96,500) (D) (16)(60)(3.0)(96,500)(2)/(63.55) 34. A power supply has lost the markings that indicate the positive and negative. A chemist suggests that the terminals be connected to a pair of platinum electrodes that dip into 0.1 M KI solution. Which correctly identifies the polarities of the terminals? (A) A gas will be evolved only at the positive electrode. (B) A gas will be evolved only at the negative electrode. (C) An amber color will appear near the negative electrode. (D) A metal will be deposited on the positive electrode. 35. In the electroplating of nickel, 0.200 faraday of electrical charge is passed through a solution of NiSO4. What mass of nickel is deposited? (A) 2.94 g (B) 5.87 g (C) 11.7 g (D) 58.7 g Practice Free Response 28. If 0.060 faraday is passed through an electrolytic cell containing a solution of In3+ ions, the maximum number of moles of In that could be deposited at the cathode is (A) 0.010 (B) 0.020 (C) 0.030 (D) 0.060 1. Questions 29-33 The spontaneous reaction that occurs when the cell below operates is 2 Ag+ + Cd 2 Ag + Cd2+ A power supply is connected to two platinum electrodes immersed in a beaker containing 1.0 M CuSO4(aq) at 25oC. As the cell operates, copper metal is deposited on the left electrode and O2(g) is produced at the right electrode. The reduction half-reactions that occur are: O2(g) + 4 H+(aq) + 4 e- 2 H2O(l) Eo = +1.23 V Cu2+(aq) + 2 e- Cu(s) Eo = +0.34 V a. Is the direction of electron flow in the wire from left to right or from right to left? b. Write a balanced net ionic equation for the electrolysis reaction that occurs. A current of 1.50 A passes through the cell for 40.0 minutes. c. Calculate the mass of the Cu(s) that is deposited. Which occurs for each of the following circumstances? (A) Voltage increases. (B) Voltage decreases but remains above zero. (C) Voltage becomes zero and remains at zero (D) No change in voltage occurs 29. A 50-mL sample of a 2-M Cd(NO3)2 solution is added to the left beaker. 30. The silver electrode is made larger. 31. The salt bridge is replaced by a platinum wire. d. 2. Calculate the dry volume, in liters, measured at 25oC and 1.16 atm, of the O2(g) that is produced. A voltaic cell is constructed with a strip of Sn in a solution of Sn(NO3)2 in one container and a strip of an unknown metal, X, in a solution of X(NO3)3 in another container. The two containers are connected by a salt bridge and the two metal strips are connected with a metal wire. The mass of the Sn electrode increases. The half-reactions are: Sn2+(aq) + 2 e- Sn(s) Eo = -0.14 V 3+ X (aq) + 3 e X(s) Eo = ? a. Which electrode is the cathode? Justify your answer. 32. Current is allowed to flow for 5 minutes. b. What directions do electrons flow in terms of the Sn and unknown metal strips. 33. The silver electrode is replaced by a copper electrode. c. If the cell potential, Eocell, is + 0.60 V, what is the standard reduction potential for X3+(aq) + 3 e- X(s)? d. Identify metal X from the chart of standard potentials. 3. e. Write a balanced net-ionic equation for the overall chemical reaction occurring in the cell. f. If the [Sn2+] = 0.50 M and [X3+] = 0.10 M. What is the cell potential, Ecell? 4. An electrochemical cell is constructed with an open switch, as shown in the diagram above. A strip of Sn and a strip of an unknown metal, X, are used as electrodes. When the switch is closed, the mass of the Sn electrode increases. An electrolytic cell contains a solution of Cr(NO3)3. Assume that chromium metal plates out at one electrode and oxygen gas is evolved at the other electrode. a. Write the anode half-reaction, the cathode halfreaction, and the overall reaction for the cell. b. How long will it take to deposit 15.0 g of chromium metal, using a current of 4.50 A? c. A current of 4.50 A is passed through the cell for 30.0 min. If we start out with 250 mL of 1.00 M Cr(NO3)3, what is the concentration of Cr3+ after electrolysis? A voltaic cell consists of two half-cells, one of which contains a Pt electrode surrounded by Cr3+ and Cr2O72ions. The other half-cell contains a Pt electrode surrounded by Mn2+ ions and MnO2(s). Assume the cell reactions, which produce a positive voltage, involves both Cr3+ and Mn2+ ions. Cr2O72- + 14 H+ + 6 e- 2 Cr3+ + 7 H2O Eored = 1.33 V + 2+ MnO2(s) + 4 H + 2 e Mn + 2 H2O Eored = 1.229 V d. Write the anode half-reaction, the cathode halfreaction, and the overall reaction for the cell. e. Calculate Eotot for the cell. f. Calculate the voltage of the cell when all ionic species except H+ have a concentration of 0.300 M and the solution has [H+] = 3.16 x 10-4 M. The half-reactions are shown below. Sn2+ + 2 e- Sn(s) Eo = -0.14 V X3+ + 3 e- X(s) Eo = ? a. In the diagram above, label the electrode that is the cathode. Justify your answer. b. In the diagram above, draw an arrow indicating the direction of the electron flow in the external circuit when the switch is closed. c. If the standard cell potential, Eocell, is + 0.60 V, what is the standard reduction potential, in volts, for the X3+/X electrode? d. Identify metal X. e. Write a balanced net-ionic equation for the overall chemical reaction occurring in the cell. f. In the cell, the concentration of Sn2+ is changed from 1.0 M to 0.50 M, and the concentration of X3+ is changed from 1.0 M to 0.10 M. (1) Substitute all the appropriate values for determining the cell potential, Ecell, into the Nernst equation. (Do not do any calculations.) (2) On the basis of your response in part (f) (1), will the cell potential, Ecell, be greater than, less than, or equal to the original Eocell? Justify your answer.
