Ch.18 - ElectrochemistryWorksheetSee all chapters
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Ch.1 - Intro to General Chemistry
Ch.2 - Atoms & Elements
Ch.3 - Chemical Reactions
BONUS: Lab Techniques and Procedures
BONUS: Mathematical Operations and Functions
Ch.4 - Chemical Quantities & Aqueous Reactions
Ch.5 - Gases
Ch.6 - Thermochemistry
Ch.7 - Quantum Mechanics
Ch.8 - Periodic Properties of the Elements
Ch.9 - Bonding & Molecular Structure
Ch.10 - Molecular Shapes & Valence Bond Theory
Ch.11 - Liquids, Solids & Intermolecular Forces
Ch.12 - Solutions
Ch.13 - Chemical Kinetics
Ch.14 - Chemical Equilibrium
Ch.15 - Acid and Base Equilibrium
Ch.16 - Aqueous Equilibrium
Ch. 17 - Chemical Thermodynamics
Ch.18 - Electrochemistry
Ch.19 - Nuclear Chemistry
Ch.20 - Organic Chemistry
Ch.22 - Chemistry of the Nonmetals
Ch.23 - Transition Metals and Coordination Compounds
Additional Problems
Refer to the tabulated values of ΔG°f in Appendix IIB to calculate E°cell for the fuel cell breathalyzer, which employs the following reaction. (ΔG°f for HC2H3O2(g) = -374.2 kJ/mol.) CH3CH2OH(g) + Oz(g) → HC2H3O2(g) + H2O(g)
For the a voltaic cell with the overall reaction:             Zn(s) + Cu2+(aq) →  Zn2+(aq) + Cu(s)           E°cell = 1.10 V Given that the standard reduction potential of Zn2+ to Zn(s) is –0.76 V, calculate the standard reduction potential for: Cu2+(aq) + 2e− → Cu(s) A.  0.76 V B.  −0.76 V C.  1.86 V D.  −1.86 V E.  0.34 V
What is the standard electrode potential for a voltaic cell constructed in the appropriate way from these two half-cells?  Standard Reduction Potentials                   E°      Cr 3+ (aq) + 3e– → Cr(s)                        – 0.74 V Co 2+ (aq) + 2e– → Co(s)                     – 0.28 V   A) – 1.02 V B) 0.46 V C) 0.64 V D) 1.02 V
Calculate the equilibrium constant for each of the reactions in Problem 66. b. O2(g) + 2H2O(l) + 2Cu(s) → 4OH–(aq) + 2Cu2+(aq)
If for the reaction, 2 Ag+(aq) + Co(s) →  2 Ag(s) + Co2+(aq), E is 0.97 V when  [Ag+] = 0.0025 M and [Co2+] = 0.032 M, then E° for this reaction is a. 1.08 V. b. 1.05 V. c. 0.89 V. d. 0.86 V.  
The reduction potentials for Au3+ and Ni2+ are +1.50 V and -0.23 V, respectively. What is ΔG° at 25°C for the reaction below? 2Au3+ + 3Ni → 3Ni 2+ + 2Au a. -5.00 x 102 kJ b. +5.00 x 102 kJ c. -2140 kJ d. +1.00 x 103 kJ e. -1.00 x 103 kJ  
What is the standard cell potention, E°, for this half-reaction? Pd2+(aq) + 2e– → Pd(s) Standard Reduction Potentials                              E°        Cu(s) + Pd 2+ → Cu 2+(aq)  + Pd(s)              + 0.650 V Cu 2+ (aq) + 2e– → Cu(s)                               + 0.337 V   A) + 0.987 V B) –0.987 V C) + 0.313 V D) – 0.313 V
Calculate the standard cell potential for each of the electrochemical cells in Problem 44. b. 2 H+(aq) + Fe(s) → H 2(g) + Fe2+(aq
Using the following standard reduction potentials, Fe3+ (aq) + e–   → Fe2+ (aq)  E°red = +0.77 V Pb2+ (aq) + 2 e– → Pb (s)      E°red = –0.13 V   Calculate the standard emf for the following cell reaction as written and determine whether or not the reaction is spontaneous or non‐spontaneous: Pb2+ (aq) + 2 Fe2+ (aq) → 2 Fe3+ (aq) + Pb (s) a) –0.90 V, non‐spontaneous b) –0.90 V, spontaneous c) +0.90 V, non‐spontaneous d) –0.64 V, non‐spontaneous e) +0.64 V, spontaneous  
Calculate the standard cell potential for each of the electrochemical cells in Problem 44. c. 2 NO3– (aq) + 8 H+(aq) + 3 Cu(s) → 2 NO(g) + 4 H 2O (/) + 3 Cu2+(aq)
Calculate the standard cell potential for each of the electrochemical cells in Problem 44. a. Ni2+(aq) + Mg(s) → Ni(s) + Mg2+(aq)
What is the standard e.m.f. for a voltaic cell based on the reactions: Sn4+ (aq) + 2e– → Sn2+ (aq)    E ̊red = +0.154 V Cr3+ (aq) + 3e–  → Cr (s)          E ̊red = –0.74 V   a) +1.94 V b) +0.89 V c) +2.53 V d) –0.59 V e) –1.02 V  
A friend wants you to invest in a new battery she has designed that produces 24 V in a single voltaic cell. Why should you be wary of investing in such a battery?
Calculate the standard cell potential of a Al/Al 3+ | Cd/Cd2+ voltaic cell.  Al3+ (aq) + 3e– → Al(s)      E ̊red = –1.66 V Cd2+ (aq) + 2e– → Cd(s)    E ̊red = –0.403V a) –2.11 V; nonspontaneous b) 2.11 V; spontaneous c) 1.26 V; spontaneous d) –1.26 V; nonspontaneous e) 1.257 V; non‐spontaneous  
What voltage can theoretically be achieved in a battery in which lithium metal is oxidized and fluorine gas is reduced? Why might such a battery be difficult to produce?
The Ksp of CuI is 1.1 x 10–12. Find Ecell for the cell: Cu(s) | Cul(s) | I–(aq)(1.0 M) || Cu+(aq)(1.0 M) | Cu(s)
The Ksp of Zn(OH)2 is 1.8 x 10-14 . Find Ecell for the half-reaction: Zn(OH)2(s) + 2 e– ⇌ Zn(s) + 2 OH– (aq)  
Calculate ΔG°rxn and K for each reaction. a. The reaction of Cr2+(aq) with Cr 2O72– (aq) in acid solution to form Cr 3+(aq). b. The reaction of Cr3+(aq) and Cr(s) to form Cr 2+(aq). [The electrode potential of Cr 2+(aq) to Cr(s) is –0.91 V.]
Consider a voltaic cell made up of Ag(s), 1M Ag(NO 3) (aq), 1M Cu(NO3)2 (aq), and Cu(s). Write out the half reaction at the anode, at the cathode, and the overall cell reaction. Give the initial voltage for the cell and the maximum amount of work that can be obtained from this battery. 
THe equilibrium constant Kb = 1.75 x 10 -5 M at 25°C for the following reaction.  Calculate ΔG (kJ) at 25°C when both [NH4+] and [OH -] = 1.0 x 10 -6 M and [NH3] = 0.050 M? NH3(aq)  +  H2O(l)  ⇌  NH4+(aq)  +  OH -(aq)] A. 27.2 B. 45.8 C. -61.1 D. -45.8 E. -33.9
Given the listed standard electrode potentials, what is E° for the cell? Ag+(aq) + Fe2+(aq)  ⇌  Ag(s) + Fe3+(aq)          Fe3+(aq) + 1 e -   → Fe2+(aq)        E° = +0.77 V          Ag+(aq) + 1 e -    → Ag(s)               E° = +0.80 V                      A. +0.03                      B. -0.03                      C. +0.74                      D. +1.57                      E. -1.57
Use the standard half-cell potentials listed below to calculate the standard cell potential, E°cell, for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.) 3 Br2(l) + 2 Fe(s) → 6Br1-(aq) + 2Fe2+(aq) Br2(l) + 2 e- → 2 Br1-(aq)                 E° = +1.09 V Fe2+(aq) + 2 e- → Fe(s)                   E° = -0.45 V A) +0.64V B) -0.64 V C) -1.54 V D) +1.54 V E) the standard cell potential is 0.0v
Calculate the standard free energy change (kJ) for the following reaction using the standard electrode potentials. Sn(s) + Ni2+(aq)  ⇌  Sn2+(aq) + Ni(s)          Ni 2+(aq) + 2 e - → Ni(s)        E° = -0.25 V          Sn2+(aq) + 2 e - → Sn(s)        E° = -0.14 V                      A. +10                      B. -10                      C. +15                      D. -21                      E. +21
Using the standard electrode potentials, determine the cell voltage for the following reaction when [Fe2+] = 0.10 M and [Cd2+] = 1.0 M. Fe(s) + Cd2+(aq)  ⇌ Fe2+(aq) + Cd(s)          Fe2+(aq) + 2 e -→ Fe(s)        E° = -0.44 V          Cd2+(aq) + 2 e -→ Cd(s)        E° = -0.40 V A. 0.04 B. 0.07 C. 0.01 D. 0.10 E. -0.02
The standard electrode potential for the following reaction is -0.13 V.  What is the potential at pH = 13.00? CrO42-(aq) + 4 H2O(l) + 3 e - → Cr(OH)3(s) + 5 OH -(aq)   A. 0.42 B. 0.031 C. -0.031 D. 0.27 E. 2.24
Consider an electrochemical cell that utilizes lithium and zinc. What is the standard voltage for this cell? Li + + e–      →       Li                                                                      E° = – 3.05 V Zn2+ + 2 e–      →        Zn                                                               E° = +0.799 V   a. 1.90 V b. 2.29 V c. 3.85 V d. 5.34 V
A spontaneous redox reaction is indicated by: a. A positive emf or cell potential. b. A negative emf or cell potential. c. A positive enthalpy. d. A negative enthalpy. 
Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. 3 Cl2(g) + 2 Fe(s) 6 Cl – (aq) + 2 Fe3+(aq) Cl2(g) + 2 e–       →      2 Cl – (aq)                                        E° = +1.36 V Fe3+(aq) + 3 e–      →         Fe(s)                                         E° = -0.04 V   A) +4.16 V B) -1.40 V C) -1.32 V D) +1.32 V E) +1.40 V
Which of the following reactions would have the smallest value of K at 298 K?   A) A + B       →       C;                E°cell = +1.22 V B) A + 2B     →       C;               E°cell = +0.98 V C) A + B      →      2 C;              E°cell = -0.030 V D) A + B      →       3 C;             E°cell = +0.15 V E) None of the above.
Consider the following reaction: 2 Fe 2+ (aq) + Cu 2+ (aq)      →       2 Fe  3+ (aq) + Cu (s) Given the following data of reduction potentials: Fe3+ + e –      →       Fe 2+                                              E °  = 0.77 Volts Cu2+ (aq) + 2 e –       →      Cu (s)                                 E °  = 0.34 Volts  Which of the following is a false statement? a) When calculating the cell potential, the coefficients are not taken into account. b) At equilibrium the cell voltage is zero. c) Cu2+ is the oxidizing agent. d) The cell potential for this reaction is 1.11 Volts. 
Calculate the maximum electrical work that can be produced by this cell. 3 Co 2+ (aq) + 2 Cr (s) 2 Cr 3+ (aq) + 3 Co (s) Given the following reduction potentials: Cr 3+ + 3 e –      →       Cr                      E ° = - 0.74 Volts  Co 2+ (aq) + 2 e –       →       Co            E ° = - 0.28 Volts  Faraday’s constant = 96,500 Coulombs / mole of electrons a) – 800 kJ      b) – 150 kJ      c) – 266 kJ      d) + 266 kJ
Calculate the value of Eocell  for the following reaction: 2 Co3+(aq) + 3 Ca(s) → 2 Co(s) + 3 Ca2+(aq) Co3+(aq) + 3e– →  Co(s)     Eo = 1.54 V Ca2+(aq) + 2e– →  Ca(s)     Eo = –2.87 V.    A)  –4.41 V    B)  –1.33 V    C)  –11.6 V    D)  +1.33 V    E)  +4.41 V  
Consider the following standard reduction potentials in acid solution.  Which half reaction is the most favored process thermodynamically?                                                                                                 E°(V)             A)        Al3+      + 3 e–                →          Al(s)                    –1.66             B)        AgBr(s) +   e–                 →          Ag(s) + Br –        +0.07             C)        Sn4+      + 2 e–               →          Sn2+                  +0.14             D)        Fe3+      + e–                  →          Fe2+                  +0.77             E)         O2        + 4 e –  + 4 H+     →          H2O                 +1.23
Consider the non–aqueous cell reaction 2 K(s) + Fe2+ (aq)  →   2 K+ (aq) + Fe(s) for which  E°cell = 1.49 V at 298 K. ΔG° at this temperature is   A)        +287.6  kJ                    D)        –143.7 kJ B)        –287.6 kJ                     E)         None of these choices is correct. C)        +143.7 kJ
What is the relation between ΔG° and E°cell for the cell reaction below? Ni2+ (aq) + Cd (s) → Cd 2+ (aq) + Ni (s)   a) ∆G ̊ = F E ̊cell b) ∆G ̊ = 2 F E ̊cell c) ∆G ̊ = –F E ̊cell d) ∆G ̊ = –2 F E ̊cell e) ∆G ̊ = –4 F E ̊cell  
A quantity of a powdered mixture of zinc and iron is added to a solution containing Fe2+ and Zn2+ ions, each at unit activity. What reaction will occur? (A) Zinc ions will oxidize Fe to Fe2+ (B) Fe2+ ions will be oxidized to Fe3+ ions (C) Zinc ions will be reduced to zinc metal (D) Zinc metal will reduce Fe2+ ions
Calculate the value of E°cell for the following reaction: 2Au(s) + 3Ca2+(aq) → 2Au3+(aq) + 3Ca(s) Au3+(aq) + 3e- → Au(s)      E° = 1.50V Ca2+(aq) + 2e- → Ca(s)      E° = -2.87V A) -4.37 V B) -1.37 V C) -11.6 V D) 1.37 V E) 4.37 V
Which combination of reactants will produce the greatest voltage based on these standard electrode potentials? (A) Cu+ and Sn2+ (B) Cu+ and Cr2+ (C) Cu and Sn2+ (D) Sn4+ and Cr2+
Consider the non-aqueous cell reaction 2Na(l) + FeCl2(s) ⇌ 2NaCl(s) + Fe(s) for which E°cell = 2.35 V at 200°C. ΔG° at this temperature is A) 453 kJ. B) -453 kJ. C) 907 kJ. D) -907kJ. E) None of these choices is correct.
What is the standard cell potential, E°, for this reaction? 3 Mn (s) + 2 AuCl4- (aq) → 3 Mn2+ (aq) + 2 Au (s) + 8 Cl- (aq) (A) -0.18 V (B) -2.18 V (C) +2.18 V (D) +5.54 V
Given Cu2+(aq) + 2e- → Cu(s)      E° = 0.337 V Al3+(aq) + 3e- → Al(s)         E° = -1.66 V Na+(aq) + e- → Na(s)          E° = -2.714 V. Which of the following reactions will occur? A) 2Na+(aq) + Cu(s) → Cu2+(aq) + 2Na(s)              E°cell = -3.051 V B) Al(s) + 3Na+(aq) → Al3+(aq) + 3Na(s)                E°cell = -1.054 V C) 2Na(s) + Cu2+(aq) → Cu(s) + 2Na+(aq)              E°cell = 3.051 V D) 2Al3+(aq) + 3Cu(s) → 3Cu2+(aq) + 2Al(s)          E°cell = -1.991 V E) 2Al3+(aq) + 3Cu(s) → 3Cu2+(aq) + 2Al(s)           E°cell = 1.991 V
The standard cell potential. E°, for this reaction is 0.79 V. 6 I- (aq) + Cr2O72- (aq) + 14H+ (aq) → 3 l2 (aq) + 2 Cr3+ (aq) + 7 H2O (l) What is the standard potential for l2 (aq) being reduced to l- (aq) given that the standard reduction potential for Cr2O72- (aq) changing to Cr3+(aq) is +1.33 V? (A) +0.54 V (B) -0.54 V (C) +0.18 V (D) -0.18 V
Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.) Sn(s) + 2Ag+(aq) → Sn2+(aq) + 2Ag(s) Sn2+(aq) + 2e- → Sn(s)                       E° = -0.14 V Ag+(aq) + e- → Ag(s)                           E° = +0.80 V A) +0.94 V B) -1.08 V C) +1.08 V D) +1.74 V E) -1.74 V
For the a voltaic cell with the overall reaction: Zn(s) + Cu 2+(aq) → Zn 2+(aq) + Cu(s) E o cell = 1.10 V Given that the standard reduction potential of Zn 2+ to Zn(s) is – 0.76 V, calculate the standard reduction potential for: Cu 2+(aq) + 2e − → Cu(s)   A. 0.76 V B. −0.76 V C. 1.86 V D. −1.86 V E. 0.34 V  
Make a sketch of the voltaic cell represented by the line notation. Write the overall balanced equation for the reaction and calculate E°cell. Mn(s) | Mn2+(aq) || ClO2–(aq) | ClO2(g) | Pt(s)    
Determine whether or not each redox reaction occurs spontaneously in the forward direction. a. Ca2+(aq) + Zn(s) → Ca(s) + Zn2+(aq)
Determine whether or not each redox reaction occurs spontaneously in the forward direction. b. 2 Ag+(aq) + Ni(s) → 2 Ag(s) + Ni 2+(aq)
Determine whether or not each redox reaction occurs spontaneously in the forward direction. c. Fe(s) + Mn2+(aq) → Fe2+(aq) + Mn(s)
What is the ΔG° (kJ/mol) for a spontaneous (voltaic) cell that uses the following half-reactions:                 Al+3 (aq) + 3 e- → Al (s)                E° = −1.66 Mn2+ (aq) + 2 e- → Mn (s)             E° = −1.18 A. −278           B. −232           C. −140           D. −46.3          E. 46.3
Determine whether or not each redox reaction occurs spontaneously in the forward direction. d. 2 Al(s) + 3 Pb 2+(aq) → 2 Al 3+(aq) + 3 Pb(s)  
Design an electrolytic cell from the following half-reactions: Cd 2+(aq) + 2 e − →  Cd(s)                   E  o = −0.40 Cr 3+(aq) + 3 e − → Cr(s)                     E   o = −0.74   The overall cell potential of the electrolytic cell is: A. −1.14           B. −0.34             C. −0.28               D. 0.28             E. 0.34  
Which metal can be oxidized with an Sn 2+ solution but not with an Fe 2+ solution?
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. a. O2(g) + 2 H2O (l) + 4 Ag(s) → 4 OH – (aq) + 4 Ag +(aq)
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. b. Br2(l) + 2 I –(aq) → 2 Br –(aq) + I 2(s)    
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. c. PbO2(s) + 4 H+(aq) + Sn(s) → Pb 2+(aq) + 2 H2O (l) + Sn2+(aq)
Use tabulated electrode potentials to calculate Δ G°rxn for each reaction at 25°C. a. 2 Fe3+(aq) + 3 Sn(s) → 2 Fe(s) + 3 Sn 2+(aq)    
Use tabulated electrode potentials to calculate Δ G°rxn for each reaction at 25°C. b. O2(g) + 2 H2O(l) + 2 Cu(s) → 4 OH–(aq) + 2 Cu 2+(aq)  
Use tabulated electrode potentials to calculate Δ G°rxn for each reaction at 25°C. c. Br2(l) + 2 I –(aq) → 2 Br – (aq) + I2(s)  
Calculate the equilibrium constant for each of the reactions in Problem 66. a. 2 Fe3+(aq) + 3 Sn(s) → 2 Fe(s) + 3 Sn 2+(aq)
Design an electrolytic cell from the following half-reactions:                         Cd2+(aq) + 2 e− → Cd(s)        E° = −0.40                         Cr3+(aq) + 3 e− → Cr(s)          E° = −0.74 The overall cell potential of the electrolytic cell is:  A. −1.14          B.  −0.34         C.  −0.28         D.  0.28           E.  0.34
Calculate the equilibrium constant for the reaction between Fe 2+(aq) and Zn(s) (at 25°C).
Calculate ΔG°rxn and E°cell for a redox reaction with n = 3 that has an equilibrium constant of K = 0.050 (at 25 °C).
Which of these changes will produce the most positive voltage for this half reaction in the direction written? Co2+ (aq) + 2e– → Co(s)      E° = –0.28 V a) increasing the amount of solid Co b) decreasing the amount of solid Co c) increasing the concentration of Co2+ d) decreasing the concentration of Co2+    
An electrochemical cell is based on these two half-reactions: Ox: Sn(s) → Sn2+(aq, 2.00 M) + 2e– Red: ClO2(g, 0.100 atm) + e– → ClO2–(aq, 2.00 M) Calculate the cell potential at 25 °C.
Calculate the equilibrium constant for each of the reactions in Problem 66. c. Br2(l) + 2l–(aq) → 2Br–(aq) + I2(s)
Cytochrome, a complicated molecule that we will represent as CyFe2 + , reacts with the air we breathe to supply energy required to synthesize adenosine triphosphate (ATP). The body uses ATP as an energy source to drive other reactions. (Section 19.7 in the textbook) At pH 7.0 the following reduction potentials pertain to this oxidation of CyFe2 + : O2 (g) + 4H+ (aq) + 4e-  →  2H2 O(l), Eredo= +0.82 V, CyFe3 + (aq) + e-  →  CyFe2 + (aq), Eredo= + 0.22 V.What is G for the oxidation of CyFe2 + by air?
Consider a redox reaction for which E is a negative number.Can an electrochemical cell based on this reaction accomplish work on its surroundings?
Using the standard reduction potentials in Appendix E, calculate the standard voltage generated by the hydrogen fuel cell in acidic solution.
Mercuric oxide dry-cell batteries are often used where a high-energy density is required, such as in watches and cameras. The two half-cell reactions that occur in the battery are l HgO(s) + H2 O(l) + 2e-  →  Hg(l) + 2OH- (aq)Zn(s) + 2OH- (aq)  →  ZnO(s) + H2 O(l) + 2e- .Why is the potential of the anode reaction different than would be expected if the reaction occurred in an acidic medium?
Can a spontaneous redox reaction be obtained by pairing any reduction half-reaction with one listed above it or with one listed below it in Table 19.1 in the textbook?
Predict whether the following reactions will be spontaneous in acidic solution under standard conditions.Oxidation of Sn to Sn2 + by I2 (to form I-).
Predict whether the following reactions will be spontaneous in acidic solution under standard conditions.Reduction of Ni2 + to Ni by I- (to form I2).
Predict whether the following reactions will be spontaneous in acidic solution under standard conditions.Reduction of Ce4 + to Ce3 + by H2 O2.
The standard reduction potentials of the following half-reactions are given in Appendix E in the textbook: l Ag+ (aq) + e-  →  Ag(s)Cu2 + (aq) + 2e-  →  Cu(s)Ni2 + (aq) + 2e-  →  Ni(s)Cr3 + (aq) + 3e-  →  Cr(s). Determine which combination of these half-cell reactions leads to the cell reaction with the smallest positive cell emf.
Predict whether the following reactions will be spontaneous in acidic solution under standard conditions.Reduction of Cu2 + to Cu by Sn2 + (to form Sn4 + ).
Given the following half-reactions and associated standard reduction potentials: AuBr4-(aq) + 3e-  →  Au(s) + 4Br-(aq) Eredo= - 0.858V Eu3+(aq) + e-  →  Eu2+(aq) Eredo= - 0.43V IO-(aq) + H2O(l) + 2e-  →  I-(aq) + 2OH-(aq) Eredo= + 0.49V Sn2+(aq) + 2e-  →  Sn(s) Eredo= - 0.14VCalculate the value of this emf.
Given the following half-reactions and associated standard reduction potentials: AuBr4-(aq) + 3e-  →  Au(s) + 4Br-(aq) Eredo= - 0.858V Eu3+(aq) + e-  →  Eu2+(aq) Eredo= - 0.43V IO-(aq) + H2O(l) + 2e-  →  I-(aq) + 2OH-(aq) Eredo= + 0.49V Sn2+(aq) + 2e-  →  Sn(s) Eredo= - 0.14VWrite the cell reaction for the combination of these half-cell reactions that leads to the largest positive cell emf.
Given the following half-reactions and associated standard reduction potentials: AuBr4-(aq) + 3e-  →  Au(s) + 4Br-(aq) Eredo= - 0.858V Eu3+(aq) + e-  →  Eu2+(aq) Eredo= - 0.43V IO-(aq) + H2O(l) + 2e-  →  I-(aq) + 2OH-(aq) Eredo= + 0.49V Sn2+(aq) + 2e-  →  Sn(s) Eredo= - 0.14VWrite the cell reaction for the combination of half-cell reactions that leads to the smallest positive cell emf.
A 1M solution of Cu ( NO3 )2 is placed in a beaker with a strip of Cu metal. A 1 M solution of SnSO4 is placed in a second beaker with a strip of Sn metal. A salt bridge connects the two beakers, and wires to a voltmeter link the two metal electrodes.What is the emf generated by the cell under standard conditions?
Using standard reduction potentials (Appendix E in the textbook), calculate the standard emf for each of the following reactions.Cl2 (g) + 2I- (aq)  →  2Cl- (aq) + I2 (s).
Using standard reduction potentials (Appendix E in the textbook), calculate the standard emf for each of the following reactions.Ni(s) + 2Ce4 + (aq)  →  Ni2 + (aq) + 2Ce3 + (aq).
Using standard reduction potentials (Appendix E in the textbook), calculate the standard emf for each of the following reactions.Fe(s) + 2Fe3 + (aq)  →  3Fe2 + (aq).
Using standard reduction potentials (Appendix E in the textbook), calculate the standard emf for each of the following reactions.2NO3- (aq) + 8H+(aq) + 3Cu(s)  →  2NO(g) + 4H2O(l) + 3Cu2 + (aq).
A voltaic cell is based on Ag+(aq)/Ag(s) and Fe3+(aq)/Fe2+(aq) half-cells.What is the standard emf of the cell?
If a standard cell potential is Eocell =+0.85 at 25 oC, is the redox reaction of the cell spontaneous?
For the half-reaction Cl2(g) + 2e-  →  2Cl-(aq), what are the standard conditions for the reactant and product?
A voltaic cell is based on the reaction Sn(s) + I2 (s)  →  Sn2 + (aq) + 2I- (aq).Under standard conditions, what is the maximum electrical work, in joules, that the cell can accomplish if 76.0 g of Sn is consumed?
Consider the following two half reactions. Mg2+ + 2 e-  →  Mg(s) O2(g) + 2 H+(aq) + 2 e-  →  H2O2(aq) Standard electrode potentials can be found here: Table 19.1Calculate the standard cell potential of the half-reactions.
What are the units for electrical potential?
A voltaic cell is constructed that uses the following reaction and operates at 298 K: Zn(s) + Ni2 + (aq)  →  Zn2 + (aq) + Ni(s).What is the emf of this cell under standard conditions?
A voltaic cell utilizes the following reaction: 4Fe2 + (aq) + O2 (g) + 4H+ (aq)  →  4Fe3 + (aq) + 2H2 O(l).What is the emf of this cell under standard conditions?
What conditions must be met for a reduction potential to be a standard reduction potential?
Why is it impossible to measure the standard reduction potential of a single half-reaction?
A voltaic cell utilizes the following reaction and operates at 298 K: 3Ce4 + (aq) + Cr(s)  →  3Ce3 + (aq) + Cr3 + (aq).What is the emf of this cell under standard conditions?
Compute the equilibrium constant for the spontaneous reaction between Ni2+(aq) and Cd(s).
The standard potential for the reduction of AgSCN(s) is +0.09 V. AgSCN(s) + e-  →  Ag(s) + SCN- (aq).Using this value and the electrode potential for Ag+ (aq), calculate the Ksp for AgSCN. Ag+(aq) + e-  →  Ag(s) Eo= +0.80 V
Using data from Appendix E in the textbook, calculate the equilibrium constant for the disproportionation of the copper(I) ion at room temperature: 2Cu+ (aq)  →  Cu2 + (aq) + Cu(s).
The Haber process is the principal industrial route for converting nitrogen into ammonia: N2 (g) + 3H2 (g)  →  2NH3 (g).Using the thermodynamic data in Appendix C in the textbook, calculate the equilibrium constant for the process at room temperature.
Given the following reduction half-reactions: Fe3+(aq) + e-  →  Fe2+(aq) Eredo= + 0.77V S2O62-(aq) + 4H+(aq) + 2e-  →  2H2SO3(aq) Eredo= + 0.60V N2O(g) + 2H+(aq) + 2e-  →  N2(g) + H2O(l) Eredo= - 1.77V VO2+(aq) + 2H+(aq) + e-  →  VO2+(aq) + H2O(l) Eredo= + 1.00VCalculate the equilibrium constant K for this reaction at 298 K.
Given the following reduction half-reactions: Fe3+(aq) + e-  →  Fe2+(aq) Eredo= + 0.77V S2O62-(aq) + 4H+(aq) + 2e-  →  2H2SO3(aq) Eredo= + 0.60V N2O(g) + 2H+(aq) + 2e-  →  N2(g) + H2O(l) Eredo= - 1.77V VO2+(aq) + 2H+(aq) + e-  →  VO2+(aq) + H2O(l) Eredo= + 1.00VCalculate the equilibrium constant K for this reaction at 298 K.
Given the following reduction half-reactions: Fe3+(aq) + e-  →  Fe2+(aq) Eredo= + 0.77V S2O62-(aq) + 4H+(aq) + 2e-  →  2H2SO3(aq) Eredo= + 0.60V N2O(g) + 2H+(aq) + 2e-  →  N2(g) + H2O(l) Eredo= - 1.77V VO2+(aq) + 2H+(aq) + e-  →  VO2+(aq) + H2O(l) Eredo= + 1.00VCalculate the equilibrium constant K for this reaction at 298 K.
Using the standard reduction potentials listed in Appendix E in the textbook, calculate the equilibrium constant for each of the following reactions at 298 K.Fe(s) + Ni2 + (aq)  →  Fe2 + (aq) + Ni(s)
Using the standard reduction potentials listed in Appendix E in the textbook, calculate the equilibrium constant for each of the following reactions at 298 K.Co(s) + 2H+ (aq)  →  Co2 + (aq) + H2 (g)
Using the standard reduction potentials listed in Appendix E in the textbook, calculate the equilibrium constant for each of the following reactions at 298 K.10Br- (aq) + 2MnO4 - (aq) + 16H+ (aq)  →  2Mn2 + (aq) + 8H2 O(l) + 5Br2 (l)
A cell has a standard emf of + 0.151 V at 298 K.What is the value of the equilibrium constant for the cell reaction if n = 1?
A cell has a standard emf of + 0.151 V at 298 K.What is the value of the equilibrium constant for the cell reaction if n = 2?
A cell has a standard emf of + 0.151 V at 298 K.What is the value of the equilibrium constant for the cell reaction if n = 3?
Consider a redox reaction for which E is a negative number.Will the equilibrium constant for the reaction be larger or smaller than 1?
An electrochemical cell has a positive standard cell potential but a negative cell potential.What must be true of Q and K for the cell?
Calculate the standard potential, E°, for this reaction from its equilibrium constant at 298 K.X(s) + Y2+(aq) ⇌ X2+(aq) + Y(s)   K= 9.49 x 105
Calculate the Ksp of silver iodide at 25°C, using data provided. AgI(s) + e - → Ag(s) + I -      E°V = -0.15 I2 (s) + 2e -  → 2I -               E°V = 0.54 Ag+ + e -  → Ag(s)               E°V = 0.80  a) 2.9 × 10–3 b) 1.9 × 10–4 c) 2.1 × 10–12 d) 8.7 × 10–17 e) 2.4 × 10–24
Consider a fuel cell that uses the reaction of ethanol with oxygen to produce electricity, CH3CH2OH(l) + 3O2(g) → 2CO2(g)+ 3H2O(l)Determine the E°cell for this cell at 25°.E°cell = v
Consider these two entries from a fictional table of standard reduction potentials. x3+ + 3e- → x(s)      E = 2.00 V y3+ + 3e- → y(s)      E = -0.49 V What is the standard potential of a cell where X is the anode and Y is the cathode?
What is the cell potential for the reactionMg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)at 67 ∘C when [Fe2+]= 3.40 M and [Mg2+]= 0.310 M .
Calculate the equilibrium constant, K, for the following reaction at 25 °C.Fe3+(aq) +   B(s) + 6H2O (l) →   Fe(s) +   H 3BO3(s) +   3H3O+(aq)The balanced reduction half-reactions for the above equation and their respective standard reduction potential values (E°) are as follows:Fe3+(aq)   +    3e-  →  Fe(s)                                           Eo = -.04 VH3BO3(s) +   3H3O+(aq) +   3e- → B(s) +   6H2O(l)      Eo = -.8698 V
Calculate the E°cell for the following equation.  Cu(s) + Ag+ (aq) → Cu+ (aq) + Ag(s)
What statement is TRUE about standard electrode potentials?A) E°cell is positive for spontaneous reactions.B) Electrons flow from the cathode to the anode.C) The electrode potential of the standard hydrogen electrode is exactly zero.D) The salt bridge allows mixing to occur efficiently in the cellE) Both A and C
Calculate the standard potential, E°, for this reaction from its equilibrium constant at 298 K.X(s) + Y4+(aq) ⇌ X4+(aq) + Y(s)  K = 3.90 x 10 5E° = ?V
The reaction below has a cell potential, Eo , of – 1.18 V. Consider the reaction below:Mn2+ (aq) + 2 e –       →      Mn (s) What is the new cell potential for this reaction?3 Mn (s)      →       3 Mn2+ (aq) + 6 e – a. + 1.18 Vb. + 3.54 Vc. + 0.39 Vd. – 3.54 V
What is the standard potential of a cell where X is the anode and Y is the cathode?
Calculate the E°cell for the following equation. Use these standard potentials . ClO4- (aq) + 6H3O+ (aq) + 6Br- (aq)  →  3Br2 (aq) + ClO- (aq) + 9H2O (l)
Calculate the standard reduction potential for the reaction of Cu(III) to Cu(II)Using the following data1. Cu3+ + 2 e- → Cu+          Eo1 = 1.28 V2. Cu2+ + e- → Cu+          Eo2 = 0.15 V3. Cu2+ + 2 e- → Cu(s)      Eo3 = 0.34 V4. Cu+ + e- → Cu(s)         Eo4 = 0.52 V
Consider the following electrochemical cell for the following questions.Fe | FeBr2 (aq. 0.010 M) || KBr (aq, 1.0 M, Br2 (l) | PtWhat is the overall reaction occurring in the cell when current flows?a) Fe(s) + FeBr2(aq) → KBr(aq) + Br2(l)b) Fe(s) + Br2(l) → FeBr2(aq)c) 2 FeBr2(aq) + Br2(l) → 2 FeBr3(aq)d) 2 FeBr3(aq) → 2 FeBr2(aq) + Br2(l)e) FeBr2(aq) → Fe(s) + Br2(l) What is ΔG° of the overall reaction taking place in the cell when current flows?a) 297 kJb) -297 kJc) 149 kJd) -149 kJe) 61.8 kJf) -61.8 kJ
Consider the reaction at 55°C, where [Fe2+] = 3.80 M and [Mg2+] = 0.310 MCalculate the standard cell potential for Mg(s) + Fe2+(aq) → Mg+2(aq) + Fe(s)
Given the following half-reactions and their respective standard reduction potentials calculate the standard reduction potential for the reduction half-reaction of Cu(lll) to Cu(ll).Cu3+ + e- → Cu2+
What statement is NOT true about standard electrode potentials?A) E°cell is positive for spontaneous reactions.B) Electrons will flow from more negative electrode to more positive electrode.C) The electrode potential of the standard hydrogen electrode is exactly zero.D) The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative to the standard hydrogen electrode and therefore has a positive E°.E) E°cell is the difference in voltage between the anode and the cathode, E°a-E°c
Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.)3 Cl2(g) + 2 Fe(s) → 6 Cl-(aq) + 2 Fe3+(aq) Cl2(g) + 2e- → 2 Cl-(aq)                      E° = +1.36 VFe3+(aq) + 3e- → Fe(s)                       E° = -0.04 VA) +4.16 VB) -1.40 VC) -1.32 VD) +1.32 VE) +1.40 V
For the a voltaic cell with the overall reaction:             Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)    E°cell = 1.10 V Given that the standard reduction potential of Zn2+ to Zn(s) is –0.76 V, calculate the standard reduction potential for: Cu2+(aq) + 2e− → Cu(s) A.  0.76 V B.  −0.76 V C.  1.86 V D.  −1.86 V E.  0.34 V
Using the following data, calculate the standard reduction potential for the reaction of Cu(III) to Cu(II).
The standard reduction potential for a substance indicates how readily that substance gains electrons relative to other substances at standard conditions. The more positive the reduction potential, the more easily the substance gains electrons. Consider the following: Fe2+ (aq) + 2e- → Fe(s),         E°red = -0.450 V Cu2+ (aq) + 2e- → Cu(s),         E°red = +0.337 V What is the standard potential, E°cell, for this galvanic cell? Use the given standard reduction potentials in your calculation as appropriate. Express your answer to three decimal places and include the appropriate units.
Calculate the equilibrium constant, K, for the following reaction at 25 °C.Fe3+(aq) + B(s) + 6H 2O(l) → Fe(s)+ H3BO3(s) + 3H3O+(aq)The balanced reduction half-reactions for the above equation and their respective standard reduction potential values (E degree) are as follows:Fe3+(aq) + 3e- → Fe(s)                                           E° = - 0.04 VH3BO3(s) + 3H3O+(aq) + 3e- → B(s) + 6H 2O(l)       E° = - 0.8698 V
Find the emf at 25°C for the following reaction? NO3-(aq) + 8H+(aq) + 3Cu(s) → 2NO(g) + 4H 2O(l) + 3Cu2+(aq).
Calculate ΔGr° for the electrochemical cell, Pt (s) | Fe2+, Fe3+ || Cu2+ | Cu (s), for the production of 1.00 g of Cu, assuming that all concentrations remain at their standard values of 1.00 M throughout the process.a) 83.0 kJb) 41.5 kJc) 3.36 kJd) 1.31 kJe) 0.66 kJ
The cell diagram for the reaction occurring in silver-zinc button batteries is Zn(s) | ZnO(s) | KOH(aq) | Ag2O(s) | Ag(s)a. Write the half-reaction equation that involve silver.b. Write the half-reaction equation that involve zinc.c. Write the balance equation for the net cell reaction equation.d. Calculate the value of E°cell.
You are given the following three half-reactions:(1) Fe3+(aq) + e− ⇌ Fe2+(aq)(2) Fe2+(aq) + 2e− ⇌ Fe(s)(3) Fe3+(aq) + 3e− ⇌ Fe(s)(a) Calculate ΔG° for (1) and (2) from their E°half-cell values.
76Exercise 42. Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine ε° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. H2O2 + 2H + + 2e- → 2H 2O    ε° = 1.78 V   O2 + 2H + + 2e- → H2O2          ε° = 0.68 VCalculate ΔG° and K at 25°C for the reaction above.
You are given the following three half-reactions:(1) Fe3+(aq) + e− ⇌ Fe2+(aq)(2) Fe2+(aq) + 2e− ⇌ Fe(s)(3) Fe3+(aq) + 3e− ⇌ Fe(s)(b) Calculate ΔG° for (3) from (1) and (2).
You are given the following three half-reactions:(1) Fe3+(aq) + e− ⇌ Fe2+(aq)(2) Fe2+(aq) + 2e− ⇌ Fe(s)(3) Fe3+(aq) + 3e− ⇌ Fe(s)(a) Calculate ΔG° for (1) and (2) from their E°half-cell values.(b) Calculate ΔG° for (3) from (1) and (2).(c) Calculate E°half-cell for (3) from its ΔG°.
Use tabulated half-cell potentials to calculate ΔG˚rxn for each of the following reactions at 25 ˚C^circ { m{C}}.Pb2+(aq) + Mg(s) → Pb(s) + Mg2+(aq)
Use tabulated half-cell potentials to calculate ΔG˚rxn for each of the following reactions at 25 ˚C.Br2(l) + 2 Cl–(aq) → 2 Br–(aq) + Cl2(g)
Use tabulated half-cell potentials to calculate ΔG˚rxn for each of the following reactions at 25 ˚C.MnO2(s) + 4 H+(aq) + Cu(s) → Mn2+(aq) + 2 H2O(l) + Cu2+(aq)
The overall reaction and standard cell potential at 25°C for the rechargeable nickel–cadmium alkaline battery isCd(s) + NiO2(s) + 2H2O(l) → Ni(OH)2(s) + Cd(OH)2(s)       E° = 1.10 VFor every mole of Cd consumed in the cell, what is the maximum useful work that can be obtained at standard conditions?
What is the maximum work that can be obtained from a hydrogen–oxygen fuel cell at standard conditions that produces 1.00 kg water at 25°C? Why do we say that this is the maximum work that can be obtained? What are the advantages and disadvantages in using fuel cells rather than the corresponding combustion reactions to produce electricity?
The overall reaction and equilibrium constant value for a hydrogen–oxygen fuel cell at 298 K is2H2 (g) + O2 (g) → 2H2O (l)                    K = 1.28 x 10   83a. Calculate E° and ΔG° at 298 K for the fuel cell reaction.
The overall reaction and equilibrium constant value for a hydrogen–oxygen fuel cell at 298 K is2H2 (g) + O2 (g) → 2H2O (l)                    K = 1.28 x 10   83b. Predict the signs of ΔH° and ΔS° for the fuel cell reaction.
The overall reaction and equilibrium constant value for a hydrogen–oxygen fuel cell at 298 K is2H2 (g) + O2 (g) → 2H2O (l)                    K = 1.28 x 10   83c. As temperature increases, does the maximum amount of work obtained from the fuel cell reaction increase, decrease, or remain the same? Explain.
Calculate ΔG˚rxn and E˚cell E_{{ m{cell}}}^circfor a redox reaction with n = 3 that has an equilibrium constant of K exttip{K}{K}= 26. (Temperature is 298 K m K.)
Consider the reaction: I2(s) + 2 Br–(aq) → 2 I–(aq) + Br2(l); ΔG˚ = 1.1 x 105 J. The calculation revealed that the reaction is not spontaneous. Based on conceptual reasoning, which of the following best explains why I2 does not oxidize Br–?(a) Br is more electronegative than I; therefore, you do not expect Br– to give up an electron to I2.(b) I is more electronegative than Br; therefore, you do not expect I2 to give up an electron to Br–.(c) Br– is in solution and I2 is a solid. Solids do not gain electrons from substances in solution.
The black silver sulfide discoloration of silverware can be removed by heating the silver article in a sodium carbonate solution in an aluminum pan. The reaction is3Ag2S (s) + 2Al (s) ⇌ 6Ag(s) + 3S 2- (aq) + 2Al 3+ (aq)a. Using data in Appendix 4, calculate ΔG°, K, and E° for the above reaction at 25°C. [For Al3+(aq), ΔG°f = -480. kJ/mol.]
76Exercise 38. Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.b. Zn(s) + Ag +(aq) ⇌ Zn 2+(aq) + Ag(s)Calculate ΔG° and K at 25°C for the reaction above.
76Exercise 42. Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine ε° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.b. Mn 2+ + 2e- → Mn                   ε° = —1.18 V    Fe 3+ + 3e- → Fe                     ε ° = —0.036 VCalculate ΔG° and K at 25°C for the reaction above.
Exercise 37. Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.b. Cu 2+(aq) + Mg(s) ⇌ Mg 2+(aq) + Cu(s)Calculate ΔG° and K at 25°C for the reaction above.
Exercise 41. Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine ε° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. Cl2 + 2e- → 2Cl -        ε° = 1.36 V     Br2 + 2e- → 2Br -       ε° = 1.09 VCalculate ΔG° and K at 25°C for the reaction above.
Exercise 41. Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine ε° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.b. MnO4 - + 8H + + 5e-  → Mn 2+ + 4H2O      ε° = 1.51 V     IO4 - + 2H + + 2e- → IO3 - + H2O              ε° = 1.60 VCalculate ΔG° and K at 25°C for the reaction above.
For the standard cell potential given here, determine the ΔG° for the cell in kJ.(a) 0.000 V, n = 2
For the standard cell potential given here, determine the ΔG° for the cell in kJ.(b) +0.434 V, n = 2
For the standard cell potential given here, determine the ΔG° for the cell in kJ.(c) −2.439 V, n = 1
For the ΔG° values given here, determine the standard cell potential for the cell.(a) 12 kJ/mol, n = 3
For the ΔG° values given here, determine the standard cell potential for the cell.(b) −45 kJ/mol, n = 1
Determine ΔG and ΔG° for each of the reactions below.(a) Hg(l) + S2−(aq, 0.10 M) + 2 Ag+(aq, 0.25 M) ⟶ 2 Ag(s) + HgS(s)(b) The galvanic cell made from a half-cell consisting of an aluminum electrode in 0.015 M aluminum nitrate solution and a half-cell consisting of a nickel electrode in 0.25 M nickel(II) nitrate solution.(c) The cell made of a half-cell in which 1.0 M aqueous bromide is oxidized to 0.11 M bromine ion and a half-cell in which aluminum ion at 0.023 M is reduced to aluminum metal. Assume the standard reduction potential for Br2(l) is the same as that of Br2(aq).
Use tabulated half-cell potentials to calculate ΔG˚rxn for each of the following reactions at 25 ˚C^circ { m{C}}.2 Fe3+(aq) + 3 Sn(s) → 2 Fe(s) + 3 Sn2+(aq)
Use tabulated half-cell potentials to calculate ΔG˚rxn for each of the following reactions at 25 ˚C.O2(g) + 2 H2O(l) + 2 Cu(s) → 4 OH–(aq) + 2 Cu2+(aq)
Use tabulated half-cell potentials to calculate ΔG˚rxn for each of the following reactions at 25 ˚C.Br2(l) + 2 I–(aq) → 2 Br–(aq) + I2(s)
Calculate ΔG˚rxn and E˚cell at 25 ˚C for a redox reaction with n exttip{n}{n}= 4 that has an equilibrium constant of K exttip{K}{K}= 5.1×10−2.
S4O62−(aq) + 2I−(aq) ⟶ I2(s) + S2O32−(aq)        ΔG° = 87.8 kJ/mol(b) Calculate E°cell.
S4O62−(aq) + 2I−(aq) ⟶ I2(s) + S2O32−(aq)        ΔG° = 87.8 kJ/mol(c) For the reduction half-reaction, write a balanced equation, give the oxidation number of each element, and calculate E°half-cell.
Hydrogen gas has the potential as a clean fuel in reaction with oxygen. The relevant reaction is2H2 (g) + O2 (g)  →  2H2 O(l).Consider two possible ways of utilizing this reaction as an electrical energy source: (i) Hydrogen and oxygen gases are combusted and used to drive a generator, much as coal is currently used in the electric power industry; (ii) hydrogen and oxygen gases are used to generate electricity directly by using fuel cells that operate at 85 oC.(a) Use data in Appendix C to calculate ΔH° and ΔS° for the reaction. We will assume that these values do not change appreciably with temperature. (b) Based on the values from part (a), what trend would you expect for the magnitude of ΔG for the reaction as the temperature increases?(c) What is the significance of the change in the magnitude of G with temperature with respect to the utility of hydrogen as a fuel (recall Equation 19.19 in the text book)?
Calculate ΔG˚rxn K for each of the following reactions.The disproportionation of Mn2+( aq ) to Mn(s) and MnO2(s) in acid solution at 25 ˚C.
Calculate ΔG˚rxn for each of the following reactions.The disproportionation of MnO2(s) to Mn2+(aq) and MnO4–(aq) in acid solution at 25 ˚C [The reduction potential of MnO4–(aq) to MnO2(s) is 1.68 V.]
A voltaic cell has one half-cell with a Cu bar in a 1.00 M Cu  2+ salt, and the other half-cell with a Cd bar in the same volume of a 1.00 M Cd2+ salt.(a) Find E°cell, ΔG°, and K.
Calculate ΔG˚rxn K for each of the following reactions.The reaction of Cr2+(aq) with Cr2O72–(aq) in acid solution to form Cr3+(aq).
Calculate ΔG˚rxn for each of the following reactions.The reaction of Cr3+(aq) and Cr(s) to form Cr2+(aq). [The reduction potential of Cr2+(aq) to Cr(s) is –0.91 V.]
Given the following reduction half-reactions:Fe3+(aq) + e-  →  Fe2+(aq)Eredo= + 0.77VS2O62-(aq) + 4H+(aq) + 2e-  →  2H2SO3(aq)Eredo= + 0.60VN2O(g) + 2H+(aq) + 2e-  →  N2(g) + H2O(l)Eredo= - 1.77VVO2+(aq) + 2H+(aq) + e-  →  VO2+(aq) + H2O(l)Eredo= + 1.00VCalculate ΔG° for each reaction at 298 K.{ m K}
The cell Pt(s) | Cu+(1 M), Cu2+(1 M) ‖ Cu+(1 M) | Cu(s) has E˚ = 0.364 V. The cell Cu(s) | Cu2+(1 M) ‖ Cu+(1 M) | Cu(s) has E˚ = 0.182 V. Calculate ΔG˚ for the first cell reaction.
The cell Pt(s) | Cu+(1 M), Cu2+(1 M) ‖ Cu+(1 M) | Cu(s) has E˚ = 0.364 V. The cell Cu(s) | Cu2+(1 M) ‖ Cu+(1 M) | Cu(s) has E˚ = 0.182 V. Calculate ΔG˚ for the second cell reaction.
A redox reaction has an equilibrium constant of K = 0.055. What is true of ΔG˚rxn and E˚cell for this reaction?
Use tabulated electrode potentials to calculate ΔG˚ for the reaction 2 Na(s) + 2 H2O(l) → H2(g) + 2 OH–(aq) + 2 Na+(aq).2{ m Na}(s)+2{ m H_2O}(l) ightarrow{ m H_2}(g)+2{ m OH^-}(aq)+2{ m Na^+}(aq)
A redox reaction has an equilibrium constant of K = 1.2 x 103. Which statement is true regarding ΔG˚rxn and E˚cell for this reaction?(a) E˚cell is positive and ΔG˚rxn is positive.(b) E˚cell is negative and ΔG˚rxn is negative.(c) E˚cell is positive and ΔG˚rxn is negative.(d) E˚cell is negative and ΔG˚rxn is positive.
If the equilibrium constant for a two-electron redox reaction at 298 K is 1.6×10−4, calculate the corresponding ΔG° and E° and E_{ m cel}^circ under standard conditions.
Chlorine dioxide (ClO2), which is produced by the reaction 2NaClO2 (aq) + Cl2 (g) → 2ClO2 (g) + 2NaCl (aq)has been tested as a disinfectant for municipal water treatment. Using data from Table 17‑1, calculate E° and ΔG° at 25°C for the production of ClO2.
The amount of manganese in steel is determined by changing it to permanganate ion. The steel is first dissolved in nitric acid, producing Mn2+ ions. These ions are then oxidized to the deeply colored MnO4- ions by periodate ion (IO4-) in acid solution.b. Calculate E° and ΔG° at 25°C for each reaction.
If the equilibrium constant for a one-electron redox reaction at 298 K is 7.3x104, calculate the corresponding ΔG° and E°.
Exercise 45. Consider the following galvanic cells:For each galvanic cell, give the balanced cell equation and determine E°. Standard reduction potentials are found in Table 17‑1.Calculate the maximum amount of work that can be obtained from the galvanic cells at standard conditions in Exercise 45.
Exercise 46. Give the balanced cell equation and determine E° for the galvanic cells based on the following half-reactions. Standard reduction potentials are found in Table 17‑1.a. Cr2O72- + 14H+ + 6e- → 2Cr 3+ + 7H2O    H2O2 + 2H+ + 2e- → 2H2Ob. 2H + + 2e- → H2    Al 3+ + 3e- → AlCalculate the maximum amount of work that can be obtained from the galvanic cells at standard conditions in Exercise 46.
The equation ΔG° = -nFE° also can be applied to half-reactions. Use standard reduction potentials to estimate ΔG°f for Fe2+(aq) and Fe3+(aq). (ΔG°f for e- = 0.)
Consider a redox reaction for which E  is a negative number.What is the sign of ΔG°  for the reaction?
Exercise 37. Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. Cr 3+(aq) + Cl 2 (g) ⇌ Cr 2O 7 2-(aq) + Cl  -(aq)Calculate ΔG° and K at 25°C for the reaction above.
76Exercise 38. Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. IO3-(aq) + Fe 2+(aq) ⇌ Fe 3+(aq) + I 2 (aq)Calculate ΔG° and K at 25°C for the reaction above.
Consider the galvanic cell based on the following half-reaction:Zn2+ + 2e- → Zn      ε° = —0.76 VFe2+ + 2e- → Fe      ε° = —0.44 Vb. Calculate ΔG° and K for the cell reaction at 25°C.
Consider the galvanic cell based on the following half-reaction:Au3+ + 3e- → Au           ε° = 1.50 V      Tl + + e- → Tl            ε° = —0.34 Vb. Calculate ΔG° and K for the cell reaction at 25°C.
Consider the following reaction.Mg(s) + SO42-(aq) + 4 H+(aq)  →  Mg2+(aq) + H2SO3(aq) + 2 H2O(l), E˚cell = 2.57 VCalculate the ΔG˚ for the reaction.
A disproportionation reaction involves a substance that acts as both an oxidizing and a reducing agent, producing higher and lower oxidation states of the same element in the products. Which of the following disproportionation reactions are spontaneous under standard conditions? Calculate ΔG° and K at 25°C for those reactions that are spontaneous under standard conditions.a. 2Cu+(aq) → Cu 2+(aq) + Cu(s)b. 3Fe2+(aq) → 2Fe 3+(aq) + Fe(s)c. HClO2(aq) → ClO3 -(aq) + HClO(aq) (unbalanced)Use the half-reactions:ClO3- + 3H +  + 2e- → HClO2 + H 2O         ε° = 1.21 VHClO2 + 2H + + 2e- → HClO + H 2O         ε° = 1.65 V
Consider a galvanic cell based on the following theoretical half-reactions:What is the value of ΔG° and K for this cell?
The following reactions are used in batteries:I. 2H2(g) + O2(g) ⟶ 2H2O(l)                                                    Ecell = 1.23 VII. Pb(s) + PbO2(s) + 2H2SO4(aq) ⟶ 2PbSO4(s) + 2H2O(l)     Ecell = 2.04 VIII. 2Na(l) + FeCl2(s) ⟶ 2NaCl(s) + Fe(s)                                Ecell = 2.35 VReaction I is used in fuel cells, II in the automobile lead-acid battery, and III in an experimental high-temperature battery for powering electric vehicles. The aim is to obtain as much work as possible from a cell, while keeping its weight to a minimum. (a) In each cell, find the moles of electrons transferred and ΔG.
The following reactions are used in batteries:I. 2H2(g) + O2(g) ⟶ 2H2O(l)                                                    Ecell = 1.23 VII. Pb(s) + PbO2(s) + 2H2SO4(aq) ⟶ 2PbSO4(s) + 2H2O(l)     Ecell = 2.04 VIII. 2Na(l) + FeCl2(s) ⟶ 2NaCl(s) + Fe(s)                                Ecell = 2.35 VReaction I is used in fuel cells, II in the automobile lead-acid battery, and III in an experimental high-temperature battery for powering electric vehicles. The aim is to obtain as much work as possible from a cell, while keeping its weight to a minimum. (b) Calculate the ratio, in kJ/g, of wmax to mass of reactants for each of the cells. Which has the highest ratio, which the lowest, and why? (Note: For simplicity, ignore the masses of cell components that do not appear in the cell as reactants, including electrode materials, electrolytes, separators, cell casing, wiring, etc.)
A voltaic cell utilizes the following reaction:2Fe3+(aq) + H2(g)  →  2Fe2+(aq) + 2H+(aq).What is the emf of this cell under standard conditions?
Calculate E˚cell for each of the following balanced redox reactions.2 Cu(s) + Mn2+(aq) → 2 Cu+(aq) + Mn(s)
Calculate E˚cell for each of the following balanced redox reactions.MnO2(s) + 4 H+(aq) + Zn(s) → Mn2+(aq) + 2 H2O(l) + Zn2+(aq)
Calculate E˚cell for each of the following balanced redox reactions.Cl2(g) + 2F–(aq) → F2(g) + 2 Cl–(aq)
Heart pacemakers are often powered by lithium-silver chromate "button" batteries. The overall cell reaction is:2Li(s) + Ag2CrO4(s)  →  Li2CrO4(s) + 2Ag(s)Choose the two half-reactions from Appendix E that most closely approximate the reactions that occur in the battery. What standard emf would be generated by a voltaic cell based on these half-reactions?
What are E°cell and ΔG° of a redox reaction at 25°C for which n = 1 and K = 5.0 × 104?
Determine whether the reaction is spontaneous as written.2 Cu(s) + Mn2+(aq) → 2 Cu+(aq) + Mn(s)
What are E°cell and ΔG° of a redox reaction at 25°C for which n = 1 and K = 5.0 × 10-6?
Determine whether the reaction is spontaneous as written.MnO2(s) + 4 H+(aq) + Zn(s) → Mn2+(aq) + 2 H2O(l) + Zn2+(aq)
What are E°cell and ΔG° of a redox reaction at 25°C for which n = 2 and K = 65?
Determine whether the reaction is spontaneous as written.Cl2(g) + 2F–(aq) → F2(g) + 2 Cl–(aq)
What are E°cell and ΔG° of a redox reaction at 25°C for which n = 2 and K = 0.065?
Mercuric oxide dry-cell batteries are often used where a high-energy density is required, such as in watches and cameras. The two half-cell reactions that occur in the battery arel HgO(s) + H2 O(l) + 2e-  →  Hg(l) + 2OH- (aq)Zn(s) + 2OH- (aq)  →  ZnO(s) + H2 O(l) + 2e-.The value of E°red for the cathode reaction is +0.098 V. The overall cell potential is +1.35 V. Assuming that both half-cells operate under standard conditions, what is the standard reduction potential for the anode reaction?
For a spontaneous reactionA(aq) + B(aq)  →  A-(aq) + B+(aq)answer the following questions:(a) If you made a voltaic cell out of this reaction, what half reaction would be occurring at the cathode, and what half-reaction would be occurring at the anode?(b) Which half-reaction from part above is higher in potential energy?
A fuel cell designed to react grain alcohol with oxygen has the following net reaction:C2H5OH (l) + 3O2 (g) → 2CO2 (g) + 3H2O (l)The maximum work that 1 mole of alcohol can do is 1.32 x 10 3 kJ. What is the theoretical maximum voltage this cell can achieve at 25°C?
For a spontaneous reactionA(aq) + B(aq)  →  A-(aq) + B+(aq)answer the following questions:What is the sign of E°cell?
A voltaic cell employs the following redox reaction: Sn2+(aq) + Mn(s) → Sn(s) + Mn2+(aq)Calculate the cell potential at 25˚C under each of the following conditions: standard conditions
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(a) Mg(s) + Ni2+(aq) ⟶ Mg2+(aq) + Ni(s)
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(b) 2Ag+(aq) + Cu(s) ⟶ Cu 2+(aq) + 2Ag(s)
Assume that you want to construct a voltaic cell that uses the following half reactions:A2 +  ( aq ) + 2e-  →  A( s )Eredo= -0.13VB2 +  ( aq ) + 2e-  →  A( s )Eredo= -1.12VYou begin with the incomplete cell pictured below, in which the electrodes are immersed in water.What additions must you make to the cell for it to generate a standard emf?
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(c) Mn(s) + Sn(NO3 )2(aq) ⟶ Mn(NO3 )2(aq) + Sn(s)
A voltaic cell that uses the reactionTl3 +  ( aq ) + 2 Cr2 +  ( aq )  →  Tl+ ( aq ) + 2Cr3 +  ( aq )has a measured standard cell potential of +1.19 V.By using data from Appendix E, determine E°red for the reduction of Tl3+(aq) to Tl+(aq).{ m Tl^+}({ m aq}
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(d) 3Fe(NO3 )2(aq) + Au(NO3 )3(aq) ⟶ 3Fe(NO3 )3(aq) + Au(s)
The black silver sulfide discoloration of silverware can be removed by heating the silver article in a sodium carbonate solution in an aluminum pan. The reaction is3Ag2S (s) + 2Al (s) ⇌ 6Ag(s) + 3S 2- (aq) + 2Al 3+ (aq)b. Calculate the value of the standard reduction potential for the following half-reaction:2e- + Ag2S (s) → 2Ag (s) + S 2- (aq)
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(a) Mn(s) + Ni2+(aq) ⟶ Mn2+(aq) + Ni(s)
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(b) 3Cu2+(aq) + 2Al(s) ⟶ 2Al 3+(aq) + 3Cu(s)
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(c) Na(s) + LiNO 3(aq) ⟶ NaNO 3(aq) + Li(s)
Calculate Eocell for the following balanced redox reaction.O2(g) + 2H2O(l) + 4Ag(s) → 4OH−(aq) + 4Ag+(aq)
For the following reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard condition.(d) Ca(NO3 )2(aq) + Ba(s) ⟶ Ba(NO3 )2(aq) + Ca(s)
A voltaic cell that uses the reactionPdCl42-(aq) + Cd(s)   →   Pd(s) + 4Cl-(aq) + Cd2+(aq)has a measured standard cell potential of +1.03 V.By using data from Appendix E, determine E°red for the reaction involving Pd{ m Pd}.
Determine the overall reaction and its standard cell potential at 25 °C for this reaction. Is the reaction spontaneous at standard conditions?Cu(s) │ Cu2+(aq) ║ Au3+(aq) │ Au(s)
Use data from the table above to calculate E˚cell for the reaction: 2 ClO2(g) + Pb(s) → 2 ClO2–(aq) + Pb2+(aq)NO3–(aq) + 4 H+(aq) + 3 e– →  NO(g) + 2 H2O(l)E˚ = 0.96 VClO2(g) + e– →  ClO2–(aq)E˚= 0.95 VCu2+(aq) +2 e– →  Cu(s)E˚= 0.34 V2 H+(aq) +2 e– →  H2(g)E˚= 0.00 VPb2+(aq) +2 e–  →  Pb(s)E˚ = –0.13 VFe2+(aq) +2 e– →  Fe(s)E˚ = –0.45 V
Determine the overall reaction and its standard cell potential at 25 °C for the reaction involving the galvanic cell made from a half-cell consisting of a silver electrode in 1 M silver nitrate solution and a half-cell consisting of a zinc electrode in 1 M zinc nitrate. Is the reaction spontaneous at standard conditions?
Consider the standard galvanic cell based on the following half-reactions:Cu2+ + 2e- → CuAg+ + e- → AgThe electrodes in this cell are Ag(s) and Cu(s). Does the cell potential increase, decrease, or remain the same when the following changes occur to the standard cell?e. The silver electrode is replaced with a platinum electrode.Pt2+ + 2e- → Pt        E° = 1.19 V
Calculate Eocell for the following balanced redox reaction.Br2(l) + 2I−(aq) → 2Br−(aq) + I2(s)Express your answer using two significant figures.
Determine the overall reaction and its standard cell potential at 25 °C for the reaction involving the galvanic cell in which cadmium metal is oxidized to 1 M cadmium(II) ion and a half-cell consisting of an aluminum electrode in 1 M aluminum nitrate solution. Is the reaction spontaneous at standard conditions?
Determine the overall reaction and its standard cell potential at 25 °C for these reactions. Is the reaction spontaneous at standard conditions? Assume the standard reduction for Br 2(l) is the same as for Br2(aq).Pt(s) │ H2(g) │ H+(aq) ║ Br2(aq), Br−(aq) │ Pt(s)
The saturated calomel electrode, abbreviated SCE, is often used as a reference electrode in making electrochemical measurements. The SCE is composed of mercury in contact with a saturated solution of calomel (Hg2Cl2). The electrolyte solution is saturated KCl. ESCE is +0.242 V relative to the standard hydrogen electrode. Calculate the potential for each of the following galvanic cells containing a saturated calomel electrode and the given half-cell components at standard conditions. In each case, indicate whether the SCE is the cathode or the anode. Standard reduction potentials are found in Table 17‑1.a. Cu2+ + 2e- → Cu
The saturated calomel electrode, abbreviated SCE, is often used as a reference electrode in making electrochemical measurements. The SCE is composed of mercury in contact with a saturated solution of calomel (Hg2Cl2). The electrolyte solution is saturated KCl. ESCE is +0.242 V relative to the standard hydrogen electrode. Calculate the potential for each of the following galvanic cells containing a saturated calomel electrode and the given half-cell components at standard conditions. In each case, indicate whether the SCE is the cathode or the anode. Standard reduction potentials are found in Table 17‑1.b. Fe3+ + e- → Fe2+
The standard reduction potentials of the following half-reactions are given in Appendix E in the textbook:Ag+ (aq) + e-  →  Ag(s)Cu2 +  (aq) + 2e-  →  Cu(s)Ni2 +  (aq) + 2e-  →  Ni(s)Cr3 +  (aq) + 3e-  →  Cr(s).Calculate the value of this emf.
Calculate Eocell for the following balanced redox reaction.PbO2(s) + 4H+(aq) + Sn(s) → Pb2+(aq) + 2H2O(l) + Sn2+(aq)Express your answer using two significant figures.
The saturated calomel electrode, abbreviated SCE, is often used as a reference electrode in making electrochemical measurements. The SCE is composed of mercury in contact with a saturated solution of calomel (Hg2Cl2). The electrolyte solution is saturated KCl. ESCE is +0.242 V relative to the standard hydrogen electrode. Calculate the potential for each of the following galvanic cells containing a saturated calomel electrode and the given half-cell components at standard conditions. In each case, indicate whether the SCE is the cathode or the anode. Standard reduction potentials are found in Table 17‑1.c. AgCl + e- → Ag + Cl -
The saturated calomel electrode, abbreviated SCE, is often used as a reference electrode in making electrochemical measurements. The SCE is composed of mercury in contact with a saturated solution of calomel (Hg2Cl2). The electrolyte solution is saturated KCl. ESCE is +0.242 V relative to the standard hydrogen electrode. Calculate the potential for each of the following galvanic cells containing a saturated calomel electrode and the given half-cell components at standard conditions. In each case, indicate whether the SCE is the cathode or the anode. Standard reduction potentials are found in Table 17‑1.d. Al 3+  + 3e- → Al
The standard reduction potentials of the following half-reactions are given in Appendix E in the textbook:Ag+ (aq) + e-  →  Ag(s)Cu2 +  (aq) + 2e-  →  Cu(s)Ni2 +  (aq) + 2e-  →  Ni(s)Cr3 +  (aq) + 3e-  →  Cr(s).Determine which combination of these half-cell reactions leads to the cell reaction with the largest positive cell emf.
The saturated calomel electrode, abbreviated SCE, is often used as a reference electrode in making electrochemical measurements. The SCE is composed of mercury in contact with a saturated solution of calomel (Hg2Cl2). The electrolyte solution is saturated KCl. ESCE is +0.242 V relative to the standard hydrogen electrode. Calculate the potential for each of the following galvanic cells containing a saturated calomel electrode and the given half-cell components at standard conditions. In each case, indicate whether the SCE is the cathode or the anode. Standard reduction potentials are found in Table 17‑1.e. Ni2+ + 2e- → Ni
Use the tabulated values of ΔG˚f in Appendix IIB in the textbook to calculate E˚cellE_{{ m{cell}}}^circ for a fuel-cell that employs the reaction between methane gas { m{(CH}}_4 ) and oxygen to form carbon dioxide and gaseous water.
A common shorthand way to represent a voltaic cell isanode | anode  solution || cathode  solution | cathodeA double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change in phase, such as from solid to solution.You may want to reference (Pages 848 - 899)Chapter 20 while completing this problem.Calculate the standard cell emf for the reaction represented by Fe | Fe2+ || Ag+ | Ag.
A common shorthand way to represent a voltaic cell isanode | anode  solution || cathode  solution | cathodeA double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change in phase, such as from solid to solution.You may want to reference (Pages 848 - 899)Chapter 20 while completing this problem.Calculate the standard cell emf for the reaction represented by Zn | Zn2+ || H+ | H2.
A common shorthand way to represent a voltaic cell isanode | anode  solution || cathode  solution | cathodeA double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change in phase, such as from solid to solution.You may want to reference (Pages 848 - 899)Chapter 20 while completing this problem.For the following reaction:ClO3-(aq) + 3 Cu(s) + 6 H+(aq) →  Cl-(aq) + 3 Cu2+(aq) + 3 H2O(l)Pt is used as an inert electrode in contact with the ClO3- and Cl-. Calculate the standard emf using data in Appendix E and given the following:ClO3-(aq) + 6 H+(aq) + 6 e-  →  Cl-(aq) + 3 H2O(l);       Eo= 1.45 V
Consider a battery made from one half-cell that consist of a copper electrode in 1 M CuSO4 solution and another half-cell that consist of a lead electrode in 1 M Pb(NO  3)2 solution.(b) What is the standard cell potential for the battery?
Consider a battery made from one half-cell that consist of a copper electrode in 1 M CuSO4 solution and another half-cell that consist of a lead electrode in 1 M Pb(NO  3)2 solution.(c) Most devices designed to use dry-cell batteries can operate between 1.0 and 1.5 V. Could this cell be used to make a battery that could replace a dry-cell battery? Why or why not.
A voltaic cell consists of a strip of cadmium metal in a solution of Cd ( NO3 )2 in one beaker, and in the other beaker a platinum electrode is immersed in a NaCl solution, with Cl2 gas bubbled around the electrode. A salt bridge connects the two beakers. Use Appendix E in the textbook to find the standard reduction potentials.What is the emf generated by the cell under standard conditions?
A voltaic cell based on the reaction 2 Eu2+(aq) + Ni2+(s)  →  2 Eu3+(aq) + Ni(s) generates Ecell = 0.07 V. Given the standard reduction potential of Ni2+ (Ered = -0.28 V) what is the standard reduction potential for the reaction Eu3+(aq) + e−  →  Eu2+(aq)?
Which of the following redox reactions do you expect to occur spontaneously in the reverse direction? [Hint: The reactions are occurring under standard conditions (1 M for the aqueous ions)]a. Ca2+(aq) + Zn(s) → Ca(s) + Zn2+(aq)b. 2 Ag+(aq) + Ni(s) → 2 Ag(s) + Ni2+(aq)c. Fe(s) + Mn2+(aq) → Fe2+(aq) + Mn(s)d. 2 Al(s) + 3 Pb2+(aq) → 2 Al3+(aq) + 3 Pb(s)
Calculate E˚cell for each of the following balanced redox reactions.O2(g) + 2 H2O(l) + 4 Ag(s) → 4 OH–(aq) + 4 Ag+(aq)
The Haber process is the principal industrial route for converting nitrogen into ammonia:N2 (g) + 3H2 (g)  →  2NH3 (g).Using the thermodynamic data in Appendix C, calculate the equilibrium constant for the process at room temperature.Calculate the standard emf of the Haber process at room temperature.
Calculate E˚cell for each of the following balanced redox reactions.Br2(l) + 2 I–(aq) → 2 Br–(aq) + I2(s)
Calculate E˚cell for each of the following balanced redox reactions.PbO2(s) + 4 H+(aq) + Sn(s) → Pb2+(aq) + 2 H2O(l) + Sn2+(aq)
Calculate E˚cell for each of the following balanced redox reactions. Which of the reactions are spontaneous as written.a. O2(g) + 2 H2O(l) + 4 Ag(s) → 4 OH–(aq) + 4 Ag+(aq)b. Br2(l) + 2 I–(aq) → 2 Br–(aq) + I2(s)c. PbO2(s) + 4 H+(aq) + Sn(s) → Pb2+(aq) + 2 H2O(l) + Sn2+(aq)
Using data in Appendix E in the textbook, calculate the standard emf for each of the following reactions.H2(g) + F2(g)   →   2H+(aq) + 2F-(aq)
Using data in Appendix E in the textbook, calculate the standard emf for each of the following reactions.Cu2+(aq) + Ca(s)   →   Cu(s) + Ca2+(aq)
Both Ti and V are reactive enough to displace H 2 from water; of the two metals, Ti is the stronger reducing agent. The difference in their E°half-cell values is 0.43 V. Given V(s) + Cu2+(aq) ⟶ V2+(aq) + Cu(s)             ΔG° = −298 kJ/moluse Appendix D to calculate the E°half-cell values for V2+/V and Ti2+/Ti half-cells.
Using data in Appendix E in the textbook, calculate the standard emf for each of the following reactions.3Fe2+(aq)   →   Fe(s) + 2Fe3+(aq)
A voltaic cell is based on Ag+(aq)/Ag(s) and Fe3+(aq)/Fe2+(aq) half-cells.Use S° values in Appendix C and the relationship between cell potential and free-energy change to predict whether the standard cell potential increases or decreases when the temperature is raised above 25°C^circ { m C}.
Using data in Appendix E in the textbook, calculate the standard emf for each of the following reactions.2ClO3-(aq) + 10Br-(aq) + 12H+(aq)   →   Cl2(g) + 5Br2(l) + 6H2O(l)
Calculate the standard cell potential for each of the following electrochemical cells. a. Ni2+(aq) + Mg(s) → Ni(s) + Mg 2+(aq) b. 2H+(aq) + Fe(s) → H 2(g) + Fe2+(aq) c. 2NO-3(aq) + 8H+(aq) + 3Cu(s) → 2NO(g) + 4H 2O(l) + 3Cu2+(aq) Express your answer using two significant figures.
A voltaic cell employs the following redox reaction: 2 Fe3+(aq) + 3 Mg(s)  →  2 Fe(s) + 3 Mg2+(aq)Calculate the cell potential at 25 ˚C under each of the following conditions: standard conditions
Two voltaic cells are to be joined so that one will run the other as an electrolytic cell. In the first cell, one half-cell has Au foil in 1.00 M Au(NO3)3, and the other half-cell has a Cr bar in 1.00 M Cr(NO3)3. In the second cell, one half-cell has a Co bar in 1.00 M Co(NO  3)2, and the other half-cell has a Zn bar in 1.00 M Zn(NO3)2.(a) Calculate E°cell for each cell.
Two voltaic cells are to be joined so that one will run the other as an electrolytic cell. In the first cell, one half-cell has Au foil in 1.00 M Au(NO3)3, and the other half-cell has a Cr bar in 1.00 M Cr(NO3)3. In the second cell, one half-cell has a Co bar in 1.00 M Co(NO  3)2, and the other half-cell has a Zn bar in 1.00 M Zn(NO3)2.(b) Calculate the total potential if the two cells are connected as voltaic cells in series.
The Ksp of Zn(OH)2 is 1.8×10−14. Find E˚cell for the following half-reaction: Zn(OH)2(s) + 2e– ⇌ Zn(s) + 2 OH–(aq)
The Ksp value for PbS(s) is 8.0 x 10 - 28.By using this value together with an electrode potential from Appendix E in the textbook, determine the value of the standard reduction potential for the reactionPbS(s) + 2e-  →  Pb(s) + S2 -  (aq).
The cell Pt(s) | Cu+(1 M), Cu2+(1 M) ‖ Cu+(1 M) | Cu(s) has E˚ = 0.364 V. The cell Cu(s) | Cu2+(1 M) ‖ Cu+(1 M) | Cu(s) has E˚ = 0.182 V. Explain the differences in E˚.
Consider the following galvanic cell:b. Determine the standard cell potential.
The Zn/Zn2+ electrode has a standard electrode potential of E˚ = –0.76 V. How does the relative potential energy of an electron at the Zn/Zn2+ electrode compare to the potential energy of an electron at the standard hydrogen electrode?
Exercise 37. Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. Cr3+ (aq) + Cl2 (g) ⇌ Cr2O72- (aq) + Cl - (aq)b. Cu2+ (aq) + Mg (s) ⇌ Mg 2+ (aq) + Cu (s)Calculate E° values for the galvanic cells in Exercise 37.
For each of the following reactions, write a balanced equation, calculate the standard emf, calculate ΔG°  at 298 K, and calculate the equilibrium constant K at 298 K.Aqueous iodide ion is oxidized to I2(s) by Hg22+(aq).
For each of the following reactions, write a balanced equation, calculate the standard emf, calculate ΔG°  at 298 K, and calculate the equilibrium constant K at 298 K.In acidic solution, copper (I) ion is oxidized to copper (II) ion by nitrate ion.
Exercise 38. Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. IO3- (aq) + Fe2+ (aq) ⇌ Fe3+ (aq) + I 2 (aq)b. Zn (s) + Ag+ (aq) ⇌ Zn2+ (aq) + Ag (s)Calculate E° values for the galvanic cells in Exercise 38.
For each of the following reactions, write a balanced equation, calculate the standard emf, calculate ΔG°  at 298 K, and calculate the equilibrium constant K at 298 K.In basic solution, Cr(OH)3(s) is oxidized to CrO42-(aq) by ClO-(aq).
Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine E° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. Cl2 + 2e- → 2Cl -                E° = 1.36 V    Br2 + 2e- → 2Br -                E° = 1.09 V
Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine E° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.b. MnO4- + 8H+ + 5e- → Mn2+ + 4H2O      E° = 1.51 V    IO4- + 2H+ + 2e- → IO3- + H2O              E° = 1.60 V
Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine E° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.a. H2O2 + 2H+ + 2e- → 2H2O        E° = 1.78 V    O2 + 2H+ + 2e- → H2O2             E° = 0.68 V
An electrode has a negative electrode potential. Which statement is correct regarding the potential energy of an electron at this electrode?(a) An electron at this electrode has a lower potential energy than it has at a standard hydrogen electrode.(b) An electron at this electrode has a higher potential energy than it has at a standard hydrogen electrode.(c) An electron at this electrode has the same potential energy as it has at a standard hydrogen electrode.
Sketch the galvanic cells based on the following half-reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine E° for the galvanic cells. Assume that all concentrations are 1.0 M and that all partial pressures are 1.0 atm.b. Mn2+ + 2e- → Mn          E° = -1.18 V    Fe3+ + 3e- → Fe            E° = -0.036 V
Consider the following galvanic cells:For each galvanic cell, give the balanced cell equation and determine E°. Standard reduction potentials are found in Table 17‑1.
Consider the following galvanic cells:For each galvanic cell, give the balanced cell equation and determine E°. Standard reduction potentials are found in Table 17‑1.
For each of the following reactions, write a balanced equation, calculate the standard emf, calculate ΔG°  at 298 K, and calculate the equilibrium constant K at 298 K.Aqueous iodide ion is oxidized to I2(s) by Hg22+(aq)
Give the balanced cell equation and determine E° for the galvanic cells based on the following half-reactions. Standard reduction potentials are found in Table 17‑1.a. Cr2O72- + 14H+ + 6e- → 2Cr 3+ + 7H2O    H2O2 + 2H+ + 2e- → 2H2O
Give the balanced cell equation and determine E° for the galvanic cells based on the following half-reactions. Standard reduction potentials are found in Table 17‑1.b. 2H + + 2e- → H2    Al 3+ + 3e- → Al
Calculate E° values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations. Standard reduction potentials are found in Table 17‑1.a. MnO4- (aq) + I - (aq) → I2 (aq) + Mn2+ (aq)
Calculate E° values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations. Standard reduction potentials are found in Table 17‑1.b. MnO4- (aq) + F - (aq) → F2 (aq) + Mn2+ (aq)
Calculate E° values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations that are not already balanced. Standard reduction potentials are found in Table 17‑1.a. H2 (g) → H+ (aq) + H - (aq)
Calculate E° values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations that are not already balanced. Standard reduction potentials are found in Table 17‑1.b. Au3+ (aq) + Ag (s) → Ag + (aq) + Au (s)
Estimate E° for the half-reaction2H2O + 2e- → H2 + 2OH -given the following values of ΔG°f:H2O (l)    = -237 kJ/molH2 (g)      = 0.0OH- (aq)  = -157 kJ/mole-             = 0.0Compare this value of E° with the value of E° given in Table 17‑1.
What is the definition of the standard cell potential (E˚cell)? What does a large positive standard cell potential imply about the spontaneity of the redox reaction occurring in the cell? What does a negative standard cell potential imply about the reaction?
Glucose is the major fuel for most living cells. The oxidative breakdown of glucose by our body to produce energy is called respiration. The reaction for the complete combustion of glucose isC6H12O6 (s) + 6O2 (g) → 6CO2 (g) + 6H2O (l)If this combustion reaction could be harnessed as a fuel cell, calculate the theoretical voltage that could be produced at standard conditions. (Hint: Use ΔG°f values from Appendix 4.)
Direct methanol fuel cells (DMFCs) have shown some promise as a viable option for providing “green” energy to small electrical devices. Calculate E° for the reaction that takes place in DMFCs:CH3OH (l) + 3/2O2 (g) → CO2 (g) + 2H2O (l)Use values of ΔG°f from Appendix 4.
How is the cell potential of an electrochemical cell (E˚cell) related to the potentials of the half-cells?
Explain why E˚cell, ΔG˚rxn, and K are all interrelated.
Will a redox reaction with a small equilibrium constant (K < 1) have a positive or a negative E˚cell? Will it have a positive or a negative ΔG˚rxn?
Electrolytic cell with an active metal electrode. Nickel dissolves from the anode to form Ni2 + ( aq). At the cathode Ni2 + (aq) is reduced and forms a nickel "plate" on the steel cathode. Standard Reduction Potentials in Water at 25°C Ered (V)Reduction Half-Reaction-0.28Ni2+(aq)+2e- → (s)+1.59NiO2(s)+4(aq)+2e- → 2+(aq)+2-(aq)+0.442(s)+2(l)+2e- → 2(s)+2-(aq)What E for this cell?
At 298 K a cell reaction has a standard emf of +0.18 V. The equilibrium constant for the cell reaction is 5.6 x 105.What is the value of n for the cell reaction?
Consider the reaction shown here occurring at 25 ˚C:Cr(s) + Cd2+(aq) →  Cr2+(aq) + Cd(s)Complete the table.[Cd2+][Cr2+]QEcellΔGrxn1.001.001.001.00 x 10-51.0 x 10-51.004.18x10-41.00Note: Cr(s)  →  Cr2+(aq) + 2 e-, Ered = –0.91 VDetermine the value of E˚cell of the reaction.
Consider the reaction shown here occurring at 25 ˚C:Cr(s) + Cd2+(aq) →  Cr2+(aq) + Cd(s)Complete the table.[Cd2+][Cr2+]QEcellΔGrxn1.001.001.001.00 x 10-51.0 x 10-51.004.18x10-41.00Note: Cr(s)  →  Cr2+(aq) + 2 e-, Ered = –0.91 VDetermine Q, Ecell, and ΔGrxn for the first row of the table.
For the reaction3 Ni2+(aq) + 2 Cr(OH)3(s) + 10 OH- (aq)  →  3 Ni(s) + 2 CrO42-(aq) + 8 H2O(l)   G = +87 kJ/molGiven the standard reduction potential of the half-reaction Ni2+(aq) + 2 e-  →  Ni(s) Ered = -0.28 V, calculate the standard reduction potential of the half-reactionCrO42-(aq) + 4 H2O(l) + 3 e-  →  Cr(OH)3(s) + 5 OH-(aq)
Consider the reaction shown here occurring at 25 ˚C:A(s) + B2+(aq) →  A2+(aq) + B(s),;;; ΔG˚rxn = –14.0 kJComplete the table.[B2+][A2+]QEcellΔGrxn1.001.001.001.00 x 10-41.0 x 10-41.03.54 x 10-31.0Determine the value of E˚cell of the reaction.
Consider the reaction shown here occurring at 25 ˚C:A(s) + B2+(aq) →  A2+(aq) + B(s), ΔG˚rxn = –14.0 kJComplete the table.[B2+][A2+]QEcellΔGrxn1.001.001.001.00 x 10-41.0 x 10-41.03.54 x 10-31.0Determine Q, Ecell, and ΔGrxn for the first row of the table.
Consider the galvanic cell based on the following half-reaction:Zn2+ + 2e- → Zn      ε° = —0.76 VFe2+ + 2e- → Fe      ε° = —0.44 Va. Determine the overall cell reaction and calculate ε°cell.
In basic solution, Se2− and SO32− ions react spontaneously:2Se2−(aq) + 2SO32−(aq) + 3H2O(l) ⟶ 2Se(s) + 6OH−(aq) + S2O32−(aq)       E°cell = 0.35 V(b) If E°sulfite is −0.57 V, calculate E°selenium.