Chem 1515 Section 2 Problem Set #15 Spring 1998 Name ALL work must be shown to receive full credit. Due Due in lecture at 1:30 p.m. Friday, May 1st. PS15.1. Balance the following oxidation-reduction reactions using the half-reaction method. (a) VO 2+ (aq) + MnO 4 (aq) VO2 + (aq) + Mn 2+ (aq) acidic solution (5e + 8H + + MnO 4 Mn 2+ + 4H 2 O) 5(H 2 O + VO 2+ VO 2 + + 2H + + e ) 8H + + 5H 2 O+ 5VO 2+ (aq) + MnO 4 (aq) 5VO2 + (aq) + Mn 2+ (aq) + 10H + + 4H 2 O H 2 O+ 5VO 2+ (aq) + MnO 4 (aq) 5VO2 + (aq) + Mn 2+ (aq) + 2H + (b) NO 3 (aq) + Cu(s) Cu 2+ (aq) + NO(g) acidic solution 2(3e + 4H + + NO 3 NO + 2H 2 O) 3(Cu Cu 2+ + 2e ) 3Cu + 8H + + 2NO 3 2NO + 4H 2 O + 3Cu 2+ (c) Fe 3+ (aq) + NH 3 OH + (g) Fe 2+ (aq) + N 2 O(g) acidic solution 4(Fe 3+ (aq) + 1e Fe 2+ (aq) ) 2NH 3 OH + (g) N 2 O(g) + H 2 O + 6H + + 4e 4Fe 3+ (aq) + 2NH 3 OH + (g) N 2 O(g) + H 2 O + 6H + + 4Fe 2+ (aq) (d) MnO 4 (aq) + I (aq) MnO2 (s) + IO 3 (aq) basic solution 2(3e + 4H + + MnO 4 MnO2 + 2H 2 O) 3H 2 O + I IO 3 + 6H + + 6e 3H 2 O + I + 8H + + 2MnO 4 2MnO2 + 4H 2 O + IO 3 + 6H + I + 2H + + 2MnO 4 2MnO2 + H 2 O + IO 3 + 2OH + 2OH I + 2H 2 O + 2MnO 4 2MnO2 + H 2 O + IO 3 + 2OH I + H 2 O + 2MnO 4 2MnO2 + IO 3 + 2OH (e) As 2 S 3 (aq) + H 2 O 2 (aq) AsO 4 3 (aq) + SO4 2 (aq) basic solution 4(2e + 2H + + H 2 O 2 2H 2 O) 20H 2 O + As 2 S 3 (aq) 2AsO 4 3 (aq) + 3SO4 2 (aq) + 20H + + 8e 8H + + 20H 2 O + As 2 S 3 (aq) + 4H 2 O 2 (aq) 2AsO 4 3 (aq) + 3SO4 2 (aq) + 8H2 O + 20H + 12H 2 O + As 2 S 3 (aq) + 4H 2 O 2 (aq) 2AsO 4 3 (aq) + 3SO4 2 (aq) + 12H + + 12OH + 12OH 12H 2 O + 12OH + As 2 S 3 (aq) + 4H 2 O 2 (aq) 2AsO 4 3 (aq) + 3SO4 2 (aq) + 12H2 O 12OH + As 2 S 3 (aq) + 4H 2 O 2 (aq) 2AsO 4 3 (aq) + 3SO4 2 (aq) May 2, 1998 1 Spring 1998
(f) S 2 O 3 2 (aq) + Cl 2 (g) SO 4 2 (aq) + Cl (aq) acidic solution 4(2e + Cl 2 (g) 2Cl (aq) ) S 2 O 3 2 (aq) + 5H 2 O 2SO 4 2 (aq) + 10H + + 8e S 2 O 3 2 (aq) + 4Cl 2 (g) + 5H 2 O 2SO 4 2 (aq) + 8Cl (aq) + 10H + (g) Bi(OH) 3 (aq) + Sn(OH) 3 (aq) Sn(OH)6 2 + Bi(s) basic solution 2(3e + 3H + + Bi(OH) 3 Bi + 3H 2 O) 3(3H 2 O + Sn(OH) 3 Sn(OH) 6 2 + 3H + + 2e ) 9H 2 O + 3Sn(OH) 3 + 6H + + 2Bi(OH) 3 2Bi + 6H 2 O + 3Sn(OH) 6 2 + 9H + 3H 2 O + 3Sn(OH) 3 + 2Bi(OH) 3 2Bi + 3Sn(OH) 6 2 + 3H + + 3OH + 3OH 3OH + 3H 2 O + 3Sn(OH) 3 + 2Bi(OH) 3 2Bi + 3Sn(OH) 6 2 + 3H 2 O 3OH + 3Sn(OH) 3 + 2Bi(OH) 3 2Bi + 3Sn(OH) 6 2 (h) H 2 O 2 (aq) + MnO4 (aq) O 2 (g) + MnO 2 (s) basic solution 3(H 2 O 2 O 2 + 2H + + 2e ) 2(3e + 4H + + MnO 4 MnO2 + 2H 2 O) 3H 2 O 2 + 8H + + 2MnO 4 2MnO2 + 4H 2 O + 3O 2 + 6H + 3H 2 O 2 + 2H + + 2MnO 4 2MnO2 + 4H 2 O + 3O 2 + 2OH + 2OH 3H 2 O 2 + 2H 2 O + 2MnO 4 2MnO2 + 4H 2 O + 3O 2 + 2OH 3H 2 O 2 + 2MnO 4 2MnO2 + 2H 2 O + 3O 2 + 2OH (i) Fe(CrO 2 ) 2 + Na 2 CO 3 + O 2 Fe 2 O 3 + Na 2 CrO 4 + CO 2 acidic solution 2(7H 2 O + 2Fe(CrO 2 ) 2 + 4Na 2 CO 3 Fe 2 O 3 + 4Na 2 CrO 4 + 4CO 2 + 14H + + 14e ) 7(4e + 4H + + O 2 2H 2 O) 28H + + 14H 2 O + 4Fe(CrO 2 ) 2 + 7O 2 + 8Na 2 CO 3 2Fe 2 O 3 + 8Na 2 CrO 4 + 8CO 2 + 28H + + 14 H 2 O 4Fe(CrO 2 ) 2 + 7O 2 + 8Na 2 CO 3 2Fe 2 O 3 + 8Na 2 CrO 4 + 8CO 2 (j) S 8 (aq) S 2 O 3 2 (aq) + S 2 (aq) basic solution (12H 2 O + S 8 4S 2 O 3 2 + 24H + + 16e ) (16e + S 8 8S 2 ) 12H 2 O + 2S 8 4S 2 O 3 2 + 24H + + 8S 2 6H 2 O + S 8 2S 2 O 3 2 + 12H + + 4S 2 + 12OH + 12OH 6H 2 O + S 8 + 12OH 2S 2 O 3 2 + 12H 2 O + 4S 2 S 8 + 12OH 2S 2 O 3 2 + 6H 2 O + 4S 2 May 2, 1998 2 Spring 1998
PS15.2. Draw a diagram of the cells in which the following reactions occur. In each case, label the anode and cathode, the anode and cathode electrode material, the half-reaction at each electrode, the ions in the anode and cathode compartments and salt bridge, the direction of electron flow, and the direction of ion movement. (a) Al(s) + Sn 2+ (aq) Al 3+ (aq) + Sn(s) Al(s) Al 3+ (aq) + 3e 2e + Sn 2+ (aq) Sn(s) (b) Al(s) + 3H + (aq) Al 3+ (aq) + 3/2H 2 (g) (c) Al(s) Al 3+ (aq) + 3e 2e + 2H + (aq) H 2 (g) NiO 2 (s) + 4H + (aq) + 2Ag(s) Ni 2+ (aq) + 2Ag + (aq) + 2H 2 O(l) May 2, 1998 3 Spring 1998
PS15.3. Write the balanced chemical equation for the overall cell reaction represented in each of the following cell notations. (a) Sn(s) Sn 2+ (aq) Ag + (aq) Ag(s) (b) Pt(s) Fe 2+ (aq), Fe 3+ (aq) (Pt(s))Cl 2 (g) Cl (aq) (c) Pb(s),PbSO 4 (s) SO 4 2 (aq) H + (aq), SO4 2 (aq) PbO2 (s),pbso 4 PS15.4. Which of the following species are reduced by Ag? Cl 2 Fe 3+ Pb 2+ I 2 NO 3 (in H + ) Ag Ag + + 1e 0.799 2e + Cl 2 2Cl +1.359 1e + Fe 3+ Fe 2+ +0.771 2e + Pb 2+ Pb -0.126 2e + I 2 2I +0.536 3e + NO 3 + 4H + NO + 2H 2 O +0.96 To determine which combination will produce a spontaneous reaction, the E for the reduction reaction must be more positive than 0.799 volts. The only species to fulfill that criteria are Cl 2 and NO 3 (in 1 M H + ). So silver will be oxidized when combined with either Cl 2 or NO 3 (in 1 M H + ). PS15.5. Which of the following species are oxidized by nitrate in acidic solution? Cl Fe 2+ Cu Au H + 3e + NO 3 + 4H + NO + 2H 2 O +0.96 2Cl 2e + Cl 2-1.359 Fe 2+ 1e + Fe 3+ 0.771 Cu Cu 2+ + 2e 0.337 Au Au 3+ + 3e 1.498 2H + no possible reaction To determine which combination will produce a spontaneous reaction, the E for the oxidation reaction must be more positive than -0.96 volts. The only species to fulfill that criteria are Fe 2+ and Cu. So nitrate in 1 M H + will be reduced when combined with either substance. May 2, 1998 4 Spring 1998
PS15.6. Select a suitable species for each of the following (a) an oxidizing agent able to convert Sn to Sn 2+, but not Cu to Cu 2+. Sn Sn 2+ + 2e +0.136 Cu Cu 2+ + 2e 0.337 So the species which is capable of oxidizing Sn, but not oxidizing Cu, must have reduction E which is between 0.136 v and +0.337 v. Some examples include Pb 2+, H + and Sn 4+. 2e + Pb 2+ Pb 0.126 2e + 2H + H 2 0.000 2e + Sn 4+ Sn 2+ +0.15 (b) a reducing agent capable of converting H + to H 2, but not Zn 2+ to Zn. 2e + 2H + H 2 0.000 2e + Zn 2+ Zn -0.763 So the species which is capable of reducing H +, but not reducing Zn 2+, must have reduction E which is between 0 v and +0.763 v. Some examples include Pb, Sn, Ni and Cr. Pb Pb 2+ + 2e +0.126 Sn Sn 2+ + 2e +0.136 Ni Ni 2+ + 2e +0.23 Cr Cr 3+ + 3e +0.74 (c) an metal capable of reacting with HNO 3 but not HCl; it displaces Ag + (aq) put not Cu 2+ (aq). 3e + 4H + + NO 3 NO + 2H 2 O +0.94 2e + 2H + H 2 0.000 2e + Cu 2+ Cu +0.337 e + Ag + Ag +0.779 So the species which is capable of reducing Ag +, but not reducing Cu 2+, must have reduction E which is between 0.337 and 0.779 v. The only example from our redox table is Hg. May 2, 1998 5 Spring 1998
PS15.7. Calculate E and determine which of the following reactions will occur in the forward direction under standard conditions? Balance the equations in acid solution. (a) Pb(s) + Cl 2 (aq) Pb 2+ (aq) + Cl (g) Pb Pb 2+ + 2e +0.126 Cl 2 + 2e 2Cl +1.36 Pb + Cl 2 Pb 2+ + 2Cl - 1.48 volts The reaction is spontaneous because the overall E is positive. (b) Mn 2+ (aq) + Cr 2 O 7 2 (aq) MnO4 (aq) + Cr 3+ (aq) 6[4H 2 O(l) + Mn 2+ (aq) MnO 4 (aq) + 8H + (aq) + 5e ] -1.491 5[6e + Cr 2 O 7 2 (aq) + 14H + (aq) 2Cr 3+ (aq) + 7H2 O(l)] +1.33 6Mn 2+ + 5Cr 2 O 7 2 + 22H + 6MnO 4 + 10Cr 3+ + 11H 2 O -0.161 volt The reaction is nonspontaneous because the overall E is negative. (c) Ag(s) + NO 3 (aq) Ag + (aq) + NO(g) 3[Ag Ag + + 1e ] 0.337 3e + NO 3 + 4H + NO + 2H 2 O +0.96 3Ag(s) + NO 3 (aq) + 4H + 3Ag + (aq) + NO(g) + 2H 2 O(l) +0.623 The reaction will occur spontaneously. PS15.8. Use standard reduction potentials to predict the spontaneous reaction, if any, that occurs between the following. If no spontaneous reaction occurs, write NR. (a) Ca 2+ (aq) + Mg(s) Mg Mg 2+ + 2e +2.37 2e + Ca 2+ Ca -2.87 Ca 2+ (aq) + Mg(s) NR The reaction is nonspontaneous because the E is negative! (b) H + (aq) + Cu 2+ (aq) 2e + Cu 2+ Cu +0.337 2e + 2H + H 2 0.000 Both reactions are reductions half-reactions. Need one oxidation half-reaction and one reduction half-reaction! No reaction. May 2, 1998 6 Spring 1998
PS15.8. (Continued) (c) Cl 2 (g) + Ag(s) Cl 2 + 2e 2Cl +1.36 Ag Ag + + 1e 0.799 Cl 2 + 2Ag 2Cl + 2Ag + +0.58 V (d) Mn 2+ (aq) + Cr 2 O 7 2 (aq) + H + (aq) 3[Mn 2+ (aq) + 2H 2 O(l) 2e + 4H + (aq) + MnO 2 (s)] -1.23 6e + Cr 2 O 7 2 (aq) + 14H + (aq) 2Cr 3+ (aq) + 7H2 O(l) +1.33 3Mn 2+ (aq) + Cr 2 O 7 2 (aq) + 2H + (aq) 2Cr 3+ (aq) + H2 O(l) + 3MnO 2 (s) +0.10 (e) Sn 4+ (aq) + Sn(s) 2e + Sn 4+ Sn 2+ +0.154 Sn Sn 2+ + 2e +0.136 Sn 4+ (aq) + Sn(s) 2Sn 2+ (aq) +0.290 PS15.9. Calculate G and K for the spontaneous reactions in problem PS15.8. (a) Ca 2+ (aq) + Mg(s) (b) H + (aq) + Cu 2+ (aq) (c) Cl 2 (g) + Ag(s) May 2, 1998 7 Spring 1998
PS15.9. (Continued) (d) Mn 2+ (aq) + Cr 2 O 7 2 (aq) + H + (aq) (e) Sn 4+ (aq) + Sn(s) PS15.10. Calculate the solubility product constant for the iron(ii) hydroxide, given the following information. Fe(OH) 2 (s) + 2e Fe(s) + 2OH (aq) E = 0.877 volts Fe 2+ (aq) + 2e Fe(s) E = 0.440 volts May 2, 1998 8 Spring 1998