Magnitude of the Cell Potential is a Measure Answer Key

Things you should know when you leave Discussion today: 1. Reduction potential E°red at standard conditions use Chemistry: Human Activity, Chemical reactivity, Mahaffy et al., 2e Appendix F for information. Joule ° 2. Cell potential E𝑐𝑒𝑙𝑙 = or Volt (V) is a measure of electrical potential difference. charge 0.05912 V ° ° ° E𝑐𝑒𝑙𝑙 = E𝑟𝑒𝑑(𝑐𝑎𝑡ℎ𝑜𝑑𝑒) − E𝑟𝑒𝑑(𝑎𝑛𝑜𝑑𝑒) = ∙ log K (@25℃) 𝑛𝑒 3. http://quantum.bu.edu/courses/ch102-spring-2016/notes/BalancingRedoxEquations.pdf 4. http://quantum.bu.edu/courses/ch102-spring-2016/notes/StandardRedcutionPotentials.pdf 𝑐ℎ𝑎𝑟𝑔𝑒 5. Gibbs free energy and potential∆𝐺 ° = −𝑛𝑒 ∙ 𝐹 ∙ 𝐸 ° where F = 96485.3 𝑚𝑜𝑙 (𝑜𝑓 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠) 6. Magnitude of the cell potential is a measure of the available energy from the reaction. 7. 0.05912 V 0.05912 V 𝑄 𝐸 = 𝐸° − ∙ log 𝑄 = − ∙ log (@25℃) 𝑛𝑒 𝑛𝑒 K a. If Q = 1 then E = E° indicating we are at standard state conditions b. If Q = K then E = 0 indicating we are at equilibrium i. E > 0 forward reaction proceeds ii. E < 0 reverse reaction proceeds iii. E = 0 no electron flow (battery is dead) 1. Choose all that must be true: a. Anode is an electrode where oxidation takes place b. Cathode is an electrode where reduction takes place c. Most positive reduction potential represents the strongest oxidizing agent (an oxidizing agent is an electron acceptor and is most likely to get reduced) d. Most negative reduction potential represents strongest reducing agent (most likely to get oxidized) e. E° is an intensive quantity as long as concentration does not change it does not depend on the size of the cell f. E° does not depend on the amount of charge transferred g. E° does not depend on stoichiometry (doubling the stoichiometry of all the species in the reaction has no effect on the E°) h. Sign of the cell potential tells the direction of electron flow i. Only reduction potentials are listed in the reference tables (oxidation potentials can be calculated from reduction potentials E°ox =−E°red) 2. Write the balanced equation for the oxidation of NO(g) to NO3-(aq) by reduction of dichromate, Cr2O72-(aq) to Cr3+(aq). Assign for each half-reaction if it is a cathodic half-reaction or an anodic half-reaction. Oxidation half-reaction: Reduction half-reaction: Net reaction in acidic media: 1 Write the line notation for the galvanic cell corresponding to the redox reaction, using platinum electrodes, Pt(s), if needed. 3. You have an electrochemical reaction that takes place in acidic solution of lead sulfate. Coming out of solution is a lead electrode and a platinum electrode. 2H+(aq) +2e― → H2(g) E°red = 0 V PbSO4(s) +2e― → Pb(s) + SO42― (aq) E°red = −0.360 V Pt(s) | H2(g) | H3O+(aq) || PbSO4(s) | SO42–(aq) | Pb(s) Anode half-reaction: Cathode half-reaction: NET reaction: Q= Eocell = Check-off one of the following: Change in the cell Increase in E 1) Increase in pH of the solution [H3O+] decreases 2) Dissolving Na2SO4(s) in the solution Ksp >>1 3) Increase in size of the Pb(s) electrode 4) Decrease in H2 gas pressure 5) Increase H2 gas pressure 6) Increase in the amount of PbSO4(s) Ksp<<1 7) Adding HCl in the solution 8) Addition of water to the solution hint: how does concentration vary with volume? 2 Decrease in E No effect on E Net reaction: 2H2O +H2(g) + PbSO4(s)  2H3O+ (aq) + Pb(s) + SO42– (aq) ne =2 (𝐻3 𝑂+ )2 (𝑆𝑂42− ) 0.05912 V 𝐸 = 𝐸° − ∙ log 𝑄 𝑄= 𝑛𝑒 (𝐻 ) 2 (𝑔) a. At Q=1 was the reaction a spontaneous proces? For the following questions [SO42–(aq)] and p(H2(g)) are kept at standard states. You can only change concentration of [H3O+]. b. What pH is needed for E cell = E°cell? c. Calculate the equilibrium constant: d. What pH is needed for E cell = 0? e. Give an example of the pH that will make this reaction spontaneous. (Explain why) f. If the cell pH=10, what is Ecell? Was this process spontaneous at pH=10? 4. Is the following expression at 25°C 𝐸 = 𝐸 ° − Always true never true never true log 𝑄 only true when Q=1 5. Is the following expression at 25°C 𝐸 ° = Always true 0.05912 V ∙ 𝑛𝑒 0.05912 V ∙ 𝑛𝑒 only true when Q=K log K only true when Q=1 only true when Q=K only true when Q=1 only true when Q=K only true when Q=1 only true when Q=K 6. Is the following expression E = 0 Always true never true 7. Is the following expression E = E° Always true never true 8. If E = – 2.00 V and E° = 1.00 V circle everything that must be true: Hint: remember what it means for Q to be equal to1? a. Q=1 Q>1 Q<1 Q < Keq Q = Keq Q > Keq Assuming the temperature is 25°C and ne= 2.00 mol calculate the values of Keq and Q to one significant figure. 3 9. For a reaction A(aq) + B(aq)  2C(aq) at 298K a. E = 2.00 V and E° = – 1.00V and ne= 2.00mol circle everything that must be true: Q=1 Q>1 Q<1 Q < Keq Q = Keq Q > Keq 𝑄𝑛𝑒𝑤 ? 𝑄𝑜𝑙𝑑 b. If the concentration of C is doubled what is the c. If the concentration of C is doubled what is the new value of E°. d. If concentration of the C is tripled how will the magnitude of E increase or decrease? e. What is new new value of E? 10. Answer the questions for the following redox reaction at 25°C. Eored(Zn2+│Zn) = –0.76 V; Eored(Ag+│Ag) = 0.80 V Zn(s) | Zn2+(aq, 0.01M) || Ag+(aq, 1M) | Ag(s) a. Cathodic reaction: b. Anodic reaction: c. Net reaction: d. Eocell = e. Q= f. Ecell = 11. Answer the questions for the following redox reaction: Cu(s) | Cu2+(aq, 0.001M) || Cu2+(aq, 1M) | Cu(s) Eored(Cu2+│Cu) = 0.34V a. Cathode reaction: b. Anode reaction: c. Net reaction: d. Eocell = e. Q= f. Ecell = 4 12. Answer the questions for the following redox reaction at 25°C. Ag(s) | Ag+(aq, 0.0010M) || Ag+(aq, 1.0M) | Ag(s) a. Cathodic reaction: b. Anodic reaction: c. Net reaction: d. Eocell = e. Q = f. Ecell = 13. Calculate the voltage (E) of a concentration cell constructed with the Cl‒ concentration difference between sea water and river water at 25°C. Assume that the Cl‒ concentration (due to dissolved NaCl) of sea water is 35 g/L and then that of river water is 1.0 mg/L. a. Cathode reaction: b. Anode reaction: c. Eocell = d. Q= e. Ecell = 14. The standard cell potential for the process of A(aq) + B(aq)  2C(aq) at 25°C in which three moles of electrons are transferred is Eocell = 3.00 V. An electrochemical cell for this process is constructed and the measured voltage is Ecell = 5.00 V. Circle all the correct statements. a. Q < 1 Q=1 Q>1 Q=K Q>K Q1, Q>Keq b. 7∙1033; 10100 9. a. Q<1, Q< Keq b. 4 c. -1 V d. 1.97 1. 2. 3. E = -0.360 V a. No b. 0 c. K=6.6∙10-13 d. pH = 6.1 e. pH>6 f. 0.24, Yes 4. Always true 5. Always true 10. Cathodic RXN: 2Ag+(aq) +2e- Ag(s) Anodic RXN: Zn(s) Zn2+(aq)+2eNet RXN: 2Ag+(aq) + Zn(s) Zn2+(aq)+ Ag(s) Eocell = 0.80V+0.76V=1.56V 11. Ecell = 1.56V-  0.05912V 0.01 log 12 =1.62V ne a. b. c. d. e. Cathode RXN Cu2+(aq, 1M)+2e-  Cu(s) Anode RXN: Cu(s) Cu2+(aq, 0.001M)+2eNet RXN: Cu (aq, 1M)  Cu2+(aq, 0.001M) Eocell =0V [Cu2+(aq,0.001M)] Q= [Cu2+(aq,1M)] =0.001 f. Ecell =0  0.05912V log0.001=0.09V; ne ne=2 12. Cathodic RXN: Ag+(aq, 1M)+ e-  Ag(s) 6 Anodic RXN: Ag(s) Ag+(aq, 0.001M)+ eNet RXN: Ag(aq, 1M)  Ag+(aq, 0.001M) Eocell = 0V Ecell = 0  0.05912V log0.001=0.18V; ne ne=1 13. Cathodic RXN : 2Cl-(aq)  Cl2(g)+2e- (sea water) Anodic RXN: Cl2(g) +2e- 2Cl- (aq) (river water) 1 𝑚𝑔 𝑔 35.5 𝑔 2 ( ∙ ∙ 0.05912V 𝐿 1000 𝑚𝑔 𝑚𝑜𝑙 ) 𝐸𝑐𝑒𝑙𝑙 = 0 − ∙ log = 0.27V; 𝑛𝑒 1 𝑚𝑜𝑙 2 ( 𝐿 ) 14. a. Q<1, Q 0 process is spontaneous: ln W final > 0; ΔS < 0 process is not spontaneous: ln W 7 initial Wfinal initial < 0; Wfinal > 1; Winitial Wfinal < 1; Winitial Wf > Win Wf < Win Look up the following concepts in Human Activity, Chemical reactivity, Mahaffy et al., 2e, pages 671-684 1. State the Second Law of Thermodynamics: 2. Define ΔSuniverse(ΔSnet) 3. What is the relationship between ΔSsystem and ΔSreaction? 4. What equation relates ΔSsystem and ΔHsystem when the system is at equilibrium? 5. What is the relationship between ΔSsurrounding and ΔSreaction? 6. What equation relates ΔSsurrounding and ΔHsurroundings? 7. What equation relates ΔSsurrounding and ΔHsystem? Answers: 1. 0.124 if (K12∙K2 2 points) 2. a. NO2 (g) decrease b. 𝐾 increase; N2 O3 (g) decrease c. NO2 (g) decrease 3. a. 0.93bar b. 0.57 4. 1.5∙10-10 5. a. Basic b. Basic 6. a. 1.36∙10-2 % b. 2.83∙10-13M 7. 2.04∙10-33 or 3.80∙10-30 8. 1.6∙10-6 9. 1∙10-5 10. 4.5∙10-13 11. a. 2.00 M b. -0.301 c. 6.7∙10-M 12. HA is a weak acid; A− is a weak base; 𝐾a > 𝐾w ; [H3 O+ ] > [OH − ]; [A− ] < [HA] 8