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Electrochemistry Generating Voltage (Potential)

Electrochemistry Generating Voltage (Potential). Historically. Historically oxidation involved reaction with O 2 . i.e., Rusting 4 Fe (s) + 3O 2 (g)  Fe 2 O 3 (s) Other example Zn (s) + Cu 2+ (aq) g Zn 2+ (aq) + Cu (s) In this reaction:

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Electrochemistry Generating Voltage (Potential)

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  1. ElectrochemistryGenerating Voltage (Potential)

  2. Historically • Historically oxidation involved reaction with O2. • i.e., Rusting • 4 Fe(s) + 3O2 (g) Fe2O3 (s) • Other example • Zn(s) + Cu2+(aq)g Zn2+(aq) + Cu(s) • In this reaction: • Zn(s)g Zn2+(aq) Oxidation • Cu2+(aq)g Cu(s) Reduction • In a redox reaction, one process can’t occur without the other. Oxidation-Reduction reaction must simultaneously occurs.

  3. Redox Between • If Zn(s) and Cu2+(aq) is in the same solution, then the electron is a transferred directly between the Zn and Cu. No useful work is obtained. However if the reactants are separated and the electrons shuttle through an external path...

  4. Zn Cu Build up of positive charge Build up of negative charge Electrochemical Cells • Voltaic / Galvanic CellApparatus which produce electricity • Electrolytic CellApparatus which consumes electricity • Consider: Initially there is a flow of e- After some time the process stops Electron transport stops because of charge build up The charge separation will lead to process where it cost too much energy to transfer electron.

  5. Completing the Circuit • Electron transfer can occur if the circuit is closed • Parts: • Two conductors • Electrolyte solution • Salt Bridge / Porous membrane 3 process must happen if e- is to flow. A. e- transport through external circuit B. In the cell, ions a must migrate C. Circuit must be closed (no charge build up) Anode (-) Black Negative electrode generates electron Oxidation Occur Cathode (+) Red Positive electrode accepts electron Reduction Occur A C B Cathode/Cation(+) Anode/Anion (-)

  6. Voltaic Cell • Electron transfer can occur if the circuit is closed • Parts: • Two conductors • Electrolyte solution • Salt Bridge / Porous membrane 3 process must happen if e- is to flow. A. e- transport through external circuit B. In the cell, ions a must migrate C. Circuit must be closed (no charge build up) Anode (-) Black Negative electrode generates electron Oxidation Occur Cathode (+) Red Positive electrode accepts electron Reduction Occur Cathode/Cation(+) Anode/Anion (-)

  7. Completing the Circuit: Salt Bridge • In order for electrons to move through an external wire, charge must not build up at any cell. This is done by the salt bridge in which ions migrate to different compartments neutralize any charge build up.

  8. Sign Convention of Voltaic Cell • @ Anode: Negative Terminal (anions). • Source of electron then repels electrons. Oxidation occurs. • Zn(s)g Zn+2(aq) + 2e- : Electron source • @ Cathode: Positive Terminal (cation) • Attracts electron and then consumes electron. Reduction occurs. • Electron target: 2e- + Cu+2(aq)g Cu(s) • Overall: • Zn(s) + Cu+2(aq)g Zn+2(aq) + Cu(s) E° = 1.10 V • Note when the reaction is reverse:Cu(s)+ Zn+2(aq)g Cu+2(aq) + Zn(s) • Sign of E ° is also reversed E° = -1.10 V • Oxidation:Zn(s)g Zn+2(aq) E° = 0.76 V • Reduction: Cu+2(aq)g Cu(s) E° = 0.34 V • 1.10 V = E°CELL • or E°CELL = E°red (Red-cathode) - E°red (Oxid-anode)

  9. Other Voltaic Cell • Zn(s) + 2H+ (aq)g Zn+2(aq) + H2 (g) E° = 0.76 V @ Anode:Negative Terminal (anions): Zn(s)g Zn+2(aq) + 2e- : Source of electron then repels electrons. Oxidation occurs. @ Cathode: Positive Terminal (cation): 2e- + 2H+(aq)g H2 (g) Attracts electron and then consumes electron. Reduction occurs. Net: Zn(s) + 2H+(aq)g Zn2+(aq) + H2(g)

  10. Other Voltaic Cell • Zn(s) + 2H+ (aq)g Zn+2(aq) + H2 (g) E° = 0.76 V @ Anode:Negative Terminal (anions): Zn(s)g Zn+2(aq) + 2e- : Source of electron then repels electrons. Oxidation occurs. @ Cathode: Positive Terminal (cation): 2e- + 2H+(aq)g H2 (g) Attracts electron and then consumes electron. Reduction occurs. Net: Zn(s) + 2H+(aq)g Zn2+(aq) + H2(g)

  11. 1 Line Notation Convention • Line notation: Convenient convention for electrochemical cell • Schematic Representation • 1. Anode g Cathode • [oxidation (-) ] [reduction (+)] • 2. “ “ phase boundary • (where potential may develop) • 3. “ “ Liquid junction • 4. Concentration of component • Zn(s) ZnSO4 (aq,1.0M) CuSO4(aq,1.0M) Cu(s) 4 3 2

  12. Line Notation Examples • Consider : Zn(s) + Cu+2(aq)g Zn+2(aq) + Cu(s • Anode: Zn g Zn+2 + 2e- • Cathode: Cu+2 + 2e- g Cu • Shorthand “Line” notation • Zn (s) Zn+2(aq)(1.0M) Cu+2(aq)(1.0M) Cu(s) • 2nd Example : Zn(s) + 2H+ (aq)g Zn+2(aq) + H2(g) • Anode: Zn g Zn+2 + 2e- • Cathode: 2H+ + 2e- g H2(g) • Shorthand “Line” notation • Zn (s) Zn+2(aq)(1.0M) H+(aq)(1.0M), H2(g, 1atm) Pt(s)

  13. Other Voltaic Cell & Their Line Notation • Zn(s) | Zn+2(aq)||H+(aq) , H2 (g,1atm)|Pt Oxidation half-reaction Cr(s)g Cr+3(aq) + 3e- Oxidation half-reaction 2I- (aq)g I2 (s) + 2e- Oxidation half-reaction Zn(s)g Zn+2(aq) + 2e- Reduction half-reaction MnO4-(aq) + 8H+(aq)+ 5e- g Mn 2+(aq) +4H2O(l) Reduction half-reaction Ag+(aq) + e- g Ag (s) Cr(s) | Cr+3(aq)||Ag+(aq) | Ag(s) C(s)| I-(aq) , I2 (g,1atm) || MnO4-(aq) , Mn+2(aq)| C(s)

  14. Line Notation Examples • Example 1: B&L 20.13 • Zn(s) + Ni2+(aq)g Zn+2(aq) + Ni(aq) • Example 2: B&L 20.19 • Tl+3(aq) + 2Cr2+(aq)g Tl+(aq) + Cr+3(aq)

  15. Voltage of Galvanic / Voltaic Cell • Transport of any object requires a net force. • Consider water flowing through pipes. This occurs because of pressure gradient. Flow (Fluid Transport) Pressure (h) Pressure (i) Or Similarly, electron are transported through wires because of the electromotive force EMF or Ecell. Object falling or transport down due to Dh (+) (-) e -

  16. EMF - ElectroMotive Force • Potential energy of electron is higher at the anode. This is the driving force for the reaction (e- transfer) e Anode (-) D P.E. = V = J e - C e- flow toward cathode (+) Cathode Larger the gap, the greater the potential (Voltage)

  17. ElectroMotive Force • EMF - Electro Motive Force • Potential energy difference between the two electrodes • The larger the DP.E. the larger EMF value. • The magnitude of P.E. for the reaction (half reaction) is an intensive property) • i.e., Size independent: r, Tbpt, Cs. • Therefore EMF is also an intensive property. • Analogy: • Size of rock not important, only the height from ground. • (Electron all have the same mass) • Unit: EMF: V - Volts : • 1V - 1 Joule / Coulomb • 1 Joule of work per coulomb of charge transferred.

  18. Stoichiometry Relationship to E° EMF - Intensive Property E°cellStandard state conditions 25°C, 1atm, 1.0 M E°cell Intensive property, Size Independent Consider: Li+ + e-g Li (s) E°Cell = -3.045 V x 2 2 Li+ + 2 e-g 2 Li (s) E°Cell = (-3.045 V) x 2 = ?? But E° = Voltage per electron E° ‘ = E° x 2 = ? g - 3.045 V • 2 = -3.045 V 2 e- \ Stoichiometry does not change E°, but reversing the reaction does change the sign of E°.

  19. Standard Reduction Potential Most spontaneous <Reduction occurs> Oxidizing Agent Written as reduction Cell Potential is written as a reduction equation. M+ + e-g M E° = std red. potential Most non-spontaneous Spontaneous in the reverse direction. <Oxidation occurs> Reducing Agent

  20. Zoom View of Std. Reduction Potential Most spontaneous Reduction Oxidizing Agent Written as reduction Most non-spontaneous Spontaneous in the reverse direction. Oxidation Reducing Agent Cell Potential is written as a reduction equation. M+ + e-g M E° = F2 (g) + 2e-g 2 F -(aq) 2.87 V Ce4++ e-g Ce3+ (aq) 1.61 V 2H++ 2e-g H2 (g) 0.00 V Li+(aq) + e-g Li(s) -3.045 V All reaction written as reduction reaction. But in electrochemistry, there can’t be just a reduction reaction. It must be coupled with an oxidation reaction.

  21. E°Cell Evaluation E°Cell Function of the reaction g Oxidation Process (Anode reaction) g Reduction Process (Cathode reaction) or E°Cell = E°Cathode & E°Anode Cathode (+) Anode (-) Most Negative Reduction reaction Therefore, E°Cell = E°red(Cathode) - E°red(anode) Neg Minus (Large negative) (Very Positive Value) Very Positive \ Very Spontaneous

  22. Standard Reduction Potential Zn g Zn+2 + 2e- E°=? Can’t determine because the reaction must be coupled How is E°red (Cathode) and E°red (Anode) determine. E° (EMF) - State Function; there is no absolute scale Absolute E° value can’t be measured experimentally The method of establishing a scale is to measure the difference in potential between two half-cells. Consider: How can a scale of reduction potential be determine ? Use a half reaction as reference and assign it a potential of zero. Electrochemical reaction more spontaneous than this reference will have positive E°, and those less spontaneous will have negative E°.

  23. Side-Bar: Relative Scale Carry the child in arms and weight both child and parent then subtract the weight of the parent from the total to yield the baby weight. Consider a baby whose weight is to be determine but will not remain still on top of a scale. How can the parents determine the babies weight?

  24. Reference Potential • Selected half reaction is: • H+ / H2 (g) couple half reaction: 2H+(aq, 1.0M) + 2e-g H2(g,1atm) • by definition c E° = 0.0 V, the reverse is also 0.0 V • H+/H2 couple - Standard Hydrogen Electrode (SHE) • To determine E° for a another half reaction, the reaction of interest needs to be coupled to this SHE. The potential measured is then assigned to the half-reaction under investigation. E°Cell = 0.76 V = E°red (Cat) - E°red (Anode) 0.0 V - (?) E°red (Anode) = - 0.76 V \Zn+2/Zn E° = -0.76 V Reduction rxn

  25. Determining Other Half-Cell Potential • Now consider the reaction: • Zn(s)|Zn+2 (1.0 M)||Cu+2(1.0 M)|Cu(s) • E°Cell = 1.10 V • E°Cell = E°red (Cat) - E°red (Anode) • recall, E° Zn+2/Zn = - 0.76 V • Therefore, • E°Cell = E°Cu+2/Cu - E° Zn+2/Zn • 1.10 V = (?) - (- 0.76 V) • E°Cu+2/Cu = + 0.34 V

  26. 1.19 V Pt Pt Tl3+g Tl+ Cr2+g Cr3+ Example: Half-Cell Potential • Example BBL20.19: • For the reaction: Tl+3 + 2Cr2+ Tl+ + 2Cr3+ E°Cell = 1.19 V • i) Write both half reaction and balance • ii) Calculate the E°Cell Tl+3 Tl+ • iii) Sketch the voltaic cell and line notation • i) Tl+3 + 2e- Tl+ • (Cr2+ 2Cr3+ + 2e- ) x 2 E° = 0.41 V • ii) E°Cell = 1.19 V = E°red (Cat) - E°red (Anode) • 1.19 V= E°red (Cat) - 0.041 V • for Tl+3 + 2e- Tl+ : • 1.19 V - 0.41 = E°red (Cat) = 0.78 V

  27. Voltaic Vs. Electrolytic Cells Electrolytic Cell Energy is absorbed to drive nonspontaneous redox reaction Voltaic Cell Energy is released from spontaneous redox reaction General characteristics of voltaic and electrolytic cells. A voltaic cell generates energy from a spontaneous reaction (DG<0), whereas an electrolytic cell requires energy to drive a nonspontaneous reaction (DG>0). In both types of cell, two external circuits provides the means or electrons to flow. Oxidation takes place all the anode, and reduction takes place at the cathode, but the relative electrode changes are opposite in the two cells. Surrounding (power supply) do work on system (cell) System does work on load (surroundings) Anode (Oxidation) Oxidation Reaction A-g A + e- Oxidation Reaction X g X+ + e- Reduction Reaction e- + Y+g Y Reduction Reaction e- + B+g B Overall (Cell) Reaction A- + B+g A + B, DG> 0 Overall (Cell) Reaction X + Y+g X+ + Y, DG = 0

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