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ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 14

ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 14. DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university. CHAPTER 14 ELECTRODE POTENTIALS. REDOX CHEMISTRY. - Electron transfer occurs in redox reactions Oxidation - Loss of electrons

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ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 14

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  1. ANALYTICAL CHEMISTRY CHEM 3811CHAPTER 14 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

  2. CHAPTER 14 ELECTRODE POTENTIALS

  3. REDOX CHEMISTRY - Electron transfer occurs in redox reactions Oxidation - Loss of electrons Reduction - Gain of electrons Oxidizing agent (oxidant) is the species reduced Reducing agent (reductant) is the species oxidized

  4. REDOX CHEMISTRY Oxidizing Agent - The species that gains electrons - The species that is reduced - Causes oxidation Aox + ne-↔ Ared Cu2+(aq) + 2e-↔ Cu(s)

  5. REDOX CHEMISTRY Reducing Agent - The species that loses electrons - The species that is oxidized - Causes reduction Bred↔ Box + ne- Fe(s) ↔ Fe2+(aq)+ 2e-

  6. REDOX CHEMISTRY The Overall Reaction - Both an oxidation and a reduction must occur in a redox reaction - The oxidizing agent accepts electrons from the reducing agent Aox + Bred↔ Ared + Box Cu2+(aq) + Fe(s) ↔Cu(s) + Fe2+(aq) - Reducing agent - Oxidized species - Electron loss - Oxidizing agent - Reduced species - Electron gain

  7. REDOX CHEMISTRY Half Reactions - Oxidation half reaction Bred↔ Box + ne- Fe(s) ↔ Fe2+(aq)+ 2e- - Reduction half reaction Aox + ne- ↔ Ared Cu2+(aq) + 2e- ↔ Cu(s)

  8. ELECTRODE - Conducts electrons into or out of a redox reaction system Examples platinum wire carbon (glassy or graphite) indium tin oxide (ITO) Electroactive Species - Donate or accept electrons at an electrode

  9. REDOX CHEMISTRY Charge (q) of an electron = - 1.602 x 10-19 C Charge (q) of a proton = + 1.602 x 10-19 C C = coulombs Charge of one mole of electrons = (1.602 x 10-19 C)(6.022 x 1023/mol) = 9.649 x 104 C/mol = Faraday constant (F) q = n x F

  10. CURRENT - The quantity of charge flowing past a point in an electric circuit per second Units Ampere (A) = coulomb per second (C/s)

  11. VOLTAGE Potential Difference (E) - Work done by or on electrons when they move from one point to another Units: volts (V or J/C) Work (J) = E (V) x q (C)

  12. CHEMICAL CHANGE Spontaneous Process - Takes place with no apparent cause Nonspontaneous Process - Requires something to be applied in order for it to occur (usually in the form of energy)

  13. ELECTROLYSIS - Voltage is applied to drive a redox reaction that would not otherwise occur Examples - Production of aluminum metal from Al3+ - Production of Cl2 from Cl-

  14. ELECTROLYSIS CELL - Nonspontaneous reaction - Requires electrical energy to occur

  15. GALVANIC CELL - Spontaneous reaction - Produces electrical energy - Can be reversed electrolytically for reversible cells Example Rechargeable batteries Conditions for Non-reversibility - If one or more of the species decomposes - If a gas is produced and escapes

  16. GALVANIC CELL - A spontaneous redox reaction generates electricity - One reagent is oxidized and the other is reduced - The two reagents must be separated (cannot be in contact) - Electrons flow through a wire (external circuit)

  17. VOLTAIC (GALVANIC) CELL Oxidation Half reaction - Loss of electrons - Occurs at anode (negative electrode) - The left half-cell by convention Reduction Half Reaction - Gain of electrons - Occurs at cathode (positive electrode) - The right half-cell by convention

  18. GALVANIC CELL Salt Bridge - Connects the two half-cells (anode and cathode) - Filled with gel containing saturated aqueous salt solution (KCl) - Ions migrate through to maintain electroneutrality - Prevents charge buildup that may cease the reaction process Preparation - Heat 3 g of agar and 30 g of KCl in 100 mL H2O - Heat until a clear solution is obtained - Pour into a U-tube and allow to gel - Store in a saturated aqueous KCl

  19. VOLTAIC (GALVANIC) CELL For the overall reaction Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) e- Voltmeter e- - + Cu electrode Zn electrode Cl- K+ Zn2+ Salt bridge (KCl) Cu2+ Anode Oxidation Zn(s) → Zn2+(aq) + 2e- Cathode Reduction Cu2+(aq) + 2e- → Cu(s)

  20. GALVANIC CELL Voltage or Potential Difference (E) - Is the voltage measured - Measured by a voltmeter (potentiometer) connected to electrodes Greater Voltage - More favorable net cell reaction - More work done by flowing electrons

  21. GALVANIC CELL Line Notation Phase boundary: represented by one vertical line Salt bridge: represented by two vertical lines Fe(s) FeCl2(aq) CuSO4(aq) Cu(s)

  22. STANDARD POTENTIALS Electrode Potentials - A measure of how willing a species is to gain or lose electrons Positive Voltage (spontaneous process) - Electrons flow into the negative terminal of voltmeter (flow from negative electrode to positive electrode) Negative Voltage (nonspontaneous process) - Electrons flow into the positive terminal of voltmeter (flow from positive electrode to negative electrode) Conventionally - Negative terminal is on the left of galvanic cells

  23. STANDARD POTENTIALS Standard Reduction Potential (Eo) - Used to predict the voltage when different cells are connected - Potential of a cell as cathode compared to standard hydrogen electrode - Species are solids or liquids - Activities = 1 - We will use concentrations for simplicity Concentrations = 1 M Pressures = 1 bar

  24. STANDARD POTENTIALS Standard Hydrogen Electrode (SHE) - Used to measure Eo for half-reactions - Connected to negative terminal - Pt electrode - Acidic solution in which [H+] = 1 M - H2 gas (1 bar) is bubbled past the electrode H+(aq, 1 M) + e- ↔ 1/2H2 (g, 1 bar) Conventionally, Eo = 0 for SHE

  25. STANDARD POTENTIALS The Eo for Ag+ + e- ↔ Ag(s) is 0.799 V Implies that if a sample of silver metal is placed in a 1 M Ag+ solution, a value of 0.799 V will be measured with S. H. E. as reference Pt(s) H2(g, 1 bar) H+(aq, 1 M ) Ag+ (aq, 1 M) Ag(s) SHE Ag+ (aq, 1 M) Ag(s)

  26. STANDARD POTENTIALS Silver does not react spontaneously with hydrogen 2H+(aq) + 2e- → H2(g) Eo = 0.000 V Ag+(aq) + e- → Ag(s) Eo = +0.799 V Reverse the second equation (sign changes) Ag(s) → Ag+(aq) + e- Eo = -0.799 V Multiply the second equation by 2 (Eo is intensive so remains) 2Ag(s) → 2Ag+(aq) + 2e- Eo = -0.799 V Combine (electrons cancel) 2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g) Eo = -0.799 V

  27. STANDARD POTENTIALS Consider Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Cu2+(aq) + 2e- → Cu(s) Eo = +0.339 V Zn2+(aq) + 2e- → Zn(s) Eo = -0.762 V Reverse the second equation (sign changes) Zn(s) → Zn2+(aq) + 2e- Eo = +0.762 V Combine (electrons cancel) Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = +1.101 V Eo is positive so reaction is spontaneous Reverse reaction is nonspontaneous

  28. STANDARD POTENTIALS - Half-reaction is more favorable for more positive Eo Formal Potential - The potential for a cell containing a specified concentration of reagent other than 1 M

  29. STANDARD POTENTIALS Eo (V) 2.890 1.507 1.280 1.229 0.799 0.339 0.000 -0.402 -0.440 -0.763 -1.659 -2.936 -3.040 Half Reaction F2 + 2e-↔ 2F- MnO4- + 5e-↔ Mn2+ Ce4+ + e- ↔ Ce3+ (in HCl) O2 + 4H+ + 4e- ↔ 2H2O Ag+ + e- ↔ Ag(s) Cu2+ + 2e- ↔ Cu(s) 2H+ + 2e- ↔ H2(g) Cd2+ + 2e- ↔ Cd(s) Fe2+ + 2e- ↔ Fe(s) Zn2+ + 2e- ↔ Zn(s) Al3+ + 3e- ↔ Al(s) K+ + e- ↔ K(s) Li+ + e- ↔ Li(s) Oxidizing agents Reducing agents Increasing reducing power Increasing oxidizing power

  30. NERNST EQUATION For the half reaction aA + ne-↔ bB The half-cell potential (at 25 oC), E, is given by

  31. NERNST EQUATION Eo = standard electrode potential R = gas constant = 8.314 J/K-mol T = absolute temperature F = Faraday’s constant = 9.649 x 104 C/mol n = number of electrons

  32. NERNST EQUATION - The standard reduction potential (Eo) when [A] = [B] = 1M - [B]b/[A]a = Q = reaction quotient - Concentration for gases are expressed as pressures in bars - Q = 1 for [ ] = 1 M and P = 1 bar logQ = 0 and E = Eo - Q is omitted for pure solids, liquids, and solvents

  33. NERNST EQUATION - When a half reaction is multiplied by a factor Eo remains the same - For a complete reaction Ecell = E+ - E- and Eo = E+o - E-o E+ = potential at positive terminal E- = potential at negative terminal

  34. NERNST EQUATION For the Cu – Fe cell at standard conditions Cu2+ + 2e- ↔ Cu(s) 0.339 V Fe2+ + 2e- ↔ Fe(s) -0.440 V Ecell = 0.779 V Galvanic Reaction Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq) Fe Fe2+ (1M) Cu2+ (1 M) Cu

  35. NERNST EQUATION - Positive E implies spontaneous forward cell reaction - Negative E implies spontaneous reverse cell reaction If cell runs for long - Reactants are consumed - Products are formed - Equilibrium is reached - E becomes 0 - Reason why batteries run down

  36. NERNST EQUATION At cell equilibrium at 25 oC E = 0 and Q = K (the equilibrium constant) Or Positive Eo implies K > 1 Negative Eo implies K < 1

  37. REFERENCE ELECTRODES Indicator Electrode - Responds directly to the analyte Reference Electrode - Provides known and constant potential Examples Silver-silver chloride electrode (Ag/AgCl) Saturated Calomel electrode (SCE)

  38. REFERENCE ELECTRODES Saturated Calomel electrode (SCE) - Saturated with KCl 1/2Hg2Cl2(s) + e- ↔ Hg(l) + Cl- E = + 0.241 V In this case, the reference is not 0.000 V (SHE) but 0.241 V (SCE)

  39. REFERENCE ELECTRODES Saturated Calomel electrode (SCE) - Different KCl concentrations can be used - 0.1 M KCl is least temperature sensitive - Saturated KCl solution is easier to make and maintain

  40. REFERENCE ELECTRODES Silver-Silver Chloride Electrode (Ag/AgCl) - Saturated with KCl AgCl(s) + e-↔ Ag(s) + Cl- E = + 0.197 V

  41. REFERENCE ELECTRODES Emeasured = Eo - 0.241 (SCE) Emeasured = Eo - 0.197 (Ag/AgCl) Eo(SHE) E(SCE) E(Ag/AgCl) Cu2+ + 2e- ↔ Cu(s) 0.339 V 0.098 V 0.142 V Fe2+ + 2e- ↔ Fe(s) -0.440 V -0.681 V -0.637 V

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