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Electrochemistry

Electrochemistry. Redox Reactions 01. Redox reaction are those involving the oxidation and reduction of species. OIL – O xidation I s L oss of electrons. RIG – R eduction I s G ain of electrons. Oxidation and reduction must occur together. They cannot exist alone.

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Electrochemistry

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  1. Electrochemistry

  2. Redox Reactions 01 • Redox reaction are those involving the oxidation and reduction of species. • OIL – Oxidation Is Loss of electrons. • RIG – Reduction Is Gain of electrons. • Oxidation and reductionmust occur together. • They cannot exist alone. Chapter 18

  3. Redox Reactions 02 Chapter 18

  4. Redox Reactions 02 • Oxidation Half-Reaction: Zn(s)  Zn2+(aq) + 2 e–. • The Zn loses two electrons to form Zn2+. Chapter 18

  5. Redox Reactions 03 • Reduction Half-Reaction:Cu2+(aq) + 2 e– Cu(s) • The Cu2+ gains two electrons to form copper. Chapter 18

  6. Redox Reactions 04 • Overall: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) Chapter 18

  7. Electrochemical Cells 01 • Electrodes: are usually metal strips/wires connected by an electrically conducting wire. • Salt Bridge: is a U-shaped tube that contains a gel permeated with a solution of an inert electrolyte. • Anode: is the electrode where oxidation takes place. • Cathode: is the electrode where reduction takes place. Chapter 18

  8. Electrochemical Cells 02 • Convention for expressing the cell: • Anode Half-Cell || Cathode Half-Cell • Electrode | Anode Soln || Cathode Soln | Electrode • Zn(s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu(s) • Pt(s) | H2 (1 atm) | H+(1 M) || Fe3+(aq), Fe2+(aq) | Pt(s) • Electrons flow from anode to cathode. Anode is placed on left by convention. Chapter 18

  9. Electrochemical Cells 03 • The standard potential of any galvanic cell is the sum of the standard half-cell potentials for the oxidation and reduction half-cells. • E°cell = E°oxidation + E°reduction • Standard half-cell potentials are always quoted as a reduction process. The sign must be changed for the oxidation process. Chapter 18

  10. Electrochemical Cells 04 • The standard half-cell potentials are determined from the differencebetween two electrodes. • The reference point is called the standard hydrogen electrode (S.H.E.) and consists of a platinum electrode in contact with H2 gas (1 atm) and aqueous H+ ions (1 M). • The standard hydrogen electrode is assigned an arbitrary value of exactly 0.00 V. Chapter 18

  11. Electrochemical Cells 05 Chapter 18

  12. Electrochemical Cells 06 Chapter 18

  13. Electrochemical Cells 07 • When selecting two half-cell reactions the more negative value will form the oxidation half-cell. • Consider the reaction between zinc and silver: • Ag+(aq) + e– Ag(s) E° = 0.80 V • Zn2+(aq) + 2 e– Zn(s) E° = – 0.76 V • Therefore, zinc forms the oxidation half-cell: • Zn(s)  Zn2+(aq) + 2 e–E° = – (–0.76 V) Chapter 18

  14. Electrochemical Cells 08 • What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr electrode in a 1.0 M Cr(NO3)2 solution? • What is the standard emf of an electrochemical cell made of a Mg electrode in a 1.0 M Mg(NO3)2 solution and a Ag electrode in a 1.0 M AgNO3 solution? Chapter 18

  15. Spontaneity of a Reaction 01 • The value of E˚cellis related to the thermodynamic quantities of ∆G˚ and K. • The value of E˚cellis related to ∆G˚ by: • ∆G˚ = –nFE˚cell • The value of Kis related to ∆G˚ by: • ∆G˚ = –RT ln K Chapter 18

  16. Spontaneity of a Reaction 02 • These are combined: • –nFE˚cell= –RT ln K • and rearranged at 25°C to give: • E˚cell= (0.0257n) lnK • orE˚cell= (0.0592n) logK • (Units = volts, n = moles, F = 96,500 coulombs) Chapter 18

  17. Spontaneity of a Reaction 03 • Calculate the standard free energy change (∆G˚) and the equilibrium constant (K) for the following reactions at 25°C: • Sn(s) + 2 Cu2+(aq) æ Sn2+(aq) + 2 Cu+(aq) • Fe2+(aq) + 2 Ag(s) æ Fe(s) + 2 Ag+(aq) • 4 Fe2+(aq) + O2(g) + 4 H+(aq)  4 Fe3+(aq) + 2 H2O(l) Chapter 18

  18. The Nernst Equation 01 • Cell potentials can be modified by temperature and composition changes according to the equation: • ∆G = ∆G° + RT lnQ •  –nFE = –nFE° + RT lnQ • Giving: Chapter 18

  19. The Nernst Equation 03 • Consider the reaction of metallic zinc with hydrochloric acid. Calculate the cell potential at 25°C when [H+] = 1.0 M, [Zn2+] = 0.0010 M, and PH2 = 0.10 atm. • Consider the reaction of metallic copper with iron(III) to give copper(II) and iron(II). What is the potential of a cell when [Fe3+] = 0.0001 M, [Cu2+] = 0.25 M, and [Fe2+] = 0.20 M? Chapter 18

  20. The Nernst Equation 04 • A particularly important use of the Nernst equation is in the electrochemical determination of pH. • Pt | H2 (1 atm) | H+ (? M) || Reference Cathode • Ecell = EH2 H++ Eref • The Nernst equation can be applied to the half-reaction: • H2(g)  2 H+(aq) + 2 e– Chapter 18

  21. The Nernst Equation 05 • For the half-reaction: H2(g)  2 H+(aq) + 2 e– • E° = 0 V for this reaction (standard hydrogen electrode). PH2 is 1 atm. Chapter 18

  22. The Nernst Equation 07 • The overall potential is given by: • Which rearranges to give an equation for the determination of pH: Chapter 18

  23. The Nernst Equation 08 • A higher cell potential indicates a higher pH, therefore we can measure pH by measuring Ecell. • A glass electrode (Ag/AgCl wire in dilute HCl) with a calomel reference is the most common arrangement. • Glass: Ag(s) + Cl–(aq)  AgCl(s) + e–E° = –0.22 V • Calomel: Hg2Cl2(s) + 2 e– 2 Hg(l) + 2 Cl–(aq) E° = 0.28 V Chapter 18

  24. The Nernst Equation 09 • The glass pH probe is constructed as follows: • Ag(s) | AgCl(s) | HCl(aq) | glass | H+(aq) || reference • The difference in [H+] from one side of the glass membrane to the other causes a potential to develop, which adds to the measured Ecell. Chapter 18

  25. The Nernst Equation 10 • The following cell has a potential of 0.55 V at 25°C: • Pt(s) | H2 (1 atm) | H+ (? M) || Cl– (1 M) | Hg2Cl2(s) | Hg(l) • What is the pH of the solution at the anode? • The following cell has a potential of 0.28 V at 25°C: • Pt(s) | H2 (1 atm) | H+ (? M) || Pb2+ (1 M) | Pb(s) • What is the pH of the solution at the anode? Chapter 18

  26. Batteries 01 • Batteries are the most important practical application of galvanic cells. • Single-cell batteries consist of one galvanic cell. • Multicell batteries consist of several galvanic cells linked in series to obtain the desired voltage. Chapter 18

  27. Batteries 02 • Lead Storage Battery: A typical 12 volt battery consists of six individual cells connected in series. • Anode: Lead grid packed with spongy lead. • Pb(s) + HSO4–(aq)  PbSO4(s) + H+(aq) + 2 e– • Cathode: Lead grid packed with lead oxide. • PbO2(s) + 3 H+(aq) + HSO4–(aq) + 2 e– PbSO4(s) + 2 H2O(l) • Electrolyte: 38% by mass sulfuric acid. • Cell Potential: 1.924 V Chapter 18

  28. Batteries 03 • Zinc Dry-Cell: Also called a Leclanché cell, uses a viscous paste rather than a liquid solution. • Anode: Zinc metal can on outside of cell. Zn(s)  Zn2+(aq) + 2 e– • Cathode: MnO2 and carbon black paste on graphite. • 2 MnO2(s) + 2 NH4+(aq) + 2 e– Mn2O3(s) + 2 NH3(aq) + 2 H2O(l) • Electrolyte: NH4Cl and ZnCl2 paste. • Cell Potential: 1.5 V but deteriorates to 0.8 V with use. Chapter 18

  29. Batteries 04 • Alkaline Dry-Cell: Modified Leclanché cell which replaces NH4Cl with NaOH or KOH. • Anode: Zinc metal can on outside of cell. • Zn(s) + 2 OH–(aq)  ZnO(s) + H2O(l) + 2 e– • Cathode: MnO2 and carbon black paste on graphite. • 2 MnO2(s) + H2O(l) + 2 e– Mn2O3(s) + 2 OH–(aq) • Electrolyte: NaOH or KOH, and Zn(OH)2 paste. • Cell Potential: 1.5 V but longer lasting, higher power, and more stable current and voltage. Chapter 18

  30. Batteries 05 • Mercury Dry-Cell: Modified Leclanché cell which replaces MnO2 with HgO and uses a steel cathode. • Anode: Zinc metal can on outside of cell. • Zn(s) + 2 OH–(aq)  ZnO(s) + H2O(l) + 2 e– • Cathode: HgO in contact with steel. • 2 HgO(s) + H2O(l) + 2 e– Hg(l) + 2 OH–(aq) • Electrolyte: KOH, and Zn(OH)2 paste. • Cell Potential: 1.3 V with small size, longer lasting, and stable current and voltage. Chapter 18

  31. Batteries 06 • Nickel–Cadmium Battery: Modified Leclanché cell which is rechargeable. • Anode: Cadmium metal. • Cd(s) + 2 OH–(aq)  Cd(OH)2(s) + 2 e– • Cathode: Nickel(III) compound on nickel metal. • NiO(OH)(s) + H2O(l) + e– Ni(OH)2(s) + OH–(aq) • Electrolyte: Nickel oxyhydroxide, NiO(OH). • Cell Potential: 1.30 V Chapter 18

  32. Batteries 07 • Nickel–Metal–Hydride (NiMH): • Replaces toxic Cd anode witha hydrogen atom impregnated ZrNi2 metal alloy. • During oxidation at the anode, hydrogen atoms are released as H2O. • Recharging reverses this reaction. Chapter 18

  33. Batteries 08 • Lithium Ion (Li–ion): The newest rechargeable battery is based on the migration of Li+ ions. • Anode: Li metal, or Li atom impregnated graphite. • Li(s)  Li+ + e– • Cathode: Metal oxide or sulfide that can accept Li+. • MnO2(s) + Li+(aq) + e– LiMnO2(s) • Electrolyte: Lithium-containing salt such as LiClO4, in organic solvent. Solid state polymers can also be used. • Cell Potential: 3.0 V Chapter 18

  34. Batteries 09 • Fuel Cell: Uses externally fed CH4 or H2, which react to form water. Most common is H2. • Anode: Porous carbon containing metallic catalysts. • 2 H2(s) + 4 OH–(aq)  4 H2O(l) + 4 e– • Cathode: Porous carbon containing metallic catalysts. • O2(s) + 2 H2O(l) + 4 e– 4 OH–(aq) • Electrolyte: Hot aqueous KOH solution. • Cell Potential: 1.23 V, but is only 40% of cell capacity. Chapter 18

  35. Batteries 10 • Fuel cells are not batteries because they are not self-contained. • Fuel cells typically have about 40% conversion to electricity; the remainder is lost as heat. • Excess heat can be used to drive turbine generators. Chapter 18

  36. Corrosion 01 • Corrosion is the oxidative deterioration of metal. • 25% of steel produced in USA goes to replace steel structures and products destroyed by corrosion. • Rusting of iron requires the presence of BOTH oxygen and water. • Rusting results from tiny galvanic cells formed by water droplets. Chapter 18

  37. Corrosion 02 Oxidation: Fe(s)  Fe2+(aq) + 2 e– Reduction: O2(g) + 4 H+(aq) + 4 e–  2 H2O(l) Overall: 2 Fe(s) + O2(g)+ 4 H+(aq)  2 Fe2+(aq) + 2 H2O(l) Chapter 18

  38. Corrosion 03 • Galvanizing: is the coating of iron with zinc. Zinc is more easily oxidized than iron, which protects and reverses oxidation of the iron. • Cathodic Protection: is the protection of a metal from corrosion by connecting it to a metal (a sacrificial anode) that is more easily oxidized. • All that is required is an electrical connection to the sacrificial anode (usually magnesium or zinc). Chapter 18

  39. Corrosion 04 • Cathodic Protection with Zinc Layer Chapter 18

  40. Corrosion 05 • Cathodic Protection with Magnesium Anode Chapter 18

  41. Electrolysis 01 Chapter 18

  42. Electrolysis 01 • Electrolysis: is the process in which electrical energy is used to drive a nonspontaneouschemical reaction. • An electrolytic cell is an apparatus for carrying out electrolysis. • Processes in an electrolytic cell are thereverseof those in a galvanic cell. Chapter 18

  43. Electrolysis 05 • Electrolysis of Water: Requires an electrolyte species, that is less easily oxidized and reduced than water, to carry the current. • Anode: Water is oxidized to oxygen gas. • 2 H2O(l)  O2(g) + 4 H+(aq) + 4 e– • Cathode: Water is reduced to hydrogen gas. • 4 H2O(l) + 4 e– 2 H2(g) + 4 OH–(aq) Chapter 18

  44. Electrolysis 02 • Electrolysis of Molten Sodium Chloride: Electrolysis of Aqueous Sodium Chloride: uses different processes to molten sodium chloride. Based on cell potentials, water (–0.83 V) would be preferentially reduced over sodium ions (–2.71 V). Based on cell potentials, water (+1.23 V) should be preferentially oxidized over chloride ions (+1.36 V). The observed product at the anode is Cl2, not O2, because of a phenomenon called overvoltage. Chapter 18

  45. Electrolysis 04 • Overvoltage: Additional voltage needed to maintain the rate of electron transfer at the electrode–solution interface. • Overvoltage is required when a half-reaction has a significant activation energy, and so a slow rate. • Overvoltage for formation of O2 or H2 is much greater than for formation of Cl2. Chapter 18

  46. Electrolysis 07 • Quantitative Electrolysis: The amount of substance produced at an electrode by electrolysis depends on the quantity of charge passed through the cell. • Reduction of 1 mol of sodium ions requires 1 mol of electrons to pass through the system. • The charge on 1 mol of electrons is 96,500 coulombs. Chapter 18

  47. Electrolysis 08 • To determine the moles of electrons passed, we measure the current and time that the current flows: • Charge (C) = Current (A) x Time (s) • Because the charge on 1 mol of e– is 96,500 C, the number of moles of e– passed through the cell is: Chapter 18

  48. Electrolysis 09 • A constant current of 30.0 A is passed through an aqueous solution of NaCl for a time of 1.00 hour. How many grams of Na and how many liters of Cl2 gas at STP are produced? • How many amperes must be passed through a Downs cell to produce sodium metal at a rate of 30.0 kg/hour? Chapter 18

  49. Electrolysis Applications 01 • Manufacture of Sodium (Downs Cell): Chapter 18

  50. Electrolysis Applications 02 • Manufacture of Cl2 and NaOH (Chlor–Alkali): Chapter 18

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