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Electrochemistry

Murray Low & Udhir Chathuri. Electrochemistry. Overview. Quiz A. Rules. ½ Cells. Quiz B. Oxidation Numbers. ½ - reactions & E 0 values. Quiz G. Corrosion. Electrochemistry. The Electrolytic Cell. The Galvanic Cell. Building a Galvanic Cell. Electrolysis in Industry. Quiz E.

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Electrochemistry

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  1. Murray Low & UdhirChathuri Electrochemistry

  2. Overview Quiz A Rules ½ Cells Quiz B Oxidation Numbers ½ - reactions & E0 values Quiz G Corrosion Electrochemistry The Electrolytic Cell The Galvanic Cell Building a Galvanic Cell Electrolysis in Industry Quiz E Quiz C General Types of Batteries Common Types of Batteries Principle Components of a battery Quiz F Al Cl2 Quiz D Zn / C Ni / Cd Pb / Acid Diaphragm Membrane Mercury

  3. Rules for Assigning Oxidation Numbers 1. The oxidation number of an elemental substance is zero. 2. The oxidation number of a monoatomic ion (simple ion) is equal to the charge of the ion. 3. The sum of the oxidation numbers of all the atoms in a species is equal to the charge of the species. 4. Assign hydrogen an oxidation number of +1 when it is combined with a non-metal and -1 when it is combined with a metal. 5. Assign fluorine an oxidation number of -1 in all its compounds. 6. Assign oxygen an oxidation number of -2. Note: these rules are in order of importance. e.g. Rule 5 is paramount over Rule 6 if there is a potential conflict between the rules.

  4. Quiz A – Oxidation Numbers 1. Assign an oxidation number to each of the elements in the following species. a) Oxygen gas, O2 b) Water, H2O c) Hydrogen peroxide, H2O2 d) The perchlorate ion, ClO4- e) Sulfurous acid, H2SO4 f) Oxygen difluoride, OF2 g) Ozone gas, O3 h) Sodium thiosulfate, Na2S2O3 i) Potassium tetrathionate, K2S4O6 j) Sodium hydride, NaH

  5. ½ - reactions & E0 values Written as reductions : Ox + ne- Red F2(g) + 2e- 2F-(aq) Zero on the relative scale. The more positive the Eo, the stronger the oxidizing agent on the LHS. The more negative the Eo, the stronger the reducing agent on the RHS.

  6. Standard Hydrogen Electrode Pt electrode PH2 (g)= 1 atm 2e- H2 2H+ 2e- 1 M H+ (H3O+) T = 25 oC (298 K) The standard hydrogen electrode is given an Eo value (by convention) of 0.00V. All Eo values are quoted relative to this zero.

  7. The two types of ½ - Cell 1. Simple metal/ion redox couple e.g. Ni2+/Ni Ni electrode 1M Ni2+ 2. Inert electrode with ionic redox couple e.g. Pt with Fe3+/Fe2+ Pt electrode 1M Fe3+ Fe2+ 1M Fe2+ Fe3+ Inert electrodes include Pt, Pd & Carbon

  8. Rules for Balancing Aqueous Redox Equations 1. Identify the elements changing oxidation numbers. 2. Separate the reaction into an oxidation ½ -reaction and reduction ½-reaction. 3. Balance the two ½-reactions thus: i) Balance the element that is changing oxidation state. ii) Insert the correct number of electrons into the ½-reaction. iii) Balance the charge by adding H+(acid) or OH-(base). iv) Balance the number of oxygen atoms by the addition of water. v) Check the hydrogens are balanced. 4. Adjust the two ½-reactions to transfer the same number of electrons. 5. Combine the two ½-reactions to give a balanced redox reaction.

  9. Quiz B- Eo values & Balancing Redox Reactions 1 a) Which is the stronger oxidizing agent : Cl2(g) or MnO4- (aq). b) Which is the stronger reducing agent : Fe(s) or Zn(s). 2. Balance the following aqueous redox reactions. a) Cr2O72- + Fe2+ Fe3+ + Cr3+ (acid) b) MnO4- + C2O42- Mn2+ + CO2 (acid) c) F2 + NaOH  OF2 + NaF d) OCl- + I-  IO3- + Cl2 (acid)

  10. The Galvanic Cell ? Cu2+ (aq) + 2e-→ Cu (s) e.g. E0 = +0.34V Zn2+ (aq) + 2e-→ Zn (s) Zn (s) → Zn2+ (aq) + 2e- E0 = -0.76V Cu2+ + Zn → Zn2+ + Cu ∆E0 = Eored – E0ox ∆E0 = +1.10V = +0.34V – (-0.76V) All galvanic cells have positive ΔE0 values All spontaneous redox reactions have postive ΔE0 values REDCat & AnOx Reduction at Cathode & Oxidation at Anode

  11. Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- K+ K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- K+ Zn Zn Cl- K+ Cl- Cu Cu K+ Cl- K+ Cl- K+ K+ = 1 [Zn2+] [Zn2+] Qinit = K = [Cu2+] [Cu2+] = 1.6 x 1037 Building a Galvanic Cell 0.00 +0.95 +1.10 K+ Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- K+ K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ K+ Cl- Cu2+ + Zn → Zn2+ + Cu Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- SO2-4 K+ Cl- K+ Cu2+ Zn2+ Cl- K+ K+ Zn2+ Zn2+ SO2-4 SO2-4 SO2-4 Cu2+ SO2-4 Zn2+ Cu2+ SO2-4 SO2-4 SO2-4 Cu2+ Zn2+ Zn2+ SO2-4 Cu2+ SO2-4 Cu2+ (aq) + 2e-→ Cu (s) Zn (s) → Zn2+ (aq) + 2e-

  12. Quiz C- Galvanic Cells • A galvanic cell involving the following two ½-cells : Mn2+/Mn (Eo = -1.18V) • and Cr3+/Cr (Eo = -0.74V), is to be constructed. a) Draw a diagram of the cell. b) Show on the diagram the direction of the electron flow. c) Show on the diagram the direction that the cations move in the salt bridge. d) Write a balanced cell equation. e) Calculate the Ecell at i) t = 0 ii) t = ∞

  13. General Types of Batteries Primary Cell Non-rechargeable Provides electricity until it dies (i.e. achieves equlibrium) Disposable as redox couple is non-reversible e.g. Zinc/Manganese battery (Dry cell) Secondary Cell Rechargeable Provides electricity until it goes flat Connect to external power source to reverse redox reactions e.g. Pb/PbSO4 battery

  14. Principle Components of a Battery Terminals Anode Current Collector Cathode Current Collector Anode active mass Cathode active mass Electrolyte Container Separator

  15. Zn/C Dry Cell Battery Zn cathode Carbon anode Carbon paste (carbon, electrolyte(NH4Cl) and MnO2) Insulation Anode Zn(s)  Zn2+(aq) + 2e- Cathode 2MnO2(s) + H2O(l) + 2e- Mn2O3(s) + 2OH-(aq) Zn(s) + 2MnO2(s) + H2O(l)  Zn2+(aq) + Mn2O3(s) + 2OH-(aq) Electrolyte : NH4Cl / ZnCl2 / MnO2 / C Powder Current collectors : Graphite & Zinc

  16. Nickel / Cadmium Battery Anode Cd(s) + 2OH-(aq) Cd(OH)2(s)+ 2e- Cathode NiO(OH)(s) + H2O(l)+ e- Ni(OH)2(s) + OH- x2 2NiO(OH)(s) + Cd(s) + 2H2O(l) 2Ni(OH)2(s) + Cd(OH)2(s) Electrolyte : KOH Current collectors : Ni & Cd

  17. Lead / Acid Battery Anode Pb(s) + SO42-(aq) PbSO4(s)+ 2e- PbO2(s) + 4H+(aq) + SO42-(aq)+ 2e- 2H2O(l) + PbSO4(s) Cathode PbO2(s) + Pb(s) + 4H+(aq) + 2SO42-(aq) 2PbSO4(s)+ 2H2O(l) Electrolyte : H2SO4 Current collectors : Both Pb

  18. Quiz D - Batteries • The lead storage battery is based upon the following ½ -cells: PbSO4(s) + 2e- Pb(s) + SO42- (aq) E0 = -0.31 V PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4 (s) + 2H2O(l) E0 = +1.70 V • Write down the overall redox reaction that occurs when the lead storage • battery is being recharged. b) Given a car battery is usually required to supply about 12 V, how many galvanic cells must be in series within a car battery?

  19. Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- K+ K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- K+ Zn Cl- K+ Cl- Cu K+ Cl- K+ Cl- K+ K+ The Electrolytic Cell Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- K+ K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ Cl- K+ K+ Cl- K+ Cl- K+ Cl- Cl- K+ K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ K+ Cl- Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- K+ Cl- Zn2+ + Cu(s)  Cu2+ + Zn(s) K+ Cl- SO2-4 K+ Cl- K+ Cl- K+ Cu2+ K+ Zn2+ SO2-4 Zn2+ SO2-4 SO2-4 SO2-4 Cu2+ Zn2+ Cu2+ SO2-4 SO2-4 SO2-4 Cu2+ Zn2+ Zn2+ SO2-4 SO2-4 Cu2+

  20. Quiz E- Electrolytic Cells • Design an electrochemical cell to produce Ce4+ and Au(gold). Ce4+(aq) + e- Ce3+(aq) E0 = +1.61 V Au3+(aq) + 3e- Au(s) E0 = +1.50 V Your diagram should include the direction of the electron flow, the direction of anion movement in the salt bridge and the minimum potential required from the battery to produce the desired products.

  21. Electrochemistry in Industry Aluminium Production + C(s) + 2O2-→ CO2 + 4e- x3 Al2O3 dissolved in molten Na3AlF6 at 1000oC Al3+ + 3e-→ Al x4 - 3C(s) + 6O2- + 4Al3+→ 3CO2(g) + 4Al(s) 3C(s) + 2Al2O3(s)→ 3CO2(g) + 4Al(s)

  22. Bauxite from mine } to Air CO2 SO2 Al2O3 purification HF + F2 removal Volatile hydrocarbons Na3AlF6 + CO + CO2 + SO2 Carbon anode plant Electrolysis cell Filter to remove Particulate matter Electrochemistry in Industry Aluminium Production Coal Tar pitch Natural Gas for heating Coke Al

  23. Chlor-Alkali Process NaCl Na2CO3 (Soda Ash) Cl2 + NaOH + H2 Fuel Feedstock NaOH Pulp & Paper Cl2 Pulp & Paper Soap Plastics Bleach Organochlorines Dyes Bleach Textiles

  24. Electrolytic Process Sodium bicarbonate Limestone & Fuel Sodium chloride Chlorine Caustic Soda Soda Ash Ammonia Carbon Dioxide Chlor-alkali industries chart Pulp & Paper Plastics Sanitation Bleach Herbicides Soap Rayon Dyes Paper Rubber Textiles Bleaching Neutralisation Soaps Glass Drugs Paper Textiles Metallurgy Petroleum Water softening Drugs Beverages Baking Powder

  25. Electrolytic Cells for Chlor-alkali Diaphragm Cell 2e- 2e- Steel gauze cathode Dimensionally stable anode Cl- Cl- Cl- Cl- Na+ Na+ Na+ Na+ Na+ NaCl O2 O2 H2O 4OH-→ 2H2O + O2 + 4e- Cl- Na+ - - - O O O H2 O O O O Permeable membrane. Both cations and anions can cross. 30% NaOHcontaining Cl- Cl2 H H H H H H H H H H H Needs Evaporation

  26. Electrolytic Cells for Chlor-alkali Membrane Cell 2e- 2e- Steel gauze cathode Dimensionally stable anode Cl- Cl- Cl- Cl- Na+ Na+ Na+ Na+ Na+ Na+ NaCl H2O - - - - O O O O H2 O O O O Cation permeable membrane. Cations can cross but not anions. 50% NaOH Cl2 H H H H H H H H H H H H Ideally!

  27. Mercury Cell Electrolytic Cells for Chlor-alkali Cl2(g) 2Cl-(aq) → Cl2(g) + 2e- Cell 35% NaCl 17% NaCl Hg(l) + 2Na+(aq)+ 2e-→ 2HgNa(l) Recycle Hg NaHg amalgam Denuder H2O H2(g) 2H2O(l) + 2HgNa(l) → 2NaOH(aq) + 2Hg(l) + H2(g) 50% NaOH

  28. Typical data for Chlor-alkali cells Table1: Typical data for recent commercial chlor-alkali cells

  29. Quiz F- Chlor-alkali data 1 a) Write down balanced half-cell reactions for the processes happening at the anode and the cathode in the membrane (and diaphragm) cell. b) Write down a balanced cell reaction for the membrane cell. 2. By consulting Table 1 answer the following questions. a) Which cell is favoured by the cell voltage? b) Which cell is favoured by the current density? c) Which cell is favoured by the current efficiency? d) Which cell is favoured by the energy consumption?

  30. Corrosion We use iron(Fe) as a building material. The trouble is it rusts (i.e. it oxidizes). Fe can exist in three different oxidation states : 0, +2 and +3. 1. Fe  Fe2+ + 2e- E = -0.44 V 2. Fe  Fe3+ + 3e- E = -0.06 V The common oxidizing agent in the environment is, of course, O2 gas. O2 + 4H+ + 4e-  2H2O E = +1.23 V Therefore, for Fe  Fe2+ E = +1.23 – (-0.44) = +1.67 V and for Fe  Fe3+ E = +1.23 – (-0.06) = +1.29 V

  31. 2 Fe(s) + O2(g) + nH2O(l)  Fe2O3.nH2O(s) Corrosion Fe is first oxidized to Fe2+ and subsequently to Fe3+. Therefore Fe  Fe2+ + 2e- will be the first thing to happen then Fe2+  Fe3+ + e- will happen subsequently . so 2Fe  2Fe2+ + 4e- (Anode) O2 + 4H+ + 4e- 2H2O (Cathode) 2Fe + O2 + 4H+  2Fe2+ + 2H2O 2Fe2+ + ½O2 + (2 + n)H2O  Fe2O3.nH2O + 4H+ then Overall : Therefore corrosion requires O2/H2O to occur and will be catalysed by H+.

  32. Fe Cu Corrosion Fe corrodes when it behaves as an anode. Cu2+ + 2e- Cu E0 = + 0.34V Fe2+ + 2e- Fe E0 = - 0.44V = Corrosion of Fe Fe + Cu2+  Fe2+ + Cu ΔE0 =+ 0.78V Spontaneous reaction Cu behaves as a cathode. Fe behaves as an anode.

  33. Fe Zn Corrosion Fe will not corrode when it behaves as a cathode. Zn2+ + 2e- Zn E0 = - 0.76 V Fe2+ + 2e- Fe E0 = - 0.44 V = Corrosion of Zn Zn + Fe2+  Zn2+ + Fe ΔE0 =+ 0.32 V Zn behaves as an anode. Fe behaves as a cathode.

  34. Quiz G- Corrosion • Which of the following metals are suitable for use as sacrificial anodes • to protect against corrosion of an underground iron tank? Explain your choices. i) Aluminium ii) Silver iii) Nickel iv) Sodium 2. Why do steel bridge-supports rust at the water line but not above or below? 3. In Scotland, the common occurrence of ice, in winter, requires the of use salt on the roads. This lowers the freezing point of water and melts the ice. Scottish cars get rusty very quickly in the winter….Why?

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