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Chapter 17: Electrochemistry

Chapter 17: Electrochemistry. Review of Redox Reactions Galvanic Cells: Using spontaneous redox reactions to generate electrical energy. Galvanic Cells Cell potential D G and work Cell potential and concentration Applications: Batteries, fuel cells, corrosion

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Chapter 17: Electrochemistry

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  1. Chapter 17: Electrochemistry • Review of Redox Reactions • Galvanic Cells: Using spontaneous redox reactions to generate electrical energy. • Galvanic Cells • Cell potential • DG and work • Cell potential and concentration • Applications: Batteries, fuel cells, corrosion • Electrolytic Cells: Using electricity to cause nonspontaneous redox reactions to occur. • Electrolytic Cells • Applications: Electrolysis of water, electrolysis of mixtures of ions, production of Al, electrorefining, metal plating, electrolysis of NaCl

  2. Example 1 Balance the following redox reactions using the half reaction method. • ClO3- + As2S3 Cl- + H2AsO4- + SO42- • K2S + KMnO4 S8 + MnO2 + KOH

  3. Figure 17.6: Cartoon of atoms reacting

  4. Figure 17.5: A standard hydrogen electrode

  5. Example 2 If a standard Cu|Cu2+ electrode is connected to a standard Al|Al3+ electrode, what reaction occurs? What is °cell? Sketch the cell, label the cathode and anode, show the direction of electron flow, and give the line notation for the cell.

  6. Example 3 Consider the following species under standard conditions: Ce4+, Ce3+, Fe2+, Fe3+, Fe, Mg, Mg2+, Ni2+, Sn • Which is the strongest oxidizing agent? • Which is the strongest reducing agent? • Will Fe dissolve in 1.0 M Ce4+? If so, will Fe3+ or Fe2+ be formed? • Which can be oxidized by H+(aq)? • Which can be reduced by H2(g)?

  7. Example 4 Select an oxidizing agent to oxidize Cl- to Cl2 without oxidizing Br- to Br2.

  8. Example 5 Calculate the maximum work available from 25.0 g of aluminum in the following galvanic cell for which the emf is 1.15 V. Note that O2 is reduced to H2O in this reaction. Al(s)|Al3+(aq)||H+(aq)|O2(g)|Pt(s)

  9. Example 6 Calculate the cell potential for the following Galvanic cell at 25°C. Ni(s)|Ni2+(1.0M)||Sn2+(1.0x10-4M)|Sn(s)

  10. Example 7 Find the potential of a Ag+(1.0x10-7M)|Ag(s) electrode at 25°C.

  11. Example 8 Calculate the equilibrium constant for the following reaction at 25°C. Ag+(aq) + Fe2+(aq)  Ag(s) + Fe3+(aq)

  12. Example 9 A concentration cell is made up of two Ag/Ag+ half cells. In the first half cell, [Ag+] = 0.010 M. In the second half cell, [Ag+] = 4.0 x 10-4 M. What is the cell potential? Which half cell functions as the anode?

  13. Figure 17.13: Lead storage battery

  14. Figure 17.14: Common dry cell battery KOH

  15. Figure 17.16: hydrogen-oxygen fuel cell

  16. Figure 17.17: Corrosion of Iron

  17. Figure 17.18: Cathodic protection of an underground pipe

  18. Example 10 Predict the products and calculate the minimum voltage required for the electrolysis of the following substances using platinum electrodes. • MgBr2(l) • 1.0 M NiCl2(aq)

  19. Example 11 Write the net ionic equation for the reaction you expect to occur when the electrolysis of NiSO4(aq) is conducted using a nickel anode and an iron cathode.

  20. Example 12 How many grams of silver are deposited at a platinum cathode in the electrolysis of an aqueous solution of AgNO3 by 1.73A of electric current in 2.5 hours?

  21. Example 13 How long will it take to produce 10.0 g of bismuth (Bi) by the electrolysis of a BiO+ solution using a current of 25.0 A?

  22. Figure 17.22: Production of Al

  23. Figure 17.23: Refining of Copper

  24. Figure 17.25: Downs cell for production of sodium and chlorine

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