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Chapter 20

Chapter 20. Electrochemistry. Chapter Goals. Electrical Conduction Electrodes Electrolytic Cells The Electrolysis of Molten Potassium Chloride The Electrolysis of Aqueous Potassium Chloride The Electrolysis of Aqueous Potassium Sulfate Faraday’s Law of Electrolysis

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Chapter 20

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  1. Chapter 20 Electrochemistry

  2. Chapter Goals • Electrical Conduction • Electrodes Electrolytic Cells • The Electrolysis of Molten Potassium Chloride • The Electrolysis of Aqueous Potassium Chloride • The Electrolysis of Aqueous Potassium Sulfate • Faraday’s Law of Electrolysis • Commercial Applications of Electrolytic Cells

  3. Chapter Goals Voltaic or Galvanic Cells • The Construction of Simple Voltaic Cells • The Zinc-Copper Cell • The Copper-Silver Cell Standard Electrode Potentials • The Standard Hydrogen Electrode • The Zinc-SHE Cell • The Copper-SHE Cell • Standard Electrode Potentials • Uses of Standard Electrode Potentials

  4. Chapter Goals • Standard Electrode Potentials for Other Half-Reactions • Corrosion • Corrosion Protection Effect of Concentrations (or Partial Pressures) on Electrode Potentials • The Nernst Equation • Using Electrohemical Cells to Determine Concentrations • The Relationship of Eocell to Go and K

  5. Chapter Goals Primary Voltaic Cells • Dry Cells Secondary Voltaic Cells • The Lead Storage Battery • The Nickel-Cadmium (Nicad) Cell • The Hydrogen-Oxygen Fuel Cell

  6. Electrochemistry • Electrochemical reactions are oxidation-reduction reactions. • The two parts of the reaction are physically separated. • The oxidation reaction occurs in one cell. • The reduction reaction occurs in the other cell. • There are two kinds electrochemical cells. • Electrochemical cells containing in nonspontaneous chemical reactions are called electrolytic cells. • Electrochemical cells containing spontaneous chemical reactions are called voltaic or galvanic cells.

  7. Electrical Conduction • Metals conduct electric currents well in a process called metallic conduction. • In metallic conduction there is electron flow with no atomic motion. • In ionic or electrolytic conduction ionic motion transports the electrons. • Positively charged ions, cations, move toward the negative electrode. • Negatively charged ions, anions, move toward the positive electrode.

  8. Electrodes • The following convention for electrodes is correct for either electrolytic or voltaic cells: • The cathode is the electrode at which reduction occurs. • The cathode is negative in electrolytic cells and positive in voltaic cells. • The anode is the electrode at which oxidation occurs. • The anode is positive in electrolytic cells and negative in voltaic cells.

  9. Electrodes • Inert electrodes do not react with the liquids or products of the electrochemical reaction. • Two examples of common inert electrodes are graphite and platinum.

  10. Electrolytic Cells • Electrical energy is used to force nonspontaneous chemical reactions to occur. • The process is called electrolysis. • Two examples of commercial electrolytic reactions are: • The electroplating of jewelry and auto parts. • The electrolysis of chemical compounds.

  11. Electrolytic Cells • Electrolytic cells consist of: • A container for the reaction mixture. • Two electrodes immersed in the reaction mixture. • A source of direct current.

  12. The Electrolysis of Molten Potassium Chloride • Liquid potassium is produced at one electrode. • Indicates that the reaction K+() + e- K() occurs at this electrode. • Is this electrode the anode or cathode? • Gaseous chlorine is produced at the other electrode. • Indicates that the reaction 2 Cl- Cl2(g) + 2 e- occurs at this electrode. • Is this electrode the anode or cathode?

  13. e- Generator-source of DC e- - electrode + electrode molten KCl K+ + e- K() cathode reaction 2Cl-Cl2(g) + 2e- anode reaction Porous barrier The Electrolysis of Molten Potassium Chloride Diagram of this electrolytic cell.

  14. The Electrolysis of Molten Potassium Chloride • The nonspontaneous redox reaction that occurs is:

  15. The Electrolysis of Molten Potassium Chloride • In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode).

  16. The Electrolysis of Aqueous Potassium Chloride • In this electrolytic cell, hydrogen gas is produced at one electrode. • The aqueous solution becomes basic near this electrode. • What reaction is occurring at this electrode? You do it! • Gaseous chlorine is produced at the other electrode. • What reaction is occurring at this electrode? You do it! • These experimental facts lead us to the following nonspontaneous electrode reactions:

  17. - pole of battery + pole of battery Battery, a source of direct current e- flow e- flow - electrode + electrode H2 gas Cl2 gas aqueous KCl The Electrolysis of Aqueous Potassium Chloride Cell diagram 2 H2O + 2e- H2(g) + 2 OH- cathode reaction 2Cl-Cl2(g) + 2e- anode reaction

  18. The Electrolysis of Aqueous Potassium Sulfate • In this electrolysis, hydrogen gas is produced at one electrode. • The solution becomes basic near this electrode. • What reaction is occurring at this electrode? You do it! • Gaseous oxygen is produced at the other electrode • The solution becomes acidic near this electrode. • What reaction is occurring at this electrode? You do it! • These experimental facts lead us to the following electrode reactions:

  19. The Electrolysis of Aqueous Potassium Sulfate

  20. - pole of battery + pole of battery Battery, a source of direct current e- flow e- flow - electrode + electrode O2 gas H2 gas aqueous K2SO4 The Electrolysis of Aqueous Potassium Sulfate Cell diagram 2 H2O + 2e- H2(g) + 2 OH- cathode reaction 2H2O O2(g) + 4H+ + 4e- anode reaction

  21. Electrolytic Cells • In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

  22. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis • Faraday’s Law - The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell. • A faraday is the amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.

  23. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis • A coulomb is the amount of charge that passes a given point when a current of one ampere (A) flows for one second. • 1 amp = 1 coulomb/second

  24. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis • Faraday’s Law states that during electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent. • This corresponds to the passage of one mole of electrons through the electrolytic cell.

  25. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis • Example 21-1: Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes.

  26. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis • Example 21-2: Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in example 21-1.

  27. Commercial Applications of Electrolytic Cells Electrolytic Refining and Electroplating of Metals • Impure metallic copper can be purified electrolytically to  100% pure Cu. • The impurities commonly include some active metals plus less active metals such as: Ag, Au, and Pt. • The cathode is a thin sheet of copper metal connected to the negative terminal of a direct current source. • The anode is large impure bars of copper.

  28. Commercial Applications of Electrolytic Cells • The electrolytic solution is CuSO4 and H2SO4 • The impure Cu dissolves to form Cu2+. • The Cu2+ ions are reduced to Cu at the cathode.

  29. Commercial Applications of Electrolytic Cells • Any active metal impurities are oxidized to cations that are more difficult to reduce than Cu2+. • This effectively removes them from the Cu metal.

  30. Commercial Applications of Electrolytic Cells • The less active metals are not oxidized and precipitate to the bottom of the cell. • These metal impurities can be isolated and separated after the cell is disconnected. • Some common metals that precipitate include:

  31. Voltaic or Galvanic Cells • Electrochemical cells in which a spontaneous chemical reaction produces electrical energy. • Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference. • Examples of voltaic cells include:

  32. The Construction of Simple Voltaic Cells • Voltaic cells consist of two half-cells which contain the oxidized and reduced forms of an element (or other chemical species) in contact with each other. • A simple half-cell consists of: • A piece of metal immersed in a solution of its ions. • A wire to connect the two half-cells. • And a salt bridge to complete the circuit, maintain neutrality, and prevent solution mixing.

  33. The Construction of Simple Voltaic Cells

  34. The Zinc-Copper Cell • Cell components for the Zn-Cu cell are: • A metallic Cu strip immersed in 1.0 M copper (II) sulfate. • A metallic Zn strip immersed in 1.0 M zinc (II) sulfate. • A wire and a salt bridge to complete circuit • The cell’s initial voltage is 1.10 volts

  35. The Zinc-Copper Cell • In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

  36. The Zinc-Copper Cell • There is a commonly used short hand notation for voltaic cells. • The Zn-Cu cell provides a good example.

  37. The Copper - Silver Cell • Cell components: • A Cu strip immersed in 1.0 M copper (II) sulfate. • A Ag strip immersed in 1.0 M silver (I) nitrate. • A wire and a salt bridge to complete the circuit. • The initial cell voltage is 0.46 volts.

  38. The Copper - Silver Cell • Compare the Zn-Cu cell to the Cu-Ag cell • The Cu electrode is the cathode in the Zn-Cu cell. • The Cu electrode is the anode in the Cu-Ag cell. • Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.

  39. The Copper - Silver Cell • These experimental facts demonstrate that Cu2+ is a stronger oxidizing agent than Zn2+. • In other words Cu2+ oxidizes metallic Zn to Zn2+. • Similarly, Ag+ is is a stronger oxidizing agent than Cu2+. • Because Ag+ oxidizes metallic Cu to Cu 2+. • If we arrange these species in order of increasing strengths, we see that:

  40. Standard Electrode Potential • To measure relative electrode potentials, we must establish an arbitrary standard. • That standard is the Standard Hydrogen Electrode (SHE). • The SHE is assigned an arbitrary voltage of 0.000000… V

  41. The Zinc-SHE Cell • For this cell the components are: • A Zn strip immersed in 1.0 M zinc (II) sulfate. • The other electrode is the Standard Hydrogen Electrode. • A wire and a salt bridge to complete the circuit. • The initial cell voltage is 0.763 volts.

  42. The Zinc-SHE Cell • The cathode is the Standard Hydrogen Electrode. • In other words Zn reduces H+ to H2. • The anode is Zn metal. • Zn metal is oxidized to Zn2+ ions.

  43. The Copper-SHE Cell • The cell components are: • A Cu strip immersed in 1.0 M copper (II) sulfate. • The other electrode is a Standard Hydrogen Electrode. • A wire and a salt bridge to complete the circuit. • The initial cell voltage is 0.337 volts.

  44. The Copper-SHE Cell • In this cell the SHE is the anode • The Cu2+ ions oxidize H2 to H+. • The Cu is the cathode. • The Cu2+ ions are reduced to Cu metal.

  45. Uses of Standard Electrode Potentials • Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials. • Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials. • Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written. • For example, the half-reaction for the standard potassium electrode is: The large negative value tells us that this reaction will occur only under extreme conditions.

  46. Compare the potassium half-reaction to fluorine’s half-reaction: The large positive value denotes that this reaction occurs readily as written. Positive E0 values denote that the reaction tends to occur to the right. The larger the value, the greater the tendency to occur to the right. It is the opposite for negative values of Eo. Uses of Standard Electrode Potentials

  47. Uses of Standard Electrode Potentials • Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously. • Example 21-3: Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn2+ ions oxidize metallic Ag to Ag+ ions? • Steps for obtaining the equation for the spontaneous reaction.

  48. Uses of Standard Electrode Potentials • Choose the appropriate half-reactions from a table of standard reduction potentials. • Write the equation for the half-reaction with the more positive E0 value first, along with its E0 value. • Write the equation for the other half-reaction as an oxidation with its oxidation potential, i.e. reverse the tabulated reduction half-reaction and change the sign of the tabulated E0. • Balance the electron transfer. • Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous.

  49. Uses of Standard Electrode Potentials

  50. Electrode Potentials for Other Half-Reactions • Example 21-4: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution? • Follow the steps outlined in the previous slides. • Note that E0 values are not multiplied by any stoichiometric relationships in this procedure.

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