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Electrochemistry

Chapter 20. Electrochemistry. Voltaic Cells. A chemical reaction can perform two types of work: Produce a gas to perform PV work Use movement of electrons from redox reactions to perform electrical work. Voltaic Cells.

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Electrochemistry

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  1. Chapter 20 Electrochemistry

  2. Voltaic Cells • A chemical reaction can perform two types of work: • Produce a gas to perform PV work • Use movement of electrons from redox reactions to perform electrical work

  3. Voltaic Cells • A voltaic (galvanic) cell is a device in which the transfer of electrons takes place through an external pathway rather than directly between reactants. • By physically separating the reduction half of a redox reaction from the oxidation half, we create a flow of electrons through an external circuit. • Used to accomplish electrical work

  4. Voltaic Cells • Two solid metals that are connected by the external circuit are called electrodes. • Anode: • Cathode: • Electrodes may or may not participate in the reaction • Zn/Cu • Pt or other conducting material

  5. Voltaic Cells • Each of the two components of a voltaic cell is called a half-cell • Oxidation half-cell • Reduction half-cell • For a voltaic cell to work, the solutions in the two half-cells must remain electrically neutral • Need migration of ions • Salt bridge or porous glass barrior

  6. Voltaic Cells • Anode  Acceptor of electrons: Oxidation • CathodeSource of electrons: Reduction • Anions always migrate toward the anode and cations toward the cathode.

  7. Voltaic (Galvanic) Cell

  8. Describing a Voltaic Cell • The following redox reaction is spontaneous: Cr2O72-(aq)+ 6I-(aq)  2Cr3+(aq) + 3I2(s) A solution containing K2Cr2O7 and H2SO4 is poured into one beaker, and a solution of KI is poured into another. A salt bridge is used to join the beakers. A metallic conductor that will not react with either solution is suspended in each solution, and the two conductors are connected with wires through a voltmeter to detect an electric current. The resultant voltaic cell generates an electric current. Indicate the reaction occurring at the anode, the reaction at the cathode, the direction of electron migration, the direction of ion migration, and the signs of the electrodes.

  9. Describing a Voltaic Cell • The two half-reactions in a voltaic cell are Zn(s)  Zn2+(aq) + 2e- ClO3-(aq) + 6H+(aq) + 6e-  Cl-(aq) + 3H2O(l) (a) Indicate which reaction occurs at the anode and which at the cathode. (b) which electrode is consumed in the cell reaction? (c) Which electrode is positive?

  10. Cell EMF Under Standard Conditions • Why do electrons transfer spontaneously during redox reactions? • Electrons flow from the anode of a voltaic cell to the cathode because of a difference in potential energy. • Potential energy higher at the anode • Electrons flow spontaneously toward the electrode with the more positive electrical potential.

  11. Cell EMF Under Standard Conditions • The difference in potential energy per electrical charge between two electrodes is measured in units of volts. 1V = 1 (J/C) Where V (volts), J (joule), and C (coulomb)

  12. Cell EMF Under Standard Conditions • The potential difference between two electrodes provides a driving force that pushes electrons through the external circuit. • Electromotive Force (emf) • Emf of a cell is denoted as Ecell(the cell potential) • For spontaneous reactions, the cell potential will be positive

  13. Standard EMF • Emf depends on • The particular cathode and anode half reactions • Concentrations of the reactants and products • Temperature • Tabulated values of standard reduction potentials denoted Eored to calculate Eocell Eocell = Eored (cathode) – Eored (anode)

  14. Standard Emf • Indirectly measure the standard reduction potential of a half-reaction • Reference point: 2H+ (aq, 1 M) + 2e-  H2 (g, 1 atm) • Assigned a standard reduction potential of exactly zero volts • Called a standardhydrogen electrode (SHE)

  15. Standard Emf • When determining standard reduction potentials from other half-reactions, write the reaction as a reduction even though it is “running in reverse” as an oxidation reaction. • Whenever an electrical potential is assigned to a half-reaction, write the reaction as a reduction. • Eored are intensive properties

  16. Calculating Eored from Eocell • For the Zn-Cu2+ voltaic cell, we have Zn + Cu2+  Zn2+ + Cu Eocell = 1.10V Given that Eored of Zn2+ to Zn is -0.76 V, calculate the Eored for the reduction of Cu2+ to Cu

  17. Calculating Eored from Eocell • A voltaic cell is based on the half-reactions: In+ In3+ + 2e- Br2 + 2e-  2Br- The standard emf for the cell is 1.46V and Eored for the reduction of bromine is +1.06V. Using this information, calculate Eored for the reduction of In3+ to In+.

  18. Calculating Eocell from Eored • Using the standard reduction potentials listed in Table 20.1, calculate the standard emf for the following voltaic cells: • Cr2O72- + 14H+ + 6I- 2Cr3+ + 3I2 + 7H2O • 2Al + 3I2 2Al3+ + 6I-

  19. Standard EMF • For each of the half-cells in a voltaic cell, the standard reduction potential provides a measure of the driving force for reduction to occur. • The more positive the value of Eored, the greater the driving force for reduction under standard conditions. • The more positive Eored value identifies the cathode

  20. Determining Half-Reactions at Electrodes • A voltaic cell is based on the following two standard half-reactions: Cd2+ + 2e-  Cd Sn2+ + 2e-  Sn • By using your chart, determine (a) the half-reaction that occurs at the cathode and the anode, and (b) the standard cell potential

  21. Strengths of Oxidizing and Reducing Agents • Use Eored values to understand aqueous reaction chemistry • The more positive the Eored value for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced and, therefore, to oxidize the other species. • Better oxidizing agent

  22. Strengths of Oxidizing and Reducing Agents • The half-reaction with the smallest reduction potential is most easily reversed as an oxidation. • The more negative the Eored, the stronger the ability to act as the reducing agent

  23. Strengths of Oxidizing and Reducing Agents

  24. Determining the Relative Strengths of Oxidizing Agents Using standard reduction potentials: • Rank the following ions in order of increasing strength as oxidizing agents: NO3-, Ag+, Cr2O72- • Rank the following species from the strongest to the weakest reducing agent:I-, Fe, Al

  25. Free Energy and Redox Reactions • Determining the spontaneity of redox reactions Eo = Eored (reduction process) – Eored (oxidation process)

  26. Spontaneous or Not? • Use standard reduction potentials to determine whether the following reactions are spontaneous under standard conditions. • Cu + 2H+ Cu2+ + H2 • Cl2 + 2I-  2Cl- + I2

  27. Spontaneous or Not? • Use standard reduction potentials to understand the activity series of metals • Activity series of metals: strongest reducing agent at the top Calculate standard emf for Ni + 2Ag+ Ni2+ + 2Ag

  28. EMF and ΔG • Relationship between G and EMF: ΔG = -nFE Where n = number of electrons transferred n the reaction, G = Gibbs free energy, E = EMF, and F = Faraday’s constant • Faraday’s constant is the quantity of electrical charge on one mole of electrons (a faraday) • 1 F = 96,485 C/mol = 96,485 J/V-mol

  29. Determining ΔGº and K (a) Use standard reduction potentials to calculate ΔGº and K at 298K for the reaction: 4Ag + O2 + 4H+ 4Ag+ + 2H2O (b) Suppose the reaction in part (a) was written: 2Ag + ½ O2 + 2H+ 2Ag+ + H2O What are values of Eº, ΔGº, and K when the reaction is written this way?

  30. Determining n and K • For the reaction 3 Ni2+ + 2Cr(OH)3 + 10OH- 3Ni + 2CrO42- + 8H2O (a) What is the value of n? (b) Given that ΔGº equals +87 kJ/mol, calculate K at a temperature of 298K

  31. Cell EMF Under Nonstandard Conditions • As a voltaic cell is discharged , the reactants of the reaction are consumed and the products are generated • The concentrations of these substances changes • EMF drops until E = 0, and the concentration of reactants and products are at equilibrium • How does cell emf depend on the concentration of reactants and products?

  32. The Nernst Equation • Dependence of cell emf on concentration • Nernst Equation: • At 298K with units of volts, the equation simplifies to:

  33. The Nernst Equation • The Nernst equation helps us understand why the emf of a voltaic cell drops as the cell discharges • Increasing the concentration of reactants or decreasing the concentration of products increasesthe driving force (higher emf) • Decreasing the concentration of reactants or increasing the concentration of the products decreases the driving force (lower emf)

  34. Voltaic Cell EMF Under Nonstandard Conditions • Calculate the emf at 298K generated by: Cr2O72-(aq)+ 14H+(aq) + 6I-(aq)  2Cr3+(aq) + 3I2(s) + 7H2O(l) When [Cr2O72-] = 2.0M, [H+] = 1.oM, [I-] = 1.0M, [Cr3+] = 1.0x10-5M

  35. Voltaic Cell EMF Under Nonstandard Conditions • Calculate the emf at 298K generated by: 2Al (s)+ 3I2(s) 2Al3+ (aq) + 6I- (aq) When [Al3+ ]= 4.0x10-3M and [I- ]0.010M

  36. Calculating Concentrations in a Voltaic cell • If the voltage of the following cell is +0.45V at 298K when [Zn2+] = 1.0M and PH2= 1atm, what is the concentration of H+? Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)

  37. Batteries and Fuel Cells • A battery is a portable, self-contained electrochemical power source that consists of one or more voltaic cells. • When cells are connected in series, the battery produces a voltage that is the sum of the emfs of the individual cells. • Multiple cells in series • Multiple batteries in series

  38. Batteries and Fuel Cells • The substances that are oxidized at the anode and reduced at the cathode determine the emf of the battery. • The usable life of the battery depends on the quantities of these substances. • Need a porous barrier between anode and cathode compartments • Primary and Secondary batteries

  39. Batteries and Fuel Cells • Lead-Acid Battery (12-v car battery, 6 voltaic cells in series that each produce 2V) • Alkaline Battery (most common primary battery) • Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion Batteries (secondary batteries)

  40. Lead-Acid Batteries

  41. Corrosion • Corrosion reactions are spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound. • Oxidation is a thermodynamically favored process in air at room temperature

  42. Corrosion • Prevent corrosion by forming a protective oxide layer that is impermeable to O2 and H2O • Examples: • Al3+ forms protective Al2O3 layer • Mg

  43. Corrosion of Iron • Rusting of iron requires both oxygen and water • pH of solution, presence of salts, contact with metals more difficult to oxidize than iron, and stress on the iron can accelerate rusting

  44. Corrosion of Iron • Cathodic protection: protecting a metal from corrosion by making it the cathode in an electrochemical cell. • The metal that is oxidized while protecting the cathode is called the sacrificial anode.

  45. Corrosion of Iron • Cathodic protection

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