Chapter 20 Notes- Electrochemistry

20.1 OxidationState and Oxidation-Reduction[p.844]

  1. Electrochemistry- study of the relationship between electricity and chemistry
  2. Redox reactions – must have an oxidation and reduction reaction occurs.
  3. Identified by assigning oxidation numbers
  4. May be two different elements gaining or losing electrons or a single element doing both
  5. Oxidizing agent (oxidant)- removes electrons from another substance, substance being reduced
  6. Reducing agent (reductant) substance that gives of electrons, substance being oxidized

20.2 Balancing Oxidation- Reduction Equations [p.846]

  1. Half-Reactions- show either the oxidation or the reduction alone
  2. For oxidation, electrons are a product (lost)
  3. For reduction , electrons are a reactant (Gain)
  4. Balancing Equations by the Method of Half-Reactions [Acidic conditions]
  5. Divide equation into 2 half-reactions
  6. Balance each half-reaction
  7. Balance elements other than H and O
  8. Balance O atoms by adding H2O
  9. Balance H atoms by adding H+
  10. Add electrons gained or lost
  11. Multiply each half-reaction by the coefficient that conserves electrons gained and lost
  12. Add coefficients form half-reactions to original equation
  13. Balance by inspection
  14. Balancing Equations for reactions Occurring in Basic Solutions
  15. Same as acid except use OH- and H2O

20.3Voltaic Cells[p. 851]

  1. Voltaic Cell (galvanic)
  2. Spontaneous reaction
  3. Transfers electrons through an external wire
  4. Electrodes – two metals connected by the external circuit
  5. Anode- electrode where oxidation occurs
  6. Disappears over time
  7. More reactive metal
  8. “-” electrode
  9. Cathode –electrode where reduction occurs
  10. Gains mass
  11. Less reactive metal
  12. “+“ electrode
  13. Each compartment of a cell is a half-cell
  14. Solutions in half-cells must remain neutral in a voltaic cell
  15. Salt bridge allows for the transfer of ions between solutions
  16. Positive ions travel through salt bridge form anode half-cell to cathode half-cell
  17. Negative ion travel in opposite direction through salt bridge
  18. A Molecular View of Electrode Processes shows that the anode is losing electrons and the cathode is gaining electrons

20.4Cell EMF Under Standard Conditions[p.855]

  1. General information
  2. Potential Difference is the “push causing the electrons to move.
  3. Movement goes form high to low
  4. Measured in volts 1V= J/C {remember one electron has a charge of 1.6 x 10-19C}
  5. Called electromotive force,EMF, or cell potential
  6. Symbol Ecell
  7. Standard EMF [Standard cell potential] E˚cell
  8. Voltage of cell under standard conditions (1M for concentrations, 1atm for gases, and 25˚C)
  9. In E˚cell, the “˚” indicates standard conditions
  10. Standard Reduction (Half –cell ) Potentials
  11. Calculated E˚cell using standard reduction half-cells (E˚red)
  12. E˚cell= E˚red (cathode) - E˚red (anode)
  13. For spontaneous reactions (voltaic cells) E˚cell > 0
  14. Standard Hydrogen Electrode (SHE) are used to get the reduction potentials because SHE”s E˚red= 0V
  15. Electrical potentials for a half-cell are always written as a reduction reaction
  16. Intensive property therefore changing the stoichiometric coefficients does not affect the standard reduction potential
  17. The more positive the valueE˚red, the greater the driving force for reduction under standard conditions
  18. Strength of Oxidizing and Reducing Agents
  19. More positive E˚red, greater the tendency for the reactant half-reaction to oxidize another species
  20. Good oxidizing agents
  21. The half-reaction with the smallest reduction potential is most easily reversed as an oxidation
  22. Group 1 and 2 metals are good reducing agents
  23. Reducing agents are difficult to store because oxygen is a good oxidizing agent

20.5Free Energy and Redox Reactions[p.862]

  1. General Information

a.Positive EMF indicate a positive reaction, negative EMF indicates a non-spontaneous reaction

b.General formula for determining the voltage potential of a redox reaction

E˚= E˚red (reduction) - E˚red (oxidation)

c.E˚ indicates standard EMF and E represents EMF under nonstandard conditions

d.metals in the activity series can be reduced by any metal below it

  1. EMF and ΔG
  2. “+” E and “-“ΔG indicate a spontaneous reaction
  3. Formula

ΔG = -nFE

ΔG˚ = -nFE˚

  1. “n” = number of electrons transferred
  2. “F” = Faraday’s Constant
  3. “E” = EMF and “E˚” = standard EMF
  4. “ΔG”= free energy and “ΔG˚”= standard free energy

20.6Cell EMF Under Nonstandard Conditions [p. 865]

  1. Nernst Equation-
  1. used to find emf of a cell under nonstandard conditions
  2. Can be used to find concentration s of a reactant or product
  3. Can be express in terms of ln and log base 10
  4. General rule
  5. Increasing the reactants and decreasing the products increases the driving force for the reaction
  6. Decreasing the reactants and increasing the product decreases the driving force
  1. Concentration Cells
  2. Cell based solely on the emf generated because of a difference in a concentration
  3. Standard emf =zero
  4. Cell operates until the concentration on both sides is equal

20.7Batteries and Fuel Cells [p. 870]

  1. General information
  2. Battery = portable, self-contained electrochemical power source that consists of one or more voltaic cells
  3. Cathode “+”
  4. Anode “-“
  5. Primary cells – batteries that cannot be recharged
  6. Secondary cells – rechargeable batteries
  7. Lead-Acid Battery
  8. Has solid lead (anode) and lead dioxide (cathode) electrodes in sulfuric acid therefore it is not a true concentration cell
  9. Rechargeable
  10. Alkaline Battery
  11. Anode – powdered zinc in contact with concentrated KOH
  12. Cathode – manganese II oxide and graphite
  13. Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium –Ion Batteries
  14. Nickel-cadmium
  15. Used in electronics
  16. Rechargeable
  17. Drawback use of Cadmium
  18. toxic metal
  19. Heavy
  20. Eventually lose charge
  21. Nickel-Metal-Hydride
  22. Used in hybrid vehicles
  23. Charges and recharges well
  24. Lasts up to 8 years
  25. Lithium -ion
  26. Lighter
  27. Has a greater energy density
  28. Used in electronics
  29. Hydrogen Fuel Cells
  30. Not batteries because they are not self-contained units
  31. Generate electricity in a highly efficient manner
  32. Used by NASA

20. 8 Corrosion [874]

  1. Corrosion –oxidation of a metal
  2. Can be damaging
  3. Some metal actually form a protective layer during oxidation
  4. Corrosion of Iron
  5. Rust often deposits where the largest supply of O2 exists. The cathode
  6. Salt increases the oxidation-reduction reaction because it provides ions to transport the electrons
  7. Preventing the Corrosion of Iron
  8. Painting protects the iron from water and oxygen
  9. Covering with tin or zinc, or creating galvanize metal
  10. Allows for the oxidation of the zinc or tin rather than iron
  11. Also use cathodic protection- used in underground pipes

20.9 Electrolysis [876]

  1. Background
  2. Electrolysis reactions- nonspontaneous reactions
  3. Electrolytic cells require energy to make reaction occur
  4. Anode- oxidation
  5. Cathode – reduction
  6. Cathode is attached to the negative terminal of the battery, anode is attached to the positive terminal of the battery
  7. Electroplating used to place thin coat of metal on surface of another metal
  8. Quantitative Aspects of Electrolysis
  9. The amount of substance that is reduced or oxidized in an electrolytic cell is directly proportional to the number of electrons passed
  10. Charge is measured in coloumbs
  11. Electrical Work
  12. ΔG= wmax
  13. wmax is negative for a voltaic cell, meaning that it is work done by the a system on its surrounding
  14. wmax is positive for anelectrolytic cell, meaning that its surrounding is doing work on the system
  15. power= work/ time which is measured in watts

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