Electrochemistry

1
Electrochemistry

• Chapter 20
• Brown, LeMay, and Bursten

2
Definition

• The study of the relationships between
  electricity and chemistry
• Review redox reactions
• Review balancing redox reactions in acid and base

3
Voltaic Cell (also called Galvanic Cell)

• Device in which the transfer of electrons takes
  place through an external pathway.
• Electrons used to do work

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Summary of Cell

• Each side is a half-cell
• Electrons flow from oxidation side to reduction
  side determine which is which
• Salt bridge allows ions to move to each terminal
  so that a charge build up does not occur.
• Assignment of sign is this
• Negative terminal oxidation (anode)
• Positive terminal reduction (cathode)
• Salt bridge allows ions to move to each terminal
  so that a charge build up does not occur. This
  completes the circuit.

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Cell EMF

• Flow is spontaneous
• Caused by potential difference of two half cells.
  (Higher PE in anode.)
• Measured in volts (V)
• 1 volt 1 Joule/coulomb
• This is the electromotive force EMF (force
  causing motion of electrons through the circuit.

6
Ecell

• Also called the cell potential, or Ecell
• Determined by reactant types, concentrations,
  temperature
• Under standard conditions, this is Ecell
• 25 C, 1 M or 1 atm pressure
• This is 1.10 V for Zn-Cu
• Shorthand Zn/Zn2//Cu2/Cu

7
Reduction Potentials

• Compare all half cells to a standard (like sea
  level)
• 2H 2e- ? H2(g) 0 volts (SHE)
• The greater the Ered, the greater the driving
  force for reduction (better the oxidizing agent)
• In a sense, this causes the reaction at the anode
  to run in reverse, as an oxidation.
• Use this equation
• Ecell Ered (cathode) - Ered (anode)

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Trends
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Spontaneity

• Positive E value indicates that the process is
  spontaneous as written.
• Activity series of Metals listed as oxidation
  reactions
• Reduction potentials in reverse
• Example, Ag is below Ni because solid Ni can
  replace Ag in a compound. Actually, Ni is losing
  electrons and thus being oxidized by Ag. Ag is
  listed very high as a reduction potential.

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Relationship to ?G

• ?G -nFE
• n number of electrons transferred
• F Faraday constant 96,500 C/mol or 96,500
  J/V-mol
• Why negative? Spontaneous reactions have E and
  ?G.
• Volts cancel, units for ?G are J/mol
• Standard conditions ?G -nFE

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Nernst Equation

• Nonstandard conditions during the life of the
  cell this is most common
• Derivation
• E E - (RT/nF)lnQ
• Consider Zn(s) Cu2 ? Zn2 Cu(s)
• What is Q?
• What is E when the ions are both 1M?
• What happens as Cu2 decreases?

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Concentration Cells

• Same electrodes and solutions, different
  molarities.
• How will this generate a voltage? Look at Nernst
  Equation. E E - (RT/nF)lnQ
• When will it stop?
• Basis for a pH meter and regulation of heartbeat
  in mammals

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EMF and equilibrium

• When cell continues to discharge, E eventually
  reaches 0. At this point, because ?G -nFE, it
  follows that ?G 0.
• Equilibrium!
• Therefore, Q Keq
• Derivation
• logKeq nE/0.0592

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Batteries

• Portable, self-contained electrochemical power
  source
• Batteries in series, voltage is added.

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Things to consider

• Size (car vs. heart)
• Amount of substances before it reaches
  equilibrium
• Toxicity (car vs. heart)
• A lot a voltage or a little (car vs. heart)
• Example alkaline camera battery
• Dry no water

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Fuel Cells

• Not exactly a battery, because it is open to the
  atmosphere
• How does the combustion of fuel generate
  electricity? heats water to steam which
  mechanically powers a turbine that drives a
  generator 40 efficient
• Voltaic cells are much more efficient
• http//www.fueleconomy.gov/feg/fuelcell8.swf

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Corrosion

• Undesirable spontaneous redox reactions
• Thin coating can protect some metals (like
  aluminum) forms a hydrated oxide)
• Iron -

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Protection

• Higher pH
• Paint surface
• Galvanize (zinc coating) why?
• Zinc is a better anode
• Called cathodic protection sacrificial metal

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More dramatic
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Electrolysis

• Cells that use a battery or outside power source
  to drive an electrochemical reaction in reverse
• Example NaCl ? Na Cl-
• Reduction at the cathode, oxidation at the anode
• Voltage source pumps electrons to cathode.

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Diagram
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Solutions

• High temperatures necessary for previous
  electrolysis (ionic solids have high MP)
• Easier for solutions, but water must be
  considered
• Example NaF
• Possible reductions are
• Na e- ? Na(s) (Ered -2.71 V)
• 2H2O 2 e- ? H2(g) 2 OH- (Ered -.83 V)
• Far easier to reduce water!
• continue

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Continued

• Look at possible oxidations
• 2F- ? F2(g) (Ered 2.87 volts)
• 2H2O ? O2(g) 4H 4e- (Ered 1.23 volts)
• Far easier to oxidize water, or even OH-!
• So for NaF, neither electrode would produce
  anything useful, and doesnt by experiment
• With NaCL, neither electrode is favored over
  water. However, the oxidation of Cl- is
  kinetically favored, and thus occurs upon
  experimentation!
• Use Ered values of two products to find Ecell
  (minimum amount of energy that must be provided
  to force cell to work)

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Active electrodes

• If electrode is not inert, it can be coated with
  a thin layer of the metal being reduced, if its
  reduction potential is greater than that of
  water.
• This is called electroplating
• Ecell 0, so a small voltage is needed to push
  the reaction.

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Quantitative relationship