Chemistry 1220

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Chemistry 1220

• Chemical Principles II

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Electrochemistry

• Is it all about batteries?
• Lithium
• Alkaline
• Carbon/Zinc
• Lead Acid
• Lithium Polymer
• Lithium Thionyl Chloride
• Ni/Cd
• Ni/MH
• https//www.batteries4everything.com/index.html

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Electrochemistry

• The conversion of a spontaneous reaction into
  useful energy is very important but only ½ of
  the picture.
• The production of a lot of materials
• Chlorine
• Bromine
• Aluminum
• Sodium / Potassium..
• And a lot of other things are accomplished by
  electrolysis.

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Galvanic Cells ( Voltaic Cells)

• Two types of electrochemistry applications.
• Galvanic process
• Electrical energy production from a spontaneous
  reaction.
• Electrolytic process
• electrical energy is consumed to (force) a
  non-spontaneous reaction to take place.
• Reversible
• Galvanic process occurs spontaneously and then is
  reversed electrolytically.

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Galvanic Cells

• Oxidation-reduction review.
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• Remember LEO / GER
• Oxidation ½ reaction
• Zn(s) ? Zn2(aq) 2 e-
• Reduction ½ reaction
• Cu2(aq) 2e- ? Cu(s)

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Galvanic Cells
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www.wantwit.com
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Galvanic Cells

• Besides the LEO the Lion says GER
• Whats a good way to remember what happens at the
  and battery terminals?
• Red Cat
• Reduction Occurs at the Cathode
• An Ox
• Oxidation Occurs at the Anode

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Galvanic Cells

• Cathode
• Positive sign (assigned)
• Reduction occurs.
• Electrons are consumed.
• Cations migrate toward it.
• Anode
• Negative sign (assigned)
• Oxidation occurs.
• Electrons are produced
• Anions migrate toward it.
• So what is the actual charge at each of these
  electrodes?

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Shorthand Notation for Galvanic Cells

• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• Salt Bridge
• Zn(s) Zn2(aq) Cu2(aq) Cu(s)
• Anode ½ cell Cathode ½ cell
• Electrons flow ? this way

Phase Boundaries
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Cell Potentials and Free-EnergyChanges for Cell
Reactions

• What drives the electrons from the anode to the
  cathode?
• Is the energy difference between the two halves
  of the process.
• It is called electromotive force (emf)
• Is also most commonly called cell potential (E)
  or cell voltage.
• 1J 1C x 1V
• ?G -nFE F Faraday constant

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Cell Potentials and Free-EnergyChanges for Cell
Reactions

• Example 18.3
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s
• Using the cell potential and the 2 e-s
  transferred at 25oC you can calculate that ?Go
  -212 kJ.
• Is there another way to calculate ?Go for this
  reaction?

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Cell Potentials and Free-EnergyChanges for Cell
Reactions

• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s
• ?G is directly related to the potential of the
  cell.

Zn(s) Cu2(aq)
?G
Zn2(aq) Cu(s
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ElectrochemistryMore Basics

• Coulomb (C) a unit of charge
• 1.6022 x 10-19 C charge of 1 electron
• Faraday (F) charge of 1 mole of electrons
• F 9.648 x 104 C/mole
• q symbol for charge
• q nF

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ElectrochemistryMore Basics

• Current quantity of charge flowing through a
  system per second.
• C/s ampere (amp, A)
• Difference in electrical potential (E) measured
  in volts (V) or J/C
• Work moving current from 1 potential to a
  higher potential.
• Work Eq where Work (J), E(V) and q(C)
• Work -?G Eq -nFE

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ElectrochemistryMore Basics

• Ohms law
• V IR
• P VI I2R
• P Power (watts, J/s)

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Standard Reduction Potentials

• Eocell Eoox Eored
• Zn(s) ? Zn2(aq) 2 e- Eoox 0.76 V
• Cu2(aq) 2e- ? Cu(s) Eored 0.34 V
• __________________________________
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• Eocell 1.1 V
• See Table 18.1

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Standard Reduction Potentials

• Zn(s) ? Zn2(aq) 2 e- Eoox 0.76 V
• This is a standard oxidation potential.
• Most tables are written to reflect a reduction.
• i.e. this Zn oxidation would be written as a
  reduction.
• Zn2(aq) 2 e- ? Zn(s) Eored - 0.76 V

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Standard Reduction Potentials

• The tables of standard reduction potentials is a
  type of activity series. Where did the base line
  come from?
• The sea level was set as the reduction ½
  reaction
• 2 H(aq) 2e- ? H2(g) Eored - 0.0 V
• The standard hydrogen electrode (SHE)

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Standard Reduction Potentials
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Standard Reduction Potentials
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Standard Reduction Potentials
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Using Standard Reduction Potentials

• A table of reduction potentials will indicate
  spontaneity of a multitude of reactions.
• Book gave an example of 100 ½ cell reactions.
• Given the table 18.1 with 25 ½ reactions.
• How many complete reactions could you calculate
  the value of Eo?
• Use table to prepare some rxns., predict reducing
  agents, select oxidizing agents..

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Interesting Sites

• http//www.chem.ox.ac.uk/vrchemistry/potential/Tex
  t/intro1.htm

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Cell Potentials and Composition of the Reaction
Mixture The Nernst Equation

• ?G ?Go RT lnQ
• -nFE -nFEo RT lnQ
• E Eo RT/nF lnQ
• At 25oC E Eo 0.0592V/n logQ

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Electrochemical Determination of pH

• This is a simple application of the Nernst
  Equation.
• pH Ecell Eref / 0.0592V _at_25oC
• Glass electrodes are used rather than the
  Pt/H2/H reference.

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Electrochemical Determination of pH

• In pH determination two half reactions are placed
  in self contained electrodes.
• Ag/AgCl reference
• Calomel electrode
• 2 x Ag(s) Cl-(aq) ? AgCl(s) e- Eo -0.22V
• Hg2Cl2(s) 2e- ? 2Hg(l) 2 Cl-(aq) Eo 0.28V
• The solution these electrodes are placed in
  become the salt bridge and the difference of H
  charge determines the Ecell. pH is electronically
  converted and displayed.

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Electrochemical Determination of pH
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Standard Cell Potentials and Equilibrium Constants

• Again manipulating ?G
• Eo 0.0592V/n log K at 25oC.
• The equation is used to determine equilibrium
  constants from cell voltage for reactions where K
  is very large or very small.
• This is necessary in these cases since the
  measurement of reactant/product concentration
  becomes very difficult.

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Standard Cell Potentials and Equilibrium Constants
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Batteries

• Lead Storage
• Dry-Cell
• Nickel-Cadmium
• Lithium
• Fuel Cells

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Batteries

• Lead Storage
• Pb PbO2 2H2SO4 ? 2PbSO4 2H2O
• Eo 1.924V

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Batteries

• Dry-Cells
• Leclanche
• Zn 2MnO2 2NH4 ? Mn2O3 2NH3 Zn2 H2O
• Alkaline
• Zn 2MnO2 ? Mn2O3 ZnO
• Mercury
• Zn HgO ? Hg(l) ZnO

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Batteries

• Nickel-Cadmium
• Cd 2NiO(OH) 2H2O ? 2Ni(OH)2 Cd(OH)2
• Lithium
• Li MnO2 ? LiMnO2
• Fuel Cells
• One or both of the ½ cells contains a standard
  fuel.
• 2 H2 O2 ? 2 H2O

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Batteries

• Location Half Cellreactions Voltage
• ____________________________________________
• Anode Zn2 2OH- Zn(OH)2 1.25
• Cathode 1/2 O2 H2O 2e 2 OH -0.4
• ____________________________________________
• Overall 2Zn O2 2H2O 2Zn(OH)2 1.65

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Batteries

• Location Half Cell reactions Voltage
• ___________________________________________
• Anode Al 4 OH- Al(OH)4- 3e -2.35
• Cathode 3/4 O2 3/2 H2O 3e 3OH -0.40
• ___________________________________________
• Overall Al 3/2 H2O 3/4 O2 Al(OH)3 2.75

http//www.powerstream.com/BatteryFAQ.html
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Batteries

• http//kahuna.sdsu.edu/hev/energy.html
• Energy studies from the engineering department of
  SanDiego State U.

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Corrosion

• The oxidation of metals by oxygen and/or water.
• Check the reduction half reaction in table 18.1
  or Appendix D for the O2/H2O ½ reaction. Also
  check the H2O/OH- ½ reaction.
• Rust deposits vs. location of pitting!

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Corrosion

• 2 H2O(l) 2 e- ? H2(g) 2 OH-(aq)
• Eo -0.83V at pH 14
• Eo -0.42V at pH 7
• O2(g) 4 H(aq) 4e- ? 2 H2O(l)
• Eo 1.23V at pH 0
• Eo 0.81V at pH 7
• Fe(s) ? Fe2(aq) 2e-
• Eo 0.45V

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Corrosion
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Corrosion

• If it is true that ¼ of the steel used in the
  U.S. is for replacement it means over 5 billion
  is spent on corrosion effects for iron alone and
  this is on the replacement cost not the damage
  cost.
• Wouldnt it be much more cost effective to find a
  protection for metals rather than replace them?

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Corrosion

• Prevention of corrosion.
• Galvanizing
• Sacrificial Anode
• Metals used chromium, tin, zinc magnesium
• By oxidizing the surface of iron with Na2CrO4 an
  oxide coating can be produced to protect the
  surface.
• Alloy of Chromium/Iron (stainless)

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Corrosion
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Corrosion
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Corrosion
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Electrolysis and Electrolytic Cells
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Electrolysis and Electrolytic Cells
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Electrolysis and Electrolytic Cells
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Electrolysis and Electrolytic Cells

• Electrolysis of Water.
• 2H2O(l) ? H2(g) O2(g)
• Anode?
• Cathode?

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Commercial Applications of Electrolysis

• Manufacturing Chemicals
• Cl2 Br2 Na H2
• Anionic Polymerization
• Purification and Plating of Metals
• Aluminum Hall-Heroult
• Copper
• Magnesium Manufacture
• Gold, Silver and Platinum purification and
  plating
• Chrome plating

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Downs Cell
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Copper Plating
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Zinc Plating ? Brass
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Electrolysis of Water
pH indicator - Bromothymol blue Yellow at low
pH Blue at High pH
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Silver Plating
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AtlantisTM NaCl Electrolysis for Swimming Pools
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Arizona HydrogenFire from Water
Hydrogen generators for welding applications
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Quantitative Aspects of Electrolysis

• Faraday
• English bookbinder who became interested in
  electricity. He obtained an assistantship in
  Davy's lab, then began to conduct his own
  experiments. He wrote a review article on current
  views about electricity and magnetism in
  1821, for which he reproduced Oersted's
  experiment. He was one of the greatest
  experimenters ever. Because he was self trained,
  however, he had no grasp of mathematics and could
  therefore not understand a word of Ampère's
  papers.
• invented the dynamo (a device capable of
  converting electricity to motion) in 1821,
  discovered electromagnetic induction in 1831, and
  devised the laws of chemical electrodeposition of
  metals from solutions in 1857.

Excerpt from http//scienceworld.wolfram.com/bio
graphy/Faraday.html
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Quantitative Aspects of Electrolysis

• Something for calculations and test.
• Is all about balanced equations and electrons.
• Charge (in coulombs) Current (amperes) x Time
  (sec.)
• C A x T
• 1 mole of electrons has a charge of 96,500 C.
• 1 Faraday.

Current x time ? Charge ? Moles of e- ? moles of
product ? mass of material