Introduction to Electrochemical Cells

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Introduction to Electrochemical Cells

• Thanks to Eric Edens 07

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Experimental Goals

• Reinforce topics from last semester
• Provide experience working with electrochemical
  cells
• Investigate dependence of cell potential on
  concentration
• Nernst Equation
• Using cell potential to determine equilibrium
  constants
• Introduce related material
• Concentration Cells
• Ionic Strength

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Two Types of Electrochemical Cells
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Information on Procedure

• Working with a partner
• Both make measurements on all half-cells to
  compare results
• Using small volumes for measurements
• Take only what you need
• Avoid making cells that form a precipitate
• What might they be
• Silver with halides
• If precipitates form combine half-cell of
  interest with different half-cell.
• Why

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Part A Potential Measurements

• Use filter crucible and beaker to construct
  electrochemical cells
• Demonstrate proper use
• Red lead is Cathode
• Black lead is Anode
• Make sure that your multimeters are set to
  measure VDC
• All half-cells will be expressed vs AgAg
  half-reaction in the end
• Need to make measurements that allow this
  conversion

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Example

• 2Ag Cu 2Ag Cu2 E1 0.46V
• Zn2 Cu Zn Cu2 E2 -1.10V
• 2Ag Zn 2Ag Zn2 E3

Under standard conditions 1atm pressure each
metal 1.0M
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Example

• 2Ag Cu 2Ag Cu2 E1 0.46V
• Zn Cu2 Zn2 Cu - E2 1.10V
• 2Ag Zn 2Ag Zn2 E3 E1 (- E2)
• E3 0.46 1.10 1.56V

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Part A Important Notes

• For Fe3/Fe2 half-reaction you will need to
  combine equal volumes of the two solutions
  immediately prior to making your measurements
• All measurements should be made three times to
  ensure that you have obtained the correct values
• I2/I- half-reaction
• What is really going on

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Remember the Nernst Eq
Demonstrates cell voltage depends on concentration
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Part B Concentration Dependence

• Exploring the Nernst Equation with real cells
• Easier to use concentration cells
• same reactions in the anode and cathode
  compartments
• DEcellEcathode-Eanode0.00V if both Cu2
• So for the generic cell
• M/M(lesser concentration) M(greater
  concentration)/M

• Ag in Beaker Cu2 dilutions in the crucible

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Part B Ionic Strength

• When salt concentrations are large enough we
  begin to see deviations from Nernstian behavior
• Knowing the ionic strength of the solution we
  can use correction factors to help make accurate
  calculations
• Beyond the scope of this experiment but it is
  important to maintain the same ionic strength for
  all measurements

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Calculating Ionic Strength
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How much KNO3 needs to be added to your dilutions

• M1V1M2V2
• The difference in ionic strengths is .3-.03 .27
• All dilutions made in 100ml volumetric flasks
• KNO3 is 0.4M so
• 0.4MV10.27M100ml
• V167.5 ml

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Parts C D Equilibrium Constants

• Using concentration cells determine the value of
  Q.
• Use Q to determine the unknown concentration

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Part C Ksp for AgCl

• Starting with 50 ml of 0.1M KCl add AgNO3
  dropwise until a precipitate forms.
• Combine this half-cell with one inside of your
  crucible containing 0.1M AgNO3
• The resulting cell is
• AgAg(unknown concentration) Ag(0.1 M)Ag

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Helpful hints
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Part D Formation Constant for Cu(NH3)42

• The overall reaction is
• Use the same type of arrangement as in Part C
  replacing Ag with Cu in the cathode half-cell
• Use a table of changes to determine the
  concentrations of each species at the end.
• The final concentration of Cu2 is the one that
  you will be determining.

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Table of Changes Using Moles
Assumes 30ml Cu(NO3)2 solution and 3 ml NH3
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Safety and Waste Disposal

• Two separate waste crocks in the lab
• One for Silver Nitrate
• One for all other solutions
• Silver nitrate stains hands so be careful!

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Before Starting

• Find a partner
• Check out kit from stockroom
• Have fun!