Electrochemical Cells and Nernst's Equation

Galvanic or Voltaic Cells

December 6, 2002

• We can form simple electrochemical cells by creating two separate electrodes
  • each represents a chemical half reaction
  • each can involve oxidation or reduction
• We will begin with the simplest electrodes
  • a metal (copper, zinc, magnesium, iron, lead, silver, cadmium, iron)
  • inserted in a solution of the corresponding cation (Cu²⁺, Zn²⁺,Mg²⁺,Pb²⁺,Ag⁺, Cd²⁺,Fe²⁺)
• We connect the two electrolytes with a salt bridge
  • this contains an inert electrolyte such as NaNO₃
  • this may be a U-tube
    • usually filled with a gelled solution so fluid cannot flow and bubbles won't form
    • you can also use absorbent paper or cloth soaked in NaNO₃ as a wick

• We will measure the voltage of these cells
  • we use digital pH meters
    • they have a simple millivolt mode

Experimental

• Select a pair of metals for your cell
  • try to avoid using the same pair of metals as others in your extended family
  • We will provide small, wide mouth bottles with various concentrations of the cations
    • make your cell directly in the bottle
    • (others will use these bottles so rinse and blot electrodes carefully and don't contaminate the solutions)
  • Insert a matching metal wire or strip of sheet metal

• Prepare the second electrode similarly
• Connect your electrodes with a salt bridge

• Set up the pH meter with electrical connectors instead of the pH electrode
  • turn the meter to the millivolte mode
    • (the millivolt mode does not have a calibration process)
  • Connect one lead to each electrode and measure the voltage

• Interpreting the sign of the voltage
  • the meter displays the voltage relative to the black lead.
  • if your meter reads a positive value, the electrode connected to the red lead is more positive than the electrode connected to the black lead
    • the electrode at the red lead is therefore consuming electrons
    • that is this (red) electrode involves reduction

• if the reading is negative, it simply means that oxidation is occurring at the electrode connected to the red lead.

• We will work examples with + voltage readings
  • You should be able to work the problem by analogy if your reading is negative
  • You could also chicken out and reverse the leads so you get a positive reading
    • If your cell voltage is relatively low (0.1-0.2 volts) the sign can easily change when we change electrode concentrations. It would be desirable to be able to handle positive and negative voltages.

Experimental:

• metals used
• solution concentrations
• which electrode is connected to the red lead
• the experimental voltage, including sign

• Repeat at two different concentrations
  • replace one of the solutions with the same cation at a different concentration
    • please be careful to rinse and blot the electrode to preserve the integrity of the solutions
  • record the new voltages

• Compare with the values expected by the Nernst Equation
• Keep the meter and one of the electrodes for the next part (workingwith Alternaivite Electrodes)

Sample calculations

• One electrode is Cadmium metal in 0.050M Cd⁺²
  • this electrode is connected to the red lead of the voltmeter
• The other electrode is Chromium metal in 0.015M Cr³⁺ solution
  • this electrode is connected to the black lead of the voltmeter
• The cell voltage is positive, meaning that
  • the Cadmium ion is undergoing reduction
    • Cd²⁺ + 2e ---->Cd
  • the Chromium electrode is being oxidized
    • Cr ----> Cr³⁺ + 3e

2 {Cr ----> Cr ³⁺ + 3e } ..........E^o=- (- 0.74)

2Cr + 3 Cd²⁺---> 2 Cr³⁺ + 3 Cd ........Eo =0.34

Nernst Equation

E = Eo -(0.0590/6) log10 { [Cr³⁺]² / [Cd²⁺]³}

E= +0.34 -0.00983 log10 ( 0.015² / 0.05³ ) = ______

E= 0.34 -0.00983 log10 (0.596) = + ______