Electrochemcial Cell Demonstration Voltaic Cell: Zinc/Copper E° = 1.10 V

A standard cell comprising of two half-cells: zinc metal electrode in 1.0 M ZnSO​4 solution, a copper metal electrode in a 1.0 M CuSO4 solution, and a connecting salt bridge.   

The cell reaction is  Zn(s) +  Cu2+(aq)  --> Zn2+(aq) + Cu(s)

The electrodes are connected to a voltmeter.  E°cell = +1.10 Volts.  The voltage of the cell can be projected using a voltmeter or a Vernier LabPro interface.

A Standard Zn/Cu Voltaic cell is used to show the E°cell generated is the difference in the standard reduction potentials of two half-reactions: Zn2++ 2e- -> Zn  and  Cu2++ 2e- -> Cu  for  Zn|Zn2+||Cu2+|Cu electrochemical cells    

cell = E°cathode (as a reduction potential) - E°Anode(as a reduction potential)  = +0.34 V - (-0.767 V) = + 1.10 V

Curriculum Notes 

 Prior to doing these electrochemical cells demonstrations, it is recommended that the demonstration showing the reaction of zinc with CuSO4(aq) and the no reaction of copper with ZnSO4(aq) be shown and discussed. This is a great demo to introduce the concept of electrochemical cells.   Pair this demonstration with computer animations showing a representation of the oxidation half-reaction occurring at the zinc anode and the reduction half-reaction occurring at the copper cathode.  

Electrochemical Cells Computer Simulation: voltaic cells Zn Cu Ag OLD FLASH-Based

©2009 Greenbowe  Chemistry Education Instructional Resources  


A new HTML5-based computer simulation is being developed for this computer simulation.

Computer animations of a standard cell comprising of two half-cells: zinc metal electrode in 1.0 M ZnSO​4 solution, a copper metal electrode in a 1.0 M CuSO4 solution, and a connecting salt bridge.   The electrodes are connected to a voltmeter.  E°cell = +1.10 Volts.  

A guided-inquiry worksheet accompanies this computer simulation.  Download the file from the menu.

  The URL is  http://introchem.chem.okstate.edu/DCICLA/voltaicCell20.html

After presenting this demonstration, the standard hydrogen electrode demonstration can be used to show the SHE|H+,1.0 M |Cu2+,1.0 M |Cu cell and the Zn|Zn2+||H+,1.0 M| SHE cell.  

Computer animations for the  Zn/Cu Voltaic cell   beta versions  (drafts)

Zn|Zn2+   oxidation half-reaction at the zinc electrode   https://vimeo.com/220550690            

Cu2+|Cu  eduction half-reaction at the copper electrode   https://vimeo.com/220550267

animation of the migration of ions in the salt-bridge  https://vimeo.com/220548484

animation of the movement of electrons in a wire   https://vimeo.com/220550589

Student Difficulties with Galvanic Cells or Electrochemical Cells

1. Cell potentials are obtained by adding individual reduction potentials

2.  Anodes, like anions, are always negatively charged and release electrons, and cathodes, like cations, are always positively charged and attract electrons.

3. The anode is positively charged because it has lost electrons. The cathode is negatively charged because it has gained electrons.

4.  Electrons flow through the salt bridge and the electrolyte solutions to complete the circuit,

Learning Objectives

1. Given a diagram of a simple electrochemical cell involving two metal electrodes and the corresponding solution of the metal ions identify: the site of oxidation reduction, the anode, the cathode, movement of electrons, migration of ions, the chemical equation representing the cell reaction.
2. Calculate the emf of a cell, given a table of standard reduction potentials. 

3. Draw a particle diagram representing the dynamic events occurring at each electrode and in the salt-bridge. 

AP Chemistry learning Objectives

3.12 The student can make qualitative or quantitative predictions about galvanic cell reactions based on half-cell reactions and electrode potentials.

3.13 The student can analyze data regarding galvanic cells to identify properties of the underlying REDOX reactions.

Lead Time 
One day of lead time is required for this project.

Due to a differences in electromotive force between zinc and copper, zinc is a more active metal compared to copper, a spontaneous oxidation-reduction process occurs when the zinc electrode is connected to the negative terminal and the copper electrode is connected to the positive terminal of the voltmeter. The spontaneous flow of electrons from anode to cathode and the migration of cations and anions in the solutions and salt bridge generates a current with a voltage near the theoretical Eo cell = 1.10 V at room temperature.

                 Eored (V)

 Anode Zn → Zn2+ + 2 e-       - 0.76

Cathode Cu2+ + 2 e--→ Cu    + 0.34

Video Links



  • 2 ea. 400 mL beakers, one containing a strip of zinc metal partially immersed in a 1.0 M ZnSO4 solution, the other containing a strip of copper metal partially immersed in a 1.0 M CuSO4 solution.
  • a salt bridge comprising a u-tube filled with 1 M KNO3 solution.  The ends of the u-tube should be tightly plugged with cotton that has been saturated in the KNO3 solution.
  • a Vernier voltage probe.
  • a Vernier "Go Link" interface.
  • an empty 250 mL beaker to hold the salt bridge upright when not in use.
  • a computer with the Vernier "Logger Lite" program that is capable of projecting the screen image.
  • Place the two beakers side by side.
  • Connect the probe leads to the metal strips where they stick up above the surface of the solutions.
  • Place the salt bridge upside down so that one arm is in each of the solutions.
  • Students should note that no voltage registers until the salt bridge is inserted into the solutions.
Safety Precautions 
  • Always wear goggles when performing chemistry demonstrations.
  • Gloves should be worn to protect your hands from the solutions.


1. Sanger, M.J. and Greenbowe, T.J. (1997).  “Student Misconceptions in Electrochemistry: Current Flow in Electrolyte Solutions and the Salt Bridge.” Journal of Chemical Education74(7), 819-823.

2. Sanger, M. J. and Greenbowe, T.J.  (1997).   “Common Student Misconceptions in Electrochemistry: Galvanic, Electrolytic, and Concentration Cells.” Journal of Research in Science Teaching34(4), 377-398.

3. Sanger, M.J. and Greenbowe, T.J. (1999).  “An Analysis of College of Chemistry Textbooks as Sources of Misconception and Errors in Electrochemistry.”  Journal of Chemical Education76(6), 853-860.

4.  Abraham, M.; Gelder, J.; Greenbowe, T. (2007).  During Class Inventions and Computer Lab Activities for First and Second Semester General Chemistry. Hayden-McNeil: Plymouth, MI.

5. Shakhashiri, B. Z. In Chemical Demonstrations: A Handbook for Teachers of Chemistry; The University of Wisconsin Press: 1992; Vol. 4, p 101-106.

© Copyright 2012 Email: Randy Sullivan, University of Oregon Chemistry Department and UO Libraries Interactive Media Group