Site for AP Chemistry students. The content of this unit is: Electrolytic and galvanic cells, Faraday's laws, standard half-cell potentials, the Nernst equation and the prediction of the direction of redox reactions. Electrochemical cells involve the transfer of electrons from one species to another. In these chemical systems, the species that loses electrons is said to be “oxidized” and the species that gain electrons is said to be “reduced”.

• Voltaic Cell Emf
• Voltaic Cell
• Voltaic Cell Measurements

20.4 Galvanic or Voltaic Cells

Chemical Concepts Demonstrated: Voltaic/galvanic cells, relative half-cell potentials

Demonstration:

• One of the dishes is filled with ZnSO₄ and the other with HCl.
• A strip of Zn metal is attached at one end to the posts of the electrochemistry template and is placed at the other end into the dish filled with Zn²⁺.
• The hydrogen electrode is attached, placed into the HCl solution, and H₂ gas is bubbled in. Insert the salt bridge.
• The Zn²⁺/Zn half-cell is replaced with a Cu²⁺/Cu half-cell.
• The H⁺/H₂ half-cell is replaced with a Zn²⁺/Zn. (picture 2)
  

Observations:
The potential in the absence of the salt bridge is 0.00 V. After the salt bridge is inserted, the potential of
the first set up is around + 0.76 V and the cell is a galvanic or voltaic cell. The Zn²⁺/Zn half-cell is the anode the H⁺/H₂ is the cathode.

In the second set up, both the magnitude and the sign of the potential change. The potential is now roughly - 0.34 V.

Picture 2 shows the third set up. The potential is now - 1.10 V. If the leads are changed, the cell potential becomes + 1.10 V and the cell becomes a galvanic or voltaic cell.

Explanations (including important chemical equations):

With the leads connected so as to produce a cell potential of + 0.76 V, the half reactions are:

If the standard-state potential for the H⁺/H₂ half-cell is assumed to be 0.00 V, and the potential for the anode half-reaction is equal in magnitude but opposite in sign to the standard-state potential for the Zn²⁺/Zn couple, then the standard-state reduction potential for the Zn²⁺/Zn half-cell must be - 0.76 V.

If the Zn²⁺/Zn half-cell is replaced with a Cu²⁺/Cu half-cell without reversing the leads to the voltmeter, the overall cell potential is - 0.34 V and the standard-state reduction potential for the Cu²⁺/Cu couple is therefore + 0.34 V.

If the H⁺/H₂ half-cell is replaced with a Zn²⁺/Zn half-cell, the overall cell potential should be - 1.10V.

To set up a voltaic cell using these half reactions, one would have to reverse the leads to the voltmeter.

Voltaic Cell Emf

1. Go to https://pages.uoregon.edu/tgreenbo/voltaicCellEMF.html Make the following voltaic cells: #1: Cu2+/Cu | | Ag+/Ag #2: Zn/Zn2+ | | Ag+/Ag #3: Zn/Zn2+ | | Cu2+/Cu
2. For each of the above, place the metal in a solution of its own ions. Make sure the cells are set up so that the cell potential is a positive value, indicating that the voltaic cell is set up correctly and the redox reaction is spontaneous. (Hint: In this simulation, the anode is black and the cathode is red.)
3. For each of the three voltaic cells, record the direction of electron flow, determine which electrode is the anode and which is the cathode, and record the cell voltage in the table on the next page.
4. For each electrode, determine whether oxidation or reduction is taking place. Record this in the table.
5. For each electrode, determine whether the electrode is dissolving away (becoming an ion and going in to solution) OR gaining mass (ions in solution are becoming neutral atoms that are deposited on the electrode). Record this in the table.
6. You must click the “Off” switch to reset for the next voltaic cell.

Voltaic Cell

Voltaic Cell Measurements

1. Go to http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/electroChem/electrolysis10.html
2. Click run experiment. pages.uoregon.edu/tgreenbo/electrolysis10.html
3. Construct a copper electroplating cell by placing a copper anode and iron cathode in a solution of Cu2+ ions. (anode is red and cathode is black)
4. Record the initial mass of the iron cathode in the data table.
5. Run the simulation at a current of 2.00 amperes at 2.00 V for 5:00 minutes. Record the final mass of the iron cathode. Record in the data table and calculate the mass of copper deposited on the iron.