Experiment: Voltaic Cells and HalfCell Potentials
In this experiment you will measure the potential of several voltaic cells and use your results to generate an activity series. Your measured cell potentials will be compared to the accepted values which can be found in the tables in your text.
Oxidationreduction reactions involve a transfer of electrons. In a spontaneous redox reaction, electrons flow from the oxidizing reactant (reducing agent) to the reducing reactant (oxidizing agent). If the two halfreactions can be separated, this flow of electrons, instead of occurring at the surface of the metal, occurs through an external wire and an electric current is generated. This is called a voltaic cell (or galvanic), and is exactly how a battery works. A battery, like the ones found in a flashlight or calculator, contains oxidizing and reducing substances. As the electrons are transferred they are tapped in order to provide the voltage necessary to power the flashlight or calculator.
A good analogy for the flow of electrons is the flow of water. Water flows spontaneously downhill. Dams and waterwheels are examples of ways that the energy of flowing water is tapped to generate power. Sometimes we want water to flow uphill. In this case we need to supply energy, in the form of a pump, to make this happen.
In order for a redox reaction to serve as a source of power, the reaction must be spontaneous. What if the reaction isnt spontaneous? In this case, we can use electricity to make the reaction go. An electrolytic cell is a device which uses electricity to drive a nonspontaneous redox reaction. For example, water can be separated into hydrogen and oxygen gas (a nonspontaneous reaction) using electricity.
In short, the field of electrochemistry has two important applications, the use of spontaneous redox reactions to generate electricity, and the use of electricity to force nonspontaneous redox reactions to occur.
In this laboratory you will construct and evaluate several voltaic cells. These cells will consist of two halfcells connected via a salt bridge. When the circuit is complete, electrons will flow spontaneously from the anode, where oxidation is occurring, to the cathode, where reduction is occurring. The potential, or EMF, of a cell is a measure of the tendency of the electrons to flow, and is dependent on the difference in oxidizing or reducing strength of the two halfcells. Continuing with our water analogy, cell potential is roughly analogous to a height difference between two reservoirs of water which are connected. The water will have a tendency (potential) to flow between the reservoirs, and that flow can be tapped to do work.
By convention one halfcell is assigned a potential of 0.00 V, and then all other halfcell potentials are measured relative to that assigned value. The halfcell which is assigned the 0.00 V potential is:
2 H+ + 2 e > H2
Note that the reaction for the halfcell is always written as a reduction reaction. The potential for other halfcells is determined by connecting an unknown halfcell to the hydrogen one (or another known halfcell), measuring the potential, and calculating the unknown halfcells potential.
As mentioned above the chemical connection between the two halves of a voltaic cell is made by a salt bridge. There are several ways to construct salt bridges; for simplicity we will use strips of filter paper soaked in an electrolyte. Because ions can travel through the filter paper the circuit is completed and current can flow, and the potential measured.
In this laboratory you will work with a partner.
Calculations and Analysis
Make a diagram of each cell studied, using Figure 21.5, p. 911 in your text, as your guide. For each cell:
Next, assume that the halfcell potential for Cu+2/Cu is 0.34 V. Based on this value, and your experimental cell potentials, compute the reduction potentials for all halfcells studied. Rank these potentials from highest to lowest (see p. A-11 in your text for the format) and compare them to the values in your text. As a part of your analysis discuss the agreement between your experimental results and the actual values in terms of both relative ranking, and measured values.
Return to the Laboratory Experiments Schedule