In an electrochemical cell, chemical energy is converted into electrical energy. This is accomplished by using a spontaneous chemical reaction to generate an electric current, which we can simply define here as electrons traveling though a wire.
To create the electrochemical cell two half-reactions will be set up in different containers. In one, an oxidation reaction will be used to generate a source of electrons. These free electrons will travel, through an external circuit, to the second container and will cause the reduction reaction to occur. The final requirement for our complete electrochemical cell will be a salt bridge that will permit ions to flow between the two half-cells, thus maintaining electrically neutral solutions.
Purpose
To build several electrochemical cells.
Safety
Follow normal lab safety guidelines. There are no specific safety hazards for this lab.
Equipment and Materials
Equipment
250-mL beakers, 2
glass U-tube for the salt bridge
cotton plugs for the salt bridge
copper wires, 2, insulated, with alligator clips
steel wool to clean electrodes
DC voltmeter
metal electrodes: copper, zinc, lead
Solutions
0.5M KNO3- for the salt bridge
Procedure
Each half-cell will be created by placing a metal electrode in an electrolytic solution containing the same metal's ions. For example the copper electrode will be placed in a copper(II) nitrate solution.
For each electrochemical cell you create you will require two half-cells. Set these cells beside each other - they will be connected by the U-tube.
Fill a beaker about two-thirds full of the electrolytic solution. Clean the electrode using the steel wool, then place the electrode in its appropriate solution.
Clip one end of each copper wire to the two electrodes using the alligator clips.
Fill the U-tube with KNO3 and stopper both ends with the cotton plugs. Turn the U-tube upside down and place one end in each half-cell.
Touch the other end of the copper wires to the voltmeter terminals. If the indicator on the voltmeter deflects in the wrong direction, switch the wires on the terminals. Read the highest voltage reading obtained - you'll need to do this quickly after connecting the wires to the voltmeter.
Repeat the experiment for other combinations of half-cells.
Results
Copy a data table similar to the one shown below into your lab notebook and use it to record your results.
Half-cells
Voltage
Cu|Cu2+
Zn|Zn2+
Cu|Cu2+
Pb|Pb2+
Zn|Zn2+
Pb|Pb2+
Questions and Conclusions
For each electrochemical cell you created:
Write out the two half-reactions for each electrochemical cell you created.
Identify each half-reaction as oxidation or reduction.
Identify each half-reaction as the anode and cathode.
Indicate the direction of the flow of electrons
Using a Table of Standard Reduction Potentials, calculate the theoretical voltage for each cell.
Compare the voltages you obtained with the theoretical voltage for each cell. What are some reasons that would account for any differences?
2.1 Introduction to Electrochemistry
During redox reactions, electrons pass from one substance to another. The flow of electrons - electric current - can be harnessed to do work. Electrochemistry is the branch of chemistry that deals with the conversion between chemical and electrical energy.
There are two major branches of electrochemistry, which we will examine in this section:
Electrochemical Cells
the energy released by a spontaneous chemical reaction is converted into electrical energy.
Examples: batteries
Electrolytic Cells
electrical energy is used to cause a non-spontaneous chemical reaction to occur.
Examples: recharging batteries & electroplating
2.2 Electrochemical Cells
The basic unit of all batteries is the electrochemical cell (also called a voltaic cell or galvanic cell). Electrochemical cells convert the energy of a spontaneous redox reaction into electricity. This will be accomplished as the electrons that are released from the oxidation half-reaction are passed to the reduction reaction which will absorb the electrons.
We will create an electrochemical cell based on the following redox reaction:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
This reactions involves two half-reactions:
In order for electrical work to be done by this reaction, we need to have the electrons travel through an external circuit. If we simply placed a piece of zinc metal in a solution containing copper(II) ions, a reaction would occur but electricity would not be created.
We'll walk through the set-up of our electrochemical cell:
1.
Begin by getting 2 beakers into which we will place metal strips in electrolytic solutions (solutions that conduct electricity due to the presence of ions). In one place a strip of zinc metal in a Zn(NO3)2 solution. In the other place a strip of copper metal in a Cu(NO3)2 solution.
Each beaker represents one of the two half cells. But because there is no way for electrons to move from one beaker to the other, our redox reaction cannot yet occur.
2.
We need to connect our two half-cells which we need to do in two ways.
First we will connect the two metal strips, our electrodes, with some wire. We'll also place a voltmeter here so we can detect the electric current once we are up and running. This will be our external circuit.
Second we add a salt bridge. A salt bridge is a U-shaped tube that contains an electrolytic solution (we'll use KNO3). This electrolytic solution will allow ions to flow between the two beakers. This is our internal circuit.
3.
An Ox Anode = Oxidation
The zinc half-cell undergoes oxidation. Here, the solid zinc electrode disintegrates, forming zinc ions and releasing electrons. By definition, the half-reaction that undergoes oxidation in an electrochemical cell is called the anode.
The anode is the source of electrons, making it the negative post of the electrochemical cell.
The copper half-cell undergoes reduction. Here, copper ions from the electrolytic solution become deposited on the copper electrode, forming more solid copper. Electrons are required for this to occur. By definition, the half-reaction that undergoes reduction in electrochemical cells is called the cathode.
The cathode is the positive post of the electrochemical cell as it consumes electrons.
Red CatReduction = Cathode
4.
It is important to understand the roles of the external circuit and the salt bridge.
External circuit - this is where the electrical work is done as electrons travel from one half-cell to the other. The electrons are produced at the zinc anode, where oxidation occurs. The electrons then travel through the wire of the external circuit to the copper cathode. The electrons are then available for the copper ions (from the Cu(NO3)2 solution) and solid copper is produced.
Internal circuit -
At the anode, Zn2+ ions are being produced and go into solution. This causes a build-up of positive ions in this solution. If this electrical imbalance is not corrected the reaction cannot continue. The excess positive charge attracts the negative NO3- ions (anions) from the salt bridge, thereby keeping the solution electrically neutral.
At the cathode the opposite occurs. As positive Cu2+ ions are removed from solution, to form solid Cu, the solution becomes overly negative. This attracts the positive K+ cations from the salt bridge, keeping this side of the cell neutral.
Once we have the entire electrochemical cell assembled - the two half-cells (the electrodes in their electrolytic solutions), the internal circuit (the salt bridge and half-cells), and the external circuit (the wire connected the two electrodes) - the cell is complete and the redox reaction will occur.
It is important to note and remember that unless the electrons can pass from one electrode to the other the reaction will not proceed.
There's a lot to remember when setting up electrochemical cells and it will take some practice to remember all of the details. Before we do that, however, there are a few more items to cover.
Original article and pictures take staff.prairiesouth.ca site
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