Galvanic cell
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The Galvanic cell, named after Luigi Galvani, is a part of a battery consisting of an electrochemical cell with two different metals connected by a salt bridge or a porous disk between the individual half-cells. It is sometimes also called a Voltaic cell.
Common usage of the word battery has evolved to include a single Galvanic cell, but the first batteries had many Galvanic cells.[1][2] The Galvanic cell should not be confused with the electrolytic cell, which decomposes chemical compounds by means of electrical energy.
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[edit] History
In 1780, Luigi Galvani discovered that when two different metals (copper and zinc for example) were connected together and then both touched to different parts of a nerve of a frog leg at the same time, they made the leg contract.[3] He called this "animal electricity". The Voltaic pile invented by Alessandro Volta in the 1800s is similar to the galvanic cell. These discoveries paved the way for electrical batteries.
[edit] Description
A Galvanic cell consists of two half-cells. Each half-cell consists of the following:
- An electrode, which in the figure are the plates of zinc (Zn) and copper (Cu).
- An electrolyte, which in the figure are aqueous solutions of zinc sulfate (ZnSO4) and copper(II) sulfate (CuSO4).
For the Daniell cell, depicted in the figure, the zinc atoms have a greater tendency to go into solution than do the copper atoms. More precisely, the electrons on the zinc electrode have a higher energy than the electrons on the copper electrode. Because the electrons have a negative charge, to pass them on the zinc electrode must have a more negative electrical potential than the copper electrode. However, in the absence of an external connection between the electrodes, no current can flow.
When the electrodes are connected externally — as in the figure, with wire and a voltmeter, the electrons tend to flow from the more negative electrode (zinc) to the more positive electrode (copper). Because the electrons have negative charge, this produces an electric current that is opposite to the electron flow. At the same time, an equal ionic current flows through the electrolyte.
For every two electrons that flow from the zinc electrode through the external connection to the copper electrode, on the electrolyte side a zinc atom must go into solution as a Zn2+ ion, at the same time replacing the two electrons that have left the zinc electrode by the external connection. By definition, the anode is the electrode where oxidation (removal of electrons) takes place, so in this galvanic cell the zinc electrode is the anode. Because the copper has gained two electrons from the external connection, it must release two electrons at the electrolyte side, where a Cu2+ ion, from the copper(II) sulfate, plates onto the copper electrode. By definition, the cathode is the electrode where reduction (gain of electrons) takes place, so the copper electrode is the cathode. Electrons will flow from the anode to the cathode.
An essential piece of a Galvanic Cell is the salt bridge. The salt bridge serves the vital role of allowing each half-cell's charge to be balanced while preventing the solutions from mixing with each other. Should the solutions be mixed, the reaction will be purely chemical and will not require electrons to flow through the wire — thus, no electricity can be harbored from the cell. Should the charges not be balanced, the anode will have an abundance of positive ions, and the cathode will have an abundance of negative ions. Electrons will not flow with this imbalance. The salt in the salt bridge, however, will ionize and neutralize the charges in each half-cell, allowing the reaction to proceed.
[edit] Notation
The galvanic cell, as the one shown in the figure, are conventionally described using the following notation:
Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s)
(anode)----------------------------------(cathode)
An alternate notation for this cell would be:
Zn(s) | Zn+2(aq) || Cu+2(aq) | Cu(s)
Where the following applies:
- (s) denotes solid.
- (aq) means aqueous solution.
- The vertical bar, |, denotes a phase boundary.
- The double vertical bar, ||, denotes a liquid junction for which the junction potential is near zero, such as a salt bridge.[4]
[edit] Corrosion
In this way the anode is consumed or corroded. When the anode material corrodes entirely away, the cell's potential drops and the current halts. The metal may be regarded as the fuel that powers the device. A similar process is used in electroplating. The ionic current in the electrolyte is equal to the current in the external circuit, so a complete circuit is formed with a path through the electrolyte.
As can be seen, electrons flow from the oxidized ion at the anode to the reduced atom, formerly an ion, at the cathode. The flow due to this redox reaction constitutes the current.
[edit] Galvanic corrosion
Galvanic corrosion is a process that degrades metals electrochemically. This corrosion occurs when two dissimilar metals are placed in contact with each other in the presence of an electrolyte, such as salt water, forming a galvanic cell. A cell can also be formed if the same metal is exposed to two different concentrations of electrolyte. The resulting electrochemical potential then develops an electric current that electrolytically dissolves the less noble material.
[edit] Electric potential
The potential of a cell can be determined by use of a standard potential table for the two half cells involved. An oxidation potential table could also be used, but the reduction table is more common. The calculation assumes that the cell operates at zero current flowing through the circuit.
The first step is to identify the two metals reacting in the cell. Then one looks up the Eo (standard electrode potential, in volts) for each of the two half reactions. The electric potential for the cell is equal to the more positive Eo value minus the more negative Eo value.
For example, in the picture above the solutions are CuSO4 and ZnSO4. Each solution has a corresponding metal strip in it, and a salt bridge or porous disk connecting the two solutions and allowing SO42− ions to flow freely between the copper and zinc solutions. In order to calculate the electric potential one looks up copper and zinc's half reactions and finds:
- Cu2+ + 2 e− → Cu (E = +0.34 V)
- Zn2+ + 2 e− → Zn (E = −0.76 V)
Thus the overall reaction is:
- Cu2+ + Zn → Cu + Zn2+
The electric potential is then +0.34 V - -0.76 V = 1.10 V under standard conditions and when no current flows in the cell.
If the cell is operated under non-standard conditions, the potentials must be adapted using the Nernst equation. If a current is allowed to flow in the circuit, the potential is going to shift towards zero in comparison with that predicted by the Nernst equation.
When measurements of the cell potential are performed, then the positive terminal of the electrometer needs to be connected to the right-hand half-cell. This is because the potential of the cell is defined as:
- ECell = ERight − ELeft
[edit] Cell types
[edit] See also
- Alessandro Volta
- Battery (electricity)
- Bio-nano generator
- Electrode potential
- Electrosynthesis
- Galvanic series
- Sacrificial anode
- Volt
- Voltaic pile
[edit] References
- ^ Merriam-Webster Online Dictionary: "battery"
- ^ "battery" (def. 4b), Merriam-Webster Online Dictionary (2008). Retrieved 6 August 2008.
- ^ Keithley, Joseph F. (1999). Daniell Cell. John Wiley and Sons. pp. 49–51. ISBN 0780311930.
- ^ Atkins, P., "Physical Chemistry", 6th edition, W.H. Freeman and Company, New York, 1997
[edit] External links
- Galvanic (Voltaic) Cells and Electrode Potential. Chemistry 115B, Sonoma.edu.
- Making and testing a simple galvanic cell. Woodrow Wilson Leadership Program in Chemistry, The Woodrow Wilson National Fellowship Foundation.
- Galvanic Cell An animation.
- Interactive animation of Galvanic Cell. Chemical Education Research Group, Iowa State University.
- Electrochemical Cells Tutorial Segment. Chemistry 30, Saskatchewan Evergreen Curriculum.
- Glossary for Galvanic Cells. Spark Notes, by Barnes & Noble. Sparknotes.com.
- Cathodic Protection 101. A basic tutorial on galvanic cells and corrosion prevention.
