- Parts of the galvanic cell
- Functioning
- Oxidation and reduction reactions
- The salt bridge
- Oxidation and reduction potentials
- Symbolic representation of a galvanic cell
- Applications
- The galvanic cell in daily life
- Construction of a homemade galvanic cell
- materials
- Process
- References
The galvanic cell or voltaic cell is a type of electrochemical cell that consists of two different metals immersed in two half cells, in which a compound in solution activates a spontaneous reaction.
Then one of the metals in one of the half cells is oxidized while the metal in the other half cell is reduced, producing an exchange of electrons through an external circuit. This makes it possible to take advantage of the electric current.
Figure 1. Scheme and parts of a galvanic cell. Source: corinto.pucp.edu.pe.
The name "galvanic cell" is in honor of one of the pioneers of experimentation with electricity: the Italian physician and physiologist Luigi Galvani (1737-1798).
Galvani discovered in 1780 that if cables of dissimilar metals were joined at one end and the free ends were brought into contact with the haunch of a (dead) frog, then a contraction occurred.
However, the first to build an electrochemical cell to produce electricity was the Italian Alessandro Volta (1745-1827) in 1800 and hence the alternative name of voltaic cell.
Parts of the galvanic cell
The parts of a galvanic cell are shown in figure 1 and are as follows:
1.- Anodic semicell
2.- Anodic electrode
3.- Anodic solution
4.- Cathode semicell
5.- Cathode electrode
6.- Cathodic solution
7.- Saline bridge
8.- Metallic conductor
9.- Voltmeter
Functioning
To explain the operation of a galvanic cell we will use the lower one:
Figure 2. Didactic model of galvanic cell. Source: slideserve.com
The fundamental idea of a galvanic cell is that the metal that undergoes the oxidation reaction is physically separated from the metal that is reduced, in such a way that the exchange of electrons occurs through an external conductor that allows to take advantage of the flow of electric current, for example to turn on a light bulb or led.
In figure 2, in the left half-cell there is a metallic copper (Cu) tape immersed in a copper sulfate solution (CuS0 4), while in the right half-cell there is a zinc (Zn) tape immersed in a solution of zinc sulfate (ZnSO 4).
It should be noted that in each half cell the metal of each one is present in two oxidation states: the neutral metal atoms and the metal ions of the salt of the same metal in solution.
If the metallic tapes are not joined by an outer conductive wire, then both metals are oxidized separately in their respective cells.
However, since they are electrically connected, it happens that oxidation will occur in Zn while there will be a reduction reaction in Cu. This is because the degree of oxidation of zinc is greater than that of copper.
The metal that is oxidized gives electrons to the metal that is reduced through the outer conductor and this flow of current can be harnessed.
Oxidation and reduction reactions
The reaction that occurs on the right side between the zinc metal electrode and the aqueous zinc sulfate solution is as follows:
Zn o (s) + Zn 2+ (SO 4) 2- → 2 Zn 2+ (ac) + (SO 4) 2- + 2 e -
A zinc atom (solid) on the surface of the anode electrode in the right half cell, stimulated by the positive ions of the zinc in solution, gives up two electrons and is detached from the electrode, passing into the aqueous solution as a double positive ion of zinc.
We realize that the net result was that a neutral zinc atom from the metal, through the loss of two electrons, became a zinc ion that adds to the aqueous solution, so that the zinc rod lost one atom and the solution acquired a positive double ion.
The released electrons will prefer to move through the outer wire towards the metal of the other positively charged half-cell (cathode +). The zinc bar is losing mass as its atoms gradually pass into the aqueous solution.
Zinc oxidation can be summarized as follows:
Zn o (s) → Zn 2+ (ac) + 2 e -
The reaction that occurs on the left side is similar, but the copper in the aqueous solution captures two electrons (coming from the other half-cell) and is deposited on the copper electrode. When an atom picks up electrons it is said to be reduced.
The copper reduction reaction is written like this:
Cu 2+ (ac) + 2 e - → Cu o (s)
The copper bar is gaining mass, as the ions of the solution pass to the bar.
Oxidation occurs at the anode (negative), which repels electrons, while reduction occurs at the cathode (positive), which attracts electrons. Electron exchange occurs through the outer conductor.
The salt bridge
The salt bridge balances the charges that accumulate in the two half cells. Positive ions accumulate in the anodic half cell, while in the cathodic cell an excess of negative sulfate ions remain.
For the salt bridge, a solution of a salt (such as sodium chloride or potassium chloride) that does not intervene in the reaction is used, which is in an inverted U-shaped tube with its ends plugged with a wall of porous material.
The sole purpose of the salt bridge is for the ions to filter into each cell, balancing or neutralizing the excess charge. In this way, a current flow is produced through the salt bridge, through the salt ions, which closes the electrical circuit.
Oxidation and reduction potentials
Standard oxidation and reduction potentials are understood to be those that occur at the anode and cathode at a temperature of 25ºC and with solutions of 1M concentration (one molar).
For zinc its standard oxidation potential is E ox = +0.76 V. While, the standard reduction potential for copper is E red = +0.34 V. The electromotive force (emf) produced by this galvanic cell is: emf = +0.76 V + 0.34 V = 1.1 V.
The global reaction of the galvanic cell can be written like this:
Zn o (s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu o (s)
Taking sulfate into account, the net reaction is:
Zn o (s) + Cu 2+ (SO 4) 2- 25ºC → Zn 2+ (SO 4) 2- + Cu o (s)
Sulfate is a bystander, while metals exchange electrons.
Symbolic representation of a galvanic cell
The galvanic cell in figure 2 is symbolically represented as follows:
Zn o (s) -Zn 2+ (aq) (1M) - Cu 2+ (aq) (1M) -Cu o (s)
By convention, the metal that oxidizes and forms the anode (-) is always placed on the left, and its ion in the aqueous state is separated by a bar (-). The anodic half cell is separated from the cathodic one by two bars (-) representing the salt bridge. On the right is placed the metal half cell that is reduced and forms the cathode (+).
In the symbolic representation of a galvanic cell, the left end is always the metal that is oxidized and the metal that is reduced is placed at the right end (in solid state). It should be noted that in Figure 2 the half cells are in reverse position with respect to the conventional symbolic representation.
Applications
Knowing the standard oxidation potentials of different metals, it is possible to determine the electromotive force that a galvanic cell built with these metals will produce.
In this section, we will apply what was stated in the previous sections to calculate the net electromotive force of a cell built with other metals.
As an example of application we consider a galvanic cell of iron (Fe) and copper (Cu). As data the following reduction reactions and their standard reduction potential are given, that is to say at 25ºC and 1M concentration:
Fe 2+ (ac) + 2 e - → Fe (s). E1 network = -0.44 V
Cu 2+ (ac) + 2 e - → Cu (s). E2 red = +0.34 V
It is asked to find the net electromotive force produced by the following galvanic cell:
Fe (s) -Fe 2+ (aq) (1M) - Cu 2+ (aq) -Cu (s)
In this cell the iron is oxidizing and is the anode of the galvanic cell, while the copper is reducing and is the cathode. The oxidation potential of iron is equal to but opposite to its reduction potential, that is, E1 oxd = +0.44.
To obtain the electromotive force produced by this galvanic cell, we add the oxidation potential of iron with the reduction potential of copper:
emf = E1 oxd + E2 red = -E1 red + E2 red = 0.44 V + 0.34 V = 0.78 V.
The galvanic cell in daily life
Galvanic cells for everyday use are very different in shape from what is used as a didactic model, but their principle of operation is the same.
The most commonly used cell is the 1.5V alkaline battery in its different presentations. The first name comes because it is a set of cells connected in series in order to increase the emf.
Lithium rechargeable batteries are also based on the same working principle as galvanic cells and are the ones used in smartphones, watches and other devices.
In the same way, lead batteries for automobiles, motorcycles and boats are 12V and are based on the same operating principle of the galvanic cell.
Galvanic cells are used in aesthetics and in muscle regeneration. There are facial treatments that consist of applying current through two electrodes in the shape of a roller or sphere that clean and tone the skin.
Current pulses are also applied to regenerate the muscles in people who are in a state of prostration.
Construction of a homemade galvanic cell
There are many ways to build a homemade galvanic cell. One of the simplest is using vinegar as a solution, steel nails, and copper wires.
materials
-Disposable plastic cups
-White vinegar
-Two steel screws
-Two pieces of bare copper wire (no insulation or varnish)
-A voltmeter
Process
-Fill ¾ parts of the glass with vinegar.
-Join the two steel screws with several turns of wire, leaving a piece of wire unwound.
The uncoiled end of the copper wire is bent into an inverted U shape so that it rests on the edge of the glass and the screws are submerged in the vinegar.
Figure 3. Homemade galvanic cell and multimeter. Source: youtube.com
Another piece of copper wire is also bent in an inverted U and is hung on the edge of the glass in a position diametrically opposite to the immersed screws, so that a portion of the copper is inside the vinegar and the other portion of the copper wire is outside. of the glass.
The free ends of the voltmeter leads are connected to measure the electromotive force produced by this simple cell. The emf of this type of cells is 0.5V. To equalize the emf of an alkaline battery, it is necessary to build two more cells and join the three in series, so that a 1.5V battery is obtained
References
- Borneo, R. Galvanic and electrolytic cells. Recovered from: classdequimica.blogspot.com
- Cedrón, J. General chemistry. PUCP. Recovered from: corinto.pucp.edu.pe
- Farrera, L. Introduction to electrochemistry. Department of Physicochemistry UNAM. Recovered from: depa.fquim.unam.mx.
- Wikipedia. Electrochemical cell. Recovered from: es.wikipedia.com.
- Wikipedia. Galvanic cell. Recovered from: es.wikipedia.com.