Chemical reactions produce a wide variety of things, such as new chemicals, heat, energy, and so forth. Some reactions also produce electricity, which at first glance appears to have a lot of potential in the environment. A reaction can produce electricity if the reaction can be made to occur in a galvanic cell, a vessel for a reaction that forces the electrons that are transferred between the reactants to do work and produce an electrical current. This article focuses on the different pieces of galvanic cells and what they do.
Most galvanic cells will consist of two reactions, one that is an oxidation reaction and one that is reduction reaction. This is because an oxidation reaction produces free electrons, and a reduction reaction requires free electrons, so the electrons can transfer between the two reactions and do electrical work and produce electricity. Both oxidation and reduction reactions are called half-reactions because free electrons are present either before or after the reaction; they have not been completely transferred from one chemical to another. Typically each half-reaction occurs in its own container, which are often made of class.
Example 1: Galvanic cells can be comprised of a huge variety of different half-reactions. For instance, the two half-reactions for a galvanic cell can be
$$Fe \rightarrow Fe^{3+} + 3e^{-}$$ $$Mn^{4+} + 4e^{-} \rightarrow Mn$$
Which chemical is oxidized and which is reduced can also be interchanged to create a different cell:
$$Fe^{3+} + 3e^{-} \rightarrow Fe$$ $$Mn \rightarrow Mn^{4+} + 4e^{-}$$
Sometimes the chemicals used in redox reactions will be polyatomic. Example 2 shows such a galvanic cell.
Example 2: Another possible pair of half-reactions for a galvanic cell could be
$$MnO_4^{-} + 8H^{+} + 3e^{-} \rightarrow Mn^{7+} + 4H_2O$$ $$Cl_2 + 3H_2O \rightarrow ClO_3^{-} + 6H^{+} + 5e^{-}$$
Most half-reactions that include polyatomic molecules/ions will also include hydrogen ions and water molecules, which come from balancing the reactions.
There are certain locations in the galvanic cell designated for the half-reactions. The reduction reaction occurs in the cathode. The oxidation reaction occurs in the anode. For a pair of chemicals, one can be selected as the cathode and the other will be the anode. Both of the chemicals that are included in the chemical equations on the sides of those equations without free electrons are called electrodes and have a wire connecting them to transmit electrons and let them use their electricity to power some device.
Example 3: Perhaps you wish to make a galvanic cell out of iron and lithium. If you want iron to be in the cathode, it will be reduced, and you could have the half-reaction
$$Fe^{3+} + 3e^{-} \rightarrow Fe$$
This means the anode contains lithium and thus the lithium gets oxidized:
$$Li \rightarrow Li^{+} + e^{-}$$
Alternatively, you could let the anode contain iron and the cathode contain lithium. In this case the reactions are
$$Li^{+} + e^{-} \rightarrow Li$$ $$Fe \rightarrow Fe^{3+} + 3e^{-}$$
However, for every cathode-anode pair for the two half-reactions, only one will generate electricity. The other half-reactions, which are collectively called nonspontaneous, require electricity to occur, and these are not galvanic cells, but electrolytic cells.
A salt bridge is a key part of a galvanic cell, because it is a source of cations and ions for the cell. Without a salt bridge, the reactants for the half-reactions would run out very quickly, and the electricity produced by the cell would be short-lived. However, by providing another source of ions, extra electrons can be transferred to the products to recreate the reactants.
Example 4a: Assume that the cathode half-reaction for a galvanic cell is
$$Mg^{2+} + 2e^{-} \rightarrow Mg$$
If magnesium can lose some of its electrons to another substance, it gets oxidized and then can be reduced again. Depending on the chemical chosen for the salt bridge, this may be possible. One such example is sodium perchlorate, where the sodium ion gives the magnesium atoms a chance to lose electrons. The following reaction occurs:
$$Mg + 2Na^{+} \rightarrow Mg^{2+} + 2Na$$
The perchlorate ions exit the salt bridge on the anode side of the galvanic cell.
Example 4b: In the galvanic cell that was partially described in Example 4a, assume that the anode reaction was
$$W \rightarrow W^{6+} + 6e^{-}$$
Without the salt bridge, the tungsten atoms would be exhausted very quickly. The perchlorate ions provided by the salt bridge let the tungsten ions reduce so they can be oxidized again. The following reaction occurs:
$$7W + 6ClO_4^{-} + 48H^{+} \rightarrow 7W^{6+} + 3Cl_2 + 24H_2O$$
The salt bridge is typically filled with an ionic compound, whose ions split apart. The cations divert to the cathode, and the anions divert to the anode (note how the two pairs of terms have the same prefixes). However, if we were to write the overall reaction for the galvanic cell, we would omit the half-reactions involving ions from the salt bridge.
Obviously, the concepts of redox reactions, half-reactions, cathodes and anodes, and the salt bridge all play crucial roles in the construction of galvanic cells. There are more technical aspects needed to design effective galvanic cells, namely determining how to choose which chemical represents the cathode and which represents the anode, but these concepts put you well in the right direction for your study of electrochemistry.