Electrochemical cells, also known as galvanic cells, are devices enabling the direct conversion of the energy of chemical bonds into electrical work. They consist of two electrodes, which are metallic conductors. They remain in constant contact with an ionic conductor – a liquid or solid electrolyte. A single electrode with the surrounding electrolyte constitutes a half-cell. Depending on the analytical method used, the electrodes may have a common electrolyte or be immersed in different electrolytes.
Such half-cells are then connected using an electrolytic key. It is used to allow the flow of electrons and thus to maintain electrical contact between the electrodes. Schematically, the construction of a galvanic cell can be described as follows:
anode | anode electrolyte || cathode electrolyte | cathode
In such notations, vertical lines indicate phase boundaries and double lines indicate the electrolytic key. Also, one must pay attention to the order in which the reactants are noted, always starting with the reduction reaction from the left, then followed by the oxidation reaction.
Energy in the cell
In galvanic cells, energy is generated as a result of spontaneous chemical reactions. A device with a similar application, but in which the reaction is forced by applying an external DC source, is the electrolyser. As the name suggests, it conducts the processes of electrolysis. All available batteries are galvanic cells. These are dry cells, mercury cells, nickel-cadmium batteries, which are used to power electrical devices. Spontaneous reactions taking place inside them occur due to the introduction of appropriate substances in the production process.
The reactions at the electrodes
During the operation of the cell, oxidation and reduction processes take place at individual electrodes. The electrons released during the oxidation, present in a single half-cell, flow towards the other half-cell, where they cause a reduction reaction. The electrode at which the reduction takes place is called the cathode, while the anode is the electrode at which the oxidation takes place. Visually, the anode always has a minus sign, and electrons from anode flow to the cathode with a positive sign. Since the positive charge corresponds to a higher potential value, the cathode shows a higher potential than the anode.
Half-cells
A half-cell can be composed of at least two phases. One of them, the electrode, conducts electrons. The second is responsible for ionic conductivity and is present in the form of an electrolyte in a solution or in a molten state. At the boundary of these phases, there is a specific arrangement of electrons, ions and dipoles determined by electrostatic interactions, sometimes also combined with the adsorption of ions and dipole molecules.
Type I half-cells
Type I half-cells include all the most common half-cells, which are formed as a result of introducing a metallic electrode into a salt solution that contains cations of the same metal. Examples of such systems are: the zinc half-cell Zn2+|Zn and the copper half-cell Cu2+|Cu. This type of half-cells is also known as cation-reversible, because the cation-mediated reaction equilibrates at its electrode surface. Gas half-cells belong to type I half-cells. In such systems, the gas is in equilibrium with its ions in the presence of a metal that is chemically inert. Its role is to transfer electrons without being a reactant in the reaction. However, it can be its catalyst. For this purpose platinum is often used. The most important example of a gas half-cell is the hydrogen half-cell. A stream of gaseous hydrogen passes through an aqueous solution containing H+ ions. The symbolic notation of the half-cell is as follows:
Pt | H2(g) | H+(c)
This is an important half-cell in a research context because its standard potential is assumed to be equal to 0 V. This is due to the activity of hydrogen and hydrogen ions equal to one. Thus, the hydrogen electrode is used as the standard reference electrode. The potentials of other half-cells are determined in relation to the potential of the hydrogen electrode. It is also a cation reversible electrode. In contrast, other gaseous electrodes can establish an equilibrium with the anion. Hence their name – anion reversible electrodes. Such half-cells include e.g.:
Cl2(g)|Cl–(c)
Type II half-cells
The next type of half-cells has a structure composed of metal, which is covered with a porous layer of a sparingly soluble salt of this metal. Such a system is immersed in a solution of a highly soluble salt having the same anion as the sparingly soluble salt. This scheme is noted as:
metal | sparingly soluble salt | common anion, e.g.:
Ag | AgCl | Cl–
These are common anion reversible electrodes and their potential depends on the activity of these ions, in this case chloride. Due to the fact that type II electrodes are characterized by reversibility, durability and constant potential, they are often used as reference electrodes when measuring the potentials of other half-cells. Two of them are most commonly used for this purpose – the already mentioned silver chloride electrode and the calomel electrode made of mercury covered with calomel paste with an admixture of mercury immersed in a solution containing chloride anions:
Hg | Hg2Cl2 | Cl–
Redox half-cells
Despite the somewhat misleading name, since all half-cells are characterized by redox reactions, this group is reserved for half-cells in which a chemically inactive metal (Pt, Au) is immersed in a solution containing a substance in both oxidized and reduced forms. An example is the quinhydrone half-cell, made of a platinum electrode immersed in a quinhydrone solution. Such a solution contains the same number of moles of quinone and hydroquinone.
Cell types
The simplest cells consist of half-cells having the same electrolyte. However, there are also those in which individual half-cells contain different solutions. An example of such a cell is the Daniell cell, which scheme can be noted as follows:
Zn | Zn2+ || Cu2+ | Cu
The anode is made of a zinc electrode immersed in an aqueous solution of zinc sulphate, while the anode is a copper electrode immersed in an aqueous solution of copper sulphate. Both half-cells are connected with an electrolytic key and they are not in direct contact with each other.
Cells can be divided into chemical and concentration cells. In chemical cells, the spontaneous process is an oxidation-reduction reaction, in which the energy of a chemical reaction is converted into electrical energy. Concentration cells are characterized by the use of the same electrodes and electrolytes of different concentrations. After such half-cells are short-circuited, a spontaneous process occurs to equalize the concentrations. The process is the source of electrical work. There are also electrode concentration cells where gaseous electrodes differ in concentration from each other, e.g. gaseous electrodes that differ in gas pressure. These can also be amalgam electrodes with different amalgam concentrations.