This section is a short explanation of the processes that lead to electrochemical potential and thus to corrosion processes.
Most readers will have a rough idea of electrochemical potentials, noble and non-noble metals. However, a small recap of this knowledge seems suitable to understand the processes happening at the electrodes or corroding metals better.
If you immerse two blank metal sheets made from different materials, e.g. iron and copper, into the same conducting solution and connect them by a device that can measure the potential difference or voltage between these two metal sheets, the device will show a characteristic voltage difference between the two metals. They have different potentials. Why is that so?
When a metal gets in contact with a solution some metal ions will be emitted by the metal. When the positive metal ions, cations, leave the metals solid phase, they leave their electrons behind. The electrons charge the metal negative. This means that the metal ions are attracted again by the metal sheet. This results in a flux of metal ions from the metal and back to the metal. A dynamic equilibrium is reached and the amount of charge in the metal is characteristic for each metal.
As we know the potential of a metal sheet can be actively influenced by a potentiostat. If the potential is changed to a more negative (cathodic) potential, the metal ions will be pulled towards the electrode. If we change the potential to a more positive (anodic) potential, the metal will release more ions and dissolve.
Anode and cathode as well as anodic and cathodic potentials are for electrochemistry beginners not always very intuitive terms. The anode is where the oxidation happens and the cathode where the reduction happens. Anodic potential are potential with more oxidizing property and cathodic potentials are potentials with more reducing properties. A good way to remember this is: Oxidation and anode start with a vowel. Cathode and Reduction start with a consonant.
If any change of the potential will lead to a change of the direction of the reaction, i.e. from reduction to oxidation or vice versa, the system is at its formal potential. If the formal potential is measured under standard conditions (activity of all components is 1, 298,15 K, 1 bar) it is called standard potential. With the standard potential and the Nernst equation (see. equation 3.4 and 3.6) the resulting potential can be calculated.
E0 is the standard potential of the redox couple Red and Ox. R is the gas constant and T the temperature. The activity of the oxidized and reduced form of the species aOx and aRed. F is the Faraday constant and z the number of electrons transferred per molecule / atom.
The two activity coefficients fOx and fRed are included in the resulting potential E0’, which is called the formal potential. Since it contains parameters that depend on the environment, such as temperature and activity coefficients, E0’ cannot be listed but needs to be determined for each experiment, if necessary. Most experiments in analytical chemistry are performed at room temperature (295 K). This makes another simplification possible. Out of convenience also the ln will be transferred to the log.
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