2.3 Corrosion processes
For many materials, degradation processes are simply one or a series of chemical reactions that act to erode or deteriorate the material. The deterioration of metals is a little more complex than that of non-metals because metals are electrical conductors. Local electrochemical cells frequently form in the exposed surfaces of metals, leading to corrosion of the metal in one part of the cell. Electron movement is an essential part of the process: as electrons are lost, metal ions are formed, and these soluble metal ions then pass into the aqueous environment, resulting in a net loss of metal. Electrochemical cells were actually used by Volta to produce electricity (in the first batteries), so you can see that corrosion can be turned to advantage to make portable power sources.
Box 1: Electrochemical cells
The most familiar type of electrochemical cell to most people is the common battery. Batteries harness the energy released by corrosion of metal components, which is why they are usually heavy. An electrochemical cell contains two electrodes made from differing metals. When the terminals are connected, each electrode reacts with a current-carrying solution known as an electrolyte and the cell provides a current.
The two electrodes in any cell are known as the cathode and the anode. At the anode, the metal reacts and releases electrons. These electrons then flow through the connection to the cathode. We can write the anodic reaction in chemical shorthand as:
M → M+ + e−
where M is a metal, and M+ is a positively charged ion formed when the metal atom loses an electron e−. Different metals will lose different numbers of electrons in a corrosion process. The resulting metal ion may be lost into a solution, or form part of a corrosion product such as rust.
At the cathode, the reverse reaction occurs. Electrons are ‘absorbed’ by ions, causing a different reaction, which might be the plating out of a metal from solution.
‘Primary cells’ are non-rechargeable: an example is the zinc-carbon battery, with a zinc anode and a carbon cathode that are separated by an electrolyte gel containing a salt (ammonium chloride). The electrolyte effectively attacks the zinc to produce electrons that can be tapped off at will. A Daniell cell uses copper and zinc as the electrodes, the copper being the cathode and the zinc the anode (Figure 13). Here the ‘salt bridge’ allows ions to travel between the two solutions, thus completing the circuit.
The life of every cell is limited by the amount of anode present, because this is attacked and effectively disappears, corroding away until little metal is left. So here is a possible tool to assess corrosive activity. Some metals are clearly more reactive than others; in other words, they have a greater electrical potential. It is possible to create a table that allows the electrical potentials of different metals to be compared. This may be done by putting two metals into a cell in order to determine which will corrode in preference to the other. Each specific cell has a characteristic electromotive force (emf), also known as the electrical potential difference, which is measured in volts. This shows how much more reactive one metal is than the other.
In order for all metals to be comparable, they must be measured against a standard point. Thus a ‘hydrogen electrode’ provides an arbitrary zero against which the other corrosion reactions are measured, to produce a list of standard electrode potentials (E0) for different metals. This is shown in Table 1 and is known as the electrochemical series. The least reactive metals are at the top of the list; the most reactive are at the bottom. So, the more positive the standard electrode potential, the less likely a material is to corrode; the more negative the value, the more likely the material is to corrode. When two dissimilar metals are in contact, it will always be the metal with lower potential that corrodes.