Why is corrosion an electrochemical process

Faraday's empirical laws of electrolysis relate the current of an electrochemical reaction to the quantity of moles of the element being reacted. Supposing that the charge needed for such a reaction was one electron per molecule, as is the case for the plating or the corrosion attack of silver, it can be shown as:. The charge carried by one mole of electrons is known as one Faraday F. When hydrogen H ions are reduced to their atomic type, they typically mix, as shown earlier, to provide hydrogen gas through a reaction with electrons at a cathodic surface.

In neutral water, the anodic corrosion of some metals, like alumunim Al , zinc Zn or magnesium Mg , creates enough energy to separate water directly, as illustrated within the following equation and figure:. The change in the concentration of H ions, or the increase in hydroxyl OH ions, may be shown by testing pH levels to find surfaces on which cathodic reactions are taking place.

There can be many cathodic reactions encountered throughout the corrosion process. They include the following:. Oxygen reduction is a common cathodic reaction because oxygen exists within the atmosphere and in solutions exposed to the environment.

Although not frequent, metal ion reduction and metal deposition will cause severe corrosion problems, for instance: the plating of copper ions, which are created upstream in a water circuit, on the inner aluminum surface of a radiator.

Therefore, the use of a copper conduit in a water-based circuit where aluminum is also present should generally be avoided. All corrosion reactions are merely combinations of one or many of the above cathodic reactions in conjunction with an anodic reaction.

Thus, each case of liquid corrosion may be reduced to those equations in most cases, either on an individual basis or in combination.

Take into account the corrosion of Zn zinc by water or wet air. By multiplying the Zn oxidization reaction by 2 and summing this with the oxygen reduction reaction, one obtains the following equation:. Likewise, the corrosion of Zn by copper sulfate represented within the following equation is simply the summation of the oxidization reaction for Zn and the metal deposition reaction involving copper II ions:.

During corrosion, more than one oxidation and one reduction reaction might take place. In the corrosion of Zn in a concentrated HCL solution containing dissolved oxygen , for example, two cathodic reactions are possible. One is an evolution of H, while the other is the reduction of oxygen. Because there are two cathodic reactions or methods that consume electrons, the general corrosion rate of zinc is overstated.

Thus, it is typically more corrosive than air-free acids, and removing oxygen from acid solutions can typically make these solutions less corrosive. This is a typical method for reducing corrosivity in many settings. Oxygen can be removed by either chemical or mechanical means. In corroding a piece of metal, the electrons created at anodic areas flow through the metal to react at cathodic areas that are equally exposed to the environment where they restore the electrical balance of the system.

The very fact that there is no net accumulation of charges on a corrosion surface is vital for understanding most corrosion processes and ways to mitigate them. However, the equality between the anodic and cathodic currents expressed within the following equation doesn't mean that the current densities for these currents are equal:. It has great implications for the occurrence of corrosion in dissimilar metals as well.

It is simple to know that the result of a particular quantity of anodic current focused on a small area of a metal surface will be much bigger than when the same quantity of current is dissipated over a much larger area. The cause of current density may be seen when two dissimilar metals are joined here it's Cu and Fe , which are shown diagrammatically in Figure 3 below.

Figure 3. The areas affected by dissimilar metals, where "a" shows rivets of steel on copper plates, while "b" shows rivets of copper on steel plates. When steel rivets are part of copper plates, the corrosion of the cathodic copper plates will be low, while the corrosion of the small anodal steel rivets will be high. On the other hand, if copper rivets are joining steel plates, corrosion on the copper will be high, while corrosion of the steel plates will hardly be noticeable.

Electrochemical reactions of corrosion are discussed in detail with the help of the Daniell cell, the anodic method, the cathodic method, Faraday's laws and surface area effects.

Because we can prevent corrosion if we thoroughly understand the electrochemistry of corrosion, manufacturers and users of corrosion products should be attentive to the electrochemical mechanisms behind corrosion. When this water encounters steel piping or a chrome-plated bathroom sink drain, the more-noble copper will plate out on the other metal, producing a new metals-in-contact corrosion cell.

Since both the cathodic and anodic steps must take place for corrosion to occur, prevention of either one will stop corrosion. The most obvious strategy is to stop both processes by coating the object with a paint or other protective coating.

Even if this is done, there are likely to be places where the coating is broken or does not penetrate, particularly if there are holes or screw threads. A more sophisticated approach is to apply a slight negative charge to the metal, thus making it more difficult for the reaction to take place:. One way of supplying this negative charge is to apply a coating of a more active metal. Thus a very common way of protecting steel from corrosion is to coat it with a thin layer of zinc; this process is known as galvanizing.

The zinc coating, being less noble than iron, tends to corrode selectively. Dissolution of this sacrificial coating leaves behind electrons which concentrate in the iron, making it cathodic and thus inhibiting its dissolution. The effect of plating iron with a less active metal provides an interesting contrast. The common tin-plated can on the right is a good example. As long as the tin coating remains intact, all is well, but exposure of even a tiny part of the underlying iron to the moist atmosphere initiates corrosion.

The electrons released from the iron flow into the tin, making the iron more anodic so now the tin is actively promoting corrosion of the iron!

You have probably observed how tin cans disintegrate very rapidly when left outdoors. A more sophisticated strategy is to maintain a continual negative electrical charge on a metal, so that its dissolution as positive ions is inhibited.

Since the entire surface is forced into the cathodic condition, this method is known as cathodic protection. The source of electrons can be an external direct current power supply commonly used to protect oil pipelines and other buried structures , or it can be the corrosion of another, more active metal such as a piece of zinc or aluminum buried in the ground nearby, as is shown in the illustration of the buried propane storage tank below.

Chem1 Virtual Textbook. Learning Objectives Make sure you thoroughly understand the following essential ideas. Electrochemical corrosion of metals occurs when electrons from atoms at the surface of the metal are transferred to a suitable electron acceptor or depolarizer. The electrons can move through the iron to the edge of the droplet, where they can reduce the atmospheric oxygen:. When they meet, they form a precipitate of iron II hydroxide.

Is corrosion an electrochemical process? Ernest Z. Sep 5, It doesn't have to be an electrochemical process, but it occurs much faster when it is. A piece of iron rusts quickly if you leave it outside in the rain.