The reaction and subsequent deterioration of metals when exposed to the environment encompasses a fundamental principle of electrochemistry and metallurgy, known as corrosion. The corrosion reaction produces a new and less desirable material from the original metal and can result in the loss of function of the component or system, a significant problem encountered everyday.

Chemical engineers are continually faced with the problem of selecting proper materials to be used for the construction of processing vessels, storage tanks, overhead receivers, valves, piping systems and whatever else that comes into contact with potentially corrosive chemicals and environments. The problem of corrosion as well as material selection for corrosion control in not limited to the chemical engineer. Many industries employ the transportation and use of potentially corrosive chemicals and/or in potentially corrosive environments. Of all industrial problems facing engineers, very few can be more economically important than the prevention of metallic corrosion and mechanical failure. Each year, corrosion has cost more than 6 billion dollars in repair and replacement.

Metals are extremely susceptible to saltwater environments. Methods of corrosion control and prevention via direct application of these methods on the various forms of corrosion are analyzed in this study.

The fundamental of corrosion lies in the flow of electricity between certain areas of a metal surface through a solution (i.e. environment) capable of conducting an electric current.

This electrochemical action causes destructive alteration or eating away of a metal at areas which are called anodes, where the electric current leaves the metal and enters the environment to which it is exposed. This is the critical step in the series associated with corrosion. This is better explained by how the metallurgical industry refines metals in their pure form, which are inherently unstable when in contact with the natural environment. Specifically, the bonds between metal ions contain an abundant amount of free potential energy, which has a tendency to be released through the process of corrosion (i.e. oxidation) converting the metal to its natural state. This difference in the binding energies between the metal atoms is what provides the driving potential towards the oxidation of metal atoms (M), resulting in the loss of one or more electrons and the production of the ionic form of the metal shown in equation (1).

M --> Mn+ + n e- ... (1)

This reaction is known as the anodic reaction in corrosion, occurring on the corroding metal’s surface (the anode). Specifically, the positively charged atoms of metal detach themselves from the solid surface of the metal and enter the solution/environment as ions while the negatively charged electrons remain behind in the metal. The negatively charged electrons travel through the bulk metal to a “location” where they undergo a secondary reaction known as the cathodic reaction, where either a non-metallic atom or ion (Nm or Nm+, respectively) or another metallic ion (P+) is reduced, resulting in the accumulation of these electrons. The process of which is shown in equations 2(a), 2(b) and/or 2(c).

Nm + n e- --->  Nmn- ...(2a)

Nmn+ + n e---->   Nm ...(2b)

P n+ + n e- --->  P ...(2c)

Essentially, both anodic and cathodic reactions behave in a manner as to balance each other out, resulting in a neutral species, similar to what is seen in an electrochemical cell or battery. Both anodic and cathodic reactions occur simultaneously at the same rates. Furthermore, the site or location of these electrodes including the site of each reaction may consist of either two different kinds of metals, or they may be on different areas of the same piece of metal. In either case, there must be a potential difference between the two electrodes, or areas, so that the release of electrons as well as formation of metal ions via oxidation of the metal at the anode and simultaneous acceptance of the released electrons such as by neutralization of positive ions or formation of negative ions at the cathode can take place (in simpler terms, the corrosion process).

The remaining fundamental requirement of ion exchange between anodes and cathodes is the presence of an electrolyte, which is essential in the corrosion process. An electrolyte is a solution which dissociates into ions in aqueous solution. It is essentially the medium which provides the ion transport mechanism between positive (cathode) and negative (anode) electrodes in the corrosion process. Water for example contains both positively charged hydrogen ions and negatively charged hydroxyl ions in equal concentration. A single drop of water can contain up to 3 million hydrogen and hydroxyl ions. Therefore, corrosive
environments can be in any form of moisture that contains an abundance of ions. Other examples include acids, bases, oils, and other solid and liquid chemicals. In particular, the electrolyte solution used in this laboratory experiment is a saltwater solution (NaCl + distilled H2O) simulating a “sea-side” corrosive vapour environment. The importance of the electrolyte solution lies within the cathodic reaction, taking place at the cathode. For example, with an acidic electrolyte solution, such as HCl (in aqueous form appears as H+ and Cl- ions), positively charged hydrogen ions within the electrolyte are responsible for neutralizing the negatively charged electrons traveling from the anode to the cathode (cathodic reaction). As a result, in losing their charge, neutral hydrogen atoms combine to form hydrogen gas, according to equation (3)

2 H+ + 2 e- --->  H2 …(3)

The hydrogen gas is then released into the atmosphere as bubbles. solution.

The saltwater solution is a neutral electrolyte, containing equal concentration of both positive and negative ions. As a result, the evolution of hydrogen ions from the metal surface is in actuality very slow. As seen below in figure 1, hydrogen ions tend to build up as a protective layer on the surface of the cathode, preventing or slowing down further corrosion, an effect known as cathodic polarization.

Saltwater environments are freely exposed to atmospheric oxygen, subsequent reactions involving the hydrogen ions on the surface of the cathode being reduced with the electrons and oxygen lead to the formation of water, as shown in equation (4).

4 H+ + O2 + 4 e- ---> 2 H2O (4)

Oxygen essentially acts as a “depolarizer”, washing away the hydrogen ion coating on the metal, allowing for the further degradation of the metal via corrosion. As such, the presence of oxygen plays an important role in the corrosion of metals. Finally, subsequent reactions of the products of both the anode and cathode form the visible products of corrosion, such as rust.

The following equations (5), (6), (7), and (8) illustrate the corrosion of iron resulting in rust formation.

At the anode, oxidation of iron occurs:
Fe(s) --->  Fe2+ + 2e-   ...(5)

At the cathode, reduction of atmospheric oxygen with water occurs:

½ O2 (g) + H2O + 2e-   --->2 OH- ...(6)

Anodic and Cathodic reaction products combine:
Fe2+ + 2 OH-  ---> Fe(OH)2  ...(7)

Subsequent oxidation reaction results in the formation of rust:

Fe(ΟΗ)2 + Ο2(g) + Η2Ο---> Fe2Ο3·2Η2Ο ... (8)

Corrosion Capabilities of Various Metals
The luster, ability to conduct heat and electricity, malleability of some and ductility of others are only some of the main characteristics which define metals. However, it is the varying ability of a metal to loose its electrons and form a positive ion that is essential in understanding the ranges of metals that are capable of corroding. It is this understanding which is used widely in industry as the number one measure for corrosion control (i.e. choosing a metal that will not corrode easily by giving up its electrons). Corrosion begins with and its extent is measured by the ability of the respective metal atom to oxidize (i.e. give up its electrons). The electrochemical field has developed a list, which details the range of most metals relative tendency to be oxidized, called an electromotive series. The list provides nearly all the information an experimenter requires to determine which metals are most susceptible to corroding. The series tells us which species behaves best as the anode undergoing oxidation, and the cathode undergoing reduction. The list begins with metals most easily oxidized, and ends with the metal least capable of oxidation, or with the greatest capability for reduction. In other words, the list begins with metals which are most easily corroded (i.e. behave as the anode in the corrosion reaction), and ends with metals which can best behave as the cathode in the corrosion reaction. The tendency of each metal to be oxidized is given in terms of the standard electrode potential or oxidizing potential and is expressed in volts, which is measured relative the standard oxidation of hydrogen gas, which is assigned an arbitrary potential of 0 volts. Table 1 provides an electromotive series list of major industrial metals.

Table 1: Electromotive force series.

From the list, metals such as potassium, magnesium, zinc and iron will oxidize easier and therefore corrode easier than platinum or gold. This list should provide you with the sufficient background needed to predict which metals during the experiment should corrode more easily and/or to a greater extent. Furthermore, as will be discussed in the sections following, the electromotive series has a large influence on a particular type of corrosion, known as galvanic corrosion.

The electromotive series serves as a corrosion map for metals in particular concentrations of their own salts. It fails in describing the behaviour, and as a result the performance, of such metals in various corrosive environments based on varying temperature, humidity, salinity, etc. Therefore, a more general table developed by electrochemists, called the galvanic series was developed in order to compensated for this realization. Table 2 provides a short galvanic series list of major industrial metals in seawater
corrosive environments.

Table 2: Galvanic Series of metals and alloys in seawater.

Just as in an electromotive series, the metals near the top of the list on the galvanic series tend to oxidize easiest (i.e. behave as the anode, and therefore undergo corrosion), while the metals near the bottom reduce easiest (i.e. behave as the cathode).

Measures to Prevent and Control Corrosion
There has been extensive and exhaustive research done on finding newer and better ways to prevent and/or control corrosion. As mentioned before, it is essential to understand the fundamental mechanism of corrosion and how it occurs, in order to control it, and with luck, prevent it. We can find the fundamentals of corrosion lies in three areas: anode, cathode, and electrolyte. These represent the “triangle” of requirement (figure 2) for corrosion. Corrosion cannot occur without the presence of all three components. In other words, one cannot survive without the other.

It is therefore natural that the preventive and controlling measures of corrosion come from breaking the linkage within this triangle. Many methods have been created to prevent corrosion, four of which are listed below and will be attempted by the student over the course of the experiment.

• Slowing down or stopping the movement of electrons by:
  
  • a. Coating the surface of the metal with a non-conductive organic coating. Essentially, organic coatings like paints and primers serve to prevent electrical contact of the metal with other pieces of metal (a requirement in galvanic corrosion), prevent contact with electrolytic environments, and create a uniform surface so as to prevent localized corrosion.
  • b. Keep the metal dry and away from corrosive moist environments. Of the corrosive environments discussed, they all share a common characteristic: abundant sources of ions in moist or liquid environments. Placing corrodible metals in dry environments can prevent most forms of corrosion. However, due to the weather conditions (especially in Canada) this is not always possible.
  • c. Apply a reverse current to the metal from an outside source, which basically reduces the corrosion rate of a metallic structure by reducing its electric potential, preventing the loss of electrons. This process is known as cathodic protection. Cathodic protection along with organic coatings has become the most widely used methods for corrosion control of metallic structures in contact with electrolytically conducting environments (i.e. environments containing enough ions to conduct electricity such as soils, seawater and natural water) [5]. Cathodic protection utilizes either a sacrificial anode or outside source of current, of which the details will not be discussed here.
• Slowing down or stopping the process of oxygen reaching the surface of the metal. The presence of oxygen is very important in the corrosion process (eqn. 4). The most common way of preventing oxygen from reaching the surface of the metal is the use of organic coatings as discussed above.
• Choosing a metal that is less willing to oxidize (i.e. give up its electrons). According to the electromotive series and shown in table 1, choosing a metal closer or above the standard electric potential for the oxidation of hydrogen is preferred. However, as you can see from the series, this can pose a significant cost problem, as noble metals such as gold, mercury, platinum and silver are very expensive.
• Choosing a metal that forms an oxide during the corrosion reaction, this thereby prevents further corrosion. The oxidation resistance of some metals, most commonly alloys, has been improved by adding alloying elements such as chromium, aluminum and silicon, which react with oxygen to form a stable, protective external oxide film on the surface of the metal when exposed to their service environments.