Introduction and Basics of Corrosion

#### **Chapter 1**

## Introductory Chapter: Corrosion

*Shumaila Masood, Anujit Ghosal, Eram Sharmin, Fahmina Zafar and Nahid Nishat*

#### **1. Introduction**

Corrosion can also be termed "metallic cancer" because it is an irreversible process that leads to the loss of billions of dollars. It involves the deterioration and consequently loss of solid metallic structure due to the chemical (dry gases, moisture, liquid, ionic solutions, microbes, etc.) or electrochemical (micro-cell formation like Daniel cell) reactions resulting from the potential difference in the structure and presence of a suitable electrolyte (salt-water) [1]. The process of corrosion is a surface phenomenon whose rate depends upon various factors such as temperature, the presence of harmful chemicals in the environment, ionic species, as well as humidity. Once the surface is breached the process of deterioration continues within, thereby making the structure vulnerable to stress and load. The combination with acid in the environment generated via., possible chemical reactions speed up the process resulting in the reduction of the life span of metal structure [2]. Corrosion can be sub-divided into many forms such as pitting, galvanic, intergranular, stress, and others. All these can affect the structural integrity by weakening the metallic units within any construction by corroding rods/wires, water pipelines (leakages), metallic bases, electrical units, etc. Few of these examples are shown in **Figure 1**.

These corrosive reactions change the microstructure of metal which results in the loss of its elasticity, mechanical and tensile strength converting them into flaky and brittle units. The losses incurred due to corrosion directly effects the gross domestic production (GDP) of all the countries. In 2018, the corrosion protection industry was estimated to be around 2.5 trillion (USD) and is expected to cross 3 trillion (USD) by the end of this year [3]. Proper use of corrosion prevention measures can help to refrain from such severe damages and can partially avoid the loss in economy, environmental pollution, direct and other indirect losses associated with corrosion [4].

The most popular remedies involved in corrosion protection are the use of cathodic protection, anodic protection, corrosion inhibitors, and protective coatings. Among them, utilization of corrosion inhibitors (chemical barrier) and protective coatings (physical barrier) is mostly focused on by a larger population around the world [5]. Coatings provide a physical barrier between the harmful environment and the metal surface which enhances the lifespan of metallic structures. A number of coatings that have been developed for corrosion protection include metallic coatings, chemical conversion coatings, organic and inorganic coatings, nanocomposite coatings [6]. However, such coatings have limited applications in complex design and high-rise infrastructures. Thereby coating with rapid curing ability, strong adhesive abilities, and stability against the harsh environment (exposure to atmospheric moisture, UV radiations, and other pollutants) are being researched. On the other hand, corrosion inhibitors are chemical additives that are added in low concentration to the

**Figure 1.** *Effect of corrosion on different metallic structures.*

destructive media, and timely they tend to suppress the corrosion progress [7]. These chemicals are classified as cathodic, anodic, or mixed inhibitors depending on the inhibition of the type of corrosion. The use of renewable resources in the field of corrosion prevention in the form of coatings and inhibitors is gaining significant importance. The chapter describes the types of corrosion, mechanism, and its control through synthetic/renewable resources based on synthetic coatings, micro to nano-coating, and corrosion inhibitors.

#### **2. Corrosion**

Corrosion of steel is a spontaneous electrochemical process that takes place in the presence of a solution containing dissolved oxygen. This process includes the delocalization of metal ions through oxidation into the solution at the anode (active area) and mobilization of electrons through reduction of metal to an acceptor such as oxidizing agents, or oxygen or hydrogen ions at cathode (less active area) [8].

The mechanism of corrosion involves the release of electrons from the metallic surface into the electrolyte in the presence of oxygen (**Figure 2**). This process occurs due to the tendency of metals to return to their natural oxidation state. The reactions occurring at both the electrodes can be expressed as follows:

#### **At anode:**

$$\text{Fe} \xrightarrow{\text{---}} \text{Fe}^{+2} + 2\text{e}^- \tag{1}$$

$$\text{Fe}^{+2} \xrightarrow{\text{---}} \text{Fe}^{+3} + \text{e}^- \tag{2}$$

**Figure 2.** *Mechanism of corrosion.*

**At cathode:**

$$\frac{1}{2}\text{O}\_2 + \text{H}\_2\text{O} + 2\text{e}^- \xrightarrow{\text{---}} 2\text{OH}^- \tag{3}$$

The deterioration of metallic surfaces depend upon the different forms of corrosion, which is caused due to different type of corrosive environment (**Figure 3**). Dry corrosion, also known as chemical corrosion, is the type of corrosion that occurs in the absence of moisture or water and metal oxidizes only due to the atmosphere. Therefore, this may also be referred to as, an oxidation process sustained by atmospheric oxygen without a liquid solution. At encompassing temperatures, most metals have slow oxidation rates [9]. Galvanic corrosion (GC), also known as bimetallic corrosion, is an electrochemical process where two different

**Figure 3.** *Different types of corrosion.*

metallic materials are connected electrically in a corrosive environment. In this type of corrosion, one metal (the anode) corrodes preferentially while the other metal (the cathode) remains protected [10]. Pitting corrosion (PC) is the most destructive type of corrosion in which the attack of corrosive ions is localized that results in the formation of pits. This type of corrosion results in the failure of machines without much of a weight loss. The process of pitting takes a longer period of time to initiate, however, once the pitting is initiated it penetrates into the section at an accelerated rate [11].

Intergranular corrosion (IC) is a localized form of corrosive attack that is preferentially along the grain boundaries or areas adjacent to them. The corrosion activity occurs at the grain boundary area since it is electrochemically different from the bulk. This process is mainly observed in stainless steel [12]. Waterline corrosion (WC) is an oxidation process that occurs when one part of the metal is submerged in water and another part is in contact with air. Water tanks are often prone to this type of corrosion [13]. Stress corrosion (SC) occurs when tensile stress and corrosive environment work together, often at elevated temperatures. In this type of corrosion stressed area of metal is anodic in respect to the unstressed area of the metal. This corrosion is not visible prior to fracture; therefore, it results in catastrophic failure [14]. Microbiologically induced corrosion (MIC) is a type of corrosion in which metal deteriorates through the metabolic activity of microorganisms. The common bacteria that cause MIC are acid-producing bacteria, sulfateproducing bacteria, and iron-reducing bacteria [15].

#### **2.1 Corrosion protective coatings: combating mechanism**

Corrosion cannot be completely eradicated but overcame by using protective coatings and corrosion inhibitors. Protective coatings are developed to retard the corrosion rate to protect metallic substrates [16]. These coatings are applied using several techniques such as roller, moving belt, or brush technique. They are functionalized with the help of various organic, inorganic, hybrid, or metallic layers to enhance their performance [17]. Highly protective coatings provide an impenetrable barrier protecting the substrate from the aggressive environment (**Figure 4**).

Modified polymeric coatings provide enhanced corrosion protection as compared to the simple polymeric coatings [18]. Various types of coating utilized for protection purposes are shown in **Figure 5**. Metallic coating means protecting the metals with the help of metal-coating. These coatings are applied to the substrates for several reasons but among all corrosion protection is major [19]. Various methods are used to apply these coatings on a substrate such as metalizing, electroplating, vapor deposition, hot dipping, cladding, etc. [20]. Depending upon the metal used to coat the substrate, metallic coatings can be divided into two categories: Anodic and cathodic coatings. In anodic coatings, anodes are made up of the alloys that are electrochemically more active than the base metal as a result of which the anodic metal depletes at a faster rate. These anodic coatings act as a physical obstruction between the corrosive environment and the base metal thereby protecting the base metal. These anodic metals can also be called "*Sacrificial anodes*". Zinc, magnesium, and cadmium are well non-sacrificial anodic metals that provide protection to steel [21]. While in cathodic coatings, the coating metal is selected in such a manner that it remains electropositive with respect to the base metal. For example, copper is used to coat steel.

Chemical Conversion coatings, also known as surface passivation, are produced through the chemical and electrochemical reaction of metal. Through these coatings, the surface of a metal is modified such that it possesses desired porosity. These types of coatings are more adhesive as there is a chemical bond and intermediate

*Introductory Chapter: Corrosion DOI: http://dx.doi.org/10.5772/intechopen.103791*

#### **Figure 4.**

*Corrosion protection mechanism of polymeric and modified polymeric coatings.*

layer between underlying metal and coating [22]. They are formed by immersing a metallic substrate in a chemical solution. Various types of conversion coatings are available. Phosphate coatings are produced on steels by dipping them in an appropriate phosphate solution. The thickness of these coatings depends upon the

porosity of the coatings as it forms. These coatings increase corrosion resistance, absorb lubricant, promote adhesion, and enhance the appearance of the substrate. There are three types of phosphate coatings *a) Iron phosphates, b) zinc phosphate* and *c) manganese chromium phosphate* [21]*.* Chromate conversion coatings (CCCs) are generally formed by chemical or electrochemical treatment of metals and their alloys in a solution containing hexavalent chromium [Cr (VI)] and trivalent chromium [Cr(III)] ions with other components. These coatings form a complex chromate film over the entire surface of the metal. They are used on aluminum, magnesium, zinc, copper, cadmium, etc., [23]. Anodized coatings are formed by converting the workpiece of metal into an anode. This is usually done in order to form an oxide coating to increase the performance of the surface [24]. Polymeric coatings are widely applied for decorating as well protecting purposes. They act as a corrosion barrier between the underlying metal and corrosive media. These coatings consist of pigment, polymer, corrosion inhibitors, additives, etc., [25]. The protection provided by these coatings depends upon their ability to form highly resistant pathways between the cathodic and anodic areas on the surface of metal [26]. Acrylic, vinylic, epoxy, polyurethane, alkyd (oil-based) coatings are some of the examples of organic coatings used in corrosion protection [27]. With the advancement of technology, the need to introduce specialized coating with highly advanced functioning continues to increase. Various types of speciality coating are being investigated these days, such as Flame retardant coatings, nano-coatings, nanocomposite coatings, organic–inorganic hybrid coatings, etc. All these coatings are tailored according to their end-use and the type of environment they will be applied in [28].

Corrosion prevention coatings highly utilize petro-based products which are high cost, toxic, and constantly depleting. Constant research is being carried out to formulate better strategies that can meet the environmental and economical requirements. Renewable resources are environment friendly, less expensive, and naturally available. Renewable resource-based coatings are called "*green".* **Figure 6** shows the different polymeric coating materials transformed from renewable resources.

**Figure 6.** *Polymers based on various renewable resource.*

*Introductory Chapter: Corrosion DOI: http://dx.doi.org/10.5772/intechopen.103791*

These green materials are highly employed in the field of corrosion protection in the form of corrosion-resistant coatings, inhibitors, pigments, composites, etc. Starch, cellulose, lignin, tannic acid, vegetable oils, and Cashew Nutshell Liquid (CNSL) are most commonly used for the synthesis of green coatings (**Figure 7**) [29].

Apart from using solely the polymeric coatings and relying on their chemical protective abilities, the formulations of nanostructures within the coating units have resulted in materials with higher corrosion protection efficiency [30–32]. These nanostructures can be formulated via., simple chemical reactions or by technologically advanced techniques like lithographical techniques [33]. The addition of inorganic or organic nanoparticles or units to improve corrosion protective performances has also been employed. Especially the inorganic nanoparticle dispersed within the polymeric matrix enhances the electrochemical stability, toughness, strength, and provides a tortuous pathway to the corrosive ions. The nanostructured surface coatings also tend to have higher surface hydrophobicity and scratch hardness [34]. Further, the *in-situ* synthesis of nanostructured components is also followed by many industries. The in-situ synthesis provides additional connectivity between the inorganic unit and organic unit, thereby generating well cross-linked high-density coatings [35, 36]. Due to all these value-added properties, nanocomposite coatings are getting more and more attention. In order to further enhance the efficacy of nanocomposite coatings, inhibitor-incorporated hybrid coatings are also considered innovative corrosion remediation systems [37].

#### **2.2 Corrosion inhibitors**

Pure metals and their alloys tend to react chemically and electrochemically with the corrosive environment to produce a stable compound. In this process, metals lose their mechanical strength and elasticity thereby becoming weak. Several strategies have been investigated which can retard or minimize or completely stop the cathodic or anodic or both reactions. Among them, utilization of corrosion inhibitors is a famous technique [7]. The chemical substances which when added in a small amount to the corrosive environment slow down/reduce the corrosion rate are called corrosion inhibitors. Inhibitors function through the adsorption of ions/molecules on the surface of the metal. They can drastically reduce the reaction rate either by decreasing/increasing the anodic and/or cathodic reactions or by decreasing the diffusion rate of reactants or decreasing the electrical resistance of the metallic surface. They can be easily applied and offer numerous advantages owing to their *in-situ* application without disturbing the process [38].

#### **3. Corrosion testing methods**

Two types of corrosion-resistant tests are generally used: *a)* salt spray test, *b)* Electrochemical impedance spectroscopy (EIS). Salt spray test (ASTM B117) is a standardized testing method conducted for the evaluation of the extent of corrosion resistance or protective coating. This test is generally carried out for 8-3000 hrs depending upon the coating in presence of 5% NaCl solution with pH between 6.5–7.2 [39]. EIS technique (ASTM G106) measures the impedance of the coating with the help of small amplitude, alternating current (AC). This AC signal is scanned at different frequencies to generate the spectrum for the electrochemical cell (the specimen) under the test [40].

#### **4. Computational fitting and programming**

Computational approaches like density functional theory (DFT), classical molecular dynamics (MD), Monte Carlo (MC) simulations, and others are becoming more and more preferred for corrosion studies. In order to improve the quality of the results, it is mostly combined with the experimental data. Particularly to verify the concept of corrosion inhibitors before performing the multiple experiments these studies were performed to find the thermodynamics, kinetics, energy levels (HOMO, LUMO, and Frontier molecular orbitals), bandgap energies, active sites, etc. All these theoretical calculation methodologies could be employed by artificial intelligence in the future to provide the platform which can create new and advanced corrosion inhibitors [41]. The complementary data with the experimental results puts the potential ground for the inhibitor to be explored further for corrosion inhibition applications. Even the small variations in the chemical structure which can result in better inhibition results can be evaluated through these computational approaches. Better pictorial view with the knowledge of possible interaction site, adsorption strength, and other quantum parameters proved to provide indepth insight about the inhibition mechanism as done by different researchers in the area [42, 43].

#### **5. Conclusion**

Corrosion is an electrochemical process that causes metals to degrade at an increasing rate. Various types of corrosion can never be stopped completely but various methods can be applied to decrease the rate of corrosion. These methods include the utilization of coatings as well as corrosion inhibitors or a combination of which tends to protect metals from the attack of the corrosive environment thereby increasing the lifetime of metals. A lot of work is going on to build corrosionresistant materials but more discussions over the topic and other advanced technologies which can improve the performance of the materials should be encouraged. The involvement of computational aided technologies can further help in reducing

*Introductory Chapter: Corrosion DOI: http://dx.doi.org/10.5772/intechopen.103791*

the search or designing of effective inhibitors without doing wet-lab experiments. The combined results of computational simulations and traditional electrochemical characterization techniques (impedance spectroscopy, variation in potential, corrosion current, rate of corrosion, and coating capacitance values) may further improve the understanding of this natural but unwanted phenomenon.

### **Author details**

Shumaila Masood<sup>1</sup> , Anujit Ghosal2,3, Eram Sharmin4 , Fahmina Zafar<sup>1</sup> \* and Nahid Nishat<sup>1</sup> \*

1 Inorganic Material Research Laboratory, Department of Chemistry, Jamia Millia Islamia, New Delhi, India

2 Department of Food and Human Nutritional Sciences, The University of Manitoba, MB, Canada

3 Richardson Centre for Functional Foods and Nutraceuticals, The University of Manitoba, MB, Canada

4 Department of Pharmaceutical Chemistry, College of Pharmacy, Umm Al-Qura University, Makkah Al-Mukarramah, Saudi Arabia

\*Address all correspondence to: fahmzafar@gmail.com; nishat\_nchem08@yahoo. com

© 2022 The Author(s). Licensee IntechOpen. This chapter is distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/ by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

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#### **Chapter 2**

## Corrosion: Favoured, Yet Undesirable - Its Kinetics and Thermodynamics

*Lukman O. Olasunkanmi*

#### **Abstract**

This chapter describes the fundamentals of metal corrosion in relation to thermodynamics and kinetics. The chapter is so titled, because corrosion of metal is thermodynamically favourable. Moreover, it impacts negatively on economy and safety. Industries expend a substantial percentage of their budgets on corrosion control, and lose revenue due to corrosion damage. Effects of corrosion on industrial and public infrastructure cannot be overemphasized. Several accidents in the transportation and recreational industries have been linked to corrosion of metallic parts of respective gadgets. Some of these accidents are utterly catastrophic and fatal. Therefore, corrosion, albeit its thermodynamic favouability, is not desired by man. Metals corrode as a way of minimizing energy contents. Active metals are more stable in combined forms such as oxides, sulphides, and hydroxides, even though these forms are less useful to man. It appears the "price" to pay for extracting the pure forms of these metals from their ores is corrosion. This chapter presents fundamentals of thermodynamics and kinetics of metal corrosion, with emphasis on aqueous medium. It promises to serve as an introductory chapter for corrosion science students and as a concise material for tutors.

**Keywords:** corrosion, metallurgy, spontaneity, passivity, electrochemical

#### **1. Introduction**

Corrosion is a spontaneous disintegration of materials owing to their reactions with chemical constituents of the surroundings. Materials in this context may include metals, polymers and ceramics [1, 2]. However, corrosion is mostly used to refer to undesirable destruction of metals and alloys due to interactions with surrounding environment. The interactions or reactions as used here could be chemical or electrochemical in nature. More concisely, for the purpose of this chapter, corrosion would be described as chemical or electrochemical reaction between a metal and constituents of its environment. Emphasis will be placed on corrosion of metal in aqueous environments. Corrosion is a "favoured" process, requiring little or no energy input for its occurrence. Despite being naturally favoured, corrosion imposes a lot of economic strain, health and safety threat on human and society. It is a naturally favoured and an unavoidable process, yet, undesirable by man.

Corrosion is a way by which a metal assumes a low energy state, by combining with some other elements such as oxygen, hydrogen and sulphur. Metals are

naturally not favoured to exist in free states. They are found in combined forms, often as ores. Pure metals are extracted from their respective ores with a lot of energy input. No wonder corrosion is also touted as "retro-extractive metallurgy" [3], as depicted in Eq. (1).

$$\begin{array}{cc} \text{Extractive Method} & \text{\(\pi\)} \\ \text{One} & \overbrace{\begin{array}{c} \text{\(\pi\)} \\ \Delta \mathbf{G} < \mathbf{0} \end{array}}^{} & \text{Metal (free state)} \\ & \text{Corrosion} \end{array} \tag{1}$$

Corrosion of metals is essentially an electrochemical process. A metal could assume an immune, active or a passive state when exposed to the environment, depending on the nature of the metal and the environment [1]. An index of the possible state of a metal in an environment is its electrochemical/redox potential (Er). The half-cell equation for the reduction of a metal (M) could be written as:

$$\text{M}^{\text{n}+}\_{\text{(aq)}} + \text{ne}^- \rightleftharpoons \text{M}\_{\text{(s)}}\text{ E} = \text{E}\_{\text{r}}\text{ (Volts)}\tag{2}$$

Metals with highly positive Er are naturally immune to oxidation, examples include Au and Pt. Those with negative Er are active in the environment. Some of them may assume an active-passive state, depending on the nature of the environment and the properties of their corrosive products.

Understanding the thermodynamics and kinetics of metal corrosion would make corrosion science students to appreciate why corrosion occurs and how its rate could be mitigated. It would also help corrosion science tutors to lay a good foundation of the course for their students.

#### **2. The fate of a metal in an environment**

The behaviour of a metal upon exposure to an environment depends on the nature of the metal and the conditions of the environment. A metal exposed to a corrosive environment could behave in one of the following ways [1]:


solutions and found that Mg exhibited the highest corrosion rate. The study also illustrates how corrosion rate could vary with immersion time.

iii. Passive behaviour: a metal that exhibits passive behaviour might corrode at some instances but assume a non-responsive state over a period. Upon exposure to the corrosive environment, a metal with passive behaviour forms insoluble protective corrosion products, which cover the surface of the metal and shield it from further exposure to the corrosive environment. The protective film slows down further reaction with the environment. The films are usually more stable (less reactive) than the metals themselves. However, if the passive film is broken or dissolved, then the metal can revert to the active state, at which instance rapid dissolution of the metal could occur. In some cases, repassivation could follow the breakdown of passive film. Various factors could be responsible for passive film breakdown or instability. Among them are film thickness, nature of corrosive ions, pH, anodic potentials and so on. Feng *et al.* [5] observed that increased strain magnitude could also increase instability of passive film on carbon steel. A more detailed overview on the passivity of metals could be found in literature [6].

It is noteworthy that the behaviour of a metal depends on its "microenvironment". The natural or real practical environments are characterized by variable factors. The conditions of the environment may change with time. There could be a change in the cell pH, fluid flow rate, and temperature. Some reactions in the environmental might also result in solid deposits. The actual environment to which a metal responds is the immediate local environment at or near its surface. It is this micro-environment that determines the behaviour of the metal. Though some metals exhibit nearly universal behaviour irrespective of the environmental conditions. Metals such as Au, Pt and Ag typically exhibit noble or immune behaviour regardless of the environment, while metals like Na, K and Mg are generally active in nearly all aqueous environments. Metals likes titanium and tantalum assume passive state in a wide range of aqueous environments, though reactive in some other environments. Aluminium and zinc are very reactive metals and often exhibit active behaviour. However, they form stable protective passive films in some environments. The behaviour of such metals could be described as "active-passive" – active in some environments at some instances, and sometimes passive. The subsequent passivity is due to the change in the local environment at the surface of the metal, being occupied by the insoluble passive corrosion products.

Active-passive behaviour of iron in nitric acid was first observed in 1790 by Keir [7]. The thickness of passive films may vary with environmental conditions. Susceptibility of the film to breakdown also depends on the thickness [6]. Sato *et al.* [8] also reported that the composition and thickness of passive films on iron immersed a borate solution could change with change in potentials. Luo *et al.* [9] observed that alloy 59 (a Ni-Cr-Mo alloy with the least Fe content and highest Cr-Mo content) develops thicker passive film in air than sulphuric acid solution, and the constituents of the film vary in the two environments.

#### **3. Corrosion: thermodynamic driving force**

As pointed out earlier, corrosion of a metal is a spontaneous process. A metal in its pure state has a considerably higher energy than its corresponding ore. Metal corrodes in an attempt to minimize its energy, while assuming a more stable state.


**Table 1.**

*Minerals or ores of some common metals, their chemical formulas and standard heats of formation.*

Corrosion is therefore a means of energy minimization as a metal tends to return to its combined form in which it exists naturally. A basic illustration of this energy minimization is the exothermic nature of the formation of metal oxides, sulphides and hydroxides (**Table 1**). Formation of hematite (the world most important ore of iron) is accompanied by the release of a huge amount of energy (825.50 kJ/mol). Meanwhile, the relative energy of a free Fe atom is 0 kJ/mol. The same applies to other metals and their minerals.

#### **3.1 Corrosion: electrochemical cell and electrode potential**

Corrosion of metals is essentially an electrochemical process, involving both anodic oxidation and cathodic reduction reactions. A micro-electrochemical cell is established on the surface of a corroding metal. Perhaps for the sake of emphasis, electrochemical corrosion cell is a galvanic (or voltaic) cell. The progress of corrosion reaction is accompanied by flow of electric current (*i*), which has to do with movement of an electric charge across a potential difference. A corroding metal in an aqueous solution sets up a galvanic cell system comprising the metal (M) in contact with its metal ion (Mn+) such that an equilibrium is established. The site on the metal surface where dissolution of metal into its ion occurs is the anodic site. The cathodic site is set up not far from the anodic site. Each site constitutes a halfelectrode reaction system, making up two half-cells, like that of a galvanic cell. The difference between the electrode potentials of the two half-cells can be expressed as:

$$E\_{cell} = E\_{cathode} - E\_{anode} \tag{3}$$

In Eq. (3), both *E*cathode and *E*anode are reduction potentials. Since we are considering corrosion, the *E*cell is equivalent to the corrosion potential, *E*corr. For a spontaneous cell reaction as we have in corrosion, *E*anode is always more negative than *E*cathode, such that *E*cell is always positive.

For a typical case of iron (Fe) corroding in an aerated aqueous solution as:

$$\text{Fe}\_{\text{(s)}} + \text{H}\_2\text{O}\_{\text{(l)}} + \text{V}\prime\text{O}\_2 \rightleftharpoons \text{Fe}(\text{OH})\_{\text{2(g)}}\tag{4}$$

The half-cell reactions equations can be expressed as:

$$\text{Anode}: \text{Fe}\_{\text{(s)}} \rightleftharpoons \text{Fe}^{2+} \text{(aq)} + 2\text{ e}^- \text{ E}^0 \text{}\_{\text{ox}} = +0.409 \text{ V} \tag{5}$$

$$\text{Cathode}: \text{\textbullet } \text{O}\_{2(g)} + \text{H}\_2\text{O}\_{(l)} + \text{2 }\text{e}^- \rightleftharpoons \text{2 }\text{OH}^-\text{\textbullet}^0\text{ }\_{\text{(aq)}}\text{ E}^0\\ \text{\textbullet } = +0.20\text{ V}\tag{6}$$

In this case, Ecell = (0.2 + 0.409) V = 0.609 V. The more positive the Ecell, the more feasible the corrosion of the metal. Practically, the Ecell or Ecorr is measured *Corrosion: Favoured, Yet Undesirable - Its Kinetics and Thermodynamics DOI: http://dx.doi.org/10.5772/intechopen.98545*

using a potentiostat. In such measurements, a half-cell must be chosen as a reference, e.g. standard hydrogen electrode (SHE). Electrode potential for SHE is set at 0V(*E*H+/H2 ). The electrode is made up of 1 M hydrogen ion (H+ ) and hydrogen gas (H2) at 1 atm, supported with a platinum plate. When coupled with an half-cell of unknown *E*, the cell potential recorded by the potentiometer is the electrode potential of the system with unknown *E*.

Potentiostats with three-electrode system are often used to measure electrochemical corrosion parameters. The three-electrode system consists of the working electrode, WE (the metal or alloy whose corrosion is being studied), reference electrode, RE (against which the corrosion potential of the metal/alloy is measured), and auxiliary/counter electrode, CE, which supports or protects the reference electrode against passage of current. Commonly used reference electrodes include calomel electrode, which composed of Hg/Hg2Cl2, sat'd. KCl, and silver-silver chloride electrode (Ag/AgCl, sat'd KCl). The latter is mostly used because of its relatively cheap cost and less toxicity compared to the mercury-based electrode.

Electrode potentials depend on concentrations of the species and temperature. Under the standard conditions of 25°C and 1 M concentration or 1 atm pressure of the species, it is referred to as the *standard electrode potential*. The standard electrode potentials (standard reduction potentials), E0 for some species are listed in **Table 2**. The dependence of electrode potentials on concentrations of the species and temperature is expressed in the form of the Nernst equation, whose general form is:

$$E\_r = E\_r^0 - \frac{RT}{nF} \ln\left(\frac{[red]}{[oxi]}\right) \tag{7}$$

where *n* is the number of electrons transferred in the redox reaction, F is the Faraday constant, R is the gas constant and T is absolute temperature; [red] and [oxi] are the concentrations of the reduced and oxidized species, respectively.

#### **3.2 Gibb's free energy and electrode potentials**

Corrosion is characterized by lowering of Gibb's free energy or increasing electrochemical cell potential. For a corroding metal, a micro-electrochemical cell is created on the surface. The progress of metal corrosion is proportional to flow of current in the electrochemical cell. Thermodynamic parameters can be expressed for electrochemical systems. Since corrosion of metals is a constant pressure process, the Gibb's free energy (ΔG) is a good thermodynamic parameter for predicting its spontaneity. A spontaneous reaction is accompanied by energy minimization, which implies a negative ΔG, while a positive ΔG connotes non-spontaneous process. A non-spontaneous reaction requires energy input to proceed. For a system at equilibrium, ΔG = 0. Corrosion being a spontaneous process has a negative ΔG. Thermodynamic favourability of metal corrosion could readily be predicted from the electrode potentials of the metal concerned. The change in Gibb's free energy (ΔG) per mole of an electrochemically reacting species is related to the electrode potentials as:

$$
\Delta \mathbf{G} = -\mathbf{n} \mathbf{F} \mathbf{E} \tag{8}
$$

where n is the valence of the species (number of electrons transferred), F is the Faraday's constant (1 F = 96,485 C) and E is the electrode potential (Volts). For a typical case of anodic dissolution of iron as it oxidizes from Fe to Fe2+ by losing two electrons (Fe ⇌ Fe2+ +2e�); *n* = 2.


**Table 2.** *Standard half-cell reduction potentials for reactions.*

#### **3.3 Impracticability of equilibrium electrochemical corrosion potentials**

Corrosion involves both anodic and cathodic reactions. Each of these reactions is reversible and has associated electrode potential (E), which tends to attain an equilibrium value. At equilibrium, ΔG = E = 0. However, attainment of this value is impracticable.

Practically, a bare (an oxide-free) metal surface releases metal ion into an aqueous solution (dissolution), leaving negatively charged electrons on the surface. This leads to an increase in the potential difference between the metal and the solution. The electrode potential becomes more negative. For the anodic dissolution of a metal, M, the half-cell reaction equation could be written as:

$$\mathbf{M} \not\equiv \mathbf{M}^{\text{n}+} + \mathbf{n} \mathbf{e}^{-} \tag{9}$$

*Corrosion: Favoured, Yet Undesirable - Its Kinetics and Thermodynamics DOI: http://dx.doi.org/10.5772/intechopen.98545*

A more negative potential tends to retard dissolution but promotes deposition, according to Eq. (8) (i.e. ΔG is more positive for a more negative E). Since the process is reversible, continuous dissolution and/or deposition might lead to a stable (reversible) potential, Er, which can be expressed in the form of Nernst Equation (for the reaction in Eq. (9)) as:

$$E\_{r, \mathcal{M}^{n+} / \mathcal{M}} = E\_{r, \mathcal{M}^{n+} / \mathcal{M}}^{0} + \frac{RT}{nF} \ln a\_{\mathcal{M}^{n+}} \tag{10}$$

If a stable *Er*,*Mn*þ*=<sup>M</sup>* is attained, dissolution would stop. However, in practice, *Er*,*Mn*þ*=<sup>M</sup>* is never attained because electrons generated in Eq. (9) are always removed from the surface by the accompanied cathodic half-cell reaction, such as:

$$2\text{H}^+\text{(aq)} + 2\text{ e}^- \rightleftharpoons \text{H}\_{2(g)}\tag{11}$$

Or

$$2\text{ }\text{Al}\_2\text{O}\_{2(g)} + \text{H}\_2\text{O}\_{(l)} + 2\text{ e}^- \rightleftharpoons 2\text{ }\text{OH}^-\text{ }^\text{(aq)}\tag{12}$$

in an acidic or basic medium respectively. The Nernst equation-type expressions for the reversible potential (Er) for the cathodic reactions in Eqs. (11) and (12) respectively are:

$$E\_{r,H^{+}/H\_{2}} = E\_{r,H^{+}/H\_{2}}^{0} - \frac{RT}{F} \ln \frac{P\_{H\_{2}}^{\frac{1}{2}}}{a\_{H^{+}}} \tag{13}$$

$$E\_{r,O\_2/OH^-} = E\_{r,O\_2/OH^-}^0 - \frac{RT}{4F} \ln \frac{a\_{OH^-}^4}{P\_{O\_2}} \tag{14}$$

where Eqs. (13) and (14) correspond to stoichiometrically adjusted forms of Eqs. (11) and (12) by multiplying the coefficients by ½ and 2 respectively for reduction of 1 mole of H+ and O2.

If Er could be attained for reactions depicted by Eqs. (11) or (12), then Er would be attained in Eq. (10). However, Er in Eq. (13) or (14) is never stable due to continuous discharge of H2 or consumption of O2. Hence, attainment of a stable *Er*,*Mn*þ*=<sup>M</sup>* is practically impossible and corrosion of metal, M is continuous.

#### **3.4 Corrosion tendency based on electrochemical potentials and pH**

Thermodynamics of electrochemical corrosion could be described as a function of electrode potential and hydrogen ions concentration (pH). This is often chatted as potential-pH diagram, popularly called the Pourbaix diagram, named after the original inventor. Pourbaix, a Belgium electrochemist and corrosion scientist invented the potential-pH diagram in 1963 for the description of thermodynamics of electrochemical corrosion. Pourbaix diagrams provide theoretical description of stability of a phase of metal/electrolyte system at a particular pH and potential. It is a kind of phase equilibrium diagram, though with different axes parameters compared to thermodynamics phase equilibrium diagram. Potential-pH diagram is often plotted at 25°C, 1 atm, and 10�<sup>6</sup> M concentration of the ionic species. A typical Pourbaix plot comprises the redox potential on the vertical axis and the pH on the horizontal axis. Potential-pH diagram for iron in aqueous environment is shown in **Figure 1**. The diagram clearly shows the stable and passive regions for iron, based

**Figure 1.** *Potential-pH diagram for iron in aqueous medium [10].*

on the combination of potential and pH. At the very top of the diagram is corrosive region, where the potential is highly positive (above 1.8 V). At this region, iron would corrode at any combination of potential and pH. The region marked immunity at the lower portion of the diagram indicates the area where iron does not corrode. This region spans over a wide range of pH, and a limited range of highly negative potentials. Within this region, iron is immune to corrosion at various combinations of potentials and pH. In-between the two extreme top and bottom regions are regions where corrosion and/or passivity could occur, depending on the operating potentials and pH. The region marked "corrosion" between lines 1 and 2 covers a large area compared to the one at lower right-hand side. This implies that corrosion of iron at intermediate potentials between 500 mV and 1000 mV would progress more favourably in acidic pH than neutral and alkaline pH. At moderately positive potentials and neutral or alkaline pH, iron forms passive oxide film on the surface, which blocks further corrosion.

Pourbaix diagram could be used as a route to first principle corrosion simulation [11], a model for optimizing corrosiveness of a medium and designing materials with desired corrosion resistance. In a recent study, Nave and Kornev [12] constructed and applied 3D Pourbaix profile to establish the conditions for thermodynamic stability of tungsten-based compounds and describe the anodic dissolution of tungsten in aqueous solutions of potassium hydroxide. Beyond pure metals, Pourbaix profiles for multielement system such as Ni-Ti alloys having different ratios of the element have been proposed [13]. In an effort to overcome the challenges associated with developing Pourbaix profiles for complex compounds, a recent study by Patel *et al.* [14] introduced a more robust algorithm for modelling Pourbaix diagram for multicomponent materials.

*Corrosion: Favoured, Yet Undesirable - Its Kinetics and Thermodynamics DOI: http://dx.doi.org/10.5772/intechopen.98545*

#### **4. Corrosion kinetics**

Having discussed the propensity of metals to corrode, it is important to also highlight the rate at which metals corrode and mechanisms of corrosion reactions.

#### **4.1 Corrosion rate**

In a general term, corrosion rate (CR) refers to the amount of metal loss to corrosion per unit time. The rate at which a metal corrodes can be monitored by various methods. This also determines the kind of CR expression and units. Both the methods and expression (together with units) are in turn dependent on the technical system and type of corrosion being investigated. These methods may be classified as chemical, electrochemical, spectroscopic, and surface analysis based methods. Basically, corrosion rate (CR) may be monitored by measuring any parameter that changes as corrosion reactions occur. For example:


Each of these methods have associated merits and demerits. Corrosion rate may expressed as weight-loss of a metal per unit time per unit area according to the equation:

$$\mathbf{C}\_{R} = \frac{\Delta w}{A t} \tag{15}$$

where Δ*w* is the weight difference of the metallic block or plate at a set time interval (exposure time), *t*, and *A* is the exposed area of the metal. Being an analytical measurement, the ideal practice is to conduct repetitive measurements of Δ*w* and utilize the average value in Eq. (15). This measurement is mainly applicable to general or uniform corrosion. However, it is the most used measurement of corrosion rate. For corrosion systems in which the metal/environment composition vary significantly with time and non-uniformity over the sites on corroding surface, measurement of CR using Eq. (15) may be deficient. The results of such measurements must at least be supported with additional information such as the type of corrosion, dependence of corrosion rate on time and other relevant factors that prevail during the experiment. Eq. (15) is the simplest form of such expression, for which the units of Δ*w* (g or nearest mg), *A* (nearest cm2 ), and t (nearest h) would give the units of CR as gcm�<sup>2</sup> h�<sup>1</sup> or mgcm�<sup>2</sup> h�<sup>1</sup> . These units may not be applicable in reporting CR in some other technical reports, for example, if CR is to be expressed in the form of penetration rate (depth per unit time).

In an electrochemical experiment, corrosion rate could be measured as corrosion current density, in the form of corrosion current (mA) per unit area (cm<sup>2</sup> ) of


**Table 3.**

*Conversion factors for various units of CR [15].*

corroding metal, i.e. mAcm<sup>2</sup> . Conversion factors from one CR unit to another are listed in **Table 3**.

where:


e.g.: 1 mA cm<sup>2</sup> = (3.28 M/nd) mm y<sup>1</sup> = (129 M/nd) mpy = (8.95 M/n) g m<sup>2</sup> day<sup>1</sup>

#### **4.2 Factors affecting corrosion rate**

Besides basic requirements for electrochemical corrosion to occur, which include the presence of anodic and cathodic sites, electrolyte and connectivity between the sites (to promote ionic conduction), there are secondary factors that affect corrosion rate. These factors are briefly discussed as follow.


corrosion rate in seawater is a function of numerous mutually dependent factors. However, according to Kirk and Pikul (1990) [16], if salinity exceeds 3%, water corrosivity decreases. This is due to decrease in oxygen solubility is in water with > 3% salinity, as posited by Weiss (1970) [17].

• *Temperature and pressure:* Generally, temperature increases the rate of metal corrosion, so does pressure. Just like every other electrochemical reactions, increase in temperature increases fluid flow and ionic mobility. Temperature can also affect scale formation and gas fugacity, which indirectly affect corrosion rate. Most corrosion models are accurate only within prescribed temperature ranges. For corrosive involving gases such as CO2 and H2S increase in operating pressure could increase the partial pressure of the corrosive gases. It should be noted that the reduction potential of the metal is dependent on fugacity of the gases present, according to the Nernst Equation. The dependence of corrosion rate (CR) on temperature could be expressed in the form of Arrhenius equation as:

$$\mathbf{C}\_{\text{R}} = \mathbf{A} \mathbf{ex} (-\mathbf{E}\_{a}/\mathbf{RT}) \tag{16}$$

where A is the pre-exponential factor, Ea is the activation energy, R is the gas constant and T is temperature.


#### **4.3 Corrosion mechanism**

Corrosion, like many other chemical reactions usually involve more than one definable step. Interests often lie in the slowest step. Electrochemical corrosion involves release of ions and movement of electrons. Corrosion requires presence and movement of ions and electrons. A typical mechanism of corrosion of Fe in acidic medium is:

> Feð Þ<sup>s</sup> ⇌ Fe<sup>2</sup><sup>þ</sup> ð Þ aq þ 2e� ð Þ Anodic dissolution of metal (17)

$$\text{2H}^+\text{(aq)} + \text{2e}^- \rightleftharpoons \text{H}\_{2(g)} \text{ (Cathodic reaction-H}\_2 \text{ gas evolution)}\tag{18}$$

The process is often more complex such that, metal ions may go into solution as complexes or even; precipitate as hydroxides, oxides, sulphides etc.

Electrochemical corrosion mechanism of an active metal in an aqueous environment can be expressed generally as [18]:

$$\mathbf{M}\_{(s)} \rightleftharpoons \mathbf{M}^{\mathrm{m}+}\mathrm{\_{(aq)}} + \mathrm{me}^- \tag{19}$$

$$\text{mH}^+\text{}\_{\text{(aq)}} + \text{me}^- \rightleftharpoons \text{m/2 H}\_{2\text{(g)}}\tag{20}$$

$$\text{Overall}: \text{M}\_{\text{(s)}} + \text{mH}^{+}\text{}\_{\text{(aq)}} \rightleftharpoons \text{M}^{\text{m}+}\text{}\_{\text{(aq)}} + \text{m/2 H}\_{\text{2(g)}}\tag{21}$$

In a deaerated acidic medium. Or.

$$\mathbf{M}\_{\mathrm{(s)}} \rightleftharpoons \mathbf{M}^{\mathrm{m}+}\mathrm{\_{(aq)}} + \mathrm{me}^- \tag{22}$$

$$\text{mH}\_2\text{O}\_{\text{(l)}} + \text{me}^- \rightleftharpoons \text{mOH}^-\text{(aq)} + \text{m/2 H}\_2\text{(g)}\tag{23}$$

$$\text{Overall}: \text{M}\_{\text{(s)}} + \text{mH}\_{2}\text{O}\_{\text{(l)}} \rightleftharpoons \text{M}^{\text{m}+}\text{(aq)} + \text{mOH}^{-}\text{(aq)} + \text{m/2 H}\_{2\text{(g)}}\tag{24}$$

In a deaerated neutral/basic medium.

In an aerated environment, oxygen plays a prominent role in the reaction and the mechanisms look like [18]:

$$\mathbf{M}\_{(s)} \rightleftharpoons \mathbf{M}^{\mathrm{m}+} + \mathbf{m} \; \mathbf{e}^- \tag{25}$$

$$\text{m/} \text{4 O}\_{2(g)} + \text{m H}^+ \text{}\_{(aq)} + \text{m e}^- \rightleftharpoons \text{m/} 2 \text{ H}\_2\text{O} \tag{26}$$

$$\text{Overall}: \text{M}\_{\text{(s)}} + \text{m}/\text{4} \,\text{O}\_{2(g)} + \text{m} \,\text{H}^+\text{}\_{\text{(aq)}} \rightleftharpoons \text{M}^{\text{m}+} + \text{m}/2 \,\text{H}\_2\text{O}\_{(l)}\tag{27}$$

In an aerated acidic environment.

Or

$$\mathbf{M}\_{(s)} \rightleftharpoons \mathbf{M}^{\mathrm{m}+} + \mathbf{m} \; \mathbf{e}^- \tag{28}$$

$$\text{m} / 4\text{ }\text{O}\_{2(g)} + \text{m} / 2\text{ }\text{H}\_2\text{O}\_{(l)} + \text{m }\text{e}^- \rightleftharpoons \text{m }\text{OH}^-\tag{29}$$

Overall: M(s) + m/4 O2(g) + m/2 H2O(l) ⇌ Mm+ + m OH� (aq). In an aerated neutral or basic environment.

#### **5. Corrosion control methods**

As contained in the previous sections, whether a material would corrode or not depends on a number of factors. The extent and rate of corrosion also varies, depending on the nature of the metal or alloy, corrosive medium, pH, temperature and so on. While it may be difficult to "work against nature", to completely stop metal corrosion, several methods of abating corrosion damage have been identified. Most of these methods reduce corrosion rate, rather than "inverting" the thermodynamics. Although, thermodynamic susceptibility might be influenced in some instances. Choosing a suitable corrosion control method requires proper understanding of corrosion type and mechanism. No particular method had been adjudged universally effective to mitigate all forms of corrosion. Common corrosion control methods are highlighted below.

• *Material selection:* this method entails careful selection of metal or alloy that is immune to corrosion in an environment. While this sounds as a good idea, such *Corrosion: Favoured, Yet Undesirable - Its Kinetics and Thermodynamics DOI: http://dx.doi.org/10.5772/intechopen.98545*

materials do not often possess the desired mechanical properties for prospective technological design. Besides, using such materials for engineering purposes could be very expensive. In other words, striking a balance between material cost, mechanical properties and corrosion resistance is not a straightforward task.


*Corrosion - Fundamentals and Protection Mechanisms*

### **Author details**

Lukman O. Olasunkanmi Department of Chemistry, Faculty of Science, Obafemi Awolowo University, Ile-Ife, Nigeria

\*Address all correspondence to: waleolasunkanmi@gmail.com

© 2021 The Author(s). Licensee IntechOpen. This chapter is distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/ by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

*Corrosion: Favoured, Yet Undesirable - Its Kinetics and Thermodynamics DOI: http://dx.doi.org/10.5772/intechopen.98545*

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### Section 2
