2.2 Anodic & Cathodic Reactions on Carbon Steel

Anodic & Cathodic Reactions on Carbon Steel

Understanding the Chemistry of Corrosion

To effectively monitor and mitigate corrosion, particularly on carbon steel structures like pipelines, tanks, and offshore platforms, one must understand the specific chemical reactions occurring at the surface of the metal. As established in the previous section, corrosion consists of coupled anodic and cathodic reactions. The anodic reaction produces electrons, and the cathodic reaction consumes them.

Because we are dealing with carbon steel, which is primarily iron (Fe), the reactions we focus on are those involving iron and the common environments it is exposed to: soil, water, and atmosphere. The specific cathodic reaction that occurs depends heavily on the pH of the electrolyte and the presence or absence of dissolved oxygen.

The Anodic Reaction (Oxidation)

The anodic reaction is relatively straightforward. When carbon steel corrodes, iron atoms at the surface of the metal give up electrons and dissolve into the electrolyte as positively charged ferrous ions. This is an oxidation process (loss of electrons).

The fundamental anodic reaction for carbon steel is:

Fe -> Fe2+ + 2e-

In this equation:

• Fe represents a neutral solid iron atom in the steel structure.
• Fe2+ represents a ferrous ion dissolved in the electrolyte.
• 2e- represents the two electrons that remain in the metal and travel through the metallic path.

This reaction is the actual physical destruction of the steel. Every time this reaction occurs, a tiny piece of the pipeline or tank is lost. The rate at which this reaction occurs dictates the corrosion rate of the structure.

The Cathodic Reactions (Reduction)

The cathodic reactions are more varied and depend on the environment. The cathode acts as an electron sink. The electrons generated by the anodic reaction must be consumed for corrosion to continue. The most common substances that consume these electrons are dissolved oxygen (O2) and hydrogen ions (H+).

1. Neutral or Alkaline Environments with Dissolved Oxygen

In most natural environments, such as typical soils, fresh water, and seawater, the pH is near neutral (pH 6-8) or slightly alkaline. In these environments, if dissolved oxygen is present, the dominant cathodic reaction is the oxygen reduction reaction. The oxygen reacts with water and the electrons from the anode to form hydroxide ions.

The reaction is:

O2 + 2H2O + 4e- -> 4OH-

In this equation:

• O2 is dissolved oxygen gas from the environment.
• 2H2O is liquid water in the electrolyte.
• 4e- are the four electrons arriving from the anode via the metallic path.
• 4OH- are four hydroxide ions produced and released into the electrolyte.

The production of hydroxide ions (OH-) increases the local pH at the cathode surface, making it more alkaline. This alkaline environment can sometimes be beneficial as it can lead to the precipitation of calcareous deposits (calcium carbonate and magnesium hydroxide) from seawater or hard water, which form a protective scale on the cathode, reducing further corrosion.

2. Acidic Environments

In acidic environments (low pH, typically below 5), there is an abundance of hydrogen ions (H+). In these conditions, the dominant cathodic reaction is often the reduction of hydrogen ions to form hydrogen gas.

The reaction is:

2H+ + 2e- -> H2

In this equation:

• 2H+ are hydrogen ions present in the acidic electrolyte.
• 2e- are the electrons from the anode.
• H2 is hydrogen gas, which often forms bubbles on the cathode surface.

This reaction is significant in environments like acid bogs, industrial chemical spills, or within deep corrosion pits where localized acidity can develop. The formation of hydrogen gas bubbles can sometimes physically disrupt protective coatings and, under certain conditions, the atomic hydrogen formed before combining into H2 gas can enter the steel structure, leading to hydrogen embrittlement.

3. Neutral or Alkaline Environments without Dissolved Oxygen (Anaerobic)

In environments devoid of oxygen (anaerobic), such as deep, waterlogged clays or beneath thick bio-films, the oxygen reduction reaction cannot occur. If the environment is neutral or alkaline, the reduction of water itself can become the primary cathodic reaction, although it proceeds at a very slow rate.

The reaction is:

2H2O + 2e- -> H2 + 2OH-

This reaction produces both hydrogen gas and hydroxide ions. However, in many anaerobic soil environments, another factor comes into play: Microbiologically Influenced Corrosion (MIC). Sulfate-reducing bacteria (SRB) can utilize the hydrogen produced at the cathode (or directly utilize electrons) to reduce sulfate ions (SO4 2-) to hydrogen sulfide (H2S), significantly accelerating the corrosion process in the absence of oxygen.

Secondary Reactions and Corrosion Products

The primary anodic and cathodic reactions describe the initial transfer of electrons. However, the products of these reactions (Fe2+ and OH-) do not remain isolated. They migrate through the electrolyte and react with each other to form secondary corrosion products, commonly known as rust.

When ferrous ions (Fe2+) from the anode meet hydroxide ions (OH-) from the cathode, they combine to form ferrous hydroxide:

Fe2+ + 2OH- -> Fe(OH)2

Ferrous hydroxide is relatively unstable. If oxygen is present in the environment, it will further oxidize the ferrous hydroxide into ferric hydroxide, which is the familiar reddish-brown rust:

4Fe(OH)2 + O2 + 2H2O -> 4Fe(OH)3

The formation of rust is a secondary process. The actual metal loss occurred during the primary anodic reaction (Fe -> Fe2+ + 2e-). The type of rust formed and how tightly it adheres to the metal surface can influence the ongoing corrosion rate by either providing a protective barrier or acting as a porous sponge that retains moisture and corrosive salts.

Mastering these reactions is critical for CP testers because cathodic protection works by intentionally driving the structure to act purely as a cathode. By supplying electrons from an external source (a rectifier or sacrificial anode), the CP system forces the cathodic reactions to occur over the entire surface of the steel, thereby suppressing the anodic reaction (metal dissolution) and halting corrosion.