Rusting of Iron: Explanation, Chemical Reaction, Prevention

The process of rusting of iron pertains to the occurrence of rust, which is a composite of iron oxides, on the exterior of iron-based items or constructions. The formation of rust occurs as a result of a Redox Reaction between oxygen and iron within a water-rich environment, such as air with elevated moisture levels. The process of iron rusting is distinguished by the development of a reddish-brown, fragmented substance that exhibits a flaky texture and readily disintegrates into a powdery form.

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One prevalent instance of corrosion involves the detachment of corrode in a metallic surface, whereby the iron undergoes a transformation into rust, characterized by an orange-red powdered form. This transformation exposes the underlying fresh metal surface, which allows for the ongoing progression of the corrosion process.

Chemistry of Iron Rusting

Rust formation is a chemical process that takes place when iron, or an iron alloy, is exposed to Oxygen (O) in the presence of moisture. The observed reaction does not exhibit an instantaneous occurrence but rather tends to unfold gradually over a considerably prolonged period of time. The phenomenon of oxygen atoms engaging in chemical reactions with iron atoms leads to the formation of compounds known as iron oxides. This phenomenon leads to a decrease in the interatomic bonding between the iron atoms that make up the object or structure.

The rusting reaction involves two key components: an elevation in the oxidation state of iron and the relinquishment of electrons. Rust primarily consists of two distinct iron oxides, exhibiting varying oxidation states of the Iron (Fe) atom. The subsequent substances can be categorized as oxides:

  • Iron (II) oxide, alternatively referred to as ferrous oxide, is a chemical compound that consists of the elements iron and oxygen. The compound FeO is distinguished by an oxidation state of +2 for the iron atom.
  • Iron (III) oxide, also known as ferric oxide, is a chemical compound that is distinguished by the iron atom having an oxidation state of +3. The chemical compound currently being discussed can be represented by the molecular formula Fe2O3.

Oxygen exhibits strong oxidizing properties, while iron demonstrates reducing capabilities. Hence, the iron atom exhibits a high tendency to donate its electrical charge upon exposure to oxygen. The following equation can represent the chemical reaction:

Fe → Fe2+ + 2e-

The presence of water induces an elevation in the oxidation state of iron through its interaction with the oxygen atom.

4 Fe2+ + O2 4 Fe3+ + 2 O2-

The subsequent acid-base reactions take place between the iron cations and the water molecules.

Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2H+

Fe3+ + 3 H2O ⇌ Fe(OH)3 + 3H+

The formation of iron hydroxides, which takes place through a direct reaction between iron cations and hydroxide ions.

The chemical equation represents the reaction of oxygen (O2) with water (H2O) in the presence of four electrons (4e) to produce hydroxide ions (OH).

O2 + H2O + 4 e → 4 OH

The chemical equation represents the reaction between Fe2+ ions and OH ions to form Fe(OH)2.

Fe2+ + 2 OH  → Fe(OH)2

The chemical equation represents the reaction between Fe3+ ions and OH ions, resulting in the formation of Fe(OH)3.

Fe3+ + 3 OH  → Fe(OH)3

The hydroxides of iron undergo a dehydration process, leading to the formation of iron oxides, which are accountable for the occurrence of rust. The process involves a variety of chemical reactions, a subset of which is presented in the following.

The chemical equation representing the equilibrium reaction between Fe(OH)2, FeO, and H2O can be expressed as follows:

Fe(OH)2 ⇌ FeO + H2O

The balanced chemical equation for the reaction is as follows:

4 Fe(OH)2 + O2 + xH2O → 2 Fe2O3.(x+4)H2O

The chemical equation representing the equilibrium between Fe(OH)3 and FeO(OH) in the presence of water can be expressed as follows:

Fe(OH)3 ⇌ FeO(OH) + H2O

The chemical equation provided represents the equilibrium reaction between 2FeO(OH) and the products Fe2O3 and H2O

2 FeO(OH) ⇌ Fe2O3 + H2O

A common characteristic shared by all the chemical reactions mentioned above is their reliance on the presence of both water and oxygen. Hence, the corrosion process of iron can be effectively managed by restricting the presence of oxygen and water in the vicinity of the metal.

Factors Contributing For Rusting of Iron

There are multiple factors that can expedite the process of iron corrosion, including the humidity level in the surrounding atmosphere and the pH level of the immediate environment. Please find below a compilation of several factors.


The presence of moisture in the environment plays a crucial role in regulating the corrosion process of iron. The primary cause of corrosion is being exposed to precipitation, particularly rain.


The acceleration of the corrosion procedure can be observed when the value of the pH of the metal’s immediate surroundings is low. The rate of iron corrosion is accelerated upon exposure to acid rain. An elevated pH level has the effect of impeding the process of iron corrosion.


The acceleration of iron corrosion in saltwater environments can be attributed to the presence of various salts. The presence of diverse ions in saline water facilitates the expedited progression of the corrosion phenomenon via electrochemical reactions.


The presence of contaminants in iron, primarily an amalgamation of metals, has been observed to accelerate the process of rusting in comparison to pure iron.

The process of oxidation can be influenced by the dimensions of the iron object. For example, it is likely that a substantial iron object would exhibit slight imperfections as a result of the smelting process. The aforementioned limitations represent a susceptible aspect that renders the metal susceptible to environmental assaults.

Impacts of Rusting Iron

The process of iron oxidation, commonly known as rusting, has the potential to significantly influence various aspects of our everyday existence. Due to its widespread utilization in contemporary society, iron and its alloys have become prevalent in the construction of various products and structures, thereby encompassing our surroundings. This discussion will address several adverse effects associated with the process of iron rusting.

Effects The Automobiles

The process of rusting leads to the development of a flaky and weakened state in iron, resulting in a degradation of its structural integrity, aesthetic qualities, and permeability. The corrosive process of iron results in the loss of its desired characteristics. The process of iron oxidation, commonly known as rusting, has the potential to cause detrimental effects on various iron-based structures that include automobiles, railings, grills, and other similar constructions.

Effects The Buildings And Structures

The occurrence of the Silver Bridge’s collapse in 1967 and the subsequent collapse of the Mianus River Bridge in 1983 were both ascribed to the corrosive deterioration of the steel and iron constituents of these bridges. The structural failures that affect many reinforced concrete buildings can occur over extended periods of time as a result of corrosion.

Effects On Health

The presence of rusted iron can serve as a conducive environment for the proliferation of bacteria responsible for the onset of tetanus. Incisions caused by these objects that penetrate the dermis can pose a risk to one’s well-being.


Due to the heightened rate of rusting in environments with high humidity, the internal surfaces of water pipes and tanks are prone to corrosion. As a result of this phenomenon, the pipes become conduits for discharging water that exhibits a discernible brown or black hue, indicating the presence of iron oxides at levels that exceed safety thresholds.

Preventive Methods To Avoid Iron Rusting

The manufacture of numerous machines, items, and structures all make extensive use of iron and its alloys. Therefore, it is vital to avoid iron from corroding. Here are a few preventative measures.

Using Rust Resistant Alloys

  • Certain alloys of iron exhibit resistance to rust formation. Illustrative instances encompass stainless steel, characterized by the presence of a chromium(III) oxide layer, as well as weathering steel.
  • The COR-TEN steel exhibits a comparatively reduced rate of rusting in comparison to conventional steel. The formation of corrosion on the exterior of this alloy leads to the formation of a protective coating, which effectively inhibits subsequent corrosion.


  • Galvanization refers to the application of a zinc coating onto a metal surface in order to provide protection. This method of prevention of iron corrosion is a widely employed technique.
  • One of two processes is capable of protecting metal:
    • (i) Hot-dip galvanization, a process that involves immersing the metal in molten zinc, (ii) electroplating.
  • Zinc, a cost-effective metallic element, exhibits a high affinity for adhering to steel surfaces. Additionally, it provides cathodic protection to the iron surface through its function as an anode. The corrosion of the zinc layer is attributed to this phenomenon, as opposed to the iron.
  • One drawback associated with galvanization is its limited ability to provide long-term protection against corrosion, as the zinc layer gradually depletes during the process. In environments that exhibit elevated levels of corrosion, the preference for employing cadmium coating arises from its superior efficacy relative to the alternative under consideration.

Cathodic Protection

  • Applying an electric charge to metals can successfully impede the generation of electrochemical reactions that result in the formation of rust.
  • The process involves converting the iron/steel into a cathode by affixing a sacrificial anode to it.
  • The sacrificial anode should possess an electrode potential that is comparatively more negative than that of iron.
  • The metals frequently employed as sacrificial anodes include magnesium, zinc, and aluminum. In order to safeguard the iron/steel, it is imperative to replace them once they have undergone corrosion.


  • There are various coating options available for protecting exposed metal surfaces from corrosion. Some commonly used coatings for corrosion prevention include paints, wax tapes, and varnish.
  • Water-displacing oils are applied to smaller objects in order to coat them and prevent rust formation on the surface of the object. In industrial settings, it is common practice to apply a layer of grease to iron-based machinery and tools. This greasing process serves two important purposes: lubrication and rust prevention. By reducing friction, the grease facilitates the smooth operation of the metal components. Additionally, it acts as a protective barrier, effectively inhibiting the formation of rust on iron surfaces.
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About Author

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Jyoti Bashyal

Jyoti Bashyal, a graduate of the Central Department of Chemistry, is an avid explorer of the molecular realm. Fueled by her fascination with chemical reactions and natural compounds, she navigates her field's complexities with precision and passion. Outside the lab, Jyoti is dedicated to making science accessible to all. She aspires to deepen audiences' understanding of the wonders of various scientific subjects and their impact on the world by sharing them with a wide range of readers through her writing.

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