Iron, symbolized as \(\text{Fe}\), is a transition metal known for its diverse and vigorous chemical behavior. It commonly exists in multiple oxidation states, primarily iron(II) (\(\text{Fe}^{2+}\)) and iron(III) (\(\text{Fe}^{3+}\)). Iron is highly reactive compared to noble metals because it has a strong tendency to lose electrons. This inherent chemical instability drives it to readily combine with other elements, seeking a more stable state through oxidation.
Reaction with Oxygen and Water (The Process of Oxidation)
The most familiar chemical reaction involving iron is the process of rusting, which is a specific type of corrosion that requires the simultaneous presence of both oxygen and water. Iron metal first loses electrons, oxidizing to form iron(II) ions (\(\text{Fe}^{2+}\)) at anodic sites on the metal surface. These released electrons then travel through the iron to cathodic sites, where they are consumed by dissolved oxygen and water to form hydroxide ions (\(\text{OH}^-\)).
The resulting iron(II) ions move through the water droplet and are further oxidized by more oxygen. They ultimately react with the hydroxide ions to form hydrated iron(III) oxide (\(\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}\)), known chemically as rust. Rust is a porous, flaky, reddish-brown substance that does not adhere tightly to the metal’s surface. Unlike the protective oxide layers formed on metals like aluminum, rust continuously flakes away, exposing fresh iron and allowing corrosion to proceed unchecked.
Reactions with Strong Acids
Iron is positioned relatively high on the metal reactivity series, meaning it readily reacts with strong, non-oxidizing acids like hydrochloric acid (\(\text{HCl}\)) and dilute sulfuric acid (\(\text{H}_2\text{SO}_4\)). In these reactions, the iron metal displaces hydrogen from the acid. This results in the formation of an iron(II) salt and the liberation of hydrogen gas (\(\text{H}_2\)).
This reactivity contrasts sharply with iron’s behavior when exposed to concentrated nitric acid (\(\text{HNO}_3\)). Instead of rapidly dissolving, the iron undergoes a process called passivation. The concentrated acid immediately oxidizes the iron surface to form an extremely thin, dense layer of iron oxide (\(\text{Fe}_2\text{O}_3\) or \(\text{Fe}_3\text{O}_4\)). This protective oxide film acts as an impenetrable barrier, halting the reaction and rendering the iron unreactive toward the concentrated acid.
Reactions with Other Non-Metals
Beyond oxygen and the hydrogen in acids, iron reacts vigorously with various other non-metals. Halogens, such as chlorine gas (\(\text{Cl}_2\)), react directly with heated iron metal in a highly exothermic process. This reaction produces iron(III) chloride (\(\text{FeCl}_3\)). Iron(III) halides are typically formed because halogens are strong oxidizing agents.
Similarly, when iron filings and elemental sulfur are heated together, a rapid reaction occurs that generates heat and light. This process forms iron(II) sulfide (\(\text{FeS}\)), a black solid. This demonstrates iron’s ability to form stable binary compounds with non-metallic elements.
Environmental Factors that Accelerate Iron Reactions
The rate at which iron reacts with its environment is significantly influenced by several kinetic factors. Increasing the temperature provides the necessary activation energy for chemical reactions to proceed more quickly, thereby speeding up the rate of corrosion. High humidity levels are also detrimental because water is a necessary component for the electrochemical process of rusting. The presence of electrolytes, such as dissolved salts or acids, dramatically increases the conductivity of the water layer on the iron surface, leading to faster rusting. Furthermore, the surface area of the iron affects the reaction speed, as smaller particles expose more atoms to the corrosive medium.