Some metals rust more quickly than others due to their chemical reactivity with oxygen and air moisture. For example, iron rusts more quickly than aluminum because it is more reactive and therefore more prone to oxidation.
Each metal has a different intrinsic chemical reactivity, meaning its natural tendency to react with oxygen and moisture. Some metals, like iron, are naturally very reactive and oxidize quickly, which causes them to rust rather rapidly under normal conditions. In contrast, others like gold or platinum are much more chemically stable, which explains why they remain shiny even after a long time. This intrinsic reactivity mainly depends on how easily a metal releases its electrons to form compounds with other elements, particularly oxygen. The more easily a metal donates its electrons, the faster it rusts.
Humidity is one of the main accomplices of rust: the more humid the air, the faster corrosion occurs. If you add salt to the equation, like that found near the sea, rust attacks even more fiercely and quickly. This acceleration is explained by the fact that salt facilitates electrical and chemical exchanges. Temperature also plays a role: moderate heat often promotes corrosion, as it speeds up chemical reactions. Conversely, extreme cold slows down these reactions. Finally, direct exposure to rain or pollutants present in urban air accelerates the formation of rust on many metals.
Common metals, even when seemingly pure, often contain impurities in small quantities that seriously change the game. Some impurities accelerate corrosion by creating mini-zones where water and air react more quickly, acting as little rust accelerators; others, on the contrary, can slow down the phenomenon. When elements are deliberately mixed to create alloys, the results can be surprising: for example, stainless steel contains chrome, which forms a thin protective invisible layer that prevents rust from progressing easily. Conversely, the addition of different metals in an alloy can sometimes cause internal electrochemical reactions, facilitating corrosion. These small chemical details determine why some metals hold up better than others against rust.
Some metals naturally protect themselves from corrosion thanks to an oxide layer that forms spontaneously on their surface. The best example? Aluminum. Just exposed to air, it quickly produces a thin, transparent, ultra-resistant layer of aluminum oxide that blocks moisture and oxygen, thus preventing deeper corrosion. The same goes for chrome in stainless steel: it forms an invisible protective layer on the surface, effectively preventing rust. In contrast, iron behaves differently: the oxide layer it forms (classic rust, orange-brown) is porous, permeable, and easily detaches, always leaving the underlying metal exposed. As a result, rust after rust, iron ends up corroded more quickly. Everything depends on the strength and stability of this natural protective layer.
Some specific treatments directly modify the rate at which a metal rusts. For example, galvanization, which involves coating a metal with a thin layer of zinc, protects steel by sacrificing the zinc instead of the iron, significantly slowing down corrosion. Conversely, if you use certain poorly adapted heat treatments, it can locally weaken the structure of the metal and make it even more vulnerable to rust. Certain operations like chemical stripping should be followed by quick protection; otherwise, the bare metal exposed to air will corrode rapidly. The same goes for certain mechanical treatments: for instance, a carefully polished metal has fewer surface irregularities, potentially reducing the attack areas for corrosion.
Did you know that corrosion can be accelerated by the presence of salt? That's why vehicles driving near the sea or exposed to the salt used on roads in winter rust more quickly.
Rust (hydrated iron oxide) occupies a greater volume than the initial iron. This is why rust tends to lift paint or coatings and accelerates the degradation of metal structures.
The formation of rust can create a spontaneous electrochemical cell: two slightly different areas on the surface of the metal form a mini-battery, thereby accelerating the rate of corrosion under certain conditions.
Unlike iron, metals such as aluminum and chromium spontaneously form a thin protective layer of oxide, thus preventing deeper corrosion. This is why aluminum appears to be much more resistant to rust.
Yes, several metals never rust: for example, gold, platinum, and silver. These metals have a very low intrinsic chemical reactivity, which prevents the corrosion phenomenon when in contact with air or water.
Unlike iron, aluminum quickly forms a compact oxide layer on its surface, which effectively blocks oxygen from accessing the interior of the metal and thus prevents further oxidation. This layer acts as a natural protective shield, something that iron does not do easily, leading to its rapid corrosion.
Rust itself is not toxic, but it can cause other problems. It weakens metal structures, poses a risk of serious mechanical issues, facilitates bacterial contamination, or causes injuries (for example, through rusty tools, which can increase the risk of infections).
Humidity, in the presence of oxygen, facilitates the chemical oxidation reaction responsible for rust. The higher the humidity, the more water particles increase the rate of oxidation by ensuring better availability of the ions needed for the chemical reaction.
Several methods exist, such as galvanization, the application of anti-corrosion paints or varnishes, protective coatings (for example, phosphating treatment), or the use of sacrificial anodes made of zinc or magnesium, which help to slow down or prevent rust.
Stainless steel contains chromium or nickel which, when exposed to air, forms a very thin but highly resistant protective layer of chromium oxide. This layer protects the steel from attacks by oxygen and moisture, effectively preserving it from rust.
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