How does metal oxidize?

corrosion

Corrosion is the destruction of metals through oxidation of these metals. The best known and most economically damaging corrosion is iron rusting. Iron reacts with oxygen and water to form various iron oxides and hydroxides. We will now take a closer look at these reactions.

First, iron is oxidized to iron (II) and reacts with hydroxide ions to form the poorly soluble iron (II) hydroxide.

\ begin {array} {llcll}
\ text {Ox .:} & {Fe} & \ quad \ rightarrow & \ quad \ {Fe ^ {2+}} \ + \ {2e ^ {-}} & \ cdot 2 \
\ text {Red .:} & \ {O_2} \ + \ {2H_2O} \ + \ {4e ^ {-}} \ & \ quad \ rightarrow & \ quad {4OH ^ {-}} \ & \
\ hline
\ text {Redox:} & \ {2Fe ^ {2 +}} \ + \ {O_2} \ + \ {2H_2O} & \ quad \ rightarrow \ quad & {2Fe ^ {2+}} \ + \ {4OH ^ {-}} & \
\ end {array}

\ begin {align *}
\ Rightarrow \ {2Fe ^ {2+}} \ + \ {4OH ^ {-}} \ quad \ rightarrow \ quad {2Fe (OH) 2}
\ end {align *}

The iron (II) oxidizes further to iron (III), which reacts again with hydroxide ions to iron (III) hydroxide - a yellow-brown solid.

\ begin {array} {llcll}
\ text {Ox .:} & {Fe ^ {2+}} & \ quad \ rightarrow & \ quad \ {Fe ^ {3+}} \ + \ {e ^ {-}} & \ cdot 4 \
\ text {Red .:} & \ {O_2} \ + \ {2H_2O} \ + \ {4e ^ {-}} \ & \ quad \ rightarrow & \ quad {4OH ^ {-}} \ & \
\ hline
\ text {Redox:} & \ {4Fe ^ {2 +}} \ + \ {O_2} \ + \ {2H_2O} & \ quad \ rightarrow \ quad & {4Fe ^ {3+}} \ + \ {4OH ^ {-}} & \
\ end {array}

\ begin {array} {cl}
\ Rightarrow & \ {Fe ^ {3+}} \ + \ {3OH ^ {-}} \ quad \ rightarrow & \ quad \ {Fe (OH) 3} \
\ Rightarrow & \ {Fe (OH) 3} \ quad \ rightarrow & \ quad {FeO (OH)} \ + \ {H_2O}
\ end {array}

The iron (III) hydroxide then reacts further to form the brown iron (III) oxide hydroxide by splitting off water. Ultimately, there is a mixture of iron (II) hydroxide, iron (III) hydroxide, iron (III) oxide hydroxide and water. The volume of rust is larger than that of the original iron, which is why the rust layer on the iron easily flakes off.

There are two different options for preventing iron from rusting: passive and active corrosion protection.

Passive corrosion protection

We can prevent a metal from corroding by ensuring that the metal does not come into contact with the reactants responsible for rusting. To do this, we cover the metal with a protective layer made of a corrosion-resistant material. This type of corrosion protection is called passive because the protective layer does not react, i.e. it is passive.

However, it becomes problematic when the protective layer is damaged. In this case, a so-called contact element is created. The protective layer, a noble metal, is (electrically conductive) connected to the iron and the damage causes water to reach both metals. In this case, water acts as an electrolyte and we have a kind of galvanic cell. On the more noble metal, water picks up electrons and hydrogen is formed. Electrons are given off at the less noble metal and thus the metal is broken down. In this case the iron will rust. However, the rusting process is accelerated by the contact element and the iron is damaged more quickly.

Active corrosion protection

In contrast to passive corrosion protection, there is also the option of active corrosion protection. Here the iron is coated with a less noble metal. First of all, it acts as a passive protective layer. The advantage here is that the iron does not rust even if the protective layer is damaged. A contact element is also created here, but with a contact element the less noble element is always degraded, which in this case is the less noble protective layer. Since the protective layer now reacts under corrosion, it is a question of one
active corrosion protection.