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The Corrosion College


In every instance of electrochemical corrosion, two reactions always take place: chemical specialists talk of oxidations and reductions, while electrochemical experts talk of anodic and cathodic part reactions. That is always linked to an exchange of electrons.

In the case of oxidation / anodic part reaction, the metal forms the anode with a neighbouring, aqueous electrolyte. Through interactions with charged particles in the electrolyte, the disintegration of the metal now gets induced. As a consequence, the metal releases its electrons and metal disintegration occurs.

Me → Mez+ + z e
Me: Any metal
e: Free electron
z: Number of electrons
z+: The metal’s (ion’s) valence

The oxidation of the metal describes only one of the two part reactions. In order to ensure electrical neutrality, one reaction partner has to absorb the electrons in a linked reduction step. This is the reduction and is described with the cathodic part reaction. An appropriate reaction partner would be dissolved oxygen, which then gets reduced to the hydroxide ion. This reaction is known by the term oxygen corrosion. Where there is a lack of oxygen, water and respectively in acidic environments, oxonium ions (H3O+) lend themselves to the role. In both cases hydrogen is produced through reduction. Experts term this type of corrosion hydrogen corrosion or acid corrosion.

Oxygen corrosion

4 e + O2 + 2 H2O → 4 OH


2 H3O+ + 2 e → H2 + 2 H2O

e: Free electron
O2: Oxygen molecule
H3O+: Oxonium ion
OH: Hydroxide ion
H2: Hydrogen molecule

Whether oxygen or hydrogen corrosion is more prevalent in an aqueous electrolyte depends on several factors. If a lot of oxygen has been dissolved in neutral water, oxygen corrosion is more likely to take place. Where corrosion happens in conditions of little oxygen, it is predominantly hydrogen corrosion that occurs.

For all reactions that take place during electrochemical corrosion it is not absolutely necessary that both partial steps occur in the local, same place. Due to the electrical conductivity of the metals and the electrolyte a certain spatial distance between the anode and cathode could be possible.


Under atmospheric conditions, noble metals such as gold or platinum react inertly to oxygen. Iron and other metals such as aluminium, nickel and zinc do this, by contrast, very quickly. In this process they become oxidised and metal oxides form, such as iron(III) oxide (Fe2O3), aluminium oxide (Al2O3), nickel oxide (NiO) and zinc oxide (ZnO). Such metals, which occur in nature as oxides, are in contrast to the noble metals described as ‘base’.

The reason for this differing reaction to oxygen is the attempt to take up an energetically stable state. In their metallic form noble metals are already in the energetically most favourable state. Base metals, by contrast, only get into their energetically most stable state by releasing electrons.

To what degree a metal is noble or base is described by experts using the term redox potential, thus providing a measure for the tendency to acquire electrons. In the elementary form the metals create a redox pair with the associated ion. The voltage of the potential difference gets measured in comparison to a H2/H+ redox pair that has been defined with 0.00 V. More noble metals have a higher electrochemical potential and base metals a lower one. The potential of gold, for instance, is + 1.40 V and that of iron – 0.41 V.

The electrochemical potentials are summarised in the electrochemical series.


(Standardpotentiale bei 25 °C, 101,3 kPa, pH=0, Ionenaktivitäten = 1)

Ele­ment im Re­dox-Paar,
des­sen Oxi­da­ti­ons­stu­fe
sich än­dert
oxi­dier­te Form+ z e−⇌ re­du­zier­te FormStan­dard­po­ten­ti­al E °
Fluor (F)F2+ 2 e⇌ 2 F+2,87 V
Schwe­fel (O)S2O82−+ 2 e⇌ 2 SO42−+2,00 V
Sau­er­stoff (O)H2O2 + 2 H3O++ 2 e⇌ 4 H2O+1,78 V
Gold (Au)Au++ e⇌ Au+1,69 V
Gold (Au)Au3++ 3 e⇌ Au+1,42 V
Gold (Au)Au2++ 2 e⇌ Au+1,40 V
Chlor (Cl)Cl2+ 2 e⇌ 2 Cl+1,36 V
Sau­er­stoff (O)O2 + 4 H++ 4 e⇌ 2 H2O+1,23 V
Pla­tin (Pt)Pt2++ 2 e⇌ Pt+1,20 V
Brom (Br)Br2+ 2 e⇌ 2 Br+1,07 V
Queck­sil­ber Hg)Hg2++ 2 e⇌ Hg+0,85 V
Sil­ber (Ag)Ag++ e⇌ Ag+0,80 V
Eisen (Fe)Fe3++ e⇌ Fe2++0,77 V
Iod (I)I2+ 2 e⇌ 2 I+0,53 V
Kup­fer (Cu)Cu++ e⇌ Cu+0,52 V
Eisen (Fe)[Fe(CN)6]3−+ e⇌ [Fe(CN)6]4−+0,361 V
Kup­fer (Cu)Cu2++ 2 e⇌ Cu+0,34 V
Kup­fer (Cu)Cu2++ e⇌ Cu++0,16 V
Zinn (Sn)Sn4++ 2 e⇌ Sn2++0,15 V
Was­ser­stoff (H2)2 H++ 2 eH20 V
Eisen (Fe)Fe3++ 3 e⇌ Fe−0,04 V
Blei (Pb)Pb2++ 2 e⇌ Pb−0,13 V
Zinn (Sn)Sn2++ 2 e⇌ Sn−0,14 V
Ni­ckel (Ni)Ni2++ 2 e⇌ Ni−0,23 V
Cad­mi­um (Cd)Cd2++ 2 e⇌ Cd−0,40 V
Eisen (Fe)Fe2++ 2 e⇌ Fe−0,41 V
Schwe­fel (S)S+ 2 e⇌ S2−−0,48 V
Ni­ckel (Ni)NiO2 + 2 H2O+ 2 e⇌ Ni(OH)2 + 2OH−0,49 V
Zink (Zn)Zn2++ 2 e⇌ Zn−0,76 V
Was­ser2 H2O+ 2 e⇌ H2 + 2 OH−0,83 V
Chrom (Cr)Cr2++ 2 e⇌ Cr−0,91 V
Niob (Nb)Nb3++ 3 e⇌ Nb−1,099 V
Va­na­di­um (V)V2++ 2 e⇌ V−1,17 V
Man­gan (Mn)Mn2++ 2 e⇌ Mn−1,18 V
Titan (Ti)Ti3++ 3 e⇌ Ti−1,21 V
Alu­mi­ni­um (Al)Al3++ 3 e⇌ Al−1,66 V
Titan (Ti)Ti2++ 2 e⇌ Ti−1,77 V
Be­ryl­li­um (Be)Be2++ 2 e⇌ Be−1,85 V
Ma­gne­si­um (Mg)Mg2++ 2 e⇌ Mg−2,38 V
Na­tri­um (Na)Na++ e⇌ Na−2,71 V
Cal­ci­um (Ca)Ca2++ 2 e⇌ Ca−2,76 V
Ba­ri­um (Ba)Ba2++ 2 e⇌ Ba−2,90 V
Ka­li­um (K)K++ e⇌ K−2,92 V
Li­thi­um (Li)Li++ e⇌ Li−3,05 V

These differing tendencies of the metals to release electrons are exploited in corrosion protection. One example: At – 0.76 V, zinc has a lower electrochemical potential than iron. Zinc is therefore less noble than iron and thus has a higher tendency to release electrons.

That has consequences: an electric voltage is created between iron and zinc. The less noble zinc assumes the role of the anode, the ‘more noble’ iron that of the cathode. Consequently zinc becomes oxidised and releases electrons. It effectively ‘sacrifices’ itself – until the ‘sacrificial anode’ has completely dissolved. What is created in the process is above all white rust (Zn(OH)2), which, as precipitation, plates the unprotected zinc or iron and thus slows down the further corrosion through a form of passive protection. (Link to 5_2_1).

The resulting potentials are dependent on many factors. For one, on the anodic and cathodic part reactions described. Also, however, on drops in the voltage of the resistors (e.g. electrolytes) inclusive of any passive layers that have formed. Many metals (e.g. aluminium) form compact oxide coats that passively increase the anticorrosive characteristic appreciably. To this extent, a practical electromotive series, which describes the typical fitted metals and alloys, enables a much better estimation of the corrosion behaviour of pairs of materials. Especially in respect of any contact corrosion occurring in the installation of two different metal pairings.