Astronomy: The Earth and Corrosion

Mars, the Red Planet - perhaps the most distinctive hue in the solar system. Compared to all other planets except Mercury, Mars' appearance is dominated by its soil, which is an alternating deep and light crimson much more colorful than Mercury's pockmarked gray. You have mostly likely heard that Mars' red color is due to "rust", or the oxidizing of metal. This oxidizing of metal by water, oxygen, and other chemical agents is called corrosion, and produces a metal oxide.

Corrosion plays an important part on Earth, as most metals are susceptible to oxidization. Audemars Piguet Royal It is estimated that about one-seventh of all iron production goes to replace the metal lost to corrosion. Rust is apparently a hydrated form of iron(III)oxide, the same compound that accounts for Mars' red color. The formula is approximately Fe₂O₃• \(\dfrac{\text{3}}{\text{2}}\) H₂O; an anhydric form appears as dusty red soil. (Note that this is about halfway between iron(III) hydroxide, Fe(OH)₃ or ½{Fe₂O₃•3H₂O], and anhydrous Fe₂O₃).

Redox reaction

Rusting requires both oxygen and water, and it is usually sped up by acids, strains in the iron, contact with less-active metals, and the presence of rust itself. In addition, observation of a rusted object, such as an iron nail from an old wooden building, shows that rust will deposit in one location (near the head of the nail) while the greatest loss of metallic iron will occur elsewhere (near the point). These facts suggest that the mechanism of rusting involves a galvanic cell. The half-equations involved are

\[\text{2Fe}(S) \rightarrow \text{2Fe}^{2+}(aq) + \text{4}e^-\label{1}\]

\[\text{4}e^- + \text{4H}^+(aq) + \text{O}_2(g)\rightarrow \text{2H}_2\text{O}\label{2}\]

yielding the full reaction:

\[\text{2Fe}(s) + \text{4H}^+(aq) + \text{O}_2(g) \rightarrow \text{2Fe}^{2+}(aq) + \text{2H}_2\text{O}\label{3}\] 

Once Fe²⁺(aq) is formed, it can migrate freely through the aqueous solution to another location on the metal surface. At that point the iron can precipitate:

\[\text{4Fe}(s) + \text{O}_2(g) + \text{7 H}_2\text{O}(l) \rightarrow \text{2Fe}_2\text{O}_3 \bullet \dfrac{3}{2} \text{H}_2\text{O}(s) + \text{8H}^+(aq)\] 

Hydrogen ions liberated by this reaction are then partially consumed by Eq. \(\ref{2}\) . The electrons required for half-equation \(\ref{2}\) are supplied from Eq. \(\ref{1}\) via metallic conduction through the iron or by ionic conduction if the aqueous solution contains a significant concentration of ions. Thus iron rusts faster in contact with salt water than in fresh.

The mechanism proposed in the preceding paragraph implies that some regions of the iron surface become cathodic, i.e., that reduction of oxygen to water occurs there. Other locations are anodic; oxidation of Fe to Fe²⁺ occurs. The chief way in which such regions may be set up depends on restriction of oxygen supply, because oxygen is required for the cathodic reaction shown in Eq. \(\ref{2}\) . In the case of the iron nail, for example, rust forms near the head because more oxygen is available. Most of the loss of metal takes place deep in the wood, however, near the point of the nail. At this location Eq. \(\ref{1}\) but not \(\ref{2}\) can occur.

A similar situation occurs when a drop of moisture adheres to an iron surface (Fig. \(\PageIndex{1}\) ). Pitting occurs near the center of the drop, while hydrated iron(III) oxide deposits near the edge.

Figure \(\PageIndex{1}\) Corrosion of iron under a drop of water. Reaction \(\ref{1}\) occurs in the anodic region, Eq. \(\ref{2}\) in the cathodic region.

A second way in which anodic and cathodic regions may be set up involves the presence of a second metal which has a greater attraction for electrons (is less easily oxidized) than iron. Such a metal can drain off electrons left behind in the iron when Fe²⁺ dissolves. This excess of electrons makes the less-active metal an ideal site for Eq. \(\ref{2}\), and so a cell is set up at the intersection of the metals. Rust may actually coat the surface of the less-active metal while pits form in the iron.

The most important technique for rust prevention is simply to exclude water and oxygen by means of a protective coating. This is the principle behind oiling, greasing, painting, or metal plating of iron. The coating must be complete, however, or rusting may be accelerated by exclusion of oxygen from part of the surface. Mars' almost exclusive iron-oxide composition is only stopped by the polar ice caps, so perhaps receding and freezing water may have locked up oxygen, creating this imbalance. In terms of protective surfaces, the main protection Martian iron has from oxygen is the superficial coating of red rust itself.

A second technique involves bringing the iron object in contact with a more active metal. This is called cathodic protection because the more active metal donates electrons to the iron, strongly inhibiting Eq. \(\ref{1}\). Both cathodic protection and a surface coating are provided by galvanizing, a process in which zinc is plated onto steel electrolytically or by dipping in the molten metal. Like many other metals, zinc is self-protective—it reacts with oxygen and carbon dioxide from air to form an adherent impervious coating of zinc hydroxycarbonate, Zn₂(OH)₂CO₃. Should there be a scratch in the zinc plate, the iron still cannot rust because zinc will be preferentially oxidized. The hydroxycarbonate formed will then cover the opening, preventing further contact of oxygen with the iron or zinc.

Example \(\PageIndex{1}\): Reduction Potentials

Small deposits of gold and platinum were found in Martian meteorites and are believed to be present in trace amounts on the Martian surface. Is cathodic protection of iron possible with these metals?

Standard Reduction Potentials:

Fe²⁺(aq) + 2e− → Fe(s) -0.44,

Au⁺(aq) + e– → Au(s) 1.83

Pt²⁺(aq) + 2e− → Pt(s) 1.188

Solution

Think about what the reduction potentials physically mean in terms of the potential of a metal to oxidize. If a metal has a high standard reduction potential, it is more likely to be reduced, accepting electrons from another metal. If its reduction potential is negative, the metal will instead donate its electrons and act as a reducing agent. Neither gold nor platinum is a better reducing agent than iron, so they will not cathodically protect the iron from rusting.

A third technique applies to situations (such as an automobile radiator) where aqueous solutions are in contact with the iron. Corrosion inhibitors include chromate salts and organic compounds such as tribntylamine, (C₄H₉)₃N. Chromates apparently form an impervious coating of FeCrO₄(s) as soon as any iron is oxidized to iron(II). Tributylamine, a derivative of ammonia, reacts with organic acids formed by decomposition of antifreeze at the high temperatures of an automobile engine. The tributylammonium salts produced are insoluble and coat the inside of the cooling system. Thus tributylamine neutralizes acid which would accelerate corrosion and provides a surface coating as well.

Few organic compounds are present on Mars, but chromium is present in amounts comparable to gold and platinum. Water is a glaring absence, but photos from NASA^[1] show that liquid water exists underground and infrequently seeps to the surface. Some scientists hypothesize that Mars' surface may have been partially covered in water in millennia past, dissolving the now anhydrous salts on the Martian surface.

From CoreChem: 22.6: Corrosion



References

1. ↑ http://www.nasa.gov/mission_pages/mars/news/mgs-20061206.html

Contributors and Attributions

• Ed Vitz (Kutztown University), John W. Moore (UW-Madison), Justin Shorb (Hope College), Xavier Prat-Resina (University of Minnesota Rochester), Tim Wendorff, and Adam Hahn.