The Detroit Post
Tuesday, 19 October, 2021

Can Rust Form Underwater

James Smith
• Friday, 13 November, 2020
• 10 min read

Only when super heated steam is based through incandescent (so hot it glows) iron tubes do you get a measurable rate of reaction. Therefore, there is no chance of iron metal reacting with liquid water.

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At liquid temperatures up to 100C, oxygen MUST be present to cause iron to rust. I'm afraid Dimwit's idea for the oxidation of iron in water without oxygen is all wet.

Given sufficient time, any iron mass, in the presence of water and oxygen, could eventually convert entirely to rust. Surface rust is commonly flaky and friable, and provides no passivation protection to the underlying iron, unlike the formation of patina on copper surfaces.

Rusting is the common term for corrosion of elemental iron and its alloys such as steel. Many other metals undergo similar corrosion, but the resulting oxides are not commonly called rust “.

Other forms of rust include the result of reactions between iron and chloride in an environment deprived of oxygen. Rebar used in underwater concrete pillars, which generates green rust, is an example.

Rapid oxidation occurs when heated steel is exposed to air Rust is a general name for a complex of oxides and hydroxides of iron, which occur when iron or some alloys that contain iron are exposed to oxygen and moisture for a long period of time. Over time, the oxygen combines with the metal forming new compounds collectively called rust.

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Although rust may generally be termed as “oxidation”, that term is much more general and describes a vast number of processes involving the loss of electrons or increased oxidation state, as part of a reaction. Many other oxidation reactions exist which do not involve iron or produce rust.

Iron or steel structures might appear to be solid, but water molecules can penetrate the microscopic pits and cracks in any exposed metal. The hydrogen atoms present in water molecules can combine with other elements to form acids, which will eventually cause more metal to be exposed.

If chloride ions are present, as is the case with saltwater, the corrosion is likely to occur more quickly. As the atoms combine, they weaken the metal, making the structure brittle and crumbly.

Iron metal is relatively unaffected by pure water or by dry oxygen. The conversion of the passivating ferrous oxide layer to rust results from the combined action of two agents, usually oxygen and water.

Under these corrosive conditions, iron hydroxide species are formed. Unlike ferrous oxides, the hydroxides do not adhere to the bulk metal.


As they form and flake off from the surface, fresh iron is exposed, and the corrosion process continues until either all the iron is consumed or all of the oxygen, water, carbon dioxide, or sulfur dioxide in the system are removed or consumed. When iron rusts, the oxides take up more volume than the original metal; this expansion can generate enormous forces, damaging structures made with iron.

O 2 + 4 e + 2 H2O 4 OH Because it forms hydroxide ions, this process is strongly affected by the presence of acid. Likewise, the corrosion of most metals by oxygen is accelerated at low pH.

Providing the electrons for the above reaction is the oxidation of iron that may be described as follows: With limited dissolved oxygen, iron(II)-containing materials are favored, including Few and black lodestone or magnetite (Fe 3 O 4).

High oxygen concentrations favor ferric materials with the nominal formulae Fe(OH) 3 x O x 2. The nature of rust changes with time, reflecting the slow rates of the reactions of solids.

Furthermore, these complex processes are affected by the presence of other ions, such as Ca 2+, which serve as electrolytes which accelerate rust formation, or combine with the hydroxides and oxides of iron to precipitate a variety of Ca, Fe, O, OH species. The onset of rusting can also be detected in the laboratory with the use of ferry indicator solution.

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Cor-Ten is a special iron alloy that rusts, but still retains its structural integrityBecause of the widespread use and importance of iron and steel products, the prevention or slowing of rust is the basis of major economic activities in a number of specialized technologies. A brief overview of methods is presented here; for detailed coverage, see the cross-referenced articles.

Interior rusts in old galvanized iron water pipes can result in brown and black waterGalvanization consists of an application on the object to be protected of a layer of metallic zinc by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap, adheres well to steel, and provides cathodic protection to the steel surface in case of damage to the zinc layer.

In more corrosive environments (such as salt water), cadmium plating is preferred. Galvanization often fails at seams, holes, and joints where there are gaps in the coating.

In some cases, such as very aggressive environments or long design life, both zinc and a coating are applied to provide enhanced corrosion protection. Typical galvanization of steel products which are to be subjected to normal day-to-day weathering in an outside environment consists of a hot-dipped 85 µm zinc coating.

Cathodic protection is a technique used to inhibit corrosion on buried or immersed structures by supplying an electrical charge that suppresses the electrochemical reaction. The sacrificial anode must be made from something with a more negative electrode potential than the iron or steel, commonly zinc, aluminum, or magnesium.

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The sacrificial anode will eventually corrode away, ceasing its protective action unless it is replaced in a timely manner. Flaking paint, exposing a patch of surface rust on sheet metal Rust formation can be controlled with coatings, such as paint, lacquer, varnish, or wax tapes that isolate the iron from the environment.

As a closely related example, iron bars were used to reinforce stonework of the Parthenon in Athens, Greece, but caused extensive damage by rusting, swelling, and shattering the marble components of the building. When only temporary protection is needed for storage or transport, a thin layer of oil, grease, or a special mixture such as Coastline can be applied to an iron surface.

Such treatments are extensively used when mothballing a steel ship, automobile, or other equipment for long-term storage. Special antiseize lubricant mixtures are available, and are applied to metallic threads and other precision machined surfaces to protect them from rust.

These compounds usually contain grease mixed with copper, zinc, or aluminum powder, and other proprietary ingredients. They are not effective when air circulation disperses them, and brings in fresh oxygen and moisture.

An example of this is the use of silica gel packets to control humidity in equipment shipped by sea. Rust removal from small iron or steel objects by electrolysis can be done in a home workshop using simple materials such as a plastic bucket filled with an electrolyte consisting of washing soda dissolved in tap water, a length of rebar suspended vertically in the solution to act as an anode, another laid across the top of the bucket to act as a support for suspending the object, baling wire to suspend the object in the solution from the horizontal rebar, and a battery charger as a power source in which the positive terminal is clamped to the anode and the negative terminal is clamped to the object to be treated which becomes the cathode.

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The Kinda Bridge in Pennsylvania was blown down by a tornado in 2003, largely because the central base bolts holding the structure to the ground had rusted away, leaving the bridge anchored by gravity alone. It is one of the most common failure modes of reinforced concrete bridges and buildings.

^ Kermit, Bart; Griesser-Stermscheg, Martina; Sewn, Indie; Sutherland, Susanne. “ Rust Never Sleeps: Recognizing Metals and Their Corrosion Products” (PDF).

^ Ramsay, Hosahalli S.; Marlette, Michele; Pastry, Sudhir; Abderrahim, Khalid (2014-02-14). CS1 main: archived copy as title (link) ^ Gupta, Lorraine Mira, Krishnakali.

This turns the point where iron dissolves into an anode, and the region around this area becomes rich in electrons, a cathode. If it was an acid, the $\CE{He+}$ would have quickly taken all the electrons liberated from the dissolution of iron earlier, but the concentration of hydrogen ions is not large enough in water, so that we get another reaction taking place at the cathode .

Now more oxygen will react to oxidize the iron (II) hydroxide... As the $\CE{Fe^2+}$ are consumed, more iron will get dissolved (Le Ch atelier's Principle) and keep the whole process going.

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$\begingroup$The dot means that the compound is a hydrate of the parent chemical; the parent compound forms a weakly-bonded structure with one or more water molecules per parent molecule, because either the hydrogen atoms of the water, or the partial positively-charged region of the water molecule caused by its “bent” molecular shape, is attracted to partially negative regions of the parent molecule. It's not a “full” traditional ionic bond but it works in somewhat of a similar way.

In solid form, this mixture of the parent compound and water molecules forms a crystal structure like ice (it doesn't have to be a well-ordered crystal like you typically get with salts), with properties that differ from a form of the substance that has more, less or no water in its structure. An everyday example of the difference between a dehydrated and hydrated compound is Portland cement, used in concrete.

In its raw form, it's “dehydrated” calcium carbonate along with a few other components, made by heating limestone until the water is released from the rock's structure. The remaining solid forms a fine powder that can 't hold much of a shape.

Add the water back in while mixing concrete, and the water molecules are re-incorporated into a solid structure with the calcium carbonate to essentially reform a rock in the desired shape. Being a polar solvent, it has some affinity for electrons, which are easily given up by the iron as a transition metal, and are attracted to the hydrogen atoms in the water.

If a hydrogen atom successfully “captures” an electron, it balances its own electric charge, and is “liberated” from the bond it has with the oxygen in the water molecule, instead pairing with another liberated hydrogen to form a diatomic gas. That second equation typically occurs “stepwise”, as you'll notice that's quite a lot of electrons and ions floating around at once.

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Various hydrate structures of this iron oxide exist, which result in different colors of the compound from dark brown to deep red to red-orange. You'll notice that the water participates in the overall reaction without really being consumed by it to form the product; it's destroyed and then re-formed in equal amounts, creating intermediates in the process.

Cold from a faucet implying possible oxygen content along with say free $\CE{Cl2}$ or even $\CE{NH2Cl}$, or natural aerated waters containing transition metal salts (including those of $\CE{Fe, Mn}$, some $\CE{Cu}$, …), along with possibly gases other than $\CE{O2}$ like $\CE{N2O}$ from nitrate (present in well water) decomposition, or perhaps boiled distilled water without any dissolved gases or minerals. The composition of the iron alloy, I will ignore and assume pure $\CE{Fe}$, along with a pH range for the $\CE{H2O}$ of between 6 and 8.

However, with the distilled water in open air contact with submerged powdered iron, I will expect a comparatively much longer induction period till any rust is observed per the reactions outlined above. This is due to the underlying electrochemical nature of the corrosion process and no electrolyte (from dissolved salts, but instead awaiting dust particles) and now dissolved $\CE{O2}$ or an acid (source of $\CE{H+}$ could arrive via air containing $\CE{CO2}$).

Iron(III) hydroxide, $\CE{Fe(OH)3}$ then dehydrates to produce $\CE{Fe2O3.nH2O(s)}$ or rust Fe2 O3 (the red oxide) conform in air at ambient temperatures.

The only example I know is fretting corrosion where the characteristic red powder forms around two rubbing steel/iron surfaces. Of course, you can make all sorts of iron oxides at elevated temperatures.

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