Rust, Part 1

An excerpt from The Perfect Edge, The Ultimate Guide to Sharpening for Woodworkers:

Rust

Any discussion about steel is incomplete without mentioning rust. Rust is the product of the oxidation of iron. Iron and oxygen are eager to combine and make rust but for this to occur, iron must be in contact with water and oxygen. Air contains both of these depending on the relative humidity. Water in the air is readily absorbed by a speck of dust on the surface of one of your tools to form a droplet. That tiny water droplet on the iron surface is all it takes to provide the electrolyte necessary to allow oxygen to combine with the iron and water creating an iron hydroxide molecule (Fe(OH)x). Additional oxygen in the water combines with the iron hydroxide to form hydrated iron oxide (Fe2O3.H2O), which we know so well as brown rust: a porous, absorbent coating that encourages yet more rust. The rust molecule is physically larger than iron alone so rust literally grows as the oxidation continues deeper into the iron. The larger iron hydroxide and iron oxide molecules push each other out of the way causing scales and flakes of rust to detach from the iron’s surface. Under the right conditions, at least as far as rust is concerned, the surface coating of brown rust will continue until all the iron is converted to its oxide and there is nothing left but a pile of rust. Musician Neil Young got it right: “Rust Never Sleeps.”

Iron is so reactive with oxygen that the only samples of pure metallic iron available on Earth are nodules buried so deep in the crust that they have yet to encounter any of oxygen’s efforts. It has been said that for every pound of iron or steel produced in a year, one-quarter pound of previously produced iron or steel is lost to rust! The National Association of Corrosion Engineers (NACE) did a study in 2002 to determine what corrosion costs the United States and estimates it to be a staggering $276 billion dollars each year. That accounts for 3.1% of our Gross Domestic Product and comes down to $970 per person in direct costs. If you include indirect costs of corrosion such as lost productivity due to failures, outages, delays, and litigation the per-person tab would be about twice that amount. So those pesky oxygen atoms, the same ones that steal the carbon from our tool steels during heat treatment, are also hell-bent on ruining our tools after hardening through a direct assault on the very iron these tools are made from.

The Chemistry of Rust

The corrosion of iron into rust is a much more complex process than it appears. Seems simple to me: iron atom gets attacked by oxygen atom and together form a new compound, rust. Yet, it’s just not that straightforward. Iron atoms are exceptionally willing to give up an electron or two at the drop of a hat. Or, rather to a drop of water. The iron atom is now positively charged (giving up electrons will do that) and will quite agreeably bond with other atoms that are negatively charged. Water absorbs oxygen from the air and our droplet has some extra electrons (from the iron atoms) that produce hydroxyl ions in the water. These negatively charged hydroxyl ions combine with the positively charged iron ions and form iron hydroxide. Because water dissolves oxygen, there is usually an excess of oxygen available to combine with the iron hydroxide to form hydrated iron oxide: red rust. I’m not qualified to explain all the chemical reactions that occur during the creation of rust so I’m grateful to
corrosion engineer Katherine Cockey for this contribution: Rust is the corrosion product of an electrochemical action. Electric potential  differences exist over the surface of the steel leading to numerous galvanic cells i.e. many minute batteries. In this case, the process begins with the transfer of electrons from iron to oxygen. The cell electrode where the loss of electrons occurs through an oxidation reaction is called the anode. Iron gives up electrons to the water:

Fe → Fe2+ + 2e-

The other electrode in the cell is called the cathode where a reduction reaction occurs. Oxygen gains the electrons and hydroxide ions are formed

O2 + 4e- + 2H2O → 4OH

Crucial to the formation of rust is the accompanying reduction/oxidation (redox) reaction between iron and oxygen in the presence of the water.

4 Fe2+ + O2 → 4 Fe3+ + 2O2-
Just like cholesterol, iron forms “good” and “bad” oxides. The ferrous (Fe2+) oxide adheres to the iron surface to form a protective layer – good. The ferric (Fe3+) oxide tends not to adhere and to flake off – bad. In the iron-oxygen-water microcosm the corrosion product is governed by the availability of oxygen and water. Limit the dissolved oxygen and this favors the ferrous (FeO) oxide in the following balance:

Fe2+ + 2 H2O ↔ Fe(OH)2 + 2HFe(OH)2 ↔ FeO + H2O

Increase the oxygen concentration and the ferric (Fe2O3) oxide wins out:

Fe3+ + 3 H2O ↔ Fe(OH)3 + 3HFe(OH)3 ↔ FeO(OH) + H2O2FeO(OH) ↔ Fe2O3 + H2O

The world is still balanced but the result is undesirable red rust and more the norm. The point to all this anode/cathode talk is that the
rust that grows on the surface may not be at the same location as the iron atoms that gave up the electrons that started the whole reaction. That’s how rust can spread under paint or plating. A pinhole in the coating can allow water to contact the metal and rust will spread  from there. Sneaky, rust.

Next week: Fighting the War on Rust

Order The Perfect Edge from our website and we’ll include a Blade Bucks coupon good for $10 off your next purchase from HOCK TOOLS.

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About Ron Hock

Owner of HOCK TOOLS (.com) and author of "The Perfect Edge, the Ultimate Guide to Sharpening for Woodworkers"
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