Sciencemadness Discussion Board

Reduction of rust to Fe(0)

WPK0129 - 20-1-2014 at 15:29

I'm looking for a way to reverse rust on some objects I've got around the house (wrench, vintage razor blade) in a non-destructive way - i.e., instead of dissolving away the rust and leaving a pit or divit, reducing it from iron (III) oxide back to elemental iron.

A caveat - my resources are quite limited, so nothing high-temperature, super-low-temperature, or high-pressure. I would ideally like to do this with something easily-obtainable, too.

Things I've considered and ruled out:
-Sodium dithionite ("Iron Out"); it only reduce the iron to iron (II).
-Lithium from batteries in liquid ammonia (can't get the ammonia cold enough to liquify; also, I'd have to distill it from a household 5% solution, which would be a pain).
-Hydrogen (storing it would be dangerous; also, the reaction might require high pressure).

Thanks in advance for any insight.

elementcollector1 - 20-1-2014 at 15:49

Clearly (and I don't mean to insult here) you don't know what you're doing. Reducing iron(III) to iron(0) is necessarily an energetic reaction - avoiding high temperatures is near impossible. What you could do if you wanted to keep roughly the same amount of iron on there is dissolve it off, and then plate it back on again using electrolysis. Personally, I would simply use naval jelly and get the rust off - is there a reason you want to 'keep' this layer of iron?

BromicAcid - 20-1-2014 at 16:24

You can use electrolysis to remove the rust. It does not back-convert it to free iron and plate it back in place though. There are some pretty quick and easy setups that can be made. Just search "Rust Removal Electrolysis" in your favorite search engine and you will find plenty of results. I think the setups in the 5 gallon buckets with rebar are pretty nifty. I tried to find a video I saw sometime back where they setup a 30 gallon tub with multiple pieces of rebar for working on large projects but really you're just limited by the size of your tub.

UnintentionalChaos - 20-1-2014 at 16:28

Quote: Originally posted by elementcollector1  
Clearly (and I don't mean to insult here) you don't know what you're doing.

Furthermore, when iron oxidizes to Fe2O3 or hydrates thereof, the crystal structure must expand. This is why rust flakes and crumbles off of iron/steel instead of forming a hard, nonreactive oxide skin like aluminum does (which would otherwise fall apart far faster than iron. Even if you could turn all the rust directly back into iron, it would not be structurally sound and would be highly prone to rusting rapidly due to high surface area.

WPK0129 - 20-1-2014 at 17:55

BromicAcid and UnintentionalChaos, thank you! You have been most helpful.

elementcollector1, you have not. When combined with a sufficiently facile oxidation half-reaction, the net Gibbs free energy of the whole process would become negative and therefore the process should be feasible. (Indeed, the oxidation of aluminum should be enough at room temperature given the standard reduction potentials involved, but I've been having problems getting this to work practically.) There is no reason why a high temperature should be necessary per se, unless the process would have an extremely high activation energy, which I would have no way of knowing off the top of my head and I don't believe you do either.

As previously stated, I'm trying to keep the steel there simply because I suspect the rust spots to be deep enough to cause pitting on the surface of the metal were I just to dissolve the rust away. This would be at best unsightly (in the case of the vintage razor blade) and at worst make the pliers (sorry, said wrench before) less functional by stripping away some of the surface of the teeth.

elementcollector1 - 20-1-2014 at 18:25

Now lay me out a reaction scheme to reduce the iron on the surface back to iron(0), with commonly available reagents, at room temperature.
There's more than Gibbs free energy to consider here - there are mechanical specifications to such redox reactions. For instance, a mixture of iron oxide and aluminum does not react to iron and aluminum oxide at room temperature because the surface area interaction of the particles is far too low. At thermite temperatures, at least one of the two reactants is molten from the heat of the reaction, and both of the products are - thus promoting a faster reaction. Similarly, iron ions in solution can be reduced by solid aluminum due to the mobility of the ions in solution - lots and lots of surface area there.

As you might imagine, thermite doesn't go off at room temperature (short of nanothermite) for a reason. That reason being, whala! Activation energy. Look it up - thermite is well-known as such for being hard to start without sufficient energy, which is often supplied in the form of a combusting boost mix. I might not know the exact value of such an energy, but I don't need to in this case: Room temperature will not provide it.

And then, there's the problem of the rust itself, as stated above: The structure isn't the same. More specifically, even if you did by some scientific miracle reduce the iron back to the metal at room temperature, it would be brittle, and unusable. Pliers have to be tempered, at least on the surface, to even work as well as they do.

Not only that, not all of your iron will be converted back by definition: Reactions in reality often have what's called a "% yield", a fraction expressed as the amount typically yielded over the theoretical.

In short, there is a whole lot more to consider here than just the analytical chemistry. There's also the metallurgy, feasibility, and reality to consider as well.

AJKOER - 29-1-2014 at 06:22

Well, there is some hope if you use a Hydrogen flame. To quote one source, "Kinetics of reduction of iron oxide with carbon monoxide and hydrogen", at :

"The oxide specimens were reduced in streams of pure hydrogen or carbon monoxide at temperatures between 700° and 1,200°C. The reduction was followed by measuring the sample weight during the reaction. Reduction rates were studied at system pressures of 1 and 2 at.
The samples reduced step by step, to magnetite, to wüstite, and finally to iron. A shell of reduction, clearly visible in sections of partially reduced specimens, moved concentrically into the core of the samples."

Per Wikipedia (see ), the temperature of a Hydrogen and air flame exceeds the required temperature cited above. So experiment with heating at a distance and moving the Hydrogen flame close in on the rusted surfaces.

No promise on whether you will find the results acceptable from an esthetics point of view. My comments relate to the art of chemistry solely. However, if the results really do look weird, you may be able to sell it as art.;)

[Edited on 29-1-2014 by AJKOER]

chornedsnorkack - 31-1-2014 at 11:51

Electrolysis would, in principle, seem the logical mechanism:
1) Some of the rust dissolves in some electrolyte - which is soaking the porous rust itself
2) The Fe cations migrate to an electrode - which is the electrically conductive surface of intact iron underlying rust
3) At the iron surface, iron cations are reduced to metal
4) As the dissolved iron is removed from solution at the metal surface and the concentration of iron in the electrolyte soaking the iron drops below saturation, more rust dissolves

So forth until there is no rust left and no iron in solution either.
Under 100 % yield? Yes. There are competing reactions. Reduction of air oxygen, and reduction of water hydrogen.
But cannot the reaction still go to completion and converting 100 % of rust to metal, only that the expenditure of electricity or of the other electrode metal exceeds the 100 % needed for stoichiometric amount of rust?

chemrox - 31-1-2014 at 13:20

What is the point in statements like " don't know what you're doing.." We ask because we don't know.