Chemistry: Corrosion & Prevention of Corrosion

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Iron is the most abundant metal on our planet, but aside from in the middle of meteorites, we don't find any iron in its elemental form. And, of course, you know the reason why, because iron rusts, just like this gutter pipe here that I am showing. We all are familiar with what iron does when it's exposed to air. And let's go ahead and take a little time and look now at the chemistry behind corrosion, specifically, corrosion of iron. Of course, the same idea would apply to other metals, but it's so significant in the case of iron that it's worth looking into in a little bit more detail. And also, we'll talk a little bit about what we can do to prevent it.
Okay, well even in absence of oxygen - like, let's suppose you're down at the bottom of a real deep lake and there's no reasonable amount of oxygen down there - iron is enough of a reducing metal that it can be oxidized just by water alone. And so under anaerobic conditions, where there's no oxygen, we have iron going to iron 2 plus plus two electrons - that half-reaction - and the other half-reaction of water taking those two electrons to become hydrogen and hydroxide. Those two half-reactions, that's a spontaneous reaction. Although it's slow, this is going to be sufficient to completely consume iron, given enough time at the bottom of a lake or something like that.
Now it's much, much worse, however, when you're in the presence of oxygen, because now we have a much more powerful oxidant than what, effectively, is H^+ up here, in oxygen. So the new half-reaction that we have worry about is oxygen combining - taking 4 electrons and protons from the water - to make hydroxide. And again, we still have the iron going to iron 2. So these two half-reactions are happening. You'll note, we're reducing the oxygen; we're oxidizing the iron. And then, the hydroxide and the iron 2 plus come together to make this insoluble iron hydroxide. That, in turn, is readily oxidized by more oxygen to make, ultimately, what we know as rust, this brown material that you're all familiar with, which has a rather undefined stoichiometry to it, but we'll approximate that, at least, as Fe[2]O[3] dot water. The reason I say that is that in the actual solid structure of rust, there are varying amounts of water incorporated into the solid of rust. So it's sometimes written in this more expanded form.
But the bottom-line. Let's look at a cartoon of what that would look like. Effectively, what we have is both the cathode and the anode are right at the surface of the metal. And, in particular, the metal acts as the conductor between the cathode and the anode. And so it's a little galvanic cell all rolled up into one nice little package. So suppose this is the surface of your metal and this represents just a raindrop, or something. And even, we could imagine, that this was painted, except for a small opening here where we have exposure to the iron. Well, what's going to happen is at the anode, which would just be where the iron is going to be oxidized - remember, oxidation occurs at the anode always - and so iron metal is converted to iron 2 plus, plus two electrons and the iron. So it's consumed, so we see pitting happening here. And at the cathode, which would just be the surface of the water, oxygen comes in. It's reduced to form hydroxide. Again, the metal acts as the conductor for those electrons. So the hydroxide that's formed then precipitates with the iron 2 plus, like I mentioned, forming iron 2 hydroxide, which is oxidized, in turn, to iron 3. Hydroxide, ultimately, continues on to form rust.
So, we're used to seeing, actually, the buildup of rust on the surface, but the damage is not actually here. The damage is in the pitting, which is actually what's, of course, removing that iron metal. So very often, you'll see the rust resulting, but you won't see the damage having to do with the pitting. And in particular, if you're talking about a bridge, or something like that, there might be surface rust that you scrape off or that falls off, but you really don't see where the real damage is which is in the more protected areas, even areas that are not exposed to oxygen at all. And see, that's the problem, that oxygen only has to get to the cathode here, but it's the anode where you're losing the iron metal, and that's where you lose your structure and bridges fall down and so on, because the iron, in areas that you can't see, is missing. And so you don't have, of course, the structural support that you think you do.
So that's the basic problem. So overall, this is the reaction that we have to worry about. This corrosion is speeded up by salt. Why would it be speeded up by salt? Well, anything that we can do to increase the transfer of charge. So salt is going to act like an electrolyte. You know, I'm from the east coast - not originally, but the last ten years, I've been out there - and the roads out there are salted and cars don't last nearly as long out there as they used to out in California, where I'm originally from, because of the salt facilitating this rusting. But also, things like acid rain are a big problem, because that H^+, formed from the acid rain, whether it's sulfur trioxide reacting with water to give sulfuring acid, or we get nitric acid from acid rain. That H^+ that we form reacts with the hydroxide to make water. And so that's a driving force - and additional driving force - for this reaction.
So those are the types of things that help, but then, of course, the biggest way to prevent rust is to protect the iron, somehow, to block oxygen from being able to get to that surface completely so we don't get any of this pitting. And they way that that can happen - there are two fundamental approaches to that. You can either physically just completely protect the iron. We're familiar with that. The most common example is just paint. Something that's properly painted won't rust and we rely on that. Our cars on the roads are there for that principle, that if they have a nice paint coating then they're not gong to rust and fall apart. And a huge amount of money has been invested into learning how to really protect cars as far as the materials that go into the paint, and so on, to prevent this kind of thing.
So that is the most common approach. And, in general, we could call this passivation - the basic idea of just simply physically blocking oxygen from getting anywhere near the iron. Now, other examples of that would be tin cans - which is an iron can that's coated with a tin coating - a chrome bumper. Again, the same idea is that if you protect it with something else, as long as you don't expose the metal, even though they're in electrical contact with each other, For instance, think of a chrome bumper. Even though the chrome is in electrical contact with the iron, if you do oxidize any of the iron, make any iron 2 plus, that ion has no place to go, because it's not on the surface anywhere, so that's not a favorable reaction and you're fine. It's only if you scratch the bumper - especially around where the bolt holes are - where the chrome is thinnest - that's going to be a problem potentially.
Now, the better approach is, instead of just physically blocking the surface of the iron - and, again, the liability there is if you scratch that surface, whether it's paint or another metal, you're in for trouble - it's to use a sacrificial anode. Now the idea of that is to use some material that you're willing to sacrifice that is in electrical contact with the iron but is more reducing than the iron is. Okay, examples of this would be magnesium or zinc or aluminum. These guys all are more reducing metals than the iron is.
And the idea behind that is - let's look at magnesium as an example. Magnesium has got a oxidation potential of 2.39 volts. Now remember, you're going to look up reduction potentials. So to get the oxidation potential, you'd switch the reaction and change the sign. So magnesium is a very powerful reducing agent. Let's suppose that we have an iron pipe that we don't want to dig up every five years because it rusts out. We need this thing to be underground for a hundred years and not worry about it. Well, what we can do is take a magnesium rod - again, a more reducing metal - have it in electrical contact with the pipe. And the idea is then as oxygen gets into the ground that, because of that electrical contact, any oxygen that does attack the iron, to make any iron 2, that iron 2 plus immediately will be converted back to iron metal, because the magnesium is in electrical contact with it and is a better reducing agent.
So the magnesium gets consumed, and thus the name "sacrificial anode." The magnesium gets oxidized, but the pipe is okay, as long as we don't run out of magnesium, at least. For long stretches of pipe, magnesium rods are used. They're connected to the iron pipe and that prevents the pipe from rusting.
Another application of this is magnesium bars are bolted to the hull of ships, for the same exact reason, to stop the ships from rusting. Slightly different idea, the same principle - a galvanized nail is a nail that is coated with zinc. And zinc is more reducing than the iron. And so there's an additional benefit here. It's just on the surface, but what happens is if the iron is exposed - if the zinc coating is partially removed, and so the iron is exposed - you're still okay because the zinc can reduce the iron - same idea as the magnesium. But zinc also acts as a passivator. It protects the surface of the iron, so it does both. It protects the nail by protecting its surface, and so no oxygen can get to the iron, but if it gets scratched, it's still okay because the zinc will protect the iron through acting as a sacrificial anode. So nails are galvanized. This pipe of steel, it's galvanized. This thing is not galvanized, because it's not made of iron.
So basically, it's a good deal as long as you're aware of the electrochemical problems with iron and you're aware, also, of ways to prevent the reaction of iron with oxygen. We can use iron, because it's a good thing because it's the most plentiful metal that we have. And so, through understanding the chemistry of iron, we're able to protect iron from the oxygen, which is all around us.
Electrochemistry
Corrosion
Corrosion and the Prevention of Corrosion Page [2 of 2]