sacrificial protection and how it works


by sgstudent
Tags: protection, sacrificial
sgstudent
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#1
Apr2-12, 08:47 AM
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How do sacrificial metals protect the less reactive metal that it protects? Some people say that it provides the iron with electrons when it starts to rust. But if it does provide electron to the iron then the magnesium block for example will become positive ions. But then how can the magnesium block exist as a block of positive ions? Similarly how does galvanisation work since they have the same principles. In the air, there are no anions that can form an ionic compound with the magnesium ions. Do electrons from the more reactive metal even give up electrons to the less reactive metal it is attached to or is there another process that I'm not seeing. Thanks for the help!
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Borek
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Apr2-12, 09:47 AM
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You can think in terms of "electron sea tide" - when you have a piece of metal it is perfectly neutral. When you connect two pieces of two different metals, "electron sea" "shifts" in the direction of the metal that has a higher reduction potential. This shift is very tiny, but it gives a small charge separation - that's where the galvanic potential comes from. While the shift is tiny, it is high enough to have an observable effect - metal that has slightly less electrons will be easier to oxidize (less electrons to remove).

That's probably the most hand wavy explanation you will ever see.
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Apr2-12, 10:08 AM
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Quote Quote by Borek View Post
You can think in terms of "electron sea tide" - when you have a piece of metal it is perfectly neutral. When you connect two pieces of two different metals, "electron sea" "shifts" in the direction of the metal that has a higher reduction potential. This shift is very tiny, but it gives a small charge separation - that's where the galvanic potential comes from. While the shift is tiny, it is high enough to have an observable effect - metal that has slightly less electrons will be easier to oxidize (less electrons to remove).

That's probably the most hand wavy explanation you will ever see.
I don't quite understand the explanation this is what I picture in my mind which is probably wrong ad that's where the loophole in the theory lies: electrons from Mg flows into Fe reducing the Fe2+ ions. So the Mg forms Mg2+ ions. However, I don't really believe this is the case as there is no electrolyte present to transfer the metal cations like in a simple cell. With the explanation you gave this is what I picture: the electrons shift towards the less reactive metal without reducing any metal ions or anything like that, all it does is to make it easier to get oxidised by the regulars means 2Mg+O2-->2MgO.

Then again, the wiki pages state that an electrolyte must be present for the sacrificial protection to work. So is there two different processes, one is the simple cell case where a electrolyte is present and it works like a simple cell. The other is where there is no electrolyte like the mentioned one. Is this correct?

I think my concept is really wrong and I'm missing something important. I just don't get how galvanized iron like a iron bin can protect it from getting rusted since there is no electrolyte in the solution. Please help me clear out these pesky confusions. Thanks for the help, Borek

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Apr2-12, 11:03 AM
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sacrificial protection and how it works


Quote Quote by sgstudent View Post
electrons from Mg flows into Fe reducing the Fe2+ ions. So the Mg forms Mg2+ ions.
I would say there are no Fe2+ ions to be reduced. Fe is not oxidized, as there is a sacrificial anode present.

However, I don't really believe this is the case as there is no electrolyte present to transfer the metal cations like in a simple cell.
In typical situation both metals are either submerged in water, or put into soil, so there is an electrolyte. When there is no electrolyte, both protected metal and sacrificial anode corrode on its own, but IMHO kinetics of the corrosion should be changed when compared with the pure and separated metal samples.

With the explanation you gave this is what I picture: the electrons shift towards the less reactive metal without reducing any metal ions or anything like that, all it does is to make it easier to get oxidised by the regulars means 2Mg+O2-->2MgO.
Yes.

Note that oxygen doesn't have to be reduced in the same place magnesium is getting oxidized (that's what you will observe in the case with electrolyte).

Then again, the wiki pages state that an electrolyte must be present for the sacrificial protection to work. So is there two different processes, one is the simple cell case where a electrolyte is present and it works like a simple cell. The other is where there is no electrolyte like the mentioned one. Is this correct?
There are many ways of skinning that cat - I don't see anything wrong with this approach.
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Apr3-12, 01:02 AM
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Quote Quote by Borek View Post
I would say there are no Fe2+ ions to be reduced. Fe is not oxidized, as there is a sacrificial anode present.



In typical situation both metals are either submerged in water, or put into soil, so there is an electrolyte. When there is no electrolyte, both protected metal and sacrificial anode corrode on its own, but IMHO kinetics of the corrosion should be changed when compared with the pure and separated metal samples.



Yes.

Note that oxygen doesn't have to be reduced in the same place magnesium is getting oxidized (that's what you will observe in the case with electrolyte).



There are many ways of skinning that cat - I don't see anything wrong with this approach.
Thanks Borek! However, the part I'm confused with is the electrolyte part. when using the sacrificial anode a must to have an electrolyte to allow the current to flow from one metal to the other. However, in the wiki link they gave an example where the iron bolts preferentialy rusted in place of the copper structure. also, in a textbook example they galvanized a iron trashcan to protect it from corroding. they said that it would protect the iron even if there was a large crack exposing a large portion of the iron. Again in this case there is no electrolyte so I'm quite confused here. so actually do they need the electrolyte for it be successfully protected?
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Apr3-12, 02:13 AM
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When the objects are dry, there is no electrolyte, but they are also (almost) not corroding. However, air contains water vapor, so any object in contact with air can be moist - and occasionally is (think temperature changes and condensation). That's enough.
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Apr3-12, 03:37 AM
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Quote Quote by Borek View Post
When the objects are dry, there is no electrolyte, but they are also (almost) not corroding. However, air contains water vapor, so any object in contact with air can be moist - and occasionally is (think temperature changes and condensation). That's enough.
But even when it is dry I thought that the iron can be oxidised by oxygen in the air? When there is a electrolyte present electrons will flow into iron like a simple cell so it cannot get oxidised at all as it has the excess electrons. However when there is no electrolyte then how can it protect the iron? Or is the limitation of this method. So this method is more effective when an electrolyte is present as it can totally stop oxidation?

Also, when the water condenses then won't it be pure water, so why is it still an electrolyte? Or the H+ and OH- is enough to allow the flow of electrons since the current flow will not be very high?
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Apr3-12, 04:05 AM
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Quote Quote by sgstudent View Post
Will the condensed water contain the ions for the sacrificial anode to transfer electrons?
Condensate is never perfectly pure - there is always some dirt on the surface. Air contains carbon dioxide, so condensate is always saturated with it - that usually means pH below 6. Plus, it is always possible there are some tiny amounts of corrosion products on the surface - to some extent they are always soluble.

Plenty of ions.
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Apr3-12, 05:17 AM
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Quote Quote by Borek View Post
Condensate is never perfectly pure - there is always some dirt on the surface. Air contains carbon dioxide, so condensate is always saturated with it - that usually means pH below 6. Plus, it is always possible there are some tiny amounts of corrosion products on the surface - to some extent they are always soluble.

Plenty of ions.
Oh okay that makes sense. Then what about what happens if both metals are dry? Will the more reactive metal have any effect in slowing the corrosion process of the less reactive metal? If so why does it happen? Thanks Borek!
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Apr3-12, 05:48 AM
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Intuition tells me that speed of corrosion would be affected in such situation, but by how much I don't know. Remember the electron cloud shift?
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Apr3-12, 07:11 AM
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Quote Quote by Borek View Post
Intuition tells me that speed of corrosion would be affected in such situation, but by how much I don't know. Remember the electron cloud shift?
Oh, but in exams when they ask to explain the sacrifcial protection in either galvanization or the sacrificial anode, how should I explain it since there is a possibility that the sacrificial anode and metal to be protected is dry. Or do I assume that there is a electrolyte present? Then what if they explicitly state that it is dry? Thanks for the help Borek!
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Apr3-12, 07:25 AM
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Assume there is electrolyte present, when they state it is dry answer it will not corrode at all.
sgstudent
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Apr3-12, 08:11 AM
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Quote Quote by Borek View Post
Assume there is electrolyte present, when they state it is dry answer it will not corrode at all.
Do you mean that the iron will not corrode st all even when stated to be dry? Sorry cos I don't understand which metal you are referring to. Thanks for the help!
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#14
Apr3-12, 10:59 AM
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Scratch it. It will corrode as usual.


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