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How does Na2Cr2O7 inhibit the corrosion of iron?


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Hi everyone, I was wondering why Na2Cr2O7 inhibits the corrosion of metallic iron and I was discussing this with some friends today but we were only able to make some educated guesses so I've been on the internet for a few hours now searching for an answer but I've had no success so far.

I then remembered a comment my Chemistry teacher told me once about this great science forum, so now I find myself here for the first time with my first post.

If anyone knows or has any idea why Na2Cr2O7 inhibits the corrosion of metallic iron then I'll be interested.

 

Thanks

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http://patft.uspto.gov/netacgi/nph-Parser?Sect1=PTO1&Sect2=HITOFF&d=PALL&p=1&u=%2Fnetahtml%2FPTO%2Fsrchnum.htm&r=1&f=G&l=50&s1=3996058.PN.&OS=PN/3996058&RS=PN/3996058

 

This patent suggests that the dichromate reacts with dissolved oxygen, thereby preventing that oxygen from corroding the steel.

 

http://en.wikipedia.org/wiki/Corrosion_inhibitor

 

This one indicates that the dichromate forms a passivation layer on the steel, protecting it. It is probably condition dependent (such as what the pipe is carrying) for which mechanism is dominant.

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http://patft.uspto.gov/netacgi/nph-Parser?Sect1=PTO1&Sect2=HITOFF&d=PALL&p=1&u=%2Fnetahtml%2FPTO%2Fsrchnum.htm&r=1&f=G&l=50&s1=3996058.PN.&OS=PN/3996058&RS=PN/3996058

 

This patent suggests that the dichromate reacts with dissolved oxygen, thereby preventing that oxygen from corroding the steel.

 

I'm not sure how it would do that, since that's pretty much the fully oxidized state of Cr as far as I know (I could be wrong, but I've never seen a more oxidized state then that). Not to mention that the article mentions a mixture of gypsum and wax, which makes me think a coating is more of the idea.

 

http://en.wikipedia.org/wiki/Corrosion_inhibitor

 

This one indicates that the dichromate forms a passivation layer on the steel, protecting it. It is probably condition dependent (such as what the pipe is carrying) for which mechanism is dominant.

 

This seems much more plausible to me.

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http://en.wikipedia.org/wiki/Corrosion_inhibitor

 

This one indicates that the dichromate forms a passivation layer on the steel, protecting it. It is probably condition dependent (such as what the pipe is carrying) for which mechanism is dominant.

 

This was one idea I did have but when I read this page: http://en.wikipedia.org/wiki/Chromate_conversion_coating

It mentions that "It cannot be applied directly to steel or iron... Phosphate coatings on iron and steel may also be treated with a chromic acid rinse to enhance the phosphate coating" so it appears that it might not be some chromate coating. Also, looking at standard reduction tables, wouldn't the solution need to be acidic to form Cr metal? Unless there is another reduction path for chromate(VI) ions to reduce without an acidic solution?

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Stainless steel contains a small percentage of cromium, and since chromium oxidizes more readily than iron, a very thin, transparent layer of cromium oxide develops on the steel, inhibiting its corrosion.

 

You say "stainless steel" but Na2Cr2O7 inhibits corrosion of steel as well... unless normal steel also has a small percentage of chromium?

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Dear Sotiacho,

 

I have had an opportunity to discuss the matter with a colleague and think about it as well. While the following answer stares at us in the face, I accept that it is simplistic, perhaps wrong and it also raises more questions.

 

If dichromate is going to reduce then it would be in proximity to a reducing agent. However, for the reduction to be successful the electron transfer would have to occur (@ 3e-s per chromium atom) and the ion would have to break up fto result free chromium III ions. (presumably due to a newfound instability). This break-up would be facilitated and occur only if there are a heap of free protons(hydrogen ions) in solution whereby the "freed" oxide ions could bond with two protons to form water. Otherwise there would be free oxide ions in solution and that just doesn't happen in aqueous solution to any appreciable extent;

 

Clumsy, simplistic and perhaps incorrect - but I find it logical for the moment.

 

The main question that is raised is: Does dichromate become unstable when it manages to accept 6 electrons? I guess so, but does the instability only cause the break up to be complete when there are hydrogen ions present to facilitate the boding to oxides as they're released?

 

See you tomorrow.:doh:

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The main question that is raised is: Does dichromate become unstable when it manages to accept 6 electrons? I guess so, but does the instability only cause the break up to be complete when there are hydrogen ions present to facilitate the boding to oxides as they're released?

 

This is an interesting premise, as the rusting of iron also releases hydroxide ions. This would suggest, then, that Le Chatelier's principle could be at work, as the oxidation of dichromate releases hydroxide ions as well. Perhaps the dichromate then is a proton sink?

 

The one problem I see with this, is why use dichromate in this application (if this really is the mechanism)? Hexavalent chromium is notoriously toxic, so why use this if it's only purpose is generating hydroxide ions? Why not sodium hydroxide, or Proton Sponge, or some other base?

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You say "stainless steel" but Na2Cr2O7 inhibits corrosion of steel as well... unless normal steel also has a small percentage of chromium?

 

stainless steel = steel with chromium

 

at least I'm pretty sure

 

search wikipedia stainless steel

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big314mp,

 

we do. We have three scenarios/conditions in our investigation that prevents the oxidation of shiny metallic iron.

 

1. iron in NaOH solution

2. iron in freshly boiled distilled water and "sealed" with a petroleum jelly cap

3. iron in non-acidified dichromate solution

 

We setted up three other scenarios to predict and then confirm corrosion of the nails.

 

4. iron in 0.1M HCl solution

5. iron in a 0.1M neutral salt solution

6. iron in distilled water(with naturally dissolved air in it)

 

The results/observations are "great" - including the much later addition of ferricyanide solution to confirm or refute the presence of Fe2+.

 

The discussion that has gained momentum is centred around the question:

 

Why doesn't an oxidising agent like dichromate (or permanganate or even others) reduce in the aqueous presence of a reducing agent( such as iron) if hydrogen ions aren't present in the solution?

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Dear Sotiacho,

 

I have had an opportunity to discuss the matter with a colleague and think about it as well. While the following answer stares at us in the face, I accept that it is simplistic, perhaps wrong and it also raises more questions.

 

If dichromate is going to reduce then it would be in proximity to a reducing agent. However, for the reduction to be successful the electron transfer would have to occur (@ 3e-s per chromium atom) and the ion would have to break up fto result free chromium III ions. (presumably due to a newfound instability). This break-up would be facilitated and occur only if there are a heap of free protons(hydrogen ions) in solution whereby the "freed" oxide ions could bond with two protons to form water. Otherwise there would be free oxide ions in solution and that just doesn't happen in aqueous solution to any appreciable extent;

 

Clumsy, simplistic and perhaps incorrect - but I find it logical for the moment.

 

The main question that is raised is: Does dichromate become unstable when it manages to accept 6 electrons? I guess so, but does the instability only cause the break up to be complete when there are hydrogen ions present to facilitate the boding to oxides as they're released?

 

Hey Comon,

Yes, I do agree with that and that was actually the first thing I thought of but a few others disagreed with that and argued that it only proves that dichromate doesn't reduce but the dichromate as a solution could have oxygen dissolved and cause the iron to corrode just as in other solutions like sodium chloride. I did also argue that iron won't corrode in basic solutions with excess OH- ions, a situation which the dichromate produces to a very small extent but the concentration of excess protons and hydroxide ions in water would be rather low initially so it seemed unlikely to be enough.

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