Sciencemadness Discussion Board

H2SO4 conc. + Fe ??

denatured - 8-1-2007 at 03:08

Hello

Why concentrated sulphuric acid cannot react with iron?

Dilute sulphuric acid reacts correctly, why not the concentrated one?

H2SO4 + Fe --> FeSO4 + H2

I assume that it reacted with the iron then the reaction stopped may be because the sulphate salt didn't dissolve in the water as the water is bonded in the concentrated acid but not in the dilute one ... and to continue the reaction as it happens in the dilute acid, the sulphate must be removed from the surface of the iron bar.. iron sulphate is soluble, right? ... so what do you think? is the sulphate is really blocking the reaction? ..

Darkblade48 - 8-1-2007 at 04:08

I'm not too sure about this, as it's still 7 am in the morning, but I thought that iron did not react in concentrated sulfuric acid because it was too strongly oxidizing.

Aqua_Fortis_100% - 8-1-2007 at 04:28

sorry about my ignorance but even HOT conc. H2SO4 cannot erode the iron???

If no this is really strange because e.g. a MUCH less reactive metal like copper can be dissoluted by hot conc. H2SO4 yielding in cupric sulphate ,H2O and SO2..

(conc.)2 H2SO4 + Cu ----heat------> CuSO4 + 2H2O + SO2

[Editado em 8-1-2007 por Aqua_Fortis_100%]

[Editado em 8-1-2007 por Aqua_Fortis_100%]

kaviaari - 8-1-2007 at 06:39

This is simply because iron does not react with sulphuric acid itself. When you a little bit of water to sulphuric acid you get a solution that has all sorts of nice ions that can lead to reactions with base metals. Because we chemists are so lazy we just write H2SO4 + Fe --> FeSO4 + H2

[Edited on 8-1-2007 by kaviaari]

Ozone - 8-1-2007 at 17:43

Dilute H2SO4 corrodes steel in a particular way (The name of which I cannot recall off-hand) leading to pitting and making the steel brittle; we had to trash a very expensive steam explosion reactor vessel on acount of this.

If the concentrated stuff will not eat Fe° then it must be an effect of there being little media, viz. water, in which ion transfer can occur (conc. is *highly* dehydrating). This is not a problem in the dilute regime.

If I have time, I dump some H2SO4 on some iron turnings tomorrow and report the results.

Cheers,

O3

guy - 8-1-2007 at 17:51

Maybe the iron oxide formed does not react fast enough to see any reaction.

The_Davster - 8-1-2007 at 18:04

Thought it was common ion effect. FeSO4 formed initially is insoluble in conc. H2SO4 preventing further reaction.

guy - 8-1-2007 at 18:06

Quote:
Originally posted by The_Davster
Thought it was common ion effect. FeSO4 formed initially is insoluble in conc. H2SO4 preventing further reaction.


Why does it work for copper then?

Ozone - 8-1-2007 at 18:16

Ah!

Passivation is a probability for sure (why did I not consider that *kicks self in ass*)! As I said, Iv'e nver tried it so what-the-hell, I'll give it a quickie tomorrow (I'll photograph the thing for your enjoyment!).

I'll bet Fe° turnings with good stirring (with some broken glass) would do the trick--if--lack of H+/Geg. mass transfer is not the limiting factor!?

Cheers all,

O3

chemoleo - 8-1-2007 at 18:44

Anhydrous FeSO4 is not particulary well soluble, not even in water. I doubt it is particularly well soluble in pure H2SO4. Thus, passivation has to be the answer... similar to lead sulfate, etc. Copper works because it is dissolved oxidatively, i.e. SO2 is produced during the reaction. It wouldnt otherwise dissolve. Different from the others above therefore.

Magpie - 8-1-2007 at 20:30

This markedly better corrosion resistance of iron (or steel) at higher concentration of sulfuric acid was a fact generally unknown to entering chemical engineering students with which our professors could amaze us. But I couldn't remember why.

According to Perry's Chemical Engineers' Handbook, 3rd ed, p 23-7, Chemoleo is correct. Here's what Perry's says:

"The corrosion resistance of steel depends on the formation of an oxide film. However, resistance to corrosion is somewhat limited. Carbon steel should not be used in contact with dilute acids. Thus it is not recommended with sulfuric acid below 90 per cent. Between 90 and 98 per cent, steel can be used up to the boiling point; between 80 and 90 per cent , steel is good up to 140F. Usually, steel is not used with hydrochloric, phosphoric, or nitric acids."

Ozone - 8-1-2007 at 20:45

Even so, It appears that the lack of transfer media (water) is the culprit, viz. FeSO4 would be solvated from the surfaces (at least in film concentrations)negating passivation. Although I generally trust Perry, I'd think that a SO4 passivation layer would be more likely than an O or O3 film, viz. the sulfate forms with water adsorbed onto the surface of the metal.

I'll check out the corrosion type with my ME friend (who diagnosed our ill-fated rxn vessel) and note the observations of the actual test.

Unfortunately, surface composition in H2SO4 (conc.) is difficult to determine, even with an entire Chemistry dept. at your disposal. I suppose that an "indirect" determination could be made by filtering the crap through glass scinter and washing under N2 or argon (crappy clothes are a must) with ice-cold water (maybe dry with acetone?). Then, I suppose, we might be able to try some conventional characterization techniques (unfortunately, using the bomb calorimeter might be the only way, viz. Fe-Fe2O3, SO4 stays SO4, etc.).

I think that the steel becomes more brittle as a result of zonal enrichment of C as the alloy is disrupted (removed); this creates "fault lines" in the metal.

Anyway thanks all, I've got to go to work tomorrow so,

Goodnight (US anyway),

O3

guy - 9-1-2007 at 01:11

Quote:
Originally posted by chemoleo
Anhydrous FeSO4 is not particulary well soluble, not even in water. I doubt it is particularly well soluble in pure H2SO4. Thus, passivation has to be the answer... similar to lead sulfate, etc. Copper works because it is dissolved oxidatively, i.e. SO2 is produced during the reaction. It wouldnt otherwise dissolve. Different from the others above therefore.


Iron is an even stronger reducing agent than copper, and the reduction of sulfate is more energetically favored over reduction of H+. Thus it would be passivated by FeO which is a substance that dissolves only slowly in acid, much like Al2O3.

woelen - 9-1-2007 at 09:40

I observed this type of phenomenon with many different combinations of chemicals. All of the following do not react, or only react very slowly. Some of them do not even react when heated. As soon as some water is added, the reaction sets in, some of them with extreme violence:

Na + Br2
Ni + conc. HNO3
Mg + Br2
Mg + ICl
Zn + H2SO4

I think that in all of these situations the water plays an important role by allowing intermediate ionic steps to be performed in the reaction. I was really surprised to see a piece of Na-metal floating on bromine without any reaction.

This topic was placed in "beginnings", but this phenomenon definitely is not something to be considered beginner's chemistry. I have the impression that this phenomenon still is not fully understood.

[Edited on 9-1-07 by woelen]

denatured - 9-1-2007 at 15:28

then the unbonded water in diluted H2SO4 dissolves the formed iron sulphate to allow furthar reaction, no?

p.s. you didn't fully read my entry, did you?;)

EDIT:I posted it here as I thought that it is a simple thing you all know...

[Edited on 9-1-2007 by alnokta]

guy - 9-1-2007 at 18:15

Quote:
Originally posted by woelen
I observed this type of phenomenon with many different combinations of chemicals. All of the following do not react, or only react very slowly. Some of them do not even react when heated. As soon as some water is added, the reaction sets in, some of them with extreme violence:

Na + Br2
Ni + conc. HNO3
Mg + Br2
Mg + ICl
Zn + H2SO4

I think that in all of these situations the water plays an important role by allowing intermediate ionic steps to be performed in the reaction. I was really surprised to see a piece of Na-metal floating on bromine without any reaction.

This topic was placed in "beginnings", but this phenomenon definitely is not something to be considered beginner's chemistry. I have the impression that this phenomenon still is not fully understood.

[Edited on 9-1-07 by woelen]


I think the answer is simple enough:

Na + Br2 <b>NaBr formed is insoluble in Br2. Unlike the reaction of Al and Br, which AlBr3 can complex with Br-.</b>
Ni + conc. HNO3 <b> Probably an oxide layer protection</b>
Mg + Br2 <b>Same as Na + Br2</b>
Mg + ICl <b>No idea :)</b>
Zn + H2SO4 <b> Check this out Pure Zinc in Acid shows no reaction</b>

12AX7 - 9-1-2007 at 20:35

Mg + ICl is a solid thing, no solvent.

Na + Br2 isn't spontaneous!? (Or hell, hypergolic would be a better word!)

Tim

guy - 9-1-2007 at 22:08

Quote:
Originally posted by 12AX7
Mg + ICl is a solid thing, no solvent.

Na + Br2 isn't spontaneous!? (Or hell, hypergolic would be a better word!)

Tim


Spontaneous is different than fast. Check out the reaction of Na with Cl2, you need to heat the Na metal first. It's because a coating of NaX, stops the reaction.

woelen - 9-1-2007 at 23:46

Quote:
Na + Br2 NaBr formed is insoluble in Br2. Unlike the reaction of Al and Br, which AlBr3 can complex with Br-.

Are you sure this is the explanation? Even if I add a very small amount of water, the reaction starts quickly and violently. Of course, if you add lots of water, all will dissolve and the Na also reacts with water, but my surprise was that just a tiny bit of water can start the reaction. So, I think there is more to say about this than just not dissolving of NaBr in Br2.

You mention the complexation of AlBr3 (formed from Al and Br2) with Br(-), but where should that Br(-) come from? From other AlBr3? But that would imply Al(3+) ions and can those dissolve. No, I do not think this is an explanation why Al+Br2 does react and Na+Br2 (and Mg+Br2) do not react.
Quote:
Mg + ICl is a solid thing, no solvent.

I did this reaction with liquid ICl. The Mg nicely floats on the ICl. I made a picture of that:

http://woelen.scheikunde.net/science/chem/exps/Cl+I/exp0013....

ICl melts at 27 C or so, but once it is liquid it does not easily solidify, it can be cooled well below 20 C before it becomes a solid. I made the ICl by passing excess Cl2 over I2 (this first makes yellow solid ICl3) and on standing, the ICl3 slowly looses Cl2 and liquid ICl remains behind. Probably the ICl was very impure and contained quite some ICl3 as well, but that does not matter for this observation.

Again, adding a small amount of water makes the reaction start immediately and very violently.

guy - 10-1-2007 at 00:14

Quote:
Originally posted by woelen
Quote:
Na + Br2 NaBr formed is insoluble in Br2. Unlike the reaction of Al and Br, which AlBr3 can complex with Br-.

Are you sure this is the explanation? Even if I add a very small amount of water, the reaction starts quickly and violently. Of course, if you add lots of water, all will dissolve and the Na also reacts with water, but my surprise was that just a tiny bit of water can start the reaction. So, I think there is more to say about this than just not dissolving of NaBr in Br2.

You mention the complexation of AlBr3 (formed from Al and Br2) with Br(-), but where should that Br(-) come from? From other AlBr3? But that would imply Al(3+) ions and can those dissolve. No, I do not think this is an explanation why Al+Br2 does react and Na+Br2 (and Mg+Br2) do not react.
Quote:
Mg + ICl is a solid thing, no solvent.

I did this reaction with liquid ICl. The Mg nicely floats on the ICl. I made a picture of that:

http://woelen.scheikunde.net/science/chem/exps/Cl+I/exp0013....

ICl melts at 27 C or so, but once it is liquid it does not easily solidify, it can be cooled well below 20 C before it becomes a solid. I made the ICl by passing excess Cl2 over I2 (this first makes yellow solid ICl3) and on standing, the ICl3 slowly looses Cl2 and liquid ICl remains behind. Probably the ICl was very impure and contained quite some ICl3 as well, but that does not matter for this observation.

Again, adding a small amount of water makes the reaction start immediately and very violently.


I meant that the NaBr protects from further oxidation. Adding water dissolves the layer.

AlBr3 can complex with Br- meaning that it will break up Br2 into Br- and Br+, and the complexation will help remove some AlBr3.

AlBr3 + Br2 ---> [AlBr4]- + Br(+)

woelen - 10-1-2007 at 00:58

Ah.. I see that explantion of NaBr. Yes, it could be that this is the explanation. I would not really call it dissolving, but it makes the layer less solid, the water makes it permeable for the Br2.

The explanation for the reaction of AlBr3 with Br2 is quite interesting. So, in solution I get Br(+) and AlBr4(-) ions? This Br(+) then in turn reacts with Al? Do you have a reference for this, it sounds quite interesting and having a good read about this would be a nice entertainment ;).

[Edited on 10-1-07 by woelen]

guy - 10-1-2007 at 14:43

Quote:
Originally posted by woelen
Ah.. I see that explantion of NaBr. Yes, it could be that this is the explanation. I would not really call it dissolving, but it makes the layer less solid, the water makes it permeable for the Br2.

The explanation for the reaction of AlBr3 with Br2 is quite interesting. So, in solution I get Br(+) and AlBr4(-) ions? This Br(+) then in turn reacts with Al? Do you have a reference for this, it sounds quite interesting and having a good read about this would be a nice entertainment ;).

[Edited on 10-1-07 by woelen]


I don't but its all the most rational hypothesis I have come up with. It came up because HX (X=halogen) and X2 always attack aluminum faster than most other compounds. This is due due to the fact that is can form complexes with halides quite easily.

The closest anaolgy to the "AlB3 + Br2" is the pseudo Friedel Craft's halogenation of benzne, where the formation of an electrophile (X+) is promoted by a lewis acid that has the ability to complex with it.

Jdurg - 11-1-2007 at 17:22

Also, with the sodium metal + halogen reactions, when you add the tiniest speck of water you can easily catalyze the reaction. If you've ever taken a plain block of sodium metal and just flicked a tiny mist of water on it, you'll notice that there is an immediate, localized heating where the reaction takes place. This heat will then kick off the reaction between the alkali metal and the halogen. (If you take a flask of chlorine and put in there some heated sodium, it will immediately inflame in the chlorine atmosphere).

In addition, the reaction between the alkali metal and water may form a tiny layer of hydroxide on the surface of the alkali which is readily attacked by the halogen as well. So either way, the addition of heat helps to get the reaction started. I wish I had gotten a video of it, but I took a small sliver of sodium metal and put it in a sealed vial with bromine. (Anhydrous bromine). As you've seen woelen, nothing happened. I then took a butane lighter and heated the vial until the bromine was nearly all vaporized. Suddenly, the reaction just took off with a "poof" and the inside was covered in NaBr. So I think that shows that the activation energy must be fairly significant for this reaction, but the addition of the tiniest drop of water can overcome that activation energy barrier and allow the reaction to proceed. (Or you can just increase the temperature of the vessel).

Ozone - 11-1-2007 at 17:53

Nevermind that once exothermia occurs with an alkali metal like Na, the metal melts. The molten blob is then mobile and can constantly reveal a fresh, non-passivated surface.

Did the reaction consume all of the metal in one "poof" or were there several? I'm betting that once it got going, the constantly refreshed surface kept it going-rapidly.

Neat experiment. I too wish you had taken some video!

Cheers,

O3

Jdurg - 12-1-2007 at 14:53

I'm tempted to try it again, but making more bromine just isn't too "fun" of an idea. What happened is that initially there was nothing going on. I put the vial over the alcohol burner and I saw the inside getting darker and darker red as more bromine vaporized inside the tube. Then, when the metal started to soften and liquify a bit, there was an instantly bright flash and the entire chunk just disappeared into the now white covered vial. I only used a piece of sodium about half the size of a pinky-nail so it all reacted pretty quickly.

I could only imagine it "poofing" repeatedly if you had a wide open container and a huge amount of bromine so that the molten metal breaks apart like it does when reacting with water.