Sciencemadness Discussion Board

about double iron salt

breeze_in_mist - 30-9-2005 at 05:31

Why double salt of iron(II) (eg. ferrous ammonium sulfate) has a better stability then normal iron(II) salt like FeSO4? And why the stability will further increase if dilute sulfuric acid is added?

woelen - 1-10-2005 at 05:41

This has to do with the redox properties of the Fe(2+) ion. This ion is VERY prone to aerial oxidation when certain ligands are attached to the ion. The presence of an hydroxo-ligand makes it very sensitive for oxidation. With two hydroxo ligands (then it precipitates as Fe(OH)2), it very easily is oxidized. The double salt with ammonium ion already is somewhat acidic and the chance of formation of hydroxo ligands is smaller. At even lower pH, the in Fe(2+) certainly will not have hydroxo-ligands and then the ion is quite stable.

This phenomenon exists for more transition metals. E.g. Co(2+) is stable with respect to oxidation when only aqua ligands are attached, but when hydroxo, ammine or nitrito ligands are attached, then it very easily is converted to Co(3+). Manganese has the same property. Mn(2+) is quite stable in acid, but it is next to impossible to obtain a pure precipitate of Mn(OH)2, it is even more sensitive than Fe(OH)2.

Some redox couples show this sensitivity on ligands:

Fe(H2O)6(3+) + e --> Fe(H2O)6(2+) 0.77 V
Fe(CN)6(3-) + e --> Fe(CN)6(4-) 0.36 V
Fe(OH)3 + e --> Fe(OH)2 + OH(-) -0.53 V

12AX7 - 1-10-2005 at 10:14

Also the general tendancy for stable valence to be somewhat proportional to pH: lower valences are more stable at low pH, higher valences are more stable at high pH. But this is more of a loose rule of thumb and has no theory behind it.

Tim

On the contrary

chloric1 - 1-10-2005 at 14:41

Talking especially about transition metal ions. As the metal assumes a higher valence it becomes more covalent. Hence the ion frequently hydrolsises and the solution it acidic ex. FeCl3. In contrast, lower valences are closer to being ionic and are more basic. FeO is a stronger base than Fe2O3.

Also, this rule applies to nontransitional elements. For example antimony trichloride is acidic but the pentachloride is much more so and is used as a lewis acid for organic chlorinations.

12AX7 - 1-10-2005 at 14:55

Yeah, which means it doesn't like being itself...or...something. Just proving my point :P

Lower states like Fe2+ and 3+, Cu+ and 2+, etc. are about evenly stable I would say... the extreme valences like Fe6+, Cr6+ and Mn7+ are easiest to make in alkaline solution and tend to be unstable in acidic solution (though CrO3 is stable by itself).

Tim

woelen - 2-10-2005 at 04:09

No, this correct partly at best.

For iron you are right. Fe(VI) only exists in strongly alkaline state or in the solid state.

For chromium, things are different. Both in acidic and in alkaline solution, Cr(VI) is stable. Look at the dichromates. Not unstable at all.

For manganese, things are even the other way around! Just for fun, add a small amount of KMnO4 to a very concentrated solution of NaOH without any reductors in it. You'll see the formation of oxygen and the liquid becomes green. At very high pH, the permanganate ion is unstable. When you have the green manganate (VI) ion and you lower the pH, this ion disproportionates to MnO2 and MnO4(-). So, manganese (VI) is stable at very high pH and managanese (VII) only is stable at low and intermediate pH.

unionised - 2-10-2005 at 11:52

"Look at the dichromates. Not unstable at all. "
The instabillity of Cr(VI) under slightly acid conditions is neatly demonstrated by the ammonium dichromate volcano.

Permanganate is a stronger oxidant in acid conditions than in neutral ones.

A really obscure example of this is that under near neutral conditions peracetic acid will oxidise Mn(II) to Mn(VII) whereas in strong acid the permanganate oxidises the peroxy species.

In general it is easier to oxidise the "metal" in the middle when it's got lots of oxide or hydroxides round it.
This isn't a great shock- for example, in order to oxidise FeII to FeIII you need to remove an electron. It's easier to do that from the neutral species Fe(OH)2 than from the positively charged Fe++(aq). You are taking a negative charge away from something that attracts it.

woelen - 3-10-2005 at 08:21

Quote:
Originally posted by unionised
"Look at the dichromates. Not unstable at all. "
The instabillity of Cr(VI) under slightly acid conditions is neatly demonstrated by the ammonium dichromate volcano.

I thought we were dealing with solutions. In the solid state, yes, ammonium dichromate is somewhat unstable, but you first need to ignite it. It can be stored safely, so its not THAT unstable.
Quote:
Permanganate is a stronger oxidant in acid conditions than in neutral ones.

A really obscure example of this is that under near neutral conditions peracetic acid will oxidise Mn(II) to Mn(VII) whereas in strong acid the permanganate oxidises the peroxy species.

Being a strong oxidizer is related to stability, but is not the same as stability. Although permanganate is a stronger oxidizer at lower pH, it is not more stable at that pH. At very high pH it is unstable and decomposes at once.

unionised - 5-10-2005 at 00:32

A couple of points.
I regularly use a solution of KMnO4 in 10% NaOH for cleaning glassware, it keeps for weeks. Sure, some of the Mn(VII) is reduced to Mn(VI) by trash in the water and/ or hydroxide, but most of it stays purple (and by extracting into acetonitrile you can prove that there is always Mn(VII) present, even when it's obscured by the green colour.)

Cr(VI) is commonly used as a laboratory oxidiser. It always gets used in acid solution because it's not much of an oxidant in basic conditions (There are exceptions to this like the pyridine complexes but ther are of limited relevance).

If you don't think Cr(VI) is an oxidant, what do you think it is used for in the lab?

The fact that you meed to heat ammonium dichromate to get it going isn't important. The fact that it keeps going shows that it's unstable. You need to look at the differennce between kinetic and thermodynamic stabillity.

Being a strong oxidiser IS the same as being unstable towards reduction.
TiO2 will react slowly with HCl - to form TiCl4 (actually, chloro/ aquo complexes).
MnO2 reacts as an oxidant and forms Cl2.

The reason for this difference is that MnCl4 isn't stable but TiCl4 is.

woelen - 5-10-2005 at 01:50

Quote:
Originally posted by unionised
A couple of points.
I regularly use a solution of KMnO4 in 10% NaOH for cleaning glassware, it keeps for weeks. Sure, some of the Mn(VII) is reduced to Mn(VI) by trash in the water and/ or hydroxide, but most of it stays purple (and by extracting into acetonitrile you can prove that there is always Mn(VII) present, even when it's obscured by the green colour.)

I did tests with concentrated solutions of NaOH in distilled water and to this I added a solution of KMnO4. The liquid quickly turned green and little bubbles of gas are produced (I think this is oxygen). Of course, small amounts of manganese (VII) may remain in the liquid, but then these are only very small amounts. How do you explain the bubbles of gas? Have a look at this observation on my website:

http://woelen.scheikunde.net/science/chem/solutions/mn.html

Look at the end of the page, where you see the formation of green manganese (VI) and bubbles of oxygen.
Quote:
Cr(VI) is commonly used as a laboratory oxidiser. It always gets used in acid solution because it's not much of an oxidant in basic conditions (There are exceptions to this like the pyridine complexes but ther are of limited relevance).

If you don't think Cr(VI) is an oxidant, what do you think it is used for in the lab?

Where did I say that Cr(VI) is not an oxidant? It is and I agree with you that it is much more so in acidic environments than in alkaline environments.
Quote:
The fact that you meed to heat ammonium dichromate to get it going isn't important. The fact that it keeps going shows that it's unstable. You need to look at the differennce between kinetic and thermodynamic stabillity.

OK, I would call it meta-stable in that case. Think of it as a ball on a high mountain, surrounded by a wall. You first have to put in some energy to roll it over the wall, before all potential energy is released by rolling downwards from the mountain.
But I would not call this 'unstable'. I have the impression that we agree here, but that our wordings are different. Using your definition, most of the world around us is 'unstable'. For example, wood in air also is 'unstable', even wood on its own is 'unstable'.
Quote:
Being a strong oxidiser IS the same as being unstable towards reduction.
TiO2 will react slowly with HCl - to form TiCl4 (actually, chloro/ aquo complexes).
MnO2 reacts as an oxidant and forms Cl2.

The reason for this difference is that MnCl4 isn't stable but TiCl4 is.

Again, I think we agree, but that it is a question of wording. I would not call a compound on its own 'unstable', such as KMnO4. What you are talking about here is a combination of compounds which is unstable (e.g. KMnO4, mixed with HCl). So, this is why I said that stability and oxidizing strength are related, but are not the same. I make difference between 'unstable' and 'unstable towards reduction'.

unionised - 6-10-2005 at 06:54

I think most of what we are disagreeing about is the exact meanings of words.
Oxygen is stable but generally a strong oxidant. However, one cannot talk about something as an oxidiser in isolation- there needs to be another reactant. For example oxygen is rapidly oxidised by PtF6.
What I didn't make clear was that I meant that oxidants are unstable in the presence of reducing agents. So CrO3 is reasonably stable - though it decomposes on heating. A mixture of CrO3 and sawdust would not last long before reaction took place.
The case for Mn (VII) vs Mn(VI) is complicated by the acid base equilibria involved.
In general it's still easier to oxidise things under alkaline conditions as with the production of chromates and manganates.
More importantly, the acidity of Mohr's salt is the reason for its stabillity- and that was the original question.

You wrote "Both in acidic and in alkaline solution, Cr(VI) is stable. Look at the dichromates. Not unstable at all." it's true as it stands but, in the presence of a reducing agent it stops being correct because acid Cr(VI) is a good oxidant. If dichromates were completely stable they wouldn't be oxidants.