iron metal

[kingjewel1]

Showing that both iron (III) and iron (II) ions in solution should react with zinc to give iron metal.

Zn=>Zn2+ e

Fe3+ +e=>Fe2+ + 2e=>Fe

 

Instead only iron (II) irons are produced.

2Fe3+ +Zn==>Fe2+ + Zn2+

 

Apparently iron metal it is not made but why not? only 1 point question, has me stumpped.

Help me with the half equations for the iron above please.

thanks guys

[YT2095]

plain and simple...

 

COST!

[kingjewel1]

How do you mean?

[YT2095]

Hmm... the current method employed is very "cost effective" and so it`s adhered to.

although your method is quite workable (and might even be used in Mil-Spec equipment), it would cost far too much using this method on a mass production scale.

 

that`s what I mean

[kingjewel1]

Interesting. I think though, the question is looking for a reason on the microscopic level if you know what i mean...

[jdurg]

I think it's all about relativity. Is pure Zinc metal really a good enough reducing agent to reduce the Iron 3+ ion back to the metal? The Zinc would rather just reduce two Fe3+ ions to Fe2+ and become Zn2+ itself. (Your first oxidation reaction is wrong. You're short an electron on the products side). In a sense, you could say that one Zn atom will reduce Fe3+ to Fe1+, but Fe1+ is not a stable ion as far as I'm concerned, so I don't believe it would exist. So now you have a solution with Fe2+ and Zn2+ ions in there. Is Zn2+ willing to give up two more electrons to become Zn4+ and have Fe2+ become iron? God no. So what happens is that no solid iron is produced. However, if you can remove all of the Fe3+ ions with some zinc metal, and then add even more zinc metal to the mixture, you'll start reducing the Fe2+ ions back down to Iron metal and you will eventually see iron forming. So the iron metal will form, but only when there's enough Zinc to remove the Fe3+ ions from the solution.

 

As for the cost, making Iron from Fe3+ is more expensive. When reducing the Fe3+ ion, for every 1.5 moles of Zinc metal you get one mole of Iron metal. When reducing the Fe2+ ion, for every mole of Zinc you get a mole of Iron. So you need to use up more zinc in order to reduce the Fe3+ ion back to the metal.

[budullewraagh]

i agree with jdurg. it's easier to reduce 3 Fe+3 cations to Fe+2 than one Fe+3 to Fe. (think reduction potentials)

 

in acidic solution (pH=0).

Fe+3 --(0.771)--> Fe+2

Fe+2 --(-0.44)--> Fe

in alkaline solution (pH=14)

FeO2- --(-0.69)--> HFeO2-

HFeO2- --(-0.8)--> Fe

[coquina]

The last time I had chemistry was in the 10th grade (40 years ago), I got a "D" then, and I haven't learned any more about it.

 

I will throw in here that Zinc is used as a "sacrificial anode" on the underwater running gear of boats. Salt water will cause galvanic corrosion between dissimilar metals. Because Zinc gives up its ions more readily that other metals, it is consumed before the steel is attacked.

 

http://www.corrosion-doctors.org/Aircraft/galvdefi.htm

 

Here's a galvanic corrosion chart, the closer two metals are to one another, the less likely corrosion will occur. Zinc and Iron are only 6 metals apart.

 

http://www.metal-mart.com/Guides/Galvanic.htm

 

This may not have anything at all to do with what you are discussing, and if it doesn't, I apologise.

[kingjewel1]

thanks guy very much!

Yeah it seems quite logical that with just 1 mole of zinc it would have to have an oxidation state of +4 and that is rediculous.

Apparently it is THERMODYNAMICALLY unstable but KINETICALLY stable.

I've never really understood the difference between the two, but i would have thought it was kinetically unstable due to such a high activation energy. :S