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Chemistry Forums for Students => High School Chemistry Forum => Topic started by: aeacfm on August 29, 2010, 03:34:04 AM
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it is known that
Fe(s) + HCl (aq) :rarrow: FeCl3 +H2
let us say that
Fe++ (iron compound in solid state) + HCl (aq) :rarrow: FeCl3 +product
but what is the product of the following :
Fe++(iron compound in solid state) + HCl (concentrated solution) :rarrow:...
will it be Fe++ compound or Fe+++ compound
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also if i want to obtain the following state :
Fe++(s) + HCl :rarrow: FeCl2 + product
how can obtain it ?
i dont mean HCl specifically i want any reagent ro dissolve the Fe++ compound as it to Fe++ liquid phase.
thanks in advance
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To answer you first question:
2H3O+(aq)+2e-→2H2O(l)+H2(g) (E0= 0.000 V)
Fe3+(aq)+e-→Fe2+(aq) (E0=0.771 V)
Fe3+(aq)+3e-→Fe(s) (E0=-0.036 V)
Fe2+(aq)+2e-→Fe(s) (E0=-0.440 V)
just try to use this data, and remember that ∆G=-nF∆E...
Talking about your second question, I don't really understand what you mean with Fe2+(s). If you're talking about a Fe (II) compound dissolved in water and you want to prevent it from taking part in possible redox reactions I think the only way is to complex it using a chelating agent, but I don't know what you could use for your purpose...
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it is known that
Fe(s) + HCl (aq) :rarrow: FeCl3 +H2
That's the first time I hear about. As far as I know H+ is not strong enough to oxidize iron to +3. Dissolving iron in acid we get Fe2+, which gets further oxidized to Fe3+ in the presence of air oxygen.
Fe3+(aq)+e-→Fe2+(aq) (E0=0.771 V)
Fe3+(aq)+3e-→Fe(s) (E0=-0.036 V)
There is some inconsistency here. I find it hard to believe that it is easier to oxidize Fe to Fe(III) than Fe(II) to Fe(III).
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To answer you first question:
2H3O+(aq)+2e-→2H2O(l)+H2(g) (E0= 0.000 V)
Fe3+(aq)+e-→Fe2+(aq) (E0=0.771 V)
Fe3+(aq)+3e-→Fe(s) (E0=-0.036 V)
Fe2+(aq)+2e-→Fe(s) (E0=-0.440 V)
just try to use this data, and remember that ∆G=-nF∆E...
Talking about your second question, I don't really understand what you mean with Fe2+(s). If you're talking about a Fe (II) compound dissolved in water and you want to prevent it from taking part in possible redox reactions I think the only way is to complex it using a chelating agent, but I don't know what you could use for your purpose...
many thanks MrTeo
so can i use hydroxyle amine ?i think it will work
the Fe++ i mean ferrous compound
also i will try H2SO4 as it will form FeSO4 which is satble!!!!!
final thing do you mean by E0 data is to see the reaction :delta:G and so elucide the rection product
it is known that
Fe(s) + HCl (aq) :rarrow: FeCl3 +H2
That's the first time I hear about. As far as I know H+ is not strong enough to oxidize iron to +3. Dissolving iron in acid we get Fe2+, which gets further oxidized to Fe3+ in the presence of air oxygen.
i mean by FeCl3 it is the final product
thanks for your kind reply
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The numbers are correct, according to the handbook. They also add up when you take the numbers of electrons into account (
3*-0.036-2*-0.44=0.772).
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final thing do you mean by E0 data is to see the reaction :delta:G and so elucide the rection product
Well, if you find out what's the ∆E0 of the reaction you can easily see if it is positive or negative (that tells you if the reaction is spontaneous or not) and eventually compare it with other spontaneous reactions (the lower the ∆G is, the more favored the reaction will be, as $$ K_{eq}=\frac{\Delta G}{RT}/$$).
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The numbers are correct, according to the handbook.
None of my handbooks (nor other sources that I use) lists this reaction, they all list separately first stage (Fe(0) -> Fe(II)) and second stage (Fe(II) -> Fe(III)).
They also add up when you take the numbers of electrons into account (
3*-0.036-2*-0.44=0.772).
Can you elaborate on what you are doing here, as I am not getting the idea behind.
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Erratum
(https://www.chemicalforums.com/proxy.php?request=http%3A%2F%2Fwww.forkosh.dreamhost.com%2Fmimetex.cgi%3F%7B+K_%7Beq%7D%3De%5E%7B-%5Cfrac%7B%5CDelta+G%7D%7BRT%7D%7D%7D&hash=35a35401090eec7c79bde4153a808f8f61852993)
@Borek:
Actually I got that value from the Net, I'll try to retrieve the site and post the URL here...
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None of my handbooks (nor other sources that I use) lists this reaction, they all list separately first stage (Fe(0) -> Fe(II)) and second stage (Fe(II) -> Fe(III)).
These potentials are listed in the electrochemical series in the CRC Handbook of Chemistry and Physics (90th ed.). Among colleagues we usually refer to it as 'the handbook', hence my lack of specification.
Can you elaborate on what you are doing here, as I am not getting the idea behind.
I just checked if the numbers add up when converting them to free energy, using :delta: G=-nFE. I left out F since it gets divided out anyway, as well as some double minus signs.
Fe3+ + 3e- :rarrow: Fe(s) :delta: G = -3*F*-0.036
Fe2+ + 2e- :rarrow: Fe(s) :delta: G = -2*F*-0.44
Subtracting the two gives :delta: G for Fe3+ + e- :rarrow: Fe2+.
And E = - :delta: G/nF then gives, -(-3F*-0.036-(-2F*-0.44))/F = 0.772V.
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OK, I see. But there is something fishy here, IMHO just because these numbers add up doesn't mean they are correct. Unfortunately my thermodynamics is so rusty I don't see what is wrong with this approach, but it is not that difficult to show that something _is_ wrong.
Imagine you have a piece of iron in an acid solution. For the sake of simplicity let's assume activities of all substances involved are 1, so we can deal with standard potentials. Iron dissolves and hydrogen evolves. In such solution, as suggested by the potential given, iron gets oxidized to Fe3+. So, what we do have now? H+ and H2 - that means half cell of potential 0, and Fe3+. But if you look at standard potentials table Fe3+ can't exist in such mixture, with its standard potential at 0.77 V it will immediately oxidize hydrogen to H+ and get reduced to Fe2+. So there will be no Fe3+ in the solution (well, some traces), only Fe2+.
I have a feeling that potential given for Fe/Fe3+ is calculated using approach similar to the one you have used to check the numbers (so in fact you have just reversed the procedure getting back to the original data - no wonder numbers add up). There is something wrong this approach, as it gives values that won't survive in solution.
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I agree it looks strange. I looked it up in "Standard Potentials in Aqueous Solutions" by Bard, Parsons and Jordan. The standard potential for Fe(III)/Fe is indeed calculated from the other two half reactions. They refer to the Pourbaix book "Atlas d'équilibres électrochimiques".
Why would iron form Fe(III) based on the potentials? The standard reduction potential of Fe(II) to Fe is -0.44V, and thus 0.44V for the reverse process, which occurs upon dissolution. For oxidation to Fe(III) it is only 0.036V. The dissolution of iron in acid should then give mostly Fe(II) and some Fe(III). The Fe(III)/Fe(II) standard potential also pushes the system in that direction. I think it is consistent even though the numbers look a bit weird.
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i can see that you Forgot heart of the matter and every body want to prove some thing to him self
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The standard reduction potential of Fe(II) to Fe is -0.44V, and thus 0.44V for the reverse process, which occurs upon dissolution.
Where this idea of reversed potential comes from? Reaction doesn't care which way it goes, for given activities of all substances involved it goes in both directions at exactly the same potential, as given by the Nernst equation. Take a look at the CV curve - http://en.wikipedia.org/wiki/Cyclic_voltammetry - one scan is reduction, other is oxidation, both go at (almost) the same potential; difference is 0.059/n and (for reversible processes) doesn't depend on standard reduction potential.
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i can see that you Forgot heart of the matter and every body want to prove some thing to him self
It is not that easy. You assumed that Fe dissolves producing Fe3+ - as I have pointed out as far as I know this assumption is wrong. However, according to CRC handbook quoted by others you are right - but numbers shown in the book lead to absurd conclusions. We are trying to find out what is really going on.
FeCl2 solution acidified with HCl will not miraculously convert itself to FeCl3 solution. Also dissolving ferrous salt in HCl will not lead to its oxidation, you will be still left with Fe2+.
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I think these last posts are actually quite relevant to your question. Although it seems to go a little off-topic the central issue is still, whether or not Fe(II) will be formed in acidic solutions.
One way to look at your question is based on pourbaix diagrams, the wikipedia page on this shows an example of the iron system. It shows that without complexing agents Fe(II) is stable in acid under 'regular' conditions, but as Borek already pointed out, the presence of oxygen can interfere, since it will lead to a more oxidising environment.
http://en.wikipedia.org/wiki/Pourbaix_diagram
You can also stabilise oxidation states using complexing agents, compare for example Fe(III)/Fe(II) (0.771V) to [Fe(CN6)]3-/[Fe(CN6)]4- (0.358V). In this case you're not dealing with just 'aquated' iron ions anymore though, but with complexes.
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Where this idea of reversed potential comes from?
I should really quit leaving out bits and pieces :-) I know redox potentials as such do not change sign, but signs do come in when taking the difference between two half reactions to get the potential of a total reaction. In this case I implicitly subtracted the reduction of Fe(II) and Fe(III) from the reduction of protons under standard conditions (0V).
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Note how Pourbaix diagram - again - contradicts the idea of Fe -> Fe3+ oxidation by H+ below 0 V.
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i can see that you Forgot heart of the matter and every body want to prove some thing to him self
FeCl2 solution acidified with HCl will not miraculously convert itself to FeCl3 solution. Also dissolving ferrous salt in HCl will not lead to its oxidation, you will be still left with Fe2+.
lets make some thing prctical here
Only seconds ago , i dissolved Ferrous sulfate in concentrated HCl and take a portion of the solution and add KI solution IMMEDIATELY ???????brown colour of iodine formed and adding chloroform the violet color of iodine appear in the oraganic layer of chloroform that shows that Fe+++ NOT Fe++ present in the solution may Fe++ present but traces
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when adding the DILUTED HCl (1:1) it takes more time to appear the iodine color than what happened with the CONCENTRATED HCl
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Ferrous sulfate or Mohr salt (Ammonium iron(II) sulfate)? Most ferrorus salts are unstable and get oxidized by the air oxygen, so they are always contaminated by Fe(III). Mohr salt is relatively stable.
Note that your solution is all the time in contact with air and it is saturated with oxygen, so small amounts of iron(II) will be oxidizied immediately. Check if using freshly boiled water (and perhaps a boiled acid) won't slow the oxidation further down.
Just because iodine forms doesn't mean your solution contains only Fe3+. Most likely you have some mixture of Fe(II)/Fe(III). If you want to be sure it is only Fe(II) you need to reduce it down, that's a standard procedure before determination of Fe(II) by any redox titration.
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If you want to be sure it is only Fe(II) you need to reduce it down, that's a standard procedure before determination of Fe(II) by any redox titration.
thats exactly what i need
how can i achieve this standard ?in other words how can i be sure that the iron II will not oxidized when i want to identify the iron III in mu sample ?
simply thats what i want ;D ;D ;D
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Determine Fe(II), reduce all, determine sum.
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Determine Fe(II), reduce all, determine sum.
thats better but lack of chemicals and time are disadvantages
then correct me if iam wrong
i make mix of hydroxyle amine +HCl (1:1) the mix was 1:1 and do the previous iodine test with ferrous sulfate and the test was -ve so do what i performed accepted practically ???