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Topic: redox titration!  (Read 10755 times)

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Offline _cheers

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redox titration!
« on: January 13, 2007, 10:34:36 PM »
with the rx:

MnO4 + 8H+ + 5Fe 2+ -> Mn2+ + 5Fe3+ + 4H20

Used to determine the %Fe in an unknown compound

KMnO4 is used to titrate an unknown Fe compound which has H2SO4 added as well as 3ml of H3PO4

Q: Why would the KMnO4 solution need to be standardized the same day as it is used to titrate the solution?

If some of the iron in the unknown sample was a ferric ion, rather than ferrous, could I use this same experiment to determine the percentage of iron?

How does the addition of H3PO4 help sharpen the end point in the titration?

Help pls! I dont understand just whats going on in this lab? What is standardization? I've read some articles, but can someone tell me in layman terms?
Is there such things as Chemistry clubs that I can go to? I'm at a college, and we dont have anything...I just need to talk Chem with other people to get whats going on. sigh

Thanks for any help. I'm sure everyone has done this lab at some point... :))

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Offline xiankai

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Re: redox titration!
« Reply #1 on: January 13, 2007, 11:08:54 PM »
Quote
If some of the iron in the unknown sample was a ferric ion, rather than ferrous, could I use this same experiment to determine the percentage of iron?

does the KMnO4 react with the Fe3+?
are u trying to find %Fe or %Fe2+?

Quote
How does the addition of H3PO4 help sharpen the end point in the titration?

look at the equation. H+ is there, but where is it supposed to come from? too little of it and the reaction will not progress to completion, what will happen then?

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What is standardization?

to ensure you know the concentration of KMnO4, which is vital to calculating the %Fe. for example it is labelled 1.0 M. but u got to make sure it is really 1.0 M, because god knows what time can do.

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Is there such things as Chemistry clubs that I can go to?

you're already in here ;)
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Offline Borek

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Re: redox titration!
« Reply #2 on: January 14, 2007, 06:04:16 AM »
Quote
How does the addition of H3PO4 help sharpen the end point in the titration?

look at the equation. H+ is there, but where is it supposed to come from? too little of it and the reaction will not progress to completion, what will happen then?

H3PO4 is too weak to ensure pH low enough for the reaction to proceed, for that we use H2SO4. H3PO4 is added to complex Fe3+ - this complex is colorless as opposed to yellow chloride complexes. Less colors in the solution - easier end point detection.

If the task is to determine total Fe and there is already some Fe3+ present in the solution before titration, it have to be reduced to Fe2+ with SnCl2/HgCl2, but I doubt you will be asked to proceed this way.
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Offline _cheers

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Re: redox titration!
« Reply #3 on: January 14, 2007, 05:18:53 PM »
does the KMnO4 react with the Fe3+?
are u trying to find %Fe or %Fe2+?

% Fe2+
Doesnt the KMnO4 react with the Fe2+ and once its all reacted, its at endpoint at turns pink? But how would it be if it were an Fe 3+ ion? I'm guessing it wouldnt be oxidized and therefor would be no reaction?

If you had a KMnO4 solution left out, would it be reacting with the air? or just degrade? but if the lid is on, how can anything happen? or am I reading too much into it?  Or because its a strong oxidizing agent, there is some kind of reaction going on, but I'm not sure what that would be in a closed container...??
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Offline xiankai

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Re: redox titration!
« Reply #4 on: January 16, 2007, 05:12:11 AM »
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H3PO4 is too weak to ensure pH low enough for the reaction to proceed, for that we use H2SO4. H3PO4 is added to complex Fe3+ - this complex is colorless as opposed to yellow chloride complexes. Less colors in the solution - easier end point detection.

interesting, in my own titrations i never used H3PO4, and the whole class including me had to try and differentiate the various shades of yellow. i'll try to ask my teacher about this, thanks. :)

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Doesnt the KMnO4 react with the Fe2+ and once its all reacted, its at endpoint at turns pink? But how would it be if it were an Fe 3+ ion? I'm guessing it wouldnt be oxidized and therefor would be no reaction?

that's right.

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If you had a KMnO4 solution left out, would it be reacting with the air? or just degrade? but if the lid is on, how can anything happen? or am I reading too much into it?  Or because its a strong oxidizing agent, there is some kind of reaction going on, but I'm not sure what that would be in a closed container...??

substances in the air may dissolve in it, there does not even need to be a reaction.

Quote
If the task is to determine total Fe and there is already some Fe3+ present in the solution before titration, it have to be reduced to Fe2+ with SnCl2/HgCl2, but I doubt you will be asked to proceed this way.

if the iron chloride sample is not dissolved beforehand, it can be weighted and thus % can be determined without need for reduction.
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Offline AWK

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Re: redox titration!
« Reply #5 on: January 16, 2007, 06:42:47 AM »
Quote
If you had a KMnO4 solution left out, would it be reacting with the air? or just degrade? but if the lid is on, how can anything happen? or am I reading too much into it?  Or because its a strong oxidizing agent, there is some kind of reaction going on, but I'm not sure what that would be in a closed containe

1. KMnO4 slowly decomposes on light
2. Glass slowly decomposes KMnO4 to MnO2+K2MnO4
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Offline Borek

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Re: redox titration!
« Reply #6 on: January 16, 2007, 07:04:32 AM »
i'll try to ask my teacher about this, thanks. :)

There is a mixture (called in German literature Reinhardt-Zimmermann reagent) that consist of phopshoric acid (to complex Fe3+), sulfuric acid (to lower pH) and manganese (II) sulfate (to lower redox potential, so that Cl- doesn't get oxidized to Cl2).

Quote
if the iron chloride sample is not dissolved beforehand, it can be weighted and thus % can be determined without need for reduction.

Not necesarilly. Fe(II) chloride is not stable if exposed to air, it gets slowly oxidised. The only reasonably stable Fe(II) salt is Mohr's salt, ferrous-ammonium sulfate, FeSO4(NH4)2SO4.6H2O.
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