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rstar
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[*] posted on 31-10-2011 at 10:34
Science project - Oxidation states


Hi Geeks,
I have been trying to make a science project for my school science exhibition. As, I love Chemistry, i have decided to make a Science project named 'Oxidation State'.

In this project I have to describe oxidation state, its uses, its importance in chemistry and finally some experiments to prove the fact :
"Different elements at different oxidation state can have different properties" ;)

I have collected few compounds of elements at different OS, and they are:
= CuSO4 (Cu at +2 OS)
= CuI (Cu at +1 OS)
= MnO2 (Mn at +4 OS)
= KMnO4 (Mn at +7 OS)
= I2 (I at 0 OS)
= K2S2O5 (S at +4 OS)
= ZnCl2 (Zn at +2 OS)
Sorry, that much only :P

<b>If you can, please give suggestions about :
>Importance and significance of OS
>Uses and problem regarding elements in their different OS.
(Like this fantastic use: Importance of Oxidation state in blood)
>any more compounds as example
>experiments to distinguish their OS </b>




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bob800
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[*] posted on 31-10-2011 at 12:12


You could also use "ferric" (+3) and "ferrous" (+2) compounds, like iron (II) chloride vs. iron (III) chloride. IIRC, potassium ferricyanide will give a light blue solution with iron (II), and a dark blue solution with iron (III) (or vice-versa, I can't remember).
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[*] posted on 31-10-2011 at 12:27


One of the nicest experiments, demonstrating different oxidation states for a single experiment is the following:

http://woelen.homescience.net/science/chem/exps/vanadium/ind...

Here, vanadium is prepared in oxidation states +2, +3, +4 and +5, each oxidation state having its own very distinctive color. The chemicals needed for this demo are easy to obtain (pottery suppliers and other fairly general chemicals).


Another experiment, which shows striking difference in oxidation states is simply making chlorine gas from chloride ions. You can do this by adding some KMnO4 to 20% hydrochloric acid. The bubbles you obtain are pure chlorine gas. Even in a test tube you can see the greenish color of this gas and you can show the exceptionally large difference between chloride ions (such as in table salt) and chlorine gas. Even better is making chlorine from swimming pool 'chlorine' chemicals and hydrochloric acid.



Also very interesting is the use of oxidation states of silver in black and white photography. Ag(+) is reduced to Ag in the gelatin layer of black and white photo-paper. Ag(+) is colorless and Ag is black, when very finely divided. The following web page from my website goes into detail about the process of developing black and white images in photography:

http://woelen.homescience.net/science/photo/developers/index...


The following is an example which I found myself. I exploit the use of oxidation states in the process of making a so-called toner, based on vanadium chemistry.

http://woelen.homescience.net/science/photo/toners/toner.pdf


There are numerous other examples of the use of redox chemistry and the effect of oxidation state on the properties of materials. I just mention some of them, you can use google to find more info on them:

- etching of printed circuit boards for making electronics (uses oxidation of copper(0) to copper(II));
- making indigo dye (used for giving jeans a blue color) and the role of oxidation of indoxyl;
- burning of wood, plants, paper and so on is a redox reaction, with oxygen acting as oxidizer;
- electrolysis of water for making hydrogen and oxygen is a redox reaction at the electrodes of the electrolysis cell;
- use of very fast and energetic redox reactions in fireworks, with chlorates or perchlorates acting as oxidizer and metals, sulphur or other flammable materials acting as reductor.




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rstar
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[*] posted on 1-11-2011 at 06:37


Sorry, no access to Vanadium :(

Which one should be the most stable Iron(II) salt and how to distinguish it from any general Iron(III) salt, like FeCl3




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[*] posted on 1-11-2011 at 07:51


Quote: Originally posted by rstar  
Sorry, no access to Vanadium :(
Have you investigated pottery suppliers? Vanadium is a readily available as a glaze-making ingredient. Find one that deals with raw materials, not just pre-formulated mixes.
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[*] posted on 1-11-2011 at 12:28


Quote: Originally posted by rstar  

Which one should be the most stable Iron(II) salt and how to distinguish it from any general Iron(III) salt, like FeCl3

The sulfate is more stable than the chloride in the +2 state, but the +3 sulfate is not very soluble in water. If you want to do tests in aqueous solutions, I would use the chlorides. They can be prepared by dissolving Fe in HCl (keeping an excess of iron to reduce any Fe+++). For the ferric salt, bubble in some chlorine or simply let it sit in open air to oxidize.
Quote: Originally posted by rstar  

how to distinguish it from any general Iron(III) salt, like FeCl3

As stated previously, ferricyanide/ferrocyanide can be used for this. More info here: http://www.lopezlink.com/Labs/Test_for_Iron_II_and_Iron_III_...

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[*] posted on 1-11-2011 at 23:49


Quote: Originally posted by bob800  
[...] but the +3 sulfate [woelen: of iron] is not very soluble in water. [...]
This is not the case, ferric sulfate is very soluble, but only the hydrated salt is. Anhydrous ferric sulfate hardly dissolves, but it SLOWLY becomes hydrated and while doing so it does dissolve. Just try it. Put some ferric sulfate in water and swirl until you get a turbid liquid with the ferric sulfate finely suspended in the water. Put the test tube aside and let it stand for a day or two days. After that period you'll see that all has dissolved. You can dissolve a lot of ferric sulfate in water, but you need to be patient.

A similar effect exists for many other anhydrous chemicals. Anhydrous NiSO4 and VOSO4 take days to be hydrated and dissolved, but finally they dissolve. CrCl3 is a similar story, it may take weeks or months to dissolve this (unless a strong reductor is present in catalytic amounts), but finally it will dissolve and lots of this material can be dissolved in water.

[Edited on 2-11-11 by woelen]




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[*] posted on 2-11-2011 at 13:07


How about the oxidation states of chlorine? If you have access to some chlorate or perchlorate, you could demonstrate the reduction of chlorine from +5 or +7 all the way down to -1.

I doubt that you have access to dichromates, but those usually go through some nice colour changes when they get reduced.

The cheapest/ easiest/ dirtiest way I can think of for demonstrating redox would be to take a solution of iron(II) acetate (colourless) and pour some H2O2 into it. The iron III acetate formed is blood red and really obvious. If you do not have hydrogen peroxide, you can also try bleach. The down side is that bleach yields a messy precipitate of iron hydroxide.




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[*] posted on 2-11-2011 at 14:54


I got these results with KMnO4 .0001M

[img]C:\Users\Public\Pictures\steve chemistry\2.jpg[/img]



2.jpg - 13kB
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[*] posted on 3-11-2011 at 01:37


Colorful demonstrations of OS in Manganese include:

KMnO4 = Mn in +7 OS, having Purple color
K2MnO4 = Mn in +6 OS, having Green color
MnO2 = Mn in +4 OS, having Black color, insoluble in Water
MnSO4 = Mn in +2 OS, having pale Pink color,

But, can anyone tell me how to make a Mn<sup>2+</sup> salt using KMnO4 as starting chemical ??? :D




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[*] posted on 3-11-2011 at 03:06


Almost any reductor does the job. E.g. dissolve KMnO4 in a solution of sodium hydroxide and add some ethanol. A precipitate of hydrous MnO2 is formed. Filter this precipitate and rinse very well to get rid of almost all sodium ions. Then add this precipitate to hydrochloric acid to which some ethanol is added as well. Heat the liquid. The precipitate dissolves and the ethanol reduces it to nearly colorless Mn(2+). Then evaporate the liquid and drive off all ethanol and acid and what remains is fairly pure MnCl2. If you want to have it really pure, then redissolve the solid in as little as possible boiling water to which one drop of hydrochloric acid is added and allow to cool and allow water to evaporate slowly. You'll get pure MnCl2.4H2O, a pink solid.

If you are happy with a mix of KCl and MnCl2 then things can be much easier. Just add solid KMnO4 to hydrochloric acid to which also some ethanol is added. Allow all of the KMnO4 to be dissolved. A nearly colorless solution is obtained. Boil off the acid and excess ethanol. A pale pink mix of KCl and MnCl2 remains.




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[*] posted on 5-11-2011 at 00:33


Quote: Originally posted by woelen  
If you are happy with a mix of KCl and MnCl2 then things can be much easier.


I will be happy, if it has the same color as MnCl2.4H2O :D

How about oxidation states of Zinc ??
The most 'easy to obtain' zinc salt for me is ZnSO4 in which Zn is in +2 OS. How about its +1 OS ??




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[*] posted on 5-11-2011 at 05:37


Zinc does exist in oxidation state +1. It exists as metal and as divalent ion.

A mix of potassium chloride and manganese chloride is pale pink, just somewhat more pale than pure manganese chloride, due to the white color of potassium chloride.




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[*] posted on 5-11-2011 at 06:08


I did a very nice series by titrating KMnO4(aq) with Na2S2O5 drop-by-drop with a buret. IIRC, I was able to observe, in reverse order, that: 2+ = pink, 3+ = purple, 4+ = brown, 5+ = blue (quite an incredible color, but unstable, and it must be observed/photographed quickly before it is oxidized by air to give...), 6+ which is a brilliant green and, of course, the bright purple pink of permangate, 7+.

I love this one,

O3




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[*] posted on 5-11-2011 at 08:18


How could Zinc in +1 OS be obtained ?

I know one way !
Dissolving Zinc in molten ZnCl2 gives, Zn2Cl2 or Zinc(I) Chloride, containing the divalent Zn<sub>2</sub><sup>2+</sup> ion.

Is there any other way ?
(And by "other way" I mean a process, not containing any molten Zinc(II) halide ;))

[Edited on 6-11-2011 by rstar]




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[*] posted on 6-11-2011 at 08:38


Aluminum compounds can be obtained in +3 OS only. They are white salts, which form colorless aqueous solutions. Is their any test for presence Aluminum ions, and to differentiate it from Zinc(II) ions ??



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[*] posted on 6-11-2011 at 10:09


You could dissolve the soluble Zn(II) and Al(III) salts in water (in separate test tubes), add NaOH solution to both until white precipitates are visible: Al(OH)3 and Zn(OH)2

Now, add excess ammonia (NH3) to both tubes:

You will find the Al(OH)3 will stay visible as a precipitate in the tube.

Zn(OH)2 however, will disappear owing to the displacement of aqua ligands by ammine ligands (formation of [Zn(NH3)4]2+).

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[*] posted on 6-11-2011 at 10:33


Quote: Originally posted by francis  
You could dissolve the soluble Zn(II) and Al(III) salts in water (in separate test tubes), add NaOH solution to both until white precipitates are visible: Al(OH)3 and Zn(OH)2

Now, add excess ammonia (NH3) to both tubes:

You will find the Al(OH)3 will stay visible as a precipitate in the tube.

Zn(OH)2 however, will disappear owing to the displacement of aqua ligands by ammine ligands (formation of [Zn(NH3)4]2+).



Careful...
In both cases, if you add too much base you will form aluminium and zinc ligands [Al(OH)4]- and [Zn(OH)4]-2 respectively; both of which are soluble.




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[*] posted on 7-11-2011 at 04:33


Any example of Zinc(I) ??



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[*] posted on 7-11-2011 at 14:22


Quote: Originally posted by rstar  
Any example of Zinc(I) ??


As far as I know the answer is no.
Cd(I) and Hg(I) are bications with a metal metal bond eg M2 2+ where M is Hg or Cd.
The mercury species is a lot more stable than the cadmium species and I doubt the zinc ion exists at all or only fleetingly.
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[*] posted on 7-11-2011 at 17:05


Quote: Originally posted by ScienceSquirrel  
Quote: Originally posted by rstar  
Any example of Zinc(I) ??

As far as I know the answer is no.


Here is a Wikipedia page, on which I saw something about Zinc(I). I think it does exist.




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[*] posted on 7-11-2011 at 17:22


While not as interesting or dramatic (colorful) as the different oxidation states of the transitional metals, an interesting and very easy to make compound is Sodium Tetrathionate wherein the Sulfur exists in two oxidation states. I only mention this because I remember it as one of my earliest experiments using easy to get chemicals.

http://en.wikipedia.org/wiki/Tetrathionate

Chromium is another possibility, as in Chromate/Dichromate (+6), Chrome Alum (+3). The reason I said Chrome Alum is that it is probably easier to get/crystalize than most Chromium III compounds.
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[*] posted on 8-11-2011 at 00:04


Quote: Originally posted by rstar  
Here is a Wikipedia page, on which I saw something about Zinc(I). I think it does exist.

I just ordered some anhydrous ZnCl2 from an eBay seller and when this arrives, then I'll add some zinc to this and heat the mixture. I never heard of this Zn(I) compound. Isn't it simply zinc metal dissolved in ZnCl2? On the other hand, if it really is yellow, then it hardly can be metallic and then it could be a zinc(I) compound.




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[*] posted on 8-11-2011 at 04:17


Quote: Originally posted by woelen  

I just ordered some anhydrous ZnCl2 from an eBay seller and when this arrives, then I'll add some zinc to this and heat the mixture. I never heard of this Zn(I) compound. Isn't it simply zinc metal dissolved in ZnCl2? On the other hand, if it really is yellow, then it hardly can be metallic and then it could be a zinc(I) compound.


I have never tried it. My Hydrochloric acid is finished, so i couldn't make any ZnCl2 .
If you get the right results then, you can post this experiment in your website, might be a new hit ;)




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[*] posted on 8-11-2011 at 04:41


Quote: Originally posted by rstar  
Quote: Originally posted by ScienceSquirrel  
Quote: Originally posted by rstar  
Any example of Zinc(I) ??

As far as I know the answer is no.


Here is a Wikipedia page, on which I saw something about Zinc(I). I think it does exist.


I stand corrected. I should have looked it up. It is what I would call a spectroscopic species and I suspect that adding water to the yellow glass will result is zinc metal or zinc hydroxide plus a solution of zinc chloride.
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