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Author: Subject: Making Ferric (3) Nitrate solution
CHRIS25
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[*] posted on 22-4-2012 at 06:55
Making Ferric (3) Nitrate solution


When I add Nitric acid (I only have 30%) to Solid Iron I have read that I will get Ferrous nitrate and not Ferric Nitrate. I am getting confused between adding an oxidiser to turn ferrous into ferric, or adding more iron until the reaction stops, or not using chunks of Iron metal but rather the Iron filings.

These are all suggestions from different chemical documents (I'll rule out adding Iron bacteria).

Can someone perhaps point me into the right direction, just want to make Ferric Nitrate.

Thankyou.
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[*] posted on 22-4-2012 at 08:11


You can do this with 30% nitric acid since HNO3 will readily oxidise ferrous to ferric ion, even in dilute solution (ever mixed solutions of FeSO4 and KNO3 and heated? It gives off copious NOx fumes!). Simply add iron to your nitric acid, and calculate the stochiometry such that the HNO3 is reduced to NO2, meaning six moles of HNO3 to one mole of Fe. This ensures that all the ferrous ion gets oxidised, but it will also mean that the solution contains a good amount of excess HNO3.



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CHRIS25
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[*] posted on 22-4-2012 at 08:55


Quote: Originally posted by garage chemist  
You can do this with 30% nitric acid since HNO3 will readily oxidise ferrous to ferric ion, even in dilute solution (ever mixed solutions of FeSO4 and KNO3 and heated? It gives off copious NOx fumes!). Simply add iron to your nitric acid, and calculate the stochiometry such that the HNO3 is reduced to NO2, meaning six moles of HNO3 to one mole of Fe. This ensures that all the ferrous ion gets oxidised, but it will also mean that the solution contains a good amount of excess HNO3.


Hallo garage Chemist, a very big thankyou, that clears everything up.
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CHRIS25
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[*] posted on 22-4-2012 at 11:09


Ok, I am driving myself a bit nuts here, my mathematics is so weak, completely unintelligent on this matter. I take 26 gram Fe and 6 x 63 grams HNO3 which is 378 grams. Good so far. BUT, I only have 30% HNO3 and the 378 assumes a 70% concentration. Because of the water in my acid I need to weigh out much more just under twice the 378 in order to get that correct amount of nitric acid added to 26 grams of FE. I have been rattling my pen and calculator trying top figure how to do this - in a nutshell - and several coffees later I simply can not work it out.
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[*] posted on 22-4-2012 at 16:26


Quote: Originally posted by CHRIS25  
Ok, I am driving myself a bit nuts here, my mathematics is so weak, completely unintelligent on this matter. I take 26 gram Fe and 6 x 63 grams HNO3 which is 378 grams. Good so far. BUT, I only have 30% HNO3 and the 378 assumes a 70% concentration. Because of the water in my acid I need to weigh out much more just under twice the 378 in order to get that correct amount of nitric acid added to 26 grams of FE. I have been rattling my pen and calculator trying top figure how to do this - in a nutshell - and several coffees later I simply can not work it out.


Equation: Fe+6HNO3===>Fe(NO3)3+3NO2+3H2O
378g nitric acid assumes a 100% concentration. So you simply multiplies the 378 by 10, then divide by 3, which would give you 1260g, or 1.260kg of 30% nitric acid to dissolve 56g, not 26g of iron, and give you a yield of 242g of ferric nitrate, 138g of NO2, and 54g H2O.

[Edited on 23-4-2012 by weiming1998]

[Edited on 23-4-2012 by weiming1998]
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[*] posted on 22-4-2012 at 17:23


If you multiply by 3 and divide by 10 you get less than 378. You mean multiply by (10/3).

It's probably good to use excess acid here, because the Fe won't dissolve very quickly in really dilute acid. Also, iron dissolves in cold, dilute nitric acid to form iron (II) nitrate and ammonium nitrate, so you might try heating the mixture or adding hydrogen peroxide to avoid this.
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[*] posted on 22-4-2012 at 22:09


Welming and Barley, thankyou for taking that moment again to help an idiot here. I almost hate even dread to ask...with trembling, where do you get the multiply by 3 divide by 10 idea? I spent 5 complete minutes staring at these figures,I am thinking Am I that stupid? It still makes no sense. I balanced the equation fine but the mathematics of the 10/3 You've lost me. Sorry Welming.

Thanks for the tip about the H2O2 Barley.
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[*] posted on 22-4-2012 at 22:11


Sorry, I made a mistake in my calculation. I have now changed it.

Anyway, although iron nitrate is soluble, if you use concentrated nitric acid (probably the commercial 65% nitric acid would cause this), then it will form a passivating layer around the iron, inhibiting further corrosion. So dilute nitric is better anyway.

You multiply the amount of nitric acid by 10/3 for the calculation because since you need 378g of 100% nitric acid, the amount of 10% nitric that you need is 10 times greater, so 378X10=3780g of 10% nitric acid. Divide that now by 3 because 30% is 3 times more concentrated than 10%, and there's you answer! We first work out the amount of 10% acid that you need because 10 is a factor of both 100 and 30, thus making this easier.

[Edited on 23-4-2012 by weiming1998]
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[*] posted on 22-4-2012 at 22:48


Aahh...thanks Welming, thankyou very much. It must be my age, Ugmm, it only affects my maths though - let's make that clear (lol). So because of the dilution would I still need to add H2O2, not now I presume?


kind Regards
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[*] posted on 23-4-2012 at 00:05


Adding some H2O2 does help. It assures that ALL Fe(2+) is oxidized to Fe(3+) very quickly. But if the acid is hot, then garage chemist's suggestion also works well. Use a large excess amount of acid and keep it hot, or use only a small excess amount of acid, but then you need to add H2O2. Keep in mind though, that the reaction with H2O2 also consumes some acid, albeit less than the reaction in which the acid alone works as oxidizer.

In order to keep things practical, first try on a small scale, e.g. a testtube to get a feeling for the reaction and allow some experimenting without wasting 100's of ml of precious chemicals.

If I wanted to make pure ferric nitrate, then I would go for 30% HNO3, to which some H2O2 (30%) is added. E.g. 5 parts acid mixed with 1 part of H2O2. This mix I would add to iron with weak heating. Use excess amounts of the mix and after all iron has dissolved, boil down in order to destroy all unreacted H2O2 and make the liquid more concentrated. It will be hard though to isolate pure ferric nitrate, because nitric acid will strongly adhere to your product and you almost certainly end up with a strongly acidic product, which still contains quite a lot of HNO3.

Good pure ferric nitrate is very pale brown/purple. It is VERY hygroscopic and if you put a crystal of the solid on a watch glass, then a few minutes later it has turned into a brown/orange droplet of liquefied salt. The brown/orange color is due to partial hydrolysis of the ferric ions. A solution of ferric ions is nearly colorless, the brown/orange and sometimes somewhat turbid solutions, which are so well-known have this color due to hydrolysis of ferric ions. Only at very low pH you will have such colorless solutions.




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[*] posted on 23-4-2012 at 02:37


In several textbooks it says that Fe<sup>3+</sup> ions are actually pale violet. Is this true at high concentration of ferric ion, or does the solution remain colorless?
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[*] posted on 23-4-2012 at 04:49


Yes, the color might be pale violet, but with pale I mean REALLY pale. Only in the solid state or in extremely concentrated solutions one can see this color.

The following may be interesting for you (see the section on iron(III)):

http://woelen.homescience.net/science/chem/solutions/fe.html




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CHRIS25
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[*] posted on 23-4-2012 at 05:47


Quote: Originally posted by woelen  
Adding some H2O2 does help. .............

Hi Woelen, apprecated. I had decided to do some experimenting at test tube size and am certainly glad things have been made clearer concerning the H2O2 and end product. By the way, like your website, especially the Photography section, spent some years developing colour and black and white film in the old fashioned way, actually it's not really old fashioned digital has come full circle in that people now try to re-create the chemical darkroom on their photos - so that says everything eh?, it is still going strong today, not so much colour developing but the black and white darkroom is alive and kicking - it will never become outdated. Now back to the ferric nitrate. thanks

[Edited on 23-4-2012 by CHRIS25]

[Edited on 23-4-2012 by CHRIS25]
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[*] posted on 23-4-2012 at 07:48


@woelen
I read that page before, so I thought my textbook might have been wrong. Thanks for clarifying that. Your website is awesome!

[Edited on 23-4-2012 by barley81]
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