MrHomeScientist
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Oxidizing Acids
I got an interesting question on one of my YouTube videos today that I can't answer, and it made me wonder. Here's the exchange:
Him: Why did you make nitric acid when you already have H2SO4?
Me: Not all acids are equal. Nitric acid is an oxidizing acid that will dissolve many things sulfuric won't, like copper.
It's also used in nitration reactions, which as you might have guessed from the name require the nitrate ions from nitric acid!
Him: I am assuming you're talking about the nitrate NO3- that does the oxidizing? If this is the case, couldn't you use something
like NaNO3(aq) for the oxidizing properties? (just curious)
Now I know sodium nitrate is still an oxidizer, but it doesn't have nearly the power of nitric acid in this regard. Being that they both separate into
ions while in solution, why does the nitrate ion act as a superior oxidant when paired with H+? For example, HNO3 dissolves
copper while NaNO3 (aq) does not. I feel like I'm missing something simple here...
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woelen
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The oxidizing power of nitric acid mainly comes from the fact that at high concentration there still is quite some undissociated acid HNO3. Dilute
nitric acid contains ions H(+) and NO3(-) and this is not strongly oxidizing. Many metals dissolve in this, with formation of H2, due to the reaction
with H(+). In more concentrated acid part of the molecules is not dissociated. This effect is quite noticeable already at 35 to 40% and it is strong
at 60% or higher concentration.
The ion NO3(-) is resonance stabilized, all three oxygen atoms are bonded to the N-atom in the same way and the electrons in the ion are distributed
over the entire ion. This makes the ion quite stable and it only has oxidizing power at high temperatures (such as when used in fireworks).
The perchlorate ion has a similar effect, but even stronger. Perchloric acid is a much stronger acid than nitric acid and even a 70% solution in water
is fully ionized. Hence, such a concentrated acid still hardly acts as oxidizer and a metal like zinc does not reduce the perchlorate ions in such a
solution. Anhydrous perchloric acid, HClO4, does not have the resonance stabilizing effect and this is an extremely strong oxidizer, which even at
room temperature ignites paper, stuff like sugar, sulphur, red phosphorus, metal powder, wood, actually everything that can burn.
So, ionic nitrates and ionic perchlorates are not strongly oxidizing at room temperature, while the free acids are. Other covalent nitrates and
perchlorates (e.g. esters with alcohols) also are much more reactive than similar looking ionic compounds.
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chornedsnorkack
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Isn´t sulphuric acid widely reported to dissolve copper, and be reduced to SO2?
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woelen
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It was (and sometimes still is), but personally, I do not think this is a nice method of making SO2. You need concentrated H2SO4 and you need to heat
the acid moderately strongy in order to get any SO2. H2SO4 has the same effect as HNO3 and HClO4, but it is less pronounced. H2SO4 acts as oxidizing
acid, but you need really high concentration and the acid must be heated.
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AndersHoveland
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Quote: Originally posted by woelen  | | The oxidizing power of nitric acid mainly comes from the fact that at high concentration there still is quite some undissociated acid HNO3. Dilute
nitric acid contains ions H(+) and NO3(-) and this is not strongly oxidizing. Many metals dissolve in this, with formation of H2, due to the reaction
with H(+). In more concentrated acid part of the molecules is not dissociated. |
Concentrated nitric acid contains the oxidizing species NO2+ in equilibrium. The reduction potential of nitric acid is equal to
the reduction potential of nitronium ions. The only difference is reaction rate. Nitrate ions become oxidizing in the presence of hydrogen ions which
are not bound to water molecules. The presence of water weakens the acidity of hydrogen ions considerably, as they actually exist in the form of the
hydronium ion, H3O+.
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papaya
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A question that may not be directly related, but if the dilute acid is less oxidizing then why organic material nitrations go well with strongly
concentrated acid and destruction occurs in case of diluted acid (OK, broad question, but still)?
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woelen
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Can you give an example of total destruction of the organic molecule with dilute nitric acid? I am inclined to believe that dilute (aqueous) nitric
acid does not do anything at all with the usual organics, subjected to nitration. If you are talking about 'dilute' in the sense of a little HNO3 in a
lot of conc. H2SO4, then things may be different. Nearly all of the nitric acid then either will exist as HNO3 or even as NO2(+).
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papaya
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I was about the presence of water, though often nitrations are done in mixture of conc. H2SO4. Well, does RDX make an example in this case? Don't
know, I just had that impression, also heard many times the synthesis of nitrate esters if there's too much H2O present in the nitration mixture
yields fountains of nitrous gases and heating..
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chornedsnorkack
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| Quote: Originally posted by woelen | | It was (and sometimes still is), but personally, I do not think this is a nice method of making SO2. |
Yes, but SO2 is not necessarily the goal - dissolution of copper may be either the goal or else something to be avoides.
| Quote: Originally posted by woelen | | You need concentrated H2SO4 and you need to heat the acid moderately strongy in order to get any SO2. H2SO4 has the same effect as HNO3 and HClO4, but
it is less pronounced. H2SO4 acts as oxidizing acid, but you need really high concentration and the acid must be heated. |
So, does it mean that cold concentrated H2SO4 has no oxidizing effects? It would be inert and safe towards reducing substances like copper, elemental
sulphur, white phosphorus etc. until heated... to which temperatures?
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woelen
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I do not say that cold H2SO4 has no oxidizing power at all, but the oxidizing power of this definitely is less than that of cold conc. HNO3 or
anhydrous HClO4. Sulphur does not react with cold H2SO4, red phosphorus also does not react. I never tried white phosphorus with cold H2SO4.
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JasonHerbalExt
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the action also require H+.
Cu + 4 H+ + 2 NO3- = Cu2+ + 2 NO2↑ + 2 H2O
or
3 Cu + 8 H+ +2 NO3- = 3 Cu2+ + 2 NO↑ + 4 H2O
Hi, I am fond of doing researches on herbal extracts and related drugs. I am happy to discuss with you from scientific groups and would like to share
my topics and views on medical herbs.
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woelen
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The reaction does not require H(+) ion, it requires undissociated HNO3. When all HNO3 were split into H(+) and NO3(-), as written in the above
reaction equations, then no reaction would occur with copper. That is the reason why dilute HNO3 does not react with copper and only gives H2 with
less noble metals.
So, the correct equations would be with HNO3 instead of H(+) and NO3(-). Part of the HNO3 will end up as nitrate ion, the rest as NOx.
[Edited on 27-9-13 by woelen]
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bfesser
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Thread Pruned
29-9-2013 at 06:44 |
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