AndersHoveland
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HNO3(dilute) to make NO2
It is both indirectly inferred several places in the chemical literature that nitric oxide can selectively reduce dilute nitric acid to nitrogen
dioxide. The reaction is:
(2)HNO3 + NO --> (3)NO2 + H2O
other related and better know equilibrium are:
(2)NO2 + H2O <==> HNO2 + HNO3
(2)HNO2 <==> H2O + NO + NO2
Here is some research I found:
"Unlike reduction of nitrite to nitric oxide and of hydroxylamine to ammonium, which have mid point potentials of
+300mV and +800mV, respectively"
the redox potential of the nitric oxide/ nitrite couple is Em>7 = +374 mV (Wood, 1978),
My question is: Is there any other common reducing agent besides nitric oxide that could be used to reduce dilute nitric acid to nitrogen
dioxide, but not any further?
I know nitrogen dioxide can also be made from moderately concentrated nitric acid on copper, but it would be much more convenient is 5% nitric acid,
for example, could be used instead.
You may be wondering what real advantage there would be, since it is well know that nitrogen dioxide can be easily prepared from the reaction of
moderately concentrated nitric acid on copper metal.
(3)Cu + (8)HNO3 --> (3)Cu[NO3]2 + (2)NO + (4)H2O
(2)NO + O2 --> (2)NO2
If the nitric acid is between a concentration of around 20 to 30%, nitric oxide gas will be produced, which then spontaneously reacts with air to form
the brown colored nitrogen dioxide. If 60% nitric acid concentration is used, then nitrogen dioxide is produced directly out of solution. There seems
to be no observable reaction, however, when 4% concentrated nitric acid is used with copper, even after several days. Even the reaction with 20%
nitric acid is somewhat inconveniently slow.
Consider, in the reaction between moderately concentrated nitric acid and copper, that 4 molar equivalents of 30% nitric acid are required to produce
each molar equivalent of nitric oxide.
By subsequently passing this [one molar equivalent] of nitric oxide into a separate dilute solution of nitric acid, three molar
equivalents of nitrogen dioxide can be obtained. And here lies the great advantage. For if the nitric oxide was simply allowed to react with air, only
a single molar equivalent of nitrogen dioxide would be obtained.
It should be noted that the nitric oxide must react with the nitric acid in a separate reaction container, which does not contain copper metal. The
precise reason for this implies substancially higher NO2 solubility, further reaction with the copper, and a complex state of equilibrium. Suffice to
say that the NO2 reaction rate with copper is significantly faster than its reaction with moderately concentrated HNO3, since the action of HNO3 on
copper must rely on the slight equilibrium formation of nitronium ions, NO2[+].
Cu + (4)NO2 --> Cu(NO3)2 + (2)NO
(3)HNO3 <==> NO2[+] + H3O[+] + (2)NO3[-]
If >60% concentrated nitric acid is used, reacting with copper to directly produce nitrogen dioxide straight out of solution, the efficiency can be
relatively improved. Consider that in the reaction using concentrated nitric acid to produce nitrogen dioxide, two molar equivalents of nitric acid
are required to produce each equivalent of nitrogen dioxide.
Cu + (4)HNO3 --> Cu(NO3)2 + (2)H2O + (2)NO2
If it is desired to produce nitrogen dioxide, the reaction of only moderately concentrated nitric acid with copper, and then the subsequent passing of
the resulting nitric oxide gas into a separate solution of dilute nitric acid, will result in much less waste. Comparing the net required reactants to
production of nitrogen dioxide, it can be seen that the last alternative is as efficient in its consumption of nitric acid as the use of concentrated
nitric acid, although 3 times as much copper is wasted.
procedure reacting NO with air, Net Cost:
(3)Cu and (8)HNO3 ==> (2)NO2
procedure directly producing NO2, Net Cost:
(1)Cu and (4)HNO3 ==> (2)NO2
procedure reacting NO with additional separate HNO3, Net Cost:
(3)Cu and (12)HNO3 ==> (6)NO2
Thus if the availability of concentrated nitric acid is much more inconvenient than the availability of moderately concentrated nitric acid, or even
of moderately concentrated sulfuric since a nitrate salt could be added with the same effect, it may likely be more desirable to utilize the special
procedure described in this post.
[Edited on 2-10-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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AJKOER
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I have a suggestion, to the dilute HNO3 add a lot of a nitrate salt, say NaNO3.
Now, it may now be this hot solution of water-stressed HNO3 has a higher 'activity level'. Perhaps, strong enough to directly form NO2 acting like
concentrated HNO3.
------------------------------------------------
An interesting question on the equation:
(2)HNO3 + NO --> (3)NO2 + H2O
and as, you noted:
(2)NO2 + H2O <==> HNO2 + HNO3
One could rewrite as:
2 HNO3 + NO --> NO2 + HNO2 + HNO3
or, cancelling out the extra HNO3:
HNO3 + NO --> NO2 + HNO2
or, multiplying by 2 and adding water to both sides:
2 HNO3 + 2 NO + H2O --> 2 NO2 + H2O + 2 HNO2
and substituting again "(2)NO2 + H2O <==> HNO2 + HNO3"
2 HNO3 + 2 NO + H2O --> HNO3 + 3 HNO2
or, cancelling out the HNO3:
HNO3 + 2 NO + H2O --> 3 HNO2
Interestingly, this reaction is also directly cited by this source, "The Encyclopaedia Britannica: latest edition. A dictionary of arts ...," Volume
5, page 513, link: http://books.google.com/books?id=14FGAQAAIAAJ&pg=PA512&a... , to quote:
"HNO3 + 2 NO + H2O = 3 HNO2"
So, dilute Nitric acid treated with NO can form Nitrous acid. Do you agree with this statement?
If yes, one implied reaction, add some Cu to dilute HNO3 in a closed container (with a safety hole) and shake. Reactions:
8 HNO3 (aq)+ 3 Cu (s) --> 3 Cu(NO3)2 + 2 NO + 4 H2O (l)
with extra HNO3:
9 HNO3 (aq)+ 3 Cu (s) --> 3 Cu(NO3)2 + 3 HNO2 + 3 H2O (l)
This reaction implies adding H2SO4 to a nitrate, diluting with water, and adding Copper could reduce nitrates to nitrites. Clearly, via the HNO2 path,
the reaction is most likely pH and temperature sensitive.
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Adas
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Probably ascorbic acid could reduce HNO3 to NO2. Even I could try it if I wanted (I don't have pure ascorbic acid, though). Another reaction, which
also yields the very useful acetaldehyde, is the reaction of hot HNO3 and ethanol. One could distill off the acetaldehyde after getting the NO2. I
doubt that this reaction would work with HNO3 with concentrations lower than 10%, but 37% H2SO4 and KNO3 are both OTC.
[Edited on 6-4-2013 by Adas]
Rest In Pieces!
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AJKOER
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Here are some interesting related reactions per an educator (see http://books.google.com/books?id=XI44wn5deMAC&pg=SA15-PA... ):
As + 3 HNO3 (dilute) --> H3ASO3 + 3 NO2
So Arsenic may be the answer to your question: "Is there any other common reducing agent besides nitric oxide that could be used to reduce dilute
nitric acid to nitrogen dioxide, but not any further?" Note also the efficiency of this reaction with no nitrogen spent on the formation of a nitrate
salt.
However, there is obvious downside (a trade-off) to employing this element. To quote one source (see http://www.chemicool.com/elements/arsenic.html ):
"Arsenic is immediately dangerous to life or health at 5 mg m-3." and further: "Our bodies do not readily absorb the element itself, hence pure
arsenic is much less dangerous than As(III) compounds such as AsH3 and As2O3 which are absorbed easily and are carcinogenic with high toxicity."
Other reactions cited:
As + 5 HNO3 (Hot Conc) --> H3ASO4 + 5 NO2 + H2O
Now these two Arsenic/HNO3 reactions are only partially confirmed by Wikipedia (http://simple.wikipedia.org/wiki/Arsenic ) to quote:
"It dissolves in concentrated nitric acid to make arsenic acid and in dilute nitric acid to make arsenious acid."
This source (see http://www.transtutors.com/chemistry-homework-help/s-and-p-b... ) appears to confirm also the formation of NO2, but not specific as to
concentration of Nitric acid.
Other related reactions forming NO or NO2 (but none apparently with dilute HNO3):
S + 2 HNO3 --> H2SO4 + 2 NO
3 H2S + 2 HNO3 --> 3 S + 2 NO + 4 H2O
P4 + 20 HNO3 --> 4 H3PO4 + 20 NO2 + 4 H2O
I2 + 10 HNO3 (Conc) --> 10 NO2 + 2 HIO3 + 4 H2O
3 I2 + 10 HNO3 (Hot Conc) --> 10 NO + 6 HIO3 + 2 H2O
6 KI + 8 HNO3 --> 6 KNO3 + 2 NO + 4 H2O + 3 I2
4 Zn + 10 HNO3 (very dilute) --> 4 Zn(NO3)2 + N2O + 5 H2O
Zn + 4 HNO3 (Conc) --> Zn(NO3)2 + 2 NO2 + 2 H2O
Also, per the second source (http://www.transtutors.com/chemistry-homework-help/s-and-p-b... ):
Sn + 4HNO3 → H2SnO3 + 4NO2 + H2O
Zn + 10HNO3 ( very dilute ) → 4 Zn (NO3)2 + NH4NO3 + 3H2O
and it is similarly claimed that Fe and Sn react with dilute nitric acid to give NH4NO3 (as NH3 produced reacts with HNO3).
[Edited on 11-5-2013 by AJKOER]
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Adas
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Quote: Originally posted by AJKOER  | 4 Zn + 10 HNO3 (very dilute) --> 4 Zn(NO3)2 + N2O + 5 H2O
Zn + 4 HNO3 (Conc) --> Zn(NO3)2 + 2 NO2 + 2 H2O
[Edited on 11-5-2013 by AJKOER] |
What??? Zinc liberates hydrogen when it reacts with acids. It is a non-precious metal! You also wrote N2O instead of NO2. Please, take more care not
to post bull****.
Rest In Pieces!
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woelen
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This time AJKOER might be right. Zinc is capable of releasing N2O from nitric acid when it is fairly dilute. It, however, is not a clean reaction,
there are other competing reactions which occur at the same time. One of them is simply formation of H2. Another is formation of NH4(+).
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AJKOER
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Per my second cited educational site (http://www.transtutors.com/chemistry-homework-help/s-and-p-b... ), to quote:
" i) Very active metals such as Mn, Mg, Ca, etc. give H2 on treatment with very dilute HNO3 (2%).
(ii) Less active metals like Cu, Hg, Ag, Pb etc. give NO with dil. HNO3 . Zinc, however, gives N2O with dil HNO3 and NH4NO3 with very dilute HNO3.
Zn + 10HNO3 ( dilute ) → 4Zn(NO3)2 + N2O + 5H2O "
while the first source states (topic page 15.12 at http://books.google.com/books?id=XI44wn5deMAC&pg=SA15-PA... ):
Zn + 10HNO3 ( very dilute ) → 4Zn(NO3)2 + N2O + 5H2O "
for the same reaction. Now, I did indicate (warn) readers that these reaction results are published by educators and sometimes, per my experience,
there are quality issues. It could be that I may miss some of these (like the example above in terms of dilution of HNO3 and Zn noted above for N2O
formation although this reaction appears to be more complex and perhaps commenting on it is for the better). However, I think it would be more
productive if you (Adas) presented any potential negative feedback for general review on this thread. Then, if confirmed and significantly serious,
perhaps we could politely contact the educator so that in future editions, any errors can be addressed given the power of the internet to educate.
-----------------------------
IMPORTANT
Came across an old reference remarking on how slowly Arsenic dissolves in HNO3 except in the presence of some HCl. The reaction is reported to
resemble the action of HNO3 on pure Iron (see "Journal of the Society of Chemical Industry", Volume 25, pages 1074 to 1076 at http://books.google.com/books?id=BBQAAAAAMAAJ&pg=PA1076&... ). I am also not clear, in the case of dilute solutions, as to whether the formed
NO2 actually escapes the solution, or simply forms HNO2 and HNO3.
2 NO2 + H2O = HNO2 + HNO3
Such reactions, of course, would be a disappointment for those seeking NO2 gas. To derive a new view of the reaction, rescale the original equation by
say 4 and add 6 H2O:
12 HNO3 (dilute) + 4 As + 6 H2O ----> 4 H3AsO3 + [12 NO2 + 6 H2O]
But the brackets quantity equals 6HNO2 + 6 HNO3, so now on net:
6 HNO3 (dilute) + 4 As + 6 H2O ----> 4 H3AsO3 + 6 HNO2
or:
3 HNO3 (dilute) + 2 As + 3 H2O ----> 2 H3AsO3 + 3 HNO2
and hence the articles possible reference to the formation of nitrous acid (and just NO, and not NO2). This would also be consistent with the limited
products citation by Wikipedia (no mention of NO2), but this new reaction equation, based on research gathered to date, is my opinion on the best form
to present the reaction. Note, if one understands the considerable presence of water, the original cited reaction isn't technically wrong, just most
likely not observed, assuming conditions favorable to the stability of HNO2 (including cold dilute solutions).
[Edited on 11-5-2013 by AJKOER]
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Adas
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So, if this is true, how to prove the presence of N2O easily? (e.g. at home, maybe?)
What does "dilute" and "very dilute" mean? "Dilute" can be anywhere from 50% to 10%...
Rest In Pieces!
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AJKOER
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| Quote: Originally posted by Adas | So, if this is true, how to prove the presence of N2O easily? (e.g. at home, maybe?)
What does "dilute" and "very dilute" mean? "Dilute" can be anywhere from 50% to 10%... |
An easy test for the home, to quote Wikipedia on N2O:
"Nitrous oxide supports combustion by releasing the dative-bonded oxygen radical, thus it can relight a glowing splint. N2O is inert at room
temperature and has few reactions."
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AJKOER
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| Quote: Originally posted by AJKOER | I have a suggestion, to the dilute HNO3 add a lot of a nitrate salt, say NaNO3.
Now, it may now be this hot solution of water-stressed HNO3 has a higher 'activity level'. Perhaps, strong enough to directly form NO2 acting like
concentrated HNO3.
.....
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A possible source for the above suggestion (see "Photodecomposition of Nitrite and Undissociated Nitrous Acid in Aqueous Solution", J. Phys. Chem.
1996, 100, 18749-18756, link: http://www.google.com/url?sa=t&rct=j&q=j.%20phys.%20... ), to quote:
"It appears that a reaction of nitrate with one of the other reactants present in the solution would be required. As an example, we have considered
the reaction:
NO + NO3- ---> NO2 + NO2-
which we estimate to be endothermic by about 40 kJ mol-1 and hence not very convincing. However, curve e in Figure 1, which was obtained by setting
k17 ) 4 104 dm3 mol-1 s-1 and neglecting reactions 11b and 13, achieves a better agreement with the experimental data than any of the other
assumptions."
To be fair, the authors do not endorse my possible explanation, but nevertheless comes to the reaction cited above as a best fit to the data. However,
if adding a cheap available nitrate salt to your dilute Nitric acid acting on Cu, for example, makes it act like the strong stuff with respect to NO2
formation, go for it!
Note, I have not yet tested this reaction mechanism myself, and it may not be an ongoing lower cost alternative unless one views the available nitrate
salt as a sunk cost. The latter, I would define in the current context, as an unrecoverable past expenditure not considered in the current cost
calculation, as opposed to investing in new, and more concentrated, Nitric acid.
[Edited on 18-5-2013 by AJKOER]
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