School textbooks provide detailed information on the ways in which Nitrous Acid is extracted. Its physical and chemical properties, so do not go beyond them to discuss some of the features of nitrogen acid. Colorless fuming nitric acid or concentric to the light gray color is caused due to the nitric acid degradation caused by NO2 in solution.
4HNO3 = 4NO2 + O2 + 2H2O.
In the conditions of warming, the rupture can take place even deeper.
2HNO 3 = N2O + 2O2 + H2O,
4HNO 3 = 2N2 + 5O2 + 2H2O.
Self-adhesion of nitrogen acid takes place in concentrated solutions.
The nitronium ion produced during dissociation is due to the strong oxidative activity of nitrogen acid and the ability to notify many organic compounds. Many aromatic dishes are colored, which explains the fact that the phenol group-containing proteins are yellowish at contact with nitrous acid, such as yellow spots on the skin falling on nitrogenous acid.
Uses of Nitrous Acid:
It should be noted that in nitration it is commonly used not pure nitrogenous acid, but a mixture of nitrogen –HNO3 and H2SO4. Increased sulfuric acid leads to increased concentration of the nitrous ion.
Nitrous Acid Properties:
Nitrogenous acid has high oxidative properties, in terms of thermodynamics, HNO3 can be recovered from various oxidation degrees of nitrogen.
10HNO3 + 3I2 = 6HIO3 + 10NO + 2H2O,
6HI + 2HNO3 = 3I2 + 2NO + 4H2O,
FeS + 12HNO3 = Fe(NO3)3 + H2SO4 + 9NO2 + 5H2O.
When combined with metals and nitrogen oxide, a complex mixture of ingredients is formed, whose composition is mainly determined by nitrogen acid density and metal nature. The more metal is active and the less nitrogenous acid it is, the deeper its recovery is.
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In fact, when all thermodynamically permissible HNO3 regeneration reactions are controlled, they are controlled only by kinetics and proceed simultaneously. During the reaction, HNO3 concentration, temperature and hence the composition of HNO3 ‘s regeneration products are changing.
The process of this or the direction of the process is determined by the difficulty of the gas phase in the solution and the solubility of N2, N2O, NO and NO2 in the reaction mixture with the potential change in the reaction.
Nitrogen (3-20%) nitrogenous acid, depending on the nature of the metal, can be restored to the water-soluble gases N2, N2O, NO, H2 or ammonium ion. It has been proven that during the interaction of manganese and magnesium and 10-20% nitric acid, a gaseous mixture (N2, N2O, NO, and H2) in which the hydrogen content is 80%, that is, it can be said that metals interact with sparse nitric acid in the following prevailing reaction:
M + 2H + = M 2+ + H2
Other metals, such as zinc, interact with the prevalence of
N 2 O or N2 5Zn + 12HNO 3 (10%) = 5Zn (NO3) 2 + N 2 + 6H 2 O.
Very oily (less than 5%) nitrous oxide restores nitrate ion to NH4+.
The reaction takes place in a few days. Ammonium ion produces a solution in the solution, so the metal salts are low oxidation levels.
4M + 10HNO 3 (3%) = 4M (NO3) 2 + NH4NO 3 + 3H 2 O. (M = Sn, Fe).
Nitric acid solutions with moderate density (20-60%) are recovered with metals mainly by NO.
4HNO + Fe 3 (30%) = Fe (NO3) 3 + 2NO + 2H 2 O, 3Hg + 8HNO 3 (30%) = 3Hg (NO 3 ) 2 + 2NO + 4H 2 O.
Concentrated nitric acid solution (60% higher) oxidic is nitronium ketone that recovers NO2’s.
So all metals, regardless of their activity, interact with dense nitrogen acid, restore it to nitrogen dioxide.
M + 4HNO 3 (68%) = M (NO3) 2 + 2NO 2 + 2H 2 O (M = Mg, Zn, Cu, Hg, etc.).
Cause nitrogen oxide IV is the fact that the degree of oxidation of all nitrogen compounds votanavor (below + 4) concentrated nitric acid to oxidize to NO2,
2HNO 3 + NO = 3NO 2 + H 2 O.
It should be noted that the concentration of metal in the concentrated nitric acid solutions and the surface activity of the metal and the salinity solubility in nitric acid are affected. Some metals (Fe, Cr, Al), nitrogen acid, at room temperature, are degraded by oxidized protective membranes.
2Al + 6HNO 3 = Al 2 O 3 + 6NO 2 + 3H 2 O.
Other metals (Ca, Ba), due to the extremely low solubility of nitrates in dense nitrogen acid, do not interact practically. In general, nitrogenous acid can interact with all metals except under certain metals of gold and platinum group, although nitric acid oxidizing properties are so strong that thermodynamically possible nitrogenic acid can even oxidize gold and platinum ( of course, very slow, almost invisible).
It is interesting to compare the oxidizing properties of the nitrogen and nitrogenic acids of the same concentration, and at first glance, it would seem that nitric acid is a stronger oxidizer because it contains +5 oxides of the nitrogen atom, but that’s not true. It turns out that HNO2 is a stronger oxidizer than HNO3 in rare solutions of the same concentration (0.1 M).
Thus, 0.05 M nitrogenous acid produced in sodium nitrite sulfate solution instantly oxidizes potassium iodide.
2NaNO3 + 2KI + 2H2SO4 = I2 + 2NO + K2SO4 + Na2SO4 + 2H2O.
And the nitrogenous acid of the same concentration does not interact with KI. Oksidich high-quality nitrous acid is kinetic in nature and can be explained by HNO2 at the oxygen atom of nitrogen atom weaker ekranatsumov (NO3- in comparison), in which a nitrogen atom HNO2 at more affordable reducing the influence of NO3- in. Also plays a role in the high thermodynamic stability of its decomposition products in relation to the thermal stability of nitrogen acid.
3HNO2 = HNO3 + 2NO + H2O.
In addition to oxidizing properties, nitric acid may also exhibit unusual behavior, such as nitric acid in the transparent solution of silver chloride ammonia complex, and then the white precipitate of silver chloride.
[Ag (NH3)2]Cl + 2HNO3 = AgCl + 2NH4NO3 .
Here, nitrogenous acid is acting in a non-specific role. it is not an oxidizer, but rather accelerates the collapse of the complex. In oxidative processes, nitrogenous acid can be an oxide but not an environment, for example.
2Mn(NO3)2 + 5PbO2 + 6HNO3 = 2HMnO4 + 5Pb(NO3)2 + 2H2O.
In the reaction, the oxidizer is PbO2.
It is important to address the thermal stability of nitrous acid salts – nitrates. As a rule, because of the crystal growth due to the Coulomb interaction, salts are more stable than their corresponding acids.
Nitrates, such as alkali, alkaline metals, and ammonium nitrates, are melted without any degradation (at high temperatures they are degraded), not to the general rule. Traditionally, it is believed that alkaline metals are degraded as a result of nitrite and oxygen.
In fact, nitrite is produced only during potassium nitrate decomposition
Table 1 Melting and decomposition temperatures of alkali metal nitrates and nitrites Liquid nitrates decompose in Li2O, while sodium nitrate breaks down in two directions at once.
It is also possible to have Na2O2 and nitrogen oxides. The second and third groups are eroding the metal nitrate metal oxide, NO2 and O2 in (their stable Nitrites 230 0 lower than C-).
Oxides also emerge during the decay of transitive and reducing metal nitrates, with the exception of mercury and silver nitrates that meltdown to metal as the nitrites and oxides of the metal at the degradation temperature are unstable.
If the nitrates in the low oxidation level are degraded, the oxygen oxidizes to the metal.
It follows from this that nitrates should not be linked with the activity of metals (traditionally coming).
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