Nitration of analine with the usual mixture of concentrated nitric/sulfuric acids is very problematic, because of both charring, and oxidation of the analine to nitrobenzene
(using a large excess of sulfuric acid helps minimize formation of nitrobenzene). It is true, however, that low yields of 3-nitroanaline (without significant formation of other isomers) can be obtained, however. Even the use of 60% nitric acid can result in some charring of analine, and this can be partially reduced by cooling the nitration bath. With 75% concentrated nitric acid added to excess analine, charring takes place fairly quickly.
Adding an acetyl group to the analine, using acetic anhydride, before nitration will make the reaction run much smoother without the byproducts, the acetyl group later being hydrolyzed off with a mildly strong base, but in this case mostly 4-nitroanaline is produced, with some lesser formation of 2-nitroanaline.
A little known fact is that nitrogen dioxide actually reacts at room temperature (20degC) with toluene to form phenylnitromethane, where a nitro goup is actually added to the methyl group of toluene. "Phenylnitromethane has been prepared by the nitration of toluene with dilute nitric acid in a sealed tube."
Konowalow, Ber. 28, 1860 (1895). (the sealed tube probably implies heating) At higher temperatures, addition of two nitro groups on the same carbon predominates. "nitration of toluene with nitrogen dioxide at a temperature between 20C to 95C yields a mixture of phenylnitromethane and phenyldinitromethane"
Thus the reaction of the nitric acid with the toluene over a longer period of time might be through the partial decomposition of the nitric acid into nitrogen dioxide. One would also expect the moderately concentrated nitric acid to very slowly attack/oxidize the toluene, not only producing small quantities of 2-nitrotoluene, but also other oxidation products together with nitrogen dioxide. The nitrogen dioxide thus formed could then react with more toluene. Of course, unless the nitric acid was extremely concentrated, the phenylnitromethane would immediately hydrolyze under the acidic conditions (through a Meyer reaction) to form benzoic acid, the hydroxylamine simultaneously formed would no doubt immediately be oxidized by additional nitric acid, since at the 60%+ concentrations it is a reactive oxidizer. Fairly complex reaction dynamics. I do not really know what your orange byproduct is, likely a mix of different compounds, which may possibly include 2-nitrobenzoic acid (this is yellowish-white in color so would not explain the brownish-orange color).
Nitrous acid (NaNO2 and acid) can be used to nitrate phenol or analine. If concentrated mineral acids are present with the nitrous acid, mostly the para-nitro will be produced (91% yield), but if glacial acetic acid (concentrated) is present with the nitrous acid, yields of 74% ortho-nitro are possible. Ortho means the nitro is in the 2-position, adjacent to whatever other group is on the ring. Para means the nitro is on the opposite end of the ring from the other group.
The theory is that a nitrosyl ion (NO+) initially nitrates the ring, and then the resulting nitroso group gets oxidized to a nitro. ("Aromatic nitration", K. Schofield. 1980)
However, this nitrosyl ion equilibrium with nitrous acid is only present under strongly acidic conditions, certainly not when nitrous acid is in an alkaline envirorment., where it is mostly in the form of nitrite.
One source mentions that reaction of analine with nitrosonium hydrogen sulfate NO[+] [-]SO4H produces nitrobenzenediazonium salts.
First, the nitro group is added to the analine. It is only after the acids have been somewhat diluted with water that the amino group can be diazotized, as diazotization cannot take place in extremely concentrated sulfuric acid.
"Interaction of Analine with diazotizing and nitrating Agents in Concentrated sulfuric Acid"
Mikhail V. Gorelik, Vera I. Lomzakova, Elena A. Khamidova
Research Institute of Organic Intermediates and Dyes, Moscow, Russian Federation
nitrosonium hydrogen sulfate is the same thing as nitrosyl sulfuric acid, which may be prepared by bubbling mixed nitric oxide/nitrogen dioxide gases into concentrated sulfuric acid.
(3)H2SO4 + NO2 + NO --> (2)NO[+]SO4H[-] + H3O[+]SO4H[-]
Whereas 3-nitroanaline can be easily diazotized with dilute nitrous acid, both 2-nitroanaline and 3,5-dinitroanaline require moderately concentrated sulfuric acid to be present, for formation of nitrosonium ions.
Treatment of a 1:1 ratio of nitrosylsulfuric acid to analine, using 92.5% concentrated sulfuric acid, leads to formation of about three times as much 4-nitrodiazonium salt as 3-nitrodiazonium salt. Using 95.5% sulfuric acid leads to a 6 to 4 ratio of the isomers, while 100% concentration leads to equal formation of both.
The treatment of previously prepared benzenediazonium salt with nitric acid does add any nitro group.
Apparently, 3-chloronitrobenzene can be prepared by the electrophilic chlorination of nitrobenzene with chlorine in the presence of ferric chloride FeCl3.
Both 2- and 4-chloronitrobenzene react with anhydrous ammonia at 200degC to form the corresponding nitroanaline, whereas 3-chloronitrobenzene did not react under these conditions.
In the presence of the iodide ion, dry NH3, dissolved in pure alcohol, reacts rapidly with 2- and 4-chloronitrobenzene at 100degC. 1-chloro-2,4-dinitrobenzene reacted with alcoholic ammonia at room temperature. V. A. Tarasevich, M. F. Rusak, A. B. Tereshko, and N. G. Kozlov, Zh. Obshch. Khim, 67, 457 (1997); Chem. Abstr., 128, 270275r (1998)
So basically, nitrobenzene would have to halogenated with bromine in the presence of ferric bromide FeBr3 (in the absence of water obviously) to 3-bromonitrobenzene, which could then be reacted with anhydrous ammonia at room temperature to form 3-nitroanaline, since bromine substitutes off much more easily.
The 3-nitroanaline is an important precursor for the preparation of diazoperchlorate-3-nitroaniline.
By using 3-4 parts (by weight) sulfuric acid to 1 part nitric acid, and a ratio of 1 molecule of benzene for every 2 molecules of nitric acid, and keeping the temperature under 80degC, it is possible to obtain a dinitrobenzene product which consists of 85% of the 3- dinitrobenzene isomer.
263 parts by weight 85% sulfuric acid and 175 parts benzene. 1315 parts (again by weight) of a previously prepared mixed acids (22% nitric acid and 71.5% sulfuric acid) are gradually added in small portions over the period of 90 minutes, maintaining the temperature at around 50degC. After all the mixed acid is added, the contents are heated to 80degC and maintained at this temperature for 40 additional minutes. After completion of the reaction, 293 parts by weight of water are cautiously and gradually added, taking care to avoid splashing caused by the heat generated from the addition of water to the concentrated acids. The reaction mixture is allowed to settle until it separates into two layers, the top layer contains the dinitrobenzene and is decanted off. The crude product still contains some dissolved nitric acid. The dinitrobenzene consists of about 95% of the 3- (meta- ) isomer. The temperature during the reaction should not be allowed to rise above 80degC, as this will lead to formation of mostly the 2- and 4- isomers.
The 3-dinitrobenzene can then be partially reduced with a limited quantity of zinc and dilute hydrochloric acid to 3-nitroaniline.
lead acetato-bromate Pb(C2H3O2)2(BrO3)2 precipitate deflagrates with a yellowish-white puff of smoke. hammer drop height sensitivity: 15-20cm. 175 g of KBrO3 are dissolved in 1.5 L hot water, and mixed with a solution of 175 g acetic acid (100%) and 260 g lead acetate hydrate in 2 L water. The solution remains clear initially. It is then filtered, cooled down and seeded with a few crystals of lead bromate while rubbing the side of the glass vessel with a glass rod. Soon after, the solution becomes turbid, and yields a heavy crystalline precipitate, which does not increase anymore after 12 hours in the cold. The solution above the precipitate is decanted off or extracted out, and discarded (possibly add some Na2CO3 to salvage the remaining lead in the waste solution, in the form of PbCO3). The crystalline precipitate is washed with cold water until it is free of acetic acid and Na/K acetate/bromate, then is hot dry air is passed over until free from moisture. Yield is only 123 g
NaNO3 + HCl
When a mixture of concentrated hydrochloric acid (32%) is mixed with some sodium nitrate and heated, brown fumes begin to be given off. This mixture is capable of slowly dissolving gold, the reaction is more rapid when heated.
It is known that there exists the following equilibrium in aqua regia:
HNO3 + (3)HCl <==> NOCl + (2)H2O + Cl2
In a mixture of muriatic acid and sodium nitrate, perhaps there is also an equilibrium, although only very small?
(2)NaNO3 + (4)HCl <==> (2)NaCl + (2)H2O + Cl2 + NO2 + NO
Theoretically, if the reaction could be made much more acidic, the equilibrium would probably shift more to the right. The reaction would not really be driven much towards the right until acid concentrations of around 70%, which is not possible for muriatic acid, since only so much HCl (which is a gas in the pure form) can dissolve in water.
One would wonder whether the sodium nitrate in the concentrated muriatic acid solution could act as a catalyst, allowing some of the hydrochloric acid to be oxidized by the air:
(4)HCl + O2 <==> + H2O + Cl2 ?
(since the oxides of nitrogen would react with oxygen to form nitric acid)
One would think that this reaction would have already been described in this forum, but a search did not reveal anything. Does anyone know what is happening in this reaction?
It is known that mixtures of nitric oxide and nitrogen dioxide together can attack ammonium salts. Perhaps letting a little ammonium nitrate dissolved in hydrochloric acid (32%conc) sit around for 24 hours will answer the question. If there is indeed any equilibrium, a portion of the ammonium nitrate should react, with the net equation:
(3)NH4NO3 + (6)HCl --> (9)H2O + (3)N2 + (3)Cl2
will do experiment and let everyone know how it went. If there is no reaction, than it would strongly suggest that there is no chlorine or nitrogen oxides in equilibrium.
(typically NCl3 will not form under strongly acidic conditions, but one should still be careful)
Dissolving Gold with Concentrated Nitric and Sulfuric Acids
Mixtures of manganese dioxide and sulfuric acid can actually dissolve gold!
The reaction is slower at room temperature, but rapid with heating.
Permanganate and sulfuric acid after a few minutes also dissolve gold.
Even a hot mixture of concentrated nitric and sulfuric acids can dissolve gold, with lower oxides of nitrogen forming. Addition of water caused the gold to precipitate back out in metallic form, but if a solution of permanganate is used instead, the gold remain dissolved.
Reynolds, later by Spiller
Chemical engineering, Volume 2, p316
The text mentions that even concentrated mixtures of nitric and phosphoric acid attacks gold at room temperature, although the reaction is very slow unless heated.
The reaction is probably:
(5)Au + (3)NO3[-] + (18)H[+] --> (5)Au[+3] + (3)NO[+] + (6)H3O[+]
The nitronium cations, NO2[+], which form in equilibrium in concentrated nitric acid solutions, probably initially attack the gold, creating nitrogen dioxide. Basically,
Au + (3)NO2[+] --> Au[+3] + (3)NO2
The nitrogen dioxide produced would likely remain in the concentrated acid,
(2)NO2 + (3)H2SO4 --> NO[+]HSO4[-] + NO2[+]HSO4[-] + H3O[+]HSO4[-]
and the nitronium ions formed from the NO2 would then attack more gold.
But ozone with sulfuric acid did not dissolve gold.
It is already known that aqueous solutions of bromine or chlorine dissolve gold.
Excess sulfuric acid needs to be used. This is an equilibrium reaction, and the gold is not going to dissolve easily. The mixture needs to be extremely acidic. Even a 1:1 ratio of 70% HNO3 to 95% H2SO4 is not going to be concentrated enough. For good results, use a 1:10 rato of 70% nitric acid to 98.5% concentrated sulfuric. Essentially, there can be no water in the reaction!
Even in the hot boiling mixed acids, the gold takes several minutes to dissolve.
The NO[+] ion hydrolyzes (reacts with) water to form nitrous acid.
NO[+] + (2)H2O --> HNO2 + H3O[+]
Nitrous acid is fairly reactive, and can act as either a reducing or an oxidizing agent. It will reduce the dissolved gold (Au+3) to elemental form (Au). This explains why the gold precipitates back out when the reaction is diluted with water.
(2)Au[+3] + (3)H2O (3)HNO2aq --> (2)Au + (6)H[+]aq + (3)HNO3aq
(note that "aq", which stands for "aqueous", means it is dissolved in water)
If fuming nitric acid is added to the reaction containing the dissolved gold, the gold will solidify out as a purple solid. The gold is probably still in its elemental form, but small particle sizes of gold are known to exhibit strong colorations, from red to purple.
Sorry for all the chemistry information, but some of you will undoubtedly be curious about the specifics of the reaction.
Nitrous acid is unstable, and only exists in the form of solutions which gradually degrade after several minutes. Solutions of nitrous acid exist in equilibrium with nitrogen dioxide and nitric oxide, the latter of which is an unstable radical which can either react with the oxygen in air to form more nitrogen dioxide, or if left on its own will disproportionate into nitrogen dioxide and nitrous oxide after several minutes.
(2)HNO2 <==> H2O + NO2 + NO
(3)NO --> N2O + NO2
In the reaction,
(2)Au + (3)NO3[-] + (18)H[+] --> (2)Au[+3] + (3)NO[+] + (6)H3O[+],
sulfate ions are not shown because they do not directly take place in the reaction. The literature even states that phosphoric acid can be used in place of the sulfuric acid.
The above reaction is in ionic form. Some of you may prefer to see it in the form:
(2)Au + (3)HNO3 + (15)H2SO4 --> (2)Au(SO4H)3 + (3)NOSO4H + (6)H2SO4*H2O
Note that the "Au(SO4H)3" only exists in the solution, it cannot be isolated. Gold trinitrate, if it even exists, would also be nearly impossible to obtain as a pure solid. Gold trinitrate only exists in highly concentrated solutions of nitric acid. When these solutions are diluted with water, auric oxide precipitates out. Similarly, auric oxide only only dissolves in very concentrated acids, since it is only very weakly basic.
Au2O3 + (9)HNO3 <==> (2)Au(NO3)2 + HNO3*H2O
The reaction is more interesting from a chemical perspective than a practical way to refine out gold. Nevertheless, the reaction may be useful to directly dissolve gold-silver alloys, without having to go to the trouble of inquartation, since aqua regia only dissolves such alloys with extreme difficulty.
Yes, it is extremely dangerous. The dangers of using concentrated mixed acids are commonly taken for granted among those that frequently perform nitrations. Obviously those unfamiliar with such procedures should think twice before handling such high concentrations of acid.
More details about the reaction. The concentrated acid mix that contains the dissolved gold should be gradually transfered into the larger bowl of water using a 10mL glass transfer pipette. You will also need a rubber pipette suction bulb. For those of you unfamiliar with this tool, it is basically like a turkey baster that is used to suck up a small quantity of liquid, then move it to another container. The pipette can be bought here:
Using the pipette to slowly add the acid mixture to the water is important for two reasons. First, safety. Water should never be added to concentrated acid, since this can result in the acid spraying up. Neither should the acid be poured into the water, because of the possibility of an accidental spill or splashing, and because it can be hard to control the rate that the liquid is poured in. Adding the acid in too fast can lead to overheating, which could result in boiling/splashing in the water. Second, it is important that each small portion of the acid quickly be diluted with as much excess water as possible. This will help prevent the gaseous nitrogen oxides (NO and NO2) from escaping. Although nitrosylsulfuric acid reacts with excess water to form a solution of nitrous acid, if not enough water is used nitrogen oxides will bubble out instead.
There will inevitibly be some loses of nitrogen oxides, in the form of some bubbling and some brown gas being given off. Unfortunately, when some of the nitrogen oxides escape, there will not be enough nitrous acid to completely reduce the gold. After neutralizing, all the gold will still precipitate out, but a small portion of it will be in the form of hydrated gold oxide, Au2O3. If the gold is going to later be melted, the gold oxide should not pose any problems, as the compound decomposes to the pure metal at 160°C, giving off oxygen gas.
One other note of warning, unless the gold oxide has been completely reduced, it should not be reacted with ammonia, as this will form the dangerous sensitive explosive known as "fulminating gold". In the event that the acid solution was previously boiled with ammonium sulfate to prevent precipitation of the gold, fulminating gold can result upon neutralization if too much ammonium sulfate was added.
More safety information:
Only use small quantities of mixed acids at a time. Be aware that with concentrated acids, even tiny drops can splash out and result in painful burns on exposed skin. To get some understanding of these dangers, try pouring cranberry into a glass, wearing a clean white long-sleaved shirt. Even with cautious pouring, you are likely to find one or two tiny little red stains on the sleeves afterwards, even though you were not aware of any splashing while the juice was being poured. If this was concentrated acid, painful burns would have been felt.
You may desire to cover your shoes with a plastic bags and a rubber band, so that if any of the acid spills onto the floor, it will not seap into your shoes. Protective shoe coverings can also be purchased:
If you choose to wear rubber boots instead, it is advised that the top of the rubber be tied tight around your legs, so that if any of the acid is spilled on you, it will not drip down into the boots and collect in a puddle. If the acid is in contact with your skin for more than a few seconds, the burns will be much more severe. http://www.amazon.com/b?ie=UTF8&node=393294011
A boiling mixture of concentrated nitric and sulfuric acids is extremely dangerous, much more so than 70% concentrated sulfuric acid, for example. The chemistry of this mixture presents several unique hazards. Extremely concentrated sulfuric is a strong dehydrating agent, that will turn anything organic, such as a strip of paper or your skin, into black char immediately on contact. A note about treating concentrated nitric acid burns, after you immediately rinse the affected area with plenty of water, and neutralize with sodium bicarbonate solution, there is special recommendation for concentrated nitric acid burns. Use a swab dipped in chlorine bleach to gently scrub the affected area. Some of the yellow color from the burn should be absorbed onto the cotton swab. Continue to scrubbing with fresh swabs until no more yellow can be absorbed onto the cotton. Then rinse well in soapy water. Doing this will help remove some of the nitro compounds which have formed. These compounds act as allergens and greatly slow the healing process. In fact nitric acid burns take much longer to heal than sulfuric acid of the same concentration. The unique effects of concentrated nitric acid are due to the formation of nitronium ions, NO2[+], in equilibrium in the solution. The addition of highly concentrated sulfuric acid greatly enhances this equilibrium, and so the special burn effect of nitric acid will be greatly exaggerated by the acid mixture. In other words, it would be very important to treat the burns in the way described above, and the healing time is likely to be much longer.
Dissolving Gold with Manganese Dioxide
about the reaction with manganese dioxide and sulfuric acid. The reaction is probably:
(2)Au + (3)MnO2 + (3)H2SO4 --> Au2O3 + Mn(SO4)2 + (2)H2O
where Mn(SO4)2 is manganese sulfate, and the gold oxide dissolves in the sulfuric acid. concentrated sulfuric acid still needs to be used, but it probably does not need to be quite so concentrated as required for the other reaction; a 70% concentration should be suitable.
One of the posters at "http://goldrefiningforum.com" stated that the "wet ash" method did indeed dissolve gold if the acids were concentrated enough, although he wrote that it was not a practical method at all.
The extremely concentrated HNO3/H2SO4 mixture might be useful for dissolving gold-silver alloys, without the need for inquartation, since aqua regia only dissolves such alloys with extreme difficulty.