Chemistry of inorganic compounds – Nitrogen or compound thereof – Oxygen containing
Reexamination Certificate
1999-06-15
2001-07-24
Langel, Wayne (Department: 1754)
Chemistry of inorganic compounds
Nitrogen or compound thereof
Oxygen containing
C423S393000
Reexamination Certificate
active
06264909
ABSTRACT:
FIELD OF THE INVENTION
The present invention relates to a method for breaking the NO to NO
2
cycle in the production of nitric acid and the recovery of nitric oxide, and more particularly to apparatus and methods for the quantitative recycle of NO
x
gas, recovery, and the production of nitric acid without the need to discharge NO
x
.
BACKGROUND OF THE INVENTION
The mathematically infinite cycle of gas phase oxidation of nitric oxide (NO) to nitric dioxide (NO
2
) and partially back to nitric oxide has dominated nitric acid production and the recovery of nitric oxide.
The nitric oxide to nitric dioxide cycle, which has been the limiting reaction for all nitric acid processes since processes based upon the catalytic oxidation of ammonia began, is described in the following equations:
4NH
3
(g)+50
2
(g)→4NO(g)+6H
2
O(g) (1)
2NO(g)+O
2
(g)→2NO
2
(g) (2)
3NO
2
(g)+H
2
O(1)→2HNO
3
(
1
)+NO(g) (3)
Equations (2) and (3) describe the cycle eliminated by the present invention. As can be seen from equations (2) and (3), the reaction of every three moles of NO
2
with water to form two moles of nitric acid return one of the three moles of oxidized nitrogen as NO. This, of course, requires a reoxidation of the NO with an additional O
2
from air or other source.
Mathematically, one-third of the oxidized nitrogen has to be reoxidized with oxygen every time that two-thirds are reacted with water to form nitric acid.
The economic and environmental problems created by the currently practiced gas phase oxidation of nitric oxide (NO) have dominated and shaped the nitric acid industry since its inception in 1908 when Professor Ostwald piloted the first production of nitric acid based on the catalytic oxidation of ammonia (NH
3
) to nitric oxide (NO).
There are three variables controlling the gas-phase oxidation or reoxidation of NO to NO
2
, Eq. (2). They are:
a. NO concentration vs. O
2
concentration;
b. Temperature: the reaction time decreases with lower temperatures;
c. Pressure: the need to achieve a termolar reaction requiring two molecules of NO and one molecule of O
2
decreases the time required for the oxidation of NO with O
2
to the third power of the pressure in which the reaction occurs.
This has resulted in the industry's development and use of expensive high pressure plants.
As currently practiced, these reactions require the return of one-third of the nitric dioxide to the gas phase as nitric oxide (NO), Eq. (3), which then re-requires the termolar reaction to nitric dioxide (NO
2
), Eq. (2), and then an additional liquid phase reaction to convert two-thirds of this NO
2
to nitric acid (HNO
3
), with one-third of the entering NO
2
again being returned to the gas phase as NO for an additional gas phase reaction, and so on.
This currently used series of reactions is an exercise in commercially striving to reach infinity. Of course, in the current practice of the art of nitric acid manufacture, economics dictate that at some point in this infinite series of reactions, whenever enough of the nitric oxides have been converted to nitric acid so that their further recycle adversely affects the economies of the further recycle process, they are wasted. This results in atmospheric discharge of NO
x
.
In U.S. Pat. No. 3,991,167, Depommier et al. of the firm Produits Chimiques Ugine Kuhlmann point out that current nitric acid from ammonia produces exhaust gases containing from 1,000 to 2,000 cm
3
of nitrogen oxides per cubic meter of effluent while recent legislation seeks to impose a limit of about 200 cm
3
of nitrogen oxides per cubic meter of tail gases.
They further add that the progressive process, Equation (3), previously described makes it “extremely difficult to absorb the last traces of nitrogen oxides in the absorption system conventionally used in manufacturing nitric acid” (column 1, lines 61-63).
Attempts to lower the amounts of released nitrogen oxides by mere extension of the absorption system are fraught with difficult technical problems. Also, the additional installations would entail considerably increased investments. (Column 2, lines 1-6).
Depommier continues (Column 2, lines 15-48) to outline many of the difficulties existing with current nitric oxide emission control processes.
Typical discharge rates from such plants are about 3.9 MT (metric ton) of gases being discharged per each MT of 100% HNO
3
produced. The discharged gases contain a typical concentration of 0.02% to 0.20% NO
x
.
With worldwide nitric acid production for 1985 estimated at greater than 30 million MT per year, this represents an atmospheric discharge of 23,400 to 234,000 MT of NO
x
per year.
Because of environmental regulations and the fact that such discharges are often marked with a disturbing reddish-brown color, most nitric acid plant discharges in the U.S.A. and in other environmentally conscious parts of the world are being treated either to obscure the discharge or to react the discharged NO
x
into another chemical form.
Earlier, the most common treatment was catalytic reaction with excess natural gas which served to reduce, dilute, and disperse the discharged nitric oxides, which made the reddish-brown fume invisible and added unreacted natural gas to the atmosphere.
Current technology is an expensive add-on consisting of ammonia gas, which is added in excess and catalytically reacted with the NO
x
to form nitrogen and ammonium nitrate, which is kept at a sufficiently high temperature to prevent a visible white fume of ammonia nitrate, as it is being discharged into the atmosphere. An excess of ammonia is usually required for this reaction which also discharges chemicals into the atmosphere.
In addition to the above economic and environmental losses caused by the nitrogen dioxide to nitric oxide cycle, there is also the economic burden caused by Eq. (2) in which two moles of NO are required to contact one mole of oxygen in order to form two moles of NO
2
. This also is repeated again and again by the cycle.
In applying Eq. (2) to obtain a sufficiently close contact to effect a reasonably economic rate of molecular reactions between nitric oxide and oxygen, plants are operated at elevated pressures. The attainment of these pressures requires expensive turbo compressor sets and expensive high pressure stainless steel construction of all equipment under pressure.
Freitag and E. Scheibler, who are experts from the Uhde Co., one ofthe largest and the oldest suppliers of nitric acid plants and processes in the world, state in
Handbook of Chemical Production Processes
, Robert A. Meyers, editor, p. 3.6-24, 1986, under their description of “Uhde Nitric Acid Processes” that such turbo compressor sets represent 25-30% of the total cost of a conventional nitric acid plant.
The elevated pressures used to increase the reaction rate also causes greater losses of the expensive platinum catalyst and lower efficiencies in ammonia oxidation and conversion to nitric acid.
Typical Performance Figures for NH
3
Combustion, provided by W. Freitag and E. Scheibler, in their report on Uhde Nitric Acid Processes, previously mentioned, give platinum catalyst consumption at 55 mg/MT HNO
3
produced at one atmosphere pressure (0.1 MPa abs.) and 280 mg/MT HNO3 produced at ten atmospheres of pressure (1.0 MPa abs.). Some of this platinum can be recovered in downstream filters but in all cases platinum costs are substantially higher at the higher operating pressures which are used in current nitric acid production.
Ammonia conversion to nitric oxide and subsequently to nitric acid suffers also. In the Freitag and Scheibler reference cited above regarding the conversion of ammonia to nitric oxide (NH
3
→NO), the percentage oxidized to nitric oxide (NO) drops from 97.5% at 0.1 MPa abs. to 94.0% at 1.0 MPa abs., a loss of 3.5%.
Experts and theory teach that dissolved trivalent nitrogen dissolved in nitric acid cannot be practically oxidized by oxygen alone or in air to additional nitric acid.
This is cogently expressed by Be
Dougherty & Clements LLP
Drinkard Metalox, Inc.
Langel Wayne
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