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Oxides of Nitrogen

Module by: Andrew R. Barron. E-mail the author

A summary of the physical properties of the oxides of nitrogen is given in Table 1.

Table 1: Physical properties of the oxides of nitrogen.
Oxide Formula Mp (°C) Bp (°C)
Nitrous oxide N2O -90.8 -88.5
Nitric oxide NO -163.6 -161.8
Dinitrogen trioxide N2O3 -100.6 3.5 (dec.)
Nitrogen dioxide (dinitrogen tetroxide) NO2/N2O4 -11.2 (NO2) 21.2 (N2O4)
Nitrogen pentoxide N2O5 30 47

Nitrous oxide

Gaseous nitrous oxide (N2O) is prepared by the careful thermal decomposition of ammonium nitrate (NH4NO2), Equation 1. Nitrous oxide is a linear molecule (Figure 1a) that is isoelectronic (and isostructural) with carbon dioxide. Despite its use as a power enhancement for automobiles, nitrous oxide is actually not very reactive and a major use is as an aerosol propellant.

graphics1.jpg
(1)
Figure 1: Structures of (a) nitrous oxide, (b) nitric oxide, (c) dinitrogen trioxide, (d) nitrogen dioxide, (e) nitrogen tetroxide, and (f) nitrogen pentoxide.
Figure 1 (FigNO.jpg)

Nitrous oxide as an anesthetic drug

Nitrous oxide is known as "laughing gas" due to the euphoric effects of inhaling it, a property that has led to its recreational use as a hallucinogen. However, it is as a anesthetic that it has a legitimate application.

The first use of nitrous oxide as anesthetic drug was when dentist Horace Wells (Figure 2) with assistance by Gardner Quincy Colton (Figure 3) and John Mankey Riggs (Figure 4), demonstrated insensitivity to pain from a dental extraction in December 1844. Wells subsequently treated 12-15 patients, and according to his own record it only failed as an anesthetic in two cases. In spite of these results, the method was not immediately adopted, probably because during his first public demonstration was only partly successful.

Figure 2: American dentist Horace Wells (1815 - 1848).
Figure 2 (graphics3.jpg)
Figure 3: American showman, lecturer, and former medical student Gardner Quincy Colton (1814 - 1898).
Figure 3 (graphics4.jpg)
Figure 4: John Mankey Riggs (1811 - 1885) was the leading authority on periodontal disease in the United States, to the point that periodontal disease was known as Riggs disease.
Figure 4 (graphics5.jpg)

The method did not come into general use until 1863, when Colton successfully used it for more than 25,000 patients. As such, the usage of nitrous oxide rapidly became the preferred anesthetic method in dentistry. Because the gas is mild enough to keep a patient in a conscious and conversational state, and yet in most cases strong enough to suppress the pain caused by dental work, it remains the preferred gas anesthetic in today's dentistry.

Nitrous: the secret to more power.

In motorsports, nitrous oxide (often referred to as nitrous or NOS) allows the engine to burn more fuel, resulting in a more powerful combustion, and hence greater horsepower. The gas itself is not flammable, but it delivers more oxygen (33%) than atmospheric air (21%) by breaking down at elevated temperatures. When N2O breaks down in during fuel combustion, the decomposition of nitrous is exothermic, contributing to the overall power increase.

Nitrous oxide is stored as a compressed liquid (Figure 5); the evaporation and expansion of liquid nitrous oxide in the intake manifold causes a large drop in intake charge temperature, resulting in a denser charge, further allowing more air/fuel mixture to enter the cylinder. Nitrous oxide is sometimes injected into (or prior to) the intake manifold, whereas other systems directly inject right before the cylinder (direct port injection) to increase power.

Figure 5: A typical NOS system for automotive use.
Figure 5 (graphics6.jpg)

One of the major problems of using nitrous oxide in a reciprocating engine is that it can produce enough power to damage or destroy the engine. Very large power increases are possible, and if the mechanical structure of the engine is not properly reinforced, the engine may be severely damaged or destroyed during this kind of operation.

Warning:

Automotive-grade liquid nitrous oxide differs slightly from medical-grade nitrous oxide in that a small amount of sulfur is added to prevent substance abuse.

Nitric oxide

Nitric oxide (NO) is formed by the high temperature oxidation of nitrogen, Equation 2, or the platinum catalyzed oxidation of ammonia at 800 °C, Equation 3.

graphics7.jpg
(2)
graphics8.jpg
(3)

Nitric oxide (Figure 1b) is electronically equivalent to dinitrogen (N2) plus an electron, and as a consequence it is paramagnetic with one unpaired electron. The location of the unpaired electron in the π* orbital (Figure 6a) results in a bond order of 2.5 rather than the triple bond observed for N2 (Figure 6b). The N-O distance of 1.15 Å is intermediate between the triple bond distance in NO+ (1.06 Å) and the typical double bond distance (ca. 1.20 Å). Furthermore, because of the location of the electron it is easy to oxidize nitric oxide to the nitrosonium ion (NO+), Equation 4.

graphics9.jpg
(4)
graphics12.jpg
(5)
Figure 6: Molecular orbital diagrams for (a) nitric oxide and (b) dinitrogen.
Figure 6 (graphics10.jpg)

Nitric oxide is unstable to heat, Equation 6, and oxidation, Equation 5. It will also react with halogens to form the nitrosyl halides, XNO.

graphics11.jpg
(6)

Dinitrogen trioxide

Dinitrogen trioxide (N2O3) is formed from the stoichiometric reaction between NO and O2 or NO and N2O4. Dinitrogen trioxide has an intense blue color in the liquid phase and a pale blue color in the solid state. Thermal dissociation of N2O3, Equation 7, occurs above -30 °C, and some self-ionization of the pure liquid is observed, Equation 8. The asymmetric structure of N2O3 (Figure 1c) results in a polar molecule (Figure 7).

graphics13.jpg
(7)
graphics14.jpg
(8)
Figure 7: The polarization of the N-N bond in dinitrogen trioxide.
Figure 7 (graphics15.jpg)

Nitrogen dioxide (and tetroxide)

Formed from the oxidation of nitric oxide, Equation 9, brown nitrogen dioxide is actually in equilibrium with its colorless dimeric form, nitrogen tetroxide (N2O4), Equation 10.

graphics16.jpg
(9)
graphics17.jpg
(10)

Nitrogen dioxide (Figure 1d) is electronically equivalent to the nitrate anion (NO2-) less one electron, and as such it is paramagnetic with one unpaired electron. The location of the unpaired electron in a nitrogen sp2 orbital, and a consequently it forms a dimer through a N-N bond (Figure 1e). Furthermore, it is easy to oxidize nitrogen dioxide to the nitronium ion (NO2+), Equation 11.

graphics18.jpg
(11)

Nitrogen dioxide dissolves in water to form a mixture of nitric and nitrous acids, Equation 12. Nitrogen dioxide acts as an oxidizing agent with the formation of nitrate anion, Equation 13.

graphics19.jpg
(12)
graphics20.jpg
(13)

The most common structural form of N2O4 (Figure 1e) is planar with a long N-N bond (1.78 Å) that is significantly longer than observed in hydrazine (1.47 Å). Rationalization of this structural effect is obtained from a consideration of the molecular orbitals, which show that the electrons in the σ-bond are actually delocalized over the whole molecule. The rotation about the N-N bond is 9.6 kJ/mol.

Nitrogen pentoxide

The dehydration of nitric acid, with P2O5, yields nitrogen pentoxide, Equation 14, which is an unstable solid at room temperature (Table 1). In the solid state nitrogen pentoxide is actually nitronium nitrate (NO2+NO3-), however, in the vapor phase it exists as a molecular species (Figure 1f) with a bent N-O-N unit. Nitrogen pentoxide is a very powerful nitrating and oxidation agent.

graphics21.jpg
(14)

Nitrogen oxides as precursors to smog and acid rain

Nitrogen oxides (NOx) emissions are estimated to be in the range of 25 - 100 megatonnes of nitrogen per year. Natural sources are thought to make up approximately 1/3 of the total. The generations of NOx (primarily a mixture of NO2 and NO) is the main source of smog and a significant contribution to atmospheric pollution; however, NOx is also responsible for much of the acidity in acid rain.

Atmospheric reactions leading to acid rain

In the dry atmosphere, nitric oxide reacts is oxidized rapidly in sunlight by ozone, Equation 15. The nitrogen dioxide reacts with the hydroxide radical, formed by the photochemical decomposition of ozone, Equation 16 and Equation 17, in the presence of a non-reactive gas molecule such as nitrogen to form nitric acid vapor, Equation 18. The conversion rate for NOx to HNO3 is approximately ten times faster than the equivalent reaction for sulfur dioxide. For example, conversion is essentially complete for a NOx plume by the time it transverses the North Sea from the UK to Scandinavia.

graphics22.jpg
(15)
graphics23.jpg
(16)
graphics24.jpg
(17)
graphics25.jpg
(18)

At night, conversion takes place via the formation of a nitrate radical, Equation 19, which subsequently photochemically unstable under sunlight forming nitrogen pentoxide, Equation 20, that reacts with water on the surface of aerosol particles to form nitric acid, Equation 21.

graphics26.jpg
(19)
graphics27.jpg
(20)
graphics28.jpg
(21)

Both NO and NO2 are only slightly soluble in water and it is therefore more probable that the nitric acid content of rain is more likely due to the dissolution of nitric acid vapor into raindrops, Equation 22, rather than a separate reaction.

graphics29.jpg
(22)

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