22.7 Nitrogen

Nitrogen was discovered in 1772 by the Scottish botanist Daniel Rutherford. He found that when a mouse was enclosed in a sealed jar, the animal quickly consumed the life-sustaining component of air (oxygen) and died. When the "fixed air" (CO2) in the container was removed, a "noxious air" remained that would not sustain combustion or life. That gas is known to us now as nitrogen.

Nitrogen constitutes 78 percent by volume of Earth's atmosphere, where it occurs as N2 molecules. Although nitrogen is a key element in living organisms, compounds of nitrogen are not abundant in Earth's crust. The major natural deposits of nitrogen compounds are those of KNO3 (saltpeter) in India and NaNO3 (Chile saltpeter) in Chile and other desert regions of South America.

Properties of Nitrogen

Nitrogen is a colorless, odorless, and tasteless gas composed of N2 molecules. Its melting point is -210°C, and its normal boiling point is -196°C.

The N2 molecule is very unreactive because of the strong triple bond between nitrogen atoms (the NN bond enthalpy is 941 kJ/mol, nearly twice that for the bond in O2; see Table 8.3). When substances burn in air, they normally react with O2 but not with N2. However, when magnesium burns in air, reaction with N2 also occurs to form magnesium nitride, Mg3N2. A similar reaction occurs with lithium.

[22.60]

[22.61]

The nitride ion is a strong Brønsted-Lowry base. It reacts with water to form ammonia, NH3:

[22.62]

The electron configuration of the nitrogen atom is [He]2s22p3. The element exhibits all formal oxidation states from +5 to -3, as shown in Table 22.6. The +5, 0, and -3 oxidation states are the most common and generally the most stable of these. Because nitrogen is more electronegative than all elements except fluorine, oxygen, and chlorine, it exhibits positive oxidation states only in combination with these three elements.

Figure 22.31 summarizes the standard reduction potentials for interconversion of several common nitrogen species. The potentials in the diagram are large and positive, which indicates that the nitrogen oxides and oxyanions shown are strong oxidizing agents.

FIGURE 22.31 Standard reduction potentials in acid solution for some common nitrogen-containing compounds. For example, reduction of NO3- to NO2 in acid solution has a standard electrode potential of 0.79 V (leftmost entry). You should be able to balance this half-reaction using the techniques discussed in Section 20.2.

Preparation and Uses of Nitrogen

Elemental nitrogen is obtained in commercial quantities by fractional distillation of liquid air. About 3.1 × 1010 kg (34 million tons) of N2 is produced annually in the United States.

Because of its low reactivity, large quantities of N2 are used as an inert gaseous blanket to exclude O2 during the processing and packaging of foods, the manufacture of chemicals, the fabrication of metals, and the production of electronic devices. Liquid N2 is employed as a coolant to freeze foods rapidly.

The largest use of N2 is in the manufacture of nitrogen-containing fertilizers, which provide a source of fixed nitrogen. We have previously discussed nitrogen fixation in the Chemistry and Life box in Section 14.6 and in the Chemistry at Work box in Section 15.1. Our starting point in fixing nitrogen is the manufacture of ammonia via the Haber process. (For more information, see Section 15.1) The ammonia can then be converted into a variety of useful, simple nitrogen-containing species, as shown in Figure 22.32. Many of the reactions along this chain of conversion are discussed in more detail later in this section.

FIGURE 22.32 Sequence of conversion of N2 into common nitrogen compounds.

Hydrogen Compounds of Nitrogen

Ammonia is one of the most important compounds of nitrogen. It is a colorless toxic gas that has a characteristic irritating odor. As we have noted in previous discussions, the NH3 molecule is basic (Kb = 1.8 × 10-5). (For more information, see Section 16.7)

In the laboratory NH3 can be prepared by the action of NaOH on an ammonium salt. The NH4+ ion, which is the conjugate acid of NH3, transfers a proton to OH-. The resultant NH3 is volatile and is driven from the solution by mild heating:

[22.63]

Commercial production of NH3 is achieved by the Haber process:

[22.64]

About 1.6 × 1010 kg (18 million tons) of ammonia is produced annually in the United States. About 75 percent is used for fertilizer.

Hydrazine, N2H4, bears the same relationship to ammonia that hydrogen peroxide does to water. As shown in Figure 22.33, the hydrazine molecule contains an NN single bond. Hydrazine is quite poisonous. It can be prepared by the reaction of ammonia with hypochlorite ion, OCl-, in aqueous solution:

[22.65]

The reaction is complex, involving several intermediates including chloramine, NH2Cl. The poisonous NH2Cl bubbles out of solution when household ammonia and chlorine bleach (which contains OCl-) are mixed. This reaction is one reason for the frequently cited warning not to mix bleach and household ammonia.

FIGURE 22.33 Structures of hydrazine, N2H4, and monomethylhydrazine, CH3NHNH2.

Pure hydrazine is an oily, colorless liquid that explodes on heating. It is a highly reactive reducing agent. Hydrazine is normally employed in aqueous solution, where it can be handled safely. The substance is weakly basic, and salts of N2H5+ can be formed:

[22.66]

The combustion of hydrazine is highly exothermic:

[22.67]

Hydrazine and compounds derived from it, such as monomethylhydrazine (Figure 22.33), are used as rocket fuels.

Sample Exercise 22.9

Hydroxylamine, NH2OH, reduces copper(II) to the free metal in acid solutions. Write a balanced equation for the reaction, assuming that N2 is the oxidation product.

SOLUTION The unbalanced and incomplete half-reactions are

Balancing these equations as described in Section 20.2 gives

Adding these half-reactions gives the balanced equation:

Practice Exercise

(a) In power plants hydrazine is used to prevent corrosion of the metal parts of steam boilers by the O2 dissolved in the water. The hydrazine reacts with O2 in water to give N2 and H2O. Write a balanced equation for this reaction. (b) Monomethylhydrazine, N2H3CH3(l), is used with the oxidizer dinitrogen tetroxide, N2O4(l), to power the steering rockets of the space shuttle orbiter. The reaction of these two substances produces N2, CO2, and H2O. Write a balanced equation for this reaction.

Answers: (a) N2H4(aq) + O2(aq) N2(g) + 2H2O(l)

(b) 5N2O4(l) + 4N2H3CH3(l) 9N2(g) + 4CO2(g) + 12H2O(g)

Oxides and Oxyacids of Nitrogen

Nitrogen forms three common oxides: N2O (nitrous oxide), NO (nitric oxide), and NO2 (nitrogen dioxide). It also forms two unstable oxides that we will not discuss, N2O3 (dinitrogen trioxide) and N2O5 (dinitrogen pentoxide).

Nitrous oxide, N2O, is also known as laughing gas because a person becomes somewhat giddy after inhaling only a small amount of it. This colorless gas was the first substance used as a general anesthetic. It is used as the compressed gas propellant in several aerosols and foams, such as in whipped cream. It can be prepared in the laboratory by carefully heating ammonium nitrate to about 200°C:

[22.68]

Nitric oxide, NO, is also a colorless gas, but, unlike N2O, it is slightly toxic. It can be prepared in the laboratory by reduction of dilute nitric acid, using copper or iron as a reducing agent, as shown in Figure 22.34.

[22.69]

It is also produced by direct reaction of N2 and O2 at high temperatures. This reaction is a significant source of nitrogen oxide air pollutants. (For more information, see Section 18.4) However, the direct combination of N2 and O2 is not presently used for commercial production of NO because the yield is low; the equilibrium constant Kc at 2400 K is only 0.05 (see the Chemistry at Work box in Section 15.6).

The commercial route to NO (and hence to other oxygen-containing compounds of nitrogen) is by means of the catalytic oxidation of NH3:

[22.70]

The catalytic conversion of NH3 to NO is the first step in a three-step process known as the Ostwald process, by which NH3 is converted commercially into nitric acid, HNO3 (Figure 22.35). Nitric oxide reacts readily with O2, forming NO2 when exposed to air (see Figure 22.34):

[22.71]

When dissolved in water, NO2 forms nitric acid:

[22.72]

Note that nitrogen is both oxidized and reduced in this reaction; it disproportionates. The reduction product NO can be converted back into NO2 by exposure to air and thereafter dissolved in water to prepare more HNO3.

FIGURE 22.35 The Ostwald process for converting NH3 to HNO3.

Recently NO has been found to be an important neurotransmitter in the human body. It causes the muscles that line blood vessels to relax, thus allowing an increased passage of blood. In fact, Science magazine named nitric oxide its "Molecule of the Year" for 1992 because of the remarkable role the molecule plays in neurotransmission!

Nitrogen dioxide is a yellow-brown gas. Like NO, it is a major constituent of smog. (For more information, see Section 18.4) It is poisonous and has a choking odor. As discussed in the Introduction of Chapter 15, NO2 and N2O4 exist in equilibrium (Figures 15.1 and 15.2):

[22.73]

The two common oxyacids of nitrogen are nitric acid, HNO3, and nitrous acid, HNO2 (Figure 22.36). Nitric acid is a colorless, corrosive liquid. Nitric acid solutions often take on a slightly yellow color (Figure 22.37) as a result of small amounts of NO2 formed by photochemical decomposition:

[22.74]

FIGURE 22.36 Structures of nitric acid and nitrous acid.

Nitric acid is a strong acid. It is also a powerful oxidizing agent, as the following standard reduction potentials indicate:

[22.75]

[22.76]

Concentrated nitric acid will attack and oxidize most metals, except Au, Pt, Rh, and Ir.

About 7.8 × 109 kg (8.6 million tons) of nitric acid is produced annually in the United States. Its largest use is in the manufacture of NH4NO3 for fertilizers, which accounts for about 80 percent of that produced. HNO3 is also used in the production of plastics, drugs, and explosives.

The development of the Haber and Ostwald processes in Germany just before World War I permitted Germany to make munitions even though naval blockades prevented access to traditional sources of nitrates. Among the explosives made from nitric acid are nitroglycerin, trinitrotoluene (TNT), and nitrocellulose. The reaction of nitric acid with glycerin to form nitroglycerin is shown in Equation 22.77:

The following reaction occurs when nitroglycerin explodes:

[22.78]

All the products of this reaction contain very strong bonds. As a result, the reaction is very exothermic. Furthermore, a tremendous amount of gaseous products form from the liquid. The sudden formation of these gases, together with their expansion resulting from the heat generated by the reaction, produces the explosion. (See the Chemistry at Work box in Section 8.9.)

Nitrous acid, HNO2 (Figure 22.36), is considerably less stable than HNO3 and tends to disproportionate into NO and HNO3. It is normally made by action of a strong acid, such as H2SO4, on a cold solution of a nitrite salt, such as NaNO2. Nitrous acid is a weak acid (Ka = 4.5 × 10-4).