Like the group 14 elements, the lightest member of group 15, nitrogen, is found in nature as the free element, and the heaviest elements have been known for centuries because they are easily isolated from their ores. Although nitrogen is the most abundant element in the atmosphere, it was the last of the pnicogens ( Group 15 elements) to be obtained in pure form. In 1772, Daniel Rutherford, working with Joseph Black (who discovered CO 2 ), noticed that a gas remained when CO 2 was removed from a combustion reaction. Antoine Lavoisier called the gas azote, meaning “no life,” because it did not support life. When it was discovered that the same element was also present in nitric acid and nitrate salts such as KNO 3 (nitre), it was named nitrogen. About 90% of the nitrogen produced today is used to provide an inert atmosphere for processes or reactions that are oxygen sensitive, such as the production of steel, petroleum refining, and the packaging of foods and pharmaceuticals. Preparation and General Properties of Nitrogen Because the atmosphere contains several trillion tons of elemental nitrogen with a purity of about 80%, it is a huge source of nitrogen gas. Distillation of liquefied air yields nitrogen gas that is more than 99.99% pure, but small amounts of very pure nitrogen gas can be obtained from the thermal decomposition of sodium azide: \[\ce{2NaN3(s) ->[\Delta] 2Na(l) + 3N2(g)} \label{Eq1} \] In contrast, Earth’s crust is relatively poor in nitrogen. The only important nitrogen ores are large deposits of KNO 3 and NaNO 3 in the deserts of Chile and Russia, which were apparently formed when ancient alkaline lakes evaporated. Consequently, virtually all nitrogen compounds produced on an industrial scale use atmospheric nitrogen as the starting material. Phosphorus, which constitutes only about 0.1% of Earth’s crust, is much more abundant in ores than nitrogen. Like aluminum and silicon, phosphorus is always found in combination with oxygen, and large inputs of energy are required to isolate it. Reactions and Compounds of Nitrogen Like carbon, nitrogen has four valence orbitals (one 2s and three 2p), so it can participate in at most four electron-pair bonds by using sp 3 hybrid orbitals. Unlike carbon, however, nitrogen does not form long chains because of repulsive interactions between lone pairs of electrons on adjacent atoms. These interactions become important at the shorter internuclear distances encountered with the smaller, second-period elements of groups 15, 16, and 17. Stable compounds with N–N bonds are limited to chains of no more than three N atoms, such as the azide ion (N 3 − ). Nitrogen is the only pnicogen that normally forms multiple bonds with itself and other second-period elements, using π overlap of adjacent np orbitals. Thus the stable form of elemental nitrogen is N 2 , whose N≡N bond is so strong (D N≡N = 942 kJ/mol) compared with the N–N and N=N bonds (D N–N = 167 kJ/mol; D N =N = 418 kJ/mol) that all compounds containing N–N and N=N bonds are thermodynamically unstable with respect to the formation of N 2 . In fact, the formation of the N≡N bond is so thermodynamically favored that virtually all compounds containing N–N bonds are potentially explosive. Again in contrast to carbon, nitrogen undergoes only two important chemical reactions at room temperature: it reacts with metallic lithium to form lithium nitride, and it is reduced to ammonia by certain microorganisms. At higher temperatures, however, N 2 reacts with more electropositive elements, such as those in group 13, to give binary nitrides, which range from covalent to ionic in character. Like the corresponding compounds of carbon, binary compounds of nitrogen with oxygen, hydrogen, or other nonmetals are usually covalent molecular substances. Few binary molecular compounds of nitrogen are formed by direct reaction of the elements. At elevated temperatures, N 2 reacts with H 2 to form ammonia, with O 2 to form a mixture of NO and NO 2 , and with carbon to form cyanogen (N≡C–C≡N); elemental nitrogen does not react with the halogens or the other chalcogens. Nonetheless, all the binary nitrogen halides (NX 3 ) are known. Except for NF 3 , all are toxic, thermodynamically unstable, and potentially explosive, and all are prepared by reacting the halogen with NH 3 rather than N 2 . Both nitrogen monoxide (NO) and nitrogen dioxide (NO 2 ) are thermodynamically unstable, with positive free energies of formation. Unlike NO, NO 2 reacts readily with excess water, forming a 1:1 mixture of nitrous acid (HNO 2 ) and nitric acid (HNO 3 ): \[\ce{2NO2(g) + H2O(l) -> HNO2(aq) + HNO3(aq)} \label{Eq2} \] Nitrogen also forms \(\ce{N2O}\) (dinitrogen monoxide, or nitrous oxide), a linear molecule that is isoelectronic with \(\ce{CO2}\) and can be represented as − N=N + =O. Like the other two oxides of nitrogen, nitrous oxide is thermodynamically unstable. The structures of the three common oxides of nitrogen are as follows: Few binary molecular compounds of nitrogen are formed by the direct reaction of the elements. At elevated temperatures, nitrogen reacts with highly electropositive metals to form ionic nitrides, such as \(\ce{Li3N}\) and \(\ce{Ca3N2}\). These compounds consist of ionic lattices formed by \(\ce{M^{n+}}\) and \(\ce{N^{3−}}\) ions. Just as boron forms interstitial borides and carbon forms interstitial carbides, with less electropositive metals nitrogen forms a range of interstitial nitrides, in which nitrogen occupies holes in a close-packed metallic structure. Like the interstitial carbides and borides, these substances are typically very hard, high-melting materials that have metallic luster and conductivity. Nitrogen also reacts with semimetals at very high temperatures to produce covalent nitrides, such as \(\ce{Si3N4}\) and \(\ce{BN}\), which are solids with extended covalent network structures similar to those of graphite or diamond. Consequently, they are usually high melting and chemically inert materials. Ammonia (NH 3 ) is one of the few thermodynamically stable binary compounds of nitrogen with a nonmetal. It is not flammable in air, but it burns in an \(\ce{O2}\) atmosphere: \[\ce{4NH3(g) + 3O2(g) -> 2N2(g) + 6H2O(g)} \label{Eq3} \] About 10% of the ammonia produced annually is used to make fibers and plastics that contain amide bonds, such as nylons and polyurethanes, while 5% is used in explosives, such as ammonium nitrate, TNT (trinitrotoluene), and nitroglycerine. Large amounts of anhydrous liquid ammonia are used as fertilizer. Nitrogen forms two other important binary compounds with hydrogen. Hydrazoic acid (\(\ce{HN3}\)), also called hydrogen azide, is a colorless, highly toxic, and explosive substance. Hydrazine (\(\ce{N2H4}\)) is also potentially explosive; it is used as a rocket propellant and to inhibit corrosion in boilers. B, C, and N all react with transition metals to form interstitial compounds that are hard, high-melting materials. Example \(\PageIndex{1}\) For each reaction, explain why the given products form when the reactants are heated. \(\ce{Sr(s) + N2O(g) ->[\Delta] SrO(s) + N2(g)}\) \(\ce{NH4NO2(s) ->[\Delta] N2(g) + 2H2O(g)}\) \(\ce{Pb(NO3)2(s) ->[\Delta] PbO2(s) + 2NO2(g)}\) Given: balanced chemical equations Asked for: why the given products form Strategy: Classify the type of reaction. Using periodic trends in atomic properties, thermodynamics, and kinetics, explain why the observed reaction products form. Solution As an alkali metal, strontium is a strong reductant. If the other reactant can act as an oxidant, then a redox reaction will occur. Nitrous oxide contains nitrogen in a low oxidation state (+1), so we would not normally consider it an oxidant. Nitrous oxide is, however, thermodynamically unstable (ΔH° f > 0 and ΔG° f > 0), and it can be reduced to N 2 , which is a stable species. Consequently, we predict that a redox reaction will occur. When a substance is heated, a decomposition reaction probably will occur, which often involves the release of stable gases. In this case, ammonium nitrite contains nitrogen in two different oxidation states (−3 and +3), so an internal redox reaction is a possibility. Due to its thermodynamic stability, N 2 is the probable nitrogen-containing product, whereas we predict that H and O will combine to form H 2 O. Again, this is probably a thermal decomposition reaction. If one element is in an usually high oxidation state and another in a low oxidation state, a redox reaction will probably occur. Lead nitrate contains the Pb 2 + cation and the nitrate anion, which contains nitrogen in its highest possible oxidation state (+5). Hence nitrogen can be reduced, and we know that lead can be oxidized to the +4 oxidation state. Consequently, it is likely that lead(II) nitrate will decompose to lead(IV) oxide and nitrogen dioxide when heated. Even though PbO 2 is a powerful oxidant, the release of a gas such as NO 2 can often drive an otherwise unfavorable reaction to completion (Le Chatelier’s principle). Note, however, that PbO 2 will probably decompose to PbO at high temperatures. Exercise \(\PageIndex{1}\) Predict the product(s) of each reaction and write a balanced chemical equation for each reaction. \(\ce{NO(g) + H2O(l) ->[\Delta]}\) \(\ce{NH4NO3(s) ->[\Delta]}\) \(\ce{Sr(s) + N2(g) ->}\) Answer \(\ce{NO(g) + H2O(l) ->[\Delta] no reaction}\) \(\ce{NH4NO3(s) ->[\Delta] N2O(g) + 2H2O(g)}\) \(\ce{3Sr(s) + N2(g) -> Sr3N2(s)}\) Summary Nitrogen behaves chemically like nonmetals, Nitrogen forms compounds in nine different oxidation states. Nitrogen does not form stable catenated compounds because of repulsions between lone pairs of electrons on adjacent atoms, but it does form multiple bonds with other second-period atoms. Nitrogen reacts with electropositive elements to produce solids that range from covalent to ionic in character. Reaction with electropositive metals produces ionic nitrides, reaction with less electropositive metals produces interstitial nitrides, and reaction with semimetals produces covalent nitrides.