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A space-filling model of a molecule shows two blue atoms labeled, “N,” bonded to one another and to three red atoms labeled, “O.” Two Lewis structures are also shown and connected by a double-headed arrow. The left image shows two nitrogen atoms that are single bonded to one another. The left nitrogen is double bonded to an oxygen atom that has two lone pairs of electrons and single bonded to an oxygen with three lone pairs of electrons. The right nitrogen has one lone pair of electrons and is double bonded to an oxygen atom with two lone pairs of electrons. The right image shows two nitrogen atoms that are single bonded to one another. The right nitrogen is double bonded to an oxygen atom that has two lone pairs of electrons and single bonded to an oxygen atom with three lone pairs of electrons. The right nitrogen has one lone pair of electrons and is double bonded to an oxygen atom with two lone pairs of electrons.
Dinitrogen trioxide, N 2 O 3 , only exists in liquid or solid states and has these molecular (left) and resonance (right) structures.

It is possible to prepare nitrogen dioxide in the laboratory by heating the nitrate of a heavy metal, or by the reduction of concentrated nitric acid with copper metal, as shown in [link] . Commercially, it is possible to prepare nitrogen dioxide by oxidizing nitric oxide with air.

Three photos are shown and connected by right-facing arrows. The left image shows a test tube in a clamp that holds a colorless solution and a wire held above it. The middle image shows a test tube in a clamp that holds a wire submerged in a pale green liquid and emitting a light brown gas. The right image shows a test tube in a clamp that holds a wire submerged in a dark green liquid and emitting a brown gas.
The reaction of copper metal with concentrated HNO 3 produces a solution of Cu(NO 3 ) 2 and brown fumes of NO 2 . (credit: modification of work by Mark Ott)

The nitrogen dioxide molecule (illustrated in [link] ) contains an unpaired electron, which is responsible for its color and paramagnetism. It is also responsible for the dimerization of NO 2 . At low pressures or at high temperatures, nitrogen dioxide has a deep brown color that is due to the presence of the NO 2 molecule. At low temperatures, the color almost entirely disappears as dinitrogen tetraoxide, N 2 O 4 , forms. At room temperature, an equilibrium exists:

2NO 2 ( g ) N 2 O 4 ( g ) K P = 6.86
Two space-filling models and two Lewis structures are shown. The left space-filling model shows a blue atom labeled, “N,” bonded to two red atoms labeled, “O,” while the right space-filling model shows two blue atoms labeled, “N,” each bonded to two red atoms labeled, “O.” The left Lewis structure shows a nitrogen atom with one lone electron single bonded to an oxygen atom with three lone pairs of electrons. The nitrogen atom is also double bonded to an oxygen atom with two lone pairs of electrons. The right structure, which is connected by a double-headed arrow to the first, is a diagram showing a similar Lewis structure, but the position of the double bond and the number of electron pairs on the oxygen atoms have switched.
The molecular and resonance structures for nitrogen dioxide (NO 2 , left) and dinitrogen tetraoxide (N 2 O 4 , right) are shown.

Dinitrogen pentaoxide, N 2 O 5 (illustrated in [link] ), is a white solid that is formed by the dehydration of nitric acid by phosphorus(V) oxide (tetraphosphorus decoxide):

P 4 O 10 ( s ) + 4HNO 3 ( l ) 4HPO 3 ( s ) + 2N 2 O 5 ( s )

It is unstable above room temperature, decomposing to N 2 O 4 and O 2 .

A space-filling model and a Lewis structure are shown. The space-filling model shows two blue atoms labeled, “N,” each bonded to two red atoms labeled, “O,” with another red atom labeled, “O,” in between them. The Lewis structure shows a nitrogen atom single bonded to an oxygen atom with three lone pairs of electrons in a downward position and double bonded to an oxygen atom with two lone pairs of electrons in an upward position. This nitrogen is single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to another nitrogen atom which is single bonded to another oxygen atom with three lone pairs of electrons in an upward position. The second nitrogen atom is also double bonded to an oxygen atom with two lone pairs of electrons in a downward position.
This image shows the molecular structure and one resonance structure of a molecule of dinitrogen pentaoxide, N 2 O 5.

The oxides of nitrogen(III), nitrogen(IV), and nitrogen(V) react with water and form nitrogen-containing oxyacids. Nitrogen(III) oxide, N 2 O 3 , is the anhydride of nitrous acid; HNO 2 forms when N 2 O 3 reacts with water. There are no stable oxyacids containing nitrogen with an oxidation state of 4+; therefore, nitrogen(IV) oxide, NO 2 , disproportionates in one of two ways when it reacts with water. In cold water, a mixture of HNO 2 and HNO 3 forms. At higher temperatures, HNO 3 and NO will form. Nitrogen(V) oxide, N 2 O 5 , is the anhydride of nitric acid; HNO 3 is produced when N 2 O 5 reacts with water:

N 2 O 5 ( s ) + H 2 O ( l ) 2HNO 3 ( a q )

The nitrogen oxides exhibit extensive oxidation-reduction behavior. Nitrous oxide resembles oxygen in its behavior when heated with combustible substances. N 2 O is a strong oxidizing agent that decomposes when heated to form nitrogen and oxygen. Because one-third of the gas liberated is oxygen, nitrous oxide supports combustion better than air (one-fifth oxygen). A glowing splinter bursts into flame when thrust into a bottle of this gas. Nitric oxide acts both as an oxidizing agent and as a reducing agent. For example:

oxidizing agent: P 4 ( s ) + 6NO ( g ) P 4 O 6 ( s ) + 3N 2 ( g )
reducing agent: Cl 2 ( g ) + 2NO ( g ) 2ClNO ( g )

Nitrogen dioxide (or dinitrogen tetraoxide) is a good oxidizing agent. For example:

NO 2 ( g ) + CO ( g ) NO ( g ) + CO 2 ( g )
NO 2 ( g ) + 2HCl ( a q ) NO ( g ) + Cl 2 ( g ) + H 2 O ( l )

Key concepts and summary

Nitrogen exhibits oxidation states ranging from 3− to 5+. Because of the stability of the N≡N triple bond, it requires a great deal of energy to make compounds from molecular nitrogen. Active metals such as the alkali metals and alkaline earth metals can reduce nitrogen to form metal nitrides. Nitrogen oxides and nitrogen hydrides are also important substances.

Chemistry end of chapter exercises

Write the Lewis structures for each of the following:

(a) NH 2−

(b) N 2 F 4

(c) NH 2

(d) NF 3

(e) N 3

(a) NH 2− :
This Lewis structure shows a nitrogen atom with three lone pairs of electrons single bonded to a hydrogen atom. The structure is surrounded by brackets. Outside and superscript to the brackets is a two negative sign. ;
(b) N 2 F 4 :
This Lewis structure shows two nitrogen atoms, each with one lone pair of electrons, single bonded to one another and each single bonded to two fluorine atoms. Each fluorine atom has three lone pairs of electrons. ;
(c) NH 2 :
This Lewis structure shows a nitrogen atom with two lone pairs of electrons single bonded to two hydrogen atoms. The structure is surrounded by brackets. Outside and superscript to the brackets is a negative sign. ;
(d) NF 3 :
This Lewis structure shows a nitrogen atom, with one lone pair of electrons, single bonded to three fluorine atoms. Each fluorine atom has three lone pairs of electrons. ;
(e) N 3 :
Three Lewis structures are shown and connected by double-headed arrows in between. The left structure shows a nitrogen atom with a lone pair of electrons triple bonded to a second nitrogen which is single bonded to a third nitrogen. The third nitrogen has three lone pairs of electrons. The entire structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign. The middle structure shows a nitrogen atom with three lone pair of electrons single bonded to a second nitrogen which is triple bonded to a third nitrogen. The third nitrogen which has one lone pair of electrons. The entire structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to a second nitrogen which is double bonded to a third nitrogen. The third nitrogen atom has two lone pairs of electrons. The entire structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign.

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For each of the following, indicate the hybridization of the nitrogen atom (for N 3 , the central nitrogen).

(a) N 2 F 4

(b) NH 2

(c) NF 3

(d) N 3

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Explain how ammonia can function both as a Brønsted base and as a Lewis base.

Ammonia acts as a Brønsted base because it readily accepts protons and as a Lewis base in that it has an electron pair to donate.
Brønsted base: NH 3 + H 3 O + NH 4 + + H 2 O
Lewis base: 2NH 3 + Ag + [H 3 N Ag NH 3 ] +

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Determine the oxidation state of nitrogen in each of the following. You may wish to review the chapter on chemical bonding for relevant examples.

(a) NCl 3

(b) ClNO

(c) N 2 O 5

(d) N 2 O 3

(e) NO 2

(f) N 2 O 4

(g) N 2 O

(h) NO 3

(i) HNO 2

(j) HNO 3

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For each of the following, draw the Lewis structure, predict the ONO bond angle, and give the hybridization of the nitrogen. You may wish to review the chapters on chemical bonding and advanced theories of covalent bonding for relevant examples.

(a) NO 2

(b) NO 2

(c) NO 2 +

(a) NO 2 :
Two Lewis structures are shown and connected by double-headed arrows in between. The left structure shows a nitrogen atom with a single electron double bonded to an oxygen atom which has two lone pairs of electrons. The nitrogen atom is also single bonded to an oxygen atom with three lone pairs of electrons. The right structure is a mirror image of the left structure.
Nitrogen is sp 2 hybridized. The molecule has a bent geometry with an ONO bond angle of approximately 120°.
(b) NO 2 :
Two Lewis structures are shown and connected by double-headed arrows in between. Each structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign. The left structure shows a nitrogen atom with a lone pair of electrons double bonded to an oxygen atom which has two lone pairs of electrons. The nitrogen atom is also single bonded to an oxygen atom with three lone pair of electrons. The right structure is a mirror image of the left structure.
Nitrogen is sp 2 hybridized. The molecule has a bent geometry with an ONO bond angle slightly less than 120°.
(c) NO 2 + :
This Lewis structure shows a nitrogen atom double bonded on both sides to an oxygen atom which has two lone pairs of electrons each. The structure is surrounded by brackets and outside and superscript to the brackets is a negative sign.
Nitrogen is sp hybridized. The molecule has a linear geometry with an ONO bond angle of 180°.

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How many grams of gaseous ammonia will the reaction of 3.0 g hydrogen gas and 3.0 g of nitrogen gas produce?

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Although PF 5 and AsF 5 are stable, nitrogen does not form NF 5 molecules. Explain this difference among members of the same group.

Nitrogen cannot form a NF 5 molecule because it does not have d orbitals to bond with the additional two fluorine atoms.

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The equivalence point for the titration of a 25.00-mL sample of CsOH solution with 0.1062 M HNO 3 is at 35.27 mL. What is the concentration of the CsOH solution?

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Practice Key Terms 1

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Source:  OpenStax, Chemistry. OpenStax CNX. May 20, 2015 Download for free at http://legacy.cnx.org/content/col11760/1.9
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