• Nahkleh Group
• Robinson Group
• Weaver Group
• Bodner Group

The Chemistry of Nonmetals

The chemistry of nitrogen is dominated by the ease with whichnitrogen atoms form double and triple bonds. A neutral nitrogenatom contains five valence electrons: 2s² 2p³.A nitrogen atom can therefore achieve an octet of valenceelectrons by sharing three pairs of electrons with anothernitrogen atom.

Because the covalent radius of a nitrogen atom is relativelysmall (only 0.070 nm), nitrogen atoms come close enough togetherto form very strong bonds. The bond-dissociation enthalpy for thenitrogen-nitrogen triple bond is 946 kJ/mol, almost twice aslarge as that for an O=O double bond.

The strength of the nitrogen-nitrogen triple bond makes the N₂molecule very unreactive. N₂ is so inert that lithiumis one of the few elements with which it reacts at roomtemperature.

6 Li(s)

+

N2(g)

2 Li3N(s)

In spite of the fact that the N₂ molecule isunreactive, compounds containing nitrogen exist for virtuallyevery element in the periodic table except those in Group VIIIA(He, Ne, Ar, and so on). This can be explained in two ways.First, N₂ becomes significantly more reactive as thetemperature increases. At high temperatures, nitrogen reacts withhydrogen to form ammonia and with oxygen to form nitrogen oxide.

N2(g)	+	3 H2(g)		2 NH3(g)
N2(g)	+	O2(g)		2 NO(g)

Second, a number of catalysts found in nature overcome theinertness of N₂ at low temperature.

The Synthesis ofAmmonia

It is difficult to imagine a living system that does notcontain nitrogen, which is an essential component of theproteins, nucleic acids, vitamins, and hormones that make lifepossible. Animals pick up the nitrogen they need from the plantsor other animals in their diet. Plants have to pick up theirnitrogen from the soil, or absorb it as N₂ from theatmosphere. The concentration of nitrogen in the soil is fairlysmall, so the process by which plants reduce N₂ to NH₃or"fix" N₂ is extremely important.

Although 200 million tons of NH₃ are produced bynitrogen fixation each year, plants, by themselves, cannot reduceN₂ to NH₃. This reaction is carried out byblue-green algae and bacteria that are associated with certainplants. The best-understood example of nitrogen fixation involvesthe rhizobium bacteria found in the root nodules of legumes suchas clover, peas and beans. These bacteria contain a nitrogenaseenzyme, which is capable of the remarkable feat of reducing N₂from the atmosphere to NH₃ at room temperature.

Ammonia is made on an industrial scale by a process firstdeveloped between 1909 and 1913 by Fritz Haber. In theHaberprocess, a mixture of N₂ and H₂gas at 200 to 300 atm and 400 to 600^oC is passed overa catalyst of finely divided iron metal.

N2(g)	+	3 H2(g)		2 NH3(g)

Almost 20 million tons of NH₃ are produced in theUnited States each year by this process. About 80% of it, worthmore than $2 billion, is used to make fertilizers for plants thatcan't fix nitrogen from the atmosphere. On the basis of weight,ammonia is the second most important industrial chemical in theUnited States. (Only sulfuric acid is produced in largerquantities.)

Two-thirds of the ammonia used for fertilizers is convertedinto solids such as ammonium nitrate, NH₄NO₃;ammonium phosphate, (NH₄)₃PO₄;ammonium sulfate, (NH₄)₂SO₄; andurea, H₂NCONH₂. The other third is applieddirectly to the soil asanhydrous (literally,"without water") ammonia. Ammonia is a gas at roomtemperature. It can be handled as a liquid when dissolved inwater to form an aqueous solution. Alternatively, it can becooled to temperatures below -33^oC, in which case thegas condenses to form the anhydrous liquid, NH₃(l).

The Synthesis ofNitric Acid

The NH₃ produced by the Haber process that is notused as fertilizer is burned in oxygen to generate nitrogenoxide.

4 NH3(g)

+

5 O2(g)

4 NO(g)

+

6 H2O(g)

Nitrogen oxideor nitric oxide, as it was once known is acolorless gas that reacts rapidly with oxygen to produce nitrogendioxide, a dark brown gas.

2 NO(g)

+

O2(g)

2 NO2(g)

Nitrogen dioxide dissolves in water to give nitric acid andNO, which can be captured and recycled.

3 NO2(g)

+

H2O(l)

2 HNO3(aq)

+

NO(g)

Thus, by a three-step process developed by Friedrich Ostwaldin 1908, ammonia can be converted into nitric acid.

4 NH3(g)	+	5 O2(g)		4 NO(g)	+	6 H2O(g)
2 NO(g)	+	O2(g)		2 NO2(g)
3 NO2(g)	+	H2O(l)		2 HNO3(aq)	+	NO(g)

The Haber process for the synthesis of ammonia combined withtheOstwald process for the conversion ofammonia into nitric acid revolutionized the explosives industry.Nitrates have been important explosives ever since Friar RogerBacon mixed sulfur, saltpeter, and powdered carbon to makegunpowder in 1245.

16 KNO3(s)

+

S8(s)

+

24 C(s)

8 K2S(s)

+

24 CO2(g)

+

8 N2(g)

Ho = -571.9 kJ/mol N2

Before the Ostwald process was developed the only source ofnitrates for use in explosives was naturally occurring mineralssuch as saltpeter, which is a mixture of NaNO₃ and KNO₃.Once a dependable supply of nitric acid became available from theOstwald process, a number of nitrates could be made for use asexplosives. Combining NH₃ from the Haber process withHNO₃ from the Ostwald process, for example, givesammonium nitrate, which is both an excellent fertilizer and acheap, dependable explosive commonly used in blasting powder.

IntermediateOxidation Numbers

Nitric acid (HNO₃) and ammonia (NH₃)represent the maximum (+5) and minimum (-3) oxidation numbers fornitrogen. Nitrogen also forms compounds with every oxidationnumber between these extremes (see table below).

Common Oxidation Numbers for Nitrogen

NegativeOxidation Numbers of Nitrogen Besides -3

At about the time that Haber developed the process for makingammonia and Ostwald worked out the process for converting ammoniainto nitric acid, Raschig developed a process that used thehypochlorite (OCl^-) ion to oxidize ammonia to producehydrazine, N₂H₄.

2 NH3(aq)

+

OCl-(aq)

N2H4(aq)

+

Cl-(aq)

+

H2O(l)

This reaction can be understood by noting that the OCl^-ion is a two-electron oxidizing agent. The loss of a pair ofelectrons and a pair of H⁺ ions by neighboring NH₃molecules would produce a pair of highly reactive NH₂molecules, which would combine to form a hydrazine molecule asshown in the figure below.

Hydrazine is a colorless liquid with a faint odor of ammoniathat can be collected when this solution is heated until N₂H₄distills out of the reaction flask. Many of the physicalproperties of hydrazine are similar to those of water.

There is a significant difference between the chemicalproperties of these compounds, however. Hydrazine burns whenignited in air to give nitrogen gas, water vapor, and largeamounts of energy.

N2H4(l)

+

O2(g)

N2(g)

+

2 H2O(g)

Ho = -534.3 kJ/mol N2H4

The principal use of hydrazine is as a rocket fuel. It issecond only to liquid hydrogen in terms of the number ofkilograms of thrust produced per kilogram of fuel burned.Hydrazine has several advantages over liquid H₂,however. It can be stored at room temperature, whereas liquidhydrogen must be stored at temperatures below -253^oC.Hydrazine is also more dense than liquid H₂ andtherefore requires less storage space.

Pure hydrazine is seldom used as a rocket fuel because itfreezes at the temperatures encountered in the upper atmosphere.Hydrazine is mixed withN,N-dimethylhydrazine, (CH₃)₂NNH₂,to form a solution that remains a liquid at low temperatures.Mixtures of hydrazine andN,N-dimethylhydrazine wereused to fuel the Titan II rockets that carried the Project Geminispacecraft, and the reaction between hydrazine derivatives and N₂O₄is still used to fuel the small rocket engines that enable thespace shuttle to maneuver in space.

The product of the combustion of hydrazine is unusual. Whencarbon compounds burn, the carbon is oxidized to CO or CO₂.When sulfur compounds burn, SO₂ is produced. Whenhydrazine is burned, the product of the reaction is N₂because of the unusually strong nitrogen-nitrogen triple bond inthe N₂ molecule.

N2H4(l)

+

O2(g)

N2(g)

+

2 H2O(g)

Hydrazine reacts with nitrous acid (HNO₂) to formhydrogen azide, HN₃, in which the nitrogen atomformally has an oxidation state of -¹/₃.

N2H4(aq)

+

HNO2(aq)

HN3(aq)

+

2 H2O(l)

Pure hydrogen azide is an extremely dangerous substance. Evendilute solutions should be handled with care because of the riskof explosions. Hydrogen azide is best described as a resonancehybrid of the Lewis structures shown in the figure below. Thecorresponding azide ion, N₃^-, is a linearmolecule, which is a resonance hybrid of three Lewis structures.

HN3
N3-

PositiveOxidation Numbers for Nitrogen: The Nitrogen Halides

Fluorine, oxygen, and chlorine are the only elements moreelectronegative than nitrogen. As a result, positive oxidationnumbers of nitrogen are found in compounds that contain one ormore of these elements.

In theory, N₂ could react with F₂ toform a compound with the formula NF₃. In practice, N₂is too inert to undergo this reaction at room temperature. NF₃is made by reacting ammonia with F₂ in the presence ofa copper metal catalyst.

NH3(g)	+	3 F2(g)		NF3(g)	+	3 HF(g)

The HF produced in this reaction combines with ammonia to formammonium fluoride. The overall stoichiometry for the reaction istherefore written as follows.

4 NH3(g)	+	3 F2(g)		NF3(g)	+	3 NH4F(s)

The Lewis structure of NF₃ is analogous to theLewis structure of NH₃, and the two molecules havesimilar shapes.

Ammonia reacts with chlorine to form NCl₃, whichseems at first glance to be closely related to NF₃.But there is a significant difference between these compounds. NF₃is essentially inert at room temperature, whereas NCl₃is a shock-sensitive, highly explosive liquid that decomposes toform N₂ and Cl₂.

2 NCl3(l)

N2(g)

+

3 Cl2(g)

Ammonia reacts with iodine to form a solid that is a complexbetween NI₃ and NH₃. This material is thesubject of a popular, but dangerous, demonstration in whichfreshly prepared samples of NI₃ in ammonia are pouredonto filter paper, which is allowed to dry on a ring stand. Afterthe ammonia evaporates, the NH₃/NI₃crystals are touched with a feather attached to a meter stick,resulting in detonation of this shock-sensitive solid, whichdecomposes to form a mixture of N₂ and I₂.

2 NI3(s)

N2(g)

+

3 I2(g)

PositiveOxidation Numbers for Nitrogen: The Nitrogen Oxides

Lewis structures for seven oxides of nitrogen with oxidationnumbers ranging from +1 to +5 are given in thetable below.

These compounds all have two things in common: they containN=O double bonds and they are less stable than their elements inthe gas phase, as shown by the enthalpy of formation data in thetable below.

Enthalpy of Formation Data for the Oxidesof Nitrogen

Compound		Hof (kJ/mol)
N2O(g)		82.05
NO(g)		90.25
NO2(g)		33.18
N2O3(g)		83.72
N2O4(g)		9.16
N2O5(g)		11.35

Dinitrogen oxide, N₂O, which is also known asnitrous oxide, can be prepared by carefully decomposing ammoniumnitrate.

NH4NO3(s)		N2O(g)	+	2 H2O(g)

Nitrous oxide is a sweet-smelling, colorless gas best known tononchemists as "laughing gas." As early as 1800,Humphry Davy noted that N₂O, inhaled in relativelysmall amounts, produced a state of apparent intoxication oftenaccompanied by either convulsive laughter or crying. When takenin larger doses, nitrous oxide provides fast and efficient relieffrom pain. N₂O was therefore used as the firstanesthetic. Because large doses are needed to produce anesthesia,and continued exposure to the gas can be fatal, N₂O isused today only for relatively short operations.

Nitrous oxide has several other interesting properties. First,it is highly soluble in cream; for that reason, it is used as thepropellant in whipped cream dispensers. Second, although it doesnot burn by itself, it is better than air at supporting thecombustion of other objects. This can be explained by noting thatN₂O can decompose to form an atmosphere that isone-third O₂ by volume, whereas normal air is only 21%oxygen by volume.

2 N2O(g)

2 N2(g)

+

O2(g)

For many years, the endings -ous and -icwere used to distinguish between the lowest and highest of a pairof oxidation numbers. N₂O is nitrous oxide because theoxidation number of the nitrogen is +1. NO is nitric oxidebecause the oxidation number of the nitrogen is +2.

Enormous quantities of nitrogen oxide, or nitric oxide, aregenerated each year by the reaction between the N₂ andO₂ in the atmosphere, catalyzed by a stroke oflightning passing through the atmosphere or by the hot walls ofan internal combustion engine.

N2(g)

+

O2(g)

2 NO(g)

One of the reasons for lowering the compression ratio ofautomobile engines in recent years is to decrease the temperatureof the combustion reaction, thereby decreasing the amount of NOemitted into the atmosphere.

NO can be prepared in the laboratory by reacting copper metalwithdilute nitric acid.

3 Cu(s)

+

8 HNO3(aq)

3 Cu(NO3)2(aq)

+

2 NO(g)

+

4 H2O(l)

The NO molecule contains an odd number of valence electrons.As a result, it is impossible to write a Lewis structure for thismolecule in which all of the electrons are paired (see table of oxides of nitrogen). When NO gas iscooled, pairs of NO molecules combine in a reversible reaction toform adimer (from the Greek, "twoparts"), with the formula N₂O₂, inwhich all of the valence electrons are paired, as shown in the table of oxides of nitrogen.

NO reacts rapidly with O₂ to form nitrogen dioxide(once known as nitrogen peroxide), which is a dark brown gas atroom temperature.

2 NO(g)

+

O2(g)

2 NO2(g)

NO₂ can be prepared in the laboratory by heatingcertain metal nitrates until they decompose.

2 Pb(NO3)2(s)

2 PbO(s)

+

4 NO2(g)

+

O2(g)

It can also be made by reacting copper metal withconcentratednitric acid,

Cu(s)

+

4 HNO3(aq)

Cu(NO3)2(aq)

+

2 NO2(g)

+

2 H2O(l)

NO₂ also has an odd number of electrons andtherefore contains at least one unpaired electron in its Lewisstructures. NO₂ dimerizes at low temperatures to formN₂O₄ molecules, in which all the electronsare paired, as shown in thetable of oxides ofnitrogen.

Mixtures of NO and NO₂ combine when cooled to formdinitrogen trioxide, N₂O₃, which is a blueliquid. The formation of a blue liquid when either NO or NO₂is cooled therefore implies the presence of at least a smallportion of the other oxide because N₂O₂ andN₂O₄ are both colorless.

By carefully removing water from concentrated nitric acid atlow temperatures with a dehydrating agent we can form dinitrogenpentoxide.

4 HNO3(aq)

+

P4O10(s)

2 N2O5(s)

+

4 HPO3(s)

N₂O₅ is a colorless solid thatdecomposes in light or on warming to room temperature. As mightbe expected, N₂O₅ dissolves in water toform nitric acid.

N2O5(s)

+

H2O(l)

2 HNO3(aq)

The Chemistry ofPhosphorus

Phosphorus is the first element whose discovery can be tracedto a single individual. In 1669, while searching for a way toconvert silver into gold, Hennig Brand obtained a white, waxysolid that glowed in the dark and burst spontaneously into flamewhen exposed to air. Brand made this substance by evaporating thewater from urine and allowing the black residue to putrefy forseveral months. He then mixed this residue with sand, heated thismixture in the presence of a minimum of air, and collected underwater the volatile products that distilled out of the reactionflask.

Phosphorus forms a number of compounds that are direct analogsof nitrogen-containing compounds. However, the fact thatelemental nitrogen is virtually inert at room temperature,whereas elemental phosphorus can burst spontaneously into flamewhen exposed to air, shows that there are differences betweenthese elements as well. Phosphorus often forms compounds with thesame oxidation numbers as the analogous nitrogen compounds, butwith different formulas, as shown in the table below.

Nitrogen and Phosphorus Compounds with theSame Oxidation Numbers but Different Formulas

The same factors that explain the differences between sulfurand oxygen can be used to explain the differences betweenphosphorus and nitrogen.

1. Nitrogen-nitrogen triple bonds are much stronger thanphosphorus-phosphorus triple bonds.

2. P-P single bonds are stronger than N-N single bonds.

3. Phosphorus (EN = 2.19) is much lesselectronegative than nitrogen (EN = 3.04).

4. Phosphorus can expand its valence shell to hold more thaneight electrons, but nitrogen cannot.

The Effect of theDifferences in the Singe and Triple Bond Strengths

The ratio of the radii of phosphorus and nitrogen atoms is thesame as the ratio of the radii of sulfur and oxygen atoms, withinexperimental error.

As a result, phosphorus-phosphorus triple bonds are muchweaker than Nitrogen-nitrogen triple bonds, for the same reasonthat S=S double bonds are weaker than O=O double bondsphosphorusatoms are too big to come close enough together to form strongbonds.

Each atom in an N₂ molecule completes its octet ofvalence electrons by sharing three pairs of electrons with asingle neighboring atom. Because phosphorus does not form strongmultiple bonds with itself, elemental phosphorus consists oftetrahedral P₄ molecules in which each atom formssingle bonds with three neighboring atoms, as shown in the figurebelow.

Phosphorus is a white solid with a waxy appearance, whichmelts at 44.1^oC and boils at 287^oC. It ismade by reducing calcium phosphate with carbon in the presence ofsilica (sand) at very high temperatures.

2 Ca3(PO4)2(s)

+

6 SiO2(s)

+

10 C(s)

6 CaSiO3(s)

+

P4(s)

+

10 CO(g)

White phosphorus is stored under water because the elementspontaneously bursts into flame in the presence of oxygen attemperatures only slightly above room temperature. Althoughphosphorus is insoluble in water, it is very soluble in carbondisulfide. Solutions of P₄ in CS₂ arereasonably stable. As soon as the CS₂ evaporates,however, the phosphorus bursts into flame.

The P-P-P bond angle in a tetrahedral P₄ moleculeis only 60^o. This very small angle produces aconsiderable amount of strain in the P₄ molecule,which can be relieved by breaking one of the P-P bonds.Phosphorus therefore forms other allotropes by opening up the P₄tetrahedron. When white phosphorus is heated to 300^oC,one bond inside each P₄ tetrahedron is broken, and theP₄ molecules link together to form apolymer(from the Greekpol, "many," andmeros,"parts") with the structure shown in the figure below.This allotrope of phosphorus is dark red, and its presence insmall traces often gives white phosphorus a light yellow color.Red phosphorus is more dense (2.16 g/cm³) than whitephosphorus (1.82 g/cm³) and is much less reactive atnormal temperatures.

The Effect ofDifferences in the Strengths of P=X and N=X Double Bonds

The size of a phosphorus atom also interferes with its abilityto form double bonds to other elements, such as oxygen, nitrogen,and sulfur. As a result, phosphorus tends to form compounds thatcontain two P-O single bonds where nitrogen would form an N=Odouble bond. Nitrogen forms the nitrate, NO₃^-,ion, for example, in which it has an oxidation number of +5. Whenphosphorus forms an ion with the same oxidation number, it is thephosphate, PO₄^3-, ion, as shown in thefigure below.

Similarly, nitrogen forms nitric acid, HNO₃, whichcontains an N=O double bond, whereas phosphorus forms phosphoricacid, H₃PO₄, which contains P-O singlebonds, as shown in the figure below.

The Effect ofDifferences in the Electronegativities of Phosphorus and Nitrogen

The difference between the electronegativities of phosphorusand nitrogen (EN = 0.85) is the same as the differencebetween the electronegativities of sulfur and oxygen (EN= 0.86), within experimental error. Because it is lesselectronegative, phosphorus is more likely than nitrogen toexhibit positive oxidation numbers. The most important oxidationnumbers for phosphorus are -3, +3, and +5 (see table below).

Common Oxidation Numbers of Phosphorus

Because it is more electronegative than most metals,phosphorus reacts with metals at elevated temperatures to formphosphides, in which it has an oxidation number of -3.

6 Ca(s)

P4(s)

2 Ca3P2(s)

These metal phosphides react with water to produce apoisonous, highly reactive, colorless gas known as phosphine (PH₃),which has the foulest odor the authors have encountered.

Ca3P2(s)

+

6 H2O(l)

2 PH3(g)

+

3 Ca2+(aq)

+

6 OH-(aq)

Samples of PH₃, the phosphorus analog of ammonia,are often contaminated by traces of P₂H₄,the phosphorus analog of hydrazine. As if the toxicity and odorof PH₃ were not enough, mixtures of PH₃ andP₂H₄ burst spontaneously into flame in thepresence of oxygen.

Compounds (such as Ca₃P₂ and PH₃)in which phosphorus has a negative oxidation number are faroutnumbered by compounds in which the oxidation number ofphosphorus is positive. Phosphorus burns in O₂ toproduce P₄O₁₀ in a reaction that gives offextraordinary amounts of energy in the form of heat and light.

P4(s)

+

5 O2(g)

P4O10(s)

Ho = -2985 kJ/mol P4

When phosphorus burns in the presence of a limited amount of O₂,P₄O₆ is produced.

P4(s)

+

3 O2(g)

P4O6(s)

Ho = -1640 kJ/mol P4

P₄O₆ consists of a tetrahedron in whichan oxygen atom has been inserted into each P-P bond in the P₄molecule (see figure below). P₄O₁₀ has ananalogous structure, with an additional oxygen atom bound to eachof the four phosphorus atoms.

P₄O₆ and P₄O₁₀react with water to form phosphorous acid, H₃PO₃,and phosphoric acid, H₃PO₄, respectively.

P4O6(s)	+	6 H2O(l)		4 H3PO3(aq)
P4O10(s)	+	6 H2O(l)		4 H3PO4(aq)

P₄O₁₀ has such a high affinity for waterthat it is commonly used as a dehydrating agent. Phosphorousacid, H₃PO₃, and phosphoric acid, H₃PO₄,are examples of a large class of oxyacids of phosphorus. Lewisstructures for some of these oxyacids and their related oxyanionsare given in the table below.

OXYACID	OXYANION

The Effect ofDifferences in the Abilities of Phosphorus and Nitrogen to ExpandTheir Valence Shell

The reaction between ammonia and fluorine stops at NF₃because nitrogen uses the 2s, 2p_x, 2p_yand 2p_z orbitals to hold valence electrons.Nitrogen atoms can therefore hold a maximum of eight valenceelectrons. Phosphorus, however, has empty 3d atomicorbitals that can be used to expand the valence shell to hold 10or more electrons. Thus, phosphorus can react with fluorine toform both PF₃ and PF₅. Phosphorus can evenform the PF₆^- ion, in which there are 12valence electrons on the central atom, as shown in the figurebelow.