INTERCHAPTER N

Phosphorus

Red and white phosphorus. White phosphorus, one of the principal allotropes

of solid phosphorus, is very reactive and must be handled with care because it

produces severe burns when it comes in contact with skin. The sample shown here

has a yellowish cast as a result of surface reactions with air. White phosphorus

is usually stored under water. Red phosphorus, on the other hand, is much less

reactive than white phosphorus and does not require special handling.

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Phosphorus (atomic number 15, atomic mass

30.973 761) was the first element whose discovery

could be attributed to a specific person. It was discovered in 1669 by the German alchemist Hennig Brandt by the unsavory process of distilling putrefied urine.

Although phosphorus constitutes less than 0.1% by

mass of the earth's crust, all living organisms contain

this element, and it is the sixth most abundant element in the human body. The energy requirements of essentially all biochemical reactions are supplied

by phosphorus compounds. Plants require phosphorus as a nutrient, and most of the phosphorus compounds that are produced are used as fertilizers.

One of the most important compounds of phosphorus is phosphoric acid, H3PO4(aq), a triprotic acid. Although the formulas of the other oxyacids

of phosphorus are usually written as H3P02 (aq) and

H3PO3(aq), we shall see that these acids are monoprotic and diprotic, respectively. Thus, the three oxyacids of phosphorus illustrate the idea of monoprotic and polyprotic acids presented in Chapter 20.

N-1 . There Are Two Principal Allotropes of Solid Phosphorus

There are several allotropes of elemental solid phosphorus, the most important of which are white phosphorus and red phosphorus (Frontispiece) . White

phosphorus is a white, transparent, waxy crystalline solid that often appears pale yellow because of impurities. It is insoluble in water and alcohol but

soluble in carbon disulfide. A characteristic property

of white phosphorus is its high chemical reactivity.

It ignites spontaneously in air at about 25 ℃. White

phosphorus is very poisonous; the lethal dose is 50 to

100 milligrams. White phosphorus should always be

kept under water and handled with forceps.

When white phosphorus is heated above 400 ℃

for several hours in the absence of air, a form called

red phosphorus is produced. Red phosphorus is a red

to violet powder that is less reactive than white phosphorus. The chemical reactions that the red form undergoes are the same as those of the white form,

Recall that allotropes are forms of an element with

different arrangements of the atoms.

Figure N.1 White phosphorus consists of tetrahedral P_4 molecules.

but they generally occur only at higher temperatures.

For example, red phosphorus must be heated to 260 ℃

before it burns in air. The toxicity of red phosphorus

is much lower than that of white phosphorus.

White phosphorus consists of tetrahedral P_4 molecules (Figure N.1), whereas red phosphorus consists of large, random aggregates of phosphorus atoms.

The structure of red phosphorus is called amorphous,

which means that it has no definite shape. Butter is another example of an amorphous substance.

Most of the phosphorus that is produced is used

to make phosphoric acid and other phosphorus compounds. Elemental phosphorus, however, is used in the manufacture of pyrotechnics, matches, rat poisons, incendiary shells, smoke bombs, and tracer bullets.

Phosphorus is not found as a free element in

nature. The principal sources are calcium phosphate

and the apatite ores (Figure N.2):

hydroxyapatite Ca10(OH)2(PO4)6(s)

fluorapatite Ca10F2(PO4)6(s)

chlorapatite Ca10C12 (PO4) 6 (s)

These ores collectively are called phosphate rock.

Large phosphate rock deposits occur in Russia,

in Morocco, and in the United States in Florida,

Tennessee, and Idaho. An electric furnace is used to

obtain phosphorus from phosphate rock. The furnace is charged with powdered phosphate rock, sand,

SiO2(s), and carbon in the form of coke. The source

of heat is an electric current that produces tempera-

tures of over 1000 ℃. A simplified version of the over-all reaction that takes place is

2Ca3(PO4)2(s) + 65i02(s) +10C(s) -->

phosphate rock sand coke

6CaSiO3(l) + 1000(g) + P4(g)

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Figure N.2 The apatite minerals.

Left to right:

hydroxyapatite, Ca10(OH)2(PO4)6(s);

fluorapatite, Ca10F2(PO4)6(s);

and chorapatite, Ca10C12(PO4)6(s).

Liquid calcium silicate, CaSiO3(l), called slag, is

tapped off from the bottom of the furnace, and the

phosphorus vapor that is produced solidifies to white

P4 (s) when the mixture of CO (g) and P4 (g) is passed

through water (carbon monoxide does not dissolve

in water) . The annual world production of elemental

phosphorus is approximately one million metric tons.

Although some phosphate rock is used to make

elemental phosphorus, most phosphate rock is used in

the production of fertilizers. Phosphorus is a required

nutrient of all plants, and phosphorus compounds

have long been used as fertilizer. In spite of its great

abundance, phosphate rock cannot be used as a fertilizer because, as the name implies, it is insoluble in water. Consequently, plants are not able to assimilate

the phosphorus from phosphate rock. To produce a

water-soluble source of phosphorus, phosphate rock is

reacted with sulfuric acid to produce a water-soluble

product called superphosphate, Ca (H2PO4) 2 (s) , one

of the world's most important fertilizers.

N-2. Phosphorus Forms Several Oxyacids

White phosphorus reacts directly with oxygen to pro-

duce the oxides P4O6(s), and P4O10(s). With excess

phosphorus present, P4O6 (s) is formed:

P4(s) + 3O2(g) -> P4O6(s)

excess

With excess oxygen present, P4O10(s) is formed:

P4(s) + 5O2(g) -* P4O10(s)

excess

In practice, a mixture of oxides is formed in each

case, but one oxide can be greatly favored over the

other by controlling the relative amounts of phosphorus and oxygen.

Before the actual molecular formulas of these

phosphorus oxides were known, the empirical formulas P2O3(s) and P2O5(s) were used. Consequently,

P4O6 (s) and P4O10 (s) are still commonly called phosphorus trioxide and phosphorus pentoxide.

It is interesting to compare the molecular structures of P4O6 and P4O10

(Figure N.3). The molecular structure of P4O6 is obtained from that of P4 by

inserting an oxygen atom between each pair of adjacent phosphorus atoms; there are six edges on a tetrahedron, and thus a total of six oxygen atoms are

required. The molecular structure of P4O10 is obtained

from that of P4O6 by attaching an additional oxygen

atom to each of the four phosphorus atoms.

The phosphorus oxides P4O6(s) and P4O10(s) react

with cold water to form the phosphorus oxyacids:

phosphorous acid, H3PO3(aq), and phosphoric acid,

H3PO4(aq). The respective equations are:

P4O6(s) + 6H2O(l) ->

P4O10(s) + 6H2O(l) ->

4 H3PO3 (aq)

4 H3PO4 (aq)

The reaction of P4O10 (s) with cold water is quite vigorous and can be explosive.

Phosphorus pentoxide, P4O10 (s), is a powerful dehydrating agent capable of removing water from concentrated sulfuric acid, which is itself a strong dehydrating

agent as we saw in Interchapter J. In a similar reaction,

N2O5 (s) can be obtained by reacting P4O10 (s) with

nitric acid. The equations for the two reactions are

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Figure N.3 Molecular structure of P4O6 and

P4O10. (a) The P4O6 molecule can be viewed as

arising from the tetrahedral P4 molecule when

an oxygen atom is inserted between each pair

of adjacent phosphorus atoms. (b) The P4O10

molecule can be viewed as arising from P4O6

when an oxygen atom is attached to each of

the four phosphorus atoms. Notice that there

are no phosphoru-phosphorus bonds in

either P4O6 or P4O10.

P4O10(s) + 6H2SO4(l) -> 6SO3(g) + 4H3PO4(l)

P4O10(s) + 12HNO3(l) -> 6N2O5(s) + 4H3PO4(l)

Phosphorus pentoxide is used as a drying agent in

desiccators and dry boxes to remove water vapor.

Hypophosphorous acid, H3PO2(aq), is prepared

by reacting P4(g) with a warm aqueous solution of

NaOH(aq), followed by addition of a strong acid:

P4(g) + 3 OH- (aq) + 3H2O(l) -> 3H2PO2(aq) + PH3(g)

H2PO(aq) + H+(aq) -* H3PO2(aq)

The Lewis formulas for the phosphate ion, PO43-, the

phosphite ion, HPO32-, and the hypophosphite ion,

H2PO2, are

phosphate ion

phosphite ion

hypophosphite ion

Using VSEPR theory (Chapter 8), we predict that

these ions are tetrahedral.

The hydrogen atoms attached to the phosphorus

atom do not dissociate in aqueous solutions. Recall

from Section 20-12 that an acid that can donate three

acidic protons in solution is called a triprotic acid, one

that can donate two acidic protons in solution is called

a diprotic acid, and one that can donate only one acidic

proton in solution is called a monoprotic acid. Thus,

phosphoric acid, H3PO4(aq), is triprotic; phosphorous

acid, H2 (HPO3) (aq), is diprotic; and hypophosphorous

acid, H(H2PO2)(aq), is monoprotic. As noted previously, the latter two formulas are generally written as

H3PO3 (aq) and H3PO2 (aq), respectively. The structures

of these three acids are shown in Figure N.4.

Almost 11.5 million metric tons of phosphoric acid

are produced annually in the United States alone. It

is produced industrially by the reaction of phosphate

rock and sulfuric acid. Commercial phosphoric acid

is sold as an 85% by mass (85 g of H3PO4 to 15 g of

H2O) solution, equivalent to 15 M. The solution is a

colorless, syrupy liquid. The principal use of phosphoric acid is in the manufacture of fertilizers. It is

also used extensively in the production of soft drinks,

and many of its salts are used in the food industry. For

example, the monosodium salt, NaH2PO4(s), is used

in a variety of foods to control acidity; and calcium

dihydrogen phosphate, Ca(H2PO4)2(s), is the acidic

ingredient in baking powder. The evolution of car-

bon dioxide that takes place when baking powder is

heated can be represented as

Ca (H2PO4) 2(S) + 2 NaHCO3 (s) 3 C

baking powder

2CO2(g) + 2H2O(g) + CaHPO4(s) + Na2HPO4(s)

The slowly evolving CO2 (g) gets trapped in small gas

pockets and thereby causes the cake or bread to rise.

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Figure N.4 The molecular structure of (a) phosphoric acid, H3PO4(aq), (b) phosphorous acid,

H3PO3 (aq) , and (c) hypophosphorous acid, H3PO2 (aq) . Note that all three hydrogen atoms of

phosphoric acid are attached to oxygen atoms. One of the hydrogen atoms in phosphorous acid is

attached directly to the phosphorus atom, and two of the hydrogen atoms in hypophosphorous acid are

attached to the phosphorus atom. Only those hydrogen atoms attached to oxygen atoms are dissociable,

so phosphoric acid is triprotic, phosphorous acid is diprotic, and hypophosphorous acid is monoprotic.

When phosphoric acid is heated gently, pyrophosphoric acid, H4P2O7(aq), (gyro- means heat) is

obtained as a result of the elimination of a water mol-

ecule from a pair of phosphoric acid molecules:

pyrophosphoric acid, H4P2O7

Pyrophosphoric acid, which is also called diphos-phoric acid, is a viscous, syrupy liquid that tends to

solidify on long standing. In aqueous solution, it

slowly reverts to phosphoric acid.

Longer chains of phosphate groups can be formed.

The compound sodium triphosphate, Na5P3O10(s),

used to be the primary phosphate ingredient of

detergents. Its role was to break up and suspend dirt

and stains by forming water-soluble complexes with

metal ions. (The formation of complexes is discussed

in Chapter 26.) In the 1960s almost all detergents

contained phosphates, sometimes as much as 50% by

mass. It was discovered, however, that the phosphates

led to a serious water pollution problem. The enormous quantity of phosphates discharged into rivers

and lakes served as a nutrient for the rampant growth

of algae and other organisms. When these organisms

died, much of the oxygen dissolved in the water was

consumed in the decay process, thus depleting the

water's oxygen supply and destroying the ecological

balance. This process is called eutrophication. As a

result of legislation in the 1970s, phosphates have

been eliminated from detergents, or at least their levels have been reduced markedly.

N-3. Phosphorus Forms a Number of Binary Compounds

Phosphorus reacts directly with reactive metals, such

as sodium and calcium, to form phosphides; for

example,

12Na(s) +P4(s) -> 4Nay(s)

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Most metal phosphides react vigorously with water to

produce phosphine, PH3 (g) :

Ca3P2(s) + 6H2O(l) - 2PH3(g) + 3Ca(OH)2(aq)

Phosphine has a trigonal pyramidal structure with an

H P H bond angle of 93.7 It is a colorless, extremely

toxic gas with an offensive odor like that of rotten

fish. Phosphine reacts violently with oxygen and the

halogens. Unlike ammonia, phosphine does not act

as a base toward water, and few phosphonium (PH4)

salts are stable. Phosphine can also be prepared by

the reaction of white phosphorus with a strong base.

The equation for the reaction is

P4(s) + 3NaOH(aq) + 3H2O(l) -

PH3 (g) + 3 NaH2PO2 (aq)

Phosphorus reacts directly with the halogens to

form halides (Figure N.5). If an excess of phosphorus

is used, then the trihalide is formed. For example,

P4(s) + 6C12(g) - 4PC13(l)

excess

Phosphorus trichloride reacts with chlorine to give

phosphorus pentachloride:

PC13(l) + C12(g) -+ PC15(s)

Figure N.5 The reaction between phosphorus and bromine.

Trigonal pyramidal

Trigonal bipyramidal

Figure N.6 The phosphorus trihalides, PX3, have a trigonal

pyramidal molecular structure in the gas phase.

Figure N.7 The phosphorus pentahalides, PX5, have a trigonal

bipyramidal molecular structure in the gas phase.

Recall from Chapter 8 that phosphorus trihalide

molecules in the gas phase have a trigonal pyramidal structure (Figure N.6) and that phosphorus

pentahalide molecules in the gas phase have a trigonal bipyramidal structure (Figure N.7) . In the solid

phase, however, X-ray diffraction studies have shown

that PC15 (s) exists as PC14 and PC16 ion pairs.

Phosphorus halides react vigorously with water:

PC13 (l) + 3 H2O (l) -> H3PO3 (aq) + 3 HC1(aq)

PC15(s) + 4H20(l) -> H3PO4(aq) + 5HC1(aq)

When phosphorus is heated with sulfur, the yellow

crystalline compound tetraphosphorus trisulfide,

P4S3(s), is formed. Matches that can be ignited by

striking on any rough surface contain a tip composed

of the yellow P4S3(s) on top of a red portion that contains lead dioxide, Pb02(s), together with antimony

sulfide, Sb2S3(s) (Figure N.8). Friction causes the

P4S3(s) to ignite in air, and the heat produced then

initiates a reaction between antimony sulfide and

lead dioxide, which produces a flame.

Safety matches consist of a mixture of potassium

chlorate, KC1O3(s), and antimony sulfide, Sb2S3(s).

The match is ignited by striking it on a special rough

surface composed of a mixture of red phosphorus,

glue, and abrasive. The red phosphorus is ignited by

friction and in turn ignites the reaction mixture in

the match head.

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Figure N.8 Phosphorus in the form of P4S3(s) is one of the

principal components of "strike-anywhere" matches.

N-4. Many Phosphorus Compounds Are Important Biologically

Ordinary tooth enamel is hydroxyapatite,

Ca10(OH) 2 (PO4) 6 (s) . If low concentrations of fluoride ion are added to the diets of children, then a

substantial amount of the tooth enamel formed

will consist of fluorapatite, Ca10F2 (PO4) 6 (s) , which is

much harder and less affected by acidic substances

than hydroxyapatite. Consequently, fluorapatite is

more resistant to tooth decay than is hydroxyapatite. Small quantities of fluoride are added to most

municipal water supplies, and the incidence of tooth

decay among children has decreased markedly over

the past few decades.

The energy requirements for many biochemical

reactions are supplied by a substance called adeno-

sine triphosphate, or simply ATP (Figure N.9). The

chain of three phosphate groups in ATP makes it an

energy-rich molecule. Under physiological condi-

tions, the reaction of one mole of ATP with water to

produce adenosine diphosphate (ADP) and a hydro-

gen phosphate ion releases 31 kilojoules of energy.

This energy is used by all living species to drive bio-

chemical reactions. We can represent the reaction of

ATP with water schematically by

ATP(aq) + H2O(l)

ADP(aq) + HPO2-4(aq)

Adenosine diphosphate is converted into ATP by

the biochemical oxidation of food molecules. The

ATP is then available to supply energy for muscular

activity, synthesis of proteins and other biochemical

molecules, production of nerve signals, and other

biological activity. In other words, ATP is a biological

fuel. The formation and utilization of ATP occur on

the average within about one minute of each other.

The amount of ATP used by the human body is truly

remarkable: at rest over a 24-hour period about 40

kilograms of ATP are utilized. For strenuous exercise,

the rate of utilization of ATP can reach 5 kilograms

in 10 minutes.

Many organic phosphates are potent insecticides

that are also toxic to humans. These insecticides act

by blocking the transmission of electrical signals in

the respiratory system, thereby causing paralysis and

death by suffocation. Fortunately, such poisons do

not last for long in the environment because they are

destroyed over a period of several days by reaction

with water. An important example of an organophosphorus insecticide is malathion, which has been used

triphosphate group

Figure N.9 The Lewis formula of ATP. ADP is similar but has two phosphate groups joined together instead of three.

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N. PHOSPHORUS

to combat the Mediterranean fruit fly infestation in

California and mosquitoes carrying the West Nile

virus in various urban communities around the United

States. Malathion is toxic to humans, but only at fairly

large doses. An enzyme in human gastric juice decom-

poses malathion (insects lack this enzyme); thus, mal-

athion is most toxic to humans when it is absorbed

directly into the bloodstream, as, for example, when it

comes into contact with a cut in the skin.

Some commercially important compounds of

phosphorus are given in Table N.1.

TABLE

N. I Some important compounds of phosphorus

QUESTIONS

N - 1 . Discuss the difference in reactivity between white phosphorus and red phosphorus.

N-2. What is the molecular structure of P4 (s) ?

N-3. What is the general formula for an apatite mineral?

N-4. What is phosphate rock? What is its most important use?

N-5. Why can't phosphate rock be used directly as a fertilizer?

N-6. Describe the molecular structures of P4O6 (s) and P4O10 (s) .

N-7. Describe the molecular structures of H3PO2(l), H3PO3(l), and H3PO4(l). How many dissociable protons are there per mole of phosphorous acid?

Of hypophosphorous acid?

N-8. Describe the action of baking powder.

N-9. Discuss the process of eutrophication.

N-10. Compare ammonia and phosphine as bases.

N-11. Describe two ways to prepare phosphine.

N-I2. Discuss the difference between safety matches and strike-anywhere matches.

N-I3. What is a desiccant? Give an example of one.

Compound Uses

phosphorus pentasulfide,

P2S10 (s)

phosphorus pentoxide,

P4010 (s)

phosphoric acid,

H3PO4(aq)

sodium phosphates:

NaH2PO4(s), Na2HPO4(s),

and Na3PO4(s)

calcium phosphates:

CaHPO4(s) and

Ca(H2PO4) 2 (s)

safety matches; oil

additive

dehydrating agent

fertilizers; soaps and

detergents; soft drinks;

soil stabilizer

synthetic detergents;

water softeners; leavening

agents

fertilizers; poultry and

animal feeds

TERMS YOU SHOULD KNOW

white phosphorus N1

red phosphorus N1

amorphous N1

superphosphate N2

adenosine triphosphate (ATP) N6

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