Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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First Thoughts

For us to answer this question, we must think about how much space a face-

centered cubic unit cell occupies. Then, if we know how much space is actually

taken up by atoms, we can determine the amount of free space in the unit cell.

Solution

Because we haven’t speciﬁed any particular example, let’s consider a generic fcc unit

cell. The volume of a fcc unit cell is

√

. We know that a face-centered cubic

unit cell has four complete atoms (each with a volume of

πr

), so the total vol-

ume of the spheres is

4 ×

πr

16πr

. The ratio of the two volumes will tell us

the percentage of the unit cell that is taken up by atoms:

% occupied

16πr

√

16π

(

√

16π

(

√

16.755

22.627

= 0.7405 ×100% = 74.05%

The percent of occupied space in the unit cell is 74.05%, which leaves 25.95% of the

unit cell vacant. That’s a lot of empty space. Note that the value of r cancels out of

the equation, so our answer is valid for any ccp crystal.

Further Insights

Is there really this much empty space in a solid? Yes, according to our calculations,

there are a lot of holes in a solid. Is the space really empty? As we’ll see later in this

chapter, it isn’t completely empty in a metal. Electrons occupying large molecular

orbitals occupy a lot of this empty space.

PRACTICE 13.2

Polonium forms the simple cubic unit cell. How much of a primitive unit cell is

empty space?

See Problems 7, 8, 11, 12, and 17.

The ease of calculating the volume of the unit cell enables us to determine the

density of a crystal without measuring it. Mathematically, the density of a solid is

the mass divided by the volume. Because we’ve already determined the volume of

the unit cell, we need to determine its mass. How do we do this calculation? If we

know how many atoms make up the unit cell and we know the mass of one atom,

then we can calculate the mass of a unit cell. For example, gold crystallizes in

the cubic closest packed structure (fcc unit cell). The atomic mass of gold is

197.0 g/mol, or 197 amu per atom. As we’ve done before, we’ll keep a couple of

extra ﬁgures in each calculation, rounding off only at the end.

197.0gAu

mol

1 mol Au

6.022 ×10

atoms

3.2713 ×10

−22

atom Au

Then

3.2713 ×10

−22

atom Au

4atomsAu

fcc unit cell

= 1.3085 ×10

−21

g/unit cell

What is the density of a block of gold? Because we can calculate the volume and

the mass of any unit cell, we can determine the density of a crystalline solid if we

know the unit cell. For instance, gold has a mass of 1.3085 × 10

–21

g / unit cell

and a unit cell volume of 6.756 × 10

. The density (mass/volume) can be

calculated:

6.756 ×10



1.0 ×10

−12





100 cm



= 6.756 ×10

−23

1.3085 ×10

−21

6.756 ×10

−23

= 19.4g/cm

548 Chapter 13 Modern Materials

The value we’ve calculated for gold (19.4 g/cm

, based solely

on physical measurements of the unit cell) agrees quite

well with the experimental value for the density of gold

(19.3 g/cm

). Table 13.2 lists the important calculations for

the face-centered cubic unit cell.

Other information about the crystalline solid can be deter-

mined by examining the unit cell. For example, the number

of species that surround a particular ion can be determined

from the packing arrangement.

EXERCISE 13.3 He Isn’t Heavy, He’s Dense

Some metals are much denser than others. What makes them so dense? The

arrangement of the atoms in the solid, coupled with the mass of the element, ac-

counts for the different densities we observe. Let’s calculate the density of an iron

bar. Iron, whose atomic radius is 124 pm, forms bcc unit cells.

Solution

We can determine the density of a block of iron by dividing the mass of the appro-

priate number of atoms by the volume of the unit cell. The body-centered cubic

unit cell has the volume



√



. For an iron body-centered cubic unit cell,

the volume is



124 pm ×

√



= (286.4)

= 2.348 ×10

2.348 ×10







100 cm



= 2.348 ×10

−23

In a bcc unit cell containing two atoms of iron, the total mass is

2atoms×

55.847 g

mol

1 mol

6.022 ×10

atoms

= 1.855 ×10

−22

The density of iron is the mass divided by the volume:

1.855 ×10

−22

2.348 ×10

−23

= 7.90 g/cm

The experimentally measured density of iron is 7.86 g/cm

PRACTICE 13.3

Lead is used as ﬁshing weights because of its malleability and density. What is the

density of lead (Pb)? Lead has an atomic radius of 175 pm and crystallizes in a fcc

arrangement.

See Problems 16, 18, 21, and 22.

Calculating the positions of the atoms in a unit cell based on an X-ray dif-

fraction pattern used to be a tedious task that could take months or years to com-

plete. Today, scientists use the X-ray diffractometer to generate X-rays and direct

them at a single crystal at a variety of different angles. Powerful computers collect

the data over a period of hours and store them until the diffraction pattern is

complete, at which time the computers go to work deciphering the pattern into a

picture of the crystal. Depending on the quality of the crystal, the location of each

of the atoms in the solid can be accurately determined overnight.

13.1 The Structure of Crystals 549

Face-Centered Cubic Unit Cell Calculations

Length of side

√

Volume of unit cell

√

Length of diagonal 4r

Mass of unit cell 4 × mass of one atom

TABLE 13.2

Liquid Crystals

We use liquid crystals in the displays of watches, pocket calculators, computers,

and privacy windows. What is a liquid crystal and how does it work? The liquid

crystal is a transitional phase between a solid and a liquid. There are some excep-

tions, but the molecules of the liquid crystal generally have long, rigid structures

with a strong dipole moment that point along a common axis (Figure 13.12) to

align into highly ordered crystalline solids. An example of a solid that forms a liq-

uid crystal is cholesteryl benzoate.

If the temperature of a solid with these characteristics is increased, the long,

rigid structures of the molecules and the strong dipole moment continue to

maintain a high degree of order within the structure. In fact, even well into

the liquid state, order is maintained despite the increase in translational freedom

(the ability to move in many directions). When the temperature gets hot enough,

the degree of order in the structure is eliminated. In other words, in the transi-

tional state between completely solid and completely liquid, the molecules are

highly ordered and yet quite ﬂuid (the

liquid crystal state).

The strong dipole moment is the key to the behavior of liquid crystals. From

Chapter 9, we know that dipole moments align themselves within an electric

ﬁeld. When such a ﬁeld is applied to the liquid crystal, its molecules orient them-

selves with the electric ﬁeld. This highly ordered structure is transparent to light

directed along the axis of orientation. If the electric ﬁeld is removed, the transla-

tional freedom of the molecules scrambles their alignment. Visually, a liquid

crystal in this state is opaque because the slight randomness scatters light. The net

effect is a material that can be made transparent or opaque with the ﬂick of a

switch.

Cholesteryl benzoate

550 Chapter 13 Modern Materials

Solid Li

uid cr

stal Li

uid

FIGURE 13.12

Liquid crystals. The liquid crystal has

properties that resemble the liquid and

properties that resemble the solid. It is

a ﬂuid, yet highly organized phase.

Application

HERE’S WHAT WE KNOW SO FAR

■

The dimensions and the density of a solid can be determined if the mass and

the radii of the components (be they ions, molecules, or atoms) are known.

■

X-ray crystallography is a useful technique that can also determine the

arrangements of atoms in the unit cell.

■

The liquid crystalline state is a transitional phase that exists between the solid

and liquid phase for some molecules. Typically, these molecules have long,

rigid structures with strong dipole moments.

13.2 Metals

To ﬁll a cavity in a tooth, a dentist cleans out the hole and packs it with a com-

posite material made of glass and plastic resin, or with a “silver” ﬁlling, to arrest

the growth of the enamel-eating bacteria. The composite ﬁlling is popular, but

because the silver ﬁlling is much harder, more durable, and relatively inexpensive,

there is still a need for it. Why did we put quotation marks around the word

silver? The ﬁlling is actually is a mixture of about 50% mercury and 50% other

metals (including silver, tin, and copper). This is one of countless applications

of metals in our lives. In this section, we will examine the properties of metals

and learn why these elements are so useful.

Using a cast iron skillet to cook your eggs and bacon used to be the rule, not

the exception. You had to have a mitten handy to pick the skillet up off the stove,

or your hand would be badly burned. Metals, such as the iron in the skillet, have

long been known to conduct heat and electricity (see Chapter 5).

Why do metals

conduct heat?

One of the theories that explains these properties (and others) is

called

band theory. This theory addresses how adjacent metal atoms interact. For

example, two lithium atoms can interact with each other by sharing their elec-

trons. As we discussed in Chapter 9, this gives rise to our postulate that dilithium

can exist. A closer look at the molecular orbitals of dilithium reveals that the va-

lence electrons come from the 2s orbital. A bonding orbital lies below the energy

of the 2s orbital, and an antibonding orbital lies above the energy of the 2s orbital

(Figure 13.13). The valence electrons are placed in the molecular orbitals in such

a way as to ﬁll the bonding orbital and leave the antibonding orbital empty.

What happens if three lithium atoms participate in bonding? The result of

mixing three atomic valence orbitals gives rise to three valence molecular or-

bitals (Figure 13.14). One of these orbitals is bonding, one is antibonding, and

one is nonbonding. A

nonbonding molecular orbital can be thought of as one that

neither enhances nor diminishes the molecular bonding. The nonbonding mole-

cular orbital arises when there is an odd number of atomic orbitals that mix in a

given energy level. The result of this mixing is an odd number of MOs in the en-

ergy level (one of the orbitals is the nonbonding MO). Three lithium atoms con-

tribute a total of three valence electrons to the valence molecular orbital diagram.

Placing them in the diagram results in a full bonding MO and a half-ﬁlled non-

bonding MO.

The overlap of four lithium atoms results in two bonding orbitals (completely

full of electrons) and two antibonding orbitals (empty), as shown in Figure 13.15.

An equal number of bonding molecular orbitals and antibonding molecular or-

bitals are formed. And the bonding orbitals reside at an energy level below that of

the antibonding orbitals.

If we increase the number of atoms that participate in forming molecular

orbitals to eight, we ﬁnd that half of the orbitals (the bonding orbitals) ﬁll with

electrons. The other half of the molecular orbitals is comprised of the empty an-

tibonding orbitals. In addition, the orbitals within the bonding and antibonding

13.2 Metals 551

2s 2s

Energy

1s 1s

FIGURE 13.13

Molecular orbital diagram of dilithium.

Two bonding and two antibonding MOs

make up the diagram for dilithium. Plac-

ing electrons in this MO diagram pro-

vides a full bonding MO and an empty

antibonding MO.

FIGURE 13.14

Molecular orbital dia-

gram of trilithium. Note

that the energy of the

MOs derived from 2s

orbitals is approaching

the energy of the MOs

derived from 1s orbitals.

Application

Video Lesson: Metallurgical

Processes

Video Lesson:

Band Theory of

Conductivity

levels begin to have noticeably different energies. This difference arises due to the

requirements of the symmetry for molecular orbitals; as the size of the molecular

orbital increases, the number of nodes within the molecular orbitals increases

and results in a decrease in the magnitude of the molecular orbital’s energy.

Overall, the energy difference between the antibonding and the bonding molec-

ular orbitals becomes smaller (bonding MOs increase in energy and antibonding

MOs decrease in energy).

As more and more atoms combine to make a metallic crystal, the difference in

energy between bonding and antibonding orbitals becomes negligible. The result

is a band of molecular orbitals constructed from the overlap of s orbitals. In met-

als, the range of energies in this band is often wide enough to overlap with simi-

larly constructed bands of p and d orbitals. Many of these bands are empty, or

only partially ﬁlled, with electrons. For any metal not at absolute zero, many

of the electrons occupy the unﬁlled portions of the higher energy bands. Because

the molecular orbitals encompass the entire metallic crystal, the electrons are

free to roam from one end of the metal crystal to the other.

Crystals made from nonmetal elements, however, have a slightly different

arrangement. As we see from Figure 13.15, the lower energy bands are accompa-

nied by an empty higher energy band that doesn’t overlap. For instance, carbon

has a band of molecular orbitals that results from the overlap of the 2s and 2p

atomic orbitals. Higher in energy lies a band of molecular orbitals that result

from the overlap of the 3s and 3p atomic orbitals. The lower-energy, partially

ﬁlled bonding molecular orbital band is known as the

valence band. This band

contains the valence electrons. The empty higher-energy molecular orbital band

is the

conduction band. Some nonmetal elements have conduction bands that are

much greater in energy than their valence band. Still others have conduction

bands that overlap or nearly overlap with the valence bands.

Band theory describes a metal as a lattice of metal cations spaced throughout

a sea of delocalized electrons. How does this theory explain why metals are duc-

tile, malleable, conductive, and shiny? Imagine, for example, deforming a metal

bar with a hammer and displacing some of the metal ions. The sea of delocalized

electrons can immediately adjust to the change, so the overall structure doesn’t

become signiﬁcantly weaker as it is bent. Metals, then, are

ductile (able to be

pulled) and

malleable (able to be pounded) to different degrees. Additionally, be-

cause electrons can easily move from one end of the molecular orbital to the

other, metals can conduct electricity.

Why can metals conduct heat? Kinetic energy can be conveyed easily across a

crystalline lattice that consists of positive ions within a sea of electrons; metals

552 Chapter 13 Modern Materials

Increasing energy

Antibonding orbitals

Bonding orbitals

FIGURE 13.15

Molecular orbital diagram of lithium metal.

As the number of lithium atoms participating

in bonding increases, the regions of anti-

bonding and bonding molecular orbitals

become continuous.

Energy

Small band gap

–

Conduction

band

(unfilled)

Valence

band

(filled)

Electrons

excited

Conduction

band

(partially filled)

Valence

band

(partially filled)

conduct heat easily. In other words, kinetic energy can be transferred from posi-

tive ion to positive ion with ease. This stands in contrast to the difﬁculty involved

in transferring kinetic energy across an ionic or molecular solid, where the ki-

netic energy must be transferred from cation to anion to cation (or from mole-

cule to molecule). This explains why metal ice cube trays feel colder than plastic

ice cube trays. The metal trays, as thermal conductors, are better able to pull the

heat from your hand. The plastic ice cube trays are made of a thermal insulator

(there isn’t enough thermal energy available to promote electrons into the con-

duction band in the plastic). Or, to say it differently, the valence electrons in plas-

tic are tied to covalent bonds speciﬁcally joined to a given pair of atoms and are not

free to move around.

Band theory also explains why metals are shiny. Because the molecular or-

bitals within the valence bond are so plentiful and close together, electrons can

absorb and release almost any wavelength of light from the infrared to the ultra-

violet. The absorption promotes the electron into an unﬁlled molecular orbital

within the band; the release of a photon returns the electron to its original loca-

tion in the band. The result is that metals reﬂect light, giving rise to the luster that

we associate with them. White paper also reﬂects light but does so diffusely (in all

directions). Some metals have a distinctive color. For example, gold has a yellow

hue. This is due to the inefﬁciency of the metal when it reﬂects certain frequen-

cies of absorbed light. These metals are still shiny; they just don’t reﬂect all wave-

lengths of light equally.

Band theory is used in determining the color of paints, identifying the insu-

lating value of materials in your home, and (remembering the opening scenario

of the chapter) in the hospital as well, in the form of heart monitors, IV pumps,

calculators, and many other instruments.

What do such instruments have in

common?

They are all controlled by small computers, and at the heart of all

computers is the element silicon. Why silicon? Its

band gap

is the key. A band gap is a small energy gap that exists be-

tween the valence band and the conduction band, as illus-

trated in Figure 13.16. The existence of a band gap in silicon

suggests that a certain amount of energy is needed in order

to promote an electron from the valence band into the con-

duction band. If the gap is relatively small, as it is in silicon,

the metal is a

semiconductor. In these cases, the amount of

energy needed to promote an electron into the conduction

band is relatively small. (See Table 13.3 for a list of some

compounds and their band gaps.) When the temperature

13.2 Metals 553

FIGURE 13.16

The band gap in silicon.

Band Gaps of Some Common Materials

Material Band Gap

Insulator Carbon (diamond) 502 kJ/mol

(>300 kJ/mol) Zinc sulﬁde 350 kJ/mol

Semiconductor Silicon 106 kJ/mol

(50–300 kJ/mol) Selenium 215 kJ/mol

Germanium 64 kJ/mol

Conductor Lithium 0 kJ/mol

(<50 kJ/mol) Tin (gray tin) 7.7 kJ/mol

TABLE 13.3

Video Lesson: Intrinsic

Semiconductors

is low, there isn’t enough kinetic energy to promote many electrons into the

conduction band, and the metal doesn’t conduct much electricity. Increase

the temperature of the semiconductor, and the conductivity increases. Silicon

(band gap = 106 kJ/mol) and germanium (band gap = 64 kJ/mol) are good

examples of semiconductors.

If the band gap is sufﬁciently large that there are barely any electrons at room

temperature in the conduction band, the material is an

insulator. Only a tiny

amount of electrical conductivity can be found in an insulator. However, as an in-

sulator becomes hotter, its conductivity increases. Our plastic ice cube tray is a

good example of an insulator. Diamond is another good insulator (the band gap

in diamond is 502 kJ/mol). In short, when the band gap is large, the material acts

as an insulator. Remove the gap, and the material is a

conductor. For example,

compare the band gap in diamond to the amount of energy required to promote

an electron into an unﬁlled orbital in a conductor such as lithium, roughly

–45

kJ/mol.

Adding small amounts of other elements to the metal can modify the behav-

ior of a semiconductor. For example, if we

dope (add an atom or other compound

to) silicon (Group IVA) with an element that has more valence electrons than sil-

icon (such as phosphorus, Group VA), the extra valence electrons must be placed

into the conduction band (Figure 13.17). As a consequence, the semiconductor

becomes capable of conducting current. The result is called an

n-type semicon-

ductor

because the addition of the phosphorus adds extra negative charges (elec-

trons) to the system. However, if we dope silicon with an element containing

fewer valence electrons (such as aluminum, Group IIIA), there will not be

enough valence electrons to ﬁll the valence band. The result: Electrons in the va-

lence band can be easily excited to move to an unoccupied molecular orbital. This

semiconductor is electrically conductive if we apply a small electric ﬁeld. We call

it a

p-type semiconductor because the addition of gallium can be thought of as

introducing positive holes in the valence band. The properties of the n- and

p-type semiconductors make them quite useful for the production of small elec-

tronic devices.

554 Chapter 13 Modern Materials

Conduction

band

(unfilled)

Valence

band

(filled)

Energy

Conduction

band

(partially filled)

p-type n-type

Valence

band

(partially filled)

FIGURE 13.17

Doping the semiconductor. An n-type semi-

conductor contains extra electrons added

to the conduction band. A p-type semicon-

ductor is missing some electrons from the

valence band.

Video Lesson: Doped

Semiconductors

EXERCISE 13.4 The Color of Paint

The mineral greenockite (CdS) is best known as cadmium yellow and is used as a

yellow pigment in paints. Cadmium sulﬁde has a medium-sized band gap (4.2 ×

–19

J/photon). Judging on the basis of the information in Table 13.3, is greenock-

ite an insulator, a semiconductor, or a conductor? What color light does cadmium

sulﬁde absorb?

Solution

The band gaps in Table 13.3 are listed in units of kilojoules per mole, whereas this

exercise gives the size of the band gap for greenockite in joules per photon. We can

use Avogadro’s number to make the units the same.

4.2 ×10

−19

photon greenockite

6.022 ×10

photons

mol greenockite

1kJ

1000 J

250 kJ

mol greenockite

This value places cadmium sulﬁde in the semiconductor region. The promotion of

an electron from the valence band to the conduction band actually requires 4.165 ×

–19

J. Using the equations E = hv and c = λv, we ﬁnd a frequency of 6.286 ×

Hz and a wavelength of 476.9 nm. This corresponds to the absorption of blue

light. A solution of cadmium sulﬁde appears yellow to the eye.

PRACTICE 13.4

What is the largest wavelength of light, in nanometers, that can be absorbed by a

particular semiconductor if the band gap is determined to be 0.95 eV? (1 eV/atom =

96.485 kJ/mol).

See Problems 27, 28, 35, and 36.

Perhaps the greatest beneﬁt that a semiconductor offers is its ability to act as

photovoltaic device, in which light energy can be used to generate an electric

current. Photovoltaic cells capable of over 30% efﬁciency have been designed.

How is the sunlight converted into electricity? When silicon is doped with two

elements such as phosphorus (Group VA) and gallium (Group IIIA), the semi-

conductor contains both electron-rich and electron-deﬁcient areas. Light energy

throughout the visible and IR range (the particular range of wavelengths depends

on the chemical composition of the semiconductor) provides enough energy for

electrons to “jump” the band gap and enter the conduction

band. This enables both negative charges (electrons) and posi-

tive holes to move so that the ﬂow of electrons can be harvested

as energy.

Photovoltaic cells have been used in many applications,

from the solar-powered calculator to solar-powered water

heaters, space stations, and satellites. The use of photovoltaic

cells to make a solar-powered car, such as that shown in Figure

13.18, is the next big hurdle. Because our world demands fossil

fuels as asourceof energyfortransportation,and because of our

limited supply of this natural resource, alternative-energy auto-

mobiles will one day be commonplace. Currently, auto makers

and researchers have been able to develop solar-powered cars

that can travel an average of 50 miles an hour relying solely on

the photovoltaic cell. However, despite the estimated 60% de-

cline in costs to implement the use of photovoltaic cells over the

last 20 years, sales of solar-powered vehicles are still dwarfed by

sales of those powered by fossil fuels; solar-powered autos made

up less than 1% of the total auto sales in 2003.

13.2 Metals 555

FIGURE 13.18

The solar-powered car is currently being tested as a replacement

for the gasoline auto that we use today. This solar-powered

prototype is not too far from what we may be driving in the

near future. Note the photovoltaic cells on the roof of the car.

Application

HEMICAL

ENCOUNTERS:

Photovoltaic

Devices

Alloys

We return now to our opening theme of materials used in the hospital by noting

that the limited properties of the metals don’t lend themselves well to any more

than routine medical tasks. For example, surgeons have a relatively limited num-

ber of choices in materials when replacing a hip joint. They need something as

strong and durable as metal, but they don’t use solid copper, silver, or gold.

Why don’t they use a coinage metal for a hip joint? Many of the pure metals

are too malleable and ductile to withstand the forces in a human body. A

gold hip joint would easily bend when a patient walked with it. Other factors,

such as the cost and toxicity of some of the pure metals, are also important.

In fact, it’s rare that we see pure metals used in the “real world.”Because there

are only a limited number of pure metals, and because the properties of

those pure metals are often not what is required, we look to mixtures of met-

als to give us the properties we need.

A homogeneous mixture of two or more metals is known as an

alloy.The

elements are typically mixed together in the molten state and then allowed

to cool to form the alloy. Alloys are useful because their properties can be

adjusted by changing the ratio in which the elements are combined. (The

Bronze Age is named after the discovery of one of the ﬁrst of these metal

blends.) Typically, alloys are less ductile and less malleable than pure metals.

Table 13.4 lists some of the common alloys, along with a typical composi-

tion for each and several uses to which each is put. Manufacturers vary the

actual composition of each type of alloy to change its properties slightly.

Two different classes of alloys exist, the

substitutional alloys and the inter-

stitial alloys

, illustrated in Figure 13.19. Substitutional alloys include metals

such as brass, sterling silver, pewter, and solder. In a substitutional alloy, spe-

ciﬁc atoms within the lattice structure of a metal are replaced with other

metal atoms. In the example of sterling silver, an average of 7 silver atoms

out of every 100 are replaced with copper atoms. Interstitial alloys are made

of metals with “impurities” trapped in the spaces of the lattice structure. By

far the most important of these is steel. Steel, which we discussed in Section 7.2,

contains carbon atoms located in the spaces between the atoms in an iron atom

lattice.

556 Chapter 13 Modern Materials

Common Alloys and Their Uses

Name Common Composition (mass %) Uses

Substitutional

Brass 90 Cu : 10 Zn Decorative ﬁttings

Pewter 85 Sn : 7 Cu : 6 Bi: 2 Sb Plates, jewelry, ﬁgurines

Medal bronze 92–97 Cu : 1–8 Sn : 0–2 Zn Medals, artwork

14-carat gold 58 Au : 30 Ag : 12 Cu Jewelry

Sterling silver 92.5 Ag : 7.5 Cu Jewelry, kitchenware

Coinage silver 90 Ag : 10 Cu Silver coins

Plumber’s solder 67 Pb : 33 Sn Pipe solder

Interstitial

Steel 98–99.9 Fe: 0–2 C Building materials

Stainless steel >10.5 Cr: <89.5 Fe Eating utensils,

structural materials

TABLE 13.4

An interstitial alloy

A substitutional alloy

FIGURE 13.19

Substitutional versus interstitial alloys.

Amalgams

An amalgam is an alloy containing the metal mercury. Dentists use an amalgam

for ﬁlling teeth by mixing a powder containing silver and tin and then adding

mercury. The amalgam immediately begins to form at room temperature, mak-

ing a complex and very hard mixture of Ag

and Sn

Hg.

14Ag

Sn + 65Hg n 21Ag

+ 2Sn

Other amalgams do exist. For instance, the amalgamation of gold with mer-

cury is one way in which gold dust was extracted from powdered gravel in the

California gold rush era of the mid- to late nineteenth century. Because of

the hazards of mercury poisoning, the use of mercury has been supplanted by

other means of extracting gold (see Chapter 4). Unfortunately, it continues to be

used by individuals and small companies in the poorer countries of the develop-

ing world. Its use has led to mercury poisoning of the workers and the local water

supplies. Other amalgams include zinc and cadmium, used in batteries to in-

crease their lifetime (see Chapter 19).

13.3 Ceramics

According to the American Academy of Orthopedic Surgeons, most of us will

have two bone fractures in our lifetime, and many of these accidents will require

immediate medical attention. Some of the casts used to hold fractured bones in

place as they heal are made of ﬁberglass or plastic, but the less expensive plaster

cast is still widely used. The plaster, also known as plaster of Paris, is a ceramic

material made from calcium sulfate hemihydrate (CaSO

⁄2 H

O) and water.

Why does plaster harden? This is the subject of the accompanying NanoWorld/

MacroWorld feature.

Ceramics are a class of nonmetallic materials that are made of inorganic com-

pounds. They can be either amorphous or crystalline in structure. Examples of

ceramics include glass, clay pottery, cement, bricks, and china dinner plates. Be-

cause ionic bonds exist within a ceramic material, the band gap is typically quite

large. These compounds therefore act as insulators with low electrical conductiv-

ity, high melting point, and a lack of malleability and ductility, but they do have

good resistance to corrosion. With the exception of glass, melting does not

reshape ceramics. Typically, a mixture of the ﬁnely ground ceramic and some

binders is used to mold the material into the desired shape. After heating of the

new shape (called ﬁring), the ceramic becomes hard. Compare the properties of

the ceramics with those of other materials (see Table 13.5 on page 559).

Research into the uses of ceramics as bone replacements or sup-

plements is currently under way. One goal of the research is the

ability to replace badly damaged bones with synthetic material that

later becomes bone, rather than using bone from elsewhere in the

patient’s body. Ceramics made from hydroxyapatite (see Chapter 8)

show some promise in this area because they contain the same

sponge-like structure as real bone. This enables new bone to grow

and integrate into or replace the ceramic material.

The large band gap and the relatively low coefﬁcient of thermal

expansion in a ceramic material allow design engineers to use

ceramics in places where insulation against extremely high tempera-

tures is needed. Of particular interest is the use of ceramics as

insulation on the space shuttle. During a typical mission, the space

shuttle’s hull is subjected to temperatures of –156°C during orbit

and up to 1650°C as the shuttle reenters the Earth’s atmosphere.

13.3 Ceramics 557

Application

HEMICAL

ENCOUNTERS:

Dental Amalgams

Hydroxyapatite

(above) can be

formed into

replacements for

bones, as seen in

the X-ray (left).

Application

Video Lesson: Ceramics and

Glass