Chapter 3
Structures of Metals and Ceramics
QUESTIONS AND PROBLEMS
Fundamental Concepts
W3.1 What is the difference between atomic structure and crystal structure?
Unit Cells
Metallic Crystal Structures
W3.2 Show for the body-centered cubic crystal structure that the unit cell edge length a and the
atomic radius R are related through a =4R/ 3 .
W3.3 For the HCP crystal structure, show that the ideal c/a ratio is 1.633.
W3.4 Show that the atomic packing factor for HCP is 0.74.
Density Computations—Metals
W3.5 Calculate the radius of a tantalum atom, given that Ta has a BCC crystal structure, a
density of 16.6 g/cm3, and an atomic weight of 180.9 g/mol.
W3.6 Titanium has an HCP crystal structure and a density of 4.51 g/cm3.
(a) What is the volume of its unit cell in cubic meters?
(b) If the c/a ratio is 1.58, compute the values of c and a.
W3.7 Niobium has an atomic radius of 0.1430 nm and a density of 8.57 g/cm3. Determine
whether it has an FCC or BCC crystal structure.
W3.8 The unit cell for uranium has orthorhombic symmetry, with a, b, and c lattice parameters
of 0.286, 0.587, and 0.495 nm, respectively. If its density, atomic weight, and atomic
radius are 19.05 g/cm3, 238.03 g/mol, and 0.1385 nm, respectively, compute the atomic
packing factor.
W3.9 Beryllium has an HCP unit cell for which the ratio of the lattice parameters c/a is 1.568.
If the radius of the Be atom is 0.1143 nm, (a) determine the unit cell volume, and (b)
calculate the theoretical density of Be and compare it with the literature value.
W3.10 Cobalt has an HCP crystal structure, an atomic radius of 0.1253 nm, and a c/a ratio of
1.623. Compute the volume of the unit cell for Co.
Ceramic Crystal Structures
W3.11 For a ceramic compound, what are the two characteristics of the component ions that
determine the crystal structure?
W3.12 Show that the minimum cation-to-anion radius ratio for a coordination number of 6 is
0.414. [Hint: Use the NaCl crystal structure (Figure 3.5), and assume that anions and
cations are just touching along cube edges and across face diagonals.]
W3.13 On the basis of ionic charge and ionic radii, predict crystal structures for the following
materials:
(a) MnS, and
(b) CsBr. Justify your selections.
W3.14 Which of the cations in Table 3.4 would you predict to form fluorides having the cesium
chloride crystal structure? Justify your choices.
W3.15 Given one set of ionic radii for cations and another set for anions, generate a spreadsheet
that will allow the user to predict which ion pairs will form the following crystal
structures: sodium chloride, cesium chloride, zinc blende, and fluorite.
Density Computations—Ceramics
W3.16 The unit cell for Al2O3 has hexagonal symmetry with lattice parameters a =0.4759 nm
and c =1.2989 nm. If the density of this material is 3.99 g/cm3, calculate its atomic
packing factor. For this computation use ionic radii listed in Table 3.4.
W3.17 Calculate the theoretical density of NiO, given that it has the rock salt crystal structure.
W3.18 (a) Using the ionic radii in Table 3.4, compute the theoretical density of CsCl. (Hint: Use
a modification of the result of Problem W3.2.)
(b) The measured density is 3.99 g/cm3. How do you explain the slight discrepancy
between your calculated value and the measured one?
W3.19 From the data in Table 3.4, compute the theoretical density of CaF2, which has the
fluorite structure.
W3.20 The unit cell for Fe3O4 (FeO-Fe2O3) has cubic symmetry with a unit cell edge length of
0.839 nm. If the density of this material is 5.24 g/cm3, compute its atomic packing factor.
For this computation, you will need to use ionic radii listed in Table 3.4.
Silicate Ceramics
W3.21 In terms of bonding, explain why silicate materials have relatively low densities.
Carbon
W3.22 Compute the theoretical density of diamond given that the C—C distance and bond angle
are 0.154 nm and 109.5°, respectively. How does this value compare with the measured
density?
Crystal Systems
W3.23 Below is a unit cell for a hypothetical metal.
(a) To which crystal system does this unit cell belong?
(b) What would this crystal structure be called?
(c) Calculate the density of the material, given that its atomic weight is 141 g/mol.
Point Coordinates
W3.24 List the point coordinates for all atoms that are associated with the FCC unit cell (Figure
3.1).
W3.25 List the point coordinates of both the zinc and sulfur atoms for a unit cell of the zinc
blende crystal structure (Figure 3.7).
W3.26 Using the Molecule Definition Utility found in both “Metallic Crystal Structures and
Crystallography” and “Ceramic Crystal Structures” modules of VMSE, located on the
book’s web site [www.wiley.com/college/callister (Student Companion Site)], generate
(and print out) a three-dimensional unit cell for β tin given the following: (1) the unit cell
is tetragonal with a =0.583 nm and c =0.318 nm, and (2) Sn atoms are located at the
following point coordinates:
000
011
100
1
2
0
3
4
110
1
2
1
3
4
010
1
1
2
1
4
001
0
1
2
1
4
101
1
2
1
2
1
2
111
Crystallographic Directions
W3.27 Draw an orthorhombic unit cell, and within that cell a [2 1 1] direction.
W3.28 What are the indices for the direction indicated by the vector in the sketch below?
W3.29 Within a cubic unit cell, sketch the following directions:
(a) [101],
(c) [ 1 1 1 ],
(b) [211],
(d) [3 1 2]
W3.30 Determine the indices for the directions shown in the following cubic unit cell:
W3.31 Convert the [110] direction into the four-index Miller–Bravais scheme for hexagonal unit
cells.
W3.32 Determine the indices for the two directions shown in the following hexagonal unit cell:
Crystallographic Planes
W3.33 Draw a monoclinic unit cell, and within that cell a (200) plane.
W3.34 What are the indices for the plane drawn in the sketch below?
W3.35 Sketch within a cubic unit cell the following planes:
(a)
(3 1 2),
(c) ( 1 1 1 ),
(b)
(301),
(d) ( 2 12).
W3.36 Determine the Miller indices for the planes shown in the following unit cell:
W3.37 Cite the indices of the direction that results from the intersection of each of the following
pair of planes within a cubic crystal: (a) (110) and (111) planes, (b) (110) and (1 1 0)
planes, and (c) (11 1 ) and (001) planes.
W3.38 Sketch the atomic packing of (a) the (100) plane for the FCC crystal structure, and (b) the
(111) plane for the BCC crystal structure (similar to Figures 3.26b and 3.27b).
W3.39 For each of the following crystal structures, represent the indicated plane in the manner
of Figures 3.26 and 3.27, showing both anions and cations:
(a) (100) plane for the cesium chloride crystal structure,
(b) (200) plane for the cesium chloride crystal structure.
W3.40 Consider the reduced-sphere unit cell shown in Problem W3.23, having an origin of the
coordinate system positioned at the atom labeled with an O. For the following sets of
planes, determine which are equivalent:
(a) (100), (0 1 0), and (001)
(b) (110), (101), (011), and ( 1 01)
(c) (111), (1 1 1), (11 1 ), and ( 1 1 1 )
W3.41 Here are three different crystallographic planes for a unit cell of a hypothetical metal. The
circles represent atoms:
(a) To what crystal system does the unit cell belong?
(b) What would this crystal structure be called?
W3.42 Convert the (111) plane into the four-index Miller–Bravais scheme for hexagonal unit
cells.
W3.43 Determine the indices for the planes shown in the hexagonal unit cells below:
W3.44 Sketch the (01 1 1) plane in a hexagonal unit cell.
Linear and Planar Densities
W3.45 (a) Derive linear density expressions for BCC [110] and [111] directions in terms of the
atomic radius R.
(b) Compute and compare linear density values for these same two directions for iron.
W3.46 (a) Derive planar density expressions for FCC (100) and (111) planes in terms of the
atomic radius R.
(b) Compute and compare planar density values for these same two planes for aluminum.
W3.47 (a) Derive the planar density expression for the HCP (0001) plane in terms of the atomic
radius R.
(b) Compute the planar density value for this same plane for titanium.
Close-Packed Crystal Structures
W3.48 The zinc blende crystal structure is one that may be generated from close-packed planes
of anions.
(a) Will the stacking sequence for this structure be FCC or HCP? Why?
(b) Will cations fill tetrahedral or octahedral positions? Why?
(c) What fraction of the positions will be occupied?
W3.49 Beryllium oxide (BeO) may form a crystal structure that consists of an HCP arrangement
of O2– ions. If the ionic radius of Be2+ is 0.035 nm, then
(a) Which type of interstitial site will the Be2+ ions occupy?
(b) What fraction of these available interstitial sites will be occupied by Be2+ ions?
Polycrystalline Materials
W3.50 Explain why the properties of polycrystalline materials are most often isotropic.
X-Ray Diffraction: Determination of Crystal Structures
W3.51 Using the data for aluminum in Table 3.1, compute the interplanar spacing for the (110)
set of planes.
W3.52 The metal rhodium has an FCC crystal structure. If the angle of diffraction for the (311)
set of planes occurs at 36.12° (first-order reflection) when monochromatic x-radiation
having a wavelength of 0.0711 nm is used, compute (a) the interplanar spacing for this
set of planes, and (b) the atomic radius for a rhodium atom.
W3.53 For which set of crystallographic planes will a first-order diffraction peak occur at a
diffraction angle of 44.53° for FCC nickel when monochromatic radiation having a
wavelength of 0.1542 nm is used?
W3.54 For an x-ray diffraction pattern (having all peaks are plane-indexed) of a metal that has a
unit cell of cubic symmetry, generate a spreadsheet that allows the user to input the x-ray
wavelength, and then determine, for each plane, (a) dhkl and (b) the lattice parameter, a.
W3.55 The diffraction peaks shown in Figure 3.39 are indexed according to the reflection rules
for FCC (i.e., h, k, and l must all be either odd or even). Cite the h, k, and l indices of the
first four diffraction peaks for BCC crystals consistent with h + k + l being even.
Noncrystalline Solids
W3.56 Would you expect a material in which the atomic bonding is predominantly ionic in
nature to be more or less likely to form a noncrystalline solid upon solidification than a
covalent material? Why? (See Section 2.6.)