the crystalline solid state chapter 7. crystalline solid state many more “molecules” in the...
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The Crystalline Solid State
Chapter 7
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Crystalline Solid State
• Many more “molecules” in the solid state.– We will focus on crystalline solids composed of
atoms or ions.
• Unit cell – structural component that, when repeated in all directions, results in a macroscopic (observable) crystal.– 14 possible crystal structures (Bravais lattices)– Discuss positions of atoms in the unit cell.
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The Cubic Unit Cell (or Primitive)
• 1 atom per unit cell (how?).
• What is the coordination number? Volume occupied?
• Let’s calculate the length of the edge. What size of sphere would fit into the hole?
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The Body-Centered Cubic
• How many atoms per unit cell?
• What is the length of the edge? This is a more complicated systems than the simple cubic.
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Close-Packed Structures
• How many atoms is each atom surrounded by in the same plane?
• What is the coordination number?• Hexagonal close packing (hcp) – discuss the third
layer (ABA).• Cubic close packing (ccp) or face-centered cubic
(fcc) – discuss the third layer (ABC).• Two tetrahedral holes and one octahedral hole per
atom. Can you see them?
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Close-Packed Structures
• The hcp has hexagonal prisms sharing vertical faces (Figure).– How many atoms per unit cell in the hcp structure?– What is the length of the cell edge?
• The unit cell for the ccp or fcc is harder to see.– Need four close-packed layers to complete the cube.– What is the length of the cell edge?
• In both close-packed structures, 74.1% of the total volume is occupied.
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Ionic Crystals
• The tetrahedral and octahedral holes can have varying occupancies.
• Holes are generally filled by smaller ions.– Tetrahedral holes– Octahedral holes
• NaCl structure
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Metallic Crystals
• Most crystalize in bcc, ccp, and hcp structures.
• Hard sphere model does not work well.– Depends on electronic structure.
• Properties– Conductivity– Dislocations
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Diamond
• Each carbon atom is bonded tetrahedrally to four nearest neighbors (Figure).– Essentially the same strength in all directions.
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Structures of Binary Compounds
• Close-packed structures are generally defined by the larger ions (usually anions). The oppositely-charged ions occupy the holes.
• Two important factors in considering the structure– Radius ratio (r+/r-)– Relative number of combining cations and anions.
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NaCl Crystal Structure
• Face-centered cubes of both ions offset by a half a unit cell in one direction.
• Many alkali metals have this same geometry.
• What is the coordination number (nearest neighbor)?
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CsCl Crystal Structure
• Chloride ions form simple cubes with cesium ions in the center (Figure 7-7).
• The cesium ion is able to fit in to center hole. How?
• Other crystal structures.
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TiO2 (the rutile structure)
• Distorted TiO6 octahedra.– Ti has a C.N. of 6,
octahedral coordination
– O has a C.N. of 3
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Rationalization of Structure of Crystalline Solids
• Predicting coordination number from radius ratio (r+/r-).– A hard sphere treatment of the ions.– Treats bonding as purely ionic.– Simply, as as the M+ ratio increases, more
anions can pack around it.• Table 7-1.
Let’s look at a few (NaCl, CaF2, and CaCl2).
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Thermodynamics of Ionic Crystal Formation
• A compound tends to adopt the crystal structure corresponding to lowest Gibbs energy.M+(g) + X-(g) MX(s)G = H - TS (standard state), 2nd term can be
ignored
• Lattice enthalpyMX(s) M+(g) + X-(g) HL (standard molar
enthalpy change)Currently, we are interested in lattice formation.
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The Born-Haber Cycle
• A special thermodynamic cycle that includes lattice formation as one step.
• The cycle has to sum up to zero if written appropriately.
• Write down values for KCl.
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The Born-Haber Cycle
• Calculate the lattice enthalpy for MgBr2.
• A discrepancy between this value and the real value may indicate the degree of covalent character.– We have assumed Coulombic interactions
between ions.
– The actual values for KCl and MgBr2 are 701 and 2406 kJ/mol (versus 720 and 2451).
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Lattice Enthalpy Calculations
• Considering only Coulombic contributions– The electrostatic potential energy between each pair.
zA, zB = ionic charges in electron units
r0 = distance between ion centerse = electronic charge
4o = permittivity of a vacuum
e2/ 4o = 2.307 10-28 J m
Calculation would be performed on each cation/anion pair (nearest neighbor).
o
2
0
BA
4
e
r
zzU
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Lattice Enthalpy Calculations
• A more accurate equation depicts the Coulombic interactions over the entire crystal.
NA = Avogadro’s constant
A = Madelung’s constant, value specific to a crystal type (in table). This is a sum of all the geometric factors carried out until the interaction become infinitesimal.
0
BA
o
2A
r
zz
4
eNU
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Lattice Enthalpy Calculations
• Repulsions between ions in close proximity term.
C’ = constant (will cancel out when finding the minimum) = compressibility constant, ~ 30 pm
• Combining terms
0r
'A eCNU
0r
A0
BA
o
2A e'CNA
r
zz
4
eNU
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Lattice Enthalpy Calculations
• Finding the minimum energy– dU/dr0 = O
• A negative of this value may be defined as the lattice enthalpy.
Ar
1r4
ezzNU
00o
2BAA
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Lattice Enthalpy Calculations
• As the polarizability of the resultant ions increase the agreement with this ionic model worsens.– Polarizibility generally indicates more covalent
character.
Calculations
NaCl and CaBr2
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Molecular Orbitals in Solids
• A very large number of atoms are used to generate molecular orbitals.– One-dimensional model.
– Creation of bands that are closely spaced.
– Factors affecting the width of the band.
This would be called an ‘s band’. A similar model can be constructed for the p-orbitals and d-orbitals.
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The Bonding Picture in Solids
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Molecular Orbitals in Solids
• Band gap – separation between bands in which no MOs exist (Figure 7-13).
• Valence band – highest energy band containing electrons.
• Conduction band – the band immediately above the valence band in energy.
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Metals and Insulators
• Metals– Partially filled valence band (e.g. s band)
• Electrons move to slightly higher energy levels by applying a small voltage. Electrons and ‘holes’ are both free to move in the metal.
– Overlapping bands (e.g. s and p bands)• If the bands are close enough in energy (or
overlapping) an applied voltage can cause the electrons to jump into the next band (conduction band).
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Density of States
• Concentration of energy levels within a band.
• Helps to describe bonding/reactivity in solids.
dE
))E(N(d
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Conductivity of Solids Versus Temperature
• Metals – decrease with temperature.
• Semiconductors – increase with temperature.
• Insulators – increase with temperature (if measurable).
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Semiconductor Types
• Intrinsic semiconductors – pure material having semiconductive properties.
• Doped semiconductors – semiconductors that are fabricated by adding a small amount of another element with energy levels close to the pure state material.– n-type semiconductors– p-type semiconductors (look at
figure)
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Semiconductors
• Fermi-level (semiconductor) – the energy at which an electron is equally likely to be in each of two levels (Figure).
• Effects of dopants on the Fermi level.– n-type and p-type.
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Diodes (creating p-n junctions)
• Migration of electrons from the n-type material to the p-type material.– Equilibrium is established due to charge transfer.
• Application of a negative potential to the n-type material and a positive potential to the p-type material.– Discuss (Figure 7-16).
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Superconductivity
• No resistance to flow of electrons.– Currents started in a loop will continue to flow indefinitely.
• Type I superconductors – expel all magnetic fields below a critical temperature, Tc (Meisner effect).
• Type II superconductors – below a critical temperature exclude all magnetic fields completely. Between this temperature and a second critical temperature, they allow partial penetration by the magnetic field.– Levitation experiment works well.
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Theory of Superconducting
• Cooper pair theory– Bardeen, Cooper, and
Schrieffer
– Electrons travel through the material in pairs.
– The formation and propagation of these pairs is assisted by small vibrations in the lattice.
• discuss
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YBa2Cu3O7 High-Temperature Superconductors
• Discovered in 1987 and has a Tc of 93 K.– N2(l) can be used
• Type II superconductor.
• Difficult to work with.
• Possesses copper oxide planes and chains.
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Bonding in Solid State Structures
• The hard-sphere model is too simplistic.– Deviations are observed in ion sizes.– Sharing of electrons (or transfer back to the
cation) can vary depending upon the polarizability.
• LiI versus NaCl (which structure would exhibit more covalent character?)
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Bonding in TiO2
• The crystal has a rutile structure.– Each titanium has ___ nearest
neighbors and each oxygen atom has ___ nearest neighbors.
• There is no effective O···O or Ti···Ti interactions (only Ti···O interactions). Why?
• The structure consists of TiO6 fragments (discuss).
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Bonding in TiO2
For a TiO6 monomer (no significant -bonding).
An approximation of the ‘bands in the solid structure.
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Bonding in TiO2
• The calculated DOS curve in 3-d space is slightly more complicated.
• The O 2s, O2p, Ti t2g, and eg bands are well separate. The separation predicts that this material has ‘insulator-like’ properties.
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Bonding in TiO
• Several of the 3d monoxides illustrate high conductivity that decreases with temperature.– TiO and VO (positioning in the
table).
• TiO adopts the rocksalt structure (NaCl).– Discuss geometry and
consequences on bonding.
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Bonding in TiO
• The titanium atoms are close enough to form a ‘conduction’ band.– Overlap of t2g orbitals of
the metal ions in neighboring octahedral sites.
– Illustrated for dxy orbitals.
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Bonding in TiO
• The calculated DOS curve for TiO reveals that the bonds aren’t well separated.– Diffuse bands indicate
more conductive behavior.
• Why is TiO2 different than TiO?
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Bonding in TiO
• MnO, FeO, CoO, and NiO do not conduct, but they have the same basic structure. Why?
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Imperfections in Solids
• All crystalline solids possess imperfections.– Crystal growth occurring at many sites causes
boundaries to form.– Vacancies and self-interstitials– Substitutions– Dislocations
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Silicates• The earth’s crustal rocks (clays, soils, and sands)
are composed almost entirely (~95%) of silicate minerals and silica (O, Si, and Al).– There exist many structural types with widely
varying stoichiometries (replacement of Si by Al is common). Consequences?
• Common to all:– SiO4 tetrahedra units
• Si is coordinated tetrahedrally to 4 oxygens
http://www.soils.wisc.edu/virtual_museum/displays.html
http://mineral.galleries.com/minerals/silicate/class.htm
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The Tetrahedral SiO4 Unit
Cheetham and Day
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Structures with the SiO4 Unit
• Discrete structural units which commonly contain cations for charge balance.
• Corner sharing of O atoms into larger units.– O lattice is usually close-packed (near)– Charge balance is obtained by presence of
cations.
Individual units, chains, multiple chains (ribbons), rings, sheets and 3-d networks.
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Structure Containing Discrete Units
• Nesosilicates – no O atoms are shared.– Contain individual SiO4
4- units.– ZrSiO4 (zircon) – illustrate with softwares
• Stoichiometry dictates 8-fold coordination of the cation.
– (Mg3 or Fe3)Al2Si3O12 (garnet) – illustrate with softwares
• 8-fold coordination for Mg or Fe and 6-fold coordination for the Al.
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Structure Containing Discrete Units
• The sorosilicates (disilicates) – 1 O atom is shared.– Contain Si2O7
6- units
– Show Epidote (Ca2FeAl2(SiO4)(Si2O7)O(OH)) with softwares.
• Epidote contains SiO44- and Si2O7
6- units
– Near linear Si-O-Si bond angle between tetrahedra.
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Cyclosilicates (discrete cyclic units)
• Each SiO4 units shares two O atoms with neighboring SiO4 tetrahedra.– Formula – SiO3
2- or [(SiO3)n]2n- (n=3-6 are the most common.
– Beryl – six-linked SiO4 tetrahedra (show with softwares). • Be3Al2(SiO3)6 – contains Si6O18
12- cyclic units• The impurities produce its colors.
– Wadeite – three-linked SiO4 tetrahedra (don’t have an actual picture)
• K2ZrSi3O9
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Silicates with Chain or Ribbon Structures
• Corner sharing of SiO4 tetrahedra (SiO3
2-)
– Very common (usually to build up more complicated silicate structures).
• Differing conformations can be adopted by linked tetrahedra.– Changes the repeat distance.
– The 2T structure is the most common (long).
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Silicates with Chain or Ribbon Structures
• The chains are usually packed parallel to provide sites of 6 and 8 coordination for the cations.– Jadeite [NaAlSi2O6]
• Illustrate the different repeat units.
• What is the repeat unit?
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Silicate Chains Linking Together
• Can form double or triple chains/ribbons linked together (or more).
• Depends on the repeat unit in the chain.
• Tremolite [Ca2Mg5(Si4O11)2(OH)2
(illustrate with softwares)
Asbestos mineral (fibrous)
• Triple chain
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Phyllosilicates (Silicates with Layer Structures)
• Clay minerals, micas, talc, soapstone.• Individual layers are formed by sharing 3 of
the 4 atoms of each tetrahedron.• Simplest structure is made up of a 2T
network of silicate chains to give a network composition of Si2O5
2-.– This is exhibited with kaolinite (illustrate the
silicon tetrahedral layer).
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Creation of Layers in the Phyllosilicates
• Can be formed by sharing the fourth O atom between pairs of tetrahedra.– Produces an SiO2
stoichiometry (neutral)
– Replacing Si with Al• Al2Si2O8
2-; requires charge balance. The cations connect the double layers.
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Creation of Layers in the Phyllosilicates
• Double layers can be produced by interleaving layers of the gibbsite Al(OH)3 or brucite Mg(OH)2 structure.– Incorporation of gibbsite produces kaolinite,
[Al2(OH)4Si2O5] (China clay); illustrate with software the different layers present.
– Placing a SiO layer on the other side of the AlO layer produces pyrophyllite, [Al2(OH)2Si4O10].
• Illustrate both with software.
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More Layered Structures
• The Al can be replaced by Mg (2:3) ratio.– Kaolinite serpentine asbestos
– Pyrophyllite talc
• Charged layers can also result by replacing the framework Si with Al or other cations. For charge balance these layers can be interleaved with M(+1) or M(+2) to give micas (illite) or by layers of hydrated cations to give montmorillonite.– Illustrate both.
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The Tectosilicates
• Each oxygen atom is shared by 2 tetrahedra (SiO2 formula).
• Silica (-quartz; one crystalline form)– Si-O-Si bond angles are ~144 degrees.
– Contains helical chains of SiO4.• Six combine to form hexagonal shape (illustrate).
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The Tectosilicates (Zeolites - aluminosilicates)
• A large fraction of the Si atoms are replaced with Al (other metals can also be used). – Charge balance will be required (Si,Al)nO2n.
• Contain cavities that allow molecules to enter.– Able to tailor electronic and physical properties.
• Pore structure and cation exchange.
• Illustrate with software.