1. CHE-30043 Materials Chemistry & Catalysis : Solid State Chemistry lecture 1
Rob Jackson
LJ1.16,
@robajackson

2. Supplementary (detailed)
Background reading
Recommended
‘Reactions and Characterisation of Solids’
Sandra Dann
(RSC Tutorial Chemistry Text No 2 (2000))
Supplementary (detailed)
‘Solid State Chemistry: An Introduction’
Lesley Smart, Elaine Moore
(4th edition, CRC Press (2012))
che lecture 1

3. Lecture 1 contents Key concepts of basic crystallography
Refer to che slides
Bonding in different structures
Metals, ionic structures, covalent structures, molecular structures
Correlation of structure & bonding with properties
Electrical conductivity in solids
che lecture 1

4. Key concepts of basic crystallography
Important concepts from che-20031:
Unit cells
Crystal class
Lattice type
Bravais lattices
Review these!
che lecture 1

5. Metal Structures - Characteristics
Most adopt close-packed structures.
Bonding is non-directional.
This is reflected in mechanical properties.
Cations are close packed and surrounded by a ‘sea’ of valence electrons – electrical conductors.
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6. Examples of metal structures
Alkali metals tend to adopt the body centred cubic (BCC) structure:
 2d view
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7. Other metal structures: hcp and ccp
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8. Ionic structures - characteristics
Structures minimise ionic repulsion.
Structures influenced by ionic radii.
Strong directional bonding.
Valence electrons are involved in ionic bonding (electron transfer).
Electrical insulators.
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9. The diversity of structures found
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10. Some ionic structures - 1
We will consider a few examples to illustrate the diversity of structures.
e.g. the metal halide structures, where MX can have at least 4 different structures depending on M and X.
NaCl and CsCl – why different?
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11. NaCl and CsCl structures compared
Note: CsCl is not BCC – why?
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12. Some ionic structures - 2
MX2 also has several forms, but an important structure is the fluorite structure, named after the mineral, fluorite, CaF2.
This is an important structure adopted by many technologically important materials, including ZrO2 and UO2.
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13. The fluorite structure: cations at alternate cube centres (shown for MF2)
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14. Covalent Structures Covalently bonded materials.
Strong bonds – can lead to structures of high strength.
Valence electrons involved in shared covalent bonds.
Electrical insulators or semiconductors.
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15. Examples of covalent solids
The diamond phase of carbon is a good example.
Each C atom is bonded to 4 others through sp3 hybrid bonding orbitals, giving a 3-dimensional network
Other examples include Ge, Si, B, P, As, Se, Te.
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16. Carbon phases – bonding and properties
Diamond – rigid 3-dimensional (sp3) covalent network with all bonds equivalent.
Graphite – layer structure – covalent (sp2) bonding within each layer, and relatively weak van der Waals forces between the layers. This gives the material its distinct properties.
What about C60?
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17. Diamond and graphite structures compared
Diamond structure
continuous 3d network
Graphite structure
2 dimensional layer structure
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18. Buckminsterfullerene (C60)
From the structure, how many C atoms is each atom bonded to?
From this information, what type of hybridisation would you expect for the C bonding orbitals?
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19. Molecular Solids
Solids formed from molecules which retain their identity and shape in the solid
Structures held together by van der Waals forces and electric multipole moments (depends on symmetry)
Bonding is weak – low melting points
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20. Examples of molecular solids
Classic example is solid CO2, ‘dry ice’
Linear O=C=O molecules in FCC lattice
Structure ‘melts’ (sublimes*) at low temperatures, indicating that the bonding is weak
Other examples – solid benzene, etc.
* goes from solid to gas phase direct
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21. (atmospheric pressure)
The solid CO2 structure
The individual CO2 molecules form an FCC lattice.
The structure is held together by van der Waals forces.
Tsub = -78.5°C
(194.5 K)
(atmospheric pressure)
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22. Example: the structure of solid CCl4
CCl4 melts at -23C (250K).
Below this temperature, it forms a simple cubic structure with a CCl4 molecule at each lattice site.
Molecules retain their identity in the solid.
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23. Bonding in Solids summarised
Metals – cations held together by delocalised valence electrons
Ionic solids – balance of repulsive and attractive interactions
Covalent solids – covalent bonds between atoms
Molecular solids – non-bonded electrostatic interactions
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24. Electrical conductivity in solids
Does a given solid state material behave as conductor, insulator or semiconductor?
How is this related to its electronic structure?
We will separately look at:
Metals
Insulators (ionic solids, covalent solids)
Semiconductors (to be defined later)
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25. Electrical Conductivity and Temperature
In a metal, conductivity decreases with temperature. Why?
In insulators and semiconductors, conductivity increases with temperature.
Why the difference between metals and insulators/semiconductors?
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26. Electron conductivity in metals – (i)
Electrons in a metal with no current flowing. Valence electrons are free to move, but randomly.
Electrons in a metal with a current flowing. Valence electrons move in the direction of the current
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27. Electron conductivity in metals – (ii)
Why does the electrical conductivity of a metal reduce with increasing temperature?
The cations move about their lattice positions and ‘get in the way’ of the electrons.
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28. Summary of lecture Key concepts from crystallography have been listed.
Some metallic, ionic, covalent and molecular structures have been reviewed.
Electrical conductivity in materials has been introduced, with metals as an example.
che lecture 1