Transcript Chemistry 1011

Chemistry 1011
TOPIC
Physical Properties of Matter
TEXT REFERENCE
Masterton and Hurley Chapter 9
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9.2 Network Covalent, Ionic, and
Metallic Solids
YOU ARE EXPECTED TO BE ABLE TO:
• Classify non-molecular solids as either
network covalent solids, ionic solids, or
metallic solids.
• Relate the physical properties of nonmolecular solids to the forces holding them
together.
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Network Covalent, Ionic, and
Metallic Solids
• Almost all substances that are gases or
liquids at room temperature and pressure
(25oC and 1 atm) are molecular
• There are three distinct types of nonmolecular solids
– Network covalent solids
– Ionic solids
– Metallic solids
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Other Materials
• There are also other materials with complex
structures, for example
– Polymers – natural and manufactured
– Biological materials
– Semiconductors
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Network Covalent Solids
• High melting points
– To melt such a solid, covalent bonds must be broken
• Insoluble in all common solvents
– For solution to occur, covalent bonds must be broken
• Poor electrical conductors
– There are no mobile electrons (EXCEPT Graphite)
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Some Examples - Carbon
• Carbon can exist as both diamond and graphite
• Both forms are network covalent solids
• This site shows diagrams of both diamond and
graphite structures:
http://www.education.eth.net/acads/chemistry/allotropic_carb
on-I.htm
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Diamond
• Diamond – pure carbon
– Three dimensional structure based on tetrahedrally
bonded carbon atoms. All single bonds
• Very hard
• Very high melting point
• Does NOT conduct electricity
– Diamond is transparent, strong, and very hard. It is a
superb cutting tool. The atoms in diamond are very
strongly held in position.
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Diamond Structure
• http://www.rensselaer.edu/dept/chemeng/WWW/faculty/cale/classes/SMP/lecture1/sld014.htm
• http://iumsc11.chem.indiana.edu/common/Minerals/diamo
nd/diamond.htm
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Carbon
• Graphite – pure carbon
– Two dimensional, layer structure based on
carbon atoms bonded in triangular planar
arrangement. Each carbon atom forms 2 single
and 1 double bond with its neighbours
– Layers held together by dispersion forces
• Very soft (layers slip)
• Very high melting point
• Conducts electricity as a result of delocalized
electrons from double bonds
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Graphite Structure
• Graphite is black, soft, and an excellent lubricant.
• It is easy to separate layers of atoms in graphite, to
make them slide past one another.
.
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Other Network Covalent Solids
• Quartz SiO2
– Three dimensional array of Si and O atoms
– Major component of sand
– Heated with CaCO3 and Na2CO3 gives glass
• Other silicon – oxygen structures include
– Mica
– Asbestos
– Talc
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Ionic Solids - Observations
•
All ionic compounds:
1. are solids at room temperature
2. have high melting points
3. have wide temperature range for liquid
phase
4. are hard, but brittle
5. conduct electricity in the liquid phase and
in solution
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Ionic Solids
• High melting and non volatile
– Strong interionic attractions must be overcome in order
to separate ions of opposite charge
• Do not conduct in solid phase
– Ions must move in order for the material to conduct
electricity. Ions can only move if the substance is
melted or dissolved in water
• Many dissolve in water, but not in organic
solvents
– Water is a highly polar substance
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Attractions Between Ions
• The force of attraction between oppositely
charged ions depends on
– The size of the charge
• Greater charge = greater force of attraction,
eg CaO vs NaCl
– The size of the ions
• Smaller ions get closer together = greater
force of attraction, eg NaCl vs KBr
http://www.scs.uiuc.edu/~chem315/solidstatestructures/naclf.
htm
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Metals - Observations
•
All metals:
1. are solids at room temperature (except for
mercury); melting points vary
2. have a wide temperature range for liquid phase
3. are ductile and malleable
4. conduct electricity in the solid and liquid phase,
but not in the gaseous phase
5. have high thermal conductivity
6. are insoluble in water or other solvents
7. are shiny
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Metal Structures
• Solid metals consist of regular arrays of
metal atoms forming a crystal lattice
• Delocalized valence electrons form a “sea
of electrons” that can move freely
throughout the metal
• Delocalization provides stability
• The “electron sea” is also present in the
liquid phase
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Explaining Metal Properties
• Variable melting points
– Melting point depends upon the strength of the bonds.
This in turn generally depends upon the number of
valence electrons. However, there is no simple
correlation
• High liquid temperature range
– Vaporizing a metal requires the valence electrons to
become localized, ie the metal bonds must be broken
• Ductility and malleability
– Metal crystal structures are flexible, layers in the crystal
lattice can slide
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Explaining Metal Properties
• Conduction of electricity
– Delocalized electrons are mobile.
• Conduction of heat
– Heat is carried through metals by colliding electrons
• Solubility
– Electrons cannot go into solution, nor cations by
themselves
• Shininess
– Delocalized electrons absorb and emit light over a wide
wavelength range
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On-line materials
• http://www.beyondbooks.com/psc92/3.asp
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