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Warren Mar
17/18/19
Chapter 16 Liquids And Solids
16.1 Intermolecular Forces
 Intramolecular bonding
 Condensed states of matter – liquid and solids
A. Dipole – Dipole Forces
 Dipole moments which line up molecules
 1% as strong as covalent or ionic bonds
 Hydrogen bonding - Hydrogen and electronegative atom, usually nitrogen, oxygen
or fluoride , strong dipole forces.
B. London Dispersion Forces
 Molecules without dipole moments exert forces on each other, even the noblest of
gases
 Instantaneous dipole can induce dipoles in other atoms
 Large atoms with many electrons exhibit a high polarizability than small atoms.
16.2 The Liquid state
 Liquids – low compressibility, lack of rigidity, and high density compared with gases.
 Minimum surface area = sphere
 Surface tension – the resistance of a liquid to increase its surface area.
 Capillary action – the spontaneous rising of a liquid in a narrow tube.
 Cohesive forces – intermolecular forces among molecules of the liquid.
 Adhesive forces – intermolecular forces b/w the molecules and the container.
 Viscosity – a measure of a liquid’s resistance to flow.
 Large intermolecular forces are highly viscous, Lot of hydrogen bonding.
 Molecular complexity also leads to higher viscosity, because of entanglement
16.3 An Introduction to Structures and Types and Solids
 Crystalline solids – regular arrangement.
 Amorphous solids – considerable disorder
A. X-ray Analysis of Solids
 Diffraction pattern to determine bond length and structure of complex biological
crystals.
 Bragg equation, n = 2dsin, d bond distance,  is the angel of incidence.
B. Types of crystalline Solids
 Atomic, Ionic, Molecular. Different lattice points.
16.4 Structure and Bonding In Metals
 Hexagonal closest packed (hcp)
 Cubic closest packed (ccp), face-centered

fv = volume occupied by spheres in the unit cell/ volume of unit cell
e
fv
Simple Cubic
2r
.520
Face Centered
8½r
.740
½
Body Centered
4r/3
.680
A. Body-Centered Cubic Packing
 Body-centered cubic (bcc)
B. Bonding in Metals
 “sea” of electrons.
 Strong bonds, non-directional.
 Band model or MO model
C. Metal Alloys
 Alloy – substance that contains a mixture of elements and has metallic properties
 Substitutional alloy – lattice points are replaced by other atoms
 Interstitial alloy – fills the holes
16.5 Carbon and Silicon: Network Atomic Solids
 Allotropes – different forms of same atoms
 Si likes forming single bonds more than double bonds, thus making it different from
Carbon.
 Most silicon compounds contain silicon and oxygen.
A. Ceramics – silica plates lock and forms network
B. Semiconductors – doped silicon to create n and p type semiconductors.
 Used as rectifier
 Conductivity increases as temperature increases
16.6 Molecular Solids
 London Dispersion forces increase as size of molecule increases
16.7 Ionic Solids
 Stable, high-melting, strong electrostatic forces
 Trigonal holes are formed by 3 spheres
 Tetrahedral holes are formed by 4 spheres
 Octahedral holes are formed b/w two sets of 3 spheres
 Trigonal < tetrahedral < octahedral , smallest to biggest
A. Octahedral Holes
 r = .414R, where R is the radius of sphere and r is the radius of hole
B. Tetrahedral Holes
 r = .225R, where R is the radius of sphere and r is the radius of hole
C. Guidelines for Filling Octahedral and Tetrahedral Holes
 Cubic hole, hole in simple cubic
+
Size of M
Type of Hole Filled
0.225R- < r+ <0.414RTetrahedral
0.414R- < r+ <0.732ROctahedral
0.732R- < r+
Cubic
16.8 Structures of Actual Ionic Solids
 Twice as many tetrahedral holes as there are packed spheres in ccp.
 Same number of octahedral holes as the number of packed spheres.
 Unit lattice must have a neutral charge
A. The Structures of the Alkali Halides
 All octahedral holes filled.
 ccp
B. The Structure of Zinc Sulfide
 ½ tetrahedral holes filled in both ccp and hcp.
C. The Structure of Calcium Fluoride
 All tetrahedral holes filled in cpp.
16.9 Lattice Defects
 Point defect – Schottky defect – missing particles
 Frenkel defects – particles migrate, mostly in ccp, ex. Silver halides
 Nonstoichiometric compounds – variations in charge.
16.10 Vapor Pressure and Changes of State
 Vaporization is endothermic
A. Vapor Pressure
 Evaporation and condensation reach equilibrium, Pressure at equilibrium is vapor
pressure.
 Patm = Pvapor + PHg column
 Liquids with high vapor pressures are said to be volatile.
 Vapor pressure determined by intermolecular forces. Large intermolecular = low
vapor pressure.
 Vapor pressure increases significantly with temperature.
 ln(Pvap) = - Hvap/R (1/T) + C, where C is a constant of the liquid.
 ln(PT1vap/PT2vap) = Hvap/R(1/T2 – 1/T1)
B. Changes of State
 Melting point is when solid and liquid have identical vapor pressures, when total
pressure = 1 atm.
 Supercooled – it has not yet achieved the degree of organization necessary to form
solid, solid rapidly forms and releases energy in an exothermic process, which
brings the temperature back up to the melting point, where the rest of the liquid
freezes.
 Superheated – bubble formation in the interior of the liquid requires that many
high-energy molecules gather in the same vicinity. Can cause bumping.
16.11 Phase Diagrams
 Represents the phases of a substance as a function of temperature and pressure.
 Triple point is where all three states can exist.
 Critical temperature – the temperature above which the vapor cannot be liquefied no
matter what pressure is applied.
 Critical pressure – the pressure required to produce liquefaction at the critical
temperature.
 Critical point is where both of the above meet. Beyond this point the transition form
one state to another involves the intermediate “fluid” region, which is unlike both
vapor and liquid.
A. Applications of the Phase Diagram for Water
B. The Phase Diagram for Carbon Dioxide
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