Chpt 10 - Condensed Phases
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Transcript of Chpt 10 - Condensed Phases
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Chpt 10 - Condensed Phases
• Condensed phases• Intermolecular forces• Special bonding - molecular
solids, network solids, metallic• Phase diagrams & Heating
curves• HW: Chpt 10 - pg. 487-496, #s 12, 14-16, 19-
21, 24, 26, 31, 32, 34, 40, 44 Due Fri Dec. 4
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States of Matter
Differences?What do the phases look like?What makes the state of matter at a given temperature?
Intermolecular forces
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Intermolecular Forces
• Intramolecular forces (chemical bonds) - forces that hold atoms together within a molecule
• Intermolecular forces - forces between molecules - aggregate or bulk material - Is it a solid, liquid or gas?– dipole-dipole force (~1% of strength of a bond) – Hydrogen bonding H and N,O,F bond– London dispersion forces
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Dipole-Dipole Force
•Dipole moment – molecules with polar bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge.•Molecules with dipole moments can attract each other electrostatically. They line up so that the positive and negative ends are close to each other.
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Hydrogen Bonding
(a) Polar water molecule (b) hydrogen bonding between water molecules - blue dotted lines
Very strong dipole-dipole force between H and N,O,F(most electro-negative elements)
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Hydrogen bonding graph of covalent hydrides
Why are these interaction forces happening?•Especially polar X-H bond•Small size of N,O, and F allow close approach of dipoles
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London Dispersion Forces
• Weakest of the intermolecular forces
• Important for atoms & non-polar molecules
•As the motion of these atoms and molecules slows (low T) the interaction becomes apparent. •Halogens Trend!!!•Occurs in all molecules even polar ones
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London Dispersion Forces - How?
Moving e- make a momentary nonsymmetric e- distribution, which produces a temporary dipole. This then can induce a similar dipole in a neighboring atom or molecule. Becomes significant for large atoms with large # of electrons.Termed polarizability of an electron cloud.
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Characteristics Intermolecular Forces
• In general, the stronger the intermolecular forces, the higher the melting and boiling points.
• Decrease rapidly with increasing intermolecular distance especially for London dispersion– Nonpolar solids (I2 and CO2) sublimate
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Dry Ice Sublimation at RTemp
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Liquids characteristics• Low compressibility, lack of rigidity, and high density
compared with gases.• Surface tension – resistance of a liquid to an increase
in its surface area: Liquids with large intermolecular forces tend to
have high surface tensions. H2O dropletsPlaying with Hg video YouTube
http://www.youtube.com/watch?v=31CE2BYicyU&feature=fvw
• Capillary action – spontaneous rising of a liquid in a narrow tube: YouTube video water special
http://www.youtube.com/watch?v=CT4pURpXkbY&feature=related
Cohesive forces – intermolecular forces among the molecules of the liquid.
Adhesive forces – forces between the liquid molecules and their container.
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Liquid - Cohesive or adhesive?
Which force dominates alongside the glass tube – cohesive or adhesive forces?
adhesive forces “Like attract like”
determines which will dominate
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Cohesive vs. Adhesive meniscus graphic
Water (polar) interaction with glass surface (polar) and mercury (non-polar) with glass surface (polar)
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Liquids characteristics - cont
• Viscosity – measure of a liquid’s resistance to flow: Liquids with large intermolecular
forces or molecular complexity tend to be highly viscous.
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Solids
• Amorphous solids– Non-uniform structure
• glasses • waxes
• Crystalline solids– Uniform lattice structure (regular
arrangement of atoms) – Unit Cell - smallest repeating unit of the
lattice
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Cubic Unit cell and lattices
X-ray diffraction (crystallography) used to determine arrangement of atoms
= 2 sin λ θn d
n = integer
lambda = wavelength of the X rays
d = distance between the atoms
theta = angle of incidence and reflection
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Bragg Diffraction graphic
= 2 sin λ θn dBragg equation
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Types of Crystalline solids
• Ionic Solids – ions at the points of the lattice that describes the structure of the solid.
• Molecular Solids – discrete covalently bonded molecules at each of its lattice points.
• Atomic Solids – atoms at the lattice points that describe the structure of the solid.
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Lattice of crystalline solids
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Structure and bonding in Metals• Closest Packing:
Assumes that metal atoms are uniform, hard spheres.
Spheres are packed in layers. Like oranges in grocery store display
abab packing - 3rd directly over 1st layer - called hexagonal closest pack (hcp)
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Structure and bonding in Metals (con’t)
abca packing - 3rd layer not directly over 1st, 4th layer is over 1st - cubic closest pack (ccp) or face centered cubic (fcc) see next slide
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Face Centered Cubic (FCC)
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Metallic Bonding Nearest Neighbors
• The Indicated Sphere Has 12 Nearest Neighbors
Each sphere in closest packed (both fcp and hcp) has 12 equivalent nearest neighbors.
What about bcc ? simple cubic ?
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Unit cell atoms
fcc and hcp8 x 1/8 spheres and 6 x 1/2 spheres = 4 total atoms in unit cellWhat about bcc? Or simple cubic?What does that say about density of metals?
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Metallic Bonding• Sea of electrons - regular array of cations
surrounded by its valence electrons
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Metallic bonding MO model
• Band Model (MO Model) - combinations of atomic orbitals.
Virtual continuum of levels, called bands. Many semiconductor applications
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Metal alloys
• Metals melted together to make a solution (homogeneous solid!!) - 2 types– Substitutional Alloy – some of the host
metal atoms are replaced by other metal atoms of similar size.
– Interstitial Alloy – some of the holes in the closest packed metal structure are occupied by small atoms.
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Metal alloys graphics
Which is a substitutional alloy?
Which is an interstitial alloy?
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Network atomic solids
2 main allotropes of carbon (3rd is buckyballs). What is hybridization on each C atom in these two structures?
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Graphite - sp2 hybridization
p-orbitals and Pi system in graphite for 1 layer (sheet). Graphite layers slide by each other because of e- repulsion. Large difference between diamond and graphite is type of bonding
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Carbon Atoms in Graphite
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Types and Properties of Solids - Table
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Vapor pressure graphic
a) Not equilibrium (pressure increasing)
b) Equilibrium (pressure constant)
Not closed --> no Pvap just Patm
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Vapor pressure rate diagram
Why does rate of condensation increase? While the rate of evaporation remain essentially constant ?
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Surface molecule interactions
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Vapor Pressure definition
• Pressure of the vapor present at equilibrium.
• The system is at equilibrium when no net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other.
• The boiling point of the liquid is when the Pvap = Patm
• Normal boiling point of liquid is at 1 atm.
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Vapor pressure trends
• Liquids in which the intermolecular forces are strong have relatively low vapor pressures.
• Vapor pressure increases significantly with temperature.
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Vapor pressure of various liquids
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Pvap rationale Temp vs. KE plot
T2 > T1, which means on average more molecules have sufficient energy to overcome liquid intermolecular forces (more evaporate --> rate faster)
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Pvap - Clausius-Clapeyron equation
Plots of In(Pvap) vs. (b) 1/T
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Clausius–Clapeyron Equation
Pvap = vapor pressure
ΔHvap = enthalpy of vaporization
R = 8.3145 J/K·mol
T = temperature (in kelvin)
1
2
vap, vap
vap, 2 1
1 1ln = ⎛ ⎞ Δ ⎛ ⎞
−⎜ ⎟ ⎜ ⎟⎜ ⎟ ⎝ ⎠⎝ ⎠
T
T
P HP R T T
For calculation: to undo ln use ex
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Vapor pressure calc. problem
The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water at 65°C.
194 torr
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Heating curve for water
Temp changing use Q = c x m x ΔT
Temp not changing use
ΔHvap liquid <--> gas ΔHfus solid <--> liquid
Why is ΔHvap > ΔHfus ?
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Phase Diagrams (P,T)
•A convenient way of representing the phases of a substance as a function of temperature and pressure:
Triple pointCritical pointPhase equilibrium lines
Phase diagram for CO2
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Phase Diagram for Water
What is different about phase diagram for water from most other substances?
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Concept check
As intermolecular forces increase, what happens to each of the following? Why? Boiling point Viscosity Surface tension Enthalpy of fusion Freezing point Vapor pressure Heat of vaporization
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