1. Intermolecular forces, Liquids and Solids
Ch. 10
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2. Assumes shape/ volume of container –particles move past one another
Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior
GAS
LIQUID
SOLID
Assumes shape/ volume of container –particles move past one another
Assumes shape of container –particles move/ slide past one another
Retains fixed volume /shape –rigid particles locked in place
Compressible-lots of free space between particles
Not easily compressible -little free space between particles
Not easily compressible- little free space between particles
Flows easily-particles move past one another
Flows easily-particles move/ slide past one another
Does not flow easily-rigid particles cannot move/slide past one another
Diffusion within gas occurs rapidly-expand easily to fill available space
Diffusion within liquid occurs slowly-does not expand to fill container
Diffusion within solid occurs extremely slowly
Not dense
More dense than gases
More dense than liquid
No significant attractive forces between molecules
Significant attractive forces between molecules
Most significant forces between molecules
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3. Intermolecular Forces
Ideal gas law used to describe gases
Volume of gas molecule is too small
99.9% of gas is empty space
Gas molecules so far apart
No significant intermolecular attraction
If gases truly ideal (zero volume/ attractive forces), couldn’t condense them to liquids
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4. To condense real gases
Intermolecular attractive forces must overcome KE of gas molecules
Increasing attractive forces by
Decreasing distance between molecules
Increasing pressure which forces gas molecules closer together
Decreasing temperature lowers average KE
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5. Three recognized types of intermolecular attractions
Intramolecular
Intramolecular
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6. Intermolecular forces (all electrostatic)
van der Waals forces (developed equation for predicting deviation of gases from ideal behavior)
Dipole-Dipole attraction: dipole molecules orient themselves so that +/ ends close to each other
London Dispersion Forces
Relatively weak forces that exist among noble gas atoms/ nonpolar molecules. (Ar, C8H18)
Caused by instantaneous dipole, in which electron distribution becomes asymmetrical
Ease with which electron “cloud” of atom can be distorted-polarizability
Hydrogen bonds: dipole-dipole attraction in which H is bound to highly electronegative atom (F/O/N)
Ion-dipole force (solutions-also electrostatic)
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7. Ion-Dipole Forces
Exists between ion and partial charge on end of polar molecule
Important for solutions of ionic substance in polar liquids
+ ions surrounded by - ends of water molecules
- ions surrounded by + ends of water molecule
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8. Dipole-Dipole Forces (about 1% strength of covalent bonds)
Weaker than ion-dipole forces
Neutral polar molecules (dipoles) attract each other with δ+ / δ- ends
Responsible for mutual solubility of polar molecules, such as NH3 in H2O
Explains how and why polar molecules may be condensed to liquid state
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9. London Forces (dispersion forces, instantaneous dipole forces, induced dipole forces)
Weakest forces of attraction
Main form of attraction in all nonpolar molecules
No other type of van der Waals forces exist
Nonpolar He/Ar can be condensed, so some kind of attractive interactions exist
Movement of electrons within electron cloud cause instantaneous or momentary dipole moment (self-polarization)
Resultant dipole induces polarization in neighboring molecules
Mutual attraction of opposite ends of neighboring dipoles
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11. An instant later, distribution of electrons shifts
As long as molecules are close together, this electron movement can occur over huge numbers of molecules
Whole lattice of molecules can be held together in solid using van der Waals dispersion forces
An instant later, distribution of electrons shifts
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12. Hydrogen Bonding special dipole-dipole attraction
Hydrogen covalently bonded to highly electronegative elements (N, O, F)
Bond strength higher than other dipole-dipole attractions
Important in bonding of molecules (water/DNA)
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14. Strength of dispersion forces
Much weaker than covalent bonds
Size of attraction varies considerably with size/shape of molecule
In comparing intermolecular forces, molar mass is 1st consideration
When molar masses comparable, differences in intermolecular attractions mainly due to differences in molecular polarity (strengths of dipole attraction)
When molar masses differ greatly, differences in intermolecular attractions mainly due to differences in molar masses themselves (strength of London forces)
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15. Magnitude of self-polarization increases with increasing numbers of electrons
Polarizability indicates ease electron “cloud” of atom can be distorted to give dipolar charge distribution
Large atoms with many electrons (electrons loosely held) exhibit higher polarizability than small atoms (electrons tightly held near nucleus)
As atomic # increases, # electrons increases
More electrons, more distance they can move
Increased chance of temporary dipole and dipole interactions
Bigger the dispersion forces
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16. In liquids and solids, molecules10x closer
Gaseous polar molecules show little attraction for each other (far apart)
Rapidly become weaker as distance between dipoles increases
Unimportant in gas phase due to distance between molecules
In liquids and solids, molecules10x closer
Boiling point (condensation point) indicative of attractive forces between molecules
Measure of how much KE must be increased so that it can overcome attractive forces in liquid
Greater attractive forces/polarity, higher BP
Low boiling point indicates low attractive forces
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17. BP/MP increase going down group
# e’s/atomic radius increase
Why I is solid/Br liquid at room temperature
Why all halogens have much higher BPs than their neighboring noble gases
More electrons = more distance they move = bigger temporary dipoles (“stickier”) = bigger dispersion forces
Because of greater temporary dipoles, xenon molecules are "stickier" than smaller neon molecules
Neon molecules will break away from each other at much lower temperatures than xenon molecules – neon has lower boiling point
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18. Shape of molecules
More linear molecules w/greater contact between molecules have higher BP than more spherically shaped molecules
Develop bigger temporary dipoles due to electron movement than short fat ones containing same numbers of electrons
Long thin molecules lie closer together-attractions most effective if molecules really close
Responsible for mutual solubility of nonpolar molecules such as Br2 in CCl4
Large electron clouds easily polarized, so they have higher BPs than neighboring noble gases
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19. The boiling point of Argon is -189.4oC.
Why is it so low?
Argon does not interact with other substances because it is so small and has a complete octet of valence electrons. Argon must be made quite cool to allow liquefication via London dispersion forces.
How does this boiling point prove that London dispersion forces exist?
If these forces did not exist, Argon would never liquefy.
The boiling point of Xenon is oC. Why is it higher than that of Argon?
Xenon is bigger and has more electrons than Argon. The likelihood of momentary dipoles is thus greater. It has a greater polarizability than Argon.
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20. F2 can only exhibit intermolecular London forces.
Put the following substances in order from lowest to highest boiling points:
C2H6
NH3 F2
F2 can only exhibit intermolecular London forces.
C2H6 is not especially polar, but it does have a very slight electronegativity difference between the carbons and the hydrogens.
NH3 exhibits hydrogen bonding, thus giving it a relatively high boiling point.
F2, C2H6, NH3
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21. The Liquid State
Properties of Liquids
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22. Surface Tension resistance of liquid to increase in surface area (polar molecules)
Increase in attractive forces between molecules at surface compared to forces between molecules in center
Interior molecules attracted to molecules from every direction
No net force on molecule (pulled in all directions)
Surface molecules don’t have molecules to attract it on top side
Molecule on surface drawn into bulk of liquid
Same total attractive force divided between fewer adjacent molecules, resulting in stronger attractive force of surface molecule toward bulk of liquid
Decreases with temperature (reduces intermolecular attractions
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23. Intermolecular attractive forces act to minimize surface area of liquid
Geometric shape w/smallest ratio of surface area to volume is sphere
Surface molecules drawn inward and area of surface of liquid reduced
Small quantities of liquid form spherical drops
As drops get larger, weight deforms them into typical tear shape
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24. Bubbles (hollow drop)
Surface tension acts to minimize surface (radius of spherical shell of liquid), but opposed by pressure of vapor within bubble
Bubbles in pure water tend to collapse
Bubbles with surfactant are stabilized
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25. Surface film
Smaller surface area causes liquid to behave as though it has skin on surface
Surface tension enables insects to walk on surface of water
High intermolecular forces = greater surface tension
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26. Cohesion-Adhesion
Liquids confined within a container have two types of forces present on molecules
Cohesive forces-intermolecular forces among like molecules
Adhesive forces-forces between unlike molecules for one another (polar molecules, oxygen in glass attracted to hydrogen in water)
Cohesion causes water to form drops, surface tension causes them to be nearly spherical, and adhesion keeps the drops in place
Picture
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27. Adhesive forces greater than cohesive forces
Water
Mercury
Adhesive forces greater than cohesive forces
Meniscus (curved upper surface) is concave (curved downward)
Cohesive forces greater than adhesive forces
Meniscus is convex (curved upward)
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28. Capillary action Instantaneous rising of liquid in narrow tube
Combination of cohesion/adhesion
Distance traveled dependent on diameter of tube
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29. Which would have a higher surface tension, H2O or C6H14. Why
Which would have a higher surface tension, H2O or C6H14? Why? Would the shape of the H2O meniscus in a glass tube be the same or different than C6H14?
Water, having large dipole moment, has relatively large cohesive forces. Hexane is essentially nonpolar so it has low cohesive forces. Water would have higher surface tension.
Water meniscus is concave because adhesive forces of water to polar constituents on surface of glass are stronger than its cohesive forces. Hexane would have a convex meniscus because it has very small adhesive forces, and the slightly larger cohesive forces would dominate.
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30. Viscosity measure of liquid’s resistance to flow
Increases with intermolecular forces
To flow, liquid’s molecules must move past each other
Move more freely in solutions with relatively low attractive forces
Nonpolar molecules attracted to each other by only London forces have lower viscosities than polar liquids like water
Increases with molecular size
Is temperature dependent
Decrease in viscosity with increased temperature
Attributed to average molecular KE of liquid which overcomes attractive forces between molecules
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31. Homework:
Read , pp
Q pp , #13, 15, 36, 38, 40, 41, 42
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32. Solids
Structure and Types
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33. Amorphous solids noncrystalline solids
Greek “without form”
No orderly structure (arrangement)
Lack well-defined faces/shapes
Many are mixtures of particles that don’ stack together well
Plastic, glass, rubber
No distinct, sharp melting point, but soften gradually over large temperature range
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34. Crystalline Solids
Atoms, ions, or molecules ordered in well-defined 3-D arrangements (lattice)
Unit cell: smallest repeating unit of lattice
Lattice: unit cells repeated in space in all three dimensions, characteristic of crystalline solid
Lattice point: part of atom in lattice
Way spheres arranged in layers determines what type of unit cell we have
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35. Primitive or Simple Cube
Shape
Stacking
Atom in each corner of cube
Atoms in contact along cell edge
Spheres in each layer lay on spheres below/above them
Stacking pattern: AAAAAAA. . .
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36. Coordination # Atoms/unit cell
# atoms/ions surrounding atom/ion in crystal lattice
6 (in red)-has 6 immediate neighbors
Value gives measure of how tightly spheres packed together
Larger coordination #, closer spheres to each other
Not close packed (least efficient method-52%)
Very rare packing arrangement for metals (ex. Polonium)
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37. Body Centered Cubic (BCC)
Shape
Stacking
Atom at each corner/in center of cube
Atoms in contact along body diagonal
2nd layer fit into depressions of 1st layer/3rd layer into 2nd
Stacking pattern: ABABABAB. . .
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38. Coordination # Atoms/unit cell
Coordination #
Atoms/unit cell
8 (each sphere in contact with 4 spheres in layer above/4 below)
Not close packed (less efficient packing-68%)
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39. Face-centered Cube (FCC)
Shape
Stacking
Atom at each corner/atom in center of each face of cube
Each layer diagonally next to each other. Alternating layer in crevices between spheres.
Stacking pattern: ABABABAB. . .
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40. Coordination # Atoms/unit cell 12 (more efficient at 74% packing)
Not closest packed (Fe, alkali metals)
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41. Determine net number of Na+ and Cl- ions in unit cell
(¼ per edge) x (12 edges) = 3
(1 per center) x (1 center) = 1
Cl-
(1/8 per corner) x (8 corners) = 1
(½ per face) x (6 faces) = 3
Correct since # Na = # Cl
(4 = 4)
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42. Bravais lattice is infinite array of discrete points with arrangement and orientation that appears exactly the same, from whichever points array viewed
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43. Introduction to structures and types of solids
X-ray analysis of solids
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44. Diffraction Scattering of light from regular array of points or lines
Spacing between points/lines/atoms, related to wavelength of light
X-rays used because their wavelengths similar to distances between atomic nuclei
Reflection of X-rays of wavelength  from pair of atoms in two different layers of crystal
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45. If incident X-ray beam hits crystal lattice, general scattering occurs
Most scattering interferes w/itself and is eliminated (destructive interference)
(b) Incident rays in phase but reflected rays out of phase: d2 (difference in distances traveled by two rays) = odd # of half wavelengths
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46. Diffraction occurs when scattering in certain direction is in phase with scattered rays from other atomic planes
Combine to form waves that reinforce each other (constructive interference)
(a) Incident/reflected rays in phase: d1 (difference in distances traveled by two rays ) = whole # of wavelengths
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47. X-ray diffraction One of best methods to determine crystal's structure
Intense X-ray beam strikes crystal
Crystal diffracts X-ray beam differently, depending on structure/orientation
Atoms in crystal interact w/x-ray waves to produce interference
Diffraction pattern consists of reflections of different intensity used to determine crystal’s structure
However, many different orientations of crystal collected before true structure of crystal determined
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48. 3/25/2017

49. Bragg Equation resolution of X-ray diffraction detector
Used for analysis of crystal structures
Each crystalline material has characteristic atomic structure, diffracts X-rays in unique characteristic pattern
n = 2d sin 
d = distance between atoms (unique for each mineral)
n = an integer
 = wavelength of x-rays
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50. X rays of wavelength 1. 54 Å were used to analyze an aluminum crystal
X rays of wavelength 1.54 Å were used to analyze an aluminum crystal. A reflection was produced at Ø = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing this reflection.
2dsinθ = nλ
2(d)(sin 19.3) = 1(1.54 Å)
2(d)(0.3305) = 1.54 Å
d = 2.33 Å = 233 pm
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51. A topaz crystal has a lattice spacing (d) of 1
A topaz crystal has a lattice spacing (d) of 1.36 Å (1 Å = 1 x m). Calculate the wavelength of X-ray that should be used if Ø = 15.0o (assume n = 1)
2dsinθ = nλ
2 (1.36 Å)(.259) = 1(λ)
0.704 Å = 70.4 pm
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52. Types of Crystalline Solids
(a) Atomic solids (b) Ionic solids (c) Molecular solids
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53. Structure and Bonding in Metals
Structural particles
Principal attractive forces between particles
Characteristics (physical behavior)
Examples
Atoms
Nondirectional covalent bond involving positive ions and mobile valence electrons delocalized throughout crystal (metallic bond)
Pure elements
All solids at 25oC except Hg
Wide range hardness (electrons move freely from atom to atom, bond metal atoms together with widely varying degrees of force)
Wide range melting point (between ionic and covalent compounds)
Excellent thermal/electrical conductors (mobile electrons quickly carry charge throughout metal)
Malleable/ductile (nondirectional, so stress alters but not destroys crystal)
Lustrous (interaction of electrons w/light) K Na Fe Mg Ca Zn
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54. Closest packing most efficient arrangement of spheres
Uniform/hard spheres most efficiently use available space
Coordination number = 12
Each sphere in contact with
6 spheres in its own layer
3 spheres in layer above
3 spheres in layer below
Stacking
2nd layer does not lie directly over those in 1st layer
3rd layer occupies dimples of 2nd layer in two ways
Each sphere in 3rd layer lies directly over sphere in 1st layer (aba)
Each sphere can occupy positions so that no sphere in 3rd layer lies over one in 1st layer (abc)
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55. Hexagonal closest packing (hcp)
aba arrangement
Each layer identical to layer two below it
Ex.: Be, Mg, Ti, Co, Zn, Cd, He (at low T)
2 atom unit cell (8 x 1/8 + 1)
Packing efficiency = 74%
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56. Cubic close-packed (ccp)
Corresponds to face-centered cube (abc)
Each layer identical to layer three below it
Ex.: Ca, Sr, Ni, Pd, Pt, Cu, Ag, Au, Pb, Pt, Ne/Ar/Kr/Xe (at low T)
Packing efficiency = 74%
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57. Coordination Numbers for Common Crystal Structures
Structure  Coordination Number  Stacking Pattern
simple cubic     AAAAAAAA. . .
body-centered cubic    ABABABAB. . .
hexagonal closest-packed  12  ABABABAB. . .
cubic closest-packed    ABCABCABC.
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58. Calculate density of LiF
4.02 Å on edge
Same arrangement of ions as NaCl
4(6.94 amu) + 4(19.0 amu) = amu
D = amu g (1 Å) = 2.65 g/cm3
4.02 Å x 1023 amu (10-8)-3 cm
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59. Calculating density of closest packed solid
Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 144 pm. Calculate the desnity of solid silver.
Textbook-length of edge of cube: d = r8 = 1448 = 407 pm
V = d3 = (407 pm)3 = 6.74 x 107 pm3
6.74 x 107 pm3 x (1 x 10-10cm)3/1pm = 6.74 x cm3
D = m = (4 atoms)(107.9 g/mol)(1 mol/6.022 x 1023 atoms)
V x cm3
10.6 g/cm3
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60. The radius of nickel atom is 1. 24 Å (1Å = 1 x 10-8 cm)
The radius of nickel atom is 1.24 Å (1Å = 1 x 10-8 cm). Nickel crystallizes with a cubic closest packed structure. Calculate the density of solid nickel.
d = r8 = 1.24 Å8 = 3.51 Å
V = d3 = (3.51 Å)3 = 43.2 Å3
43.2 Å3 x (1 x 10-8cm)3/1pm = 4.32 x cm3
D = m = (4 atoms)(58.69 g/mol)(1 mol/6.022 x 1023 atoms)
V x cm3
9.04 g/cm3
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61. Bonding Models for Metals
Electron Sea Model: regular array of metal atoms in “sea” of electrons
Band (Molecular Orbital) Model: electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms
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62. Metal Alloys
Substance containing mixture of elements/has metallic properties
Substitutional alloy
Host metal alloy atoms replaced by other atoms of similar size
Brass-Cu/Zn, sterling silver-Ag/Cu, pewter-Sn/Cu/Bi/Sb, plumber’s solder-Sn/Sb
Interstitial alloy
Holes in closest packed metal structure occupied by small atoms
Steel
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63. Network Crystal Diamond Graphite Quartz Nonmetal atoms (pure elements)
Structural particles
Principal attractive forces between particles
Characteristics (physical behavior)
Examples
Nonmetal atoms (pure elements)
Directional covalent bonds leading to giant molecules (metallic bonds)
Very hard (brittle)
Insoluble in most ordinary liquids
Sublime or melt at high temperatures (high MP)
Poor thermal and electrical conductors (most insulators)
C (diamond, graphite)
SiC BN
SiO2 (sand, quartz)
Diamond
Graphite
Quartz
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64. Semiconductors
Material w/electrical conductivity between conductor and insulator
Conductivity is enhanced by doping
In lattice, all atoms bond to 4 neighbors, leaving no free electrons to conduct current (insulator)
Intentionally introducing impurities into extremely pure semiconductor which allows electrons in bonds to move
Make electrons available for conduction (VA)
Form holes which can conduct current (IIIA)
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65. Polar Molecular Crystals
Structural particles
Principal attractive forces between particles
Characteristics (physical behavior)
Examples
Discrete polar molecules occupy lattice point
Dipole-dipole attractions
Low to moderate MP
Soluble in some polar liquids
Poor thermal/electrical conductors
HCl
CHCl3
H2S
C6H12O6
Molecules with H bonded to O, N, F
Hydrogen bonds
Soluble in some hydrogen-bonded and some polar liquids
H2O
NH3 HF
CH3OH
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66. Nonpolar Molecular Crystals
Structural particles
Principal attractive forces between particles
Characteristics (physical behavior)
Examples
Atoms (8A), nonpolar molecules
Dispersion forces
Extremely low to moderate MP
Soluble in nonpolar solvents
Poor thermal/electrical conductors Ar
Cl2 H2
CH4 I2
CO2
CCl4
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67. Ionic Crystals
Structural particles
Principal attractive forces between particles
Characteristics (physical behavior)
Examples
Positive and negative ions that give strongest attractive forces in chemistry
Ion-ion attraction
Almost all have rigid lattice
Moderate to high MP/BP (large lattice energy required to separate ions)
Hard/brittle (when hit, atoms shift position which breaks +/- attraction)
Nonconductors as solids but conductors as liquids
Many dissolve in water
KCl
CaF2
CsBr
MgO
BaCl2
NaCl
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68. Trigonal holes < tetrahedral holes < octahedral holes
Packing arrangement done to minimize anion-anion and cation-cation repulsions
Structure of most binary ionic solids explained by closest packing of spheres
Anions usually larger than cations packed as hcp or ccp arrangements with cations filling holes
Nature of holes depends on anion : cation size
Trigonal holes formed by 3 spheres in same layer
Tetrahedral holes formed by sphere sitting in dimple of three spheres in adjacent layer
Octahedral holes formed between 2 sets of 3 spheres in adjoining layers of closest packed structures
Trigonal holes < tetrahedral holes < octahedral holes
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69. Take anion to cation ratio in each case:
Would AlP have a closest packed structure which is more like NaCl or ZnS? Ionic radii:
Al3+ = 50 pm, P3- = 212 pm
Zn2+ = 74 pm, S2- = 184 pm
Na+ = 95 pm. Cl- = 181 pm
Take anion to cation ratio in each case:
S2-/Zn2+ = 2.49 (tetrahedral holes)
Cl-/Na+ = 1.91 (octahedral holes)
P3-/Al3+ = 4.24 (?)
Aluminum ions very small compared to phosphorus ion, so not much room is needed. Tetrahedral holes are adequate. So AlP is more like ZnS than NaCl.
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70. Atomic solid with metallic properties C2H6
Using table 10.7, and based on their properties, classify each of the following substances as to the type of solid it forms. Fe
Atomic solid with metallic properties
C2H6
Contains nonpolar molecules-molecular solid
CaCl2
Contains Ca2+ and Cl- ions-ionic solid
Graphite
Made up of nonpolar carbon atoms covalently bonded in directional planes-network solid F2
Nonpolar fluorine molecules-molecular solid
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71. Homework:
Read , pp
Q pp , #46, 48, 60, 68, 72
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72. Vapor Pressure
And change of state
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73. 3/25/2017

74. Vaporization (Evaporation)
Escape of surface liquid molecules to form gas
Rate depends on
Nature of liquid
Surface area (↑ SA, ↑ evaporation)
Temperature (↑ T, ↑ proportion of molecules w/ KE above escape energy)
Always endothermic (positive)
Heat of vaporization (Enthalpy of vaporization, ΔHvap) -energy required to vaporize one mole of liquid at 1 atm
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75. Equilibrium vapor pressure
Develops in gas phase above liquid when liquid placed in closed container
When evaporation occurs in closed container, gas molecules cannot escape to surroundings
As more molecules enter gas phase, pressure increases, finally stopping at level (vapor pressure) dependent only on temperature
Pressure of vapor present at equilibrium
Gas molecules collide w/container walls/liquid
Most re-condense when collide with liquid
Initially, evaporation rate greater condensation rate
As # gas molecules increase, collisions w/liquid surface increase
Rate of evaporation eventually equals to rate of condensation-equilibrium
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76. At given temperature, not all molecules moving w/same KE
At given temperature, not all molecules moving w/same KE. Small # molecules moving very slow (low KE), while few moving very fast (high KE). Vast majority somewhere in between.
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77. 3/25/2017

78. Variations in Vapor Pressure
Related to intermolecular attractive forces
Liquids w/high intermolecular attraction have relatively low vapor pressures (less volatile-liquids that evaporate readily)
Liquids w/low intermolecular attraction have relatively high vapor pressures (more volatile)
For similar-size molecules
Hydrogen-bonded substances have largest ΔHvap values (less volatile/lower vapor pressure)
Polar substances have higher values than similar-size nonpolar substances
Vapor pressure increases w/temperature
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79. Boiling Point vapor pressure of liquid = prevailing atmospheric pressure above that liquid
Increasing temperature  increases KE  increases molecular motion
Forces of attraction between molecules (H bonding) disrupted
Molecules break free of liquid and become gas
At boiling point, liquid turns into gas
Normal boiling point-BP of liquid at 1 atm
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80. Clausius-Clapeyron equation
Mathematical relationship between heat of vaporization and vapor pressure as measures of intermolecular forces that attract molecules together in liquid state
Relationship between vapor pressure and temperature
P = vapor pressure
ΔHvap = heat of vaporization
R = universal gas law constant
T = Kelvin temperature
C = constant (eliminated in 2nd equation)
Based on following assumptions (fail at high P, near critical point)
Volume of vaporized liquid negligible compared to volume of vapor
Vapor behaves as ideal gas
ΔHvao is constant over temperature interval of data
External pressure doesn’t affect vapor pressure
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81. (b) Plots of In(Pvap) versus 1/T (Kelvin temperature) Linear
(a) Vapor pressure as function of temperature
Quantitative nature of temperature dependence of vapor pressure
Nonlinear
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82. Clausius-Clapeyron Pg. 487, Sample 10.6
The vapor pressure of water at 25oC is 23.8 torr, and the heat of vaporization of water is 43.9 kJ/mol. Calculate the vapor pressure of water at 50oC.
ln (23.8 torr) = -43,900 J/mol ( – )
PvapT2 torr J/K mol K K
ln (23.8) = -1.37
PvapT2 torr
Take the antilog of both sides to get = = 93.7 torr
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83. Water has a vapor pressure of 24 mmHg at 25oC and a heat of vaporization of 40.7 kJ/mol. What is the vapor pressure of water at 67oC?
Solution: Simply use the Clausius-Clapeyron Equation to figure out the vapor pressure. We have to be a bit careful about the units of R: the units we're using are kJ, so R = 8.31x10-3 kJ/mol K. ln(P2/P1) = -DHvap/R * (1/T2- 1/T1) ln(P2/24) = kJ/8.31x10-3 kJ/mol K *(1/340- 1/298) ln(P2/24) = 2.03 P2/24 = 7.62 P2 = 182 mmHg
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84. An unknown liquid has a vapor pressure of 88mmHg at 45oC and 39 mmHg at 25oC. What is its heat of vaporization? Solution: Again, use the Clausius-Clapeyron Equation. Here, the only thing we don't know is DHvap ln(P2/P1) = -DHvap/R * (1/T2- 1/T1) ln(88/39) = -DHvap/8.31x10-3*(1/ /298) DHvap = 32.0 kJ
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85. The vapor pressure of 1-propanol at 14. 7oC is 10. 0 torr
The vapor pressure of 1-propanol at 14.7oC is 10.0 torr. The heat of vaporization is 47.2 kJ/mol. Calculate the vapor pressure of 1-propanol at 52.8oC.
ln(10.0/x) = kJ/mol/ kJ/K mol x (1/326 K – 1/287.9 K)
100.2 torr
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86. Heating curves Phase diagrams
Changes of State
Heating curves
Phase diagrams
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87. T/KE increases, PE is constant Only 1 phase present
Water starts to boil at 100oC
KE constant due to heat of vaporization-539 cal/g
PE increases until all water evaporated
Liquid/gas in equilibrium
T/KE increases and PE is constant
Only 1 phase present
Ice starts to melt at 0oC
T/KE constant due to heat of fusion-80 calories/g (bonds not broken).
PE increases until all ice melted
Solid/liquid in equilibrium
T/KE increases and PE constant
Only 1 phase present
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88. Heat capacity (J/kg-oC-1) Heat of fusion (Enthalpy of fusion, ΔHfus)
Amount of energy required to raise temperature of system by 1o
Energy required to convert mole of solid to mole of liquid at constant pressure
Determined as length of first (solid-melting) plateau, which represents heat added, divided by # moles of sample
Heat of vaporization (Enthalpy of vaporization, ΔHvap)
Heat absorbed by one mole of liquid when it changes to gas at constant pressure
Length of second (liquid-boiling) plateau divided by # moles of sample
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89. Sublimation substance goes directly from solid to gaseous state
Reasons for sublimation
Solids have vapor pressure, but it is normally very low
Solids with little intermolecular attraction may have substantial vapor pressures and be able to sublime at room conditions.
Enthalpy of sublimation (Δhsub)
Energy in solid-gas transition
State function, so ΔHsub = ΔHfus + ΔHvap
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90. Condensation point/Crystallization point-used in place of BP/MP
Cooling Curve
Reverse-all features same except start with gas and condense to get solid as heat is removed
Condensation point/Crystallization point-used in place of BP/MP
ΔHcond requires gas to give off heat-always negative (exothermic)
ΔHvap = -ΔHcond
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91. Phase Diagram
relationship between pressure and temperature and the three states of matter
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92. Fusion curve (solid-liquid line Vapor pressure curve Liquid-gas line
Sublimation curve
Solid-gas line
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93. Diagram for closed system
Pressure plotted on y-axis/temperature on x-axis
Divided into three physical states by three lines meeting at triple point
Solids-upper left
Liquids-upper right
Gases-lower part
Along each line is equilibrium mixture of two phases on two sides of that line
Liquid-solid line usually straight line-changes in pressure have very little effect on solids and liquids (only slightly compressible)
Gas-solid and liquid-gas lines are curved upward
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94. Normal melting point-temperature at which solid and liquid states have same vapor pressure under conditions where total pressure is 1 atmosphere
Normal boiling point-temperature at which vapor pressure of liquid is exactly 1 atmosphere
Supercooling -process of cooling liquid below its freezing point without its changing to solid
Superheating -process of heating liquid above its boiling point without its boiling
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95. Triple Point Critical Point
Each phase has same temperature and vapor pressure
All three phases exist together in equilibrium
Helium-only substance that doesn’t have triple point since it has no solid phase
Maximum temperature at which any liquid can exist , defined by
Critical temperature-temp above which substance can’t liquefy gas, regardless of how great pressure
Critical pressure-pressure above which substance can no longer exist as gas, no matter how high temp
For water: 374°C and 218 atm
Depends on intermolecular attraction (greater intermolecular forces, higher critical temperature)
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96. Non-compressible/high density fluid Obtained by either
Heating gas above its critical T
Compressing liquid at higher P than its critical pressure
Above critical point, differences between gases and liquids disappear
Density of gas can approach density of liquid phase
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97. Homework:
Read , pp
Q pp , #23, 29, 76 (have fun), 78, 88, 90
Do 1 additional exercise and 1 challenge problem
Submit quizzes by to me:
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