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11.1: A Molecular Comparison of Liquids 11.1: A Molecular Comparison of Liquids and Solidsand Solids
31The forces holding solids and liquids together are called
intermolecular forces.
41
• The covalent bond holding a molecule together is an intramolecular forces• The attraction between molecules is an
intermolecular force• Much weaker than intramolecular forces • Melting or boiling: the intermolecular forces are
broken (not the covalent bonds)
11.2: Intermolecular 11.2: Intermolecular ForcesForces
51
The stronger the attractive forces, the higher the boiling point of the liquid and the melting point of a solid
(low boiling point)
61
Ion-Dipole Forces• Interaction between an ion and a dipole (a polar molecule
such as water)• Strongest of all intermolecular forces (MIXTURES
ONLY!)
71
Dipole-Dipole Forces
• Between neutral polar molecules• Oppositely charged ends of molecules attract• Weaker than ion-dipole forces
• Dipole-dipole forces increase with increasing polarity• Strength of attractive forces is inversely related to
molecular volume
81
London Dispersion Forces• Weakest of all intermolecular forces• Two adjacent neutral, nonpolar molecules
• The nucleus of one attracts the electrons of the other• Electron clouds are distorted• Instantaneous dipole• Strength of forces is directly related to molecular
weight• London dispersion forces exist between all molecules
Examples: 11.11 and 11.13
London dispersion forces depend on the shape of the molecule
• The greater the surface area available for contact, the greater the dispersion forces
101
Hydrogen Bonding
• Special case of dipole-dipole forces
• H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N)
Hydrogen Bonding
Boiling point increases with increasing molecular weight. The exception is water (H bonding)
Text, P. 413
Hydrogen Bonding
Solids are usually more closely packed than liquids (solids are more dense than liquids)
Ice is ordered with an open structure to optimize H-bonding (ice is less dense than water)
Text, P. 417; example 11.9
161
Viscosity• Viscosity is the resistance of a liquid to flow
• Molecules slide over each other• The stronger the intermolecular forces, the higher the
viscosity• Viscosity increases with an increase in molecular weight
11.3: Some Properties of 11.3: Some Properties of LiquidsLiquids
171
Surface Tension• Surface molecules are only attracted inwards towards the
bulk molecules • Molecules within the liquid are all equally attracted to
each other
181
• Surface tension is the amount of energy required to increase the surface area of a liquid– Cohesive forces bind molecules to each other (Hg)
– Adhesive forces bind molecules to a surface (H2O)
– If adhesive forces > cohesive forces, the meniscus is U-shaped (water in glass)
– If cohesive forces > adhesive forces, the meniscus is curved downwards (Hg in barometer)
11.4: Phase Changes11.4: Phase Changes
Text, P. 420
(Endothermic)
(Endothermic) (Exothermic)
(Exothermic)(Endothermic)
(Exothermic)
Generally heat of fusion (melting) is less than heat of vaporization (evaporation): it takes more energy to completely
separate molecules, than to partially separate them
Text, P. 420
211
Heating Curves• Plot of temperature change versus heat added is a heating
curve• During a phase change, adding heat causes no temperature
change (equilibrium)– These points are used to calculate Hfus and Hvap
221
Text, P. 421
Added heat increases the temperature of a consistent state of matterEnergy used for
changing molecular motion, no T change
231
Critical Temperature and Pressure• Gases are liquefied by increasing pressure at some
temperature• Critical temperature: the minimum temperature for
liquefaction of a gas using pressure• A high C.T. means strong intermolecular forces
• Critical pressure: pressure required for liquefaction
241
• Examples: # 31, 33, WDP # 48
• Other WDP examples: # 44, 46, 50
251
Explaining Vapor Pressure on the Molecular Level
• Some of the molecules on the surface of a liquid have enough energy to escape to the gas phase• After some time the pressure of the gas will be
constant at the vapor pressure (equilibrium)
11.5: Vapor Pressure11.5: Vapor Pressure
261
• Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface• Vapor pressure is the pressure exerted when the liquid
and vapor are in dynamic equilibrium
Volatility, Vapor Pressure, and Temperature• If equilibrium is never established then the liquid
evaporates• Volatile substances (high VP) evaporate rapidly• The higher the T, the higher the average KE, the faster
the liquid evaporates (hot water evaporates faster than cold water)
271
• Vapor pressure increases nonlinearly with increasing temperature
• Clausius-Clapeyron Equation
281
When Temperature changes from T1 to T2, Pressure changes from P1 to P2
• Use the Clausius-Clapeyron Equation to
1. Predict the vapor pressure at a specified temperature
2. Determine the T at which a liquid has a specified VP
3. Calculate enthalpy of vaporization from measurements of VP’s at different temperatures
291
• Sample problems: # 45, WDP # 35
• Other WDP examples: # 36 & 37
301
Vapor Pressure and Boiling Point• Liquids boil when the external pressure equals the vapor
pressure• Normal BP: BP of a liquid at 1atmosphere
• Temperature of boiling point increases as pressure increases
311
• Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases• Given a temperature and pressure, phase diagrams tell
us which phase will exist
11.6: Phase Diagrams11.6: Phase Diagrams
321
Text, P. 428
Vapor Pressure curve of the liquid (increase
P, increase T)
Stable at low P and high T
Stable at low T and high P
Triple Point: all 3 phases in equilibrium
Beyond this point, liquid and gas phases are indistinguishable
Melting point curve: Increased P favors solid phase; Higher T
needed to melt the solid at higher P
331
The Phase Diagrams
of H2O and CO2
Text, P. 429; Question 49 on P. 444
Line slopes to the left: ice is less dense than water (why?) MP decreases with increased P
341
• Sample Problem: BLBB #51
351
Unit Cells• Crystalline solid: well-ordered, definite arrangements of
molecules, atoms or ions• The smallest repeating unit in a crystal is a unit cell• It has all the symmetry of the entire crystal• Three-dimensional stacking of unit cells is the crystal
lattice• Close-packed structure
11.7: Structures of 11.7: Structures of SolidsSolids
361
Unit Cells
Unit Cells
Primitive cubic: atoms at the corners of a simple cube
• each atom shared by 8 unit cells
391
Unit Cells
Body-centered cubic (bcc): atoms at the corners of a cube plus one in the center of the body of the cube
• corner atoms shared by 8 unit cells• center atom completely enclosed in 1 unit cell
401
Unit Cells
Face-centered cubic (fcc): atoms at the corners of a cube plus one atom in the center of each face of the cube
• corner atoms shared by 8 unit cells• face atoms shared by 2 unit cells
Unit Cells
Text, P. 432
2 atoms per cell
4 atoms per cell
1 atom per cell
421
The Crystal Structure of Sodium Chloride
Two equivalent ways of defining unit cell:
Cl- (larger) ions at the corners of the cell, orNa+ (smaller) ions at the corners of the cell
43111.8: Bonding in Solids11.8: Bonding in Solids
Text, P. 435
441
Covalent-Network Solids
Ionic SolidsThe structure adopted
depends on the charges and sizes of the ions
461
Metallic Solids• Various arrangements are possible• The bonding is too strong for London dispersion and
there are not enough electrons for covalent bonds• The metal nuclei float in a sea of electrons• Metals conduct because the electrons are delocalized
and are mobile• Close-packed structure
471
• Amorphous solids (rubber, glass) have no orderly structure– IMFs vary in strength throughout the sample
– No specific melting point
Sample Problems # 53, 69, 71, 73, 75