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11.1: A Molecular Comparison of Liquids and Solids

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11.1: A Molecular Comparison of Liquids and Solids. The forces holding solids and liquids together are called intermolecular forces. 11.2: Intermolecular Forces. The covalent bond holding a molecule together is an intramolecular forces

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11.2: Intermolecular Forces
  • 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)
The stronger the attractive forces, the higher the boiling point of the liquid and the melting point of a solid

(low boiling point)

Ion-Dipole Forces
  • Interaction between an ion and a dipole (a polar molecule such as water)
  • Strongest of all intermolecular forces (MIXTURES ONLY!)
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
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
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

Text, P. 413

Boiling point increases with increasing molecular weight. The exception is water (H 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)
11.3: Some Properties of Liquids
  • 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
Surface Tension
  • Surface molecules are only attracted inwards towards the bulk molecules
  • Molecules within the liquid are all equally attracted to each other
Surface tension is the amount of energy required to increase the surface area of a liquid
    • Cohesive forcesbind molecules to each other (Hg)
    • Adhesive forcesbind 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 Changes







Text, P. 420

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

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
Added heat increases the temperature of a consistent state of matter

Energy used for changing molecular motion, no T change

Text, P. 421

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
Examples: # 31, 33, WDP # 48
  • Other WDP examples: # 44, 46, 50
11.5: Vapor Pressure
  • 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)
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)
Vapor pressure increases nonlinearly with increasing temperature
  • Clausius-Clapeyron Equation
When Temperature changes from T1 to T2, Pressure changes from P1 to P2
  • Use the Clausius-Clapeyron Equation to
  • Predict the vapor pressure at a specified temperature
  • Determine the T at which a liquid has a specified VP
  • Calculate enthalpy of vaporization from measurements of VP’s at different temperatures
Sample problems: # 45, WDP # 35
  • Other WDP examples: # 36 & 37
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
11.6: Phase Diagrams
  • 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
Melting point curve: Increased P favors solid phase; Higher T needed to melt thesolid at higher P

Beyond this point, liquid and gas phases are indistinguishable

Vapor Pressure curve of the liquid (increase P, increase T)

Text, P. 428

Stable at low T and high P

Stable at low P and high T

Triple Point: all 3 phases in equilibrium

Line slopes to the left: ice is less dense than water (why?) MP decreases with increased P

Text, P. 429; Question 49 on P. 444

The Phase Diagrams

of H2O and CO2

11.7: Structures of Solids
  • 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
Unit Cells
  • Primitive cubic: atoms at the corners of a simple cube
    • each atom shared by 8 unit cells
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
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
1 atom per cell

2 atoms per cell

Unit Cells

4 atoms per cell

Text, P. 432

The Crystal Structure of Sodium Chloride
  • Two equivalent ways of defining unit cell:
    • Cl- (larger) ions at the corners of the cell, or
    • Na+ (smaller) ions at the corners of the cell
Text, P. 435

11.8: Bonding in Solids

Ionic Solids

The structure adopted depends on the charges and sizes of the ions

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
Amorphous solids (rubber, glass) have no orderly structure
    • IMFs vary in strength throughout the sample
    • No specific melting point