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Chemistry 11-21

Chapter 11

Intermolecular Forces, Liquids, and Solids

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Chapter Objectives Preview

  • Intermolecular Forces
    • Define the different types of intermolecular forces: dispersion forces, dipole-dipole, and hydrogen bonding.
    • Evaluate the effects of I.M.F. on properties of a substance.
  • Some Properties of Liquids
    • Define Viscosity
    • Define Surface Tension and explain the intermolecular bonding
    • Define boiling point
  • Vapor Pressure
    • Relate vapor pressure, volatility of a liquid, and the process of dynamic equilibrium.
    • Define boiling point normal boiling point, freezing point, and melting point.
    • Apply the Clausius-Clapeyron equation to solving various problems involving vapor pressure
  • Solids
    • Identify types of solids as molecular, network covalent, ionic, or metallic.
    • Discuss the crystal structures and the unit cell in metals and ionic crystals.
    • Define melting point
  • Ionic Bonding
    • Define lattice energy
    • Use Born-Haber cycle to show energy changes involved in ionic compound formation
  • Metallic Bonding
    • Use metallic bonding to explain the general properties of metals
  • Phase Diagrams
    • Use the phase diagram to determine the phase(s) present at a given temperature and pressure.
    • Define critical temperature and critical pressure

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States of Matter Revisited

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Particle Arrangement by State

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Properties of Molecular Substances

  • Compare to Ionic and Metallic from Ch.7

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  • Often soft and brittle
  • Generally low melting and boiling points
  • exist as solids, liquids and gases at room temperature
  • Insoluble in water but soluble in non-polar solvents
  • generally poor thermal and electrical

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  • What holds molecular substances together?

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Intermolecular Forces: Essential Questions

  • What are intermolecular forces of attraction and what are the different types?
  • How do these intermolecular forces arise?
  • How do I identify the types of intermolecular forces present in a substance based on the chemical formula and molecular structure?
  • How do I rank the relative importance/strength of the IMFs in a particular substance?
  • What properties of substances are influenced by the strength of IMFs and how?
  • What is the importance of understanding IMFs (i.e. real world applications)?

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Intermolecular Attractions

  • Covalent bonds = attractions within molecules
  • Intermolecular forces = Attractions between molecules
  • Stronger intermolecular force = higher boiling point
  • van der Waals forces
    • dispersion forces or London forces
      • Increase with increasing molar mass
        • More electrons
        • Easier to polarize
    • dipole attractions
    • hydrogen bonds (strongest)

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Dispersion (London) Forces

  • Occur between any molecules, but most important between non-polar molecules
  • Molecules collide causing temporary, weak dipoles
  • sporadic and weak attractions occur
  • force increases with increasing number of electrons (i.e. greater molar mass) in molecule

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Relative Importance of Dispersion Forces

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Dipole Attractions

  • Occur between polar molecules
  • Opposite charges attract
  • + ends of molecules attract - ends of other molecules
  • force increases with increasing electronegativity difference

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Hydrogen Bonds

  • occur in molecules that have N,O, or F attached to hydrogen
  • strongest of intermolecular forces (about 5-10% strength of covalent bond)
  • Hydrogen bonding in ice and why ice floats in water!

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Ion-Dipole Forces

Most Important when ionic compounds dissolved

in polar solvents like water.

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Effect of Molar Mass on Boiling Points of Molecular Substances

Explain the trends in boiling points based on intermolecular forces.

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Effect of Molar Mass on Boiling Points of Molecular Substances II

Explain the trends in melting and boiling points based on intermolecular forces.

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Boiling Points of Nonpolar vs. Polar Substances

Explain the trends in boiling points based on intermolecular forces.

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Boiling Point and Hydrogen Bonds

Explain the trends in boiling points based on intermolecular forces.

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Intermolecular Forces Summary

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IMF Graphic Summary

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The Nature of Liquids

  • Molecules in motion, but held by intermolecular forces
  • Intermolecular forces: attractive forces between molecules (not bonds)
  • Molecular Motions
    • Translation: move through space
    • Vibration: “stretch” and “shrink”
    • Rotation: spin around
  • Liquids maintain volume, but take shape of container

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Surface Tension

  • Energy required to increase the surface area of a liquid (i.e. spread the molecules apart against their attractions)
  • Increases with increasing imf
  • Decreases with increasing temperature
  • Explains why water “beads” up on a waxed surface

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Surface Tension (cont)

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Viscosity

  • Resistance of a liquid to flow
    • Compare acetone, water, and honey
  • Increases with
    • Increasing imf (stronger attractions mean molecules can’t slip by each other as easily
    • Molecule length (“entanglement”)
  • Decreases with
    • Increasing temperature

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Viscosity (cont.)

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Evaporation

  • Conversion of a liquid to a gas or vapor below its boiling point
    • vaporization at liquid surface
  • Fastest moving molecules have highest kinetic energy
  • Enough energy => break bonds with other molecules and “escape” from liquid
  • Has a cooling effect on liquid or surface
    • E.g. sweating

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KT and Evaporation

Why do you cool down when your body sweats?

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Evaporation

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Liquid –Vapor Equilibrium

  • Rate of Vaporization = Rate of Condensation
    • Liquid ⬄ Vapor
    • Dynamic equilibrium
  • The pressure exerted by the vapor over the liquid remains constant
  • If both liquid and vapor are present, vapor pressure is independent of container volume
    • WHY?!

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Vapor Pressure

  • Pressure exerted by vapor in equilibrium with a liquid.
  • Measure of tendency of liquid to evaporate at a given temperature
  • increases with increasing temperature
    • decreases with increasing imf
    • independent of container volume as long as both liquid and vapor are present

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Volatile

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  • A liquid is said to be volatile when it readily evaporates at room temperature
    • high vapor pressure due to weaker imf
    • examples: methanol, acetone, gasoline

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Vapor Pressure Animation

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Vapor Pressure vs. Temperature

Vapor pressure curves for various liquids

Determine boiling point for various external pressures.

Determine vapor pressure at various temperatures.

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The Boiling Point of a Liquid

  • Boiling Point: Temperature at which the vapor pressure of a liquid is equal to the external pressure
  • Normal Boiling Point: Temperature at which the vapor pressure of a liquid is equal to standard atmospheric pressure (101.3kPa)
  • Boiling
    • Vapor bubbles form (i.e. vaporization below surface/within body of liquid)
    • Temperature remains constant!
  • Stronger intermolecular forces = higher boiling point

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Boiling Graphic

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The Structure of Solids

  • Particles are held in fixed positions
    • Vibrations increase with temperature
  • Maintain shape and volume (rigid)
  • Melting Point: temperature at which solid turns to liquid
    • Vibrations become too “violent” for bonds to hold particles in place

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Network (Covalent) Solids

  • Atoms are covalently bonded together in a continuous two or three dimensional array
    • Macroscopic molecules!
  • Examples: diamond and graphite.
  • classified as 2-dimensional or 3-dimensional.
    • 2-dimensional arranged in layers, with weak attractions between the layers. Generally soft and/or slippery.
    • 3-dimensional is a giant interlocking design; exceptional hardness and a high melting and boiling point.
  • High Melting Points
  • Insoluble in common solvents
  • Poor electrical conductors

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Comparison of Diamond and Graphite

DIAMOND

GRAPHITE

Property

Diamond

Graphite

melting point

 3550 oC

3675 oC

boiling point

 4827 oC

 4200 oC

density

3.51 g/cm3

2.26 g/cm3

electrical conductivity

 Low

High

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Crystals vs. Amorphous

  • Crystal: atoms, ions, or molecules are arranged in orderly, repeating 3-d pattern
    • Unit Cell: smallest group of particles within crystal that maintains geometric shape of crystal
  • Amorphous Solids: lack ordered, internal structure
    • Rubber, plastic, asphalt
    • Glass
      • Transparent fusion products of inorganic substances (e.g.SiO2)
      • Amorphous solids or “super cooled” liquids
      • Cooled to rigid state without crystallizing
      • Do not have definite melting point: soften with heat

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Crystal vs. Amorphous

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Allotropes

  • Two or more different molecular forms of same element in same physical state
  • Examples
    • Carbon: soot (amorphous), graphite, diamond
    • Oxygen (O2) vs. Ozone (O3)

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Allotropes

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Summary of Crystalline Solids

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Summary of 4 Types of Substances

Ionic

Network Covalent

Metallic

Molecular

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Melting Point

  • Temperature at which solid and liquid phases are in equilibrium
  • Temperature at which solid changes to liquid as heat is continuously added
  • Increasing pressure favors formation of more dense phase
    • Large pressures required to see significant difference

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Heating and Cooling Curves

  • Plot of Temperature vs. Time
  • Shows melting point and boiling point
  • Temperature remains constant during phase change
  • Melting point is same as freezing point and boiling point is same as condensation point
    • Change of state depends on direction of heat flow

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Heating Curve for Water

BP

MP

vaporization

fusion

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Cooling Curve for Water

BP

MP

condensation

crystallization

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Heat and Changes of State

  • Molar heat of fusion
    • Heat absorbed in melting one mole of solid at its melting point
  • Molar heat of vaporization
    • Heat absorbed in vaporizing one mole of liquid at its boiling point
  • Molar Heat of Solution
    • Heat change caused by the dissolution of one mole of substance

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Standard Heats of Physical Change

Substance

MP (K)

ΔHfus (KJ/mol)

BP (K)

ΔHvap (KJ/mol)

Acetone

177.8

5.72

329.4

29.1

Ammonia

195.3

5.65

239.7

23.4

Argon

83.8

1.2

87.3

6.5

Benzene

278.7

9.87

353.3

30.8

Ethanol

158.7

4.60

351.5

43.5

Helium

3.5

0.02

4.22

0.08

Hydrogen

14.0

0.12

20.3

0.90

Methane

90.7

0.94

111.7

8.2

Methanol

175.5

3.16

3372

35.3

Neon

24.5

0.33

27.1

1.76

Nitrogen

63.3

0.72

77.4

5.58

Oxygen

54.8

0.44

90.2

6.82

water

273.2

6.01

373.2

40.7

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Heat and Changes of State: Example Problems

  • How many grams of ice at 0oC could be melted by the addition of 2.25 kJ of heat?
  • What is the heat of fusion (in cal/g) of ice?
  • What amount of heat is released when 145g of steam condenses to water at 100oC?
  • What is the total energy change when 35g of ice at -20oC is converted to steam at 125oC?

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Solid ⬄ Gas Transitions

  • Sublimation
    • Change of a substance from solid to gas or vapor state without passing through the liquid state
    • Example: CO2 or “dry” ice
    • Application in freeze drying
  • Deposition
    • Condensation of gas to solid state

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Summary of Changes of State

Phase Diagrams Quiz

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Phase Diagrams

  • Shows relationships among solid, liquid, and gas states in a sealed container
  • Shows state or states of matter of a substance for a given pressure and temperature
  • Lines between phases represent equilibrium conditions between those phases
  • Triple point: condition at which all three states coexist in equilibrium
    • All three lines on phase diagram meet

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Phase Diagram

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Phase Diagram

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Phase Diagram for Water

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Phase Diagrams

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Critical Temperature and Pressure

  • Critical Temperature
    • Temperature above which no amount of pressure can liquefy substance
    • Above this temperature there is no distinction between liquid and gas states
  • Critical Pressure
    • Pressure required to liquefy vapor at critical temperature
    • Vapor pressure at critical temperature

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Chapter Objectives Review

  • Intermolecular Forces
    • Define the different types of intermolecular forces: dispersion forces, dipole-dipole, and hydrogen bonding.
    • Evaluate the effects of I.M.F. on properties of a substance.
  • Some Properties of Liquids
    • Define Viscosity
    • Define Surface Tension and explain the intermolecular bonding
    • Define boiling point
  • Vapor Pressure
    • Relate vapor pressure, volatility of a liquid, and the process of dynamic equilibrium.
    • Define boiling point normal boiling point, freezing point, and melting point.
    • Apply the Clausius-Clapeyron equation to solving various problems involving vapor pressure
  • Solids
    • Identify types of solids as molecular, network covalent, ionic, or metallic.
    • Discuss the crystal structures and the unit cell in metals and ionic crystals.
    • Define melting point
  • Ionic Bonding
    • Define lattice energy
    • Use Born-Haber cycle to show energy changes involved in ionic compound formation
  • Metallic Bonding
    • Use metallic bonding to explain the general properties of metals
  • Phase Diagrams
    • Use the phase diagram to determine the phase(s) present at a given temperature and pressure.
    • Define critical temperature and critical pressure