Chapter 14

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Chapter 14 LIQUIDS AND SOLIDS

• What are the properties of the condensed states
  of matter?

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Learning Targets- Monday 5/5

• Learning Targets
• 14-2 Know the difference between intermolecular
  and intramolecular forces.
• 14-3 Identify and give examples of the following
  intermolecular forces dispersion, dipole-dipole,
  and hydrogen

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14-1 Condensed States of Matter Liquids and
Solids

• Condensed matter has much higher density
  (mass/volume) than gases.
• Unlike gases, particles of condensed matter
  experience different amounts and types of
  attractive forces.
• Kinetic-Molecular Theory can help explain the
  properties of condensed matter.
• The state of a substance at room temperature
  depends on the attractive forces between its
  particles.

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Comparing the States of Matter

• Gas
• Total disorder
• Particles free to move past each other
• Particles far apart
• Liquid
• Disorder
• Particles free to move past each other
• Particles close together
• Solid
• Ordered arrangement
• Particles vibrate, but remain in a fixed position
• Particles close together

-
5
Intermolecular Forces
Intermolecular forces attractive forces between
molecules.
Intramolecular forces hold atoms together,
attractive forces within a molecule.
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Comparing Properties of Gases, Liquids Solids
PROPERTY GAS LIQUID SOLID
COMPRESS-IBILITY
DENSITY
VOLUME
SHAPE
EXPANSION (with Heat)

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Review of Chemical Bonding

• Ionic transfer of electrons from a metal to a
  nonmetal.
• Resulting ions have opposite charge and are
  attracted to each other, forming an ionic bond.
• Metallic sharing of valence electrons of the
  metal atoms.
• Results in a network of positive ions in a sea
  of electrons.
• Covalent strong intramolecular forces from the
  sharing of valence electrons between atoms.
• Results in individual molecules with specific
  shapes.
• Intermolecular forces exist between molecules.

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Hydrogen Bonding Forces
Involves F-H, N-H or O-H bonds. How different are
the electronegativity values of these atom
pairs? What does this do to the bond
polarity? Consider water. As you will soon see,
many of waters unusual Properties result from
hydrogen bonding!





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Dipole-Dipole Forces Polar Molecules
Consider HCl (a dipole or polar molecule)
d
d-
d
d-
Cl
Cl

H
H
Compare CO2 and SO2 (covalent molecules with
polar bonds)
Polar or Nonpolar? Why?


B.P. -10C


S
B.P. -78C
O
C
O
O
O
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Dispersion Forces Boiling Points of Noble Gases
(C)
Rn
Xe
Kr
Ar
Ne
He
What causes the boiling point of helium to be so
low and that of radon to be so high? Consider
the effects shown in Figure 14-8 of text.
(Data from Figure 14-7, page 462 of your
textbook.)
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Consider Boiling Points

• As mass increases, we expect Boiling Point to
  rise.
• CH4, SiH4, GeH4, SnH4
• Observed CH4 lt SiH4 lt GeH4 lt SnH4 (follow
  trend)
• NH3, PH3, AsH3, SbH3
• Observed PH3 lt AsH3 lt SbH3 NH3 (anomaly)
• HF, HCl, HBr, HI
• Observed HCl lt HBr lt HI ltlt HF (anomaly)
• H2O, H2S, H2Se, H2Te
• Observed H2S lt H2Se lt H2Te ltlt H2O (anomaly)
• See Table on next slide.

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Consider Boiling Points (kelvins)
Formula? Period? XH4 XH3 H2X HX
Period 2 CH4 (93) NH3 (243) H2O (373) HF (303)
Period 3 SiH4 (163) PH3 (183) H2S (213) HCl (193)
Period 4 GeH4 (193) AsH3 (203) H2Se (243) HBr (213)
Period 5 SnH4 (223) SbH3 (243) H2Te (273) HI (253)
What causes these anomalies (shown in yellow)?
HYDROGEN BONDING!
Data is plotted on page 466 of text.
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Intermolecular Forces have

• a wide range of strengths.
• They are much weaker than ionic, covalent and
  metallic bonds.
• Basic types of Intermolecular Forces
• Dispersion Forces attraction between temporary
  induced dipoles.
• Consider noble gas boiling points and Fig. 14-8.

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Intermolecular Forces have

• a wide range of strengths.
• They are much weaker than ionic, covalent and
  metallic bonds.
• Basic types of Intermolecular Forces
• Dispersion Forces attraction between temporary
  induced dipoles.
• Consider noble gas boiling points and Fig. 14-8.
• Dipole-Dipole attraction between polar
  molecules (dipoles) having permanent charge
  separation.
• Consider HCl, HBr, CO2 (linear) and SO2 (bent).

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Intermolecular Forces have

• a wide range of strengths.
• They are much weaker than ionic, covalent and
  metallic bonds.
• Basic types of Intermolecular Forces
• Dispersion Forces attraction between temporary
  induced dipoles.
• Consider noble gas boiling points and Fig. 14-8.
• Dipole-Dipole attraction between polar
  molecules (dipoles) having permanent charge
  separation.
• Consider HCl, HBr, CO2 (linear) and SO2 (bent).
• Hydrogen Bonds strong attraction between H atom
  of a molecule and a very electronegative atom
  (F,O,N) of another molecule.
• See data on the next slides.

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Comparison of Intramolecular Intermolecular
Forces
FORCE ATTRACTION ENERGY, kJ/mol EXAMPLE
Ionic Anion-cation 400 4000 NaCl
Covalent Shared Electrons 150 - 1100 Cl-Cl
Metallic Cations in Sea of Electrons 75 - 1000 Cu
Ion-Dipole Ion with Dipole 40 600 Cl-H2O
Dipole-Dipole Dipole charges 5 25 Br-ClBr-Cl
H-Bond H to N, O, F 10 - 40 H2O to H2O
Ion-induced dipole Ion e- cloud of neighbor 3 15 Fe2O2
Dipole-induced dipole Dipole charge e- cloud of neighbor 2 10 H-BrBr2
Dispersion (London) Electron clouds of neighbors 0.05 40 Cl-ClCl-Cl
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Summary Intermolecular Forces have

• a wide range of strengths.
• They are much weaker than ionic, covalent and
  metallic bonds.
• Basic types of Intermolecular Forces
• Dispersion Forces attraction between temporary
  induced dipoles.
• Dipole-Dipole attraction between polar
  molecules (dipoles) having permanent charge
  separation.
• Hydrogen Bonds strong attraction between H atom
  of a molecule and a very electronegative atom
  (F,O,N) of another molecule.

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Learning Targets- Tuesday 5/6

• Learning Targets
• 14-4 Know the properties of water and other
  liquids with respect to boiling/melting point,
  surface tension, capillary action and viscosity.
• Learning Outcome
• Be able to identify the properties based on the
  IMF that are involved in the liquid.

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14-2 Properties of Liquids

• Intermolecular forces generally determine the
  physical properties of liquids, such as...
• Viscosity resistance to motion between
  molecules of a liquid as they move past each
  other.
• Water, acetone, vegetable oil!
• Surface Tension unbalanced attractive forces at
  the surface of a liquid that causes the surface
  to act like a film.
• Water bugs, and the paperclip experiment!
• Capillary Action the tendency of a liquid to
  flow through a small opening or tube.
• Wicking of fabric when it gets wet.
• Density mass/volume.
• Wide range, from lt1g/mL (butane) to gt13.6g/mL
  (mercury).

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Comparing Small Molecules
Compound Formula Mass (u) State _at_ 25C MP BP
Methane CH4 16 Gas 90 112
Ammonia NH3 17 Gas 195 240
Water H2O 18 Liquid 273 373
Nitrogen N2 28 Gas 63 77
Oxygen O2 32 Gas 55 90
MP Melting Point BP Boiling Point (kelvins).
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Water A very special compound

• The most abundant substance on Earths surface.
  (The Blue Planet.)
• Critical to life forms as we know them.
• Some unusual properties of water are
• Unexpectedly high boiling point for its size.
• Unusually high specific heat.
• Solid form (ice) has lower density than liquid.
• High surface tension.
• High heat of vaporization.
• Excellent solvent (universal solvent).
• Why? HYDROGEN BONDING!
• What are some consequences of these properties?

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Learning Targets- Wednesday 5/7/2014

• Learning Targets
• 14-5 Know the difference between amorphous and
  crystalline solids.
• Learning Outcome
• Be able to identify the structures of metallic
  solids, molecular solids, ionic solids and
  covalent-network solids.

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14-3 The Nature of Solids

• Crystalline Solids
• Highly ordered, repeating arrangement of
  particles.
• Ionic (NaCl), covalent (sugar) or metallic (Fe).
• Characterized by specific unit cells. (Fig.
  14-20)
• Generally have sharp melting points.
• Fracture occurs along definite planes when
  stressed.
• Amorphous Solids (supercooled liquids)
• Highly disordered, random arrangement of
  particles.
• Wax, plastics (PET, PE, PS, Nylon).
• Characterized by lack of organized structure.
• Generally soften over a wide temperature range.
• Fracture occurs randomly with stress.

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Bonding in Solids

• Metallic Solids
• All metallic elements.
• Molecular Solids
• Most organic compounds (contain carbon) many
  inorganic compounds (CO2, H2O, SO2).
• Ionic Solids
• Typical salts (NaCl, KBr, CaCl2)
• Covalent-Network Solids
• Diamond, graphite

See the link http//undergrad-ed.chemistry.ohio-
state.edu/chemapplets/ Crystals/ClosestPackedStruc
tures.html
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Metallic Solids

• Shiny
• Conductors
• Ductile Ability to Stretch
• Malleable Ability to be shaped
• Alloys Mixing of 2 metals (stainless steel and
  brass)

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Ionic Solids
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Covalent-Network Solids
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Properties of Crystalline Solids (Fig 14-23)
Type Particles Forces Between Particles Properties
Metallic Atoms Metallic bond Soft to hard variable melting points good conductivity malleable ductile.
Molecular Atoms or molecules H-bond, dipole-dipole, dispersion Soft variable melting points poor conductors.
Ionic Ions Electrostatic Hard brittle high melting points poor conductors.
Covalent-network Atoms Covalent bonds Very hard very high melting points poor conductivity
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Learning Targets- Thursday 5/8/2014

• Learning Targets
• 14-6 Understand phase changes including heat of
  vaporization and heat of fusion
• Learning Outcome
• Be able to analyze a phase change graph to
  determine phases of matter including heat of
  vaporization and heat of fusion.

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VAPOR PRESSURE

• Not all molecules in a sample of a substance have
  the same amount of kinetic energy.
• Molecules that move very fast may achieve escape
  velocity and leave the liquid entirely. (They
  evaporate.)
• This tendency of a liquid to become a gas at a
  given temperature is its vapor pressure.
• Vapor pressure increases with temperature.
  (Why?)

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14-4 Changes of State (Phase Changes)

• Phase Change conversion of a substance from one
  physical state of matter to another.
• Six types of phase changes to consider
• Melting (solid to liquid)
• Freezing (liquid to solid)
• Vaporization (evaporation, boiling) (liquid to
  gas)
• Condensation (gas to liquid)
• Sublimation (solid to gas)
• Deposition (gas to solid)

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Energy Changes Phase Changes

• To change a substance from one state of matter to
  another always involves a gain or loss of energy.
• Melting a solid requires energy input, but
  freezing a liquid removes energy from the liquid.
• Boiling a liquid requires energy input, but
  condensing a gas removes energy from the gas.
• Subliming a solid requires energy input, but
  deposition of a gas removes energy from the gas.

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HEAT OF VAPORIZATION is

• the amount of heat needed to vaporize a given
  amount of liquid at its boiling point.
• Molar heat of vaporization of water 40.7 kJ/mol
  or 540. cal/mol.
• Energy is added to the liquid, but the
  temperature does not change.
• The added energy is expended to overcome the
  intermolecular attractions of molecules of the
  liquid, converting the substance from liquid to
  gas.
• Liquids with strong intermolecular attractions
  (such as water) have high heats of vaporization.
• Condensation is the opposite of evaporations.
• Heat of Condensation -(Heat of Vaporization)
• Why is a burn from steam more damaging to tissue
  than that from boiling water?

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Vaporization Condensation

• According to K-M Theory, temperature is a measure
  of the average kinetic energy of the particles of
  a substance.
• Fast molecules near the surface of a liquid may
  escape the liquid (evaporation).
• This results in some vapor in the space above a
  liquid.
• Volatile liquids have a tendency to easily
  evaporate due to the poor attraction among their
  molecules, resulting in high concentrations of
  vapor around the liquid.
• Gasoline, for example, poses high hazards because
  of this property.
• What happens as fast molecules leave the liquid?
• The remaining molecules have lower average
  kinetic energy, which is observed by the lower
  temperature of the liquid.
• Evaporative cooling, sweating, wind chill are
  examples.
• Liquid-Vapor Equilibrium occurs in closed
  vessels.
• Rates of evaporation and condensation are the
  same.
• Dynamic equilibrium (rates of opposing processes
  are equal).

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Learning Targets- Thursday 5/12/2014

• Learning Targets
• 14-7 Understand how to use phase diagrams to
  identify what phase(s) is/are present at given
  pressure and temperature
• Learning Outcome
• Be able to analyze a phase diagram to determine
  the phase of a substance based on pressure and
  temperature.

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PHASE DIAGRAM SUMMARY

• A plot of pressure and temperature is determined
  experimentally to show the changes in phase that
  occur under various conditions.
• Phases in equilibrium may be identified from the
  Phase Diagram.
• Triple Point and Critical Point may be
  identified.
• Triple Point The temperature and pressure at
  which all three phases exist in equilibrium.
• Critical Point The temperature and pressure
  beyond which the substance can only exist as a
  gas.

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• Phase Diagram for CO2

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• Phase Diagram for H2O

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Ch. 14 OBJECTIVES

• Show how the Kinetic-Molecular (K-M) Theory
  accounts for the physical properties of liquids
  solids.
• Describe different types of intermolecular
  forces, and how they affect properties of liquids
  solids.
• Learn about viscosity surface tension, and
  explain their relationship to intermolecular
  forces.
• Compare crystalline amorphous substances.
• Relate the structure bonding in the four types
  of crystalline solids to the properties they
  exhibit.
• Describe the changes of state (vaporization,
  condensation, boiling, sublimation, deposition,
  melting, freezing).
• Learn how to interpret Heating/Cooling Curves
  and Phase Diagrams.