LIQUIDS AND SOLIDS by liaoqinmei

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									Chapter 14 “LIQUIDS AND SOLIDS”



    What are the properties of the
    „condensed states‟ of matter?


                          T H Witherup 02/06
                          Honors (rev 07)
                 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.”
           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.
             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
Comparing Properties of Gases, Liquids
              & Solids
PROPERTY      GAS    LIQUID    SOLID
COMPRESS-
IBILITY
DENSITY
VOLUME

SHAPE
DIFFUSION

EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids

 PROPERTY      GAS        LIQUID      SOLID

COMPRESS-      High        Slight      Low
IBILITY
DENSITY

VOLUME

SHAPE

DIFFUSION

EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids

 PROPERTY      GAS        LIQUID      SOLID

COMPRESS-      High        Slight      Low
IBILITY
DENSITY         Low        High        High

VOLUME

SHAPE

DIFFUSION

EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids

 PROPERTY       GAS       LIQUID      SOLID

COMPRESS-       High       Slight      Low
IBILITY
DENSITY         Low        High        High

VOLUME          Fills     Definite    Rigid
              container
SHAPE

DIFFUSION

EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids

 PROPERTY        GAS           LIQUID       SOLID

COMPRESS-        High           Slight      Low
IBILITY
DENSITY          Low            High        High

VOLUME            Fills        Definite     Rigid
               container
SHAPE         Of container   Of container   Fixed

DIFFUSION

EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids

 PROPERTY        GAS           LIQUID         SOLID

COMPRESS-        High           Slight          Low
IBILITY
DENSITY          Low            High           High

VOLUME            Fills        Definite        Rigid
               container
SHAPE         Of container   Of container      Fixed

DIFFUSION        Rapid          Slow        Very slow (@
                                              surface)
EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids

 PROPERTY        GAS           LIQUID         SOLID

COMPRESS-        High           Slight          Low
IBILITY
DENSITY          Low            High           High

VOLUME            Fills        Definite        Rigid
               container
SHAPE         Of container   Of container      Fixed

DIFFUSION        Rapid          Slow        Very slow (@
                                              surface)
EXPANSION        High           Low             Low
(with Heat)
              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.
        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.
                   Dispersion Forces: Boiling Points of Noble
                                  Gases (°C)
                            0
                                 0           20          40         60       80        100
                           -50
  Boiling Point (deg C)




                                                                                  -67 Rn
                          -100                                     -107 Xe

                          -150                           -153 Kr
                                               -186 Ar
                          -200

                          -250             -246 Ne
                                     -269 He
                          -300
                                                         Atomic Number

         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.)
        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).
   Dipole-Dipole Forces: Polar Molecules

Consider HCl (a dipole or polar molecule):


        δ+              δ-              δ+          δ-

         H         Cl
                             :::::::::: H      Cl
                             ::::::::::

Compare CO2 and SO2 (covalent molecules with polar bonds):

                             Polar or Nonpolar? Why?
                                                             +       +
             + +                                                         B.P. = -10°C
                                B.P. = -78°C
                                                                 S
    O        C          O                                O               O
           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.
                Consider Boiling Points
 As mass increases, we expect Boiling Point to rise.
 CH4, SiH4, GeH4, SnH4
       Observed: CH4 < SiH4 < GeH4 < SnH4 (follow trend)
   NH3, PH3, AsH3, SbH3
       Observed: PH3 < AsH3 < SbH3 ~= NH3 (anomaly)
   HF, HCl, HBr, HI
       Observed: HCl < HBr < HI << HF (anomaly)
   H2O, H2S, H2Se, H2Te
       Observed: H2S < H2Se < H2Te << H2O (anomaly)
   See Table on next slide.
          Consider Boiling Points (kelvins)

Formula         XH4           XH3     XH2          XH
Period↓
Period 2         CH4           NH3     H2O          HF
                 (93)          (243)   (373)        (303)
Period 3         SiH4          PH3     H2S          HCl
                 (163)         (183)   (213)        (193)
Period 4         GeH4          AsH3    H2Se         HBr
                 (193)         (203)   (243)        (213)
Period 5         SnH4          SbH3    H2Te         HI
                 (223)         (243)   (273)        (253)

What causes these anomalies (shown in yellow)? HYDROGEN BONDING!
Data is plotted on page 466 of text.
             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 water’s unusual
Properties result from “hydrogen bonding”!




                       *******




                                                    *******
Comparison of Intramolecular & Intermolecular Forces

FORCE           ATTRACTION           ENERGY,      EXAMPLE
                                     kJ/mol
Ionic           Anion-cation         400 – 4000   NaCl
Covalent        Shared Electrons     150 - 1100   Cl-Cl
Metallic        Cations in “Sea of   75 - 1000    Cu
                Electrons”
Ion-Dipole      Ion with Dipole      40 – 600     Cl-…H2O
Dipole-Dipole   Dipole charges       5 – 25       Br-Cl…Br-Cl
H-Bond          H to N, O, F         10 - 40      H2O to H2O
Ion-induced     Ion & e- cloud of    3 – 15       Fe2+…O2
dipole          neighbor
Dipole-induced Dipole charge & e-    2 – 10       H-Br…Br2
dipole         cloud of neighbor
Dispersion      Electron clouds of   0.05 – 40    Cl-Cl…Cl-Cl
(London)        neighbors
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.
                 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.
       Syrup, motor oil, hot fudge!
   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!
   Capillarity = 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 <1g/mL (butane) to >13.6g/mL (mercury).
Video Clip (Water)
                      Comparing Small Molecules

Compound         Formula      Mass (u)    State @ 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).
            Water: A very special compound
 The most abundant substance on Earth‟s 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 of vaporization.
       Excellent solvent (“universal solvent”).
       Why? HYDROGEN BONDING!
       What are some consequences of these properties?
             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.
                       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, silicon.

See the link: http://undergrad-ed.chemistry.ohio-state.edu/chemapplets/
Crystals/ClosestPackedStructures.html
   Properties of Crystalline Solids (Fig 14-23)

Type       Particles   Forces Between    Properties
                       Particles
Metallic   Atoms       Metallic bond     Soft to hard; variable
                                         melting points; good
                                         conductivity; malleable;
                                         ductile.
Molecular Atoms or H-bond, dipole-       Soft; variable melting
          molecules dipole, dispersion   points; poor conductors.

Ionic      Ions        Electrostatic     Hard; brittle; high melting
                                         points; poor conductors.
Covalent- Atoms        Covalent bonds    Very hard; very high
network                                  melting points; poor
                                         conductivity
        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)
                   Relative Potential Energy (P.E.) Of Matter
                   During Phase Changes (from Fig. 14-30)
                                                                                             GAS




                                                                  SUBLIMATION
                                                   CONDENSATION
                                    VAPORIZATION
POTENTIAL ENERGY




                                                                                             LIQUID




                                                                                DEPOSITION
                                  FREEZING
                        MELTING




                                                                                              SOLID
        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.
            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).
          Kinetic Energy Distribution


http://undergrad-ed.chemistry.ohio-state.edu/chemapplets/
KineticMolecularTheory/Maxwell.htmlKinetic Energy Distribution




In Chapter 13 we looked at two simulations at this website:
“Gases – Kinetic Molecular Theory.” (Maxwell distribution.)
“Maxwell distribution of a gas (or liquid) at different temperatures.”
If you need to refresh your memory, revisit the website to look
at the Maxwell distribution shapes.
The general shapes of the curves apply to liquids as well.
                     VAPOR PRESSURE
   Not all molecules in a sample of a substance have
    the same amount of kinetic energy.
       Some molecules are moving faster than others, so there
        is a distribution of kinetic energy.
       The average kinetic energy of the substance is
        proportional to its temperature.
       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,‟ and all liquids exert a
        vapor pressure against the container (or atmosphere).
       Vapor pressure increases with temperature. (Why?)
         •   Consider what happens to a puddle of water.
         •   Does the water evaporate?
         •   What if the water is in a closed jar? What would you see?
         •   This is an example of “liquid-vapor equilibrium.”
          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?
                “IN CLASS” ACTIVITY

   Explore Heating & Cooling Curves
   Explore Phase Diagrams

   To do these, get one of the laptops from the cart,
    and go to the websites shown on the following
    slides.
       Read the information.
       Perform the simulations as directed.
       Take notes as appropriate.
       Record your observations.
       Answer the questions.
             Heating/Cooling Curves
“Phase Changes” Animation and Activity:

 http://www.chm.davidson.edu/chemistryapplets/Phase
 Changes/HeatingCurve.html

In class, go to this website, read the information,
and perform the Heating Curve exercise.
Answer all questions.



Also,
Doc Brown‟s website has a nice summary of the states
of matter and “heating/cooling” curves. Check it out!
http://www.wpbschoolhouse.btinternet.co.uk/page03/3_52states.htm
    HEATING/COOLING CURVE SUMMARY
 Plots time (x-axis) vs. Temperature (y-axis).
 Identical in shape, but opposite in direction.
 Ramps indicate temperature changes.
 Plateaus indicate no change in temperature, but show
  absorption/release of energy.
       Two phases co-exist where there is a plateau.
       Length of plateau is proportional to the energy change that occurs
        during a phase change.
   During a phase change, all of the energy is used to induce
    the change, so the temperature of the mixture of phases
    stays constant.
       For example, water on a stove will continue to boil at the same
        temperature until all of the water has been converted to a gas.
       Similarly, a mixture of ice and water will remain at zero C until all of
        the ice melts, after which the liquid begins to heat.
                        Phase Diagrams
 In class, go to this animation to observe what happens
 to phases as temperature and pressure change:
  http://www.chm.davidson.edu/ChemistryApplets/PhaseChanges/
  PhaseDiagram.html

  At this website (in class) you are to explore all of the following applets:
  Phase Diagram
  Phase Diagram Part 1, Part 2, Part 3, Part 4, Part 5.
  Read the information, perform the simulations and answer all questions.



To learn more about phase diagrams, check out this website:
http://www.chemistrycoach.com/Phase_diagram.htm
           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.
       http://www.activotec.com/graphics/products/sc
        fexpl.gif
                 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.”

								
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