Chapter 11: States of matter; liquids and solids - PowerPoint PPT Presentation

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Chapter 11: States of matter; liquids and solids

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Melting (fusion): transition from a solid to a liquid ... will rise in the capillary to minimize surface tension at the inverted meniscus ... – PowerPoint PPT presentation

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Title: Chapter 11: States of matter; liquids and solids


1
Chapter 11 States of matter liquids and solids
Chemistry 1062 Principles of Chemistry II Andy
Aspaas, Instructor
2
Gases, liquids, and solids
  • Gas compressible fluid
  • Mostly empty space
  • Particles in constant random motion
  • Vapor gas form of substance normally solid or
    liquid
  • Liquid incompressible fluid
  • Tightly packed particles
  • Still in constant random motion
  • Solid incompressible and rigid
  • Tightly packed particles
  • Particles only vibrate about their fixed sites

3
Phase transitions
  • Melting (fusion) transition from a solid to a
    liquid
  • Freezing transition from a liquid to a solid
  • Vaporization change of a solid or liquid to a
    vapor
  • Sublimation change of a solid directly to a
    vapor
  • Condensation change of a gas to a liquid or a
    solid
  • Deposition change of a gas directly to a solid

4
Vapor pressure
  • Vapor pressure of a liquid partial pressure of
    vapor over the liquid, measured at equilibrium
  • Vaporization and condensation are happening
    continually
  • When the rate of condensation equals the rate of
    vaporization, equilibrium has been reached
  • Vapor pressure depends on temperature
  • More kinetic energy in the liquid makes
    vaporization occur more quickly
  • Volatile liquids and solids have high vapor
    pressure

5
Boiling point and melting point
  • Boiling point temperature at which the vapor
    pressure of a liquid equals the atmospheric
    pressure
  • Bubbles form in the liquid when its heated
    enough that its vapor pressure equals external
    pressure
  • Once boiling begins, the temperature of the
    liquid stops rising
  • Freezing point temperature at which a pure
    liquid changes to a solid
  • Identical to melting point temperature at which
    a solid changes to a liquid

6
Heat of phase transition
  • Heat of fusion amount heat required to melt a
    solid
  • H2O(s) ? H2O(l) ?Hfus 6.01 kJ/mol
  • Heat of vaporization amount of heat required to
    vaporize a liquid
  • H2O(l) ? H2O(g) ?Hvap 40.7 kJ/mol
  • Heating curveno temperaturechange during a
    phase transition

7
Phase transitions
  • Heat must be added for melting or vaporization
  • Endothermic processes
  • Positive value of ?H
  • Heat must be removed for condensation or freezing
  • Exothermic processes
  • Negative value of ?H

8
Clausius-Clapeyron equation
  • Vapor pressure of a liquid depends on temperature
  • Also depends on ?Hvap for that liquid
  • Clausius-Clapeyron equation
  • The two-point form removes the constant B from
    the equation

9
Phase diagrams
  • Phase diagram for water
  • Curves show exptl points at which phases are in
    equilibrium
  • Solid-liquid line (melting-point curve)
  • Leans left if liquid is more dense than solid (as
    in water)
  • Otherwise leans right

10
Phase diagrams
  • Curve AC gives the vapor pressure at any
    temperature
  • Triple point, A point at which all 3 phases
    exist in equilibrium
  • Sublimation occurs when heating at pressures
    below triple point
  • Critical point, C point above which gas and
    liquid are no longer distinct
  • Supercritical fluid at temperature and pressure
    above critical point

11
Liquid state surface tension
  • Surface tension energy required to increase the
    surface area of a liquid by a unit amount
  • Liquids will minimize their amount of surface
    area (since intermolecular forces pull molecules
    at the surface inward)
  • Dissolved substances reduce surface tension by
    interrupting intermolecular forces
  • Capillary rise when a liquid is attracted to a
    solid capillary (like water to glass) it will
    rise in the capillary to minimize surface tension
    at the inverted meniscus

12
Viscosity
  • Viscosity resistance to flow in a liquid or gas
  • Strong intermolecular forces increase viscosity
    and cause a liquid to flow more slowly (like
    syrup)
  • Motor oil viscosity in SAE units
  • 10W/30 has SAE 10 in winter, 30 in summer
    (because viscosity of engine oil increases as
    temperature decreases)

13
Intermolecular forces
  • Dipole-dipole force attractive interaction
    between polar molecules
  • Polar molecules have a dipole moment
  • ? and ? sides of the molecule resulting from
    electronegativity differences
  • Negative side of one molecule will be attracted
    to the positive side of another molecule
  • London dispersion forces attractive forces
    resulting from instantaneous induced dipoles
    caused by random electron groupings in an atom
  • High molecular weight higher London forces
  • Compact molecule lower London forces

14
Predicting boiling points, vapor pressure, and
viscosity
  • Polar molecules have stronger intermolecular
    forces than nonpolar molecules of similar MW
    (dipole-dipole)
  • Larger molecules have stronger intermolecular
    forces than smaller molecules of similar polarity
    (London)
  • Strong intermolecular forces
  • Low vapor pressure (less chance for molecule to
    break free)
  • High boiling point (high temperature required to
    raise vapor pressure to atmospheric pressure)
  • High viscosity (molecules cling strongly
    together)

15
Hydrogen bonds
  • Hydrogen bonds much stronger than dipole-dipole
    or London forces
  • Interaction between H and Y in XHY
  • Where X and Y are F, O, or N (electronegative
    atoms)
  • H has a strong ? when covalently bonded to an
    electronegative atom
  • Explains why CH3OH boils at 65 C and CH3F boils
    at 78 C (both have similar MW and polarities)

16
Classification of solids
  • Molecular solid solid that consists of atoms or
    molecules held together by intermolecular forces
  • H2O(s), CO2(s)
  • Metallic solid any solid metal contains a
    regular arrangement of positive metal nuclei
    surrounded by a sea of delocalized electrons
  • Ionic solid consists of cations and anions held
    together by electrostatic attractions of opposite
    charges (ionic bonds)
  • Covalent network solid consists of atoms held
    together in large networks or chains by covalent
    bonds

17
Melting point of solids
  • Molecular solid low melting point (below 300 C)
  • Only weak intermolecular forces to overcome
  • Ionic and covalent network solids high melting
    point (800 C or above)
  • Ionic and covalent bonds require much more energy
    to be broken
  • Larger charges make even higher mps in ionic
    solids
  • Metals variable (Hg 39 C W 3410 C)

18
Hardness and conductivity of solids
  • Molecular solids generally soft and brittle
  • Ionic solids generally hard and brittle
  • 3-dimensional covalent network solids are
    tremendously hard
  • Metals are malleable since the atoms can move
    past each other
  • Metals (and molten ionic solids) conduct
    electricity because electrons can move freely

19
Crystalline solids
  • Crystalline solids have an ordered 3-dimensional
    structure of particles
  • Amorphous solids have their particles locked into
    random positions
  • Crystal lattice 3-dimensional geometric
    arrangement of particles in a crystal
  • Unit cell smallest boxlike unit from which a
    crystal lattice can be constructed
  • Unit cells are stacked in 3 dimensions

20
Cubic unit cells
  • Simple cubic unit cell atoms only at the corners
    of a cubic unit cell
  • Only 1/8 of a corner atom is contained in a unit
    cell
  • Body-centered cubic atom at the center of a unit
    cell as well as the corners
  • Face-centered cubic atoms in the center of each
    of the faces as well as the corners
  • Only 1/2 of a face atom is contained in a unit
    cell

21
Cubic unit cells
22
Molecular solids closest packing
  • In simple molecular solids (such as individual
    noble gas atoms), nondirectional London forces
    govern the packing of the atoms together to form
    a solid
  • Spheres pack together most closely in a hexagonal
    honey-comb type arrangement on one layer
  • A second layer can be nested in half of the
    crevices left by the first layer
  • A third layer can be nested in the crevices in
    the second layer, but there are two possible
    positions for this layer

23
Close packing arrangements
  • Hexagonal close packing (hcp) when 3rd layer is
    identical to 1st (in x holes left by 2nd layer)
  • Cubic close packing (ccp) when 3rd layer is
    different (in y holes left by 2nd layer)

24
Cubic close-packing
  • Cubic close-packing has a face-centered cubic
    lattice structure (tilted on its side)

25
Coordination number
  • Coordination number number of nearest-neighbor
    adjacent atoms any one atom can have
  • Close packing structures CN 12
  • hcp, ccp/face-centered cubic
  • Body-centered cubic CN 8
  • Square cubic CN 4

26
Ionic solids
  • 3 common structures CsCl structure, NaCl
    structure, and zinc blende (cubic ZnS) structure
  • CsCl structure
  • 1 of eachper unit cell
  • When ions are similarin size
  • Expandedsimple cubic

27
NaCl structure
  • Expanded face-centered cubic structure
  • When cation and anion are moderately different in
    size

28
Zinc blende structure
  • When cation and anion are significantly different
    in size
  • Anion is expanded body-centered cubic
  • Cation is in 4 opposite tetrahedral quadrants
    completely inside the unit cell

29
Unit cell calculations
  • Length of side of a unit cell
  • Volume of a unit cell
  • Mass of a unit cell
  • Mass of individual atoms
  • Molar mass
  • Forward or backwards!
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