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Liquids and Solids

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Covalent network solids such as quartz where atoms are held together by 3-D networks of covalent bonds. Here the hexagonal pattern of Si (violet) and O (red) atoms in ... – PowerPoint PPT presentation

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Title: Liquids and Solids


1
Liquids and Solids
2
Intermolecular Forces
  • Forces of attraction between neighboring
    particles
  • Much weaker than bonding forces
  • Responsible for state of matter and some physical
    properties
  • e.g., The stronger the attractive forces, the
    higher the melting and boiling points
  • Also involved in change of state

3
Three Types
  • London Dispersion forces
  • Dipole-dipole forces
  • Hydrogen bonds

4
London Dispersion Forces
  • The motion of electrons can create an
    instantaneous dipole moment on an atom
  • For example, if at any one time both of a helium
    atoms electrons are on the same side of the atom
    at the same time
  • A temporary dipole on one atom can cause, or
    induce, a temporary dipole on an adjacent atom

5
London Dispersion Forces
  • These forces are significant only when molecules
    are very close together, as in a compressed gas
  • These forces are found only in nonpolar compounds
  • Molecules and atoms will lose their spherical
    shape

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  • More compact molecules have smaller surface
    areas, weaker London dispersion forces, and lower
    boiling points.
  • Flatter, less compact molecules have larger
    surface areas, stronger London dispersion forces,
    and higher boiling points.

8
Dipole-Dipole Forces
  • Polar molecules have a positive end and a
    negative end
  • Dipole-dipole forces occur when the positive end
    of one molecule is attracted to the negative end
    of another
  • Only effective when polar molecules are very
    close together
  • For molecules of about the same size, dipole
    forces increase with increasing polarity

9
If two neutral molecules, each having a permanent
dipole moment, come together such that their
oppositely charged ends align, they will be
attracted to each other.
10
Hydrogen Bonds
  • Type of dipole-dipole force
  • Not a true bond!
  • Occurs between molecules containing a hydrogen
    atom bonded to a small, highly electronegative
    atom with at least one lone pair of electrons
    (e.g., N, O F)
  • The hydrogen in one molecule will be attracted to
    the electronegative atom in another molecule

11
Hydrogen Bonds
  • Hydrogen has no inner core of electrons, so a
    dipole will expose its concentrated charge on the
    proton, its nucleus.
  • Hydrogen can approach an electronegative atom
    very closely and interact strongly with it.

12
  • Electron shell around a hydrogen atom is rather
    thin, giving the hydrogen atom a small positive
    charge.
  • Electron shell round an oxygen atom is quite
    thick, and so oxygen carries an extra bit of
    negative charge.
  • These opposite charges attract, although quite
    weakly.
  • This weak force is called a hydrogen bond. The
    hydrogen atoms of one water molecule stick to the
    oxygen atoms of nearby water molecules.

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Properties of Liquids
  • Have much greater densities than their vapors
  • Only slightly compressible not a discernable
    difference when compressed
  • Fluidity ability to flow
  • Liquids can diffuse through one another, but at a
    much slower rate than gases

15
Properties of Liquids
  • Viscosity resistance to flow
  • Determined by the type of intermolecular forces
    involved, the shape of the particle, and the
    temperature
  • The stronger the attractive forces, the higher
    the viscosity
  • The larger the particles, the higher the
    viscosity
  • Increases as temp decreases

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Properties of Liquids
  • Surface Tension the imbalance of forces at the
    surface of a liquid
  • The uneven forces make the surface behave as if
    it has a tight film stretched across it
  • The stronger the intermolecular forces, the
    higher the surface tension

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It could only happen in space A tiny bubble of
air hangs suspended inside a droplet of water.
Credit ISS Expedition 6 Flight Engineer Nikolai
Budarin
Space Picture International Space Station
Astronaut Leroy Chiao
20
Properties of Liquids
  • Surfactants compounds that lower the surface
    tension of water
  • Frequently added to detergents
  • Capillary action movement of a liquid through
    narrow spaces

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Properties of Solids
  • Have extremely strong intermolecular forces in
    order for solids to have definite shape and
    volume
  • Particle arrangement causes solids to almost
    always have higher densities than liquids
  • Ice is an exception it expands when it freezes
    because of the way the particles arrange
    themselves during the freezing process

23
Properties of Solids
  • Particle arrangements cause different types of
    solids
  • Crystalline solids
  • Molecular solids
  • Covalent network solids
  • Ionic solids
  • Metallic solids
  • Amorphous solids

24
Crystalline Solids
  • Has atoms, ions, or molecules arranged in an
    orderly, geometric, 3-D structure
  • Individual pieces of a crystalline solid are
    called crystals
  • Smallest arrangement of connected points that can
    be repeated in 3 directions to form a lattice is
    called a unit cell
  • There are 7 different crystal systems based on
    shape

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Molecular Solids
  • Held together by dispersion forces, dipole-dipole
    forces or hydrogen bonds
  • NOT held together by genuine bonds (ionic and
    covalent)
  • Most are NOT solids at room temperature
  • Poor conductors of heat and electricity (dont
    contain ions)
  • Examples are sucrose and ice

30
Molecular such as sucrose or ice whose
constituent particles are molecules held together
by the intermolecular forces.
31
Arrangement of molecules in liquid water
Arrangement of molecules in ice
32
Covalent Network Solids
  • Atoms that can form multiple covalent bonds
  • Form a network of atoms that do not have a unit
    cell
  • Most allotropes exist in this form
  • Allotropes are forms of the same element that
    have different bonding patterns of arrangement
  • Examples include diamonds and graphite, quartz

33
Graphite
Diamond
34
Covalent network solids such as quartz where
atoms are held together by 3-D networks of
covalent bonds. Here the hexagonal pattern of Si
(violet) and O (red) atoms in structure matches
the hexagonal crystal shape
35
Ionic Solids
  • Type of crystalline solid
  • Type and ratio of ions determine the structure of
    the lattice and the shape of the structure
  • The network of attractions that extend through an
    ionic compound gives these compounds their high
    melting points and hardness

36
Ionic Solids
  • Strong but brittle
  • When struck, cations and anions are shifted,
    which causes repulsion that in turn shatter the
    crystal
  • Poor conductors of heat and electricity in solid
    form

37
  • Ionic solids are an orderly pattern of one ion,
    generally the anion, with cations positioned in
    'holes' between the anions
  • The occupation of these 'holes' depends on the
    formula of the ionic compound

38
Sodium chloride
Cupric chloride
39
Metallic Solids
  • Consist of positive metal ions surrounded by a
    sea of mobile electrons
  • Mobile electrons make metals malleable, ductile,
    and good conductors of heat and electricity

40
  • A series of metals atoms that have all donated
    their valence electrons to an electron cloud that
    permeates the structure
  • This electron cloud is referred to as an electron
    sea
  • Visualize the electron sea
  • model as if it were a box of marbles that are
    surrounded by water. The marbles are the metal
    atoms and the water represents the electron sea.

41
  • The marbles can be pushed anywhere within the box
    and the water will follow them, always
    surrounding the marbles.
  • This unique property, allows metallic bonds to be
    maintained when pushed and pulled in all sorts of
    ways.
  • As a result, they are malleable and ductile.

42
Gold
Copper
Silver
43
Amorphous Solids
  • Solid in which the particles are not arranged in
    a regular, repeating pattern, but still retain
    rigidity
  • Examples include glass, rubber, many plastics,
    tar and wax
  • Particles are trapped in a disordered arrangement
    that is characteristic of liquids

44
Phase Changes
  • Always involve a change in energy
  • Energy is needed either to overcome or form
    attractive forces between particles

45
Melting and Freezing
  • Melting point/freezing point temp at which
    solid and liquid forms exist in equilibrium
  • Melting is endothermic
  • Freezing is exothermic

46
Vaporization
  • The change of state from a liquid to a gas
  • Endothermic process
  • Two methods of vaporization
  • Evaporation
  • Boiling

47
Evaporation
  • Occurs at the surface of a liquid
  • Occurs because molecules close to the surface
    have enough energy to overcome the attractions of
    neighboring molecules and escape
  • Slower molecules stay in the liquid state
  • Rate of evaporation increases as temp increases

48
Boiling
  • Occurs within the liquid
  • Boiling point temp at which vapor pressure
    equals atmospheric pressure
  • If vapor pressure is less than atmospheric
    pressure, bubbles do not form

49
Condensation
  • Change of a gas to a liquid
  • Exothermic process
  • Molecules of vapor can return to the liquid state
    by colliding with the liquid surface
  • The particles become trapped by the
    intermolecular attractions of the liquid

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Sublimation and Deposition
  • Sublimation solid goes directly to a gas
    without passing through the liquid phase
  • Deposition is the reverse process
  • Sublimation is endothermic
  • Deposition is exothermic

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Heating Curves
  • Graphic illustrations of phase changes
  • Plot of temp of a sample as a function of time
  • Notice temp remains constant during phase changes
    while amount of energy varies

54
Heating Curve of Water
A Rise in temperature as ice absorbs heat.B
Absorption of heat of fusion.C Rise in
temperature as liquid water absorbs heat.D
Water boils and absorbs heat of vaporization.E
Steam absorbs heat and thus increases its
temperature. The above is an example of a heating
curve. One could reverse the process, and obtain
a cooling curve. The flat portions of such curves
indicate the phase changes.
55
Phase Diagrams
  • Diagram that relates the states of a substance to
    temp and pressure
  • State depends on temp and pressure
  • 2 states can exist simultaneously at certain
    temps and pressures
  • Triple point the temp and pressure when all
    three states exist at the same time

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  • TRIPLE POINT - The temperature and pressure at
    which the solid, liquid, and gas phases exist
    simultaneously.
  • CRITICAL POINT - The temperature above which a
    substance will always be a gas regardless of the
    pressure.
  • FREEZING POINT - The temperature at which the
    solid and liquid phases of a substance are in
    equilibrium at atmospheric pressure.
  • BOILING POINT - The temperature at which the
    vapor pressure of a liquid is equal to the
    pressure on the liquid.
  • Normal (Standard) Boiling Point - The temperature
    at which the vapor pressure of a liquid is equal
    to standard pressure (1.00 atm 760 mmHg 760
    torr 101.325 kPa)
  • NOTE
  • The line between the solid and liquid phases is a
    curve of all the freezing/melting points of the
    substance.
  • The line between the liquid and gas phases is a
    curve of all the boiling points of the substance.
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