Liquids and Solids

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Liquids and Solids
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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

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Three Types

• London Dispersion forces
• Dipole-dipole forces
• Hydrogen bonds

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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

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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.

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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

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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.
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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

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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.

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• 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

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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
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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

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Properties of Solids

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

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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

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Molecular such as sucrose or ice whose
constituent particles are molecules held together
by the intermolecular forces.
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Arrangement of molecules in liquid water
Arrangement of molecules in ice
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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

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Graphite
Diamond
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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 structure matches
the hexagonal crystal shape
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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

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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

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• 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

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Sodium chloride
Cupric chloride
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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

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• 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.

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• 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.

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Gold
Copper
Silver
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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

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Phase Changes

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

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Melting and Freezing

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

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Vaporization

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

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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

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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

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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

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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.
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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.