Solids and Liquids

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Solids and Liquids
IMF, Properties, Changes of State
Go to this link http//www.quia.com/jg/455837.html
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Liquids and Solids The Condensed States
Gas Liquid Solid
Highly Compressible Very slightly compressible Least compressible
Low density High density High density
Fills container completely Does not expand to fill container has definite volume Rigid and retains its volume
Assumes the shape of its container Assumes the shape of its container Retains its own shape
Rapid diffusion Slow diffusion Extremely slow diffusion only at its surface
Total disorder particles have freedom of motion and are far apart from one another Disordered particles are free to move relative to one another and are close together Ordered arrangement particles can vibrate but remain fixed in position and are close together
High expansion on heating Low expansion on heating Low expansion on heating
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What causes a substance to be in one state or
another at room temp?

• All particles at room temperature have the same
  kinetic energy
• Kinetic molecular theory
• according to the kinetic molecular theory, the
  state of a substance at room temperature depends
  on the strength of the attractions between its
  particles

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Intermolecular Forces

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

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Strength of Inter vs. Intra
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Different Types of IMF

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

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Dispersion Forces

• The motion of electrons can create an
  instantaneous or temporary dipole 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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Dispersion Forces

• dispersion forces forces of attraction between
  induced dipoles
• Exist in all phases of matter
• These forces are found in ALL molecular
  compounds
• These are the only kinds of forces that effect
  nonpolar compounds.
• Increases with increasing molecular size and mass

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Dispersion Forces
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Dispersion Forces
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Recap

• Which of the following compounds will have
  dispersion forces?
• HF
• H2O
• CH4
• CH3COOH

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Recap

• Which of the following compounds will have the
  greatest dispersion force between its particles?
  Why?
• HF
• H2O
• CH4
• CH3COOH

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Dipole-Dipole Forces

• Polar covalent molecules have a positive end and
  a negative end (permanent dipoles)
• 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, but are present in all phases of
  matter
• For molecules of about the same size, dipole
  forces increase with increasing polarity

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Dipole-Dipole
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Hydrogen Bonds

• a special type of dipole-dipole force
• occurs between molecules containing a hydrogen
  atom bonded to a small, highly electronegative
  atom (H-N, H-O H-F)
• The small electronegative atom must have at least
  one lone pair of electrons
• The hydrogen in one molecule will be attracted to
  the electronegative atom in another molecule
• Strongest IMF

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Hydrogen Bonding Boiling Point
http//thestephenation.blogspot.com/2009/09/hydrog
en-bonding.html http//www.chem.ufl.edu/itl/2045/
lectures/lec_g.html
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Properties of Water

• Density of ice is less than the density of liquid
  water
• WHY??
• You get to figure it out!
• Structure of Ice Activity

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• In ice, hydrogen bonding causes hexagonal
  structures to form
• Prevents other molecules from getting inside the
  rings

Arrangement of molecules in liquid water
Arrangement of molecules in ice
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Water

• unexpectedly high boiling point causes it to be a
  liquid at room temp
• other hydrogen compounds are corrosive gases at
  room temp
• can absorb or release relatively large quantities
  of heat without large temp changes
• has relatively high surface tension
• has very high heat of vaporization
• called the universal solvent b/c it can dissolve
  so many things

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IMF Summary
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Properties of Liquids

• Only slightly compressible not a discernable
  difference when compressed
• Have much greater densities than their vapors
• 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

• Physical properties are determined mainly by the
  nature and strength of IMF present between
  molecules
• Viscosity resistance to flow
• Determined by
• 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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high viscosity thick/slow flowing low viscosity
thin/fast flowing
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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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Surface Tension
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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
• Amorphous solids

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Crystalline Solids

• particles exist in a highly ordered repeating
  pattern
• Precious stones, sugar, Ionic solids salts,
  atomic - Metallic solids
• 7 principle crystal patterns
• atoms, ions, or molecules arranged in an orderly,
  geometric, 3-D structure
• Smallest arrangement of repeated crystal is
  called a unit cell

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Examples of how particles can be arranged in a
cubic crystal.
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Amorphous Solid

• Solid arrangement of particles lacking a regular
  repeating pattern
• Like liquids that have been cooled to such a low
  temp that their viscosity becomes very high
• glass, rubber, wax, tar
• Particles are trapped in a disordered arrangement
  that is characteristic of liquids
• get softer over a wide range of temperatures
  before melting

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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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Molecular Solids

• Type of Particles atoms or molecules
• Held together by dispersion forces, dipole-dipole
  forces or hydrogen bonds
• Most are NOT solids at room temperature
• Poor conductors of heat and electricity (dont
  contain ions)
• Soft w/ low to moderate melting points
• Examples are sucrose, ice, most organics

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Sodium chloride
Cupric chloride
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Ionic Solids

• Type of particles cations anions
• Type and ratio of ions determine the shape of the
  crystaline structure
• The network of attractions that extend through an
  ionic compound gives these compounds their high
  melting points and hardness
• Hard and brittle, poor electrical and thermal
  conductors in solid state
• Examples are salts (NaCl, KBr, MgSO4)

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Covalent Network Solids-Atomic

• Atoms that can form multiple covalent bonds
• Properties very hard, very high melting point,
  often poor thermal and electrical conductors
• 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, silicon,
  quartz (SiO2)

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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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Carbon microtubules
Buckminster fullerene
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Gold
Copper
Silver
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Metallic Solids - Atomic

• 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
• Also possible to form alloys by this type of
  bonding
• substitutional vs. interstitial

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Why malleable and ductile?
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Phase Changes

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

Exothermic
Endothermic
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Vaporization

• The change of state from a liquid to a gas
• Vapor refers to the gaseous state of a
  substance that is normally a liquid or a solid at
  room temp
• Two methods of vaporization
• Evaporation
• Boiling

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Evaporation

• Occurs at the surface of a liquid
• Occurs b/c 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
• volatile evaporates easily, molecules dont
  exert a very strong attractive force upon one
  another
• evaporative cooling molecules with higher
  kinetic energy escape, the avg. kinetic energy of
  the remaining molecules decrease resulting in
  lower temperature.

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Vapor Pressure

• vapor pressure - the pressure of the vapor
  resulting from evaporation of a liquid (or solid)
  above a sample of the liquid (or solid) in a
  closed container
• Increases steadily as temperature increases
• The gas in the container is in equilibrium with
  the liquid or solid.

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Boiling

• Occurs within the liquid
• rapid vaporization of liquid
• boiling point temp at which vapor pressure
  equals atmospheric pressure
• heat of vaporization the amount of heat
  required to vaporize a given amount of liquid
• Liquids with strong intermolecular attractions
  have high heats of vaporization
• Although energy is added, temp remains constant
  during the phase change

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Condensation

• Change of a gas to a liquid
• 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
• Rate of condensation increases as the of vapor
  particles increases
• When the rate of vaporization and rate of
  condensation are equal, a state of dynamic
  equilibrium is reached (liquid-vapor equilibrium)

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Melting and Freezing

• Melting point/freezing point temp at which
  solid and liquid forms exist in equilibrium
• requires smaller potential energy changes than
  vaporization
• Particles are about the same distance apart in
  the solid and liquid forms
• not affected significantly by a change in
  external pressure
• Heat of fusion ?Hfus amount of heat required
  to convert a solid to a liquid
• depends on the strength of attractive forces
  between molecules

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When the rate of freezing is the same as the rate
of melting, the amount of ice and the amount of
water won't change. The ice and water are said to
be in dynamic equilibrium with each other. The
ice is melting, and the water is freezing, but
both are occurring at the same rate, so there is
no net change in either quantity.
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Sublimation and Deposition

• Sublimation solid goes directly to a gas
• Deposition is the reverse process
• solids exert vapor pressure
• tends to be much lower than liquid vapor pressure
• Solids with high vapor pressure sublime
  relatively easily
• Solids without strong attractive forces sublime
  readily, mostly molecular solids

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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
Gas
Temperature
Liquid
Solid
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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
• Critical point the temp and pressure
  combination at which a gas form of a substance is
  converted to a liquid

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Triple Point
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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.
• 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.
  
• 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)