Liquids, Solids and Changes of State

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Liquids, Solids andChanges of State

• Kinetic-Molecular View of Liquids and Solids
• Intermolecular Attractions
• Properties of Liquids
• Vapor Pressure and Boiling Point
• Melting Points and Freezing
• Heating and Cooling Curves
• Phase Diagrams
• Crystals

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Kinetic-molecular viewof liquids and solids

• All real gases can be condensed to liquids by
  lowering the temperature and increasing the
  pressure.
• This decreases the average speed of the
  molecules.
• When moving slow enough, they will be attracted
  to each other and form a liquid.

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Kinetic-molecular viewof liquids and solids

• If the temperature is further decreased
• Molecules can no longer move about freely.
• Motion is limited to vibration.
• Rapid temperature decrease results in a
  disorderly arrangement - amorphous.
• A slow temperature decrease allows molecules
  to form a crystalline solid.

4
Intermolecular forces

• For molecules to form liquids and solids, there
  must be attractions between the them.
• Intermolecular attractive forces
• dipole-dipole attraction including hydrogen
  bonding
• Dipole-induced dipole
• London (dispersion) forces
• Sometimes called van der Waals forces
• Relative strength
• hydrogen bonding gt dipole-dipole gt
  dipole-induced dipole gt London

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Dipole-dipole attractions

• - When electrons that make up a bond are not
  equally shared because of a difference in
  electronegativity.
• ? and ?- ends are attracted to each other.

solid
liquid
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Hydrogen bonding

• An unusually strong dipole-dipole attraction.
• Occurs when hydrogen is bound to fluorine, oxygen
  and nitrogen -- the most electronegative
  elements.
• The small sizes of the elements involved and the
  large electronegativity differences result in
  large d and d- values.
• Hydrogen bonds are usually represented using a
  dashed line.

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Hydrogen bonding
The hydrogens of one water molecule interact
with the oxygen on other water molecules.
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London forces

• Temporary dipole attractions that exist between
  molecules - also called the dispersive.
• Results from random electron motion.
• Relatively weak force.

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Properties of liquids

• Diffusion
• This takes place in both liquids and gases. It
  is the spontaneous mixing of materials that
  results from the random motion of molecules.

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Properties of liquids

• Viscosity
• Resistance to flow.
• This increases with increased intermolecular
  attractions.
• Also, liquids composed of long, flexible
  molecules can entwine, resulting in increased
  viscosity - motor oil.

Increasing viscosity
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Properties of liquids

• Surface Tension
• Force in the surface of a liquid that makes the
  area of the surface as small as possible.

Molecules at the surface interact only with
neighbors inside the liquid.
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Properties of liquids

• Capillary action
• It is the competition between two forces.
• Cohesive forces
• The attractions between molecules of a
  substance.
• Adhesive forces
• Attractions between molecules of different
  substances.

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Properties of liquids

• Capillary action.

Capillary tube
meniscus
Mercury Cohesive is larger than adhesive.
Water Adhesive is larger than cohesive.
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Properties of liquids

• Vaporization
• The formation of a gas from a liquid.
• At any temperature, at least a few of the
  molecules in the liquid are moving fast enough to
  escape.

after some time
molecules leave and re-enter liquid at the same
rate
equilibrium
initially
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Equilibrium

• A state where the forward and reverse conditions
  occur at the same rate.

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Equilibrium
A point is ultimately reached where the rates of
the forward and reverse changes are the
same. At this point, equilibrium
is reached.
Rate
Time
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Chemical equilibrium

• A dynamic process on the molecular level achieved
  when concentration of reactants and products
  remain constant over time.
• - for a physical process
• H2O(l)
  H2O(s)
• (reactant) (product)
• - the equilibrium process is indicated with an
  equilibrium arrows.

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Vapor Pressure and boiling point

• Equilibrium vapor pressure
• The pressure of a vapor in equilibrium with a
  liquid.
• It depends on
• the intermolecular forces in the liquid.
• temperature.
• It is independent of
• the volume of the liquid or vapor
• the surface area of the liquid

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Boiling Points

• Boiling point - temperature where the vapor
  pressure equals atmospheric pressure.

This is the reason that cake mixes include high
altitude baking instructions.
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Boiling point

• Boiling points are dependent on pressure.
• Normal
• boiling point
• The boiling
• point at
• standard
• atmospheric
• pressure
• (760 mmHg)

Vapor pressure of H2O
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Melting point

• Normal melting point
• Temperature at which a solid changes to a liquid
  at atmospheric pressure.
• Freezing point.
• The temperature at which a liquid changes to a
  solid.
• For the same substance, these will both be at the
  same temperature.

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Changes in state

• A substance can usually be converted to different
  states by adding or removing energy from a
  system.
• If energy must be added, the change is
• - endothermic
• If energy is given off, the change is
• - exothermic
• The same concept can also be applied to chemical
  reactions.

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Endothermicchanges of state

• Sublimation
• The direct conversion of a solid to a gas.
• Example - dry ice (solid CO2)
• Melting or fusion
• The conversion of a solid to a liquid.
• Example - melting of ice
• Evaporization or vaporization
• Converting a liquid to a gas.
• Example - boiling water
• Most materials first melt then vaporize as you
  raise the temperature.

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Endothermicchanges of state
Gas
evaporation or vaporization
sublimation
Solid
Liquid
melting or fusion
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Exothermicchanges of state

• Condensation or liquifaction
• The conversion of a gas to a liquid or solid.
• Example - steam becoming water
• Freezing or crystallization
• When a liquid becomes a solid.
• Examples - formation of ice from water
• Substances usually first condense to liquids and
  then become solids.

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Exothermicchanges of state
Gas
liquification or condensation
deposition
Solid
Liquid
freezing or crystallization
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Changes in state andattractive forces

• As the attractive forces between molecules become
  larger, more energy is needed to separate them.
• Vapor pressures become smaller, boiling points
  and melting points become larger.

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Heating and Cooling

• Changes in state involve several steps.
• Example.
• Producing 150 oC steam from -20 oC ice.
• 1. Heat ice up to 0 oC.
• 2. Convert the ice to water.
• 3. Heat the water from 0 oC to 100 oC.
• 4. Convert the water to steam.
• 5. Heat the steam to 150 oC.

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Heating and Cooling

• Heat of fusion, DHfus
• The amount of thermal energy necessary to melt
  one mole of a substance at its melting point.
• Heat of vaporization, DHvap
• The amount of thermal energy necessary to boil
  one mole of a substance at its boiling point.

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Heating and Cooling

• mp DHfus bp DHvap
• Substance oC kJ/mol oC kJ/mol
• Br2 -7.3 10.57 59.2 29.5
• CH3CH2OH -117.0 4.60 79.0 43.5
• CH3(CH2)6CH3 -56.8 20.65 125.7
  38.6
• H2O 0.0 6.01 100.0 40.7
• Na 97.8 2.60 883 98.0

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Specific heats

• Each substance requires a different amount of
  energy to increase its temperature.
• Specific heat - amount of energy needed to
  increase a substances temperature by 1oC.
• It also depends on the state of the substance.

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Phase Diagrams

• Graphs that show the states of a substance as a
  function of both pressure and temperature.

liquid
solid
pressure
gas
temperature
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Phase Diagrams

• Partial phase diagram for water.

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Phase Diagrams

• Triple point
• All three phases are in equilibrium.
  Temperature and pressure are fixed.
• The triple point for water is at 0.01 oC and
  4.58 mmHg.
• The triple point for water, 273.16 K is used to
  define the Kelvin temperature scale.

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Phase Diagrams
supercritical fluid region
Pc
liquid
solid
Critical point
pressure
Tc
gas
temperature
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Phase Diagrams

• Critical point
• The end of the vapor pressure curve.
• Critical temperature, Tc
• The temperature at the critical point.
• Critical pressure, Pc
• The pressure at the critical point.
• At temperatures above Tc, liquefying a gas is
  impossible, no matter what the pressure.

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Phase Diagrams

• At pressures and temperatures above the critical
  point, a supercritical fluid is formed.
• A supercritical fluid
• is a gas.
• has a density similar to a liquid.
• has a viscosity similar to a gas.
• Supercritical fluids have a number of uses. One
  example is their use for extractions - removal of
  caffeine from coffee.

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The solid state

• At room temperature, solids
• are not compressible
• commonly have regular repeating units
• Two types are observed
• Crystalline solids have a definite melting
  point. - ionic - covalent
• - molecular - metallic
• Amorphous solids do not have a definite melting
  point or regular repeating units.

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Ionic solids

• Ions make up the repeating units.

NaCl
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Covalent solids

• Repeating units of covalently bound atoms.

Graphite
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Molecular solids

• Repeating units are made up of molecules.

Ice
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Metallic solids

• Repeating units are made up of metal atoms,
• Valence electrons are free to jump from one atom
  to another,

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Arrangement of units in crystals

• Metals
• All atoms are spherical.
• In a crystal, they are packed to minimize the
  space they occupy.
• Coordination number
• The number of nearest neighbors that surround
  an atom in a crystal
• Unit cell
• The smallest three dimensional unit that
  describes the arrangement of the atoms

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Simple cubic crystals

• This is one of the simplest arrangements to
  visualize.
• Each atom has a coordination number of six.
• Only 52 of the space is occupied.

Single layer
Expanded model
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Close Packing

• The crystals of most metals are of this type.
• Each atom is surrounded by 6 neighbors in its
  layer and a total of 12 in three dimensions..
• This results in a high percentage of the space
  being occupied.

Single layer
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Close Packing
The atoms in a layer must rest in holes of the
two layers that touch it. Two types of
crystals can result.
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Cubic close packing

• Face-centered cubic cell
• Each face has five
• atoms the maximum
• amount of space is
• occupied by the
• atoms - 74.

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Body-centered cubic unit cell

• This unit cell is observed for all metals that do
  not crystallize in one of the two close-packed
  arrangements.
• The exception is polonium.
• The coordination
• is eight.
• 68 of the space is
• occupied.

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Body-centered cubic, GaAs
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Crystal structures

• Coordination of space
• Name number occupied Example
• Face-
• centered 12 74 Al
• cubic
• Body-
• centered 8 68 Na
• cubic
• Simple
• cubic 6 52 Po

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Ionic compounds

• Crystal structures for these compounds are
  complicated by the following
• Two or more kinds of particles are involved.
• The particles are usually differ in size and
  often in charge.
• Not all ions are spherical.
• The major attractive force is electrostatic and
  crystals should allow the largest number of
  oppositely charged particles to touch.

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Ionic compounds

• Many ionic compounds will assume a close-packed
  arrangement of anions.
• Small cations are placed in the holes.
• Because each is surrounded by four spheres, the
  smaller holes are called tetrahedral holes.

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Ionic compounds

• Many compounds will have this type of structure
  including LiCl, NaCl, NaBr, MgO, NiO, and NH4I.
• NaCl

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Ionic compounds

• CsCl

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Ionic compounds

• Al2O3

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X-ray diffraction

• This method is used to find the dimensions and
  shape of a crystal unit.
• It provides a fingerprint of a material which
  can be used
• To deduce the structure of a material
• To identify a substance
• To tell structure of a polymer
• For elemental analysis

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X-ray diffraction
Film can be used for detection of the
patterns It is now more common to rotate the
crystal and detect the x-rays with a fixed
position detector. This way, you have data that
can be processed by a computer