Chapter 10

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Chapter 10 Liquids Solids
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Intra- and Inter- molecular forces

• a. intramolecular forces
• bonds
• within molecule
• ionic or covalent

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Intra- and Inter- molecular forces

• Intermolecular Forces
• Between molecules
• Causes solids or liquids (condensed states of
  matter) to form as molecules bond together

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Intra- and Inter- molecular forces

• Determines many important properties of
  substances
• state
• boiling and melting points
• vapor pressure

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Ion-Dipole Forces

• Results from the electrostatic attraction of an
  ion and a dipole

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Dipole Dipole Forces

• Found in polar molecules
• Molecules with dipoles line up so that the
  positive end of one molecule is close to the
  negative end of another molecule
• The attraction weakens as the distance between
  molecules increases

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Dipole Dipole Forces
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Dipole Dipole Forces
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Dipole dipole attractions

• Much weaker than covalent or ionic forces
• But does explain why polar liquids are more
  soluble in polar liquids than in non-polar
  liquids
• It takes 2000 mL of H2O to dissolve 1 mL of CCl4
• It takes 50 mL of H2O to dissolve 1 mL of CH2Cl2

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London dispersion forces

• Found in non-polar molecules
• Instantaneous dipoles can be produced in
  non-polar molecules when electrons are not
  distributed evenly
• When an instantaneous dipole occurs in one
  molecule, dipoles are induced in the neighboring
  molecules.

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London dispersion forces
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London dispersion forces

• Weaker than dipole attractions and short-lived
• Increase with molecular size
• Larger molecules have greater London-dispersion
  forces
• Larger molecules have more electrons, creating a
  greater opportunity for uneven electron
  distribution

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Hydrogen Bonding

• Hydrogen bond is an especially strong
  dipole-dipole force, as shown by the trend in
  boiling points of polar molecules

HF
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Hydrogen Bonding

• H-bonding is observed for HF, H2O, NH3, but not
  CH4
• Conditions for occurrence
• H attached to a small, highly electronegative
  element in one molecule
• Small, highly electronegative element with one or
  more unshared electron pairs in the other
  molecule
• Observed for the elements
• F, O, N (rarely S and Cl)

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Hydrogen Bonding

• Which of the following molecules will
  hydrogen-bond in the pure substance?
• H2O
• H2Se
• HF
• HBr
• NH3
• PF3

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Hydrogen Bonding in Liquid Water

• H points at the electron pair on the atom in the
  other molecule
• In liquid water, each water molecule is
  surrounded by an average of 4 other water
  molecules structure is not rigid.
• Longer than covalent bond.

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Hydrogen Bonding

• Average of 4 hydrogen bonds in liquid water
  Figure 10.2
• Fluoride ion is hydrogen-bonded to water in
  solution

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Hydrogen Bonding

• Molecules hydrogen-bond to themselves or to other
  molecules.Figure 10.2

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Structure of Ice

• The water molecules in ice are fixed into a
  tetrahedral arrangement as a result of hydrogen
  bonding. Open structure makes ice less dense
  than water.

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Structure of Ice

• The open structure ofice leaveschannels
  ofempty spacethrough thecrystals.

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Identify Predominant Type of Intermolecular Forces
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Intermolecular Forces

• What types of intermolecular forces are observed
  for each of the following molecules?(A type of
  molecule may have more than one.)
• H2O HF
• HBr NH3
• PF3 CH3OH
• F2 CO
• CO2 N2

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Strengths of Intermolecular Forces

• Intermolecular forces generally increase in
  strength as
• London lt Dipole-Dipole lt H-bonding lt Ion-Dipole lt
  Ionic Bonding
• The forces are cumulative. All molecules have
  London forces. Polar molecules have both London
  and dipole-dipole forces. ...

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Trends in Intermolecular Forces

• Which member of each pair has the larger
  intermolecular forces? (we will learn later that
  this affects things such as boiling point, heat
  of vaporization)
• CH3OH, CH3SH
• F2, Kr
• F2, CO
• CO, HF
• CO2, NH3
• N2, NH3

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Structural models of liquids

• More complex than models for solids or gases for
  two reasons
• Liquids have strong intermolecular forces
• Liquids have significant molecular motion

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Surface Tension

• Resistance of a liquid to increase its surface
  area
• For the surface area of a liquid to increase,
  molecules would have to move up to the surface
• This would require internal molecules to pull
  away from their surrounding molecules, going
  against the intermolecular forces

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Surface Tension

• An uneven distribution of forces exists on
  surface molecules
• Molecules on the surface only experience
  intermolecular attractions with molecules below
  and to the side of them
• Molecules below the surface experience
  intermolecular attractions with molecules in all
  directions
• This causes molecules on the surface to be pulled
  to the interior, giving the surface a spherical
  shape

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Surface Tension

• Liquids with greater intermolecular forces have a
  greater surface tension
• Would you expect a polar liquid to have a greater
  surface tension than a nonpolar liquid? Why or
  why not?

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Surface Tension
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Surface Tension
How does a water strider stay on the top of the
water? Why does the needle float?
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Surface Tension

• Why does soap make the paper clip sink?

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Capillary Action

• Rising of a liquid in a narrow tube
• Occurs when the molecules in the container have
  polar bonds
• The polar bonds in the container attract the
  liquid, causing the liquid to try to creep up the
  sides of the container, which stretches the
  surface of the liquid

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Capillary Action

• The liquid tries to balance the attraction
  between liquid molecules (cohesive forces) and
  the attraction between the liquid and the
  container (adhesive forces)
• This causes the liquid to pull itself up the tube

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Capillary Action

• Explains the shape of the meniscus formed by a
  liquid in a tube
• The meniscus of water is concave because the
  attractions between the water molecules and the
  glass molecules are greater than the attractions
  between two water molecules

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Capillary Action

• What would you expect the meniscus to look like
  for a liquid in which the internal (liquid to
  liquid) attractions are stronger than the
  attractions between the liquid and the container?

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Capillary Action
Which forces in each system are greater?
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Viscosity

• A liquid's resistance to flow
• Can be compared to the "thickness" of a liquid
• Thicker liquids are more viscous
• Maple syrup has a greater viscosity than water

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Viscosity

• Intermolecular forces and molecular complexity
  both contribute to the viscosity of a liquid
• As intermolecular forces increase, viscosity
  increases
• As molecular complexity increases, viscosity
  increases

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Viscosity

• Why do you think that viscosity increases when
  the intermolecular forces or molecular complexity
  increases?

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Liquid Phenomenon

• In what way do you think that viscosity,
  capillary action, and surface tension are
  related?
• In other words, if a liquid had a high viscosity,
  what would you predict about its surface tension
  and capillary action?

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Phase Changes

• Physical states of a substance can co-exist under
  a variety of conditions of pressure and
  temperature.
• Phases different forms (gas, liquid, solid,
  etc.) of a substance that co-exist in a
  heterogeneous system.

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Phase Changes

• Transitions between phases are called phase
  changes
• evaporation liquid ? gas (reverse
  condensation)
• melting solid ? liquid (reverse freezing)
• sublimation solid ? gas (reverse deposition)

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What are the phase changes?
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Heat of Vaporization

• Is evaporation exothermic or endothermic?

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Heat of Vaporization

• DHvap energy needed to evaporate 1 mol of
  liquid at constant temperature
• Energy used to overcome intermolecular forces
  during evaporation
• Larger molecules have higher DHvap because of
  higher London forces
• Polar DHvap gt nonpolar DHvap if molecular size is
  similar
• H-bonded DHvap gt polar DHvap

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Heat of Fusion

• Generally heat of fusion (enthalpy of fusion) is
  less than heat of vaporization
• it takes more energy to completely separate
  molecules, than partially separate them.

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Freezing

• Cooling liquids decreases their kinetic energy
• When intermolecular forces become greater than
  kinetic energy, the liquid freezes and becomes
  solid.
• Freezing point temperature at which solid and
  liquid are in a state of equilibrium
• Normal freezing point f.p. at pressure of 1 atm.

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Vapor Pressure

• Evaporation loss of higher kinetic energy
  molecules, so the liquid cools (unless energy is
  supplied) (The process is endothermic.)
• Evaporative cooling, perspiration, alcohol bath,
  canvas water bags, wind chill factor

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Evaporation

• Open container evaporates completely
• Closed container reach a state of equilibrium

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Vapor Pressure

• Equilibrium rate of evaporation rate of
  condensation
• P at equilibrium vapor pressure
• When Pvap Patm, T boiling point
• When Pvap 1 atm, T normal boiling point
• In Phoenix, boiling point of H2O 99oC
• At sea level, b.p. of H2O 100oC
• At 9000 ft elevation, b.p. of H2O 91oC(needs a
  pressure cooker to speed up cooking)

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Vapor pressure varies with temperature and
intermolecular forces
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Phase Diagrams

• Phase diagram plot of pressure vs. temperature
  summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams
  tell us which phase will exist.
• lines equilibrium between two phases
• areas only one phase is stable
• triple point (confluence of 3 lines) equilibrium
  between three phases

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Phase Diagram Features

• Features of a phase diagram
• Triple point temperature and pressure at which
  all three phases are in equilibrium.
• Vapor-pressure curve generally as pressure
  increases, temperature increases.
• Critical point critical temperature and pressure
  for the gas.
• Melting point curve as pressure increases, the
  solid phase is favored if the solid is more dense
  than the liquid.
• Normal melting point melting point at 1 atm.

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Generic Phase Diagram
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Phase Diagram of Water

• Shows conditions of stability of phases and
  conditions of equilibriumFigure 10.45

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Phase Changes

• What phase changes are repre- sented by the
  arrows?
• Label features.

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Phase Diagram of Carbon Dioxide
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Slope of the solid-liquid boundary line

• If negative, the liquid is more dense than the
  solid and vice-versa

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The Wide World of Solids

• A concept map to organize how solids are
  classified

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Amorphous solids

• Highly disordered
• Described as a solution frozen in place
• Example Glass

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Crystalline solids

• Highly ordered arrangement
• Represented as a lattice
• Unit cell - smallest repeating unit of a lattice

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Unit Cells

• Three types of unit cells
• Simple cubic - a cube with atoms or molecules at
  each of the corners
• Body-centered cubic - a cube with atoms or
  molecules at each of the corners and one atom or
  molecule in the center of the cube.
• Face-centered cubic - a cube with atoms or
  molecules at each of the corners and atoms or
  molecules in the center of each side of the cube

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Crystalline Solids

• Structures determined by using X-ray diffraction
• This technology was helpful in determining the
  structure of DNA!
• Three main types of crystalline solids exist
  ionic, molecular, and atomic.

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Ionic Solids

• Stable
• High melting point
• Brittle
• Held together by ionic bonds
• Nonconductors when solid
• What do you think would happen to the
  conductivity of an ionic solid if it were melted
  or dissolved in water? Why?

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Ionic Solids

• Similar structural model to that of metallic
  solids
• Ions represented as hard spheres
• Ions are packed closely together and arranged in
  a way to maximize attraction and minimize
  repulsion

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Ionic Solids

• Example NaCl
• Chlorine ions are arranged in a face-centered
  cubic closest packed structure
• Sodium ions fill in the spaces in between the
  chlorine ions.

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Molecular solids

• Strong covalent bonding within the molecules
• Molecules held to each other by intermolecular
  forces
• Generally have low melting points
• Generally nonconductors of heat and electric
  current

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Molecular Solids

• Molecules are held together by intermolecular
  forces
• Polar molecules are held together by dipole
  forces
• Nonpolar molecules are held together by
  London-dispersion forces
• Intermolecular forces are generally greater when
  the molecules are polar

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Atomic solids

• Consist of atoms at each of the lattice points in
  the crystal
• Can be divided into Metallic, Network, and Group
  8A atomic solids

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Metallic solids

• High thermal conductivity
• High electrical conductivity
• Malleable
• Ductile
• Range from low melting point and soft to high
  melting point and brittle
• Strong, non-directional covalent bonding

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The Closest Packing Model

• Metal atoms are represented as hard spheres
• In a metallic solid, atoms are arranged in
  layers, packed as close together as possible, in
  a way that the space between atoms is minimized

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The Closest Packing Model

• Hexagonal closest packed structure
• Atoms arranged in the ABA form
• Creates a hexagonal unit cell
• Examples Magnesium and Zinc

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The Closest Packing Model

• Cubic closest packed structure (figure 10.15)
• Atoms arranged in the ABC form
• Creates a face-centered unit cell
• Examples Silver and Copper

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Metallic Solids

• Not all metallic solids have one of these two
  structures
• e.g. alkali metals
• Body-centered cubic unit cell
• Eight nearest neighbors (less dense)

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Bonding Models for Metallic Solids

• Electron Sea Model
• Represented as metal cations surrounded by a sea
  of electrons
• Similar to Thomson's Plum Pudding model of the
  atom

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Bonding Models for Metallic Solids

• Molecular Orbital Model
• The valence orbitals of the metals atoms
  hybridize in a way that allows valence electrons
  to move throughout the entire crystal

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Metal Alloys

• Alloy - a substance containing a mixture of
  elements and having metallic properties

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Metal Alloys

• Substitutional alloys - within a metal crystal,
  metal atoms are removed and replaced with atoms
  of similar size
• Example Brass

Cu
Cu
Cu
Zn
Zn
Cu
Cu
Cu
Zn
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Metal Alloys

• Interstitial alloys - small atoms fill in the
  holes in the metal crystal
• Example Steel
• Can make the metal stronger by creating
  directional bonds between the metal atoms and the
  smaller atoms

Fe
Fe
Fe
Fe
Fe
Fe
C
C
C
Fe
Fe
Fe
Fe
Fe
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Network Atomic Solids

• Brittle
• Poor conductors (thermal and electrical)
• Very high melting point
• Strong, directional covalent bonds
• Often referred to as "giant molecules"
• Have non-lattice type structures
• Carbon in diamond form is arranged in
  tetrahedrals connected together

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Changes of State

• Vaporization (evaporation) - molecules of a
  liquid escape the liquid's surface and form a gas
  
• Heat of Vaporization ( ?Hvap) - the energy that
  is required to vaporize one mole of liquid at one
  atmosphere of pressure

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Changes of State

• Condensation - vapor (gas) molecules going back
  into the liquid phase
• Sublimation - solid molecules escaping into the
  gas phase without going through the liquid phase
• Can you think of an example of sublimation?

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Vapor Pressure

• Vapor pressure - the pressure that is exerted by
  a vapor (gas), in a closed system
• Measured when the rate of evaporation equals the
  rate of condensation

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Vapor Pressure and Evaporation Rate

• When the vapor pressure is large, a large number
  of the liquid molecules are taking part in the
  evaporation/condensation equilibrium
• As the vapor pressure increases, the rate of
  evaporation increases.
• More of the molecules are in the gas phase

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Vapor Pressure and Intermolecular Forces

• The size of intermolecular forces has the
  greatest effect on vapor pressure
• As the intermolecular forces increase, the vapor
  pressure decreases
• Why do you think that increasing intermolecular
  forces decreases the vapor pressure?

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Vapor Pressure and Molar Mass

• Molar mass also affects the vapor pressure
• As the molar mass increases, the vapor pressure
  decreases

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Vapor Pressure and Temperature

• Vapor pressure increases with temperature
• As the temperature increases, more molecules will
  have enough energy to escape.

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Vapor Pressure and Temperature

• The relationship between vapor pressure and
  temperature is non-linear and can be represented
  by the following equation
• ln (Pvap) - DHvap 1 C
• R T
• Pvap vapor pressure Hvap heat of
  vaporization T temperature in K R
  universal gas constant C constant

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Comparing Vapor Pressures

• This equation can be rearranged to compare vapor
  pressure at two different temperatures
• ln PT1vap DHvap 1 - 1
• PT2vap R T2 T1

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The Heating Curve
Gas and Liquid
Gas
Liquid and Solid
Temperature
Phase Change No change in temperature
Liquid
Solid
Time with continuous energy input
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Heating Curves

• During a phase change, the temperature remains
  constant even though heat is being added
  continually
• If heat is still being added, why does the
  temperature not increase during a phase change?

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More State Changes

• Heat of Fusion ( DHfus) - the energy required to
  convert a solid to a liquid
• Melting point and boiling point are determined by
  the vapor pressure of the substance

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Melting Point

• The point at which the vapor pressure of the
  liquid equals the vapor pressure of the solid
  when the total pressure equals one atmosphere

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Melting Point

• Liquids and solids have characteristic vapor
  pressure at specific temperatures and pressures
• When a substance is in solid form, the
  temperature and pressure conditions do not favor
  the existence of the substance in the liquid
  phase

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Melting Point

• As the temperature is raised, the vapor pressure
  of the solid increases
• When the vapor pressure of the solid increases to
  the point that it has the same vapor pressure
  that a liquid would have under those conditions,
  this is the melting point

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Boiling Point

• The point at which the vapor pressure equals the
  pressure of the environment
• This means that when the atmospheric pressure is
  below one atmosphere, liquids will boil at a
  lower temperature

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Strange Phenomena

• Supercooling - when a liquid is cooled below the
  melting point and remains a liquid
• Occurs when the liquid is allowed to cool without
  being disturbed, causing the molecules to not be
  ordered enough to form a solid
• Temperature raises back up to melting point when
  solid begins to form

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Strange Phenomena

• Superheating - when a liquid is heated above the
  boiling point and remains a liquid
• Can occur when a liquid is heated too quickly
• The pressure in the liquid is greater than the
  pressure in the atmosphere preventing bubbles
  from forming
• Temperature drops back down to boiling point when
  gas begins to form

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Phase diagram

• Shows the relationship between phase,
  temperature, and pressure for a given substance

Pressure
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Triple point

• The temperature and pressure at which solid,
  liquid, and gas can all exist simultaneously
• This is the point at which the solid and liquid
  have the same vapor pressure
• Usually very close to the melting point

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Critical Temperature

• temperature above which the vapor cannot be
  liquefied, regardless of the pressure applied.

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Critical Pressure

• the pressure required to liquefy the vapor at
  critical temperature.

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Critical point

• Point at which the critical temperature and
  critical pressure coincide.
• The temperature and pressure at which the vapor
  can no longer be liquefied
• The vapor passes through a fluid phase before
  returning to the liquid state

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Meaning of the line segments

• The different line segments relate to the
  equilibrium points between phases at the specific
  temperatures and pressures.