Liquids and Solids

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Liquids and Solids

• Chapter 9

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Background information

• Molecules are much closer to one another in
  liquids and solids
• Gases molecules are separated by ten molecular
  diameters or more
• In liquids or solids, they touch each other
• Intermolecular forces, which are essentially
  negligible with gases, play a much more important
  role in liquids and solids

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9.1 Liquid-vapor equilibrium

• Vaporization of liquid in a closed system
• Dynamic equilibrium
• Rate of condensation rate of vaporization
• Liquid vapor
• Forward and reverse reaction are occurring at the
  same time

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Vapor Pressure

• When a liquid is placed in a closed container, it
  establishes equilibrium with its vapor.
• The pressure exerted by the vapor over the liquid
  remains constant
• liquid ? vapor

• The pressure of the vapor at equilibrium is
  referred to as the vapor pressure (VP) of the
  liquid.
• Characteristic property of a given liquid at a
  particular temperature
• Varies from one liquid to another, depending on
  the strength of the intermolecular forces

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Vapor Pressure of Ethanol
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Vapor Pressure of Various Liquids
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Vapor Pressure

• So long as both liquid and vapor are present, the
  vapor pressure is independent of the volume of
  the container.
• Small amount of liquid in the close system
  small amount of vapor before reaching equilibrium
• The greater the volume of the container, then
  more liquid will evaporate until equilibrium is
  reached at V.P.
• The vapor must obey the ideal gas law
• n/V stays constant so P nRT/V does not change
• Only if all the liquid vaporizes will the
  pressure drop below the equilibrium value

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Example

• Place 0.0100 mol of Benzene (C6H6) to 1.00 L
  flask at 25 C. (vp of benzene 92 mm Hg) How
  much evaporates?
• What if the volume was 2 L? 3 L?
• Hint make sure to use the information for the
  vapor, not the liquid
• Change mm Hg to atm by dividing by 760

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Temperature Dependence of Vapor Pressure

• The vapor pressure of a liquid always increases
  as temperature rises
• Increases steadily as temperature rises
  reflecting that more molecules vaporize at higher
  temperatures
• Increasing temperature means greater internal
  kinetic energy so a greater number of particles
  can evaporate - higher v.p.

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Page 249 two graphs

• If graph pressure versus temperature, we do not
  get a straight line
• We do if we plot P versus 1/T
• Equation for straight line
• y mx b (m slope, b y-intercept)
• Slope for this graph is
• - ?Hvap / R
• Ln P - ?Hvap / RT b

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Clausius-Clapeyron equation- page 250

• Units of Hvap and R must be consistent (recall
  table 5.1 on page 199)

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V.P. and Temperature

• ?Hvap heat of vaporization
• If ?Hvap of benzene is 30.8 kJ/mol.
• v.p. of benzene is 92 mm Hg at 25 C.
• What is v.p. at 50 C ?

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Boiling Point

• Temperature that vapor bubbles form in a liquid
  and rise to the surface where they break.
• Temperature at which a liquid boils depends on
  the pressure above it
• Vapor pressure (P2) must be at least equal to or
  greater than the applied external pressure (P1).

P2 P1
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Boiling

• A liquid boils at a temperature at which its
  vapor pressure is equal to the pressure above the
  surface
• If this pressure is 1 atm then this temperature
  is referred to as the normal boiling point
• Can increase of decrease boiling point by
  changing the pressure above it
• Water boils at 100C if P1 760 mm Hg
• If P2 1075 mm Hg, then B.P. 110 C
• If P2 5 mm Hg, then B.P. 0 C

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Critical Temperature and pressure

• A temperature that above which the liquid phase
  of a pure substance cannot exist
• Critical Temperature is where the particles are
  so mobile that they cannot condense.
• If the external pressure is greater than the
  vapor pressure then a gas can be condensed.
• Liquid Butane, Natural Gas, CO2, Air
• Called critical pressure pressure that must be
  applied to cause condensation above the critical
  temperature
• Critical Pressure is the vapor pressure at this
  critical temperature.
• Tc oxygen is -119 C so unable to liquefy at room
  temperature
• Tc propane is 98 C so able to liquefy and ship
  at room temperature

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• Table 9-1 lists the critical temperatures of
  several common substances
• Left hand side of table - cannot liquefy
  usually called permanent gases
• Have to cool the gas and lower pressure to get to
  liquefy
• When a substance is above the critical
  temperature and pressure, a substance is referred
  to as a supercritical fluid
• Have unusual solvent properties that have led to
  practical applications

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9. 2 - Phase Diagrams

• A graph showing solid, liquid, and vapor phases
  at different pressures and temperatures.

• Solids sublime below triple point

• If AD slopes negative than melting point
  decreases with pressure
• The liquid is denser than solid as in water

• If AD slopes positive than melting point
  increases with pressure

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Phase Diagram of a typical substance Solid is
denser then liquid
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Comparing DiagramsLeft Solid Floats Right
Solid Sinks
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Sublimation

• The process by which a solid changes directly to
  vapor without passing through the liquid phase
• A solid can only sublime at temperatures below
  the triple point
• Above that temperature, it would melt to a liquid
• At triple point, a solid can be made to sublime
  by reducing the pressure of the vapor above it to
  less than the equilibrium value

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Melting point

• Melting point freezing point (for a pure
  substance)
• Temperature at which solid and liquid phases are
  in equilibrium
• Usually measured in an open container at
  atmospheric pressure
• Most substances the melting point at 1 atm is
  virtually identical with the triple-point
  temperature

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When looking at the phase diagrams

• An increase in pressure favors the formation of
  the more dense phase

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9.3 molecular substances intermolecular forces

• See notes from unit 6 (bonding processes for this
  information)

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9.4 Network covalent, ionic, and metallic solids

• See notes from unit 6 (bonding processes for this
  information)

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9.5 Crystal structures

• Solids crystallize in definite geometric forms
  that be often be seen by the naked eye
• Crystals have definite geometric forms because
  the atoms or ions present are arranged in a
  definite, three-dimensional patter
• Deduced by x-ray diffraction
• Basic form that comes from this is the unit cell
• Smallest structural unit that, repeated over and
  over again in 3D, generates the crystal
• 14 different kinds of unit cells

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Metals

• Simpler unit cells found in metals
• Simple cubic cell (SC)
• Cube that consists of eight atoms whose centers
  are located at the corners of the cell
• Atoms at adjacent corners of the cube touch one
  another
• Face-centered cubic cell (FCC)
• Atoms at each corner of the cube and one in the
  center of each of the six faces of the cube
• Atoms at the corners of the cube do not touch one
  another
• Contact occurs along a face diagonal
• The atom at the center of each face touches atoms
  at the opposite corners of the face
• Body-centered cubic cell (BCC)
• A cube with atoms at each corner and one in the
  center of the cube
• Corners do not touch each other
• Contact occurs along the body diagonal
• The atom at the center of the cube touches atoms
  at opposite corners

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Table 9.6

• Number of atoms per unit cell
• Relation between side of cell (s) and atomic
  radius (r)
• Percentage of empty space

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Ionic crystals

• Like the three from before see page 270