CHM 120 CHAPTER 12 Intermolecular Forces Liquids, solids, and phase changes

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CHM 120CHAPTER 12 Intermolecular Forces -
Liquids, solids, and phase changes

• Dr. Floyd Beckford
• Lyon College

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• Solids and liquids are referred to as the
• condensed phases
• In the condensed phases, particles are close
• together and interact strongly
• Clear distinction between gases and solids and
• liquids
• It is related to the nature of the electrostatic
• interactions between particles

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INTERMOLECULAR FORCES

• Consider a water molecule two polar bonds,
• polar molecule, strong covalent O-H bonds
• The covalent bond is an intramolecular force
• Intermolecular forces the forces between two
• individual particles of a substance
• As a class they are sometimes referred to as
• van der Waals forces

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• Includes
• (i) Ion dipole
• (ii) Dipole dipole
• (iii) London dispersion forces
• (iv) Hydrogen bonding
• Electrostatic in nature they result from the
• mutual attraction of unlike charges

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ION DIPOLE FORCES

• Results from the attraction between an ion and
• a dipole
• e.g. dissolution of NaCl in water
• Magnitude of interaction depends charge on
• the ion, size of the dipole moment, ? and the
• distance between them
• E Z?/r2

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DIPOLE DIPOLE FORCES

• Occur between polar covalent molecules

• They are usually weak important only at close
• proximity

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• Factors affecting the interaction
• 1. Size of the dipole moments
• - the more polar a substance the greater
• the strength of its dipole-dipole interactions
• Result substances with strong intermolecular
• forces typically have higher physical constants
• e.g. boiling points

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LONDON DISPERSION FORCES

• Weak attractive forces important only over
• very short distances
• Exist in all types of molecules in condensed
• phases
• The only forces present in symmetrical,
• nonpolar substances
• - He, SO3, CO2, N2, etc.

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• Results from the attraction of the nucleus of
• one atom for the electron cloud of a nearby
• molecule
• Creates an instantaneous dipole in one atom ?
• induces temporary dipole in a neighbor
• - weak attractive forces develop
• Magnitude of the forces depend on the ease of
• distortion of a molecules electron cloud

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• Referred to as the polarizability of the
  molecule
• Polarizability depends on
• size of molecule and
• (ii) atomic number
• Larger molecules with more electrons are more
• polarizable
• Polarizability increases down a group
• Polarizability decreases across a period

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HYDROGEN BONDING

• Hydrogen bonds are a special case of very
• strong dipole dipole interaction
• H-bond an attractive interaction between an
• unshared electron pair on a small, highly
• electronegative atom (F, O, N) and a nearby
• hydrogen atom attached to another
• electronegative atom

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• The bonds with F, O, and N are highly polar
• large dipole moment
• Electron cloud concentrated on these atoms

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• H-bonds are quite strong (up to 40 kJ/mol)
• Has a large impact on the physical properties
• of substances
• In water an extensive 3D network of H-bonds
• exist

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• This gives ice an open structure so it floats
• H-bonds are responsible for the relatively high
• boiling points of species such as NH3, CH3OH
• e.g. CH4 lt SiH4 lt GeH4 lt SnH4

• Consider also the Group 6 hydrides
• H2O gt H2S lt H2Se lt H2Te

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THE LIQUID STATE

• Properties of liquids are explainable by the
• intermolecular forces
• VISCOSITY
• Viscosity is the resistance of a liquid to flow
• Generally the stronger the intermolecular forces
• the more viscous the liquid
• Large H-bonds ? usually high viscosities

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Glycerol vs Ethanol
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• SURFACE TENSION
• A measure of the inward forces that must be
• overcome to expand the surface area of a liquid

• Surface tension is usually higher in liquids
  that
• have high intermolecular forces

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CAPILLARY ACTION

• The movement of a liquid up a surface against
• gravity
• Results from the competition of cohesive forces
• and adhesive forces
• Cohesive forces the intermolecular forces
  within
• the liquid
• Adhesive forces intermolecular forces between
• the liquid and the surface

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PHASE CHANGES

• Phase changes processes in which the
• physical identity of a substance changes
• FUSION (melting) - solid ? liquid
• FREEZING - liquid ? solid
• VAPORIZATION - liquid ? gas
• CONDENSATION - gas ? liquid
• SUBLIMATION - solid ? gas
• DEPOSITION - gas ? solid

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• All phase changes are accompanied by a
• redistribution of energy
• Phase1 ? Phase2 ?Hphase change
• Liquid ? Gas ?Hvaporization
• Heat changes may be represented by heating
• curve

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• Points to note
• 1. No temperature change occur during the
• actual phase change
• 2. At the melting point the additional heat is
• used to break intermolecular forces
• Heat of fusion, ?Hfusion the amount of heat
• that is required to melt 1 mole of a solid
• - 6.01 kJ mol-1 for ice

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• Boiling point temperature at which the liquid
• and gas coexist in equilibrium
• 3. At boiling point, temperature is constant
  while
• the liquid changes into vapor
• Amount of energy for the process (3) is the heat
• of vaporization, ?Hvap
• - 40.67 kJ mol-1 for water
• - due to the large of H-bonds to be broken

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VAPOR PRESSURE

• The pressure above the liquid (at equilibrium)
• is called the vapor pressure

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• Evaporation (or vaporization) the process by
• which molecules on the surface of a liquid
• enters the gas phase
• Vapor pressure depends directly on
• (i) Strength of the intermolecular forces
• - Smaller intermolecular forces ? surface
• molecules held weaker ? easier they break away
• ? higher vapor pressure

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• (ii) Temperature
• - higher T ? more molecules with the necessary
• energy to leave ? more molecules in the gas
• phase ? higher vapor pressure

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• The Clausius-Clapeyron equation

• The boiling point of a liquid is the temperature
  at which its vapor pressure equals the external
  pressure

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• The normal boiling point is when the external
• pressure is exactly 1 atm
• - can see why liquids boil at lower temperatures
• at higher elevations
• - also the basis of the pressure cooker

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SOLIDS

• Divided into two broad categories
• 1. Crystalline solids
• - have particles in an extended orderly
• arrangement
• - this arrangement is seen at the macro- and
• atomic levels
• 2. Amorphous solids
• - have the particles randomly arranged

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CRYSTALLINE SOLIDS

• Subdivided into ionic, molecular, covalent or
• metallic
• Ionic solids e.g. NaCl NaCl- the crystal
  is
• an extended array of alternating Na and Cl-,
• held together by ionic bonds

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• Molecular solids Constituent particles are
• molecules e.g. CO2, ice
• - particles held together by the usual inter-
• molecular forces
• Covalent solids Consist of a large 3D network
• of covalent bonds a very large molecule
• e.g. diamond
• Metallic solids Large array of atoms

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CRYSTAL STRUCTURE

• All crystals contain 3D regularly repeating
• arrays of particles
• Unit cell the smallest repeating unit of a
  crystal
• (that shows all the characteristics of the solid)
• There are different unit cell geometries
• - most common - the cell with cubic
• geometry cells with equal sides at right
• angles

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• Three basic types of cubic unit cells
• 1. A primitive (or simple) cell has a particle
  at
• each of its eight corners
• - each atom is shared by 7 others
• 2. A body-centered cell a primitive cell with an
• additional atom in the center of the cube

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3. A face-centered cell a simple cube with a
particle in each of its six faces - the facial
particles are shared by a neighboring cube
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• NaCl, KCl have face-centered cubic unit cells
• the bigger Cl- is at the corners and face centers
• and the Na fits between them
• - each cell has 4 Cl- ions
• (1/8 x 8) 1 at corners (½ x 6) 3 faces
  4
• - also 4 Na ions
• (¼ x 12) 3 at the edges 1 center 4

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PACKING

• In crystals particles are packed so that the
  inter-
• molecular forces are maximized
• Two most efficient ways of doing this
• 1. The hexagonal closest-packed arrangement
• - have two alternating layers ABABAB..
• - each layer has a hexagonal array of touching
• particles

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• - the particle of layer B sits below the
  triangular
• depressions created by layer A
• 2. The cubic closest-packed arrangement
• - three alternating layers ABCABCABC.
• - the third layer is offset from the first two
• In both closest-packed arrangement
• - each particle has a coordination number of 12
  
• 6 in the same layer, 3 on top, 3 below

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MO BAND THEORY

• Recall that the overlap of n atomic orbitals
  give
• rise to n molecular orbitals (MO)
• - n/2 bonding MO and n/2 antibonding
• MO
• In a metal the bonding MOs overlap to form a
• band
• Lower energy MOs contain valence electrons
• and makes up the valence band

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• Higher energy MOs are empty and make up
• the conduction band
• There is free flow of electrons from one band
• to the other
• This model explains the typical physical
• properties of metals
• e.g. color, conductivity

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• Band theory can also be applied to other solids
• - useful in explaining conduction properties

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PHASE DIAGRAMS

• Phase diagrams show the equilibrium
• pressure-temperature relationships among the
• different phases in a closed system
• The lines indicate phase boundaries the two
• phases are in equilibrium
• One three-way intersection all three phases
• are in equilibrium

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• The point of the three-way intersection is
• called the triple point
• Another interesting point the critical point
• - the temperature at this point, Tc, represents
• where no more gas can be further liquefied
• - Pc, the critical pressure beyond this point
  a
• liquid cannot be vaporized

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• At critical point a liquid and a gas are
  virtually
• indistinguishable the two phases blend to give
• a supercritical fluid
• Comparing the phase diagrams for water and
• CO2 shows up an interesting point
• - the slope of the solid-liquid equilibrium
  line
• is negative for H2O

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• - a consequence of the unusual fact that solid
• water is less dense than the liquid form
• We can also see why solid CO2 (dry ice)
• sublimes
• - the triple point is at 5.11 atm CO2 cannot
  be
• a liquid below this pressure