Intermolecular forces – forces that exist between molecules
o determines many of the physical properties of molecular liquids and solids
o lead to deviations from ideal gas behavior as well
Molecular Comparisons of Liquids and Solids
 Liquids
o IMF are strong enough to hold molecules close together
o Liquids are much denser and far less compressible than gases
o Have a definite volume
o Attractive forces are not strong enough to keep molecules from moving
past one another, so they can be poured
 Solids
o IMF are so strong that molecules are virtually locked into place
o Solids are not very compressible b/c there is little space between them
o Often molecules take of positions in a highly regular pattern –
CRYSTALLINE SOLIDS
o Solids are rigid
o Units of solid vibrate in place b/c they have thermal energy
 Can change substances from one state to another by heating or cooling them
o This changes the avg. KE of the particles
o Increasing pressure will force molecules closer together
 This increases the strength of the IMF
 Propane vs. liquid propane at RT (increased pressure)
Intermolecular Forces
 Much weaker than ionic or covalent bonds
o Less E is required to vaporize a liquid or melt a solid than to break
covalent bonds in molecules
o Example – vaporize HCl requires 16 kJ/mol, Break H-Cl bond requires
431 kJ/mol
 Boiling points and melting points reflect strengths of intermolecular forces
o Boiling = bubbles of its vapor form in the liquid
o Stronger IMF between molecules = higher BP and MP
 IMF between neutral molecules:
o Dipole-dipole forces, London dispersion forces, hydrogen bonding
o Are called van der Waals forces (as a group)
o All are less than 15% as strong as a covalent bond
 Ion-dipole forces
o Between an ion and a partial charge on the end of a polar molecule
o Positive ions attract to the negative end of a dipole
o Magnitude of attraction increases as charge of ion or magnitude of dipole
moment increases
o Important for solutions of ionic substances in polar liquids
 Dipole-dipole forces
o The positive end of one dipole is near the negative end of another
o Only works when polar molecules are close together
o Molecules are free to move with respect to one another
 Sometimes in repulsive orientation, sometimes attractive
 Spend more time in attractive orientation
o
o
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For molecules of approximately equal mass and size, strengths of IMF
increase with increasing polarity
 Different dipole moments affect strength of IMF
Smaller molecules can get closer together, so have higher dipoledipole forces between them
London Dispersion Forces
o Nonpolar molecules can’t have dipole-dipole interactions
o Motion of electrons in an atom or molecule can create an instantaneous
dipole moment
o Instantaneous distribution of electrons of electrons can be different from
the average distribution
 Because e- repel one another, motions of e- influences motion of
e- on neighboring molecules
 The temporary dipole on one molecule can induce a similar dipole
on a neighboring atom
 Significant only when molecules are close together
o Polarizability – the ease with which the charge distribution in a molecule
can be distorted by an external electric field
 Larger molecules tend to have greater polarizabilites
 They have a greater # of e- and their e- are farther from the nuclei
 London dispersion forces increase with increasing MW
o Shapes of molecules influence as well
 More surface area for attractions to occur result in stronger
attraction
o Dispersion forces operate between all molecules
 Can account for most of the attraction between polar molecules
like HCl
 When molecules have similar weights and shapes, then dipoledipole interactions will determine which attraction is stronger
 When molecules have very different weights, dispersion forces
decide – most massive molecule will have the strongest
attractions
Hydrogen bonding
o A special type of IMF between the hydrogen atom in a polar bond (usually
with F, O or N) and an unshared electron pair on a nearby small
electronegative ion or atom (usually an F, O or N atom on another
molecule)
o Are unique dipole-dipole attractions – because F, N and O are so
electronegative, the bond with H is quite polar (H has no inner e-)
 H has a nearly bare p+ that can attract the negative charge of the
EN atom in a nearby molecule
 H is also very small, so it can approach the molecule more closely
o H bonds are weaker than ordinary chemical bonds, but are stronger than
other IMF
 Stabilize structures of proteins and DNA
o Responsible for the low density of ice
 When water freezes, the molecules assume a very ordered
arrangement to optimize H bonding arrangements between
molecules
 Lowers the density of ice (occupies more volume)

Comparing IMF
o Dispersion forces found in all substances
 Strength increases with increasing MW
o Dipole-dipole forces add to effect of DF and are found in polar molecules
o Hydrogen bonds add to effect of DF and are the strongest type of IMF
o No IMF is as strong as ordinary ionic or covalent bonds
o Nice summary chart – pcopy for students
Some properties of liquids
 Viscosity – the resistance of a liquid to flow
o Greater viscosity, the more slowly it flows
o Unit = poise (P) = 1 g/cms
o Reported in centipoises (cP) most often
o Depends on attractive forces between molecules and structural features
 Increases with MW
 Decreases with increasing temperature
 Surface tension
o Molecules at surface experience a net inward force (pulled inward)
 Reduces surface area (such as in a sphere)
 Molecules at surface behave like a “skin”
o Measured in energy required to increase the surface area of a liquid by a
unit amount (J/m2)
o Cohesive forces – IMF that bind similar molecules to one another
o Adhesive forces – IMF that bind a substance to a surface
 Meniscus forms b/c of this! IMF between water and glass is
greater
 Mercury meniscus points up!
o Capillary action – the rise of a liquid up a narrow tube
 Adhesive forces between liquid and wall increase the SA of the
liquid, surface tension of liquid reduces the area, pulls liquid up
tube
 Helps plants get water and nutrients
Phase Changes
 Changes of state for a chemical (from solid to liquid to gas, or vice versa)
o Every phase change is accompanied by a change in the energy of the
system
o As temperature of solid increases, units of solid vibrate with increasingly
energetic motion
o When it melts, they are freed to move with respect to one another (their
average separations increase)

Melting process is called FUSION
o Heat of fusion (enthalpy of fusion) - Hfus – the enthalpy change
required to melt a solid

As temperature of liquid phase increases, the molecules of the liquid move about
with increasing energy
o Concentration of gas molecules over the liquid increases with
temperature
o These molecules exert vapor pressure (increases with increasing
temperature)
o Once vapor increases to external pressure, the liquid boils
o The molecules of the liquid move into the gaseous state when they are
widely separated
o Heat of vaporization - Hvap – the energy required to cause the transition
from the liquid state to gaseous state (the enthalpy change for vaporizing
a liquid)

Hvap values tend to be larger because in the transition from liquid to vapor state,
the molecules must essentially sever all of their intermolecular attractive forces
Hsub – the sum of Hfus and Hvap

Practical applications
 heat of fusion of ice cools the liquid in which the ice is immersed
 heat of vaporization is drawn from our bodies as the water evaporates from our
skin (sweat/stepping out of a pool)
 refrigerator – has an enclosed gas that can be liquefied under pressure, gas
absorbs heat as it expands to a chamber where it is evaporated, and this
cools the interior of the refrigerator, then vapor is recycled through a
compressor
o what about when the vapor is condensed?
 Heat of condensation is equal in magnitude but with opposite
sign from the heat of vaporization
 The heat is dissipated through cooling coils in the back of the
refrigerator
 Heat of condensation – opposite of heat of vaporization (exothermic)
 Heat of freezing – opposite of heat of fusion (exothermic)
 Heat of deposition – opposite of heat of sublimation (exothermic)
MELTING/VAPORIZING/SUBLIMING = ENDOTHERMIC
FREEZING/CONDENSING/DEPOSITING = EXOTHERMIC

Heating curve – a graph of the temperature of the system versus the amount of
head added
o heat ice from -25°C to 0°C – heat added, temperature changes
o melt ice – heat added, no temperature change
o heat water from 0°C to 100°C – heat added, temperature changes
o vaporize water – heat added, no temperature change
o heat steam from 100°C to 125°C – heat added, temperature changes

For heating a single phase from one temperature to another, use
q = mcT
*specific heat of water is greater than that of ice
To convert from one phase to another, use Hfus or Hvap

Specific heats for H2O
Ice = 2.09 J/g°C
Water = 4.18 J/g°C
Steam = 1.84 J/g°C
Phase change constants for water
Hfus = 6.01 kJ/mol
Hvap = 40.67 kJ/mol
Critical Temperature and Pressure – teach IMF first
 A gas liquefies at some point when pressure is applied to it
o If we increase the pressure on water vapor at 55°C, then it liquefies when
the pressure equals 118 torr, and an equilibrium between the gaseous
and liquid phases exists
o If the temperature is 110°C, then the liquid phase does not form until the
pressure is 1075 torr.
o At 374°C, the liquid phase will only form at 1.655 x 105 torr (215.7 atm)
 Critical temperature – the highest temperature at which a distinct liquid phase
can form
 Critical pressure – the pressure required to bring about liquefaction at this
critical temperature
 Nonpolar, low molecular weight substances have lower critical temperatures and
pressures than polar or heavier substances
o The transition from gaseous to liquid state is determined by IMF
o For every gas, a temperature can be reached at which the motional
energies of the molecules are sufficient to overcome the attractive forces
that lead to the liquid state, regardless of increasing pressure
o Water and ammonia have high critical temperatures and pressures, due
to strong hydrogen bonding forces
 Important to engineers/gas workers, b/c they give information about conditions at
which gases liquefy
o Sometimes we want this, other times we want to avoid this
o O2 is harder to liquefy than ammonia
Vapor Pressure
 The pressure exerted by a vapor in equilibrium with its liquid or solid phase
o In a closed container, some liquid will begin to evaporate into the gaseous
phase
o After a short time, the pressure will attain a constant value
 At any instant, the molecules on the surface of the liquid possess enough KE to
overcome attractive forces and escape into the gas phase
o At any temperature, movement from liquid to gas phase occurs
continuously
 As number of gas-phase molecules increases, the probability
increases that a molecule in the gas phase will strike the surface
of the liquid and be recaptured
 Eventually the rate at which molecules return to the liquid will
exactly equal the rate at which they escape
 The number of molecules will reach a steady value, and the
pressure of the vapor at this stage becomes constant
 Called a dynamic equilibrium – evaporation and condensation
occur at equal rates
 Appears that nothing is happening!
 A great deal is happening
 Volatile – liquids that evaporate readily
o In an open container, vapor spreads away from the liquid, and there is
little chance to recapture it
o Equilibrium never occurs, and the vapor continues to form until the liquid
completely evaporate
o Substances with high vapor pressure evaporate more quickly and are
called volatile liquids
 Normal boiling point – the boiling point of a liquid at 1 atm pressure
o Bubbles of vapor form in the interior of the liquid
o Temperature of boiling increases with increasing temperature
o Maximum temperature of cooking food is boiling point of water
 Pressure cooker – causes water to boil at a higher temperature,
food cooks faster
 High altitudes – water boils at a lower temperature, so much cook
food for longer
 Clausius-Clapeyron Equation
ln P = -Hvap + C
RT
Graph ln P vs. 1/T, and Hvap = -slope x R
Phase Diagrams

Phase diagram – a graphic way to summarize the conditions under which
equilibria exist between the different states of matter
o 2D graph with pressure and temperature as the axes
o Has 3 curves – each represents the conditions of temperature and
pressure at which the phases can exist at equilbrium
 T-C = vapor-pressure curve of the liquid (eq. between liquid and
gas phases)
 Normal boiling point = T at 1 atm vapor pressure
 Ends at critical point (at critical pressure and temperature
of substance)
 Beyond this point – liquid and gas phases are
indistinguishable, called a supercritical fluid
 Solid-gas separator = change in vapor pressure of solid as it
sublimes at different temperatures
 Solid-liquid separator = change in melting point of solid with
increasing pressure
 Melting point = freezing point
*at 1 atm, this is the normal melting point
 Slopes right as pressure increases, b/c solid form is denser
than liquid form
 Increase in pressure favors solid phase (more compact
 Triple point = the temperature and pressure where all three
phases are in equilibrium
o Regions on graph = areas where that phase is stable
o H2O vs. CO2 phase diagrams
 Water is strange – melting point decreases with increasing pressure
(liquid form is more compact than solid!)
 CO2 does not have a normal boiling point – has a normal
SUBLIMATION point – makes dry ice a great coolant
Structures of Solids
 Crystalline solid – atoms, ions or molecules are ordered in well-defined 3D
arrangements
o Flat surfaces or faces that make definite angles with one another
o Examples: pyrite, fluorite, amethyst, quartz
 Amorphous solid – particles have no orderly structure
o Lack well-defined faces and shapes
o Mixtures of particles that do not stack together well
o Examples: rubber, glass (melt silicon dioxide and cool quickly)
Bonding in Solids

Molecular solids – atoms or molecules held together by intermolecular forces
(dipole-dipole, London dispersion, or hydrogen bonds)
o Relatively low melting points
o Soft
o Examples: Ar, H2O, CO2
o
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Benzene = symmetrical and planar, so can pack efficiently, so has higher
MP than toluene, which has a CH3 on the top (low symmetry prevents
efficient packing) BUT BP of toluene is higher than benzene, so attractive
forces are larger in liquid toluene than benzene (b/c it is heavier)
Covalent-network solids – atoms held together in large networks or chains by
covalent bonds
o Much harder, have higher MP
o Examples: diamond, graphite, silicon, germanium, quartz, silicon carbide,
boron nitride
o Graphite – good conductor of electricity, b/c of delocalized pi electrons
over the layers, sheets are held together w/ weak dispersion forces, so
they can slide past one another (lubricant)
Ionic solids – held together by ionic bonds
o Depends on charge of ions – higher charge = stronger attraction
o Can have different crystalline structures
Metallic solids – simple metals – entirely metal atoms, packed structures
o Bonding is too strong to be just London dispersion forces
o Strength is due to delocalized e- through solid (why they can conduct
heat and electricity)
 In general, strength of bond increases as # of e- available for
bonding increases
Packing of spheres
o Metallic solids
o Best way to maximize attractive forces between spheres – each sphere is
surrounded by 6 others in the layer
 A second layer is added above the first in the depressions
 Third layer on depressions in second
o Hexagonal close packing - Spheres in 3rd are in line with those of first –
ABAB
o Cubic close packing – Spheres in 3rd are not above 1st. 4th is above 1st –
ABCA
o Each sphere has 12 equidistant nearest neighbors – called coordination
number of 12
o Spheres occupy 74% of space, 26% empty space in both types of closepacking
o Other possibilities are large anion/small cation taking up the spaces