Chapter 10
Particle Forces
States of Matter
Solid- Particles moving about a fixed point
Liquid-Particles moving about a moving point
Gas-Particles filling the volume of the container with complete
random motions.
Particle Forces Affect
• Solubility
• Vapor Pressures
• Freezing Points
• Boiling Points
Particle Forces
• Intramolecular forces (Relative strength = 100)
 Ionic bonding
 Covalent bonding
• Interparticle forces
 Ion-dipole forces
 Dipole-dipole (Polar molecules)
(relative Strength = 1)
 London Forces (Dispersion forces)( Nonpolar
molecules)
(relative strength = 1)
 Hydrogen Bonding (Relative strength = 10)
Ion-Ion Interactions
• Coulomb’s law states that the energy (E) of
the interaction between two ions is directly
proportional to the product of the charges of
the two ions (Q1 and Q2) and inversely
proportional to the distance (d) between
them.
E 
(Q 1Q 2 )
d
Predicting Forces of Attraction
• Coulombs Law indicates the increases in the
charges of ions will cause an increase in the
force of attraction between a cation and an
anion.
• Increases in the distance between ions will
decrease the force of attraction between
them.
Size of Ions
Lattice Energy
• The lattice energy (U) of an ionic compound
is the energy released when one mole of the
ionic compound forms from its free ions in the
gas phase.
M+(g) + X-(g) ---> MX(s)
k(Q 1Q 2 )
U=
d
Comparing Lattice Energies
Lattice Energies of Common
Ionic Compounds
Compound
U(kJ/mol)
LiF
-1047
LiCl
-864
NaCl
-790
KCl
-720
KBr
-691
MgCl2
-2540
MgO
-3791
Practice
Determine which salt has the greater lattice
energy.
A. MgO and NaF
B. MgO and MgS
Lattice Energy Using Hess’s Law
Electron Affinity
• Electron affinity is the energy change
occurring when one mole of electrons
combines with one mole of atoms or ion in
the gas phase.
• Step 4 in diagram on the last slide.
Cl(g) + e-(g) ---> Cl-(g)
ΔHEa = -349 kj/mole
Calculating U
Na+(g) + e-(g) ---> Na(g)
-HIE1
Na(g) ---> Na(s)
-Hsub
Cl-(g) ---> Cl(g) + e-(g)
-HEA
Cl(g) ---> 1/2Cl2(g) -1/2HBE
Na(s) + 1/2Cl2(g) ---> NaCl(s) Hf
Na+(g) + Cl-(g) ---> NaCl(s)
U
U = Hf - 1/2HBE - HEA - Hsub - HIE1
Lattice energy for NaCl.
Interactions Involving Polar Molecules
• An ion-dipole interaction occurs between an
ion and the partial charge of a molecule with
a permanent dipole.
• The cluster of water molecules that surround
an ion in aqueous medium is a sphere of
hydration.
Illustrates of Ion-Dipole Interaction
The Solution Process
Bond Breaking Processes
•
•
Break solute particle forces (expanding
the solute), endothermic
Break solvent particle forces (expanding
the solvent), endothermic
The Solution Process
Attractive Forces
•
Energy released when solute solvent are
attracted, exothermic
Energy is released due to new attractions
•



Ion dipole if the solute is ionic and the solvent
polar.
London-Dipole for nonpolar solute and polar
solvent
Dipole-dipole for polar solute and polar solvent
The Solution Process
Theromodynamics
• Enthalpy
• Entropy (Perfect crystal, assumed to be
zero)
• Gibbs free energy
The Solution Process
Oil dissolving in water
• London forces holding the oil molecules
together are large do to the large surface area
of the oil
• The hydrogen bonds holding water molecules
together are large
• The forces of attraction of between nonpolar
oil and polar water are weak at best
• Thus the overall process is highly endothermic
and not allowed thermo chemically
The Solution Process
Oil dissolving in water
• Entropy should be greater than zero
• Free energy should be greater than zero, since
the process is highly endothermic
• Thus the overall process is nonspontaneous
The Solution Process
Sodium chloride dissolving in water
• Large amount of energy is required to break the ionic
lattice of the sodium chloride (expand solute)
• Large amount of energy is required to separate the
water molecules to expand the solvent breaking
hydrogen bonds
• Formation of the ion dipole forces releases a large
amount of energy, strong forces (why?)
• The sum of the enthalpies is about +6 kJ (slightly
endothermic), which is easily overcome by the
entropy of the solution formation.
Water as a Solvent
• Water most important solvent, important
to understand its solvent properties
• Most of the unusual solvent properties of
water stem from it hydrogen bonding
nature
• Consider the following ∆S of solution
KCl →75j/K-mole
LiF→-36j/K-mole
CaS→-138 j/K-mole
Water as a Solvent
• We would expect ∆S>0 for all solutions,
right?
• But two are negative, why?
• Obviously, something must be happening
for the increased order.
• Ion-dipole forces are ordering the water
molecules around the ions, thus causing
more order in water i.e. less positions for
water than in the pure liquid state
Water as a Solvent
• Smaller ions, have stronger ion dipole forces,
thus pulling water closer, therefore less
positions
• Also, ions with a charge greater than one will
attract to water stronger than a one plus
charge, thus more order due to less space
between particles
Dipole-Dipole Interactions
• Dipole-dipole interactions are
attractive forces between
polar molecules.
• An example is the interaction
between water molecules.
• The hydrogen bond is a
special class of dipole-dipole
interactions due to its
strength.
Dipole-Dipole Forces
Dipole-dipole (Polar molecules)
Alignment of polar molecules to two electrodes
charged + and δ–
Forces compared to ionic/covalent are about 1 in strength
compared to a scale of 100, thus 1%
δ+ δ–
H Cl
δ+
δ–
H Cl
δ+
δ–
H Cl
Dipole Dipole Interactions
Slide 28 of 35
Hydrogen Bonding
• Hydrogen bonding a stronger intermolecular
force involving hydrogen and usually N, O, F,
and sometimes Cl
–Stronger that dipole-dipole, about 10 out of
100, or 10
–Hydrogen needs to be directly bonded to the
heteroatom
–Since hydrogen is small it can get close to the
heteroatom
–Also, the second factor is the great polarity of
the bond.
Hydrogen Bonding in HF(g)
Slide 30
Hydrogen Bonding in Water
around a molecule
in the solid
in the liquid
Slide 31
Boiling Points of Binary Hydrides
Interacting Nonpolar Molecules
• Dispersion forces (London forces) are
intermolecular forces caused by the
presence of temporary dipoles in
molecules.
• A instantaneous dipole (or induced
dipole) is a separation of charge
produced in an atom or molecule by a
momentary uneven distribution of
electrons.
Illustrations
Strength of Dispersion Forces
• The strength of dispersion forces depends
on the polarizability of the atoms or molecules
involved.
• Poarizability is a term that describes the
relative ease with which an electron cloud is
distorted by an external charge.
• Larger atoms or molecules are generally
more polarizable than small atoms or
molecules.
London Forces (Dispersion)
• Induced dipoles (Instantaneous )
• Strength is surface area dependent
• More significant in larger molecules
• All molecules show dispersion forces
• Larger molecules are more polarizable
Instantaneous and Induced Dipoles
Slide 37
Molar Mass and Boiling Point
Molar Mass and Boiling Points of Common Species.
Halogen
M(g/mol)
Bp(K)
Noble Gas
M(g/mol)
Bp(K)
He
2
4
F2
38
85
Ne
20
27
Cl2
71
239
Ar
40
87
Br2
160
332
Kr
84
120
I2
254
457
Xe
131
165
Rn
211
211
London vs Hydrogen Bonding
Hydrocarbon
Molecular
Formula
CH4
Molar
Mass
Alcohol
Bp
(oC)
Molecular
Formula
Molar
Mass
Bp
(oC)
16.04 -161.5
CH3CH3
30.07
-88
CH3OH
32.04
64.5
CH3CH2CH3
44.09
-42
CH3CH2OH
46.07
78.5
CH3CH(CH)CH3 58.12
-11.7
CH3CH(OH)CH3
60.09
82
CH3CH2CH2CH3 58.12
-0.5
CH3CH2CH2OH
60.09
97
The Effect of Shape on Forces
Practice
Rank the following compound in order of increasing
boiling point. CH3OH, CH3CH2CH2CH3, and
CH3CH2OCH3
Practice
Rank the following compound in order of increasing
boiling point. CH3OH, CH3CH2CH2CH3, and
CH3CH2OCH3
MM
32.0
IM Forces
London and H-bonding
CH3CH2CH2CH3
58.0
London, only
CH3CH2OCH3
60.0
London and Dipole-dipole
CH3OH
Practice
Rank the following compound in order of increasing
boiling point. CH3OH, CH3CH2CH2CH3, and
CH3CH2OCH3
MM
32.0
IM Forces
London and H-bonding
CH3CH2CH2CH3
58.0
London, only
CH3CH2OCH3
58.0
London and Dipole-dipole
CH3OH
The order is:
CH3CH2CH2CH3 < CH3CH2OCH3< CH3OH
Polarity and Solubility
• If two or more liquids are miscible, they form
a homogeneous solution when mixed in any
proportion.
• Ionic materials are more soluble in polar
solvents then in nonpolar solvents.
• Nonpolar materials are soluble in nonpolar
solvents.
• Like dissolves like
Polarity and Solubility
• If two or more liquids are miscible, they
form a homogeneous solution when
mixed in any proportion.
• Ionic materials are more soluble in polar
solvents then in nonpolar solvents.
• Nonpolar materials are soluble in
nonpolar solvents.
Polarity and Solubility
How does polarity effect solubility?
The thermodynamic argument, is that the
lower the potential energy, the more stable the
system. If subtracting the potential energy of
the solute from the potential energy of the
original solute and solvent is negative
(exothermic) then solution is
thermodynamically favored.
Polarity and Solubility
How does polarity effect solubility?
Non polar solute and solvent: The forces holding
these particles together are London Dispersion
forces, the weakest of all of the inter-particle forces.
The strength of these forces are relative to the
surface area if solute and solvent are of similar size,
then about the same amount of energy is required to
separate solute and solvent particles from each
other. And about the same amount of energy is
released when solute and solvent are attracted to
each other forming a solution. Thus we predict non
polar solutes and solvents should dissolve
Polarity and Solubility
How does polarity effect solubility?
Non polar solute and polar solvent: Considering
solutes and solvents of similar surface area it should
be noted that more energy is required to separate the
polar solvent molecules from each other, since dipoledipole interactions are stronger. The only interaction
between a nonpolar solute and polar solvent would be
London Dispersion forces, so the energy released is
much less than required for separating the solvent and
solute. Subtracting the potential energy of the
products from reactants would give a positive
(endothermic) result and the solution would be less
stable than the dissolution.
Practice
Rank the following compound in order
of increasing boiling point. CH3OH,
CH3CH2CH2CH3, and CH3CH2OCH3
Solubility of Gases in Water
• Henry’s Law states that the solubility of a
sparingly soluble chemically unreactive gas in
a liquid is proportional to the partial pressure
of the gas.
• Cgas = kHPgas where C is the concentration of
the gas, kH is Henry’s Law constant for the
gas.
Henry’s Law Constants
Henry’s Law Constants
Gas
kH[mol/(L•atm)]
kH[mol/(kg•mmHg)]
He
3.5 x 10-4
5.1 x 10-7
O2
1.3 x 10-3
1.9 x 10-6
N2
6.7 x 10-4
9.7 x 10-7
CO2
3.5 x 10-2
5.1 x 10-5
Terms
• A hydrophobic (“water-fearing) interaction
repels water and diminishes water solubility.
• A hydrophilic (“water-loving”) interaction
attracts water and promotes water solubility.
Affects of Intermolecular Forces
• Solubility
• Vapor Pressures
• Freezing Points
• Boiling Points
• Surface tension
Vapor Pressure
• Vaporization or
evaporation is the
transformation of
molecules in the liquid
phase to the gas phase.
• Vapor pressure is the
force exerted at a given
temperature by a vapor
in equilibrium with its
liquid phase.
Vapor Pressure
Vapor Pressure
The normal boiling point of a liquid is the temperature
at which its vapor pressure equals 1 atmosphere.
Vapor Pressure of Solutions
What evaporates faster, sugar water or pure water?
swswsw
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Vapor Pressure of Solutions
What evaporates faster, sugar water or pure water?
w(g)
swswsw
w(g)
w(g)
wwwwww
Vapor Pressure of Solutions
What evaporates faster, sugar water or pure water?
w(g)
swswsw
w(g)
w(g)
wwwwww
Pure water evaporates faster, since there are more water
particles on the surface, thus lowering the average kinetic
energy. Evaporation of a solution is inversely proportional to
concentration.
Vapor Pressure of Solutions
• Raoult’s Law
 Psolution = Xsolvent (Psolvent)
• P - vapor pressure
• X - mole fraction
• Xsolute + Xsolvent = 1
Practice
A solution contains 100.0 mL of water and 0.500 mol of
ethanol. What is the mole fraction of water and the vapor
pressure of the solution at 25oC, if the vapor of pressure of
pure water is 23.8 torr?
Surface Tension
Cohesive and Adhesive Forces
Produce a Meniscus
Physical State and Phase
Transformations
• A phase diagram is a graphic
representation of the dependence of the
stabilities of the physical states of a
substance on temperature and
pressure.
Phase Diagram for Water
• Triple Point
• Critical Point
• Critical Temperature
• Critical Pressure
• Supercritical Fluid
Terms
• The triple point defines the temperature and
pressure where all three phases of a
substance coexist.
• The critical point is that specific temperature
and pressure at which the liquid and gas
phases of a substance have the same density
and are indistinguishable for each other.
• A supercritical fluid is a substance at
conditions above its critical temperature and
pressure.
Phase Diagram for CO2
Terms
• Capillary action is the rise of a liquid up
a narrow tube as a result of adhesive
forces between the liquid and the tube
and cohesive forces within the liquid.
• Viscosity is a measure of the resistance
to flow of a fluid.
Colligative Properties of Solutions
• Colligative properties
of solutions depend
on the concentration
and not the identity of
particles dissolved in
the solvent.
• Sea water boils at a
higher temperature
than pure water.
Colligative Properties
Colligative property is a physical property that
depends on the number of particles present, and not
on the nature of the particle. Since evaporation is
dependent on the number of solvent particles
present on the surface that makes evaporation and
vapor pressure colligative properties.
Colligative Properties
Is density a colligative property?
Colligative Properties
Is density a colligative property?
While density depends on the number of particles in
a given area, it is also effected by the weight of the
substance, which is a nature thing, so no density is
not a colligative property.
Calculating Changes in Boiling Point
 Tb = Kbm
 Tb is the increase in
Bp
 Kb is the boiling-point
elevation constant
 m is a new
concentration unit
called molality
n
m=
solute
kg solvent
Practice
Calculate the molality of a solution
containing 0.875 mol of glucose
(C6H12O6) in 1.5 kg of water.
Practice
Seawater contains 0.558 M Cl- at the
surface at 25oC. If the density of sea
water is 1.022 g/mL, what is the molality
of Cl- in sea water?
Practice
Cinnamon owes its flavor and odor to
cinnamaldehyde (C9H8O). Determine
the boiling-point elevation of a solution
of 100 mg of cinnamaldehyde dissolved
in 1.00 g of carbon tetrachloride (Kb =
2.34oC/m).
Freezing-point Depression
• Tf = Kfm
 Kf is the freezing-point
depression constant
and m is the molality.
Practice
The freezing point of a solution
prepared by dissolving 1.50 X 102 mg
of caffeine in 10.0 g of camphor is 3.07
Celsius degree lower than that of pure
camphor (Kf = 39.7oC/m). What is the
molar mass of caffeine?
The van’t Hoff Factor
• Tb = iKbm & Tf = iKfm

• van’t Hoff factor, i is the
number of ions in one
formula unit
Values of van’t Hoff Factors
Practice
CaCl2 is widely used to melt frozen
precipitation on sidewalks after a winter
storm. Could CaCl2 melt ice at -20oC?
Assume that the solubility of CaCl2 at
this temperature is 70.0 g/100.0 g of
H2O and that the van’t Hoff factor for a
saturated solution of CaCl2 is 2.5 (Kf for
water is 1.86 0C/m).
Osmosis
In osmosis, solvent passes through a semipermeable membrane
to balance the concentration of solutes in solution on both sides
of the membrane.
Figure 10.30
Osmosis at the Molecular Level
Osmotic Pressure
• Osmotic pressure () is the pressure that has
to be applied across a semipermeable
membrane to stop the flow of solvent form the
the compartment containing pure solvent or a
less concentrated solution towards a more
concentrated solution.
  = iMRT where i is the van’t Hoff factor, M is
molarity of solute, R is the idea gas constant
(0.00821 l•atm/(mol•K)), and T is in Kelvin
ChemTour: Lattice Energy
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PC | Mac
Students learn to apply Coulomb’s law to calculate the
exact lattice energies of ionic solids. Includes Practice
Exercises.
ChemTour: Intermolecular
Forces
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This ChemTour explores the different types of
intermolecular forces and explains how these affect the
boiling point, melting point, solubility, and miscibility of a
substance. Includes Practice Exercises.
ChemTour: Henry’s Law
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Students learn to apply Henry’s law and calculate the
concentration of a gas in solution under varying conditions
of temperature and pressure. Includes interactive practice
exercises.
ChemTour: Molecular Motion
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Students use an interactive graph to explore the relationship
between kinetic energy and temperature. Includes Practice
Exercises.
ChemTour: Raoult’s Law
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Students explore the connection between the vapor
pressure of a solution and its concentration as a gas above
the solution. Includes Practice Exercises.
ChemTour: Phase Diagrams
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Students use an interactive phase diagram and animated
heating curve to explore how changes in temperature and
pressure affect the physical state of a substance.
ChemTour: Capillary Action
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In this ChemTour, students learn that certain liquids will be
drawn up a surface if the adhesive forces between the liquid
on the surface of the tube exceed the cohesive forces
between the liquid molecules.
ChemTour: Boiling and
Freezing Points
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Students learn about colligative properties by exploring the
relationship between solute concentration and the
temperature at which a solution will undergo phase
changes. Interactive exercises invite students to practice
calculating the boiling and freezing points of different
solutions.
ChemTour: Osmotic Pressure
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Students discover how a solute can build up pressure
behind a semipermeable membrane. This tutorial also
discusses the osmotic pressure equation and the van’t Hoff
factor.
Solubility of CH4, CH2Cl2, and
CCl4
Which of the following three
compounds is most soluble in
water?
A) CH4(g)
B) CH2Cl2(λ)
C) CCl4(λ)
Solubility of CH4, CH2Cl2,
and CCl4
Consider the following arguments for each answer
and vote again:
A. A gas is inherently easier to dissolve in a liquid than
is another liquid, since its density is much lower.
B. The polar molecule CH2Cl2 can form stabilizing
dipole-dipole interactions with the water molecules,
corresponding to a decrease in ΔH°soln.
C. The nonpolar molecule CCl4 has the largest
molecular mass, and so is most likely to partially
disperse into the water, corresponding to an increase
The End