Ch 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.
(Q1Q2 )
E 
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
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 ΔS (Perfect crystal, assumed to be
zero)
• Gibbs free energy ΔG
•
•
•
•
ΔG = ΔH - T ΔS
ΔG < 0, spontaneous change
ΔG = 0, equilibrium
ΔG > 0, nonspontaneous
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 21 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 23
Hydrogen Bonding in Water
around a molecule
in the solid
in the liquid
Slide 24
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 30
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.
Phase Diagram of Water
Note: Negative
slope
Phase Diagram of Water
Note:
Posititve Slope
Phase Diagram CO2
Phase Diagram CO2
Special Properties of Water
Hydrogen bonding found in water gives it
special properties listed below:
• Surface tension
• Capillary Rise
• Viscosity
Surface Tension
Surface Tension
Surface Tension units J/m2
Surface Tension
Adhesive forces
stronger in red
aqueous solution
Cohesive forces
stronger in
mercury liquid
Capillary Rise
Viscosity
Cohesive and Adhesive Forces
Produce a Meniscus
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.
ChemTour: Lattice Energy
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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