CHAPTER 10 AND 11
AP CHEMISTRY
INTERMOLECULAR FORCES
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Molecules tend to be
Nonconductors of electricity when pure
substances
Most molecules when placed in water won’t
conduct
– EXCEPTION- highly polar molecules i.e. HCl
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Insoluble in water but soluble in nonpolar
solvents (iodine dissolves in alcohol) i.e. toluene,
CCl4
A few are very soluble in water i.e. ethyl
alcohol(C2H5OH)
Low melting and boiling points
Many are gases at room temperature
Many boiling points are directly related to the
strength of the intermolecular forces
DIPOLE FORCES
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Ion-dipole
– Ions attracted to a dipole substance i.e. NaCl in water
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Dipole-dipole
What are dipole forces?
– Attraction between polar molecules
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As the dipole moment increases the boiling point
increases
London dispersion forces (Van der Waal)
What is a dispersion force?
– Electrical temporary dipoles (POLARIZABILITY)
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All molecules have dispersion forces, the strength
depend on two factors
Number of electrons in the molecule
Ease with which electrons are dispersed to form
temporary dipoles
As the molar mass increases the strength will
increase
HYDROGEN BONDS
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Polarity has a small effect on boiling point
Compounds have to have hydrogen attached to
nitrogen, oxygen or fluorine
Hydrogen in one compound is attracted to a lone
pair on the N, O, or F of a different compound
Two reasons hydrogen bonding is stronger than
other dipole forces
Electronegative difference is large between
hydrogen(2.2) and nitrogen(3.0), oxygen(3.5), or
fluorine(4.0)
Small size of hydrogen allows the unpaired
electron pair of F, O, or N atom to approach the H
atom closely
The molar mass of H2S is double that of water,
both are polar molecules, but the boiling point of
water is 167 °C higher than H2S. Why?
PROPERTIES OF LIQUIDS
 Viscosity
–Resistance to flow
 Surface tension
–Energy required to increase surface
area of liquid by a unit amount
–Water increases surface tension
because of hydrogen bonding
PHASE CHANGES
 Vapor
pressure Pvapor = Patm - PHg
column page 460-461
 Heat of fusion
– Enthalpy change associated with
melting
 Ice
= 6.01 kJ/mol
 Heat of vaporization
– Heat required to vaporize a liquid
 Water
= 40.67 kJ/mol
CRITICAL TEMPERATURE
Highest temperature at which a substance
can exist as a liquid
 Critical pressure
– Pressure required to bring about
liquidefaction
 Sublimation
– When a substance goes from the solid
phase to the gaseous phase without
going through the liquid phase
 Deposition
– When a substance goes from the
gaseous phase to the solid phase
without going through the liquid phase
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VAPOR PRESSURE
 Dynamic
equilibrium
– Opposing processes are occurring at the
same rate
– Evaporation and condensation
 Volatile
– Liquids with high vapor pressure
evaporate quickly
 Phase
diagram
– On the overhead
STRUCTURE OF SOLIDS
 Crystalline
solid
– Geometry shaped to their arrangement
of ions
 Amorphous
solid
– No orderly arrangement
– They do not melt at a specific
temperature
– Glass
 Crystal
Lattice
 Unit cell repeated, page 432
CONTINUE
Close packing of ions cause three types of
unit cells
 Simple
Body centered

Face centered
Page 447
X-RAY DIFFRACTION
Scattering of x-rays by a regular
arrangement of atoms or ions
 Constructive interference
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– Higher intensity if waves are the same
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Destructive interference
– Cancels each other out
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Bonding in solids
– Table 10.3 page 436
– Graphite and diamonds are allotropes of
carbon (same number of carbons (60) but in
different forms
– You can convert graphite into diamonds
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Read pages 440 to 470 take notes
SOLUTION COMPOSTION
Molarity = moles of solute
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L of solution
 Percent mass = mass of solute
X 100
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Mass of solution
 Mole fraction = A = XA =
nA
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nA + nB
 nA = moles of solute
 nB = moles of solution or other substance
 Molality = mole of solute

kg of solvent
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NORMALITY
 Equivalents/L
solution
 Acid-base
– Equivalent mass of acid or base that can
get 1 mole of protons
– Equivalent mass of sulfuric acid is ½ it’s
molar mass because you get 2 protons
 Oxidation-reduction
– Equivalent amount of oxidizing or
reducing agent to accept or furnish 1
mole of electrons
– KMnO4
ENERGY SOLUTION FORMATION
Enthalpy Hsol’n = H1 + H2 + H3
 Fig 11.2 page 490 and table 11.3 page 492
 Read and take notes 492-497
 Vapor pressure
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– Nonvolatile solute lowers the vapor pressure of a
solvent
Raoult’s law
 Psol’n =Xsolvent·Posolvent
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– Po = vapor pressure of solvent
– X = mole fraction of solvent
CONTINUE
 Ptotal
= XAPoA + XBPoB
 Liquid-liquid solution that obey
Raoult’s law is an ideal solution
 Table 11.4 page 503
COLLIGATIVE PROPERTIES
 Nonvolatile
solute raises the boiling
point and lowers the freezing point
 ΔTf = kfm
m = mol solute/kg solvent
 ΔTb = kbm ΔT = change in
temperature
 kf or kb = boiling point or freezing pt.
constant
OSMOTIC PRESSURE
 Semipermeable
membrane
– Solvents can pass through but not the
solute molecule
 The
volume of solution increases and
the solvent decreases, this flow is
called OSMOSIS
 Osmotic pressure
– When you have an excess pressure on
the solution in comparison to the pure
solvent
CONTINUE

Reverse osmosis
– If an external pressure greater than the
osmotic pressure is applied the solvent
particles will go from the solution to the pure
solvent
Colligative properties of electrolytes
 Van’t Hoff factor i:
 i = moles of particles in solution
Moles of solute dissolved
 Colloids
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– Suspension of tiny particles in some medium
– Blood
– Coagulation - destroy a colloid - heating or
adding an electrolyte (ions)