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AP Chemistry
Chapter 11 Outline
11.1 A Molecular Comparison of Gases, Liquids, and Solids
11.1.1 Gases = widely separated particles in constant chaotic motion
11.1.1.1
Average attractions between molecules are much smaller than their average
kinetic energy
11.1.1.2
Gases expand to fill their container because of the lack of strong attractions
between particles
11.1.2 Liquids = intermolecular attractive forces are strong enough to keep particles close
together
11.1.2.1
Liquids are denser, far less compressible than gases
11.1.2.2
Attractive forces are strong enough to keep the particles from moving past each
other
11.1.2.3
Liquids can be poured and take the shape of their container.
11.1.3 Solids = particles are virtually locked in place because of strong intermolecular
attractions
11.1.3.1
Solids are not very compressible because particles have very little space between
them
11.1.3.2
Solids are rigid because the particles are not free to undergo long-range
movement
11.1.3.2.1 Crystalline solids = contain highly ordered structures, with a regular pattern
11.1.3.2.2 Amorphous solids = structures are not completely regular
11.1.2 Condensed phases = solids and liquids
11.1.3 The state of matter depends on two factors: average kinetic energy, and interparticle
energies of attraction
11.1.3.1
State of matter can be changed by heating or cooling, which changes the average
kinetic energy of the particles
11.2 Intermolecular Forces
11.2.1 These attractions vary in strength, but tend to be much weaker than ionic or covalent
bonds.
11.2.1.1
They take less energy to “break”, because the molecules themselves stay intact
11.2.1.2
Many properties of liquids reflect the strengths of the intermolecular attractions
11.2.2 Ion-Dipole Attractions
11.2.2.1
Exists between an ion and the partial charge on the end of a polar molecule.
11.2.2.2
The magnitude of the attraction increases as either the charge of the ion or the
magnitude of the dipole moment increases.
11.2.2.3
Important for solution process of ionic compounds in polar solvents (more in
chapter 13)
11.2.3 Dipole-Dipole Attractions
11.2.3.1
Neutral polar molecules attract each other when the positive end of one molecule
is near the negative end of another.
11.2.3.2
Through random motions, the molecules arrange themselves to maximize
attractions and minimize repulsions.
11.2.3.3
These are only significant when the molecules are close together.
11.2.3.4
For molecules of approximately equal mass and size, the strengths of the
intermolecular attractions increasing with increasing polarity.
11.2.4 London Dispersion Forces
11.2.4.1
All molecules, even non-polar, symmetric atoms, can create an instantaneous
dipole moment due to uneven electron distribution.
11.2.4.2
This instantaneous, temporary dipole can induce an adjacent atom to have a
temporary dipole, causing the two particles to be attracted to one another.
11.2.4.3
All substances exhibit dispersion forces.
11.2.4.4
Significant only when molecules are very close together.
11.2.4.5
Depends on polarizability (how easily the electron cloud is distorted)
11.2.4.5.1 Long, thin molecules are more polarizable than compact molecules, because of
increased surface area for contact between molecules.
11.2.4.5.2 Larger molecules are more polarizable than smaller molecules, because they have
more electrons, which are relatively far from the nuclei.
11.2.5 Hydrogen Bonding
11.2.5.1
A special type of dipole-dipole attraction between the hydrogen atom in a polar
bond (particularly to F, N or O) and an unshared electron pair on a nearby small
electronegative ion or atom (usually an F, O, or N atom on another molecule).
11.2.5.1.1 Because of the large electronegativity difference in these bonds, the hydrogen nucleus
is nearly exposed.
11.2.5.1.2 The small, electron-poor hydrogen can approach an electronegative atom very closely
and interact with it.
11.2.5.2
The energies of hydrogen bonds are much weaker than ordinary covalent bonds,
but stronger than other intermolecular attractions.
11.2.5.2.1 Key roles in proteins, DNA structure
Figure 11.12 is a useful flowchart for identifying the types of intermolecular attractions in a
particular system.
11.3 Some Properties of Liquids
11.3.1 Viscosity = resistance of a liquid to flow
11.3.1.1
Related to the ease with which individual molecules of the liquid can move past
each other
11.3.2 Surface Tension = the energy required to increase the surface area of a liquid by a unit
amount; a measure of the inward forces that must be overcome
11.3.2.1
Cohesive forces = intermolecular forces that bind similar molecules to one
another
11.3.2.2
Adhesive forces = intermolecular attractions that bind a substance to a surface
11.3.2.2.1 Meniscus
11.3.2.2.2 Capillary action
11.4 Phase Changes
11.4.1 Every phase change is accompanied by a change in the energy of the system.
11.4.1.1
Heat of fusion = the energy required to bring about a change from highly ordered
solid to a liquid (i.e., melting) = -heat of freezing
11.4.1.2
Heat of vaporization = the energy required to bring about a change from the
condensed liquid state to highly separated particles in the gas state = - heat of deposition
11.4.1.3
Heat of vaporization is generally several times bigger than the heat of fusion
11.4.1.3.1 In order to vaporize, the molecules must essentially sever all intermolecular
attractions
11.4.1.3.2 In melting, most of the intermolecular attractions remain unchanged
11.4.1.4
Heat of sublimation = the energy required to change a solid into a gas, without
going through the liquid state = - heat of deposition
11.4.2 Heating Curves
11.4.2.1
Phase changes occur without a change in temperature! To calculate energy, use
H = Hfus/vap * mass of substance
11.4.3 Critical temperature = the highest temperature at which a distinct liquid phase can form;
highest temperature at which a liquid can exist
11.4.3.1
Critical pressure = the pressure required to bring about liquefaction at the critical
temperature
11.5 Vapor Pressure
11.5.1 Molecules can escape from the surface of a liquid into the gas phase; will reach a state of
dynamic equilibrium
11.5.2 Vapor pressure = the pressure exerted by the vapor when at equilibrium
11.5.3 Volatile liquid = a liquid that evaporates readily
11.5.4 Boiling point = the temperature at which the vapor pressure equals the external pressure
11.5.4.1
Normal boiling point = the boiling point at 1 atm
11.6 Phase Diagrams
11.6.1 Graphical way to summarize conditions under which equilibria exist between different
states of matter
11.6.1.1
Determined for closed systems
11.6.2 Triple point = conditions at which solid, liquid and gas phases coexist
11.6.3 Critical point = critical temperature and pressure; beyond this point, gas and liquid are
indistinguishable
11.6.4 From triple point to critical point = vapor pressure curve; equilibrium between liquid and
gas phases
11.6.5 Equilibrium between solid and liquid; line gives melting points at different pressures
11.6.6 Equilibrium between solid and gas;