Chapter 12
• Average kinetic energy is related to the particles’ average speed, is proportional to the absolute temperature. We can use solids.
• Gas: low intermolecular force, little attraction between molecules. Particles are far apart. More kinetic energy causes an increase in volume or pressure as the particles move faster.
• Liquids: intermolecular forces are stronger, the particles are touching, and mobile.
Increased kinetic energy causes increased motion of particles (vs their neighbors) and increased vibration. Notice that there is a slight increase in volume (or…that’s why a mercury or alcohol thermometer works).
y
• Solids: Intermolecular forces are dominant.
Most particles don’t move. They remain in place and can vibrate, but not move.
Increased kinetic motion causes vibrations and some increase in volume.
• Intermolecular forces are the forces between atoms, molecules, and ions that hold solids and liquids together.
• Intermolecular forces also play a part in the solubilities of substances.
• Attraction of an ion to one side of a polar molecule.
• This interaction allows water to easily dissolve many ions.
• Attraction of the oppositely charged sides of polar
• May occur between molecules of the same or different compounds
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Dipole – induced dipole
• Occurs when a permanent dipole causes a non-polar species to polarize, forming a temporary dipole.
• This is a weak interaction.
• This phenomenon allows non-polar molecular to dissolve slightly in polar solvents (and vice-versa).
• The ease with which a particle’s electron cloud can be distorted.
• Smaller atoms are less polarized than larger atoms
• Polarizability increases as you go down a column and decreases left to right (big radius=more polarizable)
• Cations are less than atoms, anions are more than atoms.
Induced dipole – induced dipole:
London dispersion forces
• Dispersion forces occur between non-polar molecules.
• These forces occur when a temporary dipole forms by chance in one molecule and induces a dipole in a second.
• This is a weak interaction.
Note: greater MW=more
Deformation=more London
Bonding=higher boiling point:
• F
2 bp= -188.14 deg C
• Br
2 bp= 58.78 deg C
• I
2 bp= 184.35 deg C
Special Case: Hydrogen bonding
• A hydrogen bond is an especially strong dipole-dipole interaction. This is not a
“bond” like a covalent bond.
bound to N, O, or F (all small, very electronegative atoms) and has a dipole-dipole interaction with a highly electronegative atom that has a lone pair. (usually another N, O, or F atom in a compound) .
• Hydrogen bonding between molecules makes it a weak acid.
• Based on HCl, HBr, HI= strong acids, we
• Two other key points; it’s able to react with
SiO
2
• F- really wants to react with C 2+
2

Consequences of H-bonding
• Compounds (like water) have high boiling points, low vapor pressures, and strong cohesive and adhesive forces.
• Water: Surface tension
• Water: Capillary action
• High boiling point
Boiling Point of molecular hydrgen containing compounds.
H-bonding in DNA
Oil molecules can’t push their way between water molecules due to water-water hydrogen bonding
Intermolecular interactions and physical properties
•
Order of Strength:
Ion-dipole > Hydrogen Bonds >
Dipole-Dipole > Dispersion forces (induced
•
Effects of Intermolecular forces on Boiling
Point:
– Molar mass effects
– Intermolecular interactions
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Types of intermolecular forces:
(summary)
• Ionic cation-anion 400-4000 kj/m
• Covalent nuclei sharing electrons 150-
1100 kj/m
M t lli ti d l li d l t
1000 kj/m
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• Types of intermolecular forces:
• Ion-dipole 40-600kj/mole
• Like Na+ attraction to water.
• H-bonding alcohols)
10-40kj/mole (water,
• Dipole-dipole: 5-25kj/mole
• Ion-induced dipole 3-15
• Dipole-induced dipole 2-10
• Dispersion (London) 0.05-40
• We can easily relate the heat added to an object and its change in temperature if we know its specific heat and mass:
Δ
T
• However, when a compound changes state , there is no temperature change .
• For example, when ice (solid water) melts, it changes from solid H
2
0.0
° C.
O to liquid H
2
O all at
Phase Changes
• However, a lot of energy is required to break the intermolecular forces that hold the water molecules in their crystal structure.
• To measure this energy change, we employ the heat of fusion (
Δ
H ) for water – the amount of energy required to melt one mole of ice at 0 ° C.
• Likewise energy is required to vaporize water at
100 ° C. The amount of energy required to vaporize one mole of water is called the heat of vaporization
(
Δ
H vap
)
: q process
= n
• Δ
H process
, where n is moles.
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Phase Change Diagram (figure 6.10)
• WWII Aircraft Carrier:
• You are the new college graduate mechanical engineer assigned to the boiler room and your task is to produce the distilled water for the whole ship, using the least amount of oil to produce it.
• Cold salt water comes in, it’s heated to produce steam.
• Steam is cooled to produce water
• Energy given from cooling steam (and cooling the engines) heats up the cold salt water.
• The trick is to make the whole thing very efficient.
Specific
Heat Capacity
Solid
Liquid
J / g
•
C
°
2.1
4.184
Heat-related constants for
WATER
Phase
Changes
Δ H fus
Δ H vap kJ / mol
6.01
40.7
J / g
334
2260
Phase Changes Examples
15.What is the amount of energy required to vaporize 45.9 g of water at 100 ° C?
17.What is the energy required to change a
100.0 g ice chunk at –20.0
° C to steam at
100.0 ° C?
• So far, we have talked only about heat required to melt or vaporize a substance.
Th h Δ H l opposite signs:
Δ
H condensation
=
−Δ
H vaporization
Δ
H freezing
=
−Δ
H fusion ith
Calorimetry & Phase Changes
The q terms we discuss here can be used in 1 st
Law calorimetry problems: − q system
= q surroundings
Seven ice cubes having a combined mass of 185 g are taken from a freezer at -15.0
° C and are added to 591 mL (20 fl.oz.) of soda at 19.0 ° C in an insulated cup. When thermal equilibrium is reached the solution is at 0.0 ° C. Assuming no heat exchange with the cup or the rest of the environment, what mass of ice remains?
(Assume the density and heat capacity of the soda are the same as that of water.)
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Phase Changes
• Phase changes: go from one state to another state (involves an energy change, in order to overcome the intermolecular forces.
• Equilibrium: Below the boiling point, water will go from solid to gas or liquid to gas
We can measure the vapor pressure with a manometer. Liquid water phase changes to gas at less than it’s boiling point (why we speak of humidity). The vapor pressure is a function of temperature. At the boiling point, the vapor pressure is one atm.
Energy Required to overcome attraction; higher attraction=more energy
• If we take a liquid and overcome the intermolecular forces to produce a gas, intermolecular forces are broken or lessened for the melting to occur:
Vaporization, melting and sublimation requires energy = endothermic.
• (why sweating causes you to cool off, or pouring water over you=cooling)
• Liquid can go to gas at less than the boiling point. The boiling point is the point when the vapor pressure is 1 atm.
equilibrium, some liquid molecules obtain enough kinetic energy to leave the liquid phase, some gas molecules loose energy by dropping back to the liquid phase.
Vapor Pressure
• Molecules of a liquid are in dynamic equilibrium with molecules of gas of the same compound.
Vapor pressure increases with temperature.
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Boiling point
• The boiling point of a liquid is the temperature at which its vapor pressure equals the pressure of the atmosphere.
• “Normal” boiling point is the temperature when atmospheric pressure is 1.000 atm.
Vapor pressures of various compounds
Surface tension, capillary action, viscosity
• Intermolecular bonding happens between liquids and solids and within liquids;
• Surface tension: attraction of liquid to itself
• Capilary action: attraction of liquid to a solid
– Why we have a meniscus, why paper chromatography works
• Viscosity: attraction within a liquid
• Ever put a can of soda in the freezer?
Opps…. Solid water is less dense than liquid water. Most solids are more dense
• Result: Ice floats on water:
• Thermal moderation layer separating winter storms from the water under the ice layer: i t i d d th water near 0 deg C.
f f th
• Additional cold weather = thicker ice, while life goes on under the ice
• Also colder water is more dense=sinks, stirs up nutrients in water.
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Equilibrium nature of phase change:
• Ever hear about the biosphere? Or life on a space station. This is a completely enclosed system that only received sunlight from the outside world.
• Plants take the CO2 + H2O Æ carbohydrates.
P l th l t t t i i t CO2 d
H2O. Composting takes the excess plants and wastes and turns them into CO2 and H2O,
• The big things are growing enough food for the occupants and getting rid of the garbage.
• What must be delivered to the space station and what’s taken away?
∆
vap
• If you know the vapor pressure at two different temperatures, you can calculate the ∆ H vap pressure at the boiling point.
• Or if you know the ∆ H vap and the boiling point, you can calculate the vapor pressure at any other temperature.
(useful….)
• Remember that we looked up the vapor pressure of water in a table in the CRC? If you plot vapor pressure vs temperature, you get a curve. that:
• lnP = -( ∆ H vap compound.
/RT) + C C is a constant for each
• Rewrite:
• Ln (P
2
/P
1
) = ∆ H vap
/R [1/T
1
- 1/T
2
]
• If I know the vapor pressure at two temperatures,
I can determine the ∆ H vap
!
Phase Diagrams: Looking at solids, liquids and gases: Temperature vs pressure
• It is useful to plot temperature vs pressure and see where the element/compound is solid liquid or gas.
• Most elements/compounds show solids favored at higher pressure and liquids favored (vs solids) at higher temperature, and gases favored at higher temperatures still
• Triple point: temperature where solid, liquid gas exists
Phase Diagram of Water
• Hydrogen bonding causes water solid lattice to be ordered, larger than liquid.
• Very unusual; Net ice floats on water.
Wi t l k / id i l t d layer above the water=sealife can live under this layer.
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CO2 phase diagram
Critical point of water; The temperature where water becomes a supercritical fluid.
• If you put a gas under enough pressure, it will turn to liquid.
• The critical point is the temperature and become a critical fluid
• Metals (sea of electrons around metal ions)
• Ionic compounds (held together by
• Network molecular compounds (held together by shared electrons
We define compounds by how they pack together
• Depending on the size of the atoms/ions and their properties (ability to deform the electron clouds), they form different structures.
This is NaCl, in a cubic structure
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What determines the Lattice Forms?
• The ratio of anion size to cation size determines which lattice type is formed.
Metal crystal structu res
• Example: SiO2 is very common on the earth surface:
• Can get large quartz crystals, small
• Romans were able to use sand to make glass.
• What makes it do crystals, small crystals, agate?
• Quartz crystals can come in clear,
“smoky”, purple, yellow/tan
• Polycrystaline comes in more colors l AND th can be dyed. i
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• Ingredients (Roman Era)
• Naturally colored glass is composed of 72% silica, 15% soda, 10% lime, and 3% other impurities. Beach sand, the main ingredient in glass, contains silica and lime, the latter component coming from crushed shells. There are two steps to producing naturally colored glass:
(1) melting the ingredients at a relatively low temperature substance ( frit ) that could easily be stored for later use;
(2) re-melting the frit in a far hotter furnace (circa 1150 degrees C), thereby transforming it into fluid glass.
• http://www.museum.upenn.edu/new/research/Roman%2
0Glass/Glassmaking/glassmaking_4.html
• Ingredients (Roman Era)
• Naturally colored glass is composed of 72% silica, 15% soda, 10% lime, and 3% other impurities. Beach sand, the main ingredient in glass, contains silica and lime, the latter component coming from crushed shells. There are two steps to producing naturally colored glass:
(1) melting the ingredients at a relatively low temperature substance ( frit ) that could easily be stored for later use;
(2) re-melting the frit in a far hotter furnace (circa 1150 degrees C), thereby transforming it into fluid glass.
• http://www.museum.upenn.edu/new/research/Roman%2
0Glass/Glassmaking/glassmaking_4.html
Solids don’t always form crystals; glass is amorphous.
It is hard, as a solid, but lacks long range order.
Amorphous “solids” do not have a definite melting point, crystals do.
• A pure element/compound has a constant melting point
• A mixture has a higher boiling point than a point)
• A glass does not have a defined melting point. As the temperature increases, it
“softens”. Viscosity decreases.
– Glass can be called a solid liquid
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