Roy%20Kennedy

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Introductory Chemistry, 3rd EditionNivaldo Tro
Chapter 12 Liquids, Solids, and Intermolecular
Forces

• Roy Kennedy
• Massachusetts Bay Community College
• Wellesley Hills, MA

2009, Prentice Hall
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Interactions Between Molecules

• Many of the phenomena we observe are related to
  interactions between molecules that do not
  involve a chemical reaction.
• Your taste and smell organs work because
  molecules in the thing you are sensing interact
  with the receptor molecule sites in your tongue
  and nose.
• In this chapter, we examine the physical
  interactions between molecules and the factors
  that effect and influence them.

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The Physical States of Matter

• Matter can be classified as solid, liquid, or gas
  based on what properties it exhibits.

• Fixed Keeps shape when placed in a container.
• Indefinite Takes the shape of the container.

4
Structure Determines Properties

• The atoms or molecules have different structures
  in solids, liquids, and gases, leading to
  different properties.

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Properties of the States of MatterGases

• Low densities compared to solids and liquids.
• Fluid.
• The material exhibits a smooth, continuous flow
  as it moves.
• Take the shape of their container(s).
• Expand to fill their container(s).
• Can be compressed into a smaller volume.

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Properties of the States of MatterLiquids

• High densities compared to gases.
• Fluid.
• The material exhibits a smooth, continuous flow
  as it moves.
• Take the shape of their container(s).
• Keep their volume, do not expand to fill their
  container(s).
• Cannot be compressed into a smaller volume.

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Properties of the States of MatterSolids

• High densities compared to gases.
• Nonfluid.
• They move as an entire block rather than a
  smooth, continuous flow.
• Keep their own shape, do not take the shape of
  their container(s).
• Keep their own volume, do not expand to fill
  their container(s).
• Cannot be compressed into a smaller volume.

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The Structure of Solids, Liquids, and Gases
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Gases

• In the gas state, the particles have complete
  freedom from each other.
• The particles are constantly flying around,
  bumping into each other and their container(s).
• In the gas state, there is a lot of empty space
  between the particles.
• On average.

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Gases, Continued

• Because there is a lot of empty space, the
  particles can be squeezed closer together.
  Therefore, gases are compressible.
• Because the particles are not held in close
  contact and are moving freely, gases expand to
  fill and take the shape of their container(s),
  and will flow.

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Liquids

• The particles in a liquid are closely packed, but
  they have some ability to move around.
• The close packing results in liquids being
  incompressible.
• But the ability of the particles to move allows
  liquids to take the shape of their container and
  to flow. However, they dont have enough freedom
  to escape and expand to fill the container(s).

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Solids

• The particles in a solid are packed close
  together and are fixed in position.
• Though they are vibrating.
• The close packing of the particles results in
  solids being incompressible.
• The inability of the particles to move around
  results in solids retaining their shape and
  volume when placed in a new container, and
  prevents the particles from flowing.

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Solids, Continued

• Some solids have their particles arranged in an
  orderly geometric pattern. We call these
  crystalline solids.
• Salt and diamonds.
• Other solids have particles that do not show a
  regular geometric pattern over a long range. We
  call these amorphous solids.
• Plastic and glass.

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Why Is Sugar a Solid, ButWater Is a Liquid?

• The state a material exists in depends on the
  attraction between molecules and their ability to
  overcome the attraction.
• The attractive forces between ions or molecules
  depends on their structure.
• The attractions are electrostatic.
• They depend on shape, polarity, etc.
• The ability of the molecules to overcome the
  attraction depends on the amount of kinetic
  energy they possess.

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Properties and Attractive Forces
Phase Density Shape Volume Relative strength of attractive forces
Gas Low Indefinite Indefinite Weakest
Liquid High Indefinite Definite Moderate
Solid High Definite Definite Strongest
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Phase ChangesMelting

• Generally, we convert a material in the solid
  state into a liquid by heating it.
• Adding heat energy increases the amount of
  kinetic energy of the molecules in the solid.
• Eventually, they acquire enough energy to
  partially overcome the attractive forces holding
  them in place.
• This allows the molecules enough extra freedom to
  move around a little and rotate.

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Phase ChangesBoiling

• Generally, we convert a material in the liquid
  state into a gas by heating it.
• Adding heat energy increases the amount of
  kinetic energy of the molecules in the liquid.
• Eventually, they acquire enough energy to
  completely overcome the attractive forces holding
  them together.
• This allows the molecules complete freedom to
  move around and rotate.

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Properties of LiquidsSurface Tension

• Liquids tend to minimize their surfacea
  phenomenon we call surface tension.
• This tendency causes liquids to have a surface
  that resists penetration.
• The stronger the attractive force between the
  molecules, the larger the surface tension.

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Surface Tension

• Molecules in the interior of a liquid experience
  attractions to surrounding molecules in all
  directions.
• However, molecules on the surface experience an
  imbalance in attractions, effectively pulling
  them in.
• To minimize this imbalance and maximize
  attraction, liquids try to minimize the number of
  molecules on the exposed surface by minimizing
  their surface area.
• Stronger attractive forces between the molecules
  larger surface tension.

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Properties of LiquidsViscosity

• Some liquids flow more easily than others.
• The resistance of a liquids flow is called
  viscosity.
• The stronger the attractive forces between the
  molecules, the more viscous the liquid is.
• Also, the less round the molecules shape, the
  larger the liquids viscosity.
• Some liquids are more viscous because their
  molecules are long and get tangled in each other,
  causing them to resist flowing.

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Escaping from the Surface

• The process of molecules of a liquid breaking
  free from the surface is called evaporation.
• Also known as vaporization.
• Evaporation is a physical change in which a
  substance is converted from its liquid form to
  its gaseous form.
• The gaseous form is called a vapor.

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Evaporation

• Over time, liquids evaporatethe molecules of the
  liquid mix with and dissolve in the air.
• The evaporation happens at the surface.
• Molecules on the surface experience a smaller net
  attractive force than molecules in the interior.
• All the surface molecules do not escape at once,
  only the ones with sufficient kinetic energy to
  overcome the attractions will escape.

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Escaping the Surface

• The average kinetic energy is directly
  proportional to the Kelvin temperature.
• Not all molecules in the sample have the same
  amount of kinetic energy.
• Those molecules on the surface that have enough
  kinetic energy will escape.
• Raising the temperature increases the number of
  molecules with sufficient energy to escape.

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Escaping the Surface, Continued

• Since the higher energy molecules from the liquid
  are leaving, the total kinetic energy of the
  liquid decreases, and the liquid cools.
• The remaining molecules redistribute their
  energies, generating more high energy molecules.
• The result is that the liquid continues to
  evaporate .

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Factors Effecting the Rate of Evaporation

• Liquids that evaporate quickly are called
  volatile liquids, while those that do not are
  called nonvolatile.
• Increasing the surface area increases the rate of
  evaporation.
• More surface molecules.
• Increasing the temperature increases the rate of
  evaporation.
• Raises the average kinetic energy, resulting in
  more molecules that can escape.
• Weaker attractive forces between the molecules
  faster rate of evaporation.

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Reconnecting with the Surface

• When a liquid evaporates in a closed container,
  the vapor molecules are trapped.
• The vapor molecules may eventually bump into and
  stick to the surface of the container or get
  recaptured by the liquid. This process is called
  condensation.
• A physical change in which a gaseous form is
  converted to a liquid form.

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Dynamic Equilibrium

• Evaporation and condensation are opposite
  processes.
• Eventually, the rate of evaporation and rate of
  condensation in the container will be the same.
• Opposite processes that occur at the same rate in
  the same system are said to be in dynamic
  equilibrium.

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Evaporation and Condensation
Shortly, the water starts to evaporate. Initially
the rate of evaporation is much faster than rate
of condensation
When water is just added to the flask and it is
capped, all the water molecules are in the
liquid.
Eventually, the condensation and evaporation
reach the same speed. The air in the flask is now
saturated with water vapor.
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Vapor Pressure

• Once equilibrium is reached, from that time
  forward, the amount of vapor in the container
  will remain the same.
• As long as you dont change the conditions.
• The partial pressure exerted by the vapor is
  called the vapor pressure.
• The vapor pressure of a liquid depends on the
  temperature and strength of intermolecular
  attractions.

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Boiling

• In an open container, as you heat a liquid the
  average kinetic energy of the molecules
  increases, giving more molecules enough energy to
  escape the surface.
• So the rate of evaporation increases.
• Eventually, the temperature is high enough for
  molecules in the interior of the liquid to
  escape. A phenomenon we call boiling.

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Boiling Point

• The temperature at which the vapor pressure of
  the liquid is the same as the atmospheric
  pressure is called the boiling point.
• The normal boiling point is the temperature
  required for the vapor pressure of the liquid to
  be equal to 1 atm.
• The boiling point depends on what the atmospheric
  pressure is.
• The temperature of boiling water on the top of a
  mountain will be cooler than boiling water at sea
  level.

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Temperature and Boiling

• As you heat a liquid, its temperature increases
  until it reaches its boiling point.
• Once the liquid starts to boil, the temperature
  remains the same until it all turns to a gas.
• All the energy from the heat source is being used
  to overcome all of the attractive forces in the
  liquid.

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Energetics of Evaporation

• As it loses its high energy molecules through
  evaporation, the liquid cools.
• Then the liquid absorbs heat from its
  surroundings to raise its temperature back to the
  same as the surroundings.
• Processes in which heat flows into a system from
  the surroundings are said to be endothermic.
• As heat flows out of the surroundings, it causes
  the surroundings to cool.
• As alcohol evaporates off your skin, it causes
  your skin to cool.

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Energetics of Condensation

• As it gains the high energy molecules through
  condensation, the liquid warms.
• Then the liquid releases heat to its surroundings
  to reduce its temperature back to the same as the
  surroundings.
• Processes in which heat flows out of a system
  into the surroundings are said to be exothermic.
• As heat flows into the surroundings, it causes
  the surroundings to warm.

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Heat of Vaporization

• The amount of heat needed to vaporize one mole of
  a liquid is called the heat of vaporization.
• DHvap
• It requires 40.7 kJ of heat to vaporize one mole
  of water at
• 100 C.
• Always endothermic.
• Number is .
• DHvap depends on the initial temperature.
• Since condensation is the opposite process to
  evaporation, the same amount of energy is
  transferred but in the opposite direction.
• DHcondensation -DHvaporization

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Heats of Vaporization of Liquidsat Their Boiling
Points and at 25 C
Liquid Chemical formula Normal boiling point, C DHvap at boiling point, (kJ/mol) DHvap at 25 C, (kJ/mol)
Water H2O 100 40.7 44.0
Isopropyl alcohol C3H7OH 82.3 39.9 45.4
Acetone C3H6O 56.1 29.1 31.0
Diethyl ether C4H10O 34.5 26.5 27.1
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Example 12.1Calculate the Mass of Water that Can
Be Vaporized with 155 KJ of Heat at 100 C.
155 kJ g H2O
Given Find
1 mol H2O 40.7 kJ, 1 mol 18.02 g
Solution Map Relationships
Solution
Since the given amount of heat is almost 4x the
DHvap, the amount of water makes sense.
Check
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• Example 12.1
• Calculate the amount of water in grams that can
  be vaporized at its boiling point with 155 kJ of
  heat.

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ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Write down the given quantity and its units.
• Given 155 kJ

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ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Information
• Given 155 kJ

• Write down the quantity to find and/or its units.
  
• Find ? g H2O

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ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Information
• Given 155 kJ
• Find g H2O

• Collect needed conversion factors
• DHvap 40.7 kJ/mol ? 40.7 kJ ? 1 mol H2O
• 18.02 g H2O 1 mol H2O

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ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Information
• Given 155 kJ
• Find g H2O
• Conversion Factors
• 40.7 kJ 1 mol 18.02 g 1 mol

• Write a solution map for converting the units

kJ
mol H2O
g H2O
43
ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Information
• Given 155 kJ
• Find g H2O
• Conversion Factors
• 40.7 kJ 1 mol 18.02 g 1 mol
• Solution Map kJ ? mol ? g

• Apply the solution map

68.626 g H2O

• Significant figures and round

68.6 g H2O
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ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Information
• Given 155 kJ
• Find g H2O
• Conversion Factors
• 40.7 kJ 1 mol 18.02 g 1 mol
• Solution Map kJ ? mol ? g

• Check the solution

155 kJ of heat can vaporize 68.6 g H2O.
The units of the answer, g, are correct. The
magnitude of the answer makes sense since it is
more than one mole.
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PracticeHow Much Heat Energy, in kJ, is Required
to Vaporize 87 g of Acetone, C3H6O, (MM 58.08) at
25 ?C? (DHvap 31.0 kJ/mol)
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PracticeHow Much Heat Energy, in kJ, Is Required
to Vaporize 87 g of Acetone, C3H6O, (MM 58.08) at
25 ?C? (DHvap 31.0 kJ/mol), Continued
87 g C3H6O kJ
Given Find
1 mol C3H6O 31.0 kJ at 25 ?C, 1 mol 58.08 g
Solution Map Relationships
Solution
Since the given mass is than one mole, the answer
being greater than DHvap makes sense.
Check
47
Temperature and Melting

• As you heat a solid, its temperature increases
  until it reaches its melting point.
• Once the solid starts to melt, the temperature
  remains the same until it all turns to a liquid.
• All the energy from the heat source is being used
  to overcome some of the attractive forces in the
  solid that hold them in place.

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Energetics of Melting and Freezing

• When a solid melts, it absorbs heat from its
  surroundings, it is endothermic.
• As heat flows out of the surroundings, it causes
  the surroundings to cool.
• As heat flows out of your drink into the ice
  cubes (causing them to melt), the liquid gets
  cooler.
• When a liquid freezes, it releases heat into its
  surroundings, it is exothermic.
• As heat flows into the surroundings, it causes
  the surroundings to warm.
• Orange growers often spray their oranges with
  water when a freeze is expected. Why?

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Heat of Fusion

• The amount of heat needed to melt one mole of a
  solid is called the heat of fusion.
• DHfus
• Fusion is an old term for heating a substance
  until it melts, it is not the same as nuclear
  fusion.
• Since freezing (crystallization) is the opposite
  process of melting, the amount of energy
  transferred is the same, but in the opposite
  direction.
• DHcrystal -DHfus
• In general, DHvap gt DHfus because vaporization
  requires breaking all attractive forces.

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Heats of Fusion of Several Substances
Liquid Chemical formula Melting point, C DHfusion, (kJ/mol)
Water H2O 0.00 6.02
Isopropyl alcohol C3H7OH -89.5 5.37
Acetone C3H6O -94.8 5.69
Diethyl ether C4H10O -116.3 7.27
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Example 12.2Calculate the Mass of Ice that Can
Be Melted with 237 kJ of Heat.
237 kJ g H2O
Given Find
1 mol H2O 6.02 kJ, 1 mol 18.02 g
Solution Map Relationships
Solution
Since the given amount of heat is almost 4x the
DHvap, the amount of water makes sense.
Check
52
PracticeHow Much Heat Energy, in kJ, is Required
to Melt 87 g of Acetone, C3H6O, (MM
58.08)?(DHfus 5.69 kJ/mol)
53
PracticeHow Much Heat Energy, in kJ, Is Required
to Melt 87 g of Acetone, C3H6O, (MM 58.08)?,
Continued
87 g C3H6O kJ
Given Find
1 mol C3H6O 5.69 kJ at -94.8 ?C, 1 mol
58.08 g
Solution Map Relationships
Solution
Since the given mass is more than one mole, the
answer being greater than DHvap makes sense.
Check
54
Sublimation

• Sublimation is a physical change in which the
  solid form changes directly to the gaseous form.
• Without going through the liquid form.
• Like melting, sublimation is endothermic.

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IntermolecularAttractive Forces
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Effect of the Strength of Intermolecular
Attractions on Properties

• The stronger the intermolecular attractions are,
  the more energy it takes to separate the
  molecules.
• Substances with strong intermolecular attractions
  have higher boiling points, melting points, and
  heat of vaporization they also have lower vapor
  pressures.

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PracticePick the Substance in Each Pair with the
Stronger Intermolecular Attractions.

• sugar or water
• water or acetone
• ice or dry ice

• sugar or water.
• water or acetone.
• ice or dry ice.

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Why Are Molecules Attracted to Each Other?

• Intermolecular attractions are a result of
  attractive forces between opposite charges.
• ion to ion.
• end of one polar molecule to - end of another
  polar molecule.
• H-bonding is especially strong.
• Even nonpolar molecules will have temporary
  induced dipoles.
• Larger charge stronger attraction.

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Dispersion Forces

• Also known as London forces or instantaneous
  dipoles.
• Caused by distortions in the electron cloud of
  one molecule inducing distortion in the electron
  cloud on another.
• Distortions in the electron cloud lead to a
  temporary dipole.
• The temporary dipoles lead to attractions between
  moleculesdispersion forces.
• All molecules have attractions caused by
  dispersion forces.

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Instantaneous Dipoles
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Strength of the Dispersion Force

• Depends on how easily the electrons can move, or
  be polarized.
• The more electrons and the farther they are from
  the nuclei, the larger the dipole that can be
  induced.
• Strength of the dispersion force gets larger with
  larger molecules.

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Dispersion Force and Molar Mass
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PracticeThe Following Are All Made of NonPolar
Molecules. Pick the Substance in Each Pair with
the Highest Boiling Point.

• CH4 or C3H8
• BF3 or BCl3
• CO2 or CS2

• CH4 or C3H8.
• BF3 or BCl3.
• CO2 or CS2.

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Permanent Dipoles

• Because of the kinds of atoms that are bonded
  together and their relative positions in the
  molecule, some molecules have a permanent dipole.
• Polar molecules.
• The size of the molecules dipole is measured in
  debyes, D.

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Dipole-to-Dipole Attraction

• Polar molecules have a permanent dipole.
• A end and a end.
• The end of one molecule will be attracted to
  the end of another.

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Polarity and Dipole-to-Dipole Attraction
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Attractive Forces
Dispersion forcesAll molecules.
Dipole-to-dipole forcesPolar molecules.
-
-
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Intermolecular Attraction and Properties

• All molecules are attracted by dispersion forces.
• Polar molecules are also attracted by
  dipole-dipole attractions.
• Therefore, the strength of attraction is stronger
  between polar molecules than between nonpolar
  molecules of the same size.

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PracticeDetermine Which of the Following Has
DipoleDipole Attractive Forces.(EN C 2.5, F
4, H 2.1, S 2.5)

• CS2 Nonpolar bonds nonpolar molecule.
• CH2F2 Polar bonds and asymmetrical polar
  molecule.
• CF4 Polar bonds and symmetrical shape
  nonpolar molecule.

• CS2
• CH2F2
• CF4

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Attractive Forces and Properties

• Like dissolves like.
• Miscible Liquids that do not separate, no
  matter what the proportions.
• Polar molecules dissolve in polar solvents.
• Water, alcohol, CH2Cl2.
• Molecules with O or N higher solubility in H2O
  due to H-bonding with H2O.
• Nonpolar molecules dissolve in nonpolar solvents.
• Ligroin (hexane), toluene, CCl4.
• If molecule has both polar and nonpolar parts,
  then hydrophilic-hydrophobic competition.

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Immiscible Liquids
When liquid pentane, a nonpolar substance, is
mixed with water, a polar substance, the two
liquids separate because they are more attracted
to their own kind of molecule than to the other.
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Hydrogen Bonding

• HF, or molecules that have OH or NH groups have
  particularly strong intermolecular attractions.
• Unusually high melting and boiling points.
• Unusually high solubility in water.
• This kind of attraction is called a hydrogen bond.

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Properties and H-Bonding
Name Formula Molar mass (g/mol) Structure Boiling point, C Melting point, C Solubility in water
Ethane C2H6 30.0 -88 -172 Immiscible
Ethanol CH4O 32.0 64.7 -97.8 Miscible
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Intermolecular H-Bonding
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Hydrogen Bonding

• When a very electronegative atom is bonded to
  hydrogen, it strongly pulls the bonding electrons
  toward it.
• Since hydrogen has no other electrons, when it
  loses the electrons, the nucleus becomes
  deshielded.
• Exposing the proton.
• The exposed proton acts as a very strong center
  of positive charge, attracting all the electron
  clouds from neighboring molecules.

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H-Bonds vs. Chemical Bonds

• Hydrogen bonds are not chemical bonds.
• Hydrogen bonds are attractive forces between
  molecules.
• Chemical bonds are attractive forces that make
  molecules.

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Attractive Forces and Properties
80
Example 12.5Which of the Following Is a Liquid
at Room Temperature? (The Other Two Are Gases.)

• formaldehyde, CH2O
• 30.03 g/mol
• polar molecule ? dipoledipole attractions
  present
• polar CO bond asymmetric
• fluoromethane, CH3F
• 34.03 g/mol
• polar molecule ? dipoledipole attractions
  present
• polar C-F bond asymmetric
• hydrogen peroxide, H2O2
• 34.02 g/mol
• polar molecule ? dipoledipole attractions
  present
• polar H-O bonds asymmetric
• H-O bonds ? Hydrogen bonding present

• formaldehyde, CH2O.
• 30.03 g/mol.
• polar molecule ? dipoledipole attractions
  present.
• Polar CO bond and asymmetric.
• fluoromethane, CH3F.
• 34.03 g/mol.
• polar molecule ? dipoledipole attractions
  present.
• Polar C-F bond and asymmetric.
• hydrogen peroxide, H2O2
• 34.02 g/mol.
• polar molecule ? dipoledipole attractions
  present.
• Polar H-O bonds and asymmetric.
• H-O bonds ? hydrogen-bonding present.

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PracticePick the Compound in Each Pair Expected
to Have the Higher Solubility in H2O.

• CH3CH2OCH2CH3 or CH3CH2CH2CH2CH3.
• CH3CH2NHCH3 or CH3CH2CH2CH3.
• CH3CH2OH or CH3CH2CH2CH2CH2OH.

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PracticePick the Compound in Each Pair Expected
to Have the Higher Solubility in H2O, Continued.

• CH3CH2OCH2CH3 or CH3CH2CH2CH2CH3 contains
  polar O.
• CH3CH2NHCH3 or CH3CH2CH2CH3 contains polar N.
• CH3CH2OH or CH3CH2CH2CH2CH2OH contains less
  nonpolar parts.

83
Types of Intermolecular Forces
Type of force Relative strength Present in Example
Dispersionforce Weak, but increases with molar mass All atoms and molecules H2
Dipole Dipole force Moderate Only polar molecules HCl
Hydrogen Bond Strong Molecules having H bonded to F, O, or N HF
84
Crystalline Solids
85
Types of Crystalline Solids
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Molecular Crystalline Solids

• Molecular solids are solids whose composite units
  are molecules.
• Solid held together by intermolecular attractive
  forces.
• Dispersion, dipole-dipole, or H-bonding.
• Generally low melting points and DHfusion.

87
Ionic Crystalline Solids

• Ionic solids are solids whose composite units are
  formula units.
• Solid held together by electrostatic attractive
  forces between cations and anions.
• Cations and anions arranged in a geometric
  pattern called a crystal lattice to maximize
  attractions.
• Generally higher melting points and DHfusion than
  molecular solids.
• Because ionic bonds are stronger than
  intermolecular forces.

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Atomic Crystalline Solids

• Atomic solids are solids whose composite units
  are individual atoms.
• Solids held together by either covalent bonds,
  dispersion forces, or metallic bonds.
• Melting points and DHfusion vary depending on the
  attractive forces between the atoms.

89
PracticeClassify Each of the Following
Crystalline Solids as Molecular, Ionic, or Atomic.

• H2O(s)
• Si(s)
• C12H22O11(s)
• CaF2(s)
• Sc(NO3)3(s)

• H2O(s)molecular.
• Si(s)atomic.
• C12H22O11(s)molecular.
• CaF2(s)ionic.
• Sc(NO3)3(s)ionic.

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Types of Atomic Solids
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Types of Atomic SolidsCovalent

• Covalent atomic solids have their atoms attached
  by covalent bonds.
• Effectively, the entire solid is one giant
  molecule.
• Because covalent bonds are strong, these solids
  have very high melting points and DHfusion.
• Because covalent bonds are directional, these
  substances tend to be very hard.
• Elements found as covalent atomic solids are C,
  Si, and B.
• Compounds that occur as covalent atomic solids
  include SiO2 and SiC.

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Types of Atomic SolidsNonbonding

• Nonbonding atomic solids are held together by
  dispersion forces.
• Because dispersion forces are relatively weak,
  these solids have very low melting points and
  DHfusion.
• All the noble gases form nonbonding atomic
  solids.

93
Types of Atomic SolidsMetallic

• Metallic solids are held together by metallic
  bonds.
• Metal atoms release some of their electrons to be
  shared by all the other atoms in the crystal.
• The metallic bond is the attraction of the metal
  cations for the mobile electrons.
• Often described as islands of cations in a sea of
  electrons.

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Metallic Bonding

• The model of metallic bonding can be used to
  explain the properties of metals.
• The luster, malleability, ductility, and
  electrical and thermal conductivity are all
  related to the mobility of the electrons in the
  solid.
• The strength of the metallic bond varies,
  depending on the charge and size of the cations,
  so the melting points and DHfusion of metals vary
  as well.

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Substances with Both Bonding and Nonbonding
Attractions

• Some substances have chains or layers of bonded
  atoms that are then attracted by dispersion
  forces.
• Chain substances include grey selenium, polymeric
  SO3, and asbestos.
• Layer substances include graphite, black
  phosphorus, and mica.

96
PracticeDecide if Each of the Following Atomic
Solids Is Covalent, Metallic, or Nonbonding.

• diamond
• neon
• iron

• diamond covalent.
• neon nonbonding.
• iron metallic.

97
Water A Unique and Important Substance

• Water is found in all three states on Earth.
• As a liquid, it is the most common solvent found
  in nature.
• Without water, life as we know it could not
  exist.
• The search for extraterrestrial life starts with
  the search for water.

98
Water

• Liquid at room temperature.
• Most molecular substances that have a molar mass
  (18.02 g/mol) similar to waters are gaseous.
• Relatively high boiling point.
• Expands as it freezes.
• Most substances contract as they freeze.
• Causes ice to be less dense than liquid water.