Solids and Liquids

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Solids and Liquids

• AP Chemistry
• Chapter 11

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Kinetic Molecular Theory

• When discussing gases, we talked about the
  motions of the molecules and how they interacted.
• These interactions which were undesirable for
  ideal gases are very important with solids and
  liquids

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We said that gases

• Have minimal interactions
• Take the shape and the volume of their container
• Have extensive translational, rotational, and
  vibrational motions.

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Liquids

• Are loosely packed, so the intermolecular forces
  are important.
• Take the shape of their container but only to a
  point because they have a definite volume.
• Have some translational, rotational and
  vibrational motions.

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Solids

• Are tightly packed because of strong
  intermolecular forces.
• Have a definite shape and a definite volume.
• Have limited vibrational motions.

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Intermolecular Forces (IMF)

• In a gas, like water vapor, there are many
  molecules, but they are pretty spread out so they
  dont interact much.
• In solid ice, there are many H2O molecules, but
  each molecule attracts other molecules, forming
  bonds.
• These bonds are generally weaker than the
  chemical bonds we learned about, but are still
  strong enough to hold each other together.

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London Dispersion Forces

• This class of forces affects all molecules and
  atoms, but is the weakest IMF.
• Atoms or Nonpolar molecules can only have London
  Dispersion Forces.

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How London Dispersion Forces Work

• Electrons are always in motion. At times there
  is more charge on one side (especially when
  relatively near to a positive nucleus), causing a
  temporary pole to set up. This pole forces other
  atoms/molecules to align with it through
  electrical attractions and repulsions.

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Polarizibility

• Polarizability is the relative tendency of a
  charge distribution, like the electron cloud of
  an atom or molecule, to be distorted from its
  normal shape by an external electric field, which
  may be caused by the presence of a nearby ion or
  dipole.

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and what does that mean?

• Some atoms/molecules are more easily forced into
  dipoles than others.
• For example is we compare CH4 to C8H18, the
  noticeable size factor of octane means that it
  has more room to shift its electrons, so it is
  more polarizible.
• Since octane is more polarizible, it would have
  stronger London Dispersion forces than methane.

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Dipole - Dipole

• Polar molecules have dipole moments. The pole of
  one molecule can attract the opposite pole of
  another, causing weak bonds to form.

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Dipoles in Solids Liquids

• In solids, the dipoles tend to line up more in an
  orderly pattern.
• However liquids have too much energy for this, so
  the molecules still have attractions, but these
  forces are usually weaker.

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Hydrogen Bonding

• Hydrogen Bonding is a special case of the
  dipole-dipole IMF.
• H atoms that are bonded to either N, O, or F, are
  also attracted to the N, O, or F atoms of other
  molecules.

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Why N, O, and F?

• Look at this graph. Each line represents the
  boiling points of a family of hydrogen-nonmetallic
  compounds. We expect H2O, HF and NH3 to be the
  lowest member of their family, but they are the
  highest. Why?
• Small size and high electronegativity.

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Why Does Ice Float?

• We have already said that solids tend to be the
  most tightly packed substances. This means that
  solids have the highest densities.
• Yet we also know that ice floats in water, so
  water is more dense. Lets compare ice and water
  for a moment.

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Which is ice which is water?
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The ice was on the left

• The molecules in a solid tend to be more orderly.
  The hydrogen bonding puts the molecules into
  hexagons with holes in the middle, making the ice
  less dense.
• By contrast, the molecules of the liquid are not
  orderly, so molecules can get closer with
  stronger attractions, making the liquid more
  dense.

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Ion-dipole attraction

• In the solvation process, the polar water
  molecules separate the ions of the solute.
• The negative ends of the water attract the
  cation, and the positive ends attract the anion.

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Ion-Ion attractions

• Like an ionic bond, this attraction is between a
  cation and an anion, but the ions can be in
  different molecules.

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More on ion-ion forces

• Each anion attracts several cations (and vice
  versa). This makes ionic solids very strong.
  Ion-ion are the strongest IMFs.

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Name that IMF.

• What IMF is the major force between
• The molecules of solid CO2
• The molecules of solid Fe2O3
• The molecules of H2CO
• The molecules of a saturated solution of silver
  nitrate
• The molecules of CH3OH

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Answers and Explanations

• The molecules of solid CO2
• Since CO2 is nonpolar, it can only have London
  Dispersion Forces
• The molecules of solid Fe2O3
• Since Fe2O3 is ionic, the IMF is ion-ion
• The molecules of H2CO
• The hydrogen is not bonding to the oxygen in this
  polar molecule, so the IMF is dipole-dipole.
• The molecules of a saturated solution of silver
  nitrate
• Water is polar and silver nitrate is ionic, so
  the IMF must be ion-dipole
• The molecules of CH3OH
• Here the hydrogen is bonded to the oxygen in a
  polar molecule, so the IMF is hydrogen bonding.

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Some Important Properties that Depend on IMFs.

• Surface Tension
• Viscosity
• Adhesion
• Cohesion
• Vapor Pressure
• Boiling Point
• Melting Point

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Surface Tension

• As water molecules bond to each other, there is
  an imbalance at the surface, pulling the
  molecules together, thus creating a harder
  layer of molecules.

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Surface Tension

• The stronger the IMF of the molecules of the
  liquid, the stronger the surface tension would be.

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Viscosity

• The resistance of a fluid to flow.
• The greater the viscosity of a liquid, the harder
  it is to pour the liquid.
• The greater the IMF, the greater the viscosity.

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Adhesion and Cohesion

• Adhesion - force of attraction between molecules
  of different substances
• Cohesion- force of attraction between molecules
  of the same substance

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Beading effect

• On a waxed car, the water adheres to the wax
  somewhat, preventing the water from running off.
• The waters cohesiveness causes it to pool into
  little puddles or beads of water.

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Capillary Action

• A liquid can adhere to a glass tube, but also
  pull itself up by its cohesive properties.
• Note that narrower tubes have more capillary
  action. Why?

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Vapor Pressure

• The pressure exerted by a gas on its liquid in
  dynamic equilibrium.
• The higher the IMF, the harder it is to boil a
  substance, and so the higher its vapor pressure
  will be.

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At what Temperature does Water Boil?

• Ask the average person this, and if you are
  lucky, they will say that water boils at 212 oF,
  100 oC, or 373 K.
• Are they right?
• Sometimes yes, and many times no.

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Vapor Pressure and Boiling Point

• To boil, the vapor pressure must equal the
  atmospheric pressure. Since atmospheric pressure
  changes, so does the boiling point.
• As temperature increases, vapor pressure
  increases.

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What is the difference between evaporation and
boiling?

• Evaporation is a slow process where surface
  molecules gain enough energy to break through the
  air pressure.
• Boiling takes place from the bottom-up. As
  heating occurs, the vapor pressure builds. When
  the vapor pressure is equal to the atmospheric
  pressure, the gas bubbles up to the surface and
  is released.

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Melting

• For melting to occur, the usually strong IMFs
  have to be broken to allow the new liquid to
  rotate and to translate.
• This takes energy, usually in the form of heat.

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Phase Changes
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Endothermic Phase Changes

• We have already stated that melting (from solid
  to liquid) and boiling (from liquid to gas)
  require heat, and thus are endothermic processes.
• What if a substance goes from solid directly to
  gas?
• This process, called sublimation, is also
  endothermic

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Exothermic Phase Changes

• Logically if melting is endothermic, then
  freezing must be exothermic.
• If boiling is endothermic, condensing must be
  exothermic.
• If sublimation is endothermic, deposition must be
  exothermic.

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About Heating

• If I heat a pot of room temperature water, does
  it immediately start boiling?
• Of course not. Water has a high heat capacity,
  and can hold heat well. It will keep holding
  heat until the vapor pressure reaches the
  atmospheric pressure and then the water will
  start to boil.

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Heating/Cooling Curve
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As heat is added to a solid

• the solid warms, its temperature increasing until
  it reaches the melting point.
• At the melting point, both solid and liquid are
  present, and the temperature does not change.
• Once fully melted, any additional heat goes into
  the liquid.

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As heat is added to the liquid

• The liquid warms, its temperature increasing
  until it reaches the boiling point.
• At the boiling point, the liquid and the gas are
  both present, and the temperature does not
  change.
• Once boiling is complete, any additional heat
  goes into the gas.

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Think about this

• Lets suppose we have 200 grams of ice at -50 oC
  and we want to convert this entirely to steam at
  130 oC. Lets figure out how much heat we would
  need.

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Step One

• We have ice, but it is not at its normal melting
  point. So first we need to heat it.
• Qmcice?T
• The mass is 200 g, the specific heat capacity of
  ice is 2.02 J/gK, and the change in temperature
  would be from -50 oC to 0 oC.
• Q (200g)(2.02 J/gK)(50 K)
• Q 20200 J

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Step Two

• Now we are ready to melt the ice.
• For melting, Q m?Hfusion
• The mass is still 200 g, and the latent heat of
  fusion is 334 J/g.
• Q (200 g)(334 J/g)
• Q 66800 J

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Step Three

• Now we have water at 0 oC. The water can keep
  being heated until its temperature is 100 oC.
• Q mcwater?T
• The mass is still 200 g, the specific heat
  capacity of water is 4.18 J/gK, and the change in
  temperature would be 100 Co.
• Q (200 g)(4.18 J/gK)(100 K)
• Q 83600J

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Step Four

• Now we are ready to boil the water.
• For boiling, Q m?Hvapor
• The mass is still 200 g and the latent heat of
  vaporization of water is 2260 J/g.
• Q (200 g)(2260 J/g)
• Q 452000 J

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Step Five

• Now we finally have gas at 100 oC, but we want
  gas at 130 oC. So we heat again.
• Q mcsteam?T
• The mass is STILL 200 g, the specific heat
  capacity of steam is 2.08 J/gK, and the change in
  temperature is 30 Co.
• Q (200 g)(2.08 J/gK)(30 K)
• Q 12480 J

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Step 6

• We heated the solid ice (20200 J), we melted the
  ice (66800 J), we heated the water formed (83600
  J), we boiled this water (452000 J), and we
  heated the steam (12480 J). Now we just have to
  add this up.
• Qtotal 635,080 J or about 635 kJ

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But Remember

• Water doesnt always boil at 100 oC, nor does it
  always melt at 0 oC.
• These are the normal boiling and melting points.
  This means that we are only considering standard
  pressure.

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How do we tell what phase change will happen at
which temperature?

• If the temperature at which a substance melts,
  boils, or sublimates can vary with pressure, then
  we have to examine a pressure-temperature graph
  called a phase diagram.

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Phase Diagram
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More on the phase diagram

• Most phase diagrams take this general shape.
• The 3 curves (melting, boiling, sublimating) all
  meet at one point called the triple point.

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The Triple Point

• At a specific temperature and pressure it is
  possible for a substance to boil and freeze
  simultaneously!
• http//video.aol.com/video-detail/triple-point-of-
  water/62299915

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The Critical Point

• We have stated that if we put a liquid under
  greater pressure than it wont boil. That is
  true up until a point.
• That point is called the critical point.

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The Critical Point Explained

• As the temperature increases, the liquid becomes
  more active. Once the temperature is at the
  critical point, there is too much molecular
  motion to be contained in the liquid state,
  regardless of the pressure.
• If the pressure is still increased, a super
  critical fluid is formed.

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A Super Critical Fluid is

• a fluid that is somewhat like a gas in its
  motions and ability to penetrate other objects
  but also like a liquid in that it is a strong
  solvent.

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Of course, water is special

• Most melting curves have a positive slope, but a
  few substances, like water, have a negative slope.

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Why is water special?

• Look at 0 oC. As the pressure is increased, the
  ice melts.
• Usually an increase in pressure would not affect
  a solid, but with ice, the added pressure forces
  the molecules out of the rigid hexagonal shapes,
  causing the hydrogen bond to realign and become
  less structured.

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Helium is different too.
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Types of solids

• If the IMFs are orderly, a crystal is formed.
• If the IMFs dont line up, an amorphous solid is
  formed.

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Amorphous Solids

• Amorphous solids are generally substances like
  glass, rubber, and plastic.
• Because their IMFs are not orderly, some people
  argue that they should be considered liquids, not
  solids.

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Cystalline solids

• There are four classes of Crystalline Solids
• Metallic
• Ionic
• Covalent Network
• Molecular

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Metallic Solid

• Metallic solids are made from metals.
• Metals tend to delocalize their valence
  electrons, increasing attractions between cations
  and the wandering electrons.

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Metallic crystals

• Metals normally occur as solids (high melting
  points).
• Thus, there must be strong bonds between the
  atoms of metals causing them to bond
• Bonding in metals and alloys is different from in
  other compounds positive nuclei exist in a sea
  of electrons (this explains why metals conduct
  electricity)

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Properties of Metallic Solids

• Metals
• are good conductors as solids
• have high melting points
• are malleble and ductile
• do not dissolve in water

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Ionic Solids

• Ionic solids are made from ions
• Ionic solids are characterized by repeating
  patterns called unit cells.

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More about Ionic Solids

• Ionic Solids
• Are soluble in water
• Are not conductors as solids
• Conduct electricity in water (aq) and in their
  liquid state (l)
• Are very rigid
• Tend to have very high melting and boiling points.

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Covalent Network Solids

• Covalent Network Solids are made from carbon or
  silicon, and certain oxides.
• Very orderly with strong forces.

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Network solids (covalent crystals)

• There are some compounds that do not have
  molecules, but instead are long chains of
  covalent bonds (E.g. diamond)

• This happens in 3 dimensions, creating a crystal
• Because there are only covalent bonds, network
  solids are extraordinarily strong

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About Covalent Network Solids

• Covalent Network Solids
• Are extremely hard
• Have extremely high melting points
• Are poor conductors
• Are not soluble

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Molecular Solids

• Molecular Solids are made from polar and nonpolar
  molecules.
• Of course, polar molecules make for stronger
  molecular solids.

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About Molecular Solids

• Molecular Solids
• Are brittle
• Have lower melting points
• Are not generally soluble (although polar
  molecular solids may be slightly soluble)
• Are poor conductors

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Guess that solid!

• Identify the type of solid formed by
• dry ice (solid carbon dioxide)
• silicon
• molybdenum
• a PVC tube (or plumbing pipe)
• magnesium sulfate

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The Lucky Winners are

• CO2 is a nonpolar molecule, so it will form a
  molecular solid.
• Silicon (and carbon) form covalent network
  solids.
• Molybdenum (you know you like to say it) is a
  metal and thus a metallic solid.
• PVC (Polyvinyl chloride) is a plastic and
  plastics are amorphous.
• Magnesium sulfate is an ionic compound and an
  ionic solid.

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Unit Cells

• Unit cells are the simplest repeating units in a
  crystal (typically ionic compounds).
• Opposite faces of a unit cell are parallel.
• The edge of a unit cell connect equivalent
  points.

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Examples of Unit Cells
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Types of crystals