condensed states of matter: liquids and solids chapter 14
TRANSCRIPT
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Condensed States of Matter:
Liquids and Solids
Chapter 14
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Condensed States
• Liquids and solids = condensed states because they have significantly higher densities than gases
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Gas Liquid Solid
highly compressible
only slightly compressible
only slightly compressible
low density high density high density
fills container completely
does not expand to fill container -
has definite volume
rigidly retains its volume
assumes shape of container
assumes shape of container retains own shape
rapid diffusion slow diffusion
extremely slow diffusion - only
occurs at surfaces
high expansion on heating
low expansion on heating
low expansion on heating
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KMT of Matter
• According to the kinetic molecular theory, the state of a substance at room temperature depends on the strength of the attractions between its particles.
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Forces
• Intramolecular forces are between the atoms within a molecule = bonds
• Intermolecular forces are between molecules
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Intramolecular Forces
• Covalent Bonds – between nonmetals; sharing of electrons
• Ionic Bonds – between metals and nonmetals; transfer of electrons to form ions
• Metallic Bonds – between metals; sea of electrons
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Intermolecular Forces
1. Dipole-dipole attractionsa. Hydrogen bonding
2. Ion-dipole attractions3. London dispersion forces
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Dipole-Dipole Attractions• Attractions due to permanent
dipoles in polar molecules• Remember, a dipole is
created when positive and negative charges are separated by some distance.
• Only 1% as strong as covalent or ionic bonds and even weaker if distance between charges increases.
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Hydrogen Bonds
• Unusually strong dipole-dipole attractions that occur among molecules in which hydrogen is bonded to a highly
electronegative atom, such as F, N, or O
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Ion-Dipole Attractions
• Polar molecules surround ions based on their attraction for the charge on the ion
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London Dispersion Forces• Explain the attraction that
exists between non-polar molecules
• Even in these molecules the electrons are not uniformly distributed at every second
• Temporary dipolar arrangement of charge creates an instantaneous dipole.
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London Dispersion Forces• Instantaneous dipoles can
induce similar dipoles in neighboring atoms
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Kinetic
Kinetic EnergyIntermolecular
AttractionState at Room
Temp.
low high solids
medium medium liquids
high low gases
Kinetic energy determines if particles will overcome the intermolecular forces keeping
them together. Thus, higher attractions mean higher boiling points, because a higher kinetic
energy will be needed to overcome the attraction.
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Properties of Solids
• Intermolecular forces keep the particles of a solid packed together tightly- Particles are highly ordered
with fixed positions- Only particle movement =
vibrations• Liquids have similar distance
between particles
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Bonding in Solids
2 basic types of solids based on the nature of the particles that make them up:
1. Crystallinea. Atomicb. Ionicc. Covalent-network
2. Amorphous
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Crystalline Solids• Solid in which the
representative particles are in a highly ordered, repeating pattern called a crystal
• Can be studied as unit cells, small representative units that repeat throughout the structure
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Bonding in Crystalline Solids
The physical properties of solids, such as hardness,
electrical conductibility, and melting point, depend on the kind of particles that make up the solid and on the strength
of the attractive forces between them. (Includes ionic,
covalent, and atomic substances)
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Covalent-Network Solids
• A type of crystalline solid in which strong covalent bonds forma a network extending throughout the solid
• Have very high melting points due to strong covalent bonds
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Amorphous Solids
• Solids in which the arrangement of the representative particles lacks a regular, repeating pattern
• Also known as “supercooled liquids”- Liquids cooled until
viscosity becomes so high that no flow can occur
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Amorphous Solids
• Get softer when heated instead of reaching an abrupt melting point like crystalline solids
• Examples: glass, rubber, some plastics
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LiquidsThe physical properties of liquids are determined by the nature and strength of the intermolecular forces
present between their molecules
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Properties of Liquids
• Viscosity – resistance to motion that exists between molecules of a liquid when they move past each other– Increased intermolecular forces
yield increased viscosity
• Glycerine has lots of hydrogen bonds
• Decreased temperatures yield increased viscosity
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Properties of Liquids
• Surface tension – imbalance of attractive forces at the surface of a liquid that cause the surface to behave as if it had a film across it
»Explains “beading” of liquids»As with viscosity, increased
intermolecular forces or decreased
temperatures yield increased surface tension
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Phase Changes
• Changes of state are physical changes, not chemical.– Intermolecular forces break or
form, not intramolecular• Phase changes occur when
energy enters or leaves a compound.
• Energy in: solid liquid gas• Energy out: gas liquid solid
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Heating/Cooling Curves• Show the changes in
temperature as energy as heat is added to (or removed from) a substance and it changes states.
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Heating/Cooling Curve
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Phase Changes
• During phase changes, heat is still transferred, but no temperature change takes place.
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Phase Diagrams
• Shows the state of matter of a substance at increasing pressures and temperatures.
• Can be used to determine the state of matter of a substance at a particular pressure and temperature.
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Phase Diagram of Water
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Energy Requirements
• Molar heat of fusion: the energy required to melt one mole of a solid
• Molar heat of vaporization: the energy required to vaporize one mole of a liquid
• Molar heats of vaporization are higher because more intermolecular forces will have to be overcome to go from a liquid to a gas than a solid to a liquid.
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Vaporization
• Not all of the particles in a liquid have the same kinetic energy.
• Temperature is a measurement of the average KE.
• The particles of a liquid must be moving at a sufficient speed to overcome the intermolecular forces of the liquid to escape as a gas.
• Thus, evaporation is endothermic: it requires energy to occur.
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Vapor Pressure
• Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid phase at a certain temperature.
• For a liquid in a closed container, vaporization and condensation will occur at an equal rate once the equilibrium vapor pressure is reached.
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Vapor Pressure and Boiling Point
• When water begins to boil, bubbles of H2O gas are created as some of the molecules escape the intermolecular bonds holding the liquid together.
• The bubbles will expand and rise to the surface only if the pressure exerted by the water vapor in the bubble is greater than the atmospheric pressure.
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Vapor Pressure and Boiling Point
• Thus, the vapor pressure of the water must be equal to atmospheric pressure before boiling can occur.
• At temperatures below 100oC, the vapor pressure of water is below 1 atm, so the atmospheric pressure prevents boiling from occurring.
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Elevation and Boiling Point• What happens to
atmospheric pressure as you increase your elevation above sea level?
• What would you expect to happen to the boiling point of water at high elevations?