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Chapter 14
Liquids and Solids
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2. 14-1 Condensed States of Matter
Solids and liquids are referred to as condensed states of matter because they have much higher densities than gases.
Kinetic Molecular Theory applies to liquids and solids as well as gases.
Particles of matter are in constant motion
The collisions between them are perfectly elastic
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3. Figure 14.1: Representations of the gas, liquid, and solid states.
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4. Properties of Solids Liquids and Gases
Gas Liquid Solid
Highly compressible Slightly Compressible Slightly Compressible
Low Density High Density High Density
Fills Container Definite Volume Rigid-Definite Volume
Assumes Shape of container Assumes Shape of container Retains Own Shape
Rapid Diffusion Slow Diffusion Very Slow Diffusion
High Expansion on Heating Low Expansion on Heating Low Expansion on Heating
According to Kinetic-Molecular Theory a substances state at a given temperature depends on the strength of attraction between its particles (atoms, ions, molecules) .
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5. Intermolecular Forces
Forces between molecules of the same substance.
Dipersion Forces – (Van der Wals) Interactions between the electrons of closely spaced molecules can cause induced dipoles that cause the molecules to stick together. Large molecules have stronger dispersion forces due to more electrons.
Dipole-Dipole Forces – Polar molecules will be attracted together by the oppositely charged ends of the molecule. These are stronger than dipersion forces.
Hydrogen bonding – A special kind of dipole-dipole force that occurs when hydrogen is bonded to a highly electronegative atom. Water
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Figure 14.2: Intermolecular forces exist between molecules. Bonds exist within molecules.
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Figure 14.3: (a) Interaction of two polar molecules. (b) Interaction of many dipoles in a liquid.
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8. Figure 14.4: Polar water molecules.
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9. Figure 14.4: Hydrogen bonding among water molecules.
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10. Figure 14.5: The boiling points of covalent hydrides.
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11. Figure 14.6: Atoms with spherical electron probability.
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12. 14.6: The atom on the left develops an instantaneous dipole.
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13. Figure 14.6: Instantaneous dipole on A induces a dipole on B.
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14. 14-2 Properties of Liquids
Properties of liquids are determined mainly by the nature and strength of the intermolecular forces present.
Viscosity is the resistance of a fluid to flow.
Depends on the strength of intermolecular attractions
Decreases with increasing temperature
Surface Tension – The imbalance of forces at the surface of a liquid that creates a “surface film”.
At the surface molecules experience downward and sideways attractions but no upward forces.
Water has very unique properties due to hydrogen bonding :
high boiling point, high specific heat, high density, solid less dense than liquid, high surface tension, high heat of vaporization, universal solvent.
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14-3 The Nature of Solids
Crystalline Solids-have their particles arranged in a highly ordered repeating pattern
Unit Cell – A small repeating pattern in crystalline solids. NaCl has a cubic unit cell.
Some ionic solids can incorporate water into their crystals, this is called water of hydration and the salt is a hydrate.
Amorphous Solids – sometimes called “super-cooled liquids” because they lack the crystalline form of true solids. Glass & Rubber
Bonding in solids determines their physical properties such as hardness, conductivity, melting point, and hardness.
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Bonding in Solids
Solid Type Particle Type Forces b/w Particles Properties Examples
Metallic atoms metallic bond soft to hard Fe, Cu
good conductors Au, K
high to low melting pt.
malleable & ductile
Molecular molecules hydrogen bonds soft, low melting pt. CH4
dispersion forces poor conductors H2O
dipole-dipole C6H12O6
Ionic ions ionic bonds hard, brittle, high m.pt. NaCl
poor conductors unless molten KBr
Network-Covalent covalent bonds very hard, very high m.pt. C
atoms usually poor conductors diamond
quartz
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17. Figure 14.17: The classes of crystalline solids.
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18. Figure 14.15: Sodium and chloride ions.
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19. Figure 14.18: Molecular representation of sodium chloride.
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20. Figure 14.18: A molecular solid.
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21. Figure 14.18: Molecular representation of diamond.
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Figure 14.19: The packing of Cl¯ and Na+ ions in solid sodium chloride.
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Figure 14.21: A representation of part of the structure of solid phosphorus.
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24. Figure 14.22: Molecular representation of steel.
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25. Figure 14.22: Molecular representation of brass.
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14-4 Changes of State
Changes of state are the conversions of one phase of matter to another. They always involve energy changes but NOT temperature changes (melting/freezing condensation/vaporization sublimation/deposition)
State changes either absorb energy to break bonds between molecules (melting, vaporization, sublimation) or release energy as new bonds form between molecules (freezing, condensation, deposition)
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27. Vaporization and Condensation
Liquid to gas is vaporization
Evaporation
Boiling
Condensation is gas to liquid
Equilibrium vapor pressure depends on temperature only. The curve DC above shows the temperatures and pressures in which water is in equilibrium with its vapor .
When equilibrium vapor pressure of water at a given temperature equals the atmospheric pressure the normal boiling point is reached.
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Heat of Vaporization
The heat required per unit mass to vaporize a liquid at its boiling point is the heat of vaporization (latent heat of vaporization) LV.
LV = °C
How does this compare to the heat of fusion for water?
Why is it so high?
Where is this energy going when water boils?
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Heat of Fusion
The amount of heat needed to melt a unit mass of a substance at its melting/freezing point is called the heat of fusion (latent heat of fusion) LF.
For ice LF=334 J/g
This amount of heat energy is used to break the bonds in the solid and is converted to potential energy of the liquid molecules.
This is a constant temperature process so Kinetic energy of the molecules DOES NOT CHANGE!
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30. Sublimation and Deposition
Sublimation is the conversion of solid directly into gas. Dry Ice to CO2.
Deposition is the formation of solid directly from a gas.
Heat of sublimation and deposition is the sum of heat of fusion and heat of vaporization.
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Figure 14.7: The heating/cooling curve for water heated or cooled at a constant rate.
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Phase Diagram for H2O
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Figure 14.8: Both liquid water and gaseous water contain H2O molecules.
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34. Figure 14.9: Microscopic view of a liquid near its surface.
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35. Figure 14.10: Behavior of a liquid in a closed container.
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Figure 14.11: (a) Measuring vapor of a liquid by using a simple barometer (b) The water vapor pushed the mercury level down (c) Diethyl ether shows a higher vapor pressure than water.
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37. Figure 14.12: Water rapidly boiling on a stove.
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38. Figure 14.13: Bubble expands as H2O molecules enter.
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Figure 14.14: The formation of the bubble is opposed by atmospheric pressure.
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