Copyright© by Houghton Mifflin Company. All rights reserved. Chapter 14 Liquids and Solids Copyright© by Houghton Mifflin Company. All rights reserved.

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 Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.1: Representations of the gas, liquid, and solid states. Copyright© by Houghton Mifflin Company. All rights reserved.

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) . Copyright© by Houghton Mifflin Company. All rights reserved.

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 Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.2: Intermolecular forces exist between molecules. Bonds exist within molecules. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.3: (a) Interaction of two polar molecules. (b) Interaction of many dipoles in a liquid. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.4: Polar water molecules. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.4: Hydrogen bonding among water molecules. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.5: The boiling points of covalent hydrides. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.6: Atoms with spherical electron probability. Copyright© by Houghton Mifflin Company. All rights reserved.

14.6: The atom on the left develops an instantaneous dipole. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.6: Instantaneous dipole on A induces a dipole on B. Copyright© by Houghton Mifflin Company. All rights reserved.

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. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. 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. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. 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 Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.17: The classes of crystalline solids. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.15: Sodium and chloride ions. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.18: Molecular representation of sodium chloride. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.18: A molecular solid. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.18: Molecular representation of diamond. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.19: The packing of Cl¯ and Na+ ions in solid sodium chloride. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.21: A representation of part of the structure of solid phosphorus. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.22: Molecular representation of steel. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.22: Molecular representation of brass. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. 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) Copyright© by Houghton Mifflin Company. All rights reserved.

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. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. 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 =2260 J/g @ 100°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? Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. 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! Copyright© by Houghton Mifflin Company. All rights reserved.

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. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.7: The heating/cooling curve for water heated or cooled at a constant rate. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Phase Diagram for H2O Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.8: Both liquid water and gaseous water contain H2O molecules. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.9: Microscopic view of a liquid near its surface. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.10: Behavior of a liquid in a closed container. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. 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. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.12: Water rapidly boiling on a stove. Copyright© by Houghton Mifflin Company. All rights reserved.

Figure 14.13: Bubble expands as H2O molecules enter. Copyright© by Houghton Mifflin Company. All rights reserved.

Copyright© by Houghton Mifflin Company. All rights reserved. Figure 14.14: The formation of the bubble is opposed by atmospheric pressure. Copyright© by Houghton Mifflin Company. All rights reserved.