Chapter 10 The Liquid and Solid States

Chapter 10 The Liquid and Solid States 10-1 Questions for Consideration 1. 2. 3. 4. How do the properties of liquids and solids differ, and what happens when substances undergo a change of state? Why do atoms or molecules hold together as liquids or solids instead of existing only as gases? What are some unique properties of liquids? How do atoms or ions fit together to make solids, and how does their arrangement affect their properties? 10-2 Chapter 10 Topics: 1. 2. 3. 4. Changes of State Intermolecular Forces Properties of Liquids Properties of Solids 10-3 Introduction   How do the properties of liquid and solid substances differ? How can we predict properties based on molecular-level structure? Figure 10.3 Figure 10.1 10-4 10.1 Changes of State Solids, Liquids, and Gases: How do they differ? Figure 10.4 10-5 General Properties of the States of Matter TABLE 10.1 General Properties of the States of Matter Solid Liquid Gas Fixed shape No fixed shape No fixed shape Shape not set by container Takes shape of filled portion of the container Takes shape of container Shape remains rigid Can be poured Fills container Particles fixed in place but vibrate around a fixed position Particles move past one another Particles move through space Little or no volume change under moderate pressure Can be compressed slightly by moderate pressure Compressed under moderate pressure Little free space between particles Some free space between particles Particles are widely separated with much free space 10-6 Physical State of a Substance Figure 10.5 Aluminum melts above 2467ºC. Figure 10.6 Water exists in all three physical states under normal conditions. 10-7 Liquid-Gas Phase Changes Evaporation, or vaporization  Evaporation occurs because some surface liquid molecules have high enough kinetic energy to escape to the gas phase.  Heat is required to maintain the temperature needed for evaporation.  Endothermic Figure 10.7 Figure 10.8 10-8 Liquid-Gas Phase Changes Condensation     Transition from a gas to a liquid Occurs when gas particles cannot escape the container and thus, come into contact with a liquid An exothermic process (energy is released) The reverse of evaporation Figure 10.7 10-9 Liquid-Vapor Equilibrium  A liquid placed in a closed container will evaporate only until an equilibrium is reached.  The pressure exerted by the vapor at equilibrium Figure 10.9 is the vapor pressure. 10-10 Vapor pressure depends on molecular structure and temperature. Figure 10.10 10-11 Boiling Point    Boiling occurs when the vapor pressure equals the external pressure of the atmosphere. The temperature at which vapor pressure equals the external pressure of the atmosphere is the boiling point. If the external pressure is exactly 1 atm, the boiling temperature is called the normal boiling point. Figure 10.10 10-12 Liquid-Solid Phase Changes The temperature at which the liquid and solid states coexist in equilibrium is called the freezing point or the melting point. Figure 10.11 10-13 Melting   When ice melts, the water molecules stay intact but flow past one another. When the ionic compound NaCl melts, ionic bonds are broken and ions flow freely. Figure 10.12 10-14 Solid-Gas Phase Changes Under certain conditions, substances can go directly from the solid state to the gas state: sublimation. CO2 and I2 do this under normal conditions. Some substances can go directly from the gas state to the solid state: deposition. Figure 10.13 10-15 Activity: Changes in the State of Matter Identify the process shown in the following diagram: Figure from p. 405 10-16 Heating Curve    Heat is added to cause phase changes. Notice that temperature remains constant during a phase change, even though heat is being added. Skip the “Energy Changes” subsection (pp. 407-410) Figure 10.14B 10-17 10.2 Intermolecular Forces    Why is CO2 a gas at room temperature and H2O a liquid? Why do liquids have different boiling points? What causes molecules to stick together when in the liquid state?  Attractive forces, called intermolecular forces, exist between molecules. 10-18 Intermolecular Forces Three types of intermolecular forces:    London dispersion forces Dipole-dipole forces Hydrogen bonding 10-19 London Dispersion Forces    London dispersion forces exist in all substances. It is the only intermolecular force present in nonpolar substances. London dispersion forces are a result of temporary electron cloud distortions and temporary polarity (dipoles). Figure 10.15 10-20 London Dispersion Forces Figure 10.16 The larger the electron cloud, the more it can be distorted.  The larger the molecule, the stronger the London dispersion forces. Figure 10.15 10-21 Activity: London Dispersion Forces For each pair, determine which has the stronger London dispersion forces. 1. 2. 3. 4. He or Kr HCl or HBr CH4 or C2H6 N2 or O2 N2 O2 b.p. = 77.4 K b.p. = 90.2 K 10-22 London Dispersion Forces   Relative boiling points tell us about intermolecular force strength. The stronger the forces, the higher the temperature needed to overcome them. Figure 10.19 10-23 Dipole-Dipole Forces Dipole-dipole forces exist between polar molecules.   A polar molecule has a positive and negative end, so it is considered a dipole. SO2 is polar so it experiences dipole-dipole forces in the liquid state, in addition to London dispersion forces. Figure 10.18 10-24 Activity: Dipole-Dipole Forces Which of the following molecules experience dipole-dipole forces? Figure from p. 341 SO2 CH2Cl2 PCl3 BF3 10-25 Boiling Points Figure 10.20 10-26 Hydrogen Bonding   The irregularity in boiling point trends in Figure 10.20 is due to a particularly strong dipole-dipole force called hydrogen bonding. Hydrogen bonding occurs between molecules that contain hydrogen covalently bonded to either nitrogen, oxygen, or fluorine: N–H, O–H, or F–H   Hydrogen bonding causes boiling points to be much higher than expected. An important force in living systems, where it stabilizes molecular shapes 10-27 Hydrogen Bonding (shown by the dotted lines) Figure 10.21 10-28 Hydrogen bonding is the strong force that holds together the DNA double helix. Figure 10.22 10-29 Hydrogen Bonding Hydrogen bonding in H2O causes ice to be less dense than the liquid state. Figure 10.24 Figure 10.11 10-30 Activity: Hydrogen Bonding Which of the following hydrogen-bond in the pure liquid state? 1. 2. 3. 4. 5. CH3NH2 NF3 CH3-O-CH3 CH3CH2OH HCl 10-31 Summary of Intermolecular Forces Figure 10.23 10-32 Intermolecular Forces and Bonds Table 10.2 Intermolecular Forces and Bonds Type of Force Type of Interaction Occurrence London dispersion A temporary dipole in one molecule force induces the formation of a temporary dipole in a nearby molecule and is attracted to it. All atoms and molecules Dipole-dipole force Polar molecules Polar molecules (permanent dipoles) attract one another. Hydrogen-bonding Two dipoles, one containing hydrogen force bonded to an electronegative element and the other containing an electronegative element, attract one another. Polar molecules containing unpaired electrons and a hydrogen bonded to nitrogen, oxygen, or fluorine Covalent bond Nuclei of two atoms attract the electrons shared between them. Nonmetal-nonmetal compounds Ionic bond Cations and anions attract one another. Metal-nonmetal compounds 10-33 Activity: Trends in Boiling Point Which of the following should have the highest boiling point? 1. 2. 3. CH3OH or CH3SH I2 or SO2 CO or N2 10-34 10.3 Properties of Liquids      Vapor Pressure Density Viscosity Surface tension Capillary action 10-35 Properties of Liquids    Are related to the distance between particles and to intermolecular forces. Particles in a liquid are much closer together than the particles in a gas. Liquid particles are not fixed, as they are in a solid. 10-36 Trends in Vapor Pressure How is vapor pressure related to intermolecular force strength? Figure 10.10 10-37 Activity: Trends in Vapor Pressure Which of the following should have the highest vapor pressure (at a given temperature)? 1. CH3OH or CH3SH 2. I2 or SO2 3. CO or N2 10-38 Meniscus A meniscus is either a concave or convex curved surface of a liquid produced by intermolecular forces.   A concave surface occurs when the intermolecular forces between the liquid and the glass are greater than the intermolecular forces among the liquid molecules. A convex surface occurs when the intermolecular forces between the liquid and the glass are less than the intermolecular forces among the liquid molecules. Figure 10.30 10-39 10.4 Properties of Solids The particles in the solid state are held together in specific positions – there is no translational motion.   Crystalline solids are well ordered. Amorphous solids lack regular form. Figure 10.31 10-40 Copper Atoms in 2 Dimensions (Close Packing) Figure 10.32 10-41 Close-Packing in Nature and Commerce Pineapple Golf Ball Wasp Nest Fly Eye Bubble Wrap Chicken Wire 10-42 Close Packing in 3 Dimensions Figure 10.34 10-43 Types of Crystalline Solids (See Table 10.3)  Metallic    Ionic    Ionic compounds; composed of ions Ionic bonding Molecular    Pure metals and alloys; composed of atoms Metallic bonding Molecular compounds or nonmetal elements; composed of molecules or atoms Intermolecular forces Network   Molecular compounds or nonmetal elements; All atoms are connected by covalent bonds 10-44 Classification of Solids Table 10.3 Classification of Solids Type of Solid Fundamental Particles Attractive Forces Properties Examples Metallic Atoms Attractions between nuclei and delocalized valence electrons Low melting point and soft, or high melting point and hard; good heat and electrical conductors; malleable and ductile Hg, Na, Fe, Au, Ag, Cu, Cr, alloys such as Nitinol Ionic Cations and anions Ionic bonds High melting point, hard, brittle; nonconductors when solid; electrical conductors when melted NaCl, CsCl, CaF2, CaCO3, KNO3, Ca5(PO4)3OH, YBa2Cu3O7 Molecular Polar molecules Dipole-dipole forces Low to moderate melting point., variable hardness, may be brittle; nonconductors H2O, NH3, SO2, CH3CO2H, CH2Cl2 Nonpolar molecules London dispersion forces Low melting point, soft; poor heat conductors; electrical insulators CH4, O2, CCl4, CO2 C6,H6, I2, S8 Atoms Covalent bonds Very high melting point, very hard, somewhat brittle; nonconductors or semiconductors C (diamond, SiC (carborundum), SiO2 (silica), Al2O3 (alundum), BN (boron nitride) 10-45 Network Metallic Solids    Are ductile and malleable Valence electrons move freely through all parts of a metal. Attractions between atoms of a metal are delocalized, and therefore it is easy to move atoms by applying force. Figure 10.35 10-46 Ionic Solids      Contain cations and anions arranged in crystalline solids Electrostatic forces (ionic bonds) hold together ionic crystals. High melting points Hard and brittle Solids are not electrical conductors, but melted or dissolved, they become good conductors. Figure 10.37 10-47 Molecular Solids CO2 Figure 10.40 10-48 Network Solids         Consists of a giant molecule that forms the entire crystal Formed by metalloids or carbon Strong covalent bonds connect the atoms in a network solid Poor electrical conductors High melting points Very hard Have very stable, three– dimensional structures Many have a diamond structure, or some derivative of it Figure 10.41 10-49 Network Solids The Diamond Form of Carbon Figure 10.41 10-50 Network Solids Silicon dioxide, very high melting point Figure 10.42 10-51 Network Solid in 2 Dimensions Graphite Form of Carbon Figure 10.43 10-52 Activity: Types of Solids Predict the type of solid for each: 1. 2. 3. 4. CO2 sublimes at -78.4ºC Br2 mp = -7.2ºC C (diamond) mp = 3550ºC MgO mp = 2800ºC 10-53 Activity: Structures of Solids Identify the type of solid shown by the following molecular-level image. What types of forces hold the particles together in this solid? Figure from p. 431 10-54