Chapter 10: Liquids, Solids, and Phase Changes In-chapter exercises: 10.1 10.6, 10.11; End-of-chapter Problems: 10.26, 10.31, 10.32, 10.33, 10.34, 10.35, 10.36, 10.39, 10.40, 10.42, 10.44, 10.45, 10.66, 10.82, 10.84, 10.86 Kinetic Molecular Theory of Liquids and Solids phase (=physical state): solid, liquid, or gas Solids have the lowest kinetic energy (KE) i.e. do not move very much Highest attraction between particles particles are stuck in specific sites = very confined Liquids have slightly higher KE i.e. particles moving more than in solid Particles are still attracted and maintain contact with one another but can move past one another particles are less confined Gases have greatest KE i.e. particles move quickly and randomly Attractive forces almost (if not) completely overcome, so particles can fly freely within container particles are far away from each other = unrestricted 10.2 INTERMOLECULAR FORCES intermolecular forces: Attractive forces between 2 molecules e.g. between 2 water molecules Ion-Dipole Forces Attraction between an ion and the oppositely charged end of a polar molecule e.g. between Na + and the negative end of a H 2 O molecule (O in H 2 O) or between Cl and the positive end of a H 2 O molecule (H atoms in H 2 O) CHM 150 Chapter 10 McMurry-Fay Notes page 1 of 14
Dispersion (London or Induced-Dipole) Forces Attraction between temporary or induced dipoles in adjacent molecules Electrons constantly shift and can sometimes concentrate in one region instantaneous dipole that goes away once electrons shift again Chocolate-chip cookie dough analogy Most common type and weakest of intermolecular forces, found between all types of molecules Only type of intermolecular force between nonpolar molecules In general, larger molecules w/ more electrons are more polarizable polarizable = tendency to experience electron shifts that result in charges the larger the molecule, the stronger the dispersion forces Dispersion (Induced-dipole) Forces Electrons shift in one molecule and concentrate on one side temporary dipole (light area = +ve; dark area = ve) The temporary dipole causes electrons to shift in adjacent molecules another temporary dipole When electrons shift again, the temporary dipoles go away. Dipole-Dipole Forces: Attraction between polar molecules generally stronger than dispersion forces because attraction is due to permanent dipoles rather than temporary dipoles Note: Van der Waals forces refer to intermolecular forces due to either London dispersion forces or dipole-dipole forces. Hydrogen Bonds: CHM 150 Chapter 10 McMurry-Fay Notes page 2 of 14
Exist between molecules with following bonds: H F, H O, H N Special type of dipole-dipole force caused by small radii and large electronegativity differences between H and O, N, and F atoms strongest type of intermolecular force Responsible for the relatively high melting and boiling points for water, bending and twisting in proteins, DNA, and other important biological molecules Note: Hydrogen bonds are the strongest type of intermolecular force but ionic and covalent bonds are stronger than hydrogen bonds. How to determine type of intermolecular forces involved: yes hydrogen bonding Is the molecule polar or nonpolar polar nonpolar Are there H F, H O, or H N bonds dispersion (London) forces no dipole-dipole forces Example: For each of the following, identify the type of bond holding atoms together in the molecules and the type of intermolecular forces between the molecules. A B A: B: CHM 150 Chapter 10 McMurry-Fay Notes page 3 of 14
A B A: B: Ex. 2: Consider the following six choices: A. ionic bond D. dispersion (induced-dipole) forces B. polar covalent bond E. dipole-dipole forces C. nonpolar covalent bond F. hydrogen bond Identify the type of bond or intermolecular force described for each below: a. The C C bond in C 2 H 6. b. The bonds in NaCl. c. The bonds holding HBr molecules together. d. The bonds holding atoms together in a water molecule. e. The bonds holding two NH 3 molecules together. f. The bonds holding atoms together in a HF molecule. g. The bonds broken when KBr dissolves in water. h. The bonds formed when KBr dissolves in water. Ex. 3 Circle the molecule in each pair which experiences the stronger intermolecular forces: a. N 2 or NO b. H 2 S or H 2 O c. Cl 2 or Br 2 CHM 150 Chapter 10 McMurry-Fay Notes page 4 of 14
10.3 SOME PROPERTIES OF LIQUIDS 10.5 EVAPORATION, VAPOR PRESSURE, AND BOILING POINT evaporation: for a liquid to vaporize, the surface molecules must break the intermolecular forces with other molecules in the liquid at the boiling point, molecules have enough energy to break the intermolecular forces with other molecules and become gas Boiling Point: temperature where vapor pressure of liquid is equal to external pressure (usually atmospheric pressure) stronger intermolecular forces more energy is required to break intermolecular bonds between molecules in the liquid higher boiling point normal boiling point is the boiling point at a pressure of 1 atm e.g. water boils at 100 C at 1 atm, but it boils at ~95 C in Denver where the atmospheric pressure is ~0.85 atm Vapor Pressure: partial pressure exerted by gas molecules above the liquid varies for different liquids, varies for different temperatures more gas molecules higher vapor pressure The weaker the intermolecular forces more molecules can go from liquid to vapor higher vapor pressure In the examples below, liquid A has weaker intermolecular forces than B A B Viscosity: resistance of a liquid to flow for example, honey has high viscosity; water has low viscosity stronger intermolecular forces = stronger attraction higher viscosity CHM 150 Chapter 10 McMurry-Fay Notes page 5 of 14
Surface Tension: attraction between surface molecules in a liquid stronger intermolecular forces = stronger attraction surface molecules are held together more strongly higher surface tension Ex. 1: Water molecules experience hydrogen bonding and hexane molecules experience induced-dipole (or dispersion) forces. Which of the following statements are true/false? a. Water's intermolecular forces are weaker than hexane's. T F b. Hexane s vapor pressure is higher than water s. T F c. Hexane s boiling point is lower than water s. T F d. Water s surface tension is higher than hexane s. T F e. Water s viscosity is lower than hexane s. T F Ex. 2: Explain each the following in terms of intermolecular forces: a. Why O 2 's boiling point is -183 C while NO's boiling point is -151 C. b. Why N 2 's boiling point is -196 C while Br 2 's boiling point is 59 C. c. Why H 2 S s boiling point is 61 C and H 2 O s boiling point is 100 C. CHM 150 Chapter 10 McMurry-Fay Notes page 6 of 14
10.6 KINDS OF SOLIDS Crystalline Solids: Have an ordered arrangement extending over a long range different types of crystalline solids: molecular solids, covalent network solids, ionic and metallic solids. Molecular Solids: consist of molecules held together by intermolecular forces The Structure and Properties of Ice Ice is an example of a molecular solid. The hydrogen bonds between water molecules are responsible for many unusual properties of ice and water. The density of ice (d=0.917 g/cm 3 ) is less than the density of liquid water (d=1.00 g/cm 3 ). With all other substances, the solid is more dense than its liquid. The density of ice (d=0.917 g/cm 3 ) is less than the density of liquid water (d=1.000 g/cm 3 ) whereas for all other substances, the solid is more dense than its liquid. Because of the hydrogen bonds, the arrangement of water molecules in ice crystal has "holes" or empty space. When ice melts, the water molecules fill in the holes, so liquid water is more dense than ice. Note the hexagonal holes in the molecular-level image above. Snowflakes have hexagonal symmetry because of the hexagonal holes in the molecular-level arrangement of water molecules in ice! CHM 150 Chapter 10 McMurry-Fay Notes page 7 of 14
Ionic Crystals: lattice of metal & nonmetal ions e.g. NaCl, MgO, CaBr 2 3D network of ions held together by electrostatic attraction high melting points, hard and brittle conduct electricity when melted or dissolved in solution 10.10 STRUCTURE OF SOME COVALENT NETWORK SOLIDS Covalent Network Solids are covalently bonded atoms that form a large network of indefinite size. (a) Graphite is made up of covalently bonded carbon atoms that form layers of sp 2 hybridized carbon atoms. (b) Diamond is made up of covalently bonded carbon atoms that form such a network of sp 3 hybridized carbon atoms in 3D tetrahedral structure. Diamond is so hard because so many covalent bonds must be broken to break up the diamond crystal. CHM 150 Chapter 10 McMurry-Fay Notes page 8 of 14
Fullerenes are a class of covalently bonded carbon atoms, similar to graphite. (a) C 60 has a shape similar to a soccer ball and is often called a Buckyball. (b) Nanotubes consist of sheets of graphite rolled into tubes. up to ten times as strong as steel Amorphous Solids: solids lacking 3D arrangement of atoms Silica (SiO 2 ) makes up sand and quartz glass: optically transparent solid of inorganic materials cooled to a rigid but non-crystalline arrangement of Si-O bonded atoms called quartz glass CHM 150 Chapter 10 McMurry-Fay Notes page 9 of 14
10.7 Probing the Structure of Solids: X-Ray Crystallography X-ray diffraction: scattering of X-rays by units of a solid crystal we can construct electron density contour map, where maximum densities are near center of each atom can be used to determine the positions of nuclei bond lengths and bond angles in crystals Metallic Crystals: positive metal ions surrounded by a sea of electrons electrons are delocalized i.e. free to move around the entire metal electrons move freely throughout the metal resulting in good heat and electrical conductivity electrons act as a glue holding nuclei together, so shape of metal can be easily manipulated metals are malleable and ductile some metals will react with water but are never soluble in water or other solvents 10.4 Phase Changes: change from one physical state to another sublimation SOLID fusion freezing LIQUID vaporization condensation GAS deposition The Equilibrium Nature of Phase Changes dynamic equilibrium: rate of forward process is exactly equal to the rate of reverse process CHM 150 Chapter 10 McMurry-Fay Notes page 10 of 14
Liquid-Solid Equilibrium freezing: liquid solid melting (or fusion): solid liquid melting point: temperature at which solid and liquid phases coexist in equilibrium normal melting point is melting point at 1 atm Consider solid-liquid equilibrium of water and ice (at 0 C and 1 atm): ice water When ice cubes are placed into a glass of water, the ice cubes begin to melt, but some water between the ice cubes freezes, causing ice cubes to fuse molar heat of fusion ( H fus ): energy required to melt one mole of solid supercooling: a substance remains liquid even below its freezing point results when a liquid is cooled so rapidly that molecules don't have time to arrange themselves properly unstable condition; stirring or adding "seed" crystal causes solidification For example, supercooling can happen in your refrigerator. A bottle of carbonated soda can be supercooled, so when you pick it up, you give it enough energy to freezes instantly ice crystals form Liquid-Gas Equilibrium: liquid vapor evaporation (or vaporization): liquid vapor condensation: vapor liquid the process of a gas liquefying can result from two ways: 1. cooling sample of gas lower KE, and molecules start to aggregate to form small drops of liquid, eventually causing condensation 2. applying pressure to gas minimize space between molecules, so molecules are attracted to each other to form droplets, eventually causing condensation CHM 150 Chapter 10 McMurry-Fay Notes page 11 of 14
Solid-Gas Equilibrium: Consider the dynamic equilibrium: solid vapor For example, dry ice, CO 2 (s), sublimes at room temperature, completely skipping a liquid phase. sublimation: solid gas (with no liquid phase) deposition: gas solid (with no liquid phase) molar heat of sublimation ( H sub ): energy required to sublime one mole of solid Heat of Phase Transition heating-cooling curve: Shows the phase changes that occur when heat is added or removed from a sample Temperature ( C) Heat Added Draw a heating curve indicating the following: 1. Regions for solid only, liquid only, gas only, solid-liquid, liquid-gas 2. The relationship between melting point and phases present 3. The relationship between boiling point and phases present 4. Where the curve is flat, where the slope is positive CHM 150 Chapter 10 McMurry-Fay Notes page 12 of 14
10.11 Phase Diagrams summarize the conditions at which a substance exists as solid, liquid, or gas allow us to determine melting and boiling points at different external pressures Water s phase diagram is shown in Fig. 10.28 graph divided into three regions corresponding to each phase lines separating two regions indicate conditions when both phases exist triple point: point at which all three curves meet when all three phases exist in equilibrium with one another for water, at 0.0098 C and about 6.0 10-3 atm Phase Diagram for Water (Fig. 10.28) Phase Diagram for CO 2 (Fig. 10.29) CHM 150 Chapter 10 McMurry-Fay Notes page 13 of 14
Phase Diagram for Water versus that for CO 2 : Note that line separating solid and liquid phases has positive slope for CO 2 but negative slope for H 2 O for CO 2, the solid is more dense than the liquid, so increasing pressure converts the liquid to a solid for H 2 O, the solid is less dense than the liquid, so increasing pressure converts the solid to a liquid Critical Temperature and Pressure: critical temperature (T c ): above which its gas form cannot be made to liquefy, no matter how great the applied pressure highest temperature at which a substance can exist as a liquid intermolecular attraction is a finite quantity for given substance below T c, molecules are moving slowly enough to maintain contact above Tc, molecular motion so energetic that molecules will always break away from attraction critical pressure (P c ): minimum pressure that must be applied to liquefy sample at the critical temperature Given a Phase Diagram, be able to do the following: Determine what phase(s) is/are present at a given temperature and pressure Indicate the melting point or boiling point at a given pressure e.g. the normal melting or boiling point Describe what phase change occurs when temperature is changed at a constant pressure Describe what phase change occurs when pressure is changed at a constant temperature. CHM 150 Chapter 10 McMurry-Fay Notes page 14 of 14