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DISCLAIMER: these notes are provided to assist you in mastering the course material but they are not intended as a replacement of the lectures. Neither do they contain the comments and ancillary material of the lectures; they are just a set of points that you might bring to the lectures to annotate instead of having to write everything down and/or they may assist you in organizing the material after the lectures, in conjunction with your own notes. Chapter 5: Solids, Liquids & Phase Transitions Overview organization chart for both condensed phases forces properties phase changes 5.1 Bulk Properties of Phases & Kinetic Molecular Theory recall kinetic molecular theory (for gases): gas atoms or molecules widely-separated no forces of attraction between them atoms or molecules in continual, random, rapid motion kinetic energy determined by temperature first two above are unique to gases, latter two are largely true for liquids and solids too comparison of bulk properties, in terms of: molar volume of gases, much larger than liquids & solids compressibility of gases (to form liquids) large; not dramatic for liquids (to form solids) thermal expansion of gases much larger than liquids & solids fluidity & rigidity diffusion surface tension (Fig. 5.4; not for gases) 5.2 Types of Intermolecular Forces in order of decreasing strength, Fig. 5.9; also important for solutions (chapter 6) ion - dipole interactions (ion - ion interactions treated separately) dipole - dipole interactions dipole (or ion-)- induced dipole induced dipole - induced dipole interactions similar repulsive forces for all term: van der Waals forces used for those interactions not involving ions; define van der Waals radius of atoms based on optimum between attractive and repulsive forces when ions not involved Ion - Dipole Interactions (Fig. 5.6) dipole: charge separation, or (more likely) partial-charge separation, in a molecule i.e. - a polar molecule strength of ion - dipole interaction depends on: separation distance of ion and dipole charge on ion Chem 59-110 (’02) 2 magnitude of dipole several examples based on the polar molecule, water: Na + (g) + x H 2 O(l) [Na(H2 O)+ ]( aq); H rxn = - 405 kJ / mol x reaction is exothermic, heat given off = heat/enthalpy of hydration (hydration here; more generally, solvation) distance effect seen with alkali cations compare with H+, Hhydr = -1090 kJ/mol (i.e.- hydronium ion, H3O+) summary: Li+, -515; Na+, -405; Cs+, -263 kJ/mol similarly, charge and distance effect seen Dipole - Dipole Interactions (Fig. 5.5) energy generally released when molecules condensed, taken up when the condensed phase vaporized, i.e.- accompanying the equilibrium: gas liquid for polar molecules, this is due to dipole - dipole interactions comparison of boiling points (a measure of the heat required for vaporization) allows categorization as polar or non-polar molecules (compare pairs of similar molar masses) solubility considerations are also due to a matching of polarities of solute and solvent (“like dissolves like”) Hydrogen Bonding special class of dipole - dipole interactions due to small size and low electronegativity of H when bonded to small, electronegative atoms, especially N, O and F eg. X-- H+Y- the strongest are given in a Table (weaker ones with Cl, S, etc.) example, HF compare ethanol and dimethyl ether, both C2H6O: compare structures H-bonding in the alcohol, only dipole - dipole forces in the ether physical data: dipole moment melting pt. boiling pt. o 1.69D -114 C 78oC 1.30 -142 -25 periodic trends in Fig. 5.10 Unusual Properties of Water H-bonding to extreme!! two H’s and two lone pairs (non-bonding pairs of electrons) on each O, Fig. 5.11 form networks in the condensed phases perfect “diamond lattice” in ice, Fig. 5.12 some disruption upon melting (still 85% of H-bonds), contraction of structure hence, ice has lower density than water; water max density at 4C, Fig. 5.13 Dispersion Forces: Interactions with Induced Dipoles weakest of all intermolecular interactions Chem 59-110 (’02), ch. 5, Solids, Liquids & Phase Transitions 3 two types: between polar and non-polar molecules, Fig. 5.7 (dipole (or ion-)- induced dipole) the larger the non-polar, the greater the interaction (eg. solubility of diatomic gases in water), due to polarizability between non-polar molecules, Fig. 5.8 (induced dipole - induced dipole) groups of examples in Table, trends in boiling points Summary/Decision Tree for Intermolecular Interactions - Fig. Not in text 5.4 & 5.5 Phase Equilibria & Transitions Physical Properties of Liquids distribution of kinetic energies of molecules in a liquid sample (similar to kinetic theory of gases) equilibrium with gas phase (part of Fig. 5.17): vaporization, Hvap liquid vapor condensation, Hcond heat/enthalpy of vaporization, Hvap, is energy required (i.e.- endothermic) to escape intermolecular forces, Fig. 5.4 eg. for water, Hvap = + 40.7 kJ/mol heat/enthalpy of condensation, Hcond, is energy liberated (i.e.- exothermic) on forming intermolecular interactions (same magnitude, opposite sign) Vapor Pressure in a closed space above a liquid, Fig. 5.14, pressure in the gas phase stabilizes at a fixed value = equilibrium vapor pressure, dependent on temperature, Fig. 5.15 and Table 5.1; note: points on a line represent equilibrium pressure also a partial phase diagram - at a given T and P points to left of line represent the liquid phase region points to right of line represent the gas phase region boiling points at atmospheric pressure (see below) volatility, the tendency to escape into the gas phase ranked according to equilibrium vapor pressure practical application: “water pump” in lab better in winter than summer Boiling Point Fig. 5.15, line at 1 atm equilibrium vapor pressure equals atmospheric pressure at the boiling point in an open vessel vaporized molecules can escape note dependence on pressure, applications: cooking in Salt Lake City vacuum distillations Chem 59-110 (’02), ch. 5, Solids, Liquids & Phase Transitions 4 Physical Properties of Solids disrupt the lattice to form a liquid (part of Fig. 5.17): fusion, Hfusion solid liquid freezing, Hcryst. fusion is endothermic, freezing/crystallization is exothermic also characterized by melting point, lowest temperature at which fusion occurs grouped according to forces; note ion-size dependence (note: sublimation also possible for solid gas equilibrium; egs. H2O, I2, CO2, naphthalene; I2 on 1st page photo) 5.6 Phase Diagrams phase transitions, above influence of temperature and pressure on phases given in diagrams eg. H2O in Fig. 5.19: follow 2-phase lines (note negative slope to solid-liquid line); note triple point, freezing point, boiling point, critical point eg. CO2 in Fig. 5.21: note positive slope to solid-liquid line, critical point Critical Temperature and Pressure keep increasing temperature in Fig. 5.19 and 5.21, two phases coalesce critical temperatures and pressures (both minima) shown new “phase”: supercritical gas/fluid has a density like that of the liquid but flow properties and ability of molecules to be separate from one another like a gas applications: liquefaction of gases, eg. air conditioning, fuels (must be below critical point) supercritical fluid extraction (must be above critical point), eg. CO2 (Tc = 31C, Pc = 73 atm) used for decaffeinating coffee Surface Tension, Capillary Action & Viscosity all are phenomena due to intermolecular interactions surface tension forces different in bulk liquid than at surface, Fig. 5.4; net, inward force at surface “skin” on surface, resists spreading as a film on another surface = surface tension surface layer interactions, in some cases, counteracted by interaction with another material eg. H2O with glass (H-O-H vs. Si-O-(H)), hence meniscus in a tube (Fig. 5.20), which is extreme in a very narrow tube (capillary) application: chromatography bulk liquid flow influenced by intermolecular interactions viscosity is the resistance to flow, which increases as the intermolecular interactions do eg. ethanol (two C’s) compared to longer chain alcohol, octanol (eight C’s) or to a “polyol” such as glycerol Suggested Problems odd, 1 – 7; 13 – 17; 21 - 47 Chem 59-110 (’02), ch. 5, Solids, Liquids & Phase Transitions
"Chapter 13 Intermolecular Forces in Liquids & Solids"