Chapter 11: Intermolecular Forces, Liquids, and Solids 1l.1 A Molecular Comparison of Liquids and Solids Gases characterized by rapidly moving, widely-spaced particles Solids characterized by a regular array of closely-spaced, fixed particles Liquids are somewhere in-between, but with some special properties of their own Why do liquids and solids exist? Why is all matter not in the gaseous state? States of Matter States of Matter The forces holding solids and liquids together are called intermolecular forces. In gases, the velocity of the gas particles allows them to overcome the intermolecular forces. 11.2 Intermolecular Forces Intermolecular forces are weaker than bonds, but have profound effects on the properties of liquids: Polar liquids have higher boiling points and higher heats of vaporization than non-polar liquids. Polar liquids dissolve ionic solids and polar liquids. 431 kJ/mol 16 kJ/mol Dipole-Dipole Forces Permanent dipoles attract one another Dipole-Dipole Forces Pure substance or mixture Strength of these forces depends on the dipole moment of the molecules Dipole-Dipole Forces Which member of each set has the stronger dipole-dipole forces? SiCl4, SiHCl3, SiHF3 CO2, SO2 NF3, PF3 London Dispersion Forces London dispersion forces are present in all molecules; temporary partial charges give temporary polarity Strength increases with polarizability London Dispersion Forces Only force in non-polar molecules and in unbonded atoms London forces become stronger, the larger the atom or molecule (larger electron clouds are easier to deform) Found in mixtures or pure substances. London Forces The magnitude of London forces explains why Cl2 is a gas, Br2 is a liquid, and I2 is a solid. Explain the order of the boiling points of the halogens and noble gases. Boiling Points Boiling point increases with the size of molecules because of increases in London forces with larger electron clouds. Figure 10.3 London Dispersion Forces Which member of each pair has the stronger London forces? Ne, Kr F2, Cl2 CH4, SiCl4 N2, O2 N2 b.p. = 77.4 K O2 b.p. = 90.2 K Explain Boiling Pt Trend Hydrogen Bonding Hydrogen bond is an especially strong dipole-dipole force, as shown by the trend in boiling points of polar molecules Hydrogen Bonding H-bonding is observed for HF, H2O, NH3, but not CH4 Conditions for occurrence: H attached to a small, highly electronegative element in one molecule Small, highly electronegative element with one or more unshared electron pairs in the other molecule Observed for the elements: F, O, N (rarely S and Cl) Group Work Which of the following will hydrogen-bond in the pure liquid? CH3OH CH3SH H2 CH3NH2 CH4 Group Work Which of the following will hydrogen-bond in the pure liquid? CH3OH CH3SH H2 CH3NH2 CH4 Hydrogen Bonding in Liquid Water H points at the electron pair on the atom in the other molecule In liquid water, each water molecule is surrounded by an average of 4 other water molecules; structure is not rigid. Longer than covalent bond. Hydrogen Bonding Molecules hydrogen-bond to themselves or to other molecules. Figure 10.2 Structure of Ice The water molecules in ice are fixed into a tetrahedral arrangement as a result of hydrogen bonding. Open structure makes ice less dense than water. Structure of Ice The open structure of ice leaves channels of empty space through the crystals. Hydrogen Bonding is critical to the structure of DNA Summary of I.M. Forces in Pure Substances All substances have London-Disp. Forces LD forces increase with MW (size) Polar molecules also have Dipole-Dipole forces Some molecules have extra strong dipole-dipole forces called Hydrogen bonding. (-O-H, -N-H, H-F) Explains: H2O >> H2S < H2Se < H2Te Intermolecular Forces What types of intermolecular forces are observed in liquids composed of each of the following? (A molecule may have more than one.) H2O HF HBr NH3 PF3 CH3OH F2 CO CO2 Trends in Intermolecular Forces Which member of each pair has the stronger intermolecular forces (higher boiling point)? CH3OH, CH3SH F2, Kr F2, CO CO, HF CO2, NH3 N2, NH3 Trends in Intermolecular Forces Which member of each pair has the larger intermolecular forces (higher boiling point)? CH3OH, CH3SH F2, Kr F2, CO (-188?C, -191.5?C) CO, HF (-191.5?C, +50?C) CO2, NH3 N2, NH3 Group Work Which member of each pair has the larger intermolecular forces (higher boiling point)? CH4, CH3CH3 NH3, PH3 CO2, SO2 CO, CO2 I2, Cl2 Group Work Which member of each pair has the larger intermolecular forces (higher boiling point)? CH4, CH3CH3 NH3, PH3 CO2, SO2 CO, CO2 I2, Cl2 11.3 Some Properties of Liquids In addition to boiling point and heat of vaporization, there are some unique properties of liquids: Viscosity Surface Tension Capillary Action Vapor Pressure Heat of Vaporization Boiling Point Viscosity Viscosity - resistance to flow Values are proportional to size and intermolecular forces Water is about average (strong intermolecular forces, but small size) How does viscosity change with T? Why? Surface Tension Surface tension is the amount of energy required to increase the surface area by a given amount. Molecules must break IM forces in order to move to the surface and increase the surface area. Surface Tension Why are water droplets spherical? Surface Tension How does a water strider stay on the top of the water? Why does the needle float? Capillary Action Why do liquids rise in a capillary? Capillary Action Competing forces: Between liquid molecules Between liquid molecules and glass H2O-Glass > H2O-H2O Hg-Glass < Hg-Hg Height determined by net liquid-glass forces opposing gravity Figure 10.8 Capillary Action Which forces in each system are greater? 11.4 Phase Changes Physical states of a substance can co-exist under a variety of conditions of pressure and temperature. Phase Changes Transitions between phases are called phase changes evaporation: liquid ? gas (reverse = condensation) melting: solid ? liquid (reverse = freezing) sublimation: solid ? gas (reverse = deposition) What are the phase changes? Heat of Vaporization Is evaporation exothermic or endothermic? Patio misters Evap. Coolers Refrigerators Heat of Vaporization DHvap = heat needed to evaporate 1 mol of liquid at constant temperature Energy used to overcome intermolecular forces during evaporation DHvap ? Intermolecular Force Strength Heat of Fusion Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: it takes more energy to completely separate molecules, than partially separate them. For Water: ?Hfus = 6.01 kJ/mol ?Hvap= 40.67 kJ/mol Heating a Pure Substance such as Water Why does T become constant during melting and evaporating? How is energy used for these processes? 2.02 J/goC; 36.39 J/mol oC 6009.5 J/mol 40,700 J/mol 4.184 J/goC; 75.37 J/mol oC 2.02 J/goC; 36.39 J/mol oC 11.5 Vapor Pressure Evaporation: loss of higher kinetic energy molecules, so the liquid cools (unless energy is supplied) The process is endothermic. Evaporation Open container: evaporates completely Closed container: reaches a state of equilibrium Vapor Pressure Equilibrium: rate of evaporation = rate of condensation P at equilibrium = vapor pressure Vapor pressure varies with temperature and intermolecular forces Vapor Pressure When Pvap = Patm, T = boiling point (Pvap ) When Pvap = 1 atm, T = normal boiling point In Phoenix, boiling point of H2O = 99oC At sea level, b.p. of H2O = 100oC At 9000 ft elevation, b.p. of H2O = 91oC (needs a pressure cooker to speed up cooking) 11.6 Phase Diagrams Phase diagram: plot of pressure vs. temperature summarizing all equilibria between phases. Given a temperature and pressure, phase diagrams tell us which phase will exist. LINES: equilibrium between two phases AREAS: only one phase is stable TRIPLE PT (confluence of 3 lines): equilibrium between three phases Generic Phase Diagram Phase Diagram Features Features of a phase diagram: Triple point: temperature and pressure at which all three phases are in equilibrium. Vapor-pressure curve: generally as pressure increases, temperature required to boil increases. Critical point: critical temperature and pressure for the gas. (min. T when gas cannot be liq. by P) Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. Normal melting point: melting point at 1 atm. Phase Diagram of Water Normal Boiling Pt Normal Melting Pt Condensation Sublimation What happens when you increase pressure at 0?C? Phase Diagrams for Water and Carbon Dioxide 11.7 Structures of Solids Definite and rigid shapes and sizes Not very compressible Do not flow Two types of solids amorphous (without form) crystalline (orderly arrays) Crystals Crystal - an orderly, repeating, 3-dimensional array of particles Lattice - pattern formed by the array Unit cell - smallest repeating unit that reproduces the lattice Figure 10.12 Cubic Crystal Systems Seven crystal systems Cubic is the most common Three forms of cubic crystals: simple body-centered face-centered How many atoms in a unit cell? Coordination Number Number of nearest neighbors of an atom or ion is called its coordination number Metal Structures Most metals have a structure that involves closest-packing, the most efficient use of space One layer as close as possible; the next layer fills in the depressions in the first layer Structure of Gold What is the coordination number of gold in one layer? Close Packing Two ways to close pack spheres: Hexagonal Closest Packed (b) Cubic Closest Packed (c) Cubic Closest Packed ABCABC? If 2nd layer is in B, third can be in C or A. This layer is in C Cubic Closest Packed: ABCABC... Cubic Closest Packed CCP is the same as the face- centered cubic structure Hexagonal Closest Packed Figure 10.22 Structures of Ionic Crystals Many can be described as a simple cubic or face-centered cubic array of one ion with the second fitting into holes in the structure. Closest packed structures use 74% of the volume. The other 26% is available for occupation by another particle. The structure of an ionic crystal is often dictated by the size of the holes in the structure, which must match the size of the ion to fit into the holes. Holes in Closest Packed Arrays Tetrahedral holes are surrounded by four spheres. Octahedral holes are surrounded by six spheres. Tetrahedral and Octahedral Holes Ionic Crystal Structures Two different representations of crystal structure. Which is more accurate? Figure 10.25 Sodium Chloride Structure (NaCl) Face-centered cubic arrangement of chloride ions with sodium ions in all octahedral holes How many formula units per unit cell? Coordination number of Cl- ? ... of Na+ ? Figure 10.25 NaCl Structure Zinc Blende Structure (ZnS) CCP sulfide ions (FCC) Zinc ions in 1/2 the tetrahedral holes How many formula units per unit cell? Coordination number of S2- ? ... of Zn2+ ? Figure 10.25 Zinc Blende Fluorite Structure (CaF2) CCP calcium ions (FCC) Fluoride ions in all the tetrahedral holes How many formula units per unit cell? Coordination number of F- ? ... of Ca2+ ? Figure 10.25 Cesium Chloride Structure (CsCl) Simple cubic arrangement of chloride ions Cesium ions in the body center (too large for CCP holes) Also a simple cubic array of Cs with Cl in the body center How many formula units per unit cell? Coordination number of Cl ? ... of Cs ? What is the formula? A is located at the corners and faces of a cube (FCC) B is located in all 8 tetrahedral holes C is located at the body center What is the formula? (1/8 x 8) + (1/2 x 6) = 4 A 1 x 8 = 8 B 1 x 1 = 1 C Thus, we have A4B8C 11.8 Bonding in Solids Properties of solids correlate with the types of forces holding the fundamental particles together. Classes: Particles: Molecular Single atoms or molecules Network covalent Covalently bonded atoms in large arrays Ionic Ions Metallic Metal atoms Molecular - Dry Ice Covalent Network ? Silicon Dioxide Network - Silicon Dioxide SiO2 H2O mp = 1710?C mp = 0?C Allotropes of Carbon Ionic - Sodium Chloride Metallic - Copper Which should have the Highest melting points? Lowest melting points? Classes: Forces: molecular intermolecular forces network covalent covalent bonds ionic ionic bonds metallic metallic bonds
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