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Chapter 11 Intermolecular Forces and Liquids and Solids Scientists are interested in how matter behaves under unusual circumstances. For example, before the space station could be built, fundamental research into materials properties had to be undertaken. 1 In this chapter intermolecular forces will be used to explain (among other things): • How can carbon exist as the hardest (diamond) and one of the softest substances known (graphite) ? • Why do boiling points decrease with altitude? • Why is ice less dense than water? • What are the structures of crystals? • Why does dry ice (CO2) sublimate? Solids and Liquids Intermolecular Forces Ion-dipole, dipole-dipole and H-bonding, dipole-induced dipole, induced dipoleinduced dipole Liquids Phase Diagrams Vapor pressure and temperature, Critical T & P, Surface tension and viscosity Show relation of solid, liquid, and gas phases with change in T and P Solids Unit cells, metal structures, formulas and structures of ionic compounds, Molecular, network, and amorphous solids Properties of Solids Lattice energy, heat of fusion, melting point 2 Dispersion forces Van der Waal’s forces Summary of Intermolecular Forces 1. Ion-dipole Forces •• O δ+ H Hδ + - + - - + - + + - + - H •• O •• H δ δ•• + + δ - δ + 3 2. Dipole-Dipole Interactions - The positive and negative ends of polar molecules interact with each other, resulting in a net force of attraction. - + + + - - + + + - - The strongest type of dipole-dipole interaction involving a hydrogen on one molecule (attached to a F, O, or N) and either F, O, or N on another molecule. About 10 % as strong as an ordinary covalent bond so approximately 15-40 kJ/mol. 4 δ+ + H δ O δ- H δ+ δ+ H δ+ H O δ- H δ+ + H δ O δ- O δδ+ H H + H δ O δ- δ+ H 5 The Double Helix of DNA is held together by hydrogen bonding adenine thymine white = hydrogen blue = nitrogen black = carbon red = oxygen The polymer Nylon is also held together by hydrogen bonding 6 Why is the hydrogen bond considered a “special” dipole-dipole interaction? Decreasing molar mass Decreasing boiling point DISPERSION FORCES 3. Dipole-Induced Dipole Interactions Polarizability - a measure of the extent to which the electron cloud of an atom or molecule can be distorted by an external electric charge. In general, larger atoms or molecules are more easily polarizable than smaller ones (more shells, etc), and so experience larger dipole-induced dipole interactions. 7 8 4. Induced dipole - Induced dipole Interactions (London Forces) Even nonpolar molecules and uncombined atoms have attractive forces between them, otherwise they would never condense or solidify. Which straight-chain alkane (CnH2n+2) molecule has greater attractive intermolecular forces ? H H H H-C-C-C-H H H H H H H H - C - C - C - C - C - C -H H H H H H H H H H H H H H H H H H H H-C-C-C-H H H H H - C - C - C - C - C - C -H H H H H H H The molecule with the longer chain because there are more points of “attachment” via London Forces. 9 Straight-Chain Alkane Effect of Dispersion Forces on Melting Points of Nonpolar Compounds 10 Properties of Liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Strong intermolecular forces High surface tension Photo credit: Microsoft Encarta Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Adhesion Cohesion 11 Properties of Liquids Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity Crystals and Solids Crystal structures covered in the laboratory. Read Sections 11.4 – 11.7, p. 446-462 CsCl ZnS CaF2 12 Condensation Evaporation T2 > T1 The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation H2O (l) H2O (g) Dynamic Equilibrium Rate of Rate of = evaporation condensation 13 Enthalpy of Vaporization, ∆Hvap (Molar Heat of Vaporization) Liquid vaporization heat energy absorbed by liquid Vapor ∆Hvap = amount of energy required to evaporate 1 mole of a liquid under constant pressure Molar Heat of Vaporization and Boiling Points 14 Before Evaporation At Equilibrium The tendency for a liquid to evaporate increases as 1. the temperature rises 2. the surface area increases 3. the intermolecular forces decrease 15 Vapor Pressure and Temperature Diethyl ether - dipole-dipole interactions Water - hydrogen bonding (stronger) 16 Boiling Point “If you have a beaker of water open to the atmosphere, the mass of the atmosphere is pressing down on the surface. As heat is added, more and more water evaporates, pushing the molecules of the atmosphere aside. If enough heat is added, a temperature is eventually reached at which the vapor pressure of the liquid equals the atmospheric pressure, and the liquid boils.” Normal boiling point - the temperature at which the vapor pressure of a liquid is equal to the external atmospheric pressure of 1 atm. Increasing the external atmospheric pressure increases the boiling point Decreasing the external atmospheric pressure decreases the boiling point 17 Location Elevation (ft) San Francisco Salt Lake City Denver Mt. Everest sea level 4400 5280 29,028 Boiling Point H2O (oC) 100.0 95.6 95.0 76.5 The Clausius-Clapeyron Equation Relationship between vapor pressure, T, and ∆Hvap ln P = − ∆H vap RT R = 8.314 J/mol·K +C C = constant that depends on the compound’s volatility We can use this equation to measure ∆Hvap for a given compound if we know P and T at two different points: 18 Problem 11.88 – Estimate the molar heat of vaporization of a liquid whose vapor pressure doubles when the temperature is raised from 85oC to 95oC. The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium Freezing H2O (l) Melting H2O (s) 19 Liquid – Solid Equilibrium solid → liquid → The melting point of the solid, or the freezing point of the liquid, is the temperature at which the two phases are in equilibrium. rate melting → rate freezing (dynamic equilibrium) → ∆Hfusion = molar heat of fusion, energy in kJ required to melt one mole of solid 20 H2O (g) Molar heat of sublimation (∆Hsub) is the energy required to sublime 1 mole of a solid. Deposition Sublimation H2O (s) ∆Hsub = ∆Hfus + ∆Hvap ( Hess’s Law) Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Heating – Cooling Curves Temperature Vapor Boiling point Melting point Solid and liquid in equilibrium A C D Liquid and vapor in equilibrium Liquid B Solid Time 21 Phase Diagrams States of matter as a function of temperature and pressure Pressure → liquid solid gas Temperature → 22 Phase Diagram for Water Phase Diagram for Water 1. Curve AD is the equilibrium vapor pressure curve where the liquid and gas are in equilibium. 2. Line AC is the solid-liquid equilibrium curve. Note that the slope is negative for water - most substances have a positive slope here. A negative slope means that the substance becomes a liquid when the pressure is increased. 3. Point A is the Triple Point - all three phases are in equilibrium at once. 23 On the left is water; water is unusual because the solid is less dense than the liquid. For most substances, the solid is more dense than the liquid like for benzene (on the right). 24 25 Water has a Maximum Density at 4 oC 26 Phase Diagram for CO2 Critical point 1. Positive slope for the solid-liquid interface (normal) 2. Sublimation occurs at room temperature under 1 atm of pressure. 3. There is a “critical point” at 73 atm and 31 oC (varies from substance to substance). 4. Above the critical point the substance is known as a “supercritical fluid” (not a liquid; not a gas). Increasing the pressure normally would convert a gas to a liquid but doesn’t happen. 5. Supercritical fluids have enhanced solvation abilities, i.e. the fluid will dissolve greater amounts of solute than normal. 27