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Chapter 12 Announcements Intermolecular Forces: Need to Get Recitation Time Chapter 12 Announcements Intermolecular Forces: Need to get recitation time. Liquids, Solids, and Phase Changes Problems Chapter 12: 1,2,5,11,12,15,17,21,22,25,32,34,36,38,40,42,48,50,52,56,59, 62,64,68,75,83,89 Chapter 12: Please Read Section 12-12.5 (skip Section 12.6) There are attractive intermolecular in all solids, Intramolecular forces are attractive “bonding forces” liquids (called condensed phases) and gases. that exist within a molecule or ionic compound holding Molecules are held together by attractive and it together (internal chemical bonds). intramolecular forces (bonds within a molecule). Ionic Bonded substances have gaseous HCl molecules intramolecular bonds resulting from electrons transfer from one atom to another. Covalent molecules have intramolecular bond that result from sharing one or more electrons pairs between atoms usually non- metals. Intermolecular forces are attractive forces between Intermolecular forces play a critical role in life as we molecules that explain many important physical understand it. properties of compounds on the planet. Viscosity Boiling and Freezing Point Polar heads in aqueous exterior Embedded Proteins Wetting or Not Wetting Surface Tension capillary action Non- polar tail Aquoues Interior Intermolecular forces hold together the double Paper is “dried” network of cellulose fibers held helix of DNA. together by hydrogen bonds and Van der waals forces. entangled pulp fibers Fiber Dimensions: 10 nm to 10 mm ~0.7 mm ~0.7 Smook’s Handbook 1992 ~0.7 mm A phase is a state of matter that is homogeneous, You can think of the different phases as classes of chemically uniform and has physically distinct possible molecular motion due to different kinetic properties (density, crystal structure, index of energies (caused by temperature differences) and refraction). varying degrees of intermolecular forces. GAS Generally we LIQUID recognize 4-states CONDENSATION SOLID of matter: ΔH < 0 EVAPORATION ΔH < 0 Gas 1. Gas 2. Liquids DEPOSITION 3. Solids SUBLIMATIONΔH > 0 ΔH > 0 4. Plasma (hot ionized gas) MELTING ΔH > 0 -very high density -high density -low density -not compressibile -not compressibile -high compressibility -low Kinetic energy -mid Kinetic energy -high Kinetic energy FREEZING ΔH < 0 -small distance -molecules close but -large distance between molecules fluid between molecules Solid -IMF large -IMF large -no IMF Liquid The phase of a substance depends on the interplay between Enthalpy is just another word for heat (q) given off or KE(T), which keeps molecules far apart and moving, and IMF absorbed by a system at constant pressure. which keep molecules close together and condensed. Property Gases Liquid Solids !E = qp - P!V Internal energy definition: Molecular Expand to occupy Assumes shape of Fixed Shape container container molecules fixed Mobility qP = !E + P!V = !H Solve for qp Density Very low density High High Chemists measure the enthalpies of many reactions, tabulate them and give them “specific names”. Compressibility High Very Very low Not-compressible Diffusion high molecular Molecule slide Atoms and speeds past one another molecules are --flow fixed Kinetic Energy High Medium Low Intermolecular Few and small Many Many Forces An Exothermic reaction gives off (releases) heat to Phase changes and their enthalpy changes. Note the the surroundings. enthalpy of sublimation is the sum of the enthalpy of fusion and enthalpy of vaporization. heat written as 2H2 (g) + O2 (g) 2H2O (l) + heat a product H2O (g) H2O (l) + heat !H = <0 An Endothermic reaction absorbs heat from the surroundings into the system. heat written as a heat + 2HgO (s) 2Hg (l) + O2 (g) reactant heat + H O (s) H O (l) 2 2 !H = >0 Kinetic molecular theory states that the temperature The heat of fusion and the heat of vaporization are of collection of particles (system) is proportional to measured enthalpies for a pure solid and liquid that tell us the average kinetic energy of the collection. something about the IMF between these molecules. • Molar Heat of Fusion 3 !H˚vap (!Hfus): KE = Ek = RT !H˚fus 2 The energy required to melt one mole of solid (in kJ). Kinetic Energy Temperature • Molar Heat of Vaporization (!Hvap): The energy (in kJ) required The phase of a substance to vaporize one mole of depends on the “competition” liquid. between KE which keeps • Molar Heat of Sublimation molecules moving & far apart (!Hsub): The differences in energy required to and moving, and IMF which The energy (in kJ) required vaporize or melt a pure substance keep molecules close together to vaporize one mole of tells us something about the IMF’s that try to keep these molecules and condensed. solid to gas. together in the condensed phase. Notice how IMF’s are weaker in solid-liquid transition We can graphically show the qualitative and quantitative ("H˚fus) vs liquid-gas transition ("H˚vap) aspects of a phase change by plotting the temperature of a substance vs heat (q) added to a substance. IMF are stronger in the liquid phase. Temperature ˚C Temperature Heat Added (Joules) We can graphically show the quantitative aspects of a The specific heat (s) of a substance is the amount of phase change in a “heating curve” or a plot of heat (q) required to raise the temperature of one gram temperature vs heat added to a substance. of the substance by one degree Celsius. Heating steam past 100˚C The molar heat capacity of a substance is the amount of heat (q) required to raise the temperature of one mole of the substance by one degree Celsius. Boiling and vaporizing all water to steam at 100˚C Heat Heating water absorbed to boiling q = s m !T 100˚C (Joules) Heating Change in Temp: ˚C or K solid ice Melting solid ice to 0˚C water specific heat capacity to 0˚C # of moles J/g °C or J/mol K or grams The specific heat (s) of a substance Here’s the same curve now applying the conservation of is the amount of heat (q) required energy (sum of the heats) to raise the temperature of one q = mice ∆Hfus q = mH2O ∆Hvap gram of the substance by one q = s m ∆T q = s m ∆T q = s m ∆T stm stm degree Celsius. (How thermally ice ice H2O H2O sensitive a substance is to the Phase Phase F addition of energy!) 220 transition transition Temperature does Temperature does not change during not change during a phase transition. a phase transition. Steam Heat (q) absorbed or released: Heating Ice + Water 100 solid ice Water q = s m !T D Water + E to 0˚C mix steam mix 0 C Ice B Heating Boiling all Temperature ˚C Temperature Heating water to water to Melting steam past boiling steam -100 solid ice to 100˚C A 0˚C water 100˚C 100˚C 4.12 38 79.3 305 309 Heat Added (kJ/mole) Calculate the amount of heat required to convert 500 Calculate the amount of heat required to convert 500 grams of ice at -20.0˚C to steam at 120.˚C. The specific grams of ice at -20˚C to steam at 120˚C. The specific heat capacities of water, ice and water vapor are 4.18, heat capacities of water, ice and water vapor are 4.18, 2.06 and 1.84 J/g ˚C respectively, and the latent heat of 2.06 and 1.84 J/g ˚C respectively, and the latent heat of fusion and vaporization, "Hf and "Hv, are 6.02 and 40.7 fusion and vaporization, "Hf and "Hv, are 6.02 and 40.7 kJ/mol respectively. kJ/mol respectively. 1. Heat ice from -20˚C to ice at 0˚C = 500. g x 2.06 J/g ˚C x 20˚C n 2. Melt ice at 0˚C to water at 0˚C = 500. g/(18 g/mol) x 6.02 kJ/mol qi = 0 sum the q’s baby 3. Heat water from 0˚C to water at 100˚C= 500. g x 4.18 J/g ˚C x 100˚C 4. Evap water at 100˚C to vap at 100˚C = 500. g/(18 g/mol) x 40.7 kJ/mol i=1 5. Heat vap from 100˚C to vap at 120˚C = 500. g x 1.84 J/g ˚C x 20˚C for non-phase transitions 1. = 20.6 kJ q = si mi !T 2. = 167.2 kJ 3. = 209.0 kJ qsolid=>liquid = # moles !H˚fusion 4. = 1130.5 kJ for phase transitions 5. = 18.4 kJ qliquid=>gas = # moles !H˚vaporization Total = 1545.6 kJ.
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