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FINAL EXAM

This text discusses intermolecular forces and phase changes, including ion-dipole forces, dipole-dipole forces, and phase diagrams. It also covers topics such as specific heat, enthalpy of vaporization and fusion, and vapor pressure.

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FINAL EXAM

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  1. FINAL EXAM Wednesday,December 11, at 10:15 a.m. – 12:15 p.m. in the IC building, Room 421

  2. Intermolecular Forces • The forces holding solids and liquids together are called intermolecular forces. • The covalent bond holding a molecule together is an intramolecular forces. • The attraction between molecules is an intermolecular force. • Intermolecular forces are much weaker than intramolecular forces • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). Chapter 11

  3. Intermolecular Forces Chapter 11

  4. Intermolecular Forces Ion-Dipole Forces • Interaction between an ion and a dipole (e.g. water). • Strongest of all intermolecular forces. Chapter 11

  5. Intermolecular Forces Dipole-Dipole Forces • Exist between neutral polar molecules. • Polar molecules need to be close together. • Weaker than ion-dipole forces. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity. Chapter 11

  6. Intermolecular Forces Chapter 11

  7. Phase Changes • Sublimation: solid  gas. • Vaporization: liquid  gas. • Melting or fusion: solid  liquid. • Deposition: gas  solid. • Condensation: gas  liquid. • Freezing: liquid  solid.

  8. Phase Changes

  9. . • ENERGY ASSOCIATED WITH HEATING CURVES

  10. Topics • Vapor Pressure • Normal Boiling Point • Normal Freezing • Specific Heat • Enthalpy (Heat) of Vaporization • Enthalpy (Heat) of Fusion

  11. Vapor Pressure • THE PRESSURE OF A VAPOR IN EQUILIBRIUM WITH ITS LIQUID (OR ITS SOLID)

  12. NORMAL BOILING POINT & FREEZING POINTS • NORMAL BOILING PT. - THE TEMPERATURE @WHICH VAPOR PRESSURE = 1 atm • NORMAL FREEZING PT. – THE TEMPERATURE @ WHICH THE VAPOR PRESSURE OF THE SOLID AND THE LIQUID ARE THE SAME

  13. Heat Capacity aka Specific Heat (C) • Specific Heat (C) = the amount of energy required to raise the temperature of 1 gram of substance 1 degree celcius

  14. Specific Heat (C) aka Heat Capacity • Units for:specific heat (C) = J/g-oC where J = joules oC = temperature in oC g = mass in grams

  15. Specific Heat (C) Values(aka Heat Capacity) • Example: Water • LIQUID: CLiq = 4.18 J/ (oC . g) • LIQUID: Csol = 2.09 J/ (oC . g) • LIQUID: Cgas = 1.84 J/ (oC . g)

  16. Use of Specific Heat • q = mCDT • q = gm substance x specific heat x DT • where: • M = mass of substance in grams • q = amount of heat (energy) • C = specific heat • And DT = change in temperature

  17. Enthalpy of Vaporizationaka heat of vaporization (DHvap) • Is the amount of heat needed to convert a liquid to a vapor at its normal boiling point

  18. Enthalpy of Fusion aka heat of fusion (DHfus) • Is the amount of heat needed to convert a solid to a liquid at its normal melting (freezing) point

  19. Units for DHvap, DHfus and heat(q) Heat of fusion DHfus = kJ/mol Heat of vaporization DHvap = kJ/mol Heat (q) = Joules

  20. Therefore: • To come up with Joules which is the unit of heat, if: (1) DH is given, then: qvap = DHvap x moles and qfus = DHfus x moles (2) Specific heat (C) is given, then: q = mCDT

  21. Sample Problem • Calculate the enthalpy change upon converting 1 mole of ice at -25 oC to steam at 125 oC under a constant pressure of 1 atm? The specific heats are of ice, water and steam 2.09 J/g-K for ice, 4.18 J/g-K for water and 1.84 J/g-K for steam. For water, DHfus= 6.01 kJ/mol, and DHvap = 40.67kJ/mol. • Note: The total enthalpy change is the sum of the changes of the individual steps.

  22. Phase Changes Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions. • The sequence heat solid  melt  heat liquid  boil  heat gas is endothermic. • The sequence cool gas  condense  cool liquid  freeze  cool solid is exothermic.

  23. Phase Changes Heating Curves • Plot of temperature change versus heat added is a heating curve. • During a phase change, adding heat causes no temperature change. • Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. • Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

  24. HEATING CURVES • ENERGY ASSOCIATED WITH HEATING CURVES • During a phase change, adding heat causes no temperature change.

  25. Phase Changes Critical Temperature and Pressure • Gases liquefied by increasing pressure at some temperature. • Critical temperature: the minimum temperature for liquefaction of a gas using pressure. • Critical pressure: pressure required for liquefaction.

  26. 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. • Any temperature and pressure combination not on a curve represents a single phase.

  27. Phase Diagrams • 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 increases. • Critical point: critical temperature and pressure for the gas. • 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.

  28. Phase Diagrams

  29. Phase Diagrams The Phase Diagrams of H2O and CO2

  30. 3 Things Learned • Reading a phase diagram • Determining triple point on phase diagram • Determining critical point on phase diagram

  31. Phase Diagrams The Phase Diagrams of H2O and CO2

  32. Phase Diagrams The Phase Diagrams of H2O and CO2 • Water: • The melting point curve slopes to the left because ice is less dense than water. • Triple point occurs at 0.0098C and 4.58 mmHg. • Normal melting (freezing) point is 0C. • Normal boiling point is 100C. • Critical point is 374C and 218 atm.

  33. Phase Diagrams The Phase Diagrams of H2O and CO2 • Carbon Dioxide: • Triple point occurs at -56.4C and 5.11 atm. • Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.) • Critical point occurs at 31.1C and 73 atm.

  34. Intermolecular Forces • The forces holding solids and liquids together are called intermolecular forces. • The covalent bond holding a molecule together is an intramolecular forces. • The attraction between molecules is an intermolecular force. • Intermolecular forces are much weaker than intramolecular forces • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). Chapter 11

  35. Intermolecular Forces Chapter 11

  36. Intermolecular Forces Ion-Dipole Forces • Interaction between an ion and a dipole (e.g. water). • Strongest of all intermolecular forces. Chapter 11

  37. Intermolecular Forces Dipole-Dipole Forces • Exist between neutral polar molecules. • Polar molecules need to be close together. • Weaker than ion-dipole forces. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity. Chapter 11

  38. Intermolecular Forces Dipole-Dipole Forces Chapter 11

  39. Intermolecular Forces Dipole-Dipole Forces Chapter 11

  40. Intermolecular Forces London Dispersion Forces • Weakest of all intermolecular forces. • It is possible for two adjacent neutral molecules to affect each other. • The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). • For an instant, the electron clouds become distorted. • In that instant a dipole is formed (called an instantaneous dipole). Chapter 11

  41. Intermolecular Forces London Dispersion Forces • Polarizability is the ease with which an electron cloud can be deformed. • The larger the molecule (the greater the number of electrons) the more polarizable. • London dispersion forces increase as molecular weight increases. • London dispersion forces exist between all molecules. • London dispersion forces depend on the shape of the molecule. Chapter 11

  42. Intermolecular Forces London Dispersion Forces • The greater the surface area available for contact, the greater the dispersion forces. • London dispersion forces between spherical molecules are lower than between sausage-like molecules. Chapter 11

  43. Intermolecular Forces London Dispersion Forces Chapter 11

  44. Intermolecular Forces London Dispersion Forces Chapter 11

  45. Intermolecular Forces Hydrogen Bonding • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong. Chapter 11

  46. Hydrogen Bonding Chapter 11

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