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Intermolecular Forces, Liquids, and Solids

Intermolecular Forces, Liquids, and Solids

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Intermolecular Forces, Liquids, and Solids

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  1. Chapter 11 Intermolecular Forces, Liquids, and Solids

  2. Chapter Outline • A Molecular Comparison of Gases, Liquids and Solids • Intermolecular Forces • Some Properties of Liquids • Phase Changes • Vapor Pressure • Phase Diagrams • Structures of Solids • Bonding in Solids

  3. Lecture Outline 1) A Molecular Comparison of Gases, Liquids and Solids 2) Intermolecular Forces • Ion-Dipole Forces • Dipole-Dipole Forces • London-Dispersion Forces • Hydrogen Bonding • Comparing Intermolecular Forces

  4. 1) A Molecular Comparison of Gases, Liquids and Solids

  5. States of Matter The state of s substance depends largely on the balance between the K.E of the particle energies and V(r) (the interparticle energy of attraction) Two competing factors • K.E tend to keep particles apart • interparticle attraction forces draw particles together

  6. K.E >> V(r) K.E > V(r) K.E < V(r) States of Matter The state of s substance depends largely on the balance between the K.E of the particle energies and V(r) (the interparticle energy of attraction) Because in the solid and liquid states particles are closer together, we refer to them as condensed phases

  7. Recall From Chapter 5 The Molecular Kinetic Theory of Gases average of squares of velocities of a gas of molar mass M at a given temperature T in Kelvin average kinetic energy (K.E) per mole of a gas The average K.E is proportional to the absolute temperature The average K.E is proportional to the particle’s average speed Return to slide 4

  8. States of Matter

  9. a. gas • No comparison can be made between kinetic and attraction energy without temperature information. • The average kinetic energy is greater than the energy of attraction. • The average kinetic energy is equal to the energy of attraction. • The average kinetic energy is less than the energy of attraction.

  10. a. gas • No comparison can be made between kinetic and attraction energy without temperature information. • The average kinetic energy is greater than the energy of attraction. • The average kinetic energy is equal to the energy of attraction. • The average kinetic energy is less than the energy of attraction.

  11. b. solid • No comparison can be made between kinetic and attraction energy without temperature information. • The average kinetic energy is equal to the energy of attraction. • The average kinetic energy is greater than the energy of attraction. • The average kinetic energy is less than the energy of attraction.

  12. b. solid • No comparison can be made between kinetic and attraction energy without temperature information. • The average kinetic energy is equal to the energy of attraction. • The average kinetic energy is greater than the energy of attraction. • The average kinetic energy is less than the energy of attraction.

  13. 2) Intermolecular Forces

  14. Intermolecular Attraction Forces The intermolecular attractions between molecules are much weaker than covalent and ionic bonds that hold compounds together. 16 kJ/mol of energy is required to vaporize liquid HCl 431 kJ/mol of energy is needed to dissociate HCl into H and Cl atoms

  15. Boiling Point of HCl is -85 oC at 1 atm Why? Boiling Point of H2O is 100 oC at 1 atm Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. (Strong H-bond)

  16. van der Waals Forces: Between neutral molecules Types of Intermolecular Attraction Forces • Ion-Dipole Interactions: in ionic solutions • Dipole-dipole interactions • Hydrogen bonding • London dispersion forces All four forces are predominately electrostatic in nature All four forces are predominately electrostatic in nature

  17. Ion-Dipole Interactions • Ion-dipole interactions are in solutions of ions. • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

  18. Ion-dipole forces are encountered in both solutions. • Ion-dipole forces are encountered in neither solution. • Ion-dipole forces only are encountered in Ca(NO3)2 in water. • Ion-dipole forces only are encountered in CH3OH in water.

  19. Ion-dipole forces are encountered in both solutions. • Ion-dipole forces are encountered in neither solution. • Ion-dipole forces only are encountered in Ca(NO3)2 in water. • Ion-dipole forces only are encountered in CH3OH in water.

  20. Dipole-Dipole Interactions • Molecules that have permanent dipoles are attracted to each other. • The positive end of one is attracted to the negative end of the other and vice-versa. • These forces are only important when the molecules are close to each other. Spend more time near each other, overall effect is net attraction

  21. Dipole-Dipole Interactions Highest b.p The more polar the molecule, the higher is its boiling point. Give It Some Thoughts For which of the substances are the dipole-dipole attractive forces greatest?

  22. Notes • Dipole-dipole forces are weaker than ion-dipole forces • If two molecules have about the same mass and size, • then dipole-dipole forces increase with increasing polarity. • For molecules of similar polarity, • those with smaller volumes often have greater dipole-dipole attractions.

  23. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

  24. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

  25. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

  26. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. • These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called polarizability.

  27. London Dispersion Forces • These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called polarizability.

  28. Larger surface area Factors Affecting London Forces Molecular shape • The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). • This is due to the increased surface area in n-pentane.

  29. Factors Affecting London Forces Molecular Weight Size of the atom • The strength of dispersion forces tends to increase with increased molecular weight. • Larger atoms have larger electron clouds, which are easier to polarize.

  30. a. increasing polarizability • CH4 < CI4 < CCl4 • CCl4 < CH4 < CI4 • CH4 < CCl4 < CI4 • CI4 < CCl4 < CH4 • List the substances CCl4, CI4, and CH4 in order of increasing • polarizability • Strength of dispersion forces

  31. a. increasing polarizability • CH4 < CI4 < CCl4 • CCl4 < CH4 < CI4 • CH4 < CCl4 < CI4 • CI4 < CCl4 < CH4

  32. Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces? • If two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force. • If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

  33. Strong H-bonding in water How Do We Explain This? ? • The nonpolar series (SnH4 to CH4) follow the expected trend. • The polar series follows the trend from H2Te through H2S, but water is quite an anomaly. ? Boiling point as a function of molecular weight

  34. Hydrogen Bonding A special case of dipole-dipole interaction

  35. We call these interactions hydrogen bonds. Hydrogen Bonding • The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. • Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Examples of H bonding

  36. Water How can we explain that ice floats on water Density of ice at 0 oC is 0.917 g/mL Density of water liquid at 0 oC is 1.00 g/mL Ice is less dense than water? H-Bonding

  37. Hydrogen Bonding in water Ice hexagonal shape water dimer

  38. The densities of liquid water and ice are identical. • The densities of liquid water and ice do not vary with temperature. • Ice is more dense than liquid water. • Ice is less dense than liquid water.

  39. The densities of liquid water and ice are identical. • The densities of liquid water and ice do not vary with temperature. • Ice is more dense than liquid water. • Ice is less dense than liquid water.

  40. Which of the following molecules can hydrogen bond with itself? • 1, 2 • 2, 3 • 3, 4 • 1, 2, 3 • 1, 2, 3, 4

  41. Which of the following molecules can hydrogen bond with itself? • 1, 2 • 2, 3 • 3, 4 • 1, 2, 3 • 1, 2, 3, 4

  42. Intermolecular and intramolecular hydrogen bonds play important role in our life • They are important in stabilizing protein structure, in DNA structure and function, etc. • An interesting consequence of H-bonding is that ice floats. Ice floats, so it forms an insulating layer on top of lakes, rivers, etc. Therefore, aquatic life can survive in winter. • Water expands when it freezes. • Frozen water in pipes may cause them to break in cold weather.

  43. Summarizing Intermolecular Forces

  44. Arrange the following according to increasing melting point. Kr I2 O2 He • O2 < He < I2 < Kr • He < O2 < I2 < Kr • He < O2 < Kr < I2 • I2 < Kr < O2 < He • I2 < Kr < He < O2 Hint: All species are nonpolar then check the polarizability

  45. Arrange the following according to increasing melting point. Kr I2 O2 He • O2 < He < I2 < Kr • He < O2 < I2 < Kr • He < O2 < Kr < I2 • I2 < Kr < O2 < He • I2 < Kr < He < O2

  46. Arrange the following according to increasing melting point. MgO CO2 O2 H2O • MgO < H2O < CO2 < O2 • O2 < CO2 < H2O < MgO • O2 < H2O < CO2 < MgO • H2O < O2 < MgO < CO2 • O2 < CO2 < H2O < MgO

  47. Arrange the following according to increasing melting point. MgO CO2 O2 H2O • MgO < H2O < CO2 < O2 • O2 < CO2 < H2O < MgO • O2 < H2O < CO2 < MgO • H2O < O2 < MgO < CO2 • O2 < CO2 < H2O < MgO

  48. 3) Some Properties of Liquids

  49. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

  50. Viscosities of a series of hydrocarbons at 20 oC (F/A)t Viscosity • Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher temperature. SAE 40 SAE 10 Viscosity of motor oil The higher the number the greater the viscosity is at a given temperature SAE: Society of Automotive Engineers