1. Liquids and Solids Chapter 13

2. Key concepts 1.Know the basic intermolecular forces that occur between molecules in a substance/mixture. 2.Be able to compare and contrast the strength of intermolecular forces. 3.Learn the terms and definitions for different phase changes. 4.Know how to qualitatively identify phase changes on a heating curve. 5.Know the terms critical temperature and critical pressure. 6.Explain vapor pressure at the molecular level. 7.Know how to interpret a phase diagram. 8.Understand different types of bonding in solid structures.

3. Intermolecular forces forces between molecules (attractions and repulsions), not the forces that hold molecules together (bonds, intramolecular forces). physical properties (such as boiling point) direct result of intermolecular forces. StatevolumeexpansionCompressibilityDiffusion GasAssumes volume of container Expands to fill container CompressibleDiffusion occurs readily LiquidAssumes shape of container Does not expandVirtually incompressible Diffusion occurs slowly solidRetains its own volume Does not expandVirtually incompressible Diffusion occurs extremely slowly

4. Types of intermolecular forces ion-dipole forces: attraction between polar molecules and an ion. Often occurs in aqueous solutions of ions.

5. Types of intermolecular forces dipole-dipole forces: attraction between polar molecules. Increasing polarity of molecules increases the intermolecular force (if molecules are appx same size).

6. Hydrogen bonding Hydrogen bonding: A particularly strong dipole-dipole interaction, occurring between molecules containing H-N, H-O, or H-F bond and a nearby unshared electron pair (usually on N, O, or F). Biological processes, in particular, rely on hydrogen bonding. (hydrogen bonding responsible for helix arrangement of DNA).

7. Effects of dipole forces and H-bonds

8. Types of intermolecular forces London forces (dispersion forces): very weak electrostatic attraction between nonpolar atoms/molecules. Result of small, instantaneous changes in dipole moment. All atoms/molecules have dispersion forces between them. Noble gasMW(amu)Boiling pt.(K) He4.04.6 Ne20.227.3 Ar39.987.5 Kr83.8120.9 Xe131.3166.1

9. Comparing intermolecular forces Are ions involved? Polar molecules and ions? Ion-dipole forcesIonic bonds YESNO Polar molecules? Dipole-dipole forces Dispersion forces YESNO YESNO Ionic bonds >> ion-dipole forces > dipole-dipole forces >> dispersion forces

10. Liquid phenomena Viscosity: the resistance to flow of a liquid. Viscosity is __________ proportional to the intermolecular forces in the liquid. Viscosity is __________ proportional to the temperature. Why?

11. Liquid phenomena Surface tension: The inward forces that must be overcome to expand the surface of a liquid.

12. Liquid phenomena Capillary action: the ability of a liquid to rise in a tube. –Cohesive forces: attractive forces in a liquid –Adhesive forces: attractive forces between a liquid and another surface. For capillary action to occur, adhesive forces must be ________ _____ cohesive forces.

13. Evaporation Molecules on the surface of a liquid escape to the gas phase. An equilibrium between gas and liquid phases of the substance is established.

14. vapor pressure: dynamic equilibrium of a liquid and a gas volatile liquid—a liquid that evaporates readily vapor pressure and boiling point: –boiling point is a function of the pressure of the surroundings

15. Claussius-Clapeyron equation Given a vapor pressure at a given temperature, we can calculate the pressure at another temperature. Two conditions of the C-C equation: –1: –2: If examining changes over small temperature ranges, we usually meet these conditions. Larger temperature ranges usually require some adjustment to the model.

17. Different phase changes between states of matter liquid  gas: –vaporization –condensation solid  liquid –melting –freezing solid  gas –sublimation –deposition

18. Energy changes in phase changes Heat of fusion: Heat of vaporization: Heat of sublimation: heating curves: showing the changes between phases.

19. Heat calculations in phase changes When calculating heat flows involving phase changes, we must include the energy involved in producing the phase change. Otherwise, it is much the same as what we learned in chapter 1. Examples:

20. “cajoling” a gas into a liquid critical temperature: critical pressure: supercritical fluid: beyond the critical temperature and pressure; a “gray area”….

21. Phase diagram summarizing the equilibrium between states of matter. 1.vapor pressure curve –critical point

22. Phase diagrams sublimation curve

23. Phase diagrams melting point curve

24. Phase diagrams THE TRIPLE POINT The point at which ________________ __________.

25. Phase diagrams Solid/liquid densities and phase diagrams. –Comparisons made using ________ of melting point curve. –Most substances have a _____________ __________ melting point curve. –Density of solid is _______ ______ density of liquid.

26. Bonding in solids crystalline solid – amorphous solid –

27. Type of bondForm of unit particles Forces between particles PropertiesExamples Molecular Covalent- network Ionic Metallic

28. Band theory of metals An extension of molecular orbital theory. As orbitals combine, the energy gap between MOs of a particular AO decrease, producing a “band” of orbitals.

29. MO “bands” in sodium and magnesium Magnesium metal is still a good conductor.

30. Band Theory of Metals

31. Structures of Crystals Crystals contain regularly repeating structures. Unit cell is the smallest repeating unit of a crystal. A unit cell is the fundamental box that describes the arrangement of particles in a crystal. These unit cells are stacked in three dimensions to produce a crystal. The arrangement of these unit cells fit into one of seven crystal systems. (Table 13-9). –Crystals have the same symmetry as the unit cells since the crystals are built from multiple units of these cells.

32. Structures of Crystals Look on page 507. Notice that the structures of the unit cells are very similar to the actual structure of the crystal. Why?

33. Structures of Crystals Creating the crystal from the unit cells –Consider the corner of a unit cell as a lattice point. In three dimensions, this lattice point is shared by eight unit cells. If an object were present at this lattice point, it would be equally shared by all eight unit cells (1/8 in each). How many objects would be in each unit cell? DEMO: Use unit cell building blocks to illustrate.

34. Structures of Crystals In a simple or primitive lattice structure (discussed previously), only the corners were occupied by objects. In other crystal types, objects may also occupy other positions in the unit cell. –Simple cubic (no additional objects) –Body-centered cubic (bcc) – another object occupies the center of the unit cell –Face-centered cubic (fcc) – another object occupies the middle of each of the six square faces of the cube

35. Structures of Crystals How many objects/atoms per unit cell for each crystal type? Discuss Look at Figure 13-24.

36. Metallic Solid crystals Lattice types for metals –Body-centered cubic (bcc) –Face-centered cubic (fcc) (cubic close-packed) –Hexagonal close-packed (hcp) Examples: Li, K, Ca, and Au

37. Metallic Solids Obtaining the close-packed structures (fcc and hcp) –Hexagonal close packed structure has an ABA arrangement. The different letters correspond to different planes (Figure 13-27a). –Face-centered cubic (or cubic close-packed) structure has an ABC arrangement. Figure 13-27b For the the close-packed structures approximately 74% of the volume is occupied. The body-centered cubic structures has much less of its volume occupied by metal spheres.

38. Metallic Solids Manganese has a simple cubic unit cell. The atomic radius is 3.15 Å. What is the shortest distance between centers of neighboring Mn atoms? How many nearest neighbors does each atoms have? Nickel crystals are face-centered cubic. The radius of the nickel atom in the metal is 1.24Å. What is the distance between centers of the two closest Ni atoms. What is the length of the cell edge. How many nearest neighbors does each atoms have? Calculate the density of metallic nickel. Determine the percentage of space that is occupied by the nickel atoms.

39. Metallic Solids A group IVA element with a density of 11.35 g/cm 3 crystallizes in a face-centered cubic lattice whose unit cell edge length is 4.95 A. Calculate the element’s atomic weight. What is the atomic radius of this element?

40. Ionic Solids Most salts crystallize as ionic solids with ions occupying the unit cell. The most common example is sodium chloride, which has a face-centered cubic arrangement. –Many other salts that have the same charge on both the anion and cation have the same fcc arrangement (LiCl and MgO). Even though the solid compound possesses charge, it does not conduct electricity. Why?

41. Ionic Solids How many Cl - and Na + ions are in the fcc unit cell? Remember, a sodium ion is in the center. Sodium iodide crystallizes in the fcc structure (like NaCl). The I - ion radius is 2.20 Å. The I - ions at the corners of the unit cell are in contact with those at the centers of the faces. Determine the length of the unit cell. Calculate the radius of the Na + ion assuming anion-cation contact.