Liquids and Solids

1. Liquids and Solids Chapter 10

2. A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. 2 Phases Solid phase - ice Liquid phase - water

3. Schematic representations of the three states of matter.

5. Generally, intermolecular forces are much weaker than intramolecular forces. Intermolecular Forces Intermolecular forces are attractive forces between molecules. Intramolecular forces hold atoms together in a molecule. • Intermolecular vs Intramolecular • 41 kJ to vaporize 1 mole of water (inter) • 930 kJ to break all O-H bonds in 1 mole of water (intra) “Measure” of intermolecular force boiling point melting point DHvap DHfus DHsub

6. Intermolecular Forces Forces between (rather than within) molecules. • dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “” ends of the dipoles are close to each other. • hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

7. Orientation of Polar Molecules in a Solid Intermolecular Forces Dipole-Dipole Forces Attractive forces between polar molecules

8. (a) The electrostatic interaction of two polar molecules. (b) The interaction of many dipoles in a condensed state.

9. (a) The polar water molecule. (b) Hydrogen bonding among water molecules.

11. Decreasing molar mass Decreasing boiling point Why is the hydrogen bond considered a “special” dipole-dipole interaction? Why this sudden increase? The boiling point represents the magnitude and type of bonding

12. Ion-Dipole Interaction Intermolecular Forces Ion-Dipole Forces Attractive forces between an ion and a polar molecule

13. London Dispersion Forces • relatively weakforces that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18) • caused by instantaneous dipole, in which electron distribution becomes asymmetrical. • the ease with which electron “cloud” of an atom can be distorted is called polarizability.

14. (a) An instantaneous polarization can occur on atom A, creating an instantaneous dipole. This dipole creates an induced dipole on neighboring atom B. (b) Nonpolar molecules such as H2 also can develop instantaneous and induced dipoles.

15. Dispersion forces usually increase with molar mass. Intermolecular Forces Dispersion Forces Continued Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted. • Polarizability increases with: • greater number of electrons • more diffuse electron cloud

16. O O S What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: dispersion forces. SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

17. Properties of Liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Or The resistance to an increase in its surface area (polar molecules). Strong intermolecular forces High surface tension

18. A molecule in the interior of a liquid is attracted by the molecules surrounding it, whereas a molecule at the surface of a liquid is attracted only by molecules below it and on each side.

19. More Properties of Liquids Capillary Action: Spontaneous rising of a liquid in a narrow tube. Nonpolar liquid mercury forms a convex meniscus in a glass tube, whereas polar water forms a concave meniscus.

20. Adhesion Cohesion More Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules

21. More Properties of Liquids Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity

22. Types of Solids Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)]. Amorphous solids: considerable disorder in their structures (glass).

23. An amorphoussolid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Non-crystalline quartz glass Crystalline quartz (SiO2)

24. lattice point A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. An amorphoussolid does not possess a well-defined arrangement and long-range molecular order. A unit cell is the basic (smallest) repeating structural unit of a crystalline solid. • At lattice points: • Atoms • Molecules • Ions Unit cells in 3 dimensions Unit Cell

28. Shared by 2 unit cells Shared by 8 unit cells

29. 4 atoms/unit cell 2 atoms/unit cell 1 atom/unit cell (8 x 1/8 + 6 x 1/2 = 4) (8 x 1/8 + 1 = 2) (8 x 1/8 = 1)

30. Three cubic unit cells and the corresponding lattices.

32. Analysis of crystal structure using X-Ray Diffraction

33. Extra distance = 2d sinq = nl BC + CD = (Bragg Equation)

34. Bragg Equation Used for analysis of crystal structures. n = 2d sin  d = distance between atoms n = an integer  = wavelength of the x-rays

35. 1 x 154 pm = 2 x sin14.17 nl 2sinq X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of 14.170. Assuming that n = 1, what is the distance (in pm) between layers in the crystal? nl = 2d sin q n = 1 q = 14.170 l = 0.154 nm = 154 pm = 77.0 pm d = 11.5

36. Types of Crystalline Solids Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl). Molecular Solid: discrete covalently bondedmolecules at each of its lattice points (sucrose, ice).

37. Types of Crystals • Ionic Crystals • Lattice points occupied by cations and anions • Held together by electrostatic attraction • Hard, brittle, high melting point • Poor conductor of heat and electricity CsCl ZnS CaF2

38. carbon atoms Types of Crystals • Atomic Crystals • Lattice points occupied by atoms • Held together by covalent bonds • Hard, high melting point • Poor conductor of heat and electricity graphite diamond

39. Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors. • hexagonal closest packed (“aba”) • cubic closest packed (“abc”)

40. The closest packing arrangement of uniform spheres, (a) aba packing (b) abc packing.

41. When spheres are closest packed so that the spheres in the third layer are directly over those in the first layer (aba), the unit cell is the hexagonal prism illustrated here in red.

42. When spheres are packed in the abc arrangement, the unit cell is face-centered cubic.

43. The indicated sphere has 12 nearest neighbors.

44. d = d = m m V V x = x 1 mole Ag 107.9 g 7.17 x 10-22 g 6.022 x 1023 atoms mole Ag 6.83 x 10-23 cm3 When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver. V = a3 = (409 pm)3 = 6.83 x 10-23 cm3 4 atoms/unit cell in a face-centered cubic cell m = 4 Ag atoms = 7.17 x 10-22 g = 10.5 g/cm3

45. Types of Crystals • Molecular Crystals • Lattice points occupied by molecules • Held together by intermolecular forces • Soft, low melting point • Poor conductor of heat and electricity

46. nucleus & inner shell e- mobile “sea” of e- Types of Crystals • Metallic Crystals • Lattice points occupied by metal atoms • Held together by metallic bonds • Soft to hard, low to high melting point • Good conductors of heat and electricity Cross Section of a Metallic Crystal

47. Examples of three types of crystalline solids. (a) An atomic solid. (b) An ionic solid. (c) A molecular solid.

48. Types of Crystals

50. Bonding Models for Metals Electron Sea Model: A regular array of metals in a “sea” of electrons. Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.