1. Solids, Liquids, Gases (and Solutions)

2. Three Phases of Matter

3. Phase Differences Solid Solid – definite volume and shape; particles packed in fixed positions; particles are not free to move Liquid Liquid – definite volume but indefinite shape; particles close together but not in fixed positions; particles are free to move Gas Gas – neither definite volume nor definite shape; particles are at great distances from one another; particles are free to move

4. A Molecular Comparison of Liquids and Solids

5. Phase Changes

6. Energy Changes Accompanying Phase Changes Sublimation:  H sub > 0 (endothermic). Vaporization:  H vap > 0 (endothermic). Melting or Fusion:  H fus > 0 (endothermic). Deposition:  H dep < 0 (exothermic). Condensation:  H con < 0 (exothermic). Freezing:  H fre < 0 (exothermic). Phase Changes

7. Energy Changes Accompanying Phase Changes All phase changes possible under right conditions. heat solid  melt  heat liquid  boil  heat gas = endothermic cool gas  condense  cool liquid  freeze  cool solid = exothermic Phase Changes

9. Phase Diagram Represents phases as a function of temperature and pressure.  Triple point  Critical point  Critical temperature  Critical temperature the minimum temperature for liquefying a gas using pressure  Critical pressure  Critical pressure pressure required for liquefaction

10. Phase Changes

11. Carbon dioxide Carbondioxide

12. Water Water

13. Carbon Carbon

14. Types of Solids  Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS 2 )].

15. Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

16. Ionic Solids

17. Ions (spherical) held together by electrostatic forces of attraction. There are some simple classifications for ionic lattice types. Bonding in Solids

18. Covalent-Network Solids ALL COVALENT BONDS. Atoms held together in large networks. Examples: diamond, graphite, quartz (SiO 2 ), silicon carbide (SiC), and boron nitride (BN). In diamond: –each C atom is tetrahedral; there is a three-dimensional array of atoms. –Diamond is hard, and has a high melting point (3550  C). Bonding in Solids

19. Network Atomic Solids Some covalently bonded substances DO NOT form separate molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms

20. Amorphous solids considerable disorder in their structures (glass and plastic). considerable disorder in their structures (glass and plastic).

21. Metallic Solids Problem: the bonding is too strong for London dispersion and there are not enough electrons for covalent bonds. Resolution: the metal nuclei float in a sea of electrons. Metals conduct because the electrons are delocalized and are mobile. Bonding in Solids

23. Metals Closest Packing of Atoms

24. Metal Alloys are solid solutions  Substitutional Alloy: some metal atoms replaced by others of similar size.  brass = Cu/Zn

25. Metal Alloys (continued)  Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbonsteel = iron + carbon

26. Molecular Solids Strong covalent forces within molecules Weak covalent forces between molecules Sulfur, S 8 Phosphorus, P 4

27. Molecular Solids Intermolecular forces: dipole-dipole, London dispersion and H-bonds. Weak intermolecular forces give rise to low melting points. Room temperature gases and liquids usually form molecular solids and low temperature. Efficient packing of molecules is important (since they are not regular spheres). Bonding in Solids

28. Intermolecular Forces  Dipole-dipole attraction  Hydrogen bonds  Dispersion forces Forces of attraction between different molecules rather than bonding forces within the same molecule.

29. Hydrogen Bonding Intermolecular Forces

30. Dipole-Dipole Forces Intermolecular Forces

31. London Dispersion Forces One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). The forces between instantaneous dipoles are called London dispersion forces. Intermolecular Forces

32. London Dispersion Forces Intermolecular Forces

33. Forces and States of Matter At STP, substances with overy weak intermolecular attraction = gases ostrong intermolecular attraction = liquids overy strong intermolecular attraction or ionic attraction or ionic attraction = solids

34. Bonding in Solids

35. Classification of Matter Solutions are homogeneous mixtures

36. Solute A solute is the dissolved substance in a solution. A solvent is the dissolving medium in a solution. Solvent Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Water in salt waterWater in soda

37. Dissolution of sodium Chloride

38. Concentrated vs. Dilute

39. Some Properties of a Liquid  Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).  Capillary Action: Spontaneous rising of a liquid in a narrow tube.

40. Surface Tension

41. Some Properties of a Liquid  Viscosity: Resistance to flow  High viscosity is an indication of strong indication of strong intermolecular forces intermolecular forces