Solids, Liquids, Gases (and Solutions)

1. Solids, Liquids, Gases (and Solutions)

2. Three Phases of Matter

3. Phase Differences Solid – definite volume and shape; particles packed in fixed positions; particles are not free to move Liquid – definite volume but indefinite shape; particles close together but not in fixed positions; particles are free to move 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. Phase Changes • Energy Changes Accompanying Phase Changes • Sublimation: Hsub > 0 (endothermic). • Vaporization: Hvap > 0 (endothermic). • Melting or Fusion: Hfus > 0 (endothermic). • Deposition: Hdep < 0 (exothermic). • Condensation: Hcon < 0 (exothermic). • Freezing: Hfre < 0 (exothermic).

7. Phase Changes • 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

9. Phase Diagram Represents phases as a function of temperature &pressure. Triple point: Where all three lines meet, a unique combination of temperature and pressure where all three phases are in equilibrium together. That's why it is called a triple point. Critical point Critical temperature: above the critical temperature, it is impossible to condense a gas into a liquid just by increasing the pressure (the minimum temperature for liquefying a gas using pressure) Critical pressure : pressure required for liquefaction

10. Phase Changes

11. Carbon dioxide Carbon dioxide

12. Water Water

13. Carbon Carbon

14. Types of Solids Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)].

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

18. Bonding in Solids • Covalent-Network Solids • ALL COVALENT BONDS. • Atoms held together in large networks. • Examples: diamond, graphite, quartz (SiO2), 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).

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).

21. Bonding in Solids • 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.

23. MetalsClosest Packing of Atoms

24. Metal Alloysare 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 + carbon

26. Molecular Solids Strong covalent forces within molecules Weak covalent forces between molecules Sulfur, S8 Phosphorus, P4

27. Bonding in Solids • 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).

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. Intermolecular Forces Hydrogen Bonding

30. Intermolecular Forces Dipole-Dipole Forces

31. Intermolecular Forces • 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.

32. Intermolecular Forces London Dispersion Forces

33. Forces and States of Matter At STP, substances with very weak intermolecular attraction = gases strong intermolecular attraction = liquids very strong intermolecular 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. Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water 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. Viscosity: Resistance to flow Some Properties of a Liquid • High viscosity is an • indication of strong • intermolecular forces