1. States of Matter Gases, Liquids and Solids

2. Kinetic Molecular Theory of Gases Describes the motion of gas particles Points of Kinetic Molecular Theory: Gas particles are point masses Explains low density and compressibility Gas particles are in constant motion They move in straight lines at 100-1000m/s They change direction only when they run into something Collisions with container walls cause pressure Explains Brownian motion and diffusion/effusion

3. Brownian Motion

4. Diffusion and Effusion Diffusion is the mixing of gases Effusion is the migration of a gas through a tiny orifice into an evacuated space Graham’s Law of effusion Kinetic energy of a particle E = ½mv 2 For the same energy, a heavier particle moves more slowly [v =  (2E/m)]

5. Graham’s Law Comparing two particles with the same energy: ½m A v A 2 = ½m B v B 2 Rearrange and cancel: v A 2 /v B 2 = m B /m A or v A /v B =  (m B /m A )

6. Graham’s Law Molar mass can be used for m, and the effusion rate is directly proportional to v Example problem: Find the molar mass of a gas that effuses at a rate twice as slow as helium. Solution: v A /v B =  (m B /m A ) If gas A is helium, v A /v B = 2, so  (m B /4) = 2 (m B /4) = 4 and m B = 16g/mol

7. Kinetic Theory Points of Kinetic Molecular Theory continued Gas particles are point masses Gas particles are in constant motion All collisions between gas particles are perfectly elastic No attractive forces between particles Gases at the same temperature have the same average kinetic energy

8. Gas Pressure Pressure is force/area Units: N/m 2 (Pascal) (kPa = 1000 Pascals) psi = pounds per square inch Barometers – pressure measured as height of a column of mercury in a mercury barometer Standard pressure = 760mmHg = 1 atm = 101.325kPa = 29.92inHg = 14.7psi

9. Mercury barometer

10. Open arm manometer

11. Open arm manometer

12. Closed arm manometer

13. Dalton’s Law of Partial Pressures When different gases are mixed, every particle contributes equally to the total pressure In a gas mixture, the contribution of each component depends on the mole fraction of that component – more particles means more pressure P T = P A + P B + P C …

14. Interparticle forces London dispersion forces (van der Waal’s forces) Present in all particles Dominant attractive force in non-polar molecular compounds Weakest of all interparticle forces Due to temporary dipoles formed by electron dislocation

15. Interparticle forces

17. Magnitude of London dispersion forces depends on Size of molecule – more surface area means stronger forces Polarizability of electron cloud – larger atoms’ electron clouds are more polarizable due to shielding

18. Interparticle forces Dipole-dipole interactions – in polar molecules opposite poles attract. Stronger than LDF

19. Interparticle forces Hydrogen bonding – strong interaction between hydrogens (  +) and electronegative atoms (  -) Hydrogen must be attached to an electronegative atom (usually O or N) Strongest non-bonding interaction

20. Hydrogen bonding

21. Hydrogen bonding in ice

22. Hydrogen bonding in DNA

23. Liquids More dense than gases, less than solids (except water) Incompressible Particles are in contact but able to move past each other Liquids (and gases) are fluids – able to flow Viscosity – resistance to flow

24. Liquids Viscosity increases with molecular surface area (chain length) Viscosity decreases with increasing temperature Surface tension depends on strength of interparticle interactions Defined as the energy required to increase the surface area of a liquid by a certain amount

25. Liquids Capillary action – ability of a liquid to rise in a narrow tube Happens when adhesive forces between tube and liquid are greater than cohesive forces in liquid Height to which the liquid will rise is a measure of the difference in adhesive and cohesive forces

26. Solids Most dense state (except water) Particles vibrate in place Molecular solids Smallest particle is a molecule Molecules are composed of all nonmetals held together by covalent bonds Molecules are held next to each other by LDF, dipole-dipole interactions or H-bonds

27. Solids Molecular solids Low MP/BP Insulators Usually crystalline Examples: water, sugar, caffeine Network solids Covalent network solids Covalent bond throughout

28. Solids Highest MP/BP No smaller units Examples: C (diamond), Si, quartz Ionic solids All salts - composed of metal & nonmetal High MP/BP Simplest unit is “formula unit” Ions are held in place by attraction to oppositely charged ion Insulators unless melted or in solution

29. Solids Metals Held together by nondirectional metallic bonds “Electron sea” of shared electrons Not usually brittle like crystalline solids Conductors of heat/electricity Ductile, malleable, luster Amorphous solids No regular arrangement No sharp melting point Examples: rubber, glass

30. Crystals Have a regular, repeating pattern of atoms Unit cell is simplest repeating pattern Have sharp melting points Ionic and metallic crystals have high melting points Molecular crystals have low melting points

31. Crystal types Cubic HALITE

32. Cubic crystals FACE CENTERED CUBICBODY CENTERED CUBIC

33. Face centered cubic - halite

34. Body centered cubic - CsCl

35. Tetragonal crystals RUTILE (TiO 2 )

36. Other crystal types HEXAGONAL (QUARTZ) ORTHORHOMBIC (CALCITE)

37. Phase changes and energy Solid/liquid: melting and freezing Melting point: temperature at which vapor pressures of solid and liquid are equal Liquid/gas: vaporization and condensation Boiling point: temperature at which vapor pressure of liquid = atmospheric pressure Boiling: vaporization at BP Evaporation: vaporization at a lower temperature

38. Boiling and evaporation

39. Freezing point

40. Phase changes Solid/gas: sublimation and deposition Energy and phase changes Melting is endothermic, freezing is exothermic Boiling is endothermic, condensation is exothermic Sublimation is endothermic, deposition is exothermic

41. Phase diagrams Phase diagrams show conditions under which an element, compound or mixture will exist in a given state Variables are pressure (y) and temperature (x) Triple point: temperature and pressure where solid, liquid and gas can exist in equilibrium Critical temperature: temperature above which a substance cannot be liquified Critical pressure: pressure necessary to liquefy a substance at the critical temperature (together they make the critical point)

42. Phase diagrams

43. Phase diagrams – UF 6

44. Phase diagrams – CO 2

45. Phase diagram - water

46. Phase diagrams – water (detailed)

47. Heating curves Shows change in temperature as heat is added Slope of curve gives specific heat No change in temperature during phase changes

48. Heating curves