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Kinetic Theory

Kinetic Theory. 1.) All matter is composed of small particles. 2.) These particles are in constant motion. Kinetic Theory. 3.) Collisions between particles are perfectly elastic (no change in total kinetic energy of the system). Physical States. States of Matter Solid Liquid Gaseous

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Kinetic Theory

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  1. Kinetic Theory • 1.) All matter is composed of small particles. • 2.) These particles are in constant motion.

  2. Kinetic Theory • 3.) Collisions between particles are perfectly elastic (no change in total kinetic energy of the system).

  3. Physical States • States of Matter • Solid • Liquid • Gaseous • Plasma

  4. Liquids/Solids • Molecular Substances: • 1.)General Properties- nonconductors of electricity, often insoluble in water, low melting and boiling points.

  5. Molecular Substances • Molecules are relatively easy to separate from each other because intermolecular forces are weak.

  6. Molecular Substances • 2.) Dispersion Forces - result from temporary dipoles formed in adjacent molecules.

  7. Molecular Substances • Their strength depends on how readily electrons are dispersed and increase with increasing molecular size and mass.

  8. Molecular Substances • Ex. Ordinarily the boiling points of molecular substances will increase with increasing molar mass.

  9. Molecular Substances • Ex. F2 < Cl2 < Br2 < I2. • B.P. (in degrees Celsius) are: -188, -34, 59, 184

  10. Molecular Substances • 3.) Dipole forces- electrical attractive forces between the + end of a polar molecule and the - end of an adjacent molecule.

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  12. Molecular Substances • Ex. Compare the boiling pts of NO (bp = -151oC), N2 (bp = -196oC), and O2(bp = - 183oC).

  13. Molecular Substances • 4.) Hydrogen bonds - are unusually strong dipole forces. The H atom is very small and differs greatly in electronegativity from F, O or N.

  14. Molecular Substances • Compare the boiling points of Group 16 hydrides - • H2O = 100oC • H2S = -61oC • H2Se = -42oC • H2Te = -2oC

  15. Molecular Substances • Note that water has many unusual properties in addition to high boiling point. The open structure of ice, a result of hydrogen bonding, accounts for its low density.

  16. Nonmolecular Solids • Network Covalent Solids - ex. C, SiC, SiO2 (see. P. 246, 247) • High melting pts (covalent bonds must be broken) • Nonconducting • Insoluble in water and other common solvents.

  17. Nonmolecular Solids • Network Covalent Solids- ex. Compare the properties of diamond and graphite (allotropes)

  18. Nonmolecular Solids • Ionic Solids- ex. NaCl, KNO3 • High melting points (due to strong attractive forces between oppositely charged ions).

  19. Nonmolecular Solids • Ionic Solids- • Nonconducting as solids, conducting when molten • Often water soluble (depends on between attractive forces for each other versus for water)

  20. Nonmolecular Solids • Ionic Bond Strength- • Remember Coulomb’s Law - ex. CaO melts at 2927oC vs NaCl at 801oC • Ex NaCl slightly higher than KBr (internuclear distance)

  21. Nonmolecular Solids • Metallic Crystal/Solids- • “Electron-sea” model (d-orbital overlap); cations in mobile sea of electrons. • Conduct electricity, heat

  22. Nonmolecular Solids • Metallic Crystal/Solids- • Electrical conductivity is enhanced by the mobility of electrons • Heat is carried by the collision of electrons (which is frequent in metals)

  23. Nonmolecular Solids • Metallic Crystal/Solids- • Ductile, maleable. • Wide range of melting points, depending on the number of valence electrons.

  24. Nonmolecular Solids • Metallic Crystal/Solids- • Electrons act as a flexible glue holding the atomic nuclei together • Lowest melting pt from +1 cations

  25. Nonmolecular Solids • Metallic Crystal/Solids- • Insoluble in water and other common solvents • Cations cannot dissolve by themselves

  26. Crystal Structures • Unit Cells in Metals- • Unit Cell: smallest unit which, repeated over and over again, generates the crystal.

  27. Crystal Structures • Simple Cubic - unit cell consists of eight atoms at the corners of a cube. • 2r = s • r = atomic radius, s = edge length

  28. Crystal Structures • Face centered cubic - atoms at the corners of a cube and in the center of each face. Atoms touch along the face diagonally. • 4r = s(2)1/2

  29. Crystal Structures • Body-centered cubic: atoms at corner of cube and at center. Atoms touch along the body diagonally. • 4r = s(3)1/2

  30. Crystal Structure Sodium crystallizes into a BCC structure; the unit cell has a length of 0.429nm. What is the atomic radius? r = s(3)1/2/4 = 0.186nm

  31. Liq-Vap Equilibria • Vapor Pressure - when a liquid is introduced into a closed container, it establishes a dynamic equilibrium with its vapor. • Liquid <--> vapor

  32. Liq-Vap Equilibria • The pressure of the vapor at equilibrium is referred to as the vapor pressure of the liquid.

  33. Liq-Vap Equilibria • Vapor Pressure is independent of the volume of the container.

  34. Liq-Vap Equilibria • Add 0.0100 mol of liquid benzene to 1.00L flask at 25oC (vp benzene = 92 mm Hg). How much of the benzene vaporizes? • n vap = 0.0050 mol, n liq = 0.0050 mol

  35. Liq-Vap Equilibria • Add 0.0100 mol of liquid benzene to 2.00L flask at 25oC (vp benzene = 92 mm Hg). How much of the benzene vaporizes? • all; no equilibrium is established

  36. Liq-Vap Equilibria • Temperature dependence of vapor pressure: the vapor pressure of a liquid is always increases as temperature does.

  37. Vapor pressure vs Temperature (degrees C) Log VP vs 1/T (K)

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