States of Matter - PowerPoint PPT Presentation

states of matter n.
Skip this Video
Loading SlideShow in 5 Seconds..
States of Matter PowerPoint Presentation
Download Presentation
States of Matter

play fullscreen
1 / 93
States of Matter
Download Presentation
Download Presentation

States of Matter

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. States of Matter

  2. Kinetic Molecular Theory • Particles are always in motion. • Temperature is a measure of average kinetic energy of particles. • Intermolecular forces hold particles together. Stronger forces require more energy (higher temp.) to overcome.

  3. Intermolecular Forces • Always weaker than chemical bond • Affect structure and state of matter

  4. Dipole-Dipole Forces • Positive and negative ends of polar molecules attract each other. • About 1% as strong as covalent or ionic bonds • Weaken as distance between molecules increases

  5. Hydrogen Bonding • Especially strong dipole-dipole force • Occurs when H bonds to a strongly electronegative atom—O, N, or F • Very strong because 1) molecule is very polar & 2) small size of H

  6. H Bonding • Example—water • More pronounced in molecules formed from small atoms (dipoles can come closer) • High boiling point

  7. London Dispersion Forces • Forces that exist in all atoms and molecules but that are significant only among Noble gases and nonpolar molecules • Result from temporary dipoles formed when electrons distribute themselves unevenly—can induce a dipole in a neighboring atom • VERY WEAK

  8. London Dispersion Forces • Stronger in larger atoms or molecules due to the greater chance of the formation of instantaneous dipoles.

  9. Physical Properties • Melting and boiling points are higher when IM attractions are stronger • More energy required to separate molecules

  10. Which would have the higher boiling point & Why? • Cl2 or F2? • H2O or H2S? • SiBr2 or SBr2? • CH4 or C10H22? • O2 or NO?

  11. States of Matter • Gases—weak IM forces (like London Dispersion Forces) • Liquids—intermediate IM forces • Solids—strong IM forces

  12. Liquids & IM Forces • Surface tension—result of IM forces that resist an increase in surface area • Capillary action—result of cohesive forces within liquid and adhesive forces between liquid and tube

  13. Liquids (cont’d) • Viscosity—the ability of a liquid to resist flow (resist change in shape) • All effects are higher with more polar molecules.

  14. Solids • Amorphous—without definite structure • Crystalline—definite structures

  15. Solids (Crystals) • Ionic solids—made of charged particles; ions at lattice points • Molecular solids—made of neutral particles; molecules at lattice points • Atomic solids—made of neutral particles; atoms at lattice points; 3 types

  16. An atomic solid, an ionic solid & a molecular solid

  17. Ionic Solids • Ions at lattice points • Closest packed spheres • Arranged to minimize repulsions and maximize attractions • Conducts only when melted

  18. Molecular Solids • Lattice positions occupied by molecules • Internal covalent bonds are strong, but intermolecular forces are weak • IM force: dipole/dipole if polar covalent bond; London dispersion forces (larger in larger molecules)

  19. Atomic Solids • Network—directional covalent bonds; forms giant molecules (diamond, graphite,and silicon); highest melting points • Metallic—delocalized covalent bonds; atoms have closest packing structure; high melting points

  20. Atomic Solids • Group 8A—Noble gases—London dispersion forces only; low melting points.

  21. Network Atomic Solids • Strong, directional bonds • Form giant “molecules” • Typically brittle & poor conductors • Examples—carbon and silicon

  22. Carbon Network • Follows a molecular orbital (not atomic orbital) model

  23. Diamond • Tetrahedral--sp3 hybridized bonds stabilize structure • Large gaps exist between filled and unfilled molecular orbitals—hard for electrons to move—no conductivity

  24. Graphite • Fused carbon rings form sheets • Trigonal planar—sp2 hybridized (1 p orbital remains unhybridized) • Delocalized electrons in orbital causes graphite to be conductive

  25. Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

  26. Closest Packed Solids • aba pattern—alternating layers—atoms in 3rd layer lie directly above atoms in 1st layer—hexagonal unit cell—body centered • abca pattern—atoms in 1st and 4th layers are in line; 2nd & 5th layer; 3rd & 6th layer—face-centered cubic cell

  27. aba Packing

  28. abca Packing

  29. Density of Closest Packed Solids • To calculate density, you need to know: • MASS • VOLUME

  30. Mass • Figure out how many atoms in one unit cell • Multiply by molar mass • Divide by Avogadro’s number • You now know a mass in grams

  31. Face-Centered Cubic Unit Cell If these are atoms of calcium, what is the mass of the cell?

  32. Volume • Determine the length of one side of the cube by using the atomic radius (varies depending on type of unit cell) • Cube the side length.

  33. Simple Cubic--aaa If the atomic radius of this atom is 122 pm, what is the volume?

  34. Body-Centered Cubic--aba If the atomic radius of the atom is 246 pm, what is the volume of the cell?

  35. Face-Centered Cubic--abca If the atomic radius is 291 pm, what is the volume of the cell?

  36. Sample Problem • Silver crystallizes in a face-centered cubic closest packed structure. The radius of the silver atom is 144 pm. Calculate the density of silver.

  37. Bonding in Metals • Strong, non-directional bonds • Atoms are hard to separate but easy to move. • “Electron sea” model • Mobile electrons carry heat or electricity easily

  38. Band Model or Molecular Orbital Model • Electrons travel around metal crystal in a molecular (instead of atomic) orbitals • Result is a continuum of levels that eventually merge to form a band.

  39. Figure 10.19: The molecular orbital energy levels produced when various numbers of atomic orbitals interact.

  40. Band or MO Model • Empty orbitals close in energy exist. • Electrons are very mobile into and out of these similar-energy orbitals--CONDUCTIVITY.

  41. Semiconductors • Some electrons can cross the “energy gap” between molecular orbitals—somewhat conductive • Higher temperatures result in more electrons’ being able to reach conductive bands

  42. Doping • Adding other elements with one more or one less electron than a semiconductor can increase conductivity

  43. n-type semiconductor • An element with one more valence electron is added • More valence electrons are available to move into conduction bands • What could be used to dope Si?

  44. p-type semiconductor • An element with one less valence electron is added • The absence of a valence electron creates a hole through which electrons can travel