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METALS

METALS. Bonds and Properties. Alloys Pure Elements. LUSTER. HIGH. DEFORMABILITY. Malleable & Ductile. CONDUCTIVITY : HEAT AND ELECTRICITY. Good to Excellent. PHASE at STP. Solid (except Hg). Ion FORMATION. Lose e - to form (+) ions. IONIZATION ENERGY. Low (lowest: Fr).

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METALS

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  1. METALS Bonds and Properties Alloys Pure Elements

  2. LUSTER HIGH DEFORMABILITY Malleable & Ductile CONDUCTIVITY : HEAT AND ELECTRICITY Good to Excellent PHASE at STP Solid (except Hg) Ion FORMATION Lose e- to form (+) ions IONIZATION ENERGY Low (lowest: Fr) ELECTRONEGATIVITY Low (lowest: Fr) Some Properties of Metals

  3. Many metals have high luster Hmmm…. How much is this shiney Gold worth !? 400 ounces (27.5lbs) $1209.76 / ounce $483,904.00 per bar 12 bars = $5,806,848.00

  4. Many Metals have high melting points

  5. Metallic bonding: different from ionic bond Both bond types involve electrostatic attraction Ionic bond: transfer electrons from one atom to another Metal bonds: valence electronsROAM FREELYbetween metal atoms • this sea of mobile electrons accounts for • properties of metals Metal bonding

  6. Why can electrons in metals roam freely about?

  7. Metal bonding • Metals: form organized lattice structures similar to ionic cmpds • adjacent atoms in metal lattice are all same • close proximity of atoms allows outer electron energy levels to overlap So…

  8. electrons in outer valence shell can move freely through these overlapping energy levels results in “sea of mobile electrons”

  9. Na 3s1 Na 3s1 overlapping valence electron orbitals

  10. Metal bonding • Freely moving electrons: • called “delocalized” electrons • allows (+) metal cation to form • Delocalized electrons move freely throughout metal from one cation to next • creates“sea of mobile electrons”

  11. sea of electrons: • binds each metal cation to neighboring cation • this creates the metal bond

  12. sea of mobile electrons gives metals some unique properties since electrons move freely from place to place they: • conduct electricity (flow of electrons) • conduct heat • account for metals being: • malleable • ductile • shiney (luster) How do electrons do this? Metal Properties

  13. Electrical Conductivity - + Flow of electrons e-   e- free flow of electrons through the metal e- flow from metal through metal wire towards (+) terminal; then flow from (-) terminal back into metal

  14. Metals & non-metals behave very differently when hit with a force such as hammer: Metals DEFORM & Non-metals SHATTER Why? metals have free flowing electrons & non-metals do not! Malleability

  15. Metal Non-Metal If apply force to metal: • metal atoms shift away from force & free electrons bond newly • overlapping metal ions together • metal’s shape is deformed • but shift doesn't change metal atoms If apply force to non-metal: • like charges align (+) to (+) (-) to (-) which results in shattering due to repulsion forces

  16. Deformation of Metals Deformation of Metals

  17. As # of electrons that can be delocalized ↑ so does: Hardness and Strength Na has one valence electron that can be delocalized so: - is relatively soft ( can be cut with a butter knife) Mg has two valence electrons that can be delocalized so: - still can be cut but is much harder than Na Transition metals have varied # of e-'s that can be delocalized - Chromium (Cr+6) is very hard and has high strength

  18. Alloys • mixture of elements that have metallic properties • mixture can be adjusted to get desired properties • two types: • substitutional and interstitial alloy • (depends on size of elements – same or different size)

  19. Common alloys Brass: Cu & Zn Bronze: Cu, Sn & Al Pewter: Sn, Pb & Cu Solder: Pb & Sn Rose gold: Cu & Al White gold: Au & Ni, Pd or Pt Sterling silver: Ag & Cu Steel: C & Fe Stainless steel: Cr & Ni

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