1 / 13

Beaker Breaker Draw the Lewis structure of the following polyatomic ions:

Beaker Breaker Draw the Lewis structure of the following polyatomic ions:. nitrite ion sulfite ion. Metallic Bonding. 6-4. Do metals have “few” or “many” valence electrons?. How do they achieve “stability” ?. Metallic Bond Model. metals have very few electrons in their highest E level

mairead-flanagan
Download Presentation

Beaker Breaker Draw the Lewis structure of the following polyatomic ions:

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Beaker BreakerDraw the Lewis structure of the following polyatomic ions: • nitrite ion • sulfite ion

  2. Metallic Bonding 6-4

  3. Do metals have “few” or “many” valence electrons? • How do they achieve “stability” ?

  4. Metallic Bond Model • metals have very few electrons in their highest E level • metals frequently have many vacant d-orbitals just below the outer level • vacant orbitals of adjacent atoms overlap which allows these loosely held e-s to roam freely

  5. Metallic Bond Model (con’t) • “delocalized electrons” - e-s don’t stay in one locality like… • covalent bonding: stay in the overlapping of the shared orbitals • ionic bonding: e-s are bound to an ion within a crystal lattice • mobile electrons form a “sea of electrons”

  6. Metallic Bonding • the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons • mutual sharing of many e-s where each atom contributes its valence e-s which are then free to move about the mostly vacant outer orbitals of all the metal atoms

  7. So…why are metals…. • good electrical conductors? • good thermal conductors? • shiny? • malleable/ductile?

  8. Metallic Properties • High electrical & thermal conductivity • due to high mobility and delocalization of e-s • Luster (shine) • metals absorb E and become “excited” very easily because many of their orbitals are separated by extremely small ∆E…shine occurs when photons are emitted when excited e-s return to ground state

  9. Metallic Prop. (con’t) • Malleability (ability to be hammered/beaten into thin sheets) and ductility (ability to be drawn, pulled, or extruded to produce wire) because metallic bonding is the same in all directions and a shift in layers of atoms is inconsequential

  10. What determines if a metal is “strong” or not?

  11. Metallic Bond Strength • Expressed in the heat of vaporizationvalue where the bonded atoms in the metallic solid state are converted into indiv. metal atoms in the gaseous state (usually↑heat of vap, the ↑ the bond strength) • Determined by • strength of nuclear charge & # delocal. e-s

  12. What is the difference between…- Bond energy ??- Lattice energy ??- Heat of vaporization ?? Hint: • What kind of bonding is generally involved when this term is used? • Is energy being added or taken away?

  13. What is the difference between…- Bond energy ??- Lattice energy ??- Heat of vaporization ?? • Bond energy: E added to break a covalent bond • Lattice energy: E released when ionic cmpds are broken down into atoms • Heat of vaporization: E added when bonded, metallic, solid atoms are broken into indiv gaseous atoms

More Related