METALS

1. METALS Bonds and Properties Alloys Pure Elements Mr. Shields Regents Chemistry U09 L04

2. Recall that many metals have high luster Like GOLD! Hmmm…. How much is this shiny Gold worth ?? 400 ounces (27.5lbs) $647 / ounce* $258,800 per bar 12 bars = $3,105,600 (* 2007 price) Standard 400 oz ingots

3. Many Metals have high melting points … But not all! Where is the High Melting Pt. of W taken advantage of in your home?

4. W More Properties of Metals Approx. 78% of the Elements in the Periodic Table are Metals

5. Metal bonding The bonding in metals is very different from that of an ionic bond In an ionic bond there is a transfer of electrons from one atom to another But in metal bonds electrons ROAM FREELY from one metal atom to the next. It’s these freely roaming electrons that account for many of the properties of metals - +

6. Mobile metal electrons Why can electrons in metals roam freely about?

7. Metal bonding Metals exist in organized lattice structures similar to ionic Compounds. The difference is that Adjacent atoms in the metal Lattice are all the same. Being in close proximity, Outer Energy levels overlap. AND… Electrons in the outer Valence shell can move freely through these overlapping Energy levels.

8. Mobile Metal Electrons Na 3s1 Na 3s1 Valence electrons move from The valence shell of one atom to the next Overlapping orbitals

9. Metal bonding “Delocalized” electron leaves behind centers of temporarily Positive Metal Cations. The delocalized electron then Moves freely through the Metal from one Cation To the next. This creates what is called “the Sea of Electrons” This Sea of Electronsbinds each metal Cation to all its Neighbors. - A metal bond is the attraction of metallic cations for delocalized electrons

10. Metal Properties • It’s the sea of electrons • That give metals some • Of their unique properties. • Because they can move freely • From place to place they: • Conduct electricity • (a flow of electrons) • 2) Conduct heat • 3) Are malleable and ductile • 4) Have luster (e- absorbing then • releasing photons at the surface How do electrons lead to these properties of metals?

11. + Flow of electrons e-   e- Electrical Conductivity - Free flow of electrons through the metal Electrons flow from the metal through the metal wire towards the + charge Electrons then flow from the negative terminal back into the metal they originally came from.

12. Malleability Metals and non-metals behave very differently when they are Hit with a force such as a hammer. Metals deform and Non-metals shatter. But Why? Again it’s a Consequence of The Freely flowing Electrons in metals

13. Metal Non-Metal Malleability When a force is applied to a metal some of the metal atoms shift away from the force. But the free electrons simply bond the newly overlapping Metal ions together. The metal has Been deformed but The shift is not Of any consequence When a force is Applied to the Non-metal, positive And negative charges align. This Results in a fracture due to the force ofrepulsion

14. Hardness & Strength As the number of electrons that can be delocalized Increases so does Harness and Strength Na has one s electron that can be delocalized - its relatively soft & weak ( can be cut with a butter knife) Mg has two electrons that can be delocalized - so its much harder than sodium but not a lot of strength Transition metals have several s and d electrons that can be delocalized - Chrome is very hard and has high strength

15. Hardness & Strength In General as you move left to right across a period The strength and hardness of the metallic bond increases (as long as the # of s and d electrons that can be delocalized inc.) (for ex: Fe 2-8-14-2 is harder than Cu 2-8-18-1 because Cu’s d Orbitals are all filled so they are not available for bonding. Fe Does have d orbitals that can participate) AND In general as you move down a group the strength of the Metallic bond and the hardness decreases because the Delocalized e- are farther from the pos. cation (Cu is a harder metal than Au)