1. Unit 4: Metallic Bonding

2. Metallic Bonds are… How metal atoms are held together in the solid. Metals hold on to their valence electrons very weakly. Think of them as positive ions (cations) floating in a sea of electrons

3. Examples of Metallic Bonding ++++ ++++ ++++

4. Sea of Electrons ++++ ++++ ++++ Electrons are free to move through the solid, so..... Metals conduct electricity. Vacant “d” orbitals overlap and allow electrons to move.

5. Properties of Metals Excellent conductors of heat and electricity Solids at room temperature (except Hg) Extremely high melting points (200˚C and above) MalleableDuctile

6. Metals are Malleable Hammered into shape (bend). Also ductile - drawn into wires. Both malleability and ductility explained in terms of the mobility of the valence electrons

7. Malleable ++++ ++++ ++++ Force

8. Malleable ++++ ++++ ++++ Mobile electrons allow atoms to slide by, sort of like ball bearings in oil. Force

9. Properties of ionic compounds Crystalline structures (repeated patterns of + and – ions) Are crystalline solids at room temperature Have very high melting points Are hard and brittle – they break under pressure. -+-+ - + -+ -+-+ - + -+

10. Do they Conduct? Conducting electricity means allowing charges to move. In a solid, the ions are locked in place. Ionic solids are insulators. When melted, the ions can move around. lMelted ionic compounds conduct. –NaCl: must get to about 800 ºC. Dissolved in water, they also conduct (ions separate and are free to move in solutions)

11. - Page 198 The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

12. Ionic solids are brittle +-+- + - +- +-+- + - +- Force

13. Ionic solids are brittle + - + - + - +- +-+- + - +- Strong Repulsion breaks a crystal apart, due to similar ions being next to each other. Force

14. Structure of Ionic Compounds

15. Bond Strength Determined by Lattice Energy – the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Lattice Energy = k (Q 1 )(Q 2 ) Lattice Energy = k (Q 1 )(Q 2 )r r = bond length r = bond length Q = ion charge Q = ion charge

16. Comparing Bond Strength Higher charged ions = stronger L.E. = stronger bond = higher melting point L.E for NaF = 923 kJ L.E. for MgO = 3916 kJ The reason for this big difference is... Charge is +2 and -2 in MgO Higher Charge = Stronger Bond = Higher Melting Point! = Higher Melting Point!

17. Metallic bond Strength The higher the positive charge for the metal = smaller ions = stronger bond = higher melting point Which will have a stronger bond Mg or Al ? Why? Al will be higher. It forms a +3 charge vs. +2 for Mg Smaller ion = shorter bond = stronger bond!