METALLIC BOND - PowerPoint PPT Presentation

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  1. METALLIC BOND • Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. • Even a metal like sodium (melting point 97.8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the Periodic Table.

  2. METALLIC BOND • Happens b/ atoms w/ low electronegativities. • Metals have • Low I.E. • Many unfilled orbitals in shells, making the valence electrons delocalized(shared by) amongst all the atoms and they are free to move throughout the metal.

  3. METALLIC BOND The attraction is b/ the ions and mobile electrons. Mobile sea of valence electrons Mobile sea of valence electrons Lattice of positive ions

  4. Metallic bonding in sodium • The electrons can move freely between the positively charged ions, and so each electron becomes detached from its parent atom. • The electrons are said to be delocalised. • The metal is held together by the strong forces of attraction between the positive metal ions and the delocalised electrons.

  5. Metallic bonding in sodium • This is sometimes described as "an array of positive ions in a sea of electrons". • If you are going to use this view, beware! Is a metal made up of atoms or ions? • It is made of atoms. • Each positive centre in the diagram represents all the rest of the atom apart from the outer electron, but that electron hasn't been lost- it may no longer have an attachment to a particular atom, but it's still there in the structure. Sodium metal is therefore written as Na,not Na+ or potassium metal is written as K but not K+.

  6. The metallic bond in molten metals • In a molten metal, the metallic bond is still present, although the ordered structure has been broken down. • On melting, the bond is loosened, not broken.

  7. The physical properties of metals • Melting points and boiling points • Metals tend to have high melting and boiling points because of the strength of the metallic bond. • Group 1 metals like sodium and potassium have relatively low melting and boiling points mainly because each atom only has one electron to contribute to the bond. • They have relatively large atoms (meaning that the nuclei are some distance from the delocalised electrons) which also weakens the bond.

  8. The physical properties of metals • Heat and electrical conductivity • Metals conduct heat and electricity. The delocalised electrons are free to move and through the motion of the delocalized electrons, electricity is produced. • Liquid metals also conduct electricity, showing that although the metal atoms may be free to move, the delocalisation remains in force until the metal boils.

  9. Electrical conductivity of metals Animation showing electrons moving randomly and then the movement of electrons through a wire

  10. Thermal conductivity • Metals are good conductors of heat. • Heat energy is picked up by the electrons as additional kinetic energy (it makes them move faster). • The energy is transferred throughout the rest of the metal by the moving electrons.

  11. Thermal conductivity of metals Metals are good conductors of heat. There are two reasons for this: • the close packing of the metal ions in the lattice • the delocalised electrons can carry kinetic energy through the lattice

  12. Strength and workability • Malleability and ductility • Metals are described as malleable(can be beaten into sheets) and ductile(can be pulled out into wires). • This is because of the ability of the atoms to roll over each other into new positions without breaking the metallic bond. • If another metal is mixed, it’s not easy for the atoms roll over each other. That’s why the alloys are harder than pure metals.

  13. Because of the electron sea, metals have lustrous appearance since electrons can vibrate at the frequency of the light.

  14. REFERENCES • VIRTUAL CHEMBOOK • Gary L. Bertrand Department of Chemistry University of Missouri-Rolla • chemguideHelping you to understand Chemistry Jim Clark 2005 • Minerals