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Understanding Chemical Bonding: Ionic, Covalent, Metallic Bonds Explained

Learn how and why chemical bonds form, the difference between ionic, covalent, and metallic bonds, electron transfer, energetics, and the periodic table implications in this comprehensive guide. Watch educational videos and explore Lewis electron-dot symbols.

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Understanding Chemical Bonding: Ionic, Covalent, Metallic Bonds Explained

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  1. Chemical Bonding Why do bonds form? to lower the potential energy between positive and negative charges positive charges protons cations negative charges electrons anions non-metals metals gain e- lose e- Periodic Table

  2. + - Ionic bonding + non-metal metal Groups 1 and 2 Groups 6 and 7 high Electron Affinity low Ionization Energy lose 1 or 2 valence e- gain e- electron transfer takes place electrostatic attraction between cation and anion e- formula = ratio of anions to cations

  3. Covalent bonding non-metal + non-metal electrons shared between atoms high Ionization Energies high Electron Affinities electron density between the atoms distance between atoms = bond length formula = actual # atoms

  4. Metallic Bonding metal + metal metals valence e- well shielded low Ionization Energy low Electron Affinities share valence e- not localized between atoms delocalized move freely throughout metal Na (nucleus and core e-) e- “sea” (valence e-)

  5. Ionic bonding + non-metal metal Na (s) 2  2 NaCl (s) + Cl2 (g) http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA0/Movies/NACL1.html negative exothermic heat given off Ionization Energy Na Na+ + 496 kJ/mol Electron Affinity Cl Cl- -349 kJ/mol E = -504 kJ/mol Lattice Energy k Q1 Q2 d Coulomb’s law NaCl Na+ Cl- +

  6. Lewis electron-dot symbols element symbol = nucleus + core e- one “dot” = valence e- metals dot = e- it loses to form cation non-metal e- paired through unpaired dot = e- gain or sharing

  7. Na (s) 2  2 NaCl (s) + Cl2 (g) : : : + Cl- : Ionic bonding + non-metal metal : . : . Na Na+ + Cl : [Ne] 3s1 3s23p5 [Ar] [Ne] [Ne]  2 CaO (s) + O2(g) Ca(s) 2 . . . : : Ca O . [Ar]4s2 [He]2s2 2p4 Ca2+ O2- [Ar] [Ne] Sn, Pb, Bi and Tl exceptions

  8. + - Ionic sizes e- isoelectronic series same # electrons 46 e- ions get smaller +49 +50 +51

  9. . . H H H H . . Covalent bonding non-metal + non-metal + [He] 1s1 1s1

  10. Lewis structure : : : : : . . : : : . : : F F F F F : : : : : : . . . . H H O : : . : F + : [He]2s22p5 [Ne] lone pairs e- not used in bonding shared e- bonding pair shared equally between F : . . . H O : 1s1 [He]2s22p4 [He] [Ne] oxygen 2 lone pairs bonding pair not shared equally

  11. Electronegativity in a molecule ability of an atom to attract e- to itself Ionization Energy related to Electron Affinity Pauling scale

  12. C H C O ionic > 1.8 NaCl 2.1 ionic 801oC polar covalent 405oC BeCl2 1.5 AlCl3 1.5 polar covalent 178oC PCl3 0.9 76oC polar covalent Cl2 0.0 covalent -101oC non-polar covalent 0.4 + - polar covalent0.5-1.8 Li2O

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