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Covalent Bonding

Covalent Bonding. Covalent Bonding. Covalent bond –formed from the sharing of electrons Usually between 2 non metallic elements. Molecules. Molecule – two or more atoms covalently bound together Diatomic molecule – two of the same atom bound together. Diatomic Molecules.

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Covalent Bonding

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  1. Covalent Bonding

  2. Covalent Bonding • Covalent bond –formed from the sharing of electrons • Usually between 2 non metallic elements

  3. Molecules • Molecule – two or more atoms covalently bound together • Diatomic molecule – two of the same atom bound together

  4. Diatomic Molecules • Br I N Cl H O F or the Magnificent 7 • These atoms never exist alone. • They always come in pairs • For example: • Br  Br2 • I  I2 • N  N2 • Cl  Cl2 • H  H2 • O  O2 • F  F2

  5. Binary Molecular Compounds • Binary Compounds consist of 2 elements • Binary covalent compounds can be recognized by containing 2 nonmetals • This is different from ionic compounds that contain a metal & nonmetal, metal & a polyatomic ion, or 2 polyatomic ions

  6. Lewis Structure • Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule • There are rules for Lewis Structures that are based on the formation of a stable compound • Atoms want to achieve a noble gas configuration

  7. Octet & Duet Rules • Octet Rule – atoms want to have 8 valence electrons • Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

  8. Steps for drawing Lewis Structures • Sketch a simple structure with a central atom and all attached atoms • Add up all of the valence electrons for each individual atom • If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge • If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

  9. Steps for drawing Lewis Structures • Subtract 2 electrons for each bond drawn • Complete the octet on the central atom & subtract those electrons • Complete the octet on the surrounding atoms & subtract those electrons • Get your final number • If 0  you are done! • If +  add that many electrons to the central atom • If -  need to form multiple bonds to take away that many electrons

  10. Examples • CCl4 • Sketch a simple structure with a central atom and all attached atoms Cl │ Cl – C – Cl │ Cl

  11. Examples • Add up all of the valence electrons for each individual atom • 4 + 4(7) = 32 • Subtract 2 electrons for each bond drawn • 32-8 = 24 • Complete the octet on the central atom & subtract those electrons • Done

  12. Examples • Complete the octet on the surrounding atoms & subtract those electrons • 24 – 24 = 0 • Final number = 0…DONE! • Final structure is… • __ │Cl │ __ │ __ │Cl – C – Cl │ │ │Cl │

  13. Examples • N2

  14. Examples • NH3

  15. Examples • PO4-3

  16. Bond Types • Sigma bonds () – single covalent bond • Pi bonds () – occur when multiple bonds are formed • Single bond – sigma • Double bond – 1 sigma & 1 pi • Triple bond – 1 sigma & 2 pi

  17. Bond length & Strength • As the number of bonds increases, the bond length decreases • The shorter the bond, the stronger the bond

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