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- Lewis structures and bonds - bonding theories

- Lewis structures and bonds - bonding theories. Lewis Structures and Bonding. Use NASB to draw dot diagrams. N – electrons need ed to fill valence (8 or 2) A- electrons available in atom S – electrons shared = N – A B – formed bonds ( S divided by 2 ).

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- Lewis structures and bonds - bonding theories

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  1. - Lewis structures and bonds- bonding theories

  2. Lewis Structures and Bonding • Use NASB to draw dot diagrams. • N – electrons needed to fill valence (8 or 2) • A- electrons available in atom • S – electrons shared = N – A • B – formed bonds (S divided by 2)

  3. In Cl2, the total number of unshared pairs of electrons is 6. • Cl + ClCl―Cl or Cl:Cl • The diatomic molecule N2 contains a triple covalent bond. • N + N  N≡N or N⋮⋮N • In the N2 molecule, there is only one unshared pair of electrons in each nitrogen atom.

  4. The HI molecule contains only one single covalent bond. • H• + I  H―I or H:I • There are 2 double covalent bonds in a molecule of CO2. • C 2 O O═C═O or O::C::O • Carbon monoxide has a triple covalent bond. C O  C≡O OR C⋮⋮O

  5. Bonding Theories • According to VSEPR theory, molecules adjust their shapes to keep pairs of valence electrons as far apart as possible. • VSEPR – • Valence Shell Electron Pair Repulsion • A stereoactive set is a shared pair or an unshared pair of electrons around the central atom.

  6. The shape of a molecule of CO2 is linear.

  7. The shape of a molecule of HCN is linear.

  8. The shape of a molecule of CH4 is tetrahedral.

  9. The shape of a molecule of NH3 is trigonal pyramidal.

  10. The shape of a molecule of H2O is bent.

  11. According to VSEPR theory repulsive forces between unshared pairs of electrons causes water molecules to have their shape.

  12. Bond angle = 109.5 degrees Example: CH4 Bond angle = 106.5 degrees Example: NH3

  13. Bond angle = 104.5 degrees Example: H2O Bond angle = 120 degrees Example: CO32-

  14. Bond angle = 118.6 degrees Example: O3 Bond angle = 180 degrees Example: CO2

  15. Intermolecular Forces • Intermolecular forces – forces between 2 molecules • Van der Waals forces - weakest attractions between molecules • Dipoleinteractions – polar molecules attracted to one another • Dispersion forces – caused by the motion of electrons

  16. Hydrogen bonds – strongest intermolecular forces • Hydrogen covalently bonded to a very electronegative atom is also bonded to an unshared electron pair of another electronegative atom. • H2O is a polar molecule. • 2 H• + O H:O:H

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