2 Lewis Structures and Bonding Use NASB to draw dot diagrams.N – electrons needed to fill valence (8 or 2)A- electrons available in atomS – electrons shared = N – AB – 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:ClThe diatomic molecule N2 contains a triple covalent bond.N N N≡N or N⋮⋮NIn 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:IThere are 2 double covalent bonds in a molecule of CO2.C O O═C═O or O::C::OCarbon monoxide has a triple covalent bond.C O C≡O OR C⋮⋮O
5 Bonding TheoriesAccording to VSEPR theory, molecules adjust their shapes to keep pairs of valence electrons as far apart as possible.VSEPR –Valence Shell Electron Pair RepulsionA stereoactive set is a shared pair or an unshared pair of electrons around the central atom.
11 According to VSEPR theory repulsive forces between unshared pairs of electrons causes water molecules to have their shape.
12 Bond angle =109.5 degreesExample: CH4Bond angle =106.5 degreesExample: NH3
13 Bond angle =104.5 degreesExample: H2OBond angle =120 degreesExample: CO32-
14 Bond angle =118.6 degreesExample: O3Bond angle =180 degreesExample: CO2
15 Intermolecular Forces Intermolecular forces – forces between 2 moleculesVan der Waals forces - weakest attractions between moleculesDipole interactions – polar molecules attracted to one anotherDispersion 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