Chapter 8 – Bonding Metallic: metal bonding to metal only. ex: Mg or Na no strong bonds moving electrons Ionic: only when there are (+) and (-) ions, generally a metal and a non-metal. ex: NaCl Covalent: sharing electrons. ex: Cl₂ Diatomic Molecule: a molecule containing only 2 atoms. ex: O₂ , HF BrINClHOF = diatomic molecules, alone they are always in pairs VSEPR (valence shell electron pair repulsion) keep electron pairs as far apart as possible give lone pairs even more space Extended Octet: happens when an element has more than 8 valance electrons (only possible from phosphorus on). Isomers: Same formula, different shapes. Resonance: double bond or electron pairs could be in multiple places- different ways to draw configuration, all are correct. Polarity: determined by lone pairs/ electronegativities. Intermolecular Forces Hydrogen Bonding: must have O-H, N-H, F-H bonds in structure Dipole: between polar molecules only London Dispersion Forces: due to random nature of electron movement it creates temporary dipoles. More influence in large structures. Lewis Electron Dot Structures 1. Count all valence electrons. 2. Draw skeleton. ex: o-c-o 3. How many electrons left over after skeleton (each bond=2e-‘s). 4. Can you use leftover e-‘s to fill octets? if not make a double bond for every 2 e-‘s short. Formal Charge FC= (# of valence electrons when unbounded)-(# of valence electrons when bonded) each atom gets ½ of any shared bond Sum of FC of each atom equals overall charge of that species Lower FC means more stable molecule or ion Bonding Notes: single bonds = σ (Sigma) bond double bonds = Π (Pi) bond & σ (sigma) bond triple bonds = σ (sigma) bond & 2 Π (Pi) bonds Molecular Geometry Shapes