Chapter 8 – Bonding
Metallic: metal bonding to metal only. ex: Mg or Na
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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
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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
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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
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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