Lewis Dot Structure Rules:

1. Lewis Dot Structure Rules: • Treat ions separately (e.g. NH4Cl) • Count only valence electrons • Assemble bonding framework • Fill up non-bonding electrons on • outer atoms • Fill up non-bonding electrons on • inner atoms • Calculate Formal Charge • Minimize Formal Charge

2. To do Lewis Structures: • Must be able to recognize • polyatomic ions • Must be able to identify • valence electrons • Must be able to construct • Bond framework Periodic Table : Column numbers! More complex: H outside High c outside Formula hints Acidic Hs bond to O atoms

3. Hints on Lewis Dot Structures • Octet “rule” is the most useful guideline. • Carbon forms 4 bonds. • Hydrogen typically forms one bond to other atoms. • When multiple bonds are forming, they are usually between C, N, O or S. • Nonmetals can form single, double, and triple bonds, but not quadruple bonds. • Always account for single bonds and lone pairs before forming multiple bonds. • Look for resonance structures.

4. PCl3 5+(3*7)=26 e- Bonding Pairs Lone Pairs (a.k.a. nonbonding electrons)

5. Try Some Examples: • CH3CH2NH2 • Cl2CO • Ozone (O3) • NO2 vs. N2O SPENT LOTS O’ TIME PRACTICING…

6. Formal Charge Difference between the # of valence electrons in the free atom and the # of electrons assigned to that atom in the Lewis structure. FC = formal charge; G.N. = Group Number #BE = bonding electrons; #LPE = lone pair electrons If Step 4 leads to a positive formal charge on an inner atom beyond the second row, shift electrons to make double or triple bonds to minimize formal charge, even if this gives an inner atom with more than an octet of electrons.

7. Covalent Bonding • Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): • One shared pair of electrons = single bond (e.g. H2); • Two shared pairs of electrons = double bond (e.g. O2); • Three shared pairs of electrons = triple bond (e.g. N2). • Generally, bond distances shorten with multiple bonding. Octet in each case

9. Odd Number of Electrons… NO Number of valence electrons = 11 Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons) NO2 Number of valence electrons = 17 Molecules and atoms which are neutral (contain no formal charge) and with an unpaired electron are called Radicals O2

10. Beyond the Octet • Elements in the 3rd period or higher can have more than an octet if needed. • Atoms of these elements have valence d orbitals, which allow them to accommodate more than eight electrons.

11. More than an Octet… Elements from the 3rd period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding P : (Ne) 3s2 3p3 3d0 Number of valence electrons = 5 + (5 x 7) = 40 PCl5 S : (Ne) 3s2 3p4 3d0 Number of valence electrons = 6 + (4 x 7) = 34 SF4 The Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.

12. Less than an Octet… BCl3 Group 3A atom only has six electrons around it However, Lewis acids “accept” a pair of electrons readily from Lewis bases to establish a stable octet

13. VSEPR Definitions • Electron group –set of electrons that occupies a particular region around an atom. • Ligand – an atom or a group of atoms bonded to an inner atom • Steric number – the sum of the number of ligands plus the number of lone pairs; in other words, the total number of groups associated with that atom.

14. KNOW THESE!

15. Lone Pairs Take up a Bit More Space… Experiments show that sulfur tetrafluoride has bond angles of 86.9° and 101.5 °. Give an interpretation of these bond angles NOTE: Sizes and electronegativities of exterior atoms also affect bond angles!!!