Lewis Structures

Lewis Structures • •

One of the most important concepts in chemistry. Used as a basis for describing shape, properties and reactivity of compounds. There is no absolute set of rules to construct Lewis structures.

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Learn through practice! A Strategy for Drawing Lewis Structures of Polyatomic Molecules (used by your textbook) 1. 2.

3. 4. 5. 6. 7.

Calculate the total number of valence electrons. Write atomic symbols for the central and terminal atoms.  Hydrogen atoms are always terminal  Central atoms are generally (not always) those with the lowest electronegativity Put in single bonds between the atoms in the structure. For each bond, subtract two from the total number of valence electrons. Using the remaining valence electrons complete the octets of the terminal atoms, and then as far as possible the central atoms. If you have used all the valence electrons, you are done. If not, go to step 7. Complete the octets of central atoms by moving lone-pairs from terminal atoms to form multiple bonds.

Another strategy for drawing Lewis Structures… 1.

2.

Determine which atoms are central and terminal.  Hydrogen atoms are always terminal  Central atoms are generally (not always) those with the lowest electronegativity Write the atomic symbol(s) for the terminal atoms surrounding the central atom.

3.

Place the number valence electrons (VE) around each atom as dots. Any electrons in excess of 4 shall be paired (note: like filling 2s and 2p of H-atom). NOTE: For p-block elements, this will be the Group number minus 10. i.e. N is Group 15  5 VE – 1 lone pair; needs 3 bonds Si is Group 14  4 VE – all unpaired; needs 4 bonds Te is Group 16  6 VE – 2 lone pairs; needs 2 bonds Br is Group 17  7 VE – 3 lone pairs; needs 1 bond Unpaired electrons on the central atom are paired with unpaired electrons on the terminal atoms to form bonds. Rearrange electrons if necessary so that each atom should have an octet of electrons. This may involve forming multiple bonds using lone pairs from step 3 to achieve the octet for adjacent atoms. Hypervalence – observed for elements with n = 3 or higher; they can form expanded octets (i.e. SF6, POCl3, etc.). Phosphorus forms PCl3 and PCl5; sulfur forms SF2, SF4 and SF6. Expanded octets result from transforming a lone pair (from step 3) into two unpaired electrons and subsequently forming two additional bonds. These bonds can be single or multiple bonds (i.e. SO42- has two S-O single and two S=O double bonds).

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Other considerations… Formal Charge on an atom = number of valence electrons – 0.5x(number of bonding electrons) – number of lone pair electrons

If more than one Lewis structure may be drawn use the following rules to determine the most reasonable structure: 1. 2. 3.

Look for the structure in which formal charges are minimized. Look for the structure in which negative formal charges are on electronegative atoms. Look for structures that benefit from resonance stabilization.