Lecture 9sf

# Lecture 9sf - Lewis Structures Lewis structures show are...

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Lewis Structures Lewis structures show are representations of the structures of molecules, using the symbols of the atoms with lines and dots. Each dot represents a valence electron, and each line represents two valence electrons making a covalent bond between atoms. Double bonds are shown with two lines representing two electron pairs; triple bonds are shown with three lines representing three electron pairs. Some Lewis structures are simple enough to draw quickly at first sight. Examples of simple Lewis structures are: H 2 H-H Cl 2 HCl H 2 O NH 3 Most structures are more complicated than these examples, and require some thought and a methodical approach. Such an approach can be summarized with rules and guidelines:

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Rules for Writing Lewis Structures 1. Draw the skeleton of the molecule. 2. Count the total number of valence electrons that should be in the molecule, and subtract the electrons already in the bonds of the skeleton. This difference equals the number of additional electrons needed. 3. Add electron pairs to the molecule to complete octets around each atom, starting with the side atom and ending with the central atom, until the correct number of electrons have been added. If necessary, expand the octet around the central atom. 4. If the central atom has less than the complete octet, convert one or more pairs of electrons from the side atoms into double bonds until the central atom has its octet. 5. Check all formal charges. If the central atom is allowed to expand its octet and has a positive formal charge, you may convert more pairs of electrons from the side atoms into double bonds in order to minimize formal charge on the central atom. Check: The sum of the formal charges should equal zero for a molecule, or else equal the total charge on an ion. Let’s apply these rules to a simple molecule, NH 3 . (1) Draw the skeleton (2) Count the valence electrons: 5 for nitrogen, 1 each for hydrogen = 8 total Subtract 6 electrons for the 3 bonds, meaning you need two more electrons. (3) Add electron pairs to complete octets. Two more electrons needed: add to the central nitrogen atom, which gives it an octet—4 pairs of electrons around it: 3 bonds + one lone pair. (4) Central atom has its octet, so no further adjustment is necessary. (5) Check formal charges.
Formal Charge Formal charge is a way of checking electron distribution in a Lewis structure. Count the electrons surrounding an atom in a structure, and compare it with the valence electrons in the unbonded atom. If they are equal, the formal charge is zero. If there are extra electrons around the atom in the structure (compared with the unbonded atom), there is a negative formal charge, -1 for each extra electron. If there are fewer electrons around the atom in the structure (compared with the

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## This note was uploaded on 02/20/2012 for the course 160 161 taught by Professor Kim during the Fall '08 term at Rutgers.

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Lecture 9sf - Lewis Structures Lewis structures show are...

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