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Lecture 13

CHEM 130 Lecture 13: chem lect 13Premium

7 pages57 viewsWinter 2019

Department
Chemistry
Course Code
CHEM 130
Professor
Charles Mc Mory
Lecture
13

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Monday, February 11, 2019
LECTURE 13
Rules for drawing lewis structures
1) Count the valence electrons
2) Place the least electronegative atom in the center
3) Form a bond (2 electrons) between each pair of bond atoms
4) Place lone pairs (2 electrons) on each terminal atom to complete their octet
5) Place any remaining electrons as lone pairs on the central atom
6) If the central atom octet is incomplete, borrow electrons from surrounding atoms to make
double or triple bond
Rows 3 and below can have extended octet
Resonance;
Sometimes more than one equally good lewis structure can be written for a particular molecule
In this case: the correct structure is the average of all possible structures
Resonance: The correct description of a resonant molecule is given by the superposition of
multiple structures.
Represented by double-headed arrows
EX. NO3-
24 valence electrons
One octet does not
make full octets so
make a double bond
equally valid
structure can have
double bound on
different sides side
1
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Monday, February 11, 2019
Formal Charge:The hypothetical charge on each atom.
1) obtained by assuming:
1) That bonding electrons are equally shared between bonded atoms
2) non-bonding electrons belong completely to one atom
Formal charge= valence e- of the free atoms- 1/2 bonding electrons- non-bonding electrons
NOTE: the sum all the formal charges on the atoms in a molecule must equal the total charge
on molecule
Computing the formal charge of an atom
ex. Compare the normal charges or the following CO2 structure:
C: 4 valenece electrons, 8 bonding electrons, non lone pairs
C: 4 -1/2(8) -0 = 0
O: 6 valence electrons, 4 bonding electrons, 4 none-bonding electrons
O: 6 - 1/2(4) -4 = 0
Total charge on molecule = 0
This is what is expected because there is on charge o molecule
C: 4 valence electrons, 8 bonding electrons, no lone pairs
C: 4 -1/2(8) -0 = 0
O (on left): 6 valence electrons, 2 bonding electrons, 6 non-bonding electrons
O: 6- 1/2(2) - 6 = -1
O (on right): 6 valence electrons, 6 bonding electrons, 2 non-bonding electrons
O: 6- 1/2(6) - 2 = +1
Total charge onn molecule = 0
Draw minus one O on left and plus on O on right
First structure is more preferred
Final rule of lewis structures: we want to minimize number of formal charges
Choosing right structure with formal charges
1) Atoms in a molecules try to achieve charges as close to zero as possible
2) Any negative formal charges reside on the most electronegative atoms
Electronegative atoms like additional electron density (better at accommodating a negative
charge )
3) When possible, choose lewis structures that do not have like charges on adjacent atoms
2
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