Advanced Lewis Structures
Advanced Lewis Structures
Polyatomic Ions
Polyatomic ions are electrically charged groups of two or more chemically bonded atoms that act as a single unit.
When determining the total valence electrons for a polyatomic ion, the charge must be considered.
Example: Hydroxide (OH⁻)
Oxygen contributes 6 valence electrons.
Hydrogen contributes 1 valence electron.
Since the ion has a negative charge, it gained an electron.
Total electrons:
Lewis structure:
[ O - H ]⁻
Example: Ammonium (NH₄⁺)
Nitrogen contributes 5 valence electrons.
Each hydrogen contributes 1 valence electron, totaling 4 valence electrons.
Since the ion has a positive charge, it lost an electron.
Total electrons:
Lewis structure:
H | H - N - H | H[ ]⁺
Resonance
Resonance occurs when two or more valid Lewis structures can be drawn for the same molecule.
Example: Ozone (O₃)
Two resonance structures can be drawn.
The actual structure is a resonance hybrid, an average of all possible structures.
All structures have 2 identical O-TO-O bonds with a bond order of 1.5.
Resonance structure: A Lewis structure that is one of two or more possible Lewis structures for a molecule.
Delocalized electrons: Electrons shared between more than two atoms.
Delocalized electrons are found in resonance structures.
Example: Carbonate ion (CO₃²⁻)
Total valence electrons:
Resonance structures:
:O: || <-> :O: O - C - O <-> O - C - O | || :O: :O: -2 -2
Structural Isomers
Structural isomers have the same number and type of atoms but are arranged differently.
Structural isomers have different chemical and physical properties.
Example: C₂H₆O
Ethanol (C₂H₅OH):
H H | | H - C - C - O - H | | H HEther (CH₃OCH₃):
H H | | H - C - O - C - H | | H HEthanol and ether have different properties due to hydrogen bonding in ethanol.
Formal Charge
Formal Charge is the charge an atom would have if all atoms in a molecule had the same electronegativity (i.e., shared electrons equally).
Formal charge helps determine the most reasonable Lewis structure when multiple structures obey the octet rule.
Calculating formal charge:
All unshared electrons are assigned to the atom on which they are found.
Half of the bonding electrons are assigned to each atom in the bond.
Formula: Formal Charge = (# of valence electrons) - (# of assigned electrons in Lewis structure)
Note: Formal charges do NOT represent real charges on atoms.
General Rule: The most stable Lewis structure will have
Atoms with formal charges closest to zero.
Any negative charge residing on the more electronegative atom.
General Tips:
Carbon tends to have zero lone pairs.
Nitrogen tends to have one lone pair.
Oxygen tends to have two lone pairs.
Halogens tend to have three lone pairs.
Example: Carbonyl chloride (COCl₂)
Total valence electrons:
Determining Preferred Lewis Structure Using Formal Charge
Example: Carbonyl chloride (COCl₂)
Possible structures and formal charge calculations:
Option 1:
Cl - C = OC: 4 (valence e⁻) - 4 (assigned e⁻) = 0
O: 6 (valence e⁻) - 6 (assigned e⁻) = 0
Cl: 7 (valence e⁻) - 7 (assigned e⁻) = 0
Magnitude = 0
Option 2:
Cl - C - O - ClC: 4 (valence e⁻) - 5 (assigned e⁻) = -1
O: 6 (valence e⁻) - 5 (assigned e⁻) = +1
Cl: 7 (valence e⁻) - 7 (assigned e⁻) = 0
Magnitude = 2
Option 3:
Cl = C = OC: 4 (valence e⁻) - 6 (assigned e⁻) = -2
O: 6 (valence e⁻) - 4 (assigned e⁻) = +2
Cl: 7 (valence e⁻) - 7 (assigned e⁻) = 0
Magnitude = 4
Option 1 is preferred because all atoms have a formal charge of 0.
Another Example: NCS⁻
Total valence electrons:
Possible structures and formal charge calculations:
Option 1:
[ N - C ≡ S ]⁻¹N: 5 (valence e⁻) - 4 (assigned e⁻) = +1
C: 4 (valence e⁻) - 4 (assigned e⁻) = 0
S: 6 (valence e⁻) - 7 (assigned e⁻) = -1
Magnitude = 3
Option 2:
[ N = C = S ]⁻¹N: 5 (valence e⁻) - 6 (assigned e⁻) = -1
C: 4 (valence e⁻) - 4 (assigned e⁻) = 0
S: 6 (valence e⁻) - 6 (assigned e⁻) = 0
Magnitude = 1
Option 3:
[ N ≡ C - S ]⁻¹N: 5 (valence e⁻) - 5 (assigned e⁻) = 0
C: 4 (valence e⁻) - 4 (assigned e⁻) = 0
S: 6 (valence e⁻) - 7 (assigned e⁻) = -1
Magnitude = 1
The total formal charge in all three options is -1, matching the overall ion charge.
Option 2 is the best because it has the lowest magnitude of formal charges and places the negative charge on the more electronegative nitrogen atom.
Electronegativity values: N (3.0), C (2.5), S (2.5).
If it is not possible to form an octet, it typically indicates the presence of exceptions to the octet rule. These exceptions include:
Incomplete Octets: Some atoms, like boron (B) and beryllium (Be), are stable with fewer than eight electrons around them.
Odd Number of Electrons: Molecules with an odd number of valence electrons, such as nitrogen monoxide (NO), cannot satisfy the octet rule
If it is not possible to form an octet, it typically indicates the presence of exceptions to the octet rule. These exceptions include:
Incomplete Octets: Some atoms, like boron (B) and beryllium (Be), are stable with fewer than eight electrons around them.
Odd Number of Electrons: Molecules with an odd number of valence electrons, such as nitrogen monoxide (NO), cannot satisfy the octet rule
Formal Charge is the charge an atom would have if all atoms in a molecule had the same electronegativity (i.e., shared electrons equally). It helps determine the most reasonable Lewis structure when multiple structures obey the octet rule. To calculate it:
1. All unshared electrons are assigned to the atom on which they are found
2. Half of the bonding electrons are assigned to each atom in the bond.
Use the formula: Formal Charge = (# of valence electrons) - (# of assigned electrons in Lewis structure).
Remember, formal charges do NOT represent real charges on atoms. The most stable Lewis structure will have atoms with formal charges closest