REPRESENTING MOLECULES:
RESONANCE
STRUCTURES
Concept of Valence
• Bonding of atoms occurs in order to achieve a complete
set of valence electrons, similar to the valence electrons of
stable noble gases.
Covalent Bond:
Covalent bonds
- Double bond
- Covalent substances
• interaction between two non-metals,
• formed by sharing a pair of electrons generally a
single bond.
• Lewis structures or Lewis electron-dot structures
o two electron pairs are shared
o two lines are drawn in the Lewis
structure; for example, CO2
Covalent Bond: Lewis Structures
NCS-
No. of bonds = (Required Electrons – available electrons) / 2
Formal Charge = Valence electrons –( no. of bonds + unshared electrons)
Group Number is the number of valence electrons for each atom
Covalent Bond: Lewis Structures
Example:
1.
CO2
Rule no. 1: Identify the central atom
-least EN
Rule No. 2: put each dot (valence
electrons)
Remember: The group number of each
atom is the number of valence electrons
Connect the bonds
No. of bonds = (Required Electrons – available electrons) / 2
Formal Charge = Valence electrons –( no. of bonds + unshared electrons)
Covalent Bond: Lewis Structures
Example:
1.
CO2
Rule no. 1: Identify the central atom
-least EN
Rule No. 2: put each dot (valence
electrons)
Remember: The group number of each
atom is the number of valence electrons
Connect the bonds
No. of bonds = (Required Electrons – available electrons) / 2
Covalent Bond: Lewis Structures
Example:
No. of bonds = (Required Electrons – available electrons) / 2
1.
2.
Formal Charge = Valence electrons –( no. of bonds + unshared electrons)
CO2
O3
No. of bonds = (Required Electrons – available electrons) / 2
Covalent Bond: Lewis Structures
Example:
1.
O3
Rule no. 1: Identify the central atom
-least EN
Rule No. 2: put each dot (valence
electrons)
Remember: The group number of each
atom is the number of valence electrons
Connect the bonds
No. of bonds = (Required Electrons – available electrons) / 2
Formal Charge = Valence electrons –( no. of bonds + unshared electrons)
Covalent Bond: Lewis Structures
Example:
1. O3
We would expect the O-O bond in O3 to be longer
than the O=O bond because double bonds are
known to be shorter than single bonds.
Yet experimental evidence shows:
that both oxygen-to-oxygen bonds are equal in
length (128 pm).
We resolve this discrepancy by using both Lewis
structures to represent the ozone molecule:
No. of bonds = (Required Electrons – available electrons) / 2
RESONANCE STRUCTURE
• Resonance structures - placement of the atoms in alternative
but completely
• equivalent Lewis structures are the same, but the placement of
the electrons is different.
• A resonance structure, then, is one of two or more Lewis
structures for a single molecule that cannot be represented
accurately by only one Lewis structure.
• Resonance forms differ only in the arrangement of electrons.
RESONANCE STRUCTURE
•
Not all resonance structures are equal there are some that are
better than others.
The more resonance forms a molecule the molecule more stable.
•
The most important resonance form should have the smallest number
of formal charges.
•
The distribution of positive (+) and negative (-) FC should agree with
the electronegativity of the atom.
How can we draw the
important resonance
forms?
The concept of resonance
applies equally well to
organic systems.
Covalent Bond: Lewis Structures
Example:
Draw three resonance structures for the
molecule nitrous oxide, N2O
No. of bonds = (Required Electrons – available electrons)
/2
Formal Charge = Valence electrons –( no. of bonds +
unshared electrons)
Solution The three resonance structures are
Covalent Bond: Lewis Structures
Example:
No. of bonds = (Required Electrons – available electrons) / 2
N2O
Formal Charge = Valence electrons –( no. of bonds + unshared electrons)
Covalent Bond: Lewis Structures
No. of bonds = (Required Electrons – available electrons) / 2
Formal Charge = Valence electrons –( no. of bonds + unshared electrons)
Activity: HCN
EXCEPTIONS TO THE OCTET RULE
Recap:
o Typical Valency of Some Elements in Molecules
▪ Carbon – tetravalent; 4 bonds, no lone pairs
▪ Hydrogen – monovalent; 1 bond, no lone pairs
▪ Nitrogen – trivalent; 3 bonds, 1 lone pair
▪ Oxygen – divalent; 2 bonds, 2 lone pairs
▪ Halogen – monovalent; 1 bond, 3 lone pairs
Exceptions:
•
ionic compounds of transition metals
•
molecules and polyatomic ions containing an odd number of
electrons (NO)
•
fewer than an octet of valence electrons (BF3)
•
more than an octet of valence electrons (PF5, P being
hypervalent)
Recap:
Octet rule
▪ This rule is usually followed by elements found in the second row of the
periodic table.
▪ Organic compounds usually follow this rule in bonding.
▪ Exceptions may apply to non-second-row elements like
Hydrogen
1. Boron
2. Sulfur
3. Phosphorus
1.
BORON
BF3 (fewer octets)
2. Sulfur (more than octet)
SO3
Note:
To check the exceptions
Available electrons=
S
6
O3
6
= 24 electrons
3. Iodine
ICl4-
Note:
To check the exceptions
Available electrons=
I
7
Cl4
7
+ (ion -)
1
= 36 available electrons
4. Phosphorus (hypervalent atom)
PO43-
Note:
To check the exceptions
Available electrons=
P
5
O (4)
6
+3 (-3 Ion)
3
= 32 available electrons