1. Nature of Covalent Bonding
Part 3: Coordinate Covalent Bonds & Polyatomic Ions

2. Objectives
Demonstrate how electron dot structures represent shared electrons
Distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy
Describe how oxygen atoms are bonded in ozone

3. Coordinate Covalent Bonds
Is a covalent bond in which one atom contributes both bonding electrons
In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms
For example: carbon monoxide (CO)

4. Polyatomic Ions
A polyatomic ion is a group of 2 or more atoms joined by covalent bonds that has a positive or negative charge and behaves as a unit
Most polyatomic cations and anions contain both covalent and coordinate covalent bonds
Compounds containing polyatomic ions also include both ionic and covalent bonding

5. Ammonium Ion Formation
Forms when a positively charged hydrogen ion (H+) attaches to the unshared electron pair of an ammonia molecule (NH3)

6. Drawing Lewis Structures for Polyatomic Ions
Is different in that we must account for the charges on them
For example: SO32-
Step 2
Step 1
Step 4
Step 3

7. Examples
ClO3‒ H3O+

8. Practice Problems #3
NO3- PO43- NH4+ CO32-

9. Bond Dissociation Energy
Is the energy required to break a bond between two covalently bonded atoms
The stronger the bond the more energy is required to break it
Bond energy is measured in kJ/mol
Examples:
H2 requires 435 kJ/mol to break
C‒C requires 347 kJ/mol to break

10. Resonance
Some molecules can not be represented by a single Lewis structure
Resonance structures are two or more possible configurations for the same compound that differ in the arrangement of electrons
For example: ozone, O3
The actual bonding of oxygen atoms in ozone is a hybrid, or mixture, of the extremes represented by the resonance forms

11. Exceptions to the Octet Rule
The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number
There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons
For example: nitrogen dioxide has a total of 17 valence electrons

12. Exceptions to the Octet Rule
Several molecules with an even number of electrons fail to follow the octet rule
For example: boron trifluoride (BF3)
The boron atom within this compound is deficient by two electrons

13. Exceptions to the Octet Rule
A few atoms, especially phosphorus and sulfur, can expand the octet to include 10 or 12 electrons