1. COVALENT
BONDING
Chapter 9

2. Review….
 What is a chemical bond?
 Force that holds two atoms together
 What is an ionic bond?
 An electrostatic force that holds
oppositely charged particles together in
an ionic compound
 Compounds formed from metal &
nonmetal
 Forms when….?
 What are atoms always trying to
achieve?
 Stability
 Complete set of valence electrons…
OCTECT

3. What is a covalent bond?
 Chemical bond that results from sharing of
valence electrons
 Occurs b/w nonmetal & a nonmetal
 Balance b/w attractive and repulsive forces
2 Hydrogen Atoms
Sharing their 1 Ve-

4. Molecules
 Compound made when 2 or more atoms are bonded
covalently
 Diatomic molecules
 In nature, sometimes two atoms of the same element are
more stable when they are covalently bonded than the
individual atom alone…
 BrINClHOF (pronounced “Brinkle Hoff”)
 Br2 I2 N2 Cl2 H2 O2 F2

5. Single Covalent Bonds
 A single covalent bond –Atom shares 1 pair (2)
electrons.
 Shared pairs – both elements count the electron pair to
achieve octet
 Lone pairs – pair of electrons that are not shared b/w
the atoms
 Lewis structures- Use electron dot diagram to show
how atoms are arranged in a molecule.
..
H Cl
..
Unshared or
Lone pair (LP)
shared or Bond pair

6. In the Fluorine Molecule…..
 How many bonding pairs are there in each?
 1
 How many lone pairs are there each?
 3

7. Multiple Covalent Bonds
 Double covalent bonds share two pairs of
electrons.
 CO2 O=C=O
 Triple covalent bonds share three pairs of
electrons.
 N2 :N=N:

8. Covalent Bond Formation in Hydrogen
 Increased overlap brings
the electrons and nuclei
closer together while
simultaneously
decreasing electron-
electron repulsion.
 However, if atoms get
too close, the
internuclear repulsion
greatly raises the energy.

9. The attractive and repulsive
forces in covalent bonding must
be balanced.

10. –Bond Length - In general, the closer the electrons are held
by the atoms, the shorter the bond length and the higher
the bond energy.
–Multiple bonds result in stronger, shorter bonds.

11. –BondEnergy - The amount of energy required to break a
bond. The greater the energy, the stronger the bond.
–Bond breaking is an endothermic process, so bond breaking
enthalpies are positive.

12. Comparing Bond Length and Bond
Strength
 Using the periodic table, but not Tables 9.2 and
9.3, rank the bonds in each set in order of
decreasing bond length and bond strength:
 (a) S F, S Br, S Cl
 (b) C = O, C O, C O

13. Sigma ( ) Bonds
 Sigma bonds are characterized by
 Head-to-head overlap.
 Cylindrical symmetry of electron density about the
internuclear axis.

14. Pi ( ) Bonds
 Pi bonds are
characterized by
 Side-to-side overlap.
 Electron density above
and below the
internuclear axis.

15. Single Bonds vs. Multiple bonds
Single bonds are always  In a multiple bond one
bonds, because of the bonds is a
overlap is bond and the rest are
greater, resulting in a bonds.
stronger bond and more
energy lowering.

16. Orbital overlap

17. Lewis Dot Structures
1. Determine the number of Valence e- for all atoms in
the molecule
a. Divide the Ve- by 2 to get pairs (2 dots or 1 line)
2. Decide on central atom (least electronegative or
furthest to the left).
a. Hydrogen & halogens are terminal atoms
3. Connect all atoms to the central atom by a bonding
pair (single line)
4. Place remaining pairs around all atoms before
moving on to central atom.
5. Check for octet (not H)
a. If atom does not have an octet, move lone pairs from a
terminal atom to create a double or a triple bond (except
grp 7).