Here we will go over the covalent bond and how to name them.

1. Warm Up:
• Show me the electron transfer
between magnesium and chlorine
• What would the formula be?
• What is the name of this
compound?

2. Covalent Bonding
• Covalent Bond: Valence electrons shared
between two atoms.
• Molecule – a group of covalently bonded
atoms.
– Molecules usually consist of a nonmetal bonded
to a nonmetal or a metalloid bonded to
nonmetal.

3. • Example: H2
• s-orbital of one hydrogen atom overlaps with the s-
orbital of an other hydrogen atom.
• A dash is used to represent a shared pair of
electrons.
• Example: Cl2
• Single covalent bond – share one pair of electrons
between two atoms.

4. Drawing Molecules
• We can use electron dot notation to help
show how the atoms in a molecule are held
together.
• These drawings are known as Lewis
structures.
• All atoms seek to fill their outer energy shell,
most follow the Octet rule, where they will
have 8 electrons on their outer level.
– Hydrogen only wants 2, as it’s in the 1s energy
level.

5. Multiple Covalent Bonds
• Sometimes it is possible to form more than
one covalent bond between two atoms.
• Example: O2
• Double covalent bond – share two pairs of
electrons between the same two atoms.

6. • Example: N2
• Triple covalent bond – share three pairs of
electrons between the same two atoms.
• Double and triple covalent bonds are
referred to as multiple covalent bonds.

7. LEWIS STRUCTURES WITH MULTIPLE
COVALENT BONDS
• Example: CH2O
• If all electrons have been used and the
central atom is not stable, then consider a
multiple bond.
• Convert two dots to a dash.
• Example: CO2

8. Polyatomic Ions
• Group of covalently bonded atoms that has a charge.
(a charged molecule)
• Example: SO4
-2 (sulfate ion)
• These bond ionically with cations

9. Multiple Covalent Bonds
• Sometimes it is possible to form more than
one covalent bond between two atoms.
• Example: O2
• Double covalent bond – share two pairs of
electrons between the same two atoms.

10. • Example: N2
• Triple covalent bond – share three pairs of
electrons between the same two atoms.
• Double and triple covalent bonds are
referred to as multiple covalent bonds.

11. LEWIS STRUCTURES WITH MULTIPLE
COVALENT BONDS
• Example: CH2O
• If all electrons have been used and the
central atom is not stable, then consider a
multiple bond.
• Convert two dots to a dash.
• Example: CO2

12. Naming Binary Molecular Compounds
• Unlike ionic compounds, molecular compounds
are composed of individual covalently bonded
units, or molecules.
• The old system (prefix system) of naming
molecular compounds is based on the use of
prefixes.
– examples: CCl4 — carbon tetrachloride (tetra- = 4)
CO — carbon monoxide (mon- = 1)
CO2 — carbon dioxide (di- = 2)

13. Prefixes for Naming Covalent Compounds

14. Independent Practice
a. Give the name for As2O5.
b. Write the formula for oxygen difluoride.

15. Polar and Nonpolar Covalent Bonds
• Covalent bonding involves sharing
electrons between atoms.
• Nonpolar covalent bond – equal sharing of
electrons between two atoms.
– Both atoms have same attraction for shared pair.
– Example: H – H

16. • Polar covalent bond – unequal sharing of
electrons between atoms.
– One atom has greater attraction for shared pair.
(Electronegativity Tug – of – War)
– Example: H – Cl
• This creates partial (d) charges on each
atom in the bond.

17. • The atom that has a greater attraction for
shared electrons takes on a partial negative
charge. The atom with a weaker attraction
takes on a partial positive charge.
– How can you determine which atom has greater
attraction for electrons?
– ELECTRONEGATIVITY