1. Covalent Bonding Covalent Bond: a bond where atoms share electrons
instead of gaining or losing electrons
- A covalent bond that shares one pair of electrons
is a single covalent bond. When two pairs of
electrons are shared is a double covalent
bond…
Structural Formulas: how a formula is written to show
a covalent bond (a dash is used to represent
the two shared electrons)
Example: H-Cl

2. - Covalent compounds are called molecules
(subscripts are not necessarily in the lowest
whole number ratio like ionic compounds)
- The octet rule still applies (so each element still
wants to have eight valance electrons) except
for hydrogen.
Unshared Pairs (Lone Pairs): pairs of electrons
that are not involved in a covalent bond

3. We learned to write electron dot notation so that
electrons that share an orbital will be written
on the same side of the elemental symbol.
- When electrons are given a little bit of energy
electrons can move to higher energy
sublevels for bonding purposes
- This is why carbon makes four bonds

4. Double and Triple Covalent Bonds
Double Bonds: when two pairs of electrons are
shared between two atoms
Triple Bonds: when three pairs of electrons are
shared between two atoms
- Atoms will make double or triple bonds to
meet the requirements of the octet rule

5. Diatomic Molecules
Because of the need to have eight valance electrons or
obey the octet rule, some elements will never exist
in nature as lone atoms.
Diatomic Molecules – molecules consisting of two atoms
of the same type.
- H2, O2, F2, N2, I2, Cl2, Br2

6. Coordinate Covalent Bonds
Typically in a bond, each atom contributes one electron
to each bond.
Coordinate Covalent Bond: a covalent bond where
both electrons in the bond are contributed
from one atom.
- Atoms in polyatomic ions are coordinate
covalently bonded.

7. Bond Dissociation Energy: the energy it takes
to break a bond.
- The more energy it takes to break a bond
the stronger that bond is.

8. Resonance Structures
Sometimes it is possible to draw the dot structures
for one molecule in more than one way.
- The molecule does not flip back and forth
between the two structures; it is an average
between the two.
- The more resonance structures a molecule
has the more stable that element is.

9. Exceptions to the Octet Rule
When electrons share an orbital the have opposite
spins. If a molecule has only paired electrons
the spins will cancel out and it will be
diamagnetic
Paramagnetic: when a molecule has unshared
electrons the spin of the electrons will not
cancel each other out.
- These molecules show a strong attraction
in a magnetic field

10. Other Exceptions to the Octet Rule
Other atoms may expand there octet and take on 10
or 12 electrons.

11. Properties of Covalent Molecules
- nonmetalic
- Could be a solid, liquid, or gas
- Low melting points
- Low solubility
- Poor electrical conductors

12. Bonding Theories
VSEPR Theory:
Valence Shell Electron Pair Repulsion Theory
- Electron pairs repel
- Electron pairs are as far away from other
pairs of electrons as possible in the
geometric structure of the molecule

13. Shapes of Molecules
Because each pair of electrons wants to be as far from
another pair as possible, we can predict the
shape of the molecule.
Possible shapes:
Linear:
Pyramidal:
Tetrahedral:
Bent:
Trigonal Planar:
Trigonal Bypyramidal:

14. Hybridization: when electrons move between two
different orbitals to make “hybrid” orbitals

15. Polar Bonds
Nonpolar Covalent Bonds: the atoms pull the
electrons toward them equally
Example: the diatomic molecules
Polar Covalent Bonds: two different atoms pull
(share) the electrons unequally
- Polarity all depends on the electronegativity that
each atom has in the molecule or their
attraction for electrons.

16. Look at page 405 in your book.
- This is a table of electronegativities
- To determine if the bond between two atoms
is polar subtract the electronegativities of
the two atoms involved in the bond.

17. Polar Molecules: molecules will become polar if one
end of the molecule is more
electronegative then the other end
- All of the electrons will be pulled to the atom
that is most electronegative
- This will make the molecule negatively
charged on one end and positively
charged at the other end.

18. Rules of Thumb for Polar Molecules
- If a bond in a two atom molecule is polar then the
molecule is polar.
- If the molecule is more than two atoms, then you
must also look at the shape to determine if the
molecule is polar.
- Symmetric shapes, where the polar bonds cancel
each other out lead to nonpolar molecules.

19. Attractions Between Molecules
Molecules have attractions towards other molecules.
- These attractions are not as strong as bonds
Van der Waals Forces:
1. Dispersion Forces: the attractions that occur
with nonpolar molecules
- This happens because the nucleus of
each atom is attracted to the electrons
of other molecules

20. 2. Dipole Attractions: the attractions between
polar molecules.
- The slightly positive end of one molecule
is attracted to the slightly negative
end of another molecule.

21. Hydrogen Bonds (a type of dipole attraction):
Attractive forces that occur between the
hydrogen of one molecule and the highly
electonegative element of another
molecule when hydrogen is covalently
bonded to a highly electronegative
element.

22. This is why water has such an unexpected high
boiling point and why it gets larger when it
freezes.
- To boil water you have to add more energy to
break all of the hydrogen bonds.
- When water freezes the molecules of water
are all organized with hydrogen bonds. This
makes water expand when freezing.
Hydrogen bonding is the strongest of the
intermolecular forces.

23. Although most molecular compounds are unstable
there is an exception:
Network Solids (Network Crystals): solids in which
all the atoms are covalently bonded to
eachother.
- These are very hard substances such as
diamonds