1. Chapter 8
Covalent Bonding

2. Bond Type Using Electronegativities
Nonpolar
(0 – 0.3)
Covalent
Polar
(0.3 – 1.7)
Ionic
Ionic
(1.7 – 3.3)
Covalent Bonds
atoms are joined by sharing of electrons.
Nonpolar-covalent: an equal sharing of electrons/an even distribution of electron clouds.
Polar-covalent: an unequal sharing of electrons/an uneven distribution of electron clouds.

3. “molecule”: a neutral group of atoms held together by covalent bonds.
Ionic Bonds: a bond that results from a mutual electrical attraction between a cation and an anion.
Na + Cl
“molecule”: a neutral group of atoms held together by covalent bonds.
“diatomic molecule”: a molecule containing only two atoms.
ex: O2, H2, N2

4. Electron Dot Diagrams
shows how an atom’s valence electrons become involved in bonding.
Hydrogen (1 valence electron)
Beryllium (2 valence electron)
Nitrogen (5 valence electron)
Carbon (4 valence electron)
Oxygen (6 valence electron)
Neon (8 valence electron)

5. Bonding site
Unshared
Pairs/
Lone Pairs
Unshared pairs (lone pairs): a pair of electrons that
are not involved in bonding (exclusive only to that atom).
Lewis Structures
Shows how atoms are bonded in a molecule.
Draw the dot diagrams for each element first,
then “connect the dots”.

6. Ex: H2O
Ex: BF3
Ex: C2H4

7. Atoms of metals lose valence electrons when forming an ionic compound.
Atoms of nonmetals gain electrons.
Enough atoms of each element must be used in the formula so that electrons lost equal electrons gained.

8. Solve Revision of ionic bond formation2
a. Start with the atoms.
• • •
K and O
• • • •

9. Solve Apply the concepts to this problem.2
a. In order to have a completely filled valence shell, the oxygen atom must gain two electrons. These electrons come from two potassium atoms, each of which loses one electron. K •
• • O + K+ 2–

10. Bonding and Interactions
The electrostatic forces between the oppositely charged ions hold the cations and anions together in an ionic compound.
Ionic compounds generally have high melting points and can conduct an electric current in solution and in the molten state.

11. O O B Cl Molecular Shapes
VSEPR Theory (Valence Shell Electron Pair Repulsion)
pairs of electrons are arranged as far apart from each other as possible.
Linear: 180o between bonds
Ex: O2
Linear
O O
2. Trigonal Planar: “flat triangle” – one central atom surrounded by three atoms with NO LONE PAIRS.
Ex: BCl3 B Cl
Trigonal
Planar

12. C H N H 3. Tetrahedral: four atoms bonded to a central atom.
4. Trigonal pyramidal: three atoms bonded to a central atom with one lone pair of electrons about the central atom. N H

13. 5. Bent: a central atom bonded to two other atoms AND one or two LONE PAIRS.
Hybridization
see overheads

14. Molecular Polarity
Polar molecules occur because of their uneven distribution of charge. The negative end will be at the more electronegative atom.
If bond polarities extend equally and symmetrically in different directions, then the molecule as a whole will be nonpolar
polarity of a molecule can be determined by the arrangement of the atoms in a molecule, better known as…………….
MOLECULAR SHAPE!

15. Nonpolar Linear Polar Polar Bent Nonpolar Polar Polar Pyramidal
Atoms are
Identical
Not identical
Nonpolar
Polar
Bent
Polar
Trigonal
Planar
Atoms are
Identical
Not identical
Nonpolar
Polar
Pyramidal
Polar
Tetrahedral
Atoms are
Identical
Not identical
Nonpolar
Polar

16. Hybridization: The blending of Orbitals+ =
Poodle +
Cocker Spaniel =
Cockapoo + =
s orbital +
p orbital =
sp orbital

17. What Proof Exists for Hybridization?
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent
bonds.
Lets look at a molecule of methane, CH4.
Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen atoms covalently bonded around it.

18. Carbon ground state configuration
What is the expected orbital notation of carbon
in its ground state?
Can you see a problem with this?
(Hint: How many unpaired electrons does this
carbon atom have available for bonding?)

19. Carbon’s Bonding Problem
You should conclude that carbon only has TWO
electrons available for bonding. That is not not enough!
How does carbon overcome this problem so that
it may form four bonds?

20. Carbon’s Empty Orbital
The first thought that chemists had was that carbon promotes one of its 2s electrons…
…to the empty 2p orbital.

21. However, they quickly recognized a problem with such
an arrangement…
Three of the carbon-hydrogen bonds would involve
an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

22. This would mean that three of the bonds in a methane
molecule would be identical, because they would involve
electron pairs of equal energy.
But what about the fourth bond…?

23. The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.

24. This bond would be slightly different in character than
the other three bonds in methane.
This difference would be measurable to a chemist
by determining the bond length and bond energy.
But is this what they observe?

25. The simple answer is, “No”.
Measurements show that
all four bonds in methane
are equal. Thus, we need
a new explanation for the
bonding in methane.
Chemists have proposed an explanation – they call it
Hybridization.
Hybridization is the combining of two or more orbitals
of nearly equal energy within the same atom into
orbitals of equal energy.

26. In the case of methane, they call the hybridization
sp3, meaning that an s orbital is combined with three
p orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy than
the 2s orbital…
… and slightly LESS energy than the 2p orbitals.

27. sp3 Hybrid Orbitals
Here is another way to look at the sp3 hybridization
and energy profile…

28. sp Hybrid Orbitals While sp3 is the hybridization observed in methane,
there are other types of hybridization that atoms
undergo.
These include sp hybridization, in which one s
orbital combines with a single p orbital.
This produces two hybrid orbitals, while leaving two normal p orbitals

29. sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.
One p orbital remains unchanged.

30. Intermolecular Forces
the higher the force of attraction between molecules, the harder it is to pull them apart.
high BP = large attractive forces
low BP = small attractive forces
the strongest intermolecular forces act between polar molecules.
Dipole: formed due to differences in electronegativity.
A dipole is represented by an arrow pointing towards the negative pole (more electronegative element).

31. dipole-dipole forces: the negative region of one molecule is attracted to the positive region of another molecule.
Balloon-
Water
Explanation
a polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons  “induced dipole”.
Induced
Dipole

32. Hydrogen Bonding
a strong type of dipole-dipole force that accounts for the unusually high BP of hydrogen containing compounds. O H
London Dispersion Forces
a weak intermolecular attraction resulting from instantaneous dipoles.
all molecules experience London Forces, however they are the only force of attraction among noble-gases, nonpolar, and slightly polar molecules.
London
Forces