1. Covalent Bonding and Molecular Compounds
Chemical Bonding
Covalent Bonding and Molecular Compounds

2. A molecule is a neutral group of atoms that are held together by covalent bonds.
It can consist of two or more atoms of the same element (like oxygen) or the atoms of two or more elements combined (like water).

3. Any compound whose simplest units are molecules is called a molecular compound.
A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound.
H2O
C6H12O6

4. A diatomic molecule is a molecule containing only two atoms.
There are 7 elements which form diatomic molecules:
Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine

5. A covalent bond is formed when potential energy is at a minimum
Bonding occurs because most atoms have lower potential energy when they are bonded to other atoms.
A covalent bond is formed when potential energy is at a minimum
That minimum occurs where there is a balance between attractive and repulsive forces

6. Characteristics of covalent bonds
The distance between two bonded atoms is the bond length
Bond energy is the energy required to break a chemical bond
In general, shorter bond lengths have higher bond energy.

7. The Octet Rule
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.

8. As with most rules, there are exceptions to the octet rule.
Hydrogen forms bonds in which it is surrounded by only 2 electrons
Boron tends to form bonds in which it is surrounded by 6 electrons

9. Sometimes an element can be surrounded by more than 8 electrons
In these cases of expanded valence, bonding involves electrons in d orbitals as well as the s and p orbitals
Some examples are PF5 and SF6

10. Electron Dot Notation
Covalent bond formation usually involves only the valence electrons
Electron dot notation shows just the valence electrons of an atom with the electrons pictured as dots surrounding the symbol:

11. Lewis Structures
Electron dot notation can be used to represent molecules
A pair of dots between two elements is the shared electron pair of the covalent bond
The bond can also be represented as a dash. Each dash counts as two electrons

12. Lewis structures are formulas in which:
atomic symbols represent nuclei and inner-shell electrons
dot-pairs or dashes between two symbols represent electron pairs in covalent bonds
dots adjacent to only one atomic symbol represent unshared electrons

13. Lewis structures can be drawn if you know the composition of the molecule and which atoms are bonded to each other.

14. Drawing Lewis structures
1. Determine the type and number of atoms in the molecule
2. Count up the total number of valence electrons in the molecule
3. Arrange the atoms and draw covalent bonds
If carbon is present, it always goes in the middle
Hydrogen never goes in the middle
If carbon is not present, the least electronegative element goes in the middle

15. 4. Add unshared pairs of electrons so that each is surrounded by an octet (except H which only gets 2 electrons)
5. Count the electrons to be sure that the number in the structure is the same as the total number of valence electrons that you calculated in step 2
If not, you need a double or triple bond in your structure

16. H2O

17. Cl2

18. CH3I

19. A double bond is a covalent bond in which two pairs of electrons are shared between two atoms.
A double bond is represented by two parallel dashes

20. A triple bond is a covalent bond in which three pairs of electrons are shared between two atoms
A triple bond is represented by three parallel dashes

21. Double and triple bonds are called multiple bonds
Multiple bonds have shorter bond lengths and higher bond energies than single bonds
They are most commonly found in molecules that contain carbon, nitrogen, or oxygen

22. O2

23. CO2

24. N2

25. Resonance Structures
Some molecules and ions cannot be represented by just one Lewis structure
These molecules have a multiple bond which can be placed in two or more places in the molecule.
Example - ozone

26. O3

27. Chemists used to think that ozone spent its time alternating between the two structures, or resonating between them.
Now chemists know that the bonds in ozone are identical to each other
Ozone’s structure is actually an average of the two possible Lewis structures

28. Together, the structures are called resonance structures or resonance hybrids
To represent resonance, draw both Lewis structures and put a double-headed arrow between them.

29. SO2