1. Chapter 8
Covalent bonding

2. Valence Electrons
Elements with similar chemical behavior have the same number of valence electrons.
For the representative elements (1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A) the group number corresponds to the number of valence electron in each group (with the exception of He)
When examining electron configurations, the electrons that are present in the highest principle energy level represent the valence electrons of those atoms.
Br: [Ar] 4s2 3d10 4p Bromine has 7 valence electrons

3. Valence Electrons and Electron Dot Structures
Valence electrons are the electrons that participate in chemical bonds
Electron dot structures consist of the atom symbol and its valence electrons represented as dots.
Br: [Ar]4s2 3d10 4p5

4. Covalent Bonding Covalent bonds occur between two or more non-metals
Unlike ionic bonds where electrons are transferred from one atom to another, electrons are shared between atoms in a covalent bond.
Atoms joined together by covalent bonds are called molecules
A compound composed of molecules is called a molecular compound

5. Molecular and Structural Formulas
A molecular formula indicates the types and numbers of each atom in a molecule
The structural formula indicate the arrangement of the atoms in the molecule
H2O

6. Covalent Bonds and the Octet Rule
Atoms share electrons in a covalent bond so that each atom has enough electrons to satisfy the octet rule

7. Varieties of Covalent Bonds
Single bonds (sigma bonds)
One pair of electrons is shared between two atoms
Overlap of orbitals occurs between the two nuclei
Lone pair

8. Varieties of Covalent Bonds
Double Bonds (1sigma bond, 1 pi bond)
Atoms share two pairs of electrons
Overlap of orbitals in the pi bond occur off to the side of adjoining nuclei
Triple Bonds (1sigma bond, 2 pi bonds)
Atoms share three pairs of electrons

9. Coordinate Covalent Bonds
A covalent bond in which one atom contributes both bonding electrons.
Carbon

10. Resonance Structures
A condition when more than one valid Lewis structure can be written for a molecule or ion.

11. Exceptions to the Octet Rule
Too few electrons surrounding the central atom (ex: BH3)
Boron will not have a full octet, only 6 electrons. It can only achieve a full octet when another atom shares an entire pair of electrons with it (Coordinate covalent bonding)
Too many electrons surrounding the central atom (ex: PCl5)
An odd number of electrons

12. How to Draw a Lewis Structure for Molecules
Predict the location of atoms
If there are more than two atoms, place the least electronegative atom in the center and surround it by the remaining atoms.
Hydrogen is always terminal (outside) because it can only make one bond
Determine the total number of electrons if each atom had a full set of valence electrons (2 for H, 8 for all others)
Add up the number of valence electron that you have to work with
Subtract total valence electrons from total electrons and divide by two. This is the number of bonding pairs that are needed to put together the molecule.
Connect the atoms with the number of bonds that you calculated above
Add lone pairs where needed so that each atom has a full octet (except for hydrogen which can only have two electrons)
Molecule
Total Electrons
Valence Electrons
Bonding Pairs
HCN

13. Polyatomic Ions
Polyatomic ions are a cluster of non-metals that carry a charge.
To draw the structure of a polyatomic ion, follow the procedure for drawing ordinary molecules but add or subtract the number of electrons gained or lost to the total number of valence electrons in your structure as indicated by the charge on the ion.
Molecule
Total Electrons
Valence Electrons
Bonding Pairs
IO3-

14. Molecular Shape (VSEPR)
Valence Shell Electron Pair Repulsion – minimizes the repulsion of shared and unshared pairs of electrons around the central atom.
The shape of a molecule determines many of its physical and chemical properties.
The VSEPR is based on the arrangement of bonding and lone electrons around a central atom to minimize repulsion.
The repulsion of electrons creates a specific bond angle between a central atom and two terminal atoms.
Lone pairs of electrons occupy more space than bonding pairs of electrons

15. Molecular Geometry

16. Electronegativity and Polarity
Recall: Electronegativity is the ability of an atom to attract an electron.

17. Chemical bonding is like “Tug-o-War”
Electronegativity
Difference
Bond Type
Non-polar Covalent 0-0.4
Polar Covalent
Ionic >2.0

18. Molecular Polarity Linear Trigonal Planar Tetrahedral
Molecules are either polar or non-polar
Both polar and non-polar molecules may contain polar bonds. What determines whether a molecule is polar or non-polar is the symmetry of the molecule
PolarBonds
Present
Symmetry
Polar/
Non-Polar
Examples No
NO2
Yes
Non-polar
SiH4
Polar
NH3
CO2
VSEPR shapes that can demonstrate symmetry are:
Linear Trigonal Planar Tetrahedral

19. Naming Binary Covalent Compounds
If there is more than one of the 1st atom, precede the atom name by the appropriate prefix (di, tri, tetra, penta, hexa, hepta, octa, nona, deca)
Example: C6O2 hexacarbon dioxide
If there is only one of the first atom, do not precede the atom name by mono.
CO2 = monocarbon dioxide CO2 = carbon dioxide
Precede the second atom name by the appropriate prefix, including mono if there is only one of that atom. Drop the last syllable (or 2) and add –ide to the element name
C2O Dicarbon monoxide
2nd Element
Name
2nd Element C
Carbide S
Sulfide N
Nitride Cl
Chloride O
Oxide Se
Selenide F
Fluoride Br
Bromide P
Phosphide I
Iodide