Why Compounds Form
• Nobel gases do not form bonds because they have completely filled energy levels and do not need to gain or lose electrons.
• Other elements do form bonds and they do it so they too can have nobel gas electron configurations.
Valence
• Valence electrons = electrons in the highest energy level. These are the electrons that form chemical bonds.
• Sulfur– 1s2 2s2 2p6 3s2 3p4
• Oxygen– 1s2 2s2 2p4
Ga = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
Determining Valence Electrons
• 1. Write down the electron configuration• 2. Count how many electrons are in the highest s and
p orbitals. You should get a # between 1 and 8. • 3. These are the valence electrons.
2 + 1 = 3
Lewis Dot Diagrams• Represent the valence e- for an atom using dots.• Start at the top and place e- on each side going
clockwise. Once there are e- on each side, begin to pair them up.
Sodium 1 valence e-
Magnesium2 valence e-
Aluminum 3 valence e-
Silicon 4 valence e-
Phosphorus5 valence e-
Sulfur 6 valence e-
Chlorine 7 valence e-
Argon 8 valence e-
Bonding
• When chemical bonds form, some valence electrons are either shared or transferred between atoms.
• Only unpaired electrons participate in chemical bonds.
• Each unpaired electron can form 1 “covalent” bond. – Electrons are shared between them to complete
their outer shell.
The Octet Rule
• Elements share or transfer electrons in chemical bonds to reach a stable configuration of 8 valence electrons.
Bonding Using Lewis Dot Diagrams
and Share a pair of e- to form
So that each atom has a full outer energy level
A pair of shared e- can also be represented with a dashed line.
Exceptions to the Octet Rule• In covalent bonds, atoms always share e- to reach
a full valence shell of 8 valence e-…except…– Hydrogen only needs 2 e- in its outer energy level.– Boron only needs 6 e- in its outer energy level.– Some elements can form an expanded octet using
empty d-orbitals to form bonds and have more than 8 valence e-.
5 Steps for Drawing Lewis Structures1. Count the total number of valence electrons
for all atoms.2. Attach each atom to the central atom with a
single bond (single bond = 2 shared electrons)
3. Complete the octet for the attached atoms by adding pairs of non-bonding electrons.
4. Complete the octet for the central atom by adding pairs of non-bonding electrons
5. Count the total number of electrons in your structure and compare to step one. – If the number of e- is the same, it is correct.– If you used too many e-, add double
bond(s) and check your total again (usually add one double bond for each two electrons that you are over the total).
– If there are extra electrons left over, add them as non-bonding electrons on the central atom. This is called an expanded octet.
Practice Exercises
• http://chemsite.lsrhs.net/bonding/flashLewis.html
• Let’s find the Lewis Dot Diagrams for:– H2
– H2O– CH4
– SO– NH3
– CO2
Valence Shell Electron Pair Repulsion Theory
• Abbreviated “VSEPR”• Pairs of e- around an atom repel each other
and will form an arrangement that minimizes this repulsion (i.e. spread as far apart from each other as possible). As a result, molecules tend to form predictable shapes.
• Lone pairs of non-bonding e- have greater repulsion than bonded pairs of e-.
Basic VSEPR Geometries Molecular Geometry ABE Notation Atoms bonded
to the central atom
Non-bonding pairs on the central
atom
Tetrahedral AB4 4 0
Trigonal Pyramid AB3E 3 1
Bent AB2E2 or AB2E 2 1 or 2
Linear AB2 2 0
Trigonal Planar AB3 3 0
Expanded Octet GeometriesMolecular Geometry ABE Notation Atoms bonded to
the central atomNon-bonding pairs
on the central atom
Octahedral AB6 6 0
Square Pyramid AB5E 5 1
Square Planar AB4E2 4 2
Trigonal Bipyramid
AB5 5 0
Seesaw AB4E 4 1
T-Shaped AB3E2 3 2
Linear AB2E4 2 4