Lewis Structures & Molecular Geometries

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-

Elements in the same family have the same number of valence e- and the same Lewis dot arrangement

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

Practice as a Class• CH4

• NH3

• H2O

• CO2

• SF6

• I3-

• AB4

• AB3E

• AB2E2

• AB2

• AB6

• AB2E3