1. Drawing Lewis Structures
Find the total # of valence electrons from all atoms
Anion: add e- (NO3- : add 1 e-) Cation: subtract e- (NH4+: minus 1 e-)
Predict the arrangement of the atoms
Place least electronegative atom in center
Usually written first in the formula (often C, never H)
Make a single bond (2 e-) between each pair of atoms
Use remaining electrons to satisfy octets (8 e- around each)
Place electrons around atoms in pairs (lone pairs)
b. Try forming multiple bonds between atoms:
double bond (4 e-) and triple bond (6 e-)
Check your structure by verifying that all e- have been used and that all atoms have a satisfied octet
Exceptions: Remember that H only needs 2e- for its “octet”

2. Lewis Structure Trends
Certain trends can be used to help predict the bonding behavior of particular atoms
C group
Forms combination of 4 bonds (4 singles, 2 doubles, etc.) and no LP (i.e. CH4)
N group
Forms combo of 3 bonds and 1 LP (i.e. NH3)
O group
Forms combo of 2 bonds and 2 LP (i.e. H2O)
F group (halogens)
Forms 1 bond and has 3 LP (i.e. HCl)
Note that these are NOT always true!

3. Lewis Structures and Formal Charge
Formal charge = (# valence e-) – (# lone e- + ½ # bonded e-)
or = (# valence e-) – (dots + bonds)
Sum of formal charges = overall charge on molecule/ion
Formal Charges
Water
H2O
H: 1 – 1 = 0
O: 6 – (4 + 2) = 0
Total charge = 0
Hydronium
H3O+
H: 1 – 1 = 0
O: 6 – (2 + 3) = +1
Total charge = +1

4. Resonance structures differ only in the position of the electrons
Show resonance
Show movement of e-
The actual structure is a hybrid (average) of the resonance structures
All three oxygen’s share part of the double bond
In reality, there are three bonds of equal length (in between a double and single bond)
Arrow formalism: curved arrows show electron movement

5. Predicting Molecular Shape: VSEPR
(Valence Shell Electron Pair Repulsion)
Electrons repel each other
The molecule will adopt a 3-D shape that keeps the electrons (lone pairs and bonded e-) as far apart as possible
Different arrangements of bonds/lone pairs result in different shapes
Using Lewis structures, the shapes will depend on the total number of bonds/lone pairs (“things”) and how many lone pairs are around the central atom

7. Carbon Dioxide: CO2 Lewis Structure Two “things” (bonds or lone pairs)
Linear geometry
0 LP → Linear Shape
180o Bond angle C O

8. Formaldehyde: CH2O Lewis Structure Three “things”
Trigonal planar geometry
0 LP → Trigonal planar shape
120° bond angles

9. Sulfur Dioxide: SO2 Lewis Structure Three “things”B A
Lewis Structure
Three “things”
Trigonal planar geometry
1 LP → Bent shape
120° bond angles

10. Methane: CH4 Lewis Structure Four “things” (bonds/LP)
Tetrahedral geometry
0 LP → Tetrahedral shape
109.5o bond angles

11. Ammonia: NH3 Lewis Structure Four “things” (bonds/LP)
Tetrahedral geometry
1 LP → Trigonal pyramid shape
107o bond angles

12. Water: H2O Lewis Structure 4 “things” (bonds/LP) Tetrahedral Geometry
2 LP → Bent Shape
104.5o bond angle

13. Hydrogen Chloride: HCl
Lewis Structure
Four “things” (bonds/LP)
Tetrahedral geometry
3 LP → Linear Shape
180o Bond angle Cl H

14. Electronegativity and Bond Type
The electronegativity difference between two elements helps predict what kind of bond they will form. H
2.1 Li
1.0 Be
1.5 B
2.0 C
2.5 N
3.0 O
3.5 F
4.0 Na
0.9 Mg
1.2 Al Si
1.8 P S Cl Br
2.8 I K
0.8 Ca
Electronegativity difference
≤ 0.4 
0.5 – 1.8
> 1.8
Bond type
Covalent 
Polar covalent
Ionic
Sample Bonds
NaCl
Cl-Cl
C-O
C-H
Electronegativities
3.0 – 0.9 = 2.1
3.0 – 3.0 = 0
3.5 – 2.5 = 1.0
2.5 – 2.1 = 0.4
Bond Type
Ionic
Covalent
Polar covalent

15. Dipole Moment
Occurs in polar covalent bonds where e- are unequally shared
One atom will attract more e- density than the other
One atom becomes partially – charged, and the other partially + charged
“+” charged
end
Arrow points toward “-” end δ+ δ-

16. Carbon Dioxide CO2 Lewis Structure Linear geometry
180o Bond angle
O is more electronegative than C
e- more attracted to O in each bond
Creates polar bonds (unequal sharing of e-)
O atoms pulling e- equally in opposite directions, so no net polarity or dipole formed
Intermolecular forces:
London dispersion forces

17. Trigonal planar geometry
Formaldehyde CH2O
Lewis Structure C H O
Trigonal planar geometry
120° bond angles
Polar C=O bond:
Net dipole moment
Intermolecular forces:
dipole-dipole

18. Methane CH4 Lewis Structure Non-polar covalent bonds No net dipole
Intermolecular forces:
London dispersion forces
Tetrahedral geometry
109o bond angles

19. Ammonia: NH3 Lewis Structure Trigonal pyramid shape 107o bond angles
No Polar Bonds
Shape is unbalanced
e- around LP, creating dipole

20. Intermolecular forces:
Water: H2O
Lewis Structure
Polar bonds
Net dipole moment
Intermolecular forces:
Dipole-dipole
H-bonding
Bent

21. Intermolecular Forces
Ion-ion (ionic bonds)
Ion-dipole
Dipole-dipole
Hydrogen bonding
London dispersion forces − +
 − − +
 −

22. Bonding and Electrons Each atom contains electrons in atomic orbitals
Each orbital can hold a max of 2 electrons E 1s 2s 2p
 Valence electrons
 Core electrons
Atomic orbitals are contained in energy levels (1, 2 etc.)
Inner energy levels = core electrons (a.k.a kernel e-)
Outermost energy level = valence electrons
Valence electrons are used for bonding
Ionic bonding: transfer of electrons
Covalent bonding: sharing of electrons

23. Atomic Orbitals and Bonding
Bonds between atoms are formed by electron pairs in overlapping atomic orbitals E 1s 1s 1s
Example: H2 (H-H)
Use 1s orbitals for bonding :
Example: H2O
From VSEPR: bent, 104.5° angle between H atoms
Use two 2p orbitals for bonding? E 2s 2p 2p 1s
How do we explain the structure predicted by VSEPR using atomic orbitals?
90°

24. Bonding Single bonds Double bonds
Overlap of bonding orbitals on bond axis
Termed “sigma” or σ bonds
Double bonds
Sharing of electrons between 2 p orbitals perpendicular to the bonding atoms
Termed “pi” or π bonds 2p 2p
Bond Axis of σ bond
One π bond

25. Hybrid Orbitals
Mix starting atomic orbitals together to form hybrid orbitals
Lone pairs
Bonds
with H
One s and three p orbitals on O mix to form four identical sp3 hybrid orbitals that point towards the corners of a tetrahedron
Two sp3 orbitals contain O lone pairs
Two form  bonds with H 1s orbitals

26. Hybrid Orbitals for Trigonal Planar Bonding
 bonds
One s and two p orbitals on C mix to form three identical sp2 hybrid orbitals pointing towards the corners of a triangle.
One p orbital remains unhybridized.
sp2 orbitals form  bonds with O and H atoms
Leftover p orbital forms  bond with O

27. Hybrid Orbitals for Linear Bonding
 bonds
 bonds
One s and one p orbital on C mix to form two identical sp hybrid orbitals pointing 180° away from each other
Two p orbitals remain unhybridized.
Each sp orbital forms a  bond with O
Each leftover p orbital forms a  bond with O

28. Determining Hybridization
To determine an atom’s hybridization, count “things” around the atom
“Things” = bonded atoms and lone pairs
Two things: sp hybrid orbitals
Three things: sp2 hybrid orbitals
Four things: sp3 hybrid orbitals
Methane
Acetonitrile sp
sp3
4 “things”
sp3 hybridization