1. Unit 5B: Covalent Bonding

2. Bonding Review C F Covalent Bonds (2 nonmetals)
…atoms share e– to get a full valence shell
C 1s2 2s2 2p2
F 1s2 2s2 2p5
*Both need 8 v.e – for a full outer shell (octet rule)!*
4 valence e-
7 valence e- o x x C x F o o x x x x o

3. Draw the Lewis dot structure for the following elements (write e- config first):Si O P B Ar Br
1s2 2s2 2p6 3s2 3p2
4 valence e-
1s2 2s2 2p4
6 valence e-
1s2 2s2 2p6 3s2 3p3
5 valence e-
1s2 2s2 2p1
3 valence e-
1s2 2s2 2p6 3s2 3p6
8 valence e-
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
7 valence e-

4. Notice any trends…? TRANSITION METALS H He
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar
TRANSITION METALS
K Ca Se Br Kr
Rb Sr Te I Xe
Cs Ba
The group # corresponds to the # of valence e–

5. F2 Let’s bond two F atoms together…
Each F has 7 v. e– and each needs 1 more e– F F F F2
Now let’s bond C and F atoms together…
carbon tetrafluoride (CF4) F C F F C F F

6. Lewis Structures: 2D Structures
NH3
CH2O
CO2
SO2
CH4

7. Drawing Lewis Structures
Sum the # of valence electrons from all atoms
Anions: add e– (CO32- : add 2 e– ) Cation: subtract e– (NH4+: minus 1 e– )
Predict the arrangement of the atoms
Usually the first element is in the center (often C, never H)
Make a single bond (2 e–) between each pair of atoms
Arrange remaining e– to satisfy octets (8 e– around each)
Place electrons in pairs (lone pairs)
Too few? Form multiple bonds between atoms:
double bond (4 e– ) and triple bond (6 e– )
Check your structure!
All electrons have been used
All atoms have 8e-
Exceptions: Remember that H only needs 2e– !

8. Lewis Structure Practice
Draw a Lewis Structure for the following compounds:
CH4
H2O
NF3
HBr
OF2
HCN
NO3-
CO32- N H C

9. N H C

10. Lewis Structure Trends
Here are some useful trends…
C group
Forms a combo of 4 bonds and no LP (Lone Pairs)
i.e. CO2
N group
Forms a combo of 3 bonds and 1 LP
i.e. NH3
O group
Forms a combo of 2 bonds and 2 LP
i.e. CH2O
F group (halogens)
Forms 1 bond and 3 LP
i.e. OF2
Note that these are NOT always true!

11. Carbonite
CO22-
Carbonate? CO32-

12. 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
Technically NOT two single bonds and one double bond
All 3 Oxygen atoms share the double bond
3 equal bonds (somewhere between a double and single)
Arrow formalism: curved arrows show electron movement

13. Predicting Molecular Shape: VSEPR
(Valence Shell Electron Pair Repulsion)
Electrons repel each other
The molecule adopts a 3-D shape to keep the electrons (lone pairs and bonded e-) as far apart as possible
Different arrangements of bonds/lone pairs result in different shapes
Shapes depend on # of bonds/lone pairs (“things”) and LP around the central atom

14. Selected Shapes and Geometries using VSEPR
“Things” 14

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

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

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

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

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

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

21. Hydrogen Chloride: HCl
Lewis Structure
Four “things” (bonds/LP)
Tetrahedral geometry
3 LP → Linear Shape
No Bond angle Cl Cl H

22. Cl N O H Br A special note… For any molecule having only two atoms…
e.g. N2, CO, O2, Cl2, HBr, etc.
Geometry = Linear
Shape = Linear
Bond Angle(s)? = None
It is much like geometry…
what is formed by connecting two points?
…a line. Cl N O H Br

23. You will need to commit these to memory!
“Things” 23

24. VSEPR Practice (w/o aid of yellow sheet)
CO2 G: S:
Angle:
ClO2-
NO2-
CH3COO- G: S:
Angle:
PBr3
AsO43-

25. Electronegativity and Bond Type
The electronegativity difference between two elements helps predict what kind of bond they will form.
Definition
e- are evenly shared
e- are unevenly shared
e- are exchanged (gained or lost)
Electronegativity difference
≤ 0.4 
0.5 – 1.8
> 1.8
Bond type
Covalent 
Polar covalent
Ionic

26. Practice with Bond Types
Sample Bonds
NaCl
Cl-Cl
C-O
C-H
Electronegativity Difference
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 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

27. Dipole Moments and Polarity
Occurs in polar covalent bonds
Uneven distribution of e-
Atoms become partially charged
Partially “+” charged
end
Arrow points toward partially “-” end δ+ δ-

28. Polarity Examples HCN CO2 CO32- CH2O SO2 Polar CH4 CH3F C3H8 CO NH3
Check molecule for dipole moments (polar bonds)
When determining overall polarity, an imbalanced structure will likely be polar (at least partially)
Even with polar bonds, a balanced structure is non-polar overall
Any structure with lone pairs on the central atom is automatically polar!
Try these with your neighbors…
HCN
CO2
CO32-
CH2O
SO2
Polar
CH4
CH3F
C3H8 CO
NH3
Non-polar
Non-polar
Polar
Non-polar
Non-polar
Polar
Polar
Polar
Polar