1. Chapter 5
Molecular Compounds

2. Chemical Bonds Three basic types of bonds: Ionic Covalent Metallic
Electrostatic attraction between ions
Covalent
Sharing of electrons
Metallic
Metal atoms bonded to several other atoms

3. Attraction Between Two or More Ions
Ionic Bonding
Attraction Between Two or More Ions

4. The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6
[Ne] Li
Li+ + e-
e- + F - F -
Li+ +
Li+
9.2

5. Sharing electrons within a Molecule!
Covalent Bonding
Sharing electrons within a Molecule!

6. Valence Electrons
Valence electrons are the outer shell s and p electrons of an atom.
The valence electrons are the electrons that participate in chemical bonding.
We show the valence electrons in Lewis dot structures of atoms and molecules.

7. 9.1

8. Molecular Compounds
Molecular compound: a compound in which all bonds are covalent
Naming binary molecular compounds
the less electronegative element is named first
prefixes “di-”, tri-”, etc. are used to show the number of atoms of each element; the prefix “mono-” is omitted when it refers to the first atom, and is rarely used with the second atom. Exception: carbon monoxide
NO is nitrogen monoxide (nitric oxide)
SF2 is sulfur difluoride
N2O is dinitrogen monoxide (laughing gas)

9. Forming a Covalent Bond
A covalent bond is formed by sharing one or more pairs of electrons
the pair of electrons is shared by both atoms and, at the same time, fills the valence shell of each atom
example: in forming H2, each hydrogen contributes one electron to the single bond

10. Why do two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.
Why do two atoms share electrons?
7e-
7e-
8e-
8e- F F + F
Lewis structure of F2
lone pairs F
single covalent bond
single covalent bond F
9.4

11. Lewis structure of water
single covalent bonds
2e-
8e-
2e- H + O + H O H or
Double bond – two atoms share two pairs of electrons
8e-
8e-
8e-
double bonds O C or O C
double bonds
Triple bond – two atoms share three pairs of electrons
8e-
triple bond N
8e- or N
triple bond
9.4

12. Lengths of Covalent Bonds
Bond Type
Bond Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple < Double < Single
9.4

13. Comparison of Ionic and Covalent Compounds
9.4

14. Drawing Lewis Structures
1. Set aside the single bonders (H, F, Cl, Br, I).
2. Put together the 2, 3, and 4 bonders with single bonds (May require thinking).
3. Count the holes left
A. If holes = single bonders, put them on.
B. If holes > single bonders, make double or triple bonds, or a ring until A is satisfied.

15. Lewis Structures Examples
draw a Lewis structure for hydrogen peroxide, H2O2
draw a Lewis structure for methanol, CH3OH
draw a Lewis structure for acetic acid, CH3COOH

16. Lewis Structures Draw a Lewis Structure and Make a model for C2H6O
BF3
HNO3
HCN
C2H3Cl

17. Lewis Structures
You will build models and draw Lewis Structures in Lab.
You will have to identify the polarity in a molecule as well.
Link to polarity idea

18. Polar Covalent Bonds
Polar covalent bond or polar bond is a covalent bond with uneven electron density
electron rich
region
electron poor
region H F
e- poor
e- rich F H d+ d-
9.5

19. Electronegativity
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electronegativity - relative, F is highest
9.5

20. Electronegativity
9.5

21. Increasing difference in electronegativity
Classification of bonds
Difference
Bond Type
Covalent
 2
Ionic
0.4 < and <2
Polar Covalent
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
9.5

22. Polar Covalent Bonds
The greater the difference in electronegativity, the more polar is the bond.

23. Solve The Following
Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the CN bond in HCN; and the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
Polar Covalent
C – 2.5
N – 3.0
3.0 – 2.5 = 0.5
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Covalent
9.5

24. Write the Lewis structure of nitrogen trifluoride (NF3).
9.6

25. Write the Lewis structure of the carbonate ion (CO32-).
9.6

26. Formal Charge and Lewis Structures
Lewis structures with large formal charges are less plausible than those with small formal charges.
Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.
Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O
9.7

27. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
Ions or molecules with an odd number of electrons.
Ions or molecules with less than an octet.
Ions or molecules with more than eight valence electrons (an expanded octet).

28. Exceptions to the Octet Rule
The Incomplete Octet
BeH2 H Be F B
BF3
9.9

29. Exceptions to the Octet Rule
Odd-Electron Molecules NO N O
The Expanded Octet (central atom with principal quantum number n > 2) S F
SF6
9.9