1. Chapter 8 – Covalent Bonding
The unspoken hero: “Covalent Bond” 

2. Review of Chapter 7
In Chapter 7, we learned about electrons being transferred (“given up” or “stolen away”)
This type of “tug of war” between a METAL and NONMETAL is called an IONIC BOND, which results in a SALT being formed

3. Chapter 8.1 – Molecular Compounds
Covalent Bonds - atoms held together by SHARING electrons between NONMETALS

4. Salt versus Molecules
Molecule - A group of atoms joined together by a covalent bond
Compound - a group of two or more elements bonded together (Ionic or Covalent).

5. Monatomic vs. Diatomic Molecules
Most molecules can be monatomic or diatomic
Diatomic Molecule - molecule consisting of two atoms
There are 7 diatomic molecules (SUPER 7) –
H2, O2, N2, Cl2, Br2, I2, F2

6. Properties of Molecular Compounds
Liquids or gases at room temperature
Lower Melting Points than Ionic Compounds (means weaker bonds than ionic)

7. Molecular Formulas Molecular Formula – formula of a molecular compound
Shows how many atoms of each element a molecule contains
Example
H2O contains 3 atoms (2 atoms of H, 1 atom of O)
C2H6 contains 8 atoms (2 atoms of C, 6 atoms of H)

8. Practice
How many atoms total and of each do the following molecular compounds contain? H2 CO
CO2
NH3
C2H6O 2 2 3 4 9

9. Practice: True or False
All molecular compounds are composed of atoms of two or more elements.
All compounds are molecules.
Molecular compounds are composed of two or more nonmetals.
Atoms in molecular compounds exchange electrons.
Molecular compounds have higher melting and boiling points than ionic compounds.
Share
Lower

11. Ionic versus Covalent IONIC COVALENT Bonded Name Salt Molecule
Bonding Type
Transfer e-
Share e-
Types of Elements
Metal & Nonmetal
Nonmetals
Physical State
Solid
Solid, Liquid, or Gas
Melting Point
High (above 300ºC)
Low (below 300 ºC)
Solubility
Dissolves in Water
Varies
Conductivity
Good
Poor

12. Chapter 8.2 – Covalent Bonding
Remember that ionic compounds transfer electrons in order to attain a noble gas electron configuration
Covalent compounds form by sharing electrons to attain a noble gas electron configuration
Regardless of the type of bond, the Octet Rule still must be obeyed (8 valence electrons)
All elements need 8 except Hydrogen, which needs 2.

13. Single Covalent Bond
A Single Covalent Bond consists of two atoms held together by sharing 1 pair of electrons (2 e-)

14. Electron Dot Structure

15. Shared versus Unshared Electrons
A Shared Pair is a pair of valence electrons that is shared between atoms
An Unshared Pair is a pair of valence electrons that is not shared between atoms

16. Practice Lewis Dot Structures
Chemical Formula
# of Valence Electrons
Single Line Bond Structure
# of Remaining Electrons
Lewis Dot Structure
Octet Check All Atoms=8
Hydrogen=2 F2
H2O
NH3
CH4 14
F-F 12
F, F = 8
O = 8
H, H = 2 8
H-O-H 4
N = 8
3 H = 2 8 2
C = 8
4 H = 2 8

17. Double Covalent Bonds
Sometimes atoms need to share 2 or 3 pairs of electrons
Hydrogen will NEVER form double or triple bonds.
Double Bond - bond that involves 2 shared pairs of electrons (4 e-)

18. Triple Covalent Bond
Triple Bond - bond that involves 3 shared pairs of electrons (6 e-)

19. Covalent Bonds

20. Practice Lewis Dot Structure
Chemical Formula
# of Valence Electrons
Single Line Bond Structure
# of Remaining Electrons
Lewis Dot Structure
Octet Check All Atoms=8
Hydrogen=2 O2
CO2 N2
HCN 12
O-O 10
2 O = 8 16
O-C-O 12
C = 8
2 O = 8
2 N = 8 10
N-N 8
C = 8
N = 8
H = 2 10
H-C-N 6

21. Drawing Lewis Structures of Molecules
Chapter 8: Basic Concepts of Chemical Bonding
Drawing Lewis Structures of Molecules
If the compound contains more than 2 atoms:
how are the atoms bonded and,
if there are nonbonding electron, where are they?

22. Molecules with a central atom : NH3, PCl3, CHCl3
Chapter 8: Basic Concepts of Chemical Bonding
Molecules with a central atom : NH3, PCl3, CHCl3
Central atom is generally the first in the molecular formula and the most electronegative one.

23. …unless the first element is Hydrogen :
Chapter 8: Basic Concepts of Chemical Bonding
…unless the first element is Hydrogen :
H2O
HCN
(same order as in formula)

24. How to find the # of bonds in a lewis structure
How to find the # of bonds in a lewis structure **Doesn’t work for molecule over an octet
Find the total # of valence electrons.
2. Use the formula to find the number of bonds.
# of val e- needed (all have 8 or 2 e-)
- # of val e- available
= ____/2 to find the # of bonds
(# val e- – # val e- available) =
# of bonds 2

25. a.) CO b.) C2F4 c.) C2H6 3 6 7 Find the total # of valence electrons.
2. Use the formula to find the number of bonds.
# of val e- needed (all have 8 or 2 e-)
- # of val e- available
= ____/2 to find the # of bonds
Ex: Find the number of bonds for each molecule or compound and write the lewis dot structure # of Bonds LDS
a.) CO
b.) C2F4
c.) C2H6
**Doesn’t work for molecule over an octet 3 6 7

26. (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding
Rules for Drawing Lewis Structures
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed 8
NH3
(2) Connect atoms by covalent bonds: count electrons left 2
(3) Complete "octets" of atoms around central atom
n/a
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
2 leftovers
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds
n/a

27. CO 10 8 C-O C-O C-O 2 left (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding CO
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed 10
(2) Connect atoms by covalent bonds: count electrons left 8
C-O
(3) Complete "octets" of atoms around central atom
C-O
(treat C as central)
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
C-O
2 left
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds

28. SF2 20 F-S-F 16 4 left (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding
SF2
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed 20
(2) Connect atoms by covalent bonds: count electrons left
F-S-F 16
(3) Complete "octets" of atoms around central atom
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
4 left
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds

29. Lewis Structure for Ions
If a molecule has a positive charge, subtract that many electrons from the total valence electrons available.
𝑁𝐻 has 8 electrons available.
If a molecule has a negative charge, add that many electrons to the total valence electrons available.
𝑁𝑂 3 − has 24 electrons available.
𝑆𝑂 4 −2 has 32 electrons available.

30. square brackets and the sign
For ions, the charge is generally indicated by
square brackets and the sign

31. NH4+ 8 0 left n/a n/a (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding
Ions
NH4+
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed 8
(2) Connect atoms by covalent bonds: count electrons left
0 left
(3) Complete "octets" of atoms around central atom
n/a
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
n/a
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds

32. Chapter 8: Basic Concepts of Chemical Bonding
Ions
ClO2-
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed 20
(2) Connect atoms by covalent bonds: count electrons left
O-Cl-O
16 left
(3) Complete "octets" of atoms around central atom
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
4 left
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds

33. If the molecule has an odd number of valence electrons
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule
On occasion, an atom in a molecule does not have an octet of valence electrons:
If the molecule has an odd number of valence electrons
an atom may have less than an octet [mainly Be, B]
an atom may have more than an octet [periods 3-7]

34. NO2 17 O-N-O 17 1 left 13 (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule: odd number of electrons
NO2
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed 17
(2) Connect atoms by covalent bonds: count electrons left
O-N-O 17
(3) Complete "octets" of atoms around central atom
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
1 left
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds 13

35. BF3 24 18 left 0 left (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule: less than an octet (B or Be)
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
BF3 24
(2) Connect atoms by covalent bonds: count electrons left
18 left
(3) Complete "octets" of atoms around central atom
0 left
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds

36. More Than Eight Electrons
The only way PCl5 can exist is if phosphorus has 10 electrons around it.
Periods 3-7 can expand its orbitals.

37. More Than Eight Electrons
Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

38. More Than Eight Electrons
This eliminates the charge on the phosphorus and the charge on one of the oxygens.
The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

39. BrF5 42 32 (1) Sum valence electrons from all atoms:
Chapter 8: Basic Concepts of Chemical Bonding
Exceptions to the Octet rule: more than an octet
(1) Sum valence electrons from all atoms:
these are the ones that need to be distributed
BrF5 42
(2) Connect atoms by covalent bonds: count electrons left 32
(3) Complete "octets" of atoms around central atom
2 left
(4) Place any leftover electrons on the central atom.
Check that central atom has octet
(5) If there are not enough electrons to give the central atom an octet, try multiple bonds

40. Bond Dissociation Energy
Bond Dissociation Energy - energy required to break a bond between two atoms
A large bond dissociation energy corresponds to a strong bond which makes it unreactive
Carbon has strong bonds, which makes carbon compounds stable and unreactive

41. Chapter 8.3 - Bonding Theories
Determining shape through bonds

42. VSEPR Theory VSEPR Theory predicts the 3D shape of molecules
According to VSEPR, the shape of the molecule adjusts so that the electrons are far apart

43. A Few VSEPR Shapes

44. Nine possible molecular shapes

45. VSEPR Theory
Unshared pairs of electrons are important in predicting the shapes of molecules
Each bond (single, double, or triple) and unshared pair is considered a steric number
Use the steric number to predict the molecular geometry
VSEPR can only be used with the central atom

46. Bent 4 2 Linear 6 Molecule Lewis Dot Structure Steric # Shape H2O CO2
CO2
XeF4
Bent 4 2
Linear
Square
Planar 6

47. Hybrid Orbitals
VSEPR is good at describing the molecular shapes, but not the types of bonds formed
Orbital hybridization provides information about both molecular bonding and molecular shape
In hybridization, several atomic orbitals mix to form hybrid orbitals

48. Bond Hybridization Hybridization Involving Single Bonds – sp3 orbital
Ethane (C2H6)
Hybridization Involving Double Bonds – sp2 orbital
Ethene (C2H4)
Hybridization Involving Triple Bonds – sp orbital
Ethyne (C2H2)

49. Gets another orbital added for each atom and lone pair around atom of interest.
Start with sp…
…then sp2…
(what is the orbital
hybridization of C?)
…then sp3

52. Chapter 8.4 – Polar Bonds and Molecules
There are two types of covalent bonds
Nonpolar Bonds (share equally)
Polar Bonds (share unequally)

53. Polar Covalent
Polar Bond - unequal sharing of electrons between two atoms (ex: HCl)
In a polar bond, one atom typically has a negative charge, and the other atom has a positive charge

54. Nonpolar Covalent Bond
Nonpolar Bond - equal sharing of electrons between two atoms (Cl2, N2, O2)

56. Classification of Bonds
You can determine the type of bond between two atoms by calculating the difference in electronegativity values between the elements
Type of Bond
Electronegativity Difference
Nonpolar Covalent
0  0.4
Polar Covalent
0.5  1.9
Ionic
2.0  4.0

57. Practice Your Turn To Practice
What type of bond is HCl? (H = 2.1, Cl = 3.1)
Difference = 3.1 – 2.1 = 1.0
Therefore it is polar covalent bond.
Your Turn To Practice
N(3.0) and H(2.1)
H(2.1) and H(2.1)
Ca(1.0) and Cl(3.0)
Mg(1.2) and O(3.5)
H(2.1) and F(4.0)

58. Polar molecules –
All molecules with lone pairs (unless it’s linear)
All molecules surrounded by different atoms.
Nonpolar molecule –
Linear molecules with the same atoms
All molecules surrounded by the same atoms.

59. Dipole When there is unequal sharing of electrons, a dipole exists
Dipole - a molecule with two poles with opposite charges
Represented by an arrow pointing towards the more negative end

60. Practice Drawing Dipoles
P- Br
P = 2.1
Br = 2.8
P –Br
 -
Practice
H(2.1) – Cl(3.0)
C(2.5) - F(4.0)
H(2.1)– F(4.0)

61. Attractions Between Molecules
There are also attractions between molecules
Intermolecular attractions are weaker than ionic, covalent, and metallic bonds
There are 2 main types of attractions between molecules: Van der Waals and Hydrogen

62. Van der Waals Forces
Van der Waals forces consists of the two weak attractions between molecules
2. dispersion forces – caused by the motion of electrons (weakest of all forces)
1. dipole interactions – polar molecules attracted to one another

63. Hydrogen Bond
Hydrogen Bonds - forces where a hydrogen atom is weakly attracted to an unshared electron pair of another atom
Strongest of all intermolecular forces