Transcript Chapter 8 Covalent bonds

Ch. 8- Covalent Bonding
II. Molecular Compounds
I
II
III
IV
A. Energy of Bond Formation
Potential Energy
based on position of an object
low PE =
high stability
A. Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
no interaction
increased
attraction
A. Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
increased
repulsion
balanced attraction
& repulsion
A. Energy of Bond Formation
Bond Energy
Energy required to break a bond
Bond
Energy
Bond
Length
A. Energy of Bond Formation
Bond Energy
Short bond = high bond energy
B. Lewis Structures
Electron Dot Diagrams
show valence e- as dots
distribute dots like arrows
in an orbital diagram
4 sides = 1 s-orbital, 3 p-orbitals
EX: oxygen
2s
2p
O
X
B. Lewis Structures
Octet Rule
Most atoms form bonds in order to
obtain 8 valence eFull energy level stability ~ Noble
Gases
Ne
B. Lewis Structures
Nonpolar Covalent - no charges
Polar Covalent - partial charges
+
+
C. Molecular Nomenclature
Prefix System (binary compounds)
1. Less e-neg atom
comes first.
2. Add prefixes to indicate # of atoms.
Omit mono- prefix on first element.
3. Change the ending of the
second element to -ide.
C. Molecular Nomenclature
PREFIX
monoditritetrapentahexaheptaoctanonadeca-
NUMBER
1
2
3
4
5
6
7
8
9
10
C. Molecular Nomenclature
CCl4
carbon tetrachloride
N2O
dinitrogen monoxide
SF6
sulfur hexafluoride
C. Molecular Nomenclature
arsenic trichloride
AsCl3
dinitrogen pentoxide
N2O5
tetraphosphorus decoxide
P4O10
C. Molecular Nomenclature
The Seven Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2
H
N O F
Cl
Br
I
I. Structural Formulas
 Show a structural formula for PH3.
 Structural formulas show shared pairs of
electrons as lines (bonds) and lone pairs
of electrons as dots.
H–P–H
H
Molecular Structure
I. Lewis Diagrams
I
II
III
A. Octet Rule

Remember…
 Most atoms form bonds in order to
lower energy and gain stability
 Want to have 8 valence electrons.
A. Octet Rule

F
F
 Hydrogen  2 valence e
F
B
F
 Groups F
1,2,13S
get 2,4,6
valence e
F
H
N
O
O
H
 Expanded octet

more
than
8
F
Very
unstable!!
valence
e
(e.g.
S,
P,
Xe)
F
F
Exceptions:
-
-
-
 Radicals  odd # of valence e-
Lewis Structures

Formulas in which atomic symbols
represent nuclei and inner shell
electrons. Dot-pairs or lines between
two atomic symbols represent
electron pairs in covalent bonds, and
dots adjacent to only one atomic
symbol represent unshared electrons.
C. Drawing Lewis Diagrams

Find total # of valence e-.

Arrange atoms - singular atom is
usually in the middle.

Form bonds between atoms (2 e-).

Distribute remaining e- to give each
atom an octet (recall exceptions).

If there aren’t enough e- to go around,
form double or triple bonds.
B. Drawing Lewis Diagrams
 CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
F
F C F
F
B. Drawing Lewis Diagrams
 BeCl2
1 Be × 2e- = 2e2 Cl × 7e- = 14e16e- 4e12e-
Cl Be Cl
B. Drawing Lewis Diagrams
 CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e-
O C O
C. Polyatomic Ions

To find total # of valence e-:
 Add 1e- for each negative charge.
 Subtract 1e- for each positive
charge.

Place brackets around the ion and
label the charge.
C. Polyatomic Ions
 ClO4-
1 Cl × 7e- = 7e4 O × 6e- = 24e31e+ 1e32e- 8e24e-
O
O Cl O
O
C. Polyatomic Ions
 NH4+
1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 8e0e-
H
H N H
H
D. Resonance Structures

Molecules that can’t be correctly
represented by a single Lewis
diagram.
Referred to as an average of all the
possibilities.
 Show possible structures separated
by a double-headed arrow.

D. Resonance Structures
 SO3
O
O S O
O
O S O
O
O S O