CHM 138
BASIC CHEMISTRY
Chapter 6
Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
participate in chemical bonding.
Group
e- configuration
# of valence e-
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
Lewis Dot Symbols for the Representative Elements
& Noble Gases
• Lewis dot symbol consists of the symbol of an element
and one dot for each valence electron in an atom of the
element.
The Ionic Bond
• Ionic bond: the electrostatic force that holds ions
together in an ionic compound.
• Atoms of the elements with low ionization energies tend
to form cation – alkali metals and alkaline earth metals
•Atoms of the elements with high electron affinities tend
to form anion – halogens and oxygen
Li + F
Li+ F -
(LiF)
1s22s1 1s22s22p5 1s2 1s22s22p6
Li+ + e-
Li
e- +
Li+ +
F
F -
F Li+ F -
A covalent bond is a chemical bond in which two or more
electrons are shared by two atoms.
Why should two atoms share electrons?
F
+
7e-
F
F F
7e-
8e- 8e-
Lewis structure of F2
single covalent bond
lone pairs
F
F
single covalent bond
lone pairs
F F
lone pairs
lone pairs
A Lewis structure:
- a representation of covalent bonding in which
shared electron pairs are shown either as lines or as
pairs of dots between two atoms, and lone pairs are
shown as pairs of dots an individual atoms.
Only valence electrons are shown.
The formation of the molecules illustrates the octet
rule.
- Octet rule: An atom other than hydrogen tends to
form bonds until it is surrounded by eight valence
electrons.
Lewis structure of water
H
+
O +
H
single covalent bonds
H O H
or
H
O
H
2e-8e-2eDouble bond – two atoms share two pairs of electrons
O C O
or
O
O
C
double bonds
8e- 8e- 8edouble bonds
Triple bond – two atoms share three pairs of electrons
N N
8e-8etriple bond
or
N
N
triple bond
Lengths of Covalent Bonds
Bond Lengths
Triple bond < Double Bond < Single Bond
COMPARISON OF GENERAL PROPERTIES
OF IONIC COMPOUND AND COVALENT
COMPOUND
IONIC COMPOUND
COVALENT COMPOUND
Solid at room temperature,
high melting points
Gases, liquids, low melting
solids
Soluble in water
Insoluble in water
Aqueous solution conduct
electricity
Aqueous solution do not
conduct electricity
Strong electrolytes
Nonelectrolytes
Writing Lewis Structures
1. Count total number of valence e-. Add 1 for each
negative charge. Subtract 1 for each positive
charge.
2. Draw skeletal structure of compound showing what
atoms are bonded to each other. Put least
electronegative element in the center.
3. Complete an octet for all atoms except hydrogen
4. If structure contains too many electrons, form
double and triple bonds on central atom as needed.
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
F
N
F
F
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
O
C
O
O
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24
Write the Lewis structure for:
i) NO -2
ii) CS2
iii) SO3
FORMAL CHARGE
An atom’s formal charge is the difference between the
number of valence electrons in an isolated atom and the
number of electrons assigned to that atom in a Lewis structure.
formal charge
on an atom in
a Lewis
structure
=
total number
total number
of valence
electron assigned
electrons in to atom
the free atom
The sum of the formal charges of the atoms in a molecule
or ion must equal the charge on the molecule or ion.
Examples:
H
C
H
H
O
C
H
O
formal charge
on C
= -1
formal charge
on O
=+1
formal charge
on C
=0
formal charge
on O
=0
Formal Charge and Lewis Structures
1. For neutral molecules, a Lewis structure in which there are no
formal charges is preferable to one in which formal charges
are present.
2. Lewis structures with large formal charges are less plausible
than those with small formal charges.
3. 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
-1
+1
C
O
H
H
H
0
C
0
O
Resonance Structure
A resonance structure is one of two or more Lewis structures
for a single molecule that cannot be represented accurately by
only one Lewis structure.
O
O
+
-
-
O
O
+
O
O
What are the resonance structures of the carbonate (CO32-) ion?
-
O
C
O
O
-
O
C
O
O
-
-
-
O
C
O
O
-
Exceptions to the Octet Rule
The Incomplete Octet
BeH2
BF3
B – 3e3F – 3x7e24e-
Be – 2e2H – 2x1e4e-
F
B
F
H
F
Be
H
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24
Exceptions to the Octet Rule
Odd-Electron Molecules
NO
N – 5eO – 6e11e-
N
O
The Expanded Octet (central atom with principal quantum number n > 2)
SF6
S – 6e6F – 42e48e-
F
F
F
S
F
F
F
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
Dative Covalent Bond /
Coordinate Covalent Bond
A covalent bond in which one of the atoms
donates both electrons.
Examples:
- NH4+, NH3AlCl3, NH3BF3
Hybridization
Hybridization – mixing of two or more atomic orbitals to
form a new set of hybrid orbitals.
1. Mix at least 2 nonequivalent atomic orbitals (e.g. s and p).
Hybrid orbitals have very different shape from original
atomic orbitals.
2. Number of hybrid orbitals is equal to number of pure atomic
orbitals used in the hybridization process.
3. Covalent bonds are formed by:
a. Overlap of hybrid orbitals with atomic orbitals
b. Overlap of hybrid orbitals with other hybrid orbitals
Formation of sp3 Hybrid Orbitals
Formation of Covalent Bonds in CH4
sp3-Hybridized N Atom in NH3
Formation of sp Hybrid Orbitals
Formation of sp2 Hybrid Orbitals
How to predict the hybridization of the central atom?
1. Draw the Lewis structure of the molecule.
2. Count the number of lone pairs AND the number of
atoms bonded to the central atom
# of Lone Pairs
+
# of Bonded Atoms
Hybridization
Examples
2
sp
BeCl2
3
sp2
BF3
4
sp3
5
sp3d
PCl5
6
sp3d2
SF6
CH4, NH3, H2O
Bonding in Ethylene, C2H4
Sigma bond (s) – covalent bonds formed by orbitals overlapping end –
to-end , with the electron density between the nuclei of the bonding
atoms
Pi bond (p) – a covalent bond formed by sideways overlapping
orbitals with electron density concentrated above and below plane of
nuclei of the bonding atoms
Another View of p Bonding in Ethylene, C2H4
Bonding in Acetylene, C2H2
Describe the bonding in CH2O
H
H
C
O
C – 3 bonded atoms, 0 lone pairs
C – sp2
Sigma (s) and Pi Bonds (p)
Single bond
1 sigma bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid (vinegar)
molecule CH3COOH?
H
C
H
O
H
C
O
H
s bonds = 6 + 1 = 7
p bonds = 1
Intermolecular Forces
Intermolecular forces: attractive forces between
molecules.
Van der Waals forces:
- the attractive or repulsive force between molecules
due to covalent bonds or to the electrostatic
interaction of ions with one another or with neutral
molecules.
The term includes:
- permanent dipole–permanent dipole forces
- instantaneous induced dipole-induced dipole
(London dispersion force). Examples: interaction
between H2, Cl2, F2, CH4
Hydrogen Bond
The hydrogen bond is a special dipole-dipole interaction
between they hydrogen atom in a polar N-H, O-H, or F-H
bond and an electronegative O, N, or F atom.
A
H…B
or
A
A & B are N, O, or F
H…A
Why is the hydrogen bond considered a “special”
dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point