covalent bonding_I.ppt

Chemical Bonding I:
Covalent Bonding
Chapter 9
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2
Trends in Metallic Character

Effective Nuclear Charge (Zeff)

Effective nuclear charge (Zeff) is the “positive charge” felt by an
electron.
Na
Mg
Al
Si
11
12
13
14
10
10
10
10
1
2
3
4
186
160
143
132
Zeff
Core
Z Radius (pm)
Zeff = Z - s 0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons

Zeff = Zactual – No. of Core Electrons or Electron shielding

6
increasing Zeff
decreasing
Z
eff

Electronegativity is the ability of an atom to attract
toward itself the electrons in a chemical bond.
Electron Affinity - measurable, Cl is highest
Electronegativity - relative, F is highest
X (g) + e- X-
(g)
9.5

9.5
The Electronegativities of Common Elements

10
Ionization energy is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground
state.
I1 + X (g) X+
(g) + e-
I2 + X+
(g) X2+
(g) + e-
I3 + X2+
(g) X3+
(g) + e-
I1 first ionization energy
I2 second ionization energy
I3 third ionization energy
I1 < I2 < I3
Ionization Energy

1st Ionization Energy
0
100
200
300
400
500
600 1
4
7
10
13
16
19
22
25
28
31
34
Atomic number
kcal/mol
1s
2s
3s
You can see each subshell
and spin direction if you
look closely
2p
2p
3p
3p
4s
3d
3d
4p
4p

12
General Trends in First Ionization Energies
Increasing First Ionization Energy
Increasing
First
Ionization
Energy

Electron affinity is the negative of the energy change that occurs
when an electron is accepted by an atom in the gaseous state to form
an anion.
X (g) + e- X-
(g)
F (g) + e- F-
(g)
O (g) + e- O-
(g)
DH = −328 kJ/mol EA = +328 kJ/mol
DH = −141 kJ/mol EA = +141 kJ/mol
Electron Affinity
O-
(g) + e- O2-
(g) DH = +780 kJ/mol EA = −780 kJ/mol
 EA of O− is highly –ve; i.e. O2− formation is unfavorable,
 electron-electron repulsion (g) > stability from noble gas configuration;
 O2− ion stabilized by neighboring cation, e.g. Li2O, Mg2O
?

Periodic Properties of the Elements

Nonpolar Covalent
share e- equally
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Increasing difference in electronegativity
Classification of bonds by difference in electronegativity
Approximate Difference Bond Type
0 to 0.4 Nonpolar Covalent
2 or greater Ionic
0.5 to 1.9 Polar Covalent
9.5

Three Major Types of Bonding
• Ionic Bonding
– forms ionic compounds
– transfer of valence e-
• Metallic Bonding
• Covalent Bonding
– forms molecules
– sharing of valence e-
– This is our focus this chapter
19

Ionic Bonding
• Always formed between metal cations and
non-metals anions
• The oppositely charged ions stick like
magnets
[METALS ]+ [NON-METALS ]
-
Lost e-
Gained e-
20

Metallic Bonding
• Always formed between 2 metals (pure
metals)
– Solid gold, silver, lead, etc…
21

Covalent Bonding
• Pairs of e- are
shared
between 2 non-
metal atoms to
acquire the electron
configuration of a
noble gas.
molecules
22

9.1
Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
particpate in chemical bonding.
1A 1
ns1
2A 2
ns2
3A 3
ns2np1
4A 4
ns2np2
5A 5
ns2np3
6A 6
ns2np4
7A 7
ns2np5
Group # of valence e-
e- configuration

9.1
Lewis Dot Symbols for the Representative Elements &
Noble Gases

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

9.3
Lattice energy (E) increases
as Q increases and/or
as r decreases.
cmpd lattice energy
MgF2
MgO
LiF
LiCl
2957
3938
1036
853
Q= +2,-1
Q= +2,-2
r F- < r Cl-
Electrostatic (Lattice) Energy
E = k
Q+Q-
r
Q+ is the charge on the cation
Q- is the charge on the anion
r is the distance between the ion
centers
Lattice energy (E) is the energy required to completely separate
one mole of a solid ionic compound into gaseous ions.

9.3
Born-Haber Cycle for Determining Lattice Energy
DHoverall = DH1 + DH2 + DH3 + DH4 + DH5
o o
o
o
o
o

Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2O; and
the NN bond in H2NNH2.
Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic
H – 2.1 O – 3.5 3.5 – 2.1 = 1.4 Polar Covalent
N – 3.0 N – 3.0 3.0 – 3.0 = 0 Nonpolar
Covalent
9.5

H2, the Hydrogen Molecule
MasterCard
Sigma National Bank

H2, the Hydrogen Molecule
P+
P+
e-
e-
Each hydrogen has a noble gas
configuration through sharing
Each hydrogen has a noble gas
configuration through sharing

Diatomic Halogen Molecules
 The halogens all form diatomic molecules
containing a sigma bond
F2 Cl2 Br2 I2
 By sharing one e- apiece they each achieve
a noble gas configuration
F - F

Diatomic Halogen Molecules
 The halogens all form diatomic molecules
containing a sigma bond
F2 Cl2 Br2 I2
 By sharing one e- apiece they each achieve
a noble gas configuration
F - F
s

Lewis Structures
 Sharing of electron pairs can be
diagrammed in detail with Lewis Electron
Dot Structures
 Each atom begins with its valence electrons
 A dash represents two shared electrons in a
bond
 Unbonded electrons are kept in lone pairs

Three types of covalent bonds
 A Single Bond – only a pair of electrons are
shared between two atoms
 A Double Bond – two pairs of electrons are
 A Triple Bond – three pairs of electrons are
 Again un-bonded electrons are kept in lone
pairs

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 F
+
7e- 7e-
F F
8e- 8e-
F F
F F
Lewis structure of F2
lone pairs
lone pairs
lone pairs
lone pairs
single covalent bond
single covalent bond
9.4

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

The O2 Molecule
O O
An octet for the left
oxygen An octet for the right
oxygen

The O2 Molecule
O O
The traditional Lewis
structure for O2

The Double Bond
 There are 2 pairs of electrons shared
 The first pair is in a s-bond
 Second pair is a new type of covalent bond -
a pi bond (p-bond)

The Double Bond
O O
s-orbital
p--orbital
p+-orbital

http://en.wikipedia.org/wiki/Image:Pi-bond.jpg
Better depictions of a double bond
Forming the two lobes of a p-orbital

The Double Bond
O O
s-orbital
p+-orbital
p--orbital
If I slice the double bond in half...

The Double Bond
s-orbital
p+-orbital
p--orbital
…it looks like this

Your text’s (poor) depiction of a double bond

The O2 Molecule
O O
An official Lewis structure
one s-bond
one p-bond
Four lone pairs
s
p

Double Bonding Elements
7 N
6 C 8 O
15P 16S
Only five elements
participate in double
bonds

The N2 Molecule
N N
The left nitrogen has
achieved an octet The right nitrogen has
achieved an octet

The N2 Molecule
N N
The traditional
Lewis structure for
N2

The Triple Bond
 There are 3 pairs of electrons shared
 The first pair is in a s-bond
 The second pair is in a pi-bond (p-bond)
 The third pair is in a second pi-bond (p-
bond)

The Triple Bond
If I slice the triple bond like this...

The Triple Bond
s-orbital
p1
+-orbital
p1
--orbital
…it looks like this
p2
--orbital
p2
+-orbital

The N2 Molecule
So a triple bond is
a s-bond
and 2 p-bonds
In N2 there are two
lone pairs, too
N N
s
p2
p1

Triple Bonding Elements
7 N
6 C 8 O
Only three elements
participate in triple
bonds

Lengths of Covalent Bonds
Bond Lengths
Triple bond < Double Bond < Single Bond 9.4

H F F
H
Polar covalent bond or polar bond is a covalent
bond with greater electron density around one of the
two atoms
electron rich
region
electron poor
region e- rich
e- poor
d+ d-
9.5

1. Find total # of valence e-’s in molecule or ion
2. Connect the skeleton with single bonds
3. Place lone pairs to give non-H terminal atoms an octet
Check to see how many e-’s are left and if all octets
(duets) are satisfied
If e-’s still left, go to 4
If e-’s are short, go to 5
If e-’s just right, you’re done
4. Place remaining e-’s on central atom (but check)
5. If short e-’s, look for multiple bonders or e- deficients
General Rules for Lewis Structures

Simple Lewis Structures
Element Bonds Lone pairs
H 1
C 4
N 3-4 1- 0
P 3-4 1- 0
O 2 2
S 2 2
Halogens 1 3

Example 1: Water
 First, you need the molecule’s skeletal
structure
 I’ll provide this or give you enough
information to deduce it
 For example, in water, oxygen is the central
atom

Example 1: Water
H H
O
Central atom

Example 1: Water
H H
O
Terminal atoms

Example 1: Water
H H
O
Now start checking each atom
for its typical bonding pattern

Example 1: Water
H H
O
Each hydrogen forms one
covalent bond

Example 1: Water
H H
O
Oxygen forms two covalent
bonds and has 2 lone pairs

Example 1: Water
H H
O
In this molecule there are

Example 1: Water
H H
O
s s
2 s-bonds

Example 1: Water
H H
O
s s
2 s-bonds
2 lone pairs

Example 2: Carbon Dioxide
 In carbon dioxide, CO2, carbon is the
central atom

O O
C

O O
C
Each oxygen forms 2
bonds and has 2 lone
pairs

O O
C
Carbon forms 4
bonds

O O
C

O O
C
s s
2 s-bonds

O O
C
s s
p p
2 s-bonds
2 p-bonds

O O
C
s s
p p
2 s-bonds
2 p-bonds
4 lone pairs

Example 3: Ethyl Alcohol
C O
C
H H
H H
H H

C O
C
1 bond per hydrogen
H H
H H
H H

C O
C
4 bonds per carbon
H H
H H
H H

C O
C
2 bonds and 2 lone pairs
for oxygen
H H
H H
H H

C O
C
The condensed structural formula for ethyl alcohol is
CH3CH2OH
H H
H H
H H

C O
C
H H
H H
H H

C O
C
s
H H
s s
H H
H H
s
s s
s s
8 s-bonds

C O
C
s
H H
s s
H H
H H
s
s s
s s
8 s-bonds
2 lone pairs

Example 4: Hydrogen Cyanide
 Hydrogen cyanide, HCN, is a toxic gas
because it sits on the oxygen-binding site of
hemoglobin
 Carbon is the central atom

H N
C

H N
C
1 bond for the hydrogen

H N
C
4 bonds for the carbon

H N
C
3 bonds and 1 lone pair
for nitrogen

H N
C

H N
C
s
s
2 s-bonds

H N
C
s
s
p
p
2 s-bonds
2 p-bonds

H N
C
s
s
p
p
2 s-bonds
2 p-bonds
1 lone pair

Naming Binary Molecular
Compounds
 Different schemes for ionic compounds vs.
covalent compounds
 You can’t name ‘em if you can’t classify
‘em
 Ionics start with a metal or ammonium
 Covalents start with a nonmetal or metalloid

Covalent Compound Names
 Least electronegative element generally first
(Table 3.5, page 74)
 Most electronegative element generally
second -- with “-ide” ending
 Need to use numerical prefixes because no
charges to help determine ratios

Numerical Prefixes
1 mono-
2 di-
3 tri-
4 tetra-
5 penta-
6 hexa-
7 hepta-
8 octa-
9 nona-
10 deca-

(prefix) element (prefix) element
with “-ide” ending
(space)
What is the name of BF3?

(space)
boron

(space)
boron
Mono
prefix
rarely
used

(space)
boron tri-

(space)
boron trifluoride

(space)
What is the name of N2O5?

(space)
(Note: the “a” of the prefix
is dropped with oxide)
dinitrogen pentoxide

(space)
What is the name of NI3?

Example 1: Laughing Gas
 Let’s try a couple of examples
 For starters, let’s consider laughing gas
It’s an anesthetic used by dentists
It’s nonsystematic name is nitrous oxide
It’s formula is N2O

2 x N = 2 x 5e- = 10 e-
1 x O = 1 x 6e- = 6 e-
species charge = 0 e-
______________________________
Total 16 e-
1. Find total # of valence e-’s in molecule or ion

2. Connect the skeleton with single bonds
N
N O
16 e-

3. Place lone pairs to give non-H terminal atoms an octet
N
N O
16 e-

3a. Check to see how many e-’s are left and all octets
N
N O
8 e-’s, octet okay
16 e-

N
N O
8 + 8 = 16 e-’s, octet okay
16 e-

N
N O
16 e-’s used, octet NOT okay
16 e-

N
N O
Uh-oh! We’re short e-’s!
16 e-

If e-’s are short, go to 5
N
N O
16 e-

N
N O
16 e-

N
N O
We’re short 4 e-’s, so need 2 multiple bonds
16 e-

N
N O
Move a lone pair into a shared position
16 e-

N
N O
Move another lone pair into a shared position
16 e-

N
N O
Now last octet is full; 16 e-’s preserved!
16 e-

N
N O
We’re done!
16 e-

N
N O
NOTE: This would have been okay, too...
16 e-

N
N O
…but it’s not the way nature prefers.
16 e-

N
N O
It wouldn’t affect geometry anyway
16 e-

Laughing Gas
N
N O
Where did these e-’s come from?

Laughing Gas
N O
Left N
N
Middle N Right O

Laughing Gas
N O
Left N
N
Middle N Right O
It’s a coordinate covalent bond!

Laughing Gas
N O
Left N
N
Middle N Right O
Both e-’s came from the N

Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
F N F
F
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
9.6

Write the Lewis structure of the carbonate ion (CO3
2-).
Step 1 – C is less electronegative than O, put C in center
O C O
O
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e-
4 + (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
9.6
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24

9.7
Two possible skeletal structures of formaldehyde (CH2O)
H C O H
H
C O
H
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
=
1
2
total number
of bonding
electrons
( )
total number
of valence
electrons in
the free atom
-
total number
of nonbonding
electrons
-
The sum of the formal charges of the atoms in a molecule
or ion must equal the charge on the molecule or ion.

H C O H
C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
1 double bond = 4
Total = 12
formal charge
on C
= 4 -2 -½ x 6 = -1
formal charge
on O
= 6 -2 -½ x 6 = +1
formal charge
on an atom in
a Lewis
structure
=
1
2
total number
of bonding
electrons
( )
total number
of valence
electrons in
the free atom
-
total number
of nonbonding
electrons
-
-1 +1
9.7

C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
1 double bond = 4
Total = 12
H
C O
H
formal charge
on C
= 4 -0 -½ x 8 = 0
formal charge
on O
= 6 -4 -½ x 4 = 0
formal charge
on an atom in
a Lewis
structure
=
1
2
total number
of bonding
electrons
( )
total number
of valence
electrons in
the free atom
-
total number
of nonbonding
electrons
-
0 0
9.7

Formal Charge and Lewis Structures
9.7
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 C O H
-1 +1 H
C O
H
0 0

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
+
-
O C O
O
- -
O C O
O
-
-
O
C
O
O
-
- 9.8
What are the resonance structures of the
carbonate (CO3
2-) ion?

Exceptions to the Octet Rule
The Incomplete Octet
H H
Be
Be – 2e-
2H – 2x1e-
4e-
BeH2
BF3
B – 3e-
3F – 3x7e-
24e-
F B F
F
Total = 24
9.9

Exceptions to the Octet Rule
Odd-Electron Molecules
N – 5e-
O – 6e-
11e-
NO N O
The Expanded Octet (central atom with principal quantum number n > 2)
SF6
S – 6e-
6F – 42e-
48e-
S
F
F
F
F
F
F
Total = 48
9.9

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