Chemistry 1A
General
Chemistry
Ch. 10: The
Shape of
Molecules
Instructor:
Dr. Orlando E. Raola
Santa Rosa Junior College
Chemical bonding
Chemical bonding
Problems and questions —
How is a molecule or polyatomic ion held
together?
Why are atoms distributed at strange
angles?
Why are molecules not flat?
Can we predict the structure?
How is structure related to chemical and
physical properties?
Models
Models are attempts to explain
how nature operates on the
microscopic level based on
experiences in the macroscopic
world.
Fundamental Properties of Models
•
A model does not equal reality.
•
Models are oversimplifications, and are therefore
often wrong.
•
Models become more complicated as they age.
•
We must understand the underlying assumptions
in a model so that we don’t misuse it.
Localized Electron Model
A molecule is composed of
atoms that are bound together
by sharing pairs of electrons
using the atomic orbitals of the
bound atoms.
Structure and bonding
NN triple bond. Molecule is unreactive
Phosphorus is a
tetrahedron of P
atoms. Very
reactive!
Red
phosphorus, a
polymer. Used
in matches.
Forms of chemical bonds
• There are 2 extreme forms of
connecting or bonding atoms:
• Ionic—complete transfer of 1
or more electrons from one
atom to another
• Covalent—some valence
electrons shared between atoms
• Most bonds are
somewhere in between.
Forms of chemical bonds
Ionic Compounds
Metal of low IE
Nonmetal
of high
EA
2 Na(s) + Cl2(g) 
2 Na+ + 2 Cl-
Covalent Bonding
The bond arises from the mutual attraction
of 2 nuclei for the same electrons.
Electron sharing results.
Bond is a balance of attractive and repulsive
forces.
Bond formation
A bond can result from a “head-tohead” overlap of atomic
orbitals on neighboring atoms.
••
H
+
••
•
•
Cl
••
H
•
•
Cl
••
Overlap of H (1s) and Cl (2p)
Note that each atom has a single,
unpaired electron.
Chemical Bonding:
Objectives
Objectives are to understand:
1. valence e- distribution in
molecules and ions.
2. molecular structures
3. bond properties and their
effect on molecular properties.
Electron
Distribution in
Molecules
• Electron distribution
is depicted with
Lewis
electron dot
structures
• Valence electrons
are distributed as
shared or BOND
PAIRS and
unshared or LONE
G. N. Lewis
1875 - 1946
PAIRS.
Lewis Electron-Dot Symbols
For main group elements Place one dot per valence electron on each
of the four sides of the element symbol.
Pair the dots (electrons) until all of the
valence electrons are used.
Example: Nitrogen, N, has 5 valence electrons.
:
. N. .
.
. N:
.
.
. N.
:
.
:N .
.
Lewis electron-dot symbols for
elements in Periods 2 and 3
1
2
13
14
15
16
17
18
Figure 9.6
The Born-Haber cycle for lithium fluoride.
Energy balance in the formation of
an ionic compound
Formation of an ionic solid
1.
2.
3.
4.
5.
Sublimation of the solid metal
M(s)  M(g) [endothermic]
Ionization of the metal atoms
M(g)  M+(g) + e [endothermic]
Dissociation of the nonmetal
1/2X (g)  X(g)
[endothermic]
2
Formation of X ions in the gas phase:
X(g) + e  X(g) [exothermic]
Formation of the solid MX
M+(g) + X(g)  MX(s) [quite exothermic]
Lattice energy calculations
Q1, Q2 = charges on the ions
r = shortest distance between centers of
the cations and anions
Born-Haber cycle for MgO
Process
Mg(s) ® Mg(g)
Mg(g) ® Mg 2+ (g) + 2e 1
2
O2 (g) ® O(g)
O(g) + 2e- ® O2- (g)
Mg 2+ (g) + O2- (g) ® MgO(s)
Mg(s) + 21 O2 (g) ® MgO(s)
H0 (kJ·mol-1)
148
736 (1st)
1450 (2nd)
249
-141
878
?
-601
Born-Haber cycle for MgO
Process
Mg(s) ® Mg(g)
Mg(g) ® Mg 2+ (g) + 2e 1
2
O2 (g) ® O(g)
O(g) + 2e- ® O2- (g)
Mg 2+ (g) + O2- (g) ® MgO(s)
Mg(s) + 21 O2 (g) ® MgO(s)
H0 (kJ·mol-1)
148
736 (1st)
1450 (2nd)
249
-141
878
-3921
-601
Compare lattice energy
of LiF and MgO
Similar ionic radii:
Mg2+: 72 pm
Li+: 76 pm
O2-: 140 pm
F-: 133 pm
Hlattice MgO = 3929 kJ·mol-1
Hlattice LiF = 1050 kJ·mol-1
Trends in lattice energy.
Bond and Lone Pairs
• Valence electrons are distributed
as shared or BOND PAIRS and
unshared or LONE PAIRS.
••
H
•
•
Cl
••
shared or
bond pair
lone pair (LP)
This is called a LEWIS
ELECTRON DOT structure.
Valence Electrons
Electrons are divided between core and
valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5
Valence electrons
Number of valence electrons:
Main group elements: (1,2,13-18):
- total of s and p electrons in the outer shell
Transition elements: (3-12):
- ns and (n-1) d
Octet rule
Each atom has a tendency to
be surrounded by 8 electrons
in molecules and polyatomic
ions (H is surrounded by 2
electrons).
Electronegativity
The ability of an atom in a molecule
to attract shared electrons to itself.
(H  X)expected = ½(H-H + X-X)
 = (H  X)actual  (H  X)expected
The Pauling electronegativity (EN) scale
Electronegativity and atomic size
Lewis Dot Diagrams
•
Shows how valence electrons are
arranged among atoms in a molecule.
•
Reflects central idea that stability of a
compound relates to noble gas electron
configuration.
Rules for forming
Lewis dot diagrams
1. Determine the total number of valence electrons in the
molecule or ion
2. Determine the arrangement of atoms within a molecule
(the least electronegative is central, halogens are central
only with oxygen, oxygen is central only in water,
hydrogen is always terminal.
3. Use a pair of electrons to form a bond between each pair
of atoms
4. Arrange the remaining electrons as lone pairs to satisfy
the duet rule for hydrogen or the octet rule for other
terminal atoms
5. If the central atom has fewer than eight electrons, move
lone pairs from the terminal atoms to form multiple bonds.
Comments About the Octet Rule
•
2nd row elements C, N, O, F observe the octet rule.
•
2nd row elements B and Be often have fewer than 8
electrons around themselves - they are very reactive.
•
3rd row and heavier elements CAN exceed the octet
rule using empty valence d orbitals.
•
When writing Lewis structures, satisfy octets first,
then place electrons around elements having
available d orbitals.
Writing Lewis structures
nitrogren trifluoride NF3
1.Number of valence electrons: 5+(3x7)=26
2.Use bond pairs to form bonds between
atoms
3.Arrange remaining electrons to satisfy
octet rule (duet for H)
More Lewis structures
hydrazine N2H4
chloroethane C2H5Cl
phosgene COCl2
chlorate ion ClO3perchlorate ion ClO4-
Building a dot structure
Ammonia, NH3
1. Decide on the central atom; never H.
Central atom is atom of lowest affinity for
electrons.
Therefore, N is central
2. Count valence electrons
H = 1 and N = 5
Total = (3  1) + 5
= 8 electrons / 4 pairs
Building a dot structure
3. Form a single bond between the
central atom and each
surrounding atom
4.Remaining electrons form LONE
PAIRS to complete octet as needed.
3 BOND PAIRS and 1 LONE
PAIR.
H N H
H
••
H N H
H
Note that N has a share in 4 pairs (8 electrons),
while H shares 1 pair.
Sulfite ion, SO32Step 1. Central atom = S
Step 2. Count valence electrons
S= 6
3 x O = 3 x 6 = 18
Negative charge = 2
TOTAL = 26 e- or 13
pairs
Step 3. Form bonds
Sulfite ion, SO32Remaining pairs become lone
pairs, first on outside atoms and
then on central atom.
••
•
•
O
•
•
••
•
•
O
••
••
S
••
O
••
•
•
Each atom is surrounded by an
octet of electrons.
Carbon Dioxide, CO2
1. Central atom = _______
2. Valence electrons = __ or __
pairs
3. Form bonds. This leaves 6 pairs.
4. Place lone pairs on outer atoms.
Carbon Dioxide, CO2
4. Place lone pairs on outer atoms.
5. So that C has an octet, we shall form
DOUBLE BONDS between C and O.
The second bonding pair forms a pi
(π) bond.
Double and even
triple bonds are
commonly
observed for C, N,
P, O, and S
H2CO
SO3
C2F4
Sulfur Dioxide, SO2
1. Central atom = S
2. Valence electrons = 18 or 9 pairs
3. Form double bond so that S has an octet
— but note that there are two ways of doing
this.
bring in
left pair
••
•
•
O
••
••
S
OR bring in
right pair
••
•
•
O
••
Sulfur Dioxide, SO2
This leads to the following structures.
These equivalent structures are called
RESONANCE STRUCTURES. The true
electronic structure is a HYBRID of the
two.
Urea, (NH2)2CO
Urea, (NH2)2CO
1. Number of valence electrons = 24
e2. Draw sigma bonds.
Urea, (NH2)2CO
3. Place remaining electron pairs in
the molecule.
Urea, (NH2)2CO
4. Complete C atom octet with
double bond.
Formal Atom
Charges
• Atoms in molecules often bear a charge (+
or -).
• The predominant resonance structure of a
molecule is the one with charges as close
to 0 as possible.
• Formal charge
= Group number
– 1/2 (no. of bonding electrons)
- (no. of LP electrons)
Carbon Dioxide,
CO2
+6 - ( 1 / 2 ) ( 4 ) - 4
••
•
•
O
••
C
+4 - ( 1 / 2 ) ( 8 ) - 0
O
=
•
•
0
= 0
Calculated Partial Charges
in CO2
Yellow = negative & red = positive
Relative size = relative charge
Thiocyanate Ion,
SCN
6 - (1/2)(2) - 6 = -1
5 - (1/2)(6) - 2 = 0
••
•
•
S
C
N
•
•
••
4 - (1/2)(8) - 0 = 0
Thiocyanate Ion,
SCN
••
••
•
•
S
C
N
•
•
•
•
••
S
C
N
•
•
••
••
•
•
S
C
N
•
•
••
Which is the most important resonance form?
Calculated Partial
Charges in SCN-
All atoms negative, but
most on the S
••
•
•
S
••
C
N
•
•
Violations of the Octet
Rule
Usually occurs with B and
elements of higher periods.
BF3
SF4
Boron
Trifluoride
• Central atom = _____________
• Valence electrons =
__________ or electron pairs =
__________
The B atom has a
• Assemble dotshare
structure
in only 6 pairs
of electrons (or 3
pairs). B atom in
many molecules is
electron deficient.
Boron Trifluoride, BF3
F
+1
B
-1
•
•
•
•
••
•
•
F
••
•
•
•
•
F
••
What if we form a B—F double
bond to satisfy the B atom
octet?
Is There a B=F Double Bond in
BF3
Calc’d partial charges in BF3
F is negative
and B is
positive
Sulfur Tetrafluoride,
SF
• Central atom = 4
• Valence electrons = ___ or ___
pairs.
• Form sigma bonds and
5 pairs
around the S
distribute electron
pairs.
atom. A common
occurrence outside the
2nd period.