Molecular Structure
Structure and bonding
To understand molecular structure, we
will need an understanding of molecular
bonding
Valence Electrons
Valence electrons are the electrons that
participate in bonding to other atoms

Core electrons do not
Every element in the same group has
the same number of valence electrons

This is why they have similar properties
Octet Rule
Each noble gas atom has a filled outermost shell, which results in its stability
All noble gases have eight valence
electrons – hence the octet rule
The octet rule provides a way of
predicting the results of the most
common reactions
Lewis dot symbols
Picture used to represent the atomic
nucleus and its valence electrons


The elements symbol represents the
nucleus and the core electrons
Dots surrounding the symbol represent the
valence electrons
Chemical Bond Formation
Covalent bond – The sharing of
electrons between atoms



Electrons shared by two nuclei
Electrons evenly distributed
Usually occur between nonmetals
Ionic bond – if the electrons involved
are strongly displaced toward one atom
and away from the other

Involve metals interacting with nonmetals
Covalent bonds and drawing Lewis dots
1. determine the number of electrons to be
used to connect the atoms


Consider carbon dioxide CO2
carbon (C) has four valence electrons x 1 carbon = 4 eoxygen (O) has six valence electrons x 2 oxygens = 12 eThere are a total of 16 e- to be placed in the Lewis
structure.
2. Connect the central atom to the other atoms
in the molecule with single bonds. Carbon is
the central atom, the two oxygens are
bound to it and electrons are added to fulfill
the octets of the outer atoms.
Drawing Lewis dots cont.
3. Complete the valence shell of the outer atoms
in the molecule
4. Place any remaining electrons on the central
atom



There are no more electrons available in this
example
If the valence shell of the central atom is
complete, you have drawn an acceptable Lewis
structure
In this example, the valence shell of carbon has
only four electrons – so its not right
Lewis dots
use a lone pair on one of the outer
atoms to form a double bond on each
side
This is the correct Lewis structure for
CO2
More examples
f
Practice
Draw the Lewis structures for the following
compounds (5 points)






CO
Cl2CO
H2O
C2H4
SCl2
N2
Homework: 1-3 page 202 (5 points)
Resonance Structures
For some molecules, you will find that
there are two or more possible ways to
draw the Lewis structures
Linus Pauling proposed the theory of
resonance to deal with this problem
Paulings theory combines the structures
into a composite or resonance hybrid, a
single structure formed by the
combination of equivalent contributing
structures
Resonance Structures
Ozone, O3
Practice drawing resonance structures

Draw the resonance structures for NO3-
Assignment
7-9 page 204
10 points
Exceptions to the Octet Rule
Some compounds have fewer than four
pairs of electrons around the central
atom

Compounds with hydrogen
Some compounds have atoms with
more than eight valence electrons

Elements of period 3 or higher
 PF5
 Compounds of noble gases
Draw the Lewis structure for [ClF2]-
Molecules with an odd number
of electrons
Place the odd electron on the central
atom

Draw the Lewis dot structure for NO
Any atom or molecule with an unpaired
electron is called a Free Radical



Free radicals are very reactive
Are central to formation of air pollutants
Can combine with themselves to form
dimers
Bond properties – Bond order
The order of a bond is the number of
bonding electron pairs shared by two
atoms in a molecule
Bond orders will usually be between 1
and 3
A fractional bond order is possible in
molecules and ions having resonance
structures
Bond order = number of shared pairs
divided by number of links
Bond Length
Bond length is the distance between the
nuclei of two bonded atoms
Is largely determined by the size of the
atoms
Bonds become shorter as the bond
order increases

A triple bond is shorter than a double bond
is shorter than a single bond
Problems
1. Arrange the following in order of
decreasing bond distance
C=N
C=N
C-N
2. Draw resonance structures for NO2What is the NO bond order for this
ion?
Bond Energy
The bond dissociation energy, D, is
the enthalpy (H=U + PV) change for
breaking a bond in a molecule with the
reactants and products in the noble gas
phase under standard conditions
The process of breaking bonds in a
molecule is always endothermic
The process of forming bonds is always
exothermic
Bond Polarity and
Electronegativity
When a bond pair is not equally shared
by two atoms, the bonding electrons
are displaced towards one of the atoms


The atom toward the displaced electron
acquires a negative charge and the other
end becomes positive
This results in the molecule having electric
poles, and is called a polar bond
Bond Polarity and
Electronegativity
If the displacement is complete, the
bond is ionic
If the displacement is less than
complete, the bond is a polar covalent
bond
If no displacement occurs, the bond is
nonpolar covalent
Bond Polarity and
Electronegativity
The electronegativity () of an atom in
a molecule is a measure of the ability of
the atom to attract electrons to itself
Problems
For each of the following, tell which
bond pair is more polar and indicate
the negative and positive poles
1. Li-F and Li-I
2. C-S and P-P
3. C=O and C=S
Oxidation Numbers
Why do atoms have the oxidation
numbers that they do?
Formal Charges on Atoms
The formal charge of an atom is an
estimate of atom charges in a molecule
Atom formal charge = group number –
number of lone pair electrons – ½
(number of bonding electrons)
Atoms in molecules should have formal
charges as small as possible (principle
of electroneutrality)
A molecule is most stable when any
negative charge resides on the most
electronegative atom
Formal Charges on Atoms
Use your knowledge of formal charges
to decide which resonance structures
of the following molecules are most
stable
1. CO2
2. HOCN
Molecular Shape
3-dimensional analysis
Use the Valence Shell Electron-pair
Repulsion (VSEPR) model
We will look at these trends



Central atoms with only bond pairs
Central atoms with bond pairs and lone
pairs
Multiple bonds
VSEPR Model
Valence Shell Electron Pair Repulsion
is a simple but effective model for
predicting molecular geometry.
The first assumption of VSEPR is:

A molecule adopts the geometry that
minimizes the repulsive force among a
given number of electron pairs.
Applying VSEPR
Draw the Lewis structure of the
molecule
Count the number of electron pairs
around the central atom. Multiple bonds
count as one electron pair
The arrangement of electron pairs that
minimizes repulsion is called the
electron-pair geometry
The arrangement of atoms is called the
molecular geometry
Central Atoms with Only Bond
Pairs
AX2 – Linear


BeF2
180°
AX3 – Triangular-planar


BF3
120 °
Central Atoms with Only Bond
Pairs
AX4 – Tetrahedral


CH4
109.5°
AX5 – Triangular-bipyramidal

PCl5
Central Atoms with Only Bond
Pairs
AX6 – Octahedral

SF6
Central Atoms with Bond Pairs
and Lone Pairs
Lone pairs of electrons around the
central atom occupy spacial positions
even though they are not included in
the description of the molecule or ion
N02
H2O
Multiple Bonds and Molecular
Geometry
Although double bonds and triple bonds
are shorter than single bonds, they do
not affect predictions of overall
molecular shape
Exercise
Use VSEPR rules to predict the shapes
and geometries of the central atoms of
the molecules on the handout
Molecular Polarity
For each of the following molecules
decide whether the molecule is polar
and which side is positive and which is
negative
1.
2.
3.
BFCl2
NH2Cl
SCl2
That’s It!
STUDY FOR YOUR TEST!!!