When a chemical bond forms, energy is __________________________. When a chemical bond breaks, energy is
__________________________.
Example: Approximately 1652 kJ of energy is required to break a mole of methane, ________, into separate C and H atoms.
Therefore, 1652 kJ is ______________________when 1 mole of methane is formed from gaseous C and H atoms.
Bond energy –
Example: Methane has four identical ___________ bonds. We can calculate the average bond energy for __________.
single bond- double bond - triple bond –
Table 8.4 (p 373) gives average bond energies
Table 8.5 ( p 374) gives bond lengths for selected bonds
A relationship exists between the number of shared electron pairs and the bond length. In general, as the number of shared electrons_____________________, the bond length __________________, and bond energy _________________.
Bond Energy and Enthalpy
Enthalpy Change –
Bond energy can be used to calculate ∆ H for a reaction.
Using bond energies: ∆H = ∑ D reactants
(bonds broken) - ∑ D products
(bonds formed)
Example 1: Calculate the ∆H, enthalpy, using bond energies for the following reaction: H
2
(g) + F
2
(g)

2HF (g)
Example 2: Calculate the ∆H for the following reaction using bond energies: H-C=C-H(g) + H
2
(g)

CH
2
=CH
2

Examples
1. PH
3
3. CO
2
5. carbonate ion
8.9 THE LOCALIZED ELECTRON BONDING MODEL
Localized Electron (LE) Model – assumes a molecule is a compound of atoms that are bound together by
_______________ _______________ of electrons using the_____________ _____________ of the bound atoms.
Lone pairs – pairs of electrons ________________ on an atom.
Bonding pairs – pairs of electrons found in the space _________________ atoms.
8.10 LEWIS STRUCTURES
Lewis Structures
– shows how the ___________________ electrons are arranged among the atoms in the molecule/polyatomic ion.
Every period 1 and 2 element (with the exception of H, He, B, and Be) can form compounds of lowest energy if their highest energy levels are filled (s 2 p 6 ). This is called the _______________ rule.
Hydrogen follows a ____________rule, it needs _________ electrons to be stable. (He already has a _________)
We will discuss Be and B later.
RULES FOR DRAWING LEWIS STRUCTURES
1. Choose the center atom. It is usually the least electronegative atom. C is always the center atom, H is never the
center atom.
2. Draw a skeletal structure – symmetrically arrange the other atoms around the center atom.
3. To determine number and types of bonds - one strategy you can use is S = N – A
S= shared electrons (those involved in bonding)
N = needed electrons. This is the total number of electrons needed for an atom to be stable (either 8 or 2 – we
will discuss exceptions to this later)
A = available valence electrons
4. Complete the structure by adding lone pairs to complete octets for all atoms
5. Double check the # of electrons used in the structure – must equal A, number of available valence electrons.
2. HCN
4. phosphite ion

Examples:
1. PCl
5
3. RnCl
2
8.11 Exceptions to the Octet Rule
The second row elements B and Be often have fewer than 8 electrons (electron deficient) around them in compounds.
Usually, Be needs (N=) _____ electrons, and B needs (N=) __________ electrons.
Examples:
1. BeCl
2
2. BH
3
Odd number of electrons: In a few molecules and polyatomic ions, the number of valence electrons is odd and an octet around each atom cannot be achieved.
Ex. NO
The third row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals (expanded octet)
When using S=N-A, if S is less than the number needed to bond all atoms to the central atom, then an expanded octet is needed around the center atom.
When writing the Lewis structure for a molecule, bond all atoms to the central atom and satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied then place them on the element having available d orbitals (elements in period 3 or beyond) – this will involve the center atom.
2. ClF
3

The arrows do not mean the structure “flips” from one resonance structure to another. They simply show the structure is the __________________ of the resonance structures. Measurements in bond length suggest all 3 N-O bonds lengths are ____________________________.
Bond order – The number of electron pairs involved in a covalent bond. A single bond has a BO of ______, a double bond ________, and a triple ______. For resonance structures, the BO is the average of the bonds.
Example: For the nitrate ion, the bond order for nitrogen-oxygen bond is
Formal Charge
Resonance usually involves equivalent Lewis structures – contain same number of single and multiple bonds.
Nonequivalent Lewis structures contain different numbers of single and multiple bonds.
When we assign oxidation numbers to atoms involved in a covalent bond, we always count both the shared electrons as belonging to the more ____________________ atom in a bond. It is useful for keeping track of electrons in a redox reaction, but not a realistic estimate of charges of atoms in individual molecules. Another definition of charge on an atom in a molecule, the formal charge, can be used to evaluate Lewis structures and select the most stable structures when nonequivalent structures are present.
Formal charge – the hypothetical charge on an atom in a molecule or polyatomic ion. It is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule.
To calculate the formal charge on an atom:
1) Determine the number of valence electrons on the free, neutral atom.
2) Take the sum of the lone pairs electrons and one-half the shared electrons. This is the number of valence
electrons assigned to the atom in the molecule. Subtract this number from the number of valence e on the
free atom.
The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species.
If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion.
Example: Draw all resonance structures and select the most stable one for the thiocyanate ion, SCN .

8.13 MOLECULAR STRUCTURE: THE VSEPR THEORY
Molecular structure – the 3D arrangement of the atoms in a molecule
VSEPR (valence shell electron-pair repulsion ) model is useful is predicting the geometries of molecules formed from nonmetals. The main postulate is that the structure around a given atom is determined principally be minimizing the electron-pair repulsions – bonding and nonbonding pairs around a given atom will be positioned as far apart as possible.
In VSEPR notation the molecule is represented by a formula using the letter A for the central atom, X for terminal atoms and E for lone pairs of electrons. The table below uses this notation and shows the geometry of various combinations.
Determining the Shape of a Molecule
The best way to determine the architecture of a molecule is to:
1. Determine what the central atom is.
2. Draw the Lewis structure of the molecule.
3. Determine the number of bonding pairs and lone pairs around the central atom.
4. Refer to the following chart and determine the shape of the molecule.
# of
Bonding
Groups on
Central
Atom
2
# of
Lone
Pairs on
Central
Atom
0
VSEPR notation
AX
2
Molecular
Geometry
Linear
Bond
Angle
180
Diagram
2
2
2
1 AX
2
E
AX
2
E
2
Bent
Bent less than
120
Less than
109.5
Electron Pair
Geometry
Linear
Trigonal
Planar
Tetrahedral
2 3 AX
2
E
3
Linear 180
Trigonal
Bipyramidal
3 0 AX
3
Trigonal
Planar
120 Trigonal
Planar

3
3
1 AX
3
E Trigonal pyramidal
Less than
109.5
2 AX
3
E
2
T-shaped 90
4 0 AX
4
Tetrahedral 109.5
4
4
5 0
2
1 AX
4
E
AX
4
E
2
See-saw
Square planar
90 and
120
90
AX
5
Trigonal bipyramidal
90 and
120
5 1 AX
5
E Square pyramidal
90
6 0 AX
6
Octahedral 90
Tetrahedral
Trigonal
Bipyramidal
Tetrahedral
Trigonal
Bipyramidal
Octahedral
Trigonal
Bipyramidal
Octahedral
Octahedral

PRACTICE : MOLECULAR GEOMETRY
1) Draw the Lewis structure and predict the molecular structure (geometry) – including
bond angles for each of the following:
(a) XeCl
2
(b) SeO
3
(c) TeF
4
(d) SCl
2
Dipole moment – molecule has a center of ______________ charge and a center of ____________________ charge.
Some molecules have polar bonds but do not have a dipole moment. This occurs when the individual bond polarities are arrange in such a way that they ____________________ each other out.
Example: CO
2
If the individual bond polarities do not cancel each other out – then a dipole moment exists.
Example: NH
3
Summary
If a molecule has nonpolar bonds, then the molecule itself must be ______________________.
If a molecule has polar bonds and an asymmetrical shape, the molecule will be _______________________.
Examples of asymmetrical geometries:
If a molecule has a symmetrical shape and the bonds are polar, the molecule may be _______________ or _____________.
Examples of symmetrical geometries:
It’s nonpolar if all the atoms bonded to the central atom are the same. Ex:
It’s polar if all the atoms bonded to the central atom are not the same. Ex:

HOMEWORK 10/22 : BOND ENERGY
Use bond energy values (Table 8.4) to estimate ∆H for each of the following reactions:
1. H-C ≡ N + 2 H
2
 CH
3
NH
2
2. N
2
H
4
+ 2 F
2

N≡ N + 4 HF
HOMEWORK 10/23 : LEWIS STRUCTURES
Draw Lewis Structures for the following molecules/ions:
1. nitrogen trifluoride 2. sulfur dioxide
3. chlorite ion 4. perchlorate ion
HOMEWORK 10/24 :LEWIS STRUCTURES
Draw Lewis Structures for the following molecules/ions:
1. BeF
2
3. SF
4
2. XeF
4, Br
3
-
4
HOMEWORK 10/25 : LEWIS STRUCTURES, RESONANCE, AND FORMAL CHARGE
1. Which of the following have center atoms that obey the octet rule? Draw Lewis Structures
for each.
(a) AsF
3
(b) ICl
4
(c) XeO
4
2. (a) Draw 2 possible Lewis structures that obey the octet rule for nitrosyl chloride, NOCl.
(b) Using formal charges, explain which Lewis structure is more likely to be correct.
3. (a) Draw 2 resonance structures for sulfur dioxide.
(b) What is the bond order of the sulfur- oxygen bond?
HOMEWORK 10/26: MOLECULAR GEOMETRY
1) Draw the Lewis structure and predict the molecular structure (geometry) – including
bond angles for each of the following:
(a) SiF
4
(b) ICl
3
(c) PCl
3
(d) PCl
5
2) Draw the Lewis structure and predict the molecular structure (geometry) – including
bond angles for each of the following:
(a) ICl
5
(b) XeCl
4
(c) SeCl
6
HOMEWORK 10/29: MOLECULAR POLARITY
Write Lewis Structures and predict both the molecular structure and polarity for the following:
(a) SO
3
(b) NF
3
(c) IF
3
(d) CF
2
Cl
2
(d) COS
(e) SeF
6