CHAPTER 9:
BONDING AND
MOLECULAR STRUCTURE
"I have attempted to
give you a
glimpse...of what
there may be of soul
in chemistry.”
-G.N. Lewis
9.0 Objectives:






Understand the basic process of ionic bonding: identify ionic
compounds and describe their internal structure and properties,
understand how size of ions affects ionic properties, calculate
lattice energy, and draw Lewis diagrams of ionic structures.
Identify covalent compounds and characterize their properties.
Draw Lewis structures of covalent substances including exceptions
such as reduced and expanded octets, radicals, and resonance
structures.
Define and predict trends in bond order, bond length, and bond
dissociation energy. Use bond energy to predict enthalpy of a
reaction.
Understand the concept of electronegativity and how it is used to
predict polarity of individual bonds and entire molecules.
Use VSEPR Theory to predict the shapes of simple covalent
molecules.
Homework:

HW#1 – 29, 33, 38, 39, 40, 41, 43, 45
– Valence e-, LEWIS STRUCTURES!

HW#2 – 47, 49, 51, 53, 95
– Formal Charge, Polarity, Electonegativity

HW#3 – 69, 93, 109
– Bond Energy

HW#4 – 73, 75, 77, 79, 81, 89, 99, 103
– Molecular Geometry

HW#5 – 83, 85, 97
– Polarity
9.1 VALENCE ELECTRONS

1. Bonding – definition
– “forces that hold atoms together”
– What role do the e- play?

2. Valence electrons vs. Core electrons
– MAIN GROUP Elements

Outermost “s” and “p” e-
– Transition Metals

Outermost “s” and “p” e- as well as (n-1) “d” e-
When in doubt, write the noble gas e- configuration
9.1 VALENCE ELECTRONS

3. Lewis dot diagrams of elements
– Diagrams that showcase valence e– Lewis says, “Place the first four dots separately!”

Ex. Li, Be, B, C, N, O, F, Ne
9.2 CHEMICAL BOND
FORMATION

1. Ionic bonding – definition and Lewis
representation
– Bond between metal and nonmetal due to
“electrostatic interactions”


Metal donates eNonmetal accepts e-
– Ex. NaCl and Na2S
9.2 CHEMICAL BOND
FORMATION

2. Covalent bonding and Lewis representation
of Cl2
– Bond in which e- are shared
– Overlap of e- density between 2 orbitals
– Ex. Cl2
9.2 CHEMICAL BOND
FORMATION

3. Continuum
– Complete ionic or complete sharing of e- is a bit
extreme; most bonding has uneven sharing of e(sometimes ionic, sometimes covalent)

4. Other bond types
– Metallic bonding

Ex. Alloys
9.3 BONDING IN IONIC
COMPOUNDS

1. Steps in formation of NaCl
– 1. Na(g)  Na+(g) + e-
E = +496 kJ/mol
– 2. Cl(g) + e-  Cl-(g)
E = -349 kJ/mol
– 3. Na+(g) + Cl-(g)  [Na+, Cl-] E = -498 kJ/mol
Eoverall = -351 kJ/mol
9.3 BONDING IN IONIC
COMPOUNDS

2. Lattice energy
– “energy for the formation of 1 mol of solid
crystalline ionic compound when ions in the gas
phase combine”
9.3 BONDING IN IONIC
COMPOUNDS

3. Formula units
– RECALL: Smallest repeating unit of an ionic
compound
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

1. Diagram of H2
– Single Hydrogens
H
H
Both want 1s2
– 1 shared pair
H:H
– Bonds are represented as single lines
H—H
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

2. Orbital overlap diagrams of H2, HCl, Cl2
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

3. Terminology – single, double, and triple
bonds, bonding pairs and nonbonding or lone
pairs of electrons
– Single Bond: 2 e- shared between 2 atoms

Ex. H2
– Double Bond: 4 e- shared between 2 atoms

Ex. O2
– Triple Bond: 6 e- shared between 2 atoms

Ex. N2
9.4 COVALENT BONDING
AND LEWIS STRUCTURES
– Bonding Pairs: e- involved in bonding

(See preceding examples)
– Nonbonding (lone) pairs: e- that are not involved
with bonding but help provide the octet for an atom

Ex. Cl2
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

4. Octet Rule
– The “INFAMOUS” noble-gas configuration
– “tendency for molecules/polyatomic ions to have
structures in which 8 e- surround each atom”
– H, He have a “duet”
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

5. Rules for drawing Lewis structures
– a. Choose a central atom



Usually the atom with the lowest e- affinity
Usually makes a lot of bonds
Halogens are generally terminal atoms
– b. Count the total number of valence electrons



Neutral Molecule: sum of valence e- for each atom
Anions: sum of valence e- and negative charge
Cations: valence e- minus the total positive charge
9.4 COVALENT BONDING
AND LEWIS STRUCTURES
– c. Draw a skeleton structure

Use one pair of electrons to form a bond between each pair
of bound atoms
– d. Place the remaining electrons to fulfill the octet
rule


Do this for each atom
Hydrogen gets a duet
9.4 COVALENT BONDING
AND LEWIS STRUCTURES
– e. Lack of electrons:


Requires multiple bonds (double, triple)
Could be more than one multiple bond
– f. Too many electrons:


Verify that your structure is correct (octets for all?)
Watch anions!
9.6 Lewis Structures of
Some Simple Molecules
O-V=S
S: Shared e- in bonds
O: total # e- required for an Octet
V: Valence e- for all elements
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

6. Diagrams of H2 F2 CH4 NH3 H2O HF
OH- NH4+
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

H2 F2 CH4 NH3 H2O HF OH- NH4+
9.4 COVALENT BONDING
AND LEWIS STRUCTURES

7. Isoelectronic species: NO+ N2
CO
CN-
9.5 RESONANCE

1. Definition
– Alternative and equivalent Lewis structure
“created” by shifting the e- in a structure
– Spinning Rim Analogy
9.5 RESONANCE

2. Examples: NO3- and NO2-
9.5 RESONANCE

3. Experimental evidence says:
– “It’s a combination of both”
– There are however, MORE PREVALENT resonance
structures for some molecules
– Benzene is the most classic of all resonance
structures
9.5 RESONANCE
9.6 EXCEPTIONS
TO THE OCTET RULE

1. Reduced octets for H, B and Be
– Ex. BeCl2, BCl3 (Be = 4 e-, B = 6e-)
9.6 EXCEPTIONS
TO THE OCTET RULE

2. Expanded octets: PF5 SF6
ClF4- XeF2
– Watch these elements (and some others) for
expanded octets: P, S, Cl, As, Se, Br, Kr, Xe
9.6 EXCEPTIONS
TO THE OCTET RULE

3. Radicals (paramagnetic): NO and NO2
– Structure that has unpaired e– Extremely Reactive
9.6 EXCEPTIONS
TO THE OCTET RULE

3. Problems with Lewis structures
– Only show 2-D view life (chemistry) is 3-D
– Works for most molecules, but not all
– Doesn’t show how evenly/unevenly e- are being
shared
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

1. Definitions: polar and nonpolar bonds
– Nonpolar bonds: 2 e- in a bond are “evenly” shared
between the 2 atoms
– Polar bonds: 2 e- in a bond are unevenly shared;
one atom is taking more of the e- density; atoms
have a partial charge
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

2. Electronegativity
– a. definition

(EN): ability of an atom to attract bonding e- to itself when
the atom is in a molecule
– b. Table and Periodic trends


See Pg.10 in Reference Booklet
Increases going left to right and bottom to top
– (Fluorine greatest at 4.0)
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

3. EN – parameters
– Prediction of “Ionic Character”
Pure Covalent
0
.5
In General:
Pure Ionic
1
0.0 < 0.45
0.45 ≤ 1.8
> 1.8
1.5
2
Nonpolar
Polar Covalent
Ionic
2.5
3
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

4. Ex9.1 Arrange the following bonds in order
of increasing polarity: F-Cl, F-F, F-Na
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

5. Central atom in Lewis structure:
– Many times has a formal charge
– Making more/less bonds than it “normally” does
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

6. Formal Charge
– a. Definition and Use

Charge for an atom in a molecule based on premise that
bonding e- are evenly shared
– b. Calculating – equation

Formal Charge = Group # - [Lone Pair e- + ½ Bonding e-]
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES
– c. Examples: OH- and NO3-
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

Ex9.2 Calculate the formal charge on each
atom in CO32- and NH4+
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

7. Electroneutrality – Definition
– The e- in a molecule are distributed so that the
formal charge is minimal
– Most Probable Lewis Structure = one with minimal
FC; minimal FC is more important than symmetry
– Negative charge should reside on the most
electronegative element
– Formal charge > +/- 2 is not likely
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES
– a. Example: CO2
9.7 CHARGE DISTRIBUTION IN
COVALENT BONDS AND
MOLECULES

b. Ex9.3 Use formal charge and the electroneutrality concept to
determine the most likely structures for N2O and OCN1-
9.8 BOND PROPERTIES

1. Bond order
– a. definition and examples




number of bonding e- pairs shared between 2 atoms
Usually an integer (1, 2, or 3)
BOND ORDER =
(# shared pairs linking X-Y)_____
(number of X-Y links in the molecule)
Ex. CH4, CO2
– b. resonance structures



Bond order are fractions
e- residing over both locations evenly
Ex. O3
9.8 BOND PROPERTIES

2. Bond Length – definition and examples
– Bond length: distance between nuclei in a covalent
bond

More Polar bonds = shorter length
More bonds = shorter length

Ex. C-C

1.54Å
C=C
C≡C
1.34Å
1.20Å
9.8 BOND PROPERTIES

1. Bond dissociation energy – definition and
examples
– Bond Dissociation Energy (D):

energy needed to break a covalent bond in the gas phase
– Higher Bond Order Higher D
9.8 BOND PROPERTIES
– a. Estimating Enthalpy of reaction from bond
energies – equation
 Hrxn
=  D (bonds broken) -  D (bonds
formed)
 Energy
is required to break bonds
 Energy is released when bonds are
formed
9.8 BOND PROPERTIES
9.8 BOND PROPERTIES
– b. Example: Estimate the HRxn for the synthesis
reaction between gaseous hydrogen and chlorine.
9.8 BOND PROPERTIES
– c. Ex9.4 Estimate the enthalpy of reaction for the
combustion of methane, CH4, to produce gaseous
carbon dioxide and water vapor.
9.9 MOLECULAR SHAPES

1. Gumdrops and toothpicks
– A tasty way to practice
chemistry!
9.9 MOLECULAR SHAPES

2. VSEPR Theory and importance of shapes
– VSEPR: Valence Shell Electron Pair Repulsion
– VSEPR gives our 2-D Lewis structures LIFE (3-D)
– Geometry has HUGE impact on properties
– Based on idea pairs of e- in bonded atoms repel one another

Want to be as far apart as possible Gives shape
– Electron Group

Any collection of valence e- around an atom that repel other e– Single unpaired e– Lone pair e– Bonding pairs of e- (1, 2, 3)
9.9 MOLECULAR SHAPES

NOTATION:
AXnEm
A = central atom
X = terminal atoms
E = lone pair e- on central atom
9.9 MOLECULAR SHAPES

3. Single bonds, no unshared pairs of
electrons
– Lewis structure, geometry and bond angles of:
– a. BeH2
9.9 MOLECULAR SHAPES


BH3
CH4
9.9 MOLECULAR SHAPES
9.9 MOLECULAR SHAPES

4. Unshared pairs of electrons on the central
atom
– a. NH3
– b. H2O
9.9 MOLECULAR SHAPES
– c. analogs (H2S, PCl3, etc.)
9.9 MOLECULAR SHAPES

Ex9.5 Predict the molecular geometry and bond
angles of HOCl and SiO44-
9.9 MOLECULAR SHAPES

6. Multiple Bonds
– a. CO2
– b. H2CO
9.9 MOLECULAR SHAPES
– c. HCN
9.9 MOLECULAR SHAPES

7. Ex9.6 Predict the molecular geometry and bond angles in the
following species: C2H2 C2H4 ClO31- NO31- N2O ONCl
9.9 MOLECULAR SHAPES

7. Expanded octets
– a. definition and recognizing


More than 4 e- groups around a central atom
Use formal charge to help guide Lewis
Structure/Geometry
9.9 MOLECULAR SHAPES

b. Examples:
– PCl5
– SF6
9.9 MOLECULAR SHAPES
– ClF5
– XeF4
9.9 MOLECULAR SHAPES

8. Ex9.7 Determine the molecular geometry and bond angles
of ICl2-1 IF3 XeOF4 How does the electronic geometry differ
from the molecular geometry for these species?
9.10 MOLECULAR
POLARITY

1. Nonpolar molecules – definition and
examples, CH4 CO2
– Molecule consisting entirely of nonpolar bonds
OR
– Molecule with polar bonds that cancel one another
out
9.10 MOLECULAR
POLARITY

2. Dipoles – definition and examples,
H2O
NH3
– Molecule with separate centers of (+) and (-)
charge
– Polar bonds present with no canceling out
9.10 MOLECULAR
POLARITY

4. Dipole moments and vectors ()

=d
 = magnitude of charge
d = distance

Measured in debye (D)
– Nonpolar:  = 0
– Polar:  ≠ 0
9.10 MOLECULAR
POLARITY

5. Rules for determining polarity of a molecule. A
molecule is a dipole if:
– (uneven balance of e- density)
– RULES:




0. Draw Lewis Structure
1. Use VSEPR to predict molecular shape
2. Electronegativity to predict bond dipoles
3. Determine whether bond dipoles cancel to produce nonpolar
or combine to give polar molecule
– KEY CLUES:



1. Polar Bonds
2. Lone Pair e3. 2 different atoms bonded to central atom
9.10 MOLECULAR
POLARITY

6. Ex9.8 Determine which of the following are dipoles: SO2
BF3 CO2 N2O ClO3- ONCl NCl3 BFCl2
SCl2
9.10 MOLECULAR
POLARITY

6. Exceptions: XeF4
9.10 MOLECULAR
POLARITY

7. Properties of dipoles
– Special properties observed due to interactions
between molecules (intermolecular forces)
– Attraction to other dipoles
END OF CHAPTER 9!