Chapter 9: Bonding and Molecular Structure; Fundamental Concepts
Chapter 9 Problem Set
Pages 430-434
29, 33, 37, 39, 43, 47, 50, 54, 63, 65, 67, 69, 81, 87, 93
9.1
VALENCE ELECTRONS
Valence electrons are in the outmost energy level. The remaining electrons are core electrons.
Valence electrons determine the chemical properties of an element
For main group elements, (A columns) valence electrons are the s and p electrons in the outermost
energy levels
For main group elements, the number of valence electrons is equal to the group number.
Lewis Symbols for Atoms/Lewis electron dot symbols
Symbol surrounded by dots representing valence electrons
Octet of electrons is considered to be a stable configuration
Exceptions are H and He
EXAMPLE 9L.1 Valence Electrons
Give the number of valence electrons for Ga and Br. Draw the Lewis electron dot symbol for each
element
9.2
CHEMICAL BOND FORMATION
Ionic bonds form when one or more valence electrons are transferred from one atom to another,
creating positive and negative ions. The “bond” is the attractive force between positive and negative
ions
Each atom in the compound has an octet of electrons.
For example:
Covalent bonding involves, in contrast, involves sharing of valence electrons between atoms.
Neither atom has an octet
For example:
Most cmpds are somewhere in between. READ bottom of page 376 and top of page 377
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9.3
BONDING IN IONIC COMPOUNDS
The tendency to gain noble gas configurations by losing and gaining electrons is important
Even more important is the favorable energetics of ionic compound formation.
 A chem reaction occurs when the products have a lower PE than the reactants ie: product favored
 The structure of a cmpd, ionic or covalent is the one having the lowest potential energy, greatest
thermodynamic stability
Energy of Ion pair Formation
The overall energy of the formation of a [Na+, Cl-] ion pair can be considered to be the sum of 3
individual steps:
 The ionization of Na
 The electron affinity of Cl
 The formation of an ion pair
The energy of attraction cannot be measured directly but can be calculated from Coulomb’s Law
E ion pair 
( n  e) ( n  e)
C
d
C is a constant
d = distance between ion centers
e = electron charge
n= number of positive or negative charges on the ion
The energy of attraction depends on 2 factors
 The magnitude of the ion charges
 The distance between the ions
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Lattice Energy
Ionic compounds exist as solids because…
The lattice energy, Elatticeis the energy of formation of 1 mol of a solid crystalline ionic compound
when….
EXERCISE 9.2 Page 381 Using Lattice Energies
Calculate the molar enthalpy of formation,  f0 of solid sodium iodide using Table 9.3 and
Appendices F and L
Compounds such as NaCl2 and NaNe do not exist because….
9.4
COVALENT BONDING
Sharing of electrons (each atom “thinks” it has an octet)
Polyatomic ions have covalent bonds
For Example: NH4+
EXERCISE 9L.3
All of the compounds contain both ionic bonds and covalent bonds except
NaCN,
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KSCN,
CsNO3,
LiBr,
MgCO3
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Lewis Electron Dot Structures
Write the Lewis structure for Cl2 showing both the bonding pair and nonbonding (lone) pairs of
electrons
Show Lewis structures for CO2 and N2. Tell how many bonding and nonbonding pairs for each atom
What is the octet rule?
Lewis structures for molecules
1. Determine arrangement of atoms within molecule
(central atom). Central atom usually one with
lowest E.A. (usually C, N, P, S)
2. Determine total number of valence electrons
in molecule or ion
3. Place 1 pair of electrons between each pair of
bonded atoms (to form a single bond)
4. Use any remaining pairs as lone pairs around
each terminal atom ( except H) so that each
atom is surrounded by 8 electrons
5. If the central atom has fewer that 8 electrons at
this point, move one or more of the lone pairs on
terminal atoms to form multiple bonds
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How many valence electrons are in the acetate ion, CH3COO-
EXAMPLE 9.2 Page 384 Drawing Lewis Structures
Draw Lewis structures for NH3, ClO- , NO2+ , and PO431.
2.
3.
4.
Central atom
Valence electrons
Form single covalent bonds
Place the remaining pair of
electrons on the central atom
Predicting Lewis Structures
Number of expected bonds
EXAMPLE 9.3 Page 386 Drawing Lewis Structures
Oxo Acids and their anions Table 9.5 pg 389
NO3 and HNO3 H is attached to single bonded O
Isoelectronic species (isostructural) Table 9.6 pg 391
Molecules and ions that have the same number of valence electrons and the same Lewis structrues
*Problem Solving Tips pg 391
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Resonance
In certain cases, the Lewis structure does not adequately describe the properties of the ion or
molecule that it represents.
A molecule can actually be a hybrid of 2 or more possible structures, proven by bond lengths.
Ex: ozone page 392
Ex: benzene page 393
*Problem solving tips and Ideas pg 393
EXERCISE 9.7 Page 394 Drawing Resonance Structures
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Exceptions to the octet rule
Compounds in Which an Atom has Fewer than Eight Valence Electrons
BCl3, BF3, BBr3, BI3, B(OH)3
May contain coordinate covalent bonds
Compounds in Which an Atom has More than Eight Valence Electrons
SF6, PF5, ClF3
Only elements of the 3rd and higher periods in the periodic table form compounds in which an octet
is exceeded.
Central atom usually bonded to F, Cl, or O
Molecules with an Odd Number of Electrons NO and NO2
Two Nitrogen oxides violate the octet rule
These compounds are called as free radicals
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9.5
BOND PROPERTIES
Bond Order
The order of a bond is the number of bonding electron pairs shared by 2 atoms
Can have fractional bond orders when there are resonance structures
Ex: O3
Bond Length
Bond length is the distance between the nuclei of 2 bonded atoms
Factors affecting bond length
size of atoms
number of electron pairs shared
Bond Energy
The bond dissociation energy (D) is the enthalpy change for breaking a bond in a molecule with the..
Process of breaking bonds is always endothermic
Formation of bonds from atoms or radicals in the gas phase is always exothermic.
The enthalpy change for a reaction depends upon the strength of the bonds broken compared to the
strength of the bonds formed. The enthalpy change can be estimated using the equation:
H0rxn = D (bonds broken) - D (bonds formed)
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EXAMPLE 9.6 Page 405 Using Bond Energies
Acetone, a common industrial solvent, can be converted to isopropanol, rubbing alcohol, by
hydrogenation. Calculate the enthalpy change for the reaction using bond energies.
Exercise 9.10 Page 406 Using Bond Energies
Using the bond energies In Table 9.9, estimate the heat of combustion of gaseous methane, CH4. That
is, estimate the H0rxn for the reaction of methane with O2 to give water vapor and carbon dioxide gas.
9.6
CHARGE DISTRIBUTION IN COVALENT COMPOUNDS
Some bonding electrons aren’t shared equally. As a result, some atoms have partial charges
Formal Charges on Atoms
Formal charge of an atom or ion = group number of the atom
– [number of unshared electrons belonging to a given atom + ½ (number of bonding electrons)]
The sum of the formal charges on the atoms in a molecule or ion always equals its net charge.
EXAMPLE 9.7 Page 408 Calculating Formal Charges
Calculate formal charges for the atoms in NH4+ and in one resonance structure of CO32-.
EXERCISE 9.11 Page 408 Calculating Formal Charges
Calculate formal charges on each atom in CN-1 and SO3.
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Bond Polarity and Electronegativity
Pure covalent bond occurs only when 2 identical atoms are bonded and electrons are shared equally.
Cl2
Unequal sharing results in a polar covalent bond, forming a dipole or dipolar molecule.
The polarity of a bond is determined by an element’s electronegativity, , the ability of an atom in a
molecule to attract electrons to itself.
Largest is F 4.0, smallest is Fr.
Periodic trend electronegativity increases left to right, decreases down a column
Electronegativity values Fig 9-10 pg 410
EXAMPLE 9.8 Page 410 Estimating Bond Polarities
For each of the following bond pairs, decide which is more polar and indicate the negative and
positive poles.
1. Li – F and Li – Cl
2. Si – O and P – P
3. C = O and C = S
EXERCISE 9.12 Page 411 Bond Polarity
For each of the following bond pairs, decide which is more polar and indicate the negative and
positive poles. Make predictions first using periodic table. Then calculate .
1. H – F and H – I
2. B – C and B – F
3. C – S and C – Si
Combining Formal Charge and Bond Polarity
Some dilemmas
Electroneutrality principle
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EXERCISE 9.13 Page 413 Formal Charge, Bond Polarity, and Electronegativity
Consider all possible resonance structures for SO2. What are the formal charges on each atom in
each resonance structure? What are the bond polarities? Do they agree with the formal charges?
9.7
MOLECULAR SHAPES
The physical and chemical properties of cmpds are related to the structures of the molecules.
The valence shell electron-pair repulsion model (VSEPR) model was developed by Gillespie and
Nyholm. The central idea is that the bond and lone electron pairs in the valence shell of an atom
repel each other and seek to be as far apart as possible.
Central Atoms (A) Surrounded Only by Bond Pairs (X)
4 single bonds or 4 electron pairs, AX4
tetrahedral
3 single bonds or 3 electron pairs, AX3
trigonal planar
2 single bonds or 2 electron pairs, AX2
linear
5 single bonds or 5 electron pairs, AX5
trigonal-bipyramidal
6 single bonds or 6 electron pairs, AX6
octahedral
Central Atoms with Bond Pairs and Lone Pairs
Electron pair geometry
Molecular geometry
Effect of Lone Pairs on Bond Angles
From consideration of the relative strength of repulsions:
Lone pair-lone pair > lone pair-bond pair > bond pair –bond pair
Account for bond angles in CH4, NH3, and H2O
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EXAMPLE 9.11 Page 418 Finding the Shapes of Molecules
What are the molecular shapes of H3O+ and ClF2+
EXERCISE 9.15 VSEPR and Molecular Shape
Give the electron-pair geometry and molecular shape for BF3 and BF4- . What is the effect on the
molecular geometry of adding and F- ion to BF3 to give BF4- ?
Central Atoms with More than 4 Valence Electron Pairs
Consider the following species
AX4·E
AX4·2E
AX3·2E
AX5·E
Multiple Bonds and Molecular Geometry
Multiple bonds occupy the same region in space as a single bond and contribute to the molecular
geometry the same as a single bond.
Ex: CH2O
EXAMPLE 9.13 Page 421 Finding the Shapes of Molecules and Ions
What are the shapes fo the nitrate ion, NO3-, and XeOF4?
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EXERCISE 9.17 Page 422 Determining Molecular Shapes
Use Lewis structures and VSEPR model to determine the electron-pair and molecular geometries
for
1. the phosphate
2. the sulfite ion,
3. IF5
ion, PO43SO32-
9.8
MOLECULAR POLARITY
Some molecules with polar bonds are polar and some are nonpolar
Polar molecules have a dipole moment.
BF3 is a nonpolar molecule, whereas NF3 is a polar molecule
H2O is a polar molecule, but CS2 is a nonpolar molecule
EXAMPLE 9.14 Page 428 Molecular Polarity
Are nitrogen trifluoride (NF3), dichloromethane (CH2Cl2), and sulfur tetrafluoride (SF4) polar or
nonpolar? If polar indicate positive and negative sides of the molecule
EXERCISE 9.18 Page 428 Molecular Polarity
For each of the following molecules, decide whether the molecule is polar and which side is
positive and which negative: BFCl2, NH2Cl, and SCl2.
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