Unit 6
Chemical Bonding
AP Chemistry
Nikki Chamberlain
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Chapter 9
Bonding Review
 Lewis Dot Structures -
Example 9.1
Give Lewis symbols for magnesium, silicon, and
phosphorus.
 Octet Rule –
 Valence electrons play a fundamental role in
chemical bonding.
 In losing, gaining, or sharing electrons to form
chemical bonds, atoms tend to acquire the electron
configurations of __________________________.
 Forces called __________________________ hold
atoms together in molecules and keep ions in place in
solid ionic compounds.
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 Chemical bonds are _____________________ forces;
they reflect a balance in the forces of attraction and
repulsion between electrically charged particles.
 In this unit we will look at intramolecular forces and
intermolecular forces.
o Intra- means
 Examples:
o Inter- means
 Examples:
 Ionic bonding –
 When atoms lose or gain electrons, they may acquire
a _____________________________, but do not
become noble gases.
 Because the two ions formed in a reaction between a
metal and a nonmetal have opposite charges, they
are strongly attracted to one another and form an
___________________.
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 The net attractive electrostatic forces that hold the
cations and anions together are _________________.
 The highly ordered solid collection of ions is called an
_____________________________.
 Lewis symbols can be used to represent ionic
bonding between nonmetals and:
 Instead of using complete electron configurations to
represent the loss and gain of electrons, Lewis
symbols can be used.
Example 9.2 Use Lewis symbols to show the formation of ionic
bonds between magnesium and nitrogen. What are the name and
formula of the compound that results?
Try It Out!
Use Lewis symbols to show the formation of ionic bonds
between sodium and phosphorus. What are the name and formula of
the compound that results?
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 Lattice energy –
 Stability of ionic compounds (high melting point,
brittle) because of large lattice energy.
 Coulomb’s Law –
 Says that the force of attraction or repulsion between
two point charges is _______________
_______________________ to the product of
magnitude of each charge and_______________
_______________________ to the square of the
distance between them.
 Simply put:
o Because the force is proportional to the charge
on each ion, ________________ charges lead
to __________________ interactions.
o Because the force is inversely proportional to
the square of the distance between the centers
of the ions (nuclei), ___________________
ions lead to __________________ interactions.
o Since the charges of cations and anions are
____________________, they are attracted to
one another.
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 Examples:
o Which has higher lattice energy: LiBr or CaO?
o Which has more lattice energy: NaCl or CsCl?
•
The melting point of MgO is higher than that of NaF. Explanations
for this observation include which of the following?
I.
Mg+2 is more positively charged than Na+1
II.
O-2 is more negatively charged than F-1
III.
The O-2 is smaller than the F-1 ion.
A.
B.
C.
D.
E.
II only
I and II only
I and III only
II and III only
I, II, and III
 Covalent bond –
 Single covalent bond –
 Some molecules require more than single bonds to
provide each atom with the required octet (formed
primarily by __________________)
 Double bond –
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 Triple bond –
The Lewis Theory of Chemical Bonding
 A _______________________________ is a
combination of Lewis symbols that represents the
formation of covalent bonds between atoms.
 In most cases, a Lewis structure shows the bonded
atoms with the electron configuration of a noble gas;
 The shared electrons can be counted for each atom
that shares them, so each atom may have a…
Electronegativity & Bond Type
 _________________________________ is a measure
of the ability of an atom to attract its bonding
electrons to itself.
 EN is related to ______________________________
and ______________________________.
 The greater the EN of an atom in a molecule, the
more ____________________ the atom attracts the
electrons in a covalent bond.
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 Electronegativity generally _________________ from
left to right within a period, and it generally
_________________ from the bottom to the top
within a group.
Example 9.4 Referring only to the periodic table inside the front
cover, arrange the following sets of atoms in the expected order
of increasing electronegativity.
(a) Cl, Mg, Si
(b) As, N, Sb
(c) As, Se, Sb
Drawing Lewis Structures
 The skeletal structure shows the arrangement of
atoms.
 Hydrogen atoms are terminal atoms (bonded to only
one other atom).
 The central atom of a structure usually has the lowest
electronegativity.
 In oxoacids (HClO4, HNO3, etc.) hydrogen atoms are
usually bonded to oxygen atoms.
 Molecules and polyatomic ions usually have compact,
symmetrical structures.
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A Method:
1. Determine the total number of valence electrons.
2. Write a plausible skeletal structure and connect the
atoms by single dashes (covalent bonds).
3. Place pairs of electrons as lone pairs around the
terminal atoms to give each terminal atom (except H)
an octet.
4. Assign any remaining electrons as lone pairs around
the central atom.
5. If necessary (if there are not enough electrons), move
one or more lone pairs of electrons from a terminal
atom to form a multiple bond to the central atom.
Some Handy Rules:
 Hydrogen and the halogens bond ______________.
 The family oxygen is in can bond _______________.
 The family nitrogen is in can bond _____ times. So
can boron.
 The family carbon is in can bond _____ times.
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Example 9.6
Write the Lewis structure of nitrogen trifluoride, NF3.
Example 9.7
Write a plausible Lewis structure for phosgene, COCl2.
Example 9.6 A
Write the Lewis structure of hydrazine, N2H4.
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Example 9.7 A
Write a plausible Lewis structure for carbonyl sulfide,
COS.
Example 9.7 B
Write a plausible Lewis structure for nitrosyl chloride,
NOCl.
What if an element has charge?
 Negative =
 Positive =
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Example 9.8
Write a plausible Lewis structure for the chlorate ion,
ClO3-1.
Example 9.8 B
Write a plausible Lewis structure for the nitronium ion,
NO2+1.
Formal Charge
o Formal charge can be used as a criterion for
determining which of several possible valid Lewis
diagrams provides the best model for predicting
molecular structure and properties (in some cases).
Formal Charge =
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o The best formula:
 Is one with formal charges of 0 on all atoms
 If they are not all 0, they should be as small as
possible. Negative FC should appear on the
most electronegative atoms
 Adjacent atoms in a structure should not carry
formal charges of the same sign
 The total of formal charges must = 0 for a
neutral molecule and must equal the net
charge for a polyatomic ion
Which of the following structures for COCl2 is best?
Resonance:
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 When a molecule or ion can be represented by two
or more plausible Lewis structures that differ ______
___________________________________, the true
structure is a composite, or hybrid of them.
 The different plausible structures are called
____________________________.
 The actual molecule or ion that is a hybrid of the
resonance structures is called a resonance hybrid.
 Electrons that are part of the resonance hybrid are
spread out over several atoms and are referred to as
being __________________________.
Example 9.10 Write three equivalent Lewis structures
for the SO3 molecule that conform to the octet rule, and
describe how the resonance hybrid is related to the
three structures.
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Example 9.10 A
Write three Lewis structures for the nitrate ion, NO3-1.
Molecules that DON’T follow the Octet Rule
 Molecules with an _______________________ of
valence electrons have at least one of them unpaired
and are called ________________________.
 Some molecules have _________________________.
These are usually compounds of ________________;
they generally have some unusual bonding
characteristics, and are often quite reactive.
 Some compounds have _______________________,
which means that the central atom has more than
eight electrons around it.
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 A central atom can have expanded valence if it is in
the ________________________ or lower (i.e.,
_____________________)
Example 9.11
Write the Lewis structure for bromine pentafluoride,
BrF5.
Example 9.11 A
Write the Lewis structure of phosphorus trichloride.
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Example 9.11 B
Write the Lewis structure of chlorine trifluoride.
Example 9.11 C
Write the Lewis structure of sulfur tetrafluoride.
 Keep in mind, as with any model, there are
limitations to the use of the Lewis structure model,
particularly in cases with an odd number of valence
electrons.
 It is important to note that Lewis diagrams have
limitations.
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Electronegativity Difference & Bond Type
 Identical atoms have the same electronegativity and
share a bonding electron pair _______________. The
bond is a ___________________________________.
 When electronegativities differ significantly, electron
pairs are shared _________________________.
 The electrons are drawn closer to the atom of higher
electronegativity; the bond is a __________________
____________________________.
 With still larger differences in electronegativity,
electrons may be completely transferred from metal
to nonmetal atoms to form ____________________.
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 Electronegativity Difference:
Example 9.5 Use electronegativity values to arrange the
following bonds in order of increasing polarity:
Br—Cl,
Cl—Cl,
Cl—F,
H—Cl,
I—Cl
Try It Out – Indicate the type of bond expected to form.
If it is polar covalent, place a δ- to indicate the negative
end of any dipole formed.
a) H – H
b) H – O
c) H – N
d) H – C
e) H – F
f) Na – Cl
g) Cl – F
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Bond Type
 Difference in electronegativity is not the only factor
in determining if a bond is ionic or covalent.
Examination of the properties of a compound is the
best way to determine the type of bonding
Type
Example
Melting Solubility
Electric
Physical
Point
in Water Properties Properties
Ionic
Covalent
Molecular
Covalent
Atomic
Metallic
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Bond Order and Bond Length
 Bond order is
 A single bond has a BO = ____, a double bond has a
BO = _____, etc.
 Bond length is the
 Bond length depends on the…
•
The length of a covalent bond is the sum of the
covalent radii of the two atoms.
– Bigger atoms will have longer bond lengths
– As bond order increases, bond length
decreases (triple bonds < double bonds < single
bonds)
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Example 9.13 Estimate the length of (a) the nitrogento-nitrogen bond in N2H4 and (b) the bond in BrCl.
Example 9.13 A Estimate the oxygen to fluorine bond
length in OF2.
Bond Energy
 Bond-dissociation energy (D) is
 Energies of some bonds can differ from compound to
compound, so we use an ________________ bond
energy.
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Trends in Bond Lengths & Energies
 The _______________ the order (for a particular
type of bond), the _________________ and the
_________________ (higher energy) the bond.
 A ___________________________ is shorter and
stronger than a ______________________________.
 There are ___________ electrons between the two
positive nuclei in ___________. This produces more
electrostatic attraction than the two electrons
between the nuclei in _____________.
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Calculations Involving Bond Energies:
 It _______________________ energy to break a
bond.
o This means that breaking a bond is an
 Energy is _______________________ when a bond is
formed.
o This means that forming a bond is an
 A reaction is ___________________ if it requires
more energy to break bonds than is given off when
they are formed.
 A reaction is ___________________ if it requires less
energy to break bonds than is given off when they
are formed.
Example 9.14 Use bond energies from Table 9.1 to estimate the
enthalpy of formation of gaseous hydrazine. Compare the result
with the value of ΔH°f [N2H4(g)] from Appendix C.
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Example 9.14 A Estimate ∆H for the reaction:
C2H6 + Cl2  C2H5Cl + HCl
Graph of PE vs. Distance
 Bond energies and bond lengths for nonpolar
covalent compounds can be determined by plotting
potential energy vs. distance between atoms
 _____________ atoms are often used as an example.
 The relative strengths of attractive and repulsive
forces as a function of distance determine the
________________ of the graph.
 The _______________________ is the distance
between the bonded atoms’ nuclei, and is the
distance of minimum potential energy where the
attractive and repulsive forces are balanced.
 The _______________________ is the energy
required for the dissociation of the bond compared
to the two separated atoms. Typically, bond energy
is given on a per mole basis.
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 Remember, as bond length ________________, bond
energy ______________________
 _________________________ are stronger bonds
 Bonds become weaker moving down the group.
Molecular Geometry
 _______________________________ is simply the
shape of a molecule.
 Molecular geometry is described by the geometric
figure formed when the atomic nuclei are joined by
(imaginary) straight lines.
 Molecular geometry is found _____________ the
Lewis structure, but the Lewis structure itself does
NOT necessarily represent the molecule’s shape.
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Valence Shell Electron-Pair Repulsion (VSEPR)
• Basis:
• These mutual repulsions push electron pairs as far
from one another as possible.
• An electron group is a collection of valence electrons,
localized in a region around a central atom.
• One electron group:
 The repulsions among electron groups lead to an
orientation of the groups that is called…
 These geometries are based on the ______________
of electron groups.
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Electron Groups
Electron Group
Geometry
2
3
4
5
6
 VSEPR Notation
A=
E=
X=
• The H2O molecule would therefore carry the
designation __________________.
• For structures with no lone pairs on the central
atom…
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 When there are lone pairs, the molecular geometry is
____________________________ the electron-group
geometry.
 In either case, the electron-group geometry is the
tool we use to obtain the molecular geometry.
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Some Examples: Give the electron-group geometry and
VSEPR (molecular) geometry for each structure below:
Example 10.1 Use the VSEPR method to predict the
shape of the nitrate ion.
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Example 10.2
Use the VSEPR method to predict the
molecular geometry of XeF2.
Example 10.3 Use the VSEPR method to describe, as
best you can, the molecular geometry of the nitric acid
molecule, HNO3.
Polar Molecules and Dipole Moments:
•
A polar bond has separate centers of positive and
negative charge.
• A molecule with separate centers of positive and
negative charge is a ______________ molecule.
• The _____________________ (µ) of a molecule is the
product of the magnitude of the charge (δ) and the
distance (d) that separates the centers of positive
and negative charge.
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• Formula:
• A unit of dipole moment is the ______________ (D).
• One debye (D) is equal to ___________________ C
m.
• A polar covalent bond has a ___________________;
a separation of positive and negative charge centers
in an individual bond.
• Bond dipoles have both a __________________ and
a _________________ (they are ______________
quantities).
• Ordinarily, a polar molecule must have polar bonds,
BUT…
• A molecule may have polar bond and be a
_________________ molecule – IF…
• Bond dipoles and Molecular Dipoles
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• To predict molecular polarity:
1. Use electronegativity values to predict bond
dipoles.
2. Use the VSEPR method to predict the molecular
shape.
3. From the molecular shape, determine whether
bond dipoles cancel to give a nonpolar molecule,
or combine to produce a resultant dipole moment
for the molecule.
• Note:
Example: Would the following molecules be polar or nonpolar?
Example 10.4 Explain whether you expect the following
molecules to be polar or nonpolar.
(a) CHCl3
(b) CCl4
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Example 10.5 A Conceptual Example Of the two compounds
NOF and NO2F, one has m = 1.81 D and the other has m = 0.47 D.
Which dipole moment do you predict for each compound?
Explain.
Atomic Orbital Overlap
• Valence Bond (VB) theory states that a covalent bond
is formed…..
• In the overlap region, electrons with opposing spins
produce a high electron charge density.
• In general, the more extensive the overlap between
two orbitals, the stronger is the bond between two
atoms.
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Important Points of VB Theory
• Most of the electrons in a molecule remain in the
same orbital locations that they occupied in the
separated atoms.
• Bonding electrons are _________________ in the
region of atomic orbital overlap.
• For atomic orbitals with directional lobes
______________________, maximum overlap occurs
when the AOs overlap ________________________.
Hybridization of Atomic Orbitals
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• sp3 hybridization:
• sp2 hybridization:
• sp hybridization:
• Easy Rule for Hybridization:
Example What is the central atom hybridization for each of the
structures below:
Hydrid Orbitals and Multiple Covalent Bonds
• Covalent bonds formed by the end-to-end overlap of
orbitals are called ________________________.
• All single bonds are ___________________.
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• A bond formed by parallel, or side-by-side, orbital
overlap is called _____________________.
• A double bond is made up of one ____________ and
one _________________.
• A triple bond is made up of _________________
bond and ______________________________.
• Sigma bonds are ___________________ than pi
bonds.
• When bonds are broken, pi bonds are always broken
_________________.
• A triple bond has _________________ and
______________________. It is stronger than a
single bond because each ______ bond has to be
broken and then a __________ bond has to be
broken.
Ex 10.7 - Formic acid, HCOOH, is the simplest carboxylic acid.
(a) Predict a plausible molecular geometry for this molecule.
(b) Propose a hybridization scheme for the central atoms that
is consistent with that geometry.
(c) Sketch a bonding scheme for the molecule.
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Molecular Orbital Theory
 MO Theory describes covalent bonding in a manner
that can capture a wider array of systems and
phenomena than the Lewis or VSEPR models.
 MO diagrams showing the correlation between
atomic and molecular orbitals, are a useful
qualitative tool related to MO theory.
Intermolecular Attractions
 The attractions _______________ molecules are
much weaker than the covalent bonds ____________
molecules.
 These attractions vary depending on the nature of
the molecules.
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 Dipole – dipole
o Many molecules have ____________________
__________________.
o These polar covalent molecules form
attractions between the positive end of one
molecule and the negative end of the adjacent
molecule.
o Dipole-dipole attractions are ______________
over short distances between molecules and
_______________ as the distance between
molecules increases.
o Dipole-dipole interactions can be represented
by diagrams of attraction between the positive
and negative ends of polar molecules trying to
maximize attractions and minimize repulsions
in the liquid or solid state.
 Dipole-induced dipole interactions
o These are present between a ___________ and
a ________________ molecule.
o The strength of these forces increases with…
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 Hydrogen bonding
o Strong dipole-dipole attractions occur in
molecules containing hydrogen covalently
bonded to…
o The hydrogen in these bonds has a partial
______________ charge, since there are no
inner core electrons and the shared electrons
are strongly attracted to the small, very
_________________________ atoms.
o The attractions in hydrogen bonds are
___________________________ stronger than
other dipole-dipole attractions.
o Hydrogen bonding is responsible for…
o Hydrogen bonding between molecules or
different parts of molecules is represented by a
diagram:
Solubility
 Ionic interaction with dipoles is important in
solubility of ionic compounds in __________
solvents.
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 These interactions are why salt and other ionic
compounds dissolve in water.
 London dispersion forces
o ______________________ molecules do not
have permanent dipoles.
o At any given instant, the electrons may be
unevenly distributed within an atom or
molecule giving it a partial ______________
charge (instantaneous dipole)
• Coulombic interaction of the temporary dipole
with the electron distribution in neighboring atom
or molecule results in atoms being attracted to
one another for an instant (can even happen in
__________________________).
• Dispersion forces increase with contact area
between molecules and with increasing
polarizability of the molecules.
• Polarizability increases with:
– Number of ________________in the
molecule (molar mass)
– Presence of _______ bonding
 Melting point and boiling point…
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 Intermolecular forces depend on…
 General Order of Strength:
 Specifically, the stronger the force, the higher the
melting and boiling point.
 If the forces are the same, larger molecules with less
compact shapes have higher melting and boiling
points.

Draw the Lewis structures for each of these compounds and
indicate which intermolecular forces (if any) are involved.
a) NI3
b) PI3
c) BF3
d) CH3COOH
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Based on the structures and intermolecular forces you indicated on the
previous problem, which compound will have the highest melting point.
Consider mass as well.
AP Exam Questions:
Use principles of atomic structure, bonding and/or intermolecular forces to
respond to each of the following. Your response must include specific
information about all substances referred to in each question.
• At a pressure of 1 atm, the boiling point of NH3 (l) is 240 K, whereas the
boiling point of NF3 (l) is 144 K
• Identify the intermolecular force(s) in each substance.
•
Account for the differences in the boiling points of the substances.
Which of the following describes the changes in forces of attraction that occur
as H2O changes phase from a liquid to a vapor?
a) H – O bonds break as H – H and O – O bonds form.
b) Hydrogen bonds between H2O molecules are broken
c) Covalent bonds between H2O molecules are broken.
d) Ionic bonds between H+ ions and OH- ions are broken.
e) Covalent bonds between H+ ions and H2O molecules become more
effective.
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