Chapter 8 Bonding

Chapter 8 Bonding
8.1 Types of Bonds
• There are lots of experiments we can do
to determine the nature of materials
Melting point
Charge distribution in an electric field
• Bond Energy – The energy required to
break a bond
– Tells us the strength of a bonds interactions.
Ionic Compounds
• We all know that ionic bonding is a result
of electrostatic attractions of oppositely
charged ions.
• And ionic compounds are formed when a
nonmetal and a metal react.
Coulomb’s law
• Coulomb’s law is used to measure the
energy of interaction between a pair of
 Q1Q2 
E  (2.31x10 J  nm) 
 r 
• Where E is energy in joules, r is the
distance between the ion centers in nm,
and Q1 and Q2 are the charges of the
Coulomb’s law in use
• Let’s look at salt with ions
at the given distance:
E  (2.31x10
  1 1 
J  nm) 
 0.276nm 
E  8.37 x1019 J
• Notice the E is negative (indicating an
attractive force) which means the ion pair
has LOWER energy than the separated
Repulsive forces
• Let’s look at H-H
bonds. When will a
H2 molecule be
A bond will form if
the energy of the
aggregate (whole
created from its parts)
is lower than that of
the separated ions!!
An energy profile
Bond length
• The distance where the energy is
MINIMAL is called the bond length.
• The type of bond we seen in an H2
molecule is called a covalent bond. (where
electrons are shared) A mutual attraction
of the two nuclei for the shared electrons.
Polar Covalent Bonds
• In between the two extreme types of
bonding (ionic and covalent) is an
intermediate case; called polar covalent
• These bonds are between atoms that are
not so different that electrons are
completely transferred but are different
enough that there is unequal sharing.
• Look at the HF molecule below. The
symbol δ (lowercase delta) indicates a
fractional charge.
electron poor
electron rich
• This indicates that fluorine has a stronger
attraction for the shared e- than hydrogen
8.2 Electronegativity
• Electronegativity – is the ability of an
atom in a molecule to attract shared
electrons to itself.
• Linus Pauling’s model is what we used to
assign a value to electronegativity.
• The trend is that electronegativity
generally increases across a period and
decreases down a group.
Relationship between
electronegativity and bond type.
Using electronegativity to
determine bond polarity
• Using electronegativity values, arrange
the following bonds in order of increasing
polarity: H – H, O – H, Cl – H, S – H, and F
– H.
Remember bond polarity is the difference of the
electronegativities of the atoms forming the
The answer!!
In increasing polarity
Bond type
Cl – H
8.3 Bond Polarity & Dipole
• When a molecule has a center of
positive charge and center of negative
charge it is said to be dipolar or have a
dipole moment.
• The dipolar character of a molecule is
often represented by an arrow that
points toward the negative charge
center and the tail indicates the
positive center of charge.
Another way to look at it…
• Electrostatic potential maps also show the
charge distribution of a molar molecule.
• Red shows the electron rich
area and the violet shows the
electron poor region.
• Any diatomic molecule with a polar bond
will show a molecular dipole moment
• Polyatomic molecules exhibit dipolar
behavior. And in an electric field, water
acts as if it has 2 centers of charge.
Other times…
• Individual bond polarities are arranged
in such a way that they “cancel” each
other out. As seen below in the CO2
• Polar bonds but NO dipole moment.
• There are many more of these cases!!
Types of Molecules with
Polar Bonds but No Resulting
Dipole Moment
Try these.
• Show the direction of the bond polarities and
indicate which ones have a dipole moment:
Cl2 has NO bond
polarity because
the e- are shared
equally. Therefore,
there is NO dipole
8.4 Ions: Electron
Configuration & Sizes
• Generalizations:
– When 2 nonmetals react to form a covalent
bond, they share e- so that the valence
electron configuration of both atoms is
complete. (i.e. both nonmetals have attained
noble gas electron configurations)
– When a nonmetal and a representative metal
react to form a binary ionic compound, ions
form so that the valence e- configuration of
the nonmetal achieves the e- config of the
next noble gas and the valence orbitals of
the metal are emptied. (both ions attain
noble gas e- configurations)
Predicting formulas of
ionic compounds
• When the term ionic compound is
used, typically, the state of matter
being referred to is a SOLID.
• The + and – ions are packed
together in such a way so that
the + + and - - repulsions are
minimized and the + - attractions
are maximized.
Sizes of Ions
• The size of an ion plays an
important role in the stability of an
ionic solid and the properties of the
ions in solution.
• Most of the time ionic radii are
determined by measuring the
distances between ions in ionic
compounds… BUT this assumes how
the distance is divided between the
two ions.
• This method of determining ion size
creates a controversy…sooo….we
will concentrate on the trends!
• Consider relative ion size and the
size of the parent atom. + ions are
formed by removing an e- so the
resulting cation is SMALLER than its
parent atom. And the opposite is
true for anions.
Depends on parents
• Remember the generalizations we
talked about…
• Size depends on the “parent’s”
position in the periodic table. A
given period has both elements that
give up and gain electrons. So some
elements get larger (to emulate the
next noble gas) and some elements
get smaller (to “look like” the
previous noble gas)
Isoelectronic ions.
• Isoelectronic ions are ions that contain
the same number of electrons.
– For example O2-, F-, Na+, Mg2+, and
Al3+ each have 10 electrons ([Ne]) so
the amount of repulsions should be
the same.
• Let’s look at how their sizes vary.
What other factor should we
look at for size?
• When looking at isoelectronic ions, the
other factor to consider is the number of
protons in the nucleus. The number of
protons increases from 8 to 13 from O2to Al3+.
• So as the number of protons increases,
the ATTRACTION to the 10 electrons is
GREATER and it causes the ions to
become smaller.
8.5 Energy Effect in Binary
Ionic Compounds
• We know that metals and nonmetals react
by transferring e- and that the result is a
solid ionic compound due to the
oppositely charged ions having a lower
energy than the original elements.
• As always ENERGY tells us how strongly
the ions are attracted to each other. In a
solid ionic compound the energy is called
lattice energy.
Lattice energy defined
• Lattice energy is defined as the change
in energy that takes place when
separated gaseous ions are packed
together to form an ionic solid.
• It is the energy released when an ionic
solid forms from its ions!
– Remember the sign is determined from the
system’s P.O.V! Exo = negative: since the
energy is leaving the system.
Li(s)  F2 (g)  LiF(s)
Let’s look at the energy changes involved
with the formation of this ionic solid
Sublimation – solid to gas phase
Ionization of Li
Dissociation of fluorine
Formation of F- ions
Formation of solid LiF from gaseous ions
This is an energy diagram!
Notice the exothermic reaction…more energy
is released than absorbed in the process
Lattice energy calculations
• This leads us to a modified version of
Coulomb’s law:
 Q1Q2 
Lattice energy=k 
 r 
• Where k is a proportionality constant
that depends on the structure of the
solid and the electron configurations
of the ions and r is the shortest
distance between the ions.
 Q1Q2 
Lattice energy=k 
 r 
• Looking at the formula, can you see
that the process becomes MORE
exothermic as the ionic charges
increase AND as the distance between
the ions decreases
• Let’s look at how charges affect energy
by comparing MgO and NaF energy
8.6 Partial Ionic Character
of Covalent Bonds
• Let’s think of polar covalent bonds and
ionic. We understand these bonds to
be between atoms that have different
• They either share e- unequally or
transfer one or more e-. So how do
we tell if it is polar covalent or ionic?
• The REAL answer is there are NO
completely ionic compounds. Let’s
A complication
• The previous slide defies the idea that
we know many of those compounds
(that are above 50%) as ionic solid.
• We must consider that the compounds
in the table are in the gas phase. And
that these results can not be assumed
to also apply to the solid phase.
Another complication
• Another complication in identifying
ionic compounds is that many contain
polyatomic ions.
• Polyatomic ions are held together by
covalent bonds
• So calling NH4Cl or Na2SO4 ionic is
What will we call ionic?
• So from now on…we define ionic
compound as any compound that
conducts an electric current when
8.7 The Covalent Chemical
Bond: A Model
• Bonding is a model proposed to
explain molecular stability.
• The bond concept is a human
invention to provide a method for
dividing up the energy evolved when a
stable molecule is formed from its
component atoms.
• It is physically sensible and makes
sense that atoms can form stable
groups by sharing e- since shared egive a lower energy state because they
are simultaneously attracted by 2
8.8 Covalent Bond Energies
and Chemical Reactions
• From the stepwise decomposition of
methane it is determined that the
energy to break a C—H bond varies in
a nonsystematic way. This also shows
that bond strength varies significantly
with its environment.
• So bond energy as an average is what
is useful to chemists and given in a
table format.
Single, double, triple
and I don’t mean baseball
• Single bonds – 1 pair of e-shared.
• Double bonds – 2 pair of e- shared.
• Triple bonds – 3 pair of e-shared.
• There is a relationship between bond
length and # of shared e-. The more
e- shared, the shorter the bond.
Bond Energy and Enthalpy
• Using bond energy, the energies of
reactions can be approximated.
• Breaking bonds requires energy so the
sign is +
• Formation of bonds releases energy so
the sign is -
The “formula”
• ∆H = sum of the energies required to
break old bonds plus the sum of the
energies released in the formation of
new bonds.
H  D(bonds broken)-D(bonds formed)
energy required
energy released
∑ is the sum of the terms and D represents
the bond energy per mole of bonds. (D is
always positive)
Let’s try with…
H2(g) + F2(g)  2HF(g)
We need to break 1) H—H bond
1) F—F bond
We need to form 1) H—F bond
H  DH H  DF F  DH F
 1mol 
 1mol 
 2mol 
 544KJ
If we compare this with ∆H from standard
enthalpy for HF (-271kJ/mol):
∆Ho = 2mol • -271kJ/mol = -542 kJ
So bond energies work well to find ∆H
8.9 The localized Electron
Bonding Model
• A “localized electron” assumes that a
molecule is composed of atoms that
are bound together by sharing pairs of
electrons using the atomic orbitals of
the bound atoms.
• Lone pairs are the e- that are localized
on the atom (or in space)
• Bonding pairs are those e- between
the atoms.
Localized electron model
when applied tells us…
1. The description of the arrangement of
the valence electrons (Lewis
2. A prediction of the molecule’s
geometry (VSEPR)
3. A description of the atomic orbital
types used to share electrons or hold
long pairs. (we’ll talk about this in
Chapter 9)
8.10 Lewis Structures
• Shows how valence electrons are
arranged among atoms in a molecule.
• Reflects central idea that stability of a
compound relates to noble gas
electron configuration.
Steps for writing/drawing
Lewis Structures
1. Sum the valence electrons. (Add for –
ions, subtract for + ions)
2. Use a pair of electrons to form a
bond between each pair of bound
3. Atoms usually have noble gas
configurations. Exceptions are
Hydrogen (duet rule) and more we’ll
talk about later.
8.11 Exceptions to the
Octet Rule
• We know that hydrogen is an
exception as hydrogen’s outer most
orbital is the 1s so when it is full it is
1s2 therefore it follows the duet rule
not the octet rule
Another e- poor exception
• Boron – Boron is electron deficient therefore
it often only bonds with 6 electrons around it
NOT 8. This is indicated often by its violent
B – 3e3F – 3x7e24e-
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24
• Electron poor atoms are typically found in period
2 (this makes sense, because they don’t have 8
e- to begin with). They will NEVER exceed 8
since their 2s and 2p orbitals can’t exceed 8
Electron rich exceptions
• There are some occasions where an
atom will exceed the octet rule. This
is only observed for some elements in
period 3 and beyond. Common ones
are S and P, but there are more too!
• SF6
S – 6e6F – 42e48e-
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
Should we try?
Draw a Lewis structure for each of the
following molecules:
• H2
• N2
• O2
• F2
• H2O
• NH3
• CO
• They get harder as
you go
8.12 Resonance
• Resonance is when there is more than
one valid Lewis structure for a
particular model
• Reality is that the actual structure is an
average of the resonance structures.
Here’s one example
• Resonance is necessary because the LE
model assumes that e- are localized
between atoms, however, we know
that electrons are really delocalized
and they can move around the entire
• Resonance “compensates” for this
defective assumption.
Odd electron molecules
• There are very FEW of them!!!
• Formed from nonmetals.
• The most common example is NO
N – 5eO – 6e11e-
Formal Charge
• How do we decide which Lewis
Structure best describes the actual
bonding in a molecule.
• One method is to estimate the charge
on each atom in the various possible
Lewis structures. We will use formal
charge to do this.
Things to know:
Formal charge is the difference between
the # of valence e- on the free atom
and the # of valence e- assigned to
the atom in the molecule
To use formal charge we need to know 2
1. The number of valence electrons on the free
neutral atom
2. The number of e- belonging to the atom in
the molecule.

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